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The Alternative Nitrogenases
1. Vanadium Nitrogenase
The "essentiality" of molybdenum for nitrogen fixation was first reported by Bortels in 1930.308 This finding led ultimately to the characterization of the molybdenum nitrogenases discussed in the preceding section. Bortels' work has been cited many times, and is often referred to without citation. Following this seminal work, many other Mo-containing enzymes were subsequently sought and found.25,309 At present more than a dozen distinct Mo enzymes are known, and new ones are continually being discovered.
In addition to the classic 1930 paper, Bortels310 reported in 1935 that vanadium stimulated nitrogen fixation. In contrast to the 1930 paper, the 1935 paper languished in obscurity. Then, starting in the 1970s, attempts were made to isolate a vanadium nitrogenase. In 1971, two groups reported isolating a vanadium-containing nitrogenase from A. vinelandii.311,312 The interesting notion at this time was that V might substitute for Mo in nitrogenase, not that there was a separate system. The isolated enzyme was reported to be similar to the Mo enzyme, but had a lower activity and an altered substrate specificity. One of the groups carefully reinvestigated their preparation, and found small amounts of molybdenum, which were presumed to be sufficient to account for the low activity, although the altered selectivity was not addressed.313 The vanadium was suggested to play a stabilizing role for [FeMo], allowing the small amount of active Mo-containing protein to be effectively isolated. Apparently the possibility was not considered that a truly alternative nitrogenase system existed, whose protein and metal centers both differed from that of the Mo nitrogenase.
The unique essentiality of molybdenum for nitrogenase fixation went unchallenged until 1980, when it was demonstrated314 that an alternative nitrogenfixation system could be observed in A. vinelandii when this organism was starved for molybdenum.315 Despite skepticism from the nitrogenase research community, it was eventually shown that even in a mutant from which the structural genes for the Mo nitrogenase proteins (nif H, D, and K) had been deleted, the alternative system was elicited upon Mo starvation. In 1986, two groups316-321 isolated the alternative nitrogenase component proteins from different species of Azotobacter, and demonstrated unequivocally that one component contained vanadium and that neither component contained molybdenum.
One of the two components of the V-nitrogenase system is extremely similar to the Fe protein of nitrogenase. This similarity is evident in the isolated proteins from A. vinelandii316 and in the genetic homology between nif H (the gene coding for the subunit of the Fe protein in the Mo-nitrogenase system) and nif H* (the corresponding gene in the V-based system). Both Fe proteins have an $\alpha_{2}$ subunit structure, and contain a single Fe4S4 cluster that is EPR-active in its reduced state.
The FeV proteins from Azotobacter vinelandii and Azotobacter chroococcum each have an $\alpha_{2} \beta_{2} \delta_{2}$ subunit structure.322 Metal composition and spectroscopic comparisons between the FeMo and FeV proteins are shown in Table 7.9. Although there is the major difference involving the presence of V instead of Mo in the FeV protein and in the probable presence of the small $\delta$ subunits (13 kDa), the two nitrogenase systems are otherwise quite similar.322 In each, a system of two highly oxygen-sensitive proteins carries out an ATP-dependent N2 reduction with concomitant H2 evolution. The Fe proteins have the same subunit structure and cluster content, and are spectroscopically very similar. The V versions of the larger protein have somewhat lower molecular weights than their Mo analogues, and by MCD spectroscopy seem to contain P-like clusters.318 The FeV site still may be an S = $\frac{3}{2}$ center (by EPR, although its EPR differs significantly from that of the FeMo center).323 The V-S and V-Fe distances as measured by EXAFS324,325 are similar to those in thiocubane VFe3S4 clusters and to Mo-S and Mo-Fe distances like those in [FeMo], which are in turn similar to those in MoFe3S4 thiocubanes. Likewise, XANES324,325 indicates VS3O3 type coordination in [FeV] nitrogenase similar to the MoS3O3 coordination suggested by XANES for FeMoco. The "FeV cofactor" is extractable into NMF, and can reconstitute the nif B-, FeMoco-deficient mutant of the Mo system.326 Despite the substitution of V for Mo, the proteins and their respective M-Fe-S sites do not differ drastically. However, the compositional changes do correlate with altered substrate reactivity.
A major difference between the V and Mo enzymes lies in substrate specificity and product formation.321 As is clearly shown in Table 7.9, the FeV nitrogenase has a much lower reactivity toward acetylene than does the Mo system. Furthermore, whereas the FeMo system exclusively produces ethylene from acetylene, the FeV system yields significant amounts of the four-electron reduction product, ethane.321 The detection of ethane in the acetylene assay may prove a powerful technique for detecting the presence of the V nitrogenase in natural systems.322 Moreover, this reactivity pattern is found in the nif B- mutant reconstituted with FeVco, indicating that the pattern is characteristic of the cofactor and not the protein.326 The reactivity change upon going from Mo to V in otherwise similar protein systems clearly adds weight to the implication of the M-Fe-S center (M = V or Mo) in substrate reduction.
Table 7.9 - Comparison of alternative nitrogenase proteinsa
a) Av1 is the FeMo protein of Azotobacter vinelandii, Av1* is the FeV protein of A. vinelandii and Ac1* is the FeV protein of A. chroococcum. Data from References 317,377, and 319, respectively.
b) Atoms per molecule.
c) nmol product/min/mg of protein.
Property Av147 Av1*47 Ac1*50
Molecular Weight 240,000 200,000 210,000
Molybdenumb 2 < 0.05 < 0.06
Vanadiumb 0.7 2
Ironb 30-32 9.3 23
Activityc
H+ 2200 1400 1350
C2H2 2000 220 608
N2 520 330 350
EPR g values 4.3 5.31 5.6
3.7 4.34 4.35
2.01 2.04 3.77
1.93 1.93
2. The All-iron Nitrogenase322
The first sign that there is yet another alternative nitrogenase again came from genetic studies. A mutant of A. vinelandii was constructed with deletions in both nif HDK and nif H*D*K*, i.e., the structural genes for the Mo and V nitrogenases, respectively. Despite lacking the ability to make the two known nitrogenases, the mutant strain nevertheless was able to fix nitrogen, albeit poorly. Moreover, this mutant strain's nitrogenase activity was clearly inhibited when either Mo or V was present in the culture medium. Preliminary studies indicate that the nitrogenase proteins produced by this organism are closely related to those previously isolated. A 4Fe-4S Fe-protein nif H$\dagger$ and a protein due to nif D$\dagger$ was produced. The latter appeared to contain no stoichiometric metal other than iron. Symmetry of nomenclature would suggest calling this the FeFe protein and its cofactor FeFeco. Interestingly, this nitrogenase seems to be the poorest of the set in reducing N2 and makes ethane from ethylene. The finding of the all-iron nitrogenase, if fully confirmed, will add significantly to the comparative biochemistry of nitrogen fixation. Speculatively, one might suggest that the concomitant absence of V and Mo suggests that nitrogen fixation need not directly involve the noniron heterometal in the cofactor cluster. This result may explain the lack of direct implication of Mo in the nitrogen fixation mechanism, despite many years of intense effort by workers in the field. (The above discussion should be taken cum grana salis until the existence of the all-iron nitrogenase is confirmed.)
3. Model Systems
Three types of model systems for nitrogenase may be considered. First, there are transition-metal sulfide clusters that resemble the FeMoco or FeVco centers of the active proteins. Although there has been significant progress, there are not yet any definitive models (as there are for Fe2S2 and Fe4S4). A second approach uses the reactions of N2 and related substrates or intermediates with metal centers in order to gain insights into the way in which transition-metal systems bind N2 and activate it toward reduction. Here to date the most reactive systems bear little direct chemical resemblance to the nitrogenase active sites. Nevertheless, these systems carry out bona fide nitrogen fixation from which one may learn the various ways in which N2 can be activated. Finally, there are other inorganic systems that display some of the structural and possibly some of the reactivity characteristics of the nitrogenase active sites without binding or reducing N2 or precisely mimicking the active center. We may nevertheless be able to learn effectively about nitrogenase reactivity from these interesting chemical systems.
a. Transition-metal Sulfide Models for Nitrogenase Sites
Although there has been great activity in synthetic Fe-S cluster chemistry, there is to date no example of a spectroscopic model for the P-cluster sites in nitrogenase. If the P-clusters are indeed asymmetrically bound high-spin Fe4S4 clusters, then the recent work on high-spin versions of Fe4S4 clusters327 and site-selectively derivatized Fe4S4 centers143 may hint that appropriate model systems are forthcoming.
b. Fe-M-S Cluster Models for FeMoco
Despite the importance of P-clusters, the modeling of the FeMoco center has properly received the most attention. The significant structural parameters that any model must duplicate are the Mo-S and Mo-Fe distances detennined by EXAFS. Spectroscopically, the S = $\frac{3}{2}$ EPR signal provides a stringent feature that model systems should aspire to mimic.
Many FeMoS clusters have been prepared in the quest to duplicate the FeMoco center, but none of the chemically synthesized clusters can reactivate the (UW-45 or Nif B-) cofactor-less mutants, perhaps because of their lack of homocitrate, which only recently has been discovered as a key component of FeMoco. Undoubtedly, new FeMoS clusters containing homocitrate will be prepared, and perhaps these will activate the mutant proteins, thereby revealing a close or full identity with FeMoco.
Despite the absence of homocitrate, some interesting model systems have been investigated. It is beyond the scope of this chapter to give a comprehensive account of FeMoS chemistry. We concentrate on the so-called "thiocubane" model systems. Heterothiocubane models were first synthesized using self-assembly approaches analogous to those used for the simpler Fe-S model systems. The reaction328-330a
$MoS_{4}^{2-} + Fe^{3-} + SR^{-} \rightarrow (MoFe_{3}S_{4})_{2}(SR)_{9}^{3-} \; and\; (MoFe_{3}S_{4})_{2}Fe(SR)_{12}^{3-,4-} \tag{7.17}$
uses tetrathiomolybdate, MoS42-,as the source of Mo, and leads to the double cubane structures shown in Figure 7.32A,B. The Fe7Mo2S8 structure proved particularly interesting, since it was possible to complex the central ferric iron atom with substituted catecholate ligands331,332 and eventually isolate a single thiocubane unit (Figure 7.32C). Significantly, the single unit has S = $\frac{3}{2}$ and Mo-S and Mo-Fe distances that match precisely those found by EXAFS for the M center of nitrogenase. Single cubes with VMo3S4 cores have also been prepared.160,160a Although the single thiocubanes display spectroscopic similarity and distance identity with FeMoco, they are not complete models. They are stoichiometrically Fe and S deficient, lack homocitrate, and most importantly, fail to activate the UW-45 and Nif B- mutants.
Other interesting FeMoS (and FeWS) clusters with structurally distinct properties are shown in Figure 7.33. These include the "linear" (MoS4)2Fe3- ion, the linear (WS4)2Fe[HCON(CH3)2]22- ion, the linear Cl2FeS2MS2Fe2CI22- (M = Mo, W), the "linear" (MoS4)2Fe2S24- ion, the trigonal (WS4)3Fe3S24-, the capped thioprismane Fe6S6X6[M(CO)3]23- (X = CI, Br, I; M = Mo, W), and the organometallic clusters MoFe6S6(CO)162-, MoFe3S6(CO)6(PEt3)3, and MoFe3S6(CO)62-. Structures suggested for FeMoco based on these and other chemically synthesized transition metal sulfides and on spectroscopic studies of the enzyme are shown in Figure 7.31. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/07%3A_Ferrodoxins_Hydrogenases_and_Nitrogenases_-_Metal-Sulfide_Proteins/7.09%3A_Multisite_Redox_Enzymes_%28Part_5%29.txt |
N2 and Related Complexes
The triple bond of N2 has one $\sigma$ and two $\pi$components. Each nitrogen atom has a lone pair oriented along the N-N direction. The two lone pairs allow N2 to bind in an end-on fashion in either a terminal or a bridging mode. Both modes of binding are illustrated in the binuclear zirconium complex333 shown in Figure 7.34. In this and in many other N2 complexes, the N-N bond is not significantly lengthened and is therefore presumed to be insignificantly weakened in the complex. Interestingly, the complex in Figure 7.34, despite not having long N-N distances, forms hydrazine quantitatively upon protonation. Only one of the three N2 molecules is reduced, and all four electrons required come from the two Zr(III) by presumed internal electron transfer. The related $\mu])-N2 complex [W(\(\eta^{5}$-C5Me5)Me2(SC6H2Me3)]2($\mu$-N2) is one of the few dinitrogen complexes to contain an S donor ligand.333a
In addition to the N lone pairs, the $\pi$ components of the N$\equiv$N triple bond can serve as donor-acceptor orbitals in the Dewar-Chatt-Duncanson (olefin binding) manner. This less-common mode of N2 binding is illustrated by the structure of the Ti complex334 shown in Figure 7.35. Here, as in the few other known side-on bound N2 complexes,335 the N-N bond is significantly lengthened. The lengthened bond at 1.30 Å is presumed to be sufficiently weakened [v(N-N) = 1280 cm-1] that it is susceptible to further lengthening and reduction. As the N-N distance lengthens, it is more appropriate to consider the ligand as a deprotonated diimide or hydrazine.
Complexes that have proven particularly useful are bis(dinitrogen)phosphines of Mo(0) and W(0) such as M(N2)2(Ph2PCH2CH2PPh2)2 and M(N2)2(PPh2Me)4. As shown in Figure 7.36, treatment of the complexes336-337a with acid leads to the formation of the diazenido(-H) and hydrazido(-2H) complexes, and sometimes to the production of ammonia. The finding of a bound N2H22- species is consistent with the proposed presence of similar bound species in nitrogenase. The complexes of reduced dinitrogen intermediates are stabilized by multiple M-N binding. Further protonation of these intermediates or treatment of the original complex with strong acid leads to the formation of NH3 from the bound nitrogen. Here the Mo(0) starting complex has enough electrons [six from the Mo(0) → Mo(VI) conversion] to reduce one N2 molecule in conjunction with its protonation from the external solution.
In a general sense this reaction may be telling us something about nitrogenase. The enzyme may be able to deliver six reducing equivalents to N2, and protonation, perhaps carefully orchestrated by neighboring amino-acid or homocitrate groupings, may facilitate the process. However, it is virtually certain that the Mo in nitrogenase is not able to change its oxidation state by six units. In the enzyme the multimetal, multisulfur FeMoco site may serve the equivalent function, by providing multiple sites at which reduced intermediates can simultaneously bind.
Only a few of the known N2 complexes contain S-donor ligands. One of these, Mo(N2)2(S(CH2C(CH3)2CH2S)3), shown in Figure 7.37, has four thioether S-donor atoms bound to Mo(0). This Mo(0) complex shows reactivity reminiscent of the related phosphine complexes.337a A remarkable complex (Figure 7.38) has been isolated338 in which two lone pairs of trans-diimide bind to two Fe, concomitantly with H-binding of the two diimide hydrogen atoms to coordinated sulfur atoms. The ability of an Fe-S system to stabilize the very reactive trans-N2H2 grouping adds support to the notion that similar metal-sulfide sites of nitrogenase may stabilize related intermediates along the N2 → 2NH3 reaction path.
Most of the model systems involving N2 do not lead to NH3 formation. Moreover, many systems that do form NH3 are not catalytic. However, certain V-based and Mo-based systems can catalytically reduce N2 to N2H4 or NH3 using strong reducing agents.339 Although kinetic studies indicate the possibility of intermediates, little structural information is available at present on these interesting systems.
Insights from Relevant Inorganic Reactivity
Certain studies on inorganic systems that do not model the nitrogen-fixation process can nevertheless potentially give insight into nitrogenase action. Two categories of relevant chemistry are acetylene binding/reactivity and dihydrogen binding/activation. Modes of dihydrogen activation on sulfide systems have previously been discussed in the section on hydrogenase.
Acetylene has long been known to bind to metal centers using its $\pi$and $\pi$* orbitals as, respectively, $\sigma$-donor and $\pi)-acceptor orbitals. Even when the metal is predominantly sulfur-coordinated,340,341 such side-on bonding of RC2R is well known340,341 as in MoO(S2CNR2)2(RC\(\equiv$CR) and Mo(S2CNR2)2(RC$\equiv$CR)2. The direct interaction of acetylene with the metal center must be considered as a potential binding mode for nitrogenase substrates.
A totally different, sulfur-based mode of acetylene binding is now also well established. For example, (Cp')2Mo2S4 reacts with acetylene342,225 to produce
$\tag{7.18}$
containing a bridging ethylene-1,2-dithiolate (dithiolene). The acetylene binds directly to the sulfur atoms by forming S—C bonds. Acetylenes or substituted (activated) acetylenes are able to displace ethylene from bridging or terminal 1,2-dithiolate ligands225,341 to produce the 1,2-dithiolenes. In these reactions the sulfur rather than the metal sites of the cluster are reactive toward these small unsaturated molecules. Clearly, for nitrogenase, where we do not know the mode of binding, sulfur coordination might be a viable possibility. The (Cp')2Mo2S4 systems that bind H2 and C2H2, wherein bound C2H2 can be reduced to C2H4 and displaced by C2H2, are potential models for substrate reduction by nitrogenase.225,342
The versatility of transition-metal sulfur systems is further illustrated by the observation that activated acetylene can insert into a metal-sulfur bond in Mo2O2S2(S2)22-, forming a vinyl-disulfide-chelating ligand
It has recently been suggested343 that the presence of a dihydrogen complex is required for H2 to be displaced by N2 to form a dinitrogen complex. This reaction would explain the required stoichiometry of N2 reduction and H2 evolution. Such an explanation had been suggested previously with dihydride complexes acting as the N2-binding and N2-displacing site.182 Clearly, this new suggestion is an interesting embellishment of potential N2/H2 relationships.
At present, the activation process that is at work in the enzyme is unknown. We need greater structural definition of the active site, which should be forthcoming through the continued application of sophisticated diffraction and spectroscopic probes. Diffraction alone, however, will be incapable of locating protons and possibly other low-molecular-weight ligands. Therefore, spectroscopic probes such as ENDOR10 and ESEEM,277-Z79,344 which are based on EPR spectroscopy, and x-ray-based techniques, such as EXAFS and XANES, will remain crucial in elucidating mechanistically significant structural details. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/07%3A_Ferrodoxins_Hydrogenases_and_Nitrogenases_-_Metal-Sulfide_Proteins/7.10%3A_Multisite_Redox_Enzymes_%28Part_6%29.txt |
A significant breakthrough has occurred in the crystallographic analysis of the iron-molybdenum protein of nitrogenase. The overall distribution of the metal clusters in the protein is shown in Figure 7.40. The distance between the two FeMoco units is fully consistent with each cofactor acting as an independent active site. On the other hand, the closeness of the P cluster and FeMoco centers in each unit is indicative of their likely cooperation in the N2 fixation reaction.
The proposed structure of the P cluster, shown in Figure 7.41, involves a doubly bridged, double cubane unit consisting of one normally bound Fe4S4 cluster with all cysteine ligands and one Fe4S4 cluster that contains an unusual cysteine/serine (S/O) ligand pair on one of its two nonbridged Fe positions. Such five-coordinate iron in an Fe4S4 cluster is not unprecedented.138 The two Fe4S4 clusters are disposed to produce a face-sharing arrangement with two cysteine ligands bridging the two sets of Fe atoms. An interesting feature of the structure is a disulfide unit linking the two clusters; this unit potentially could be redox-active during nitrogenase turnover.
Most striking of the new results is the proposed structure of FeMoco shown in Figure 7.42. The cluster core of composition Fe7MoS8 can be viewed as two halves bridged by two S2- ions and an unknown ligand (designated Y in the figure). The MoFe3S3 half of the core is in the shape of a thiocubane fragment missing one $\mu_{3}$-S2- ion. The Mo is six coordinate; the ligands are three $\mu_{2}$-S2- ions, which bridge to the three Fe ions, an $\alpha$-His-442 nitrogen, and two oxygen donors (the hydroxyl and central carboxylate) of the homocitrate ligand. Interestingly, the second half of FeMoco is a similar thiocubane fragment, Fe4S3, also missing a $\mu_{3}$-S2- ion. This unit has a single noncore ligand, $\alpha$-Cys-275, which is bound to the terminal Fe atom of the cluster. The two thiocubane fragments (MoFe3S3 and Fe4S3) are bridged by three ligands in a face-sharing mode with the two Fe3 faces eclipsed with respect to each other. The eight metal ions display a bis(end-capped) trigonal prismatic arrangement with three bridges on the edges of the prism, which connect the two thiocubane fragments. The two sulfide bridges between the thiocubane halves are clearly defined in the structure, but the third bridge is not, suggesting the possibility that this is in fact part of the N2-binding site. Interestingly, $\alpha$-His-195, identified as essential for N2 fixation by mutagenesis and ESEEM studies, does not appear to be covalently bound, although it is close to the FeMoco unit.
Clearly, this structure is not the same as any of those previously proposed (Figure 7.31), although it does possess many features that were identified in model studies. While it is tempting to speculate that the central bridge of the cluster (the Y ligand) is the site of N2 reduction, this is in no way established at present. The structural definition of the nitrogenase proteins is now progressing at a rapid rate. Many of the physical measurements will have to be reexamined in light of the new data. Through further experimentation involving physical methods, mutagenesis, and kinetic/mechanistic studies, much more information about the role of ATP, the activation of hydrogen, and the binding, activation, and reduction of N2 and other nitrogenase substrates should be obtained. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/07%3A_Ferrodoxins_Hydrogenases_and_Nitrogenases_-_Metal-Sulfide_Proteins/7.11%3A_Report_on_the_Nitrogenase_Crystal_Structure%28378-381%29.txt |
The physiological role of rubredoxins (sometimes abbreviated as Rd) is not always known with certainty. In particular, although rubredoxin was first identified26 in the anaerobe Clostridium pasteurianum, its role in anaerobic metabolism remains obscure. Some rubredoxins, such as that from the aerobe Pseudomonas oleovorans, participate in fatty acid $\omega$-hydroxylation, i.e., hydroxylation at the end of the hydrocarbon chain farthest from the carboxylic acid.27 Like the Fe2S2 proteins putidaredoxin28 and adrenodoxin,29 the rubredoxin provides electrons to the hydroxylase, which acts as a monooxygenase forming the w-alcohol product and water (see Figure 7.3). In a reaction catalyzed by rubredoxin reductase, rubredoxin is reduced by NADH to the ferrous state and reoxidized by the w-hydroxylase to the ferric form during the catalytic cycle.
Most rubredoxins contain a single Fe atom, which can exist in the ferrous or ferric state. For the rubredoxin from Clostridium pasteurianum,26 the E°' value is -57 mV, which is much more positive than that of ferredoxins from the same organism (see below). The 6-kDa clostridial protein has only 54 amino acids in its polypeptide chain, and has a very low isoelectric point of 2.93. The rubredoxin from P. oleovorans27 has one or two iron atoms in a single polypeptide chain of MW ~ 20 kDa. Its redox potential is -37 mV for the Fe3+/2+ couple. Rubredoxins as a class show considerable sequence identity, and the larger 2Fe members of the class show evidence, involving internal-sequence homology, that they may have. evolved through gene duplication.
A protein from Desulfovibrio gigas, called desulforedoxin,30,31 appears to resemble rubredoxins in some respects, but the two Fe atoms in the 7.6-kDa protein appear to be spectroscopically and structurally distinct from the Fe atoms in rubredoxins.31 A protein from Desulfovibrio vulgaris called ruberythrin has a single rubredoxin site as well as a strongly coupled 2Fe site resembling that of hemerythrin. Its physiological function is unknown. Table 7.1 lists some of the known rubredoxins and their properties.
Table 7.1 - Properties of some iron-sulfur proteins.
a) g'-tensors for ± $\frac{1}{2}$ and ±$\frac{3}{2}$ Kramers doublets, respectively, of the S = $\frac{5}{2}$ system. The values of 0.9 and 1.25 are calculated (not observed)44
b) The fully reduced protein has a complex spectrum due to magnetic coupling between the two identical Fe4S4 clusters. The g-values are those for partly reduced samples, and represent a magnetically isolated cluster.
c) The reported spectrum is complex because of magnetic interaction with the reduced Fe3S4 cluster.
d) Recent evidence suggests that Thermus thermophilus and Thermus aquaticus are actually the same species.362 EPR parameters of the homologous Thermus thermophilus ferredoxin estimated from computer simulations.361 In this protein a signal originating from the Fe3S4 cluster at g' $\simeq$ 12, attributable to $\Delta$Ms = ± 4 transitions, is observed for the reduced (S = 2) cluster.
Protein Source
Molecular Weight
(subunits)
Fe-S Composition
Redox Potention
mV (pH)
EPR
g values
References
Rubredoxins
Clostridium pasteurianum 6,000 Fe -58 (7) 9.3 4.3 26
Pseudomonas oleovorans 6,000 Fe
9.42
4.02
0.9
4.77
1.25a
4.31
44
Fe2S2 Proteins
Spinach ferredoxin 11,000 [2Fe-2S] -420 (7.0) 2.05 1.96 1.89 350, 351
Parsley ferredoxin 11,000 [2Fe-2S] 2.05 1.96 1.90 352
Euglena ferredoxin 11,000 [2Fe-2S] 2.06 1.96 1.89 352
Adrenal cortex ferredoxin (pig)
[Adrenodoxin]
16,000 [2Fe-2S] -270 (7.0) 2.02 1.93 1.93 352, 353
Pseudomonas putida ferredoxin
[Putidaredoxin]
12,500 [2Fe-2S] -240 (7.0) 2.02 1.93 1.93 352, 353
Clostridium pasteurianum 25,000 [2Fe-2S] -300 (7.5) 2.00 1.96 1.94 354
Xanthine Oxidase 280,000 (2)
2 x [2Fe-2S] I
2 x [2Fe-2S] II
-343 (8.2)
-303 (8.2)
2.02
2.12
1.94
2.01
1.90
1.91
355, 356
Thermus thermophilus Rieske 20,000 2 x [2Fe-2S] +150 (7.8) 2.02 1.90 1.80 93, 357
Fe4S4 Proteins
Clostridium pasteurianum 6,000 2 x [4Fe-4S] -420 (8.2) 2.06 1.92 1.89b 115
Bacillus stearothermophilus 9,100 [4Fe-4S] -280 (8.0) 2.06 1.92 1.89 358
Desulfovibrio gigas ferredoxin I 18,000 (3) [4Fe-4S] -455 (8.0) 2.07 1.94 1.92 359
Aconitase (beef heart) [active]
81,000
[4Fe-4S] 2.06 1.93 1.86
Chromatium vinosum HiPIP
10,000
[4Fe-4S] +356 (7.0) 2.12 2.04 2.04 353
Paracoccus sp. 10,000 [4Fe-4S] +282 (7.0) 353
Azotobacter vinlandii Fd I 14,500
[3Fe-4S]
[4Fe-4S]
-645 (8.3)
2.06
1.93
1.89c
360
Thermus aquaticus 10,500
[3Fe-4S]
[4Fe-4S]
-550 (9.0)
2.06
1.93
1.92c
353, 361
Fe3S4 Proteins
Desulfovibrio gigas Fd II 6,000 (4) [3Fe-4S] -130 (8.0) 2.02 359
Azotobacter vinlandii Fd I 14,500
[3Fe-4S]
[4Fe-4S]
-450 (8.3) 2.01 360
Thermus aquaticus 10,500
[3Fe-4S]
[4Fe-4S]
-260 (9.0) 2.02 1.99 1.94d 353, 361
Aconitase (beef heart) [inactive] 81,000 [3Fe-4S] 2.01
The x-ray crystal structures of the rubredoxins from C. pasteurianum32 and D. vulgaris33 have been determined.33a The C. pasteurianum protein structure is known to a resolution of 1.2 Å, placing it among the metalloproteins whose structures are known with greatest precision. The individual Fe and S atoms are clearly resolvable. As shown in Figure 7.4, the single iron is coordinated by four S ligands provided by Cys-6, Cys-9, Cys-39, and Cys-42. The sequence Cys-x-y-Cys is a common one in Fe-S proteins, because it allows both cysteine residues to bind to the same metal site or cluster. The Fe-S distances and angles in the clostridial rubredoxin are shown in Table 7.2. The range of distances and angles reveals a slightly distorted tetrahedral structure.
Table 7.2 - Bond distances and bond angles around Fe in rubredoxin from Clostridium pasteurianum (W1).32
Distance (Å)
Fe-S[Cys(6)] 2.33(11)
Fe-S[Cys(9)] 2.288(15)
Fe-S[Cys(39)] 2.300(15)
Fe-S[Cys(42)] 2.235(12)
Angle (°)
S-Fe-S[Cys(6)-Fe-Cys(9)] 113.8 (4)
S-Fe-S[Cys(6)-Fe-Cys(39)] 109.0 (4)
S-Fe-S[Cys(6)-Fe-Cys(42)] 103.8 (4)
S-Fe-S[Cys(9)-Fe-Cys(39)] 103.7 (4)
S-Fe-S[Cys(9)-Fe-Cys(42)] 114.3 (5)
S-Fe-S[Cys(39)-Fe-Cys(42)] 112.4 (5)
The initial structural results on C. pasteurianum rubredoxin were reported at a slightly lower resolution than those displayed in Table 7.2. In fact, the early study34 reported a range of Fe-S distances from 2.05 to 2.34 Å. Prior to the higher-resolution refinement, a synchrotron-radiation x-ray-absorption spectroscopy study of the iron-absorption edge of rubredoxin was reported.35,36 Using the technique of Extended X-ray Absorption Fine Structure,* EXAFS, the average Fe-S distance was found 35,36 to be 2.26 Å, in agreement with the average distance from the x-ray crystallographic study. However, the EXAFS indicated a much narrower permissible range of Fe-S distances than did the early crystallographic study. The later, more highly refined crystallographic treatment32 agreed nicely with the EXAFS result, illustrating the importance of applying more than one technique to the elucidation of key parameters. Here, as with the 3Fe proteins we will discuss later, EXAFS proved a useful complementary technique to x-ray crystallography.
* X-ray absorption spectroscopy is most commonly (and conveniently) used with the K-edges of transition-metal ions, such as Fe or Mo. It can be split up into two distinct types; X-ray Absorption Near-Edge Structure (or XANES), and Extended X-ray Absorption Fine Structure (or EXAFS) spectroscopy. The former consists of features near the absorption edge itself, which are due to transitions of the photoelectron to bound states and also to other, more complex, phenomena (e.g., the so-called shape resonances). Although the spectra are highly dependent on the nature of the site, they are quite difficult to interpret, and most analyses are based upon simple comparisons with spectra from model compounds. The EXAFS are oscillations of the absorption coefficient at rather higher x-ray energies, and arise from scattering of the emitted photoelectron by surrounding atoms. In contrast to the XANES, EXAFS spectra are relatively simple to interpret in a quantitative manner, yielding a local radial structure. With proper interpretation of the spectra, very accurate interatomic distances (e.g., to ± 0.02 Å), plus more approximate ligand coordination numbers and atomic numbers can be obtained.
The tetrahedral iron sites in rubredoxins offer an interesting glimpse of ligand-field theory in action, and illustrate the use of various physical methods in deducing electronic structure and coordination geometry. The four sulfur ligands are expected to split the iron 3d orbitals into e and t2 sets, with the e set lower as shown in Figure 7.5. The small tetrahedral splitting causes the 3d5 Fe3+ ion to have five unpaired electrons, (e)2(t2)3, 6A1. Consistent with this configuration, the magnetic susceptibility of rubredoxin gives a $\mu_{eff}$ of 5.85 Bohr magnetons.37 No spin-allowed ligand-field transitions are expected, and the red color is caused by S → Fe charge-transfer transitions in the visible region.38,39
In contrast, the 3d6 Fe2+ state, with one additional electron, has four unpaired electrons, as confirmed by its magnetic moment of 5.05 Bohr magnetons. In exact tetrahedral symmetry, a single, low-energy, low-intensity d-d absorption of designation 5E → 5T [(e)3(t2)3 → (e)2(t2)4] is expected for the high-spin ferrous site (Figure 7.5). Indeed, reduced rubredoxin displays a band in the near-infrared region at 6,250 cm-1 that arises as a component of the 5E → 5T2 transition.40 This band stands out particularly vividly in the low-energy circular dichroism (CD) spectrum of reduced rubredoxin.41 Moreover, magnetic circular dichroism (MCD) has proven valuable in dissecting electronic transitions in several rubredoxins and metal-sulfide proteins.38,39,42,43
The EPR spectrum of oxidized rubredoxin (Figure 7.6) shows characteristic peaks at g = 4.31 and 9.42 (P. oleovorans), which have been assigned44 to transitions within excited and ground-state Kramers doublets, respectively, of a nearly completely rhombic S = $\frac{5}{2}$ site, with D = 1.8 and E = 0.5 cm-1. These values for the mononuclear Fe3+ ion stand in sharp contrast to those for other iron-sulfur proteins, which are usually S = $\frac{1}{2}$ (when reduced) and have g values close to 2. The even-electron Fe2+ state (S = 2) in reduced rubredoxin has no detectable EPR when conventional instruments are used.*
Mössbauer spectroscopy has proven to be a particularly powerful complementary tool to EPR in probing the iron sites in Fe-S proteins.3,37,51,52 It is a nuclear spectroscopy that can give valuable information not available from other techniques.$\dagger$ Unlike EPR, where only paramagnetic centers are "seen," every 57Fe atom in the sample will contribute to the Mössbauer spectrum. For rubredoxin, the high-spin nature of the ferric and ferrous sites are clearly seen in the Mössbauer spectra.53 The high-spin Fe3+ sites show a small quadrupole splitting of roughly 0.7-0.8 mm/s due to the almost spherical distribution of the five d electrons in the five d orbitals (Figure 7,7A). In contrast, the high-spin Fe2+ ion with an additional d electron has a significant asymmetry, and thus displays large and quite characteristic quadrupole splitting of 3.1-3.4 mm/s (Figure 7.7B). The isotope shift also distinguishes between Fe2+ and Fe3+, although not as dramatically.37 Finally, the observation53 of magnetic hyperfine interaction in the Mössbauer spectrum at low temperature in the Fe3+ state directly reveals the presence of unpaired electrons, i.e., magnetic coupling with a hyperfine field of 370 ± 3 kG. Although in rubredoxins with a single Fe atom, this observation of magnetic coupling does not reveal any new information, similar magnetic coupling is particularly useful in unraveling the Fe sites in more complex multiiron proteins.
* But see Reference 45. EPR spectroscopy uses magnetic fields to split the electron spin states into levels that differ by energy in the microwave region of the spectrum. For an S = $\frac{1}{2}$ system, the g value (and its anisotropy) and the a values (hyperfine splitting from various nuclei and their anisotropy) are the major parameters reported. EPR spectroscopy has played a role in the development of Fe-S biochemistry akin to the role played by optical spectroscopy in the development of the biochemistry of the cytochromes,46-49 particularly for mitochondria47 and chloroplasts,50 where the g = 1.9 EPR signal has facilitated the monitoring of electron flow through these redox systems. Although EPR has been a powerful tool, it does have some important limitations. A necessary but not sufficient condition for EPR is that the center to be observed must be in a paramagnetic state. Fortunately, this condition is met for at least one member of each one-electron redox couple, i.e., the odd-electron species. However. even when the even-electron species is paramagnetic, it is usually not observed in the EPR, because of the presence of large zero-field splittings. Moreover, relaxation effects and/or the population of excited states often cause the EPR of proteins to be unobservable at room temperature. necessitating the use of liquid N2 or liquid He temperatures to observe the signals in the frozen state. The need to freeze samples prior to observation can lead to artifacts involving the observation of nonphysiological states and processes. On the positive side, the low temperature increases the signal intensity by altering the Boltzmann distribution of the spin population, and allows various quenching techniques to be used with EPR to evaluate kinetic and electrochemical parameters. Nevertheless, one cannot usually observe real-time kinetics or be certain that one is observing a physiologically relevant state. Despite these caveats, EPR has proven a valuable and, in some cases, indispensable tool for identification and monitoring of Fe-S sites. Recently, the advanced EPR techniques ENDOR (Electron Nuclear Double Resonance) and ESEEM (Electron Spin Echo Envelope Modulation) have allowed the extraction of additional information from the EPR signal.
$\dagger$ Mossbauer spectroscopy measures nuclear absorption of light at $\gamma$-ray energies, and can be used to probe nuclear energy levels (usually of 57Fe). The splitting of these levels is influenced by the (s) electron density at the nucleus, and by the electric-field gradient that is set up by nearby atoms. These factors affect the isomer shift and the quadrupole splitting of the Mössbauer spectrum, respectively. Information on nuclear hyperfine couplings is also available when experiments are conducted in the presence of an external (usually applied) magnetic field. Fortunately, the nucleus most commonly (and easily) studied by this technique is present in all the proteins discussed in this chapter, although the level of 57Fe (2 percent natural abundance) must be increased by isotopic enrichment to achieve a high-enough signal-to-noise ratio. For spectra containing one type of site, the spectra are relatively straightforward to interpret. For multisite systems deconvolution is required to get data on individual centers. When possible, selective labeling of sites with 57Fe is extremely helpful in the deconvolution process.
NMR studies on both the oxidized and the reduced states of rubredoxins have been reported. The strongly paramagnetic iron atoms have a profound effect on the NMR spectra of protons in the vicinity of the iron. The iron drastically affects the relaxation behavior of such protons, causing line-broadening, sometimes so much that the protons become nonobservable. If observed, the protons are shifted far from the values found in diamagnetic proteins by combinations of Fermi contact (through overlap/through bond) and pseudo-contact (through space/dipolar) coupling.54,55 In the rubredoxins, the reduced state shows resolved spectra,37,56 which can be assigned56 with the help of data from model systems.*
Resonance Raman spectroscopy provides information involving molecular vibrations that is not dependent on either nuclear or magnetic properties. Electronic excitation of bands involving S → Fe charge transfer often leads to resonance enhancement of Fe-S stretching modes. In rubredoxin,57,57a the Fe-S stretching vibrations are located between 300-400 cm-1. Deviations of the expected two-band tetrahedral pattern (T2 and A1 modes) are attributable to coupling of the Fe-S vibrations with S-C-C bending modes. This coupling makes for greater variability, and the detailed vibrational assignment is thus more difficult for bands involving the cysteinyl sulfur atoms. In contrast, for sites containing inorganic S2-, the Fe-S vibrations involving the inorganic core are less variable and therefore more characteristic of the core type.57
In theory, each of the spectroscopic techniques applied to rubredoxins can give useful information about the other iron-sulfur proteins. In practice, some techniques have proven more useful than others in particular situations, and combined use of several techniques is necessary to draw meaningful conclusions.
Chemical studies of rubredoxins have led to the replacement of the Fe2+ with an Fe4S4 center,58 with Co2+, and with Ni2+. The Co2+ replacement of Fe2+, in P. oleovorans rubredoxin, leads to a stable protein that displays reduced (but not trivial) reactivity in the $\omega$-hydroxylation reaction.59,60 The spectral properties of the cobalt(II) site show the expected changes in d-d bands and the expected shifts in charge-transfer transitions.59 Interestingly, when Ni2+ is substituted into rubredoxins from desulfovibrio species, the resultant proteins show hydrogenase activity.61
* NMR is a technique whose great utility in the study of low-molecular-weight proteins and model systems has not (yet) carried over to the study of larger proteins. Slower tumbling rates, rapid electronic relaxation, multiple paramagnetic sites, large numbers of protons, and more dilute solutions conspire to make the observation and/or interpretation of NMR spectra a daunting task in multisite redox proteins of > 50 kDa. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/07%3A_Ferrodoxins_Hydrogenases_and_Nitrogenases_-_Metal-Sulfide_Proteins/7.12%3A_Rubredoxin-_A_Single-Fe_Tetrathiolate_Protein.txt |
I. Introduction
The interest of the bioinorganic community in the field of metal/nucleic-acid interactions has burgeoned in the last decade. This interest and the resulting progress have come about primarily because of the tremendous advances that have occurred in nucleic-acid technology. We can now isolate, manipulate, and even synthesize nucleic acids of defined sequence and structure, as we would other molecules that chemists commonly explore. Furthermore, as may be evident already in other chapters of this book, bioinorganic chemistry has itself been evolving from a field focused on delineating metal centers in biology to one that includes also the application of inorganic chemistry to probe biological structures and function. In the past decades it has become clear that nucleic acids, structurally, functionally and even remarkably in terms of catalysis, play active and diverse roles in Nature. Transition-metal chemistry, both in the cell and in the chemist's test tube, provides a valuable tool both to accomplish and to explore these processes.
There are also many practical motivations behind the study of how metal ions and complexes interact with nucleic acids. Heavy-metal toxicity in our environment arises in part from the covalent interactions of heavy-metal ions with nucleic acids. In addition, these heavy metals interfere with metalloregulatory proteins and in so doing disrupt gene expression. We need to understand the functioning of the natural metalloregulators of gene expression and we need to design new metal-specific ligands, which, like the proteins themselves, capture heavy metals before their damage is done. Heavy-metal interactions with nucleic acids indeed have provided the basis also for the successful application of cisplatin and its derivatives as anticancer chemotherapeutic agents (see Chapter 9). The design of new pharmaceuticals like cisplatin requires a detailed understanding of how platinum and other metal ions interact with nucleic acids and nucleic-acid processing. Furthermore, we are finding that metal complexes can be uniquely useful in developing spectroscopic and reactive probes of nucleic acids, and hence may become valuable in developing new diagnostic agents. Finally, Nature itself takes advantage of metal/nucleic-acidchemistry, from the biosynthesis of natural products such as bleomycin, which chelates redox-active metal ions to target and damage foreign DNA, to the development of basic structural motifs for eukaryotic regulatory proteins, the zinc-finger proteins, which bind to DNA and regulate transcription. In all these endeavors, we need first to develop an understanding of how transition-metal ions and complexes interact with nucleic acids and how this chemistry may best be exploited.
In this chapter we first summarize the "basics" needed to consider the interactions of metal ions and complexes with nucleic acids. What are the structures of nucleic acids? What is the basic repertoire of modes of association and chemical reactions that occur between coordination complexes and polynucleotides? We then consider in some detail the interaction of a simple family of coordination complexes, the tris(phenanthroline) metal complexes, with DNA and RNA to illustrate the techniques, questions, and applications of metal/nucleic-acid chemistry that are currently being explored. In this section, the focus on tris(phenanthroline) complexes serves as a springboard to compare and contrast studies of other, more intricately designed transition-metal complexes (in the next section) with nucleic acids. Last we consider how Nature uses metal ions and complexes in carrying out nucleic-acid chemistry. Here the principles, techniques, and fundamental coordination chemistry of metals with nucleic acids provide the foundation for our current understanding of how these fascinating and complex bioinorganic systems may function.
Contributors and Attributions
• Jacqueline K. Barton (California Institute of Technology, Division of Chemistry and Chemical Engineering)
08: MetalNucleic Acid Interactions
Nucleic-acid Structures1
Figure 8.1 displays a single deoxyribonucleotide and the four different nucleic-acid bases. As may be evident, each mononucleotide along a nucleic-acid polymer contains a variety of sites for interactions with metal ions, from electrostatic interactions with the anionic phosphate backbone to soft nucleophilic interactions with the purine heterocycles. The different nucleic-acid bases furthermore offer a range of steric and electronic factors to exploit. Coordination of a metal complex to the N7 nitrogen atom of a purine, for example, would position other coordinated ligands on the metal center for close hydrogen bonding to the O6 oxygen atom of guanine, but would lead to clashes with the amine hydrogen atoms of adenine.
The monomeric units strung together in a polynucleotide furthermore provide an array of polymeric conformers. Figure 8.2A (See color plate section, pages C-14, C-15.) shows three crystallographically characterized structures of double-helical DNA oligonucleotides,2-4 Figure 8.2B a schematic illustration of other conformations of DNA, and Figure 8.2C the crystal structure5 of yeast tRNAPhe. In double-helical DNA,1 the two antiparallel polynucleotide strands are intertwined in a helix, stabilized through Watson-Crick hydrogen bonding between purines and pyrimidines, and through $\pi - \pi$ stacking interactions among the bases arranged in the helical column. There are electrostatic repulsions between the anionic phosphate backbones of the polymer, causing a stiffening; each double-helical step has two formal negative charges. An atmosphere of metal ions condensed along the sugar-phosphate backbone serves partially to neutralize these electrostatic interactions. In the B-DNA conformation, the bases are stacked essentially perpendicular to the helical axis, and the sugars are puckered in general, with a C2'-endo geometry (the C2' carbon is to the same side as the C5' position relative to a plane in the sugar ring defined by the C1', C4', and O atoms). This conformer yields a right-handed helix with two distinct, well-defined grooves, termed the major and minor. The A-form helix, while still right-handed, is distinctly different in structure. The sugar rings are puckered generally in the C3'-endo conformation, causing the bases to be pushed out from the center of the helix toward the minor groove, and tilted relative to the helix perpendicular by almost 20°. What results is a shorter and fatter helix than the B-form; the helical pitch is 28.2 Å in A-DNA for an 11-residue helix and 33.8 Å for a 10-residue helix in B-DNA. The A-form helical shape is best characterized by the very shallow minor groove surface; what was the major groove in the B-form has been pulled deeply into the interior of the A-conformer and is really not accessible to binding by small molecules in solution. Transitions to the A-conformation are promoted by hydrophobic solvents or solutions of high ionic strength. The Z-conformation is perhaps most distinctive, owing to its left-handed helicity.4 The conformer was dubbed Z-DNA because of the zig-zag in the helix. Alternations both in sugar puckering, between C2'-endo and C3'-endo, and in the rotation of the base about the glycosidic bond, anti or syn relative to the sugar, are evident, and lead to a dinucleoside repeating unit versus a mononucleoside repeat in the A- and B-helices. Alternating purinepyrimidine sequences have the highest propensity to undergo transitions into the Z-form. It is actually this syn conformation of purines that leads to the lefthanded helicity of the polymer. But it is not only its left-handedness that distinguishes the Z-conformation. The polymer is long and slender (the pitch is 45 Å for a 12-residue helix), and the major groove is a shallow and wide, almost convex, surface, whereas the minor groove is narrowed into a sharp and small crevice.
These crystal structures, shown in Figure 8.2A (see color plate section, page C-15), in fact each represent a family of conformations. The bases in a base pair often do not lie in the same plane, but are instead propeller-twisted with respect to one another. The local unwinding of the helix and tilting of the base pairs furthermore tend to vary with the local nucleic-acid sequence so as to maximize stacking or hydrogen-bonding interactions among the bases. Hence there is a variety of structures within each conformational family. Our understanding of these structural variations as a function of solution conditions and importantly of local sequence is still quite poor. But surely these structural variations affect and are affected by the binding of metal ions and complexes.
Even less defined structurally are other conformations of DNA, some of which are illustrated schematically in Figure 8.2B (see color plate section, page C-15). Double-helical DNA can bend,6 form loops and cruciforms,7 and fold back on itself into intramolecular triple helices, termed H-DNA.8 At the ends of chromosomes, four strands may even come together in a unique conformation. These structures, characterized thus far by means of biochemical techniques, arise because of sequence and local torsional stress, or supercoiling. Many of these structures are stabilized by the binding of highly charged metal ions, probably because the highly charged metal center in a small volume can neutralize the electrostatic repulsions between polyanionic strands that are bundled together. Metal complexes can furthermore be extremely useful in targeting and characterizing these structures, as we will see. In chromosomes the DNA is packaged by histone proteins into even tighter bundles, with helical segments wrapped about the basic proteins to form superhelical nucleosomal units which are then arranged like beads on a string of more loosely packed DNA.9
This complexity in DNA structure is in fact small compared to that of RNA. Figure 8.2C (see color plate section, page C-15) shows the first crystallographically characterized structure5 of an RNA polymer, yeast tRNAPhe. Ostensibly single-stranded RNAs do not exist as random coils, but instead fold up into well-defined three-dimensional structures, much like proteins. The structural variety, of course, bears some resemblance to that found in DNAs. Double-helical regions in the tRNA are A-like in conformation; helices fold together as one might imagine to occur in cruciforms, and even triple-helical segments are evident where three strands fold together in the polymer. But overall our ability to characterize structures of RNA thus far is lower than that with DNAs. RNAs are less stable in solution than is DNA, and fewer chemical as well as enzymatic tools are available for structural characterization. Yet the recent discovery of ribozymes,10 the finding that RNAs can indeed catalyze nucleolytic reactions, makes our need to understand these structures even greater. Again transition-metal chemistry may participate in stabilizing, promoting, and probing these structures.
Fundamental Interactions with Nucleic Acids
Metal ions and complexes associate with DNA and RNA in a variety of ways, as illustrated in Figure 8.3. Both strong covalent interactions and weak noncovalent complexes are observed.11 Each may yield a significant perturbation in the nucleic acid and/or may be exploited to obtain a site-specific response. Clearly there are some general guidelines, based on principles of coordination chemistry, that may be helpful in sorting out these interactions.
1. Coordination
Most prevalent among covalent complexes with DNA are those involving coordination between soft metal ions and nucleophilic positions on the bases. The structure12 of cis-(NH3)2Pt-dGpG is an example: its platinum center coordinates to the N7 position of the guanine bases. In terms of interactions with the full polynucleotide, it is likely that the cis-diammineplatinum center, with two coordination sites available, would yield an intrastrand crosslink between neighboring guanine residues on a strand (see Chapter 9). Other nucleophilic sites targeted by soft metal ions on the bases include the N7 position of adenine, the N3 position on cytosine, and the deprotonated N3 position on thymine and uracil.12,13 Some additional covalent binding to the N1 positions of the purines has also been observed. Indeed, coordination by the metal to one site on the heterocyclic base lowers the pKa and increases the metal-binding affinity to secondary sites. It is noteworthy, however, that in base-paired double-helical DNA only the N7 positions on the purines are easily accessible in the major groove of the helix. Base binding at the purine N7 position is, of course, not limited to soft metal ions such as Pt(II), Pd(II), and Ru(II). Coordination at these sites has been evident also with first-row transition-metal ions such as Cu(II) and Zn(II).13 For these, as is consistent with basic coordination chemistry, the lability of complexes formed is higher.
Transition-metal ions with decreasing softness are capable of coordinating also to the phosphate oxygen atoms. The ionic versus covalent character of these complexes clearly depends on the metal ions involved. In a classic study, examining the melting temperature of double-helical DNA in the presence of different metal ions and as a function of their concentration, Eichhorn and coworkers established the preference of the metal ions for base versus phosphate binding (Figure 8.4).14 The preference for phosphate over base association was found to decrease in the order Mg(II) > Co(II) > Ni(II) > Mn(II) > Zn(II) > Cd(II) > Cu(II). This series arises from examination of DNA helix-melting temperatures, since base interactions in general should destabilize the helical form [except where interstrand crosslinking occurs, as may happen with Ag(I)], whereas phosphate coordination and neutralization would increase the helix stability and hence the melting temperature.
Also of interest, but less common, are covalent interactions with the sugar moiety.15,16 Although the pentose ring in general provides a poor ligand for metal ions, osmate esters can form quite easily across the C2'-C3' positions in ribose rings. This particular interaction has been suggested as a basis for heavymetal staining of RNA. In fact, OsO4 is not restricted in its reactivity with the sugar positions. Cisoid osmate esters form as well upon reaction of OsO4 across the electron-rich C5-C6 double bonds of accessible pyrimidines on DNA. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/08%3A_MetalNucleic_Acid_Interactions/8.01%3A_The_Basics.txt |
2. Intercalation and Hydrogen Bonding
But important interactions of metal complexes with polynucleotides are not restricted to those involving direct coordination of the metal center to the polymer. Instead, an abundance of highly selective interactions arise from an ensemble of weaker noncovalent interactions between the ligands of coordinatively saturated metal complexes and the nucleic acid. Two primary examples of noncovalent association are given by metallointercalation and hydrogen-bonding interactions of coordinated ligands.17,18 Planar aromatic heterocyclic ligands such as phenanthroline and terpyridine can stack in between the DNA base pairs, stabilized through dipole-dipole interactions. Here, depending on the complex and its extent of overlap with the base pairs, the free energy of stabilization can vary from ~2 to 10 kcal. Nonintercalative hydrophobic interactions of coordinated ligands in the DNA grooves also can occur, as we will see. Hydrogenbonding interactions of coordinated ligands with the polynucleotide are quite common, and arise in particular with the phosphate oxygen atoms on the backbone. With cobalt hexaammine, for example, hydrogen bonding to an oligonucleotide occurs between the ammine hydrogens and both phosphate oxygen atoms and purine bases.19
A mix of covalent and noncovalent interactions is also possible. With cis-diammineplatinum(II) coordinated to the guanine N7 position, the ammine ligands are well-poised for hydrogen-bonding interactions with the phosphate backbone.12 The steric constraints on the molecule must be considered, however. With Pt(terpy)CI+, both intercalation of the terpy ligand and direct coordination of the platinum center (after dissociation of the coordinated chloride) are available, but not simultaneously; coordination of the platinum to the base would likely position the terpyridyl ligand away from the base stack in the DNA major groove, precluding intercalation.20 Sigel and coworkers21 have studied the thermodynamics of noncovalent interactions coupled to direct coordination of simple first-row transition-metal complexes with mononucleotides, and these results illustrate well the interplay of weak noncovalent interactions and direct coordination in generating geometric specificity in complex formation.
Fundamental Reactions with Nucleic Acids
The reactions of transition-metal complexes with polynucleotides generally fall into two categories: (i) those involving a redox reaction of the metal complex that mediates oxidation of the nucleic acid; and (ii) those involving coordination of the metal center to the sugar-phosphate backbone so as to mediate hydrolysis of the polymer. Both redox and hydrolytic reactions of metal complexes with nucleic acids have been exploited with much success in the development of tools for molecular biology.
1. Redox Chemistry
The simplest redox reaction with polynucleotides one might consider as an illustration is the Fenton reaction, which indirectly promotes DNA strand scission through radical reactions on the sugar ring. The reaction with Fe(EDTA)2- is shown in Figure 8.5A. As do other redox-active divalent metal ions, ferrous ion, in the presence of hydrogen peroxide, generates hydroxyl radicals, and in the presence of a reductant such as mercaptoethanol, the hydroxyl radical production can be made catalytic. Although ferrous ion itself does not appear to interact appreciably with a nucleic acid, especially when chelated in an anionic EDTA complex and repelled by the nucleic-acid polyanion, the hydroxyl radicals, produced in appreciable quantities catalytically, attack different sites on the sugar ring, indirectly yielding scission of the sugar-phosphate backbone. One such reaction that has been characterized in some detail is that involving hydroxyl radical reaction at the C4' position, the position most accessible to the diffusible radical in the minor groove of the helix.22 As illustrated in Figure 8.5B, the products of this reaction include a 5'-phosphate, a mixture of 3'-phosphate and phosphoglycolates, and a mixture of free bases and base propenalso Reactions of the hydroxyl radical at other sites on the sugar ring are now being identified as well by isotope-labeling studies. Comparable reactions with RNA have also been described.23
The application of this Fenton chemistry to promote site-specific or sequence-neutral cleavage of DNA was first demonstrated24 by Dervan and coworkers, and has provided the basis for the design of a tremendous range of new and valuable DNA cleavage agents. The development of this chemistry was originally based on modeling Fe-bleomycin, a natural product with antitumor and antibiotic activity, which binds and cleaves DNA.25 The chemistry mediated by Fe-bleomycin, as we will discuss later, is likely to be far more complex, however, involving direct reaction of an intimately bound ferryI intermediate species with the nucleic acid, rather than net oxidation of the sugar mediated by a diffusing hydroxyl radical. Other metal ions such as Cu(Il) can also promote redox-mediated cleavage of DNA26,27 through reactions on the sugar ring; whether the oxidizing radical is still coordinated to the metal or is a dissociated and diffusing species is a topic of much debate.26
Metal ions can also be used to generate other oxidizing intermediates in aerated aqueous solution, such as superoxide ion and singlet oxygen. DNA strand-cleavage reactions mediated by superoxide have not thus far been demonstrated, however. Singlet oxygen may be produced by photosensitization of Ru(phen)32+, and indeed photolysis of Ru(phen)32+ bound to DNA yields oxygen-dependent, alkaline-sensitive strand cleavage.28,29 For singlet oxygen, the oxidation occurs on the nucleic-acid base rather than on the sugar ring. As such, the reaction varies with base composition; guanine residues are most reactive. Furthermore, since the primary lesion is that of a base modification, piperidine treatment, or other weakly basic conditions, are needed to convert the base lesion into a strand-scission event.
Another scheme for oxidative cleavage of DNA mediated by metal complexes involves formation of a coordinated ligand radical bound to the helix that directly abstracts a hydrogen atom from the sugar ring. The photoreaction of Rh(phen)2phi3+ (phi = 9,10-phenanthrenequinone diimine) exemplifies this strategy.30 Here photolysis promotes a ligand-to-metal charge transfer with formation of a phi-centered radical. Isotope-labeling studies and product analysis have shown that this phi radical bound intercalatively in the major groove of DNA directly abstracts the C3'-H (which sits in the major groove of the helix);31 subsequent hydroxylation or dioxygen addition at this position promotes DNA strand scission without base treatment. Some potent photooxidants can also produce outer-sphere electron transfer from the DNA. Here it is the guanine bases, likely those stacked with neighboring purines, that are most easily oxidized and hence most susceptible to attack. Again, this base modification requires alkaline treatment to convert the lesion to a strand breakage.11b,17The DNA double helix can furthermore also mediate electron-transfer reactions between bound metal complexes. The DNA polymer has, for example, been shown to catalyze photoinduced electron-transfer reactions between Ru(phen)32+ and Co(phen)33+ bound along the DNA strand.32 Table 8.1 summarizes different redox reactions of metal complexes bound to DNA.
Table 8.1 - Examples of metal complexes that cleave DNA through redox chemistry.
a) DNA may be modified by attack either at the sugar or at the nucleotide base position.
b) The reactive species involved in DNA cleavage, if known.
c) Some reactive species are diffusible, producing broad patterns of DNA damage along the strand. Others are nondiffusible, resulting in cuts at single discrete sites.
d) The site of metal complex binding to DNA, if known.
e) The sites cleaved by the metal complex.
f) Not known.
* Indicates an excited-state reaction requiring photoactivation.
Complex Targeta Chemistryb Diffusibilityc DNA Bindingd Site Selectivitye
Fe(EDTA)2- sugar OH•, Fenton diffusible none none
MPE-Fe(II) sugar (C4'-H) OH•, Fenton diffusible sequence-neutral none
Co(NH3)63+* base phtoelectron transfer f hydrogen-bonding 5'-G-pur-3'
Cu(phen)2* sugar Cu2+-OH• slight AT-rich AT-rich
Mn-Porphyrin sugar M=O none AT-rich At-rich
U(O2)(NO3)2* f f diffusible f none
Ru(TMP)32+* base 1O2 diffusible A-form A-form, G
Ru(phen)32+* base 1O2 diffusible sequence-neutral G
Co(DIP)33+* sugar ligand radical none Z-form (non-B) Z-form (non-B)
Rh(DIP)33+* sugar ligand radical none Z, cruciforms Z, cruciforms
Rh(phen)2phi3+* sugar (C3'-H) ligand radical none open major groove 5'-pyr-pyr-pur-3'
Rh(phi)2bpy3+* sugar (C3'-H) ligand radical none sequence-neutral none
2. Hydrolytic Chemistry
Hydrolysis reactions of nucleic acids mediated by metal ions are important elements in natural enzymatic reactions; chemists would like to exploit them in the design of artificial restriction endonucleases.33 Hydrolysis reactions of the phosphodiester linkage of polynucleotides appear preferable to redox-mediated cleavage reactions, since in the hydrolytic reaction all information is preserved. In redox cleavage by sugar oxidation, for example, both a sugar fragment and free nucleic-acid base are released from the polymer, and, in contrast to hydrolytic chemistry, the direct religation of the fragments becomes practically impossible.
Metal ions can be effective in promoting hydrolysis of the phosphodiester, since they can function as Lewis acids, polarizing the phosphorus-oxygen bond to facilitate bond breakage, and can also deliver the coordinated nucleophile to form the pentacoordinate phosphate intermediate. Figure 8.6 illustrates one crystallographically characterized model system developed by Sargeson and coworkers, where hydrolysis of a model phosphodiester was enhanced dramatically by taking advantage of both the acidic and the nucleophilic characteristics of the bound cobalt(III) species.34 A whole series of model systems utilizing both cobalt and zinc ions has been designed to explore the hydrolytic reactions of simple phosphodiesters.35 This strategy coupled to a DNA binding functionality has also been exploited, albeit inefficiently, in the hydrolytic cleavage of doublehelical DNA by Ru(DIP)2Macro with Zn2+, Cd2+, or Pb2+ added in situ.36 In this complex (see Figure 8.6), the central portion of the molecule, held together by the ruthenium(II), is responsible for DNA binding. Tethered onto the coordinatively saturated ruthenium complex are two diethylenetriamine functionalities (in the Macro ligand), however, and these serve to coordinate hydrolytically active metal ions such as Zn(II) and Co(II), which promote DNA hydrolysis once delivered to the sugar-phosphate backbone by the DNA-binding domain.
Perhaps simpler and certainly better understood are the hydrolytic reactions of RNAs mediated by metal ions. More than twenty years ago Eichhorn and coworkers showed that simple metal ions such as Zn(II) and Pb(II) promote the hydrolysis of RNA.37 Figure 8.6 illustrates also the crystallographically characterized site-specific hydrolyis in tRNA by plumbous ion.38 In tRNA, Pb(II) occupies three quite specific high-affinity binding sites, and at one of these sites, the metal ion becomes poised to promote strand cleavage. The crystal structure with bound Pb2+ suggests that the lead-coordinated hydroxide ion deprotonates the 2'-hydroxyl of one residue, so that the resulting 2'-oxygen nucleophile may attack the phosphate to give a pentavalent intermediate that decays to form the 2',3'-cyclic phosphate and, after reprotonation, the 5'-hydroxide. This very specific cleavage reaction is already being used by biologists as a tool in probing structures of mutant tRNAs, since the reaction is exquisitely sensitive to the stereochemical alignment of the nucleic-acid residues, phosphate backbone, and associated metal ion. In hydrolytic reactions on RNA, it is commonly considered, though certainly not established, that the job of the metal ion may be simpler than with DNA, since the ribose provides a nearby nucleophile already in the 2'-hydroxide. The reaction of tRNA with Pb(II) nonetheless illustrates how a metal ion may be utilized in promoting highly specific chemistry on a nucleic-acid polymer.
Last, it must be mentioned that metal coordination to the purine N7 position can also indirectly promote strand cleavage, although not through direct hydrolytic reaction on the sugar-phosphate backbone. Metal ions such as Pd2+ and Cu2+, through coordination at N7, promote depurination. The depurinated site then becomes easily susceptible to hydrolysis upon treatment with mild base. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/08%3A_MetalNucleic_Acid_Interactions/8.02%3A_The_Basics_%28Part_2%29.txt |
Now we may examine in detail the interaction of one class of metal complexes with nucleic acids, how these complexes bind to polynucleotides, the techniques used to explore these binding interactions, and various applications of the complexes to probe biological structure and function. Tris(phenanthroline) metal complexes represent quite simple, well-defined examples of coordination complexes that associate with nucleic acids. Their examination should offer a useful illustration of the range of binding modes, reactivity, techniques for study, and applications that are currently being exploited and explored. In addition, we may contrast these interactions with those of other transition-metal complexes, both derivatives of the tris(phenanthroline) family and also some complexes that differ substantially in structure or reactivity.
Binding Interactions with DNA
Tris(phenanthroline) complexes of ruthenium(II), cobalt(III), and rhodium(III) are octahedral, substitutionally inert complexes, and as a result of this coordinative saturation the complexes bind to double-helical DNA through a mixture of noncovalent interactions. Tris(phenanthroline) metal complexes bind to the double helix both by intercalation in the major groove and through hydrophobic association in the minor groove.11b,40 Intercalation and minor groove-binding are, in fact, the two most common modes of noncovalent association of small molecules with nucleic acids. In addition, as with other small molecules, a nonspecific electrostatic interaction between the cationic complexes and the DNA polyanion serves to stabilize association. Overall binding of the tris(phenanthroline) complexes to DNA is moderate (log K = 4).41
The extent of intercalative versus groove binding is seen to depend upon environmental conditions, such as temperature and ionic strength, the charge of the metal center, and the DNA base sequence; groove binding is favored at AT-rich sequences.41 Second-generation mixed-ligand derivatives of the tris(phenanthroline) series have been prepared, and their interactions with DNA have provided useful insight into the factors important for promoting either intercalation or groove binding.42 Aromatic heterocyclic ligands with increased surface areas that are planar bind DNA with increasing avidity through intercalation, irrespective of the charge on the metal center. Intercalative binding constants greater than 107 M-1 can be easily achieved with planar heterocyclic ligands that jut out from the metal center. Not surprisingly, complexes containing ligands of increasing hydrophobicity that are not planar favor minor-groove binding.28
Critically important as well in determining the binding mode is the chirality of the metal complex.40 Intercalation into the right-handed helix favors the $\Delta$-isomer, whereas groove binding favors the $\Lambda$-isomer. Figure 8.7 illustrates these symmetry-selective interactions. In intercalation, we consider that one phenanthroline inserts and stacks in between the base pairs, essentially perpendicular to the helix axis. For the $\Delta$-isomer, once intercalated, the ancillary non-intercalated ligands are aligned along the right-handed groove of the helix. For the $\Lambda$-isomer, in contrast, with one ligand intercalated, the ancillary ligands are aligned in opposition to the right-handed groove, and steric interactions become evident between the phenanthroline hydrogen atoms and the phosphate oxygen atoms. Increasing the steric bulk on these phenanthrolines furthermore increases the enantioselective preference for intercalation of the $\Delta$-isomer.40,43 For intercalation, then, the chiral discrimination depends on matching the symmetry of the metal complex to that of the DNA helix. For groove binding, where the metal complex is thought to bind against the helix, instead it is a complementary symmetry that is required. In our model for groove binding of the tris(phenanthroline) metal complex, two phenanthroline ligands are likely bound against the right-handed helical groove, stabilized through hydrophobic association. For the $\Lambda$-isomer, bound in this fashion, the ligands lie against and complement the right-handed groove; with the $\Delta$-isomer, the ligands oppose the groove, and no close surface contacts are made.
Intercalation of metal complexes in DNA is not uncommon. Lippard and coworkers first established metallointercalation by Pt(II) complexes in the 1970s.18,20 Square-planar platinum(II) complexes containing the terpyridyl ligand were shown to intercalate into DNA. In an elegant series of x-ray diffraction experiments on DNA fibers, Lippard illustrated the requirement for planarity in the complex.18,44 Although (phen)Pt(en)2+ and (bpy)pt(en)2+ were shown to intercalate into the helix, (pyr)2Pt(en)2+, with pyridine ligands rotated out of the coordination plane, could not. Complex planarity is in itself insufficient to promote intercalation, however. Cis-(NH3)2PtCI2 or even cis-(NH3)Pt(en)2+ does not appear to intercalate into a helix, despite full planarity. Instead, aromatic heterocyclic ligands must be included in order to promote dipole-dipole interactions with the heterocyclic bases stacked in the helix. Indeed, planarity of the full complex is not required. Intercalation is not restricted to coordination complexes that are square planar. The tris(phenanthroline) complexes represented the first examples of "three-dimensional intercalators" and illustrated that octahedral metal complexes could also intercalate into the helix.40,45,46 Here one can consider the partial intercalation of one ligand into the helix, providing the remaining ligands on the complex an opportunity to enhance specificity or reactivity at a given site.
Curiously, one unique and apparently general characteristic of metallointercalators is their preference for intercalation from the major groove of the helix. Most small molecules associate with DNA from the minor groove, but metallointercalators, both those that are square planar, such as (terpyridyl)platinum(II) complexes, and those that are octahedral, such as the tris(phenanthroline) metal complexes, appear to intercalate into the major groove. This then mimics quite well the association of much larger DNA-binding proteins with the helix; DNA regulatory proteins generally appear to target the major groove. The reason why metallointercalators favor major groove association is still unclear.
Transition-metal complexes with aromatic ligands also generally associate by minor-groove binding or through the mix of intercalative and groove-bound interactions. Cu(phen)2+, a tetrahedral complex, appears to favor minor-groove binding over intercalation.26 Perhaps the tetrahedral coordination does not permit appreciable overlap of the phenanthroline ring with the bases in an intercalative mode. Metalloporphyrins, despite their large expanse and the presence commonly of nonplanar substituents, appear to bind to double-helical DNA both by intercalation and by minor-groove binding at AT-rich sequences.47 Occupation of the porphyrins by transition-metal ions, such as Cu(II), which bind axial ligands, leads to the favoring of groove binding over intercalation. Figure 8.8 illustrates some of the complexes that bind DNA noncovalently.
The tris(phenanthroline) metal complexes themselves do not offer an illustration of hydrogen-bonding interactions with the helix, since these ligands lack hydrogen-bonding donors and acceptors, but as mentioned already, hydrogen bonding of coordinated ligands to the helix can add some measure of stabilization, comparable to, but likely no greater in magnitude than, that provided by intercalative stacking, hydrophobic, or dispersive interactions. Indeed, mixed-ligand derivatives of the phenanthroline complexes have been prepared that include hydrogen-bonding groups (amides, hydroxyls, and nitro substituents) on the ancillary phenanthroline ligands, and these have shown no greater avidity for double-helical DNA than their counterparts with hydrophobic substituents.42 A large number of weak hydrogen-bonding interactions to DNA by one complex can be stabilizing, however, as with, for example, hexaamminecobalt(III) or hexaaquoterbium(III).
Tris(phenanthroline) metal complexes also do not offer an opportunity to explore covalent binding interactions with the helix in greater detail, but these interactions are, in fact, a major focus of Chapter 9, concerned with the mode of action of cisplatin. One derivative of the tris(phenanthroline) series, Ru(phen)2Cl2, has been shown to bind to DNA covalently.48 In aqueous solution the dichlororuthenium(II) complex undergoes hydrolysis to form an equilibrium mixture of bis(phenanthroline) diaquo and chloroaquo species. These species bind covalently to DNA, with preferential reactivity at guanine sites. It is interesting that the same structural deformations in the DNA evident upon binding cis-diammineplatinum units become apparent upon coordination of bis(phenanthroline)ruthenium(II). It is also noteworthy that the chiral preference in coordination is for the $\Lambda$-isomer. As with groove binding, direct coordination to base positions requires a complementary symmetry, with the the $\Lambda$-isomer binding against the right-handed groove. This preference for the $\Lambda$-isomer reaffirms that, rather than noncovalent intercalation (which would favor the $\Delta$-isomer), covalent binding dominates the interaction. The energetic stabilization in direct coordination of the ruthenium(II) center is certainly more substantial than the weaker stabilization derived from intercalation. Rh(phen)2Cl2+ and its derivatives have also been shown to bind covalently to DNA but only upon photoactivation, since light is needed to promote dissociation of the coordinated chloride and substitution of the nucleic acid base as a ligand.49
Techniques to Monitor Binding
Many of the same techniques employed in studying the basic chemistry of coordination complexes can be be used in following the binding of transition-metal complexes to nucleic acids, but biochemical methods, with their often exquisite sensitivity, become valuable aids as well in delineating specific binding interactions. Tris(phenanthroline) metal complexes are particularly useful to illustrate this point, since here the metal center in the complex is selected in terms of the technique used for examination.
Coordination complexes are often visibly colored, and these colorations provide a useful and sensitive spectroscopic handle in following fundamental reactions. This notion holds as well with tris(phenanthroline) metal complexes in their interactions with nucleic acids. Ru(phen)32+ and its derivatives are highly colored because of an intense metal-to-ligand charge-transfer band ($\lambda_{max}$ = 447 nm, $\epsilon$ = 1.9 x 104 M-1cm-1). Furthermore, the complexes are highly photoluminescent ($\lambda_{cm}$ = 610 nm, $\tau$ = 0.6 $\mu$s in aerated aqueous solution). On binding to nucleic acids these transitions are perturbed. Hypochromism is observed in the charge-transfer band, and intercalation leads to an increase in lifetime of the charge-transfer excited state.43,46 Indeed, single-photon counting experiments show a biexponential decay in emission from Ru(phen)32+ bound to double-helical DNA. The longer-lived component ($\tau$ = 2 $\mu$s) has been assigned as the intercalated component and the shorter-lived 0.6 $\mu$s component has been attributed to a mixture of free and groove-bound species. These spectroscopic perturbations permit one to define equilibrium-binding affinities for the different components of the interaction as a function of metal-center chirality and under different solution conditions.41 One can also follow the polarization of emitted light from the complexes after excitation with polarized light, and these studies have been helpful in describing the dynamics of association of the complexes on the helix.41,43 Mixed-ligand complexes of ruthenium(II) show similar spectroscopic perturbations, and these have been used to characterize binding affinities and chiral preferences, as well as the extent of intercalation versus groove binding as a function of ligand substitution on the metal center.42 The spectroscopic handle of the metal center therefore affords a range of experiments to monitor and characterize the binding of the metal complexes to polynucleotides.
Binding interactions of metal complexes with oligonucleotides can also be followed by NMR, and here as well the metal center offers some useful characteristics to exploit. As with organic DNA-binding molecules, shifts in the 1H-NMR resonances of both the DNA-binding molecule and the oligonucleotide become apparent as a function of increased association with the helix. These shift variations can be used empirically to watch the dynamics of association and to gain some structural insights into the binding modes of the complexes on the helix. These kinds of experiments have been performed with tris(phenanthroline) complexes of ruthenium(II) and rhodium(III), where it was observed that the double-helical oligonucleotide is an exceedingly good chiralshift reagent to separate resonances in an enantiomeric mixture of the tris(phenanthroline) complexes.50 For covalent binding molecules, such as cis-diammineplatinum(II), furthermore, the lowering of the pKa of purine positions and therefore shifting of resonances as a function of coordination to an alternate site has been helpful as well in assigning the sites of covalent binding on the oligonucleotide.51 But also in NMR experiments, special advantage can be taken of the metal center. For tris(phenanthroline) metal complexes, 1H-NMR experiments52 were performed on the paramagnetic analogues, Ni(phen)32+ and Cr(phen)33+. It was reasonable to assume the binding characteristics would be identical with their respective diamagnetic analogues, Ru(II) and Rh(III); yet paramagnetic broadening by the metal complexes of nearby resonances on the oligonucleotide would allow one to deduce where along the helix the complexes associate. Using this method the groove-binding interaction of the complexes was identified as occurring in the minor groove of the helix. Figure 8.9 illustrates the monitoring of DNA binding by tris(phenanthroline) metal complexes using both the luminescence characteristics of ruthenium(II) complexes and the paramagnetic characteristics of nickel(II).
There are numerous other classic techniques of inorganic chemistry that have been or could be applied in studying the binding of metal complexes to nucleic acids. Coordination complexes have invariably been used in x-ray diffraction experiments because of the high electron density of the metal center. The tris(phenanthroline) metal complexes have not yet been applied in this context, but, as mentioned already, platinum metallointercalators were examined by fiber diffraction to delineate intercalation requirements. In fact, many nucleic-acid crystal structures have required specific metal ion additions for isomorphous heavy-metal derivatives to solve the structure. Such has certainly been true for the crystal structure of tRNAPhe, where heavy-metal ions such as platinum, osmium, and mercury were targeted to specific base positions, and lanthanide ions were used to label phosphate positions around the periphery of the molecule.53 Other techniques can also be exploited to monitor and characterize binding. A recent novel illustration is one from electrochemistry, which has been applied in monitoring the binding of Co(phen)33+ to DNA.54 Surely other techniques, from EXAFS to scanning tunneling microscopy, will be exploited in the future.
Biochemistry also provides very sensitive techniques that have been invaluable in characterizing interactions of metal complexes with nucleic acids. First are simply gel electrophoresis experiments, which permit an assessment of changes in the nucleic-acid conformation, through its changes in gel mobility, as a function of metal binding. A classic illustration is that of the unwinding of superhelical DNA as a function of intercalation. Closed circular DNA has much the same topological constraints on it as does a rope or a telephone cord; the DNA helices can wind up in coils. We define the duplex turning in a double helix as the secondary helical turns, and turns of the helices about one another as the supercoils or tertiary turns. As long as a DNA double helix is closed in a circle (form I), the total winding, that is, the total number of secondary and tertiary turns, is fixed. Molecules with differing extents of winding have different superhelical densities. In a circular molecule with one strand scission, what we call form II (nicked) DNA, the topological constraints are relaxed, and no supercoils are apparent. The same, by analogy, can be said of a telephone cord off the phone receiver, which can turn about itself to relax its many supercoils. Now let us consider a DNA unwinding experiment, monitored by gel electrophoresis. Supercoiled form I DNA can be distinguished from nicked DNA (form II) in an agarose gel because of their differing mobilities; the wound-up supercoiled molecule moves easily through the gelatinous matrix to the positive pole, whereas the nicked species is more floppy and thus is inhibited in its travels down the gel. A closed circular molecule with no net supercoils (form I0) comigrates with the nicked species. Consider now the supercoiled molecule in the presence of an intercalator. Since the intercalator unwinds the DNA base pairs, the number of secondary helical turns in the DNA is reduced. In a negatively supercoiled, closed circular DNA molecule, the number of supercoils must be increased in a compensatory fashion (the total winding is fixed); hence the total number of negative supercoils is reduced, and the molecule runs with slower mobility through the gel. As the intercalator concentration is increased still further, the mobility of the supercoiled species decreases until no supercoils are left, and the species comigrates with the nicked form II DNA. Increasing the bound intercalator concentration still further leads to the positive supercoiling of the DNA and an increase in mobility. Figure 8.10 illustrates the experiment with tris(phenanthroline)ruthenium(II) isomers.46 This kind of unwinding experiment is an example of the sensitivity with which DNA structural changes can be monitored using biochemical methods; only low quantities (< $\mu$g) of materials are needed to observe these effects.
DNA strand scission can also be sensitively monitored, and even more importantly the specific nucleotide position cleaved can be pinpointed by biochemical methods. This methodology has been applied successfully in monitoring both the efficiency of DNA strand scission by metal complexes and the specific sites cleaved, and hence where the complexes are specifically bound on the helical strand.
Relative extents of cleavage of DNA by different metal complexes can be easily assayed in an experiment that is an extension of the unwinding experiment described above. One simply measures the conversion of supercoiled form I DNA to nicked form II species. One strand cleavage on the DNA circle releases the topological constraints on the circular molecule and relaxes the supercoils. Two cleavage events within 12 base pairs on opposite strands will convert the DNA to a linear form (III), which also has a distinguishable gel mobility. Photoactivated cleavage of DNA by tris(phenanthroline) complexes of cobalt(lIl) and rhodium(III) was first demonstrated using this assay.55,56 Given the high sensitivity of this assay, redox-mediated cleavage of DNA by a wide range of metal complexes can be easily demonstrated. However, other techniques are required to analyze whether appreciable and significant cleavage results, and, if so, what products are obtained. Since the assay can monitor, in a short time using little sample, a single nick in a full 4,000-base-pair plasmid, reactions of very low, almost insignificant yield can be detected. The assay provides, however, a simple scheme to assess relative extents of cleavage by different metal complexes, as well as a first indication that a cleavage reaction by a given metal complex occurs at all.
More informative are experiments on 32P-labeled DNA fragments using high-resolution polyacrylamide gel electrophoresis, since these experiments allow one to find the exact nucleotide where the complexes break the sugar-phosphate backbone. Consider a cleavage reaction by a given metal complex on a DNA fragment of 100 base pairs in length that has been labeled enzymatically with 32P on one end of one strand. If the metal complex cleaves the DNA at several different sites, then one can arrive at conditions where full cleavage is not obtained, but instead a population of molecules is generated where single cleavage events per strand are obtained, and cleavage at each of the sites is represented. After denaturation of the fragment, electrophoresis through a high-density polyacrylamide gel, and autoradiography, only fragments that are radioactively end-labeled are detected, and hence the population of sites cleaved is determined. The denatured cleaved fragments move through the gel according to their molecular weight. By measuring their length, using molecular-weight markers, one can find the specific position cleaved relative to the end of the fragment. By this route the specific sites cleaved by a molecule that binds and cleaves DNA, or end-labeled RNA, at unique positions may be identified. In a complementary experiment, using footprinting, where a molecule cleaves DNA nonspecifically at all sites along a fragment, one can find the binding positions of other molecules such as DNA-binding proteins. In this experiment, cleavage with the sequence-neutral cleaving reagent is carried out both in the presence and in the absence of the other binding molecule. In the absence of the protein, cleavage ideally occurs at all sites; hence a ladder of cleaved fragments is apparent on the autoradiograph. If cleavage is carried out in the presence of the protein, however, those sites that are bound by the protein are protected from cleavage by steric considerations, producing a shadow or footprint of the protein-binding site on the gel. Both the site-specific and footprinting experiments are illustrated schematically in Figure 8.11. This very powerful technology was first applied by Dervan and coworkers in demonstrating the application of methidium-propyl-FeEDTA (MPE-Fe) as a chemical footprinting reagent.22,57 Tris(phenanthroline) metal complexes have been shown to cleave DNA nonspecifically, and their derivatives have been applied either as sensitive photofootprinting reagents, or as site-specific cleaving molecules, as we will see. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/08%3A_MetalNucleic_Acid_Interactions/8.03%3A_A_Case_Study-_Tris%28phenanthroline%29_Metal_Complexes.txt |
Both the spectroscopy and the chemical reactivity of transition-metal complexes, coupled to biochemical assays, can therefore be exploited to obtain a wide range of useful reagents to probe nucleic acids. Here some specific applications are described.
Spectroscopic Probes
As discussed previously, the tris(phenanthroline)ruthenium(II) complexes offer a novel spectroscopic probe of nucleic acids, since their luminescence is increased upon intercalation into the double helix. As a result the complexes provide a simple luminescent stain for DNA in fluorescent microscopy experiments. More interesting, perhaps, is the conformational selectivity of derivatives of tris(phenanthroline)ruthenium. Ru(DIP)32+ (DIP = 4,7-diphenyl-1,10-phenanthroline) shows enantiospecificity in binding to B-form DNA.40 Because of the steric bulk of the phenyl rings, detectable binding is seen only with the \(\Delta\)-isomer in a righthanded helix; no binding is evident with the \(\Lambda\)-isomer. But with the left-handed Z-form helix, both isomers bind avidly.40,58 The shallow left-handed major groove can accomodate the two enantiomers., A left-handed but more B-like helix shows selectivity instead for the \(\Lambda\)-isomer. Spectroscopic experiments that measure the chiral selectivity of Ru(DIP)32+ isomers in binding to a given DNA then provide a novel probe for helical handedness. Indeed, \(\Lambda\)-Ru(DIP)32+ was the first spectroscopic probe for Z-DNA (or other alternate conformations that are sufficiently unwound to permit binding by the bulky left-handed isomer).58
Another set of derivatives of the tris(phenanthroline) metal complexes that may become exceedingly useful as spectroscopic probes are Ru(bpy)2dppz2+ and Ru(phen)2dppz2+ (dppz = dipyridophenazine).59 In these complexes the metal-to-ligand charge transfer is preferentially to the electron-accepting dppz ligand. In nonaqueous solutions, the complexes luminesce. However, in aqueous solution at pH 7, no luminescence is observed, likely because hydrogen bonding by water to the phenazine nitrogen atoms quenches the charge-transfer excited state. But the dppz ligand is also an expansive, aromatic heterocyclic ligand, and as a result both Ru(bpy)2dppz2+ and Ru(phen)2dppz2+ bind avidly to DNA by intercalation. Once intercalated, the phenazine ligand is protected from water. Therefore these complexes are luminescent when intercalated into DNA, whereas no luminescence is apparent from the complexes in the absence of DNA in aqueous solution. The enhancement factor is > 104 with DNA. One might consider the ruthenium complexes as true "molecular light switches" for DNA.
Both simpler bipyridyl and phenanthroline derivatives as well as dppz complexes of ruthenium are currently being tethered onto other DNA binding moieties, in particular onto oligonucleotides, so as to develop new, nonradioactive luminescent probes for DNA sequences. These transition-metal complexes may provide the basis for the development of new families of DNA diagnostic agents, and many industrial laboratories are currently exploring routes to accomplish these goals. Figure 8.12 illustrates \(\Lambda\)-Ru(DIP)32+ and Ru(bpy)2dppz2+, two complexes whose luminescence properties can be employed to probe nucleic acids.
Other transition-metal complexes besides those of ruthenium have shown some promise in spectroscopic applications with nucleic acids. Lanthanide ions have been applied both in NMR experiments and in luminescence experiments to probe tRNAs, and more recently with synthetic DNAs of differing sequence and structure.60 Lanthanide ions have been exceedingly useful in probing Ca2+ binding sites in proteins, and one would hope that a parallel utility would be achieved with nucleic acids. Their poor absorptivity have made luminescent experiments difficult, however, requiring relatively high concentrations of material. Nonetheless, the sensitivity of luminescent lifetimes to coordination and indeed solvation is providing a novel spectroscopic handle to explore binding sites and structures of the macromolecules. Another quite novel luminescent handle has been phenanthroline and diphenylphenanthroline complexes of copper(I).61 These complexes are extremely valuable cleavage probes, as we will see later; to characterize better their interactions with the helix, luminescence experiments are being explored. A problem here has been the nonphysiological conditions necessary to achieve detectable luminescence. Nonetheless, studies with the copper complexes demonstrate how the whole range of transition-metal chemistry and spectroscopy is beginning to be applied in sorting through nucleic-acid interactions.
Metallofootprinting Reagents
Probably the most widespread application of metal nucleic-acid chemistry in the biology community has been the utilization of metal complexes for chemical footprinting. The footprinting technique (Figure 8.11) was developed by biologists62 as a means of locating protein-binding sites on DNA. 32P-end-labeled double-stranded DNA fragments could be digested with a nuclease, such as DNAse, in the presence or absence of DNA-binding protein. After electrophoresis of the denatured digests and autoradiography, one would find a "footprint," that is, the inhibition of cleavage by DNAse, at the spot bound by protein, in comparison to a randomly cleaved pattern found on the DNA in the absence of binding protein. Although DNAse is still widely used, this footprinting reagent has some disadvantages: (i) the nuclease is not sequence-neutral in its cleavage, resulting in lots of noise in the footprinting background; and (ii) since the nuclease is itself a large protein, its ability to provide high-resolution footprinting patterns of smaller molecules is quite limited.
Several metal complexes now serve as high-resolution, sequence-neutral chemical footprinting reagents. Some of these reagents are shown in Figure 8.13. The first, as mentioned previously, was MPE-Fe(II).57 The complex cotains a sequence-neutral DNA binding moiety, the intercalator methidium, and a tethered DNA redox cleaving moiety, Fe(EDTA). The methidium, in binding nonspecifically to DNA, delivers the hydroxyl radicals, generated via Fenton chemistry at the Fe(II) center in the presence of peroxide and a reducing agent, to the DNA backbone in a random manner. Since the complex is small, high resolution can be achieved. Indeed, MPE-Fe(II) has been shown to footprint small natural products that bind to DNA, in addition to footprinting much larger DNA-binding peptides and proteins.
Perhaps simpler still and now very widely used as a footprinting reagent is Fe(EDTA)2- itself.63 The concept here is that Fe(EDTA)2- , as a dianion, is unlikely to associate at all with the DNA polyanion. Hence hydroxyl radicals, generated via Fenton chemistry at a distance from the helix, would likely diffuse to the helix with a uniform concentration along the helix and provide a completely sequence-neutral pattern of cleavage. Tullius and coworkers have demonstrated63 this to be the case. The resolution is furthermore extremely high since the hydroxyl radical is sufficiently small that it can even diffuse within the DNA-binding protein to delineate binding domains. Some difficulties are found, however, with the high concentrations of activating reagents needed to activate a cleavage reagent that does not bind to the helix, and problems of course arise in trying to footprint metalloproteins. Nonetheless, Fe(EDTA)2-, a reagent easily found on the biologist's shelf, is now finding great utility in labs as a chemical substitute for DNAse.
Other transition-metal complexes are also finding applications in chemical footprinting. Both Cu(phen)2+ and manganese porphyrins have been used to footprint DNA-binding proteins.64,65 These complexes likely cleave DNA through either Fenton chemistry or direct reaction of a coordinated metal-oxo intermediate with the sugar-phosphate backbone. The complexes, however, appear to bind DNA predominantly along the surface of the DNA minor groove, and with some preference for AT-rich regions. The patterns obtained are actually quite similar to those found with DNAse, and thus the lack of high sequence neutrality is somewhat limiting. Furthermore, the complexes are most sensitive to binding moieties in the minor groove, rather than those in the major groove, where proteins bind. Intercalators such as MPE-Fe(II) can sense binding species in both grooves. Cu(phen)2+ has nonetheless proved to be quite effective in detecting hyperreactivities in the minor groove, owing to DNA structural perturbations that arise from protein binding in the major groove. Whether this sensitivity emanates from the intimate interaction of Cu(phen)2+ in the minor groove of the helix, or because of the characteristics of the reactive radical formed, is not known.
Inorganic photochemistry has also been applied in developing metal complexes as photofootprinting reagents. Uranyl acetate, for example, at high concentrations, upon photolysis, promotes DNA cleavage.66 It is thought that the ions interact with the phosphates, generating some excited-state radical chemistry, although no detailed characterization of this chemistry has been undertaken. The cleavage reaction is nonetheless remarkably sequence-neutral and therefore shows some promise for photofootprinting applications. In fact, the applicability of uranyl acetate typifies how simple coordination chemistry and now even photochemistry may be helpful in the design of a variety of reagents that interact and cleave DNA, both nonspecifically and specifically. The biochemical techniques used to monitor such processes are sufficiently sensitive that even quite inefficient reactions in solution can be harnessed in developing useful reagents. The better our understanding of the chemistry of the coordination complex, the more effectively it may be utilized.
The best derivative of a tris(phenanthroline) metal complex currently being applied in footprinting experiments is Rh(phi)2bpy3+, a second-generation derivative of the tris(phenanthroline) series67 that binds DNA avidly by intercalation and in the presence of light promotes direct strand cleavage by hydrogenatom abstraction at the C3'-position on the sugar.31 Since no diffusing intermediate is involved in this photocleavage reaction, the resolution of the footprinting pattern is to a single nucleotide. Here the excited-state transition-metal chemistry involves a ligand-to-metal charge transfer, producing a phi cation radical that directly abstracts the hydrogen from the sugar at the intercalated site. The high efficiency of this photoreaction and high sequence-neutral binding of the complex to double-stranded DNA add to the utility of this reagent in footprinting studies. Indeed, both DNA-binding proteins, bound in the major groove, and small natural products, associated with the minor groove, have been footprinted with Rh(phi)2bpy3+ to precisely that size expected based upon crystallographic results. One may hope that this and other photofootprinting reagents will soon find applications for footprinting experiments in vivo. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/08%3A_MetalNucleic_Acid_Interactions/8.04%3A_Applications_of_Different_Metal_Complexes_that_Bind_Nucleic_Acids.txt |
Conformational Probes
Metal complexes are also finding wide application in probing the local variations in conformation that arise along nucleic-acid polymers. X-ray crystallography has been critical in establishing the basic conformational families of double-helical DNA, and to some extent how conformations might vary as a function of nucleic-acid sequence. Yet many conformations have still not been described to high resolution, and only a few oligonucleotides have been crystallized. Other techniques are therefore required to bridge the small set of oligonucleotide crystal structures that point to plausible structures and the large array of structures that arise as a function of sequence on long helical polymers. Furthermore, only a very small number of RNA polymers has been characterized crystallographically; hence other chemical methods have been needed to describe the folding patterns in these important biopolymers. Metal complexes, mainly through specific noncovalent interactions, appear to be uniquely useful in probing the structural variations in nucleic acids.
1. Nonspecific Reactions of Transition-metal Complexes
Hydroxyl radical cleavage with Fe(EDTA)2- illustrates again how simple metal complexes can be used in characterizing nucleic acids. One example involves efforts to describe the local structural variations in "bent" DNA. Biochemists had found that DNA fragments containing runs of adenines, such as in the tract dAAAAAA, possessed unusual gel-electrophoretic mobilities. Indeed, kinetoplast DNA isolated from mitochondria of trypanosomes showed a remarkable lacework pattern of structure, with loops and circles of DNA; these structures were found to be governed by the placement of these d(A)6 tracts. By constructing a series of oligonucleotides with adenine runs positioned either in or out of phase relative to one another, researchers found that the adenine tracts caused a local bending of the DNA toward the minor groove.6 But what were the detailed characteristics of these bent sites? Using hydroxyl radical cleavage of DNA, generated with Fe(EDTA)2-, Tullius and coworkers found a distinctive pattern of cleavage across the adenine tracts, consistent with a locally perturbed structure.68Here the notion again was that Fe(EDTA)2- in the presence of peroxide would generate hydroxyl radicals at a distance from the helix, and thus careful densitometric analysis of the cleavage across 32P-end-Iabeled DNA fragments would reveal any differential accessibility of sugar residues to cleavage mediated by the radicals caused by the bending. The cleavage patterns suggested a smooth bending of the DNA across the tract and indicated furthermore an asymmetry in structure from the 5'- to 3'-end of the adenine run.
The reactivities of other transition-metal reagents have also been used advantageously in probing nucleic-acid structures. As described in Section II, OsO4 reacts across the 5,6 position of accessible pyrimidines to form a cis-osmate ester. Upon treatment with piperidine, this base modification can yield scission of the sugar-phosphate backbone. Hence DNAs containing unusual local conformations with prominent solvent-accessible pyrimidines can be probed with OsO4. The junction regions of Z-DNA, the single-stranded loops in cruciform structures, and a segment of the dangling third strand in H-DNA, have all been probed by means of the differential reactivity of osmium tetroxide with DNA sites dependent upon their accessibility.7,8,16,69 Surely other transition-metal chemistry will become similarly applicable.
2. Transition-metal Complexes as Shape-selective Probes
Transition-metal complexes have also been designed with three-dimensional structures that target complementary structures along the helical polymer. This recognition of DNA sites, based upon shape selection, has proved to be extremely useful both in demarcating and in characterizing structural variations along the polymer and in developing an understanding of those factors important to the recognition of specific polynucleotide sites. Complexes, basically derivatives of the tris(phenanthroline) metal series, have been designed that specifically target A- and Z-form helices, cruciforms, and even subtle variations such as differential propeller twisting within B-form DNA.11c By appropriate substitution of the metal at the center of the coordinatively saturated complex, complexes that cleave the DNA at the binding site are obtained. Figure 8.14 shows some of these shape-selective conformational probes.
One example of this shape-selective cleavage is apparent in reactions of Ru(TMP)32+ (TMP = 3,4,7,8-tetramethylphenanthroline), a probe of the A-conformation.28,29 The complex was designed by incorporating methyl groups about the periphery of each phenanthroline ligand to preclude intercalative binding of the complex to the helix, owing to the bulkiness of the methyl groups, and at the same time to promote hydrophobic groove binding, Importantly, however, this hydrophobic groove binding could not occur against the minor groove of B-DNA, given the width and depth of the groove versus the size of the complex. Instead, the shape of the complex was matched well to the shallow minor-groove surface of an A-form helix. Binding studies with synthetic polynucleotides of A, B, and Z-form were consistent with this scheme. Photolysis of the ruthenium complex, furthermore, as with Ru(phen)32+, leads to the sensitization of singlet oxygen, and hence, after treatment with piperidine, to strand cleavage. Thus, photocleavage reactions with Ru(TMP)32+ could be used to delineate A-like regions, with more shallow minor grooves, along a helical polymer. At such sites, Ru(TMP)32+ would bind preferentially, and upon photolysis, generate locally higher concentrations of singlet oxygen to mediate cleavage of the sugar-phosphate backbone. This scheme revealed that homopyrirnidine stretches along the helix adopt a more A-like conformation.29
The targeting of altered conformations such as Z-DNA has been described earlier58 in the context of a spectroscopic probe, A-Ru(DIP)32+. Substitution of a photoredox-active metal into the core of the tris(diphenylphenanthroline) unit leads also to a complex that both binds and, with photoactivation, cleaves at the altered conformation.55 Both Co(III) polypyridyl and Rh(III) polypyridyl complexes have been shown to be potent photooxidants. Coupled to site-specific DNA binding, these metal complexes, with photoactivation, become conformationally selective DNA cleavage agents. Co(DIP)33+, for example, has been shown to cleave specifically at Z-form segments inserted into DNA plasmids.55,70 Perhaps even more interesting, on both natural plasmids and viral DNAs, the various sites cleaved by Co(DIP)33+, corresponding both to Z-form sites and to other locally altered non-B-conformations, coincide with functionally important regions of the genome, e.g., regulatory sites, gene termination sites, and intron-exon joints.70,71 The altered structures recognized by the metal complexes, therefore, appear to mark biologically important sites, those presumably recognized also by cellular proteins. Cleavage studies with these metal complexes, therefore, are providing some insight also into how Nature specifically targets and accesses the sequence information encoded along the DNA polymer, sequence information encoded indirectly through local structure.
The most striking example of the specificity to be derived from shape-selective targeting has been given by the double-stranded cleavage induced by Rh(DIP)33+ at cruciforms.72 Rh(DIP)33+, like its Co(III) and Ru(II) congeners, binds to locally unwound, non-B-conformations such as Z-DNA, but interestingly this potent photooxidant yields the specific cleavage of both DNA strands at cruciform sites. Lacking any crystallographic information, our understanding of the local structure of a cruciform is poor. In these palindromic sites, a torsionally strained DNA extrudes two intrastrand hydrogen-bonded helices from the main helix (see Figure 8.2B). Clearly the structure is grossly altered and locally unwound. Rh(DIP)33+ appears to bind into a pocket generated by the folding of the extruded helix onto the main helix. The recognition is of this intricately folded structure, not of the sequence used to generate the cruciform. Studies with the transition-metal complex on different cruciforms should be useful in helping to characterize this interesting tertiary DNA structure.
Shape-selective transition-metal probes have also been useful in delineating more subtle variations in structure, such as the propeller twisting and tilting evident in B-form DNA.30,73 Rh(phen)2phi3+ was found to target preferentially sites in the major groove where the DNA base pairs are more open; this preferential recognition arises from the steric constraints at more-closed intercalation sites because of the bulkiness of the ancillary phenanthroline ligands above and below the intercalation plane. Two straightforward structural perturbations that lead to an opening of the major groove involve the propeller twisting of bases with respect to one another and the tilting of base pairs along the helix. Chiral discrimination in cleavage by Rh(phen)2phi3+ is now being used quantitatively to discriminate among these structural parameters. The goal in these studies is to use cleavage results with the shape-selective metal complexes to describe the three-dimensional structure of a long double-helical DNA sequence in solution. Probing structurally well-defined sequences with a whole family of shape-selective metal complexes may provide a route to this goal.
As mentioned above, describing the three-dimensional structures of RNAs is an even more complicated task than it is for double-stranded DNAs. Only a few tRNAs have been characterized crystallographically to high resolution, and for other larger RNA structures, such as 5S RNA, or any of the catalytic intervening-sequence RNAs, little is known about their folding characteristics. To understand the regulation and catalytic functions of these biopolymers, we need to develop chemical tools to explore these structures. Figure 8.15 shows the results of cleavage studies using the variety of transition-metal probes on tRNAPhe. Hydroxyl-radical cleavage mediated by Fe(EDTA)2- reveals the protection of solvent-inaccessible regions, the "inside" of the molecule.23 MPE-Fe(II) appears to demarcate the double-helical regions,74 Cu(phen)2+ shows the loopedout single-stranded segments,75 and Rh(phen)2phi3+ seems to delineate those regions involved in triple-base interactions, the sites of tertiary folding.76 Taken together, the full structure of the tRNA can be described based upon cleavage data with transition-metal complexes. It therefore seems as if this full family of coordination complexes might be generally useful in delineating RNA structures. Still more work is needed quantitatively to compare the patterns obtained with the few well-characterized structures. Nonetheless, an important role for these and possibly other transition-metal reagents is indicated.
Other Novel Techniques
Transition-metal ions can also be used advantageously tethered onto peptides, proteins, oligonucleotides, and other natural products, to provide a chemical probe for their binding interactions with nucleic acids. This strategy, termed affinity cleavage, was developed by Dervan and coworkers in preparing and characterizing distamycin-Fe(II)EDTA.24 Distamycin is a known natural product that binds in the minor groove of DNA at AT-rich sequences. By tethering Fe(II)EDTA onto distamycin, the researchers converted the DNA-binding moiety into a DNA-cleaving moiety, since, as with MPE-Fe(1I), in the presence of peroxide and a reductant, hydroxyl radical chemistry would be delivered to the distamycin binding site. Unlike MPE-Fe(II), however, the distamycin moiety shows preferential binding at some sites along the polymer, and hence only at those sites would the local hydroxyl-radical concentration be increased and cleavage be obtained. As a result the tethered Fe(EDTA)2- could be used as a cleavage probe, marking sites of specific binding.
Affinity cleaving has been generalized so that now Fe(EDTA)2- can be tethered onto both oligonucleotides and peptides to follow their interactions with nucleic acids. The sequence-specific binding of oligonucleotides to double-helical DNA through triple-helix formation is but one of many examples where the tethering of Fe(EDTA)2- has been applied advantageously.77
Other redox-active metals can be incorporated into DNA-binding moieties as well. Schemes have been developed to functionalize accessible lysine residues on DNA-binding repressor proteins with phenanthrolines, so that in the presence of copper ion, peroxide, and a reductant, the phenanthroline-bound copper on the protein would induce DNA strand cleavage. Through this scheme, again the conversion of a DNA-binding moiety into a cleaving moiety by incorporation of a redox-active metal, the specific binding sites of repressor proteins can be readily identified (far more quickly on large DNA than through footprinting).78
Another scheme, which perhaps takes advantage more directly of bioinorganic chemistry, involves engineering redox metal-binding sites into DNA-binding proteins and peptides. The DNA-binding domain of the protein Hin recombinase was synthesized chemically, and first, to examine the folding of the peptide on the DNA helix, EDTA was tethered onto the peptide for Fe(II) cleavage experiments.79 But as is illustrated repeatedly in these chapters, Nature has already provided amino-acid residues for the chelation of metal ions into proteins. Thus the DNA-binding domain of Hin recombinase was synthesized again, now including at its terminus the residues Gly-Gly-His, a known chelating moiety for copper(II).80 This chemically synthesized peptide, with now both DNA-binding and DNA-cleaving domains, as illustrated in Figure 8.16, specifically promotes cleavage at the Hin recombinase binding site in the presence of bound copper and ascorbate. Interestingly, the addition of nickel(II) also leads to specific strand cleavage, without diffusible intermediates. Using this approach, taking advantage of the chelating abilities of amino acids and the cleaving abilities of different metal ions, one may prepare new synthetic, functional metalloproteins that bind and react with DNA. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/08%3A_MetalNucleic_Acid_Interactions/8.05%3A_Applications_of_Different_Metal_Complexes_that_Bind_Nucleic_Acids_%28Part_2%29.txt |
In the context of what we understand about the fundamental interactions and reactions of metal ions and complexes with nucleic acids, and also in comparison to how chemists have been exploiting these interactions in probing nucleic acids, we can also consider how Nature has taken advantage of metal ions in the construction of metalloproteins, nucleic-acid assemblies, and smaller natural products containing metal ions that interact with DNA and RNA.
A Structural Role
One of the chief functions attributed to metal ions in biological systems is their ability to provide a structural center to direct the folding of a protein. Just as shape-selective recognition has been helpful in targeting metal complexes to specific sites on DNA, it appears that one element of the recognition of sites by DNA regulatory proteins may also involve the recognition of complementary shapes. Furthermore, metal ions appear to be used in these proteins to define the shape or folding pattern of the peptide domain that interacts specifically with the nucleic acid.
The DNA-binding metalloproteins that have received the greatest attention recently have been the "zinc-finger" regulatory proteins. It was discovered in 1983 that zinc ions played a role in the functioning of the nucleic acid-binding transcription factor IlIA (TFIIIA) from Xenopus laevis, which binds specifically both DNA, the internal control region of the 5S rRNA gene, and RNA, the 5S RNA itself.81 The protein (actually the 7S storage particle) was found to contain two to three equivalents of zinc ion. Dialysis removed both the associated zinc ions and the nucleic-acid-binding ability of the protein. Importantly, treatment with zinc ion, or in later studies with higher concentrations of Co2+, restored the specific binding ability. Hence, zinc ion was shown to be functionally important in these eukaryotic regulatory proteins.
The notion of a "zinc-finger structural domain" was first provided by Klug and coworkers, after examination of the amino-acid sequence in TFIIIA.82 It was found that TFIIIA contained nine imperfect repeats of a sequence of approximately 30 amino acids, and furthermore that each repeat contained two cysteine residues, two histidine residues, and three hydrophobic residues, in conserved positions. In addition, subsequent metal analyses were revealing higher zinc contents (7 to 11 equivalents) associated with the protein, and protein-digestion experiments indicated that several repeated structural domains existed in the protein. The two cysteine thiolates and two histidine imidazoles in each repeated domain could certainly serve to coordinate a zinc ion. Thus it was proposed that each peptide repeat formed an independent nucleic-acid-binding domain, stabilized in its folded structure through coordination of a zinc ion. The peptide unit was termed a "zinc finger," which is illustrated schematically in Figure 8.17. TFIIIA was therefore proposed to contain nine zinc fingers, which would cooperatively bind in the internal control region of the 5S RNA gene.
An enormous number of gene sequences from a variety of eukaryotic regulatory proteins was then found to encode strikingly similar amino-acid sequences,83 and many were dubbed zinc-finger proteins. The bioinorganic chemist, however, should be aware that chemical analyses supporting such assignments are first required. Nonetheless, several legitimate examples of eukaryotic nucleic-acid-binding zinc-finger proteins containing multiple zinc-binding peptide domains have emerged since the first study of TFIIIA, including the proteins Xfin from Xenopus, the Kruppel protein from Drosophila, the Sp1 transcription factor, and human testes-determining factor. It has therefore become clear that the zinc-finger domain represents a ubiquitous structural motif for eukaryotic DNA-binding proteins.84
What is the structure of a zinc finger, and how is this structure important for binding a specific nucleic-acid site? Based on a search of crystallographic databases for metalloproteins and an examination of the consensus sequence emerging for zinc fingers (that is, which residues were truly conserved and common to the different putative zinc fingers), Berg proposed a three-dimensional structure for a zinc finger, shown schematically in Figure 8.17.85 The proposed structure included the tetrahedral coordination of zinc by the two cysteine and histidine residues at the base of the finger and an $alpha$-helical region running almost the length of the domain. EXAFS studies also supported the tetrahedral zinc site. Since this proposal, two detailed two-dimensional NMR studies have been reported that are consistent with the tetrahedral zinc coordination and the a-helical segment.86 More recently, a crystal structure of a three-finger binding domain associated with an oligonucleotide was determined.87 The zinc fingers lie in the major groove of DNA, the $\alpha$-helical region being within the groove. Not surprisingly, given basic coordination chemistry, the zinc does not interact directly with the nucleic acid. Instead, the zinc ion must serve a structural role, defining the folding and three-dimensional structure of the protein scaffolding about it. This structure, defined by the metal at its center, like other coordination complexes, is able to recognize its complementary structure on the nucleicacid polymer.
It should also be noted that this zinc-finger structural motif is not the only metal-containing or even zinc-containing structural motif important in nucleic-acid-binding proteins.88 A clearly different domain is evident in the protein GAL4, a transcription factor required for galactose utilization in S. cerevisiae.88a A recent crystal structure of the protein bound to an oligonucleotide shows the protein to bind to DNA as a dimer; each monomer contains a binuclear zinc cluster with two zinc ions tetrahedrally coordinated by six cysteines (two cysteines are bridging), not dissimilar from proposed structures in metallothionein. Still another structural motif is found in the glucocorticoid receptor DNA-binding domain. Crystallography89 has revealed that this domain also binds DNA as a dimer; here each monomer contains two zinc-nucleated substructures of distinct conformation. The zinc ions are each tetrahedrally coordinated to four cysteine residues. Likely this too represents another structural motif for proteins that bind nucleic acids, and one again in which the metal serves a structural role.
Lastly, one might consider why zinc ion has been used by Nature in these nucleic-acid binding proteins. Certainly, the natural abundance of zinc is an important criterion. But also important is the absence of any redox activity associated with the metal ion, activity that could promote DNA damage [as with Fe(II) or Cu(II), for example]. In addition, other softer, heavier metal ions might bind preferentially to the DNA bases, promoting sequence-specific covalent interactions. Zinc ion, therefore, is clearly well-chosen for the structural center of these various nucleic-acid-binding proteins.
A Regulatory Role
Metalloregulatory proteins, like the transcription factors described above, affect the expression of genetic information through structural interactions that depend upon the metal ions, but unlike the zinc-finger proteins, metalloregulatory proteins act as triggers, repressing or activating transcription given the presence or absence of metal ion. In some respects, even more than zinc fingers, these systems resemble the Ca2+-activated proteins described in Chapter 3.
Consider the biological system that must respond to changing intracellular metal concentrations. At high concentrations many metal ions become toxic to the cell; hence, a full system of proteins must be synthesized that will chelate and detoxify the bound-metal-ion pool. In order to actively engage these proteins, the genes that encode them must be rapidly transcribed. But at the same time, the DNA itself must be protected from the high concentrations of metal ion. Hence the need for these metalloregulatory proteins, which bind DNA in the absence of metal ion, usually repressing transcription, but in the presence of metal ion bind the metal ion tightly and specifically, and as a consequence amplify transcription.
Perhaps the best-characterized metalloregulatory system thus far is the MerR system, regulating mercury resistance in bacteria.90 An inducible set of genes arranged in a single operon is under the control of the metal-sensing MerR protein, and it is this system that mediates mercury resistance. Mercury resistance depends upon the expression of these genes to import toxic Hg(lI), reduce Hg(II) to the volatile Hg(0) by NADPH, and often additionally to cleave organomercurials to their corresponding hydrocarbon and Hg(II) species. The MerR protein regulates mercury resistance both negatively and positively. As illustrated in Figure 8.18, MerR in the absence of Hg(II) binds tightly and sequence-specifically to the promoter. In so doing, MerR inhibits binding to the promoter by RNA polymerase. When Hg2+ is added at low concentrations, the metal ion binds specifically and with high affinity to the DNA-bound MerR, and causes a DNA conformational change detectable by using other metal reagents as conformational probes. This conformational change now facilitates the binding of RNA polymerase and hence activates expression of the gene family.
What are the structural requirements in the metal-binding site? Certainly one requirement is Hg(II) specificity, so that other metal ions will not also trigger transcriptional activation. Another is high metal-binding affinity to protect the DNA from direct coordination of the Hg(lI). The MerR protein is dimeric, and contains four cysteine residues per monomer. Site-directed mutagenesis studies91 have indicated that three of these four cysteine residues are needed for Hg(II) binding, and EXAFS studies92 have been consistent with tricoordinate ligation, clearly a well-designed system for Hg(II) specificity. Perhaps even more interesting, the site-directed mutagenesis studies91 on heterodimers (a mixture of mutant and wild-type monomers) have indicated that the coordinated Hg(II) bridges the dimer, ligating two cysteines of one monomer and one cysteine of the other. This scheme may provide the basis also for the kinetic lability needed in a rapidly responsive cellular system.
Model systems are also being constructed to explore metal modulation of DNA binding. One system involves the assembly of two dipeptides linked by a central acyclic metal-binding polyether ligand, with Fe(EDTA)2- tethered to one end to mark site-specific binding.93 In the presence of alkaline earth cations, which induce a conformational change that generates a central macrocycle, the linked peptides become oriented to promote sequence-specific binding in the minor groove. In the absence of the alkaline earth ion, no site-specific binding, or cleavage of DNA, is evident. One might consider this system as a simple, first-order synthetic model for the metalloregulatory proteins.
MerR is clearly only one natural metalloregulatory system. Other metal ions bind regulatory factors to mediate the regulation of gene expression in a metalspecific manner. Two examples include the Fe(II)-binding Fur protein from enteric bacteria94 and the copper-binding protein ACE1/CUP2 from S. cerevisiae.95 Both copper and iron are essential trace elements for which high concentrations are toxic; for nucleic acids this toxicity is certainly the result of redox-mediated strand damage. Other metal-specific regulatory systems are surely present as well. Both the MerR and the synthetic system may exemplify how these various systems function, how Nature might construct a ligand system to facilitate toxicmetal-specific binding in the presence of DNA that then alters or triggers how other moieties bind and access the nucleic acid.
A Pharmaceutical Role
With the exception of cisplatin (see Chapter 9), most pharmaceuticals currently being used as DNA-binding agents were first isolated as natural products from bacteria, fungi, plants, or other organisms. For the most part they represent complex organic moieties, including peptide and/or saccharide functionalities, and often a unique functionality, such as the ene-diyne in calichimycin. These natural products bind DNA quite avidly, through intercalation, groove binding, or a mixture thereof. Often the efficacy of these antitumor antibiotics stems from subsequent alkylation or DNA strand-cleavage reactions that damage the DNA.
Among the various natural products used clinically as antitumor antibiotics are bleomycins, a family of glycopeptide-derived species isolated from cultures of Streptomyces.25,96 The structure of bleomycin A2 is shown schematically in Figure 8.19. The molecular mode of action of these species clearly involves binding to DNA and the promotion of single-stranded cleavage at GT and GC sequences. Importantly, as demonstrated by Horwitz, Peisach, and coworkers, this DNA cleavage requires the presence of Fe(II) and oxygen.97 Thus, one might consider Fe-bleomycin as a naturally occurring inorganic pharmaceutical.
What is the role of the metal ion in these reactions? As one might imagine, based upon our earlier discussions of metal-promoted DNA cleavage, the iron center is essential for the oxidative cleavage of the strand through reaction with the sugar moiety. The reaction of Fe(II)-bleomycin can, however, clearly be distinguished from the Fe(EDTA)2- reactions discussed earlier in that here no diffusible intermediate appears to be involved. Instead of generating hydroxyl radicals, the Fe center must be positioned near the sugar-phosphate backbone and activated in some fashion to promote strand scission directly.
Despite extensive studies, in fact little is known about how Fe(II)-bleomycin is oriented on the DNA. Indeed, the coordination about the metal is the subject of some debate. The structure of Cu(II)-P-3A,98 a metallobleomycin derivative, is also shown in Figure 8.19. On the basis of this structure and substantive spectroscopic studies on Fe-bleomycin itself, it is likely that, as with Cu(II), in the Fe(II)-bleomycin complex the metal coordinates the $\beta$-hydroxyimidazole nitrogen, the secondary amine of $\beta$-aminoalanine, and the N1 nitrogen of the pyrimidine. Whether in addition the primary amines of the amino alanine and of the histidine coordinate the metal is still not settled. Possibly bithiazole coordination or some coordination of the sugar moieties is involved. Nonetheless, given five different coordination sites to the bleomycin, the sixth axial site is available for direct coordination of dioxygen. How is this Fe—O2 complex oriented on the DNA? It is likely that at least in part the complex binds against the minor groove of the helix. There is some evidence that suggests that the bithiazole moiety intercalates in the helix. It is now becoming clear, however, that the structure of the metal complex itself, its three-dimensional shape, rather than simply the tethered bithiazole or saccharide, is needed for the sequence selectivity associated with its mode of action.
Although the coordination and orientation of the metal complex are still not understood, extensive studies have been conducted concerning the remarkable chemistry of this species.96,99 The overall mechanism of action is described in Figure 8.19. In the presence of oxygen, the Fe(Il) O2 species is formed and is likely rapidly converted to a ferric superoxide species. The one-electron reduction of this species, using either an organic reductant or another equivalent of Fe(II)-bleomycin, leads formally to an Fe(III)-peroxide, which then undergoes O—O bond scission to form what has been termed "activated bleomycin." This species might be best described as Fe(V)=O (or [Fe O]3+). This species is comparable in many respects to activated cytochrome P-450 or perhaps even more closely to the Fe center in chloroperoxidase (see Chapter 5). Like these systems, activated bleomycin can also epoxidize olefins and can generally function as an oxo transferase. In contrast to these systems, Fe-bleomycin clearly lacks a heme. How this species can easily shuttle electrons in and out, forming and reacting through a high-valent intermediate, without either the porphyrin sink or another metal linked in some fashion, is difficult to understand. In fact, understanding this process, even independently of our fascination with how the reaction is exploited on a DNA helix, forms the focus of a substantial effort of bioinorganic chemists today.
What has been elucidated in great detail is the reaction of activated bleomycin with DNA. It has been established that the activated species promotes hydrogen abstraction of the C4'-H atom, which is positioned in the minor groove of the helix (Figure 8.19). Addition of another equivalent of dioxgen to this C4'-radical leads to degradation of the sugar to form a 5'-phosphate, a 3'-phosphoglycolate, and free base propenal. Alternatively, oxidation of the C4'-radical followed by hydroxylation in the absence of oxygen yields, after treatment with base, a 5'-phosphate, an oxidized sugar phosphate, and free base.
Other metals such as copper and cobalt can also activate bleomycins, although their mechanistic pathways for strand scission are clearly different from that of Fe(II)-bleomycin. Whether other natural products that bind DNA also chelate metal ions and exploit them for oxidative strand cleavage is not known, but several systems provide hints that they do. Furthermore, such a fact would not be surprising given our understanding of the utility of metal ions in promoting this chemistry. An even more detailed understanding of this chemistry might lead to the development of second-generation synthetic transition-metal pharmaceuticals that specifically and efficiently target and cleave DNA sites.
A Catalytic Role
In addition to serving structural and modulating roles in proteins which bind nucleic acids, metal ions also appear to be essential to the functioning of various complex enzymes that act on nucleic acids. At this stage our understanding of the participation of the metal ion in the catalytic chemistry of these enzymes is somewhat sketchy, and we are relying more on our current understanding of the possible roles where metal ions may prove advantageous. These remain areas of biochemical focus where the inorganic chemist could make a major contribution.
For example, zinc ion appears to be essential to the functioning of both RNA polymerases and DNA topoisomerases.100-102 These multisubunit enzymes perform quite complex tasks. RNA polymerase must bind site-specifically to its DNA template, bind its nucleotide and primer substrates, and form a new phosphodiester bond in elongating the growing RNA. Two zinc ions appear to be involved. One may be involved in orienting the nucleotide substrate, and the other structurally in template recognition. It would not be surprising, however—indeed, it might be advantageous—if one or both metal ions also participated in the polymerization step. Our mechanistic understanding of how topoisomerases function is even more cursory. These complex enzymes bind supercoiled DNA, sequentially break one strand through hydrolytic chemistry, move the strand around the other (releasing one tertiary turn), and religate the strand. Again, the zinc ion might participate in the hydrolytic chemistry, the ligation step, or both; alternatively, the metal might again serve a structural role in recognition of the site of reaction.
We do have some understanding of the role of metal ions in several endonucleases and exonucleases. As discussed in Section II.C, metal ions may effectively promote phosphodiester hydrolysis either by serving as a Lewis acid or by delivering a coordinated nucleophile. Staphylococcal nuclease103 is an extracellular nuclease of Staphylococcus aureus that can hydrolyze both DNA and RNA in the presence of Ca2+. The preference of the enzyme is for single-stranded DNA, in which it attacks the 5'-position of the phosphodiester linkage, cleaving the 5'-P-O bond to yield a 5'-hydroxyl and 3'-phosphate terminus. Ca2+ ions are added as cofactors and are strictly required for activity. The structure of staphylococcal nuclease, determined by x-ray crystallography and crystallized in the presence of Ca2+ and the enzyme competitive inhibitor pdTp, as well as subsequent NMR and EPR studies on mutant enzymes using Mn2+ as a substitute for the Ca2+ ion, have provided the basis for a detailed structural analysis of the mechanism of this enzyme. In this phosphodiester hydrolysis, the metal ion appears to function primarily as an electrophilic catalyst, polarizing the P-O bond, and stabilizing through its positive charge the evolving negative charge on the phosphorus in the transition state. The base is thought here not to be directly coordinated to the metal; instead, action of a general base is invoked.
Metal ions also participate in the functioning of other nucleases, although the structural details of their participation are not nearly as established as those for staphylococcal nuclease. DNAse I also requires Ca2+ for its catalytic activity.104 S1 endonuclease, mung bean nuclease, and Physarum polycephalem nuclease require zinc ion either as cofactors or intrinsically for nuclease activity, and the restriction enzyme EcoRI may also require intrinsically bound zinc ion.33 In terms of how the zinc ion might function in these enzymes, one can look both to staphylococcal nuclease and to bacterial alkaline phosphatase105 for some illustrations. One would expect that this metal ion could serve both as an electrophilic catalyst and also in the delivery of a zinc-coordinated hydroxide, as it does in alkaline phosphatase, directly attacking the phosphate ester. More work needs to be done to establish the mechanisms by which zinc ion promotes phosphodiester hydrolysis in these enzymatic systems.
Probably most intriguing and mysterious at this stage is the metal participation in the very complex DNA-repair enzyme endonuclease III from E. coli (similar enzymes have also been isolated from eukaryotic sources). This enzyme is involved in the repair of DNA damaged by oxidizing agents and UV irradiation, and acts through an N-glycosylase activity to remove the damaged base, and through an apurinic/apyrimidinic endonuclease activity to cleave the phosphodiester bond adjacent to the damaged site. Although more complex in terms of recognition characteristics, this enzyme functions in hydrolyzing the DNA phosphodiester backbone. What is so intriguing about this enzyme is that it contains a 4Fe-4S cluster (see Chapter 7) that is essential for its activity!106 We think generally that Fe-S clusters best serve as electron-transfer agents. In the context of this repair enzyme, the cluster may be carrying out both an oxidation and a reduction, to effect hydrolysis, or alternatively perhaps a completely new function for this metal cluster will emerge. (Fe-S clusters may represent yet another structural motif for DNA-binding proteins and one which has the potential for regulation by iron concentration.) Currently the basic biochemical and spectroscopic characterization of the enzyme is being carried out. Understanding this very novel interaction of a metal center and nucleic acid will require some new ideas, and certainly represents one new challenge for the bioinorganic chemist. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/08%3A_MetalNucleic_Acid_Interactions/8.06%3A_Nature%27s_Use_of_MetalNucleic-acid_Interactions.txt |
Metal ions are required for many critical functions in humans. Scarcity of some metal ions can lead to disease. Well-known examples include pernicious anemia resulting from iron deficiency, growth retardation arising from insufficient dietary zinc, and heart disease in infants owing to copper deficiency. The ability to recognize, to understand at the molecular level, and to treat diseases caused by inadequate metal-ion function constitutes an important aspect of medicinal bioinorganic chemistry.
Metal ions can also induce toxicity in humans, classic examples being heavy-metal poisons such as mercury and lead. Even essential metal ions can be toxic when present in excess; iron is a common household poison in the United States as a result of accidental ingestion, usually by children, of the dietary supplement ferrous sulfate. Understanding the biochemistry and molecular biology of natural detoxification mechanisms, and designing and applying ion-specific chelating agents to treat metal overloads, are two components of a second major aspect of the new science that is evolving at the interface of bioinorganic chemistry and medicine.
Less well known than the fact that metal ions are required in biology is their role as pharmaceuticals. Two major drugs based on metals that have no known natural biological function, Pt (cisplatin) and Au (auranofin), are widely used for the treatment of genitourinary and head and neck tumors and of rheumatoid arthritis, respectively. In addition, compounds of radioactive metal ions such as 99mTc and complexes of paramagnetic metals such as Gd(III) are now in widespread use as imaging agents for the diagnosis of disease. Many patients admitted overnight to a hospital in the U.S. will receive an injection of a 99mTc compound for radiodiagnostic purposes. Yet, despite the obvious success of metal complexes as diagnostic and chemotherapeutic agents, few pharmaceutical or chemical companies have serious in-house research programs that address these important bioinorganic aspects of medicine.
This chapter introduces three broad aspects of metals in medicine: nutritional requirements and diseases related thereto; the toxic effects of metals; and the use of metals for diagnosis and chemotherapy. Each area is discussed in survey form, with attention drawn to those problems for which substantial chemical information exists. Since there is only a primitive understanding at the molecular level of the underlying biochemical mechanisms for most of the topics, this field is an important frontier area of bioinorganic chemistry. The major focus of this chapter is on the platinum anticancer drug cisplatin, which is presented as a case study exemplifying the scope of the problem, the array of methodologies employed, and the progress that can be made in understanding the molecular basis of a single, if spectacular, metal complex used in medicine today.
VI. Restrospective
The topics discussed in this chapter are helping to expand bioinorganic chemistry from a subject that arose chiefly from spectroscopic analysis of metal centers in proteins, because they were uniquely convenient functional groups, to a discipline where fundamental knowledge about metal functions and the application of metals as diagnostic and chemotherapeutic agents are making important contributions to medicine. As the case study of cisplatin is intended to demonstrate, progress in understanding how metals function in chemotherapy can be made only by the combined efforts of many disciplines, including synthetic and physical inorganic and organic chemistry, molecular and cell biology, immunology, pharmacology, toxicology, and clinical medicine. Although we have not yet reached the day where chemotherapeutic agents can be rationally designed from knowledge of a molecular mechanism, such a concept does not seem that farfetched. If nothing else, knowledge of fundamental bioinorganic processes related to metal-macromolecule interactions will continue to grow enormously through efforts to achieve this ultimate goal.
Contributors and Attributions
• Stephen J. Lippard (Massachusetts Institute of Technology, Department of Chemistry)
Thumbnail: Cisplatin, \(PtCl_2(NH_3)_2\) A platinum atom with four ligands. Image used with permission (Public Domain; Benjah-bmm27).
09: Metals in Medicine
Essential Metals
Four main group (Na, K, Mg, and Ca) and ten transition (V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Mo, and Cd) metals are currently known or thought to be required for normal biological functions in humans. Table 9.1 lists these elements, their relative abundances, and the medical consequences of insufficient quantities where known. The nutritional requirements for selected members of the essential metals are discussed in the following sections.
Table 9.1 - Essential metals and medical consequences resulting from their deficiency.a
a) Data taken from E.-i. Ochiai, Bioinorganic Chemistry, Allyn & Bacon, 1977, p. 6
Metal Abundance
Sea Water
mg/1 (ppm)
Abundance
Earth's Crust
mg/1 (ppm)
Diseases Resulting from Metal Deficiency
Na 1.05 x 104 2.83 x 104
K 380 2.59 x 104
M- 1.35 x 104 2.09 x 104
Ca 400 3.63 x 104 bone deterioration
V 2 x 10-3 135
Cr 5 x 10-3 100 glucose tolerance (?)
Mn 2 x 10-5 950
Fe 1 x 10-2 5.00 x 104 anemia
Co 1 x 10-4 25 anemia
Ni 2 x 10-4 75
Cu 3 x 10-3 55 brain disease, anemia, heart disease
Zn 1 x 10-2 70 growth retardation, skin changes
Mo 1 x 10-2 1.5
Cd 1.1 x 10-4 0.2
Anemia and Iron2
Anemia results from insufficient oxygen supply, often because of a decrease in hemoglobin (Hb) blood levels. Approximately 65 to 70 percent of total body iron resides in Hb. In the U.S., many foods, especially those derived from flour, are enriched in iron. In third-world countries, however, scarcity of dietary iron is a major contributor to anemia. This information illustrates one important fact about disease that results from metal deficiency, namely, the need for an adequate supply of essential metals in food. A related aspect, one of greater interest for bioinorganic chemistry, is the requirement that metals be adequately absorbed by cells, appropriately stored, and ultimately inserted into the proper environment to carry out the requisite biological function. For iron, these tasks, among others, are performed by specific iron-chelating agents, the storage protein ferritin and the transport protein transferrin, the bioinorganic chemistry of which is extensively discussed in Chapter 1.
Another cause of anemia exists in individuals who have a mutant variety of hemoglobin, HbS, in which valine has been substituted for glutamic acid in the sixth position of the $\beta$ subunits.3 Interestingly, extensive studies have shown that this phenomenon, which leads to sickling of the red blood cells, does not result from failure of the protein to bind heme or from changes in the O2 binding constant of the iron atom. Rather, deoxy HbS polymerizes into soluble, ordered fibrous structures that lower the ability of blood to carry oxygen effectively to the tissues. These results illustrate the importance of structural features remote from the metal-binding domain in determining the functional characteristics of a metalloprotein.
Causes and Consequences of Zinc Deficiency4-6
The average adult contains ~ 2 g of zinc and requires a daily intake of 15 to 20 mg, only half of which is absorbed, to maintain this level. Although food in many technologically advanced societies contains sufficient zinc to afford this balance, zinc deficiencies occur in certain populations where there is either an unbalanced diet or food that inhibits zinc absorption. An especially interesting example of the latter phenomenon is found in certain villages in the Middle East where phytates, organic phosphates present in unleavened bread, chelate zinc ion and render it inaccessible. Zinc deficiency produces growth retardation, testicular atrophy, skin lesions, poor appetite, and loss of body hair. Little is known about the biochemical events that give rise to these varied consequences, although the three most affected enzymes are alkaline phosphatase, carboxypeptidase, and thymidine kinase. About 30 percent of zinc in adults occurs in skin and bones, which are also likely to be affected by an insufficient supply of the element. Zinc deficiency is readily reversed by dietary supplements such as ZnSO4, but high doses (>200 mg) cannot be given without inducing secondary effects of copper, iron, and calcium deficiency.
Copper Deficiency7
More copper is found in the brain and heart than in any other tissue except for liver, where it is stored as copper thionein and released as ceruloplasmin or in the form of a complex with serum albumin. The high metabolic rate of the heart and brain requires relatively large amounts of copper metalloenzymes including tyrosinase, cytochrome c oxidase, dopamine-$\beta$-hydroxylase, pyridoxal-requiring monamine oxidases, and Cu-Zn superoxide dismutase. Copper deficiency, which can occur for reasons analogous to those discussed above for Fe and Zn, leads to brain disease in infants, anemia (since cytochrome oxidase is required for blood formation), and heart disease. Few details are known about the molecular basis for copper uptake from foods.
Summary
From the above anecdotal cases, for which similar examples may be found for the other metals in Table 9.1, the biological consequences of metal deficiency are seen to result from a breakdown in one or more of the following steps: adequate supply in ingestible form in foodstuffs; absorption and circulation in the body; uptake into cells; insertion into critical proteins and enzymes requiring the element; adequate storage to supply needed metal in case of stress; and an appropriate mechanism to trigger release of the needed element under such circumstances. Only for iron, and to a lesser extent copper and zinc, is there a reasonably satisfying picture of the molecular processes involved in this chain of events. The elucidation of the detailed mechanisms of these phenomena, for example, the insertion of iron into ferritin, remains an exciting challenge for the bioinorganic chemist (see Chapter 1). | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/09%3A_Metals_in_Medicine/9.01%3A_Metal_Deficiency_and_Disease.txt |
Two Classes of Toxic Metal Compounds
As intimated in the previous section, the presence of excess quantities of an essential metal can be as deleterious as insufficient amounts. This situation can arise from accidental ingestion of the element or from metabolic disorders leading to the incapacitation of normal biochemical mechanisms that control uptake and distribution phenomena. These possibilities constitute one major class of metal toxicity. The other broad class results from entry of nonessential metals into the cell through food, skin absorption, or respiration. The toxicities associated with this latter class have received much recent attention because of the public health risks of chemical and radioisotopic environmental pollutants.
In this section, we survey examples of both categories, and discuss ways in which bioinorganic chemistry can contribute to the removal of toxic metals and restoration of normal function. One way involves chelation therapy, in which metal-specific chelating agents are administered as drugs to complex and facilitate excretion of the unwanted excess element. The use of desferrioxamine to treat iron poisoning is one example of this approach. A second role of bioinorganic chemistry is to identify fundamental biological mechanisms that regulate metal detoxification, and to apply the principles that emerge to help control the toxic effects of metal ions in the environment. Recent studies of mercury resistance and detoxification in bacteria provide an elegant example of the way in which biochemistry and molecular biology can be used to elucidate events at the molecular level. This work, which has uncovered the existence of metalloregulatory proteins, is described in some detail in Section III.F below. It represents a benchmark by which other investigations into the mechanisms of metal-detoxification phenomena may be evaluated.
Copper Overload and Wilson's Disease8
Wilson's disease results from a genetically inherited metabolic defect in which copper can no longer be tolerated at normal levels. The clinical manifestations are liver disease, neurological damage, and brown or green (Kayser-Fleischer) rings in the cornea of the eyes. Patients suffering from Wilson's disease have low levels of the copper-storage protein ceruloplasmin; the gene and gene products responsible for the altered metabolism have not yet been identified. Chelation therapy, using K2Ca(EDTA), the Ca2+ ion being added to replenish body calcium stores depleted by EDTA coordination, 2,3-dimercaptopropan-1-ol (BAL, British Anti-Lewisite), or d-penicillamine to remove excess copper, causes the symptoms to disappear. The sulfhydryl groups of the latter two compounds presumably effect removal of copper as Cu(I) thiolate complexes. Wilson's disease offers an excellent opportunity for modem methodologies to isolate and clone the gene responsible for this altered Cu metabolism, ultimately providing a rational basis for treatment.
Iron Toxicity9
Chelation therapy is also used to treat iron overload. Acute iron poisoning, such as that resulting from accidental ingestion of FeSO4 tablets, results in corrosion of the gastrointestinal tract. Chronic iron poisoning, or hemochromatosis, arises from digestion of excess iron usually supplied by vessels used for cooking. A classic case of the latter is siderosis induced in members of the Bantu tribe in South Africa, who consume large quantities of beer brewed in iron pots and who suffer from deposits of iron in liver, kidney, and heart, causing failure of these organs. The chelating agent of choice for iron toxicity is the siderophore desferrioxamine, a polypeptide having a very high affinity for Fe(III) but not for other metals. Ferrioxamine chelates occur naturally in bacteria as iron-transport agents. Attempts to mimic and improve upon the natural systems to provide better ligands for chelation therapy constitutes an active area of bioinorganic research (see Chapter 1).
Toxic Effects of Other Essential Metals10,11
When present in concentrations above their normal cellular levels, most of the other metals listed in Table 9.1 are toxic. Calcium levels in the body are controlled by vitamin D and parathyroid hormones. Failure to regulate Ca2+ leads to calcification of tissue, the formation of stones and cataracts, a complex process about which little is understood (see Chapter 3). Chronic manganese poisoning, which can occur following ingestion of metal-oxide dust, e.g., among miners in Chile, produces neurological symptoms similar to Parkinson's disease. Neuron damage has been demonstrated. Although Zn toxicity is rare, it can lead to deficiencies in other essential metals, notably calcium, iron, and copper. Cobalt poisoning leads to gastrointestinal distress and heart failure. Metal poisoning by those elements has been treated by chelating agents, most frequently CaNaiEDTA), but the selectivity offered by the ferrioxamine class of ligands available for iron has not even been approached. Fortunately, there are few cases involving these metals.
Plutonium: A Consequence of the Nuclear Age12
Some of the chelating agents developed to treat iron toxicity have found application as therapeutics for plutonium poisoning. Diethylenetriaminepentaacetic acid (DTPA) salts and siderophores are especially effective. Some improvement over the naturally occurring chelates has been made by tailoring the ligand to encapsulate completely the eight-coordinate Pu(IV) center. Although few individuals have been affected, ingestion of 239Pu, for example, as small particles of PuO2, at nuclear-reactor sites can have dire consequences. 239Pu emits high energy ex particles, leading to malignancies of bone, liver, lung and lymph nodes, to which tissues it is transported by transferrin. With a maximum tolerated dose of only 1.5 $\mu$g, plutonium is among the most toxic metals known. We turn now to other, more classic examples of such industrial pollutants.
Mercuy Toxicity13 and Bacterial Resistance14-17
Mercury is released into the environment as Hg(II) ions through weathering of its most common ore, HgS, red cinnabar. Organomercurials of general formula RHgX used in agriculture have also entered the environment as toxic waste. Both RHgX and HgX2 compounds bind avidly to sulfhydryl groups in proteins, which can lead to neurological disease and kidney failure. Metallothionein is a favored protein target, which may help to limit mercury toxicity. A highly publicized case occurred in 1953 at Minimata, Japan, where 52 people died after eating mercury-contaminated fish and crustaceans near a factory waste outlet. The volatile, elemental form of mercury, Hg(0), is reportedly nontoxic, but its conversion to alkylmercury compounds by anaerobic microorganisms utilizing a vitamin B-12 biosynthetic pathway constitutes a serious health hazard.
Because of the high affinity of mercury for sulfur-donor ligands, mercury poisoning is treated by BAL; N-acetylpenicillamine has also been proposed. Recently, a very interesting natural detoxification system has been discovered in bacteria resistant to mercury; this system, when fully elucidated, might provide important strategies for treating heavy-metal poisoning in humans.
Presumably under environmental pressure, bacteria have developed mechanisms of resistance to HgX2 and RHgX compounds in which mercury is recycled back to Hg(0). At least five gene products are involved in the bacterial mercury-resistance mechanism. MerT and MerP mediate the specific uptake of mercury compounds. MerB, organomercury lyase, and MerA, mercuric reductase, catalyze two of the reactions, given in Equations (9.1) and (9.2). Plasmids encoding the genes for these two proteins have been isolated. A typical arrangement of genes in the mer operon
$RHgX + H^{+} + X^{-} \xrightarrow[lyase]{organomercury} HgX_{2} + RH \tag{9.1}$
$Hg(SR)_{2} + NADPH + H^{+} \xrightarrow[reductase]{mercuric} Hg(0) + NADP^{+} + 2RSH \tag{9.2}$
region of these plasmids is shown in Figure 9.1. The most thoroughly studied gene product is MerR, a metalloregulatory protein that controls transcription of the mer genes. In the absence of Hg(II) the MerR protein binds to DNA as a repressor, preventing transcription of the merT, P, A, and B genes (Figure 9.1) and negatively autoregulating its own synthesis. When Hg(II) is present, transcription of these genes is turned on. Interestingly, the MerR protein remains bound to the same site on DNA whether acting as an activator in the presence of Hg(II) or as a repressor in its absence. Random and site-specific mutagenesis studies implicate several cysteine residues in the carboxyl terminal region of the protein as candidates for the mercury-binding site.
Organomercury lyase, encoded by the merB gene, achieves the remarkable enzymatic step of breaking Hg-C bonds (Equation 9.1). It is a 22-kDa protein with no metals or cofactors. Two cysteine-sulfhydryl groups on the protein have been postulated to effect this chemistry, as depicted in Equation (9.3). Stereochemical studies of the Hg-C bond cleavage revealed retention of configuration, indicating that cleavage of the Hg-C bond probably does not proceed by a radical pathway. A novel concerted SE2 mechanism has been suggested. The enzyme turnover numbers, ranging from 1 min-1 for CH3HgCl to 240 min-1 for butenylmercuric chloride, although slow, are ~105-108-fold faster than the nonenzymatic rate.
$\tag{9.3}$
Mercuric ion reductase, the FAD-containing merA gene product, has several pairs of conserved cysteines. From site-specific mutagenesis studies, cysteine residues in the sequence 134-Thr-Cys-Val-Asn-Val-Gly-Cys-140 are known to comprise a redox-active disulfide group; in addition, a redox-inactive pair of cysteines near the carboxyl terminus is also required for the selective reduction of Hg(II). Exactly how the enzyme achieves the chemistry shown in Equation (9.2) is currently uncertain, but the redox activities of the flavin and disulfide/ thiol centers are undoubtedly involved. This enzyme serves both to detoxify mercury supplied directly from the environment as Hg(II) salts and to complete clearance of Hg2+ generated by the MerB protein from RHgX compounds. Clearly, Nature has invented a remarkable system to detoxify mercury in this fascinating class of Hg-resistant bacteria.
Cadmium and Lead Toxicity18
Gastrointestinal, neurological, and kidney toxicity are among the symptoms experienced by acute or chronic exposure to these heavy metals. The use of unleaded gasoline and the removal of lead-containing pigments from paint have substantially diminished the quantity of this element released to the environment each year. Cadmium sources include alkaline batteries, pigments, and plating. Lead poisoning can be treated by chelation therapy using CaNa2(EDTA) (acute) or penicillamine (chronic). Although both Cd(II) and Pb(II) bind to sulfhydryl groups in thionein, we have little information at the molecular level on the mechanisms by which these elements induce toxicity.
Metals as Carcinogens19,20
Although most metal ions have been reported to be carcinogenic, the three most effective cancer-causing metals are Ni, Cr, and, to a lesser extent, Cd. Nickel subsulfide, Ni2S3, found in many nickel-containing ores, has been extensively studied and shown to be carcinogenic in humans and other animals. In short-term bioassays including mutagenesis, enhanced infidelity of gene replication in vitro and altered bacterial DNA repair were observed. Chromium is most carcinogenic as chromate ion (CrO42-), which enters cells by the sulfate uptake pathway and is ultimately reduced to Cr(III) via a Cr(V)-glutathione intermediate species. The latter complex binds to DNA to produce a kinetically inert and potentially damaging lesion. Despite the fact that much information is available about metal-DNA interactions, molecular mechanisms of metal-induced carcinogenesis have not been elucidated. Two aspects of the problem are tumor initiation and tumor development, which are likely to involve different pathways. As new methods become available for studying the molecular events responsible for cancer (oncogenesis), it should be possible for bioinorganic chemists to unravel details of how metals act as carcinogens and as mutagens. Since cancer has genetic origins, metal/nucleic-acid chemistry is likely to be prominent in such mechanisms. As discussed later, metal-DNA interactions are an important aspect of the antitumor drug mechanism of cis-[Pt(NH3)2Cl2].
Summary
Toxicity can arise from excessive quantities of either an essential metal, possibly the result of a metabolic deficiency, or a nonessential metal. Both acute and chronic exposure can be treated by chelation therapy, in which hard-soft acid-base relationships are useful in the choice of chelating agent. Since chelates can also remove essential metals not present in toxic amounts, ligands with high specificity are greatly desired. The design and synthesis of such ligands for chelation therapy remains an important objective for the medicinal bioinorganic chemist. Until recently, studies of the toxic effects of metals and their removal, sometimes categorized under "environmental chemistry," have been empirical, with little insight at the molecular level. Application of the new tools of molecular biology to these problems has the potential to change this situation, as illustrated by rapid progress made in cloning the genes and studying the gene products of the mercury-resistance phenotype in bacteria. The discovery of such resistance phenomena in mammalian cells, and even the remote prospect of transferring Hg-resistant genes from bacteria to humans, are exciting possibilities for the future. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/09%3A_Metals_in_Medicine/9.02%3A_Toxic_Effects_of_Metals.txt |
Given that DNA is a major target of platinum binding in cells, it is incumbent upon the bioinorganic chemist to investigate the nature of these interactions and their biological consequences. Of all the ligands studied in coordination chemistry, DNA is surely among the most complex. In the ensuing discussion, we first present experiments that delineate the chemical steps involved in cis- and trans-DDP binding to DNA as well as the chemical consequences of the adducts formed. We next describe the physical changes in the double helix that accompany platinum binding, and then we discuss the biological consequences that attend the platination of DNA. Subsequent sections describe the major adducts formed, in other words the regiospecificity of the drug, the three-dimensional structures of the adducts, and the way in which different structures within DNA can modulate platinum binding. Finally, we consider the response of the cell to Pt-DNA adducts, including studies with site-specifically modified DNA, and speculate about how this chemistry might relate to the antitumor drug mechanism.
a. Kinetics of Platinum Binding to DNA
The binding of cis- and trans-DDP to DNA has been studied67 by 195Pt NMR spectroscopy with the use of isotopically enriched 195Pt, which has a nuclear spin I = $\frac{1}{2}$. The DNA used in this experiment was obtained from chicken red blood cell chromosomes that had been enzymatically degraded to relatively small pieces ranging from 20 to 60 base pairs in length (molecular-weight range 13 to 30 kDa). Since the 195Pt chemical shifts are very sensitive to chemical environment, this NMR study provided important details about the kinetics and mechanism of platinum binding to the biopolymer. The rate-determining step in platination of the DNA is loss of chloride ion (Equation 9.5) to form the monoaqua complex, which rapidly coordinates to a nitrogen donor on the nucleic acid. The identification of the coordinating atom as nitrogen was possible because the 195Pt chemical shift is characteristic of species having one chloride and three nitrogen ligands bound to Pt(II).67 The spectroscopic changes that accompany the formation of the family of monofunctional adducts are shown in Figure 9.8. Subsequent hydrolysis of the second chloride ion leads to the formation of a second bond with DNA. This sequence of events affords bifunctional adducts and is similarly accompanied by discrete 195Pt spectral changes (Figure 9.8). From the 195Pt chemical-shift range of the final products, it was apparent that the cis-{Pt(NH3)2}2+ moiety is bound primarily to two nitrogen donors on the nucleic acid. This chemistry is summarized in Equation (9.10), together with the half-lives for the mono-functional adducts. The half-lives were calculated from a kinetic analysis of the time-dependence of the 195Pt spectral changes. As can be seen, the rates of closure of mono- to bifunctional adducts for the two isomers are quite similar, suggesting that their different biological properties are not a consequence of the kinetics of binding to DNA.
$\tag{9.10}$
The next logical question to address is what donor atoms on DNA are coordinating to platinum in the mono- and bifunctional adducts. This important issue is discussed in considerable detail in Sections V.D.4 and V.D.5. As will be shown, the N7 atoms of the purine bases adenine and guanine are the principal binding sites. Alkylation of DNA at these positions facilitates depurination. Platinum binding to N7 atoms of purines (Figure 9.9), however, stabilizes the glycosidic (N9-C1') linkage.68-70 Presumably the positive charge is better distributed over the platinum atom and its ligands in the adduct than over a purine alkylated at N7. On the other hand, platinum binding to N7 of guanine does perturb the charge distribution in the purine ring, as evidenced by the lowering of the pKa of N1 by ≈ 2 units from its value in the unplatinated nucleotide (usually from pKa ≈ 10 to pKa ≈ 8).71,72 This effect has been used to assign platinum binding sites in DNA fragments, as discussed below.
What are the chemical changes at the platinum center when cis-DDP binds to DNA? Both chloride ions are lost from the coordination sphere, as already indicated. Platinum EXAFS studies of calf-thymus DNA modified with cis-DDP revealed no chlorine backscattering features characteristic of Pt-CI bonds.73 The spectra were consistent with the presence of four Pt-N/O linkages, since the technique is unable to distinguish between the two low-Z elements oxygen and nitrogen. Various studies reveal that, under most circumstances, the NH3 ligands are not lost from DNA upon the binding of platinum ammine halides. For example, when 14C-Iabeled cis-[Pt(NH2CH3)2Cl2] was allowed to bind to T7 (47 percent GC content) or M. luteus (73 percent GC content) DNA, no loss of radiolabel was found to accompany platinum binding.74 In vivo, however, loss of amine ligands has been observed. Injection of 195mPt and 14C doubly labeled [Pt(en)CI2] into tumor-bearing mice resulted in unequal distribution of the two labels in various biochemical fractions, but there is no reason to believe that this result is relevant to the antitumor mechanism.75 Metabolic inactivation of the drug could occur in a variety of ways unrelated to anticancer activity. The best evidence that ammine loss does not occur at the critical biological target of cisplatin is the finding, by using antibodies specific for cis-{Pt(NH3)2}2+ nucleotide complexes (see Section V.D.4.c), that DNA, extracted from cells in culture or from human cancer patients treated with cis-DDP and subsequently degraded, contains intact Pt-NH3 Iinkages.76
Once bonds are made between Pt and its targets on DNA, they are relatively inert kinetically. Platinum-DNA complexes can be subjected to various physical methods of separation and purification, including gel electrophoresis, ethanol precipitation, centrifugation, and chromatography, as well as to enzymatic and even chemical degradation procedures that digest the DNA, without releasing the platinum. Platinum can be removed, however, either by use of cyanide ion, to form the very stable (K ~ 1041) [Pt(CN)4]2- complex, or by excess thiourea.77,78 These properties have proved to be extremely valuable in facilitating localization and characterization of the major cis- and trans-DDP binding sites on DNA.
Although Pt-DNA linkages are, generally speaking, kinetically inert, sometimes a particular adduct will rearrange into a more stable linkage isomer. One interesting example is the product of the reaction of trans-DDP with the dodecanucleotide 5'-d(TCTACGCGTTCT).79 Initially the platinum coordinates to the two guanosine residues, forming a trans-[Pt(NH3)2{d(GCG)}] 1,3-intrastrand crosslink. This complex rearranges to a more stable trans-[Pt(NH3)2{d(CGCG)}] 1,4-intrastrand crosslink with a half-life of 129 h at 30 °C or 3.6 h at 62 °C. In this rearrangement product the platinum is coordinated to a cytosine and a guanosine residue.
As just described, the binding of bifunctional platinum complexes to DNA proceeds in a stepwise fashion. The second step is sufficiently slow (a few hours), however, that various reagents such as NH3, nucleobases, and low concentrations of thiourea can coordinate to the fourth site and trap the monoadducts. Generally speaking, however, given sufficient time both cis- and trans-DDP will bind DNA in a bifunctional manner. As such, they bear some resemblance to organic alkylating agents, such as the nitrogen mustards, which have been employed as anticancer agents.80
b. Crosslinking Reactions of Platinum Complexes
There are three broad classes of DNA adducts that can be made by bifunctional platinum complexes. As illustrated for cis-DDP in Figure 9.10, they are DNA-protein crosslinks, interstrand DNA-DNA crosslinks, and intrastrand crosslinks,81 A fourth possibility for platinum complexes is bidentate chelate ring formation utilizing two donor atoms on a nucleotide. For many years, a favored such postulated mode of binding was chelation by the N7-O6 positions of the guanine base (Figure 9.9), since this structure could be formed only by cis- and not by trans-DDP.82,83
Such a structure has never been observed for cis-DDP binding to DNA, however. DNA-protein and interstrand crosslinks formed by platinum complexes have been the focus of many attempts to explain cytotoxicity and antitumor behavior.84 The technique of alkaline elution, in which crosslinked DNA-DNA strands or DNA-protein molecules bind to filter paper following denaturation under basic conditions, sensitively and easily reveals such adducts. trans-DDP forms such adducts more rapidly than the cis isomer, perhaps because of its faster chloride-ion hydrolysis rates (see above) and a more favorable geometry, but they also seem to be repaired more rapidly in cells. As will be shown, interstrand and DNA-protein crosslinks are a small minority of adducts formed by cisplatin, and their contribution to the cytotoxic and anticancer properties of the drug remains to be established. In studies of SV40 replication in vivo, DNA-protein crosslinking by cis- and trans-DDP was shown not to be correlated with the inhibition of DNA replication.85
What proteins form crosslinks to DNA? One possibility is the histones that make up the spools around which DNA is wound when packaged into chromatin in the nucleus. Studies86 of cis- and trans-DDP binding to nucleosome core particles (each particle made up of eight histone proteins; around each particle is wound a 146-bp piece of DNA in a shallow superhelix of 1.75 turns) revealed the DNA binding to be little affected by the protein core. Both DNA-protein and specific histone crosslinked species were observed; from the latter it was suggested that DDP complexes might be useful crosslinking probes of biological structures. Other proteins likely to form crosslinks to DNA in the presence of platinum complexes are DNA-processing enzymes, or enzymes requiring a DNA template for normal function. In the in vivo SV40 study, for example, T antigen was one of the proteins found to be crosslinked to SV40 DNA by cisplatin.85 Other nuclear proteins such as the high-mobility group (HMG) class are also crosslinked to DNA in the presence of cis-DDP. In all cases so far, DNA-protein crosslinking has occurred when platinum was added to cells. There is as yet no evidence that transfection of platinated DNA into cells results in such crosslinking or that crosslinks form during in vitro enzymatic digestions of platinated DNAs.
c. Physical Effects of Platinum-DNA Binding
(i). Unwinding, Shortening, and Bending of the Double Helix
Early studies of cis- and trans-DDP binding to DNA employed closed and nicked circular plasmids.87 As was described in more detail in Chapter 8, closed circular DNAs are topologically constrained such that any change in the number of helical turns must result in an equal and opposite number of superhelical turns. Any reagent that unwinds the double helix reduces the number of helical turns. Consider, for example, a stretch of DNA that is 360 base pairs (bp) long. Normal B-DNA has ≈ 10.5 bp per turn or a helical winding angle of ≈ 34.3° per bp. Suppose the DNA is unwound, so that there are now 12 bp per turn or a winding angle of ≈ 30°. Instead of 34.3 helical turns (360 ÷ 10.5), the DNA now has only 30 (360 ÷ 12). If this DNA molecule were in the form of a covalently closed circle, the helical unwinding of -4.3 turns would be accompanied by a superhelical winding of +4.3 turns.
Planar organic dyes such as ethidium bromide (EtdBr) and inorganic complexes such as [Pt(terpy)(HET)]+ (Figure 9.11) bind to DNA by intercalation, inserting between the base pairs and unwinding the double helix by ~ 26° per molecule bound (Figure 9.12).88 This unwinding can be measured by monitoring changes in the superhelicity of closed circular DNA. This kind of DNA is subjected to certain topological constraints that lead to the formation of supercoils and superhelical winding that dramatically alter the hydrodynamic properties of the DNA. Either gel electrophoresis or analytical ultracentrifugation can be used to measure this phenomenon. The platinum complexes cis- and trans-DDP also produce changes in the superhelix density when bound to closed circular DNA.87 As shown in Figure 9.13, increasing concentrations of platinum bound per nucleotide on the DNA first retard its mobility and then increase its mobility through the gel. These interesting alterations in gel mobility occur because the negatively coiled superhelix unwinds first into an open, or untwisted, form and then into a positively supercoiled form. The conformational changes, which are depicted in Figure 9.14, are directly proportional to the drug-per-nucleotide, or (D/N)b, ratio. In addition to superhelical winding, both platinum complexes increase the mobility of nicked circular DNA in the gels (Figure 9.13). Nicked DNA has one or more breaks in the sugar-phosphate backbone, which relieve the topological constraint and prohibit the DNA from twisting into superhelical structures.
What could be the cause of these physical changes in the DNA structure upon cis- or trans-DDP binding? Intercalation can be excluded, not only because the compounds do not have the aromatic character normally associated with intercalators (Figure 9.11), but also through studies of the manner by which these and other platinum complexes inhibit the intercalative binding of EtdBr to DNA.89,90 Platinum metallointercalators such as [Pt(terpy)(HET)]+ are competitive inhibitors of EtdBr binding, as measured by fluorescence Scatchard plots, whereas the non-intercalators cis- and trans-DDP are not. Moreover, intercalation tends to lengthen and stiffen the double helix, whereas the mobility changes of nicked circular DNAs upon binding of cis- or trans-DDP were shown by electron microscopy experiments to arise from a pronounced shortening of the DNA with increased Pt binding.
One manner by which cis- or trans-DDP might produce these physical alterations in DNA structure is by kinking the double helix at or near the binding site. Such an effect could be produced by the bidentate attachment of platinum; the monofunctional [Pt(dien)Cl]+ complex does not have these pronounced effects on DNA secondary structure.91 Recently, it has been demonstrated that cis-DDP binding to DNA does indeed produce a pronounced bend in the helix axis.92,93 The proof employed a gel-electrophoretic method of analysis that had previously been used to study DNA bending at naturally occurring specific sequences called A-tracts, consisting of five or six adenosine nucleosides in a row followed by about the same number of thymidine residues.94 When these d(A5T5)2 sequences are positioned in the center of a DNA restriction fragment of, say, 150 bp, the mobility of the DNA through polyacrylamide electrophoresis gels is greatly retarded compared to that of a similar DNA fragment where the A-tract is at the end. For the former fragment, the bent molecules presumably cannot snake their way through the pores of the polyacrylamide as well as the molecules whose bends are at the ends and have little effect on the linear structure. It was further shown that A-tracts bend the duplex toward the minor groove of the DNA. Moreover, in a DNA containing multiple A-tracts, the bends must be separated by integral numbers of helical turns (~ 10.5 bp) or else the effect will cancel and the gel mobility will be that of normal DNA of similar length. This latter phenomenon has been referred to as phasing.
With this background information in mind, we can now discuss the experiments with cis-DDP that demonstrated bending.92,93 By methods described in Section V.D.8, a 22-bp oligonucleotide (22-mer) containing self-complementary overhanging ends ("sticky ends") was synthesized with a single cis-diammineplatinum(II) moiety linking adjacent guanosine residues (Figure 9.15A). A 22-mer was chosen since it has approximately two helical turns, accounting for some platinum-induced unwinding, and will thus have phased bends when polymerized. This platinated DNA was then labeled with 32P and treated with the enzyme DNA ligase, which seals the ends, producing oligomers of the 22-mer having lengths 22, 44, 66, 88, 110, etc., bp. In these oligomers, the platinum atoms are spaced apart approximately by integral numbers of helical turns. As shown in Figure 9.15, studies of this family of oligomers by gel electrophoresis revealed a pronounced retardation compared to the mobility of unplatinated DNA oligomers of comparable size (line P22 in Figure 9.15B). The plots in this figure show the relative mobilities (RL) of the different length multimers, compared to a control in which the top strand is not platinated, as a function of the length in base pairs. From the resulting curves may be extracted the extent of cooperative bending. When oligomers of a platinated DNA fragment in which the metal atoms were spaced apart by 27 bp were examined, their relative mobilities were found to be nearly the same as unplatinated control molecules (line P27 in Figure 9.15B). These experiments unequivocally established that platinum kinks the double helix. As with A-tract-induced bends, the platinum atoms must be phased in order to induce cooperative bending. Comparison of the magnitudes of the gel mobility changes made it evident that cis-{Pt(NH3)2}2+ binding produces a bend comparable to that of two A-tracts, ≈ 34°.
In a related series of experiments,93 the platinated 22-mer was copolymerized with various A-tract-containing Il-mers to produce ladders of oligomers in which the phasing of Pt with respect to the center of the A-tract was varied, but the Pt atoms were always in phase. The results of these studies showed that maximum gel-mobility retardation occurred when the Pt and A-tract center were spaced apart by half-integral numbers of helical turns (Figure 9.15C). Since A-tracts bend the DNA into the minor groove, this result implies that platinum bends the DNA into the major groove. Only when phased by n/2 (n = integer) turns will copolymers of species situated alternatively in the major and minor grooves of DNA exhibit such cooperative bending. It will be shown later that helix bending of cis-DDP-DNA adducts into the major groove is in accord with their known structures.
The ability to prepare site-specifically platinated oligonucleotides (see Section V.D.8) has provided a means for measuring the extent to which cis-DDP produces local unwinding of the double helix.95 When the platinum atoms are positioned with respect to one another, or phased, by exactly integral numbers of helical turns, the retardation of the DNA multimers in the gel is maximized. This phenomenon is illustrated in Figure 9.16, where the RL values are plotted as a function of the interplatinum spacing for oligonucleotides containing the cis-{Pt(NH3)2d(GpG)} intrastrand crosslink. When the resulting curve was analyzed, the maximum was found to occur at 21.38 bp. Since normal B-DNA has a helical repeat of 10.5 bp, one can compute the effect of platination from the expression [(21.38 - 2(10.5)] bp = 0.38 bp. From the fact that one helical turn of DNA comprises 360° and 10.5 bp, the unwinding of the DNA double helix due to the presence of a single cis-{Pt(NH3)2d(GpG)} intrastrand crosslink can be calculated as
$\left(\dfrac{0.38}{10.5}\right) \times 360 = 13^{o} \ldotp$
Similar studies of DNA platinated with trans-DDP have been carried out. In these, oligonucleotides containing the 1,3-trans-{Pt(NH3hd(GpNpG)} intrastrand crosslink were examined. The electrophoresis gels of polymerized 15-mers and 22-mers containing this adduct showed cooperative bending. This result indicates that bends at the sites of platination by trans-DDP are not phase sensitive, and has been interpreted to imply the formation of a "hinge joint" at these positions.92,95 The directed bends and local unwinding of DNA produced by cisplatin could be an important structural element that triggers a response by cellular proteins. This subject is discussed in greater detail in Section V.D.7.d.
d. Biological Consequences of Platinum-DNA Binding
(i). Inhibition of Replication
Binding of cis-DDP to DNA inhibits replication both in vivo and in vitro, as shown by a variety of assays. Inhibition of replication of SV40 viral DNA in African green monkey cells as a function of the concentration of added cis-DDP is shown in Figure 9.17. When SV40 virus infects monkey cells, it does not integrate its DNA into the genome of the host. Instead, it forms its own chromosomes in the cell nucleus. These so-called mini-chromosomes consist of ≈ 20 nucleosomes, fundamental chromosome building blocks. SV40 has its own life cycle, using virally encoded and cellular proteins to replicate and, ultimately, reassemble virus particles before lysing the cell and departing to infect neighboring cells.
In the experiment shown in Figure 9.17, the SV40-infected cells were treated with cisplatin. After 24 h, [3H]thymidine was added and, after 24 more hours, the cells were harvested, and SV40 DNA was isolated; the amount of DNA synthesis was recorded by comparing incorporated radiolabel with results from control experiments where no platinum was present. The data show that, when 25 $\mu$M platinum was present, SV40 DNA replication was reduced to about 5 percent of control. Quantitation reveals that, at ≈ 2 platinum atoms bound per thousand nucleotides (drug-per-nucleotide, or (D/N)b, = 0.002), synthesis is only 10 percent that of control.
Recently, a related series of experiments has been carried out that can monitor DNA synthesis from templates platinated in vitro.96 In this work, DNA plasmids containing the SV40 origin of replication are added to cellular extracts prepared from human kidney cells previously infected with adenovirus. In the presence of large T antigen, a virally encoded protein required for replication, SV40 DNA is synthesized from the plasmid templates. Synthesis can be conveniently monitored by [32P]dATP incorporation. At a (D/N)b ratio of only 1.7 x 10-3, DNA synthesis is about 5 percent of control, in agreement with the results of the in vivo study.
The binding of cis-DDP to DNA has also been measured for normal and tumor cells implanted in nude mice and in cells obtained from the ascites fluid of patients with ovarian carcinoma 24 h after their last dose.97 The data for mouse bone marrow and a human pancreatic tumor xenograft show that, at a dose of 10 mg/kg, (D/Nh platinum binding levels of 3.3 x 10-6 and 1.82 x 10-6 reduce survival to 20 and 10 percent of control, respectively. These ratios are in good accord with platinum levels required to inhibit DNA synthesis in mammalian cells, as revealed by various studies, but substantially less than that needed for replication inhibition in the SV40 experiments described above. The difference can be readily explained, illustrating an important point. The SV40 genome, like most other DNAs of viral or plasmid origin, consists of only 15,000 nucleotides whereas the nuclear DNA of mammals has about 109 nucleotides. Thus, (D/N)b levels of ~ 10-6 would leave 99 out of 100 SV40 DNA molecules with no platinum at all, and replication would hardly be affected. For the mammalian genome, (D/N)b values of 10-6 place 103 platinum atoms on each DNA genome, sufficient to inhibit replication and reduce cell survival. Platinum-DNA binding levels of this magnitude are found for ovarian ascites cells taken from patients receiving cisplatin chemotherapy.97
(ii). Mutagenesis and Repair
Apart from inhibition of DNA synthesis, what are the other biological consequences of cisplatin binding to DNA? One such consequence is mutagenesis, in which a normal base in the sequence is replaced by a different base. This phenomenon has been demonstrated for cis-DDP-treated cells in a variety of studies. What brings about such mutagenesis? There are several possibilities. One is that errors are introduced in DNA strands during attempts of the replication apparatus to synthesize past a platinum lesion. Another is that the platinum-damaged DNA is recognized by cellular repair systems that, in attempting to eliminate the platinated stretch of DNA, incorporate one or more incorrect nucleotides. Platinum-induced mutagenesis can lead to deleterious long-term health problems in patients treated with cisplatin. It is therefore important to understand the mechanism by which cellular DNA becomes mutated following platination, and to devise strategies for minimizing or eliminating this mutagenesis.
The foregoing considerations bring up another biological consequence of cis-DDP binding, namely, DNA repair. Removal of platinum from DNA by cellular repair mechanisms has been demonstrated by several groups. For example, in studying cis-DDP-treated human fibroblast cells in culture, it was found that the amount of bound platinum per nucleotide decreased according to first-order kinetics, from (D/N)b of 2.3 x 10-5 to 3.3 x 10-6 over a six-day period. Since Pt-DNA adducts are stable with respect to dissociation from DNA under physiological conditions (see above discussion), loss of platinum was attributed to DNA repair.98
How does the cell remove platinum from DNA? One mechanism is by a process known as excision repair, whereby the sugar-phosphate backbone on the platinated strand is hydrolyzed ("nicked") on either side of the damage and the remaining, unplatinated strand is used as a template for new DNA synthesis. The platinated oligonucleotide is displaced and the resulting gap filled in. In support of this picture is the fact that, in xeroderma pigmentosum (XP) human fibroblast cells, known to be deficient in excision repair, there is very little removal of platinum during post-treatment incubation.99 Recent studies of in vitro repair of cisplatin-DNA adducts by a defined enzyme system, the ABC excision nuclease of E. coli, have provided some details at the molecular level about the process.100,101 [32P]-Labeled double-stranded DNA fragments containing {Pt(NH3)2}2+ or {Pt(en)}2+ adducts at random or defined sites were incubated with the enzyme. Cleavage of the platinated strand occurred at the 8th phosphodiester bond 5', and the 4th phosphodiester bond 3', to the GG or AG intrastrand crosslink. Further details about the identification and construction of such specific crosslinks will be given later in this chapter.
(iii). Drug Resistance
Another biological consequence of DNA-platinum interactions, probably related to the repair phenomenon, is resistance. Resistance of a cell to a chemotherapeutic agent, which can be inherent or acquired, is a phenotypical ability of the cell to tolerate doses of a drug that would be toxic to normal, or parent, cells.102 Resistance is often acquired by prolonged exposure of cells in culture to the drug or, in patients, to repeated doses of drug therapy. There is not yet any direct proof that platinum-DNA interactions are responsible for acquired resistance to cisplatin. Studies of sensitive and resistant tumors in rats have shown, however, that after intravenous injection of 10 mg/kg of the drug, the platinum levels were the same after an hour, but after 24 hours a larger proportion of adducts had been removed in the resistant cells.103 Similar results have been found for studies of Pt-DNA adducts in cultured L1210 cells of varying levels of resistance to cisplatin where, in the 18 h following a 6-hour incubation with the drug, the resistant cells had up to fourfold more platinum removed than the sensitive cells.104
Experiments have also been carried out showing that cis-DDP binding to DNA inhibits transcription, the formation of RNA from a gene, and that this phenomenon is less efficiently reversed for parent versus resistant L1210 cells in culture.105 The assay involves transfection (the process whereby free or viral DNA or RNA is taken into a cell) of pRSVcat plasmid DNA into L1210 cells. The plasmid contains the bacterial cat gene in a position that permits its expression in mammalian cells. The cat gene encodes the enzyme chloroamphenicol acyltransferase (CAT), an activity readily measured following lysis of the cells. Transfection of the cis-DDP-damaged plasmid into resistant L1210 cells showed that up to eight times the amount of platinum was required in the resistant versus sensitive cells to produce a mean lethal hit (63 percent reduction in activity). This result is consistent with greater repair of platinum-DNA adducts in the resistant cells.
These results should not be construed to mean that DNA repair is the only mechanism of cisplatin resistance. There is evidence that relative amounts of glutathione are increased in cisplatin-resistant cells.106 Glutathione presumably uses its thiol moiety to coordinate platinum and diminish the amount that can bind to DNA. Reduced influx or increased efflux of a drug constitutes additional mechanisms by which cells become resistant. Further studies are required to ascertain which of these possibilities is most important for the cisplatin resistance phenomenon.
The discovery that cells can become resistant to cisplatin by repairing DNA lesions suggests a way to explain the selectivity of the drug for certain tumor tissue, and even the selective cytotoxicity of the drug for tumor versus normal cells of the same tissue. Tumor cells that cannot repair platinum-DNA adducts would be most affected by cis-DDP. This idea forms one of the central hypotheses about the molecular mechanism of action of cis-DDP, details of which can be probed by bioinorganic chemists. Specifically, it is important to inquire what DNA adducts formed by cis-DDP are both cytotoxic and repairable, what enzymes are responsible for such repair in mammalian cells, by what mechanisms these enzymes operate, and how this knowledge can be used to design better metal-based antitumor drugs and chemotherapeutic protocols.
(iv). DNA-protein Interactions
Most of the phenomena discussed in this section, inhibition of replication, DNA repair, drug resistance, and mutagenesis, probably involve interaction of a protein or group of proteins with platinated DNA. These interactions are clearly important in determining the biological consequences of DNA templates containing bound platinum. Very recent experiments have uncovered the existence of proteins from a variety of mammalian sources that bind specifically to DNA platinated with cis- but not trans-DDP.107 Identifying the nature and function of these factors may provide important clues about the mechanisms of antitumor activity, drug resistance, or repair. Study of protein-DNA-drug interactions is an essential feature of the bioinorganic chemistry of platinum chemotherapeutic agents. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/09%3A_Metals_in_Medicine/9.03%3A_Aspects_of_Platinum_Binding_to_DNA.txt |
Radiodiagnostic Agents21,22
Metal complexes having radioactive nuclei find many applications in medicine, such as in organ, and tissue imaging, Early detection of cancer, for example, by selective and imaging of the tumor using a radioactive metal compound can facilitate surgical removal or chemotherapeutic treatment before the disease reaches an advanced stage. radioisotopes used for diagnostic purposes be emit low-energy $\gamma$ and no $\alpha$ or $\beta$ particles. Table 9.2 lists the radionuclides most commonly employed for purpose in nuclear medicine. Among these, 99mTc is perhaps the most desirable,23 for it gives off a 140-keV $\gamma$ ray that is detected scintillation cameras and produces clear images. radionuclide is prepared from an alumina column loaded with 99MoO42-, which decays to form 99mTcO4-, which in turn may be selectively eluted from the column with saline owing to its lower charge. treatment with a reducing agent in the presence of the appropriate ligands produces radiopharmaceuticals with desired water stability, and properties. Such complexes may be injected at concentrations of 10-6-10-8 M. For example, isocyanide complexes such as [Tc(CNR)6]+= t-Bu, CH2CO2But, etc.) have been found to be taken up selectively into heart tissue and thus have the potential to be used as heart-imaging agents. Figure 9.2 displays bone as imaged a 99mTc bone agent. The dark correspond to surface areas of metabolic which can be used to diagnose or disease. One goal of research in this field is to images of myocardial infarcts or clogged arteries for physicians who can watch the patient's heart on a video surgery. Although chemical details responsible for the selective tissue of Tc isocyanide, and other complexes are largely synthetic modifications are and have many new compounds for evaluation.
Table 9.2 - Radionuclides most commonly employed in diagnostic nuclear medicine.a
a) Data are from Table of the Isotopes in D. R. Lide, ed., CRC Handbook of Chemistry and Physics, CRC Press, 71 st ed., 1990-91, pp. 11-33 ff.
Radionuclide Half-Life Energy (keV)
57Co 271 d 836
67Ga 78 h 1,001
99mTc 6 h 140
111In 67 h 172, 247
113mIn 104 m 392
123I 13 h 1,230
169Yb 32 d 207
197Hg 64 h 159
201Tl 72 h 135, 167
Among the few known to be absorbed selectively tumor cells is the antitumor antibiotic bleomycin (BLM),24,25 the structure of which is portrayed in Figure 9.3. binds most radioactive metal ions, but the 57Co(III) complex has the best tumor-to-blood ratio. Unfortunately, the long 57Co (Table 9.2) has limited its clinical utility. Attempts to prepare 99mTc complexes of BLM with selective uptake properties approaching that of the cobalt have not yet been successful, although the target molecule would be a most valuable radiodiagnostic agent. One imaginative solution26 to this problem was achieved by covalent attachment of an EDTA moiety to the terminal thiazole ring of BLM (Figure 9.3). The resulting Co(III) BLM-EDTA molecule was radiolabeled with 111In3+ and found to be useful for diagnosis of cancer in humans. Also used for tumor imaging are 99mTc and 67Ga citrate complexes, the latter being the agent of choice for many applications. Again, there is little known at the molecular level about the mechanism of tumor-cell specificity
An alternative approach to radionuclide-based tumor-imaging agents for diagnosis of disease is to modify, with metal chelating agents, antibodies raised against a biological substance, such as a tumor-cell antigen, hormone, or other target. Antibodies are proteins that are synthesized by specialized cells of the immune system in response to an external stimulant, or antigen. The high specificity and affinity of antibodies for the antigen can be used to target the antibody to a particular biological site, such as a site on the membrane of a particular cell type. Chelating agents are now routinely attached to antibodies and used to bind radioactive metal ions. The resulting radionuclide-labeled products are currently under extensive study in diagnostic medicine.26
Magnetic Resonance Imaging (MRI)27
Nuclear magnetic resonance (NMR) spectroscopy can be used to image specific tissues of biological specimens because of differences in the relaxation times of water proton resonances, usually brought about by paramagnetic metal ions. An early, pioneering example was the demonstration that Mn(lI) salts localize in normal heart-muscle tissue in dogs rather than in regions affected by blocked coronary arteries. Since the paramagnetism of the d5 Mn(lI) ions alters the relaxation rate of nearby water protons, the normal and diseased tissue could be distinguished. Of the various metal ions surveyed in attempts to provide clinically useful NMR images in humans, Gd(III), Fe(III) , and Mn(lI) were found to give the best proton-relaxation enhancements. The gadolinium complex [Gd(DTPA)(H2O)]2-,an agent currently used in the clinic, has been successfully employed to image brain tumors. Ferric chloride improves gastrointestinal tract images in humans and, as already mentioned, manganous salts can be used for heart imaging. NMR imaging methodologies have advanced to the stage where increases as small as 10 to 20 percent in T1-1, the inverse nuclear-spin relaxation time, can be detected. As with radionuclide labeling, the complexes must be soluble and stable in biological fluids and relatively nontoxic, and are of greatest value when able to target a specific tissue. Even more important than targeting, however, is that proton relaxivity be maximally enhanced, an objective that depends not only upon the local binding constant but also upon large magnetic moments, long electron-spin relaxation (Tle) values, access to and the residence lifetime in the inner and outer coordination spheres by water molecules, and the rotational correlation time of the complex at its binding site. An obvious advantage of paramagnetic NMR over radioisotopic imaging agents is that there is no possibility of radiation damage; on the other hand, the need for 10-100 $\mu$M concentrations at the site of imaging is a distinct drawback. Both methods are likely to continue to be used in the future, and both will benefit from the design of new stable chelates that are selectively absorbed by the tissue to be diagnosed.
Lithium and Mental Health28-31
One in every 1,000 people in the United States currently receives lithium, as Li2CO3, for the treatment and prophylaxis of manic-depressive behavior. Doses of 250 mg to 2 g per day are administered in order to maintain a 0.5 to 2.0 mM concentration window, outside of which the drug is either toxic or ineffective. The detailed molecular mechanism by which Li+ ion brings about its remarkable chemotherapeutic effects is largely unknown, but there are various theories. One theory proposes that lithium binds to inositol phosphates, inhibiting their breakdown to inositol, and so reducing inositol-containing phospholipids. A consequence of this chain of events would be disruption of the neurotransmission pathway based on inositol 1,4,5-triphosphate and 1,2-diacylglycerol, reducing neuronal communication, which is most likely hyperactivated in the manic state. This theory does not account for the antidepressive action of the drug, however. An alternative explanation is that lithium inhibits cyclic adenosine monophosphate (AMP) formation, again interfering with neurotransmission by intercepting this key intracellular signaling molecule. Recent experiments indicate that lithium affects the activation of G-proteins, a class of guanosine triphosphate (GTP)-binding proteins involved in information transduction. Possibly these effects result from displacement by Li+ of Mg2+ from GTP and/or from protein-binding sites normally required for activation. Use of 7Li NMR spectroscopy to study lithium transport in human erythrocytes suggests that it might be possible to apply this method to unravel details of the bioinorganic chemistry of lithium associated with the management of manic depression.
Gold and Rheumatoid Arthritis23,32,33
Gold compounds have been used in medicine for centuries, an application known as chrysotherapy. Since 1940, however, complexes of gold have been used most successfully to treat arthritic disorders in humans and other animals. Au(I) compounds are currently the only class of pharmaceuticals known to halt the progression of rheumatoid arthritis.
Until recently, gold compounds used to treat arthritis were painfully administered as intramuscular injections. Included were colloidal gold metal, colloidal gold sulfides, Na3[Au(S2O3)2] (Sanocrysin), gold thiomalate and its sodium and calcium salts (Myochrisin), and polymeric gold thioglucose (Solganol, approved by the FDA). It was discovered, however, that triethylphosphinegold(l) tetra-O-acetylthioglucose (auranofin, Figure 9.4, approved by the FDA) was equally effective against rheumatoid arthritis and could be orally administered. The availability of this compound has sparked many studies of its biodistribution, stability, and possible metabolism that lead to antiarthritic activity. The mode of action of antiarthritic gold drugs is largely unknown, but it may involve binding of Au(I) to protein thiol groups, a process that inhibits the formation of disulfide bonds, and could lead to denaturation and subsequent formation of macroglobulins.
Anticancer Drugs
1. Platinum Ammine Halides34,35
The discovery that cis-diamminedichloroplatinum(Il), cis-DDP or cisplatin (Figure 9.4), has anticancer activity in mice, and its subsequent clinical success in the treatment of genitourinary and head and neck tumors in humans, constitutes the most impressive contribution to the use of metals in medicine. Given in combination chemotherapy as an intravenous injection together with large amounts of saline solution to limit kidney toxicity, cisplatin treatment results in long-term (>5 yr) survival for more than 90 percent of testicular cancer patients. In a typical course, ~ 5 mg/kg body weight of the drug is administered once a week for four weeks. Extensive studies of platinum ammine halide analogues led to a series of empirical rules governing their chemotherapeutic potential. Specifically, it was concluded that active compounds should:
1. be neutral, presumably to facilitate passive diffusion into cells;
2. have two leaving groups in a cis configuration;
3. contain nonleaving groups with poor trans-Iabilizing ability, similar to that of NH3 or organic amines;
4. have leaving groups with a "window of lability" centered on chloride.
These early structure-activity relationships have had to be modified somewhat, however, since chelating dicarboxylate ligands such as 1,1-dicarboxylatocyclobutane can replace the two chloride ions, and since cationic complexes with only one labile ligand, specifically, cis-[Pt(NH3)2CI(4-X-py)]+, where X = H, Br, CH3, etc., showed activity in some tumor screens. The two compounds shown in Figure 9.4, cisplatin and carboplatin (Figure 9.4), were the first to be approved for clinical use. Of particular interest to the bioinorganic chemist is that complexes having a trans disposition of leaving groups are inactive in vivo. This difference suggests the presence of a specific cellular receptor that, when identified, should facilitate the design of new, metal-based anticancer drugs. Present evidence strongly points to DNA as being the relevant cellular target molecule. Section V of this chapter expands on this topic in considerable detail.
2. Metallocenes and Their Halides: Ti, V, Fe36,37
Several compounds in this category, including [(C5H5)2TiX2] (X = CI, Br, O2CCl3), [(C5H5)2VCl2], [(C5H5)2NbCl2], [(C5H5)2MoCl2], and [(C5H5)2Fe]+ salts, exhibit significant activity against experimental animal tumors. Higher quantities (200 mg/kg) of these compounds than of cis-DDP can be tolerated with fewer toxic side effects, but their failure in two mouse leukemia screens commonly used to predict the success of platinum anticancer agents appears to have delayed their introduction into human clinical trials. Studies of Ehrlich ascites tumor cells treated with [(C5H5)2VCl2] in vitro revealed selective inhibition of incorporation of radiolabeled thymidine, versus uridine or leucine, indicating that the complex blocks DNA replication. Unlike cisplatin, however, metallocene halides undergo rapid hydrolysis reactions in aqueous media, forming oxobridged and aqua complexes that may have a higher affinity for phosphate oxygen atoms than the heterocyclic nitrogen atoms of the bases in DNA.38 Exactly how the ferrocenium ion might bind to DNA is even more obscure, although partial metallointercalation and groove binding are more likely than covalent attachment of the chemically unmodified cation. From the limited information available, metallocenes and their halides appear to behave fundamentally differently from platinum antitumor compounds. As a class, they provide a promising new opportunity to expand the scope of metal complexes used in cancer chemotherapy.
3. Gold and Other Metal Phosphines39
Following the successful entry of the soluble gold-phosphine complex auranofin (Figure 9.4) into the metal-based pharmaceuticals industry, several gold-phosphine complexes were examined for possible anticancer activity. Although auranofin itself was active in only a small fraction of the mouse tumor models tested, biological activity approaching that of cisplatin was discovered for many analogues, most notably the diphosphine bridged complex [ClAu(PPh2CH2CH2 PPh2)AuCI]. Attempts to replace the phosphine with As or S donor ligands, to increase or decrease the length of the 2-carbon bridge, or to replace the phenyl with alkyl groups all led to diminished activity. Most noteworthy is that the diphosphine ligands themselves have activity very similar to that of their gold complexes, and that Ag(I) and Cu(I) analogues are also effective. These results strongly imply that the phosphine ligands are the chemical agents responsible for the anticancer properties of these compounds. Coordination to a metal presumably serves to protect phosphines against oxidation to the phosphine oxides, which independent investigations have proved to be ineffective. A possible role for the metal in the cytotoxicity of the compounds cannot be ruled out, however.
4. Other Main Group and Transition-metal Compounds36,40,41
Several main group metal complexes exhibit anticancer activity. Gallium(III) nitrate is active against human lymphomas, but with dose-limiting side effects on the kidneys and gastrointestinal tract. Tin complexes of general formula R2L2SnX2, where R = alkyl or phenyl, L2 = py2, bpy, or phen, and X2 = two cis-oriented halide or pseudohalide leaving groups, are active against the mouse P388 leukemia tumor. The cis disposition of the leaving groups suggests a possible mechanism analogous to that of cisplatin (see below). Organo-gennanium compounds are also active, notably the derivative spirogermanium shown in Figure 9.4. Nothing is known about the mechanism of action of any of these compounds.
Following the discovery of activity for cisplatin, several thousand platinum and nearly 100 other transition-metal complexes have been screened in various tumor model systems in the hope of achieving better activity against a broader range of tumors. Among the classes of nonplatinum compounds showing some activity are ruthenium complexes cis-[RuCl2(DMSO)4], [Ru(NH3)5(Asc)](CF3SO3), where Asc is ascorbate dianion, and fac-[Ru(NH3)2Cl4] , all of which are believed to bind to DNA; binuclear rhodium complexes [Rh2(O2CR)4L2]; octahedral Pd(IV) complexes such as cis-[Pd(NH3)2CI4]; and such miscellaneous molecules as the iron(II) complex of 2-formylpyridine thiosemicarbazone, the site of action of which is thought to be ribonucleotide reductase. These examples illustrate the broad scope encompassed by this field, which has a potential for developing fundamental infonnation about metal-biomolecule interactions as well as novel anticancer drugs. Much remains to be explored.
Miscellaneous Metals in Medicine
Numerous other anecdotal and some fairly elaborate studies have been reported for metal complexes as medicinal agents. The use of zinc applied topically to promote the healing of wounds dates back to around 1500 B.C., and silver is now commonly applied to prevent infection in bum patients.42,43 Osmium carbohydrate polymers have been reported to have antiarthritic activity.44 Transition-metal complexes have a long history of use as antibacterial and antiviral agents; for example, Zn2+ is used to treat herpes, possibly by inhibiting the viral DNA polymerase.45 Early transition-metal (e.g., tungsten) polyoxoanions have been employed to treat AIDS patients.46 Numerous reports have appeared detailing the anti-inflammatory, antiulcer, and analgesic activities of copper carboxylate complexes.7 As in the previous section, these reports and others like them require more serious attention from bioinorganic chemists to elucidate the molecular events responsible for such a fascinating menu of biologically active metal complexes.
Summary and Prospectus
The clinical successes of platinum anticancer and gold antiarthritic drugs have changed the attitudes of many who doubted that heavy-metal compounds, notorious for their deleterious effects on human health, would ever playa serious role in chemotherapy. Indeed, we have seen that Hg2+, Pb2+, and Cd2+ are toxic elements. Even essential metals can be highly toxic if present in excess, either because of chronic or acute poisoning or because of metabolic defects that deregulate their control in the cell. An important common theme running throughout this discussion is selectivity. For a drug to be effective, it must be selectively toxic to diseased tissue while leaving nonnal tissue alone; or it must selectively kill harmful microorganisms at levels where it fails to deplete helpful ones. For a chelating agent to be useful in the toxic effects of metals, it must bind as selectively as possible to the deleterious ion while coordinating only weakly, if all, to others. For a diagnostic metal complex to be it must be taken up (or excluded) selectively from diseased cells relative to normal ones, or to one tissue type versus another. Rarely has such selectivity been designed in advance of the discovery of a useful metal-based pharmaceutical, although spectacular advances in biology, such as monoclonal antibodies, may be hastening the day when such an objective be common. Interestingly, the successes of such unlikely as cis-[Pt(NH3)2Cl2] and [(Et3P)Au(OAc4-thioglucose)] in chemotherapy were driven by the personal involvement of individuals like B. Rosenberg for the former and B. Sutton for the latter. Like Hollywood producers, these men mustered every conceivable resource to promote the compounds for testing, introduction into human clinical trials, and eventually approval by the FDA. Such zeal requires years, usually more than a decade, of sustained personal effort, and may be the reason why other metal complexes, such as those mentioned above, have not had the impact of a cisplatin or an auranofin. On average, only one of 7,000 such compounds makes it from the laboratory bench to the patient, at an average cost of 250 million dollars and a time interval of 13 years.
Another component of the evolving field of metals in medicine, however, is that, once a has proved its in the clinic, how does it work? This question is deceptively for coordination chemistry in vivo, and the of cells to respond to unnatural external stimuli such as metal complexes, are matters about which we are beginning to learn. As progress is made in this latter area, it should become possible to design drugs in a rational way to achieve the required selectivity.
The remainder of this chapter focuses on a case study where some progress in unraveling the molecular mechanism of a metal-based drug, cisplatin being made. If nothing else, this discussion will elucidate strategic guidelines that may be employed to attack similar questions about other chemotherapeutic metal compounds discussed earlier in this section. Unfortunately, there is very little information available about the molecular mechanisms of these other complexes. At this transition in our discussion, we move from general considerations to a specific, analysis. The reader must here take time to become familiar with the biological aspects of the new material. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/09%3A_Metals_in_Medicine/9.04%3A_Survey_of_Metals_Used_for_Diagnosis_and_Chemotherapy.txt |
History of the Discovery47
cis-[Pt(NH3hCh], a molecule known since the mid-19th century, has been a subject of considerable importance in the of inorganic stereochemistry and substitution reaction kinetics.48 Its biological was discovered by accident. In the mid-1960s, biophysicist Barnett Rosenberg, at Michigan State University, was studying the effects of electric fields on the growth of Escherichia coli cells in culture. They had hypothesized that, during cell division, the orientation of the mitotic be affected by local electric which they hoped to perturb. Instead, they observed growth without cell division, the result being elongated, spaghetti-like bacterial filaments approaching 1 cm in length. After much detective work, they realized that small amounts of platinum from the electrodes used to apply the electric fields had reacted with NH4CI in their buffer to produce various platinum ammine halide compounds. Two of these, cis-[Pt(NH3)2Cl2] and cis-[Pt(NH3)2CI4], were capable of inducing filamentous growth in the absence of any electric field. Since chemicals that produce filamentation in bacteria had been known to exhibit antitumor activity, Rosenberg was eager to have his platinum compounds tested. Unable to convince existing agencies like the National Cancer Institute (who to their credit later spearheaded the development of cisplatin) that a heavy-metal complex could actually be beneficial to animals, the Michigan State group set up their own animal-tumor screens. The results were nothing short of spectacular. Injection of cis-DDP directly into the abdominal cavity of mice into which a solid Sarcoma-180 tumor had been implanted led within a few days to a blackening (necrosis), reduction in size, and eventual disappearance of the tumor (Figure 9.5). The cured mouse enjoyed a normal lifespan. From these and other animal studies, physicians became convinced that administering platinum compounds to cancer patients might be worthwile, and a new field involving bioinorganic chemistry and cancer chemotherapy was born. The drug, marketed as Platinol with the generic name cisplatin, received FDA approval in 1979 and is today one of the leading anticancer agents.
Principles that Underlie Drug Development
1. Strategic Considerations
There are two general routes to the development of inorganic complexes, and indeed most chemical compounds, as drugs. One, illustrated by cisplatin, arises from an empirical observation of biological activity followed by attempts to optimize the efficacy through investigations of structure-activity relationships (SAR). The goals are to minimize toxicity, to develop cell culture and animal screens for testing related compounds, and ultimately to elucidate the molecular mechanism. Knowledge of the molecular mechanism might evenlead to a rational strategy for designing better drugs.
The second general approach to drug design begins with known biochemistry. For example, ribonucleotide reductase is required in the first committed step in the biosynthesis of DNA, the conversion of ribo- to deoxyribonucleoside dir)hclsphat.es. The mammalian enzyme contains a binuclear non-heme iron center required for activity. Compounds that would selectively inhibit this enzyme by destroying this center are potentially useful as antiviral or antitumor agents. Another example is the enzyme reverse Iranscriptase, encoded by the HIV (AIDS) virus and required for its integration into the genome of the host cell. Compounds like 3'-azidothymidine (AZT) are accepted by the enzyme as substrates but, when added to the growing DNA chain, cannot be linked to the next nucleotide. Chain termination therefore occurs and the replication process becomes permanently interrupted. Attempts to find organic molecules or inorganic complexes that are more effective chain terminators than AZT constitute a rational strategy for developing new anti-AIDS drugs.
In the remainder of this chapter we describe research that has evolved following the discovery of biological activity for cisplatin. Although the initial breakthrough was serendipitous, subsequent studies have revealed many aspects of the molecular mechanism. From this known biochemistry we may one day be in a position to design more effective anticancer drugs and therapies based upon the fundamental bioinorganic chemistry of cisplatin.
2. Pre-clinical and Clinical Trials49
Predicting the chemotherapeutic potential of an inorganic compound such as cis-DDP prior to its introduction into human cancer patients is an important objective. Compounds are most easily tested for their cytotoxic effects on bacterial or mammalian cells in culture. Shown in Figure 9.6 are results for the survival of cultured L1210 cells in the presence of increasing amounts of cis- or trans-DDP.50,51 The data reveal the markedly greater toxicity of the cis isomer, which is a much better anticancer agent than its stereoisomer. Unfortunately, no single assay has yet been found that can predict the chemotherapeutic potential of platinum compounds in humans. The best that can be obtained are results relative to those for cis-DDP, in which case toxicity at low dose is usually scored positive.
The next level of testing, often employed directly without first examining cell-culture results, involves animal (usually mammals, excluding human) screens.49 Among the most popular measures of the chemotherapeutic activity of platinum compounds has been their ability to prolong the survival of mice bearing the L1210 or P388 leukemia. A suspension of cells is inoculated intraperitoneally (i.p., in the abdominal cavity), producing a leukemia that eventually progresses to the generalized disease. In one commonly used protocol, platinum compounds are dissolved in physiological saline (0.85 percent NaCl) or sterile H2O and injected i.p. 24 h, 5, 9, and 13 days after inoculation of the leukemia cells. Several indices of antitumor activity and toxicity have been defined. The percent I.L.S., or increased lifespan measures the mean survival of treated versus control animals that were given no platinum drug. A related index is the median survival, percent T/C (Test/Control), which is 100 + percent I.L.S. The LD50 value measures toxicity as mean lethal dose,the amount of drug (usually in mg/kg body weight) required to kill half the animals. Potency is defined by lD90, the inhibiting dose at which 90 percent of the tumor cells are killed. From these values, a therapeutic index (TI) = LD50/lD90 is sometimes defined, which should be substantially greater than one. Typical values for cis-DDP are 85 percent I.L.S. at 8 mg/kg for the L1210 tumor, 13.0 mg/kg LD50, and 1.6 mg/kg lD90 resulting in a TI of 8.1.
In addition to being tested in mice, cisplatin and related compounds have been screened in other mammals, specifically dogs and monkeys, mainly to look for possible dose-limiting side effects. Severe vomiting, once thought to be an insurmountable obstacle, was monitored by using ferrets. None of the animal screens can substitute for the ultimate test, however, which is human clinical trials. In 1972, such trials were initiated using terminally ill cancer patients. It was determined that intravenous (i.v.) injection, rather than i.p. or oral administration, was the preferred method for giving the drug. Further details of the clinical development of cisplatin are discussed in a later section.
From the animal screens emerged the set of structure-activity relationships enumerated earlier (Section IV.E.1). Both cisplatin and carboplatin conform to these rules, and to date no compounds with demonstrably better antitumor activity have been tested in humans. The decision to move an experimental drug into the clinic is a difficult one, however, and it may be that molecules such as cis- [Pt(NH3)2(4-Br-py)Cl]Cl (see Section V.D.7.c) would be effective for tumors that are refractive to cisplatin chemotherapy. In any case, the foregoing chain of events, from studying the effects of a compound on cells in culture through animal screens and eventually to humans, constitutes the principal route for introducing a new anticancer drug. The process can take more than a decade.
3. Mechanism of Action Atudies
Once a class of compounds has been identified as biologically active, studies to elucidate the molecular mechanism of action can be undertaken. A first step is to identify the major cellular target or targets responsible for the chemotherapeutic properties of the drug. These investigations must also focus on chemical transformations that might take place in the solutions being administered and in the biological fluids that transport the drug to its ultimate target site. The next major step is to characterize the adduct or family of adducts made with the biological target molecule. The structure and kinetic lifetime of these adducts need to be investigated. Once this information is in hand, the effect of the adducts on the structure, stability, and function of the biological target molecule must be studied. Here many powerful new methodologies of modern molecular biology, genetics, and immunology can be brought to bear on the problem. The ultimate goals are to translate the molecular events elucidated into a realistic mechanism for how the drug molecule brings about its toxic effects selectively at the sites responsible for the disease and to use this information to design even better drugs.
Having progressed this far, we next need to bridge the gap between fundamental knowledge gained in studies of the mechanism of action and the SAR gleaned through pre-clinical and clinical trials. Whether such a happy situation can be reached for cisplatin remains to be seen, but there are encouraging signs, as we hope to demonstrate in the following discussion.
Clinical Picture for Cisplatin and Carboplatin49,52
1. Responsive Tumors and Combination Chemotherapy
It was an early observation that the best responses to cisplatin occurred in patients with genitourinary tumors. For testicular cancer, once a leading cause of death for males of age 20-40, cisplatin cures nearly all patients with stage A (testes alone) or B (metastasis or retroperitoneal lymph nodes) carcinomas. Platinum is usually given in combination with other drugs, commonly vinblastine and bleomycin for testicular cancer. This combination chemotherapy, as it is known, has several objectives. Some tumors have a natural or acquired resistance to one class of drugs and, by applying several, it is hoped that an effective reduction in tumor mass can be achieved. In addition, various drugs are known to affect different phases of the cell cycle, so several are applied simultaneously to allow for this possibility. Finally, synergism, where the response is greater than expected from simple additive effects, can occur, although it is rare. In addition to testicular cancer, platinum chemotherapy has produced responses in patients with ovarian carcinomas (>90 percent), head and neck cancers, non-small-cell lung cancer, and cervical cancer. Cisplatin is also effective when combined with radiation therapy.
2. Dose-limiting Problems; Toxicology
An early and quite worrisome adverse side effect of cisplatin was kidney toxicity. This problem, not commonly encountered with the older cancer drugs, nearly prevented its widespread use and eventual FDA approval. The major breakthrough here was made by E. Cvitkovic, who, while working at Sloan-Kettering Memorial Hospital in New York, administered large quantities of water by intravenous injection to patients, together with an osmotic diuretic agent such as D-mannitol. The rationale was that such hydration could ameliorate kidney toxicity by flushing out the heavy-metal complex. This simple idea worked, and the dose of the platinum compound could be increased threefold without accompanying nephrotoxicity. Hydration therapy is now commonly employed when cisplatin is administered. Among the other toxic effects encountered in cisplatin chemotherapy are nausea and vomiting, but this problem has also been controlled by use of antiemetic agents. Patients have also been known to experience bone-marrow suppression, a ringing in the ears, and occasional allergic reactions.
More recently, attempts have been made to extend cisplatin treatment to other broad classes of tumors by raising the dose above the ~5 mg/kg body weight levels given by i.v. injection every few weeks. Direct injection into the peritoneal cavity has been employed for refractory ovarian tumors. These more aggressive therapeutic protocols have been frustrated by drug resistance, a phenomenon by which cells learn to tolerate a toxic agent and for which many mechanisms exist, and by the return of the usual cisplatin side effects, most notably kidney toxicity and neurotoxicity. In order to combat toxic effects to the kidneys, chemoprotector drugs have been introduced. Based on the known affinity of platinum(II) complexes for sulfur-donor ligands, sodium diethyldithiocarbamate (DDTC) has been given both to experimental animals and to humans by i.v. infusion over about an hour following cisplatin administration.53 DDTC inhibits many of the toxic side effects, particularly to the kidneys and bone marrow, without itself producing long-term side effects or apparently inhibiting the antitumor properties of cis-DDP. Similar efforts have been made to reduce the toxic effects of cisplatin with other sulfur-containing compounds including thiosulfate and the naturally occurring biomolecules glutathione, cysteine, and methionine. The relative amounts of the latter three molecules can be controlled by drugs that affect their normal cellular concentrations.
Another approach to reducing cisplatin toxicity is to develop new classes of platinum drugs or different routes of their administration. Carboplatin (Figure 9.4) is one result of these efforts. The bidentate chelating dicarboxylate leaving group in carboplatin presumably retards the rates of reactions leading to toxicity, but does not adversely interfere with the chemistry required for antitumor activity. Recently, promising platinum compounds for oral administration have been developed.54 In Pt(IV) complexes of the kind cis, trans, cis-[Pt(NH3)(C6H11NH2)•(O2CCH3)2Cl2], where C6H11NH2 is cyclohexylamine, have been found to be effective in preclinical screens. The greater kinetic inertness of these complexes apparently renders them sufficiently stable to the chemically harsh environment of the gastrointestinal tract. Once absorbed into the bloodstream, these compounds are metabolized to the Pt(II) analogues, cis-[Pt(NH3)(C6H11NH2)Cl2], which are presumed to be the active form of the drug. The Pt(IV) compound has recently entered clinical trials.
Although impressive inroads have been made in the management of human tumors by platinum chemotherapy, the fact remains that, apart from testicular and to a lesser extent ovarian cancer, the median survival times are measured in months. Clearly, there is much room for improvement.
3. Pharmacology49,52
Solutions of cisplatin are usually given in physiological saline (NaCl), since hydrolysis reactions occur that can modify the nature of the compound and its reactions in vivo (see below). Cisplatin is rapidly cleared from the plasma after injection, 70-90 percent of the platinum being removed within the first 15 minutes. It has been found that more than half the platinum binds to serum proteins and is excreted. Most of the platinum exits the body via the urine within a few days. These results account for the use of multiple-dose chemotherapy at intervals of several weeks. Animal studies employing cis-DDP labeled with 195mPt, a 99 keV $\gamma$-emitter with a 4.1-day half-life, reveal retention half-times in various tissues of 8.4 (kidney), 6.0 (ileum), 4.1 (liver), 2.8 (tumor), and 1.9 (serum) days following a single dose. Platinum distributes widely to all tissue, with kidney, uterus, liver, and skin having the most, and muscles, testes, and brain the least amount of the compound. There is no evidence for selective uptake into normal versus tumor cells. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/09%3A_Metals_in_Medicine/9.05%3A_Platinum_Anticancer_Drugs-_A_Case_Study.txt |
As we have seen, the antitumor activity of cisplatin is most likely the result of its DNA-binding properties. But what are the adducts? The human genome has more than a billion nucleotides. Does platinum recognize any special regions of the DNA or any particular sequences? In other words, is binding simply random or is there at least a regioselectivity? In this section, we discuss the best strategies for answering these questions, strategies that evolved in pursuit of learning how cis-DDP binds to DNA. We also illustrate their power in elucidating the DNA-binding properties of other metal complexes of interest to bioinorganic chemists.
a. Early Strategic Approaches
The first experiments to imply the sequence preferences of cis-DDP binding to DNA employed synthetic polymers.108,109 Specifically, the buoyant density of poly(dG)•poly(dC), poly(dG•dC), and their cis-DDP adducts was studied in the analytical ultracentrifuge. The greatest shift in buoyant density was seen for the platinum adducts of poly(dG)•poly(dC), from which it was concluded that platinum forms an intrastrand crosslink between two neighboring guanosine nucleosides on the same strand. This interpretation was suggested by the known preference of metal ions, and especially platinum, for binding at the N7 position on the guanine base (Figure 9.9), information available from model studies of metal-nucleobase chemistry. Although other interpretations of the buoyant-density shift were possible, especially since the amount of platinum bound was not quantitated, the conclusion proved to be correct, as confirmed by later investigations. Interestingly, trans-DDP did not selectively increase the buoyant density of poly(dG)•poly(dC).
Following these initial experiments, the regioselectivity of cis-DDP binding was investigated by studying the inhibition of enzymatic digestion of platinated DNA. For example, the platinum complex inhibits the cleavage of DNA by restriction enzymes that recognize specific sequences and cut both strands of the double helix.110 The resulting fragments are readily identified on electrophoresis gels. One such restriction enzyme is Bam HI. As shown by the arrows in Scheme (9.11), Bam HI cleaves a six-bp palindromic sequence at the phosphodiester bonds between two guanosine nucleosides. Formation of an intrastrand crosslink between the two adjacent guanosine nucleosides inhibits digestion by the enzyme. Another method, termed exonuclease mapping, involves digestion of the strands of duplex DNA from its 3'-ends.111,112 When the enzyme encounters a bound platinum atom, it is unable to proceed further. Analysis of the digestion products by gel electrophoresis reveals the presence of discrete bands caused by the inhibition of digestion by bound platinum at specific sequences. Results from experiments of this kind were the most definitive at this time in demonstrating the profound regioselectivity of cisplatin for adjacent guanosines, and strongly supported the earlier conclusion that the drug was making an intrastrand d(GpG) crosslink.
In yet another approach to the problem, DNA containing cis- or trans-DDP adducts was electrostatically coupled to bovine serum albumin, to enhance its antigenicity, and injected into rabbits.113,114 The resulting antisera and antibodies were then studied for their ability to recognize and bind specifically to platinated DNAs having defined sequences, such as poly(dG)•poly(dC) and poly[d(GC)]•poly[d(GC)].
From experiments of this kind, the major cis-DDP adduct recognized by the antibody was found to be cis-[Pt(NH3)2{d(GpG)}], in accord with the findings of the enzymatic mapping experiments. Unplatinated DNA was not recognized, nor was DNA platinated with trans-DDP. On the other hand, the antibody recognized DNA platinated with antitumor-active compounds [Pt(en)Cl2] and [Pt(DACH)(CP)], where DACH = 1,2-diaminocyclohexane and CP = 4-carboxyphthalate. This result revealed that the antibody recognized the structural change in DNA that accompanies formation of d(GpG) intrastrand crosslinks, irrespective of the diamine ligand in the coordination sphere of the platinum atom. The antibody is also capable of distinguishing adducts formed by active versus inactive platinum complexes. Most importantly, DNA isolated from the cells of mice bearing the L1210 tumor five hours after cisplatin injection, was recognized.113,115 Subsequent studies116 revealed that these antibodies could detect cisplatin-DNA adducts formed in the white blood cells of patients receiving platinum chemotherapy. Thus, the antibody work linked the regiospecificity of platination chemistry in vitro with that occurring in vivo and in a clinically relevant manner.
Additional studies with monoclonal antibodies generated using DNA platinated with cis- or trans-DDP further confirmed and extended these results.117 This later work indicated that intrastrand crosslinked d(ApG) and d(GpG) sequences possess a common structural determinant produced by cis-DDP platination, and that carboplatin is also capable of inducing the same DNA structure. For trans-DDP-platinated DNA, a monoclonal antibody was obtained that appeared to have the intrastrand d(GpTpG) adduct as its major recognition site. In all these studies, the primary structural determinant appears to be DNA duplex opposite the site of platination, since fairly major stereochemical changes could be made in the amine ligands with no appreciable effect on antibody binding.
b. Degradation, Chromatographic Separation, and Quantitation of DNA Adducts
Experiments in which DNA platinated with cis-DDP is degraded to chromatographically separable, well-defined adducts have been invaluable in revealing the spectrum of products formed. In a typical experiment, platinated DNA is digested with DNAse I, nuclease P1, and alkaline phosphatase. These enzymatic digestions degrade DNA into nucleosides that can be readily separated by high-performance liquid chromatography (HPLC). Detection of the adducts can be accomplished by the UV absorption of the nucleoside bases at 260 nm or, for platinum complexes containing a radioactively labeled ligand such as [14C]ethylenediamine,118 by monitoring counts. In addition to peaks corresponding to dA, dC, dG, and dT, the chromatographic trace contains additional peaks corresponding to specific platinum nucleobase adducts such as cis-[Pt(NH3)2(dG)2]. The precise nature of these adducts was established by comparison with chemically synthesized compounds structurally characterized by NMR spectroscopy.118-121 An alternative method for identifying the adducts employed antibodies raised against specific platinum-nucleobase complexes.122
This approach has revealed the relative amounts of various adducts formed by a variety of platinum complexes; selected results are summarized in Table 9.4. Usually, for cisplatin, the relative amounts of the various adducts formed varies according to the series cis-[Pt(NH3)2{d(pGpG)}] > cis-[Pt(NH3)2{d(pApG)}] > cis-[Pt(NH3)2{d(GMP)}2] > monofunctional adducts. Only when the total incubation time was short, less than an hour, were the monofunctional adducts more prevalent, as expected from the kinetic studies of cis-DDP binding to DNA discussed previously. It is noteworthy that no d(pGpA) adducts were detected. This result, which is consistent with information obtained by enzymatic mapping, can be understood on stereochemical grounds.123 If the guanosine nucleoside N7 position is the most-preferred binding site on DNA, closure to make an N7,N7 intrastrand crosslink between two adjacent purine nucleotides is more feasible in the 5' direction along the helix backbone (N7•••N7 distance of ≈ 3 Å) than in the 3' direction (N7•••N7 distance ≈ 5 Å). In addition, molecular-mechanics modeling studies124 indicate that a highly unfavorable steric clash occurs between the 6-amino group of the 3'-adenosine residue in a d(pGpA) crosslink and the platinum ammine ligand, whereas in the platinated d(pApG) sequence, the 6-oxo group forms a stabilizing hydrogen bond to this ligand. A 28 kJ mol-1 preference of cis-DDP for binding d(pApG) over d(pGpA) was calculated.
Table 9.4 - Geometric features of the platinum coordination spheres of cis- [Pt(NH3)2{d(pGpG)}].
a) Bond distances are in Angstroms and angles are in degrees.
b) Conventions used for assigning Base/Base and Base/PtN4 dihedral angles can be found in J. D. Orbell, L. G. Marzilli, and T. J. Kistenmacher, J. Am. Chem. Soc. 103 (l981), 5126. The numbers in square brackets refer to the corresponding N(ammine)•••O6 distance, in Å (see text).
Bond Distances and Anglesa
Molecule 1
Molecule 2
Molecule 3
Molecule 4
Pt-N1 2.03(2) 2.01(2) 2.08(2) 2.08(2)
Pt-N2 2.03(3) 2.09(2) 2.04(3) 2.06(3)
Pt-N7A 2.01(2) 2.02(2) 1.91(3) 1.93(3)
Pt-N7B 2.05(2) 1.95(3) 2.00(3) 2.06(3)
N7A-Pt-N1 88.6(9) 90.3(9) 91.0(1) 88.4(9)
N7A-Pt-N2 179(1) 173.3(8) 178(1) 177(1)
N7A-Pt-N7B 89.1(9) 90.0(1) 85(1) 89(1)
N1-Pt-N2 92(9) 90.8(9) 91(1) 93(1)
qN1-Pt-N7B 176.5(9) 179.0(1) 173(1) 175(1)
N2-Pt-N7B 90.3(9) 89.0(1) 93(1) 89(1)
Dihedral Anglesb
Molecule
3'-Gua/5'-Gua
5'-Gua/PtN4
3'-Gua/PtN4
1 76.2(5) 110.6(5) [3.30(3)] 86.1(5)
2 81.0(5) 110.8(5) [3.49(3)] 95.5(5)
3 86.8(6) 81.0(6) 58.0(6) [3.11(4)]
4 80.6(5) 76.6(6) 59.6(6) [3.18(4)]
There are two likely sources of cis-[Pt(NH3)2{d(GMP)}2] in the spectrum of adducts. This species could arise from long-range intrastrand crosslinks, where the two coordinated guanosines are separated by one or more nucleotides. In support of this possibility is the fact that digestion of chemically synthesized cis-[Pt(NH3)2{d(GpNpG)}], where N = C or A, led to cis-[Pt(NH3)2{d(Gua)}-{d(GMP)}] and mononucleotides.118,119,121 The other source of this product is interstrand crosslinked DNA, known to occur from the alkaline elution studies.
As indicated in Table 9.4, in all the experiments there was platinum that was unaccounted for in the quantitation procedures, which employed either antibodies, platinum atomic absorption spectroscopy, or a radiolabeled ethylenediamine ligand. Some of this material was assigned to oligonucleotides having high platinum content, resistant to enzymatic degradation.
Two important points emerge from the quantitation of adducts by this method. One is that intrastrand d(GpG) and d(ApG) crosslinks constitute the major adducts (>90 percent of total platination) made by cisplatin on DNA in vivo. Because they were identified by an antibody specific for their structures, no chemical change brought about by cellular metabolism has occurred. Secondly, the preponderance of these adducts far exceeds the frequency of adjacent guanosine or guanosine/adenosine nucleosides in DNA. This latter result implies a kinetic preference for, or recognition of, d(pGpG)- and d(pApG)-containing sequences by cisplatin.
c. Postscript: A Comment on Methodologies
With few exceptions, none of the experimental studies described in this section could have been carried out in 1969, when Rosenberg first demonstrated the anticancer activity of cis-DDP. The techniques of DNA sequencing, monoclonal antibody formation, oligonucleotide synthesis, HPLC, FPLC, and many of the higher resolution gel electrophoresis methodologies employed were the result of later developments driven by rapid advances in the fields of molecular biology and immunology. Future progress in elucidating the molecular mechanisms of action of cisplatin and other inorganic pharmaceuticals will no doubt benefit from new technological discoveries and inventions of this kind yet to come. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/09%3A_Metals_in_Medicine/9.06%3A_Mapping_the_Major_Adducts_of_cis-_and_trans-DDP_on_DNA_Sequence_Specificity.txt |
a. NMR Studies of Platinated Oligonucleotides
Once the major spectrum of adducts formed by cis- and trans-DDP with DNA began to emerge, it was of immediate interest to learn to what positions on the nucleobases the platinum atom was coordinated. Proton NMR spectroscopy soon proved to be an invaluable tool for obtaining this information.71,125,126 Several ribo- and deoxyribooligonucleotides containing GG, AG, or GNG sequences were synthesized, and allowed to react with cis-DDP or its hydrolysis products, and the resulting complexes were purified by chromatography. All GG-containing oligomers formed intrastrand crosslinks with the {Pt(NH3)2}2+ moiety coordinated to the N7 atoms. This structure was deduced from several criteria. Most frequently studied were the nonexchangeable base protons H8 and H2 of adenine, H6 of thymine, H8 of guanine, and H5 and H6 of cytosine (Figure 9.9). Coordination of platinum to N7 of guanine causes a downfield shift of the H8 proton resonance. More importantly, however, it also lowers the pKa of the N1 proton by ~ 2 units, because platination adds positive charge to the base. Thus, titration of the platinated oligonucleotide over a pH range, and comparison of the results to those obtained for the unplatinated oligomer, reveals a difference in the midpoint of the transition in chemical shift of the H8 proton by 2 pH units if coordination occurs at N7. This effect is illustrated in Figure 9. 19 for the adduct cis- [Pt(NH3)2{d(ApGpGpCpCpT)}N7-G(2),N7-G(3)], where the pKa of N1 is seen to shift from ~ 10 to ~ 8 upon platination.71 The pH titration in this example also reveals the pH-dependent chemical shift of the cytosine 1H resonances at a pH of ~ 4.5, corresponding to protonation of the N3 atoms. The protonation of adenine N7 (pKa ~ 4) is also frequently observed in these studies. These results conclusively demonstrate platinum coordination at N7 of the two guanosine nucleosides.
Although several of the oligonucleotides studied have self-complementary sequences, such that they can form a double helix when unplatinated, in no such case was a duplex observed for their platinated forms. The presence of the platinum-induced crosslink presumably decreases the stability of the doublestranded form of the oligonucleotide. Another interesting result is that all intrastrand {Pt(NH3)2}2+ adducts of d(GpG) or d(ApG) have an altered deoxyribosesugar ring conformation. In normal, unplatinated form, these single-stranded or duplex oligonucleotides have a C2'-endo sugar pucker (Figure 9.9). Upon platination, the 5'-nucleotide switches to C3'-endo. This change is readily monitored by the ring proton-coupling constants JH1'-H2' and JH1'-H2". These protons constitute an ABX spin system such that the sum, $\Sigma$3J = 3J1'2' + 3J1'2', is most easily measured as the separation between the outermost peaks in the multiplet. For the C2'-endo conformation, a pseudotriplet occurs with $\Sigma$3J = 13.6 Hz, and for C3'-endo, $\Sigma$3J = 7.5 Hz. The 3'-guanosines in the adducts show greater conformational flexibility, having ~ 70 to 80 percent C2'-endo sugar puckers, depending upon the temperature.
Another conformational feature that could be deduced from 1H NMR studies of all cis-DDP-platinated oligonucleotides containing an embedded d(GpG) sequence is that both guanosine nucleosides retain the anti orientation of the base around the C1'-N9 glycosidic linkage (Figure 9.9). This result was deduced from the lack of a pronounced nuclear Overhauser effect (NOE) between H8 and the H1' protons, such as would occur in the syn conformation. An NOE between H8 resonances on the two coordinated nucleosides was observed for adducts of d(ApTpGpG) and d(CpGpG), indicating that the two bases are in a head-to-head orientation with respect to the platinum coordination plane. In other words, both O6 atoms lie on the same side of that plane. Two oligonucleotides containing cis-[Pt(NH3)2(ApG)] adducts have been examined; their structural properties closely resemble those of the (GpG) adducts, with platinum coordinated to N7 of both purine bases.
In order to study double-stranded DNAs platinated on one strand, it was necessary to adopt a special strategy. First, the desired oligonucleotide is synthesized. It is preferable that the DNA strands not be self-complementary, since the affinity of such an oligomer for itself is so much greater than that for its platinated form that the desired, singly platinated duplex will not form. After the platinated single strand is synthesized and purified, the complementary strand is added. Several duplex oligonucleotide-containing cis-[Pt(NH3)2{d(pGpG)}]-embedded adducts prepared in this manner have been studied by 1H NMR spectroscopy. With the use of two-dimensional and temperature-dependent techniques, both the nonexchangeable base and sugar protons as well as the exchangeable (guanine N1 and thymine N3) N-H (imino) proton resonances were examined. The last are useful, since they give some measure of the extent to which the double helix remains intact. When not base-paired to their complements in the other strand, these protons exchange more rapidly with solvent (water) protons, leading at moderate exchange rates to broadening of the resonances and, at high exchange rates (>107 s-1), disappearance of the signals. Several interesting results were obtained in these studies. In all of them, platination of the d(GpG) sequence brought about the same C2'-endo → C3'-endo sugar-ring pucker switch for the 5'-guanosine as seen in the single-stranded adducts. Head-to-head, anti conformations were also observed. At low temperatures, below the melting transition temperature, above which the duplex becomes single-stranded, the imino proton resonances were observed. This result was interpreted to mean that normal, Watson-Crick base pairs can still exist between the cis-DDP d(GpG) adduct and the d(CpC) sequence on the complementary strand. In the case of [d(TpCpTpCpG*pG*pTpCpTpC)]•[d(GpApGpApCpCpGpApGpA)], where the asterisks refer to the sites of platination, the imino proton resonances were assigned with the assistance of NOE experiments.125 Temperature-dependent studies showed that, in the range - 4° < T < 46 °C, the imino resonances of the coordinated guanosine nucleosides broadened first with increasing temperature. Apparently the base pairs of the intrastrand crosslinked, platinated duplex DNA are disrupted, or "melted," outward from the point of platination as well as from the ends. Since the amino hydrogen atoms involved in base pairing were not observed in this study, a completely definitive structural analysis was not possible. Nevertheless, the authors proposed that the duplex would be kinked by an angle of ~ 60° at the cis-DDP binding site in order to preserve full duplex character.
Another useful NMR nucleus for monitoring cis-DDP-DNA interactions is 195Pt, which is 34 percent abundant with I = $\frac{1}{2}$. When used in conjunction with 15N (I = $\frac{1}{2}$) enriched NH3 ligands, 195Pt NMR resonances provide a powerful means for characterizing complexes in solution. The 195Pt and 15N chemical shifts are both sensitive to the ligand trans to the NH3 group, as is the 195Pt-15N coupling constant.127 195Pt NMR studies of cis-DDP binding have been carried out using nucleobases, small oligonucleotides, and even double-stranded fragments of 20 to 40 bp in length, as previously described (Section V.D.l). The major contribution of this method is to show whether platinum coordinates to a nitrogen or an oxygen donor atom on the DNA, since the 195Pt chemical shift is sensitive to this difference in ligation.
b. X-ray Structural Studies
In recent years several oligonucleotide duplexes have been crystallized and characterized by x-ray diffraction methods. The probability of forming suitable single crystals of DNA fragments is disappointingly low, however, with only 1 in 10 such attempts being successful. Correspondingly, it has been difficult to crystallize platinated oligonucleotides. An alternative approach has been to soak nucleic-acid crystals of known structure with the platinum reagent in the hope of forming an isomorphous derivative, the structure of which could be obtained by using the changes in phases from the native material. In attempts to characterize a cis-DDP nucleic acid adduct, crystals of tRNAPhe and the self-complementary dodecamer d(CpGpCpGpApApTpTpCpGpCpG) were soaked with cisplatin solutions in the hope of obtaining useful metric information.123,128,129 These efforts have thus far failed to produce a high-resolution structure, although they confirm the predilection for platinum to coordinate to the N7 position of purine rings. Addition of cis-DDP tends to disorder the crystal, with platinum going to several sites of partial occupancy.
A more fruitful approach has been to crystallize a purified oligonucleotide containing the coordinated cis-{Pt(NH3)2}2+ moiety. The first x-ray structure to be deciphered through such a strategy was that of cis-[Pt(NH3)2{d(pGpG)}].130 This compound crystallizes with water solvent and glycine buffer molecules in the lattice. The crystals were grown at pH 3.8, where the terminal phosphate is monoprotonated in order to provide a neutral complex of diminished solubility. Two crystalline forms have been obtained, and both structures solved, one to 0.94 Å resolution. The latter contains four crystallographically independent molecules, which, although complicating the structure solution, afforded four independent views of the major adduct formed by cis-DDP with DNA. The four molecules form an aggregate, held together by hydrogen bonding and intermolecular base-base stacking interactions (Figure 9.20). There are two conformationally distinct classes that comprise molecules 1 and 2, and molecules 3 and 4; within each class, the molecules are related by an approximate C2 symmetry axis.
The molecular structure of molecule 1 is displayed in Figure 9.21; geometric information about all four molecules is contained in Table 9.4. As expected from the NMR studies, platinum coordinates to N7 atoms of the guanine bases, which are completely destacked (dihedral angles range from 76.2 to 86.7°), to form a square-planar geometry. The bases have a head-to-head configuration and conformational angles $\chi$ (Table 9.4 and Figure 9.9) that fall in the anti range. The sugar puckers of the 5'-deoxyribose rings for all four molecules have a C3'-endo conformation, and some of the 3'-sugar carbon atoms exhibit large thermal parameters suggestive of a less well-ordered structure. These results further demonstrate the similarity of the structure as detected in the solid state by x-ray diffraction and in the solution state by NMR spectroscopy.
An interesting additional feature of the cis-[Pt(NH3)2{d(pGpG)}] crystal structure is a hydrogen bonding interaction between an ammine ligand and the oxygen atom of the terminal phosphate group (OP1A•••N1, Figure 9.21). This intramolecular hydrogen bond is prominent in three of the four molecules in the asymmetric unit. Although the relevance of this hydrogen bonding interaction to the solution structure and molecular mechanism of cisplatin is presently unknown, it is interesting to note that the antitumor activity of platinum amine halide complexes is reduced when protons on coordinated NH3 are replaced by alkyl groups.34
A second cis-DDP-oligonucleotide adduct characterized by x-ray crystallography is the neutral molecule cis-[Pt(NH3)2{d(CpGpG)}].131 Here again, there are several (three) molecules in the asymmetric unit. Although determined at lower resolution, the structure is similar in most respects to that of cis-[Pt(NH3)2{d(pGpG)}] except for the presence of some weak NH3•••O6(guanosine) intramolecular hydrogen bonding interactions and a few unusual sugar-phosphate backbone torsional angles. Also, no NH3(H)••• phosphate(O) hydrogen bonds were observed.
From the foregoing discussion, it is apparent that adequate x-ray structure information is available for the cis-{Pt(NH3)2}2+/d(pGpG) intrastrand crosslink. What is needed now are structures of the minor adducts and, most importantly, of adducts in double-stranded DNA. Very recently, dodecanucleotide duplexes containing cis-{Pt(NH3)2}2+/d(pGpG) adducts have been crystallized, the structures of which are currently being investigated.132
c. Molecular Mechanics Calculations on Platinated Duplexes
As a supplement to x-ray structural information on double-stranded oligonucleotides containing an embedded cis-[Pt(NH3)2{d(pGpG)}] adduct, several models have been constructed by using a molecular mechanics approach.133 In this work, a set of coordinates was first obtained by amalgamation of structural information about standard double-helical DNAs and the platinated d(pGpG) fragment. Various starting structures were assumed, both linear and bent. The models were then refined according to various charge and stereochemical constraints built into the calculation. The results, which can reveal only what is feasible and not necessarily what actually happens, for both linear and bent structures are depicted in Figure 9.22 for two of the duplex sequences studied. In the linear model, the 5'-coordinated guanosine is rotated out of the stack, and its hydrogen bonding to the cytosine on the complementary strand is seriously disrupted. The imino N-H group is still involved in H-bonding, however; so this structure is not inconsistent with the NMR results. Two classes of kinked platinated duplex structures were encountered, with bending angles of 61 and 50°. In one of these, all Watson-Crick hydrogen bonds remain intact. These kinked structures are supported most strongly by the gel-electrophoresis experiments discussed in Section V.D.3.b.v.
Molecular mechanics and the related molecular dynamics calculations are a potentially valuable tool for the bioinorganic chemist interested in how metal complexes might perturb the structures of biopolymers. Analysis of the results for cisplatin-DNA binding reveals that, compared with the sum of all contributions from the biopolymer, the Pt-DNA interactions constitute a small part of the overall energy. For the most accurate results, it is important to know the charge distributions on the metal and its ligands as well as the effects of solvent interactions. Much work needs to be done in these areas before the results of molecular mechanics and dynamics calculations can be used reliably to predict or analyze structures. At present, however, they are far superior to examination of space-filling molecular models, for example, and produce quantitatively revealing structural diagrams.
d. Platinum-Nucleobase Model Complexes
Several studies have been carried out of the cis-diammineplatinum(II) moiety coordinated to nucleobases in which the N9 (purine) or N3 (pyrimidine) positions either have been alkylated, to simulate the glycosidic linkages, or in which the actual nucleotide (AMP, dGMP, etc.) is employed.134 These investigations are in many respects analogous to the synthesis and characterization by bioinorganic chemists of model complexes for the active site of a metalloenzyme. Their purpose is to simplify the problem, revealing kinetic, thennodynamic, and structural preferences of the primary building blocks involved in the metallodrug-biopolymer interaction, without the profound stereochemical constraints of the latter. Early studies of cis- and trans-DDP adducts with nucleobases (i) revealed the kinetic preferences for platinum binding to GMP and AMP, (ii) mapped out the preferred sites of platination (N7 of A and G; N1 of A; N3 of C; no N7-O6 chelate; no ribose or deoxyribose binding; only rare binding to phosphate oxygen atoms), (iii) demonstrated that Pt-N7 binding to G lowered the pKa of N1-H by ~ 2 units, and (iv) led to the discovery of interesting new classes of coordination complexes such as the cis-diammineplatinum pyrimidine blues and metal-metal bonded diplatinum(III) complexes.
Attempts to model the intrastrand d(GpG) crosslink with nucleobases have met with only moderate success. Usually the O6 atoms of the two guanosine rings are on opposite sides of the platinum coordination plane ("head-to-tail" isomer). Only for cis-[Pt(NH3)2(9-EtG)2]2+ was the correct isomer obtained. Nucleobase complexes of the cis-diammineplatinum(II) moiety have been valuable for testing the controversial proposal of N7,O6 chelate formation, which to date has not been observed. Several interesting discoveries of metal-nucleobase chemistry are that metal binding can stabilize rare tautomers, for example, the 4-imino, 2-oxo form of cytosine, through N4 binding, that coordination of platinum often produces unusual base pairing, and that metal migration from one donor site to another on an isolated nucleobase can occur. These model studies will continue to provide valuable insights into the possible chemistry of platinum antitumor drugs with DNA.
e. trans-DDP-DNA Adducts
Because trans-DDP is biologically inactive, it has received less attention than the cis isomer. Nevertheless, knowledge of its binding to DNA is important to have as a reference point for mechanistic comparison with the active compounds. Shortly after replication mapping experiments established that trans-DDP binds preferentially to d(GpNpG) and d(ApNpG) sequences,135 several synthetic oligonucleotides containing such sequences were prepared and used to investigate reactions with the trans isomer.136-138 Kinetic studies of trans-DDP with d(GpCpG) and d(ApGpGpCpCpT) revealed the presence of, presumably monofunctional, intermediates that closed to form both intra- and interstrand products. In the reaction with d(GpCpG), the 1,3-intrastrand G-G chelate accounted for 70 percent of the product, and 21 percent of the remaining material was unreacted oligonucleotide. Proton NMR studies of purified trans-[Pt(NH3)2{d(GpCpG)}] as well as the d(GpTpG) analog established platinum binding to N7 positions of the two trans guanosine nucleosides. As with the cis-[Pt(NH3)2{d(pGpG)}] adducts, the 5'-guanosine residue no longer retained the normal B-DNA type conformation; instead, the sugar ring pucker switched to C3'-endo. A fairly detailed 1H NMR characterization of trans- [Pt(NH3)2{d(ApGpGpCpCpT)-N7-A(1),N7-G(3)}] revealed very similar features. This example nicely illustrates the different stereoselectivity of cis- and trans-DDP binding to DNA. The cis isomer forms exclusively an intrastrand d(GpG) crosslink, whereas the trans isomer makes a 1,3-d(A*pGpG*) adduct. A schematic depiction of the trans-{Pt(NH3)2}2+ adduct is shown in Figure 9.23. As can be seen, the two purine rings enclose a large, 23-membered ring, the central guanosine residue is "bulged out," and the 5'-residue has a C3'-endo sugar pucker. This structure may be compared with that of cis-[Pt(NH3)2{d(pGpG)}] (Figure 9.21), where the platinum is part of a smaller, 17-membered ring. Both space-filling model building studies and molecular mechanics calculations reveal that it would be stereochemically very unfavorable for the trans-{Pt(NH3)2}2+ fragment to replace the cis analogue in an intrastrand crosslinked d(GpG) structure of the kind shown in Figure 9.21. Thus, for bidentate adducts, it seems clear that the important difference between cis- and trans-DDP binding to single-stranded DNA is revealed by the structures shown in Figures 9.21 and 9.23, respectively.
Information about trans-DDP binding to double-stranded DNA is scanty, but very recent studies indicate that the trans-[Pt(NH3)2{d(GpApG)-N7-G(1),N7-G(3)}] intrastrand crosslinked fragment can be embedded in duplex dodecamers.139 Interestingly, for one sequence the melting temperature (TM) of this duplex is not reduced over that of the unplatinated DNA fragment, in contrast to results for cis-DDP intrastrand d(GpG) adducts. This intriguing result, which agrees with earlier TM studies of DNA platinated by trans-DDP, does not yet have a structural rationale. It is possibly relevant to the processing of bifunctional trans-DDP-DNA adducts in vivo.
f. Effects of Platination on DNA Structure
It is valuable to summarize at this stage all that has been learned concerning the changes in DNA structure that occur upon cis- or trans-DDP binding. cis-DDP intrastrand crosslinks result in unstacking of neighboring bases and a switch in the sugar pucker of the 5'-nucleoside from C2'-endo, the standard B-DNA conformation, to C3'-endo, a conformation encountered in A-DNA. These various forms of DNA have already been introduced in the previous chapter. Watson-Crick base pairing, although weakened, is probably maintained. Evidence that base pairing is altered comes from studies with antinucleoside antibodies that bind appreciably better to DNA platinated with cis-DDP than to unmodified DNA. These antibodies recognize the nucleobases much better in platinated than in unplatinated DNA, presumably because platination disrupts the double helix. Additional support for base-pair disruption comes from gradient gel-denaturation experiments using sitespecifically platinated DNA (see Section V.D.8.b). Intrastrand crosslinking by cis-DDP also bends the helix by about 34° and unwinds the duplex by 13°. When trans-DDP forms 1,3-intrastrand crosslinks, the nucleotides situated between the platinated residues may be bulged out; consistent with this picture is the fact that they present an especially good target for antinucleoside antibodies. In 1,3-intrastrand d(GpNpG) or d(ApNpG) adducts, the 5'-nucleoside sugar pucker is altered to C3 I-endo. Intrastrand crosslink formation by trans-DDP also leads to DNA bending, but the platinum serves as the locus for a hinge joint and not for cooperative bending. These different effects of platination on DNA structure brought about by the two isomers are likely to be related to their different biological activities. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/09%3A_Metals_in_Medicine/9.07%3A_Structure_of_Platinum-DNA_Complexes.txt |
a. The Problem
Much of the information obtained about the mechanism of action of cisplatin has been derived from experiments where Pt-DNA binding has occurred in vivo or in vitro, with the use of random-sequence DNA having all available targets for the drug. In these studies, platination is controlled by the inorganic chemistry of cis-DDP in the medium and the accessibility of target sites on the DNA, as already discussed in considerable detail. As such, this situation best represents drug action as it actually occurs in the tumor cell. On the other hand, the resultant spectrum of DNA adducts makes it difficult, if not impossible, to understand the structural and functional consequences of any specific adduct. In order to address this problem, a methodology has been developed in which a single platinum adduct is built into a unique position in the genome. This approach is powerful and has the potential to be extended to the study of many other metal-based drugs. In this section, we discuss the strategy used to construct such site-specifically platinated DNA molecules and the information obtained thus far from their study. Some uses have already been discussed.
b. Synthesis and Characterization
Figure 9.29 displays the map of a genome constructed by insertion of platinated or unplatinated dodecanucleotide duplexes d(pTpCpTpApGpGpCpCpTpTpCpT)•d(pApGpApApGpGpCpCpTpApGpA) into DNA from bacteriophage M13mp18. This genome was constructed in the following manner.154 Double-stranded DNA from M13mp18 was first digested with Hinc II, a restriction enzyme that recognizes a unique six-base-pair sequence in the DNA and cleaves the double helix there, leaving a blunt-ended (no overhanging bases) cleavage site. The unplatinated dodecamer duplex was next ligated into the Hinc II site, and the DNA amplified in vivo. The dodecamer can insert into the genome in two different orientations, the desired one of which, termed M13-12A-Stu I, was identified by DNA sequencing. The presence of the insert in the new DNA was checked by its sensitivity to the restriction enzyme Stu I, which cleaves at the d(AGGCCT) sequence uniquely situated in the dodecamer insert, and the absence of cleavage by Hinc II the site for which was destroyed. Next, Hinc II-linearized M13mp18 replicative form (RF) DNA was allowed to form a heteroduplex with excess viral DNA (which has only the + strand in the presence of the denaturant formamide which was dialyzed away during the experiment. The resulting circular DNA has a gap in the minus strand into which the platinated d(TCTAG*G*CCTTCT) was ligated (Figure 9.30). The latter material was prepared by the methods described in Section V.D.5.a and characterized by 1H NMR spectroscopy. The resulting site-specifically platinated DNA contains a single cis-[Pt(NH3)2{d(pGpG)}] intrastrand crosslink built into a unique position. The methodology is general, and has been used to create other known platinum-DNA adducts site-specifically in M13mp18.
The chemical properties of the platinated DNA, termed M13-12A-Pt(-)-Stu I, were investigated by enzymatic, digestion and gel electrophoresis experiments. Platinum completely inhibits cleavage of the DNA by Stu I, as expected from the earlier restriction enzyme mapping studies. In addition, the cis- [Pt(NH3)2{d(pGpG)}] and cis-[Pt(NH3)2{d(pApG)}] intrastrand crosslinks were found to inhibit a variety of DNA polymerases, with only a small amount of bypass of the platinum lesion.149 These results indicate that the most abundant adducts of cisplatin on DNA are able to block replication efficiently.
c. Biological Properties
When M13-12A-Pt(-)-Stu I DNA was introduced into E. coli cells by transformation, DNA synthesis was uninterrupted, because the cell can both repair the damage and use the unmodified (+) strand for synthesis. Consequently, a slightly different strategy was used to construct single-stranded M13-12A-Pt(+)-Stu I DNA, the details of which are available elsewhere.154 This platinated template, in which the damage can neither be repaired nor bypassed by known mechanisms in vivo, was then transformed into E. coli cells co-plated with GW5100 cells. Under these conditions, viral DNA replication is detected by the expression of the $\beta$-galactosidase gene, which, in the presence of appropriate reagents in the medium, leads to formation of blue plaques on a clear background. The results clearly indicate that many fewer plaques appear when M13-12A-Pt(+) is introduced into the cells than when M13-12A-u(+) was employed, where u stands for unmodified DNA. In three repeats of this experiment, survival of DNA containing only a single cis-[Pt(NH3h{d(pGpG)}] crosslink was only 11 ± 1 percent.
These data provide unambiguous proof that the most frequent DNA adduct formed by cisplatin is toxic, capable of inhibiting replication when only a single such lesion is present on a natural DNA template of 7,167 nucleotides. The fact that as many as 10 percent of the transformed cells can bypass or repair the lesion is also of interest, and parallels the results found in vitro. In related work, it was found that the cis-[Pt(NH3)2{d(pGpG)}] intrastrand crosslink is not very mutagenic, but that cis-[Pt(NH3)2{d(pApG)}] intrastrand adducts are considerably more so. This finding is important, since mutations could lead to long-term secondary tumor production in patients treated with cisplatin. The methodology affords a way to screen new compounds that one would like to be equally effective at inhibiting replication but less mutagenic. In addition, by using repairdeficient mutant cell lines, as well as cisplatin resistant cells, one can study the effects of varying the properties of the host cells. Incorporation of site-specifically platinated DNA sequences into appropriate shuttle vectors will also facilitate investigation of toxicity, repair, and resistance in mammalian cells.
d. Prospectus
The foregoing discussion illustrates the power of site-specifically platinated DNAs as a probe of the molecular mechanism of the drug. We recall that similar strategies were employed to obtain uniquely modified DNA in the bending92,93 and unwinding95 experiments discussed previously. In principle, this technique can be applied to examine other aspects of the molecular mechanism of other metallochemotherapeutic agents. The requirements are a synthetic route to the uniquely modified genome, for which both the inorganic coordination chemistry and molecular biology must be amenable, an adduct stable to the biological conditions for DNA synthesis, and a method (usually genetic) for scoring the biological effects being investigated. Site-specifically platinated DNAs allow the bioinorganic chemist to have maximal control over the genetics and should continue to provide valuable information about the molecular mechanism of action of cisplatin. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/09%3A_Metals_in_Medicine/9.08%3A_Site-specifically_Platinated_DNA_%28154%29.txt |
The material in this section constitutes the major portion of this chapter. One important goal of the discussion is to illustrate, by means of an in-depth analysis of a single case history, the questions that must be addressed to elucidate the molecular mechanism of an inorganic pharmaceutical. Another is to introduce the techniques that are required to answer these questions, at least for the chosen case. The inorganic chemist reading this material with little or no biological background may find the experience challenging, although an attempt has been made to explain unfamiliar terms as much as possible. It is strongly advised that material in Chapter 8 be read before this section. Toward the end of this section, the results obtained are used to speculate about a molecular mechanism to account for the biological activity of the drug. Experiments directed toward evaluating the various hypotheses are delineated. Once the mechanism or mechanisms are known, it should be possible to design new and better antitumor drugs which, if successful, would be the ultimate proof of the validity of the hypotheses. This topic is discussed in the next and final section of the chapter. Such an analysis could, in principle, be applied to probe the molecular mechanisms of the other metals used in medicine described previously. In fact, it is hoped that the approach will prove valuable to students and researchers in these other areas, where much less information is currently available at the molecular level.
The material in this section has been organized in the following manner. First we discuss the relevant inorganic chemistry of platinum complexes in biological media. Next we summarize the evidence that DNA is a major target of cisplatin in the cancer cell, responsible for its antitumor activity. The chemical, physical, and biological consequences of damaging DNA by the drug are then described, followed by a presentation of the methodologies used to map its binding sites on DNA. The detailed structures of the DNA adducts of both active and inactive platinum complexes are then discussed, together with the way in which the tertiary structure of the double helix can modulate these structures. Finally, the response of cellular proteins to cisplatin-damaged DNA is presented, leading eventually to hypotheses about how tumor cells are selectively destroyed by the drug. Together these events constitute our knowledge of the' 'molecular mechanism," at least as it is currently understood.
1. Reactions of cis-DDP and Related Compounds in Aqueous, Biological, and Other Media
cis-Diamminedichloroplatinum(II) is a square-planar d8 complex. As such, it belongs to a class of compounds extensively investigated by coordination chemists.55 Typically, such compounds are relatively inert kinetically, do not usually expand their coordination numbers, and undergo ligand substitution reactions by two independent pathways with the rate law as given by Equation (9.4). The rate constants k1 and k2 correspond to first-order (solvent-assisted) and second-order (bimolecular) pathways;
$Rate = (k_{1} + k_{2}[Y]) [complex] \tag{9.4}$
[Y] is the concentration of the incoming ligand. Usually, k1 < < k2 by several orders of magnitude. In biological fluids, however, the concentration of a potential target molecule could be ~10-6 M, in which case k1 $\simeq$ k2[Y]. Substitution of ligands in cis-DDP, required for binding to a cellular target molecule, is therefore likely to proceed by the solvent-assisted pathway. Such a pathway is assumed in the ensuing discussion.
For the hydrolysis of the first chloride ion from cis- or trans-DDP,
$[Pt(NH_{3})_{2}Cl_{2}] + H_{2}O \rightleftharpoons [Pt(NH_{3})_{2}Cl(OH_{2})]^{+} + Cl^{-} \tag{9.5}$
the k1 values at 25 °C are 2.5 x 10-5 and 9.8 x 10-5 s-1, respectively.55 These hydrolyzed complexes can undergo further equilibrium reactions, summarized by Equations (9.6) to (9.9).
$[Pt(NH_{3})_{2}Cl(OH_{2})]^{+} \rightleftharpoons [Pt(NH_{3})_{2}Cl(OH)] + H^{+} \tag{9.6}$
$[Pt(NH_{3})_{2}Cl(OH_{2})]^{+} + H_{2}O \rightleftharpoons [Pt(NH_{3})_{2}(OH_{2})_{2}]^{2+} + Cl^{-} \tag{9.7}$
$[Pt(NH_{3})_{2}(OH_{2})_{2}]^{2+} \rightleftharpoons [Pt(NH_{3})_{2}(OH_{2})(OH)]^{+} + H^{+} \tag{9.8}$
$[Pt(NH_{3})_{2}(OH_{2})(OH)]^{+} \rightleftharpoons [Pt(NH_{3})_{2}(OH)_{2}] + H^{+} \tag{9.9}$
The formation of dimers such as [Pt(NH3)2(OH)]22+ and higher oligomers can also occur,56,57 but such reactions are unlikely to be important at the low platinum concentrations encountered in biological media, Reactions (9.5) to (9.9), which depend on pH and the chloride-ion concentrations, have been followed by 195Pt (I = $\frac{1}{2}$, 34.4 percent abundance) and 15N (using enriched compounds) NMR spectroscopy. The latter method has revealed for the cis-diammine complexes pKa values of 6.70 ± 0.10 at 25 °C for Reaction (9.6) and of 5.95 ± 0.1 and 7.85 ± 0.1 at 5 °C for Reactions (9.8) and (9.9), respectively.58
The effects of pH and Cl- ion concentration on the species distribution of platinum compounds have been used to fashion the following plausible argument for the chemistry of cis-DDP in vivo.59 With the use of thermodynamic data for the ethylenediamine (en) analogue [Pt(en)Cl2] the relative concentrations of hydrolyzed species at pH 7.4 were estimated (see Table 9.3) for blood plasma and cytoplasm (Figure 9.7). The higher chloride ion concentration in plasma preserves the complex as the neutral molecule cis-DDP, which passively diffuses across cell membranes. The lower intracellular chloride ion concentration facilitates hydrolysis reactions such as Equations (9.5) to (9.9), thereby activating the drug for binding to its biological target molecules. There is, of course, a reasonable probability that cis-DDP and species derived from it will encounter small molecules and macromolecules in vivo that divert it from this route to the target. We have already seen such cases; cisplatin binds to serum proteins, and there is good evidence that intracellular thiols react with the drug.60 Glutathione, for example, is present in millimolar concentrations in cells. How, one might ask, does cisplatin swim through such a sea of sulfur donors to find its target in the tumor cell? Is it possible that a modified form of the drug, in which a Pt-CI bond has been displaced by thiolate to fonn a Pt-S bond, is the actual species responsible for its activity? Although these questions have not yet been satisfactorily answered, there is reason to believe that such reactions are not directly involved in the molecular mechanism of action. As evident from structure-activity relationship studies, the most active compounds have two labile ligands in cis positions. If Pt-S bonds were required, then compounds already having such linkages would be expected to exhibit activity and they do not. Rather, it seems most likely that the antitumor activity of cisplatin results from surviving species of the kind written in Equations (9.5) to (9.9) that find their way to the target molecule, and that the induced toxicity must arise from a significantly disruptive structural consequence of drug binding. Since only cis complexes are active, it is reasonable for the coordination chemist to infer that the stereochemistry of this interaction is of fundamental importance.
Table 9.3 - Distribution of various adducts formed between cis-DDP or [3H][Pt(en)Cl2]a and DNA in vitro and in vivo.118-122
a) A radiolabeled analogue of cis-DDP, [3H]dichloroethylenediamineplatinum(II).
b) A2 represents either (NH3)2 or ethylenediamine.
c) By difference.
d) Percentage of adducts based on total amount of platinum eluted from the separation column.
e) Percentage of adducts based on total amount of radioactivity eluted from the separation column.
f) Not given.
g) Chinese hamster ovary cells treated with 83 $\mu$M cis-DDP.
h) Results from ELISA.
i) Results from AAS. Where the signal was too weak for reliable quantitation, the maximal amount possible is given. Adapted from Table 1 in Reference 81.
D/N Ratio Total Incubation Time Adducts
cis-[PtA2{d(pGpGl}]b
cis-[PtA2{d(pApG}]b
Formed
cis-[PtA2{d(GMP)}2]b
Mono-functional Adducts Remaining Platinumc
In vitro
0.055c 5 h (50 °C) 47-50% 23-28% 8-10% 2-3% 10%
0.022d 5 h (50 °C) 60-65% 20% ~4% ~2% 9-14%
0.01e 16 h (37 °C) 62% 21% 7% 10%
ef 30 m (37 °C) 36% 3% 8% 40% 13%
ef 2 h (37 °C) 54% 9% 9% 14% 14%
ef 3 h (37 °C) 57% 15% 9% 4% 15%
In vivo
dg 1 h (37 °C) 35.9 ± 4.7%h <34%i 3.1 ± 1.6%h 38.5%i ~22%
dg 25 h (37 °C) 46.6 ± 6.8%h <48%i 3.0 ± 0.9%h <14.5%i ~50%
Reactions of platinum compounds with components in media used to dissolve them can give and undoubtedly have given rise to misleading results, both in fundamental mechanistic work and in screening studies. A particularly noteworthy example is dimethylsulfoxide (DMSO), which even recently has been used to dissolve platinum compounds, presumably owing to their greater solubility in DMSO compared to water. As demonstrated by 195Pt NMR spectroscopy, both cis- and trans-DDP react rapidly (t1/2 = 60 and 8 min at 37 °C, respectively) to form [Pt(NH3)2CI(DMSO)]+ complexes with chemical and biological reactivity different from those of the parent ammine halides.61
2. Evidence that DNA is the Target
Two early sets of experiments pointed to interactions of cisplatin with DNA, rather than the many other possible cellular receptors, as an essential target responsible for cytotoxicity and antitumor properties.62,63 Monitoring the uptake of radiolabeled precursors for synthesizing DNA, RNA, and proteins, showed that [3H]thymidine incorporation was most affected by therapeutic levels of cisplatin for both cells in culture and Ehrlich ascites cells in mice. Since independent studies showed that cis-DDP binding to DNA polymerase does not alter its ability to synthesize DNA, it was concluded that platination of the template and not the enzyme was responsible for the inhibition of replication.
In a second kind of experiment demonstrating that DNA is a target of cisplatin, hydrolyzed forms of the drug in low concentrations were added to a strain of E. coli K12 cells containing a sex-specific factor F.64,65 After free platinum was removed, these F+ cells were conjugated with a strain of E. coli K12 cells lacking this factor that had previously been infected with lambda bacteriophage. Addition of cis-DDP, but not trans-DDP, directly to the latter infected F- cells had been shown in a separate study to accelerate cell lysis. Conjugation with the platinum-treated F+ cells produced the same effect, strongly suggesting that Pt had been transferred from the F+ to the F- cells. Since only DNA is passed between the F+ and F- strains, it was concluded that Pt was attached to the DNA and that this modification was essential for the observed lysis of the cell. Further studies showed a good correlation between cell lysis by platinum compounds and their antitumor properties.
Various other observations are consistent with the notion that platinum binding to DNA in the cell is an event of biological consequence.66 The filamentous bacterial growth observed in the original Rosenberg experiment is one such piece of evidence, since other known DNA-damaging agents, for example, alkylating drugs and x-irradiation, also elicit this response. Another is the greater sensitivity toward cis-DDP of cells deficient in their ability to repair DNA. Finally, quantitation of the amount of platinum bound to DNA, RNA, and proteins revealed that, although more Pt was bound to RNA per gram biomolecule, much more Pt was on the DNA when expressed as a per-molecule basis. In the absence of any selective interaction of Pt with a specific molecule, only one out of every 1,500 protein molecules (average M.W. ~60 kDa) in a cell will contain a single bound platinum atom, whereas hundreds or thousands of Pt atoms are coordinated to DNA (M.W. ~1011). If the replication apparatus cannot bypass these lesions, then cell division will not occur, and tumor growth is inhibited.
Although these and other results all point to DNA as an important cellular target of cisplatin, most likely responsible for its anticancer activity, this information does not explain why tumor cells are more affected by cis-DDP than non-tumor cells of the same tissue. Moreover, why is trans-DDP, which also enters cells, binds DNA, and inhibits replication, albeit at much higher doses (see discussion below), not an active anticancer drug? What causes cisplatin to kill cells and not merely to arrest tumor growth? The latter can be explained by DNA synthesis inhibition, but not necessarily the former. Very recent studies have begun to address these questions using powerful new methodologies of molecular and cell biology, as described in subsequent sections of this chapter. The results, although preliminary, continue to point to DNA as the most important cellular target of cisplatin. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/09%3A_Metals_in_Medicine/9.09%3A_Bioinorganic_Chemistry_of_Platinum_Anticancer_Drugs-_How_Might_They_Work.txt |
6. Effects of DNA Structure on Platinum Binding
a. A-, B-, and Z-DNA140
As discussed in more detail in Chapter 8, double-helical DNA can adopt different polymorphic forms depending on the conditions in solution or polycrystalline fiber. Even within a given DNA molecule, there can be sequence-dependent local secondary and tertiary structural differences that constitute important signals for cellular DNA binding and processing molecules. An example already discussed is the recognition of palindromic sequences by type II restriction endonucleases. As shown in Figure 9.24, three such DNA polymorphs are the right-handed A- and B- and the left-handed Z-forms. Most commonly encountered in solution is B-DNA, characterized by well-classified major and minor grooves designated by arrows in Figure 9.24. The targets of platinum binding, guanine N7 atoms, are situated in the major groove.
To what extent do sequence-dependent local structural modulations affect platinum binding? Although no general answer to this question can be given, there are several interesting anecdotal pieces of information worth mentioning. Z-DNA, a form favored by alternating purine-pyrimidine sequences such as in poly d(GC), does not constitute a particularly good target for cis-DDP binding. For one thing, it lacks the preferred d(GpG) or d(ApG) sequences. The monofunctional [Pt(dien)CI]+ complex, however, facilitates the B-DNA → Z-DNA conformational transition, as demonstrated by circular dichroism and 31P NMR spectroscopic data.141 In Z-DNA, the guanosine nucleoside adopts the syn conformation (Figure 9.9), which is presumably favored by placing a bulky {Pt(dien)}2+ moiety on N7. Moreover, the local charge density on DNA is greater in Z- than B-DNA, owing to the closer proximity of the phosphate groups, and the former is presumably stabilized by the +2 charge on the platinum complex.
b. Effects of Local Sequence and of Free and Linked Intercalators on Platinum Binding
Of more interest perhaps to anticancer drug-DNA interactions is the fact that some d(GpG), d(ApG), and even d(GpA) targets for cis-DDP binding are very sensitive to the sequences in which they are embedded. This phenomenon was first discovered during exonuclease III mapping studies of cisplatin binding to a 165-bp restriction fragment from pBR322 DNA.142 Although cis-DDP binding stops the enzyme at G3, G5, and GAGGGAG sequences, at a (D/N)b ratio of 0.05 there is little evidence for coordination to an apparently favored G6CG2 sequence (Figure. 9.25, lane 8). When platination was carried out in the presence of the DNA intercalator EtdBr (Figure 9.11), however, the G6CG2 sequence became a Pt binding site (Figure 9.25, lanes 9-12). A more extensive exonuclease III mapping study of this phenomenon suggested that d(CGG)-containing sequences in general are less well platinated by cis-DDP.143,144 Moreover, only EtdBr, and not other acridine or phenanthridinium type intercalators, was able to promote an enzyme-detectable cis-DDP binding to these sequences. A suggested explanation for these results is that local d(purCGG) sequences might have an A-DNA-type structure (Figure 9.24) in which the major groove is narrow, inhibiting access of platinum to N7 of guanosine nucleosides. In the presence of the intercalator EtdBr, the local DNA structure might be altered in such a manner as to permit binding.143,144
In accord with this interpretation, and further to delineate a possible reason why acridines and deaminated ethidium cations do not promote cisplatin binding to d(purCGG) sequences, NMR studies were performed that revealed the mean residence time of EtdBr on DNA to be 6 to 21 times longer than that of any of the other intercalators examined.144 Thus, for these latter intercalators, the local DNA structure presumably can relax back to one unfavored for cisplatin binding before it can diffuse to the site. Moreover, when acridine orange (AO), one of the five intercalators studied that does not promote cisplatin binding to excluded sites, was covalently attached to dichloroethylenediamineplatinum(ll) (Figure 9.26) via a hexamethylene linker chain, the resulting AO-Pt molecule was able to bind to all d(CpGpG) sites, as determined by exonuclease III mapping. In the tethered molecule, the high local concentration near the intercalator binding site facilitates attachment of the {Pt(en)}2+ moiety to DNA before the acridine orange fragment can diffuse away and the structure can relax to reform the excluded site.
Subsequently, the excluded site phenomenon was found for cis-DDP binding, as assayed by the 3'-5'-exonuclease activity of T4 DNA polymerase.145 Enzyme stopping sites were observed at all d(GpG) sequences, but only weakly when at a d(GTGGTC) site. Similarly, d(ApG) was not modltled when embedded in pyGAGCpy and pyGAGCA sequences. Although most d(GpA) sequences were not platinated, as detected by T4 mapping, a few were. These results further underscore the importance of local sequence modulation of cisbinding to DNA.
c. DNA-Promoted Reaction Chemistry
In the EtdBr-enhanced binding of cis-DDP to DNA, a small fraction (< 5 percent) of the intercalator is strongly bound and can be dialyzed out only very slowly.145 The detailed structure of this DNA-cisplatin-EtdBr ternary complex has been established, and involves cis-{Pt(NH3)2}2+ binding to the exocyclic amino groups of ethidium as well as to donor sites on DNA.146 This assignment was proved by synthesizing cis- [Pt(NH3)2CEtd)CI]2+ complexes in dimethylformamide solution and then allowing them to react with DNA. The optical spectra of the resulting adducts were identical to that of the ternary complex. The reaction of cis-DDP, and DNA to form the ternary complex is promoted by the favorable orientation of the exocyclic amino group of intercalated Etd with respect to the coordination plane of bound to the double helix. The N-8 exocyclic amino group of ethidium, bound intercalatively at a site adjacent to a purine N-7 coordinated cis-{Pt(NH3)2CI}+ moiety, is positioned above the platinum atom in a structure resembling the transition state for a square-planar substitution reaction. The structure of this transition state has been modeled in a molecular mechanics calculation (Figure 9.27 See color plate section, page C-16.),147 and evidence has been obtained that indicates selective binding of platinum to the Etd N-8 amino position.148
Since cisplatin is usually administered in combination chemotherapy with other drugs, many of which contain intercalating functionalities, strong covalent, DNA-promoted interactions between drug molecules at a target site must be considered as possibly relevant to the molecular mechanism of action. In such a situation, there must be a strong binding preference for both drug molecules for the same target sequences, since on probability grounds alone it is unlikely that both would migrate to the same site by random diffusion at the low concentrations found in vivo.
d. Effects of DNA Function on Platinum Binding
Although there is yet little known about this topic for cisplatin, it is worth pointing out that other DNA-targeted drugs, such as bifunctional alkylating agents, bind preferentially to actively transcribed genes. It is therefore possible that platinum exhibits such preferences, for example, to single-stranded DNA at or beyond the transcription fork, compared to duplex DNA in chromatin structures. Or, perhaps, it too binds selectively to actively transcribed DNA. Investigation of these possibilities seems worthwhile.
7. Speculations About the Molecular Mechanism
a. Is There a Single Mechanism?
Most investigators now agree that DNA is the cytotoxic target of cisplatin. We have seen that the drug inhibits DNA replication by binding to the template and halting the processive action of DNA polymerase. Less well-studied is the inhibition of transcription by platinum-DNA adducts, but recent evidence clearly indicates that they can do so. Studies of the effects of platinum on cells growing in culture reveal that DNA replication and cell growth can continue without cell division in the presence of low levels (1 $\mu$g/mL) of cis-DDP cells are arrested at the G2 phase, the stage of cell growth just preceding division.105 G2 arrest was reversible, but at higher cisplatin levels (8 $\mu$g/mL), cell death occurred. These observations led to the speculation that, perhaps, post-replication DNA repair can handle the toxicity associated with a platinum-damaged template, at least for DNA synthesis, but that there is no known pathway by which transcription can circumvent Pt-DNA lesions. Possibly, inhibition of transcription is ultimately a more lethal event than inhibition of replication. This idea is inconsistent with the well-established fact that thymidine incorporation into DNA is more affected by low levels of cisplatin than is uridine incorporation into RNA. Might there be more than one biochemical pathway by which cisplatin manifests its anticancer activity? Further work is necessary to address this intriguing question.
b. Is There a "Critical Lesion"?
We now have an excellent understanding of the major DNA adducts made by cis-DDP, their structures, and the corresponding DNA distortions. Information about adducts made by the inactive trans isomer, though not as complete, is also substantial. During the period when this knowledge was being accumulated, it was of interest to learn whether a "critical lesion," a specific DNA adduct with a unique molecular structure might be responsible for the antitumor activity of the drug. At present, it appears that all bidentate adducts made by cis- and trans-DDP can inhibit replication, although they may not be equally efficient at doing so.149 Even monofunctional adducts of the kind formed by cis-[Pt(NH3)2(4-Br-py)CI]+ can block replication.96 Thus, it be better to think about the concept of "critical lesion" in a functional sense, where the rates of adduct formation, removal, and enzyme inhibition together determine which family of adducts will exhibit antitumor activity and which will not. Here the biochemistry of the host cell will also be an important determinant. Clearly, more studies are required to delineate these possibilities.
c. Replication and Repair in the Tumor Cell150
If the anticancer activity of cisplatin arises from damaged DNA templates, then the drug could be selectively toxic to cancer versus normal cells of the same tissue if repair of DNA damage occurred more efficiently in the latter. The best way to study this phenomenon would be to measure the platinum-DNA levels and list the spetrum of adducts formed in tumor versus normal biopsy tissue obtained from patients undergoing cisplatin chemotherapy. As described previously, methodologies are now reaching the point where such experiments can be carried out in order to test the key hypotheses about the mechanism of action of cisplatin. In addition, powerful new methods have recently been developed to screen for DNA bmdlnlg proteins. If one could identify proteins that bind selectively to cis-DDP-platinated DNA and determine function, further insights into cellular replication and repair phenomena would be forthcoming. Such cellular factors that bind selectively to DNA containing cisplatin adducts have, in fact, recently been discovered.107 The experiments that led to this finding and their possible implications for the molecular mechanism of cisplatin are described in the next section.
d. Structure-Specific (or Damage) Recognition Proteins150
If selective repair of platinum-DNA adducts in cells of different origin is an integral part of the anticancer mechanism of cis-DDP, then it is important to identify the cullular factors associated this phenomenon. In bacteria, cis-DDP adducts on DNA are removed by excision repair, a process in which the lesion is first identified and then excised by the uvrABC excinuclease system.151 In this process, the uvrA protein first binds to the adducted DNA. Subsequently, the uvrB and C proteins excise the damaged strand, which additional cellular proteins rebuild by copying the genetic information from the remaining strand.
The repair of cis-DDP intrastrand crosslinks in mammalian cells is much less well understood. Under the assumption that an analogue of the uvrA protein might exist in such cells, experiments were carried out to try to isolate and clone the gene for such a protein. In particular, the mobility of platinated DNA restriction fragments of defined length was found to be substantially retarded in electrophoresis gels following incubation extracts from human HeLa cells.107 This gel-mobility shift was attributed to the binding of factors termed "damage recognition proteins" (or DRPs). Subsequent studies with site-specifically platinated oligonucleotides (see V.D.8) revealed that the cisplatin DRP binds specifically to DNA containing the intrastrand cis-[Pt(NH3)2{d(pGpG)}] or cis-[Pt(NH3)2{d(pApG)}] crosslink. In parallel work, the gene encoding for a DRP was cloned152 and used to demonstrate the occurrence of such a in nearly all eukaryotic cells. Since binding of the DRP to platinated DNA is not specific for the ammine ligands opposite the crosslinked nucleobases, the interaction is thougllt to involve recognition of local changes in the twist and bending of the double helix. Figure 9.28 depicts one possible structure for the complex formed between cis-DDP platinated DNA and a DRP. More recently, the cloned proteins were found to contain a high mobility group (HMG) protein box, and even HMG 1 itself binds to cisplatin-modified DNA.152 The class of proteins was renamed "structure-specific recognition proteins" (SSRPs).
The discovery of SSRPs that bind specifically to cisplatin-modified DNA raises several questions that are the subjects of current study. The first is to determine whether the proteins are an integral component in the mechanism of action of the drug. Although it has not yet been possible to induce the proteins by treating cells with cisplatin, nor have elevated or suppressed levels been found in platinum-resistant cells, deletion of an SSRP gene in yeast has re(~enltlv afforded a mutant strain less sensitive to cisplatin than wild-type cells.153 This result links a yeast SSRP with cellular sensitization to the drug. Such a protein could contribute to the molecular mechanism in one of several ways (Figure 9.28). It might be the analogue of uvrA, which, as mentioned above, recognizes damage and signals the cell to perform excision repair. If so, then one would like to depress the levels of the protein in cancer cells to make them more sensitive to the drug. A second possibility is that the true role of the SSRP is to serve as a tumor-cell activator, and that cisplatin lesions titrate it away from its functionally active sites on the DNA. Alternatively, binding of the protein could protect cisplatin adducts from repair, preserving their lethality at the time of cell division and leading to the arrest of tumor growth. This last hypothesis would require more of the SSRP in cells sensitive to the drug. Studies are currently in progress to delineate these three and other hypotheses, and to learn whether the discovery of the SSRPs has heralded the final chapter in the quest for the molecular mechanism of cisplatin or merely been an entertaining sidelight.
e. Drug Resistance: What Do We Know?150
Perhaps the most serious problem for successful chemotherapy of tumors is drug resistance.102 In most tumors there exists a subpopulation of cells that are naturally resistant to a given drug; as the sensitive cells are killed, these refractory clones take over. In addition, resistance can be acquired by tumor cells following repeated application of the drug. Attempts to identify mechanisms responsible for cisplatin resistance have therefore been the subjects of considerable research activity. Other DNA-damaging agents sometimes amplify genes as a mechanism of drug resistance. An example is the multidrug resistance phenomenon, in which a gene encoded for a P-glycoprotein is amplified in cells resistant to agents such as daunomycin. This protein is believed to increase efflux of the drug through the cell membrane by an ATP-dependent, energy-driven pump. There is currently an intensive search underway to see whether the cisplatin resistance phenomenon has a genetic origin. If a cisplatin resistance gene could be cloned and its phenotype identified, a powerful new avenue would be opened to overcome drug resistance. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/09%3A_Metals_in_Medicine/9.10%3A_Bioinorganic_Chemistry_of_Platinum_Anticancer_Drugs_How_Might_They_Work_%28Part_2%29.txt |
1. Objectives
Although chemotherapy has made significant contributions to cancer treatment, the effect of cisplatin on testicular cancer being a showcase example, early detection and surgical removal of all neoplastic tissue still remain the preferred means of combating most forms of the disease. What steps need to be taken to devise better chemotherapeutic agents? One answer is to understand the biochemical mechanisms that underlie the transformation of normal into neoplastic cells and to attack the disease on the basis of that knowledge. The value of this approach is indisputable, but it need not be the only one. We have seen that cis-[Pt(NH3)2Cl2], a simple third-row transition-metal complex containing no carbon atoms, can contribute significantly to cancer chemotherapy. This example alone should lead us to search for improved inorganic drugs based on the evolving knowledge of the mechanism of action of cis-DDP. What then should our objectives be? Three answers are immediately apparent. First, we need to find compounds that are active against resistant cells. Such compounds are termed "second-generation" platinum drugs, and are the focus of much activity in the pharmaceutical industry. Their development will be facilitated by understanding the fundamental biochemistry of cisplatin drug resistance, designing complexes to circumvent the cellular resistance mechanisms. Second, there needs to be an improved spectrum of activity, to be provided by the so-called "third-generation" compounds. The major cancers of the colon, breast, and lung are not effectively diminished by cisplatin chemotherapy. Finally, cisplatin toxicity is often dose-limiting, and there is a need for agents with a greater chemotherapeutic index-to-toxicity ratio. Some of these objectives may ultimately be met by modifying the mode of delivery of cisplatin, for example, by encapsulating the drug in a tumor-seeking liposome or attaching it to a tissue-specific monoclonal antibody. A major step in alternative delivery has recently been taken with the development of a class of oral platinum complexes that have just entered clinical trials.54 These complexes are platinum(IV) cycloalkylamine species of the kind cis, trans, cis-[Pt(NH3)(C6H5NH2)(O2CCH3)2Cl2]. The prospects are reasonably good that new platinum and other metal anticancer drugs can be designed in a bioinorganic chemical approach to the problem.
2. Strategies for Drug Development
a. Can We Build on Our Knowledge About Cisplatin?
If we consider what is known about the molecular mechanism of cisplatin, what properties are desirable in the design of new metal complexes for testing? The molecules should be reasonably stable kinetrically and soluble in biological fluids, cross the cell membrane, bind covalently to DNA, and inhibit gene function. As described previously in this section of the chapter, powerful methods are now available to screen compounds for these properties in a relatively short time. But there are additional factors required for metallodrug anticancer activity, above and beyond these criteria; trans-DDP, after all, has all five of the above properties and is not active. Probably one should add to the list the requirement that the complex have two substitutionally labile cis sites for intrastrand crosslinking of adjacent DNA nucleotides; such a criterion would, of course, rule out molecules like trans-DDP. Recall, however, that cis-[Pt(NH3)2(4-X-py)CI]+ complexes (X = Br, Me) are active. These cations have the five properties listed above, but, as far as is currently known, bind only monofunctionally to DNA. The pyridine ring moiety of a covalently attached platinum atom could possibly intercalate into a neighboring interbase pair site on the DNA, making a pseudointrastrand crosslinked adduct structurally similar to the cis-DDP-d(pGpG) structure. Further information is required about these active, monofunctional cations before any firm conclusions can be drawn. Nevertheless, it is useful to remember that if the requirement of two substitutionally labile cis ligands had been rigorously followed, this new class of monofunctional platinum complexes would not have been discovered.
Another rationale for designing new platinum or other metal antitumor drugs could emerge with a better understanding of the SSRPs in the mechanism of action of cisplatin. For example, if they serve to protect cisplatin lesions on DNA from repair, one would want to design complexes that form adducts that bind even more strongly to the purified protein. The strength of this binding interaction, having been a serendipitous discovery, surely cannot have been maximized. A tighter SSRP-platinated DNA complex would require the use of less platinum, and thus afford lower toxicities.
b. Is Platinum Uniquely Suited?
Given the above criteria, platinum is the only metal to be chosen for further drug development? The answer to this question is "probably not" but a few points need to be kept in mind. Given the assumption that the geometry of the cis-[Pt(NH3)2{d(GpG)}] intrastrand crosslink was important for the antitumor activity of cisplatin, computer graphics methods were employed to probe the stereochemical consequences of modifying this structure.155 Addition of axial chloride or water ligands in fifth and sixth coordination positions to form pseudo-octahedral adducts, for example, introduces several steric clashes with the guanosine O6 atoms. An octahedral complex, for example cis,cis,cis-[Pt(NH3)2Cl2(OH)2], bifunctionally coordinated to DNA either would not form an intrastrand d(GpG) crosslink or would form an adduct structurally different from that made by cis-DDP. This octahedral complex Pt(IV), known as "tetraplatin" in the pharmaceutical industry, is active, but is believed to be reduced in vivo to platinum(II) before coordinating to DNA.156,157 These considerations might imply that the best strategy for inorganic drug development would be to employ square-planar d8 complexes. Clearly there are as yet no definitive answers. Nevertheless, the criteria derived from the mechanism of action studies of cisplatin represent an excellent starting point for designing new antitumor metallodrugs.
c. How Important are Amine Ligands?
Here again, the answer is not unequivocal, but amines (including NH3) are probably ideally suited ligands for covalent DNA-binding metal complexes. Even completely inert complexes such as [Co(NH3)6]3+ show sequence and DNA polymorph binding preferences,158 suggesting that the N-H bonds orient toward the phosphate and heterocyclic nitrogen atoms in the major groove, forming hydrogen-bonding interactions. This chemistry is analogous to the binding and recognition of organic amines and polyamines, such as spermine and spermidine, by nucleic acids. Apart from amines, hydrophobic groove-binding and/or intercalating ligands such as o-phenanthroline and its derivatives should be considered. Molecules such as [Rh(DIP)3]3+ , where DIP = 4,7-diphenyl-1,10penanthroline, bind to DNA and have proved to be useful structural probes (Chapter 8). Recent work has shown that [Rh(DIP)2Cl2]+ binds preferentially to d(GpG) sequences, like cisplatin, although its antitumor properties have not yet been investigated.159
3. Second and Third-generation Platinum Anticancer Drugs
Improvements over cisplatin have been made, most notably the molecule carboplatin (Figure 9.4), which is less nephrotoxic and has been reported to be effective in some patients where cisplatin chemotherapy has failed. These properties come solely from the dicarboxylate leaving group, which is kinetically more inert to substitution. Studies with monoclonal antibodies have shown the DNA adducts of carboplatin to be identical with those formed by cis-DDP.117Other platinum compounds that have undergone clinical trials are close analogues of cis-[PtA2X2], or tetraplatin, cis,cis,cis-[PtA2X2Y2], that obey the classic structure-activity relationships. The activity of cationic triamines, cis-[Pt(NH3)2LCI]CI, where L is pyridine, a substituted pyridine, pyrimidine, or purine, against S180 ascites and L1210 tumors in mice opens a new vista of possible structures to be tried. The intercalator-linked complex AO-Pt (Figure 9.26) has also been found to show activity in the S180 ascites system, suggesting a further class of complexes that could be studied. The oral compounds, cis,trans,cis-[Pt(NH3)(C6H5NH2)(O2CCH3)2Cl2], survive the digestive tract, are taken across the gastrointestinal mucosa, and metabolize to cis-[Pt(NH3)(C6H5NH2)Cl2], a cisplatin analogue.54 As such they are effective pro-drugs that could become the major platinum agent in clinical use. Until these recent advances, there was a general impression that, by chance, the best compound discovered was the first one, cisplatin. There is now sufficient reason to expect that innovative experimentation will lead to improved drugs, bearing in mind the comment made earlier (Section IV.G.) that sustained individual effort for up to a decade can be required to move a compound from the laboratory bench into the clinic.
4. Nonplatinum Antitumor Metal Complexes
a. Soft Metals
As mentioned in Section IV.E., some compounds of Pd(II), Au(I), Rh(II), and Ru(II or III) have been screened for antitumor activity, but much more work needs to be done in this arena. The higher metal-ligand exchange rates of Pd(II), ~105 faster than those of Pt(Il), make these complexes potentially more toxic, as some preliminary animal studies have shown. By use of chelating or organometallic complexes, however, this problem might be avoided. The properties of Ru, Rh, and to a lesser extent Au amine and polypyridine complexes would seem to make them attractive candidates, and indeed there appears to be renewed interest in these molecules.160 Inorganic chemists interested in pursuing drug development with these metals need to forge alliances with biological colleagues equipped to do the necessary animal screening and to develop in-house expertise for cell culture and related biochemical work. The techniques are not all that difficult and it is actually fun to undertake studies of the biological consequences of metallodrug chemistry.
b. Metallocenes and Metallocene Dihalides36-37
Although complexes such as [(C5H5)2TiCl2] are superficially analogous to cis-DDP, in being potentially bifunctional DNA crosslinking agents, their hydrolytic reactions are sufficiently different to cast doubt on the value of this comparison. The fact that antitumor activity has been found for this very different class of inorganic compound, however, suggests that perhaps bioinorganic chemists have explored only a very small sample of possible metallodrugs. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/09%3A_Metals_in_Medicine/9.11%3A_Design_of_New_Inorganic_Anticancer_Drugs.txt |
The oxidation state of an element is defined as the formal charge on the atom if all bonds were assumed to be fully ionic.
In an ionic compound the oxidation state is equal to the charge on the ion, e.g., in NaCl the charge on the sodium is +1 and the oxidation state is also +1. In contrast, the charge on an atom in a covalent compound never approaches the charge implied from the oxidation number, e.g., in CCl4, the oxidation state of carbon is +4 but the charge on the carbon atom is signicantly less.
Thus, while oxidation state is a simple formulism it is very useful in a number of ways:
1. Classication of compounds of elements, especially those of the transition metals.
2. Enables the balance of reduction-oxidation (redox) reactions.
To determine the oxidation state of an element within a particular compound the following steps act as a guideline:
1. The oxidation state of any atom of any element in its elemental form is zero.
2. The oxidation state of a monatomic ion is equal to its charge.
3. Any homoleptic X-X bonds (e.g., the C-C bond in H3C-CH3) are assumed to be non-ionic and do not contribute to the oxidation state of X.
4. Fluorine always has oxidation state of 1
5. Elements of Group 1 (IA) (except hydrogen) have an oxidation state of +1 in compounds.
6. Elements of Group 2 (IIA) have an oxidation state of +2 in compounds.
7. Elements of Group 17 (VIIA) have an oxidation state of -1 when they combine with elements below or to the left of their position in the periodic table
8. Oxygen is usually assigned the oxidation state of 2, except in compounds with uorine, oxygen has a positive oxidation number.
9. Hydrogen is assigned the oxidation state of +1 with non-metals and 1 with metals.
10. In any compound the sum of the oxidation states (oxidation numbers) is equal to the overall charge.
Exercises
Exercise \(1\)
What is the oxidation state of sulfur in Na2SO4?
Solution
The two sodium atoms each have an oxidation state of +1, while the oxygen atoms have an oxidation state of 2, and the overall charge is 0.
1. overall charge = sum of oxidation states
2. 0 = (2 x oxidation state of Na) + (oxidation state of S) + (4 x oxidation state of O)
3. 0 = (2 x +1) + (oxidation state of S) + (4 x -2)
4. oxidation state of S = 0 (2 x +1) - (4 x 2)
5. oxidation state of S = +6
Exercise \(2\)
What is the oxidation state of sulfur in Na2SO3?
Solution
The two sodium atoms each have an oxidation state of +1, while the oxygen atoms have an oxidation state of 2, and the overall charge is 0.
1. overall charge = sum of oxidation states
2. 0 = (2 x oxidation state of Na) + (oxidation state of S) + (3 x oxidation state of O)
3. 0 = (2 x +1) + (oxidation state of S) + (3 x -2)
4. oxidation state of S = 0 (2 x +1) - (3 x 2)
5. oxidation state of S = +4
Exercise \(3\)
What is the oxidation state of sulfur in H2S?
Solution
The two hydrogen atoms each have an oxidation state of +1 and the overall charge is 0.
• overall charge = sum of oxidation states
• 0 = (2 x oxidation state of H) + (oxidation state of S)
• 0 = (2 x +1) + (oxidation state of S)
• oxidation state of S = 0 (2 x +1)
• oxidation state of S = -2
In writing the oxidation state of an element within a compound it is common to use a Roman numeral, rather than the charge, i.e., Al(III) rather than Al3+. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/01%3A_General_Concepts_and_Trends/1.01%3A_Fundamental_Properties_-_Oxidation_State.txt |
The ionization energy (IE), or more properly the ionization enthalpy, is defined as the energy required to loose an electron from a gaseous atom or ion.
$M^{n+}_{(g)} \rightarrow M^{(n+1)+}_{(g)} + e^-$
Each subsequent ionization energy is greater than the previous one because of the increase in charge on the ion. For example, for any given atom or ion the 1st ionization energy is less than the 2nd ionization energy, and so on. This is shown in Table $1$.1.
Ionization Ionization energy (kJ/mol)
$Al^0 \rightarrow Al^+ + e^-$ 548
$Al^+ \rightarrow Al^{2+} + e^-$ 1823
$Al^{2+} \rightarrow Al^{3+} + e^-$ 2751
Table $1$.1: The first three ionization energies for aluminum.
How does the ionization energy vary with elements in the periodic table? If we consider the 1st ionization potential of the elements in a particular group of the periodic table we note that there is a decrease in ionization potential as you go down the Group. The reason for this trend is due to the increased shielding of the outer shell electrons (ns1) by the completed (filled) inner shells. The ns1 electron thus exhibits a lower effective nuclear charge and makes it easier to remove. For example, Table $1$.2 shows the 1st ionization potential for the Group 1 alkali metals.
Element
Ionization energy (kJ/mol)
Li 526
Na 502
K 425
Rb 409
Cs 382
Table $1$.2: Variation of the first ionization potential (M0 → M+) for the Group 1 (IA) elements.
In contrast, to individual Groups, moving across a particular Period results in a general increase in the ionization potential as is shown in Figure $1$.1. The lack of additional screening of filled shells across the Period means that the ionization energy for the outer shell electrons is dominated by the increase in nuclear charge (number of protons) with increased atomic number.
From Figure $1$.1 it can be seen that the Periodic trend is not linear, there are significant steps in the plot. Boron, for example, has a lower first ionization potential than beryllium, why? A consideration of the electron configuration for the elements provides in answer (Figure $1$.2). Beryllium has a 2s2 outer shell configuration, while boron has a 2s2 2p1 configuration. The 2p1 electron is easy to remove because it exhibits increased shielding from the nucleus due to the filled 2s orbital.
As we move from, boron to nitrogen, the 2p shell is filled (Figure $1$.2) without additional shielding and the effect of the increased nuclear charge dominates. Finally, the 2p4 configuration for oxygen (Figure $1$.2) results in an electron pair, which repel each other, thus making it easier to remove an electron (lower ionization potential) that expected from the increased nuclear charge. From Figure $1$ we can see that the effect of electron pairing is less than that of a filled shell.
The trend for the 2nd ionization potential is similar, but different, to that of the 1st ionization potential. As may be seen in Figure $1$.3 the steps observed for the 1st ionization energy plot (i.e., between Be/B and N/O have moved one element to the right. A view of the electron conguration for the E+ ions (Figure $1$.4) shows that the rational for the trend in the 1st ionization potential trends still applies but to the ion of the element to the right in the Periodic table. Now B+ has a 2s2 outer shell conguration, while C+ has a 2s2 2p1 configuration. A similar plot for the 3rd ionization energy would move the steps another element to the right.
The other observation to be made from Figure $1$.3 is the very large 2nd ionization potential for lithium associated with the ionization of Li+ to Li2+. The large increase is due to the removal of an electron from the filled 1s2 shell. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/01%3A_General_Concepts_and_Trends/1.02%3A_Fundamental_Properties_-_Ionization_Energy.txt |
The electron affinity (EA) of an element is defined as the energy given of when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion.
$X^{n-}_{(g)} + e^- \rightarrow X^{(n+1)-}_{(g)}$
Electron anities are more difficult to measure than ionization energies and are usually less accurately known. Electron affinities are large and negative for elements such as fluorine and oxygen, and small and positive for metals.
Electron affinities generally become smaller as you go down a Group of the periodic table (Table $1$.3). This is because the electron being added to the atom is placed in a larger orbital, where it spends less time near the nucleus of the atom, and also the number of electrons on an atom increases as we go down a column, so the force of repulsion between the electron being added and the electrons already present on a neutral atom becomes larger. Electron anities are further complicated since the repulsion between the electron being added to the atom and the electrons already present on the atom depends on the volume of the atom. Thus, for the nonmetals in Groups 6 (VIA) and 7 (VIIA), this force of repulsion is largest for the very smallest atoms in these columns: oxygen and uorine. As a result, these elements have a smaller electron anity than the elements below them in these columns as shown in Table $1$.3.
Element Electron affinity (kJ/mol)
F -322
Cl -349
Br -325
I -295
Table $1$.3: The electron anity for the non-metallic halogens.
Although there is a general trend that for Group 1 (IA) to Group 17 (VIIA) elements the electron affinity increases across the Periodic table from left to right, the details of the trend are more complex. As may be seen from Figure $1$.5, there is a cyclic trend. The explanation of this is a consequence of the unusually stable electron configurations exhibited by atoms with filled or half filled shells, i.e., helium, beryllium, nitrogen and neon (see Table $1$.4). These configurations are so stable that it actually takes energy to force one of these elements to pick up an extra electron to form a negative ion.
Element Electron affinity (kJ/mol) Electron configuration
H -72.8 1s1
He +ve
1s2
Li -59.8 [He] 2s1
Be +ve [He] 2s2
B -27 [He] 2s22p1
C -122.3 [He] 2s22p2
N +ve [He] 2s22p3
O -141.1 [He] 2s22p4
F -328.0 [He] 2s22p5
Ne +ve
[He] 2s22p6
Table $1$.4: Electron affinities of the elements hydrogen to neon. N.B. Values for helium, beryllium, nitrogen, and neon are not known with any accuracy but are all positive.
1.04: Fundamental Properties - Electronegativity
An issue with ionization potential and electron affinity is that they are defined and measured as reactions in the gas phase. Although values have been determined for molecular fragments it is still difficult to correlate with reaction trends in solution. To overcome this issue the concept of electronegativity was developed.
Electronegativity is defined as the tendency of an atom in a molecule to attract electrons to itself. Although several electronegativity scales have been developed, that by Linus Pauling (Figure \(1\).6) is the most often used. Table \(1\).5 provides selected Pauling electronegativity values (unit less).
Element Pauling Scale
F 4.0
O 3.5
Cl 3.0
N 3.0
S 2.5
C 2.5
H 2.1
B 2.0
Na 0.9
Table \(1\).5: Selected Pauling electronegativity values.
The advantage of the Pauling electronegativity scale is that it allows the prediction of general behavior. For example, the larger the dierence in electronegativity between two elements the more ionic character or more polar the bonding interaction. Thus, a H-O bond (3.5 - 2.1 = 1.4) is more polar than a H-S bond (2.5 2.1 = 0.4). | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/01%3A_General_Concepts_and_Trends/1.03%3A_Fundamental_Properties_-_Electron_Affinity.txt |
The idea of a correlation between molecular geometry and the number of valence electrons was first presented in 1940 by Sidgwick and Powell; however, in 1957, Ronald Gillespie (Figure \(1\).7) and Sir Ronald Nyholm (Figure \(1\).8) refined this concept to build a more detailed theory. It is their work that provides the basis of the valence shell electron pair repulsion (VSEPR) theory, and as such it is also known as the Gillespie-Nyholm theory.
One attribute of VSEPR is that with the ability to predict the shape of a molecule for a compound comes the ability to predict some of the physical and chemical properties of that compound.
The formal definition that is the basis for VSEPR is as follows: Pairs of electrons in the valence shell of a central atom of a molecule repel each other and take up positions as far apart as possible.
Within this definition it is implicitly assumed that the core shells are not polarized and therefore take no part in bonding, and therefore can be ignored. Since the maximum repulsion of the electron pairs (be they associated with a bonding interaction or a lone pair) control the shape of a molecule, for each number of electron pairs we can dene the geometric optimum position that maximizes the distance between the electron pairs (Table \(1\).6). The number of electron pairs surrounding an atom, both bonding and nonbonding, is called its steric number.
Number of central atom electron pairs Bonding pairs Non-bonding pairs Shape Example
2 2 0 Linear BeCl2
3 3 0 Triangular BF3
3 2 1 Bent SnCl2
4 4 0 Tetrahedral CCl4
4 3 1 Pyramidal NH3
4 2 2 Bent H2O
5 5 0 Trigonal bipyramidal (tbp) PF5
5 4 1 Pseudo-tdp BrF4-
5 3 2 T-Shaped BrF3
5 2 3 Linear XeF2
6 6 0 Octahedral SF6,PF6-
6 5 1 Square pyramidal IF5
6 4 2 Square planar XeF4,IF4-
Table \(1\).6: Shapes of molecules and ions
The steps for dening the molecular shape are as follows:
1. Draw a simple Lewis structure including single, double, and triple bonds where appropriate.
2. Count the number of electrons on the central atom assuming it is neutral.
3. Add one electron for each σ-bond.
4. Subtract one electron for each π-bond.
5. Subtract one electron for each positive (+) charge.
6. Add one electron for each negative (-) charge.
7. Divide the number of electrons by two to give the number of electron pairs.
8. Use the list in Table \(1\).6 to predict the structure of the molecule.
Example \(1\)
What is the shape of NH4+
Solution
1. Nitrogen has 5 valence electrons
2. Add 4 electrons for the four σ-bonds: 5 + 4 = 9
3. Subtract one electron for the positive (+) charge: 9 1 = 8
4. Divide the number of electrons by two to give the number of electron pairs: 8/2 = 4
5. Four bonding pairs and no lone pairs = tetrahedral geometry
Exercises
Exercise \(1\)
What is the shape of HgCl2?
Answer
1. Mercury has 2 valence electrons
2. Add 2 electrons for the two σ-bonds: 2 + 2 = 4
3. Divide the number of electrons by two to give the number of electron pairs: 4/2 = 2
4. Two bonding pairs and no lone pairs = linear geometry
Exercise \(2\)
What is the shape of H2CO?
Answer
1. Carbon has 4 valence electrons
2. Add 3 electrons for the three σ-bonds: 4 + 3 = 7
3. Subtract one electron for each π-bond: 7 1 = 6
4. Divide the number of electrons by two to give the number of electron pairs: 6/2 = 3
5. Two bonding pairs and no lone pairs = triangular geometry
Lone pairs versus bonding pairs
The prediction of the detailed molecular structure (including bond angles) is not as simple as shown in Table \(1\).6. In molecules with either lone pair electrons or multiple (double or triple) bonds the angles about the central atom are distorted due to the increased electron repulsion (Figure \(1\).9). The differences in repulsion caused by a lone pair or a bonding pair may be rationalized in a simple manner by a lone pair taking up more space than a bonding pair (Figure \(1\).10).
Water is one of the classic cases in considering the issue of non-bonding (unshared) electron pairs.
1. Oxygen has 6 valence electrons
2. Add 2 electrons for the two σ-bonds: 6 + 2 = 8
3. Divide the number of electrons by two to give the number of electron pairs: 8/2 = 4
4. Two bonding pairs and two lone pairs = tetrahedral geometry
From this an idealized tetrahedral geometry would give the H-O-H angle as 109.5, however, from Figure \(1\).9 we know that the lone pair...lone pair repulsion is greater than the lone pair...bonding pair repulsion which is greater than the bonding pair...bonding pair repulsion, and thus, the H-O-H angle should be decreased from the ideal tetrahedral. The experimentally determined H-O-H angle in water is in fact 104.5.
Ethylene is a good case in considering the issue of multiple bonds. Ethylene contains both σ-bond and π-bond between the carbon atoms. This combination can be thought of as a super bond, and as such its effect is similar to a lone pair.
Carbon has 4 valence electrons
Add 3 electrons for the two σ-bonds: 4 + 3 = 7
Subtract one electron for each π-bond: 7 1 = 6
Divide the number of electrons by two to give the number of electron pairs: 6/2 = 3
Three bonding pairs and no lone pairs = triangular geometry
From this an idealized tetrahedral geometry would give the H-C-H angle as 120, however, the π-bond repulsion is greater than the σ-bond repulsion, and thus, the H-C-H angle should be decreased from the ideal tetrahedral. The experimentally determined H-O-H angle in water is in fact 118.3.
Resonance structures
If a molecule or ion has two or more resonance forms it is necessary to consider each form before angles are predicted. For example, the carbonate anion, CO3-, can be drawn as a single structure from which it would be predicted that two groups of O-C-O angles would result (Figure \(1\).11).
However, CO3- should actually be drawn in each of its resonance forms as in Figure \(1\).12.
From Figure \(1\).12 it is clear that the real structure will be an average of the three resonance forms, and hence there will be a single O-C-O angle = 120 (Figure \(1\).13).
Atom electronegativities
In an A-X bond where the atom electronegativities are very different the bonding pair is assumed to occupy less space than in bond between two atoms of similar electronegativities. As the bonding pair occupies less space it will repel neighboring electron pairs less.
For example, based on the above, a comparison of the tetrahedral compounds H2O and F2O would suggest that the F-O-F angle be smaller than the H-O-H angle since fluorine has a higher electronegativity than hydrogen (4.0 and 2.1, respectively). This is indeed observed (Figure \(1\).14).
In many cases, inter-ligand steric interactions can also be used to explain the dffierence in angle. For example, while fluorine is more electronegative than chlorine (4.0 versus 3.0, respectively) it is also significantly smaller (Table \(1\).7). Thus, a larger Cl-X-Cl angle than a F-X-F angle in the homologous compound, can be attributed to a greater steric interactions, rather than difference in electronegativities.
Ligand Covalent Radius (Å) Van der Waal radius (Å)
F 0.57 1.35
Cl 1.02 1.80
Table \(1\).7: Comparison of size between fluoride and chloride ligands.
Bent's rule
There is the potential for more than one isomer for molecules that adopt structures in which there is a symmetry difference between at least two of the ligand positions. For example a trigional bipyramidal compound of the formula PXY4 has two possible structures. One where the X occupies an axial position (Figure \(1\).15a) and the other where it occupies an equatorial position (Figure \(1\).15b).
Through the consideration of structures Henry Bent suggested a rule: More electronegative substituents prefer hybrid orbitals with less s character, and conversely, more electropositive substituents prefer hybrid orbitals with greater s character.
For example, in PFCl5 the fluorine is the most electronegative substituent and will therefore occupy the axial (p-character only) position.
Bibliography
• H. A. Bent, J. Chem. Educ., 1960, 37, 616.
• H. A. Bent, Chem. Rev., 1961, 61, 275.
• R. J. Gillespie, J. Chem. Educ., 1970, 47, 18.
• R. J. Gillespie, Chem. Soc. Rev., 1992, 21, 59.
• R. J. Gillespie and R. S. Nyholm, Quart. Rev., 1957, 11, 339. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/01%3A_General_Concepts_and_Trends/1.05%3A_Structure_and_Bonding_-_Valence_Shell_Electron_Pair_Repulsion_%28VSEPR%29_Theory.txt |
The idea of a correlation between molecular geometry and the number of valence electrons was first presented in 1940 by Sidgwick and Powell; however, in 1957, Ronald Gillespie (Figure \(1\).7) and Sir Ronald Nyholm (Figure \(1\).8) refined this concept to build a more detailed theory. It is their work that provides the basis of the valence shell electron pair repulsion (VSEPR) theory, and as such it is also known as the Gillespie-Nyholm theory.
One attribute of VSEPR is that with the ability to predict the shape of a molecule for a compound comes the ability to predict some of the physical and chemical properties of that compound.
The formal definition that is the basis for VSEPR is as follows: Pairs of electrons in the valence shell of a central atom of a molecule repel each other and take up positions as far apart as possible.
Within this definition it is implicitly assumed that the core shells are not polarized and therefore take no part in bonding, and therefore can be ignored. Since the maximum repulsion of the electron pairs (be they associated with a bonding interaction or a lone pair) control the shape of a molecule, for each number of electron pairs we can dene the geometric optimum position that maximizes the distance between the electron pairs (Table \(1\).6). The number of electron pairs surrounding an atom, both bonding and nonbonding, is called its steric number.
Number of central atom electron pairs Bonding pairs Non-bonding pairs Shape Example
2 2 0 Linear BeCl2
3 3 0 Triangular BF3
3 2 1 Bent SnCl2
4 4 0 Tetrahedral CCl4
4 3 1 Pyramidal NH3
4 2 2 Bent H2O
5 5 0 Trigonal bipyramidal (tbp) PF5
5 4 1 Pseudo-tdp BrF4-
5 3 2 T-Shaped BrF3
5 2 3 Linear XeF2
6 6 0 Octahedral SF6,PF6-
6 5 1 Square pyramidal IF5
6 4 2 Square planar XeF4,IF4-
Table \(1\).6: Shapes of molecules and ions
The steps for dening the molecular shape are as follows:
1. Draw a simple Lewis structure including single, double, and triple bonds where appropriate.
2. Count the number of electrons on the central atom assuming it is neutral.
3. Add one electron for each σ-bond.
4. Subtract one electron for each π-bond.
5. Subtract one electron for each positive (+) charge.
6. Add one electron for each negative (-) charge.
7. Divide the number of electrons by two to give the number of electron pairs.
8. Use the list in Table \(1\).6 to predict the structure of the molecule.
Example \(1\)
What is the shape of NH4+
Solution
1. Nitrogen has 5 valence electrons
2. Add 4 electrons for the four σ-bonds: 5 + 4 = 9
3. Subtract one electron for the positive (+) charge: 9 1 = 8
4. Divide the number of electrons by two to give the number of electron pairs: 8/2 = 4
5. Four bonding pairs and no lone pairs = tetrahedral geometry
Exercises
Exercise \(1\)
What is the shape of HgCl2?
Answer
1. Mercury has 2 valence electrons
2. Add 2 electrons for the two σ-bonds: 2 + 2 = 4
3. Divide the number of electrons by two to give the number of electron pairs: 4/2 = 2
4. Two bonding pairs and no lone pairs = linear geometry
Exercise \(2\)
What is the shape of H2CO?
Answer
1. Carbon has 4 valence electrons
2. Add 3 electrons for the three σ-bonds: 4 + 3 = 7
3. Subtract one electron for each π-bond: 7 1 = 6
4. Divide the number of electrons by two to give the number of electron pairs: 6/2 = 3
5. Two bonding pairs and no lone pairs = triangular geometry
Lone pairs versus bonding pairs
The prediction of the detailed molecular structure (including bond angles) is not as simple as shown in Table \(1\).6. In molecules with either lone pair electrons or multiple (double or triple) bonds the angles about the central atom are distorted due to the increased electron repulsion (Figure \(1\).9). The differences in repulsion caused by a lone pair or a bonding pair may be rationalized in a simple manner by a lone pair taking up more space than a bonding pair (Figure \(1\).10).
Water is one of the classic cases in considering the issue of non-bonding (unshared) electron pairs.
1. Oxygen has 6 valence electrons
2. Add 2 electrons for the two σ-bonds: 6 + 2 = 8
3. Divide the number of electrons by two to give the number of electron pairs: 8/2 = 4
4. Two bonding pairs and two lone pairs = tetrahedral geometry
From this an idealized tetrahedral geometry would give the H-O-H angle as 109.5, however, from Figure \(1\).9 we know that the lone pair...lone pair repulsion is greater than the lone pair...bonding pair repulsion which is greater than the bonding pair...bonding pair repulsion, and thus, the H-O-H angle should be decreased from the ideal tetrahedral. The experimentally determined H-O-H angle in water is in fact 104.5.
Ethylene is a good case in considering the issue of multiple bonds. Ethylene contains both σ-bond and π-bond between the carbon atoms. This combination can be thought of as a super bond, and as such its effect is similar to a lone pair.
Carbon has 4 valence electrons
Add 3 electrons for the two σ-bonds: 4 + 3 = 7
Subtract one electron for each π-bond: 7 1 = 6
Divide the number of electrons by two to give the number of electron pairs: 6/2 = 3
Three bonding pairs and no lone pairs = triangular geometry
From this an idealized tetrahedral geometry would give the H-C-H angle as 120, however, the π-bond repulsion is greater than the σ-bond repulsion, and thus, the H-C-H angle should be decreased from the ideal tetrahedral. The experimentally determined H-O-H angle in water is in fact 118.3.
Resonance structures
If a molecule or ion has two or more resonance forms it is necessary to consider each form before angles are predicted. For example, the carbonate anion, CO3-, can be drawn as a single structure from which it would be predicted that two groups of O-C-O angles would result (Figure \(1\).11).
However, CO3- should actually be drawn in each of its resonance forms as in Figure \(1\).12.
From Figure \(1\).12 it is clear that the real structure will be an average of the three resonance forms, and hence there will be a single O-C-O angle = 120 (Figure \(1\).13).
Atom electronegativities
In an A-X bond where the atom electronegativities are very different the bonding pair is assumed to occupy less space than in bond between two atoms of similar electronegativities. As the bonding pair occupies less space it will repel neighboring electron pairs less.
For example, based on the above, a comparison of the tetrahedral compounds H2O and F2O would suggest that the F-O-F angle be smaller than the H-O-H angle since fluorine has a higher electronegativity than hydrogen (4.0 and 2.1, respectively). This is indeed observed (Figure \(1\).14).
In many cases, inter-ligand steric interactions can also be used to explain the dffierence in angle. For example, while fluorine is more electronegative than chlorine (4.0 versus 3.0, respectively) it is also significantly smaller (Table \(1\).7). Thus, a larger Cl-X-Cl angle than a F-X-F angle in the homologous compound, can be attributed to a greater steric interactions, rather than difference in electronegativities.
Ligand Covalent Radius (Å) Van der Waal radius (Å)
F 0.57 1.35
Cl 1.02 1.80
Table \(1\).7: Comparison of size between fluoride and chloride ligands.
Bent's rule
There is the potential for more than one isomer for molecules that adopt structures in which there is a symmetry difference between at least two of the ligand positions. For example a trigional bipyramidal compound of the formula PXY4 has two possible structures. One where the X occupies an axial position (Figure \(1\).15a) and the other where it occupies an equatorial position (Figure \(1\).15b).
Through the consideration of structures Henry Bent suggested a rule: More electronegative substituents prefer hybrid orbitals with less s character, and conversely, more electropositive substituents prefer hybrid orbitals with greater s character.
For example, in PFCl5 the fluorine is the most electronegative substituent and will therefore occupy the axial (p-character only) position.
Bibliography
• H. A. Bent, J. Chem. Educ., 1960, 37, 616.
• H. A. Bent, Chem. Rev., 1961, 61, 275.
• R. J. Gillespie, J. Chem. Educ., 1970, 47, 18.
• R. J. Gillespie, Chem. Soc. Rev., 1992, 21, 59.
• R. J. Gillespie and R. S. Nyholm, Quart. Rev., 1957, 11, 339. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/01%3A_General_Concepts_and_Trends/1.05%3A_Structure_and_Bonding_-_Valence_Shell_Electron_Pair_Repulsion_(VSEPR)_Theory.txt |
Introduction
In any sort of discussion of crystalline materials, it is useful to begin with a discussion of crystallography: the study of the formation, structure, and properties of crystals. A crystal structure is defined as the particular repeating arrangement of atoms (molecules or ions) throughout a crystal. Structure refers to the internal arrangement of particles and not the external appearance of the crystal. However, these are not entirely independent since the external appearance of a crystal is often related to the internal arrangement. For example, crystals of cubic rock salt (NaCl) are physically cubic in appearance. Only a few of the possible crystal structures are of concern with respect to simple inorganic salts and these will be discussed in detail, however, it is important to understand the nomenclature of crystallography.
Crystallography
Bravais lattice
The Bravais lattice is the basic building block from which all crystals can be constructed. The concept originated as a topological problem of finding the number of different ways to arrange points in space where each point would have an identical atmosphere. That is each point would be surrounded by an identical set of points as any other point, so that all points would be indistinguishable from each other. Mathematician Auguste Bravais discovered that there were 14 different collections of the groups of points, which are known as Bravais lattices. These lattices fall into seven different "crystal systems, as differentiated by the relationship between the angles between sides of the unit cell and the distance between points in the unit cell. The unit cell is the smallest group of atoms, ions or molecules that, when repeated at regular intervals in three dimensions, will produce the lattice of a crystal system. The lattice parameter is the length between two points on the corners of a unit cell. Each of the various lattice parameters are designated by the letters a, b, and c. If two sides are equal, such as in a tetragonal lattice, then the lengths of the two lattice parameters are designated a and c, with b omitted. The angles are designated by the Greek letters α, β, and γ, such that an angle with a specific Greek letter is not subtended by the axis with its Roman equivalent. For example, α is the included angle between the b and c axis.
Table $1$.8 shows the various crystal systems, while Figure $1$.16 shows the 14 Bravais lattices. It is important to distinguish the characteristics of each of the individual systems. An example of a material that takes on each of the Bravais lattices is shown in Table $1$.9.
System Axial lengths and angles Unit cell geometry
cubic a = b = c, α = β = γ= 90
tetragonal a = b c, α = β = γ= 90
orthorhombic a b c, α = β = γ= 90
rhombohedral a = b = c, α = β = γ 90
hexagonal a = b c, α = β = 90, γ = 120
monoclinic a b c, α = γ = 90, β 90
triclinic a b c, α β γ
Table $1$.8: Geometrical characteristics of the seven crystal systems.
Crystal system Example
triclinic K2S2O8
monoclinic As4S4, KNO2
rhombohedral Hg, Sb
hexagonal Zn, Co, NiAs
orthorhombic Ga, Fe3C
tetragonal In, TiO2
cubic Au, Si, NaCl
Table $1$.9: Examples of elements and compounds that adopt each of the crystal systems.
The cubic lattice is the most symmetrical of the systems. All the angles are equal to 90, and all the sides are of the same length (a = b = c). Only the length of one of the sides (a) is required to describe this system completely. In addition to simple cubic, the cubic lattice also includes body-centered cubic and face-centered cubic (Figure $1$.16). Body-centered cubic results from the presence of an atom (or ion) in the center of a cube, in addition to the atoms (ions) positioned at the vertices of the cube. In a similar manner, a face-centered cubic requires, in addition to the atoms (ions) positioned at the vertices of the cube, the presence of atoms (ions) in the center of each of the cubes face.
The tetragonal lattice has all of its angles equal to 90, and has two out of the three sides of equal length (a = b). The system also includes body-centered tetragonal (Figure $1$.16).
In an orthorhombic lattice all of the angles are equal to 90, while all of its sides are of unequal length. The system needs only to be described by three lattice parameters. This system also includes body-centered orthorhombic, base-centered orthorhombic, and face-centered orthorhombic (Figure $1$.16). A base-centered lattice has, in addition to the atoms (ions) positioned at the vertices of the orthorhombic lattice, atoms (ions) positioned on just two opposing faces.
The rhombohedral lattice is also known as trigonal, and has no angles equal to 90, but all sides are of equal length (a = b = c), thus requiring only by one lattice parameter, and all three angles are equal (α = β = γ).
A hexagonal crystal structure has two angles equal to 90, with the other angle ( γ) equal to 120. For this to happen, the two sides surrounding the 120 angle must be equal (a = b), while the third side (c) is at 90 to the other sides and can be of any length.
The monoclinic lattice has no sides of equal length, but two of the angles are equal to 90, with the other angle (usually defined as β) being something other than 90. It is a tilted parallelogram prism with rectangular bases. This system also includes base-centered monoclinic (Figure $1$.16).
In the triclinic lattice none of the sides of the unit cell are equal, and none of the angles within the unit cell are equal to 90. The triclinic lattice is chosen such that all the internal angles are either acute or obtuse. This crystal system has the lowest symmetry and must be described by 3 lattice parameters (a, b, and c) and the 3 angles (α, β, and γ).
Atom positions, crystal directions and Miller indicies
Atom positions and crystal axes
The structure of a crystal is defined with respect to a unit cell. As the entire crystal consists of repeating unit cells, this definition is sufficient to represent the entire crystal. Within the unit cell, the atomic arrangement is expressed using coordinates. There are two systems of coordinates commonly in use, which can cause some confusion. Both use a corner of the unit cell as their origin. The first, less-commonly seen system is that of Cartesian or orthogonal coordinates (X, Y, Z). These usually have the units of Angstroms and relate to the distance in each direction between the origin of the cell and the atom. These coordinates may be manipulated in the same fashion are used with two- or three-dimensional graphs. It is very simple, therefore, to calculate inter-atomic distances and angles given the Cartesian coordinates of the atoms. Unfortunately, the repeating nature of a crystal cannot be expressed easily using such coordinates. For example, consider a cubic cell of dimension 3.52 Å. Pretend that this cell contains an atom that has the coordinates (1.5, 2.1, 2.4). That is, the atom is 1.5 Å away from the origin in the x direction (which coincides with the a cell axis), 2.1 Å in the y (which coincides with the b cell axis) and 2.4 Å in the z (which coincides with the c cell axis). There will be an equivalent atom in the next unit cell along the x-direction, which will have the coordinates (1.5 + 3.52, 2.1, 2.4) or (5.02, 2.1, 2.4). This was a rather simple calculation, as the cell has very high symmetry and so the cell axes, a, b and c, coincide with the Cartesian axes, X, Y and Z. However, consider lower symmetry cells such as triclinic or monoclinic in which the cell axes are not mutually orthogonal. In such cases, expressing the repeating nature of the crystal is much more difficult to accomplish.
Accordingly, atomic coordinates are usually expressed in terms of fractional coordinates, (x, y, z). This coordinate system is coincident with the cell axes (a, b, c) and relates to the position of the atom in terms of the fraction along each axis. Consider the atom in the cubic cell discussion above. The atom was 1.5 Å in the a direction away from the origin. As the a axis is 3.52 Å long, the atom is (1.5/3.52) or 0.43 of the axis away from the origin. Similarly, it is (2.1/3.52) or 0.60 of the b axis and (2.4/3.5) or 0.68 of the c axis. The fractional coordinates of this atom are, therefore, (0.43, 0.60, 0.68). The coordinates of the equivalent atom in the next cell over in the a direction, however, are easily calculated as this atom is simply 1 unit cell away in a. Thus, all one has to do is add 1 to the x coordinate: (1.43, 0.60, 0.68). Such transformations can be performed regardless of the shape of the unit cell. Fractional coordinates, therefore, are used to retain and manipulate crystal information.
Crystal directions
The designation of the individual vectors within any given crystal lattice is accomplished by the use of whole number multipliers of the lattice parameter of the point at which the vector exits the unit cell. The vector is indicated by the notation [hkl], where h, k, and l are reciprocals of the point at which the vector exits the unit cell. The origination of all vectors is assumed defined as [000]. For example, the direction along the a-axis according to this scheme would be [100] because this has a component only in the a-direction and no component along either the b or c axial direction. A vector diagonally along the face defined by the a and b axis would be [110], while going from one corner of the unit cell to the opposite corner would be in the [111] direction. Figure $1$.17 shows some examples of the various directions in the unit cell. The crystal direction notation is made up of the lowest combination of integers and represents unit distances rather than actual distances. A [222] direction is identical to a [111], so [111] is used. Fractions are not used. For example, a vector that intercepts the center of the top face of the unit cell has the coordinates x = 1/2, y = 1/2, z = 1. All have to be inversed to convert to the lowest combination of integers (whole numbers); i.e., [221] in Figure $1$.17. Finally, all parallel vectors have the same crystal direction, e.g., the four vertical edges of the cell shown in Figure $1$.17 all have the crystal direction [hkl] = [001].
Crystal directions may be grouped in families. To avoid confusion there exists a convention in the choice of brackets surrounding the three numbers to differentiate a crystal direction from a family of direction. For a direction, square brackets [hkl] are used to indicate an individual direction. Angle brackets <hkl> indicate a family of directions. A family of directions includes any directions that are equivalent in length and types of atoms encountered. For example, in a cubic lattice, the [100], [010], and [001] directions all belong to the <100> family of planes because they are equivalent. If the cubic lattice were rotated 90, the a, b, and c directions would remain indistinguishable, and there would be no way of telling on which crystallographic positions the atoms are situated, so the family of directions is the same. In a hexagonal crystal, however, this is not the case, so the [100] and [010] would both be <100> directions, but the [001] direction would be distinct. Finally, negative directions are identified with a bar over the negative number instead of a minus sign.
Crystal planes
Planes in a crystal can be specified using a notation called Miller indices. The Miller index is indicated by the notation [hkl] where h, k, and l are reciprocals of the plane with the x, y, and z axes. To obtain the Miller indices of a given plane requires the following steps:
1. The plane in question is placed on a unit cell.
2. Its intercepts with each of the crystal axes are then found.
3. The reciprocal of the intercepts are taken.
4. These are multiplied by a scalar to insure that is in the simple ratio of whole numbers.
For example, the face of a lattice that does not intersect the y or z axis would be (100), while a plane along the body diagonal would be the (111) plane. An illustration of this along with the (111) and (110) planes is given in Figure $1$.18.
As with crystal directions, Miller indices directions may be grouped in families. Individual Miller indices are given in parentheses (hkl), while braces {hkl} are placed around the indices of a family of planes. For example, (001), (100), and (010) are all in the {100} family of planes, for a cubic lattice.
Description of crystal structures
Crystal structures may be described in a number of ways. The most common manner is to refer to the size and shape of the unit cell and the positions of the atoms (or ions) within the cell. However, this information is sometimes insufficient to allow for an understanding of the true structure in three dimensions. Consideration of several unit cells, the arrangement of the atoms with respect to each other, the number of other atoms they in contact with, and the distances to neighboring atoms, often will provide a better understanding. A number of methods are available to describe extended solid-state structures. The most applicable with regard to elemental and compound semiconductor, metals and the majority of insulators is the close packing approach.
Close packed structures: hexagonal close packing and cubic close packing
Many crystal structures can be described using the concept of close packing. This concept requires that the atoms (ions) are arranged so as to have the maximum density. In order to understand close packing in three dimensions, the most efficient way for equal sized spheres to be packed in two dimensions must be considered.
The most efficient way for equal sized spheres to be packed in two dimensions is shown in Figure $1$.19, in which it can be seen that each sphere (the dark gray shaded sphere) is surrounded by, and is in contact with, six other spheres (the light gray spheres in Figure $1$.19). It should be noted that contact with six other spheres the maximum possible is the spheres are the same size, although lower density packing is possible. Close packed layers are formed by repetition to an infinite sheet. Within these close packed layers, three close packed rows are present, shown by the dashed lines in Figure $1$.19.
The most efficient way for equal sized spheres to be packed in three dimensions is to stack close packed layers on top of each other to give a close packed structure. There are two simple ways in which this can be done, resulting in either a hexagonal or cubic close packed structures.
Hexagonal close packed
If two close packed layers A and B are placed in contact with each other so as to maximize the density, then the spheres of layer B will rest in the hollow (vacancy) between three of the spheres in layer A. This is demonstrated in Figure $1$.20. Atoms in the second layer, B (shaded light gray), may occupy one of two possible positions (Figure $1$.20a or b) but not both together or a mixture of each. If a third layer is placed on top of layer B such that it exactly covers layer A, subsequent placement of layers will result in the following sequence ...ABABAB.... This is known as hexagonal close packing or hcp.
The hexagonal close packed cell is a derivative of the hexagonal Bravais lattice system (Figure $1$.16) with the addition of an atom inside the unit cell at the coordinates (1/3,2/3,1/2). The basal plane of the unit cell coincides with the close packed layers (Figure $1$.21). In other words the close packed layer makes-up the {001} family of crystal planes.
The packing fraction in a hexagonal close packed cell is 74.05%; that is 74.05% of the total volume is occupied. The packing fraction or density is derived by assuming that each atom is a hard sphere in contact with its nearest neighbors. Determination of the packing fraction is accomplished by calculating the number of whole spheres per unit cell (2 in hcp), the volume occupied by these spheres, and a comparison with the total volume of a unit cell. The number gives an idea of how open or filled a structure is. By comparison, the packing fraction for body-centered cubic (Figure $1$.16) is 68% and for diamond cubic (an important semiconductor structure to be described later) is it 34%.
Cupic close packed: face-centered cubic
In a similar manner to the generation of the hexagonal close packed structure, two close packed layers are stacked (Figure $1$.19) however, the third layer (C) is placed such that it does not exactly cover layer A, while sitting in a set of troughs in layer B (Figure $1$.22), then upon repetition the packing sequence will be ...ABCABCABC.... This is known as cubic close packing or ccp.
The unit cell of cubic close packed structure is actually that of a face-centered cubic (fcc) Bravais lattice. In the fcc lattice the close packed layers constitute the {111} planes. As with the hcp lattice packing fraction in a cubic close packed (fcc) cell is 74.05%. Since face centered cubic or fcc is more commonly used in preference to cubic close packed (ccp) in describing the structures, the former will be used throughout this text.
Coordination number
The coordination number of an atom or ion within an extended structure is defined as the number of nearest neighbor atoms (ions of opposite charge) that are in contact with it. A slightly different definition is often used for atoms within individual molecules: the number of donor atoms associated with the central atom or ion. However, this distinction is rather artificial, and both can be employed.
The coordination numbers for metal atoms in a molecule or complex are commonly 4, 5, and 6, but all values from 2 to 9 are known and a few examples of higher coordination numbers have been reported. In contrast, common coordination numbers in the solid state are 3, 4, 6, 8, and 12. For example, the atom in the center of body-centered cubic lattice has a coordination number of 8, because it touches the eight atoms at the corners of the unit cell, while an atom in a simple cubic structure would have a coordination number of 6. In both fcc and hcp lattices each of the atoms have a coordination number of 12.
Octahedral and tetrahedral vacancies
As was mentioned above, the packing fraction in both fcc and hcp cells is 74.05%, leaving 25.95% of the volume unfilled. The unfilled lattice sites (interstices) between the atoms in a cell are called interstitial sites or vacancies. The shape and relative size of these sites is important in controlling the position of additional atoms. In both fcc and hcp cells most of the space within these atoms lies within two different sites known as octahedral sites and tetrahedral sites. The difference between the two lies in their coordination number, or the number of atoms surrounding each site. Tetrahedral sites (vacancies) are surrounded by four atoms arranged at the corners of a tetrahedron. Similarly, octahedral sites are surrounded by six atoms which make-up the apices of an octahedron. For a given close packed lattice an octahedral vacancy will be larger than a tetrahedral vacancy.
Within a face centered cubic lattice, the eight tetrahedral sites are positioned within the cell, at the general fractional coordinate of (n/4,n/4,n/4) where n = 1 or 3, e.g., (1/4,1/4,1/4), (1/4,1/4,3/4), etc. The octahedral sites are located at the center of the unit cell (1/2,1/2,1/2), as well as at each of the edges of the cell, e.g., (1/2,0,0). In the hexagonal close packed system, the tetrahedral sites are at (0,0,3/8) and (1/3,2/3,7/8), and the octahedral sites are at (1/3,1/3,1/4) and all symmetry equivalent positions.
Important structure types
The majority of crystalline materials do not have a structure that ts into the one atom per site simple Bravais lattice. A number of other important crystal structures are found, however, only a few of these crystal structures are those of which occur for the elemental and compound semiconductors and the majority of these are derived from fcc or hcp lattices. Each structural type is generally defined by an archetype, a material (often a naturally occurring mineral) which has the structure in question and to which all the similar materials are related. With regard to commonly used elemental and compound semiconductors the important structures are diamond, zinc blende, Wurtzite, and to a lesser extent chalcopyrite. However, rock salt, β-tin, cinnabar and cesium chloride are observed as high pressure or high temperature phases and are therefore also discussed. The following provides a summary of these structures. Details of the full range of solid-state structures are given elsewhere.
Diamond Cubic
The diamond cubic structure consists of two interpenetrating face-centered cubic lattices, with one offset 1/4 of a cube along the cube diagonal. It may also be described as face centered cubic lattice in which half of the tetrahedral sites are filled while all the octahedral sites remain vacant. The diamond cubic unit cell is shown in Figure $1$.23. Each of the atoms (e.g., C) is four coordinate, and the shortest interatomic distance (C-C) may be determined from the unit cell parameter (a).
$C-C = a\dfrac{\sqrt{3}}{4} \approx 0.442a$
Zinc blende
This is a binary phase (ME) and is named after its archetype, a common mineral form of zinc sulfide (ZnS). As with the diamond lattice, zinc blende consists of the two interpenetrating fcc lattices. However, in zinc blende one lattice consists of one of the types of atoms (Zn in ZnS), and the other lattice is of the second type of atom (S in ZnS). It may also be described as face centered cubic lattice of S atoms in which half of the tetrahedral sites are filled with Zn atoms. All the atoms in a zinc blende structure are 4-coordinate. The zinc blende unit cell is shown in Figure $1$.24. A number of inter-atomic distances may be calculated for any material with a zinc blende unit cell using the lattice parameter (a).
$Zn-S = a\dfrac{\sqrt{3}}{4} \approx 0.442a$
$Zn-Zn = \dfrac{a}{\sqrt{2}} \approx 0.707a$
Chalcopyrite
The mineral chalcopyrite CuFeS2 is the archetype of this structure. The structure is tetragonal (a = b c, α = β = γ = 90, and is essentially a superlattice on that of zinc blende. Thus, is easiest to imagine that the chalcopyrite lattice is made-up of a lattice of sulfur atoms in which the tetrahedral sites are filled in layers, ...FeCuCuFe..., etc. (Figure $1$.25). In such an idealized structure c = 2a, however, this is not true of all materials with chalcopyrite structures.
Rock salt
As its name implies the archetypal rock salt structure is NaCl (table salt). In common with the zinc blende structure, rock salt consists of two interpenetrating face-centered cubic lattices. However, the second lattice is offset 1/2a along the unit cell axis. It may also be described as face centered cubic lattice in which all of the octahedral sites are filled, while all the tetrahedral sites remain vacant, and thus each of the atoms in the rock salt structure are 6-coordinate. The rock salt unit cell is shown in Figure $1$.26. A number of inter-atomic distances may be calculated for any material with a rock salt structure using the lattice parameter (a).
$Na-Cl = \dfrac{a}{2} \approx 0.5a$
$Na-Na = Cl-Cl = \dfrac{a}{\sqrt{2}} \approx 0.707a$
Cinnabar
Cinnabar, named after the archetype mercury sufilde, HgS, is a distorted rock salt structure in which the resulting cell is rhombohedral (trigonal) with each atom having a coordination number of six.
Wurtzite
This is a hexagonal form of the zinc sulfide. It is identical in the number of and types of atoms, but it is built from two interpenetrating hcp lattices as opposed to the fcc lattices in zinc blende. As with zinc blende all the atoms in a wurtzite structure are 4-coordinate. The wurtzite unit cell is shown in Figure $1$.27. A number of inter atomic distances may be calculated for any material with a wurtzite cell using the lattice parameter (a).
$Zn-S = a\sqrt{\dfrac{3}{8}} \approx 0.612a = \dfrac{3c}{8} = 0.375c$
$Zn-Zn = S-S = a = 1.632c$
However, it should be noted that these formulae do not necessarily apply when the ratio a/c is different from the ideal value of 1.632.
Cesium Chloride
The cesium chloride structure is found in materials with large cations and relatively small anions. It has a simple (primitive) cubic cell (Figure $1$.16) with a chloride ion at the corners of the cube and the cesium ion at the body center. The coordination numbers of both Cs+ and Cl-, with the inner atomic distances determined from the cell lattice constant (a).
$Cs-Cl = \dfrac{a\sqrt{3}}{2} = 0.866a$
$Cs-Cs = Cl-Cl = a$
β-Tin
The room temperature allotrope of tin is β-tin or white tin. It has a tetragonal structure, in which each tin atom has four nearest neighbors (Sn-Sn = 3.016 Å) arranged in a very flattened tetrahedron, and two next nearest neighbors (Sn-Sn = 3.175 Å). The overall structure of β-tin consists of fused hexagons, each being linked to its neighbor via a four-membered Sn4 ring.
Defects in crystalline solids
Up to this point we have only been concerned with ideal structures for crystalline solids in which each atom occupies a designated point in the crystal lattice. Unfortunately, defects ordinarily exist in equilibrium between the crystal lattice and its environment. These defects are of two general types: point defects and extended defects. As their names imply, point defects are associated with a single crystal lattice site, while extended defects occur over a greater range.
Point defects: too many or too few or just plain wrong
Point defects have a significant effect on the properties of a semiconductor, so it is important to understand the classes of point defects and the characteristics of each type. Figure $1$.28 summarizes various classes of native point defects, however, they may be divided into two general classes; defects with the wrong number of atoms (deficiency or surplus) and defects where the identity of the atoms is incorrect.
Interstitial Impurity
An interstitial impurity occurs when an extra atom is positioned in a lattice site that should be vacant in an ideal structure (Figure $1$.28b). Since all the adjacent lattice sites are filled the additional atom will have to squeeze itself into the interstitial site, resulting in distortion of the lattice and alteration in the local electronic behavior of the structure. Small atoms, such as carbon, will prefer to occupy these interstitial sites. Interstitial impurities readily diffuse through the lattice via interstitial diffusion, which can result in a change of the properties of a material as a function of time. Oxygen impurities in silicon generally are located as interstitials.
Vacancies
The converse of an interstitial impurity is when there are not enough atoms in a particular area of the lattice. These are called vacancies. Vacancies exist in any material above absolute zero and increase in concentration with temperature. In the case of compound semiconductors, vacancies can be either cation vacancies (Figure $1$.28c) or anion vacancies (Figure $1$.28d), depending on what type of atom are missing.
Substitution
Substitution of various atoms into the normal lattice structure is common, and used to change the electronic properties of both compound and elemental semiconductors. Any impurity element that is incorporated during crystal growth can occupy a lattice site. Depending on the impurity, substitution defects can greatly distort the lattice and/or alter the electronic structure. In general, cations will try to occupy cation lattice sites (Figure $1$.28e), and anion will occupy the anion site (Figure $1$.28f). For example, a zinc impurity in GaAs will occupy a gallium site, if possible, while a sulfur, selenium and tellurium atoms would all try to substitute for an arsenic. Some impurities will occupy either site indiscriminately, e.g., Si and Sn occupy both Ga and As sites in GaAs.
Antisite Defects
Antisite defects are a particular form of substitution defect, and are unique to compound semiconductors. An antisite defect occurs when a cation is misplaced on an anion lattice site or vice versa (Figure $1$.28g and h). Dependent on the arrangement these are designated as either AB antisite defects or BA antisite defects. For example, if an arsenic atom is on a gallium lattice site the defect would be an AsGa defect. Antisite defects involve tting into a lattice site atoms of a different size than the rest of the lattice, and therefore this often results in a localized distortion of the lattice. In addition, cations and anions will have a different number of electrons in their valence shells, so this substitution will alter the local electron concentration and the electronic properties of this area of the semiconductor.
Extended Defects: Dislocations in a Crystal Lattice
Extended defects may be created either during crystal growth or as a consequence of stress in the crystal lattice. The plastic deformation of crystalline solids does not occur such that all bonds along a plane are broken and reformed simultaneously. Instead, the deformation occurs through a dislocation in the crystal lattice. Figure $1$.29 shows a schematic representation of a dislocation in a crystal lattice. Two features of this type of dislocation are the presence of an extra crystal plane, and a large void at the dislocation core. Impurities tend to segregate to the dislocation core in order to relieve strain from their presence.
Epitaxy
Epitaxy, is a transliteration of two Greek words epi, meaning "upon", and taxis, meaning "ordered". With respect to crystal growth it applies to the process of growing thin crystalline layers on a crystal substrate. In epitaxial growth, there is a precise crystal orientation of the lm in relation to the substrate. The growth of epitaxial films can be done by a number of methods including molecular beam epitaxy, atomic layer epitaxy, and chemical vapor deposition, all of which will be described later.
Epitaxy of the same material, such as a gallium arsenide lm on a gallium arsenide substrate, is called homoepitaxy, while epitaxy where the lm and substrate material are different is called heteroepitaxy. Clearly, in homoepitaxy, the substrate and lm will have the identical structure, however, in heteroepitaxy, it is important to employ where possible a substrate with the same structure and similar lattice parameters. For example, zinc selenide (zinc blende, a = 5.668 Å) is readily grown on gallium arsenide (zinc blende, a = 5.653 Å). Alternatively, epitaxial crystal growth can occur where there exists a simple relationship between the structures of the substrate and crystal layer, such as is observed between Al2O3 (100) on Si (100). Whichever route is chosen a close match in the lattice parameters is required, otherwise, the strains induced by the lattice mismatch results in distortion of the lm and formation of dislocations. If the mismatch is signicant epitaxial growth is not energetically favorable, causing a textured lm or polycrystalline untextured lm to be grown. As a general rule of thumb, epitaxy can be achieved if the lattice parameters of the two materials are within about 5% of each other. For good quality epitaxy, this should be less than 1%. The larger the mismatch, the larger the strain in the lm. As the lm gets thicker and thicker, it will try to relieve the strain in the lm, which could include the loss of epitaxy of the growth of dislocations. It is important to note that the <100> directions of a lm must be parallel to the <100> direction of the substrate. In some cases, such as Fe on MgO, the [111] direction is parallel to the substrate [100]. The epitaxial relationship is specified by giving first the plane in the film that is parallel to the substrate [100].
Bibliography
• International Tables for X-ray Crystallography. Vol. IV; Kynoch Press: Birmingham, UK (1974).
• B. F. G. Johnson, in Comprehensive Inorganic Chemistry, Pergamon Press, Vol. 4, Chapter 52 (1973).
• A. R. West, Solid State Chemistry and its Applications, Wiley, New York (1984). | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/01%3A_General_Concepts_and_Trends/1.06%3A_Structure_and_Bonding_-_Crystal_Structure.txt |
Stereo isomers
Stereo isomers have the same empirical formula or molecular formula but different structural formulas. A typical example is butane, C4H10, which can have two different possible structures: n-butane (Figure \(1\).30a) and iso-butane also known as 1-methylpropane (Figure \(1\).30b).
Geometric isomers
Geometric isomers have the same empirical formula or molecular formula and also the same structural formula, but have a different relative arrangement of the substituent groups. For example, the two geometric isomers of 1,2-dichloroethene (Figure \(1\).31) have the molecular formula of C2H2Cl2, and the same structural formula of Cl(H)C=C(H)Cl, but the relative position of the two chlorine atoms can either be the same side of the C=C double bond (i.e., cis, see Figure \(1\).31a) or on opposite sides of the C=C double bond (i.e., trans, see Figure \(1\).31b). The use of cis and trans is not limited to organic compounds such as olefins, but can also be used in metal complexes, e.g., Figure \(1\).32.
When it is not possible to describe geometric isomers by the terms cis or trans, the terms facial (fac, Figure \(1\).33a) or meridinal (mer, Figure \(1\).33b) are commonly employed.
Optical isomers
If the mirror object and mirror image of a molecule are not the same (i.e., they are not superimposable) they are known as enantiomers. If they also have optical activity they are called chiral and are described as having chirality. Enantiomers of a particular compound have the same general properties with two exceptions:
1. Their behavior to polarized light.
2. Their reaction with other chiral molecules
It is possible to determine if a molecule is chiral or not from its symmetry. Chiral molecules will have no symmetry or axis of rotation. The optical activity of a chiral molecule will turn the plane of polarized light either to the right (+) or the left (-). The former is known as dextrorotatory (D), while the latter is known as levorotatory (L).
Configuration
The configuration of a chiral molecule may be represented in a number of ways. Lactic acid (C3H6O3) provides a suitable and simple example. The two optical isomers of lactic acid can be represented in 3-dimensional (3D) form as shown in Figure \(1\).34; however, it is also possible to draw the same molecules in 2D using the Fischer projection (Figure \(1\).35). While the Fisher projection does not appear to dene the geometry about the central carbon atom, the convention is that the side chains are out of the plane of the page (towards the observer), while the top and bottom groups are into the plane of the page (away from the observer). Thus, Figure \(1\).35a is a representation of Figure \(1\).34a, and Figure \(1\).35b is a representation of Figure \(1\).34b.
Fischer-Rosenoff convention
The Fischer-Roseno convention defines chirality of a molecule using the Fischer projection. The convention is based upon D-(+)-glyceraldehydes (Figure \(1\).36).
To determine the chiral label for a molecule, the Fischer projection is drawn with the longest carbon chain pointing away from the observer, i.e., into the plain of the page. The number 1 carbon (that of the highest substitution) is positioned at the top of the diagram. If the functional group is on the right of the diagram the isomer is given the label D (from the dextrorotatory enantiomer of glyceraldehydes rather than the molecule being dextrorotatory per se), i.e., Figure \(1\).36. On the other hand if the functional group is on the left of the diagram the isomer is given the label L (from the levorotatory enantiomer of glyceraldehyde rather than the molecule being levorotatory per se).
Rules
Where molecules are large and complex a number of additional rules are useful:
1. Make any alcohol substituent to be nearest to the top, e.g., a butan-2-ol (Figure \(1\).37a) rather than a butan-3-ol (Figure \(1\).37b).
2. Aldehydes and ketones are the number 1 carbon atom.
3. If there are both sets of substituents, rule 1 has precedence over rule 2.
Nomenclature in terms of R and S
The R/S nomenclature is a more widely used alternative to the D/L nomenclature, and is based upon the hierarchical order of the substituents. From the Fischer projection, the groups are ordered with the largest atom first, as shown in Figure \(1\).38. The lowest numbered substituent is then oriented away from the observer (i.e., H in Figure \(1\).38). The order of the remaining substituents 1, 2, 3 is then traced: if it goes clockwise the molecule is labeled R, if it goes anti-clockwise (counter-clockwise) the molecule is labeled S.
The advantage of the R/S methodology is that it can also be used for geometric isomers. Thus, while cis and trans are useful when the groups either side of a C=C double bond are the same (e.g., cis-1,2dichloroethene, Figure \(1\).39a); however, when four different groups are present (e.g., A, B, X, and Y as seen in Figure \(1\).39b) then an alternative must be used. By taking either group A/B or X/Y and placing in order of precedence, the relative order defines the isomer. Thus, if in Figure \(10\)b A>B and X>Y the label Z (zusammen from the German for together) is used. Conversely, if in Figure \(10\)b A<B and X>Y the label E (entgegen from the German for opposite) is used.
More complex molecules
As molecules become more complex with multiple functional groups and multiple chiral centers the D/L nomenclature has the potential to become confusing. Erythrose has two chiral centers, and based upon the Fischer projections it may be seen that since the alcohol substituents in the 2 and 3 positions are on the same side of the backbone in the Fischer projection irrespective if it is D-(-)-erythrose or L-(+)-erythrose (Figure \(1\).40a and b). Thus, the naming is simplified despite the presence of multiple side groups. However, in threose (Figure \(1\).40c) the substituents are on different sides of the Fischer projection raising the issue as to whether this should be L or D. However, in such cases the substituent at the bottom of the side chain takes priority, thus Figure \(1\).40c is D-(-)-threose.
Note
D-(-)-erythrose and L-(+)-erythrose are enantiomers, and D-(+)-threose and L-(+)-threose are also enantiomers, but threose and erythrose are diasteriomers. Enantiomers have the same physical properties (just different behavior to polarized light), while diasteriomers may have different physical properties such as melting points and solubility.
Amino Acids
Amino acids (or more properly α-amino acids) are compounds containing both an amine (NH2) group and a carboxylic acid (CO2H) attached to the same carbon atom. The presence of both acid and base groups can result in the formation of the Zwitterionic form with ammonium (NH3+) and carboxylate (CO2-) groups.
In defining the labels for amino acids, the basis is D-(-)-allothreonine (Figure \(1\).41), whereas for an amino acid such as threonine the top substituent is used for a label, e.g., Figure \(1\).42.
Similar chiral centers
Just because a compound has more than one chiral center does not mean that it is optically active. A consideration of erythritol shows the presence of a mirror plane of symmetry (Figure1.43). As such, erythritol is not optically active, i.e., there is no effect on polarized light. In contrast, while threitol has the same molecular and geometric formula, the lack of a mirror plane of symmetry (Figure \(1\).44) means that the L- and D+ forms are optically active.
Compounds with a D:L ratio of 1:1 are called racemic compounds and are totally optically inactive. However, it is also possible for a racemic solution of a compound to crystallize to form crystals of pure D or L, which may be manually separated.
Optical isomers that are not tetrahedral
Pyrimidal molecules
Pyramidal molecules can also exhibit chirality when the three substituents are distinct, e.g., PR1R2R3 (Figure \(1\).45). Unfortunately, most pyramidal compounds undergo an inversion of their isomers (in a similar manner to turning an umbrella inside out), such that the chiral forms interconvert rapidly. In such a case, it is not possible to resolve (separate) the different forms. Phosphines may generally be resolved because the barrier to their inversion is suficiently large (ca. 132 kJ/mol)
The notation used for pyramidal molecules is the same as that for tetrahedral molecules, in that the individual substituents are ordered with the largest atom first. The lone pair is defined as having the lowest numbering and thus is oriented away from the observer. The order of the remaining substituents 1, 2, and 3 is then traced: if it goes clockwise the molecule is labeled R, if it goes anti-clockwise (counter-clockwise) the molecule is labeled S.
Chirality in octohedral complexes
The presence of a chelate ligand on an octahedral complex can induce chirality in the complex.
Note
A chelate ligand is a molecule or ion that is bonded to at least two points of a central atom or ion. The term chelate is from the Greek chelè, meaning claw.
Unlike tetrahedral based compounds, chiral octahedral compounds have their nomenclature based on the structure of a helix. For example, in the case of a bis-chelate complex (Figure \(1\).46), with one of the chelate ligands pointing straight up the helix, the direction of the other defines the chirality of the helix, i.e., if it is a left or right hand helix. If the ligand points up to the left the complex is assigned the symbol shown in Figure \(1\).46a, while if the ligand points up to the right the complex is assigned the symbol shown in Figure \(1\).46b.
In complexes with three chelate ligands (i.e., a tris-chelate complex, Figure \(1\).47), the same methodology applies, in that any two ligands can be chosen and the same rules applied.
Methods of resolution of racemic mixtures
A racemic mixture, or racemate, is one that has equal amounts of left- and right-handed enantiomers of a chiral molecule. The first known racemic mixture was "racemic acid," which Louis Pasteur (Figure \(1\).48) found to be a mixture of the two enantiomeric isomers of tartaric acid (Figure \(1\).49).
In nature, it is common that only one of an optical isomers is naturally produced, but in laboratory synthesis it is more common that both isomers are made in equal amounts. The separation of a racemate into its components (the pure enantiomers) is called a chiral resolution.
Mechanical seperation
The crystals of enantiomerically pure compounds often have different appearances. Thus, just as Pasteur did in 1848, it is often possible to look under a microscope and physically separate the two different enantiomers.
Resolution by formation of diasteriomers
One of the differentiating properties of a chiral molecule is that each enantiomer reacts with another chiral molecule to form a diastereomeric pair of compounds. For example, a racemic acid, (+)HA and (-)HA, will react with a chiral base, (-)B, to form a mixture of diasteriomers, [(-)BH+.(+)A-] and [(-)BH+.(-)A-]. Because diasteriomers have different physical properties they can be separated by recrystallization. Once separated by recrystallization, the addition of excess acid will liberate the enantiomerically pure compound, i.e., (+)HA or (-)HA. A typical chiral base would be a natural alkaloid base to ensure that it is pure + or -. If it is a base that is needed to be separated, then (-)malic acid is a suitable acid.
Resolution by chromatography
In a related method to resolution by the formation of diasteriomers, a chiral column material will allow for the chromatographic separation of an enantiomeric mixture. Thus, the retention rate for (+)X will be different from (-)X on a column in which the stationary phase is chiral, i.e., (-)A. Many common chiral stationary phases are based on oligosaccharides such as cellulose or cyclodextrin (in particular with β-cyclodextrin).
Labile stereo isomers and racemization
It is only possible to separate (resolve) racemic mixtures if the molecule stays as one form for a long time. In other words, if there is a mechanism by which the two forms are interconverted, then resolution cannot be achieved.
Intramolecular rearrangement
Intramolecular rearrangements involve no bonds being broken. The classic example is the inversion of an amine, (1.12), which has a low energy barrier (24.7 kJ/mol).
Another example of an intramolecular rearrangement is the conversion of a square-based pyramidal geometry via a trigonal bipyramidal geometry to the isomeric square-based pyramidal geometry, or the alternative isomerization of one trigonal bipyramidal geometry to another via a square-based pyramidal transition state. Such a process is known as a Berry rotation. The Berry mechanism is a pseudorotation process for simultaneously interchanging two equatorial groups with the two axial groups, while the third equatorial group (called the pivot group) remains an equatorial group (Figure \(1\).50).
Intermolecular processes
Intermolecular processes involve bond breaking (and bond formation). For example a tetrahedral compound could loose one ligand, which creates a labile pyramidal compound. Rapid inversion of the pyramidal compound is followed by re-attachment of the ligand. Such a process is often so fast that resolution cannot be achieved.
Intramolecular processes
Intramolecular processes with bond breaking (and making) also lead to racemization. For example, in an octahedral complex with chelate ligands, if one of the ends of a chelate ligand detaches, it can reattach in the same configuration, or attach differently to change the chirality. In this manner the chirality can be changed from one form to another.
Conformation
The conformation of a molecule arises from the rotation of a single bond (Figure \(1\).51). However, even though there is rotation about the bond, there are energy barriers due to steric interactions of the substituents. In order to understand (and predict) these interactions, it is necessary to visualize the molecule in such a manner as to highlight the across-bond interactions; this is done by using a Newman projection.
A Newman projection, useful in alkane stereochemistry, visualizes chemical conformations of a carbon-carbon chemical bond from front to back, with the front carbon represented by a dot and the back carbon as a circle (Figure \(1\).52). The front carbon atom is called proximal, while the back atom is called distal. This type of representation is useful for assessing the torsional angle between bonds. Using ethane as an example, the Newman projection along the C-C single bond results in two basic conformations: eclipsed (Figure \(1\).52a) and staggered (Figure \(1\).52b). The staggered conformation will be energetically favored since the substituents are the most distant from each other. Conversely, the staggered will be the highest in energy. The difference in energy between the two conformations will define the barrier to rotation. In the case of ethane this is very small (12.5 kJ/mol).
Although we often only consider the two extreme conformations, in reality there is a continuum around 360◦ rotation of the C-C bond. Figure \(1\).53 shows the relative energy as a function of the dihedral angle for ethane. Since each carbon in ethane has equivalent substitution (i.e., three H atoms) the energy for each staggered conformation is the same. This is not true for more complex molecules such as butane (Figure \(1\).54).
As may be seen from Figure \(1\).54, the staggered conformation in which the two methyl groups (represented by the black circles) are as far away from each other (anti) is the most energetically favored. The other two staggered conformations (gauche) are mirror images of each other and are hence conformation enantiomers. It should also be noted from Figure \(1\).54 that the eclipsed conformations vary in energy since the presence of methyl...methyl near neighbors is clearly less energetically favorable than methyl...H near neighbors.
Generally the rotation about C-C bonds has a low barrier to rotation, however, if substituents are suciently bulky the molecule will not twist around the bond, e.g., substituted bi-phenyl with sterically bulky substituents (Figure \(1\).55).
Free rotation about a C-C bond is not fully possible when you have a ring system, e.g., in a cyclic compound such as cyclohexane, C6H12. The limited rotation about the C-C bonds results in the ipping of the ring conformation from the chair form (Figure \(1\).56a) to the boat form (Figure \(1\).56b). Since the boat form has steric hindrance between the hydrogen atoms, the chair form is the more stable.
Conformation of compounds with lone pairs
Lone pairs often behave in a different manner to substituents in regard to conformations. Thus, methylamine (CH3NH2) would be predicted to have the staggered conformation shown in Figure \(1\).57a. However, the lower steric bulk of the lone pair results in the nitrogen being 0.09 Å away from the true center of the CH3 projection (Figure \(1\).57b).
In compounds with more than one lone pair, the lowest energy form is not always the anti conformation. For example, at 20 C hydrazine (H2NNH2) is 100% gauche (Figure \(1\).58a); but for substituted hydrazines (i.e., a diamine, R2NNR2) if the substituent groups are suficiently large then the anti conformation will dominate (Figure \(1\).58b).
The conformation of hydrogen peroxide (H2O2) is dominated by the lone pairs rather than the hydrogen atoms. Instead of the expected anti conformation (c.f., Figure \(1\).54 where the black circles would represent the hydrogen atoms) in the free state the dihedral angle is 94 (Figure \(1\).59). The conformation of hydrogen peroxide is therefore neither eclipsed nor staggered but an intermediate structure. When in the solid state hydrogen bonding will cause the shape and angles to change.
Bibliography
• R. S. Berry, J. Chem. Phys., 1960, 32, 933.
• A. Keys, S. G. Bott, and A. R. Barron, J. Chem. Cryst., 1998, 28, 629. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/01%3A_General_Concepts_and_Trends/1.07%3A_Structure_and_Bonding_-_Stereochemistry.txt |
The choice of solvent is an important parameter for any chemical reaction. The following provides a guide to some of the consideration to be made in choosing a solvent to ensure the desired reaction occurs.
Solvation
Solvation may be defined as the interaction between the solvent and the solute, however, two general classes of solvation have different consequences to the stability of either reagents or products in a chemical reaction, and hence the potential of a reaction to occur
• Specific solvation is where the solvent interacts with one of the ions (or molecules) in solution via a covalent interaction. Furthermore, there will be a specific number of solvent molecules bound to each ion (or molecule), e.g., [Cu(NH3)4]2+ and [Mg(H2O)6]2+ (Figure $1$.60a).
• Non-specific solvation is as a result of van der Waals or dipole-dipole forces between the solvent and an ion (or molecule). There will be no defined number of interactions and the solvent...ion interaction will be highly fluxional, e.g., while water solvates the chloride ion (Figure $1$.60b) the number of water molecules around each anion is not fixed
Table $1$.10 shows the ability of three solvents to act with specific and non-specific solvation. The relative solvation ability of each solvent results in three different products from the dissolution of iron(III) chloride (FeCl3).
Solvation DMSO (Me2SO) Pyridine (C5H5N) Acetonitrile (MeCN)
Specific Good Very good Poor
Non-specific Good Poor Moderately good
Dissolution of FeCl3 in DMSO results in the dissociation of a chloride ligand, (1.8.1), due to both the specific solvation of the FeCl2+ cation and the non-specific solvation of the Cl- anion. In fact, the good solvation properties of DMSO means that depending on the concentration (and temperature) a series of dissociations may occur, (1.8.2).
$FeCl_3 \overset{DMSO}{\rightleftharpoons} [FeCl_2(DMSO)_4]^+ Cl^-_{(solv)}$
$[FeCl_2(DMSO)_4]^+ \xleftrightarrow[-Cl^-_{(solv)}]{+DMSO}[FeCl(DMSO)_5]^{2+} \xleftrightarrow[-Cl^-_{(solv)}]{+ DMSO} [Fe(DMSO)_6]^{3+}$
In contrast, if FeCl3 is dissolved in pyridine (py) the neutral Lewis acid-base complex is formed, (1.8.3), because while pyridine is a very good at specific solvation (Table $1$.10), it is poor at solvating the chloride anion.
$FeCl_3 + py \rightarrow FeCl_3(py)$
In a similar manner, FeCl3(MeCN)3 will be formed by the dissolution in acetonitrile, because although it is not good at specific solvation, it is not sufficiently good at non-specific solvation to stabilize the chloride anion. However, since the FeCl4- anion has a lower charge density that Cl-, it can be supported by the non-specific solvation of acetonitrile and thus a disproportionation reaction occurs, (1.8.4).
$2 FeCl_3(MeCN)_3 \rightleftharpoons [FeCl_2(MeCN)_4]^+ + FeCl_4^-$
Interference by the solvent
Rather than solvating a molecule or ion, the solvent can take an active and detrimental role in the synthesis of a desired compound.
Solvolysis
The archetypal solvolysis reaction is the reaction with water, i.e., hydrolysis, (\ref{1.8.5}). However, solvolysis is a general reaction, involving bond breaking by the solvent. Thus, the reaction with ammonia is ammonolysis, (\ref{1.8.6}), the reaction with acetic acid is acetolysis, (\ref{1.8.7}), and the reaction with an alcohol is alcoholysis, (\ref{1.8.8}) where Et = $\ce{C2H5}$. In each case the same general reaction takes place yielding the cation associated with the solvent.
$\ce{SO_2Cl_2 + 4 H_2O \rightarrow SO_2(OH)_2 + 2 H_3O^+ + 2 Cl^-} \label{1.8.5}$
$\ce{SO_2Cl_2 + 4 NH_3 \rightarrow SO_2(NH_2)_2 + 2 NH_4^+ + 2 Cl^-} \label{1.8.6}$
$\ce{ SO_2Cl_2 + 4 MeCO_2H \rightarrow SO_2(O_2CMe)_2 + 2 MeCO_2H_2^+ + 2 Cl^-} \label{1.8.7}$
$\ce{SO_2Cl_2 + 4 EtOH \rightarrow SO_2(OEt)_2 + 2 EtOH_2^+ + 2 Cl^-} \label{1.8.8}$
Competition reactions
Where more than one reaction could occur the reaction involving the solvent can often compete with the desired reaction.
If it is desired to synthesis the Lewis acid-base complex between diethylether (Et2O) and boron triuoride (BF3) it is important to choose a solvent that will not compete with the complex formation. For example, pyridine is a poor choice because the nitrogen donor is a stronger Lewis base than the diethylether, (1.8.9), and thus no reaction would occur between diethylether and boron trifluoride. In contrast, since acetonitrile (MeCN) is a poor Lewis base, then the reaction will occur.
$BF_3(Et_2O) + py \rightarrow BF_3(py) + Et_2O$
If the synthesis of GeCl62- from germanium tetrachloride (GeCl4) and a source of chloride anion, then water would be a poor choice of solvent since hydrolysis of GeCl4 would result. Liquid hydrogen chloride would be equally poor solvent since strong Cl-...H-Cl hydrogen bonding would stabilize the chloride anion and preclude reaction. In contrast, nitromethane (CH3NO2) would be a polar enough solvent to solvate the GeCl4, but it will be displaced by the chloride anion, which would be only weakly solvated.
Salt formation
The formation of a salt via a double displacement reaction, (1.8.10), can be promoted by the choice of solvent by shifting the equilibrium by stabilization of one or more reagent/product.
$MX + M'X' \rightleftharpoons MX' + M'X$
Salt stabilization through relative acidity
The attempted formation of nitronium perchlorate from nitric acid and perchloric acid, (1.8.11), in water will result in the decomposition of the NO2+ cation, (1.24). However, if the reaction is carried out in a stronger acid, i.e., sulfuric acid, the NO2+ cation is stable, and the resulting salt can be recystallized, (1.8.12).
$HNO_3 + HClO_4 \rightarrow NO_2^+ + ClO_4^- + H_2O$
$NO_2^+ + 2 H_2O \rightarrow HNO_3 + H_3O^+$
In a similar manner, the weak basic character of water means the equilibrium reaction, (1.25), has a very small equilibrium constant, K. However, if the reaction is carried out in a strongly basic solvent such as ammonia the uride anion is stabilized, (1.26), and can be precipitated by cation exchange.
Salt stabilization through solvation
The following observations may be explained by a consideration of the solvation ability of the solvent.
1. The reaction of silver nitrate with barium chloride in water yields silver chloride and barium nitrate, (1.8.13).
2. The reaction of barium nitrate with silver chloride in ammonia yields barium chloride and silver nitrate, (1.8.14).
$2 AgNO_3 + BaCl_2 \xrightarrow{H_2O} 2 AgCl\downarrow + Ba^{2+} + 2 NO_3^-$
$2 AgCl + Ba(NO_3)_2 \xrightarrow{NH_2} BaCl_2\downarrow + 2 Ag^+ + 2 NO_3^-$
Since silver nitrate and barium nitrate are soluble in both solvents, the differences must be due to differences in the solubility of the chlorides in each solvent. A consideration of the relative stability of solid silver chloride versus the solvated species (Figure $1$.61) shows that the enthalpy of solvation in water is less than the lattice energy. Thus, if silver chloride were present as Ag+ and Cl- in water it would spontaneously precipitate. In contrast, the enthalpy of solvation in ammonia is greater than the lattice energy, thus solid AgCl will dissolve readily in liquid ammonia. The reason for the extra stabilization from the specific solvation of the silver cation by the ammonia, i.e., the formation of the covalent complex [Ag(NH3)2]+.
As may be seen from Figure $1$.62, the opposite effect occurs for barium chloride. Here the enthalpy of solvation in ammonia is less than the lattice energy. Thus, if barium chloride were present as Ba2+ and Cl- in ammonia it would spontaneously precipitate. In contrast, the enthalpy of solvation in water is greater than the lattice energy, thus solid BaCl2 will dissolve readily in water. The stabilization of Ba2+(aq) occurs because water will have a larger sphere of non-specific solvation as a consequence of having two lone pairs, allowing interaction with the Ba2+ as well as other water molecules (Figure $1$.63).
Electron transfer reaction
A consideration of the oxidation, (1.8.15), and reduction, (1.8.16), reactions that occur for pure water at neutral pH (where [H+] = 10-7) would suggest that water will not tolerate oxidants whose E0 is greater than 0.82 V nor tolerate reductants whose E0 is less than -0.41 V.
$2 H^+ + \dfrac{1}{2} O_2 + 2 e^- \rightarrow H_2O \space\space\space\space\space\space E_0=+0.82V$
$H^+ + e^- \rightarrow \dfrac{1}{2} H_2 \space\space\space\space\space\space E_0 = -0.41V$
Thus, while water has a fair range to support redox reactions it is not very good at the extremes with strong reducing agents or strong oxidizing agents. Liquid ammonia is an excellent solvent for very strong reducing agents because of the stabilization of solvated electrons, i.e., [e-(NH3)6]. In contrast, hydrochloric acid is a good solvent for reactions involving very strong oxidizing agents. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/01%3A_General_Concepts_and_Trends/1.08%3A_Acids_Bases_and_Solvents_-_Choosing_a_Solvent.txt |
When discussing the combustion of a compound it is ordinarily referring to the reaction of an organic compound (hydrocarbon) with oxygen, in which the carbon is converted to carbon dioxide (CO2) and the hydrogen forms water (H2O) as a vapor, e.g., (1.9.1).
$CH_4 + 2 O_2 \rightarrow CO_2 + 2 H_2O$
However, this only a narrow view of combustion, and a more general denition should be that combustion or burning is the sequence of exothermic chemical reactions between a fuel and an oxidant accompanied by the production of heat and conversion of chemical species. Based upon this definition methane can combust in the presence of fluorine (F2) as strong oxidant, (1.9.2).
$CH_4 + 4 F_2 \rightarrow CF_4 + 4 HF$
In considering the combustion of any flammable compound, for example gasoline, it should be noted that the compounds that make up gasoline are quiet stable in the absence of a source of oxygen (usually from the air). Furthermore, some form of energy input (heat, flame, or spark) must be provided. Thus, the combustion of gasoline provides the archetypal example of the three component explosive system typical of a traditional chemical explosive: fuel (something that will burn), an oxidizer (usually a source of oxygen), and energy (ignition). In this regard, combustion also includes the exothermic reactions of many metals with oxygen, (1.9.3).
$6Al + 3 O_2 \rightarrow 2 Al_2O_3$
It is not just reactive metals that can be used as the fuel component of combustion, but many of their compounds as well. The formation of water from the hydrogen in organic compounds, in combination with an oxygen source, releases significant energy. It stands to reason therefore that any compound comprising of hydrogen and an element can be a potential fuel: a compound of hydrogen and another element is known as a hydride. This is especially true for the hydrides of reactive metals such as aluminum and sodium (and metalloids such as boron, (1.9.4)), but is also true for the hydrides of silicon and phosphorus. These hydride compounds react with an oxidizer in a manner analogous to that of a hydrocarbon, as may be seen by a comparison of (1.9.5) and (1.9.1).
$2BH_3 + 3 O_2 \rightarrow B_2O_3 + 3 H_2O$
$SiH_4 + 2 O_2 \rightarrow SiO_2 + 2 H_2O$
Figure $1$.64 shows a comparison of the heat of combustion of various inorganic fuels with commonly used jet fuels (JP-10 and JP-8). Hydrocarbons heat of combustion is limited by the C:H ratio, while graphite is the limiting case. Boron and boron compounds have greater volumetric and gravimetric energy density than hydrocarbons, and therefore were studied as potential high-energy fuels.
Note
Ammonia has been proposed as a practical alternative to fossil fuel for internal combustion engines. The energy value of ammonia is 22.5 MJ/kg, which is about half that of diesel. In a normal engine, in which the water vapor is not condensed, the caloric value of ammonia will be about 21% less than this value; however, it can be used in existing engines with only minor modifications to carburetors/fuel injectors.
Oxygen (O2) from the air does not have to be the source of oxidizer. Alternatives such as hydrogen peroxide (H2O2), nitrous oxide (N2O), and nitrates (e.g., ammonium nitrate, NH4(NO3)) are all sources of oxygen for combustion or explosions. However, this does not mean that every compound containing oxygen can be an oxidizer. For example, alcohols such as methanol will burn (in the presence of additional oxygen), but they will not act as an oxidizer. The oxygen within any compound must be reactive. By this we mean that it must be able to be released, preferentially as the more reactive oxygen atom (O) rather than O2, or be attached (chemically bonded to an element that wants to get rid of the oxygen (i.e., an element that is readily reduced). The most likely element in this case is nitrogen, with sulfur and phosphorus also potential candidates. Almost all compounds containing nitrogen bonded to oxygen can act as an oxidizer. Generally, the more oxygen atoms attached to nitrogen the more reactive the compound.
The efficiency of a chemical reaction such as combustion is dependent on how well the fuel and oxidizer are mixed at the molecular scale. Obviously the best situation is if both components are in the same molecule. Self-oxidizers are compounds containing oxygen in a reactive form as well as a suitable fuel (carbon or hydrogen). The most common self-oxidizers are organic nitrates. It should be pointed out that in spite of the presence of reactive oxygen, self-oxidizers may still require an external source of oxygen to ensure complete reaction, and some form of energy input (ignition) is still required.
Oxygen balance
While many compounds contain oxygen that does not mean they will combust efficiently in the absence of an external oxidizer, or whether they have sufficient oxygen to completely self-combust, or whether they can act as an oxidizer for other compounds. The simplest test for a compound's potential to fulfill these roles is its oxygen balance. The oxygen balance for a chemical is the amount of oxygen needed or produced to ensure the complete oxidation of all the carbon, hydrogen, or other elements.
Compounds such as trinitrotoluene (TNT, Figure $1$.65a) have a negative oxygen balance since extra oxygen is needed for complete formation of all the CO2 and H2O possible, (1.9.6). Thus, despite its reputation as an explosive, TNT is only efficient in the presence of an external oxidant, which may be air or another compound that provides a positive oxygen balance.
$2 C_7H_5N_3O_6 + 21 "O" \rightarrow 14 CO_2 + 5 H_2O + 3N_2$
In contrast to TNT, performic acid (Figure $1$.65b) is an example of a compound with a zero oxygen balance: it has all the oxygen it needs for complete combustion, (1.9.7), and hence only requires energy to detonate making at a much more dangerous compound per se than TNT.
$2 CH_2O_3 \rightarrow CO_2 + H_2O$
A positive oxygen balance means that the compound liberates oxygen surplus to its own needs, for example the decomposition of ammonium nitrate provides one atom of oxygen per molecule, (1.9.8). Clearly, any compound with a positive oxygen balance makes a good oxidizer and is highly incompatible with combustible chemicals.
$NH_4(NO_3) \rightarrow 2 H_2O + N_2 + "O"$
Quantification of oxygen balance
A quantification of oxygen balance allows for the determination of approximate ratio of reagents to optimize combustion/explosion. In this regard oxygen balance is defined as the number of moles of oxygen (excess or deficient) for 100 g of a compound of a known molecular weight (Mw), (1.8.9) where x = number of atoms of carbon, y = number of atoms of hydrogen, z = number of atoms of oxygen, and m = number of atoms of metal oxide produced.
$OB\% = \dfrac{-1600}{M_w} \times (2x + \dfrac{y}{2} + m - z)$
Example $1$
In the case of TNT, the Mw = 227.1 g/mol, the number of carbon atoms (x) = 7, the number of hydrogen atoms (y) = 5, the number of oxygen atoms (z) = 6, and the number of atoms of metal oxide produced (m) = 0. Therefore:
$OB\% = \dfrac{-1600}{227.1} \times (14 + 2.5 + 0 - 6) = -74\%$
A summary of selected oxygen balance values is given in Table $1$.11.
Element or Compound Oxygen balance (%)
Carbon -266.7
Sulfur -100
Aluminum powder -89
Trinitrotoluene -74
Nitroglycerine +3.5
Ammonium nitrate +20
Ammonium perchlorate +27
Potassium chlorate +26
Sodium chlorate +45
Sodium nitrate +37
Tetranitromethane +57
Lithium perchlorate +45
Table $1$.11: Oxygen balance of selected compounds. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/01%3A_General_Concepts_and_Trends/1.09%3A_Chemical_Reactivity_-_The_Basics_of_Combustion.txt |
Within the main group (s- and p-block) elements of the Periodic Table (Figure $1$.66) there are some general trends that we can observe for the elemental form, as well as the hydrides, oxides, and halides.
Periodic trends for the main group elements
Within the main group (s- and p-block) elements there are some general trends that we can observe.
• The further down a given Group the elements have increased metallic character, i.e., good conductors of both heat and electricity, and exhibit delocalized bonding.
• Moving from left to right across a Period the elements have greater non-metallic character, they are insulators with localized bonding.
• Within the p-block at the boundary between the metallic elements (Figure $1$.66, grey elements) and nonmetal elements (Figure $1$.66, green elements) there is positioned boron and silicon that are metalloid in character (Figure $1$.66, pink elements), i.e., they have low electrical conductivity but it increases with temperature.
As an example of these changes Table $1$.12 shows the trends across one Period.
Element Na Mg Al Si P S Cl Ar
Properties Electropositive metal Electropositive metal Metal but forms covalent bonds Metalloid semiconductor metal/non-metal characteristics E-E bonding in elements E-E bonding in elements Simple molecule Monoatomic gas
Table $1$.12: Summary of trends for elements across the Periodic Table.
Periodic trends for the main group hydrides
The properties of main group hydrides are dependant on the difference in electronegativity between the element and hydrogen (Table $1$.13). Elements on the left of the Periodic Table are highly electropositive and form ionic hydrides, while those of the center and right are covalent in character. However, of those with covalent E-H bonds, there is a change from polymeric hydrides to molecular compounds. For example, the Group 13 element hydrides (i.e., BH3) form hydrogen-bridged oligomers (i.e., B2H6). In contrast, HCl is a diatomic molecule.
Hydride Element electronegativity Hydrogen electronegativity E-H polarity Structure Comments
NaH 0.9 2.1 M+H- Ionic Reacts with H2O to liberate H2
BH3 2.0 2.1 Bδ+-Hδ- Oligomeric and polymeric Reacts slowly with H2O
CH4 2.4 2.1 Cδ--Hδ+ Molecular Insoluble in H2O
HCl 3.0 2.1 Clδ--Hδ+ Molecular Dissolves in H2O to form H+ and Cl-
Table $1$.13: Summary of properties of selected main group hydrides as a function of the relative electronegativities.
Periodic trends for the main group oxides
As with hydrides the properties of main group oxides are dependant on the difference in electronegativity between the element and oxygen. Highly electropositive metals for ionic oxides, while other elements for covalent bonds (albeit polar in character) with oxygen. In addition, the aggregation of covalent oxides decreased across the Period from left to right (Table $1$.14). As may also be seen from Table $1$.14, oxides of elements on the left of the Periodic Table dissolve in water to form basic solutions, while those on the right form acidic solutions. There is a class of oxides (especially those of Group 13 and 14) that can react as either an acid or a base. These are known as amphoteric substances.
Note
The word is from the Greek prefix ampho meaning "both"
Oxide Bonding Reactivity with H2O
Description
Na2O Ionic Dissolves to give a strong base Basic
Al2O3 Covalent polymeric Dissolves in both acidic and basic solution Amphoteric
SiO2 Covalent polymeric Dissolves in both acidic and basic solution Amphoteric
CO2 Covalent molecular Dissolves to give a weak acid Acidic
SO3 Covalent molecular Dissolves to give a strong acid Aci
Table $1$.14: Comparison of oxides across the Periodic Table.
In summary, oxides of the main group elements show two trends.
1. From left to right across a Period, the oxides change from ionic → oligomeric/polymeric covalent → molecular covalent
2. From left to right across a Period, the oxides change from ionic → oligomeric/polymeric covalent → molecular covalent.
Periodic trends for the main group chlorides
The trend between ionic and non-ionic/covalent in moving across a Period is also true for the chlorides of the main group elements. Those on the left (i.e., Group 1 and 2) are ionic and soluble in water, while those to the right tend to give acidic solutions due to reactions with the water and the formation of hydrochloric acid, e.g., (1.10.1).
$SiCl_4 + 2 H_2O \rightarrow SiO_2 + 4 HCl_{(aq)}$ | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/01%3A_General_Concepts_and_Trends/1.10%3A_Periodic_Trends_for_the_Main_Group_Elements.txt |
Thumbnail: Zeppelin the Hindenburg on fire at the mooring mast of Lakehurst (United States of America) 6 May 1937. (Public Domain; Sam Shere via Wikipedia)
02: Hydrogen
Hydrogen gas, H2, was first artificially synthesized by Phillip von Hohenheim (known as Paracelsus, Figure \(2\).1) by mixing metals with strong acids. He was unaware that the flammable gas produced by this chemical reaction was a new chemical element. In 1671, Robert Boyle (Figure \(2\).2) rediscovered the reaction between iron lings and dilute acids, which results in the production of hydrogen gas. He noted that these fumes were highly flammable and that the flame gave o a lot of heat but not much light.
In 1766, Henry Cavendish (Figure \(2\).3) was the first to recognize hydrogen gas as a discrete substance, by identifying the gas from a metal-acid reaction as flammable air. One of the richest men in Britain at the time, he lived in London and spent his time in his private laboratory at his home. In 1781 he was the first person to nd that the gas produces water when burned. This was a key experiment in disproving the Aristotelian theory of the four elements. As a consequence of his work he is usually given credit for its discovery as an element. However, it was Antoine Lavoisier (Figure \(2\).4) who in 1783 named the element hydrogen (from the Greek hydro meaning water and genes meaning creator) after he reproduced Cavendish's findings.
The four Aristotalian elements were Earth, Fire, Air and Water. A fifth element, Aether, was ascribed as a divine substance that made up the heavenly spheres and heavenly bodies (stars and planets).
Using his invention, the vacuum flask, James Dewar (Figure \(2\).5) was the first to liquefy hydrogen in 1898. He produced solid hydrogen the next year.
2.02: The Physical Properties of Hydrogen
Chemical Symbol H
Atomic number 1
Electron configuration 1s1
Atomic weight 1.00794
Melting point -259 C
Boiling point -253 C
Density (gas) 0.090 g/L
Density (liquid) 0.70g/L
Ionization potential
13.598 eV
Pauling electronegativity 2.1
Ionic radius (H-) 1.54 Å
van der Waal radius 1.2 Å
Explosive limit 4 - 75%
Ignition temperature 585 C
Table \(2\).1: Physical properties of hydrogen.
2.03: Synthesis of Molecular Hydrogen
Although hydrogen is the most abundant element in the universe, its reactivity means that it exists as compounds with other elements. Thus, molecular hydrogen, H2, must be prepared from other compounds. The following outlines a selection of synthetic methods.
Steam reforming of carbon and hydrocarbons
Many reactions are available for the production of hydrogen from the reaction of steam with a carbon source. The choice of reaction is guided by the availability of raw materials and the desired purity of the hydrogen. The simplest reaction involves passing steam over coke at high temperatures (1000C).
$\ce{C(s) + H2O(g) \rightarrow H2(g) + CO(g)} \nonumber$
Coke is a grey, hard, and porous carbonaceous material derived from destructive distillation of low-ash, low-sulfur bituminous coal. As an alternative to coke, methane may be used at a slightly higher temperature (1100 C).
$\ce{ CH4(g) + H2O(g) \rightarrow 3 H2(g) + CO(g)} \nonumber$
In each case the carbon monoxide formed in the reaction can react further with steam in the presence of a suitable catalyst (usually iron or cobalt oxide) to generate further hydrogen.
$\ce{CO(g) + H2O(g) \leftrightharpoons H2(g) + CO2(g)}\nonumber$
This reaction is known as the water gas-shift reaction, and was discovered by Italian physicist Felice Fontana (Figure $2$.6) in 1780.
The dominant industrial process for hydrogen production uses natural gas or oil refinery feedstock in the presence of a nickel catalyst at 900C.
$C_2H_{8(g)} + 3 H_2O_{(g)} \rightarrow 7 H_{2(g)} + 3 CO_{(g)} \nonumber$
Electrolysis of water
Electrolysis of acidified water in with platinum electrodes is a simple (although energy intensive) route to hydrogen.
$2 H_2O_{(l)} \rightarrow 2 H_{2(g)} + O_{2(g)}\nonumber$
On a larger scale hydrolysis of warm aqueous solutions of barium hydroxide can yield hydrogen of purity greater than 99.95%. Hydrogen is also formed as a side product in the production of chlorine from electrolysis of brine (NaCl) solutions in the presence of a mercury electrode.
$2NaCl_{(aq)} + 2 Hg_{(l)} \rightarrow Cl_{2(g)} + 2 NaHg_{(l)}\nonumber$
The sodium mercury amalgam reacts with water to yield hydrogen.
$NaHg_{(l)} + 2 H_2O_{(l)} \rightarrow H_{2(g)} + 2NaOH_{(aq)} + 2Hg_{(l)}\nonumber$
Thus, the overall reaction can be written as:
$NaCl_{(aq)} + H_2O_{(l)}\rightarrow H_{2(g)} + 2 NaOH_{(aq)} + Cl_2\nonumber$
However, this method is being phased out for environmental reasons.
Reaction of metal with acid
Hydrogen is produced by the reaction of highly electropositive metals with water, and less reactive metals with acids, e.g.,
$Fe_{(s)} + 2H_3O^+_{(aq)} \rightarrow H_{2(g)} + 2 H_2O_{(aq)} + Fe^{2+}_{(aq)}\nonumber$
This method was originally used by Henry Cavendish (Figure $2$.7) during his studies that led to the understanding of hydrogen as an element (Figure $2$.8).
The same method was employed by French inventor Jacques Charles (Figure $2$.9) for the first ight of a hydrogen balloon on 27th August 1783. Unfortunately, terrified peasants destroyed his balloon when it landed outside of Paris.
Hydrolysis of metal hydrides
Reactive metal hydrides such as calcium hydride (CaH2) undergo rapid hydrolysis to liberate hydrogen.
$CaH_{2(s)} + 2H_2O_{(l)} \rightarrow 2 H_{2(g)} + 2 OH^-_{(aq)} + Ca^{2+}_{(aq)}\nonumber$
This reaction is sometimes used to inflate life rafts and weather balloons where a simple, compact means of generating H2 is desired.
2.04: Atomic Hydrogen
Atomic hydrogen has the electron conguration of 1s1 and as such represents the simplest atomic conguration. However, as a consequence there is dispute as to its proper position within the Periodic Table. Its electron conguration is similar to the valence electron conguration of the alkali metals (ns1) suggesting it be listed at the top of Group 1 (1A). However, its reaction chemistry is dissimilar to the alkali metals. Hydrogen is also one electron short of a Nobel gas conguration, and therefore it is possible to think of its relationship to the halogens.
Vapor phase
Atomic hydrogen (H.) is highly reactive and consequently has a short lifetime due to its reaction chemistry. Consequently, in order to generate and observe the reactivity they must be generate at low pressures.
Thermolysis of hydrogen compound (commonly halide) or photolysis at an energy above the bond dissociation energy results in the homoleptic cleavage of the H-X bond to generate the appropriate radical species.
$H-X \xleftrightarrow{\Delta \space or \space h\nu} H + X$
Alternatively, atomic hydrogen can be generated from elemental hydrogen.
$H-H \xleftrightarrow{\Delta \space or \space h\nu} H^. + H^.$
The reverse reaction (recombination of two hydrogen atoms) is highly exothermic (-434 kJ.mol-1) and forms the basis of the heat generated in arc welding.
Solution
Atomic hydrogen may be generated in aqueous solution through the solvation of electrons.
$e^-_{(aq)} + H_3O^+ \leftrightharpoons H^. + H_2O$
The formation equilibrium constant (Keq) is very small resulting in very low concentrations being generated (10-5 M). As expected solvated atomic hydrogen is a strong reducing agent.
$Ag^+ + H^. \rightarrow Ag + H^+$
$2I^- + 2H^. \rightarrow 2H^+ + I_2$
Solid state
Hydrogen atoms may be trapped in the solid state lattice upon generation by photolysis of HX. Observation by electron spin resonance (esr) of a signal split by s = 1/2 nucleus (i.e., 1H) results in a doublet with a coupling of 1428 MHz. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/02%3A_Hydrogen/2.01%3A_Discovery_of_Hydrogen.txt |
The proton, H+, is the name given to hydrogen in the +1 oxidation state.
Gas phase
The proton can be formed from the photolysis of atomic hydrogen in the vapor phase at low pressure.
$H^._{(g)} + h\nu \rightarrow H^+_{(g)} + e^-_{(g)}$
The proton is more reactive than the hydrogen atom because of its high charge density. In addition, the proton's small ionic radius, 1.5 x 10-15 cm, means that it can get close to other atoms and hence form strong bonds.
$H^+_{(g)} + NR_3 \rightarrow HNR^+_{3(g)}$
The strength of the bonding interaction is such that it is very hard to measure directly. Instead the relative bond strength between the proton and an appropriate base, B1, is measured in the presence of a competing base, B2.
$B_1H^+_{(g)} + B_{2(g)} \leftrightharpoons B_2H^+_{(g)} + B_{2(g)}$
In measuring the exchange reaction, the relative proton affinity of B1 and B2 is measured. This is also known as the gas phase acidity, and as such it is a measure of the inherent acidity of a species X-H because it obviates any solvent effects.
Liquid and solution
The high reactivity of the proton means that it does not exist free in solution. There are however many H+ containing species. These are generally classified as acids.
$B_1H^+_{(sol)} + B_{2(sol)} \leftrightharpoons B_2H^+_{(sol)} + B_{1(sol)} \ \text{acid base acid base}$
The reaction between the acid and the base is a proton transfer reaction. While the proton travels from B1 to B2, it is never free in solution. Instead a bridged transition state or intermediate is formed, B1...H+...B2.
Acidity and pH
The most common solvent for H+ is water. The acid form is usually defined as the hydronium ion or H3O+, (2.5.5). The terms oxonium, hydroxonium and oxidanium are also used for the H3O+. Although we commonly use H3O+ it is known from spectroscopy that larger complexes are formed such as H9O4+ (Figure $2$.10).
$H_2O + H^+ \rightarrow H_3O^+$
Acids and bases have been characterized in a number of different ways. In 1680 Robert Boyle (Figure $2$.11) defined an acid as a compound that dissolved many other compounds, had a sour taste, and reacted with alkali (base).
Boyle's simple observational description was rationalized by Danish physical chemist Johannes Brønsted (Figure $2$.12). Brønsted proposed that acids are proton donors, and bases are proton acceptors. An acid-base reaction is one in which a proton is transferred from a proton donor (acid) to a proton acceptor (base). Based upon Brønsted's proposal simple acids contain an ionizable proton. Examples of simple acids include neutral molecules (HCl, H2SO4), anions (HSO4-, H2PO4-), and cations (NH4+). The most common Brønsted bases include metal hydroxides (MOH).
Brønsted noted that when an acid donates a proton it forms a conjugate base. The following are examples of an acid and its conjugate base.
$H_2O \space\space \rightarrow H^+ + OH^- \ \text{acid conjugate base}$
$H_2SO_4 \space\space \rightarrow H^+ + HSO^-_4 \ \text{acid conjugate base}$
$NH^+_4 \space\space\space\space \rightarrow H^+ + NH_3 \ \text{acid conjugate base}$
Exercises
Exercise $1$\
What is the conjugate base of HCl?
Answer
Cl-
Exercise $2$\
What is the conjugate base of HSO4-
Answer
SO42-
Exercise $3$\
What is the conjugate base of [Al(H2O)6]3+
Answer
[Al(H2O)5(OH)]2+
The same occurs when a base accepts a proton it forms a conjugate acid. The following are examples of a base and its conjugate acid.
$H_2O \space\space + \space\space H^+ \rightarrow H_3O^+ \ \text{base conjugate acid}$
$HCO^-_3 \space\space + \space\space H^+ \rightarrow H_2CO_3 \ \text{base conjugate acid}$
$F^- \space\space + \space\space H^+ \rightarrow HF \ \text{base conjugate acid}$
Exercises
Exercise $4$\
What is the conjugate acid of NH3?
Answer
NH4+
Exercise $5$\
What is the conjugate acid of S2-?
Answer
HS-
Exercise $6$\
What is the conjugate acid of CO32-?
Answer
HCO3-
Thus, the reaction between an acid and a base results in the formation of the appropriate conjugate base and conjugate acid.
$acid_1 + base_1 \leftrightharpoons acid_2 + base_2$
A specific example is as follows:
$HNO_3 + NH_3 \leftrightharpoons NH^+_4 + NO^-_3 \ acid_1 \space\space\space\space base_1 \space\space\space\space\space\space\space acid_2 \space\space\space\space base_2$
Exercise $7$
What is the conjugate acid and base formed from the reaction of NH4+ with S2+?
Answer
$NH_4^+ + S^{2-} \leftrightharpoons HS^- + NH_3 \ acid_1 \space\space\space\space base_1 \space\space\space\space\space\space\space acid_2 \space\space\space\space base_2$
In the equilibrium reactions shown in (2.12) and (2.13) there is a competition between the two bases for the proton. As would be expected the strongest base wins.
When a strong acid is added to (dissolved in) water it will react with the water as a base:
$HCl + H_2O \leftrightharpoons H_3O^+ + Cl^- \ acid \space\space\space\space base \space\space\space\space\space\space\space acid \space\space\space\space base$
In contrast, when a strong base is added to (dissolved in) water it will react with the water as an acid:
$H_2O + NH_3 \leftrightharpoons NH_4^+ + OH^- \ acid \space\space\space\space base \space\space\space\space\space\space\space acid \space\space\space\space base$
pH a measure of acidity
The acidity of a water (aqueous) solution depends on the concentration of the hydronium ion, i.e., [H3O+]. The acidity of a solution is therefore the ability of the solution to donate a proton to a base. The acidity or pH of a solution is defined as:
$pH = -log [H_3O^+_a]$
It is important to note that the value is the activity of H3O+ and not the concentration.
Note
Activity is a measure of the eective concentration of a species in a mixture. The dierence between activity and other measures of composition such as concentration arises because molecules in non-ideal gases or solutions interact with each other, either to attract or to repel each other.
The activity of the H3O+ ion can be measured by
(a) A gas electrode
(b) Acid-base indicators
Proton Transfer Reactions
The proton transfer reaction is one of the simplest reactions in chemistry. It involves no electrons and low mass transfer/change, giving it a low energy of activation. For proton transfer between O-H or N-H groups and their associated bases the reaction is very fast. The proton transfer occurs across a hydrogen-bonded pathway during which the proton is never free. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/02%3A_Hydrogen/2.05%3A_The_Proton.txt |
The combination of hydrogen with another element produces a hydride, ExHy. The formal charge or oxidation state of the hydrogen in these compounds is dependant on the relative electronegativity of the element in question.
Ionic hydrides
Hydrogen compounds with highly electropositive metals, i.e., those in which the metal has an electronegativity of less than 1.2, are ionic with the hydrogen having a s2 configuration (H-). Typical ionic metal hydrides are those of the Group 1 (IA) metals and the heavier Group 2 (IIA) metals.
The ionic radius of the hydride ion is in between that of fluoride and chloride and the same as oxide (Table $2$.2). As a consequence, in the solid state the hydride ion replicates that of a halide ion (e.g., Cl-), and as such similar solid state structures are observed (Table $2$.3).
Ion Ionic Radius (Å)
H- 1.40
F- 1.36
Cl- 1.81
O2- 1.40
Table $2$.2: Selected ionic radii.
Metal Hydride Fluoride Chloride
Li 4.085 4.0173 5.129
Na 4.880 4.620 5.640
K 5.700 5.347 6.292
Rb 6.037 5.640 6.581
Cs 6.376 6.008 7.020
Table $2$.3: Lattice parameters (Å) for hydrides, and halides of the Group 1 metal salts with cubic rock salt structures.
Unlike the halide ions that are soluble in water, the hydride ion reacts with water, (2.32), and consequently NaH and CaH2 are commonly used as drying agents. The liberation of hydrogen was used as a commercial source of hydrogen for small-scale applications.
$H^- + H_2O \rightarrow OH^- + H_{2 (g)}$
WARNING
Group 1 and 2 metal hydrides can ignite in air, especially upon contact with water to release hydrogen, which is also flammable. Hydrolysis converts the hydride into the analogous hydroxide, which are caustic bases. In practice, most ionic hydrides are dispensed as a dispersion in oil, which can be safely handled in air.
Covalent Hydrides
The most common binary compounds of hydrogen are those in which hydrogen bonds have covalent bond character. The E-H bond is usually polar ranging from those in which the hydrogen is polarized positively (e.g., those with non-metals such as F, O, S, and C) to where it is negative (e.g., those with metals and metalloids such as B, Al, etc). Magnesium hydride is intermediate between covalent and ionic since it has a polymeric solid similar to AlH3, but reacts rapidly with water like ionic hydrides. Table $2$.4 lists the important covalent hydrides of p-block elements. It should be noted that all organic hydrocarbons can be thought of as simply the hydrides of carbon!
Group 13 Group 14 Group 15 Group 16 Group 17
B2H6 CnH2n+2, CnH2n, CnH2n-2 NH3, N2H4 H2O, H2O2 HF
(AlH3)n SinH2n+2 PH3, P2H4 H2S, H2Sn HCl
Ga2H6 GenH2n+2 AsH3 H2Se HBr
SnH4 SbH3 H2Te HI
The elements of Group 14 to 17 all form hydrides with normal covalent bonds in which the hydrogen is bonded by a single bond to the element in question. In contrast, the elements of Group 13 (as typified by boron) all exhibit a second type of covalent bond: the electron deficient hydrogen bridged bond. In this type of bond the hydrogen nucleus is embedded in a molecular orbital that covers more than two atoms to create a multi-center two-electron bond. Diborane (Figure $2$.13) represents the archetypal compound containing the hydrogen bridge bond. Hydrides are not limited to terminal (E-H) or those bridging two atoms (E-H-E), but are also known where the hydrogen bridges (or caps) more than two atoms, i.e., Figure $2$.14.
Synthesis of covalent hydrides
Covalent hydrides can be made by a range of synthetic routes. The simplest is direct combination of the elements (in a similar manner to that used with ionic hydrides).
$2H_2 + O_2 \rightarrow 2H_2O$
The use of a hydride as a reagent to reduce a halide or oxide of the desired element
$SiCl_4 + LiAlH_4 \rightarrow SiH_4 + LiAlCl_4$
Metal phosphides, carbides, silicides, and borides result in the formation of the hydride.
$Ca_3P_2 + 6 H_2O \rightarrow 2 PH_3 + 3 Ca(OH)_2$
Hydride compounds can be interconverted in the presence of a catalyst, heat, or an electrical discharge. This is the basis of catalytic cracking of petroleum mixtures.
Interstitial hydrides
Many transition metal, lanthanides and actinides absorb hydrogen to give a metallic hydride which retain the properties of a metal, although the presence of hydrogen does result in embrittlement of the metal. As such these hydrides are best considered as alloys since they do not have defined stoichiometries. For example, vanadium absorbs hydrogen to form an alloy with a maximum hydrogen content of VH1.6. In a similar manner palladium forms PdH0.6. The hydrogen atoms are present in interstitial sites in the metal's lattice; hence interstitial hydride. Interstitial hydrides show certain promise as a way for safe hydrogen storage.
Bibliography
• A. R. Barron, M. B. Hursthouse, M. Motevalli, and G. Wilkinson, J. Chem. Soc., Chem. Commun., 1986, 81. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/02%3A_Hydrogen/2.06%3A_Hydrides.txt |
Hydrogen bonds are formed between a species with a polar Xδ--Hδ+ bond and a species with a lone pair (Yδ-), i.e., Xδ--Hδ+...Yδ-. The most common species for X are oxygen and nitrogen, and to a lesser extent carbon, fluorine, and sulfur. However, as long as the X-H bond is polar then hydrogen bonding is possible. Similarly, the most common Lewis bases that hydrogen bond involve oxygen, nitrogen, and fluorine as the donor atom. Again there are many examples of other atoms, but as long as the atom has a lone pair that is chemically active, hydrogen bonding can occur.
The majority of hydrogen bonds are asymmetrical, that is the hydrogen is closer to one atom than the other (Figure \(2\).15), even when X and Y are the same element, i.e., O-H...O. While the typical hydrogen bond involves one Lewis base (lone pair donor), there are many examples where the hydrogen interacts with two Lewis base lone pairs (Figure \(2\).16).
Hydrogen bonds are mostly electrostatic attractions, and as such they are weaker than covalent bonds, but stronger than van der Waal interactions. With bond strengths generally covering the range of 5 50 kJ/mol, the energy required to break a hydrogen bond is comparable to that of thermal motion within the temperature range of 0 200C. As a consequence the number of groups involved in hydrogen bonding decreases with increasing temperature, until few hydrogen bonds are observed in the vapor phase. One noted exception is the hydrogen bridged anion [F-H-F]-, in which the strong interaction (243 kJ/mol) is covalent in character involving a three-center molecular orbital bond.
Classes of hydrogen bond
Although hydrogen bonds may be characterized with respect to the X and Y atom, it is more useful to classify them as either intramolecular or intermolecular hydrogen bonds. This is due to the difference in physical and chemical properties between these two classes.
Intramolecular
Intramolecular hydrogen bonds (X-H...Y) arise where the X and Y atoms are in the same molecule (Figure \(2\).17).
Intermolecular
If the hydrogen bond (X-H...Y) involves X and Y being from different molecules this is an intermolecular hydrogen bond. Within the range of intermolecular hydrogen bonded compounds there are two sub-categories: those involving discrete molecular species (oligomers) and those resulting in polymeric species.
Carboxylic acids are a typical example of a discrete oligomeric species that are held together by intermolecular hydrogen bonds (Figure \(2\).18a). A wide range of structurally analogous compounds also form head-to-tail hydrogen bonded dimers (e.g., Figure \(2\).19). In a polymeric hydrogen bonded species every molecule hydrogen bonds but in a random form. As an example, liquid primary alcohols form extended hydrogen networks (Figure \(2\).18b). Such an arrangement is labile and as such it is difficult to determine definitive speciation. Liquids that form this type of hydrogen-bonded network are known as associated liquids. In the solid state the networks generally adopt a more ordered structure. For example as is seen in the structure of ice.
Methods of study
The study of the structure arising from hydrogen bonding and the properties exhibited due to the presence of hydrogen bonds is very important.
Diffraction methods
X-raydiffraction of single crystals is the most common structural method employed to determine the presence, effect, and strength of a hydrogen bond. Unfortunately, in order for the location of the hydrogen to be determined with some degree of accuracy, diffraction data of a high quality is needed and/or low temperature (e.g., -196 C) data collection is required. Neutron scattering can be used where very accurate data is required because hydrogen atoms scatter neutrons better than they do X-rays. Figure \(2\).20 summarizes the key parameters that are obtained from X-ray (and neutron) diffraction experiments.
Given the electrostatic nature of a hydrogen bond between a polar X-H bond and a Lewis base it is reasonable that the X-H...Y angle (θ) is roughly linear (i.e., 180). However, it is not always so and nonlinear interactions are known where steric or conformational restrictions limit the orientation of the X-H bond with respect to Y.
The distance between X and Y, d(X...Y), is less than the sum of the van der Waal radii of X and Y (Table \(2\).5). This is in line with the relative strength of these interactions. As would be expected the shorter the X...Y distance the stronger the hydrogen bond.
X Y Sum of van der Waal radii (Å) Typical X...Y distance (Å)
O O 2.8 2.50-2.69
O N 2.9 2.75-2.85
N N 3.0 2.69-2.98
Table \(2\).5: Comparison of the X...Y distance in hydrogen bonded species with the sum of the van der Waal radii.
The bond distance to hydrogen, d(X-H), is often longer in hydrogen bonded species. For example the O-H distance for an alcohol in the absence of hydrogen bonding is typically 0.97 Å. In contrast, the value typically seen for a hydrogen-bonded analog is 1.05 Å.
Spectroscopy
Spectroscopy is a simple method of comparing hydrogen-bonded systems in particular in the solution or liquid phase.
Infra red and Raman
The X-H stretching frequency in the IR (and Raman) spectrum is dependant on the identity of X, i.e., O-H = 3610 - 3640 cm-1 and N-H = 3400 - 3500 cm-1. However, the ν(X-H) is shifted to lower energy (lower frequency) as a consequence of hydrogen bonding. In addition, while non-hydrogen bonded X-H stretches are sharp, the presence of hydrogen bonding results in the peak being broadened. Figure X demonstrates both these effects. The O-H stretch for dilute nBuOH in CCl4 is a sharp peak at 3650 cm-1 due to the lack of hydrogen bonding between the two components (Figure \(2\).21a), and the presence of hydrogen bonding between nBuOH and nBuOH is limited by the dilution. By contrast, a dilute solution of nBuOH in Et2O results in a shift to lower frequency and a significant increase of peak width (Figure \(2\).21b) as a result of fairly strong O-H...O bonds. Finally, a dilute solution of nBuOH in NMe3 results in a further shift to 3250 cm-1 and a very broad peak (Figure \(2\).21c). The broadening of the peaks is due to the distribution of X-H distances within a X-H...Y hydrogen bond.
NMR
The presence of hydrogen bonding results the shift to higher ppm (lower frequency) of the 1H NMR resonance for the proton. This shift is due to the decrease in shielding of the proton. A dilute solution of nBuOH in CCl4 shows a resonance typical of a non-hydrogen bonded compound (Figure \(2\).22a), while that for nBuOH in NMe3 (Figure \(2\).22b) shows a significant low field shift. Very strong intra or intermolecular hydrogen bonded species show a very large 1H NMR shift (e.g., Figure \(2\).22c).
Effects of hydrogen bonding
Physical effects
The presence of intermolecular hydrogen bonding provides additional attractive forces between molecules. Thus, properties that depend on intramolecular forces are affected.
Liquids with significant hydrogen bonding exhibit higher boiling points, higher viscosity, and higher heat of vaporization (∆Hv) as compared to analogous compounds without extensive hydrogen bonding. For solids the presence of hydrogen bonding results in an increase in the melting point of the solid and an increase in the associated heat of fusion (∆Hf).
The archetypal case for the effect of hydrogen bonding is the melting and boiling points of the hydrides of the Group 16 elements, i.e., H2E. For a series of analogous compounds with the same molecular structure it would be expected that the boiling points would be related to the molecular mass. However, as can be seen from Table \(2\).6, the melting and boiling points of water are anomalously higher than those of its heavier analogs. In fact from Figure \(2\).23 it is clear that just considering H2S, H2Se, and H2Te, the expected trend is observed, and it is similar to that for the Group 14 hydrides (CH4, SiH4, etc). Therefore, water must have additional intermolecular forces as compared to its heavier homologs. This observation is consistent with the strong hydrogen bonding in water, and the very weak if nonexistent hydrogen bonding in the sulfur, selenium, and tellurium analogs.
Compound Molecular weight (g/mol) Melting point (C) Boiling Point (C)
H2O 18.01 0 100
H2S 24.08 -85.5 -60.7
H2Se 80.98 -60.4 -41.5
H2Te 129.62 -49 -2
Table \(2\).6: Summary of physical properties for the hydrides of the Group 16 elements.
A similar but not as pronounced trend is observed for the Group 15 hydrides, where ammonia's higher values are associated with the presence of significant hydrogen bonding (Table \(2\).7).
Compound Molecular weight (g/mol) Melting point (C) Boiling Point (C)
NH3 17.03 -77.7 -33.35
PH3 34.00 -133.5 -87.4
AsH3 77.95 -113.5 -55
ShH3 124.77 -88.5 -17
Table \(2\).7: Summary of physical properties for the hydrides of the Group 15 elements.
Exercise \(1\)
Would you expect H2S2 to have a higher or lower boiling point that H2O2? Why?
Answer
The boiling point for H2O2 is 150.2 C, while that of 70.7 C. The difference in boiling point is due to the stronger intermolecular hydrogen bonding in H2O2 than in H2S2.
The types of hydrogen bond can also have a signicant effect on the physical properties of a compound. For example, the cis isomer of hydroxybenzaldehyde melts at 1 C, while the trans isomer has melting and boiling points of 112 C. Both compounds exhibit strong hydrogen bonding in the solid state, however, as may be seen from Figure \(2\).24a, cis-hydroxybenzaldehyde (salicylaldehyde) has a configuration that allows strong intramolecular hydrogen bonding, which precludes any intermolecular hydrogen bonding. The melting point of cis-hydroxybenzaldehyde is going to be controlled by the van der Waal forces between adjacent molecules. In contrast, since intramolecular hydrogen bonding is precluded in the trans isomer (Figure \(2\).24b) it can form strong intermolecular hydrogen bonds in the solid state, and thus, it is these that define the melting point. The boiling points are controlled in a similar manner.
Exercise \(2\)
4-Hydroxybenzoic acid melts at 213 C, while 2-hydroxybenzoic acid melts at 158 C. Explain this observation.
Answer
2-Hydroxybenzoic acid exhibits strong intramolecular hydrogen bonding while 4-hydroxybenzoic acid has strong intermolecular hydrogen bonding.
Melting and boiling are not the only physical properties that are aected by hydrogen bonding. Solubility can also be aected. Consider two isomers of C4H10O: nBuOH and Et2O. The n-butanol is much more soluble in water than diethyl ether. The reason for this is that while both compounds can hydrogen bond to water, those between nBuOH and water are much stronger than those between Et2O and water, and thus, dissolution of nBuOH in water does not disrupt the very strong hydrogen bonding in water as much as Et2O does.
Acid strength
The acidity of a protic species can be affected by the presence of hydrogen bonding. For example, consider the di-carboxylic acid derivatives of ethylene (fumaric acid). Each of the carboxylic acid groups has sequential equilibria that may be defined by the pK values, (2.36). Table \(2\).8 lists the pK values for the cis and trans isomers. The acidity of the first and second carboxylic group for the trans isomer is similar, the difference being due to the increased charge on the molecule. In contrast, the second proton in the cis isomer is much less acidic than the first proton: why? A consideration of the structure of the mono anion of the cis isomer (Figure \(2\).25) shows that a very strong intramolecular hydrogen bond is formed once the first proton is removed. This hydrogen bond makes the second acidic proton much more difficult to remove and thus lowers the acidity of the proton.
\[ HO_2C-HC=CH-CO_2H \xleftrightarrow{pK_1} O_2^-C-HC=CH-CO_2H \xleftrightarrow{pK_2} O_2^-C-HC=CH-CO_2^- \]
Isomer pK1 pK2 Ratio
Trans 3 4.5 25:1
Cis 1.9 6.2
10,000:1
Table \(2\).8: Acidity equilibrium constants for cis and trans isomers of fumaric acid.
Solid state structure
In the solid state, molecules that can form hydrogen bonds will tend to arrange themselves so as to maximize the formation of linear hydrogen bonds. The structure of ice is a typical example, where each water molecule's hydrogen bonds to four other molecules creating a diamond-like lattice (Figure \(2\).26). However, ice isn't the only compound whose solid state structure is defined by its hydrogen bonding. Cyanuric acid forms six strong intermolecular hydrogen bonds to three other molecules in the plane of the heterocyclic ring and thus creates a graphite-like structure (Figure \(2\).27).
Conformational stabilization
The presence of hydrogen bonding can stabilize certain conformations over others that in the absence of hydrogen bonding would be more favored. For example, based on steric considerations H2NCH2CH2C(O)H would be expected to adopt a staggered conformation as is typical for compounds with free rotation about a C-C bond. However, due to the presence of a strong intramolecular hydrogen bond it adopts a sterically disfavored eclipsed conformation (Figure \(2\).28).
This conformational stabilization is even more important in biopolymers such as peptides. If we assume that a peptide chain (Figure \(2\).29) made from a sequence of amino acids would not adopt any eclipsed conformations on steric grounds, and that the amide group (H-N-C=O) is near planar due to delocalization, then there will be 32 possible conformations per nitrogen in the peptide chain, i.e., three each for the C-C and C-N bonds (Figure \(2\).29). Assuming a modest peptide has 50 amino acids, there will be 32x50 potential conformations of the polymer chain, i.e., over 5 x 1047 conformations. However, in reality there are limits to the number of conformations observed since particular ones are stabilized by intramolecular hydrogen bonds. The most important of these are the α-helix (Figure \(2\).30) and β-sheet structures.
Bibliography
J. T. Leman, J. Braddock-Wilking, A. J. Coolong, and A. R. Barron, Inorg. Chem., 1993, 32, 4324.
K. Nakamoto, M. Margoshes, and R. E. Rundle, J. Am. Chem. Soc., 1955, 77, 6480.
M. B. Power, A. R. Barron, S. G. Bott, E. J. Bishop, K. D. Tierce and J. L. Atwood, J. Chem. Soc., Dalton Trans., 1991, 241.
R. Taylor and O. Kennard, Acc. Chem. Res., 1984, 17, 320. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/02%3A_Hydrogen/2.07%3A_The_Hydrogen_Bond.txt |
Physical effects
The presence of intermolecular hydrogen bonding provides additional attractive forces between molecules. Thus, properties that depend on intramolecular forces are affected.
Isotope Hydrogen-1 Hydrogen-2 Hydrogen-3
Special name Hydrogen Deuterium Tritium
Symbol H D T
Atomic number
1
1 1
Number of neutrons 0 1 2
Mass number 1 2 3
Natural abundance 99.9844% 0.0156% very small
Table $2$.9: Summary of isotopes of hydrogen.
Synthesis of deuterium compounds
Electrolysis of water
The electrolysis of hydrogen-1 water (H2O) in the presence of an alkali results in the formation of hydrogen and oxygen.
$2 H_2O_{(l)} \rightarrow 2H_{2(g)} + O_{2(g)}$
In a similar manner the hydrolysis of deuterated water (D2O) yields deuterium and oxygen.
$2D_2O_{(l)} \rightarrow 2 D_{2(g)} + O_{2(g)}$
However, the rate of electrolysis of D2O is slightly slower than that of H2O. Thus, the partial hydrolysis of water with a mixture of natural isotopes results in the slight enrichment of the water with D2O. The level of enrichment in one step is less than 1%. In order to obtain high levels of D2O (e.g., ca. 30%) it is necessary to reduce the original volume of water by 1/100,000th.
Chemical equilibrium
Proton exchange reactions can be used to enrich compounds in deuterium. For example, the reaction of HSD with water shown in (2.8.3) has a slight preference for the formation of H2S, i.e., Keq = 1.012. Thus, bubbling HSD through water results in the enrichment of the water in HOD. However, about 30% enrichment is about the best that can be achieved by this method.
$H_2O_{(l)} + HSD_{(g)} \leftrightharpoons HOD_{(l)} + H_2S_{(g)}$
Fractional distillation
The boiling point of H2O is (by definition) 100C, in contrast the boiling point of D2O is 101.4 C. Thus, it is possible to separate H2O from D2O by fractional distillation. This method provides the most suitable route to high isotopic enrichment and D2O of 99.8% can be produced this way.
Note
The term heavy water is used for D2O of greater than 99.8% enrichment.
Uses of deuterium compounds
Possible nuclear fusion
The largest use of D2O is as a moderator and heat exchanger for fission nuclear reactors, however, the biggest potential application will be if nuclear fusion is realized as a commercial process.
The fusion of two deuterium atoms to form a helium atom and energy would be one source of energy, (2.8.4), however, deuterium-tritium fusion is the most promising, (2.8.5).
$^2_1H + ^2_1H \rightarrow ^3_2He + ^1_0n +energy$
$^2_1H + ^3_1H \rightarrow ^4_2He + ^1_0n +energy$
The deuterium part of the fuel does not pose a great problem because about 1 part in 5000 of the hydrogen in seawater is deuterium. This amounts to an estimate that there is over 1015 tons of deuterium in the oceans. The tritium part of the fuel is more problematic since there is no significant natural source (Table $2$.9), and the tritium would have to be obtained by breeding the tritium from lithium.
$^6_3Li + ^1_0n \rightarrow ^4_2He + ^3_1H$
Since a gallon of seawater could produce as much energy as 300 gallons of gasoline, there is clearly a large amount of energy that can potentially be realized through nuclear fusion. Unfortunately, this advantage is also a disadvantage since the temperatures attained are similar to the surface of the sun, which would vaporize any conventional container. Fusion experiments therefore use a magnetic field to contain the reaction. The shape of the eld is like a bottle, hence the term "magnetic bottle".
One demonstrated fusion process is the so-called hydrogen bomb or thermonuclear bomb in which a fission atom bomb is used to initiate a fusion reaction. The atomic bomb is surrounded by a layer of lithium deuteride. Neutrons from the atomic explosion (fission) cause the lithium to be converted into helium, tritium, and energy, (2.8.6). The atomic explosion also supplies the 50,000,000 C temperature needed for the subsequent fusion of deuterium with tritium, (2.8.5). So in-fact the hydrogen bomb is misnamed and it should be called a deuterium bomb.
Note
The original calculations to model the hydrogen bomb were performed using ENIAC (short for Electronic Numerical Integrator And Computer) that was originally designed to generate tables of trajectories of shells red from large artillery. The artillery ring tables were made by women mathematicians who were called calculators hence the name used today. Built in 1946 ENIAC is often assumed to be the first programmable electronic computer, however, it was predated by the six Colossus machines that were used to successfully crack the German Enigma code as early as 1944. However, the existence of the Colossus machines was kept secret until 1975.
Spectroscopy
In the chemical laboratory deuterium compounds are commonly used in spectroscopy for:
a) The assignment of resonances in IR, Raman, and NMR spectroscopy.
b) As a non-proton containing solvent in 1H NMR spectroscopy.
A description of these applications is given below.
Reaction mechanism and rate determination
Given the larger mass of deuterium over hydrogen there is an associated dierence in the rate of reactions (see below) and therefore investigations using hydrogen and deuterium analogs can provide information as to reaction mechanisms.
The spectroscopic dierences between hydrogen and deuterium can also be used as a tracer to uniquely determine the source of particular substituents. For example, the magnesium (or Grignard) reduction of a ketone yields upon hydrolysis the secondary alcohol. If the reaction is carried out in a deuterated solvent and H2O used for hydrolysis then the secondary carbon is deuterated, (2.43). In contrast, if the reaction is carried out in a non-deuterated solvent and hydrolysis is accomplished with D2O then the deuterated alcohol is formed, (2.44). These experiments define that the initial reduction occurs at the ketone's α-carbon.
Exercise $1$
Given the following reactions and the isotope distribution of the products suggest the reaction mechanism.
Answer
Differences between hydrogen and deuterium
Properties that depend on nuclei properties
The nuclear magnetic moment of an atomic nucleus arises from the spins of the protons and neutrons within the nucleus. As a consequence the magnetic moment for hydrogen and deuterium are very different and hence the conditions for detection by NMR are very different. Thus, in observing the 1H NMR spectrum of a compound not only are the deuterium atoms not observed, but the coupling is now H-D rather than H-H (Figure $2$.31).
Deuterium is better at scattering neutrons than hydrogen. The H and D cross sections are very distinct and different in sign, which allows contrast variation in such experiments. Hydrogen's low electron density makes it difficult to determine its position by X-ray diffraction methods, neutron diffraction methods allow for highly accurate structure determination. Hydrogen can be seen by neutron diffraction and scattering, however, it has a large incoherent neutron cross-section. This is nil for deuterium and thus delivers much clearer signals may be obtained for deuterated samples. Neutron scattering of deuterated samples is indispensable for many studies of macromolecules in biology.
Properties that depend on mass
The difference in mass between hydrogen and deuterium obviously results in a difference in molecular mass of their analogous compounds. This difference can be used for analysis by mass spectrometry, but it also results in different densities of compounds. For example, the density of H2O at 25 C is 0.997 g/cm3, while the density of D2O at 25 C is 1.104 g/cm3.
The vibrational frequency for a diatomic molecule, H-X, can be defined by the equation:
$v_{H-X} = \dfrac{1}{2\pi} \sqrt{\dfrac{f_{H-X}}{\mu_{H-X}}}$
where fH-X is the H-X bond force constant, and µH-X is the reduced mass.
$\mu_{H-X} = \dfrac{m_h \cdot m_X}{m_H + m_X}$
If substitute H for D the D-X force constant is the same as the H-X force constant, but the reduced mass is twice the value for the H-X bond. As a result the ratio of the vibrational frequency of an H-X bond to that of the analogous D-X bond is given by the following equation.
$\dfrac{\mu_{D-X}}{\mu_{H-X}} = (\dfrac{\mu_{H-X}}{\mu_{D-X}})^{\dfrac{1}{2}} = \dfrac{1}{\sqrt{2}}$
With the change in vibrational energy there is concomitant change in the bond strength.
$E_{D-X} > E_{H-X}$
Thus, the rate of reactions will be faster for hydrogen derivative than the deuterium analog. The ratio of the rate constants will be dependant on the involvement of H-X bond breaking or forming in the rate limiting step (the slowest reaction step within the overall reaction mechanism). When an H-X bond is made or broken in the rate limiting step, then the ratio of the rate constants upon deuterium substitution will be:
$\dfrac{k_{H-X}}{k_{D-X}} \approx 7$
This is known as the primary isotope effect. In this case where H-X bond breaking or forming is not part of the rate limiting step, then the isotope effect will be much smaller and is known as a secondary isotope effect.
The position of equilibrium reactions that involve hydrogen exchange, (2.8.12), will be effected by the presence of deuterium to favor the deuterium being concentrated in the more stable bond. This is the basis of the concentration of HOD from HSD and water, (2.8.3).
$\text{X-H} + \text{Y-D} \leftrightharpoons \text{X-D} + \text{Y-H}$
Bibliography
• M. B. Power, S. G. Bott, J. L. Atwood, and A. R. Barron. J. Am. Chem. Soc., 1990, 112, 3446.
• A. S. Borovik and A. R. Barron, Main Group Chem. 2005, 4, 135. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/02%3A_Hydrogen/2.08%3A_Isotopes_of_Hydrogen.txt |
The process of nuclear fission and radioactive decay are both associated with the conversion of an atom with a large nucleus to an atom (or atoms) with a smaller nucleus. In the process, mass is lost and energy is produced. However, what happens if two atoms with small nuclei are combined to give a single atom with a larger nucleus? In such a process the nuclei would be fused together, and this process is called nuclear fusion.
One of the simplest fusion processes involves the fusing of two hydrogen-2 (deuterium) atoms:
$^2_1H + ^2_1H \rightarrow ^4_2He$
The mass of each deuterium atom is 2.0140 amu, while the mass of the resulting helium is 4.0026 amu. The mass defect of the reaction is 0.0254 amu or 0.63% of the original mass. While this percentage of the original mass may not seem much, it should be noted that the mass defect for the conversion of uranium-238 to lead-206 is only 0.026%, and that for splitting uranium-235 is 0.056%. Based upon these comparisons it is clear that fusion of hydrogen produces 24x the energy kg/kg than natural radioactivity and 11x that of nuclear fission.
In addition to being a plentiful source of energy, fusion is actually the most important process in the universe. Since the statement in 1847 of the Law of conservation of energy (the total amount of energy in an isolated system remains constant) scientists had wondered how the sun works. No source of energy was known in the 19th century that could explain the sun. Based upon the age of the earth it is known that the sun is 4,550,000,000 years old, and it was marveled at the continued source of energy over that span of time. By the 1920s nuclear energy was defined as being the most powerful source of energy, and British astrophysicist Arthur Eddington (Figure $2$.32) suggested that the sun's energy arises from the fusion of hydrogen into helium.
In 1929 American Henry Norris Russell (Figure $33$) studied the spectrum of the sun. Based upon his studies he calculated the composition of the sun to be 90% hydrogen, 9% helium, and 1% of all other elements up to iron (but nothing of higher atomic number). Given the composition it became clear that the only reaction possible to account for the sun's energy was the fusion of hydrogen. In 1938 Hans Bethe (Figure $2$.34) demonstrated a model for how the sun worked. The energy source is the fusion of hydrogen to give helium (Figure $2$.35), while in massive starts the presence of heavier elements such as carbon, oxygen, nitrogen, neon, silicon, and iron are the result of the fusion of helium.
Why is it that uranium, thorium and other radioactive elements undergo radioactive decay, but hydrogen does not undergo spontaneous fusion? The reason for this difference is that any change that uranium undergoes occurs inside a nucleus that is already formed, while fusion requires that two nuclei must come together. This process results in extremely large repulsive forces. So why does it happen in a star? The temperatures inside a star approach 15,000,000 C, while the high pressure results in a density of approximately 160 g/cm3 which for comparison is 8x that of gold. Under these conditions the nuclei are free to move in a sea of electrons. Nuclear fusion takes place in the core of the sun where the density is at the highest but an explosion does not result due to the extreme gravity of the sun: 333,000x that of Earth.
Research into controlled fusion, with the aim of producing fusion power for the production of electricity, has been conducted for over 50 years. It has been accompanied by extreme scientific and technological difficulties, but has resulted in progress. At present, break-even (self-sustaining) controlled fusion reactions have not been demonstrated, however, work continues since the fuel for such a fusion reaction is hydrogen (in its compounds including water) is the 3rd most abundant element on the Earth. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/02%3A_Hydrogen/2.09%3A_Nuclear_Fusion.txt |
Note
This module is based upon the Connexions course Methods of Hydrogen Storage for Use as a Fuel Case Study by Christian Cioce.
Introduction
Dihydrogen is a colorless and odorless gas at room temperature which is highly ammable, releasing a large amount of energy when combusted. As compared with combustion of the current fuels which operate automobiles, for example petrol or diesel, the energy released when hydrogen is combusted is more than three times greater. The heat of combustion for hydrogen is 141.9 kJ/mol as compared to 47.0 kJ/mol and 45.0 kJ/mol for gasoline and diesel, respectively.
Furthermore, the combustion of hydrocarbons releases the greenhouse gas carbon dioxide (CO2) into the atmosphere, and is therefore not a "clean" fuel. When hydrogen is combusted in the presence of oxygen (from air) the only product is water, (2.52). Both its clean reactivity and the large chemical energy make H2 extremely appealing for use as a fuel in automobiles.
$2H_{2(g)} + O_{2(g)} \rightarrow 2 H_2O_{(g)}$
If hydrogen has such a potential as a fuel why has it not been widely implemented? Dihydrogen is a gas at room temperature. Gases, compared to the other states of matter (liquid and solid), occupy the most volume of space, for a given number of molecules. Octane and the other hydrocarbons found in gasoline are liquids at room temperature, demanding relatively small fuel tanks. Liquids are therefore easier to store than compressed gases.
Hydrogen has a high energy content per weight (more than three times as much as gasoline), but the energy density per volume is rather low at standard temperature and pressure. Volumetric energy density can be increased by storing the gaseous hydrogen under increased pressure or storing it at extremely low temperatures as a liquid. Hydrogen can also be adsorbed into metal hydrides and highly porous materials (Table $2$.10). The current available methods of storing hydrogen include compressed hydrogen and liqueed hydrogen, however many promising methods exist, namely metal organic materials (MOMs), metal hydrides and carbon nanostructures.
Material H-atoms per cm3 (x 1022) % of weight that is H2
H2 gas, 200 bar (2850 psi) 0.99 100
H2 liquid, 20 K (-253 C) 4.2 100
H2 solid, 4.2 K (-269C) 5.3 100
MgH2 6.5 7.6
Mg2NiH4 5.9 3.6
FeTiH2 6.0 1.89
LaNi5H6 5.5 1.37
Table $2$.10: Comparison of hydrogen storage ability of metal hydrides.
Liquid hydrogen
Liquid hydrogen is made possible by cryogenically cooling it to below its boiling point, -253 C. As a liquid, the same amount of gaseous hydrogen will require much less volume, and therefore is feasible to individual automobile use. A refrigeration system is required to keep the liquid cooled, for if the system temperature rises above hydrogen's critical point (-241 C), the liquid will become a gas. There must exist a vacuum insulation between the inner and outer walls of liquid hydrogen tank system, for heat cannot travel through a vacuum. There is a tradeo, however, because the tank must be an open system to prevent overpressure. This will lead directly to heat loss, though minimal.
The relative tank size has a broad range, with small tanks having a volume of 100 L, and large spherical tanks sizing all the way up to 2000m3. Refrigerationsystemsarenotalikelyfeatureforeveryautomobile, and open systems may pose a hazard should an accident occur. Cooling hydrogen down to a liquid is a convenient method of storage, however, and its implementation most likely will be limited to large stationary tanks as well as mobile multi-axle trucks.
Compressed hydrogen
Compressing gas is the process of applying an external force which minimizes the distance between gas particles, therefore forcing the system to occupy less volume. This is attractive since many particles can exist in a reasonably sized tank. At room temperature and atmospheric pressure, 4 kg of hydrogen occupies a volume of 45 m3, which corresponds to a balloon with a diameter of 5 m. Clearly compression is required to store and transport the gas. When it comes to individual mobility however, these tanks are still far too large for the average sized automobile.
Compressed tanks are regularly lled to 200 atmospheres in most countries. Storing 4 kg of hydrogen still requires an internal volume of 225 L (about 60 gallons). This amount can be divided into 5 tanks with 45 L internal volume.
Metal hydrides for storage
Metal hydrides are coordinated complexes and/or crystal systems which reversibly bind hydrogen. The hydrogen is favorably incorporated into the complex and may be released by applying heat to the system. A major method to determine a particular complex's eectiveness is to measure the amount of hydrogen that can be released from the complex, rather than the amount it can store (Table $2$.10).
Some issues with metal hydrides are low hydrogen capacity, slow uptake and release kinetics, as well as cost. The rate at which the complex accepts the hydrogen is a factor, since the time to fuel a car should ideally be minimal. Even more importantly, at the current stage of research, the rate at which hydrogen is released from the complex is too slow for automobile requirements. This technology is still a very promising method, and further research allows for the possibility of highly binding and rapid reversal rates of hydrogen gas. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/02%3A_Hydrogen/2.10%3A_Storage_of_Hydrogen_for_Use_as_a_Fuel.txt |
The Group 1 metals have a particular name: the alkali metals. This is due to the formation of alkali (basic) solutions upon their reaction with water. Table $3$.1 lists the derivation of the names of the alkali metals.
Element Symbol Name
Lithium Li Greek lithos meaning stone
Sodium Na Latin natrium or Arabic natrun meaning soda
Potassium K From the Latin kalium, and from Arabic al-qali meaning plant ashes
Rubidium Rb Latin rubidus meaning deepest red
Caesium Cs Latin caesius meaning blueish grey
Francium Fr
Named after France
Table $3$.1: Derivation of the names of each of the alkali metal elements.
Note
Caesium is the international spelling standardized by the IUPAC, but in the United States it is more commonly spelled as cesium.
Discovery
Lithium
Petalite (Li2O.Al2O3.8SiO2) was first discovered in 1800 by José Bonifácio de Andrade e Silva (Figure $3$.1), who discovered the mineral in a Swedish mine on the island of Utö. However, it was not until 1817 that Johan August Arfwedson (Figure $3$.2) working in the laboratory of Jöns Jakob Berzelius (Figure $3$.3), discovered the presence of a new element while analyzing petalite ore. Named from the Greek lithos meaning stone reflected its discovery in a mineral, as opposed to sodium and potassium, which had been discovered in plant tissue; its name was later standardized as lithium. The element was not isolated until 1821, when William Brande (Figure $3$.4) isolated the element by performing electrolysis on lithium oxide, a process previously employed by Sir Humphry Davy to isolate potassium and sodium.
Sodium
3.1.1.2 Sodium Elemental sodium was first isolated by Sir Humphry Davy (Figure $3$.5) in 1806 by passing an electric current through molten sodium hydroxide.
Potassium
The name kalium was taken from the word alkali, which came from Arabic al qali meaning the calcined ashes. The name potassium was made from the English word potash, meaning an alkali extracted in a pot from the ash of burnt wood or tree leaves. Potassium metal was discovered in 1807 by Sir Humphry Davy (Figure $3$.5), who derived it from caustic potash (KOH), by the use of electrolysis of the molten salt.
Rubidium
Rubidium was discovered using spectroscopy in 1861 by Robert Bunsen (Figure $3$.6) and Gustav Kirchho (Figure $3$.7) in the mineral lepidolite. The first rubidium metal was produced by Bunsen from the reaction of rubidium chloride (RbCl) with potassium.
Caesium
Like Rubidium, caesium was discovered spectroscopically by Bünsen (Figure $3$.6) and Kirchho (Figure $3$.7) in 1860 in mineral water from Drkheim, Germany. The residues of 44,000 liters of mineral water yielded several grams of a caesium salt. Its identification was based upon the bright blue lines in its spectrum and it was the first element discovered by spectral analysis. The first caesium metal was subsequently produced in 1882 by electrolysis of caesium chloride.
Francium
Originally known as eka-caesium, francium was discovered in 1939 by Marguerite Perey (Figure $3$.8) of the Curie Institute in Paris, France when she purified a sample of actinium-227, which had been reported to have a decay energy of 220 keV. However, Perey noticed decay particles with an energy level below 80 keV. Perey thought this decay activity was caused by a previously unidentified element, which exhibited chemical properties of an alkali metal. This led Perey to believe that it was element 87, caused by the alpha decay of actinium-227. Perey named the new isotope actinium-K (now referred to as francium-223) and in 1946, she proposed the name catium for her newly discovered element, however, she subsequently suggested francium, after France.
Abundance
The abundance of the alkali metals is given in Table $3$.2. Sodium's high abundance is due mainly to large underground deposits of rock salt (NaCl). However, sodium is also more abundant in seawater (10,800 ppm) as compared to potassium (380 ppm).
Element Terrestrial abundance (ppm)
Li 20 (Earth's crust), 40 (soil), 0.17 (sea water)
Na 23,000 (Earth's crust), 10,500 (sea water)
K 21,000 (Earth's crust), 14 (soil), 380 (sea water)
Rb 90 (Earth's crust), 30 - 250 (soil), 0.1 (sea water)
Cs 3 (Earth's crust), 0.0001 (soil), 0.0003 (sea water)
Fr Essentially nil
Table $3$.2: Abundance of alkali metal elements.
Isotopes
The naturally abundant isotopes of the alkali metals are listed in Table $3$.3. All of the isotopes of francium are radioactive. Lithium-7 and sodium-23 are both useful NMR nucleus having I = 1/2.
Isotope Natural abundance (%)
Lithium-6 7.5
Lithium-7 92.5
Sodium-23 100
Potassium-39 93
Potassium-40 0.0118
Potassium-41 6.9
Caesium-133 100
Table $3$.3: Abundance of the major isotopes of the alkali metals.
Potassium has three isotopes (Table $3$.3), of which potassium-40 is radioactive and provides the basis for the determination of the age of rocks between 105 and 109 years old, i.e., those formed in proterozoic and cenozoic periods of geological time. The decay of potassium-40 occurs with a half life of 1.31 x 109 years, by two routes. That associated with a beta particle decay accounts for 89% of the decay:
${40}_{19}K \rightarrow ^{97}_{20}Ca + ^0_{-1}e$
While that associated with an electron capture and by positron emission decay accounts for 11% of the decay to give argon-40. Since many rocks contain potassium containing minerals the decay of potassium-40 after solidification of the rock will result in the formation of argon trapped in the rock. The argon-40 content is determined by mass spectrometry, while the potassium content is determined by flame spectrophotometry. The ratio of the two, (3.2), will allow for the determination of the elapsed time since the rock solidified.
$\dfrac{[^{40}_{18}Ar]}{[^{40}_{19}K]}$
Caesium has at least 39 known isotopes (more than any other element except francium) ranging from caesium-112 caesium-151; however, caesium-133 is the only naturally occurring stable isotope. The other isotopes have half-lives from a few days to fractions of a second. The radiogenic isotope caesium-137 is produced from the detonation of nuclear weapons and is produced in nuclear power plants, and was released to the atmosphere most notably from the 1986 Chernobyl accident.
Physical properties
Many of the physical properties of the alkali metals (Table $3$.4) are typical of metals, e.g., thermal and electrical conductivity. However, due to the relatively weak inter-atomic forces (weak M-M bonding) they are soft and readily cut with a knife.
Element Melting point (C) Boiling point (C) Density (g/cm3) Electrical resistivity (Ω·cm)
Li 453 1615 0.534 12.17 @ 86 C
Na 370 1156 0.968 5.23 @ 29 C
K 336 1032 0.89 7.01 @ 22.8 C
Rb 312 961 1.532 12.52 @ 53 C
Cs 201 944 1.93 37.38 @ 28.1 C
Table $3$.4: Selected physical properties of the alkali metal elements.
Reactivity
All the alkali metals are highly reactive and are as a consequence of the stability of the M+ ion are strong reducing agents (Table $3$.5). The metals react readily with hydrogen and oxygen.
$\ce{ M + 1/2 H2 \rightarrow M^{+}H^{-}}$
$\ce{ M + O2 -> M^{+}(O_2)^{-} }$
Reduction Reduction potential (V)
Li+ + e- → Li -3.045
Na+ + e- → Na -2.7109
K+ + e- → K -2.924
Rb+ + e- → Rb -2.925
Cs+ + e- → Cs -2.923
Table $3$.5: Electrochemical reduction potential for alkali metals.
All of the alkali metals react with water to liberate hydrogen.
$\ce{ M + H2O -> M(OH) + 1/2H2}$
WARNING
The reactions of alkali metals with water are exothermic and the heat generated is sucient to ignite the hydrogen. In addition the solutions formed are highly alkaline. Caution should be taken when handling alkali metals and storage should always be under mineral oil.
A similar, but less violent, reaction is also observed with ammonia when catalyzed by transition metal ions.
$\ce{M + NH3 -> M(NH2) + 1/2 H2}$
In the absence of a catalyst, the Group 1 metals dissolve in liquid ammonia to form solutions with characteristic properties.
• Highly reducing.
• Blue color.
• ESR signal due to solvated electrons.
As an example, the dissolution of sodium in liquid ammonia results in the formation of solvated Na+ cations and electrons.
$Na_{(s)} \rightarrow Na_{(solv)} \leftrightharpoons Na^+_{(solv)} + e^-_{(solv)}$
The solvated electrons are stable in liquid ammonia and form a complex: [e-(NH3)6]. It is this solvated electron that gives the strong reducing properties of the solution as well as the characteristic signal in the ESR spectrum associated with a single unpaired electron. The blue color of the solution is often ascribed to these solvated electrons; however, their absorption is in the far infra-red region of the spectrum. A second species, Na-(solv), is actually responsible for the blue color of the solution.
$2 Na_{(solv)} \leftrightharpoons Na^+_{(solv)} + Na^-_{(solv)}$
The formation of the sodium anion is confirmed by complexation of the cation with a cryptan ligand (C) such as a crown ether.
$Na^+_{(solv)} + Na^-_{(solv)} + "C" \rightarrow [Na(C)]^+ + Na^-$
The resulting complex is found to be isostructural to the iodide analog in the solid state.
$Na^+_{(solv)} + I^-_{(solv)} + "C" \rightarrow [Na(C)]^+ + I^-$
Vapor phase
All the alkali metals form M2 dimers in the vapor phase in an analogous manner to hydrogen. As with dihydrogen the bonding is associated with the molecular orbital combination of the two valence s-orbitals (Figure $3$.9).
Sodium vapor is commonly used for lighting in a gas discharge lamp, which uses sodium in an excited state to produce light (Figure $3$.10). There are two varieties of such lamps: low pressure and high pressure. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/03%3A_Group_1_-_The_Alkali_Metals/3.01%3A_The_Alkali_Metal_Elements.txt |
The chemistry of the alkali metals is dominated by the stability of the +1 oxidation state and the noble gas configuration of the M+ cation. The alkali metals all have low first ionization energies (Table $3$.6) but very high second ionization energies.
Element 1st ionization energy (kJ/mol)
Li 526
Na 502
K 425
Rb 409
Cs 382
Table $3$.6: First ionization potentials for the alkali metals.
As a consequence of the stability of M+, the Group 1 metals have the least variation in chemistry of the any Group in the periodic table. The only exceptions are the subtle trends that exist for lithium due to its small size (Table $3$.7). All of the metals are more electropositive than hydrogen (Table $3$.8).
Element Atomic radius (Å) Ionic radius (Å) Covalent radius (Å) Van der Waals radius (Å)
Li 1.52 0.68 1.52 1.82
Na 1.86 0.97 1.53 2.27
K 2.31 1.33 1.90 2.75
Rb 2.44 1.47 2.47 -
Cs 2.62 1.67 2.65 -
Fr - 1.80 2.70 -
Table $3$.7: Radii of alkali metals. N.B. Some values are unknown.
Element Electronegativity
H 2.20
Li 0.98
Na 0.93
K 0.82
Rb 0.82
Cs 0.79
Fr 0.70
Table $3$.8: Pauling electronegativities of Group 1 elements.
Solid State
In the solid state, the compounds of the alkali metals generally form ionic lattices, e.g., $\ce{Na^{+}Cl^{-}}$. These structures are essentially electrostatic in nature and the lattice energy is usually defined as the enthalpy of formation of the ionic compound from gaseous ions and as such is invariably exothermic.
$M^+_{(vap)} + X^-_{(vap)} \rightarrow MX_{(s)}$
In all cases the lattice energies is high and is found to be proportional to the ratio of the charges on the ions and the sum of the ionic radii ($r$).
$U \propto \dfrac{z^+ z^-}{\displaystyle \sum(r)}$
The ionic radii for alkali metal cations are given in Table $3$.7; those for common anions are given in Table $3$.9.
Anion Ionic radius (Å)
F- 1.33
Cl- 1.81
Br- 1.96
I- 2.20
H- 1.54
O2- 1.32
Table $3$.9: Ionic radii of common anions.
The ratio of the ionic radii (r+/r-) neatly defines the structural type observed for alkali metal salts (Table $3$.10). The unit cells for ZnS (zinc blende), NaCl (rock salt), and CsCl are shown in Figure $3$.11, Figure $3$.12, and Figure $3$.13, respectively.
r+/r- Structural type Metal coordination number
0.225 - 0.414 ZnS (zinc blende) 4
0.414 - 0.732 NaCl (rock salt) 6
0.732 - CsCl 8
Table $3$.10: Defining MX structural types for alkali metal salts.
As an example, the structure of KBr can be predicted from the data in Table $3$.7 and Table $3$.9. The ionic radius of K+ is 1.33 Å, while that for Br- is 1.96 Å. The ratio of the ionic radii (r+/r-) is 0.67. Hence, KBr has a NaCl (rock salt) structure.
Exercise $1$: Sodium Hydride
What is the structure of NaH?
Answer
The ionic radius of Na+ is 0.97 Å, while that for H- is 1.54 Å. The ratio of the ionic radii (r+/r-) is 0.67. Hence, NaH has a NaCl (rock salt) structure
Exercise $2$: Rubidium Fluoride
What is the structure of RbF?
Answer
The ionic radius of Rb+ is 1.47 Å, while that for F- is 1.33 Å. The ratio of the ionic radii (r+/r-) is 1.10. Hence, RbF has a CsCl structure.
Complexes
The coordination complexes of the alkali metal cations (M+) involve electrostatic, or ion-dipole, interactions (Figure $3$.14) that have no preferred direction of interaction. Thus, the ionic radius of the cation (Table $3$.7) controls the coordination numbers of the metal in its complexes (Table $3$.11).
Aqua ion n
[Li(H2O)n]+
4
[Na(H2O)n]+ 6
[K(H2O)n]+ 6
[Rb(H2O)n]+ 6-8
[Cs(H2O)n.]+ 8
Table $3$.11: Coordination number for alkali metal ions in aqueous complexes.
In general the alkali metal ions form complexes with hard donor such as oxygen (H2O, ROH, RCO2-, etc.) or nitrogen (e.g., NH3, NR3, etc.). The aquo complexes readily exchange the water for other ligands, (3.13); however, the equilibrium constants are small when the ligand is similar in size to water.
$[M(H_2O)_n]^+ + L \leftrightharpoons [M(H_2O)_{n-1}L]^+ + H_2O$
As a consequence of the low equilibrium constants for monodentate ligands, the alkali metal cations, M+, favor coordination to polydentate ligands such as ethylenediaminetetraacetic acid (EDTA, Figure $3$.15), polyethers, and even natural polyesters or polypeptides. In each case the polydentate ligand wraps itself around the cation.
Macrocyclic ligands
Macrocyclic ligands represent a special class of polydentate ligand. They are defined as being a cyclic compound with nine or more members including all the heteroatoms and with three or more donor atoms. and are given the special name of cryptands when they are synthetic bi- and poly-cyclic multidentate ligands. Crown ethers is the name applied to macrocyclic ligands that consist of a ring containing several ether groups. Figure $3$.16 shows several common macrocyclic ligands.
Macrocyclic ligands are generally oxygen or nitrogen donor ligands, and they form highly stable 1:1 complexes with alkali metal ions in which all or most of the coordination sites on the metal are occupied, i.e., [M(L)]+ rather than [M(L)(H2O)n]+. Since the external surface of the macrocyclic ligands comprises of organic residue (e.g., CH2 groups) the complexes are soluble in organic solvents. Thus, crown ethers are commonly used to solubilize salts (e.g., NaCl) in organic solvents. They have also been used to create solutions of nanoparticles, such as carbon nanotubes (Figure $3$.17), that are ordinarily highly insoluble in common organic solvents (Table $3$.12).
Solvent SWNT concentration (mg/L)
CH2Cl2 14.05
DMF 11.42
hexane 4.65
toluene 3.79
EtOH 2.37
MeOH 2.37
CHCl3 0.30
H2O 0.10
Table $3$.12: The concentration of Na/dibenzo-18-crown-6 solubilized reduced single walled carbon nanotubes (SWNTs) in various solvents.
The most important factor in the coordination of various macrocyclic ligands to alkali metal ions is the relative size of the ion and the ligand cavity. For example, 4,7,13,16,21,24-hexaoxa-1,10diazabicyclo[8,8,8]hexacosane is a potentially octa-dentate ligand; however, the binding eciency is very dependant on the identity of the M+ ion (Figure $3$.18). The low binding constant for lithium is probably as a consequence that the ligand would have to distort to coordinate to the small cation. Conversely, the lower equilibrium for caesium is due to its being to large to t completely into the ligand cavity.
One application of the size effect for macrocyclic ligands is the ability to selectively bind different metals. For example, the 4,7,13,16,21-pentaoxa-1,10-diazabicyclo[8,8,5]tricosane and 4,7,13,16,21,24-hexaoxa-1,10diazabicyclo[8,8,8]hexacosane ligands shown in Figure $3$.19 have very different binding constants to Na+ and K+, as a consequence of the relative size of the cation and the ligand cavity (Table $3$.13).
Cation [2,2,2] [2,2,1]
Na+ 800 250,000
K+ 250,000 700
Table $3$.13: Equilbrium constants for the M+ complexes for macrocyclic ligands with different cavity sizes. The structures of the ligands are shown in Figure $3$.19.
Bibliography
1. R. E. Anderson and A. R. Barron, J. Nanosci. Nanotechnol., 2007, 7, 3436.
2. J. M. Lehn, Supramolecular Chemistry: Concepts and Perspectives. VCH (1995). | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/03%3A_Group_1_-_The_Alkali_Metals/3.02%3A_Compounds_of_the_Alkali_Metals.txt |
While lithium shows many properties that are clearly consistent with its position in Group 1, it also has key differences to the other alkali metals. In fact, in many ways it is more similar to its diagonal neighbor magnesium (Mg) than the other Group 1 metals - a phenomenon known as the diagonal effect.
Charge/radius
The ionic radius for the +1 cation of lithium is very small in comparison with its next highest homolog, sodium (Table $3$.14). This results in a correspondingly high value for the charge density (z/r). As may be seen from Table $3$.14 the charge density for lithium is significantly higher than that of its Group 1 relations.
Element z r (Å) z/r (Å-1)
Li +1 0.68 1.47
Na +1 0.97 1.03
K +1 1.33 0.75
Mg +2 0.66 3.03
Table $3$.14: Comparison of charge densities for lithium, sodium, potassium, and magnesium.
As a result of the high charge density, the Li+ ion is a highly polarizing ion. One of the main consequences of this is that lithium tends to form polar covalent bonds rather than ionic interactions. For example, alkyl lithium compounds (RLi) contain covalent Li-C bonds in a similar manner to the Mg-C bonds in Grignards (RMgX, where X = Cl, Br)
Lattice energy
Lithium compounds have high lattice energies as compared to the other Group 1 metals (Table $3$.15). As a consequence Li2O, Li3N, and LiF are all insoluble in water, whereas their sodium compounds are highly soluble.
Compound Lattice energy (kJ/mol)
LiF -1046
NaF -923
KF -821
MgF2 -2957
Table $3$.15: Comparison of lattice energies for compounds of lithium, sodium, potassium, and magnesium.
Coordination number
The small size of lithium results in a lower coordination number (4) for compounds and complexes than observed for the other Group 1 metals. However, lithium and magnesium complexes and organometallic compounds both have most commonly four-coordinate metal centers (in the absence of large steric constraints).
Chemical reactivity
A review of some of the reactions of lithium, magnesium and the other Group 1 metals shows the anomalous behavior of lithium and its similarity to magnesium. Both lithium and magnesium reacts with carbon or nitrogen to form the corresponding carbide and nitride. Whereas sodium and the other Group 1 metals show no reaction under ambient conditions. The combustion of either lithium or magnesium in air results in the formation of the oxides, Li2O and MgO, respectively. In contrast, sodium forms the peroxide, Na2O2.
It is not only in the reactivity of the elements that this relationship between lithium and its diagonal neighbor exists. Many of the compounds of lithium have a similar reactivity to those of magnesium rather than sodium. For example, the carbonates of lithium and magnesium decompose under thermolysis to yield the oxides, (3.14) and (3.15), in contrast, sodium carbonate (Na2CO3) is stable to thermolysis.
$Li_2CO_3 \rightarrow Li_2O + CO_2$
$MgCO_3 \rightarrow MgO + CO_2$
3.04: Organolithium Compounds
One of the major uses of lithium is in the synthesis of organolithium compounds, RLi. They have great importance and utility in industry and chemical research. Their reactivity resembles that of Grignard reagents, but they are generally more reactive.
Synthesis
The best general method for RLi synthesis involves the reaction of an alkyl or aryl chloride with lithium metal in benzene or an aliphatic hydrocarbon (e.g., hexane), (3.4.1).
$RCl + 2 Li \rightarrow RLi + LiCl$
While it is possible to use diethyl ether (Et2O), the solvent slowly attack the resultant alkyl lithium compound, (3.4.2).
$Et_2O + ^nBuLi \rightarrow EtOLi + H_2C\text{=}CH_2 + ^nBuH$
Metal-hydrogen exchange, (3.4.3), metal-halogen exchange, (3.19), and metal-metal exchange can also be used, (3.4.4).
$RH + R'Li \rightarrow R'H + RLi$
$2 Li + R_2Hg \rightarrow 2 RLi + Hg$
All organolithium compounds are produced as solutions and are hence used in synthetic protocols by volume of solution. It is therefore important to know the exact concentration of RLi in solution. The simplest approach to quantify the amount of organolithium is to react a known volume with water, (3.4.5), and then titrate (with acid) the resultant base that is formed.
$RLi + H_2O \rightarrow LiOH + RH$
However, while the concentration of freshly prepared samples of organolithium reagents can the theoretically measured in this way, real samples always contain some amount of LiOH or other bases. A simple titration inevitably results in an over estimation of the organolithium reagent. To overcome this a double titration method is used.
Gillman double titration method
The careful addition of a known volume of an organolithium reagent solution (between 0.5 and 1.5 mL) to an excess of water yields a solution of LiOH that can be titrated with a standardized solution of hydrochloric acid, using phenolphthalein as the indicator. The presence of any LiOH in the original organolithium solution will be incorporated into this titration, and thus the result will be a measure of the total base content in the solution, i.e., (3.4.6).
$\text{Total base content} = \text{LiOH formed from the reaction of RLi with H}_2\text{O} + \text{LiOH present as impurity in the RLi solution}$
In order to determine the amount of LiOH present as impurity in the organolithium solution it is necessary to react the RLi without the formation of base, then titrate the resulting solution. To do this, an aliquot (the same amount as used before) of the organolithium is reacted slowly with 1,2-dibromoethane (BrCH2CH2Br) dissolved in dry diethyl ether (Et2O). After 5 min of stirring, the solution is diluted with an excess of water and then titrated with a standardized solution of hydrochloric acid, again using phenolphthalein as the indicator. The dierence of the two titrations gives the exact concentration of the organolithium.
Example $1$
An aliquot of nBuLi in hexanes (0.50 mL) was added to degassed water (20 mL). After any visible reaction had ceased, a few drops of a phenolphthalein solution in water/methanol are added resulting in a pink color indicative of a basic pH. The resulting mixture is titrated with standardized hydrochloric acid ([HCl] = 0.1034 N) until complete disappearance of the pink color (7.90 mL).
A second aliquot of nBuLi in hexanes (0.50 mL) is added to 1,2-dibromoethane (0.20 mL, Et2O). After 5 min of stirring, the mixture was diluted with water (20 mL) and after addition of the phenolphthalein indicator titrated (with vigorous stirring due to the biphasic nature of the system) with standardized hydrochloric acid ([HCl] = 0.1034 N) until complete disappearance of the pink color (0.25 mL).
The concentration of nBuLi is calculated as follows:
Step 1. $\text{[total base]} = \dfrac{\text{volume HCl } \times \text{ [HCl]}}{\text{volume } ^n \text{BuLi}} = \dfrac{7.90 \times 0.1034}{0.50} = 1.633$
Step 2. $\text{[residual base]} =\dfrac{\text{volume HC l} \times \text{ [HCl]}}{\text{volume } ^n \text{BuLi}} = \dfrac{0.25 \times 0.1034}{0.50} = 0.013$
Step 3. $\text{[}^n\text{BuLi]} = \text{[total base]} - \text{[residual base]} = 1.633 - 0.013 = 1.620M$
Properties
Alkyl lithium compounds are either low melting solids or liquids, and often with high volatility (depending on the substituent) due to the covalent nature of the bonding. They are soluble in aliphatics, aromatics, and ethers. However, while the reaction with ethers is generally slow, (3.17), alkyl lithium compounds can polymerize tetrahydrofuran (THF).
Organolithium compounds react rapidly with air and water (both vapor and liquid). The reaction with water is the basis of the Gillman double titration method for determining the concentration of organolithium reagents in solution.
Structure
The structure of organolithium compounds is dominated by their highly oligomeric nature as a result of 3-center 2-electron bridging bonds. In all cases the extent of oligomerization is dependant on the identity of the alkyl (or aryl) group. The alkyl-bridged bond is similar to those found for beryllium and aluminum compounds.
In the vapor phase any particular organolithium derivative show a range of oligomeric structures. For example, the mass spectrum of EtLi shows ions associated with both tetramers (e.g., [Et3Li4]+) and hexamers (e.g., [Et5Li6]+). The structures of the different oligomers have been predicted by molecular orbital calculations (Figure $3$.20).
Solution molecular weight measurements indicate the oligomerization is present (in the absence of a coordinating ligand such as Et2O or an amine). The extent of oligomerization depends on the steric bulk of the alkyl group (Table $3$.16). Oligomerization and solution structures have also been investigated by 7Li and 13C NMR spectroscopy.
R [RLi]n R [RLi]n
Me 4 Et 6
nBu 6 tBu 4
Table $3$.16: Extent of oligomerization (n) for organolithium compounds [RLi]n in benzene solution.
There are a large number of X-ray crystallographically determined structures for organolithium derivatives. The archetypal example is MeLi, which exists as a tetramer in the solid state (Figure $3$.21). The lithium atoms are arranged as a tetrahedron and the carbon atoms are positioned on the center of the facial planes, i.e., the carbon is equidistant from each of the lithium atoms. In contrast, EtLi has a similar tetrahedral structure, but the α-carbon of the ethyl groups are asymmetrically arranged such that it is closer to one lithium atom than the other two. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/03%3A_Group_1_-_The_Alkali_Metals/3.03%3A_The_Anomalous_Chemistry_of_Lithium.txt |
Thumbnail: A crystal of strontianite. Both strontianite, one of the most important strontium ores, and strontium are named after the town of Strontian, Scotland, the location of one of the first mines for strontium ores.
04: Group 2 - The Alkaline Earth Metals
The Group 2 metals have a particular name: the alkaline earth metals. The name is derived from the observation that they have such high melting points (Table $4$.1) that they remain solids (earths) in a fire. Table $4$.2 lists the derivation of the names of the alkali metals.
Alkali metal Melting point (C) Alkaline earth metal Melting point (C)
Li 180.54 Be 1287
Na 97.72 Mg 650
K 63038 Ca 842
Rb 39.31 Sr 777
Cs 28.44 Ba 727
Fr 27 (estimated) Ra 700
Table $4$.1: Melting points of the alkaline earth metals in comparison with the alkali metals.
Element
Symbol Name
Beryllium Be
From the Greek berullos meaning to become pale, in reference to the pale semiprecious gemstone beryl
Magnesium Mg From the Magensia district in Greece
Calcium Ca From the Latin calcis meaning lime
Strontium Sr From the mineral strontianite named after the Scottish village of Strontian
Barium Ba From the Greek bary, meaning heavy
Radium Ra From the Latin radius meaning ray
Table $4$.2: Derivation of the names of the alkaline earth metals.
Discovery
Beryllium
Beryllium was discovered by Louis-Nicolas Vauquelin (Figure $4$.1) in 1798 as a component in beryls and emeralds; however, Fredrich Wöhler (Figure $4$.2) and Antoine Bussy (Figure $4$.3) independently isolated the metal in 1828 by reacting potassium with beryllium chloride, (4.1.1).
$BeCl_2 + 2K \rightarrow Be + 2 KCl$
Magnesium
Magnesium is found in large deposits of magnesite and dolomite, and in mineral waters where the Mg2+ ion is soluble. In 1618 a farmer at Epsom in England attempted to give his cows water from a well. The farmer noticed that the water seemed to heal scratches and rashes. These Epsom salts were recognized to be hydrated magnesium sulfate, MgSO4. The metal was first isolated in 1808 by Sir Humphry Davy (Figure $4$.4) via the electrolysis of a mixture of magnesia and mercury oxide.
Antoine Bussy (Figure $4$.3) subsequently prepared magnesium by heating magnesium chloride and potassium in a glass tube, (4.1.2). When the potassium chloride was washed out, small globules of magnesium remained.
$MgCl_2 + 2K \rightarrow Mg + 2KCl$
Calcium
Calcium oxide or lime was known in ancient Rome, while even in 975 AD, Plaster of Paris (calcium sulphate) was reported to be useful for setting broken bones. The element itself was not isolated until 1808 when Sir Humphry Davy (Figure $4$.4) electrolyzed a mixture of lime and mercuric oxide (HgO). His work was based upon prior work by Jöns Berzelius (Figure $4$.5) who had prepared calcium amalgam (an alloy of calcium and mercury) by electrolyzing lime in mercury.
Strontium
Discovered in lead mines in 1787 the mineral strontianite was named after the Scottish village of Strontian. Although it was realized that this mineral was different from others that contained barium, it wasn't until 1798 that Thomas Hope (Figure $4$.6) suggested the presence of a new element. As with calcium, metallic strontium was first isolated by Sir Humphry Davy (Figure $4$.4) in 1808 using electrolysis of a mixture containing strontium chloride and mercuric oxide.
Barium
Barium minerals were known by alchemists in the early Middle Ages. Stones of the mineral barite found in Bologna, Italy (also known as Bologna stones), were known to glow after exposure to light. Carl Scheele (Figure $4$.7) identified barite in 1774, but did not isolate barium. Barium was first isolated as Ba2+ in solution by Sir Humphry Davy (Figure $4$.4) in 1808. The oxidized barium was at first called barote, by Guyton de Morveau, (Figure $4$.8) which was changed by Antoine Lavoisier (Figure $4$.9) to baryta, from which barium was derived to describe the metal.
Radium
Radium was discovered by Marie Curie (Figure $4$.10) and her husband Pierre (Figure $4$.11) in 1898 while studying pitchblende. After removing uranium they found that the remaining material was still radioactive. They then separated out a radioactive mixture consisting mostly of barium and an element that gave crimson spectral lines that had never been documented before. In 1910, radium was isolated as a pure metal by Curie and André-Louis Debierne (Figure $4$.12) through the electrolysis of a radium chloride solution by using a mercury cathode and distilling in an atmosphere of hydrogen gas.
Abundance
The abundance of the alkaline earth elements is given in Table $4$.3. Beryllium is rare, but found in the mineral beryl (Be3Al2Si6O18). While magnesium is widespread within the Earth's crust, commercial sources tend to be from sea water as well as the carbonate minerals magnesite (MgCO3) and dolomite [(Ca,Mg)CO3]. Calcium is also commonly found as the carbonate, however, strontium and barium are present as the sulfates celestine (SrSO4) and barites (BaSO4), respectively.
Element Terrestrial abundance (ppm)
Be 2.6 (Earth's crust), 6 (soil), 2 x 10-7 (sea water)
Mg 23,000 (Earth's crust), 10,000 (soil), 1,200 (sea water)
Ca 41,000 (Earth's crust), 20,000 (soil), 400 (sea water)
Sr 370 (Earth's crust), 200 (soil), 8 (sea water)
Ba 500 (Earth's crust), 500 (soil), 0.001 (sea water)
Ra 6 x 10-7 (Earth's crust), 8 x 10-7 (soil), 1 x 10-10 (sea water)
Table $4$.3: Abundance of alkaline earth elements.
Calcium is a key element for living. Not only is it present as the skeletal material for shell sh and crabs (CaCO3) its phosphate derivative, hydroxyapetite [Ca5(OH)(PO4)3] is the structural material of bones and teeth. Calcium is also present in soft tissue at a level of ca. 22g/kg. Calcium is a vital metal for the following:
• Links large molecules together.
• Used in the activation of muscles
• Used in enzyme activation by stabilization of particular conformations of proteins to be acted upon by enzymes.
Isotopes
The naturally abundant isotopes of the alkaline earth elements are listed in Table $4$.4. All of the 25 isotopes of radium are radioactive, and while radium-223, radium-224, and radium-228 are found in nature as decay products of either uranium or thorium, they are only present in trace amounts.
Isotope Natural abundance (%)
Beryllium-9 100
Magnesium-24 78.99
Magnesium-25 10
Magnesium-26 11.01
Calcium-40 96.941
Calcium-42 0.647
Calcium-43 0.135
Calcium-44 2.086
Calcium-46 0.004
Calcium-48 0.187
Strontium-84 0.56
Strontium-86 9.86
Strontium-87 7.0
Strontium-88 82.58
Barium-130 0.106
Barium-132 0.101
Barium-134 2.417
Barium-135 6.592
Barium-136 7.854
Barium-137 11.23
Barium-138 71.7
Radium-226 100
Table $4$.4: Abundance of the major isotopes of the alkaline earth elements.
Although beryllium-7 and beryllium-10 are found as trace isotopes, they are so rare beryllium is considered mononuclidic element (a chemical element which is found essentially as a single nuclide, of only one atomic mass). Calcium has four stable isotopes plus two more isotopes (calcium-46 and calcium-48) that have such long half-lives (2.8 x 1015 and 4 x 1019 years, respectively) that for all practical purposes they can be considered stable.
Measurement of the 87Sr/86Sr ratio allows for geological dating of minerals and rocks. Strontium-90 (half-life = 28.9 years) is a by-product of nuclear fission and found in nuclear fallout. For example, the 1986 Chernobyl nuclear accident contaminated released a large amount of strontium-90. Since strontium substitutes for calcium in bones it prevents excretion from the body and thus presents a significant health risk, however, strontium-89 is a short-lived artificial radioisotope that is used in the treatment of bone cancer.
Naturally occurring barium is a mix of seven stable isotopes (Table $4$.4), but there are a total twenty-two isotopes known, most of which are highly radioactive and have half-lives in the several millisecond to several day range. The only notable exception is barium-133 which has a half-life of 10.51 years.
Industrial production
The industrial production of beryllium is usually from the reaction of beryl (Be3Al2Si6O18) with Na2(SiF6) which yields the beryllium fluoride, Na2(BeF4). Subsequent reactions with base give the hydroxide.
$Na_2(BeF_4) + 2 NaOH \rightarrow Be(OH)_2 + 4 NaF$
Reaction of the hydroxide with the ammonium salt of HF2-, followed by thermolysis gives beryllium fluoride (BeF2). Finally, reduction of the fluoride with magnesium yields beryllium.
$BeF_2 + Mg \rightarrow MgF_2 + Be$
Although magnesium is an abundant metal in dozens of mineral the majority of commercial production comes from sea water, where it is present at about a level of 12% that of sodium. Calcium hydroxide is added to seawater to form magnesium hydroxide precipitate.
$MgCl_2 + Ca(OH)_2 \rightarrow Mg(OH)_2 + CaCl_2$
Subsequent reaction with hydrochloric acid yields concentrated magnesium chloride solution.
$Mg(OH)_2 + 2 HCl \rightarrow MgCl_2 + 2 H_2O$
Electrolysis of the magnesium chloride produces magnesium. At the cathode, the Mg2+ ion is reduced to magnesium metal, (4.1.7), while at the anode chlorine gas is formed, (4.1.8).
$Mg^{2+} + 2 e^- \rightarrow Mg$
$2 Cl^- \rightarrow Cl_{2(g)} + 2 e^-$
Strontium metal is produced in an analogous manner; however, it may also be prepared from strontium oxide by reduction with aluminum in vacuum at a temperature at which strontium distills off.
$3 SrO + 2 Al \rightarrow 3 Sr + Al_2O_3$
Reactivity and toxicity
The chemistry of the Group 2 elements is dominated by the +2 oxidation state and the noble gas configuration of the M2+ cation.
Calcium, strontium, and barium react with water on contact to produce the hydroxide and hydrogen gas. Although the lighter alkaline earth metals do not react violently with water, they do burn in air.
Magnesium burns with a very bright white flame such that caution should be taken not to look at the ame directly. Magnesium is capable of reducing water, (4.1.10), and as a result, water cannot be used to extinguish magnesium res; the hydrogen gas produced will only intensify the re. In addition, magnesium also reacts with carbon dioxide, (4.1.11), precluding the use of carbon dioxide re extinguishers. Class D dry chemical re extinguisher or sand are used for magnesium res.
$Mg_{(s)} + 2 H_2O \rightarrow Mg(OH)_{2(s)} + H_{2(g)}$
$2 Mg_{(s)} + CO_2 \rightarrow 2 MgO_{(s)} + C_{(s)}$
Strontium and barium burn in air to produce both the oxide and the nitride, but since the metals do not react with nitrogen except at high temperatures they only form the oxide spontaneously at room temperature.
Beryllium is a class 1 carcinogen, i.e., it is carcinogenic to both animals and humans. Beryllium is harmful if inhaled; if the concentration in air is high enough (greater than 100 µg/m3) an acute condition can result, called acute beryllium disease, which resembles pneumonia. Acute beryllium disease was reported as being associated with the manufacture of fluorescent lighting tubes (a practice that ceased in 1949).
The human body absorbs strontium as if it were calcium, and while stable isotopes have pose no signicant health threat, the uptake of radioactive strontium-90 can lead to various bone disorders and diseases, including bone cancer. All water or acid soluble barium compounds are extremely poisonous. At low doses, barium acts as a muscle stimulant, while higher doses affect the nervous system. Radium is highly radioactive and its decay product, radon gas, is also radioactive. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/04%3A_Group_2_-_The_Alkaline_Earth_Metals/4.01%3A_The_Alkaline_Earth_Elements.txt |
Calcium represents the typical Group 2 metal. The chemical properties of its compounds can be applied to all the heavier homologs.
Solid state
In the solid state the compounds of the alkaline metals generally form ionic lattices. In fact (except for beryllium and to a lesser extend magnesium) the lattice parameters calculated from the ionic radii of Group 2 metals are within 1% of the experimentally determined values, indicating the highly ionic character.
Due to the increased ionic charge the cations of the alkaline earths are less than that of their alkali metal neighbors. Thus, the ionic radius of Ca2+ (0.99 Å) is less than that of K+ (1.33 Å), but more similar to Na+ (0.97 Å). This diagonal relationship is seen for the other metals in Group 2.
Oxides
Combustion of the Group 2 metals gives the monoxide, MO. In the case of SrO and BaO further reaction occurs by the absorption of oxygen under pressure to give the peroxides, MO2. The peroxide and superoxides are not stable for the lighter homologs because the smaller M2+ ions are more polarizing and cause the peroxide and superoxides to decompose to the monoxide. Calcium peroxide can be prepared by the reaction of the hydroxide with hydrogen peroxide.
$Ca(OH)_2 + H_2O_2 + 6 H_2O \rightarrow CaO_2 \cdot 8 H_2O$
As typical for the Group 2 oxides, calcium monoxide is basic and reacts with water to give the hydroxide, (4.2.2). In fact, thin films of the Group 2 oxides of calcium, barium and strontium will readily absorb water and form the hydroxides.
$CaO + H_2O \rightarrow Ca(OH)_2$
Carbonate
Calcium carbonate (CaCO3) is a common substance found in rock in all parts of the world, and is the main component of shells of marine organisms, snails, pearls, and eggshells. Calcium carbonate is usually the principal cause of hard water. Most calcium carbonate used is extracted by quarrying or mining. However, pure calcium carbonate (e.g., for pharmaceutical use) can be produced from a pure quarried source (usually marble) or manufactured by the sequential reaction involving the thermal decomposition of the carbonate to the monoxide, (4.2.3), followed by the reaction with water to give the hydroxide, (4.2.2), and finally, the reaction with carbon dioxide to reform the carbonate, (4.2.4).
$CaCO_3 \rightarrow CaO + CO_2$
$Ca(OH)_2 + CO_2 \rightarrow CaCO_3 + H_2O$
Calcium carbonate crystallizes as a variety of mineral forms.
• Aragonite
• Calcite (Figure $4$.13)
• Vaterite
• Chalk
• Travertine
• Limestone
• Marble
Calcium carbonate is one of the most widely used mineral materials, the following represents a list of some of the main applications.
• Construction industry as a building material (marble) or as an ingredient of cement.
• Purification of iron from iron ore in a blast furnace.
• A drilling fluids in the oil industry.
• Filler material for latex gloves.
• Filler (extender) in paints.
• Filler in plastics.
• Babies' diapers.
• DIY adhesives, sealants, and decorating fillers.
• Whiting in ceramics/glazing application.
• The filler in glossy paper.
• The production of blackboard chalk (CaSO4).
• An abrasive in household cleaning products.
• Dietary calcium supplement.
• Inert filler for tablets and pharmaceuticals.
• Toothpaste.
Halides
Calcium chloride, bromide, and iodide are all ionic, water-soluble salts. In contrast, due to its high lattice energy for the small fluoride ion, CaF2 is only slightly soluble (Table $4$.5). The fluorite structure typified for CaF2 (Figure $4$.14) is found for most of the MX2 ionic solids.
Compound Solubility @ 20 °C (g/100 mL) Solubility @ 100 °C (g/100 ml)
CaF2 0.0016 0.0017
CaCl2 74.5 159
CaBr2 142 312
CaI2 209 426
Table $4$.5: Solubility of calcium halides in water.
Complexes
The coordination complexes of the alkaline earth metal cations (M2+) involve electrostatic, or ion-dipole, interactions that have no preferred direction of interaction. Calcium will tolerate a wide range of coordination numbers, however, 6 and 8 are the most common. Complexes generally form with oxygen donors in a wide range of ligands, including: ROH, R2O, R2C=O, RCO2-, and PO42-. While complexes with nitrogen donor ligands are known they are usually present with oxygen donors as well. As with the Group 1 metals, the aquo complex readily exchanges the water for other ligands, however, since the bond energy is larger for Ca2+ than for Na+, the equilibrium constants are larger for analogous complexes.
Bibliography
• J. M. Buriak, L. K. Cheatham, R. G. Gordon, J. J. Graham, and A. R. Barron, Euro J. Solid State Inorg. Chem., 1992, 29, 43.
• R. E. Anderson and A. R. Barron, Main Group Chem., 2005, 4, 279.
• C, Lupu, R. S. Arvidson, A. Lüttge and A. R. Barron, Chem. Commun., 2005, 2354.
Calcium Carbide: From Gaslight to Fertilizer
Calcium carbide has an interesting role in the societal and commercial changes that took place in the late 19th and early 20th centuries. However, in order to understand the effects of calcium carbide it is important to realize the state of the art of lighting in the late 18th century.
It was in 1792 that William Murdoch (Figure $4$.15) first began experimenting with the use of gas, derived from the heating of coal and other materials, for lighting. Murdoch produced this coal gas, or manufactured gas and conveyed it through metal pipes, lighting his cottage and offices in Redruth, Cornwall (Figure $4$.16).
In 1802 as part of the public celebrations of the Peace of Amiens (between England and France) Murdoch made a public exhibition of his lighting by illuminating the exterior of the Soho Foundry in Birmingham, England. Then in 1807 an entrepreneur, Fredrick Winsor (originally Friedrich Albrecht Winzer) displayed gaslights along the top of the wall between Carlton House and the Mall in London. This demonstration for city use was a revelation. By 1823 Britain had 300 miles of gas pipe and by 1850 it was 2000 miles. Gaslight had a profound impact on society. Walking the streets at night was safer, and it allowed for longer working hours. It also made evening activities easier. As a consequence reading and evening schools became popular pastimes. Unfortunately, gaslight was rather dull orange in color (Figure $4$.17), but it was due to another area of chemical research that a brighter alternative was discovered.
In 1895 the Frenchman Henry Moissan (Figure $4$.18) was trying to make diamonds by the reaction of carbon (graphite) with almost anything he could lay his hands on. Although highly unsuccessful, one of his experiments did prove useful. By reacting carbon with lime, the common name for calcium oxide (CaO) at 2000 C (in an electric arc furnace that he had helped develop) he produced calcium carbide (CaC2).
Pure calcium carbide is colorless, but most samples have a color ranging from black to grayish-white, depending on the grade. As an ionic salt it has a high melting point (2160 C). While the structure of calcium carbide (Figure $4$.19) has a tetragonal lattice, it is related to that of a cubic rock salt structure, but where the anion is the linear C22- moiety.
Although now used extensively, at the time of its discovery calcium carbide itself did not prove very interesting, its reaction with water had a profound effect on illumination. The reaction of calcium carbide with water yields acetylene, (4.2.5).
$\text{Ca}_2\text{C} + \text{H}_2\text{O} \rightarrow \text{Ca(OH)}_2 + \text{HC=CH}$
Unlike coal gas, acetylene burns with a very bright white ame. Although electricity was starting to become more commonly used it was very expensive and acetylene offered a cheaper alternative for domestic lighting. Thus, by 1899 there were over 250,000 acetylene gas jets in Germany alone. The reaction of calcium carbide to form acetylene was used in a variety of portable lamps. These so-called carbide lamps were used in slate, copper and tin mines, and were also used extensively as headlights in early automobiles and bikes (Figure $4$.20).
Unfortunately for acetylene there was one discovery and one economic change that brought the use of acetylene as a light source to an end. In 1893 Auer von Welsbach (Figure $4$.21) invented the gas mantel (Figure $4$.22). By the impregnation of silk or cotton with a mixture of thorium dioxide and cerium(IV) oxide (99:1), and using this in combination with gas he was able to produce a very white ame.
The other impact on acetylene, was that by 1905 the cost of electricity was significantly lower, and as a consequence the price of CaC2 dropped to 30%. There were stockpiles of calcium carbide all over Europe and America. This may have been the end of calcium carbide's usefulness, however, in 1895 Heinrich Caro (Figure $4$.23) and Adolf Frank (Figure $4$.24), at the German chemical giant Badische Anilin- und Soda-Fabrik (BASF), were trying to make hydrogen cyanide (HCN) to use in its color dye business. In 1898, one of their colleagues demonstrated that what was actually produced during the reaction at temperatures exceeding 1000 °C was not cyanide, as they had hoped. It turned out that what Caro and Frank had found was that that when calcium carbide is reacted with nitrogen at 1000 ◦C it forms calcium cyanamide (CaCN2), (4.2.6). The cyanamide anion has the structure [N=C=N]2-.
$CaC_2 + N_2 \rightarrow CaCN_2 + C$
In contact with water calcium cyanamide decomposes and liberates ammonia:
$CaCN_2 + 3 H_2O \rightarrow 2 NH_3 + CaCO_3$
As such, CaCN2 is an excellent solid fertilizer that is readily plowed into the soil. By 1908 calcium cyanamide was also found to be a plant protection agent, which, at a time when all weed control was performed mechanically, represented a great step forward. Consequently, the output of calcium cyanamide grew enormously. In 1910, 30,000 tons were produced, but in 1928, global production had reached 1.2 million tons. After a temporary decline, demand has again risen in recent years owing to the ban on several pesticides due to the environmental damage they cause. Even after 100 years of use, no harmful long-term effects to the earth or environment have been observed, nor have weeds or pests developed a resistance to calcium cyanamide.
Bibliography
• N.-G. Vannerberg, Acta Chem. Scand., 1962, 16, 1212.
• M. A. Bredig, J. Am. Chem. Soc., 1942, 64, 1730.
• Y. Yamamoto, K. Kinoshita, K. Tamaru, and T. Yamanaka, Bull. Chem. Soc. Japan, 1958, 31, 501.
Portland Cement
Chemical Composition of Portland Cement
There are four chief minerals present in a Portland cement grain: tricalcium silicate (Ca3SiO5), dicalcium silicate (Ca2SiO4), tricalcium aluminate (Ca3Al2O5) and calcium aluminoferrite (Ca4AlnFe2-nO7). The formula of each of these minerals can be broken down into the basic calcium, silicon, aluminum and iron oxides (Table $4$.6). Cement chemists use abbreviated nomenclature based on oxides of various elements to indicate chemical formulae of relevant species, i.e., C = CaO, S = SiO2, A = Al2O3, F = Fe2O3. Hence, traditional cement nomenclature abbreviates each oxide as shown in Table $4$.6.
Mineral Chemical formula Oxide composition Abbreviation
Tricalcium silicate (alite) Ca3SiO5 3CaO.SiO2 C3S
Dicalcium silicate (belite) Ca2SiO4 2CaO.SiO2 C2S
Tricalcium aluminate Ca3Al2O4 3CaO.Al2O3 C3A
Tetracalcium aluminoferrite Ca4AlnFe2-nO7 4CaO.AlnFe2-nO3 C4AF
Table $4$.6: Chemical formulae and cement nomenclature for major constituents of Portland cement. Abbreviation notation: C = CaO, S = SiO2, A = Al2O3, F = Fe2O3.
The composition of cement is varied depending on the application. A typical example of cement contains 5070% C3S, 1530% C2S, 510% C3A, 515% C4AF, and 38% other additives or minerals (such as oxides of calcium and magnesium). It is the hydration of the calcium silicate, aluminate, and aluminoferrite minerals that causes the hardening, or setting, of cement. The ratio of C3S to C2S helps to determine how fast the cement will set, with faster setting occurring with higher C3S contents. Lower C3A content promotes resistance to sulfates. Higher amounts of ferrite lead to slower hydration. The ferrite phase causes the brownish gray color in cements, so that white cements (i.e., those that are low in C4AF) are often used for aesthetic purposes.
The calcium aluminoferrite (C4AF) forms a continuous phase around the other mineral crystallites, as the iron containing species act as a fluxing agent in the rotary kiln during cement production and are the last to solidify around the others. Figure $4$.25 shows a typical cement grain.
It is worth noting that a given cement grain will not have the same size or even necessarily contain all the same minerals as the next grain. The heterogeneity exists not only within a given particle, but extends from grain to grain, batch-to-batch, plant to plant.
Bibliography
• H. F. W. Taylor, Cement Chemistry, 2nd Ed., Academic Press, London (1997).
Manufacture of Portland Cement
Portland Cement is manufactured by heating calcium carbonate and clay or shale in a kiln. During this process the calcium carbonate is converted to calcium oxide (also known as lime) and the clay minerals decompose to yield dicalcium silicate (Ca2SiO4, C2S) and other inorganic oxides such as aluminate and ferrite. Further heating melts the aluminate and ferrite phases. The lime reacts with dicalcium silicate to from tricalcium silicate (Ca3SiO5, C3S). As the mixture is cooled, tricalcium aluminate (Ca3Al2O6, C3A) and tetracalcium aluminoferrite (Ca4AlnFe2-nO7, C4AF) crystallize from the melt and tricalcium silicate and the remaining dicalcium silicate undergo phase transitions. These four minerals (C3S, C2S, C3A, and C4AF) comprise the bulk of most cement mixtures. Initially Portland cement production was carried out in a furnace, however, technological developments such as the rotary kiln have enhanced production capabilities and allowed cement to become one of the most widely used construction materials.
Cement plants generally produce various grades of cement by two processes, referred to as either the wet or dry process. The dry process uses a pneumatic kiln system which uses superheated air to convert raw materials to cement, whereas the wet process slurries the raw materials in water in preparation for conversion to cement. Cement manufacturers due to its higher energy efficiency generally favor the dry process, but the wet process tends to produce cement with properties more palatable to the energy services industry. The American Petroleum Institute (API) Class H cement used in energy service applications is produced by the wet process, and thus will be the focus of the following discussion.
The cement manufacturing process begins at the quarry (Figure $4$.26), where limestone formations are ripped and crushed in two crushers to a mean particle size of 4". The quarry formation is not entirely limestone, and no attempt is made to isolate the limestone from the other minerals. On the contrary, the rippers act to blend in the impurity minerals as evenly as is feasible while still maintaining an acceptable limestone content so as not to waste the formation. This is accomplished by ripping the formation face at a 45 (Figure $4$.27). The rock is quality controlled via mobile X-ray fluorescence (XRF) spectroscopy (Figure $4$.28) at the starting point of a mobile covered conveyor belt system (Figure $4$.29), which transports the material to a dome storage unit (Figure $4$.30).
The dome storage unit has a capacity of 60 kilotons, and is filled by dispensing the rock from the conveyor at the top of the dome into a pile built in a circular pattern (Figure $4$.31). The rock is reclaimed from storage via a raking device (Figure $4$.32) that grates over the pile at the natural angle of material slide. The raked material slides to the base of the raking unit, where a second conveyor system transfers material to either of two limestone buffer bins (Figure $4$.33), each of which is dedicated to a particular kiln process. There is an additional buffer bin for mill scale from a nearby steel plant, as well as a buffer bin for sand. It is worth noting at this point that the mill scale from the steel plant contains significant levels of boron, which acts as an innate retarder and seems to affect adversely, though not overly severely, the early compressive strength development when compared to cement from other plants.
Material leaving the buffer bins is monitored for elemental composition via XRF and feed rates are adjusted for maintaining proper ow of calcium, silicon, aluminum, and iron. The raw materials are carried to ball mills (Figure $4$.34) for grinding to ne powder, which is then mixed with water. The resulting slurry is then sent to the rotary kiln for burning (Figure $4$.35), or transformation into cement clinker.
The kilns are red to an internal material temperature of 2700 F (Figure $4$.36) with a fuel of nely ground coal, natural gas, and/or various waste materials. Around fifty percent of the energy expenditure in the wet kiln process is dedicated to evaporating the water from the slurry, in contrast to the dry process, which spends most of its energy on the calcining process. Since the dry process only requires approximately half the energy of the wet process, it is generally more attractive to cement manufacturers. Unfortunately, the dry process in current use produces poor API Class H cement. A fuller understanding of the differences in cement synthesis via the two processes could lead to the development of a more effective dry synthesis of Class H cements, but that is beyond the scope of this work.
After the clinker leaves the kiln, it enters a cooler that uses pressurized air to cool the clinker. The energy absorbed by the air in the cooler serves to pre-heat the air for feed into the kiln. The cooled clinker is then taken to storage to await nal grinding with approximately ve percent gypsum by weight. After grinding to the specied fineness, the nal cement powder is pneumatically transferred to storage silos until it is shipped to the customer.
Quality control of the clinker and nal powder is handled via an automated X-ray diffraction/X-ray fluorescence (XRD/XRF) system, simple wet chemical analyses, simple optical microscopy, and periodic performance tests, including compressive strength and thickening time. This entire process results in the heterogeneous nanocomposite of calcium silicate and aluminate particles, among other materials, which make up a typical cement grain.
Hydration of Portland Cement
The addition of water to dry cement powder results in a thin cement slurry that can be easily manipulated and cast into different shapes. In time, the slurry sets and develops strength through a series of hydration reactions. Hydration of cement is not linear through time, it proceeds very slowly at first, allowing the thin mixture to be properly placed before hardening. The chemical reactions that cause the delay in hardening are not completely understood; however, they are critical to developing a rational methodology for the control of cement setting.
Tri- and di-calcium silicates
The tri- and di-calcium silicates (C3S and C2S, respectively) comprise over 80% by weight of most cement. It is known that C3S is the most important phase in cement for strength development during the first month, while C2S reacts much more slowly, and contributes to the long-term strength of the cement. Both the silicate phases react with water as shown below to form calcium hydroxide and a rigid calcium-silicate hydrate gel, CSH, (4.2.7) and (4.2.8).
$2 (CaO)_3(SiO_2) + 7 H_2O \rightarrow (CaO)_3(SiO_2)_24(H_2O) + 3 CA(OH)_2$
$2 (CaO)_2(SiO_2) + 5 H_2O \rightarrow (CaO)_3(SiO_2)_24(H_2O) + Ca(OH)_2$
The detailed structure of CSH is not completely known, however it is generally agreed upon that it consists of condensed silicate tetrahedrally sharing oxygen atoms with a central, calcium hydroxide-like CaO2 layer. Calcium hydroxide consists of hexagonal layers of octahedrally coordinated calcium atoms and tetrahedrally coordinated oxygen atoms. Taylor has proposed that the structure is most similar to either Tobermorite or Jennite, both of which share a skeletal silicate chain Figure $4$.37.
Although the precise mechanism of C3S hydration is unclear, the kinetics of hydration is well known. The hydration of the calcium silicates proceeds via four distinct phases as shown in Figure $4$.38. The first 15-20 minutes, termed the pre-induction period (Figure $4$.38a), is marked by rapid heat evolution. During this period calcium and hydroxyl ions are released into the solution. The next, and perhaps most important, phase is the induction period (Figure $4$.38b), which is characterized by very slow reactivity. During this phase, calcium oxide continues to dissolve producing a pH near 12.5. The chemical reactions that cause the induction period are not precisely known; however, it is clear that some form of an activation barrier must be overcome before hydration can continue. It has been suggested that in pure C3S, the induction period may be the length of time it takes for CSH to begin nucleation, which may be linked to the amount of time required for calcium ions to become supersaturated in solution. Alternatively, the induction period may be caused by the development of a small amount of an impermeable calcium-silicon-hydrate (CSH) gel at the surface of the particles, which slows down the migration of water to the inorganic oxides. The initial Ca/Si ratio at the surface of the particles is near 3. As calcium ions dissolve out of this CSH gel, the Ca/Si ratio in the gel becomes 0.8-1.5. This change in Ca/Si ratio corresponds to a change in gel permeability, and may indicate an entirely new mechanism for CSH formation. As the initial CSH gel is transformed into the more permeable layer, hydration continues and the induction period gives way to the third phase of hydration, the acceleratory period (Figure $4$.38c).
After ca. 3 hours of hydration, the rate of CSH formation increases with the amount of CSH formed. Solidification of the paste, called setting, occurs near the end of the third period. The fourth stage (Figure $4$.38d) is the deceleratory period in which hydration slowly continues hardening the solid cement until the reaction is complete. The rate of hydration in this phase is determined either by the slow migration of water through CSH to the inner, unhydrated regions of the particles, or by the migration of H+ through the CSH to the anhydrous CaO and SiO2, and the migration of Ca2+ and Si4+ to the OH- ions left in solution.
Calcium aluminate and ferrite
In spite of the fact that the aluminate and ferrite phases comprise less than 20% of the bulk of cement, their reactions are very important in cement and dramatically affect the hydration of the calcium silicate phases, see below. Relative to C3S, the hydration of C3A is very fast. In the absence of any additives, C3A reacts with water to form two intermediate hexagonal phases, C2AH8 and C4AH13, (4.2.9). The structure of C2AH8 is not precisely known, but C4AH13 has a layered structure based on the calcium hydroxide structure, in which one out of every three Ca2+ is replaced by either an Al3+ or Fe3+ with an OH- anion in the interlayer space to balance the charge. All of the aluminum in C4AH13 is octahedral. C2AH8 and C4AH13 are meta-stable phases that spontaneously transform into the fully hydrated, thermodynamically stable cubic phase, C3AH6, (4.2.10). In C3A, aluminum coordination is tetrahedral. The structure consists of rings of aluminum tetrahedrally linked through bridging oxygen atoms, which slightly distorts the aluminum environment. In C3AH6, aluminum exists as highly symmetrical, octahedral Al(OH)6 units.
$2 (CaO)_3(Al_2O_3) + 21 H_2O \rightarrow (CaO)_4(Al_2O_3) \cdot 13(H_2O) + (CaO)_2(Al_2O_3) \cdot 8(H_2O)$
$(CaO)_4(Al_2O_3)\cdot 13(H_2O) + (CaO)_2(Al-2O_3) \cdot 8(H_2O) \rightarrow 2 (CaO)_3(Al_2O_3) \cdot 6(H_2O) + 9 H_2O$
If the very rapid and exothermic hydration of C3A is allowed to proceed unhindered in cement, then the setting occurs too quickly and the cement does not develop strength. Therefore, gypsum [calcium sulfate dihydrate, CaSO4·2(H2O)] is added to slow down the C3A hydration. In the presence of gypsum, tricalcium aluminate forms ettringite, [Ca3Al(OH)6.12(H2O)]2.(SO4)3.2(H2O), (4.2.11), which can also be written as C3A.3(CaSO4).32(H2O). Ettringite grows as columns of calcium, aluminum and oxygen surrounded by water and sulfate ions, as shown in Figure $4$.39.
$(CaO)_4(Al_2O_3) + 3 CaSO_4 \cdot 2(H_2O) + 26 H_2O \rightarrow (CaO)_3(Al_2O_3)(CaSO_4)_3 \cdot 32(H_2O)$
Tetracalcium aluminoferrite (C4AF) reacts much like C3A, i.e., forming ettringite in the presence of gypsum. However, hydration the ferrite phase is much slower than hydration of C3A, and water is observed to bead up on the surface of C4AF particles. This may be due to the fact that iron is not as free to migrate in the pastes as aluminum, which may cause the formation of a less permeable iron rich layer at the surface of the C4AF particles and isolated regions of iron hydroxide. In cement, if there is insufficient gypsum to convert all of the C4AF to ettringite, then an iron-rich gel forms at the surface of the silicate particles which is proposed to slow down their hydration.
Portland Cement
The hydration of cement is obviously far more complex than the sum of the hydration reactions of the individual minerals. The typical depiction of a cement grain involves larger silicate particles surrounded by the much smaller C3A and C4AF particles. The setting (hydration) of cement can be broken down into several distinct periods. The more reactive aluminate and ferrite phases react first, and these reactions dramatically affect the hydration of the silicate phase. Scrivener and Pratt used TEM to develop the widely accepted model depicted in Figure $4$.40.
In the first few minutes of hydration (Figure $4$.40b), the aluminum and iron phases react with gypsum to form an amorphous gel at the surface of the cement grains and short rods of ettringite grow. After this initial period of reactivity, cement hydration slows down and the induction period begins. After about 3 hours of hydration, the induction period ends and the acceleratory period begins. During the period from 3 to 24 hours, about 30% of cement reacts to form calcium hydroxide and CSH. The development of CSH in this period occurs in 2 phases. After ca. 10 hours hydration (Figure $4$.40c), C3S has produced outer CSH, which grows out from the ettringite rods rather than directly out from the surface of the C3S particles. Therefore, in the initial phase of the reaction, the silicate ions must migrate through the aluminum and iron rich phase to form the CSH. In the latter part of the acceleratory period, after 18 hours of hydration, C3A continues to react with gypsum, forming longer ettringite rods (Figure $4$.40d). This network of ettringite and CSH appears to form a hydrating shell about 1 µm from the surface of anhydrous C3S. A small amount of inner CSH forms inside this shell. After 13 days of hydration, reactions slow down and the deceleratory period begins (Figure $4$.40e). C3A reacts with ettringite to form some monosulfate. Inner CSH continues to grow near the C3S surface, narrowing the 1 µm gap between the hydrating shell and anhydrous C3S. The rate of hydration is likely to depend on the diusion rate of water or ions to the anhydrous surface. After 2 weeks hydration (Figure $4$.40f), the gap between the hydrating shell and the grain is completely filled with CSH. The original, outer CSH becomes more fibrous.
Bibliography
• H. F. W. Taylor, Cement Chemistry, 2nd Ed., Academic Press, London (1997).
• H. F. W. Taylor, J. Am. Ceram. Soc., 1986, 69, 464.
• V. S. Ramanchandran, R.F. Feldman, and J. J. Beaudoin, Concrete Science, Heyden and Son Ltd., Philadelphia, PA, 1981.
• H. N. Stein and J. Stevels, J. App. Chem., 1964, 14, 338.
• M. Grutzeck, S. Kwan, J. Thompson, and A. Benesi, J. Mater. Sci. Lett., 1999, 18, 217.
• V. S. Ramachandran, Concrete Admixtures Handbook, 2nd Edition, Noyes Publications, New Jersey (1995).
• K. L. Scrivener and P. L. Pratt, Mater. Res. Soc. Symp. Proc., 1984, 31, 351.
Hydration Inhibition of Portland Cement
In the oil industry, Portland cement supports boreholes of ever increasing depth. This application requires a high degree of control over the setting kinetics to allow the cement to be pumped down in a liquid form. A number of chemical inhibitors are employed to delay the setting time. The ideal inhibitor for oil well cementing would predictably delay the setting of cement, and then suddenly allow hydration to continue at a rapid rate.
A wide range of compounds show set inhibition of the hydration of Portland cement. Some common examples include, sucrose, tartaric acid, gluconic acid δ-lactone, lignosulfonate, and organic phosphonic acids, in particular nitrilo-tris(methylene)phosphonic acid (H6ntmp). The structures of these retarders are shown in.
In spite of the fact that the science of cement hydration inhibition has been investigated for over 40 years, the mechanistic details are still the subject of much speculation. There are ve primary models for cement hydration inhibition: calcium complexation, nucleation poisoning, surface adsorption, protective coating/osmotic bursting, and dissolution-precipitation. A summary of the characteristic behavior of selected retarders is shown in Table $4$.7.
Retarder
Characteristic behavior
sucrose Ca binding, acts directly on silicates, accelerates ettringite formation
tartaric acid acts via calcium complexation and calcium tartrate coating, inhibits ettringite formation
lignosulfonate accelerates ettringite formation, calcium becomes incorporated into the polymer matrix during hydration, forms a diffusion barrier
nitrilo-tris(methylene)phosphonic acid (H6ntmp)
promotes Ca dissolution, forms [Ca(H6ntmp)], heterogeneous nucleation on aluminates creates a protective coating around the grain
Table $4$.7: Summary of the behavior of various hydration retarders.
Calcium complexation
Inhibition by calcium complexation relies largely on the requirement that small calcium oxide/hydroxide templates must form in the pore water of cement pastes before silicate tetrahedra can condense into dimeric and oligomeric silicates to form CSH. Calcium complexation involves either removing calcium from solution by forming insoluble salts, or chelating calcium in solution. Calcium complexation lowers the amount of calcium effectively in solution, delaying the time to Ca(OH)2 super-saturation and preventing precipitation of the necessary templates. Simple calcium complexation should dramatically increase the amount of Si(OH)4 tetrahedra in solution, and indeed this is observed with most retarders. However, if the retarder were acting solely by calcium complexation, then one molecule of retarder would be required per calcium ion in solution, and good inhibitors are used in much smaller quantities, on the order of 0.1-2% by weight of cement. In addition, there is no simple correlation between either calcium binding strength or calcium salt solubility and retarding ability. Yet it has been shown that in pure systems, i.e., of C3S and C2S, that the lime concentration in solutions is the most important factor in determining the precipitation of CSH. Therefore, although calcium complexation must play some role in inhibition, other mechanisms of inhibition must be at work as well. An example of a retarder that operates primarily through calcium complexation is tartaric acid, however, the formation of insoluble calcium tartrate on cement grains suggest that dissolution/precipitation occurs in addition (Figure $4$.42).
Nucleation poisoning
As with calcium complexation, nucleation poisoning must rely on the formation of small calcium oxide/hydroxide templates in the pore water of cement pastes before silicate tetrahedra can condense into dimeric and oligomeric silicates to form CSH. Inhibition by nucleation poisoning is where the retarder blocks the growth of CSH or Ca(OH)2 crystals through inhibiting agglomerates of calcium ions from forming the necessary hexagonal pattern. Nucleation inhibitors act on the surface of small clusters, therefore, less than one molecule of retarder per calcium ion is required to produce dramatic results. This type of inhibition also results in an increase in the amount of silicate ions in solution, as condensation of silicate chains onto calcium oxide templates to form the CSH is inhibited. As a retarder sucrose acts via nucleation poisoning/surface adsorption.
Surface adsorption
Surface adsorption of inhibitors directly onto the surface of either the anhydrous or (more likely) the partially hydrated mineral surfaces blocks future reactions with water. In addition, if such inhibitors are large and anionic, then they produce a negative charge at the surface of the cement grains, causing the grains to repel each other thereby reducing inter-particle interactions. Lignosulfonates are typical of retarders that act via surface adsorption.
Protective coating/osmotic bursting
The formation of a protective coating with its subsequent bursting due to the build up of osmotic pressure was originally posited to explain the existence of the induction period in C3S and cement hydration, however it may be applied to inhibition in general. In this mechanism, a semi-permeable layer at the surface of the cement grain forms and slows down the migration of water and lengthens the induction period. Osmosis will drive water through the semi-permeable membrane towards the unhydrated mineral, and eventually the ow of water creates higher pressure inside the protective coating and the layer bursts. Hydration is then allowed to continue at a normal rate.
Dissolution-precipitation
A detailed study of several retarders (in particular the organic phophonates) has shown that the actually accelerate certain stages of the hydration process. This is unexpected since the phosphonates have been termed super retarders, due to their increased effect on cement hydration relative to the effect of conventional retarders. So how is it that a retarder can be so efficient at hydration inhibition at the same time as accelerating the process? The ability of phosphonates to retard cement setting is due to the lengthening the induction period, without slowing down the time it takes for setting to occur (once the acceleratory period has begun).
Phosphonates are known to complex calcium and other M2+ cations, poison the nucleation and growth of barium sulfate crystals, and inhibit the hydration of Fe2O3 and Al2O3 surfaces via direct surface adsorption, thus it was assumed that with regard to cement hydration inhibition occurred by one of these mechanism. However, the mechanism by which phosphonates inhibit cement hydration consists of two steps. First dissolution, whereby calcium is extracted from the surface of the cement grains (Figure $4$.43a) exposing the aluminum rich surface to enhanced (catalyzes) hydration and ettringite formation (Figure $4$.43b). Second precipitation, whereby the soluble calcium-phosphonate oligomerizes either in solution or on the hydrate surface to form an insoluble polymeric Ca-phosphonate (Figure $4$.43c). The Ca-phosphonate material binds to the surface of the cement grains inhibiting further hydration by acting as a diffusion barrier to water as well as a nucleation inhibitor.
Bibliography
• N. Thomas and J. Birchall, Cement and Concrete Research, 1983, 13, 830.
• D. Double, Phil. Trans. R. Soc. London, 1983, A30, 53.
• M. Bishop, S. G. Bott, and A. R. Barron, Chem. Mater., 2003, 15, 3074.
• M. Bishop and A. R. Barron, Ind. Eng. Chem. Res., 2006, 45, 7042. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/04%3A_Group_2_-_The_Alkaline_Earth_Metals/4.02%3A_Calcium_the_Archetypal_Alkaline_Earth_Metal.txt |
While the chemistry of strontium, barium (and radium) is similar to that of calcium, magnesium and beryllium show marked differences. In both cases these differences are due to the small size of the ions.
Beryllium
Beryllium can be thought of as being even more covalent than magnesium. The small size (ca. 0.3 Å) results in a very high charge density of Be2+. In addition, the ionization energy for beryllium is a large positive value (1st ionization energy = 899.5 kJ/mol, 2nd ionization energy = 14,848.7 kJ/mol). Both of these factors means that the free ion does not exist. Instead, beryllium forms covalent compounds in a similar manner to its diagonal analog aluminum. Both beryllium and aluminum form covalent compounds or strongly solvated cations, and both form polymeric hydrides, chlorides, and alkyls.
Beryllium chloride is not a lattice structure with a concomitantly high melting and boiling point as observed for the other Group 2 metals (Table4.8). Instead BeCl2 is a polymer in the solid state (Figure \(4\).44a), and an equilibrium between a monomer (Figure \(4\).44b) and dimer (Figure \(4\).44c) in the vapor phase.
M Structure
Be Polymer (4-coordinate Be)
Mg Cadmium chloride structure (6-coordinate Mg)
Ca Deformed rutile structure (6-coordinate Ca)
Ca Deformed rutile structure (6-coordinate Sr)
Ba PbCl2 structure (9-coordinate Ba) or fluorite structure (8-coordinate Ba)
Table \(4\).8: Summary of structures for alkaline earth chlorides (MCl2).
Magnesium
The ionic radius for the +2 cation of magnesium is fairly small (0.65 Å). As a consequence the charge density (z/r) is high, which results in a high polarizing power of the Mg2+ ion. Thus, magnesium tends to form polar covalent bonds rather than ionic complexes. As with lithium there is a wide range of organometallic derivatives of magnesium, especially the Grignards (RMgX, where X = Cl, Br).
A further consequence of the covalent character of the bonding is that magnesium tends to form either 4-coordinate (tetrahedral) or 6-coordinate (octahedral) complexes with well-defined geometries.
4.04: Organometallic Compounds of Magnesium
While beryllium makes a range of organometallic compounds, their hazardous nature has limited their study. In contrast, the ionic nature of calcium, strontium, and barium compounds limits the number of stable organometallic derivatives. However, the organometallic chemistry of magnesium is rich and extensive. The importance of Grignards (RMgX, where X = halide) and dialkyl magnesium compounds (R2Mg) is due to their use in organic synthesis and as synthons for a range of organometallic compounds.
Grignard reagents
Grignard reagents (and the Grignard reaction using these compounds) are named after Victor Grignard (Figure $4$.45). After studying mathematics at Lyon he transferred to chemistry, becoming a professor at the University of Nancy in 1910. During World War I, he was involved in the eld of chemical warfare; however, it is for his major contribution to organic chemistry he is remembered.
Preparation
The general the synthesis of a Grignard reagent involves the reaction of an alkyl halide (RX, where X = Cl, Br, I) with magnesium metal in a suitable ether solvent, (4.4.1).
$RX + Mg \rightarrow RMgX$
While diethyl ether (Et2O) and tetrahydrofuran (THF) are commonly used as solvents, other polar nonprotic solvents are suitable, including: triethylamine (NEt3), dimethylsulphide (Me2S), dimethylselenide (Me2Se), and dimethyltelluride (Me2Te).
In general the alkyl halide is added to an excess of magnesium suspended in the solvent. In most cases it is necessary to activate the magnesium, by the addition of iodine (I2), 1,2-dibromoethane, or sonication. If the halide is very inert reaction can be promoted by the co-condensation of magnesium and THF under vacuum.
There is often an induction period after the initial addition of alkyl halide. However, since the reaction, (4.4.1), is highly exothermic care should be taken to ensure that the reaction does not run-away. For this reason it is normal to initially add a small quantity of the alkyl halide to ensure the reaction initiates. Once reaction is initiated, the addition of alkyl halide is maintained at a suitable rate to ensure the reaction is maintained until all the alkyl halide is consumed. The excess reaction magnesium is removed from the reaction mixture by filtration.
It is not always necessary to use a liquid or solid halide dissolved in the solvent. Bubbling methyl chloride (MeCl) through an Et2O suspension of magnesium yields MeMgCl. The advantage of a gaseous alkyl halide is that the reaction is very clean as all the magnesium is consumed and the excess alkyl halide is bubbled away.
The purity of the magnesium is very important. For his original experiments Grignard used magnesium of a purity of 99.2%. However, it is now more typical to use 99.8% pure magnesium. It is important that the magnesium not be too pure since it is thought that the transition metal impurities catalyze the reaction.
The relative order of reactivity of the alkyl halide follows the trend:
$I > Br > Cl > F$
In fact alkyl uorides are suciently inert that highly coordinating polar solvents such as THF or dimethylformamide (DMF) must be used.
If the reaction is allowed to get too hot then several possible side reactions can occur. In THF reaction with the solvent occurs:
$RMgX + THF \rightarrow RH + H_2C=CH_2 + H_2C=C(H)MgX$
Alternatively, a transition metal catalyzed radical coupling between the Grignard and unreacted alkyl halide is observed irrespective of the identity of the solvent, (4.4.4).
$RMgX + RX \rightarrow R-R + MgX_2$
The mechanism for Grignard formation is thought to be radical in nature; however, a study of the surface of the magnesium during the reaction has shown the presence of corrosion pits. It is generally agreed that initiation occurs at surface dislocations, but the major reaction occurs at a polished surface.
The kinetics of the reaction is 1st order with respect to the alkyl halide concentration, but it has also been claimed to be 1st order with respect to the solvent concentration. It has therefore been concluded that the rate-determining step involves the metal solvent interface.
The reaction of magnesium with aryl bromides has been studied and is proposed to occur by two reactions. The first involves electron transfer between the aryl halide and the metal, while the second involves aryl radical formation.
A number of alternative synthetic routes are used with polyhalogenated hydrocarbons, (4.4.5) and (4.4.6), and where the alkyl radical is unstable, (4.4.7).
$X_3CH + ^iPrMgX \rightarrow (X_3C)MgX + ^iPrH$
$C_6Br_6 + EtMgX \rightarrow (C_6Br_5)MgX + EtBr$
$RX + R'MgX' \rightarrow RMgX' + R'X$
Structure
The solid state structure of Grignard reagents is controlled by the presence and identity of the solvent used in the synthesis. In this regard the size and the basicity of the solvent is important. For example, the structure of EtMgBr crystallized from diethyl ether exists as a 4-ccordinate monomer (Figure $4$.46a), while the use of the sterically less demanding THF results in a 5-coordinate monomeric structure (Figure $4$.46b). In contrast, the use of triethylamine yields a dimeric bromide bridged structure (Figure $4$.46c), and the use of a chelate bidentate amine gives a structure (Figure $4$.46d) similar to that observed with diethyl ether (Figure $4$.46a).
In solution, Grignards are fluxional such that no single defined structure is present. The series of exchange reactions are known as an extended Schlenk equilibrium (Figure $4$.47).
It is observed that Grignard solutions are also slightly conducting, and magnesium is deposited at both the anode and cathode suggesting the formation of RMg+ and [RMgX2]-. The alkyl/halide exchange is thought to occur through a bridging intermediate (Figure $4$.48).
Dialkyl magnesium (R2Mg)
Dialkyl magnesium compounds are involatile white solids. They generally have similar reactivity to their Grignard analogs.
Synthesis
The most common synthesis of R2Mg is by the reaction of a Grignard with dioxane (C4H8O2), (4.32), where the precipitation of the dihalide is the reaction driving force.
This method is useful for the synthesis of cyclic compounds, (4.33).
An alternative synthesis that does not require dioxane involves the metal exchange reaction between magnesium metal and a dialkyl mercury compound.
$R_2Hg + Mg \rightarrow R_2Mg + Hg$
Finally, in selected cases, magnesium will react with acidic hydrocarbons such as cyclopentadienyl at high temperatures (600C).
Structure
In the vapor phase dialkyl magnesium compounds are generally monomeric linear compounds. In solution, in the absence of coordinating solvents R2Mg form a variety of oligomers (Figure $4$.49a-c) in solution as determined by molecular weight measurements. In the presence of coordinating solvents 4-coordinate monomers predominate (Figure $4$.49d).
As similar trend is observed in the solid state, where polymers have been characterized in the absence of coordinating solvents (Figure $4$.50a), while monomers or dimmers are generally observed when crystallized from a coordinating solvent (Figure $4$.50b and c).
The use of organomagnesium compounds in organic synthesis
Hydrolysis and related reactions
Grignard compounds react with water to give the hydrocarbon, (4.4.9), they also react with other hydroxylic compounds such as alcohols and carboxylic acids. One important use of the hydrolysis reaction is specifically deuteration, (4.4.10).
$CH_3MgBr + H_2O \rightarrow CH_4 + BrMgOH$
$CH_3CH_2(CH_3)_2CMgBr + D_2O \rightarrow CH_3CH_2(CH_3)_2CD + BrMgOD$
The hydrogen atom on a terminal alkyne is suciently acidic that the reaction with Grignards occurs in an analogous manner to that of hydrolysis.
$C_6H_5C=CH + C_2H_5MgBr \rightarrow C_6H_5C=CMgBr + C_2H_6$
Once formed the alkynyl Grignard undergoes the same hydrolysis reaction.
$C_6H_5C=CMgBr + D_2O \rightarrow C_6H_5C=CD + BrMgOD$
Reaction with CO2
Grignards react readily with carbon dioxide to form the carboxylate, which yields the associated carboxylic acid upon hydrolysis, (4.4.13).
$RMgX + CO_2 \rightarrow RCO_2MgX \xrightarrow{H_2O} RCO_2H + HOMgX$
Reaction with carbonyls
Organomagnesium compounds react with organic carbonyls (aldehydes, ketones, and esters) to yield the alcohol on hydrolysis, (4.4.14). This synthetic route is useful for the formation of primary, secondary and terminal alcohols.
$RMgX + R'_2C=O \rightarrow R'_2(R)COMgX \xrightarrow{H_2O} R'_2(R)COH + HOMgX$
Unfortunately, for some carbonyls there is a competing side reaction of enolization, where the starting ketone is reformed upon hydrolysis.
When the Grignard reagent has a β-hydrogen another side reaction occurs in which the carbonyl group is reduced and an alkene is formed.
$R_2C=O + (CH_3CH_2)MgX \rightarrow R_2(H)COH + H_2C=CH_2$
Both the enolization and reduction occur via similar 6-membered cyclic transition states (Figure $4$.51).
Grignards react with α,β-unsaturated ketones to give either the 1,2-addition product or the 1,4-addition product, or both.
$\text{Ph(H)C=C(H)-C(O)Me + EtMgBr} \xrightarrow{H_2O} \text{Ph(H)C=C(H)-C(OH)EtMe} + \text{Ph(H)EtC-CH}_2\text{-C(O)Me}$
Reaction with acyl halides
Acyl halides react with Grignards to give ketones (4.4.16). Best results are obtained if the reaction is carried out at low temperature and in the presence of a Lewis acid catalysts (e.g., (FeCl3).
$\text{CH}_3\text{C(O)Cl + RMgX} \rightarrow \text{CH}_3\text{C(O)R + XMgCl}$
Reaction with epoxides
Oxirane (epoxide) rings are opened by Grignards in a useful reaction that extends the carbon chain of the Grignard by two carbon atoms. This reaction is best performed with ethylene oxide since the magnesium halide formed is a Lewis acid catalyst for further reactions in the case of substituted oxiranes.
Reaction with salts
One of the most useful methods of preparing organometallic compounds is the exchange reaction of one organometallic compound with a salt of a different metal, Equation. This is an equilibrium process, whose equilibrium constant is defined by the reduction potential of both metals. In general the reaction will proceed so that the more electropositive metal will form the more ionic salt (usually chloride).
$\text{RM + M'X} \leftrightharpoons \text{RM' + MX}$
Grignard reagents are particularly useful in this regard, and may be used to prepare a wide range of organometallic compounds. For example:
$\text{2 CH}_3\text{CH}_2\text{MgCl + CdCl}_2 \rightarrow \text{Cd(CH}_2\text{CH}_3\text{)}_2 \text{ + 2 MgCl}_2$
$\text{4 CH}_3 \text{MgCl + SiCl}_4 \rightarrow \text{Si(CH}_3\text{)}_4 \text{ + 4 MgCl}_2$
The reaction with a Grignard is milder than the analogous reaction with lithium reagents, and leads to a lower incident of side-products.
Bibliography
• H. Bader and N. M. Smyth, J. Org. Chem., 1964, 29, 953.
• C. L. Hill, J. B. Vander Sande, G. M. Whitesides, J. Org. Chem., 1980, 45, 1020.
• E. Weiss, J. Organomet. Chem., 1964, 2, 314.
• A. R. Barron, J. Chem. Soc., Dalton Trans., 1989, 1625. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/04%3A_Group_2_-_The_Alkaline_Earth_Metals/4.03%3A_Differences_for_Beryllium_and_Magnesium.txt |
Although the Group 12 metals (Table $1$) are formally part of the d-block elements from their position in the Periodic Table, their electronic configuration in both their elemental form (d10s2) and the vast majority of their compounds (d10) is that of the main group elements. The common oxidation state for all the Group 12 elements is +2, and the chemistry of zinc and cadmium compounds in particular is very similar to the analogous magnesium derivatives.
Note
The IUPAC (International Union of Pure and Applied Chemistry) definition of a transition metal (or transition element) states that a transition metal is "an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell." Thus, Group 12 elements are not transition metals.
Element Symbol Name
Zinc Zn From German zinke, meaning tooth-like, pointed or jagged (metallic zinc crystals have a needle-like appearance), or meaning tin-like because of its relation to German word zinn meaning tin, or from Persian seng meaning stone.
Cadmium Cd From the Latin cadmia, meaning calamine
Mercury Hg From the Latin hydrargyrum, meaning watery or liquid silver
Table $1$: Derivation of the names of the Group 12 metals.
Discovery
Zinc
Artifacts with a high zinc content (as much as 90%) have been fond to be over 2500 years old, and possibly older. As such it is clear that several cultures had the knowledge of working with zinc alloys, in particular brass (a zinc/copper alloy). Zinc mines at Zawar, near Udaipur in India, have been active since the late 1st millennium BC. However, the smelting of metallic zinc appears to have begun around the 12th century AD.
The isolation of purified metallic zinc was reported concurrently by several people. The extraction of zinc from its oxide (ZnO) was reported as early as 1668, while John Lane is supposed to have smelted zinc in 1726. The first Patent for zinc smelting was granted to English metallurgist William Champion in 1738; however, the credit for discovering pure metallic zinc is often given to Andreas Marggraf in 1746.
Cadmium
Cadmium was discovered in 1817 by Friedrich Stromeyer as an impurity in calamine (zinc carbonate, ZnCO3). Stromeyer observed that impure samples of calamine changed color when heated but pure calamine did not. Eventually he was able to isolate cadmium metal by roasting and reduction of the sulfide.
Mercury
Mercury was known to the ancient Chinese and was found in Egyptian tombs that date from 1500 BC. In China and Tibet, mercury use was thought to prolong life, heal fractures, and maintain generally good health. The ancient Greeks used mercury in ointments; the ancient Egyptians and the Romans used it in cosmetics that sometimes deformed the face.
Abundance
The Group 12 elements mainly occur in sulfide ores, however, as with their Group 2 analogs, carbonate are known, but not as economically viable. The major zinc containing ore is zinc blende (also known as sphalerite), which is zinc sulfide (ZnS). Other important ores include, wurtzite (ZnS), smithsonite (zinc carbonate, ZnCO3), and hemimorphite (calamine, Zn2SiO4). The basic form of zinc carbonate (hydrozincite, Zn5(CO3)2(OH)6) is also mined where economically viable. The main source of cadmium is as an impurity in zinc blende; however, there are several other ores known, e.g., cadmoselite (cadmium selenide, CdSe) and otavite (CdCO3). Mercury sulfide (cinnabar, HgS) is the major source of mercury, and in fact metallic liquid mercury droplets are often found in the ore. The terrestrial abundance of the Group 12 elements is given in Table $2$.
Element Terrestrial abundance (ppm)
Zn 75 (Earth’s crust), 64 (soil), 30 x 10-6 (sea water)
Cd 0.1 (Earth’s crust), 1 (soil), 1 x 10-6 (sea water)
Hg 50 x 10-6 (Earth’s crust), 2 x 10-8 (soil), 40 x 10-12 (sea water)
Table $2$: Abundance of Group 12 elements.
Isotopes
The naturally abundant isotopes of the Group 12 metals are listed in Table $3$.
Isotope Natural abundance (%)
Zinc-64 48.6
Zinc-66 27.9
Zinc-67 4.1
Zinc-68 18.8
Zinc-70 0.6
Cadmium-106 * 1.25
Cadmium-108 * 0.89
Cadmium-110 12.49
Cadmium-111 12.8
Cadmium-112 24.13
Cadmium-113 * 12.22
Cadmium-114 * 28.73
Cadmium-116 * 7.49
Mercury-196 0.15
Mercury-198 9.97
Mercury-199 16.87
Mercury-200 23.1
Mercury-201 13.18
Mercury-202 29.86
Mercury-204 6.87
Table $3$: Abundance of the major (non-synthetic) isotopes of the Group 12 metals. Isotopes labeled with * are radioactive.
Many radioisotopes of zinc have been characterized. Zinc-65 that has a half-life of 244 days, is the most long-lived isotope, followed by 72Zn with a half-life of 46.5 hours. The most common decay mode of an isotope of zinc with a mass number lower than 64 is electron capture, producing an isotope of copper, (Equation 5.1.1).
$^{30}_n\text{Zn} + \text{e}^- \rightarrow ^{29}_n\text{Cu}$
T he most common decay mode of an isotope of zinc with mass number higher than 64 is beta decay (β–), which produces an isotope of gallium, (Equation 5.1.2).
$^{30}_n\text{Zn} \rightarrow ^{31}_n\text{Ga} + \text{e}^- + \nu_e$
Naturally occurring cadmium is composed of 8 isotopes. For two of them, natural radioactivity was observed, and three others are predicted to be radioactive but their decay is not observed, due to extremely long half-life times. The two natural radioactive isotopes are 113Cd (half-life = 7.7 x 1015 years) and 116Cd (half-life = 2.9 x 1019 years).
There are seven stable isotopes of mercury with the longest-lived radioisotopes being 194Hg (half-life = 444 years) and 203Hg (half-life = 47 days). 199Hg and 201Hg are the most often studied NMR-active nuclei, having spins of 1/2 and 3/2 respectively.
Properties
A summary of the physical properties of the Group 12 metals is given in Table $4$. Because of the ns electron in the Group 12 metals are tightly bound, and hence relatively unavailable for metallic bonding, the metals are volatile with low boiling points, as compared to the Group 2 metals.
Element Mp (°C) Bp (°C) Density (g/cm3)
Zn 419.53 907 7.14
Cd 321.07 767 8.65
Hg -38.83 356.73 13.534 (liquid)
Table $4$: Selected physical properties of the Group 12 metals.
The most notable anomaly in the Group 12 metals is the low melting point of mercury compared to zinc and cadmium. In order to completely understand the reasons for mercury’s low melting point quantum physics is required; however, the key point is that mercury has a unique electronic configuration, i.e., [Xe] 5d 6s. The stability of the 6s shell is due to the presence of a filled 4f shell, because an f shell poorly screens the nuclear charge that increases the attractive coulomb interaction of the 6s shell and the nucleus. Such a configuration strongly resists removal of an electron and as such mercury behaves similarly to noble gas elements, which form weakly bonded and thus easily melting solids.
Industrial production
The vast majority (95%) of zinc is mined from of the zinc sulfide ores. The zinc is most often mixed with copper, lead, and iron. Zinc metal is produced by extraction, in which the ore is ground and then the minerals are separated from the gangue (commercially worthless mineral matter) by froth flotation (a process for selectively separating hydrophobic materials from hydrophilic). Roasting converts the zinc sulfide concentrate produced to zinc oxide (Equation 5.1.3). Reduction of the zinc oxide with carbon (5.1.4) or carbon monoxide (5.1.5) at 950 °C into the metal is followed by distillation of the metal. Since cadmium is a common impurity in zinc ores, it is most often isolated during the production of zinc. Cadmium is isolated from the zinc metal by vacuum distillation if the zinc is smelted, or cadmium sulfate is precipitated out of the electrolysis solution.
$\text{2 ZnS + 3 O}_2 \rightarrow \text{2 ZnO + 2 SO}_2$
$\text{2 ZnO + C} \rightarrow \text{2 Zn + CO}_2$
$\text{2 ZnO + 2 CO} \rightarrow \text{2 Zn + 2 CO}$
Mercury is extracted by heating cinnabar (HgS) in a current of air, Equation, and condensing the vapor.
$\text{HgS + O}_2 \rightarrow \text{Hg + SO}_2$ | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/05%3A_Group_12/5.01%3A_The_Group_12_Elements.txt |
The most studied non-oxide semiconductors are cadmium chalcogenides (CdE, with E = sulfide, selenide and telluride). CdE nanocrystals were probably the first material used to demonstrate quantum size effects corresponding to a change in the electronic structure with size, i.e., the increase of the band gap energy with the decrease in size of particles. These semiconductors nanocrystals are commonly synthesized by thermal decomposition of an organometallic precursor dissolved in an anhydrous solvent containing the source of chalcogenide and a stabilizing material (polymer or capping ligand). Stabilizing molecules bound to the surface of particles control their growth and prevent particle aggregation.
Although cadmium chalcogenides are the most studies semiconducting nanoparticles, the methodology for the formation of semiconducting nanoparticles was first demonstrated independently for InP and GaAs, e.g. Equation 5.1.1. This method has been adapted for a range of semiconductor nanoparticles.
$\text{InCl}_3 \text{ + P(SiMe}_3\text{)}_3 \rightarrow \text{InP + 3 Me}_3\text{SiCl}$
In the case of CdE, dimethylcadmium Cd(CH3)2 is used as a cadmium source and bis(trimethylsilyl)sulfide, (Me3Si)2S, trioctylphosphine selenide or telluride (TOPSe, TOPTe) serve as sources of selenide in trioctylphosphine oxide (TOPO) used as solvent and capping molecule. The mixture is heated at 230-260 °C over a few hours while modulating the temperature in response to changes in the size distribution as estimated from the absorption spectra of aliquots removed at regular intervals. These particles, capped with TOP/TOPO molecules, are non-aggregated and easily dispersible in organic solvents forming optically clear dispersions. When similar syntheses are performed in the presence of surfactant, strongly anisotropic nanoparticles are obtained, e.g., rod-shaped CdSe nanoparticles can be obtained.
Because Cd(CH3)2 is extremely toxic, pyrophoric and explosive at elevated temperature, other Cd sources have been used. CdO appears to be an interesting precursor. CdO powder dissolves in TOPO and HPA or TDPA (tetradecylphosphonic acid) at about 300 °C giving a colorless homogeneous solution. By introducing selenium or tellurium dissolved in TOP, nanocrystals grow to the desired size.
Nanorods of CdSe or CdTe can also be produced by using a greater initial concentration of cadmium as compared to reactions for nanoparticles. This approach has been successfully applied for synthesis of numerous other metal chalcogenides including ZnS, ZnSe, and Zn1-xCdxS. Similar procedures enable the formation of MnS, PdS, NiS, Cu2S nanoparticles, nano rods, and nano disks.
Bibliography
• C. R. Berry, Phys. Rev., 1967, 161, 848.
• M. D. Healy, P. E. Laibinis, P. D. Stupik, and A. R. Barron, J. Chem. Soc., Chem. Commun., 1989, 359.
• L. Manna, E. C. Scher, and A. P. Alivisatos, J. Am. Chem. Soc., 2000, 122, 12700.
• C. B. Murray, D. J. Norris, and M. G. Bawendi, J. Am. Chem. Soc., 1993, 115, 8706.
• Z. A. Peng and X. Peng, J. Am. Chem. Soc., 2002, 12, 3343.
• R. L. Wells, C. G. Pitt, A. T. McPhail, A. P. Purdy, S. R. B. Shafieezad, and Hallock Chem. Mater., 1989, 1, 4.
• X. Zong, Y. Feng, W. Knoll, and H. Man, J. Am. Chem. Soc., 2003, 125, 13559.
5.03: Organometallic Chemistry of Zinc
The first dialkyl zinc derivatives, Me2Zn and Et2Zn, were prepared in 1848 by Edward Franklin. He also prepared the monoalkyl derivatives, RZnX.
Initially alkyl zinc compounds were used in organic synthesis, however, their use diminished significantly once Grignard reagents had been discovered. Little further was investigated of their chemistry until their use in the growth of electronic materials developed in the 1980’s.
RZnX
The monoalkyl derivatives are not widely used, but were historically the first to be prepared (Equation 5.3.1).
$\text{RX + Zn} \rightarrow \text{RZnX}$
While the iodide derivatives can be isolated as unsolvated derivatives, the chloride and bromides need to be prepared in the presence of dimethylformamide (DMF) or dimethyl sulfoxide (DMSO). Alternative methods of synthesis involve the reaction with a Grignard reagent, (Equation 5.3.2), or electrochemical synthesis with a zinc electrode.
$\text{ZnX}_3 \text{ + RMgX} \rightarrow \text{RZnX + MgX}_2$
In the solid state RZnI exists as cage oligomers or polymeric chain structures. These structures are broken by the addition of strong Lewis bases to form Lewis acid-base complexes. In solution there exists a Schlenk equilibrium (equation 5.3.3) whose presence has been determined by IR and Raman spectroscopy (Table $1$).
$\text{RZnI} \leftrightharpoons \text{R}_2\text{Zn + ZnI}_2$
Stretch IR (cm-1) Raman (cm-1)
symmetric C-Zn-C Not observed 477
asymmetric C-Zn-C 551 Not observed
C-Zn-I 510 510
Table $1$: IR and Raman spectroscopic characterization of the components of the Schlenk equilibrium,
R2Zn
Dialkyl zinc compounds are prepared via the monoalkyl derivatives (Equation 5.3.4). The Schlenk equilibrium is shifted at high temperatures by the distillation of the volatile R2Zn derivative. The zinc is usually alloyed with copper (10%) to improve the reaction rate.
$\text{2 Zn + 2 RX} \rightarrow \text{2 RZnX} \xleftrightarrow{\Delta} \text{R}_2\text{Zn + 2 MgX}_2$
Alternative preparation methods include the reaction of ZnX2 with a Grignard (Equation 5.3.5) or by metal-metal exchange (Equation 5.3.6).
$\text{2 RMgX + ZnX}_2 \rightarrow \text{R}_2\text{Zn + 2 MgX}_2$
$\text{R}_2\text{Hg + Zn} \rightarrow \text{Hg + R}_2\text{Zn}$
Bibliography
• E. von Frankland. Justus Liebigs Ann. Chem., 1849, 71, 171.
• J. J. Habeeb, A. Osman, and D. G. Tuck. J. Organomet. Chem., 1980, 185, 117 | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/05%3A_Group_12/5.02%3A_Cadmium_Chalcogenide_Nanoparticles.txt |
Synthesis
The most common routes to organomercury compounds involve the direct reaction of mercury with an alkyl iodide (Equation 5.4.1) to form the mercury analog of a Grignard reagent.
$\text{Hg + RX} \rightarrow \text{RHgX}$
The subsequent reaction of RHgI with potassium cyanide yields the appropriate dialkyl mercury derivative (Equation 5.4.2).
$\text{2 RHgI + 2 KCN} \rightarrow \text{R}_2\text{Hg + Hg + 2 KI + (NC)}_2$
Solvomercuration
The general process of solvomercuration involves the addition of HgX2 across an alkene double (or alkyne triple) bond in the presence of a solvent. Solvomercuration applies where HY, from Equation 5.4.3, are part of the solvent system, e.g., in water the process can be described as hydroxymercuration.
The addition to the alkene occurs with Markovnikoff addition, i.e., through the formation of the most stable carbonium ion resulting with the mercury adding to the less substituted carbon. The Hg-C bond can be cleaved by the addition of NaBH4 to yield the C-H bond. It is common to employ mercury acetate, Hg(OAc)2 (OAc = O2CCH3), with subsequent reaction with NaCl or NaI to form the halide, rather than the mercury halide directly. However, the acetate can also act as a neucleophile resulting in a mixture of products. The order of reactivity for alkenes follows the trend:
$\text{R}_2\text{C=CH}_2\text{ > R(H)C=CH}_2 \text{ >} cis \text{ R(H)C=C(H)R > } trans \text{ R(H)C=C(H)R > R}_2\text{C=C(H)R > R}_2\text{C=CR}_2$
This is due to a combination of steric and electronic effects. The addition occurs in a trans/anti fashion.
Mercury(II) halides dissociate in polar solvents (Equation 5.4.4) and this species is commonly associated with the initial reaction step.
$\text{HgX}_2 \leftrightharpoons \text{HgX}^+ \text{ + X}^-$
The mechanism of solvomercuration is best described by the reaction shown in.
Evidence for the mechanism is twofold:
1. The addition is exclusively trans.
2. No rearrangement is observed even for tBu(H)C=C(H)tBu whose carbonium ion is known to undergo rearrangement.
Isotope studies indicate that the C-O bond formation is present in the transition state. The mercury bridge in the transition state may not be symmetrical, but the structure is similar to the addition of Br+ and AuX+ to alkenes.
The actual reaction was originally carried out with Hg(OAc)2 in benzene at 110 °C for several hours in acetic acid solution. It was found that the reaction was catalyzed by the presence of HClO4, H2SO4, and HNO3, which were found to replace the acetate ion. The reaction rate is also increased by 690,000 times by the use of Hg(O2CCF3)2 in HO2CCF3.
The solvomercuration of alkynes gives alkenylmercury compounds, but the reaction is more sluggish than for the reaction with alkenes, and the product is always the trans isomer.
Mercuration of aromatic compounds
Mercuration is an electrophilic aromatic substitution reaction that is possible for most 2n+2 π-electron species. Evidence for the π-complex intermediate is indicated by UV spectroscopy, which shows an increase in the region 280 – 320 nm for the reaction of aromatic compounds with Hg(O2CCF3)2 in HO2CCF3. The σ-complex has detected in liquid SO2.
As a preparative method mercuration suffers from alack of selectivity, including an isomerization from para substitution at low temperature to meta substitution at high temperatures.
Structure
Dialkyl mercury compounds, R2Hg, are generally air stable and show little Lewis acid behavior. They are monomeric colorless liquids or low melting solids, e.g., Bp = 92.5 °C for Me2Hg. No solubility is observed in water, except for (F3C)2Hg (Equation 5.4.5).
$\text{(F}_3\text{C)}_2\text{Hg + H}_2\text{O} \leftrightharpoons \text{(F}_3\text{C)}_2\text{Hg(OH)}^- \text{ + H}^+$
The hybridization at mercury involves the 6s and 6p orbitals; however, the 5d may be involved. X-ray crystal structures for both R2Hg and RHgX show linear structures. In general R2Hg compounds are very weak Lewis acids, but adduct are formed if the alkyl group is sufficiently electron withdrawing, e.g., (C6F5)2Hg. The geometry of the Lewis acid-base complex is not triangular as predicted from VSEPR theory, but a T-shape. The distortion from a linear C-Hg-C unit is minor. For example, in (C6F5)2Hg the C-Hg-C angle is 176.2°, while in [(C6F5)2Hg]2(diars), where diars = Ph2AsCH2CH2AsPh2, the same angle is 173°. Furthermore, the donor atom-Hg bond distance (e.g., As-Hg = 3.4 Å) is only slightly shorter than the sum of the van der Waal radii (e.g., 3.5 Å).
Bibliography
• L. C. Damude and P. A. W. Dean. J. Organomet. Chem., 1979, 181, 1.
• L. C. Damude and P. A. W. Dean. J. Chem. Soc., Chem. Commun., 1978, 1083.
• J. L. Courtneidge, A. G. Davies, P. S. Gregory, D. C. McGuchan, and S. N. Yazdi. J. Chem. Soc., Chem. Commun., 1987, 1192.
• P. B. Hitchcock, J. M. Keates, and G. A. Lawless, J. Am. Chem. Soc., 1998, 120, 599. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/05%3A_Group_12/5.04%3A_Organomercury_Compounds.txt |
Mercury has a reputation for being a dangerous element, but is its reputation deserved? Given the large-scale use of mercury today it is important to understand the risks and issues related to mercury exposure. Nowhere is this now important than with the use of mercury for small low energy fluorescent lights that are being advocated by everyone from the electricity companies to Greenpeace.
When considering the issue of mercury toxicity it is important to separate the effects of mercury metal (as a liquid or vapor) from the compounds of mercury.
Mercury metal
It was found very early on that people who worked with mercury, in mining for example, had very bad health. Other jobs that exposed people to mercury were mirror makers and hatters (people who manufactured hats). The problems in this latter occupation will forever live on with one of the central characters in Lewis Carroll's Alice’s Adventures in Wonderland; the Mad Hatter.
Hats were made from felt, which is a non-woven textile of animal hair. Wool interlocks naturally due to the surface texture of the individual hairs, but rabbit and beaver have to be artificially roughened. This process was usually accomplished with nitric acid (HNO3). It was found that if mercury was added to the nitric acid, a better quality of felt was produced. Unfortunately, when the felt was dried a fine dust was formed containing mercury. The hatters who shaped the felt inhaled large quantities of this dust were found to suffer from excessive salivation, erethism (presenting with excessive shyness, timidity and social phobia), and shaking of the limbs, which became known as hatter’s shakes. The madness that was observed is the derivation of the phrase “mad as a hatter”.
Note
It is interesting that while Carroll's Mad Hatter is mad, he does not show the classic symptoms of mercury poisoning. In particular he can be in no way described as shy!
Hatters were not the only people that mercury caused a problem for. Chemists doing research using large quantities of mercury were also affected. They were given to violent headaches, tremors of the hands, “socially troublesome inflammation of the bladder”, loss of memory, and slow mental processes. In 1926 Alfred Stock (Figure) and his research group all suffered from symptoms. However, when the lab was cleaned of mercury the symptoms went away.
Many other notable scientists have also suffered from mercury poisoning. Faraday (Figure), Pascal (Figure), and most probably Sir Isaac Newton (Figure) were affected. As part of his research studies, Newton boiled several pounds of mercury a day just before his period of insanity between 1692 and 1693. It is likely that the mercury vapor was the cause of his malady. However, in each case, the symptoms (and insanity) abated once the source of mercury was removed.
Note
It is important to remember that in all the cases described above it is the inhalation of the mercury vapor that was the cause of the trouble. Solid alloys of mercury such as those found in dental fillings have never been shown to cause any medical issues directly. Despite this the US banned the use of Cu/Hg dental amalgams until 1850! More recently, it has been suggested that dentists are exposed at higher levels during the placing and removal of fillings. Dentists as a group have higher mercury levels than those associated with people with amalgam restorations, but experience no increase in disease or death rates, and in fact tend to be healthier than the general population.
Although elemental mercury was clearly toxic, this did not stop its use in pharmacy for hundreds of years. In the 1500’s mercury was used in the treatment (albeit ineffective) of syphilis. Syphilis was a new disease in Europe; it had been brought back from America by Columbus’ sailors, and was promptly spread through Europe by the French army, amongst others! Syphilis was much more fatal and had more dramatic symptoms than today.
Initially mercury was used as an ointment, but the patients often got worse. Then there was the tub, which was a mercury vapor bath, and even calomel (Hg2Cl2) was used, but with little effect. These treatments were used for over four centuries, but none provided a cure, despite claims at the time. For example, John Hunter, a doctor who gave himself syphilis by mistake (!) claimed he had been cured, but he actually died of a heart attack during an argument, so it is unlikely the mercury worked. Despite this it became known that “a night with Venus results in a lifetime with Mercury”.
The reasons that mercury was thought erroneously to cure syphilis are twofold:
1. Until 1906 it was difficult to diagnose syphilis. It was often confused with gonorrhea, and therefore it is likely that some people did not have the far more deadly syphilis.
2. Syphilis occurs in three phases, each with remission between the phases. The period of remission between secondary and tertiary phases can be two to three years, and therefore it may appear that a cure is found. Especially as may patients (like John Hunter) died other deaths during this remission phase.
The prevalent use of mercury and its presence in many cadavers, led some doctors to assume that mercury was a natural part of the body. It was not just humans that were treated with mercury, cattle were also treated, and one druggist sold 25 tons of mercury to a single farmer in one year!
The density of mercury and its liquid state at room temperature led to another unusual application that was somewhat more successful, although equally dangerous: constipation. In medical texts of the time it was noted that “mercury is given in the disease called Miserere, unto two or three pounds, and is voided again by siege to the same weight; it is better to take a great deal of it that a little, because a small quantity might be apt to stop in the circumvolutions of the guts, and if some acid humors should happen to join with it, a sublimate corrosive would be made; but when a large quantity of it is taken, there’s no need to fearing this accident, because it passes through by its own weight.”
It is interesting that the mention of the corrosive sublimate; this is in fact mercury(II) dichloride (HgCl2) which unlike mercury(I) chloride (Hg2Cl2), is a very violent poison. Death is caused by renal failure. So while there is no evidence for elemental mercury itself causing fatalities, its compounds are another matter to be considered.
Organomercury compounds
In 1953 it was noted that the fishing village of Minamata in Japan had an epidemic in which a large number of people died. Initial thoughts of either an infectious disease or malnutrition were discounted; then it was found that the fish eaten by the villagers was highly contaminated by mercury.
It was found that the mercury came from the Chisso Corporation chemical plant that made acetaldehyde from acetylene using a mercury catalyst. The plant was loosing 1 Kg of mercury metal for every ton of acetaldehyde being produced. As a consequence it was originally assumed that the poisoning of the village was due to inorganic mercury. Based upon prior incidents, an obvious response was to ban consumption of all fish and shellfish. As a consequence no new cases were reported, however, people already effected continued to die. This was unlike any previous mercury poisoning.
Further analysis showed small quantities of water soluble methyl mercury (MeHg+) was present and sequestered by the shellfish to give MeHgSMe. While the lethal effects of organomercury compounds were known, the source of the methyl mercury was a mystery. A group of Swedish researchers showed that the bacterial action in river sediment or rotting fish converted inorganic mercury to either volatile Me2Hg or water soluble MeHg+. With this discovery, it was understood how the anaerobic mud of the estuary near Minamata could perform this methylation.
Of course Minamata was not the first report or an organomercury compound, but it was the first time that it was shown that mercury metal could be converted to a highly toxic organometallic derivative in the environment. The hazardous nature of organomercurials was found almost as soon as the first compounds were reported!
While working in Bunsen’s research group in Marburg, Edward Franklin discovered the synthesis of the zinc analog of a Grignard reagent. Subsequently, in 1851 Franklin moved to Owens College in Manchester where he extended his work to mercury. In his publication he noted that these organomercury compounds had a “nauseous taste”, but didn’t realize they were toxic.
$\mathrm{Zn}+\mathrm{Mel} \rightarrow \mathrm{MeZnl}$
$\mathrm{Hg}+\mathrm{RX} \rightarrow \mathrm{RHgX}$
In 1858 George Buckton working at the then Royal College of Chemistry (now Imperial College) reported the synthesis of dimethyl mercury as a volatile liquid.
\[2 \mathrm{MeHgl}+2 \mathrm{KCN} \rightarrow \mathrm{Me}_{2} \mathrm{Hg}+\mathrm{Hg}+2 \mathrm{KI}+(\mathrm{NC})_{2}]
When Frankland moved his research to St Bartholomew’s Hospital (“Barts”) London he started looking into the chemistry of R2Hg with an assistant called Bill Odling in collaboration with Dr Carl Ulrich.
Ulrich died in 1865 as a consequence of exposure to Me2Hg. In his own statement, he had inhaled a large quantity of the volatile compound without having taken the proper precautions. The following day “his countenance had attained a dull, anxious, and confused expression” and he was admitted to the hospital in a weak condition on 3rd February. On the 9th he became noisy and had to be put under mechanical restraint. The next day his breath and body began to smell offensively and he was in a coma. He would rise from the coma periodically to utter incoherent howls. He died on the 14th of February.
A technician from the same research group (who is only identified as ‘T. C.’) was also admitted to the hospital on 28th March of the same year. His symptoms were initially milder than Ulrich’s, but soon developed. By that summer he was completely demented, with no control over his body functions. He stayed in this state for many months, only dying on 7th April 1866. Records indicate that a third assistant was also taken ill, but there was no further mention of him, so it is unknown if he died.
Summary and the “green” future
Metallic mercury causes severe symptoms, but all records show that if the patient is removed from the source they recover. Thus, short term exposure to metallic mercury, while dangerous, is not fatal if proper precautions are taken. In contrast, mercury compounds offer different risks. As a general rule, inorganic mercury(I) compounds are far less toxic than their mercury(II) analogs, however, all should be treated with care.
Where mercury compounds offer the greatest risk of fatality is their organometallic derivatives. There is no known cure for exposure to even modest doses of organomercury compounds. Furthermore, the ability of elemental mercury to be transformed into water-soluble organomercury compounds such as MeHg+, offers a future threat to public health.
The new generation of low energy consumption light bulbs contain mercury vapor. While they last longer than a traditional tungsten filament light bulb (Figure), they do have a lifetime. The presence of mercury means that they should be disposed-off separately from household waste to ensure that when the glass is broken the mercury is not released; however, this is unlikely. Most will be disposed off along with household waste which may be subsequently land filled. The lesson from Minamata should be that the bacterial action under anaerobic condition allows for the formation of water soluble MeHg+, that can diffuse into the water table. Although the amount of mercury in each bulb is very small, the highly lethal nature (low LD50) of orgnaomercury compounds should be considered in efforts to conserve energy by the use of the low energy light bulbs. At the very minimum protocols for their efficient disposal and recycling should be in place.
Bibliography
• S. Jensen and A. Jernelöv, Nature, 1969, 223, 753.
• J. C. Clifton II, Pediatr. Clin. North Am., 2007, 54, 237.
• H. A. Waldron, Br. Med. J., 1983, 287, 24:
• C. Naleway, R. Sakaguchi, E. Mitchell, T. Muller, W. A. Ayer, and J. J. Hefferren, J. Amer. Dent. Assoc., 1985, 111, 37.
• D. McComb, J. Can. Dent. Assoc., 1997, 63, 372.
• T. G. Duplinsky1 and D. V. Cicchetti, Int. J. Stats. Med. Res., 2012, 1, 1.
• M. Aucott, M. McLinden, and M. Winka, J. Air. Waste Manag. Assoc., 2003, 53, 143. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/05%3A_Group_12/5.05%3A_The_Myth_Reality_and_History_of_Mercury_Toxicity.txt |
Thumbnail: Crystals of 99.999% gallium. (CC-SA-BY 3.0; Foobar)
06: Group 13
Table $1$ lists the derivation of the names of the Group 13 (IIIA) elements.
Element Symbol Name
Boron B From the Arabic word buraq or the Persian word burah for the mineral borax
Aluminium (Aluminum) Al From alum
Gallium Ga From Named after the Latin word for France (Gaul) Gallia
Indium In Latin rubidus meaning deepest red
Thalium Tl From the Latin thallus meaning sprouting green twig
Table $1$: Derivation of the names of each of the alkali metal elements.
Note
Aluminium is the international spelling standardized by the IUPAC, but in the United States it is more commonly spelled as aluminum.
Discovery
Boron
Borax (a mixture of Na2B4O7.4H2O and Na2B4O7.10H2O) was known for thousands of years. In Tibet it was known by the Sanskrit name of tincal. Borax glazes were used in China in 300 AD, and the writings of the Arabic alchemist Geber appear to mention it in 700 AD. However, it is known that Marco Polo brought some borax glazes back to Italy in the 13th century. In 1600 Agricola in his treatise De Re Metallica reported the use of borax as a flux in metallurgy.
Boron was not recognized as an element until its isolation by Sir Humphry Davy, Joseph Louis Gay-Lussac and Louis Jacques Thénard in 1808 through the reaction of boric acid and potassium. Davy called the element boracium. Jöns Jakob Berzelius identified boron as an element in 1824.
Aluminum
Ancient Greeks and Romans used aluminum salts as dyeing mordants and as astringents for dressing wounds; alum (KAl(SO4)2.12H2O) is still used as a styptic (an antihaemorrhagic agent). In 1808, Sir Humphry Davy identified the existence of a metal base of alum, which he at first termed alumium and later aluminum.
The metal was first produced in 1825 (in an impure form) by Hans Christian Ørsted by the reaction of anhydrous aluminum chloride with potassium amalgam. Friedrich Wöhler repeated the experiments of Ørsted but suggested that Ørsted had only isolated potassium. By the use of potassium (Equation 6.1.1), he is credited with isolating aluminum in 1827. While Wöhler is generally credited with isolating aluminum, Ørsted should also be given credit.
$\text{AlCl}_3 \text{ + 3 K} \rightarrow \text{Al + 3 KCl}$
In 1846 Henri Deville improved Wöhler's method, and described his improvements in particular the use of sodium in place of the expensive potassium
$\text{AlCl}_3 \text{ + 3 Na} \rightarrow \text{Al + 3 NaCl}$
Gallium
The element gallium was predicted, as eka-aluminum, by Mendeleev in 1870, and subsequently discovered by Lecoq de Boisbaudran in 1875; in fact de Boisbaudran had been searching for the missing element for some years, based on his own independent theory. The first experimental indication of gallium came with the observation of two new violet lines in the spark spectrum of a sample deposited on zinc. Within a month of these initial results de Boisbaudran had isolated 1 g of the metal starting from several hundred kilograms of crude zinc blende ore. The new element was named in honor of France (Latin Gallia), and the striking similarity of its physical and chemical properties to those predicted by Mendeleev did much to establish the general acceptance of the periodic Law; indeed, when de Boisbaudran first stated that the density of Ga was 4.7 g cm-3 rather than the predicted 5.9 g/cm3, Mendeleev wrote to him suggesting that he redetermine the value (the correct value is 5.904 g/cm3).
Indium
While testing ores from the mines around Freiberg, Saxony, Ferdinand Reich and Hieronymous Theodor Richter when they dissolved the minerals pyrite, arsenopyrite, galena and sphalerite in hydrochloric acid, and since it was known that ores from that region contained thallium they searched for the green emission lines by spectroscopy. Although the green lines were absent, a blue line was present in the spectrum. As no element was known with a bright blue emission they concluded that a new element was present in the minerals. They named the element with the blue spectral line indium. Richter went on to isolate the metal in 1864.
Thallium
After the publication of their improved method of flame spectroscopy by Robert Bunsen and Gustav Kirchhoff this method became an accepted method to determine the composition of minerals and chemical products. Two chemists, William Crookes and Claude-Auguste Lamy, both started to use the new method and independently employed it in their discovery of thallium.
Crookes was making spectroscopic determinations on selenium compounds deposited in the lead chamber of a sulfuric acid production plant near Tilkerode in the Harz mountains. Using a similar spectrometer to Crookes', Lamy was determining the composition of a selenium-containing substance that was deposited during the production of sulfuric acid from pyrite. Using spectroscopy both researchers both observed a new green line the atomic absorption spectrum and assigned it to a new element. Both set out to isolate the new element. Fortunately for Lamy, he had received his material in larger quantities and thus he was able to isolate sufficient quantities of thallium to determine the properties of several compounds and prepare a small ingot of metallic thallium. At the same time Crookes was able to isolate small quantities of elemental thallium and determine the properties of a few compounds. The claim by both scientists resulted in significant controversy during 1862 and 1863; interestingly this ended when Crookes was elected Fellow of the Royal Society in June 1863.
Abundance
The abundance of the Group 13 elements is given in Table $2$. Aluminum is the most abundant metal in the Earth’s crust and is found in a wide range of minerals. While boron is not as common it is also found in a range of borate minerals. In contrast, gallium, indium, and thallium are found as impurities in other minerals. In particular indium and thallium are found in sulfide or selenide mineral rather than oxides, while gallium is found in both sulfides (ZnS) and oxides (bauxite). Although indium and thallium minerals are known, they are rare: indite (FeIn2S4), lorandite (TlAsS2), crookesite (Cu7TlSe4).
Table $2$: Abundance of Group 13 elements.
Element Terrestrial abundance (ppm)
B 10 (Earth’s crust), 20 (soil), 4 (sea water)
Al 82,000 (Earth’s crust), 100,000 (soil), 5 x 10-4 (sea water)
Ga 18 (Earth’s crust), 28 (soil), 30 x 10-6 (sea water)
In 0.1 (Earth’s crust), 0.01 (soil), 0.1 x 10-6 (sea water)
Tl 0.6 (Earth’s crust), 0.2 (soil), 10 x 10-6 (sea water)
Isotopes
The naturally abundant isotopes of the Group 13 elements are listed in Table $3$. Thallium has 25 isotopes that have atomic masses that range from 184 to 210. Thallium-204 is the most stable radioisotope, with a half-life of 3.78 years.
Table $3$: Abundance of the major isotopes of the Group 13 elements.
Isotope Natural abundance (%)
Boron-10 19.9
Boron-11 80.1
Aluminum-27 100
Gallium-69 60.11
Gallium-71 39.89
Indium-113 4.3
Indium-115 95.7
Thallium-203 29.52
Thallium-205 70.48
The Group 13 elements offer potential as NMR nuclei (Table $4$). In particular 11B and 27Al show promise for characterization in both solution and solid state.
Table $4$: Isotopes of Group 13 elements for NMR spectroscopy.
Isotope Spin Natural abundance (%) Quadrupole moment (10-30 m2) NMR frequency (MHz) at a field of 2.3488 T Reference
Boron-10 3 19.58 8.459 -10.746 BF3.Et2O
Boron-11 3/2 80.42 4.059 -32.084 BF3.Et2O
Aluminum-27 5/2 100 14.66 -26.057 Al(NO3)3
Gallium-69 3/2 60.4 17.1 -24.003 Ga(NO3)3
Gallium-71 3/2 39.6 10.7 -30.495 Ga(NO3)3
Indium-113 9/2 4.28 79.9 -21.866 In(NO3)3
Indium-115 9/2 95.72 81.0 -21.914 In(NO3)3
Industrial production
Borax is mined as a mixture of Na2B4O7.4H2O and Na2B4O7.10H2O. Acidification gives boric acid, B(OH)3, which can be reduced with sodium amalgam (Na/Hg) to give amorphous boron. Pure boron can be prepared by reducing boron halides (e.g., BF3 and BCl3) with hydrogen at high temperatures. Ultrapure boron, for the use in semiconductor industry, is produced by the decomposition of diborane (B2H6) and then further purified with the zone melting or Czochralski processes.
The only two economic sources for gallium are as byproduct of aluminum and zinc production. Extraction during the Bayer process followed by mercury cell electrolysis and hydrolysis of the amalgam with sodium hydroxide leads to sodium gallate. Electrolysis then gives gallium metal.
The lack of indium mineral deposits and the fact that indium is enriched in sulfides of lead, tin, copper, iron and zinc, makes the zinc production the main source for indium. The indium is leached from slag and dust of zinc production. Up until 1924, there was only about a gram of isolated indium on the planet, however, today worldwide production is currently greater 476 tons per year from mining and a 650 tons per year from recycling. This massive increase in demand is due to applications in LCD displays and solar cell applications.
Aluminum
Due to aluminum’s position as the most abundant metallic element in the Earth's crust (7.5 - 8.1%) and its enormous industrial importance warrants detailed discussion of its industrial production. Aluminum only appears in its elemental form in nature in oxygen-deficient environments such as volcanic mud. Ordinarily, it is found in a variety of oxide minerals.
In comparison to other metals aluminum is difficult to extract from its ores. Unlike iron, aluminum oxides cannot be reduced by carbon, and so purification is only possible on an economic scale using electrolysis. Prior to electrolysis purified aluminum oxide is obtained by refining bauxite in the process of developed by Karl Bayer.
Bauxite, the most important ore of aluminum, contains only 30-50% alumina, Al2O3, the rest being a mixture of silica, iron oxide, and titanium dioxide. Thus, the alumina must be purified before it can be used as the oxide or refined into aluminum metal. In the Bayer process, bauxite is digested in hot (175 °C) sodium hydroxide (NaOH) solution (Figure). This converts the alumina to aluminum hydroxide, Al(OH)3, which dissolves in the hydroxide solution.
$\text{Al}_2\text{O}_3 \text{ + 2 OH}^- \text{ + 3 H}_2\text{O} \rightarrow \text{2 [Al(OH)}_4\text{]}^-$
The other components do not dissolve and are filtered off. The hydroxide solution is cooled, and the dissolved aluminum hydroxide precipitates, which when heated to 1050 °C is calcined into alumina.
$\text{2 Al(OH)}_3 \rightarrow \text{Al}_2\text{O}_3 \text{ + 3 H}_2\text{O}$
Once a pure alumina is formed, it is dissolved in molten cryolite (Na3AlF6) and reduced to the pure metal at elevated temperatures (950 - 980 °C) using the Hall-Héroult process, developed independently Charles Hall and Paul Héroult.
Both of the electrodes used in the electrolysis of aluminum oxide are carbon. The reaction at the cathode involves the reduction of the Al3+. The aluminum metal then sinks to the bottom and is tapped off, where it is usually cast into large blocks called aluminum billets.
$\text{Al}^{3+} + \text{3 e}^- \rightarrow \text{Al}$
At the anode oxygen is formed, (6.1.6) , where it reacts with the carbon anode is then oxidized to carbon dioxide, (6.1.7). The anodes must be replaced regularly, since they are consumed. While the cathodes are not oxidized they do erode due to electrochemical processes and metal movement.
$\text{2 O}_2^- \rightarrow \text{O}_2 + \text{4 e}^-$
$\text{C + O}_2 \rightarrow \text{CO}_2$
Although the Hall-Héroult process consumes a lot of energy, alternative processes have always found to be less economically and ecologically viable.
Physical properties
Table summarizes the physical properties of the Group 13 elements. While, aluminum, indium, and thallium have typical metal properties, gallium has the largest liquid range of any element. Boron exists as a molecular compound in the solid state, hence its high melting point.
Element Mp (°C) Bp (°C) Density (g/cm3)
B 2300 3658 2.3
Al 661 2467 2.7
Ga 30 2403 5.9 (solid), 6.1 (liquid)
In 156 2080 7.3
Tl 304 1457 11.9
Table $5$: Selected physical properties of the Group 13 elements.
Bibliography
• C. M. Hall. Process of reducing aluminum from its fluoride salts by electrolysis. US Patent 400,664 (1886).
• L. B. Alemany, S. Steuernagel, J.-P. Amoureux, R. L. Callender, and A. R. Barron. Solid State Nuclear Magnetic Resonance, 1999, 14, 1.
• L. B. Alemany, R. L. Callender, A. R. Barron, S. Steuernagel, D. Iuga, and A. P. M. Kentgens, J. Phys. Chem., B., 2000, 104, 11612.
• M. Bishop, N. Shahid, J. Yang, and A. R. Barron, Dalton Trans., 2004, 2621. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/06%3A_Group_13/6.01%3A_The_Group_13_Elements.txt |
Boron is a non-metal with metalloidal tenancies. The higher ionization energies for boron than for its other Group homologs are far more than would be compensated by lattice energies, and thus, the B3+ ion plays no part in the chemistry of boron, and its chemistry is dominated by the formation of covalent compounds. In contrast, the elements aluminum through thallium each has a low electronegativity and the chemistry of their compounds reflects this characteristic. Each of the Group 13 metals forms both covalent compounds and ionic coordination complexes.
All of the Group 13 (IIIA) elements have a valence shell electron configuration of ns2np1. As a consequence all of the Group 13 elements for compounds in which they adopt a +3 oxidation state. While the lighter elements do form compounds with lower oxidation state, they are not the norm; however, the +1 oxidation state is more prevalent for the heavier elements in particular thallium. The rational for this is described as the inert pair effect. The inert pair effect is usually explained by the energy of the ns orbital is lower making it harder to ionize and stabilizing a ns2np0 valence shell. However, as may be seen from Table \(1\) the sum of the second and third ionization enthalpies is lower for indium (4501 kJ/mol), than for gallium (4916 kJ/mol), but with thallium intermediate (4820 kJ/mol). The true source of the inert pair effect is that the lower bond strengths observed for the heavier elements (due to more diffuse orbitals and therefore less efficient overlap) cannot compensate for the energy needed to promote the ns2 electrons. For example, the bond energies for gallium, indium, and thallium in MCl3 are 242, 206, and 153 kJ/mol, respectively. It has also been suggested that relativistic effects make a contribution to the inert pair effect.
Table \(1\): Summary of first three ionization enthalpies for the Group 13 metals.
Ionization enthalpy (kJ/mol) Al Ga In Tl
1 576.4 578.3 558.1 589.0
2 1814.1 1969.3 1811.2 1958.7
3 2741.4 2950.0 2689.3 2868.8
In summary, it may be stated that while the chemistry of gallium, indium and thallium is very similar, that of aluminum is slightly different, while boron’s chemistry is very different from the rest of the Group.
A second effect is noticed in the transition from aluminum, to gallium, to indium. Based upon their position in the Group it would be expected that the ionic radius and associated lattice parameters should follow the trend:
However, as may be seen from Table \(2\) the values for gallium are either the same as, or smaller, than that of aluminum. In a similar manner, the covalent radius and covalent bond lengths as determined by X-ray crystallography for a range of compounds (Table \(3\)).
Table \(2\): Lattice parameter (a) for zinc blende forms of the Group 13 phosphides and arsenides. Data from Semiconductors: Group IV Elements and III-V Compounds, Ed. O. Madelung, Springer-Verlag, Berlin (1991).
Element Phosphide lattice parameter (Å) Arsenide lattice parameter (Å)
Al 5.4635 5.6600
Ga 5.4505 5.6532
In 5.8687 6.0583
Table \(3\): Comparative crystallographically determined bond lengths.
Element M-C (Å) M-N (Å) M-O (Å) M-Cl (Å)
Al 1.96 – 2.02 2.03 – 2.19 1.74 – 1.93 2.09 – 2.11
Ga 1.97 – 2.01 1.95 – 2.12 1.89 - 1.94 2.12 – 2.23
In 2.14 – 2.17 2.23 – 2.31 2.19 – 2.20 2.39 – 2.47
Gallium is significantly smaller than expected from its position within the Group 13 elements (Table \(4\)). The rational for this may be attributed to an analogous effect as seen in the lanthanide contraction observed for the lanthanides and the 3rd row of transition elements. In multi-electron atoms, the decrease in radius brought about by an increase in nuclear charge is partially offset by increasing electrostatic repulsion among electrons. In particular, a “shielding effect” results when electrons are added in outer shells, electrons already present shield the outer electrons from nuclear charge, making them experience a lower effective charge on the nucleus. The shielding effect exerted by the inner electrons decreases in the order s > p > d > f. As a sub-shell is filled in a period the atomic radius decreases. This effect is particularly pronounced in the case of lanthanides, as the 4f sub-shell is not very effective at shielding the outer shell (n = 5 and n = 6) electrons. However, a similar, but smaller effect should be observed with the post-transition metal elements, i.e., gallium. This is indeed observed (Table \(4\)).
Table \(4\): Comparison of the covalent and ionic radii of Group 13 elements.
Element Covalent radius (Å) Ionic radius (Å)
Aluminum 1.21 0.53
Gallium 1.22 0.62
Indium 1.42 0.80
Iron (low spin) 1.32 0.55
Iron (high spin) 1.52 0.64
The anomalous size of gallium has two positive effects.
1. The similarity in size of aluminum and gallium means that their Group 15 derivatives have near identical lattice parameters (Table \(3\)). This allows for both epitaxial growth of one material on the other, and also the formation of ternary mixtures (i.e., AlxGa1-xAs) with matched lattice parameters. The ability to grow hetrojunction structures of Group 13-15 compounds (III-V) is the basis for the fabrication of a wide range of important optoelectronic devices, including: LEDs and laser diodes.
2. The similarity in size of gallium(III) to iron(III) (Table \(4\)) means that gallium can substitute iron in a range of coordination compounds without alteration of the structure. Because of a similar size and charge as Fe3+, Ga3+ is widely used as a non-redox-active Fe3+ substitute for studying metal complexation in proteins and bacterial populations.
Bibliography
• K. S. Pitzer, Acc. Chem. Res., 1979, 12, 271.
• K. D. Weaver, J. J. Heymann, A. Mehta, P. L. Roulhac, D. S. Anderson, A. J. Nowalk, P. Adhikari, T. A. Mietzner, M. C. Fitzgerald, and A. L. Crumbliss, J. Biol. Inorg. Chem., 2008, 13, 887.
• Semiconductors: Group IV Elements and III-V Compounds, Ed. O. Madelung, Springer-Verlag, Berlin (1991). | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/06%3A_Group_13/6.02%3A_Trends_for_the_Group_13_Compounds.txt |
The non-metallic nature of boron means that it makes a number of binary compounds with elements more electropositive than itself (i.e., metals). These compounds are called, borides, and some are also formed with metalloid elements as well (e.g., arsenic). In this regard, borides may be considered similar to carbides, silicides, and some phosphides.
Borides are prepared in a number of ways, however, direct combination of the elements, (6.3.1), is the simplest. Other routes include, electrolysis of the fused salts, and the reduction of the metal oxide with a mixture of carbon and boron carbide.
$\text{M + x B} \rightarrow \text{MB}_\text{x}$
Metal borides are generally refractory in character and chemically inert, while they often have properties better that of the constituent elements. For example, the thermal conductivity of TiB2 is about ten times greater than that of titanium, and the melting point is significantly higher (Table $1$).
Table $1$: The melting points of Group 4 metals and their borides.
Element Melting point (°C) Boride Melting point (°C)
Ti 1725 TiB2 3225
Zr 1855 ZrB2 2990
Hf 2233 HfB2 3100
The structures of metal borides depends on the M:B ratio. Borides with an isolated boron atom have a low B:M ratio: M4B, M3B, M2B, M5B2, and M7B3. In such compounds the boron atom is normally in a triangular-prismatic or square-antiprismatic hole in a metal lattice. Borides with equal or near equal metal and boron ratio have structures with either pairs of boron atoms (as in V3B2), single boron chains (seen in all MB compounds), or double boron chains (observed for many M3B4 compounds). Increasing the boron content results in two-dimensional structures. For example, MB2 usually consists of alternate hexagonal layers of metal and boron (Figure $1$). Finally, boron rich borides (e.g., MB4, MB6, and MB12) all have three-dimensional structures.
6.04: Boron Hydrides
Borane and diborane
Borane (BH3) formed in the gaseous state from decomposition of other compounds, (6.4.2) but cannot be isolated except as a Lewis acid-base complex (6.4.1). As such many borane adducts are known.
$\text{BH}_3 \text{ + PPh}_3 \rightarrow \text{H}_3\text{B-PPh}_3$
In the absence of a Lewis base the dimeric diborane (B2H6) is formed. Diborane is generally synthesized by the reaction if BF3 with a hydride source, such as NaBH4, (6.4.4), or LiAlH4, (6.4.3).
$\text{H}_3\text{B-PF}_3 \xrightarrow{\Delta} \text{BH}_3 \text{ + PF}_3$
$\text{3 LiAlH}_4 \text{ + 4 BF}_3 \rightarrow \text{2 B}_2\text{H}_6 \uparrow \text{ + 3 LiAlF}_4$
The structure of diborane (Figure $1$ a) is considered to be electron deficient, and has been confirmed by IR spectroscopy and electron diffraction. The four terminal B-H bonds are normal covalent bonds, however, the bridging B-H-B unit consists of two three-centered two-electron bonds, each ordinarily considered to be formed by the combination of two boron sp3 orbitals and one hydrogen s orbital (Figure $1$ b). However, a consideration of the H-B-H bond angle associated with the terminal hydrides (120°) it is perhaps better to consider the BH2 fragment to be sp2 hybridized, and the B-H-B bridging unit to be a linear combination of one sp2 orbital and one p orbital from each boron atom with the two hydrogen s orbitals. Diborane represents the archetypal electron deficient dimeric compound, of which Al2Me3 is also a member of this class of electron deficient molecules.
B2H6 is spontaneously and highly exothermically inflammable above 25 °C (ΔH = -2137.7 kJ/mol), (6.4.5). It is often used as one of a range of solvate forms for safety for both its flammability and toxicity.
$\text{3 NaBH}_4 \text{ + 4 BF}_4 \rightarrow \text{2 B}_2\text{H}_6 \uparrow \text{ + 3 NaBF}_4$
$\text{B}_2\text{H}_6 \text{ + 3 O}_2 \rightarrow \text{B}_2\text{O}_3 \text{ + 3 H}_2\text{O}$
Most reactions of diborane involve the cleavage of the dimeric structure. Hydrolysis of diborane yields boric acid, (6.4.6), while alcoholysis yields the appropriate borate ester, (6.4.7). Diborane reacts with Lewis bases to form the appropriate Lewis acid-base complex, (6.4.8).
$\text{B}_2 \text{H}_6 \text{ + 6 H}_2\text{O} \rightarrow \text{2 B(OH)}_3 \text{ + 6 H}_2 \uparrow$
$\text{B}_2 \text{H}_6 \text{ + 6 ROH} \rightarrow \text{2 B(OR)}_3 \text{ + 6 H}_2 \uparrow$
$\text{B}_2 \text{H}_6 \text{ + NR}_3 \rightarrow \text{2 H}_3\text{B-NR}_3$
Borohydride
The borohydride anion (or more properly the tetrahydridoborate anion), BH4-, can be considered as the Lewis acid-base complex between borane and H-. A typical synthesis involves the reaction of a borate ester with a hydride source, (6.4.9).
$\text{4 NaH + B(OMe)}_3 \rightarrow \text{NaBH}_4 \text{ + 3 NaOMe}$
Sodium borohydride is a stable white crystalline solid that is stable in dry air and is non-volatile. The boron in borohydride (BH4-) is tetrahedral. Although it is insoluble in Et2O, it is soluble in water (in which it reacts slowly), THF, ethylene glycol, and pyridine. Interestingly, NaBH4 reacts rapidly with MeOH, but dissolves in EtOH. Sodium borohydride has extensive uses in organic chemistry as a useful reducing agent in which it donates a hydride (H-).
Higher Boranes
Higher boron hydrides contain, in addition to the bridging B-H-B unit, one or more B-B bonds. The higher boranes are usually formed by the thermal decomposition of diborane, (6.4.10) and (6.4.11).
$\text{2 B}_2\text{H}_6 \rightarrow \text{B}_4\text{H}_{10} \text{ + H}_2$
$\text{5 B}_2\text{H}_6 \rightarrow \text{2 B}_5\text{H}_9 \text{ + 6 H}_2$
These higher boranes have ‘open’ cluster structures, e.g., Figure $2$ - Figure $4$. Tetraborane, or to be more precise tetraborane(10) or arachno-B4H10, is a foul-smelling toxic gas. Pentaborane (9) is a toxic liquid (with a distinctive garlic odor) that can detonate in air, and like decaborane(14) was at one time considered as a potential rocket fuel. Because simple boron compounds burn with a characteristic green flame, the nickname for these fuels in the US military was Green Dragon. Problems with using boranes as a fuel included their toxicity and the characteristic of bursting into flame on contact with the air; furthermore, the exhaust would also be toxic. The US program resulted in a stockpile of borane fuels, in particular pentaborane(9), which was not destroyed until 2000. The system for the destroying the boranes was appropriately known as Dragon Slayer.
Bibliography
• D. F. Gaines, Acc. Chem. Res., 1973, 6, 416.
• C. E. Housecroft, Boranes and metalloboranes: structure, bonding and reactivity, Ellis Horwood, Chichester (1990).
• C. F. Lane, Chem. Rev., 1976, 76, 773. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/06%3A_Group_13/6.03%3A_Borides.txt |
Ken Wade (Figure \(1\)) developed a method for the prediction of shapes of borane clusters; however, it may be used for a wide range of substituted boranes (such as carboranes) as well as other classes of cluster compounds.
Wade’s rules are used to rationalize the shape of borane clusters by calculating the total number of skeletal electron pairs (SEP) available for cluster bonding. In using Wade’s rules it is key to understand structural relationship of various boranes (Figure \(2\)).
The general methodology to be followed when applying Wade’s rules is as follows:
1. Determine the total number of valence electrons from the chemical formula, i.e., 3 electrons per B, and 1 electron per H.
2. Subtract 2 electrons for each B-H unit (or C-H in a carborane).
3. Divide the number of remaining electrons by 2 to get the number of skeletal electron pairs (SEP).
4. A cluster with n vertices (i.e., n boron atoms) and n+1 SEP for bonding has a closo structure.
5. A cluster with n-1 vertices (i.e., n-1 boron atoms) and n+1 SEP for bonding has a nido structure.
6. A cluster with n-2 vertices (i.e., n-2 boron atoms) and n+1 SEP for bonding has an arachno structure.
7. A cluster with n-3 vertices (i.e., n-3 boron atoms) and n+1 SEP for bonding has an hypho structure.
8. If the number of boron atoms (i.e., n) is larger than n+1 SEP then the extra boron occupies a capping position on a triangular phase.
What is the structure of B5H11?
1. Total number of valence electrons = (5 x B) + (11 x H) = (5 x 3) + (11 x 1) = 26
2. Number of electrons for each B-H unit = (5 x 2) = 10
3. Number of skeletal electrons = 26 – 10 = 16
4. Number SEP = 16/2 = 8
5. If n+1 = 8 and n-2 = 5 boron atoms, then n = 7
6. Structure of n = 7 is pentagonal bipyramid (Figure \(2\)), therefore B5H11 is an arachno based upon a pentagonal bipyramid with two apexes missing (Figure \(3\)).
What is the structure of B5H9?
1. Total number of valence electrons = (5 x B) + (9 x H) = (5 x 3) + (9 x 1) = 24
2. Number of electrons for each B-H unit = (5 x 2) = 10
3. Number of skeletal electrons = 24 – 10 = 14
4. Number SEP = 14/2 = 7
5. If n+1 = 7 and n-1 = 5 boron atoms, then n = 6
6. Structure of n = 6 is octahedral (Figure \(2\)), therefore B5H9 is a nido structure based upon an octahedral structure with one apex missing (Figure \(4\)).
Example \(1\)
What is the structure of B6H62-?
Solution
Table \(1\) provides a summary of borane cluster with the general formula BnHnx- and their structures as defined by Wade’s rules.
Table \(1\): Wade’s rules for boranes.
Type Basic formula Example # of verticies # of vacancies # of e- in B + charge # of bonding MOs
Closo BnHn2- B6H62- n 0 3n + 2 n + 1
Nido BnHn4- B5H9 n + 1 1 3n + 4 n + 2
Arachno BnHn6- B4H10 n + 2 2 3n + 6 n + 3
Hypho BnHn8- B5H112- n + 3 3 3n + 8 n + 4
Bibliography
• R. W. Rudolph, Acc. Chem. Res., 1976, 9, 446.
• K. Wade, Adv. Inorg. Chem. Radiochem., 1976, 18, 1.
6.06: Trends for the Oxides of the Group 13 Elements
All of the Group 13 elements form a trivalent oxide (M2O3). The chemical properties of the oxides follow the trend acidic to basic going down the Group (Table \(1\)). The physical properties are consistent with the electronegativities and covalent character in the M-O bonds. Thallium oxide is unique in that it decomposes above 100 °C to yield the thalium(I) oxide, Tl2O. The other oxides are all stable to high temperatures.
Table \(1\): Properties of the Group 13 oxides.
Oxide Color Chemical property Melting point (°C)
B2O3 White/colorless Weak acid 450 (trigonal), 510 (tetrahedral)
Al2O3 White/colorless Amphoteric 2072 (α)
Ga2O3 White/colorless Amphoteric 1900 (α), 1725 (β)
In2O3 Yellow Weakly basic 1910
Tl2O3 Brown-black Basic, oxidizing 100 (decomposes)
Bibliography
• G. E. Jellison, Jr., L. W. Panek, P. J. Bray, and G. B. Rouse, Jr., J. Chem. Phys., 1977, 66, 802. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/06%3A_Group_13/6.05%3A_Wades_Rules.txt |
Oxides
Boron oxide, B2O3, is made by the dehydration of boric acid, (6.7.1). It is a glassy solid with no regular structure, but can be crystallized with extreme difficulty. The structure consists of infinite chains of triangular BO3 unit (Figure $1$). Boron oxide is acidic and reacts with water to reform boric acid, (6.7.1).
$\text{2 B(OH)}_3 \xrightleftharpoons{\Delta} \text{B}_2\text{O}_3 \text{ + H}_2\text{O}$
The reaction of B2O3 with hydroxide yields the metaborate ion, (6.7.2), whose planar structure (Figure $2$) is related to metaboric acid. Boron oxide fuses with a wide range of metal and non-metal oxides to give borate glasses.
$\text{B}_2\text{O}_3\text{ + 6 KOH} \rightarrow \text{2 K}_3\text{[B}_3\text{O}_6\text{] + 3 H}_2\text{O}$
Boric acid
Boric acid, B(OH)3, usually obtained from the dissolution of borax, Na2[B4O5(OH)4], is a planar solid with intermolecular hydrogen bonding forming a near hexagonal layered structure, broadly similar to graphite (Figure $3$). The inter layer distance is 3.18 Å.
A summary of reactions of the reactivity of boric acid is shown in Figure $4$.
Upon dissolution of boric acid in water, boric acid does not act as a proton acid, but instead reacts as a Lewis acid, (6.7.3).
$\text{B(OH)}_3 \text{ + 2 H}_2\text{O} \rightleftharpoons \text{B(OH)}_4^- \text{ + H}_3\text{O}^+$
The reaction may be followed by 11B NMR spectroscopy from the change in the chemical shift (Figure $5$).
Since the 11B NMR shift is directly proportional to the mole fraction of the total species present as the borate anion (e.g., [B(OH)4]-) the 11B NMR chemical shift at a given temperature, δ(obs), may be used to calculate both the mole fraction of boric acid and the borate anion, i.e., (6.7.4) and (6.7.5), respectively. Using these equations the relative speciation as a function of pH may be calculated for both boric acid (Figure $6$). The pH at which a 50:50 mixture of acid and anion for boric acid is ca. 9.4.
$\chi_{\text{(acid)}} = \dfrac{\delta_{\text{(obs)}} - \delta_{\text{(anion)}}}{\delta_{\text{(acid)}} - \delta_{\text{(anion)}}}$
$\chi_{\text{(acid)}} = \dfrac{\delta_{\text{(acid)}} - \delta_{\text{(obs)}}}{\delta_{\text{(acid)}} - \delta_{\text{(anion)}}}$
In concentrated solutions, the borate ion reacts further to form polyborate ions. The identity of the polyborate is dependent on the pH. With increasing pH, B5O6(OH)4-, (6.7.5), B3O3(OH)4-, (6.7.6), and B4O5(OH)42-, (6.7.7), are formed. The structure of each borate is shown in Figure $7$. Once the ratio of B(OH)3 to B(OH)4- is greater than 50%, only the mono-borate is observed.
$\text{4 B(OH)}_3 \text{ + B(OH)}^-_4 \rightleftharpoons \text{B}_5\text{O}_6\text{(OH)}_4^- \text{ + 6 H}_2\text{O}$
$\text{4 B(OH)}_3 \text{ + B(OH)}^-_4 \rightleftharpoons \text{B}_3\text{O}_3\text{(OH)}_4^- \text{ + 3 H}_2\text{O}$
$\text{2 B(OH)}_3 \text{ + 2 B(OH)}^-_4 \rightleftharpoons \text{B}_4\text{O}_5\text{(OH)}_4^{2-} \text{ + 5 H}_2\text{O}$
Borax, the usual mineral form of boric acid, is the sodium salt, Na2[B4O5(OH)4], which upon dissolution in water re-equilibrates to B(OH)3.
Enough to make your hair curl
In 1906, a German hairdresser, Charles Nessler who was living in London, decided to help his sister who was fed-up with having to put her straight hair in curlers. While looking for a solution, Nessler noticed that a clothesline contracted in a wavy shape when it was wet. Nessler wound his sister’s hair on cardboard tubes; then he covered the hair with borax paste. After wrapping the tubes with paper (to exclude air) he heated the entire mass for several hours. Removing the paper and tubes resulted in curly hair. After much trial and error (presumably at his sisters discomfort) Nessler perfected the method by 1911, and called the process a permanent wave (or perm). The process involved the alkaline borax softening the hair sufficiently to be remodeled, while the heating stiffened the borax to hold the hair into shape. The low cost of borax meant that Nessler’s methods was an immediate success.
Metaboric acid
Heating boric acid results in the partial dehydration to yield metaboric acid, HBO2, (6.7.8). Metaboric acid is also formed from the partial hydrolysis of B2O3.
$\text{B(OH)}_3 \xrightleftharpoons[\text{H}_2\text{O}]{\Delta} \text{HBO}_2 \xrightleftharpoons[\text{H}_2\text{O}]{\Delta} \text{B}_2\text{O}_3$
If the heating is carried out below 130 °C, HBO2-III is formed in which B3O3 rings are joined by hydrogen bonding to the hydroxide on each boron atom (Figure $8$). Continued heating to 150 °C results in HBO2-II, whose structure consists of BO4 tetrahedra and B2O5 groups chain linked by hydrogen bonding. Finally, heating above 150 °C yields cubic HBO2-I with all the boron atoms tetrahedral.
Borate esters
Boric acid reacts with alcohols in the presence of sulfuric acid to form B(OR)3 (Figure $4$). This is the basis for a simple flame test for boron. Treatment of a compounds with methanol/sulfuric acid, followed by placing the reaction product in a flame results in a green flame due to B(OMe)3.
In the presence of diols, polyols, or polysaccharides, boric acid reacts to form alkoxide complexes. In the case of diols, the mono-diol [B(OH)2L]- (Figure $9$a) and the bis-diol [BL2]- (Figure $9$b) complexes.
Originally it was proposed that diols react with the anion rather than boric acid, (6.7.9). In contrast, it was suggested that the optimum pH for the formation of carboxylic acid complexes is under the condition where pKa (carboxylic acid) < pH < pKa (boric acid). Under these conditions it is the boric acid, B(OH)3, not the borate anion, [B(OH)4]-, that reacts to form the complex.
$\text{[B(OH)}_4\text{]}^- \xrightleftharpoons{\text{+ LH}_2} \text{[B(OH)}_2\text{L]}^- \xrightleftharpoons{\text{+ LH}_2} \text{[BL}_2\text{]}^-$
The conversion of boric acid to borate (6.7.3) must occur through attack of hydroxide or the deprotonation of a coordinated water ligand, either of which is related to the pKa of water. The formation of [B(OH)2L]- from B(OH)3 would be expected therefore to occur via a similar initial reaction (attack by RO- or deprotonation of coordinated ROH) followed by a subsequent elimination of H2O and the formation of a chelate coordination, and would therefore be related to the pKa of the alcohol. Thus the pH at which [B(OH)2L]- is formed relative to [B(OH)4]- will depend on the relative acidity of the alcohol. The pKa of a simple alcohol (e.g., MeOH = 15.5, EtOH = 15.9) are close to the value for water (15.7) and the lowest pH at which [B(OH)2L]- is formed should be comparable to that at which [B(OH)4]- forms. Thus, the formation of [B(OH)2L]- as compared to [B(OH)4]- is a competition between the reaction of B(OH)3 with RO- and OH- (Figure $10$), and as with carboxylic acids it is boric acid not borate that reacts with the alcohol, (6.7.10).
$\text{B(OH)}_3 \xrightleftharpoons{\text{+ LH}_2} \text{[B(OH)}_2\text{L]}^- \xrightleftharpoons{\text{+ LH}_2} \text{[BL}_2\text{]}^-$
A key issue in the structural characterization and understanding of the diol/boric acid system is the assignment of the 11B NMR spectral shifts associated with various complexes. A graphical representation of the observed shift ranges for boric acid chelate systems is shown in Figure $11$.
Boric acid cross-linking of guar gum for hydraulic fracturing fluids
Thick gels of guar gum cross-linked with borax or a transition metal complex are used in the oil well drilling industry as hydraulic fracturing fluids. The polysaccharide guaran (Mw ≈ 106 Da) is the major (>85 wt%) component of guar gum, and consists of a (1→4)-β-D-mannopyranosyl backbone with α-D-galactopyranosyl side chain units attached via (1→6) linkages. Although the exact ratio varies between different crops of guar gum the general structure is consistent with about one galactose to every other mannose (Figure $12$).
The synthesis of a typical boric acid cross-linked guar gel, with viscosity properties suitable for use as a fracturing fluid, involves mixing guar gum, water and boric acid, B(OH)3, in a 10:2000:1 wt/wt ratio. Adjusting the pH to between 8.5 and 9 results in a viscous gel. The boron:guaran ratio corresponds to almost 2 boron centers per 3 monosaccharide repeat units. Such a large excess of boric acid clearly indicates that the system is not optimized for cross-linking, i.e., a significant fraction of the boric acid is ineffective as a cross-linking agent.
The reasons for the inefficiency of boric acid to cross-link guaran (almost 2 borate ions per 3 monosaccharide repeat unit are required for a viscous gel suitable as a fracturing fluid): the most reactive sites on the component saccharides (mannose and galactose) are precluded from reaction by the nature of the guar structure; the comparable acidity (pKa) of the remaining guaran alcohol substituents and the water solvent, results in a competition between cross-linking and borate formation; a significant fraction of the boric acid is ineffective in cross-linking guar due to the modest equilibrium (Keq).
Bibliography
• F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 4th Ed. Wiley Interscience (1980).
• M. Bishop, N. Shahid, J. Yang, and A. R. Barron, Dalton Trans., 2004, 2621.
• M. Van Duin, J. A. Peters, A. P. G. Kieboom, and H. Van Bekkum, Tetrahedron, 1984, 40, 2901.
• Zonal Isolation, “Borate Crosslinked Fluids,” by R. J. Powell and J. M Terracina, Halliburton, Duncan, OK (1998).
• R. Schechter, Oil Well Stimulation, Prentice-Hall Inc., Englewood Cliffs, NJ (1992).
• S. Kesavan and R. K. Prud'homme, Macromolecules, 1992, 25, 2026.
6.08: Aluminum Oxides Hydroxides and Hydrated Oxides
The many forms of aluminum oxides and hydroxides are linked by complex structural relationships. Bauxite has the formula Alx(OH)3-2x (0 < x < 1) and is thus a mixture of Al2O3 (α-alumina), Al(OH)3 (gibbsite), and AlO(OH) (boehmite). The latter is an industrially important compound that is used in the form of a gel as a pre-ceramic in the production of fibers and coatings, and as a fire-retarding agent in plastics.
Knowledge of microstructural evolution in ceramic systems is important in determining their end-use application. In this regard alumina has been the subject of many studies in which the phase, morphology, porosity and crystallinity are controlled by physical and chemical processing. The transformation from boehmite [γ-Al(O)(OH)] o corundum (α-Al2O3) has been well characterized and is known to go through the following sequence:
$\gamma\text{-Al(O)(OH)} \xrightarrow{\approx\text{500 °C}} \gamma\text{-Al}_2\text{O}_3 \xrightarrow{\approx\text{1000 °C}} \theta\text{-Al}_2\text{O}_3 \xrightarrow{\text{>1100 °C}} \alpha\text{-Al}_2\text{O}_3$
The phase changes from boehmite through θ-Al2O3 are known to be topotactic (i.e., changes in crystal structure are accomplished without changes in crystalline morphology), however, each phase change is accompanied by a change in porosity. The θ- to α-Al2O3 phase transition occurs through nucleation and growth of the θ-Al2O3 crystallites. The α-Al2O3 phase transition temperature can be altered by the addition of certain additives. For example, because the α-Al2O3 phase occurs by nucleation, the addition of small seed crystals can lower the transition temperature between 100 and 200 °C. The addition of certain transition metals (chromium, manganese, iron, cobalt, nickel, and copper) has also been shown to decrease the transition temperature, while lanthanum or rare earth metals tend to increase the temperature. Finally, the addition of metal oxides has also shown to affect the growth rate in α-Al2O3.
A third form of Al2O3 forms on the surface of the clean aluminum metal, (6.8.2). This oxide skin is rapidly self-repairing because its heat of formation is so large (ΔH = 3351 kJ/mol). The thin, tough, transparent oxide layer is the reason for much of the usefulness of aluminum.
$\text{4 Al + 3 O}_2 \rightarrow \text{2 Al}_2\text{O}_3$
Bibliography
• K. Wefers and C. Misra, Oxides and Hydroxides of Aluminum, Alcoa Laboratories (1987).
• H. L. Wen and F. S. Yen, J. Cryst. Growth, 2000, 208, 696.
• G. K Priya, P. Padmaja, K. G. K. Warrier, A. D. Damodaran, and G. Aruldhas, J. Mater. Sci. Lett., 1997, 16, 1584.
• E. Prouzet, D. Fargeot, and J. F. Baumard, J. Mater. Sci. Lett., 1990, 9, 779. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/06%3A_Group_13/6.07%3A_Boron_Oxides_Hydroxides_and_Oxyanions.txt |
Introduction
While aluminum is the most abundant metal in the earth's crust (ca. 8%) and aluminum compounds such as alum, K[Al(SO4)2].12(H2O), were known throughout the world in ancient times, it was not until the isolation of aluminum in the late eighteenth century by the Danish scientist H. C. Öersted that research into the chemistry of the Group 13 elements began in earnest. Initially, metallic aluminum was isolated by the reduction of aluminum trichloride with potassium or sodium; however, with the advent of inexpensive electric power in the late 1800's, it became economically feasible to extract the metal via the electrolyis of alumina (Al2O3) dissolved in cryolite, Na3AlF6, (the Hall-Heroult process). Today, alumina is prepared by the Bayer process, in which the mineral bauxite (named for Les Baux, France, where it was first discovered) is dissolved with aqueous hydroxides, and the solution is filtered and treated with CO2 to precipitate alumina. With availability of both the mineral and cheap electric power being the major considerations in the economical production of aluminum, it is not surprising that the leading producers of aluminum are the United States, Japan, Australia, Canada, and the former Soviet Union.
Aluminum oxides and hydroxides
The many forms of aluminum oxides and hydroxides are linked by complex structural relationships. Bauxite has the formula Alx(OH)3-2x (0 < x < 1) and is thus a mixture of Al2O3 (α-alumina), Al(OH)3 (gibbsite), and AlO(OH) (boehmite). The latter is an industrially important compound which is used in the form of a gel as a pre-ceramic in the production of fibers and coatings, and as a fire retarding agent in plastics.
Heating boehmite and diaspore to 450 °C causes dehydration to yield forms of alumina which have structures related to their oxide-hydroxide precursors. Thus, boehmite produces the low-temperature form γ-alumina, while heating diaspore will give α-alumina (corundum). γ-alumina converts to the hcp structure at 1100 °C. A third form of Al2O3 forms on the surface of the clean aluminum metal. The thin, tough, transparent oxide layer is the reason for much of the usefulness of aluminum. This oxide skin is rapidly self-repairing because its heat of formation is so large (ΔH = -3351 kJ/mol).
$\text{4 Al + 3 O}_2 \rightarrow \text{2 Al}_2\text{O}_3$
Ternary and mixed-metal oxides
A further consequence of the stability of alumina is that most if not all of the naturally occurring aluminum compounds are oxides. Indeed, many precious gemstones are actually corundum doped with impurities. Replacement of aluminum ions with trace amounts of transition-metal ions transforms the formerly colorless mineral into ruby (red, Cr3+), sapphire (blue, Fe2+/3+, Ti4+), or topaz (yellow, Fe3+). The addition of stoichiometric amounts of metal ions causes a shift from the α-Al2O3 hcp structure to the other common oxide structures found in nature. Examples include the perovskite structure for ABO3 type minerals (e.g., CeTiO7 or LaAlO3) and the spinel structure for AB2O4 minerals (e.g., beryl, BeAl2O4).
Aluminum oxide also forms ternary and mixed-metal oxide phases. Ternary systems such as mullite (Al6Si2O13), yttrium aluminum garnet (YAG, Y3Al5O12), the β-aluminas (e.g., NaAl11O17) and aluminates such as hibonite (CaAl12O19) possessing β-alumina or magnetoplumbite-type structures can offer advantages over those of the binary aluminum oxides.
Applications of these materials are found in areas such as engineering composite materials, coatings, technical and electronic ceramics, and catalysts. For example, mullite has exceptional high temperature shock resistance and is widely used as an infrared-transparent window for high temperature applications, as a substrate in multilayer electronic device packaging, and in high temperature structural applications. Hibonite and other hexaluminates with similar structures are being evaluated as interfacial coatings for ceramic matrix composites due to their high thermal stability and unique crystallographic structures. Furthermore, aluminum oxides doped with an alkali, alkaline earth, rare earth, or transition metal are of interest for their enhanced chemical and physical properties in applications utilizing their unique optoelectronic properties.
Synthesis of aluminum oxide ceramics
In common with the majority of oxide ceramics, two primary synthetic processes are employed for the production of aluminum oxide and mixed metal oxide materials:
1. The traditional ceramic powder process.
2. The solution-gelation, or "sol-gel" process.
The environmental impact of alumina and alumina-based ceramics is in general negligible; however, the same cannot be said for these methods of preparation. As practiced commercially, both of the above processes can have a significant detrimental environmental impact.
Traditional ceramic processing
Traditional ceramic processing involves three basic steps generally referred to as powder-processing, shape-forming, and densification, often with a final mechanical finishing step. Although several steps may be energy intensive, the most direct environmental impact arises from the shape-forming process where various binders, solvents, and other potentially toxic agents are added to form and stabilize a solid ("green") body (Table $1$).
Table $1$: Typical composition of alumina green body
Function Composition Volume (%)
Powder alumina (Al2O3) 27
Solvent 1,1,1-trichloroethane/ethanol 58
Deflocculant menhaden oil 1.8
Binder poly(vinyl butyrol) 4.4
Plasticizer poly(ethylene glycol)/octyl phthalate 8.8
The component chemicals are mixed to a slurry, cast, then dried and fired. In addition to any innate health risk associated with the chemical processing these agents are subsequently removed in gaseous form by direct evaporation or pyrolysis. The replacement of chlorinated solvents such as 1,1,1-trichloroethylene (TCE) must be regarded as a high priority for limiting environmental pollution. The United States Environmental Protection Agency (EPA) included TCE on its 1991 list of 17 high-priority toxic chemicals targeted for source reduction. The plasticizers, binders, and alcohols used in the process present a number of potential environmental impacts associated with the release of combustion products during firing of the ceramics, and the need to recycle or discharge alcohols which, in the case of discharge to waterways, may exert high biological oxygen demands in the receiving communities. It would be desirable, therefore, to be able to use aqueous processing; however, this has previously been unsuccessful due to problems associated with batching, milling, and forming. Nevertheless, with a suitable choice of binders, etc., aqueous processing is possible. Unfortunately, in many cast-parts formed by green body processing the liquid solvent alone consists of over 50 % of the initial volume, and while this is not directly of an environmental concern, the resultant shrinkage makes near net shape processing difficult.
Sol-gel
Whereas the traditional sintering process is used primarily for the manufacture of dense parts, the solution-gelation (sol-gel) process has been applied industrially primarily for the production of porous materials and coatings.
Sol-gel involves a four stage process: dispersion, gelation, drying, and firing. A stable liquid dispersion or sol of the colloidal ceramic precursor is initially formed in a solvent with appropriate additives. By changing the concentration (aging) or pH, the dispersion is "polymerized" to form a solid dispersion or gel. The excess liquid is removed from this gel by drying and the final ceramic is formed by firing the gel at higher temperatures.
The common sol-gel route to aluminum oxides employs aluminum hydroxide or hydroxide-based material as the solid colloid, the second phase being water and/or an organic solvent, however, the strong interactions of the freshly precipitated alumina gels with ions from the precursor solutions makes it difficult to prepare these gels in pure form. To avoid this complication, alumina gels are also prepared from the hydrolysis of aluminum alkoxides, Al(OR)3.
$\text{Al(OR)}_3 \text{ + H}_2\text{O/H}^+ \rightarrow \text{Al-gel}$
$\text{Al-gel} \xrightarrow{\Delta} \text{Al}_2\text{O}_3$
The exact composition of the gel in commercial systems is ordinarily proprietary, however, a typical composition will include an aluminum compound, a mineral acid, and a complexing agent to inhibit premature precipitation of the gel, e.g., Table $2$.
Table $2$: Typical composition of an alumina sol-gel for slipcast ceramics.
Function Composition
Boehmite precursor ASB [aluminum sec-butoxide, Al(OC4H9)3]
Electrolyte HNO3 0.07 mole/mole ASB
Complexing agent glycerol ca. 10 wt.%
The principal environmental consequences arising from the sol-gel process are those associated with the use of strong acids, plasticizers, binders, solvents, and sec-butanol formed during the reaction. Depending on the firing conditions, variable amounts of organic materials such as binders and plasticizers may be released as combustion products. NOx’s may also be produced in the off-gas from residual nitric acid or nitrate salts. Moreover, acids and solvents must be recycled or disposed of. Energy consumption in the process entails “upstream” environmental emissions associated with the production of that energy.
Bibliography
• Advances in Ceramics, Eds. J. A. Mangels and G. L. Messing, American Ceramic Society, Westville, OH, 1984, Vol. 9.
• Adkins, J. Am. Chem. Soc., 1922, 44, 2175.
• A. R. Barron, Comm. Inorg. Chem., 1993, 14, 123.
• M. K. Cinibulk, Ceram. Eng. Sci., Proc., 1994, 15, 721.
• F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 5th Ed., John Wiley and Sons, New York (1988).
• N. N. Greenwood and A. Earnshaw, Chemistry of the Elements, Pergamon Press, Oxford (1984).
• P. H. Hsu and T. F. Bates, Mineral Mag., 1964, 33, 749.
• W. D. Kingery, H. K. Bowen, and D. R. Uhlmann, Introduction to Ceramics, 2nd Ed. Wiley, New York (1976).
• H. Schneider, K. Okada, and J. Pask, Mullite and Mullite Ceramics, Wiley (1994).
• R. V. Thomas, Systems Analysis and Water Quality Management, McGraw-Hill, New York (1972).
• J. C. Williams, in Treatise on Materials Science and Technology, Ed. F. F. Y. Wang, Academic Press, New York (1976). | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/06%3A_Group_13/6.09%3A_Ceramic_Processing_of_Alumina.txt |
The “B-N” unit is isoelectronic (3 + 5 valence electrons) to the “C-C” unit (4 + 4 valence electrons). The two moieties are also isolobal, and as such there are many of the compound types formed by carbon have analogous derivatives in the chemistry of boron-nitrogen.
Lewis acid-base addition compounds
Boron compounds, BX3, are strong Lewis acids and as such form stable addition compounds with Lewis bases, in particular those with nitrogen donor ligands.
$\text{BF}_3 \text{ + NMe}_3 \rightarrow \text{F}_3\text{B-NMe}_3$
In principle these Lewis acid-base complexes should be similar to their isolobal hydrocarbon analogs, however, whereas the dipole in ethane is zero (by symmetry) the dipole in H3NBH3 is 5.2 D as a consequence of the difference in the Pauling electronegativities (i.e., B = 2.04 and N = 3.04). It is this dipole that generally differentiates the B-N compounds from their C-C analogs.
Homolytic cleavage of the C-C bond in ethane will yield two neutral methyl radicals, (6.7.2). In contrast, heterolytic cleavage will result in the formation of two charged species, (6.7.3). Thus, the products either have a net spin, (6.7.2), or a net charge, (6.7.3). By contrast, cleavage of the B-N bond in H3N-BH3 either yields products with both spin and charge, (6.7.5), or neither, (6.7.4). Heterolytic cleavage of the B-N bond yields neutral compounds, (6.7.4), while hemolytic cleavage results in the formation of radical ions, (6.7.5).
$\text{H}_3\text{C-CH}_3 \rightarrow \cdot\text{CH}_3 \text{ + } \cdot\text{CH}_3$
$\text{H}_3\text{C-CH}_3 \rightarrow \text{CH}^+_3 \text{ + } \cdot\text{CH}^-_3$
$\text{H}_3\text{N-BH}_3 \rightarrow \text{NH}_3 \text{ + BH}_3$
$\text{H}_3\text{N-BH}_3 \rightarrow \text{NH}^+_3\cdot \text{ + BH}^-_3\cdot$
The difference in bond strength between H3N-BH3 and ethane is reflected in the difference in bond lengths (Table $1$).
Table $1$: A comparison of bonding in H3E-E’H3.
Compound Bond length (Å) Bond strength (kcal/mol)
H3C-CH3 1.533 89
H3N-BH3 1.658 31
Aminoboranes
The group R’2N-BR2 is isoelectronic and isolobal to the olefin sub-unit R’2C=CR2, and there is even appreciable π-bonding character (Figure $1$). A measure of the multiple bond character can be seen from a comparison of the calculated B-N bond in H2NBN2 (1.391 Å) as compared to a typical olefin (1.33 Å). It is interesting that a consideration of the possible resonance form (Figure $1$) suggest the dipole in the σ-bond is in the opposite direction of that of the π-bond.
Unlike olefins, borazines oligomerize to form dimmers and trimers (Figure $2$) in the absence of significant steric hindrance. Analogous structures are also observed for the other Group 13-15 homologs (R2AlNR’2, R2GaPR’2, etc.).
Borazines
The condensation of boron hydride with ammonia results in the formation of a benzene analog: borazine, (6.7.6). Substituted derivatives are formed by the reaction with primary amines.
$\text{BH}_3 \text{ + NH}_3 \rightleftharpoons \text{H}_3\text{B-NH}_3 \xrightarrow[\text{- H}_2]{\Delta} \text{[H}_2\text{B-NH}_2\text{]}_n \xrightarrow[\text{- H}_2]{\Delta} \text{[HBNH]}_6$
Despite the cyclic structure (Figure $3$), borazine is not a true analog of benzene. Despite all the B-N bond distances being equal (1.44 Å) consistent with a delocalized structure, the difference in electronegativity of boron and nitrogen (2.04 and 3.04, respectively) results in a polarization of the bonds (i.e., Bδ+-Nδ-) and hence a limit to the delocalization. The molecular orbitals of the π-system in borazine are lumpy in appearance (Figure $4$ a) compared to benzene (Figure $4$ b). This uneven distribution makes borazine prone to addition reactions, making it as a molecule less stable than benzene.
Iminoboranes: analogs of acetylene
Iminoboranes, RB=NR’, are analogs of alkynes, but like aminoboranes are only isolated as monomers with sterically hindered subsistent. In the absence of sufficient steric bulk oligomerization occurs, forming substituted benzene analogs.
Boron nitrides: analogs of elemental carbon
The fusion of borax, Na2[B4O5(OH)4] with ammonium chloride (NH4Cl) results in the formation of hexagonal boron nitride (h-BN). Although h-BN has a planar, layered structure consisting of six-member rings similar to graphite (Figure $5$), it is a white solid. The difference in color is symptomatic of the more localized bonding in BN than in graphite.
As is found for its carbon analog, hexagonal boron nitride (h-BN or α-BN) is converted at high temperatures (600 – 2000 °C) and pressures (50 – 200 kbar) to a cubic phase (c-BN or β-BN). In a similar manner to diamond, cubic-BN is very hard being actually able to cut diamond, and as a consequence its main use is as an industrial grinding agent. The cubic form has the sphalerite crystal structure (Figure $6$). Finally, a wurtzite form of boron nitride (w-BN) is known that has similar structure as lonsdaleite, rare hexagonal polymorph of carbon. Table $2$ shows a comparison of the properties of the hexagonal and cubic phases of BN with their carbon analogs.
Table $2$: Comparison of structural and physical properties of carbon and boron nitride analogs.
Phase Carbon Boron nitride
Cubic Colorless, hard, mp = 3550 °C. C-C = 1.514 Å Colorless, hard, B-N = 1.56 Å
Hexagonal Black solid, planar layers, conductor, mp = 3652 – 3697 °C (sublimes), C-C = 1.415 Å White solid, planar layers, semiconductor (Eg = 5.2 eV), mp = 2973 °C (sublimes), B-N = 1.45 Å
The partly ionic structure of BN layers in h-BN reduces covalency and electrical conductivity, whereas the interlayer interaction increases resulting in higher hardness of h-BN relative to graphite.
Bibliography
• K. M. Bissett and T. M. Gilbert, Organometallics, 2004, 23, 850.
• P. Paetzold, Adv. Inorg. Chem., 1987, 31, 123.
• L. R. Thorne, R. D. Suenram, and F. J. Lovas, J. Chem. Phys., 1983, 78, 167. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/06%3A_Group_13/6.10%3A_Boron_Compounds_with_Nitrogen_Donors.txt |
Gallium: the element
The element gallium was predicted, as eka-aluminum, by Mendeleev in 1870, and subsequently discovered by Lecoq de Boisbaudran in 1875; in fact de Boisbaudran had been searching for the missing element for some years, based on his own independent theory. The first experimental indication of gallium came with the observation of two new violet lines in the spark spectrum of a sample deposited on zinc. Within a month of these initial results de Boisbaudran had isolated 1 g of the metal starting from several hundred kilograms of crude zinc blende ore. The new element was named in honor of France (Latin Gallia), and the striking similarity of its physical and chemical properties to those predicted by Mendeleev (Table \(1\)) did much to establish the general acceptance of the periodic Law; indeed, when de Boisbaudran first stated that the density of Ga was 4.7 g/cm3 rather than the predicted 5.9 g/cm3, Mendeleev wrote to him suggesting that he redetermine the value (the correct value is 5.904 g/cm3).
Table \(1\): Comparison of predicted and observed properties of gallium.
Property Mendeleev's prediction (1871) for eka-aluminum, M Observed properties of gallium (discovered 1875)
Atomic weight ca. 68 69.72
Density, g.cm-3 5.9 5.904
Melting point Low 29.78
Vapor pressure Non-volatile 10-3 mmHg, 1000 °C
Valence 3 3
Oxide M2O3 Ga2O3
Density of oxide (g/cm3) 5.5 5.88
Properties of metal M should dissolve slowly in acids and alkalis and be stable in air Ga metal dissolves slowly in acids and alkalis and is stable in air
Properties of hydroxide M(OH)3 should dissolve in both acids and alkalis Ga(OH)3 dissolves in both acids and alkalis
Properties of salts M salts will tend to form basic salts; the sulfate should form alums; M2S3 should be precipitated by H2S or (NH4)2S; anhydrous MCl3 should be more volatile than ZnCl2 Ga salts readily hydrolyze and form basic salts; alums are known; Ga2S3 can be precipitated under special conditions by H2S or (NH4)2S, anhydrous GaCl3 is more volatile than ZnCl2.
Gallium has a beautiful silvery blue appearance; it wets glass, porcelain, and most other surfaces (except quartz, graphite, and Teflon®) and forms a brilliant mirror when painted on to glass. The atomic radius and first ionization potential of gallium are almost identical with those of aluminum and the two elements frequently resemble each other in chemical properties. Both are amphoteric, but gallium is less electropositive as indicated by its lower electrode potential. Differences in the chemistry of the two elements can be related to the presence of a filled set of 3d orbitals in gallium.
Gallium is very much less abundant than aluminum and tends to occur at low concentrations in sulfide minerals rather than as oxides, although gallium is also found associated with aluminum in bauxite. The main source of gallium is as a by-product of aluminum refining. At 19 ppm of the earth's crust, gallium is about as abundant as nitrogen, lithium and lead; it is twice as abundant as boron (9 ppm), but is more difficult to extract due to the lack of any major gallium-containing ore. Gallium always occurs in association either with zinc or germanium, its neighbors in the periodic table, or with aluminum in the same group. Thus, the highest concentrations (0.1 - 1%) are in the rare mineral germanite (a complex sulfide of Zn, Cu, Ge, and As); concentrations in sphalerite (ZnS), bauxite, or coal, are a hundred-fold less.
Gallium pnictides
Gallium's main use is in semiconductor technology. For example, GaAs and related compounds can convert electricity directly into coherent light (laser diodes) and is employed in electroluminescent light-emitting diodes (LED's); it is also used for doping other semiconductors and in solid-state devices such as heterojunction bipolar transistors (HBTs) and high power high speed metal semiconductor field effect transistors (MESFETs). The compound MgGa2O4 is used in ultraviolet-activated powders as a brilliant green phosphor used in Xerox copying machines. Minor uses are as high-temperature liquid seals, manometric fluids and heat-transfer media, and for low-temperature solders.
Undoubtedly the binary compounds of gallium with the most industrial interest are those of the Group 15 (V) elements, GaE (E = N, P, As, Sb). The compounds which gallium forms with nitrogen, phosphorus, arsenic, and antimony are isoelectronic with the Group 14 elements. There has been considerable interest, particularly in the physical properties of these compounds, since 1952 when Welker first showed that they had semiconducting properties analogous to those of silicon and germanium.
Gallium phosphide, arsenide, and antimonide can all be prepared by direct reaction of the elements; this is normally done in sealed silica tubes or in a graphite crucible under hydrogen. Phase diagram data is hard to obtain in the gallium-phosphorus system because of loss of phosphorus from the bulk material at elevated temperatures. Thus, GaP has a vapor pressure of more than 13.5 atm at its melting point; as compared to 0.89 atm for GaAs. The physical properties of these three compounds are compared with those of the nitride in Table \(2\). All three adopt the zinc blende crystal structure and are more highly conducting than gallium nitride.
Table \(2\): Physical properties of 13-15 compound semiconductors. a Values given for 300 K. b Dependent on photon energy; values given for 1.5 eV incident photons. c Dependent on temperature; values given for 300 K.
Property GaN GaP GaAs GaSb
Melting point (°C) > 1250 (dec) 1350 1240 712
Density (g/cm3) ca. 6.1 4.138 5.3176 5.6137
Crystal structure Würtzite zinc blende zinc blende zinc blende
Cell dimen. (Å)a a = 3.187, c = 5.186 a = 5.4505 a = 5.6532 a = 6.0959
Refractive indexb 2.35 3.178 3.666 4.388
k (ohm-1cm-1) 10-9 - 10-7 10-2 - 102 10-6 - 103 6 - 13
Band gap (eV)c 3.44 2.24 1.424 0.71
Gallium arsenide versus silicon
Gallium arsenide is a compound semiconductor with a combination of physical properties that has made it an attractive candidate for many electronic applications. From a comparison of various physical and electronic properties of GaAs with those of Si (Table \(3\)) the advantages of GaAs over Si can be readily ascertained. Unfortunately, the many desirable properties of gallium arsenide are offset to a great extent by a number of undesirable properties, which have limited the applications of GaAs based devices to date.
Table \(3\): Comparison of physical and semiconductor properties of GaAs and Si.
Properties GaAs Si
Formula weight 144.63 28.09
Crystal structure zinc blende diamond
Lattice constant 5.6532 5.43095
Melting point (°C) 1238 1415
Density (g/cm3) 5.32 2.328
Thermal conductivity (W/cm.K) 0.46 1.5
Band gap (eV) at 300 K 1.424 1.12
Intrinsic carrier conc. (cm-3) 1.79 x 106 1.45 x 1010
Intrinsic resistivity (ohm.cm) 108 2.3 x 105
Breakdown field (V/cm) 4 x 105 3 x 105
Minority carrier lifetime (s) 10-8 2.5 x 10-3
Mobility (cm2/V.s) 8500 1500
Band gap
The band gap of GaAs is 1.42 eV; resulting in photon emission in the infra-red range. Alloying GaAs with Al to give AlxGa1-xAs can extend the band gap into the visible red range. Unlike Si, the band gap of GaAs is direct, i.e., the transition between the valence band maximum and conduction band minimum involves no momentum change and hence does not require a collaborative particle interaction to occur. Photon generation by inter-band radiative recombination is therefore possible in GaAs. Whereas in Si, with an indirect band-gap, this process is too inefficient to be of use. The ability to convert electrical energy into light forms the basis of the use of GaAs, and its alloys, in optoelectronics; for example in light emitting diodes (LEDs), solid state lasers (light amplification by the stimulated emission of radiation).
A significant drawback of small band gap semiconductors, such as Si, is that electrons may be thermally promoted from the valence band to the conduction band. Thus, with increasing temperature the thermal generation of carriers eventually becomes dominant over the intentionally doped level of carriers. The wider band gap of GaAs gives it the ability to remain 'intentionally' semiconducting at higher temperatures; GaAs devices are generally more stable to high temperatures than a similar Si devices.
Carrier density
The low intrinsic carrier density of GaAs in a pure (undoped) form indicates that GaAs is intrinsically a very poor conductor and is commonly referred to as being semi-insulating. This property is usually altered by adding dopants of either the p- (positive) or n- (negative) type. This semi-insulating property allows many active devices to be grown on a single substrate, where the semi-insulating GaAs provides the electrical isolation of each device; an important feature in the miniaturization of electronic circuitry, i.e., VLSI (very-large-scale-integration) involving over 100,000 components per chip (one chip is typically between 1 and 10 mm square).
Electron mobility
The higher electron mobility in GaAs than in Si potentially means that in devices where electron transit time is the critical performance parameter, GaAs devices will operate with higher response times than equivalent Si devices. However, the fact that hole mobility is similar for both GaAs and Si means that devices relying on cooperative electron and hole movement, or hole movement alone, show no improvement in response time when GaAs based.
Crystal growth
The bulk crystal growth of GaAs presents a problem of stoichiometric control due the loss, by evaporation, of arsenic both in the melt and the growing crystal (> ca. 600 °C). Melt growth techniques are, therefore, designed to enable an overpressure of arsenic above the melt to be maintained, thus preventing evaporative losses. The loss of arsenic also negates diffusion techniques commonly used for wafer doping in Si technology; since the diffusion temperatures required exceed that of arsenic loss.
Crystal Stress
The thermal gradient and, hence, stress generated in melt grown crystals have limited the maximum diameter of GaAs wafers (currently 6" diameter compared to over 12" for Si), because with increased wafer diameters the thermal stress generated dislocation (crystal imperfections) densities eventually becomes unacceptable for device applications.
Physical strength
Gallium arsenide single crystals are very brittle, requiring that considerably thicker substrates than those employed for Si devices.
Native oxide
Gallium arsenide's native oxide is found to be a mixture of non-stoichiometric gallium and arsenic oxides and elemental arsenic. Thus, the electronic band structure is found to be severely disrupted causing a breakdown in 'normal' semiconductor behavior on the GaAs surface. As a consequence, the GaAs MISFET (metal-insulator-semiconductor-field-effect-transistor) equivalent to the technologically important Si based MOSFET (metal-oxide-semiconductor-field-effect-transistor) is, therefore, presently unavailable.
The passivation of the surface of GaAs is therefore a key issue when endeavoring to utilize the FET technology using GaAs. Passivation in this discussion means the reduction in mid-gap band states which destroy the semiconducting properties of the material. Additionally, this also means the production of a chemically inert coating which prevents the formation of additional reactive states, which can effect the properties of the device.
Bibliography
• S. K. Ghandhi, VLSI Fabrication Principles: Silicon and Gallium Arsenide. Wiley-Interscience, New York, (1994).
• Properties of Gallium Arsenide. Ed. M. R. Brozel and G. E. Stillman. 3rd Ed. Institution of Electrical Engineers, London (1996). | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/06%3A_Group_13/6.11%3A_Properties_of_Gallium_Arsenide.txt |
Introduction
The synthesis and purification of bulk polycrystalline semiconductor material represents the first step towards the commercial fabrication of an electronic device. This polycrystalline material is then used as the raw material for the formation of single crystal material that is processed to semiconductor wafers.
In contrast to electronic grade silicon (EGS), whose use is a minor fraction of the global production of elemental silicon, gallium arsenide (GaAs) is produced exclusively for use in the semiconductor industry. However, arsenic and its compounds have significant commercial applications. The main use of elemental arsenic is in alloys of Pb, and to a lesser extent Cu, while arsenic compounds are widely used in pesticides and wood preservatives and the production of bottle glass. Thus, the electronics industry represents a minor user of arsenic. In contrast, although gallium has minor uses as a high-temperature liquid seal, manometric fluids and heat transfer media, and for low temperature solders, its main use is in semiconductor technology.
Isolation and purification of gallium metal
At 19 ppm gallium (L. Gallia, France) is about as abundant as nitrogen, lithium and lead; it is twice as abundant as boron (9 ppm), but is more difficult to extract due to the lack of any major gallium-containing ore. Gallium always occurs in association either with zinc or germanium, its neighbors in the periodic table, or with aluminum in the same group. Thus, the highest concentrations (0.1-1%) are in the rare mineral germanite (a complex sulfide of Zn, Cu, Ge, and As), while concentrations in sphalerite (ZnS), diaspore [AlO(OH)], bauxite, or coal, are a hundred-fold less. Industrially, gallium was originally recovered from the flue dust emitted during sulfide roasting or coal burning (up to 1.5% Ga), however, it is now obtained as side product of vast aluminum industry and in particular from the Bayer process for obtaining alumina from bauxite.
The Bayer process involves dissolution of bauxite, AlOxOH3-2x, in aqueous NaOH, separation of insoluble impurities, partial precipitation of the trihydrate, Al(OH)3, and calcination at 1,200 °C. During processing the alkaline solution is gradually enriched in gallium from an initial weight ratio Ga/Al of about 1/5000 to about 1/300. Electrolysis of these extracts with a Hg cathode results in further concentration, and the solution of sodium gallate thus formed is then electrolyzed with a stainless steel cathode to give Ga metal. Since bauxite contains 0.003-0.01% gallium, complete recovery would yield some 500-1000 tons per annum, however present consumption is only 0.1% of this about 10 tons per annum.
A typical analysis of the 98-99% pure gallium obtained as a side product from the Bayer process is shown in Table $1$. This material is further purified to 99.99% by chemical treatment with acids and O2 at high temperatures followed by crystallization. This chemical process results in the reduction of the majority of metal impurities at the ppm level, see Table $1$. Purification to seven nines 99.9999% is possible through zone refining, however, since the equilibrium distribution coefficient of the residual impurities k0 ≈ 1, multiple passes are required, typically > 500. The low melting point of gallium ensures that contamination from the container wall (which is significant in silicon zone refining) is minimized. In order to facilitate the multiple zone refining in a suitable time, a simple modification of zone refining is employed shown in Figure $1$. The gallium is contained in a plastic tube wrapped around a rotating cylinder that is half immersed in a cooling bath. A heater is positioned above the gallium plastic coil. Thus, establishing a series of molten zones that pass upon rotation of the drum by one helical segment per revolution. In this manner, 500 passes may be made in relatively short time periods. The typical impurity levels of gallium zone refined in this manner are given in Table $1$.
Table $1$: Typical analysis of gallium obtained as a side product from the Bayer process.
Element Bayer process (ppm) After acid/base leaching (ppm) 500 zone passes (ppm)
aluminum 100-1,000 7 < 1
calcium 10-100 not detected not detected
copper 100-1,000 2 < 1
iron 100-1,000 7 < 1
lead < 2000 30 not detected
magnesium 10-100 1 not detected
mercury 10-100 not detected not detected
nickel 10-100 not detected not detected
silicon 10-100 ≈ 1 not detected
tin 10-100 ≈ 1 not detected
titanium 10-100 1 < 1
zinc 30,000 ≈ 1 not detected
Isolation and purification of elemental arsenic
Elemental arsenic (L. arsenicum, yellow orpiment) exists in two forms: yellow (cubic, As4) and gray or metallic (rhombohedral). At a natural abundance of 1.8 ppm arsenic is relatively rare, however, this is offset by its presence in a number of common minerals and the relative ease of isolation. Arsenic containing minerals are grouped into three main classes: the sulfides realgar (As4S4) and orpiment (As2S3), the oxide arsenolite (As2O3), and the arsenides and sulfaresenides of the iron, cobalt, and nickel. Minerals in this latter class include: loellinginite (FeAs2), safforlite (CoAs), niccolite (NiAs), rammelsbergite (NiAs2), ansenopyrite or mispickel (FeAsS), cobaltite (CoAsS), enargite (Cu3AsS4), gerdsorfite (NiAsS), and the quarturnary sulfide glaucodot [(Co,Fe)AsS]. Table $2$ shows the typical impurities in arsenopyrite.
Table $2$: Typical impurities in arsenopyrite.
Element Concentration (ppm) Element Concentration (ppm)
silver 90 nickel < 3,000
gold 8 lead 50
cobalt 30,000 platinum 0.4
copper 200 rhenium 50
germanium 30 selenium 50
manganese 3,000 vanadium 300
molybdenum 60 zinc 400
Arsenic is obtained commercially by smelting either FeAs2 or FeAsS at 650-700 °C in the absence of air and condensing the sublimed element (Tsub = 613 °C), (6.12.1).
$\text{FeAsS} \xrightarrow{\text{650-700 °C}} \text{FeS + As(vapor)} \xrightarrow{\text{<613 °C}} \text{As(solid)}$
The arsenic thus obtained is combined with lead and then sublimed (Tsub = 614 °C) which binds any sulfur impurities more strongly than arsenic. Any residual arsenic that remains trapped in the iron sulfide is separated by forming the oxide (As2O3) by roasting the sulfide in air. The oxide is sublimed into the flue system during roasting from where it is collected and reduced with charcoal at 700-800 °C to give elemental arsenic. Semiconductor grade arsenic (> 99.9999%) is formed by zone refining.
Synthesis and purification of gallium arsenide.
Gallium arsenide can be prepared by the direct reaction of the elements, (6.12.2). However, while conceptually simple the synthesis of GaAs is complicated by the different vapor pressures of the reagents and the highly exothermic nature of the reaction. Furthermore, since the synthesis of GaAs at atmospheric pressure is accompanied by its simultaneous decomposes due to the loss by sublimation, of arsenic, the synthesis must be carried out under an overpressure of arsenic in order to maintain a stoichiometric composition of the synthesized GaAs.
$\text{Ga(liquid) + As(vapor)} \xrightarrow{ \text{>1240 °C}} \text{GaAs(solid)}$
In order to overcome the problems associated with arsenic loss, the reaction is usually carried out in a sealed reaction tube. However, if a stoichiometric quantity of arsenic is used in the reaction a constant temperature of 1238 °C must be employed in order to maintain the desired arsenic overpressure of 1 atm. Practically, it is easier to use a large excess of arsenic heated to a lower temperature. In this situation the pressure in the tube is approximately equal to the equilibrium vapor pressure of the volatile component (arsenic) at the lower temperature. Thus, an over pressure of 1 atm arsenic may be maintained if within a sealed tube elemental arsenic is heated to 600-620 °C while the GaAs is maintained at 1240-1250 °C.
Figure $2$ shows the sealed tube configuration that is typically used for the synthesis of GaAs. The tube is heated within a two-zone furnace. The boats holding the reactants are usually made of quartz, however, graphite is also used since the latter has a closer thermal expansion match to the GaAs product. If higher purity is required then pyrolytic boron nitride (PBN) is used. One of the boats is loaded with pure gallium the other with arsenic. A plug of quartz wool may be placed between the boats to act as a diffuser. The tube is then evacuated and sealed. Once brought to the correct reaction temperatures (Figure $2$), the arsenic vapor is transported to the gallium, and they react to form GaAs in a controlled manner. Table $3$ gives the typical impurity concentrations found in polycrystalline GaAs.
Table $3$: Impurity concentrations found in polycrystalline GaAs.
Element Concentration (ppm) Element Concentration (ppm)
boron 0.1 silicon 0.02
carbon 0.7 phosphorus 0.1
nitrogen 0.1 sulfur 0.01
oxygen 0.5 chlorine 0.08
fluorine 0.2 nickel 0.04
magnesium 0.02 copper 0.01
aluminum 0.02 zinc 0.05
Polycrystalline GaAs, formed in from the direct reaction of the elements is often used as the starting material for single crystal growth via Bridgeman or Czochralski crystal growth. It is also possible to prepare single crystals of GaAs directly from the elements using in-situ, or direct, compounding within a high-pressure liquid encapsulated Czochralski (HPLEC) technique.
Growth of gallium arsenide crystals
When considering the synthesis of Group 13-15 compounds for electronic applications, the very nature of semiconductor behavior demands the use of high purity single crystal materials. The polycrystalline materials synthesized above are, therefore, of little use for 13-15 semiconductors but may, however, serve as the starting material for melt grown single crystals. For GaAs, undoubtedly the most important 13-15 (III - V) semiconductor, melt grown single crystals are achieved by one of two techniques: the Bridgman technique, and the Czochralski technique.
Bridgman growth
The Bridgman technique requires a two-zone furnace, of the type shown in Figure $3$. The left hand zone is maintained at a temperature of ca. 610 °C, allowing sufficient overpressure of arsenic within the sealed system to prevent arsenic loss from the gallium arsenide. The right hand side of the furnace contains the polycrystalline GaAs raw material held at a temperature just above its melting point (ca. 1240 °C). As the furnace moves from left to right, the melt cools and solidifies.
If a seed crystal is placed at the left hand side of the melt (at a point where the temperature gradient is such that only the end melts), a specific orientation of single crystal may be propagated at the liquid-solid interface eventually to produce a single crystal.
Czochralski growth
The Czochralski technique, which is the most commonly used technique in industry, is shown in Figure $4$. The process relies on the controlled withdrawal of a seed crystal from a liquid melt. As the seed is lowered into the melt, partial melting of the tip occurs creating the liquid solid interface required for crystal growth. As the seed is withdrawn, solidification occurs and the seed orientation is propagated into the grown material. The variable parameters of rate of withdrawal and rotation rate can control crystal diameter and purity. As shown in Figure $4$ the GaAs melt is capped by boron trioxide (B2O3). The capping layer, which is inert to GaAs, prevents arsenic loss when the pressure on the surface is above atmospheric pressure. The growth of GaAs by this technique is thus termed liquid encapsulated Czochralski (LEC) growth.
While the Bridgman technique is largely favored for GaAs growth, larger diameter wafers can be obtained by the Czochralski method. Both of these melt techniques produce materials heavily contaminated by the crucible, making them suitable almost exclusively as substrate material. Another disadvantage of these techniques is the production of defects in the material caused by the melt process.
Bibliography
• S. K. Ghandhi, VLSI Fabrication Principles: Silicon and Gallium Arsenide. Wiley-Interscience, New York, (1994).
• J. Krauskopf, J. D. Meyer, B. Wiedemann, M. Waldschmidt, K. Bethge, G. Wolf, and W. Schültze, 5th Conference on Semi-insulating III-V Materials, Malmo, Sweden, 1988, Eds. G. Grossman and L. Ledebo, Adam-Hilger, New York (1988).
• W. G. Pfann, Zone Melting, John Wiley & Sons, New York (1966).
• R. E. Williams, Gallium Arsenide Processing Techniques. Artech House (1984).
• Properties of Gallium Arsenide. Ed. M. R. Brozel and G. E. Stillman. 3rd Ed. Institution of Electrical Engineers, London (1996). | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/06%3A_Group_13/6.12%3A_Electronic_Grade_Gallium_Arsenide.txt |
The only stable chalcogenides of aluminum are Al2S3 (white), Al2Se3 (grey), and Al2Te3 (dark grey). They are each prepared by the direct reaction of the elements (100 °C) and hydrolyze rapidly in aqueous solution, (6.13.1). All the chalcogenides have a hexagonal ZnS structure in which 2/3 of the metal sites are occupied.
$\text{Al}_3\text{E}_3 \text{ + 6 H}_2\text{O} \rightarrow \text{2 Al(OH)}_3 \text{ + 3 H}_2\text{E}$
The chalcogenides of gallium and indium are more numerous than those of aluminum, and are listed in Table $1$ and Table $2$ along with selected physical properties.
Table $1$: Stoichiometries, structures and selected physical properties of the crystalline chalcogenides of gallium. a dir = direct, ind = indirect, opt = optical.
Compound Structural type Crystallographic system Cell parameters (Å, °) Band Gap (eV)a
GaS Hexagonal a = 3.587, c = 15.492 3.05 (dir.), 2.593 (ind.)
GaS ZnS or NaCl Cubic a = 5.5 4.0 (opt.)
β-GaSe GaS Hexagonal a = 3.742, c = 15.919 2.103 (dir.), 2.127 (ind.)
γ-GaSe GaS Rhombohedral a = 3.755, c = 23.92
δ-GaSe GaS Hexagonal a = 3.755, c = 31.99
β-GaTe GaS Hexagonal a = 4.06, c = 16.96
GaTe GaS Monoclinic a = 17.44, b = 4.077, c = 10.456, β = 104.4 1.799 (dir.)
α-Ga2S3 Wurtzite Cubic a = 5.181
α-Ga2S3 Wurtzite Monoclinic a = 12.637, b = 6.41, c = 7.03, β = 131.08 3.438 (opt.)
β-Ga2S3 Defect wurtzite Hexagonal a = 3.685, c = 6.028 2.5 - 2.7 (opt.)
α-Ga2Se3 Sphalerite Cubic a = 5.429 2.1 (dir.), 2.04 (ind.)
α-Ga2Te3 Sphalerite Cubic a = 5.886 1.22 (opt.)
Table $2$: Stoichiometries, structures and selected physical properties of the crystalline chalcogenides of indium. a dir = direct, ind = indirect, opt = optical. b High pressure phase.
Compound Structural type Crystallographic system Cell parameters (Å, °) Band gap (eV)a
β-InS GaS Orthorhombic a = 3.944, b = 4.447, c = 10.648 2.58 (dir.), 2.067 (ind.)
InSb Hg2Cl2 Tetragonal
InSe GaS Rhombohedral a = 4.00, c = 25.32 1.3525 (dir.), 1.32 (ind.)
β-InSe GaS Hexagonal a = 4.05, c = 16.93
InTe TlSe Tetragonal a = 8.437, c = 7.139 Metallic
InTeb NaCl Cubic a = 6.18
α-In2S3 γ-Al2O3 Cubic a = 5.36
β-In2S3 Spinel Tetragonal a = 7.618, c = 32.33 2.03 (dir.), 1.1 (ind.)
α-In2Se3 Defect wurtzite Hexagonal a = 16.00, c = 19.24
β-In2Se3 Defect wurtzite Rhombohedral a = 4.025, c = 19.222 1.2 - 1.5 (ind.)
α-In2Te3 Sphalerite Cubic a = 6.158 0.92 - 1.15 (opt.)
In6S7 Monoclinic a = 9.090, b = 3.887, c = 17.705, β = 108.20 0.89 (dir.), 0.7 (ind.)
In6Se7 In6S7 Monoclinic a = 9.430, b = 4.063, c = 18.378, β = 109.34 0.86 (dir.), 0.34 (ind.)
In4Se3 Orthorhombic a = 15.297, b = 12.308, c = 4.081 0.64 (dir.)
In4Te3 In4Se3 Orthorhombic a = 15.630, b = 12.756, c = 4.441 0.48 (dir.)
The hexagonal β-form of Ga2S3 is isostructural with the aluminum analogue; however, while the α-phase was proposed to be hexagonal it was later shown to be monoclinic. A cubic α-phase has been reported. Cubic Sphalerite structures are found for Ga2Se3, Ga2Te3, and In2Te3, in which the structure is based on a cubic close packing of the chalcogenides and the metal atoms occupying 1/3 of the tetrahedral sites. These structures are all formed with rapid crystallization; slow crystallization and/or thermal annealing leads to ordering and the formation of more complex structures. The indium sulfides, and selenides derivatives are spinel (γ-Al2O3), and defect Würtzite, respectively.
Unlike the chalcogenides of aluminum, those of gallium and indium also form subvalent compounds, i.e., those in which the metal is formally of an oxidation state less than +3. Of these subvalent chalcogenides the (formally) divalent materials are of the most interest. The thermodynamically stable phase of GaS has a hexagonal layer structure (Figure $1$) with Ga-Ga bonds (2.48 Å). The compound can, therefore, be considered as an example of Ga(II). Each Ga is coordinated by three sulfur atoms and one gallium, and the sequence of layers along the z-axis is ...S-Ga-Ga-S...S-Ga-Ga-S....
The structures of β-GaSe, and β-InSe are similar to hexagonal GaS. The layered structure of GaTe is similar in that it consists of ...TeGaGaTe... layers, but is monoclinic, while InS is found in both a (high pressure) tetragonal phase (Figure $2$a) as well as an orthorhombic phase (Figure $2$b). By contrast to these M-M bonded layered compounds InTe (Figure $3$) has a structure formalized as In(I)[In(III)Te2]; each In(III) is tetrahedrally coordinated to four Te and these tetrahedra are linked via shared edges; the In(I) centers lying between these chains.
Further sub-chalcogenides are known for indium, e.g.; In4Se3, which contains [In(III)3Se2]5+ groups (Figure $4$). While the formally In(I) molecule In2S has been detected in the gas phase, it is actually a mixture of In and InS in the solid state.
Bibliography
• W. J. Duffin and J. H. C. Hogg, Acta Crystallogr., 1966, 20, 566.
• J. Goodyear and G. Steigman, Acta Crystallogr. 1963, 16, 946.
• H. Hahn and G. Frank, Z. Anorg. Allgem. Chem., 1955, 278, 340.
• S. Kabalkina, V. G. Losev, and N. M. Gasanly, Solid State Commun., 1982, 44, 1383.
• A. Keys, S. G. Bott, and A. R. Barron, Chem. Mater., 1999, 11, 3578.
• A. N. MacInnes, M. B. Power, and A. R. Barron, Chem. Mater., 1992, 4, 11.
• A. N. MacInnes, W. M. Cleaver, A. R. Barron, M. B. Power, and A. F. Hepp, Adv. Mater. Optics. Electron., 1992,1, 229.
• K. Schubert, E. Dörre, and E. Günzel, Naturwissenschaften, 1954, 41, 488. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/06%3A_Group_13/6.13%3A_Chalcogenides_of_Aluminum_Gallium_and_Indium.txt |
Trihalides, MX3
As shown in Table $1$ all the combinations of Group 13 element (M) and halogen (X) exist for the trihalides (MX3), except thallium(III) iodide. It should be noted that while there is a compound with the general formula TlI3, it is actually a thallium(I) compound of I3-.
Table $1$: Selected physical properties of the Group 13 trihalides, MX3.
Element Mp (°C) Bp (°C)
BF3 -126.8 -100.3
BCl3 -107.3 12.6
BBr3 -46.3 91.3
BI3 49.9 210
AlF3 1291 -
AlCl3 192.4 (anhydrous), 0.0 (hexahydrate) 120 (hexahydrate)
AlBr3 97.8 265
AlI3 189.4 (anhydrous) 185 dec. (hexahydrate) 300 subl.
GaF3 800 1000
GaCl3 77.9 201
GaBr3 121.5 278.8
GaI3 212 345
InF3 1172 -
InCl3 586 800
InBr3 220 -
InI3 210 subl. -
TlF3 300 dec. -
TlCl3 40 dec. -
TlBr3 40 dec. -
The trihalides of boron are all monomers with a coordination number of 3 (Table $2$), as evidence from their low melting points (Table $1$). In contrast, the fluorides and chlorides of the heavier Group 13 elements (except GaCl3) are generally ionic or have a high ionic character, with a coordination number of 6 (Table $2$, Figure $1$ and Figure $2$). The bromides and iodides (except InBr3) are generally dimeric with a coordination number of 4 (Table $2$) and have molecular structures involving halide bridging ligands (Figure $3$ and Table $3$). AlCl3 is unusual in that in the solid state it has an ionic structure, but it is readily sublimed, and in the vapor phase (and liquid phase) it has a dimeric structure (Figure $3$).
Table $2$: The Group 13 element coordination numbers for the trihalides, MX3.
Element Fluoride Chloride Bromide Iodide
B 3 3 3 3
Al 6 6 (4) 4 4
Ga 6 4 4 4
In 6 6 6 4
Tl 6 6 4 -
Table $3$: Selected bond lengths and angles for dimeric M2X6 compounds. aXt and Xb indicate terminal and bridging halides, respectfully.
Compound M-Xt (Å)a M-Xb (Å)a Xt-M-Xt (°)a Xb-M-Xb (°)a M-X-M (°)
Al2Br6 2.21 2.33 115 93 87
In2I6 2.64 2.84 125.1 93.7 86.3
Synthesis
Boron trifluoride (BF3) is manufactured commercially by the reaction of boron oxides with hydrogen fluoride, (6.14.1). The HF is produced in-situ from sulfuric acid and fluorite (CaF2). On smaller scales, BF3 is prepared by the thermal decomposition of diazonium salts, (6.14.2).
$\text{B}_2\text{O}_3 \text{ + 6 HF} \rightarrow \text{2 BF}_3 \text{ + 3 H}_2\text{O}$
$\text{PhN}_2\text{[BF}_4\text{]} \rightarrow \text{PhF + BF}_3 \text{ + N}_2$
Boron trichloride is also made from boron oxide, but in the presence of carbon, (6.14.3).
$\text{B}_2\text{O}_3 \text{ + 3 C + 3 Cl}_2 \rightarrow \text{2 BCl}_3 \text{ + 3 CO}$
Many of the trihalides are readily prepared by either the direct reaction of the metal with the appropriate halogen, (6.14.4) - (6.14.6), or the acid, (6.14.7) and (6.14.8). Thallium tribromide can be prepared in CH3CN by treating a solution of the monobromide with bromine gas, (6.14.9).
$\text{2 Al + 3 Cl}_2 \rightarrow \text{2 AlCl}_3$
$\text{2 Al + 3 Br}_2 \rightarrow \text{2 AlBr}_3$
$\text{2 Al + 3 I}_2 \rightarrow \text{2 AlCl}_3$
$\text{2 Al + 6 HCl} \rightarrow \text{2 AlCl}_3 \text{ + 3 H}_2$
$\text{2 Al + 6 HBr} \rightarrow \text{2 AlBr}_3 \text{ + 3 H}_2$
$\text{TlBr + Br}_2 \rightarrow \text{TlBr}_2$
Reactivity
The reaction chemistry of the Group 13 trihalides tends to fall into two categories:
• Lewis acid-base complex formation.
• Hydrolysis.
There are, however, a number of reactions involving halide exchange reactions. Aluminum tribromide reacts with carbon tetrachloride at 100 °C to form carbon tetrabromide, (6.14.10), and with phosgene yields carbonyl bromide and aluminum chlorobromide, (6.14.11).
$\text{4 AlBr}_3 \text{ + 3 CCl}_4 \rightarrow \text{4 AlCl}_3 \text{ + 3 CBr}_4$
$\text{AlBr}_3 \text{ + COCl}_2 \rightarrow \text{COBr}_2 \text{ + AlCl}_2\text{Br}$
Group 13 halides are used as synthons for their organometallic derivatives, (6.14.12) and (6.14.13).
$\text{MX}_3 \text{ + 3 RMgX} \rightarrow \text{MR}_3 \text{ + MgX}_2$
$\text{2 MR}_3 \text{ + MX}_3 \rightarrow \text{3 MXR}_2$
Lewis acid-base complexes
All of the trihalides are strong Lewis acids, and as such react with Lewis base compounds to form Lewis acid-base complexes, (6.14.12). The extent of the equilibrium is dependant on the Lewis acidity of the trihalide and the basicity of the Lewis base. For example, with BCl3, oxygen donor ligands result in approximately 50:50 ratio of BCl3 and BCl3L, while for nitrogen donor ligands the equilibrium is shifted to the formation of the complex.
$\text{MX}_3 \text{ + L} \rightleftharpoons \text{X}_3\text{M-L}$
The general structure of the Lewis acid-base complexes is such that the Group 13 element is close to tetrahedral (Figure $4$). However, for aluminum and the heavier Group 13 elements, more than one ligand can coordinate (Figure $5$) up to a maximum of six.
It should be noted that the dimeric form of MX3 (Figure $3$) can be thought of as mutual Lewis acid-base complexes, in which a Lewis basic lone pair of a halide on one MX3 unit donates to the Lewis acidic metal on another MX3 unit.
Hydrolysis
Generally the fluorides are insoluble in water while the heavier halides are more soluble. However, BF3, BCl3, and BBr3 all decompose in the presence of water, (6.14.13). In the case of the fluoride, the HF formed reacts with BF3 to form fluoroboric acid, (6.14.14). However, there is also a minor equilibrium (2-3%) resulting in the formation of the BF3 complex of OH- and H3O+, (6.14.15).
$\text{BX}_3\text{ + H}_2\text{O} \rightarrow \text{B(OH)}_3\text{ + 3 HX}$
$\text{HF + BF}_3 \rightarrow \text{HBF}_4$
$\text{BF}_4 \text{ + H}_3\text{O} \rightleftharpoons \text{F}_3\text{B-OH}_2 \rightleftharpoons \text{[BF}_3\text{OH][BF}_3\text{(OH}_3\text{)]}$
While the boron compounds (and AlBr3) decompose even in moist air, AlCl3 reacts more slowly to make aluminum chlorohydrate (ACH) which has the general formula AlnCl3n-m(OH)m. While ACH has been proposed to exist as a number of cluster species, it is actually a range of nanoparticles.
ACH is also known as polyaluminum chloride (PAC). The latter name is often used in water purification, where ACH is preferred over alum derivatives (Al2(SO4)3). The combination of ACH and a high molecular weight quaternized ammonium polymer (e.g., dially dimethyl ammonium chloride (DADMAC)), has been known as an effective combination as a flocculant in water treatment process to remove dissolved organic matter and colloidal particles present in suspension.
Aluminum chlorohydrate (ACH) and aluminum-zirconium compounds, are frequently used as the active ingredient in antiperspirants. The mode of action of most aluminum-based compounds involves the dramatic change in the particle size from nano to micro as a function of pH and electrolyte changes on the skin (as compared to the antiperspirant stick or suspension) and hence forming a gel plug in the duct of the sweat gland. The plugs prevent the gland from excreting liquid and are removed over time by the natural sloughing of the skin. A further mechanism of action involves the absorption of 3-methyl-2-hexenoic acid (Figure $6$). Human perspiration is odorless until bacteria ferment it. Bacteria thrive in hot, humid environments such as the underarm. When adult armpits are washed with alkaline pH soaps, the skin loses its acid mantel (pH = 4.5 - 6), raising the skin pH and disrupting the skin barrier. The bacteria thrive in the basic environment, and feed on the sweat from the apocrine glands and on dead skin and hair cells, releasing 3-methyl-2-hexenoic acid, which is the primary cause of body odor. As with all carboxylic acids, 3-methyl-2-hexenoic acid, reacts in a facile manner with the surface of the alumina nanoparticles. Aluminum chloride salts also have a slight astringent effect on the pores; causing them to contract, further preventing sweat from reaching the surface of the skin.
Boron trihalides: a special case
The three lighter boron trihalides, BX3 (X = F, Cl, Br) form stable adducts with common Lewis bases. Their relative Lewis acidities can be evaluated in terms of the relative exothermicity of the adduct-forming reaction:
$\text{BF}_3 \text{ < BCl}_3 \text{ < BBr}_3\text{(strongest Lewis acid)}$
This trend is opposite to that expected based upon the electronegativity of the halogens. The best explanation of this trend takes into account the extent of π-donation that occurs between the filled lone pair orbital on the halogens and the empty p-orbital on the planar boron (Figure $7$). As such, the greater the π-bonding the more stable the planar BX3 configuration as opposed to the pyramidalization of the BX3 moiety upon formation of a Lewis acid-base complex, (6.4.12).
The criteria for evaluating the relative strength of π-bonding are not clear, however, one suggestion is that the F atom is small compared to the Cl atom, and the lone pair electron in pz of F is readily and easily donated and overlapped to empty pz orbital of boron (Figure $7$a). In contrast, the overlap for the large (diffuse) p-orbitals on the chlorine is poorer (Figure $7$b). As a result, the π-donation of F is greater than that of Cl. Interestingly, as may be seen from Table $4$, any difference in B-X bond length does not follow the expected trend associated with shortening of the B-X bond with stronger π-bonding. In fact the B-Br distance is the most shortened as compared to that expected from the covalent radii (Table $4$).
Table $4$: The B-X bond distances in the boron trihalides, BX3, as compared to the sum of the covalent radii. aCovalent radius of B = 0.84(3) Å.
Compound B-X (Å) X covalent radius (Å) Sum of covalent radii (Å)a Δ (Å)
BF3 1.313 0.57(3) 1.41 0.097
BCl3 1.75 1.02(4) 1.86 0.11
BBr3 1.898 1.20(3) 2.04 0.142
BI3 2.125 1.39(3) 2.23 0.105
At the simplest level the requirements for bonding to occur based upon the molecular or atomic orbital are:
• Directional relationship of the orbitals.
• Relative symmetry of the orbitals.
• Relative energy of the orbitals.
• Extent of orbital overlap
In the case of the boron trihalides, the direction (parallel) and symmetry (p-orbitals) are the same, and the only significant difference will be the relative energy of the donor orbitals (i.e., the lone pair on the halogen) and the extent of the overlap. The latter will be dependant on the B-X bond length (the shorter the bond the greater the potential overlap) and the diffusion of the orbitals (the less diffuse the orbitals the better the overlap). Both of these factors will benefit B-F over B-Cl and B-Br. Thus, the extent of potential overlap would follow the order: (6.14.16). Despite these considerations, it is still unclear of the exact details of the rationalization of the low Lewis basicity of BF3 as compared to BCl3 and BBr3.
Anionic halides
The trihalides all form Lewis acid-base complexes with halide anions, (6.14.17), and as such salts of BF4-, AlCl4-, GaCl4-, and InCl4- are common.
$\text{MX}_3 \text{ + X}^- \rightarrow \text{MX}_4^-$
In the case of gallium the Ga2Cl7- anion (Figure $8$) is formed from the equilibrium:
$\text{2 GaCl}_4^- \rightleftharpoons \text{Ga}_2\text{Cl}_7^- \text{ + Cl}^-$
As a consequence of its larger size indium forms a wide range of anionic halides addition compounds with trigonal bipyramidal, square pyramidal, and octahedral coordination geometries. For example, salts of InCl52-, InBr52-, InF63-, InCl63- and InBr63- have all been made. The InCl52- ion has been found to be square pyramidal in the salt [NEt4]2InCl5, but is trigonal bipyramidal in the acetonitrile solvate of [Ph4P]2InCl5. The oligomeric anionic halides In2X7- and In2X93- (X = Cl and Br) contain binuclear anions with tetrahedral and octahedrally coordinated indium atoms, respectively (Figure $8$).
Low valent halides
Oxidation state +2
Boron forms a series of low oxidation halides containing B-B bonds a formal oxidation state of +2. Passing an electric discharge through BCl3 using mercury electrodes results in the synthesis of B2Cl4, (6.14.19). An alternative route is by the co-condensation of copper as a reducing agent with BCl3, (6.14.20).
$\text{2 BCl}_3\text{ + 2 Hg} \rightarrow \text{B}_2\text{Cl}_4 \text{ + Hg}_2\text{Cl}_2$
$\text{2 BCl}_3\text{ + 2 Cu} \rightarrow \text{B}_2\text{Cl}_4 \text{ + 2 CuCl}$
B2F4 has a planar structure (Figure $9$) with D2h symmetry, while B2Cl4 has the same basic structure it has a staggered geometry (Figure $10$). The energy for bond rotation about the B-B bond is very low (5 kJ/mol) that can be compared to ethane (12.5 kJ/mol). The bromide, B2Br4, is also observed to be staggered in the solid state. The staggered conformation is favorable on steric grounds, however, for B2F4 the planar geometry is stabilized by the smaller size of the halide, and more importantly the presence of strong delocalized π-bonding.
Oxidation state +1
Boron forms a number of halides with cluster structures, BnCln where n = 4 (Figure $11$), 8, 9, 10, 11, and 12. Each compound is made by the decomposition of B2Cl4. For gallium, none of the monohalides are stable at room temperature, but GaCl and GaBr have been produced in the gas form from the reaction of HX and molten gallium. The stability of thallium(I) as compared to thallium(III) results in the monohalides, TlCl, TlBr, and TlI being stable. Each compound is insoluble in water, and photosensitive.
Intermediate halides
The dihalides (MX2) of gallium, indium, and thallium do not actually contain the metal in the +2 oxidation state. Instead they are actually mixed valence compound, i.e., M+[MX4]-. The dihalides of gallium are unstable in the presence of water disproportionating to gallium metal and gallium(III) entities. They are soluble in aromatic solvents, where arene complexes have been isolated and the arene is η6-coordinated to the Ga+ ion. InBr2 and InI2 are greenish and yellow crystalline solids, respectively, which are formulated In(I)[In(III)X4]. TlCl2 and TlBr2 both are of similar formulations.
Ga2X3 (X = Br, I) and In2Br3 are formulated M(I)2[M(II)2X6]. Both anions contain a M-M bond where the metal has a formal oxidation state of +2. The Ga2Br62- anion is eclipsed like the In2Br62- anion, whereas the Ga2I62- anion is isostructural with Si2Cl6 with a staggered conformation. In2Cl3 is colorless and is formulated In(I)3[In(III)Cl6].
Ga3Cl7 contains the Ga2Cl7- ion, which has a structure similar to the dichromate, Cr2O72-, ion with two tetrahedrally coordinated gallium atoms sharing a corner (Figure). The compound can be formulated as gallium(I) heptachlorodigallate(III), Ga(I)[Ga(III)2Cl7].
In4Br7 is light sensitive (like TlCl and TlBr) decaying to InBr2 and In metal. It is a mixed salt containing the InBr4- and InBr63- anions balanced by In+ cations. It is formulated In(I)5[In(III)Br4]2[In(III)Br6]. In5Br7 is a pale yellow solid formulated as In(I)3[In(II)2Br6]Br. The In(II)2Br62- anion has an eclipsed ethane like structure with an In-In bond length of 2.70 Å. In5Cl9 is formulated In(I)3[In(III)2Cl9], with the In2Cl92- anion having two 6 coordinate indium atoms with 3 bridging chlorine atoms, face sharing bioctahedra. Finally, In7Cl9 and In7Br9 have a structure formulated as InX6[In(III)X6]X3.
Bibliography
• P. M. Boorman and D. Potts, Can. J. Chem., 1974, 52, 2016.
• A. Borovik and A. R. Barron, J. Am. Chem. Soc., 2002, 124, 3743.
• A. Borovik, S. G. Bott, and A. R. Barron, J. Am. Chem. Soc., 2001, 123, 11219.
• C. S. Branch, S. G. Bott, and A. R. Barron, J. Organomet. Chem., 2003, 666, 23.
• W. M. Brown and M. Trevino, US Patent 5,395,536 (1995).
• S. K. Dentel, CRC Critical Reviews in Environmental Control, 1991, 21, 41.
• D. E. Hassick and J. P. Miknevich, US Patent 4,800,039 (1989).
• M. D. Healy, P. E. Laibinis, P. D. Stupik and A. R. Barron, J. Chem. Soc., Chem. Commun., 1989, 359.
• K. Hedberg and R. Ryan, J. Chem. Phys., 1964, 41, 2214.
• Y. Koide and A. R. Barron, Organometallics, 1995, 14, 4026.
• G. Santiso-Quiñones and I. Krossing, Z. Anorg. Allg. Chem., 2008, 634, 704. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/06%3A_Group_13/6.14%3A_Group_13_Halides.txt |
The group was once also known as the tetrels (from Greek tetra meaning four), stemming from the earlier naming convention of this group as Group IVA. Table lists the derivation of the names of the Group 14 elements.
Table $1$: Derivation of the names of each of the Group 14 elements.
Element Symbol Name
Carbon C From the Latin carbo meaning coal
Silicon Si From the Latin silicis meaning flints
Germanium Ge From the Latin Germania for Germany
Tin Sn From the Anglo-Saxon and from the Latin stannum meaning melts easily
Lead Pb From the Anglo-Saxon, and from the Latin plumbum meaning soft metal
Discovery
Carbon
Carbon was known in prehistory in the form of soot; while charcoal was made in Roman times (by heating wood while exclude air) and diamonds were known as early as 2500 BC in China. In 1772, Antoine Lavoisier (Figure $1$) showed that diamonds were a form of carbon, when he burned samples of carbon and diamond and showed that both formed the same amount of carbon dioxide per gram of material. Carl Scheele (Figure $2$) showed that graphite was a form of carbon rather a form of lead.
A new allotrope of carbon, fullerene, was discovered in 1985 by Robert Curl, Harry Kroto, and Richard Smalley (Figure $3$) who subsequently shared the Nobel Prize in Chemistry in 1996. Fullerenes have been reveled to include nanostructured forms such as buckyballs and nanotubes. The renewed interest in new forms lead to the discovery of further exotic allotropes, including glassy carbon, and the realization that amorphous carbon is not amorphous.
Silicon
Silicon was first identified by Antoine Lavoisier (Figure $1$) in 1787 as a component of flints, and was later mistaken by Humphry Davy (Figure $4$) for a compound rather than an element. In 1824, Berzelius (Figure $5$) prepared amorphous silicon by the reaction of potassium with silicon tetrafluoride, (7.1.1).
$\text{SiCl}_4 \text{ + 4 K} \rightarrow \text{Si + 4 KCl}$
Germanium
In 1869 Dmitri Mendeleev (Figure $6$) predicted the existence of several unknown elements, including ekasilicon (Es) between silicon and tin.
In 1885 a new mineral (named argyrodite because of its high silver content) was found in a mine near Freiberg, Saxony. Clemens Winkler (Figure $7$) isolated Mendeleev’s missing element. He originally was going to name neptunium because like this element, because like ekasilicon, the planet Neptune had been preceded by mathematical prediction of its existence. However, the name neptunium had already been given to an element and so Winkler named the new metal germanium in honor of his fatherland.
Winkler was able to isolate sufficient germanium from 500 kg of ore to determine a number properties, including an atomic weight of 72.32 g/mol by analyzing pure germanium tetrachloride (GeCl4). Winkler prepared several new compounds of germanium, including the fluorides, chlorides, sulfides, germanium dioxide, and tetraethylgermane (Ge(C2H5)4). The physical data from these compounds, corresponded with Mendeleev's predictions (Table $2$).
Table $2$: Properties predicted for ekasilicon compared those determined for germanium.
Property Ekasilicon Germanium
Atomic mass 72 72.59
Density (g/cm3) 5.5 5.35
Melting point (°C) High 947
Color Gray Gray
Oxide type Refractory dioxide Refractory dioxide
Oxide density (g/cm3) 4.7 4.7
Oxide activity Feebly basic Feebly basic
Chloride boiling point (°C) Under 100 86 (GeCl4)
Chloride density (g/cm3) 1.9 1.9
Tin
Tin is one of the earliest metals known. When the addition of about 5% tim to molten copper produced an alloy (bronze) that was easier to work and much harder than copper, it revolutionized civilization. The widespread use of bronze to make tools and weapons became part of what archaeologists call the Bronze Age. The Bronze Age arrived in Egypt, Mesopotamia and the Indus Valley culture by around 3000 BC.
Lead
Lead has been commonly used for thousands of years because of its ease of extraction, and its ease of smelting. Lead beads dating back to 6400 BC have been found in Çatalhöyük in modern-day Turkey, while lead was used during the Bronze Age.
Abundance
Carbon and silicon are amongst the most abundant elements (Table $3$). Silicon is the second most abundant element (after oxygen) in the Earth’s crust, making up 28% of the crust. Carbon is the fourth most abundant chemical element in the universe after hydrogen, helium, and oxygen. In combination with oxygen in carbon dioxide, carbon is found in the Earth's atmosphere (in quantities of approximately 810 gigatonnes) and dissolved in all water bodies (approximately 36,000 gigatons). Around 1,900 gigatons are present in the biosphere. Hydrocarbons (such as coal, petroleum, and natural gas) contain carbon amounts to around 900 gigatons. Natural diamonds occur in the rock kimberlite, found in ancient volcanic "necks," or "pipes". Most diamond deposits are in Africa but there are also deposits in Canada, the Russian Arctic, Brazil, and Australia.
Table $3$: Abundance of Group 14 elements.
Element Terrestrial abundance (ppm)
C 480 (Earth’s crust), 28 (sea water), 350 (atmosphere CO2), 1.6 (atmosphere, CH4), 0.25 (atmosphere, CO)
Si 28,000 (Earth’s crust), 2 (sea water)
Ge 2 (Earth’s crust), 1 (soil), 5 x 10-7 (sea water)
Sn 2 (Earth’s crust), 1 (soil), 4 x 10-6 (sea water)
Pb 14 (Earth’s crust), 23 (soil), 2 x 10-6 (sea water)
Isotopes
Table $4$ summarizes the naturally occurring isotopes of the Group 14 elements.
Table $4$: Abundance of the major isotopes of the Group 14 elements.
Isotope Natural abundance (%)
Carbon-12 98.9
Carbon-13 1.1
Carbon-14 trace
Silicon-28 92.23
Silicon-29 4.67
Silicon -30 3.1
Germanium-70 21.23
Germanium-72 27.66
Germanium-73 7.73
Germanium-74 35.94
Germanium-76 7.44
Tin-112 0.97
Tin-114 0.66
Tin-115 0.34
Tin-116 14.54
Tin-117 7.68
Tin-118 24.22
Tin-119 8.59
Tin-120 32.58
Tin-122 4.63
Tin-124 5.79
Lead-204 1.4
Lead-24.1 24.1
Lead-207 22.1
Lead-208 52.4
Although radioactive, carbon-14 is formed in upper layers of the troposphere and the stratosphere, at altitudes of 9–15 km. Thermal neutrons produced by cosmic rays collide with the nuclei of nitrogen-14, forming carbon-14 and a proton. Because of its relatively short half-life of 5730 years, carbon-14 is absent in ancient rocks, but is incorporated in living organisms.
Carbon dating
Carbon dating is a process whereby the age of a material that contains carbon can be determined by comparing the decay rate of that material with that of living material.
Carbon-14 has a half life (t1/2) of 5.73 x 103 years for its decay to nitrogen-14 by the loss of a β particle, (7.1.2).
$^{14}_6C \rightarrow ^{14}_7N + ^0_{-1}e$
The rate of radioactive decay can be expressed as a rate constant (k):
$\text{k = } \dfrac{\text{ln[2]}}{\text{t}_{1/2}} \text{ = } \dfrac{\text{0.693}}{\text{t}_{1/2}}$
For carbon-14, using (7.1.3),
$\text{k = } \dfrac{\text{0.693}}{\text{ 5.73 x 10}^3} \text{ = 1.21 x 10}^{-4} \text{ year}^{-1}$
In 1947 samples of the Dead Sea Scrolls were analyzed by carbon dating. It was found that the carbon-14 present had an activity of d/min.g (where d = disintegration); by contrast in living material the activity is 14 d/min.g. Thus,
$\text{ln}\dfrac{\text{14}}{\text{11}} \text{ = (1.21 x 10}^{-4}\text{) t}$
$\text{t = } \dfrac{\text{ln 1.272}}{\text{1.21 x 10}^{-4}} \text{ = 2.0 x 10}^3\text{ years}$
From the measurement performed in 1947 the Dead Sea Scrolls were determined to be 2000 years old giving them a date of 53 BC, and confirming their authenticity. This discovery is in contrast to the carbon dating results for the Turin Shroud that was supposed to have wrapped Jesus’ body. Carbon dating has shown that the cloth was made between 1260 and 1390 AD. Thus, the Turin Shroud is clearly a fake having been made over a thousand years after its supposed manufacture.
Industrial production
Due to the industrial importance of carbon and silicon, as well as the wide range of fullerene materials, the production of these elements is discussed elsewhere. However, the diamond supply chain is controlled by a limited number of commercial concerns, the largest of which is DeBeers in London (Figure). Diamonds make up only a very small fraction of ore bearing rock. The ore is crushed and subsequently the particles are sorted by density. Diamonds are located in the diamond-rich fraction by X-ray fluorescence, after which the final sorting steps are done by hand.
Germanium ore concentrates are mostly sulfidic, e.g., as an impurity in zinc blende. They are converted to the oxides by heating under air (roasting), (7.1.7).
$\text{GeS}_2\text{ + 3 O}_2 \rightarrow \text{GeO}_2 \text{ + 2 SO}_2$
Part of the germanium ends up in the dust produced during this process, while the rest is converted to germanates which are leached together with the zinc by sulfuric acid. After neutralisation the germanium and other metals are precipitated (leaving the Zn2+ in solution). Germanium dioxide is obtained as a precipitate and converted with chlorine gas or hydrochloric acid to germanium tetrachloride,( 7.1.8) (7.1.9), which has a low boiling point and can be purified by distillation.
$\text{GeO}_2\text{ + 4 HCl} \rightarrow \text{GeCl}_4\text{ + 2 H}_2\text{O}$
$\text{GeO}_2\text{ + 2 Cl}_2 \rightarrow \text{GeCl}_4\text{O}_2$
The germanium tetrachloride is hydrolyzed to the give pure oxide (GeO2), which is then converted to germanium glass for the semiconductor industry, (7.1.10). Germanium used in steel production and other applications that do not require the high purity is produced by reduction with carbon, (7.1.11).
$\text{GeO}_2\text{ + 4 H}_2 \rightarrow \text{Ge + H}_2\text{O}$
$\text{GeO}_2\text{ + C} \rightarrow \text{Ge + CO}_2$
Tin is mined and subsequently smelted, and its production has changed little. In contrast, lead-rich ores contain less than 10% lead, but ores containing as little as 3% lead can be economically exploited. Ores are crushed and concentrated to 70%. Sulfide ores are roasted, producing lead oxide and a mixture of sulfates and silicates of lead. Lead oxide from the roasting process is reduced in a coke-fired blast furnace. This converts most of the lead to its metallic form. Metallic lead that results from the roasting and blast furnace processes still contains significant contaminants of arsenic, antimony, bismuth, zinc, copper, silver, and gold. The melt is treated in a reverberatory furnace (Figure $9$) with air, steam, and sulfur, which oxidizes the contaminants except silver, gold, and bismuth.
Physical properties
Table $5$ provides a summary of the physical properties of the Group 14 elements.
Table $5$: Selected physical properties of the Group 14 elements.
Element Mp (°C) Bp (°C) Density (g/cm3)
C 642 (sublimes) 2.267 (graphite), 3.515 (diamond), 1.8 - 2.1 (amorphous)
Si 1414 3265 2.3290
Ge 938 2833 5.323
Sn 232 2602 7.365 (white), 5.769 (gray)
Pb 327 1749 11.34
Cubic structure
The elements carbon through tin (in its α form) all exist in diamond cubic structure (Figure $10$a), while lead crystallizes in a cubic close packed structure (Figure $10$b). As expected the lattice parameter (a) increases with increased atomic radius (Table $6$). The switch from diamond cubic to close packed cubic may be rationalized by the relative atomic sizes. The diamond cubic structure comprises of two interpenetrating cubic close packed lattices. As the atomic size increases large interstitial vacancies would result, resulting in an unfavorable low-density structure.
Table $6$: Lattice parameter and crystal density for Group 14 elements.
Element Structure a (Å) Atomic radius (Å)
C diamond cubic 3.566 0.70
Si diamond cubic 5.431 1.10
Ge diamond cubic 5.657 1.25
α-Sn (gray) diamond cubic 6.489 1.45
Pb cubic close packed 4.951 1.80 | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/07%3A_Group_14/7.01%3A_The_Group_14_Elements.txt |
It is a commonly held fallacy that James Watt (Figure \(1\)) was the inventor of the steam engine. This actually honor belongs to Thomas Newcomen (Figure \(2\)). Watt’s actual achievement was to improve Newcomen’s design of a steam pump. While Watt was working in the repair shop at the University of Glasgow he was fixing a Newcomen pump when he developed several key improvements on the original design.
The reason for Watt’s success was that Britain was going through a major industrial boom and was in need of significant quantities of raw material including coal. Many of the coalmines, especially those in Devon were prone to flooding. Unfortunately, Newcomen’s engine (which was actually a steam pump) could not pump the water out fast enough, whereas Watt’s engine was powerful enough to drain the mines. As a result Watt’s career as a manufacturer was assured. However, Watt did not see the full potential of his invention, this was left to an employee of his, William Murdock (Figure \(3\)), whose invention of the gearing to allow the steam engine to be used for powering machinery.
One outcome that Watt found from his increased business was an increase in paperwork! While he was living in Redruth, Cornwall, close to where many of the mines were situated he told a friend that he was having “excessive difficulty in finding intelligent managing clerks”. In 1780 Watt solved his paperwork problems by inventing the first method of making copies. This was the subject of a Patent entitled “A new method of copying letters and writing expeditiously”. Watt’s invention involved making ink out of gum Arabic and carbon black.
Note
Gum arabic, also known as gum acacia, is a natural gum made of the hardened sap taken from two species of the acacia tree; acacia senegal and acacia seyal. Carbon black is a form of amorphous carbon that has a high surface area to volume ratio it is commonly produced by the incomplete combustion of heavy petroleum products such as coal tar.
Watt’s ink would stay wet for 24 hours. Writing with the ink and then pressing the result against another piece of paper created a copy. Initially there was great resistance to the copy paper. In particular banks believed that this form of copying could result in forgeries. However, by the end of its first year on sale, Watt sold over 200 examples. The real commercial push came after Watt demonstrated his process to the House of Parliament. The resultant consternation amongst the Members of Parliament resulted in their having to be reminded that Parliament was still in session! By 1785 copies were in general use, however, Cyrus P. Dalkin made the biggest advance in 1823. By using a mixture of carbon black and hot paraffin wax the back of a piece of paper was coated. The carbon black was transferred to another piece of paper underneath by the pressure of the pen. Dalkin had created carbon copies. Both Ralph Wedgwood in England and Pellegrino Turri in Italy had developed forms of carbon paper between 1806 and 1808, but it was Dalkin’s version of carbon copy paper that found usage.
Initially there was almost no market for this new carbon copy paper, until in 1868 when American Lebbeus Rogers was talking part in a balloon ascent. Rogers was an owner of biscuit and greengrocer companies, and was intrigued when a reporter for the Associated Press who was interviewing him used Dalkin’s carbon paper. Rogers gave up his biscuit business and founded a firm to sell carbon paper. In 1873 he talked with E. Remington and Sons the typewriter manufacturer, and it was this application that created a success out of the carbon paper business.
Note
The term “CC” commonly used in e-mail programs today grew from the use of carbon paper and means carbon copy.
Interestingly, carbon black is still used today in modern photocopiers and laser printers. One of the attributes that allow its use is the ability for carbon black particles to become charged. This is related to another application of carbon black in early communications: its ability to conduct electric, and the changes that occur as a function of external pressure.
While Alexander Graham Bell was the first to be awarded a US patent for the electric telephone in March 1876, it was an invention by Thomas Edison (Figure \(5\)) that provided significant improvements. Early telephones relied on a metal diaphragm that was attached to an electromagnet. Speaking into the metal diaphragm caused its vibration, which in tern vibrated the electromagnetic and hence created a current. Unfortunately, the current was very low and subsequent clarity was poor. Between 1877 and 1878 Edison investigated methods to improve the clarity of the signal. The key development was the carbon microphone (Figure \(6\)).
The carbon microphone, also known as a carbon button microphone consisting of two metal plates separated by granules of carbon. One plate faces outward and acts as a diaphragm. When sound waves strike this plate, the pressure on the granules changes, which in turn changes the electrical resistance between the plates. A direct current is passed from one plate to the other, and the changing resistance results in a changing current, which can be passed through a telephone system. The carbon microphone was used in all telephones until the 1980s.
It was one of Edison’s researchers, Edward Acheson (Figure \(7\)) whose discovery allowed for the extension of communication on a global scale. While trying to make artificial diamonds, Acheson began mixing clay and coke (carbon) at very high temperatures. In an electric furnace at high temperatures he found hexagonal crystals of silicon carbide (SiC) attached to the carbon electrode. He called this material carborundum. When, by mistake, he overheated the mixture to 4150 °C he found that the silicon evaporated (boiling point 3265 °C) leaving pure and highly crystalline carbon: graphite.
Graphite is a natural mineral normally associated with other minerals, although an enormous deposit of graphite was discovered in Cumbria, England, which the locals found very useful for marking sheep! Graphite was named by Abraham Werner (Figure \(8\)) in 1789 from the Greek graphein meaning to draw/write due to its use in pencils. However, its use was limited due to the cost. The ability to manufacture high purity graphite resulted in its use in electrodes, dynamo brushes, and batteries. The most terrifying application of graphite was as a result of a material problem associated with the V-2 rocket built by Germany during the Second World War.
The V-2 rockets or Vergeltungswaffen-2 (vengeance weapon-2) was 47 feet long, and reached 3,600 mph with an altitude of 300,000 feet (Figure \(9\)). In order to control the direction of flight the V-2 was guided by four external rudders on the tail fins, and four internal vanes at the exit of the jet (Figure \(10\)). These vanes were made of graphite, it being the only material that would survive the extreme temperatures.
Of course it was as a result of the same German scientists who worked on the V-2, working for NASA, that allowed rockets of sufficient power to escape the Earth’s gravitational pull and send man to the moon, and position many of the satellites that are now vital for global communication. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/07%3A_Group_14/7.02%3A_Carbon_Black-_From_Copying_to_Communication.txt |
Fullerenes and Nanotubes
Introduction
Although nanomaterials had been known for many years prior to the report of C60 the field of nanoscale science was undoubtedly founded upon this seminal discovery. Part of the reason for this explosion in nanochemistry is that while carbon materials range from well-defined nano sized molecules (i.e., C60) to tubes with lengths of hundreds of microns, they do not exhibit the instabilities of other nanomaterials as a result of the very high activation barriers to their structural rearrangement. As a consequence they are highly stable even in their unfunctionalized forms. Despite this range of carbon nanomaterials possible they exhibit common reaction chemistry: that of organic chemistry.
The previously unknown allotrope of carbon: C60, was discovered in 1985, and in 1996, Curl, Kroto, and Smalley were awarded the Nobel Prize in Chemistry for the discovery. The other allotropes of carbon are graphite (sp2) and diamond (sp3). C60, commonly known as the “buckyball” or “Buckminsterfullerene”, has a spherical shape comprising of highly pyramidalized sp2 carbon atoms. The C60 variant is often compared to the typical soccer football, hence buckyball. However, confusingly, this term is commonly used for higher derivatives. Fullerenes are similar in sheet structure to graphite but they contain pentagonal (or sometimes heptagonal) rings that prevent the sheet from being planar. The unusual structure of C60 led to the introduction of a new class of molecules known as fullerenes, which now constitute the third allotrope of carbon. Fullerenes are commonly defined as “any of a class of closed hollow aromatic carbon compounds that are made up of twelve pentagonal and differing numbers of hexagonal faces.”
The number of carbon atoms in a fullerene range from C60 to C70, C76, and higher. Higher order fullerenes include carbon nanotubes that can be described as fullerenes that have been stretched along a rotational axis to form a tube. As a consequence of differences in the chemistry of fullerenes such as C60 and C70 as compared to nanotubes, these will be dealt with separately herein. In addition there have also been reports of nanohorns and nanofibers, however, these may be considered as variations on the general theme. It should be noted that fullerenes and nanotubes have been shown to be in flames produced by hydrocarbon combustion. Unfortunately, these naturally occurring varieties can be highly irregular in size and quality, as well as being formed in mixtures, making them unsuitable for both research and industrial applications.
Fullerenes
Carbon-60 (C60) is probably the most studied individual type of nanomaterial. The spherical shape of C60 is constructed from twelve pentagons and twenty hexagons and resembles a soccer ball (Figure \(1\)a). The next stable higher fullerene is C70 (Figure \(1\)b) that is shaped like a rugby or American football. The progression of higher fullerenes continues in the sequence C74, C76, C78, etc. The structural relationship between each involves the addition of six membered rings. Mathematically (and chemically) two principles define the existence of a stable fullerene, i.e., Euler’s theorem and isolated pentagon rule (IPR). Euler’s theorem states that for the closure of each spherical network, n (n ≥ 2) hexagons and 12 pentagons are required while the IPR says no two pentagons may be connected directly with each other as destabilization is caused by two adjacent pentagons.
Although fullerenes are composed of sp2 carbons in a similar manner to graphite, fullerenes are soluble in various common organic solvents. Due to their hydrophobic nature, fullerenes are most soluble in CS2 (C60 = 7.9 mg/mL) and toluene (C60 = 2.8 mg/mL). Although fullerenes have a conjugated system, their aromaticity is distinctive from benzene that has all C-C bonds of equal lengths, in fullerenes two distinct classes of bonds exist. The shorter bonds are at the junctions of two hexagons ([6, 6] bonds) and the longer bonds at the junctions of a hexagon and a pentagon ([5,6] bonds). This difference in bonding is responsible for some of the observed reactivity of fullerenes.
Synthesis of fullerenes
The first observation of fullerenes was in molecular beam experiments at Rice University. Subsequent studies demonstrated that C60 it was relatively easy to produce grams of fullerenes. Although the synthesis is relatively straightforward fullerene purification remains a challenge and determines fullerene’s commercial price. The first method of production of measurable quantities of fullerenes used laser vaporization of carbon in an inert atmosphere, but this produced microscopic amounts of fullerenes. Laboratory scales of fullerene are prepared by the vaporization of carbon rods in a helium atmosphere. Commercial production ordinarily employs a simple ac or dc arc. The fullerenes in the black soot collected are extracted in toluene and purified by liquid chromatography. The magenta C60 comes off the column first, followed by the red C70, and other higher fullerenes. Even though the mechanism of a carbon arc differs from that of a resistively heated carbon rod (because it involves a plasma) the He pressure for optimum C60 formation is very similar.
A ratio between the mass of fullerenes and the total mass of carbon soot defines fullerene yield. The yields determined by UV-Vis absorption are approximately 40%, 10-15%, and 15% in laser, electric arc, and solar processes. Interestingly, the laser ablation technique has both the highest yield and the lowest productivity and, therefore, a scale-up to a higher power is costly. Thus, fullerene commercial production is a challenging task. The world's first computer controlled fullerene production plant is now operational at the MER Corporation, who pioneered the first commercial production of fullerene and fullerene products.
Endohedral fullerenes
Endohedral fullerenes are fullerenes that have incorporated in their inner sphere atoms, ions or clusters. Endohedral fullerenes are generally divided into two groups: endohedral metallofullerenes and non-metal doped fullerenes. The first endohedral metallofullerenes was called La@C60. The @ sign in the name reflects the notion of a small molecule trapped inside a shell.
Doping fullerenes with metals takes place in-situ during the fullerene synthesis in an arc reactor or via laser evaporation. A wide range of metals have been encased inside a fullerene, i.e., Sc, Y, La, Ce, Ba, Sr, K, U, Zr, and Hf. Unfortunately, the synthesis of endohedral metallofullerenes is unspecific because in addition a high yield of unfilled fullerenes, compounds with different cage sizes are prepared (e.g., La@C60 or La@C82). A characteristic of endohedral metallofullerenes is that electrons will transfer from the metal atom to the fullerene cage and that the metal atom takes a position off-center in the cage. The size of the charge transfer is not always simple to determine, but it is usually between 2 and 3 units (e.g., La2@C80) but can be as high as 6 electrons (e.g., Sc3N@C80). These anionic fullerene cages are very stable molecules and do not have the reactivity associated with ordinary empty fullerenes (see below). This lack of reactivity is utilized in a method to purify endohedral metallofullerenes from empty fullerenes.
The endohedral He@C60 and Ne@C60 form when C60 is exposed to a pressure of around 3 bar of the appropriate noble gases. Under these conditions it was possible to dope 1 in every 650,000 C60 cages with a helium atom. Endohedral complexes with He, Ne, Ar, Kr and Xe as well as numerous adducts of the He@C60 compound have also been proven with operating pressures of 3000 bars and incorporation of up to 0.1 % of the noble gases. The isolation of N@C60, N@C70 and P@C60 is very unusual and unlike the metal derivatives no charge transfer of the pnictide atom in the center to the carbon atoms of the cage takes place.
Chemically functionalized fullerenes
Although fullerenes have a conjugated aromatic system all the carbons are quaternary (i.e., containing no hydrogen), which results in making many of the characteristic substitution reactions of planar aromatics impossible. Thus, only two types of chemical transformations exist: redox reactions and addition reactions. Of these, addition reactions have the largest synthetic value. Another remarkable feature of fullerene addition chemistry is the thermodymics of the process. Since the sp2 carbon atoms in a fullerene are paramidalized there is significant strain energy. For example, the strain energy in C60 is ca 8 kcal/mol, which is 80% of its heat of formation. So the relief of this strain energy leading to sp3 hybridized C atoms is the major driving force for addition reactions (Figure \(2\)). As a consequence, most additions to fullerenes are exothermic reactions.
Cyclic voltammetry (CV) studies show that C60 can be reduced and oxidized reversibly up to 6 electrons with one-electron transfer processes. Fulleride anions can be generated by electrochemical method and then be used to synthesize covalent organofullerene derivatives. Alkali metals can chemically reduce fullerene in solution and solid state to form MxC60 (x = 3 - 6). C60 can also be reduced by less electropositive metals like mercury to form C60- and C602-. In addition, salts can also be synthesized with organic molecules, for example [TDAE+][C60-] possesses interesting electronic and magnetic behavior.
Geometric and electronic analysis predicted that fullerene behaves live an electro-poor conjugated polyolefin. Indeed C60 and C70 undergo a range of nucleophilic reactions with carbon, nitrogen, phosphorous and oxygen nucleophiles. C60 reacts readily with organolithium and Grignard compounds to form alkyl, phenyl or alkanyl fullerenes. Possibly the most widely used additions to fullerene is the Bingel reaction (Figure \(3\)), where a carbon nucleophile, generated by deprotonation of α-halo malonate esters or ketones, is added to form a cyclopropanation product. The α-halo esters and ketones can also be generated in situ with I2 or CBr4 and a weak base as 1,8-diazabicyclo[5.4.0]unde-7ene (DBU). The Bingel reaction is considered one of the most versatile and efficient methods to functionalize C60.
Cycloaddition is another powerful tool to functionalize fullerenes, in particular because of its selectivity with the 6,6 bonds, limiting the possible isomers (Figure \(4\)). The dienophilic feature of the [6,6] double bonds of C60 enables the molecule to undergo various cycloaddition reactions in which the monoadducts can be generated in high yields. The best studies cycloadditon reactions of fullerene are [3+2] additions with diazoderivatives and azomethine ylides (Prato reactions). In this reaction, azomethine ylides can be generated in situ from condensation of α-amino acids with aldehydes or ketones, which produce 1,3 dipoles to further react with C60 in good yields (Figure \(5\)). Hundreds of useful building blocks have been generated by those two methods. The Prato reactions have also been successfully applied to carbon nanotubes.
The oxidation of fullerenes, such as C60, has been of increasing interest with regard to applications in photoelectric devices, biological systems, and possible remediation of fullerenes. The oxidation of C60 to C60On (n = 1, 2) may be accomplished by photooxidation, ozonolysis, and epoxidation. With each of these methods, there is a limit to the isolable oxygenated product, C60On with n < 3. Highly oxygenated fullerenes, C60On with 3 ≤ n ≤ 9, have been prepared by the catalytic oxidation of C60 with ReMeO3/H2O2.
Carbon nanotubes
A key breakthrough in carbon nanochemistry came in 1993 with the report of needle-like tubes made exclusively of carbon. This material became known as carbon nanotubes (CNTs). There are several types of nanotubes. The first discovery was of multi walled tubes (MWNTs) resembling many pipes nested within each other. Shortly after MWNTs were discovered single walled nanotubes (SWNTs) were observed. Single walled tubes resemble a single pipe that is potentially capped at each end. The properties of single walled and multi walled tubes are generally the same, although single walled tubes are believed to have superior mechanical strength and thermal and electrical conductivity; it is also more difficult to manufacture them.
Single walled carbon nanotubes (SWNTs) are by definition fullerene materials. Their structure consists of a graphene sheet rolled into a tube and capped by half a fullerene (Figure \(6\)). The carbon atoms in a SWNT, like those in a fullerene, are sp2 hybridized. The structure of a nanotube is analogous to taking this graphene sheet and rolling it into a seamless cylinder. The different types of SWNTs are defined by their diameter and chirality. Most of the presently used single-wall carbon nanotubes have been synthesized by the pulsed laser vaporization method, however, increasingly SWNTs are prepared by vapor liquid solid catalyzed growth.
The physical properties of SWNTs have made them an extremely attractive material for the manufacturing of nano devices. SWNTs have been shown to be stronger than steel as estimates for the Young’s modulus approaches 1 Tpa. Their electrical conductance is comparable to copper with anticipate current densities of up to 1013 A/cm2 and a resistivity as low as 0.34 x 10-4 Ω.cm at room temperatures. Finally, they have a high thermal conductivity (3000 - 6000 W.m/K).
The electronic properties of a particular SWNT structure are based on its chirality or twist in the structure of the tube which is defined by its n,m value. The values of n and m determine the chirality, or "twist" of the nanotube. The chirality in turn affects the conductance of the nanotube, its density, its lattice structure, and other properties. A SWNT is considered metallic if the value n-m is divisible by three. Otherwise, the nanotube is semi-conducting. The external environment also has an effect on the conductance of a tube, thus molecules such as O2 and NH3 can change the overall conductance of a tube, while the presence of metals have been shown to significantly effect the opto-electronic properties of SWNTs.
Multi walled carbon nanotubes (MWNTs) range from double walled NTs, through many-walled NTs (Figure \(7\)) to carbon nanofibers. Carbon nanofibers are the extreme of multi walled tubes (Figure \(8\)) and they are thicker and longer than either SWNTs or MWNTs, having a cross-sectional of ca. 500 Å2 and are between 10 to 100 μm in length. They have been used extensively in the construction of high strength composites.
Synthesis of carbon nanotubes
A range of methodologies have been developed to produce nanotubes in sizeable quantities, including arc discharge, laser ablation, high pressure carbon monoxide (HiPco), and vapor liquid solid (VLS) growth. All these processes take place in vacuum or at low pressure with a process gases, although VLS growth can take place at atmospheric pressure. Large quantities of nanotubes can be synthesized by these methods; advances in catalysis and continuous growth processes are making SWNTs more commercially viable.
The first observation of nanotubes was in the carbon soot formed during the arc discharge production of fullerenes. The high temperatures caused by the discharge caused the carbon contained in the negative electrode to sublime and the CNTs are deposited on the opposing electrode. Tubes produced by this method were initially multi walled tubes (MWNTs). However, with the addition of cobalt to the vaporized carbon, it is possible to grow single walled nanotubes. This method it produces a mixture of components, and requires further purification to separate the CNTs from the soot and the residual catalytic metals. Producing CNTs in high yield depends on the uniformity of the plasma arc, and the temperature of the deposit forming on the carbon electrode.
Higher yield and purity of SWNTs may be prepared by the use of a dual-pulsed laser. SWNTs can be grown in a 50% yield through direct vaporization of a Co/Ni doped graphite rod with a high-powered laser in a tube furnace operating at 1200 °C. The material produced by this method appears as a mat of “ropes”, 10 - 20 nm in diameter and up to 100 μm or more in length. Each rope consists of a bundle of SWNTs, aligned along a common axis. By varying the process parameters such as catalyst composition and the growth temperature, the average nanotube diameter and size distribution can be varied. Although arc-discharge and laser vaporization are currently the principal methods for obtaining small quantities of high quality SWNTs, both methods suffer from drawbacks. The first is that they involve evaporating the carbon source, making scale-up on an industrial level difficult and energetically expensive. The second issue relates to the fact that vaporization methods grow SWNTs in highly tangled forms, mixed with unwanted forms of carbon and/or metal species. The SWNTs thus produced are difficult to purify, manipulate, and assemble for building nanotube-device architectures for practical applications.
In order to overcome some of the difficulties of these high-energy processes, the chemical catalysis method was developed in which a hydrocarbon feedstock is used in combination with a metal catalyst. The catalyst is typically, but not limited to iron, colbalt, or iron/molybdenum, it is heated under reducing conditions in the presence of a suitable carbon feedstock, e.g., ethylene. This method can be used for both SWNTs and MWNTs; the formation of each is controlled by the identity of the catalyst and the reaction conditions. A convenient laboratory scale apparatus is available from Nanotech Innovations, Inc., for the synthesis of highly uniform, consistent, research sample that uses pre-weighed catalyst/carbon source ampoules. This system, allows for 200 mg samples of MWNTs to be prepared for research and testing. The use of CO as a feedstock, in place of a hydrocarbon, led to the development of the high-pressure carbon monoxide (HiPco) procedure for SWNT synthesis. By this method, it is possible to produce gram quantities of SWNTs, unfortunately, efforts to scale beyond that have not met with complete success.
Initially developed for small-scale investigations of catalyst activity, vapor liquid solid (VLS) growth of nanotubes has been highly studied, and now shows promise for large-scale production of nanotubes. Recent approaches have involved the use of well-defined nanoparticle or molecular precursors and many different transition metals have been employed, but iron, nickel, and cobalt remain to be the focus of most research. The nanotubes grow at the sites of the metal catalyst; the carbon-containing gas is broken apart at the surface of the catalyst particle, and the carbon is transported to the edges of the particle, where it forms the nanotube. The length of the tube grown in surface supported catalyst VLS systems appears to be dependent on the orientation of the growing tube with the surface. By properly adjusting the surface concentration and aggregation of the catalyst particles it is possible to synthesize vertically aligned carbon nanotubes, i.e., as a carpet perpendicular to the substrate.
Of the various means for nanotube synthesis, the chemical processes show the greatest promise for industrial scale deposition in terms of its price/unit ratio. There are additional advantages to the VLS growth, which unlike the other methods is capable of growing nanotubes directly on a desired substrate. The growth sites are controllable by careful deposition of the catalyst. Additionally, no other growth methods have been developed to produce vertically aligned SWNTs.
Chemical functionalization of carbon nanotubes
The limitation of using carbon nanotubes in any practical applications has been its solubility; for example SWNTs have little to no solubility in most solvent due to the aggregation of the tubes. Aggregation/roping of nanotubes occurs as a result of the high van der Waals binding energy of ca. 500 eV per mm of tube contact. The van der Waals force between the tubes is so great, that it take tremendous energy to pry them apart, making it very to make combination of nanotubes with other materials such as in composite applications. The functionalization of nanotubes, i.e., the attachment of “chemical functional groups” provides the path to overcome these barriers. Functionalization can improve solubility as well as processibility, and has been used to align the properties of nanotubes to those of other materials. The clearest example of this is the ability to solubilize nanotubes in a variety of solvents, including water. It is important when discussing functionalization that a distinction is made between covalent and non-covalent functionalization.
Current methods for solubilizing nanotubes without covalent functionalization include highly aromatic solvents, super acids, polymers, or surfactants. Non-covalent “functionalization” is generally on the concept of supramolecular interactions between the SWNT and some macromolecule as a result of various adsorption forces, such as van der Waals’ and π-stacking interactions. The chemical speciation of the nanotube itself is not altered as a result of the interaction. In contrast, covalent functionalization relies on the chemical reaction at either the sidewall or end of the SWNT. As may be expected the high aspect ratio of nanotubes means that sidewall functionalization is much more important than the functionalization of the cap. Direct covalent sidewall functionalization is associated with a change of hybridization from sp2 to sp2 and a simultaneous loss of conjugation. An alternative approach to covalent functionalization involves the reaction of defects present (or generated) in the structure of the nanotube. Defect sites can be the open ends and holes in the sidewalls, and pentagon and heptagon irregularities in the hexagon graphene framework (often associated with bends in the tubes). All these functionalizations are exohedral derivatizations. Taking the hollow structure of nanotubes into consideration, endohedral functionalization of SWNTs is possible, i.e., the filling of the tubes with atoms or small molecules. It is important to note that covalent functionalization methods have one problem in common: extensive covalent functionalization modifies SWNT properties by disrupting the continuous π–system of SWNTs.
Various applications of nanotubes require different, specific modification to achieve desirable physical and chemical properties of nanotubes. In this regard, covalent functionalization provides a higher degree of fine-tuning the chemistry and physics of SWNTs than non-covalent functionalization. Until now, a variety of methods have been used to achieve the functionalization of nanotubes (Figure \(9\)).
Taking chemistry developed for C60, SWNTs may be functionalized using 1,3 dipolar addition of azomethine ylides. The functionalized SWNTs are soluble in most common organic solvents. The azomethine ylide functionalization method was also used for the purification of SWNTs. Under electrochemical conditions, aryl diazonium salts react with SWNTs to achieve functionalized SWNTs, alternatively the diazonium ions may be generated in-situ from the corresponding aniline, while a solvent free reaction provides the best chance for large-scale functionalization this way. In each of these methods it is possible to control the amount of functionalization on the tube by varying reaction times and the reagents used; functionalization as high as 1 group per every 10 - 25 carbon atoms is possible.
Organic functionalization through the use of alkyl halides, a radical pathway, on tubes treated with lithium in liquid ammonia offers a simple and flexible route to a range of functional groups. In this reaction, functionalization occurs on every 17 carbons. Most success has been found when the tubes are dodecylated. These tubes are soluble in chloroform, DMF, and THF.
The addition of oxygen moieties to SWNT sidewalls can be achieved by treatment with acid or wet air oxidation, and ozonolysis. The direct epoxidation of SWNTs may be accomplished by the direct reaction with a peroxide reagent, or catalytically. Catalytic de-epoxidation (Figure \(10\)) allows for the quantitative analysis of sidewall epoxide and led to the surprising result that previously assumed “pure” SWNTs actually contain ca. 1 oxygen per 250 carbon atoms.
One of the easiest functionalization routes, and a useful synthon for subsequent conversions, is the fluorination of SWNTs, using elemental fluorine. Importantly, a C:F ratios of up to 2:1 can be achieved without disruption of the tubular structure. The fluorinated SWNTs (F-SWNTs) proved to be much more soluble than pristine SWNTs in alcohols (1 mg/mL in iso-propanol), DMF and other selected organic solvents. Scanning tunneling microscopy (STM) revealed that the fluorine formed bands of approximately 20 nm, while calculations using DFT revealed 1,2 addition is more energetically preferable than 1,4 addition, which has been confirmed by solid state 13C NMR. F-SWNTs make highly flexible synthons and subsequent elaboration has been performed with organo lithium, Grignard reagents, and amines.
Functionalized nanotubes can be characterized by a variety of techniques, such as atomic force microscopy (AFM), transmission electron microscopy (TEM), UV-vis spectroscopy, and Raman spectroscopy. Changes in the Raman spectrum of a nanotube sample can indicate if functionalization has occurred. Pristine tubes exhibit two distinct bands. They are the radial breathing mode (230 cm-1) and the tangential mode (1590 cm-1). When functionalized, a new band, called the disorder band, appears at ca.1350 cm-1. This band is attributed to sp3-hybridized carbons in the tube. Unfortunately, while the presence of a significant D mode is consistent with sidewall functionalization and the relative intensity of D (disorder) mode versus the tangential G mode (1550 – 1600 cm-1) is often used as a measure of the level of substitution. However, it has been shown that Raman is an unreliable method for determination of the extent of functionalization since the relative intensity of the D band is also a function of the substituents distribution as well as concentration. Recent studies suggest that solid state 13C NMR are possibly the only definitive method of demonstrating covalent attachment of particular functional groups.
Coating carbon nanotubes: creating inorganic nanostructures
Fullerenes, nanotubes and nanofibers represent suitable substrates for the seeding other materials such as oxides and other minerals, as well as semiconductors. In this regard, the carbon nanomaterial acts as a seed point for the growth as well as a method of defining unusual aspect ratios. For example, silica fibers can be prepared by a number of methods, but it is only through coating SWNTs that silica nano-fibers with of micron lengths with tens of nanometers in diameter may be prepared.
While C60 itself does not readily seed the growth of inorganic materials, liquid phase deposition of oxides, such as silica, in the presence of fullerenol, C60(OH)n, results in the formation of uniform oxide spheres. It appears the fullerenol acts as both a reagent and a physical point for subsequent oxide growth, and it is C60, or an aggregate of C60, that is present within the spherical particle. The addition of fullerenol alters the morphology and crystal phase of CaCO3 precipitates from aqueous solution, resulting in the formation of spherical features, 5-pointed flower shaped clusters, and triangular crystals as opposed to the usual rhombic crystals. In addition, the meta-stable vaterite phase is observed with the addition of C60(OH)n.
As noted above individual SWNTs may be obtained in solution when encased in a cylindrical micelle of a suitable surfactant. These individualized nanotubes can be coated with a range of inorganic materials. Liquid phase deposition (LPD) appears to have significant advantages over other methods such as incorporating surfacted SWNTs into a preceramic matrix, in situ growth of the SWNT in an oxide matrix, and sol-gel methods. The primary advantage of LPD growth is that individual SWNTs may be coated rather than bundles or ropes. For example, SWNTs have been coated with silica by liquid phase deposition (LPD) using a silica/H2SiF6 solution and a surfactant-stabilized solution of SWNTs. The thickness of the coating is dependent on the reaction mixture concentration and the reaction time. The SWNT core can be removed by thermolysis under oxidizing conditions to leave a silica nano fiber. It is interesting to note that the use of a surfactant is counter productive when using MWNTs and VGFs, in this case surface activation of the nanotube offers the suitable growth initiation. Pre-oxidation of the MWNT or VGF allows for uniform coatings to be deposited. The coated SWNTs, MWNTs, and VGFs can be subsequently reacted with suitable surface reagents to impart miscibility in aqueous solutions, guar gels, and organic matrixes. In addition to simple oxides, coated nanotubes have been prepared with minerals such as carbonates and semiconductors.
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Graphene
Introduction
Graphene is a one-atom-thick planar sheet of sp2-bonded carbon atoms that are densely packed in a honeycomb crystal lattice (Figure \(11\)). The name comes from “graphite” and “alkene”; graphite itself consists of many graphene sheets stacked together.
Single-layer graphene nanosheets were first characterized in 2004, prepared by mechanical exfoliation (the “scotch-tape” method) of bulk graphite. Later graphene was produced by epitaxial chemical vapor deposition on silicon carbide and nickel substrates. Most recently, graphene nanoribbons (GNRs) have been prepared by the oxidative treatment of carbon nanotubes and by plasma etching of nanotubes embedded in polymer films.
Physical properties of graphene
Graphene has been reported to have a Young’s modulus of 1 TPa and intrinsic strength of 130 GP; similar to single walled carbon nanotubes (SWNTs). The electronic properties of graphene also have some similarity with carbon nanotubes. Graphene is a zero-bandgap semiconductor. Electron mobility in graphene is extraordinarily high (15,000 cm2/V.s at room temperature) and ballistic electron transport is reported to be on length scales comparable to that of SWNTs. One of the most promising aspects of graphene involves the use of GNRs. Cutting an individual graphene layer into a long strip can yield semiconducting materials where the bandgap is tuned by the width of the ribbon.
While graphene’s novel electronic and physical properties guarantee this material will be studied for years to come, there are some fundamental obstacles yet to overcome before graphene based materials can be fully utilized. The aforementioned methods of graphene preparation are effective; however, they are impractical for large-scale manufacturing. The most plentiful and inexpensive source of graphene is bulk graphite. Chemical methods for exfoliation of graphene from graphite provide the most realistic and scalable approach to graphene materials.
Graphene layers are held together in graphite by enormous van der Waals forces. Overcoming these forces is the major obstacle to graphite exfoliation. To date, chemical efforts at graphite exfoliation have been focused primarily on intercalation, chemical derivatization, thermal expansion, oxidation-reduction, the use of surfactants, or some combination of these.
Graphite oxide
Probably the most common route to graphene involves the production of graphite oxide (GO) by extremely harsh oxidation chemistry. The methods of Staudenmeier or Hummers are most commonly used to produce GO, a highly exfoliated material that is dispersible in water. The structure of GO has been the subject of numerous studies; it is known to contain epoxide functional groups along the basal plane of sheets as well as hydroxyl and carboxyl moieties along the edges (Figure \(12\)). In contrast to other methods for the synthesis of GO, the the m-peroxybenzoic acid (m-CPBA) oxidation of microcrystalline synthetic graphite at room temperature yields graphite epoxide in high yield, without significant additional defects.
As graphite oxide is electrically insulating, it must be converted by chemical reduction to restore the electronic properties of graphene. Chemically converted graphene (CCG) is typically reduced by hydrazine or borohydride. The properties of CCG can never fully match those of graphene for two reasons:
1. Oxidation to GO introduces defects.
2. Chemical reduction does not fully restore the graphitic structure.
As would be expected, CCG is prone to aggregation unless stabilized. Graphene materials produced from pristine graphite avoid harsh oxidation to GO and subsequent (incomplete) reduction; thus, materials produced are potentially much better suited to electronics applications.
A catalytic approach to the removal of epoxides from fullerenes and SWNTs has been applied to graphene epoxide and GO. Treatment of oxidized graphenes with methyltrioxorhenium (MeReO3, MTO) in the presence of PPh3 results in the oxygen transfer, to form O=PPh3 and allow for quantification of the C:O ratio.
Homogeneous graphene dispersions
An alternate approach to producing graphene materials involves the use of pristine graphite as starting material. The fundamental value of such an approach lies in its avoidance of oxidation to GO and subsequent (incomplete) reduction, thereby preserving the desirable electronic properties of graphene. There is precedent for exfoliation of pristine graphite in neat organic solvents without oxidation or surfactants. It has been reported that N,N-dimethylformamide (DMF) dispersions of graphene are possible, but no detailed characterization of the dispersions were reported. In contrast, Coleman and coworkers reported similar dispersions using N-methylpyrrolidone (NMP), resulting in individual sheets of graphene at a concentration of ≤0.01 mg/mL. NMP and DMF are highly polar solvents, and not ideal in cases where reaction chemistry requires a nonpolar medium. Further, they are hygroscopic, making their use problematic when water must be excluded from reaction mixtures. Finally, DMF is prone to thermal and chemical decomposition.
Recently, dispersions of graphene has been reported in ortho-dichlorobenzene (ODCB) using a wide range of graphite sources. The choice of ODCB for graphite exfoliation was based on several criteria:
1. ODCB is a common reaction solvent for fullerenes and is known to form stable SWNT dispersions.
2. ODCB is a convenient high-boiling aromatic, and is compatible with a variety of reaction chemistries.
3. ODCB, being aromatic, is able to interact with graphene via π-π stacking.
4. It has been suggested that good solvents for graphite exfoliation should have surface tension values of 40 – 50 mJ/m2. ODCB has a surface tension of 36.6 mJ/m2, close to the proposed range.
Graphite is readily exfoliated in ODCB with homogenization and sonication. Three starting materials were successfully dispersed: microcrystalline synthetic, thermally expanded, and highly ordered pyrolytic graphite (HOPG). Dispersions of microcrystalline synthetic graphite have a concentration of 0.03 mg/mL, determined gravimetrically. Dispersions from expanded graphite and HOPG are less concentrated (0.02 mg/mL).
High resolution transmission electron microscopy (HRTEM) shows mostly few-layer graphene (n < 5) with single layers and small flakes stacked on top (Figure \(13\)). Large graphitic domains are visible; this is further supported by selected area electron diffraction (SAED) and fast Fourier transform (FFT) in selected areas. Atomic force microscope (AFM) images of dispersions sprayed onto silicon substrates shows extremely thin flakes with nearly all below 10 nm. Average height is 7 - 10 nm. The thinnest are less than 1 nm, graphene monolayers. Lateral dimensions of nanosheets range from 100 – 500 nm.
As-deposited films cast from ODCB graphene show poor electrical conductivity, however, after vacuum annealing at 400 °C for 12 hours the films improve vastly, having sheet resistances on the order of 60 Ω/sq. By comparison, graphene epitaxially grown on Ni has a reported sheet resistance of 280 Ω/sq.
Covalent functionalization of graphene and graphite oxide
The covalent functionalization of SWNTs is well established. Some routes to covalently functionalized SWNTs include esterification/ amidation, reductive alkylation (Billups reaction), and treatment with azomethine ylides (Prato reaction), diazonium salts, or nitrenes. Conversely, the chemical derivatization of graphene and GO is still relatively unexplored.
Some methods previously demonstrated for SWNTs have been adapted to GO or graphene. GO carboxylic acid groups have been converted into acyl chlorides followed by amidation with long-chain amines. Additionally, the coupling of primary amines and amino acids via nucleophilic attack of GO epoxide groups has been reported. Yet another route coupled isocyanates to carboxylic acid groups of GO. Functionalization of partially reduced GO by aryldiazonium salts has also been demonstrated. The Billups reaction has been performed on the intercalation compound potassium graphite (C8K), as well as graphite fluoride, and most recently GO. Graphene alkylation has been accomplished by treating graphite fluoride with alkyllithium reagents.
ODCB dispersions of graphene may be readily converted to covalently functionalize graphene. Thermal decomposition of benzoyl peroxide is used to initiate radical addition of alkyl iodides to graphene in ODCB dispersions.
Additionally, functionalized graphene with nitrenes generated by thermal decomposition of aryl azides
Bibliography
• P. Blake, P. D. Brimicombe, R. R. Nair, T. J. Booth, D. Jiang, F. Schedin, L. A. Ponomarenko, S. V. Morozov, H. F. Gleeson, E. W. Hill, A. K. Geim, and K. S. Novoselov, Nano Lett., 2008, 8, 1704.
• J. Chattopadhyay, A. Mukherjee, C. E. Hamilton, J.-H. Kang, S. Chakraborty, W. Guo, K. F. Kelly, A. R. Barron, and W. E. Billups, J. Am. Chem. Soc., 2008, 130, 5414.
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There are a myriad of organic compounds containing carbon-nitrogen bonds, including: amines, imines, and nitriles. However, here we are concerned with the simplest carbon-nitrogen compounds.
Cyanogen
Cyanogen, (CN)2, may be considered the smallest molecular fragment containing carbon and nitrogen (Figure $1$a). The reaction chemistry of cyanogen is related to that of the halogens, i.e., F2, Cl2, etc. Consequently, cyanogen is called a pseudo halogen.
As shown in Table $1$ the bonding in cyanogen is consistent with localization of the π-bonding between carbon and nitrogen given the similarity of the C-N bond distance in cyanogens and acetonitrile. However, there is clearly some π-delocalization associated with the C-C distance given its shortening as compared to ethane.
Table $1$: A comparison of the bond distances in selected carbon nitrogen compounds.
Compound Formula C-C bond distance (Å) C-N bond distance (Å)
Cyanogen (CN)2 1.393 1.163
Hydrogen cyanide HCN - 1.154
Acetonitrile CH3CN 1.46 1.16
Ethane C2H6 1.535 -
Ethylene C2H4 1.339 -
Cyanogen is produced by the reaction of a mixture of the cyanide and chloride of mercury, (7.4.1).
$\text{Hg(CN)}_2 \text{ + HgCl}_2 \rightarrow \text{(CN)}_2 \text{ + 2 Hg + Cl}_2$
Alternatively, the decomposition of unstable copper(II) cyanide, formed from a copper(II) salts with a Group 1 cyanide, (7.4.2), yields cyanogens, (7.4.3).
$\text{CuSO}_4\text{ + 2 KCN} \rightarrow \text{Cu(CN)}_2\text{ + K}_2\text{SO}_4$
$\text{2 Cu(CN)}_2 \rightarrow \text{(CN)}_2 \text{ + 2 CuCN}$
Cyanogen is a flammable gas (Mp = -28 °C and Bp = -21 °C) that produces the second hottest flame natural flame (after carbon subnitride, C4N2) with a temperature of over 4525 °C when burnt in oxygen. Heating cyanogen in the absence of oxygen results self polymerizes.
Hydrolysis of cyanogen results in addition across the carbon-nitrogen triple bonds and the formation of oxamide. Cleavage of the C-C bond does occur in the presence of base (e.g., KOH), with the formation of cyanide (CN-) and cyanate (CNO-) salts, (7.4.4).
$\text{(CN)}_2\text{ + 2 OH}^- \rightarrow \text{CN}^-\text{ + CNO}^- \text{ + H}_2\text{O}$
Dicyanoacetylene
Dicyanoacetylene (also known as carbon subnitride or by its IUPAC name but-2-ynedinitrile) has the structure shown in Figure $1$b, and may be thought of as a dicyanaide substituted acetylene.
At room temperature, dicyanoacetylene is a clear liquid, however, solid dicyanoacetylene has been detected in the atmosphere of Titan (the largest moon of the planet Saturn) by infrared spectroscopy. Dicyanoacetylene is an entropic explosive giving carbon powder and nitrogen gas. In the presence of oxgyen it burns with a bright blue-white flame at a temperature of 4990 °C.
Hydrogen cyanide
Hydrogen cyanide (HCN) is a colorless, highly poisonous, gas (Mp = -13.5 °C and Bp = 25.6 °C). Due to its original isolation from Prussian blue (hydrated ferric ferrocyanide), hydrogen cyanaide is also known by the name of prussic acid.
The synthesis of hydrogen cyanide is accomplished commercially by the partial oxidation of methane in the presence of ammonia, (7.4.5), using a platinum catalyst. The heat to activate the reaction is derived from the partial combustion of the methane and ammonia. The resulting aqueous solution is dried by distilled from phosphorus pentoxide (P2O5) to yield anhydrous hydrogen cyanide.
$\text{2 CH}_4 \text{ + 3 O}_2\text{ + 2 NH}_3 \rightarrow \text{ 2 HCN + 6 H}_2\text{O}$
Hydrogen cyanide may also be formed in the absence of oxygen, (7.4.6); however, in this case the reaction must be heated externally.
$\text{CH}_4 \text{ + NH}_3 \rightarrow \text{ HCN + 3 H}_2$
Small quantities of hydrogen cyanide for laboratory use may be prepared by the reaction of an acid with a cyanide salt (either potassium or sodium), (7.4.7).
$\text{KCN + H}^+ \rightarrow \text{HCN + K}^+$
The structure of hydrogen cyanide is shown in Figure $2$ along with its isomeric form, hydrogen isocyanide (HNC). While hydrogen cyanide is present in the pits of many fruits, and is generated by burnet moths and some millipedes, hydrogen isocyanide is only found in interstellar space. It is postulated, however, that along with HCN, HNC is an important building block for amino acids and hence life.
In the liquid state hydrogen cyanide forms strong hydrogen bonds (Figure $3$). Hydrogen cyanide is a good solvent for polar compounds due to its high permittivity (∈r) and high dipole moment (2.98 D).
In aqueous solution hydrogen cyanide is a weak acid, (7.4.7), and several salts are known. However, HCN also reacts with water to give ammonium formate via formamide, (7.4.8).
$\text{HCN + H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+ \text{ + CN}^-$
$\text{HCN + H}_2\text{O} \rightarrow \text{HC(O)NH}_2 \xrightarrow{\text{H}_2\text{O}} \text{HCO}_2^- \text{ + NH}_4^+$
In a similar manner to cyanogen’s relationship to the halogens, the cyanide anion (CN-) is considered a pseudo halide (i.e., F-, Cl-, etc), and as such forms many coordination compounds, e.g., [Fe(CN)6]3- and [Ag(CN)2]-.
Assassination, execution, and the Holocaust
Hydrogen cyanide is fatal to humans due to its inhibition of the enzyme cytochrome c oxidase by the cyanide ion (CN-), which results in the halting of cellular respiration. A concentration of 300 mg/m3 will kill within 10 minutes, while 3200 mg/m3 (ca. 3500 ppm) will be fatal in about 1 minute.
The symptoms of cyanide poisoning appear similar to a heart attack and this has led to it being the poison of choice for both fictional murder mystery writers as well as the former KGB (Konitet gosudarstvennoy bezopasnosti or Committee for State Security) and its predecessor SMERSH (from the contraction smert shpionam meaning death to spies) in real life. Possibly the most famous use of hydrogen cyanide for assassination was the use of an atomizer mist gun by KGB agent Bohdan Stashynsky for the killing of the Ukranian political writer and anti-communist Lev Rebet in 1957, and later in 1959 that of fellow Ukranian, Stepan Bandera. In both cases the intention was to induce cardiac arrest and make it look like the victim had died of a heart attack.
Without doubt the most notorious use of hydrogen cyanide is in the form of the product Zyklon B, which was originally developed as an insecticide. Dr. Walter Heerdt found that hydrogen cyanide could be absorbed onto substrates such as absorbent pellets (e.g., silica), fibers, or diatomaceous earth (Figure $4$). While stable in an airtight container, once opened the hydrogen cyanide is released. The “B” in the trade name comes from the German name for prussic acid (the common name for hydrogen cyanide), i.e., Blausäure meaning blue acid.
The first wholesale use of Zyklon B was actually in the US where it was used as early as 1929 to disinfect the freight trains and cloths of Mexican immigrants entering the US. The first use of Zyklon B in the concentration camps during World War II was for a similar purpose (in particular for delousing to control typhus); however, for its use as an insecticide, Zyklon B contained a warning odorant. The deliberate manufacture of Zyklon B without the odorant resulted in the material that was used on a group of 250 gypsy children at the Buchenwald concentration camp in early 1940. Subsequently in September 1940, a similar number of sick Polish prisoners of war and 600 Soviet prisoners of war were killed at Auschwitz (Figure $5$). Once these horrific tests were performed, the systematic murder of millions of people, including Jews, gypsies, and homosexuals, was accomplished using Zyklon B at Auschwitz, Majdanek, Sachsenhausen, and one of the Operation Reinhard camps. It is only fitting that many of the architects of the Holocaust themselves died from cyanide, including: Adolf Hitler (in addition to a bullet), Joseph Goebbles, Hermann Göring, and Heinrich Himmler.
Despite the horror associated with the use of hydrogen cyanide for the Holocaust, it was used by 11 US states for the death penalty (Figure $6$). Arizona, Maryland, and Missouri retain the gas chamber as a secondary method of execution though they have lethal injection as the primary method. Potassium cyanide (KCN) pellets are placed into a compartment directly below the chair in the gas chamber . The condemned person is then strapped into the chair, and the airtight chamber sealed. Concentrated sulfuric acid (H2SO4) is then poured down a tube onto the cyanide pellets to generate hydrogen cyanide, (7.4.7). Execution by gas chamber is especially unpleasant for the witnesses to the execution due to the physical responses exhibited during the process of dying, including: convulsions and excessive drooling.
Bibliography
• J. Wu and N. J. Evans, Astrophys. J., 2003, 592, L79. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/07%3A_Group_14/7.04%3A_Nitrogen_Compounds_of_Carbon.txt |
Carbon monoxide (CO) is iso-electronic with nitrogen (N2) and formed through the incomplete combustion of carbon, (7.5.1), or hydrocarbon compounds.
$\text{2 C + O}_2 \rightarrow \text{ 2 CO}$
Carbon monoxide may also be made from steam and coal as part of synthesis gas, (7.4.2). A convenient laboratory preparation of CO is the dehydration of formic acid by sulfuric acid, (7.4.3).
$\text{C + H}_2\text{O} \rightarrow \text{CO + H}_2$
$\text{HCO}_2\text{H} \rightarrow \text{CO + H}_2\text{O}$
Hazards and toxicity
Carbon monoxide is flammable, (7.4.4), and has an explosive limit of 12.5 – 74% with an auto ignition temperature of 609 °C.
$\text{2 CO + O}_2 \rightarrow \text{2 CO}_2$
Carbon monoxide is also very toxic; however, it is colorless, odorless, tasteless and non-irritating all of which increase its danger. The incomplete combustion of hydrocarbon (natural gas or heating oil) or carbon sources (coal or charcoal) is a common hazard in the home. In a closed environment (e.g., charcoal grill in room with no, or poor, ventilation) as carbon uses up oxygen in room the formation of toxic carbon monoxide results instead of carbon dioxide (CO2).Typical sources of CO in the home are shown in Figure $1$.
The toxicity of CO is due to its competition with oxygen at the heme binding site in hemoglobin. The binding affinity for CO is 200 times greater than that for oxygen, meaning that just small amounts of CO dramatically reduces hemoglobin's ability to transport oxygen around the body. The bright red color of the CO-heme complex is the reason that chronic exposure results in the skin adopting a bright red color. The symptoms of CO poisoning include headache, nausea, weakness, and eventually death. When air contains CO levels as low as 0.02%, headache and nausea occur; if the CO concentration is increased to 0.1%, unconsciousness will follow.
Note
Cigarette smoke containing large amounts of carbon monoxide and as a result heavy smokers can have up to 20% of the oxygen-active sites in their blood blocked by CO. However, despite this hazard, cigarettes would be even more deadly if they were not burnt. A key (and often engineered) ingredient in cigarettes is the addictive drug nicotine. Nicotine is an alkaloid (Figure $2$a) as is caffeine (Figure $2$b) that is in coffee and tea. The lethal dose of caffeine is approximately 10 g, which relates to approximately 70 – 100 cups of coffee (assuming a concentration of 100- 150 mg per cup). Alternatively, a lethal dose of caffeine from cola would require approximately 180 – 280 12 oz bottles, each containing 35 – 44 mg. In comparison nicotine has a lethal dose of 50 mg. This means that 12 cigarettes can provide a lethal does if consumed. The only reason smoking 12 cigarettes do not kill immediately is that the majority of the nicotine is burnt in the smoking of the cigarette. If this were not true smokers would be killed before they could develop a habit!
Structure and bonding
The bonding in CO involves 1 σ-bond and 2 sets of π-bonds (Figure $3$). The C-O bond distance in carbon monoxide is 1.1 Å and therefore consistent with a triple bond. In comparison a typical C-O single bond is ca. 1.43 Å, and an average C-O double bond is ca. 1.23 Å. The major absorption band in the infra red spectrum for CO is 2143 cm-1, while for 13CO it is 2099.2 cm-1. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/07%3A_Group_14/7.05%3A_Carbon_Monoxide.txt |
Carbon dioxide (CO2) is the most stable oxide of carbon and is formed from the burning of carbon or carbon containing compounds in air or an excess of oxygen, (7.6.1). For industrial applications it is usually prepared from the decomposition of calcium carbonate (limestone), (7.6.2), rather than separation from combustion products.
$\text{C + O}_2 \rightarrow \text{CO}_2$
$\text{CaCO}_3 \rightarrow \text{CaO + CO}_2$
Phase chemistry of carbon dioxide
Carbon dioxide does not exists as a liquid under normal atmospheric pressure, but solid CO2 (also known as dry ice) sublimes at -78.5 °C (Figure $1$). Dry ice (Figure $2$) is commonly used as a refrigerant for food or biological sample preservation.
Note
When dry ice is placed in water (especially heated) sublimation is accelerated, and a low-sinking dense cloud of fog (smoke-like) is created. This is used in fog machines, at theaters, concerts, haunted houses, and nightclubs for dramatic effects (Figure $3$). Fog from dry ice hovers above the ground unlike other artificial fog machines (that use partial combustion of oil) where the fog rises like smoke.
Supercritical carbon dioxide
As noted above carbon dioxide usually behaves as a gas in air at standard temperature and pressure (STP = 25 °C and 1 atm) or as a solid when frozen. However, if the temperature and pressure are both increased from STP to be at or above the critical point (Figure $1$), carbon dioxide adopts properties midway between a gas and a liquid (Tc = 31.1 °C and Pc = 72.9 atm).
Supercritical CO2 has become an important industrial solvent due to its role in chemical extraction in addition to its low toxicity and environmental impact. In this regard it is seen as a promising green solvent. One of the biggest applications is the decaffeination of coffee and tea without leaving any residue and allowing the caffeine to be separated and used in other beverage products.
Structure and bonding
Carbon dioxide is a linear molecule due to π-localization. The bonding in CO2 involves 2 σ-bond and 2 sets of 3 center π-bonds (Figure $4$). The C-O bond length of 1.2 Å should be compared to the value observed for organic carbonyls (e.g., ketones, esters, aldehydes) of 1.2 – 1.3 Å.
Dissolution and reaction with water
Although CO2 has no dipole moment it is very polar (dielectric constant = 1.60 at 0 °C, 50 atm) and consequently dissolves in polar solvents such as water up to a concentration of 18% (0.04 M). Most of it (+99%) is present as solvated CO2 (Figure $5$), and only ca. 0.2% is reacted to form carbonic acid, (7.6.1), with subsequent equilibria resulting in the formation of bicarbonate (HCO3-) and carbonate (CO32-).
$\text{H}_2\text{O + CO}_2 \rightarrow \text{H}_2\text{CO}_3$
The overall reaction involves a series of equilibria. The first equilibrium is the formation of carbonic acid, (7.6.4). The reaction rates, (7.6.5), are on the magnitude of 1 second (i.e., slow), and as a consequence when carbon dioxide is carried in the body an enzyme is present to speed up the reaction.
$\text{CO}_{\text{2(solv)}} \text{ + H}_2\text{O} \xrightleftharpoons[\text{k}_{\text{H}_2\text{CO}_3}]{\text{k}_{\text{CO}_2}} \text{H}_2\text{CO}_3$
$\text{K = } \dfrac{\text{k}_{\text{H}_2\text{CO}_3}}{\text{k}_{\text{CO}_2}} \text{ = } \dfrac{25}{0.04} \text{ = 600}$
The 2nd equilibrium is as a consequence of first ionization of carbonic acid to form bicarbonate (HCO3-), (7.6.6). In contrast to the first reaction, (7.6.4), this reaction is very fast with a Keq = 1.6 x 10-4 @ 25 °C.
$\text{H}_2\text{CO}_3 \rightleftharpoons \text{HCO}_3^- \text{ + H}^+$
The 3rd equilibrium involves the formation of the carbonate ion (7.6.7), and has a Keq = 4.84 x 10-11. Carbonate (CO32-) is a delocalized ligand, which can act as a mono or bidentate or bridging group.
$\text{HCO}_3^- \rightleftharpoons \text{CO}_3^{2-} \text{ + H}^+$
The formation of carbonic acid is the reason that even in the absence of pollutants (such as SO2) natural rain water is slightly acidic due to dissolved CO2. The equilibrium associated with carbonic acid is also responsible for the buffering of the pH in blood.
Reaction chemistry
Photosynthesis in plants reduces CO2 to organic matter but similar reactions have yet to be developed in non-living systems.
Grignards react readily with carbon dioxide to form the carboxylate, which yields the associated carboxylic acid upon hydrolysis, (7.6.8). Similar reactions occur with other organometallic compounds. In addition, CO2 reacts with alkali metal salts of phenols (phenolates) to yield the hydroxy-carboxylate.
$\text{RMgX + CO}_2 \rightarrow \text{RCO}_2\text{MgX} \xrightarrow{\text{H}_2\text{O}} \text{RCO}_2\text{H + HOMgX}$
A number of complexes of CO2 with transition metals are known in which the coordination can occur via the central carbon (Figure $6$a) or the C=O bond (Figure $6$b). Alternatively, CO2 can bridge two metal centers.
Global warming and carbon dioxide
Global warming is the process of the observed increase in the average temperature of the Earth's near-surface air and oceans since the mid-20th century. Global surface temperature increased 0.74 °C (1.33 °F) between the start and the end of the 20th century (Figure $7$). It is generally agreed that the majority of this temperature increase has occurred since the middle of the 20th century and was caused by increasing concentrations of greenhouse gases resulting from burning fossil fuels (the generation of additional CO2) and deforestation (the loss of a mechanism for the consumption of CO2), see Figure $7$. While it is appreciated that natural phenomena (including solar radiation and volcanoes) produced most of the warming from pre-industrial times, the magnitude of the changes brought on by global industrialization is more significant.
The Earth’s atmosphere has two functions. First, the ozone (O3) in the upper atmosphere screens harmful UV from reaching the surface of the Earth. Second, as solar radiation penetrates the atmosphere a portion of the heat is then retained as a consequence of the CO2 in the atmosphere. It is this process that modulates the surface temperature and provides a stable environment for life. The failure or alteration of either of these processes can have a dramatic effect on the livability of a planet.
Consider the relative position of Venus, Earth, and Mars to the Sun (Figure $8$). The closer a planet is to the sun the greater the UV radiation and the greater the heating of the planet; however, the temperature is also greatly modulated by the atmosphere. Venus has an atmosphere comprising 95% CO2 and has a surface temperature of approximately 450 °C. In contrast, while Mars’ atmosphere is also 95% CO2, it is only 1% as dense as that of Earth’s, and thus the surface temperatures range from 40 °C during the day (due to radiative heating) to –80 °C at night (due to the lack of retained heat because of the thin atmosphere). These should be compared to Earth’s atmosphere which is 0.038% CO2, which allows for the correct amount of heat to be retained to sustain life. Clearly any significant change in the CO2 content of the atmosphere will change the global temperatures of a planet.
Bibliography
• N. Stern, The Stern Review: The Economics of Climate Change, HM Treasury, London.
• R. B. Gupta and J.-J. Shim, Solubility in Supercritical Carbon Dioxide, CRC Press (2006).
• Carbon Dioxide Capture and Storage: Special Report of the Intergovernmental Panel on Climate Change, Cambridge University Press (2005).
7.07: Suboxides of Carbon
Carbon suboxide
Carbon suboxide is the third oxide of carbon, C3O2. It is made from the dehydration of malonic acid, (7.7.1), with P4O10 above 140 °C Like carbon dioxide, the C3O2 molecule is linear, with pπ-pπ bonding.
$\text{CH}_2\text{(CO}_2\text{H)}_2 \rightarrow \text{O=C=C=C=O + 2 H}_2\text{O}$
Gaseous carbon suboxide has an evil smell and while stable at –78 °C it polymerizes at 25 °C. Photolysis of C3O2 yields the unstable C2O. As expected from its synthesis, carbon suboxide reacts slowly with water to form malonic acid, i.e., the reverse of (7.7.1); however, the reaction with stronger nuceophiles such as amines is rapid, (7.7.2).
$\text{C}_3\text{O}_2\text{ + 2 HNR}_2 \rightarrow \text{R}_2\text{NC(O)CH}_2\text{C(O)NR}_2$
Mellitic acid anhydride
The anhydride of mellitic acid (Figure $1$a) may be considered as an oxide of carbon since its chemical formula contains only carbon and oxygen, i.e., C12O9 (Figure $1$b). | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/07%3A_Group_14/7.06%3A_Carbon_Dioxide.txt |
There are two general classes of carbon halides.
1. Homoleptic halides, e.g., CCl4, CCl2F2, C6Cl6, etc.
2. Carbonyl halides, e.g., Cl2C=O.
A summary of some simple carbon halides is given in Table $1$.
Table $1$: Physical properties of simple halogen compounds of carbon.
Compound Mp (°C) Bp (°C) Remarks
CF4 -185 -128 Very stable gas
CCl4 -23 76 Colorless liquid, stable
CBr4 93 190 Pale yellow solid, decomposes upon boiling
CI4 171 - Bright red solid, decomposes prior to boiling, sublimed at low pressure
F2C=O -114 -83 Decomposed by H2O
Cl2C=O -118 8 Phosgene, highly toxic
Br2C=O - 65 Fumes in air
Carbon tetrahalides
The carbon tetrahalides are generally prepared by the direct (thermal) reaction of carbon with the appropriate halogen, (7.8.1); however, specific syntheses are possible for each derivative.
$\text{C + 2 X}_2 \rightarrow \text{CX}_4$
In addition to the direct reaction of fluorine with carbon, CF4 can be prepared from SiC, (7.8.2). The SiF4 side product is removed by passing the reaction mixture through NaOH solution, in which SiF4 reacts to form silicate. The difference in reactivity of SiF4 and CF4 is attributable to the lack of an energetically accessible five-coordinate intermediate required for the associative mechanism.
$\text{SiC + F}_2 \rightarrow \text{SiF}_4\text{ + CF}_4$
Carbon tetrabromide can be obtained by bromination of CH4 with HBr or Br2, or by the reaction of CCl4 with AlBr3, (7.8.3). Carbon tetraiodide (CI4) can be made by the Lewis acid catalyzed halogen exchange reaction, (7.8.4).
$\text{3 CCl}_4\text{ + 4 ALBr}_3 \rightarrow \text{3 CBr}_4\text{ + 4 AlCl}_3$
$\text{CCl}_4\text{ + 4 C}_2\text{H}_5\text{I} \rightarrow \text{CI}_4 \text{ + 4 C}_2\text{H}_5\text{Cl}$
CF4 is very stable. In fact, it is so stable that it does not even react with molten sodium. In contrast to CF4, carbon tetrachloride (CCl4) reacts readily with alkali metals (K and Na) or other strong reducing agents (e.g., F2, Al, Ba, Be, and Zn). While CCl4 is thermodynamically unstable with respect to hydrolysis, it is kinetically stable, and thus finds extensive use as a solvent. Photolysis can result in the transfer of a chloride radical to various substrates. It is also used in the conversion of metal oxides to the chlorides. Carbon tetrabromide (CBr4) is insoluble in water and other polar solvents, but soluble in benzene. Carbon tetraiodide (CI4) decomposes thermally, (7.8.5).
$\text{2 CI}_4 \rightarrow \text{2 I}_2 \text{ + I}_2\text{C=CI}_2$
The decreasing stability of CX4, from fluorine to iodine, is directly related to the C-X bond energy.
Table $2$: Bond energies for carbon-halide bonds.
C-X Bond energy (kJ/mol)
C-F 485
C-Cl 327
C-Br 285
C-I 213
Hazards
Despite its use as a solvent CCl4 has significant hazardous effects. Inhalation of carbon tetrachloride vapor can cause headaches, mental confusion, depression, fatigue, loss of appetite, nausea, vomiting, and coma. The symptoms can take many hours to appear. The vapor and liquid irritate the eyes, and internal irritation, nausea, and vomiting are caused when taken orally. Chronic effects from prolonged inhalation include bronchitis and jaundice, while skin exposure can cause dermatitis.
Carbon tetrabromide is toxic by inhalation, and the vapor is narcotic if taken in high concentrations. As with CCl4, CBr4 can react explosively with alkali metals.
Higher homoleptic halides
Organic compounds that contain only carbon and a halogen are called halocarbons, and these include fluorocarbons and chlorocarbons. The easiest route to fluorocarbons involves the reaction of a hydrocarbon with a high valent fluoride (e.g., CoF3) or the reaction of a chlorocarbon with SbF3. In general, chlorocarbons with sp3 carbon atoms are more stable than those with sp2 carbon centers. The exception to this is aromatic compounds such as C6Cl6.
The physical properties of fluorocarbons range from inert to toxic. Thus, poly(tetrafluoroethylene), (C2F4)n, known by either its acronym (PTFE) or its trade name (Teflon), is chemically inert and has a low coefficient of friction (Table $3$). As a consequence its uses include coatings on armor-piercing bullets (to stop the wear on the gun barrel), laboratory containers and magnetic stirrers, tubing for corrosive chemicals, and thread seal tape in plumbing applications (plumbers tape). A summary of the physical properties of PTFE is given in Table $3$. PTFE is synthesized by the emulsion polymerization of tetrafluoroethylene monomer under pressure through free radical catalyst.
Table $3$: Physical properties of poly(tetrafluoroethylene) (PTFE).
Property Value
Density 2.2 g/cm3
Melting point 327 °C
Young's modulus 0.5 GPa
Yield strength 23 MPa
Coefficient of friction 0.05 - 0.10
Dielectric constant 2.1
Dielectric strength (1 MHz) 60 MV/m
In contrast with PTFE, octafluoroisobutylene, (CF3)2C=CF2, is highly toxic, while perfluorodecahydronapthalene (C10F8, Figure $1$) is used as a blood substitute component.
Mixed halides
Mixed halides are an important class of halocarbon compound. They are synthesized by halide exchange, (7.8.6). The high cost of SbF3 means that the reaction is generally run with an excess of the chloride.
$\text{3 CCl}_4 \text{ + 2 SbF}_3 \rightarrow \text{2 CCl}_2\text{F}_2 \text{ + 2 SbCl}_2$
The ordinary name for mixed carbon halide is halon or Freon, although Freon is actually a Du Pont trademark. A list of selected Freon compounds are given in Table $4$. Halons are non-toxic, non-flammable, and have no odor. However, it is their very lack of reactivity that has caused a problem.
Table $4$: Selected Freons and their applications.
Freon Formula Uses
12 CCl2F2 Refrigerant
11 CCl3F Refrigerant
114 ClF2C-CClF2 Refrigerant
113 Cl3C-CF3 Solvent
13B1 CBrF3 Fire extinguisher
1211 CBrClF2 Fire extinguisher
Environmental impact of chlorofluorcarbon compounds (CFCs)
Chlorofluorcarbon compounds (CFCs) are very stable and are not degraded in the environment. As a consequence they are transported to the stratosphere where they decomposed upon photolysis, (7.8.7). The resulting chloride radical is a catalyst for the decomposition of ozone, (7.8.8), as well as a catalyst for the reaction of ozone with molecular oxygen, (7.8.9).
$\text{CCl}_2\text{F}_2\text{ + h}\nu \rightarrow \text{CClF}_2\text{ + Cl}\cdot$
$\text{2 O}_3 \xrightarrow{\text{Cl}\cdot} \text{3 O}_2$
$\text{O}_3 \text{ + O} \xrightarrow{\text{Cl}\cdot} \text{2 O}_2$
The widespread use of CFCs as refrigerants and propellants meant that by 1986 there were 2.5 billion pounds of CFC being liberated to the atmosphere. This was equivalent to 1/2. lb per person on the planet. Since the ozone layer provides the vital protection to life on the Earth’s surface from high energy UV radiation the release of CFC (along with other chemicals) caused a dramatic change in the ozone layer, including the increase in the polar hole in the ozone layer. As a result of the EU called for a complete ban of CFCs (which was followed by other countries). In their place new chemicals with similar refrigerant properties were developed. These compounds contained C-H bonds (e.g., C2HCl2F3 and C2H3Cl2F) that are readily broken in the lower atmosphere, thus limiting the transport to the stratosphere.
Carbonyl halides
All the carbonyl halides (X2C=O, X = F, Cl, Br, I) are known (Table $1$). Phosgene (Cl2C=O) was first synthesized by John Davy (Figure $2$) in 1812 by exposing a mixture of carbon monoxide and chlorine to sunlight, (7.8.10). He named it phosgene from the Greek, phos (light) and gene (born), in reference to use of light to promote the reaction. The fluoride is also prepared by the reaction of carbon monoxide with the halogen, while the bromide is prepared by the partial hydrolysis of CBr4 with sulfuric acid.
$\text{CO + Cl}_2 \rightarrow \text{Cl}_2\text{C=O}$
The synthesis of isocyanates from alkyl or aryl amines illustrates the electrophilic character of phosgene and its ability to introduce the equivalent of "CO2+", (7.8.11). This reaction is conducted in the presence of a base such as pyridine that absorbs the hydrogen chloride. Phosgene may also be used to produce acyl chlorides from carboxylic acids, (7.8.12). However, thionyl chloride is more commonly and more safely used in this reaction.
$\text{RNH}_2 \text{ + Cl}_2\text{C=O} \rightarrow \text{RN=C=O + 2 HCl}$
$\text{RCO}_2\text{H + Cl}_2\text{C=O} \rightarrow \text{RC(O)Cl + HCl + CO}_2$
Phosgene as a weapon of war
Phosgene is a toxic gas with the smell of “new-mown hay” and was used in chemical warfare during the First World War (Figure $3$) where it was a more potent weapon than chlorine. While chlorine was potentially deadly it caused the victim to violently cough and choke (the bodies natural defense to limiting inhalation), in contrast, phosgene caused much less coughing with the result that more of it was inhaled. Phosgene often had a delayed effect; apparently healthy soldiers were taken down with phosgene gas poisoning up to 48 hours after inhalation. A fatal dose of phosgene eventually led to shallow breathing and retching, pulse up to 120, an ashen face and the discharge of four pints of yellow liquid from the lungs each hour for the 48 of the drowning spasms.
Although phosgene’s boiling point (7.6 °C) meant that is was a vapor, the so-called "white star" mixture of phosgene and chlorine was commonly used on the Somme, because the chlorine supplied the necessary vapor with which to carry the phosgene. A summary of the casualties inflicted by chemical warfare agents during the Great War is shown in Table $5$.
Table $5$: Casualties from gas attacks during the First World War (including chlorine, phosgene, and mustard gas). British Empire includes troops from United Kingdom, Australia, Canada, India, New Zealand, and South Africa.
Country Total casualties Deaths
Russia 419,340 56,000
Germany 200,000 9,000
France 190,000 8,000
British Empire 188,706 8,109
Austria-Hungary 100,000 3,000
USA 72,807 1,462
Italy 60,000 4,627
Others 10,000 1,000 | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/07%3A_Group_14/7.08%3A_Carbon_Halides.txt |
An understanding of the differences between carbon and silicon is important in understanding the relative chemistry of these Group 14 elements.
Size
As expected silicon is larger than carbon due to the presence of a second shell: i.e., C = 1s2 2s2 2p2 while Si = 1s2 2s2 2p6 3s2 3p2. A comparison of the relative sizes of carbon and silicon are given in Table \(1\).
Table \(1\): Atomic, covalent, and van der Waals radii of carbon and silicon.
Element Atomic radius (Å) Covalent radius sp3 (Å) van der Waal radius (Å)
C 0.70 0.75 1.70
Si 1.10 1.14 2.10
Covalent and van der Waal radii from Royal Society of Chemistry Online Periodic TableWebelements has a more detailed discussion of all three types of radii, the atomic radii quoted here are empirical.
Coordination number
Carbon is known to have a coordination number of 2, 3, and 4 in its compounds depending on the hybridization. A coordination number of 1 can also be considered for CO and CN-. Four-coordinate carbon may be considered to be coordinatively saturated. In contrast, in the absence of overwhelming steric bulk, silicon is observed to have coordination numbers of 3, 4, 5, and 6. Examples of five and six-coordinate silicon include Si(acac)2Cl and SiF62-, respectively. Coordination numbers of higher than 4 have been ascribed to the use of low-lying d orbitals; however, calculations show these are not significant. Instead, hypervalent silicon is better described by the formation of 3-center molecular orbitals, e.g., Figure \(1\).
Note
A hypervalent molecule is a molecule that contains one or more typical elements (Group 1, 2, 13-18) formally bearing more than eight electrons in their valence shells.
Electronegativity
The electronegativities of silicon and carbon are given in Table along with hydrogen. Since carbon is more electronegative than hydrogen the C-H bond is polarized towards carbon resulting in a more protic hydrogen (Figure \(2\)a). In contrast, the lower electronegativity of silicon results in a more hydridic hydrogen (Figure \(2\)b). This difference is reflected in the reaction chemistry of SiH4 versus CH4.
Table \(2\): Selected Pauling electronegativity values. Royal Society of Chemistry Online Periodic Table
Element Pauling scale
C 2.55
H 2.20
Si 1.90
Bond energies
The E-E and E-O bond energies for carbon and silicon are given in Table \(3\). The bond energy for a C-C bond is slightly greater than for a C-O bond, while the Si-O bond is significantly stronger than the Si-Si bond. This difference is reflected in the chemistry of silicon versus carbon compounds. The chemistry of carbon is dominated by catenation: the ability of a chemical element to form a long chain-like structure via a series of covalent bonds. Although silicon does form Si-Si bonds, they are far more reactive than their C-C analogs, and polymers of silicon are predominantly comprised of Si-O chains (as a result of the very strong bond).
Table \(3\): Selected bond energies for carbon and silicon. Royal Society of Chemistry Online Periodic Table
Element E-E bond energy (kJ/mol) E-O bond energy (kJ/mol)
C 345.6 357.7
Si 222 462
Multiple bonds
While unsaturated compounds for carbon (i.e., alkenes and alkynes) are common, the analogous silicon compounds (disilenes) were only reported in 1981, and disilynes in 2004. The Si=Si double bond lengths are 2.14 - 2.29 Å which is 5 - 10% shorter than the Si-Si single bond lengths. This bond shortening is less than ca. 13% in carbon compounds.
Note
The traditional lack of multiple bonds for the Period 3 elements and lower led to the formulation of the double bond rule which states that chemical elements with a principal quantum number greater than 2 do not form multiple bonds (e.g., double bonds and triple bonds) with themselves or with other elements. This rule was made obsolete starting from 1981 with the discovery of silicon and phosphorus double bonds. Double bonds that would ordinarily not form can be stabilized with proper functional groups through kinetic stabilization, i.e., either electronically or sterically.
Bibliography
• R. West, M. J. Fink, and J. Michl, Science, 1981, 214, 1343.
• A. Sekiguchi, R. Kinjo, and M. Ichinohe, Science, 2004, 305, 1755. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/07%3A_Group_14/7.09%3A_Comparison_Between_Silicon_and_Carbon.txt |
Introduction
The synthesis and purification of bulk polycrystalline semiconductor material represents the first step towards the commercial fabrication of an electronic device. This polycrystalline material is then used as the raw material for the formation of single crystal material that is processed to semiconductor wafers. The strong influence on the electric characteristics of a semiconductors exhibited by small amounts of some impurities requires that the bulk raw material be of very high purity (> 99.9999%). Although some level of purification is possible during the crystallization process it is important to use as high a purity starting material as possible.
Following oxygen (46%), silicon (L. silicis flint) is the most abundant element in the earth's crust (28%). However, silicon does not occur in its elemental form, but as its oxide (SiO2) or as silicates. Sand, quartz, amethyst, agate, flint, and opal are some of the forms in which the oxide appears. Granite, hornblende, asbestos, feldspar, clay and mica, etc. are a few of the numerous silicate minerals. With such boundless supplies of the raw material, the costs associated with the production of bulk silicon is not one of abstraction and conversion of the oxide(s), but of purification of the crude elemental silicon. While 98% elemental silicon, known as metallurgical-grade silicon (MGS), is readily produced on a large scale, the requirements of extreme purity for electronic device fabrication require additional purification steps in order to produce electronic-grade silicon (EGS). Electronic-grade silicon is also known as semiconductor-grade silicon (SGS). In order for the purity levels to be acceptable for subsequent crystal growth and device fabrication, EGS must have carbon and oxygen impurity levels less than a few parts per million (ppm), and metal impurities at the parts per billion (ppb) range or lower. Table $1$ and Table $2$ give typical impurity concentrations in MGS and EGS, respectively. Besides the purity, the production cost and the specifications must meet the industry desires.
Table $1$: Typical impurity concentrations found in metallurgical-grade silicon (MGS).
Element Concentration (ppm) Element Concentration (ppm)
aluminum 1000-4350 manganese 50-120
boron 40-60 molybdenum < 20
calcium 245-500 nickel 10-105
chromium 50-200 phosphorus 20-50
copper 15-45 titanium 140-300
iron 1550-6500 vanadium 50-250
magnesium 10-50 zirconium 20
Table $2$: Typical impurity concentrations found in electronic-grade silicon (EGS).
Element Concentration (ppb) Element Concentration (ppb)
arsenic < 0.001 gold < 0.00001
antimony < 0.001 iron 0.1-1.0
boron ≤ 0.1 nickel 0.1-0.5
carbon 100-1000 oxygen 100-400
chromium < 0.01 phosphorus ≤ 0.3
cobalt 0.001 silver 0.001
copper 0.1 zinc < 0.1
Metallurgical-grade silicon (MGS)
The typical source material for commercial production of elemental silicon is quartzite gravel; a relatively pure form of sand (SiO2). The first step in the synthesis of silicon is the melting and reduction of the silica in a submerged-electrode arc furnace. An example of which is shown schematically in Figure $1$, along with the appropriate chemical reactions. A mixture of quartzite gravel and carbon are heated to high temperatures (ca. 1800 °C) in the furnace. The carbon bed consists of a mixture of coal, coke, and wood chips. The latter providing the necessary porosity such that the gases created during the reaction (SiO and CO) are able to flow through the bed.
The overall reduction reaction of SiO2 is expressed in (7.10.1), however, the reaction sequence is more complex than this overall reaction implies, and involves the formation of SiC and SiO intermediates. The initial reaction between molten SiO2 and C, (7.10.2), takes place in the arc between adjacent electrodes, where the local temperature can exceed 2000 °C. The SiO and CO thus generated flow to cooler zones in the furnace where SiC is formed, (7.10.3), or higher in the bed where they reform SiO2 and C, (7.10.2). The SiC reacts with molten SiO2, (7.10.4), producing the desired silicon along with SiO and CO. The molten silicon formed is drawn-off from the furnace and solidified.
$\text{SiO}_2\text{(liquid) + 2 C(solid)} \rightarrow \text{Si(liquid) + 2 CO(gas)}$
$\text{SiO}_2\text{ + 2 C} \xrightleftharpoons[\text{<1600 °C}]{\text{>1700 °C}} \text{SiO + CO}$
$\text{SiO + 2C} \rightarrow \text{SiC + CO (1600 - 1700 °C)}$
$\text{SiC _ SiO}_2 \rightarrow \text{Si + SiO + CO}$
The as-produced MGS is approximately 98-99% pure, with the major impurities being aluminum and iron (Table $1$), however, obtaining low levels of boron impurities is of particular importance, because it is difficult to remove and serves as a dopant for silicon. The drawbacks of the above process are that it is energy and raw material intensive. It is estimated that the production of one metric ton (1,000 kg) of MGS requires 2500 - 2700 kg quartzite, 600 kg charcoal, 600 - 700 kg coal or coke, 300 - 500 kg wood chips, and 500,000 kWh of electric power. Currently, approximately 500,000 metric tons of MGS are produced per year, worldwide. Most of the production (ca. 70%) is used for metallurgical applications (e.g., aluminum-silicon alloys are commonly used for automotive engine blocks) from whence its name is derived. Applications in a variety of chemical products such as silicone resins account for about 30%, and only 1% or less of the total production of MGS is used in the manufacturing of high-purity EGS for the electronics industry. The current worldwide consumption of EGS is approximately 5 x 106 kg per year.
Electronic-grade silicon (EGS)
Electronic-grade silicon (EGS) is a polycrystalline material of exceptionally high purity and is the raw material for the growth of single-crystal silicon. EGS is one of the purest materials commonly available, see Table $2$. The formation of EGS from MGS is accomplished through chemical purification processes. The basic concept of which involves the conversion of MGS to a volatile silicon compound, which is purified by distillation, and subsequently decomposed to re-form elemental silicon of higher purity (i.e., EGS). Irrespective of the purification route employed, the first step is physical pulverization of MGS followed by its conversion to the volatile silicon compounds.
A number of compounds, such as monosilane (SiH4), dichlorosilane (SiH2Cl2), trichlorosilane (SiHCl3), and silicon tetrachloride (SiCl4), have been considered as chemical intermediates. Among these, SiHCl3 has been used predominantly as the intermediate compound for subsequent EGS formation, although SiH4 is used to a lesser extent. Silicon tetrachloride and its lower chlorinated derivatives are used for the chemical vapor deposition (CVD) growth of Si and SiO2. The boiling points of silane and its chlorinated products (Table $3$) are such that they are conveniently separated from each other by fractional distillation.
Table $3$: Boiling points of silane and chlorosilanes at 760 mmHg (1 atmosphere).
Compound Boiling point (°C)
SiH4 -112.3
SiH3Cl -30.4
SiH2Cl2 8.3
SiHCl3 31.5
SiCl4 57.6
The reasons for the predominant use of SiHCl3 in the synthesis of EGS are as follows:
1. SiHCl3 can be easily formed by the reaction of anhydrous hydrogen chloride with MGS at reasonably low temperatures (200 - 400 °C);
2. it is liquid at room temperature so that purification can be accomplished using standard distillation techniques;
3. it is easily handled and if dry can be stored in carbon steel tanks;
4. its liquid is easily vaporized and, when mixed with hydrogen it can be transported in steel lines without corrosion;
5. it can be reduced at atmospheric pressure in the presence of hydrogen;
6. its deposition can take place on heated silicon, thus eliminating contact with any foreign surfaces that may contaminate the resulting silicon; and
7. it reacts at lower temperatures (1000 - 1200 °C) and at faster rates than does SiCl4.
Chlorosilane (Seimens) process
Trichlorosilane is synthesized by heating powdered MGS with anhydrous hydrogen chloride (HCl) at around 300 °C in a fluidized-bed reactor, (7.10.5).
$\text{Si(solid_ + 3 HCl(gas)} \xrightleftharpoons[\text{>900 °C}]{\text{ca. 300 °C}} \text{SiHCl}_3\text{(vapor) + H}_2\text{(gas)}$
Since the reaction is actually an equilibrium and the formation of SiHCl3 highly exothermic, efficient removal of generated heat is essential to assure a maximum yield of SiHCl3. While the stoichiometric reaction is that shown in (7.10.5), a mixture of chlorinated silanes is actually prepared which must be separated by fractional distillation, along with the chlorides of any impurities. In particular iron, aluminum, and boron are removed as FeCl3 (b.p. = 316 °C), AlCl3 (m.p. = 190 °C subl.), and BCl3 (b.p. = 12.65 °C), respectively. Fractional distillation of SiHCl3 from these impurity halides result in greatly increased purity with a concentration of electrically active impurities of less than 1 ppb.
EGS is prepared from purified SiHCl3 in a chemical vapor deposition (CVD) process similar to the epitaxial growth of Si. The high-purity SiHCl3 is vaporized, diluted with high-purity hydrogen, and introduced into the Seimens deposition reactor, shown schematically in Figure $2$. Within the reactor, thin silicon rods called slim rods (ca. 4 mm diameter) are supported by graphite electrodes. Resistance heating of the slim rods causes the decomposition of the SiHCl3 to yield silicon, as described by the reverse reaction shown in (7.10.5).
The shift in the equilibrium from forming SiHCl3 from Si at low temperature, to forming Si from SiHCl3 at high temperature is as a consequence of the temperature dependence, (7.10.6), of the equilibrium constant, (7.10.7) where ρ = partial pressure, for (7.10.5). Since the formation of SiHCl3 is exothermic, i.e., ΔH < 0, an increase in the temperature causes the partial pressure of SiHCl3 to decrease. Thus, the Siemens process is typically run at ca. 1100 °C, while the reverse fluidized bed process is carried out at 300 °C.
$\text{lnK}_{\text{p}} \text{ = } \dfrac{ \text{-}\Delta\text{H}}{\text{RT}}$
$\text{K}_{\text{p}} \text{ = } \dfrac{^{\rho}\text{SiHCl}_3 \text{ } ^{\rho}\text{H}_2}{^{\rho}\text{HCl}}$
The slim rods act as a nucleation point for the deposition of silicon, and the resulting polycrystalline rod consists of columnar grains of silicon (polysilicon) grown perpendicular to the rod axis. Growth occurs at less than 1 mm per hour, and after deposition for 200 to 300 hours high-purity (EGS) polysilicon rods of 150 - 200 mm in diameter are produced. For subsequent float-zone refining the polysilicon EGS rods are cut into long cylindrical rods. Alternatively, the as-formed polysilicon rods are broken into chunks for single crystal growth processes, for example Czochralski melt growth.
In addition to the formation of silicon, the HCl coproduct reacts with the SiHCl3 reactant to form silicon tetrachloride (SiCl4) and hydrogen as major byproducts of the process, (7.10.8). This reaction represents a major disadvantage with the Seimens process: poor efficiency of silicon and chlorine consumption. Typically, only 30% of the silicon introduced into CVD reactor is converted into high-purity polysilicon.
$\text{HCl + SiHCl}_3 \rightarrow \text{SiCl}_4\text{ + H}_2$
In order to improve efficiency the HCl, SiCl4, H2, and unreacted SiHCl3 are separated and recovered for recycling. Figure $3$ illustrates the entire chlorosilane process starting with MGS and including the recycling of the reaction byproducts to achieve high overall process efficiency. As a consequence, the production cost of high-purity EGS depends on the commercial usefulness of the byproduct, SiCl4. Additional disadvantages of the Seimens process are derived from its relatively small batch size, slow growth rate, and high power consumption. These issues have lead to the investigation of alternative cost efficient routes to EGS.
Silane process
An alternative process for the production of EGS that has begun to receive commercial attention is the pyrolysis of silane (SiH4). The advantages of producing EGS from SiH4 instead of SiHCl3 are potentially lower costs associated with lower reaction temperatures, and less harmful byproducts. Silane decomposes < 900 °C to give silicon and hydrogen, (7.10.9).
$\text{SiH}_4\text{(vapor)} \rightarrow \text{Si(solid) + 2 H}_2\text{(gas)}$
Silane may be prepared by a number of routes, each having advantages with respect to purity and production cost. The simplest process involves the direct reaction of MGS powders with magnesium at 500 °C in a hydrogen atmosphere, to form magnesium silicide (Mg2Si). The magnesium silicide is then reacted with ammonium chloride in liquid ammonia below 0 °C, (7.10.10).
$\text{Mg}_2\text{Si + 4 NH}_4\text{Cl} \rightarrow \text{SiH}_4\text{ + 2 MgCl}_2\text{ + 5 NH}_3$
This process is ideally suited to the removal of boron impurities (a p-type dopant in Si), because the diborane (B2H6) produced during the reaction forms the Lewis acid-base complex, H3B(NH3), whose volatility is sufficiently lower than SiH4, allowing for the purification of the latter. It is possible to prepare EGS with a boron content of ≤ 20 ppt using SiH4 synthesized in this manner. However, phosphorus (another dopant) in the form of PH3 may be present as a contaminant requiring subsequent purification of the SiH4.
Alternative routes to SiH4 involve the chemical reduction of SiCl4 by either lithium hydride, (7.10.11), lithium aluminum hydride, (7.10.12), or via hydrogenation in the presence of elemental silicon, (7.10.13) - (7.10.16). The hydride reduction reactions may be carried-out on relatively large scales (ca. 50 kg), but only batch processes. In contrast, Union Carbide has adapted the hydrogenation to a continuous process, involving disproportionation reactions of chlorosilanes, (7.10.14) - (7.10.16), and the fractional distillation of silane, Table $3$.
$\text{SiCl}_4\text{ + LiH} \rightarrow \text{SiH}_4\text{ + 4 LiCl}$
$\text{SiCl}_4\text{ + LiAlH}_4 \rightarrow \text {SiH}_4 \text{ + LiCl + AlCl}_4$
$\text{SiCl}_4\text{ + 2 H}_2\text{ + Si(98%)} \rightarrow \text{4 SiHCl}_3$
$\text{2 SiHCl}_3 \rightarrow \text{SiH}_2\text{Cl}_2\text{ + 2 SiHCl}_3$
$\text{3 SiH}_2\text{Cl}_2 \rightarrow \text{SiH}_2\text{Cl + 2 SiHCl}_3$
$\text{2 SiH}_3\text{Cl} \rightarrow \text{SiH}_4\text{ + SiH}_2\text{Cl}_2$
Pyrolysis of silane on resistively heated polysilicon filaments at 700 - 800 °C yields polycrystalline EGS. As noted above, the EGS formed has remarkably low boron impurities compared with material prepared from trichlorosilane. Moreover, the resulting EGS is less contaminated with transition metals from the reactor container because SiH4 decomposition does not cause as much of a corrosion problem as is observed with halide precursor compounds.
Granular polysilicon deposition
Both the chlorosilane (Seimens) and silane processes result in the formation of rods of EGS. However, there has been increased interest in the formation of granular polycrystalline EGS. This process was developed in 1980’s, and relies on the decomposition of SiH4 in a fluidized-bed deposition reactor to produce free-flowing granular polysilicon.
Tiny silicon particles are fluidized in a SiH4/H2 flow, and act as seed crystal onto which polysilicon deposits to form free-flowing spherical particles. The size distribution of the particles thus formed is over the range from 0.1 to 1.5 mm in diameter with an average particle size of 0.7 mm. The fluidized-bed seed particles are originally made by grinding EGS in a ball (or hammer) mill and leaching the product with acid, hydrogen peroxide, and water. This process is time-consuming and costly, and tended to introduce undesirable impurities from the metal grinders. In a new method, large EGS particles are fired at each other by a high-speed stream of inert gas and the collision breaks them down into particles of suitable size for a fluidized bed. This process has the main advantage that it introduces no foreign materials and requires no leaching or other post purification.
The fluidized-bed reactors are much more efficient than traditional rod reactors as a consequence of the greater surface area available during CVD growth of silicon. It has been suggested that fluidized-bed reactors require 1/5 to 1/10 the energy, and half the capital cost of the traditional process. The quality of fluidized-bed polysilicon has proven to be equivalent to polysilicon produced by the conventional methods. Moreover, granular EGS in a free-flowing form, and with high bulk density, enables crystal growers to obtain the high, reproducible production yields out of each crystal growth run. For example, in the Czochralski crystal growth process, crucibles can be quickly and easily filled to uniform loading with granular EGS, which typically exceed those of randomly stacked polysilicon chunks produced by the Siemens silane process.
Zone refining
The technique of zone refining is used to purify solid materials and is commonly employed in metallurgical refining. In the case of silicon may be used to obtain the desired ultimate purity of EGS, which has already been purified by chemical processes. Zone refining was invented by Pfann, and makes use of the fact that the equilibrium solubility of any impurity (e.g., Al) is different in the solid and liquid phases of a material (e.g., Si). For the dilute solutions, as is observed in EGS silicon, an equilibrium segregation coefficient (k0) is defined by k0 = Cs/Cl, where Cs and Cl are the equilibrium concentrations of the impurity in the solid and liquid near the interface, respectively.
If k0 is less than 1 then the impurities are left in the melt as the molten zone is moved along the material. In a practical sense a molten zone is established in a solid rod. The zone is then moved along the rod from left to right. If k0 < 1 then the frozen part left on the trailing edge of the moving molten zone will be purer than the material that melts in on the right-side leading edge of the moving molten zone. Consequently the solid to the left of the molten zone is purer than the solid on the right. At the completion of the first pass the impurities become concentrated to the right of the solid sample. Repetition of the process allows for purification to exceptionally high levels. Table $4$ lists the equilibrium segregation coefficients for common impurity and dopant elements in silicon; it should be noted that they are all less than 1.
Table $4$: Segregation coefficients for common impurity and dopant elements in silicon.
Element k0 Element k0
aluminum 0.002 iron 8 x 10-6
boron 0.8 oxygen 0.25
carbon 0.07 phosphorus 0.35
copper 4 x 10-6 antimony 0.023
Bibliography
• K. G. Baraclough, K. G., in The Chemistry of the Semiconductor Industry, Eds. S. J. Moss and A. Ledwith, Blackie and Sons, Glasgow, Scotland (1987).
• L. D. Crossman and J. A. Baker, Semiconductor Silicon 1977, Electrochem. Soc., Princeton, New Jersey (1977).
• W. C. O’Mara, Ed. Handbook of Semiconductor Silicon Technology, Noyes Pub., New Jersey (1990).
• W. G. Pfann, Zone Melting, John Wiley & Sons, New York, (1966).
• F. Shimura, Semiconductor Silicon Crystal Technology, Academic Press (1989). | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/07%3A_Group_14/7.10%3A_Semiconductor_Grade_Silicon.txt |
Note
This module was developed as part of the Rice University course CHEM-496: Chemistry of Electronic Materials. This module was prepared with the assistance of Andrea Keys.
Introduction
In the fabrication of integrated circuits (ICs), the oxidation of silicon is essential, and the production of superior ICs requires an understanding of the oxidation process and the ability to form oxides of high quality. Silicon dioxide has several uses:
1. Serves as a mask against implant or diffusion of dopant into silicon.
2. Provides surface passivation.
3. Isolates one device from another (dielectric isolation).
4. Acts as a component in MOS structures.
5. Provides electrical isolation of multi-level metallization systems.
Methods for forming oxide layers on silicon have been developed, including thermal oxidation, wet anodization, chemical vapor deposition (CVD), and plasma anodization or oxidation. Generally, CVD is used when putting the oxide layer on top of a metal surface, and thermal oxidation is used when a low-charge density level is required for the interface between the oxide and the silicon surface.
Oxidation of silicon
Silicon's surface has a high affinity for oxygen and thus an oxide layer rapidly forms upon exposure to the atmosphere. The chemical reactions which describe this formation are:
$\text{Si}_{\text{(s)}} \text{ + O}_{\text{2(g)}} \rightarrow \text{SiO}_{\text{2(s)}}$
$\text{Si}_{\text{(s)}} \text{ + 2 H}_2\text{O}_{\text{(g)}} \rightarrow \text{SiO}_{\text{2(s)}} \text{ + 2 H}_{\text{2(g)}}$
In the first reaction a dry process is utilized involving oxygen gas as the oxygen source and the second reaction describes a wet process which uses steam. The dry process provides a "good" silicon dioxide but is slow and mostly used at the beginning of processing. The wet procedure is problematic in that the purity of the water used cannot be guaranteed to a suitable degree. This problem can be easily solved using a pyrogenic technique which combines hydrogen and oxygen gases to form water vapor of very high purity. Maintaining reagents of high quality is essential to the manufacturing of integrated circuits, and is a concern which plagues each step of this process.
The formation of the oxide layer involves shared valence electrons between silicon and oxygen, which allows the silicon surface to rid itself of "dangling" bonds, such as lone pairs and vacant orbitals. These vacancies create mid-gap states between the valence and conduction bands, which prevents the desired band gap of the semiconductor. The Si-O bond strength is covalent (strong), and so can be used to achieve the loss of mid-gap states and passivate the surface of the silicon.
The oxidation of silicon occurs at the silicon-oxide interface and consists of four steps:
1. Diffusive transport of oxygen across the diffusion layer in the vapor phase adjacent to the silicon oxide-vapor interface.
2. Incorporation of oxygen at the outer surface into the silicon oxide film.
3. Diffusive transport across the silicon oxide film to its interface with the silicon lattice.
4. Reaction of oxygen with silicon at this inner interface.
As the Si-SiO2 interface moves into the silicon its volume expands, and based upon the densities and molecular weights of Si and SiO2, 0.44 Å Si is used to obtain 1.0 Å SiO2.
Pre-oxidation cleaning
The first step in oxidizing a surface of silicon is the removal of the native oxide which forms due to exposure to open air. This may seem redundant to remove an oxide only to put on another, but this is necessary since uncertainty exists as to the purity of the oxide which is present. The contamination of the native oxide by both organic and inorganic materials (arising from previous processing steps and handling) must be removed to prevent the degradation of the essential electrical characteristics of the device. A common procedure uses a H2O-H2O2-NH4OH mixture which removes the organics present, as well as some group I and II metals. Removal of heavy metals can be achieved using a H2O-H2O2-HCl mixture, which complexes with the ions which are formed. After removal of the native oxide, the desired oxide can be grown. This growth is useful because it provides: chemical protection, conditions suitable for lithography, and passivation. The protection prevents unwanted reactions from occurring and the passivation fills vacancies of bonds on the surface not present within the interior of the crystal. Thus the oxidation of the surface of silicon fulfills several functions in one step.
Thermal oxidation
The growth of oxides on a silicon surface can be a particularly tedious process, since the growth must be uniform and pure. The thickness wanted usually falls in the range 50 - 500 Å, which can take a long time and must be done on a large scale. This is done by stacking the silicon wafers in a horizontal quartz tube while the oxygen source flows over the wafers, which are situated vertically in a slotted paddle (boat). This procedure is performed at 1 atm pressure, and the temperature ranges from 700 to 1200 °C, being held to within ±1 °C to ensure uniformity. The choice of oxidation technique depends on the thickness and oxide properties required. Oxides that are relatively thin and those that require low charge at the interface are typically grown in dry oxygen. When thick oxides are required (> 0.5 mm) are desired, steam is the source of choice. Steam can be used at wide range of pressures (1 atm to 25 atm), and the higher pressures allow thick oxide growth to be achieved at moderate temperatures in reasonable amounts of time.
The thickness of SiO2 layers on a Si substrate is readily determined by the color of the film. Table $1$ provides a guidline for thermal grown oxides.
Table $1$: Color chart for thermally grown SiO2 films observed under daylight fluorescent lighting.
Film thickness (μm) Color Film thickness (μm) Color
0.05 tan 0.63 violet-red
0.07 brown 0.68 "bluish"
0.10 dark violet to red-violet 0.72 blue-green to gree
0.12 royal blue 0.77 "yellowish"
0.15 light blue to metallic blue 0.80 orange
0.17 metallic to light yellow-green 0.82 salmon
0.20 light gold 0.85 light red-violet
0.22 gold 0.86 violet
0.25 orange to melon 0.87 blue violet
0.27 red-violet 0.89 blue
0.30 blue to violet blue 0.92 blue-green
0.31 blue 0.95 yellow-green
0.32 blue to blue-green 0.97 yellow
0.34 light green 0.99 orange
0.35 green to yellow-green 1.00 carnation pink
0.36 yellow-green 1.02 violet red
0.37 green-yellow 1.05 red-violet
0.39 yellow 1.06 violet
0.41 light orange 1.07 blue-violet
0.42 carnation pink 1.10 green
0.44 violet-red 1.11 yellow-green
0.46 red-violet 1.12 green
0.47 violet 1.18 violet
0.48 blue-violet 1.19 red-violet
0.49 blue 1.21 violet-red
0.50 blue green 1.24 carnation pink to salmon
0.52 green 1.25 orange
0.54 yellow-green 1.28 "yellowish"
0.56 green-yellow 1.32 sky blue to green-blue
0.57 "yellowish" 1.40 orange
0.58 light orange to pink 1.46 blue-violet
0.60 carnation pink 1.50 blue
High pressure oxidation
High pressure oxidation is another method of oxidizing the silicon surface which controls the rate of oxidation. This is possible because the rate is proportional to the concentration f the oxide, which in turn is proportional to the partial pressure of the oxidizing species, according to Henry's law, (7.11.3), where C is the equilibrium concentration of the oxide, H is Henry's law constant, and pO is the partial pressure of the oxidizing species.
$\text{C = H}_{\text{(pO)}}$
This approach is fast, with a rate of oxidation ranging from 100 to 1000 mm/h, and also occurs at a relatively low temperature. It is a useful process, preventing dopants from being displaced and also forms a low number of defects, which is most useful at the end of processing.
Plasma oxidation
Plasma oxidation and anodization of silicon is readily accomplished by the use of activated oxygen as the oxidizing species. The highly reactive oxygen is formed within an electrical discharge or plasma. The oxidation is carried out in a low pressure (0.05 - 0.5 Torr) chamber, and the the plasma is produced either by a DC electron source or a high-frequency discharge. In simple plasma oxidation the sample (i.e., the silicon wafer) is held at ground potential. In contrast, aniodization systems usually have a DC bias between the sample and an electrode with the sample biased positively with respect to the cathode. Platinum electrodes are commonly used as the cathodes.
There have been at least 34 different reactions reported to occur in an oxygen plasma, however, the vast majority of these are inconsequential with respect to the formation of active species. Furthermore, many of the potentially active species are sufficiently short lived that it is unlikely that they make a significant contribution. The primary active species within the oxygen plasma are undoubtedly O- and O2+. Both being produced in near equal quantities, although only the former is relevant to plasma aniodization. While these species may be active with respect to surface oxidation, it is more likely that an electron transfer occurs from the semiconductor surface yields activated oxygen species, which are the actual reactants in the oxidation of the silicon.
The significant advatage of plasma processes is that while the electron temperature of the ionized oxygen gas is in excess of 10,000 K, the thermal temperatures required are significantly lower than required for the high pressure method, i.e., < 600 °C. The advantages of the lower reaction temperatures include: the minimization of dopant diffusion and the impediment of the generation of defects. Despite these advantages there are two primary disadvantages of any plasma based process. First, the high electric fields present during the processes cause damage to the resultant oxide, in particular, a high density of interface traps often result. However, post annealing may improve film quality. Second, the growth rates of plasma oxidation are low, typically 1000 Å/h. This growth rate is increased by about a factor of 10 for plasma aniodization, and further improvements are observed if 1 - 3% chlorine is added to the oxygen source.
Masking
A selective mask against the diffusion of dopant atoms at high temperatures can be found in a silicon dioxide layer, which can prove to be very useful in integrated circuit processing. A predeposition of dopant by ion implantation, chemical diffusion, or spin-on techniques typically results in a dopant source at or near the surface of the oxide. During the initial high-temperature step, diffusion in the oxide must be slow enough with respect to diffusion in the silicon that the dopants do not diffuse through the oxide in the masked region and reach the silicon surface. The required thickness may be determined by experimentally measuring, at a particular temperature and time, the oxide thickness necessary to prevent the inversion of a lightly doped silicon substrate of opposite conductivity. To this is then added a safety factor, with typical total values ranging from 0.5 to 0.7 mm. The impurity masking properties result when the oxide is partially converted into a silica impurity oxide "glass" phase, and prevents the impurities from reaching the SiO2-Si interface.
Bibliography
• M. M. Atalla, in Properties of Elemental and Compound Semiconductors, Ed. H. Gatos, Interscience: New York (1960).
• S. K. Ghandhi, VLSI Fabrication Principles, Silicon and Gallium Arsenide, Wiley, Chichester, 2nd Ed. (1994).
• S. M. Sze, Physics of Semiconductor Devices, 2nd Edition, John Wiley & Sons, New York (1981).
• D. L. Lile, Solid State Electron., 1978, 21, 1199.
• W. E. Spicer, P. W. Chye, P. R. Skeath, and C. Y. Su, I. Lindau, J. Vac. Sci. Technol., 1979, 16, 1422.
• V. Q. Ho and T. Sugano, IEEE Trans. Electron Devices, 1980, ED-27, 1436.
• J. R. Hollanhan and A. T. Bells, Techniques and Applications of Plasma Chemistry, Wiley, New York (1974).
• R. P. H. Chang and A. K. Sinha, Appl. Phys. Lett., 1976, 29, 56. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/07%3A_Group_14/7.11%3A_Oxidation_of_Silicon.txt |
Introduction
While the physical properties of silica make it suitable for use in protective and optical coating applications, the biggest application of insulating SiO2 thin films is undoubtedly in semiconductor devices, in which the insulator performs a number of specific tasks, including: surface passivation, field effect transistor (FET) gate layer, isolation layers, planarization and packaging.
The term insulator generally refers to a material that exhibits low thermal or electrical conductivity; electrically insulating materials are also called dielectrics. It is in regard to the high resistance to the flow of an electric current that SiO2 thin films are of the greatest commercial importance. The dielectric constant (ε) is a measure of a dielectric materials ability to store charge, and is characterized by the electrostatic energy stored per unit volume across a unit potential gradient. The magnitude of ε is an indication of the degree of polarization or charge displacement within a material. The dielectric constant for air is 1, and for ionic solids is generally in the range of 5 - 10. Dielectric constants are defined as the ratio of the material’s capacitance to that of air, i.e., (7.12.1). The dielectric constant for silicon dioxide ranges from 3.9 to 4.9, for thermally and plasma CVD grown films, respectively.
$\epsilon \text{ = } \dfrac{\text{C}_{\text{material}}}{\text{C}_{\text{air}}}$
An insulating layer is a film or deposited layer of dielectric material separating or covering conductive layers. Ideally, in these application an insulating material should have a surface resistivity of greater than 1013 Ω/cm2 or a volume resistivity of greater than 1011 Ω.cm. However, for some applications, lower values are acceptable; an electrical insulator is generally accepted to have a resistivity greater than 105 Ω.cm. CVD SiO2 thin films have a resistivity of 106 - 1016 Ω.cm, depending on the film growth method.
As a consequence of its dielectric properties SiO2, and related silicas, are used for isolating conducting layers, to facilitate the diffusion of dopants from doped oxides, as diffusion and ion implantation masks, capping doped films to prevent loss of dopant, for gettering impurities, for protection against moisture and oxidation, and for electronic passivation. Of the many methods used for the deposition of thin films, chemical vapor deposition (CVD) is most often used for semiconductor processing. In order to appreciate the unique problems associated with the CVD of insulating SiO2 thin films it is worth first reviewing some of their applications. Summarized below are three areas of greatest importance to the fabrication of contemporary semiconductor devices: isolation and gate insulation, passivation, and planarization.
Device isolation and gate insulation
A microcircuit may be described as a collection of devices each consisting of "an assembly of active and passive components, interconnected within a monolithic block of semiconducting material". Each device is required to be isolated from adjacent devices in order to allow for maximum efficiency of the overall circuit. Furthermore within a device, contacts must also be electrically isolated. While there are a number of methods for isolating individual devices within a circuit (reverse-biased junctions, mesa isolation, use of semi-insulating substrates, and oxide isolation), the isolation of the active components in a single device is almost exclusively accomplished by the deposition of an insulator.
In Figure is shown a schematic representation of a silicon MOSFET (metal-oxide-semiconductor field effect transistor). The MOSFET is the basic component of silicon-CMOS (complimentary metal-oxide-semiconductor) circuits which, in turn, form the basis for logic circuits, such as those used in the CPU (central processing unit) of a modern personal computer. It can be seen that the MOSFET is isolated from adjacent devices by a reverse-biased junction (p+-channel stop) and a thick oxide layer. The gate, source and drain contact are electrically isolated from each other by a thin insulating oxide. A similar scheme is used for the isolation of the collector from both the base and the emitter in bipolar transistor devices.
As a transistor, a MOSFET has many advantages over alternate designs. The key advantage is low power dissipation resulting from the high impedance of the device. This is a result of the thin insulation layer between the channel (region between source and drain) and the gate contact, see Figure $1$. The presence of an insulating gate is characteristic of a general class of devices called MISFETs (metal-insulator-semiconductor field effect transistor). MOSFETs are a subset of MISFETs where the insulator is specifically an oxide, e.g., in the case of a silicon MISFET device the insulator is SiO2, hence MOSFET. It is the fabrication of MOSFET circuits that has allowed silicon technology to dominate digital electronics (logic circuits). However, increases in computing power and speed require a constant reduction in device size and increased complexity in device architecture.
Passivation
Passivation is often defined as a process whereby a film is grown on the surface of a semiconductor to either (a) chemically protect it from the environment, or (b) provide electronic stabilization of the surface.
From the earliest days of solid state electronics it has been recognized that the presence or absence of surface states plays a decisive role in the usefulness of any semiconducting material. On the surface of any solid state material there are sites in which the coordination environment of the atoms is incomplete. These sites, commonly termed "dangling bonds", are the cause of the electronically active states which allow for the recombination of holes and electrons. This recombination occurs at energies below the bulk value, and interferes with the inherent properties of the semiconductor. In order to optimize the properties of a semiconductor device it is desirable to covalently satisfy all these surface bonds, thereby shifting the surface states out of the band gap and into the valence or conduction bands. Electronic passivation may therefore be described as a process which reduces the density of available electronic states present at the surface of a semiconductor, thereby limiting hole and electron recombination possibilities. In the case of silicon both the native oxide and other oxides admirably fulfill these requirements.
Chemical passivation requires a material that inhibits the diffusion of oxygen, water, or other species to the surface of the underlying semiconductor. In addition, the material is ideally hard and resistant to chemical attack. A perfect passivation material would satisfy both electronic and chemical passivation requirements.
Planarization
For the vast majority of electronic devices, the starting point is a substrate consisting of a flat single crystal wafer of semiconducting material. During processing, which includes the growth of both insulating and conducting films, the surface becomes increasingly non-planar. For example, a gate oxide in a typical MOSFET (see Figure $1$) may be typically 100 - 250 Å thick, while the isolation or field oxide may be 10,000 Å. In order for the successful subsequent deposition of conducting layers (metallization) to occur without breaking metal lines (often due to the difficulty in maintaining step coverage), the surface must be flat and smooth. This process is called planarization, and can be carried out by a technique known as sacrificial etchback. An abrupt step (Figure $2$a) is coated with a conformal layer of a low melting dielectric, e.g., borophosphorosilicate glass, BPSG (Figure $2$b), and subsequently a sacrificial organic resin (Figure $2$c). The sample is then plasma etched such that the resin and dielectric are removed at the same rate. Since the plasma etch follows the contour of the organic resin, a smooth surface is left behind (Figure $2$d). The planarization process thus reduces step height differentials significantly. In addition regions or valleys between individual metallization elements (vias) can be completely filled allowing for a route to producing uniformly flat surfaces, e.g., the BPSG film shown in Figure $1$.
The processes of planarization is vital for the development of multilevel structures in VLSI circuits. To minimize interconnection resistance and conserve chip area, multilevel metallization schemes are being developed in which the interconnects run in 3-dimensions.
Bibliography
• J. L. Vossen and W. Kern, Phys. Today, 1980, 33, 26.
• S. K. Ghandhi, VLSI Fabrication Principles, Silicon and Gallium Arsenide, Wiley, Chichester, 2nd Ed. (1994).
• S. M. Sze, Physics of Semiconductor Devices, 2nd Edition, John Wiley & Sons, New York (1981).
• W. E. Beadle, J. C. C. Tsai, R. D. Plummer, Quick Reference Manual for Silicon Integrated Cuircuit Technology, Wiley, Chichester (1985).
• A. C. Adams and C. D. Capio, J. Electrochem. Soc., 1981, 128, 2630. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/07%3A_Group_14/7.12%3A_Applications_for_Silica_Thin_Films.txt |
The elements
The Group 15 elements have a particular name pnictogens. Despite the modern IUPAC notation, the Group 15 elements are still referred to as Group V elements in particular by the semiconductor industry. Table $1$ lists the derivation of the names of the Group 15 elements.
Table $1$: Derivation of the names of each of the Group 15 (V) elements.
Element Symbol Name
Nitrogen N Latin nitrogenium, where nitrum (derived from Greek nitron) means saltpetre
Phosphorus P From the Greek phosphoros meaning bringer of light
Arsenic As Derived from Syriac zarniqa and Persian zarnikh, meaning yellow orpiment
Antimony Sb Greek anti and monos meaning not alone. The symbol Sb from Latin stibium
Bismuth Bi New Latin bisemutum from German Wismuth, meaning white mass
Note
According to the Oxford English Dictionary, the correct spelling of the element is phosphorus. The word phosphorous is the adjectival form of the P3+ valence. In the same way that sulfur forms sulfurous and sulfuric compounds, phosphorus forms phosphorous compounds (e.g., phosphorous acid) and P5+ valency phosphoric compounds (e.g., phosphoric acids and phosphates).
Discovery
Nitrogen
Nitrogen was discovered by Rutherford (Figure $1$) in 1772. He called it noxious air or fixed air because there it had been known since the late 18th century that there was a fraction of air that did not support combustion. Nitrogen was also studied by Scheele (Figure $2$), Cavendish (Figure $3$), and Priestley (Figure $4$), who referred to it as burnt air or phlogisticated air.
Phosphorus
German alchemist Hennig Brand (Figure $5$) was experimenting with urine (which contains dissolved phosphates) in 1669. While attempting to create the fabled philosopher's stone (the legendary alchemical substance capable of turning base metals, such as lead, into gold) by the distillation of salts from urine, he produced a white material that glowed in the dark and burned with a brilliant light. He gave the substance the name phosphorus mirabilis (miraculous bearer of light). His process involved letting the urine stand for days then boiling it down to a paste which led to a white waxy substance, white phosphorus.
Brand sold the recipe for 200 thaler (a silver coin from whose name the dollar is derived) to D Krafft who toured much of Europe showing it. During his journeys he met Robert Boyle (Figure $6$) who without learning the details of the synthesis recreated and improved it by using sand in the reduction of the phosphate, (8.1.1).
$\text{4 NaPO}_3 \text{ + 2 SiO}_2\text{ + 10 C} \rightarrow \text{2 Na}_2\text{SiO}_3\text{ + 10 CO + P}_4$
Arsenic
Arsenic sulfides and oxides were known since ancient times. Zosimos (ca. 300 AD) describes roasting sandarach (realgar, α-As4S4) to obtain cloud of arsenious oxide (As2O3) that he reduced to metallic arsenic (Figure $7$).
Antimony
Antimony(III) sulfide, Sb2S3 was known as early as 3000 BC. Pastes of Sb2S3 powder in fat were used as eye cosmetics in the Middle East. An artifact made of antimony dating to about 3000 BC was found at Tello (part of present-day Iraq), and copper objects plated with antimony from 2500 - 2200 BC have been found in Egypt. The first European description of a procedure for isolating antimony is in the book De la pirotechnia by Vannoccio Biringuccio (1480 - 1539).
Bismuth
Since bismuth was known in ancient times, no one person is credited with its discovery. However, the French chemist Claude François Geoffroy (1729 - 1753) demonstrated in 1753 that this metal is distinct from lead and tin.
Abundance
The abundance of the Group 15 elements is given in Table $2$.
Table $2$: Abundance of the Group 15 elements.
Element Terrestrial abundance (ppm)
N 25 (Earth’s crust), 5 (soil), 0.5 (sea water), 78 x 104 (atmosphere)
P 1000 (Earth’s crust), 0.65 (soil), 60 x 10-3 (sea water), trace (atmosphere)
As 1.5 (Earth’s crust), 10 (soil), 16 x 10-3 (sea water), trace (atmosphere)
Sb 0.2 (Earth’s crust), 1 (soil), 0.3 x 10-3 (sea water)
Bi 48 x 10-3 (Earth’s crust), 0.25 (soil), 400 x 10-6 (sea water)
Isotopes
The naturally abundant isotopes of the Group 15 elements are listed in Table $3$.
Table $3$: Abundance of the non-synthetic isotopes of the Group 15 elements.
Isotope Natural abundance (%)
Nitrogen-14 99.634
Nitrogen-15 0.0366
Phosphorus-31 100
Arsenic-75 100
Antimony-121 57.36
Antimony-123 42.64
Bismuth-209 100%
Two radioactive isotopes of phosphorus (32P and 33P) have half-lives that make them useful for scientific experiments (14.262 and 25.34 days, respectively). 32P is a β-emitter (1.71 MeV) and is used to produce radiolabeled DNA and RNA probes. Due to the high energy of the β particles which can penetrate skin and corneas, and because any 32P ingested, inhaled, or absorbed is incorporated into bone and nucleic acids extreme care needs to be taken in handling. The lower energy β particles emitted from 33P (0.25 MeV) make it useful for applications such as DNA sequencing.
While bismuth is traditionally regarded as the element with the heaviest stable isotope, 209Bi, it had long been suspected to be unstable on theoretical grounds. In 2003 researchers at the Institut d'Astrophysique Spatiale in Orsay, France, measured the alpha emission half-life of 209Bi to be 1.9 × 1019 years, over a billion times longer than the current estimated age of the universe!
Industrial production of the elements
Nitrogen is the largest constituent of the Earth's atmosphere (78.082% by volume, 75.3% by weight). It is created by fusion processes in stars, and is estimated to be the 7th most abundant element by mass in the universe. Industrial gas produced is by the fractional distillation of liquid air, or by mechanical means using gaseous air (i.e., pressurized reverse osmosis membrane or pressure swing adsorption). Commercial nitrogen is often a byproduct of air processing for industrial concentration of oxygen for steelmaking, etc.
White phosphorus was originally made commercially for the match industry in the 19th century, by distilling off phosphorus vapor from precipitated phosphates, mixed with ground coal or charcoal, (8.1.2). The precipitated phosphates were made from ground-up bones that had been de-greased and treated with strong acids. This process is, however, obsolete due to the submerged-arc furnace for phosphorus production was introduced to reduce phosphate rock. Calcium phosphate (phosphate rock) is heated to 1200 - 1500 °C with SiO2 and coke (impure carbon) to produce vaporized tetraphosphorus, P4.
$\text{Ca}_{10}\text{(PO}_4\text{)}_6\text{F}_2\text{ + 15 C + 9 SiO}_2\rightarrow \text{6 P}_{\text{n(g)}}\text{ + 9 [(CaO.SiO}_2\text{)] + CaF}_2\text{ + 15 CO}_{\text{(g)}}$
Physical properties
The physical properties of the Group 15 elements (Table $3$) encompasses a gas (N2), a non-metallic solid (P4), metalloids (As and Sb), and a metal (Bi).
Table $3$: Selected physical properties of the Group 15 elements.
Element Mp (°C) Bp (°C) Density (g/cm3)
N -210.00 -195.79 1.251 g/L (0 °C @ 101.325 kPa)
P 44.2 (white), 610 (black) 280.5 (white), 416 - 590 (sub., red), 620 (sub, violet) 1.823 (white), 2.2 - 2.34 (red), 2.36 (violet), 2.69 (black)
As 817 615 (sub.) 5.727
Sb 630.63 1587 6.697 (solid), 6.53 (liquid)
Bi 271.5 1564 9.78 (solid), 10.05 (liquid)
Vapor phase
Nitrogen forms a dimer in the vapor phase with a triple bond (Figure $8$). In the vapor phase above 800 °C tetraphosphorus (P4) is partially dissociated to P2.
Solid state
Phosphorus forms a number of allotropes with very different properties (Figure $9$). Red phosphorus is an intermediate phase between the white and violet forms. Scarlet phosphorus is obtained by allowing a solution of white phosphorus in carbon disulfide to evaporate in sunlight. Black phosphorus is formed by heating white phosphorus under high pressures (ca. 12,000 atmospheres).
White phosphorus has two forms, low-temperature β form and high-temperature α form; both of which contain the P4 tetrahedron (Figure $10$). White phosphorus is the least stable, the most reactive, most volatile, less dense, and most toxic of the allotropes.
The structural relationship between white and red phosphorus involves breaking one of the P-P bonds in the P4 unit and forming a bond with a neighboring tetrahedron to give a chain structure (Figure $11$). Red phosphorus is formed by heating white phosphorus to 250 °C or by exposing white phosphorus to sunlight. Actually red phosphorus is not a single allotrope, but rather an intermediate phase between the white and violet phosphorus, and most of its properties have a range of values (Table $3$).
Violet phosphorus (Figure $12$) is the thermodynamic stable form of phosphorus that is produced by heating red phosphorus above 550 °C. Due to the synthesis being developed by Johann Hittorf (Figure $13$) it is sometimes known as Hittorf's phosphorus.
Black phosphorus is the least reactive allotrope and the thermodynamic stable form below 550 °C. It is also known as β-metallic phosphorus and has a structure somewhat resembling that of graphite (Figure $14$).
In a similar manner to phosphorus, arsenic has several allotropes some of which a structurally related to those of phosphorus. Grey arsenic has a structure similar to black phosphorus (Figure $10$). Yellow arsenic (As4) is soft and waxy with a structure similar to too P4 (Figure $14$). Finally, black arsenic is similar in structure to red phosphorus (Figure $11$). Antimony and bismuth are both traditional metals and have trigonal hexagonal structures (a = 4.299, c = 11.25 Å, and a = 4.537, c = 11.838 Å, respectively).
Bibliography
• V. Biringuccio, The Pirotechnia of Vannoccio Biringuccio: The Classic Sixteenth-Century Treatise on Metals and Metallurgy, Dover Publications (1990).
8.02: Reaction Chemistry of Nitrogen
Despite nitrogen being the inert component of the Earth’s atmosphere, dinitrogen undergoes a range of reactions, although it only reacts with a few reagents under standard temperature and pressure. Nitrogen reacts with oxygen in an electric arc, (8.2.1), both in the laboratory and within lightening strikes.
$\text{N}_2\text{ + O}_2\rightarrow\text{2 NO}$
The synthesis of ammonia is accomplished by the Harber process, using an iron oxide (Fe3O4) catalyst, (8.2.2), at about 500 °C and 200 atmospheres pressure.
$\text{N}_2\text{ + 3 H}_2 \rightarrow \text{2 NH}_3$
Nitrogen reacts with lithium metal at room temperature to form the nitride, (8.2.3). Magnesium also burns in nitrogen, forming magnesium nitride, (8.2.4).
$\text{6 Li + N}_2 \rightarrow \text{2 Li}_3\text{N}$
$\text{3 Mg + N}_2 \rightarrow \text{Mg}_3\text{N}_2$
Nitrogen forms complexes with transition metals yielding nitrogeno complexes, (8.2.5). Under some conditions these complexes react to give ammonia, (8.2.6), and as such may give a hint to the action of nitrogenase in which molybdenum in the active site.
$\text{[(NH}_3\text{)}_5\text{Ru}^{II}\text{(H}_2\text{O)]}^{2+}\text{ + N}_2 \rightarrow \text{[(NH}_3\text{)}_5\text{Ru}^{II}\text{(N}_2\text{)]}^{2+}$
$\text{[W(N}_2\text{)}_2\text{(PR}_3\text{)}_4\text{]} \xrightarrow{\text{H}_2\text{SO}_4\text{/MeOH}} \text{2 NH}_3\text{ + N}_2 \uparrow \text{ + W}^{VI}\text{ compounds}$ | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/08%3A_Group_15_-_The_Pnictogens/8.01%3A_The_Group_15_Elements-_The_Pnictogens.txt |
Trihydrides of the Group 15 Elements
All five of the Group 15 elements form hydrides of the formula EH3. Table $1$ lists the IUPAC names along with those in more common usage.
Table $1$: Traditional and IUPAC (International Union of Pure and Applied Chemistry) names for the Group 15 hydrides.
Compound Traditional name IUPAC name
NH3 Ammonia Azane
PH3 Phosphine Phosphane
AsH3 Arsine Arsane
SbH3 Stibine Stibane
BiH3 Bismuthine Bismuthane
The boiling point and melting point increase increases going down the Group (Table $2$) with increased molecular mass, with the exception of NH3 whose anomalously high melting and boiling points (Figure $1$) are a consequence of strong N-H...H hydrogen bonding. A similar (and stronger) effect is observed for the Group 16 hydrides (H2E).
Table $2$: Selected physical properties of Group 15 hydrides.
Compound Mp (°C) Bp (°C) ΔHf (kJ/mol) E-H bond energy (kJ/mol) H-E-H bond angle (°)
NH3 -77.7 -33.35 -46.2 391 107
PH3 -133 -87.7 9.3 322 93.5
AsH3 -116.3 -55 172.2 247 92
SbH3 -88 -17.1 142.8 255 91.5
The E-H bond strengths decrease down the group and this correlates with the overall stability of each compound (Table $2$). The H-E-H bond angles (Table $2$) also decrease down the Group. The H-E-H bond angle is expected to be a tetrahedral ideal of 109.5°, but since lone pairs repel more than bonding pairs, the actual angle would be expected to be slightly smaller. Two possible explanations are possible for the difference between NH3 and the other hydrides.
1. The N-H bond is short (1.015 Å) compared to the heavier analogs, and nitrogen is more electronegative than hydrogen, so the bonding pair will reside closer to the central atom and the bonding pairs will repel each other opening the H-N-H angle more than observed for PH3, etc.
2. The accessibility of the 2s and 2p orbitals on nitrogen allows for hybridization and the orbitals associated with N-H bonding in NH3 are therefore close to sp3 in character, resulting in a close to tetrahedral geometry. In contrast, hybridization of the ns and np orbitals for P, As, etc., is less accessible, and as a consequence the orbitals associated with P-H bonding in PH3 are closer to p in character resulting in almost 90° H-P-H angle. The lower down the Group the central atom the less hybridization that occurs and the closer to pure p-character the orbitals on E associated with the E-H bond.
Ammonia
Ammonia (NH3) is a colorless, pungent gas (bp = -33.5 °C) whose odor can be detected at concentrations as low 20 – 50 ppm. Its high boiling point relative to its heavier congeners is indicative of the formation of strong hydrogen bonding. The strong hydrogen bonding also results in a high heat of vaporization (23.35 kJ/mol) and thus ammonia can be conveniently used as a liquid at room temperature despite its low boiling point.
WARNING
Ammonia solution causes burns and irritation to the eyes and skin. The vapor causes severe irritation to the respiratory system. If swallowed the solution causes severe internal damage.
Synthesis
Ammonia is manufactured on the industrial scale by the Haber process using the direct reaction of nitrogen with hydrogen at high pressure (102 – 103 atm) and high temperature (400 – 550 °C) over a catalyst (e.g., α-iron), (8.3.1).
$\text{N}_2\text{ + 3 H}_2\rightarrow \text{2 NH}_3$
On the smaller scale ammonia is prepared by the reaction of an ammonium salt with a base, (8.3.2), or hydrolysis of a nitride, (8.3.3). The latter is a convenient route to ND3 by the use of D2O.
$\text{NH}_4\text{X + OH}^- \rightarrow \text{NH}_3\text{ + H}_2\text{O + X}^-$
$\text{Mg}_3\text{N}_2\text{ + H}_2\text{O} \rightarrow \text{3 Mg(OH)}_2 \text{ + 2 NH}_3$
Structure
The nitrogen in ammonia adopts sp3 hybridization, and ammonia has an umbrella structure (Figure $2$) due to the stereochemically active lone pair.
The barrier to inversion of the umbrella is very low (Ea = 24 kJ/mol) and the inversion occurs 100’s of times a second. As a consequence it is not possible to isolate chiral amines in the same manner that is possible for phosphines.
In a similar manner to water, (8.3.4), ammonia is a self-ionizing, (8.3.5); however, the equilibrium constant (K = 10-33) is much lower than water (K = 10-14). The lower dielectric constant of ammonia (16.5 @ 20 °C) as compared to water (80.4 @ 20 °C) means that ammonia is not as good as water as a solvent for ionic compounds, but is better for covalent organic compounds.
$\text{2 H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+\text{ + OH}^-$
$\text{2 NH}_3 \rightleftharpoons \underset{\text{ammonium}}{\textext{NH}_4^+} \t{ + } \und\terset{ext{amide}}{\text{NH}_2^- }$
Reactions
The similarity of ammonia and water means that the two compounds are miscible. In fact, ammonia forms a series of solid hydrates, analogous to ice in which hydrogen bonding defines the structures (Figure $3$). Several hydrates of ammonia are known, including: NH3.2H2O (ammonia dihydrate, ADH), NH3.H2O (ammonia monohydrate, AMH), and 2NH3.2H2O (ammonia hemihydrate, AHH).
It should be noted that these hydrates do not contain discrete NH4+ or OH- ions, indicating that ammonium hydroxide does not exist as a discrete species despite the common useage of the name. In aqueous solution, ammonia is a weak base (pKb = 4.75), (8.3.6).
$\rm NH_{3(aq)} + H_2O_{(l)} \rightleftharpoons NH_{4(aq)}^+ + OH^-_{(aq)}$
Note
Ammonia solutions commonly used in the laboratory is a 35% solution in water. In warm weather the solution develops pressure and the cap must be released with care. The 25% solution sold commercially (for home use) is free from this problem.
Ammonia is a Lewis base and readily forms Lewis acid-base complexes with both transition metals, (8.3.7), and main group metals (Figure $4$).
$\rm [Ni(H_2O)_6]^{2+} + 6 NH_3 \rightleftharpoons [Ni(NH_3)_6]^{2+} + 6 H_2O$
The formation of stable ammonia complexes is the basis of a simple but effective method of detection: Nessler’s reagent, (8.3.8). Using a 0.09 mol/L solution of potassium tetraiodomercurate(II), K2[HgI4], in 2.5 mol/L potassium hydroxide. A yellow coloration indicates the presence of ammonia: at higher concentrations, a brown precipitate may form. The sensitivity as a spot test is about 0.3 μg NH3 in 2 μL.
$\rm NH_4^+ + 2 [HgI_4]^{2-} + 4 OH^- \rightarrow HgO\cdotHg(NH_2)I + 7 I^- + 3 H_2O$
Ammonia forms a blue solution with Group 1 metals. As an example, the dissolution of sodium in liquid ammonia results in the formation of solvated Na+ cations and electrons, (8.3.9) where solv = NH3. The solvated electrons are stable in liquid ammonia and form a complex: [e-(NH3)6].
$\rm Na_{(s)} \rightarrow Na_{(solv)} \rightleftharpoons Na^+_{(solv)} + e_{(solv)}^-$
It is this solvated electron that gives the strong reducing properties of the solution as well as the characteristic signal in the ESR spectrum associated with a single unpaired electron. The blue color of the solution is often ascribed to these solvated electrons; however, their absorption is in the far infra-red region of the spectrum. A second species, Na-(solv), is actually responsible for the blue color of the solution.
$\rm 2 Na_{(solv)} \rightleftharpoons Na^+_{(solv)} + Na^-_{(solv)}$
The reaction of ammonia with oxygen is highly favored, (8.3.11), and the flammability limit of ammonia is 16 – 25 vol%. If the reaction is carried out in the presence of a catalyst (Pt or Pd) the reaction can be limited to the formation of nitric oxide (NO), (8.3.12).
$\rm 4 NH_3 + 3 O_2 \rightarrow 2 N_2 + 6 H_2O$
$\rm 4 NH_3 + 5 O_2 \rightarrow 4 NO + 6 H_2O$
Ammonium salts
The ammonium cation (NH4+) behaves in a similar manner to the Group 1 metal ions. The solubility and structure of ammonium salts particularly resembles those of potassium and rubidium because of their relative size (Table $3$). One difference is that ammonium salts often decompose upon heating, (8.3.13).
$\rm NH_4Cl_{(s)} \rightarrow NH_{3(g)} + HCl_{(g)}$
Table $3$: Ionic radius of the ammonium ion compared to those of potassium and rubidium.
Cation Ionic radius (Å)
K+ 1.33
NH4+ 1.43
Rb+ 1.47
The decomposition of ammonium salts of oxidizing acids can often be violent to highly explosive, and they should be treated with care. For example, while ammonium dichromate, (NH4)2Cr2O7, decomposes to give a volcano (Figure $5$), ammonium permanganate, NH4[MnO4], is friction sensitive and explodes at 60 °C. Ammonium nitrate, NH4[NO3], can cause fire if contacted with a combustible material and is a common ingredient in explosives since it acts as the oxygen source due to its positive oxygen balance, i.e., the compound liberates oxygen surplus to its own needs upon decomposition, (8.3.14).
$\rm NH_4[NO_3] \rightarrow H_2O + N_2 + O$
Bibliography
• A. R. Barron, The Detonator, 2009, 36, 60.
• A. D. Fortes, E. Suard, M. -H. Lemée-Cailleau, C. J. Pickard, and R. J. Needs, J. Am. Chem. Soc., 2009, 131, 13508.
• M. D. Healy, J. T. Leman, and A. R. Barron, J. Am. Chem. Soc., 1991, 113, 2776.
• A. I. Vogel and G. Svehla, Textbook of Macro and Semimicro Qualitative Inorganic Analysis, Longman, London (1979).
Liquid Ammonia as a Solvent
Ammonia has a reasonable liquid range (-77 to –33 °C), and as such it can be readily liquefied with dry ice (solid CO2, Tsub = -78.5 °C), and handled in a thermos flask. Ammonia’s high boiling point relative to its heavier congeners is indicative of the formation of strong hydrogen bonding, which also results in a high heat of vaporization (23.35 kJ/mol). As a consequence ammonia can be conveniently used as a liquid at room temperature despite its low boiling point.
Liquid ammonia is a good solvent for organic molecules (e.g., esters, amines, benzene, and alcohols). It is a better solvent for organic compounds than water, but a worse solvent for inorganic compounds. The solubility of inorganic salts is highly dependant on the identity of the counter ion (Table $4$).
Table $4$: General solubility of inorganic salts in liquid ammonia as a function of the counter ion.
Soluble in liquid NH3 Generally insoluble in liquid NH3
SCN-, I-, NH4+, NO3-, NO2-, ClO4- F-, Cl-, Br-, CO32-, SO42-, O2-, OH-, S2-
The difference in solubility of inorganic salts in ammonia as compared to water, as well as the lower temperature of liquid ammonia, can be used to good advantage in the isolation of unstable compounds. For example, the attempted synthesis of ammonium nitrate by the reaction of sodium nitrate and ammonium chloride in water results in the formation of nitrogen and water due to the decomposition of the nitrate, (8.3.15). By contrast, if the reaction is carried out in liquid ammonia, the sodium chloride side product is insoluble and the ammonium nitrate may be isolated as a white solid after filtration and evaporation below its decomposition temperature of 0 °C, (8.3.16).
$\rm NaNO_2 + NH_4Cl \xrightarrow{H_2O} NaCl + NH_4(NO_2) \rightarrow N_2 + 2 H_2O$
$\rm NaNO_2 + NH_4Cl \xrightarrow{NH_3} NaCl \downarrow + NH_4(NO_2)$
Ammonation
Ammonation is defined as a reaction in which ammonia is added to other molecules or ions by covalent bond formation utilizing the unshared pair of electrons on the nitrogen atom, or through ion-dipole electrostatic interactions. In simple terms the resulting ammine complex is formed when the ammonia is acting as a Lewis base to a Lewis acid, (8.3.17) and (8.3.18), or as a ligand to a cation, e.g., [Pt(NH3)4]2+, [Ni(NH3)6]2+, [Cr(NH3)6]3+, and [Co(NH3)6]3+.
$\rm SiF_4 + 2 NH_3 \rightarrow SiF_4(NH_3)_2$
$\rm BF_3 + NH_3 \rightarrow BF_3(NH_3)$
Ammonolysis
Ammonolysis with ammonia is an analogous reaction to hydrolysis with water, i.e., a dissociation reaction of the ammonia molecule producing H+ and an NH2- species. Ammonolysis reactions occur with inorganic halides, (8.3.18) and (8.3.19), and organometallic compounds, (8.3.20). In both case the NH2- moiety forms a substituent or ligand.
$\rm P(O)Cl_3 + 6 NH_3 \rightarrow P(O)(NH_2)_3 + 3 NH_4Cl$
$\rm BCl_3 + 6 NH_3 \rightarrow B(NH_2)_3 + 3 NH_4Cl$
$\rm Al(CH_3)_3 + NH_3 \rightarrow \dfrac{1}{n}[(H_3C)_2Al(NH_2)]_n + CH_4$
The reaction of esters, (8.3.21), and aryl halides, (8.3.22), are also examples of ammonolysis reactions.
$\rm RC(O)OR' + NH_3 \rightarrow RC(O)NH_2 + R'OH$
$\rm C_6H_5Cl + 2 NH_3 \rightarrow C_6H_5NH_2 + NH_4Cl$
Homoleptic amides
A homoleptic compound is a compound with all the ligands being identical, e.g., M(NH2)n. A general route to homoleptic amide compounds is accomplished by the reaction of a salt of the desired metal that is soluble in liquid ammonia (Table $4$) with a soluble Group 1 amide. The solubility of the Group 1 amides is given in Table $5$. Since all amides are insoluble (except those of the Group 1 metals) are insoluble in liquid ammonia, the resulting amide may be readily isolated, e.g., (8.3.23) and (8.3.24).
$\rm Mn(SCN)_2 + 2 KNH_2 \rightarrow Mn(NH_2)_2 \downarrow + 2 KSCN$
$\rm Cr(NO_3)_3 + 3 KNH_2 \rightarrow Cr(NH_2)_3 \downarrow + 3 KNO_3$
Table $5$: Solubility of Group amides in liquid ammonia.
Amide Solubility in liquid ammonia
LiNH2 Sparingly soluble
NaNH2 Sparingly soluble
KNH2 Soluble
RbNH2 Soluble
CsNH2 Soluble
Redox reactions
Ammonia is poor as an oxidant since it is relatively easily oxidized, e.g., (8.3.25) and (8.3.26). Thus, if it is necessary to perform an oxidation reaction ammonia is not a suitable solvent; however, it is a good solvent for reduction reactions.
$\rm 4 NH_3 + 5 O_2 \rightarrow 4 NO + 6 H_2O$
$\rm 2 NH_3 + 3 CuO \rightarrow N_2 + 3 H_2O + 3 Cu$
Liquid ammonia will dissolve Group 1 (alkali) metals and other electropositive metals such as calcium, strontium, barium, magnesium, aluminum, europium, and ytterbium. At low concentrations (ca. 0.06 mol/L), deep blue solutions are formed: these contain metal cations and solvated electrons, (8.3.27). The solvated electrons are stable in liquid ammonia and form a complex: [e-(NH3)6].
$\rm Na_{(s)} \rightarrow Na_{(solv)} \rightleftharpoons Na^+_{(solv)} + e^-_{(solv)}$
The solvated electrons provide a suitable and powerful reducing agent for a range of reactions that are not ordinarily accomplished, e.g., (8.3.28) and (8.3.29).
$\rm [Ni(CN)_4]^{2-} + 2 e^-_{(solv)} \rightarrow [Ni(CN)_4]^{4-}$
$\rm Co_2(CO)_8 + 2 e^-_{(solv)} \rightarrow 2 [Co(CO)_4]^- The rapid increase in the German population also put a strain on the countries food resources. What compounded this issue was that the aristocratic Junkers families of East Prussia who owned much of the land in what was known as Germany’s breadbasket. Junkers grew rye on their estates because the soil was too light for wheat, and since rye was fertilized with potash (potassium oxide, K2O) of which Germany had vast resources. However, in 1870 grain from the US was becoming cheaper and thus competitive with German rye. To protect their profits the Junkers demanded both subsidies for the export of rye and tariffs for the import of wheat. The result of this was that all the German rye was leaving the country and there was not enough wheat being produced to satisfy the needs of the local population. Sufficient wheat could be grown in Germany if a suitable fertilizer was available. Sodium nitrate (NaNO3), also known as Chile saltpeter, was the most common fertilizer. Unfortunately, by 1900 the deposits looked to be depleted and an alternative was needed. The alternative was found as a component in coal tar. It was known that one of the chemicals that caused the stink associated with coal gas and coal tar was ammonia (NH3). Chemist George Fownes (Figure $8$) had suggested that ammonia be turned into a salt and used as a fertilizer. Unfortunately, the amount of ammonia that could be separated from coal tar was still insufficient, so if ammonia could be made on a large enough scale, then large-scale manufacture of a fertilizer could be realized. In 1909 Fritz Haber (Figure $9$) presented a method of ammonia synthesis to BASF. His work in collaboration with Carl Bosch (Figure $10$) resulted in the process known as the Haber-Bosch process in which nitrogen and hydrogen are mixed at high temperature (600 °C) under pressure (200 atm) over an osmium catalyst, (8.3.30). \[ \rm N_2 + 3 H_2 \xrightarrow{\text{Os cat}} 6 NH_3$
It is interesting to note that the realization of the Haber-Bosch process required not only high-pressure vessels to be constructed by the steel industry, but also the liquid forms of nitrogen and hydrogen. As it turned out ammonia was a necessary component for enabling the production of liquid nitrogen and hydrogen, and involved a false hypothesis of what caused malaria, which led to a desire to keep drinks cold.
Long before it was understood the real cause of malaria, John Gorrie (Figure $11$), a doctor working in Apalachicola on the Gulf Coast of Florida, was obsessed with finding a cure for malaria. The term malaria originated from Medieval Italian: mala aria (bad air), and it was associated with swamps and marshlands. Gorrie noticed that malaria was connected to hot humid weather so he began hanging bowl of ice in wards and circulating the air with a fan. However, ice was cut from frozen lakes and rivers in the North East of the US, stored and then shipped all over the world, and Apalachicola was so small that ice was seldom delivered. Gorrie started looking into methods of making ice. It was well known that when a compressed gas expands it takes heat from its surroundings. Gorrie made a steam engine that compressed air in a piston, which when the piston retracted the air cooled. On the next compression stroke the cold air was pushed out across a brine solution (saturated aqueous NaCl) cooling it. When he brought water in contact with the cold brine, it froze creating the first man made ice. On 14 July 1850 Gorrie produced ice for the French Consul to cool the champagne for the celebration of Bastille Day. Just before he died, Gorrie suggested that his (by then) Patented process could be used for cooling food for transport, and it was this application that was used extensively by British merchants to bring meat from Australia to Britain. However, in Germany, Gorrie’s invention was more useful for beer.
Whereas the British traditionally brewed beer in which the yeast ferments on the surface (top fermentation) at a temperature of 60 – 70 °F, in Germany beer was made using bottom fermentation. This style of fermentation requires a temperature just above freezing. Traditionally, cold cellars were used to store the fermenting beer, and it is from here the name lager is derived from the German verb largern: to store. There had been a law in Germany preventing brewing in the summer, but with Gorrie’s process the possibility was to be able to brew beer all year. Carl von Linde (Figure $12$) was asked to develop a refrigeration system. He used ammonia instead of air in Gorrie’s system, and in 1879 he set up a company to commercialize his ideas. The success of his refrigerator was such that by 1891 there were over 12,000 fridges being used, and more importantly there was now a convenient method of liquefying gases such as hydrogen and nitrogen; both of which were needed for the Haber-Bosch process.
As a consequence of the use of ammonia as a refrigerant, it was possible to prepare ammonia on a large industrial scale. Ammonia prepared by the Haber-Bosch process can be converted to nitric acid by the Ostwald process developed by Wilhelm Ostwald (Figure $13$). Treatment of ammonia with air over a platinum catalyst yields initially nitric oxide, (8.3.31), and subsequently to nitrogen dioxide, (8.3.32), which dissolves in water to give nitric acid, (8.3.33).
$\rm 4 NH_3 + 5 O_2 \xrightarrow{\text{Pt cat}} 4 NO + 6 H_2O$
$\rm 2 NO + O_2 \xrightarrow{\text{Pt cat}} 2 NO_2$
$\rm 3 NO_2 + H_2O \rightarrow 2 HNO_3 + NO$
Addition of soda (sodium hydroxide, NaOH) to nitric acid results in the formation of sodium nitrate, (8.3.34), which was the same fertilizer produced from the deposits in Chile.
$\rm HNO_3 + NaOH \rightarrow NaNO_3 + H_2O$
Unfortunately, for the Haber-Bosch and Ostwald processes, an even cheaper form of fertilizer was synthesized around the same time using calcium carbide to prepare calcium cyanamide (CaCN2), (8.3.35). As a consequence, the Haber-Bosch process was forgotten until the outbreak of the First World War in 1914.
$\rm CaC_2 + N_2 \rightarrow CaCN_2 + C$
Within weeks of the outbreak Germany realized it had only enough explosives for about a year of conflict. This was because the main source of explosives, sodium nitrate was the same source that gave fertilizer, i.e., Chile. Realizing this, the Royal Navy effectively blockaded the supply lines. If Germany did not find another source of Great War would have been over early in 1916, however, it was remembered that the Haber-Bosch process in combination with the Ostwald processes would allow the synthesis of nitric acid, which when mixed with cotton, made nitrocellulose (Figure $14$), also known as gun cotton, an explosive, (8.3.36).
$\rm HNO_3 + C_6H_{10}O_5 \rightarrow C_6H_7(NO_2)_3O_5 + 3 H_2O$
As a result of the industrial synthesis of ammonia Germany was able to manufacture sufficient explosives to fight until 11 November 1918, by which time almost 10 million were dead, almost 7 million missing, and over 21 million were wounded (Figure $15$).
Hydrazine
Hydrazine (N2H4) is a colorless liquid with an odor similar to ammonia. Hydrazine has physical properties very close to water, with a melting point of 2 °C and a boiling point of 114 °C. The similarity in its chemistry to water is as a result of strong intermolecular hydrogen bonding.
WARNING
Hydrazine is highly toxic and dangerously unstable, and is usually handled as aqueous solution for safety reasons. Even so, hydrazine hydrate causes severe burns on the skin and eyes. Contact with transition metals, their oxides (e.g., rust), or salts cause catalytic decomposition and possible ignition of evolved hydrogen. Reactions with oxidants are violent.
Synthesis
Hydrazine is manufactured on the industrial scale by the Olin Raschig process using the reaction of a sodium hypochlorite solution with ammonia at 5 °C to form chloramine (NH2Cl) and sodium hydroxide, (8.3.37). The chloramines solution is the reacted with ammonia under pressure at 130 °C, (8.3.38). Ammonia is used in a 33 fold excess.
$\rm NH_3 + OCl^- \rightarrow OH^- + NH_2Cl$
$\rm NH_2Cl + OH^- + NH_3 \rightarrow N_2H_4 + Cl^- + H_2O$
If transition metals are present then decomposition occurs, (8.3.38), and therefore, ethylenediaminetetraacetic acid (EDTA, Figure $16$) is added to complex the transition metal ions. The as produced hydrazine solution can be concentrated by distillation to give a 65% solution. Anhydrous hydrazine is formed by the distillation from NaOH.
$\rm 2 NH_2Cl + N_2H_4 \rightarrow 2 NH_4^+ + 2 Cl^- + N_2$
Alternative routes to hydrazine include the oxidation of urea with sodium hypochlorite, (8.3.40), and the reaction of ammonia and hydrogen peroxide, (8.3.41).
$\rm (H_2N)_2\text{C=O} + NaOCl + 2 NaOH \rightarrow N_2H_4 + H_2O + NaCl + Na_2CO_3$
$\rm 2 NH_3 + H_2O_2 \rightarrow N_2H_4 + 2 H_2O$
Structure
The nitrogen atoms in hydrazine adopt sp3 hybridization (Figure $17$a), and molecule adopts a gauche conformation in the vapor, liquid and solid states (Figure $17$b).
In a similar manner to ammonia, (8.3.42), hydrazine is a self-ionizing, (8.3.43). While there is a wide range of salts of the N2H5+ cation, only the sodium and potassium salts of N2H3- are stable.
$\rm 2 NH_3 \rightleftharpoons \underset{ammonium}{NH_4^+} + \underset{amide}{NH_2^-}$
$\rm 2 N_2H_4 \rightleftharpoons N_2H_5^+ + N_2H_3^-$
Reaction chemistry and uses
Hydrazine is polar, highly ionizing, and forms stable hydrogen bonds, and its resemblance to water is reflected in the formation of aqueous solutions and hydrates. In the solid state the monohydrate is formed, i.e., N2H4.H2O. In solution hydrazine acts as a base to form the hydrazinium ion, (8.3.44) where Kb = 8.5 x 10-7. The presence of a second Lewis base site means that hydrazine can be protonated twice to form the hydrazonium ion, (8.3.45) where Kb = 8.9 x 10-16. Salts of the cation N2H5+ are stable in water; however, the salts of the dication are less stable.
$\rm 2 NH_3 \rightleftharpoons NH_4^+ + NH_2^-$
$\rm N_2H_5^+ + H_2O \rightleftharpoons N_2H_5^{2+} + OH^-$
The reaction of hydrazine with oxygen is highly favored, (8.3.46), and the explosive limit is 1.8 – 100 vol%.
$\rm N_2H_4 + O_2 \rightarrow N_2 + 2 H_2O$
Hydrazine is useful in a number of organic reactions for the synthesis of a wide range of compounds used in pharmaceuticals, textile dyes, and in photography, including:
• Hydrazone formation, (8.3.47) and (8.3.48).
• Alkyl-substituted hydrazine synthesis via direct alkylation with alkyl halides.
• Reaction with 2-cyanopyridines to form amide hydrazides, which can be converted using 1,2-diketones into triazines.
• Use in the Wolff-Kishner reduction that transforms the carbonyl group of a ketone or aldehyde into a methylene (or methyl) group via a hydrazone intermediate, (shown below).
• As a building block for the preparation of many heterocyclic compounds via condensation with a range of difunctional electrophiles.
• Cleavage N-alkylated phthalimide derivatives.
• As a convenient reductant because the by-products are typically nitrogen gas and water.
$\rm 2 (CH_3)_2\text{C=O} + N_2H_4 \rightarrow 2 H_2O + [(CH_3)_2\text{C=N}]_2$
$\rm [(CH_3)_2\text{C=N}]_2 + N_2H_4 \rightarrow 2 (CH_3)+2\text{C=N}NH_2$
Messerschmitt Me 163 Komet
Designed by Alexander Lippisch (Figure $18$), the Messerschmitt Me 163B Komet (Figure $19$) was the first rocket-powered fighter plane. With a top speed of around 596 mph (Mach 0.83) and a service ceiling of 40,000 ft, the Komet’s performance of the Me 163B far exceeded that of contemporary piston engine fighters. However, despite its impressive performance, it was only produced in limited numbers (ca. 370 as compared to the 1,430 built of its jet powered compatriot the Me 262) and was not an effective combat airplane.
The Komet was powered by the HWK 109-509 hot engine (Figure $20$) that used a mixture of a fuel and an oxidizer. The fuel was a mixture of hydrazine hydrate (30%), methanol (57%), and water (13%) that was designated by the code name, C-Stoff, that burned with the oxygen-rich exhaust from hydrogen peroxide (T-Stoff) used as the oxidizer. The C-Stoff was stored in a glass tank on the plane, while the T-Stoff was stored in an aluminum container. An oxidizing agent cocktail of CaMnO4 and/or K2CrO4 was added to the T-Stoff generating steam and high temperatures, this in tern reacted violently with the C-Stoff. The flow of reagents was controlled by two pumps, to regulate the rate of combustion and thereby the amount of thrust. The violent combustion process resulted in the formation of water, carbon dioxide and nitrogen, and a huge amount of heat sending out a superheated stream of steam, nitrogen and air that was drawn in through the hole in the mantle of the engine, thus providing a forward thrust of approximately 3,800 lbf. Because of the potential hazards of mixing the fuels, they were stored at least 1/2 mile apart, and the plane was washed with water between fueling steps and after missions.
Phosphine and Arsine
Because of their use in metal organic chemical vapor deposition (MOCVD) of 13-15 (III-V) semiconductor compounds phosphine (PH3) and arsine (AsH3) are prepared on an industrial scale.
Synthesis
Phosphine (PH3) is prepared by the reaction of elemental phosphorus (P4) with water, (8.3.49). Ultra pure phosphine that is used by the electronics industry is prepared by the thermal disproportionation of phosphorous acid, (8.3.50).
$\rm 2 P_4 + 12 H_2O \rightarrow 5 PH_3 + 3 H_3PO_4$
$\rm 4 H_3PO_4 \rightarrow PH_3 + 3 H_3PO_4$
Arsine can be prepared by the reduction of the chloride, (8.3.51) or (8.3.52). The corresponding syntheses can also be used for stibine and bismuthine.
$\rm 4 AsCl_3 + 3 LiAlH_4 \rightarrow 4 AsH_3 + 3 LiAlCl_4$
$\rm 4 AsCl_3 + 3 NaBH_4 \rightarrow 4 AsH_3 + 3 NaCl + 3 BCl_3$
The hydrolysis of calcium phosphide or arsenide can also generate the trihydrides.
Structure
The phosphorus in phosphine adopts sp3 hybridization, and thus phosphine has an umbrella structure (Figure $21$a) due to the stereochemically active lone pair. The barrier to inversion of the umbrella (Ea = 155 kJ/mol) is much higher than that in ammonia (Ea = 24 kJ/mol). Putting this difference in context, ammonia’s inversion rate is 1011 while that of phosphine is 103. As a consequence it is possible to isolate chiral organophosphines (PRR'R"). Arsine adopts the analogous structure (Figure $21$b).
Reactions
Phosphine is only slightly soluble in water (31.2 mg/100 mL) but it is readily soluble in non-polar solvents. Phosphine acts as neither an acid nor a base in water; however, proton exchange proceeds via the phosphonium ion (PH4+) in acidic solutions and via PH2- at high pH, with equilibrium constants Kb = 4 x 10-28 and Ka = 41.6 x 10-29, respectively.
Arsine has similar solubility in water to that of phosphine (i.e., 70 mg/100 mL), and AsH3 is generally considered non-basic, but it can be protonated by superacids to give isolable salts of AsH4+. Arsine is readily oxidized in air, (8.3.53).
$\rm 2 AsH_3 + 3 O_2 \rightarrow As_2O_3 + 3 H_2O$
Arsine will react violently in presence of strong oxidizing agents, such as potassium permanganate, sodium hypochlorite or nitric acid. Arsine decomposes to its constituent elements upon heating to 250 - 300 °C.
Gutzeit test
The Gutzeit test is the characteristic test for arsenic and involves the reaction of arsine with Ag+. Arsine is generated by reduction of aqueous arsenic compounds, typically arsenites, with Zn in the presence of H2SO4. The evolved gaseous AsH3 is then exposed to silver nitrate either as powder or as a solution. With solid AgNO3, AsH3 reacts to produce yellow Ag4AsNO3, while with a solution of AgNO3 black Ag3As is formed.
Hazards
Pure phosphine is odorless, but technical grade phosphine has a highly unpleasant odor like garlic or rotting fish, due to the presence of substituted phosphine and diphosphine (P2H4). The presence of P2H4 also causes spontaneous combustion in air. Phosphine is highly toxic; symptoms include pain in the chest, a sensation of coldness, vertigo, shortness of breath, and at higher concentrations lung damage, convulsions and death. The recommended limit (RL) is 0.3 ppm.
Arsine is a colorless odorless gas that is highly toxic by inhalation. Owing to oxidation by air it is possible to smell a slight, garlic-like scent when arsine is present at about 0.5 ppm. Arsine attacks hemoglobin in the red blood cells, causing them to be destroyed by the body. Further damage is caused to the kidney and liver. Exposure to arsine concentrations of 250 ppm is rapidly fatal: concentrations of 25 – 30 ppm are fatal for 30 min exposure, and concentrations of 10 ppm can be fatal at longer exposure times. Symptoms of poisoning appear after exposure to concentrations of 0.5 ppm and the recommended limit (RL) is as low as 0.05 ppm.
Bibliography
• R. Minkwitz, A. Kornath, W. Sawodny, and H. Härtner, Z. Anorg. Allg. Chem., 1994, 620, 753. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/08%3A_Group_15_-_The_Pnictogens/8.03%3A_Hydrides.txt |
Oxides of Nitrogen
A summary of the physical properties of the oxides of nitrogen is given in Table $1$.
Table $1$: Physical properties of the oxides of nitrogen.
Oxide Formula Mp (°C) Bp (°C)
Nitrous oxide N2O -90.8 -88.5
Nitric oxide NO -163.6 -161.8
Dinitrogen trioxide N2O3 -100.6 3.5 (dec.)
Nitrogen dioxide (dinitrogen tetroxide) NO2/N2O4 -11.2 (NO2) 21.2 (N2O4)
Nitrogen pentoxide N2O5 30 47
Nitrous oxide
Gaseous nitrous oxide (N2O) is prepared by the careful thermal decomposition of ammonium nitrate (NH4NO2), (8.4.1). Nitrous oxide is a linear molecule (Figure $1$a) that is isoelectronic (and isostructural) with carbon dioxide. Despite its use as a power enhancement for automobiles, nitrous oxide is actually not very reactive and a major use is as an aerosol propellant.
$\rm NH_4NO_2 \xrightarrow{\Delta} N_2O + 2 H_2O$
Nitrous oxide as an anesthetic drug
Nitrous oxide is known as "laughing gas" due to the euphoric effects of inhaling it, a property that has led to its recreational use as a hallucinogen. However, it is as a anesthetic that it has a legitimate application.
The first use of nitrous oxide as anesthetic drug was when dentist Horace Wells (Figure $2$) with assistance by Gardner Quincy Colton (Figure $3$) and John Mankey Riggs (Figure $4$), demonstrated insensitivity to pain from a dental extraction in December 1844. Wells subsequently treated 12-15 patients, and according to his own record it only failed as an anesthetic in two cases. In spite of these results, the method was not immediately adopted, probably because during his first public demonstration was only partly successful.
The method did not come into general use until 1863, when Colton successfully used it for more than 25,000 patients. As such, the usage of nitrous oxide rapidly became the preferred anesthetic method in dentistry. Because the gas is mild enough to keep a patient in a conscious and conversational state, and yet in most cases strong enough to suppress the pain caused by dental work, it remains the preferred gas anesthetic in today's dentistry.
Nitrous: the secret to more power.
In motorsports, nitrous oxide (often referred to as nitrous or NOS) allows the engine to burn more fuel, resulting in a more powerful combustion, and hence greater horsepower. The gas itself is not flammable, but it delivers more oxygen (33%) than atmospheric air (21%) by breaking down at elevated temperatures. When N2O breaks down in during fuel combustion, the decomposition of nitrous is exothermic, contributing to the overall power increase.
Nitrous oxide is stored as a compressed liquid (Figure $5$); the evaporation and expansion of liquid nitrous oxide in the intake manifold causes a large drop in intake charge temperature, resulting in a denser charge, further allowing more air/fuel mixture to enter the cylinder. Nitrous oxide is sometimes injected into (or prior to) the intake manifold, whereas other systems directly inject right before the cylinder (direct port injection) to increase power.
One of the major problems of using nitrous oxide in a reciprocating engine is that it can produce enough power to damage or destroy the engine. Very large power increases are possible, and if the mechanical structure of the engine is not properly reinforced, the engine may be severely damaged or destroyed during this kind of operation.
Automotive-grade liquid nitrous oxide differs slightly from medical-grade nitrous oxide in that a small amount of sulfur is added to prevent substance abuse.
Nitric oxide
Nitric oxide (NO) is formed by the high temperature oxidation of nitrogen, (8.4.2), or the platinum catalyzed oxidation of ammonia at 800 °C, (8.4.3).
$\rm N_2 + O_2 \rightleftharpoons 2 NO$
$\rm 4 NH_3 + 5 O_2 \xrightarrow{Pt} 4 NO + 6 H_2O$
Nitric oxide (Figure $6$b) is electronically equivalent to dinitrogen (N2) plus an electron, and as a consequence it is paramagnetic with one unpaired electron. The location of the unpaired electron in the π* orbital (Figure $6$a) results in a bond order of 2.5 rather than the triple bond observed for N2 (Figure $6$b). The N-O distance of 1.15 Å is intermediate between the triple bond distance in NO+ (1.06 Å) and the typical double bond distance (ca. 1.20 Å). Furthermore, because of the location of the electron it is easy to oxidize nitric oxide to the nitrosonium ion (NO+), (8.4.4).
$\rm NO \rightarrow NO^+ + e^-$
$\rm 2 NO + O_2 \rightarrow 2 NO_2$
Nitric oxide is unstable to heat, (8.4.6), and oxidation, (8.4.5). It will also react with halogens to form the nitrosyl halides, XNO.
$\rm 3 NO \rightarrow N_2O + NO_2$
Dinitrogen trioxide
Dinitrogen trioxide (N2O3) is formed from the stoichiometric reaction between NO and O2 or NO and N2O4. Dinitrogen trioxide has an intense blue color in the liquid phase and a pale blue color in the solid state. Thermal dissociation of N2O3, (8.4.7), occurs above -30 °C, and some self-ionization of the pure liquid is observed, (8.4.8). The asymmetric structure of N2O3 (Figure $1$c) results in a polar molecule (Figure $7$).
$\rm N_2O_3 \rightleftharpoons NO + NO_2$
$\rm N_2O_3 \rightleftharpoons NO^+ + NO_2^-$
Nitrogen dioxide (and tetroxide)
Formed from the oxidation of nitric oxide, (8.4.9), brown nitrogen dioxide is actually in equilibrium with its colorless dimeric form, nitrogen tetroxide (N2O4), (8.4.10).
$\rm 2 NO + O_2 \rightleftharpoons 2 NO_2$
$\rm \underset{brown\paramagnetic}{2NO_2} \rightleftharpoons \underset{colorless\diamagnetic}{N_2O_4}$
Nitrogen dioxide (Figure $1$d) is electronically equivalent to the nitrate anion (NO2-) less one electron, and as such it is paramagnetic with one unpaired electron. The location of the unpaired electron in a nitrogen sp2 orbital, and a consequently it forms a dimer through a N-N bond (Figure $1$e). Furthermore, it is easy to oxidize nitrogen dioxide to the nitronium ion (NO2+), (8.4.11).
$\rm NO_2 \rightarrow NO_2^+ + e^-$
Nitrogen dioxide dissolves in water to form a mixture of nitric and nitrous acids, (8.4.12). Nitrogen dioxide acts as an oxidizing agent with the formation of nitrate anion, (8.4.13).
$\rm 2 NO_2 + H_2O \rightarrow NO_3^- + NO_2^- + 2H^+$
$\rm 2 NO_2 + 2 I^- \rightarrow I_2 + 2 NO_2^-$
The most common structural form of N2O4 (Figure $1$e) is planar with a long N-N bond (1.78 Å) that is significantly longer than observed in hydrazine (1.47 Å). Rationalization of this structural effect is obtained from a consideration of the molecular orbitals, which show that the electrons in the σ-bond are actually delocalized over the whole molecule. The rotation about the N-N bond is 9.6 kJ/mol.
Nitrogen pentoxide
The dehydration of nitric acid, with P2O5, yields nitrogen pentoxide, (8.4.14), which is an unstable solid at room temperature (Table $1$). In the solid state nitrogen pentoxide is actually nitronium nitrate (NO2+NO3-), however, in the vapor phase it exists as a molecular species (Figure $1$f) with a bent N-O-N unit. Nitrogen pentoxide is a very powerful nitrating and oxidation agent.
$\rm 2 HNO_3 \rightleftharpoons N_2O_5 + H_2O$
Nitrogen oxides as precursors to smog and acid rain
Nitrogen oxides (NOx) emissions are estimated to be in the range of 25 - 100 megatonnes of nitrogen per year. Natural sources are thought to make up approximately 1/3 of the total. The generations of NOx (primarily a mixture of NO2 and NO) is the main source of smog and a significant contribution to atmospheric pollution; however, NOx is also responsible for much of the acidity in acid rain.
Atmospheric reactions leading to acid rain
In the dry atmosphere, nitric oxide reacts is oxidized rapidly in sunlight by ozone, (8.4.15). The nitrogen dioxide reacts with the hydroxide radical, formed by the photochemical decomposition of ozone, (8.4.16) and (8.4.17), in the presence of a non-reactive gas molecule such as nitrogen to form nitric acid vapor, (8.4.18). The conversion rate for NOx to HNO3 is approximately ten times faster than the equivalent reaction for sulfur dioxide. For example, conversion is essentially complete for a NOx plume by the time it transverses the North Sea from the UK to Scandinavia.
$\rm NO + O_2 \rightarrow NO_2 + O_2$
$\rm O_3 + h\nu \rightarrow O* + O_2$
$\rm O* + H_2O \rightarrow 2HO \cdot$
$\rm NO_2 + \cdot OH + N_2 \rightarrow HNO_3 + N_2$
At night, conversion takes place via the formation of a nitrate radical, (8.4.19), which subsequently photochemically unstable under sunlight forming nitrogen pentoxide, (8.4.20), that reacts with water on the surface of aerosol particles to form nitric acid, (8.4.21).
$\rm NO_2 + O_3 \rightarrow NO_3 + O_2$
$\rm NO_2 + NO_3 + N_3 \rightarrow N_2O_5 + N_2$
$\rm N_2O_5 + H_2O \rightarrow 2 HNO_3$
Both NO and NO2 are only slightly soluble in water and it is therefore more probable that the nitric acid content of rain is more likely due to the dissolution of nitric acid vapor into raindrops, (8.4.22), rather than a separate reaction.
$\rm N_2O_5 + H_2O_{(l)} \rightarrow 2 HNO_{3(aq)}$
Oxoacids of Nitrogen
Nitrous acid
In the gas phase nitrous acid can be made by the following reaction:
$\rm NO + NO_2 + H_2O \rightleftharpoons 2 HNO_2$
The gas phase structure as determined by IR spectroscopy is shown in Figure $8$a, in which the nitrogen is planar with sp2 hybridization.
In basic aqueous solution the same reaction results in the formation of the nitrite ion, (8.4.24), which can be precipitated as the Ba2+ salt. After separation the addition to sulfuric acid yields a solution of nitrous acid. However, it is not possible to concentrate by heating since decomposition occurs, (8.4.25).
$\rm NO + NO_2 + 2 OH^- \rightleftharpoons 2 NO_2^- + H_2O$
$\rm 3 HNO_2 \rightarrow H_3O^+ + NO_3^- + 2 NO$
Nitrous acid is a fairly weak acid in water (pK = 5.22); however, many salts are known of the nitrite ion, NO2- (Figure $8$b). Complexes of the nitrite ion can be monodentate with bonding via nitrogen (nitro) or an oxygen (nitrito). Both isomers can be isolated in the case of an inert metal, i.e., substitutionally inert d6 octahedral complexes (Figure $9$). Bidentate and bridging modes of coordination are also known for the nitrite anion (Figure $10$).
Nitrite can act as either an oxidizing agent, (8.4.26), or a reducing agent, (8.4.27).
$\rm NO_2^- + 2H^+ + Fe_{2+} \rightarrow NO + H_2O + Fe^{3+}$
$\rm NO_2^- + H_2O + 2 Ce^{4+} \rightarrow NO_3^- + 2 H^+ + 2 Ce^{3+}$
The most important use of nitrous acid is in the diazotization reactions in which nitrous acid is generated by acidifying nitrite solution.
Nitric acid
Nitric acid (HNO3), also known as aqua fortis and spirit of nitre, is made by the dissolution of nitrogen dioxide in water, (8.4.28). Nitric acid can be concentrated by distillation from concentrated sulfuric acid.
$\rm 3 NO_2 + H_2O \rightarrow 2 HNO_3 + NO$
The structure of HNO3 in the gas phase is planar at the sp2 nitrogen (Figure $11$a).
Pure 100% nitric acid is a very corrosive liquid that is strongly acidic and protonates and dissolves organic species. In the liquid phase it is slightly dissociated, (8.4.29). It is also a powerful oxidizing agent, converting non-metal elements to either the oxide or oxoacid. In contrast with metals it forms either salts or complexes in which the metal is in its highest oxidation state. It is unstable and decomposes upon heating or photolysis.
$\rm 3 HNO_3 \rightleftharpoons H_3O^+ + NO_2^+ + 2 NO_3^-$
The pure acid has the highest self ionization of pure liquid acids, (8.4.30). However, the loss of water results, (8.4.31), such that the overall reaction can be described by (8.4.32).
$\rm 2 HNO_3 \rightleftharpoons H_2NO_3^+ + NO_3^+$
$\rm H_2NO_3^+ \rightleftharpoons H_2O + NO_2^+$
$HNO_3 \rightleftharpoons NO_3^+ + NO_2^+ + H_2O$
The common concentration of nitric acid is 70%. While the pure acid is colorless, samples often take on a yellow color due to the photochemical decomposition of nitric acid to give brown NO2.
$\rm 4 HNO_3 \xrightarrow{h\nu} 4 NO_2 + 2 H_2O + O_2$
Note
The term the “acid test” is derived from the medieval practice of debasing of gold and silver currencies (often by the Monarchs who issued them) by debasing with copper. With ducats (Milan), livres (France), florins (Florence), maravedies (Spain) and bezants (Constantinople) in widespread use, and each with a nominal gold or silver content, it was important for a merchant to be able verify the worth of any particular coin. If a drop of dilute nitric acid was placed onto a silver coin adulterated with copper, it turned green, due to the formation of copper(II) nitrate. Conversely, if a gold coin reacted in any way with the nitric acid it was not pure. In both cases the coins failed the “acid test”.
Aqua regia (so called because it dissolves gold) is a mixture of 70% nitric acid and hydrochloric acid in a 1:3 ratio. Aqua regia is a very powerful oxidizing agent (it contains Cl2) and stablizes some metals as their chloro complexes (Table $2$). If HF is added in place of HCl, tantalum may be dissolved with the formation of [TaF2]-.
Table $2$: Metal chloride salts formed by the dissolution of metals in aqua regia.
Metal Chloride salt formed
Au [AuCl4]-
Pt [PtCl6]2-
Fuming nitric acid (100% nitric) is exceedingly corrosive and should not be used. In 100% sulfuric acid, nitric acid acts as a base and gets protonated, (8.4.34) and (8.4.35), and acts as a powerful nitrating agent.
$\rm HNO_3 + H_2SO_4 \rightarrow ON(OH)_2^+ + HSO_4^-$
$\rm ON(OH)_2^+ + HSO_4^- + H_2SO_4 \rightarrow H_3O^+ + NO_2^+ + 2 HSO_4^-$
In water dilute nitric acid is fully ionized, (8.4.36).
$\rm HNO_3 + H_2O \rightarrow H_3O^+ + NO_3^-$
The nitrate ion is planar (Figure $11$b) and forms many salts and complexes. The nitrate anion is most commonly a monodentate ligand, but can also be a bidentate ligand (Figure $12$a) or a bridging ligand (Figure $12$b and 12c).
Phosphorous Oxides
Phosphorous trioxide
The stoichiometric oxidation of white phosphorus yields phosphorous trioxide, (8.4.37).
$\rm P_4 + 3 O_2 \rightarrow 2 P_2O_3$
Several structural forms of phosphorous trioxide are known, but P4O6 is a stable molecule structure (Figure $13$), while the rest are polymers.
Phosphorous pentoxide
The reaction of white phosphorus with excess oxygen (or the oxidation of phosphorous trioxide) yields phosphorous pentoxide, (8.4.38).
$\rm P_4 + 5 O_2 \rightarrow 2 P_2O_5$
The structure of hexagonal phosphorous pentoxide is actually that of the dimeric form, P4O10, and is based upon the structure of P4O6, but with P=O units instead of lone pairs (Figure $14$). The structure is maintained in the vapor phase; however, other crystalline and glassy forms comprise of a sheet-like structure (Figure $15$).
Phosphorous pentoxide is an excellent drying agent below 100 °C. It reacts with water to form various phosphoric acids, and it will extract water from other ‘drying agents’, e.g., (8.4.39) and (8.4.40). Phosphorous pentoxide will dehydrate amides to give nitriles, (8.4.41).
$\rm 2 HNO_3 \xrightarrow[-H_2O]{P_4O_{10}}N_2O_5$
$\rm H_2SO_4 \xrightarrow[-H_2O]{P_4O_{10}}SO_3$
$\rm RC(O)NH_2 \xrightarrow[-H_2O]{P_4O_{10}} RC\equiv N$
Oxoacids of Phosphorus
Phosphorous pentoxide (P2O5) is an excellent drying agent, and its action is a result of the formation of a range of oxoacids.
Hypophosphorous acid
Hypophosphorous acid, H3PO2, is easily prepared pure buy the reaction of white phosphorous with base, followed by acidification, (8.4.42). The pure acid is a solid (Mp = 27 °C) and very soluble in water.
$\rm P_4 + 4 OH^- + 4 H_2O \rightarrow \underset{\underset{H_3PO_2}{\downarrow \text{ acidify}}}{4 H_2PO_2^-} + 2 H_2$
The structure of H3PO2 is determined by X-ray crystallography to be tetrahedral with two hydride ligands and a hydroxide (Figure $16$a). The presence of two hydrides is confirmed by NMR spectroscopy. The 1H NMR resonance shows an OH line and a doublet from the P-H with a large one-bond coupling constant to the 31P nucleus. The non-decoupled 31P NMR spectrum shows a triplet (δ = 13 ppm, JP-H = 530 Hz) due to the two hydrides.
In water hypophosphorous acid is a monobasic acid (pK = 1.2), (8.4.43), and it forms a wide range of salts. It is also a powerful reducing agent, but its reaction kinetics is slow.
$\rm H_2P(O)OH + H_2O \rightleftharpoons H_2P(O)_2^- + H_3O^+$
Phosphorous acid
The reaction of P4O6 or PCl3 with water yields phosphorous acid, H3PO3; which like hypophosphorous acid is a solid (Mp = 70.1 °C) and very soluble in water. The structure is shown by X-ray crystallography to be comprised of a tetrahedral phosphorus with one hydride and two hydroxides (Figure $16$b). 31P NMR spectroscopy demonstrates the presence of a single hydride by the presence of a doublet as a consequence of the phosphorous center being split by a single hydride (δ = 4 ppm, JP-H = 700 Hz). The 1H NMR spectrum shows a doublet for the hydride and a single resonance of twice the intensity for the hydroxide.
As expected, in water phosphorous acid is dibasic, (8.4.44). The acid (and the anions) are strong reducing agents, yielding phosphoric acid. They actually react very slowly, and it is thought that this may be due to the reaction being in the tautomeric form, P(OH)3. Although this has not been isolated, the trialkyl derivatives exist in both forms, i.e., esters of phosphoric acids (Figure $17$a) and trialkyl phosphates (Figure $17$b).
$\rm HP(O)(OH)_2 \xrightleftharpoon[pK_1=1.8] [HP(O)_2OH]^- \xrightleftharpoons[pK_2=6.7] [HP(O)_3]^{2-}$
Ortho phosphoric acid
Orthophosphoric acid (H3PO4) is the most common oxoacid of phosphorus (Figure c). The term acid phosphoric acid is commonly used. It is made from the reaction of phosphates with sulfuric acid (H2SO4), or from the hydrolysis of P4O10. The pure acid is a colorless crystalline solid (Mp = 42.35 °C) with extensive hydrogen bonding. The inter-phosphoric acid hydrogen bonding is partially maintained in aqueous solutions above 50% solutions.
Phosphoric acid is very stable and shows no oxidation chemistry below 350 °C. As expected, phosphoric acid is tribasic, (8.4.45). The anions H2PO4- and HPO42- have particular names, dihydrogen phosphate and monohydrogen phosphate, respectively.
$\rm PO(OH)_3 \xrightleftharpoons[pK_1=2.15] [PO_2(OH)_2]^- \xrightleftharpoons[pK_2=7.1] [PO_3(OH)]^{2-} \xrightleftharpoons[pK_3=12.4] [PO_4]^{3-}$
Many salts of are known for all three anions; those with phosphate (PO42-) are often insoluble in water. Many coordination complexes are known, especially with M3+ and M4+ ions.
Phosphorous esters
The alkyl and aryl homologs of phosphoric acid (i.e., the phosphate triester O=P(OR)3) are prepared by the reaction of P4O10 with the appropriate alcohol. The triesters are good solvents and Lewis basic ligands with coordination via the oxide moiety. The diesters and monoesters can also be made, and they are important in biochemical processes.
Phosphite triesters, P(OR)3, may be made by the reaction of PCl3 with alcohols or phenols in the presence of an organic base as the acceptor for the HCl formed. Alternatively, they can be prepared directly from white phosphorous, (8.4.46).
$\rm P_4 + 6 OR^- + 6 CCl_4 + 6 ROH \rightarrow 4 P(OR)_3 + 6 CHCl_3 + 6 Cl^-$
Phosphite triesters are readily oxidized to the appropriate phosphate triester, (8.4.47). They also react with alkyl halides to form the dialkyl phosphonate via the Michaelis-Arbusov reaction (8.4.48).
$\rm 2 P(OR)_3 + O_2 \rightarrow 2 \text{O=P}(OR)_3$
$\rm P(OR)_3 + R'X \rightarrow \text{O=P}(OR)_2R' + RX$
Polyphosphates
Polyphosphates contain the PO4 unit, and the simplest example (pyrophosphate or diphosphate, (Figure $18$a) can be considered as condensation products of monohydrogen phosphate, (8.4.49). Longer chains can be formed, e.g., the triphosphate P3O105- (Figure $18$b). The formation and reverse (hydrolysis) reaction are slow, but are readily catalyzed, e.g., by enzymes.
$\rm 2 HPO_4^{2-} \xrightarrow{\Delta} [O_3P\text{-O-}PO_3]^{4-} + H_2O$
The monoesters of the diphosphate and triphosphate are very important in biological processes. In particular the conversion of the triphosphate ATP (adenosine triphosphate) to the diphosphate ADP (adenosine diphosphate) by the transfer of a phosphate group is important in the energetics of biological reactions.
Cyclic polyphosphates
Also known as meta phosphates, cyclic phosphates are also made from fused PO4 units. The simplest, P3O93-, is shown in Figure $19$. Slow addition of water to P4O10 results in the formation of the tetrameric polyphosphate, P4O124- (Figure $20$), which is known as calgon due to its ability to complex Ca2+ as well as other metals. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/08%3A_Group_15_-_The_Pnictogens/8.04%3A_Oxides_and_Oxoacids.txt |
Phosphorous trihalides
Phosphorous trihalides, PX3, are produced from the direct reaction of phosphorous with the appropriate halogen, (8.5.1). The fluoride is readily made from the halide exchange reaction of PCl3 with fluoride salts, (8.5.2). Mixed trihalides are formed by halide exchange, (8.5.3).
$\rm P_4 + 6 Cl_2 \rightarrow 4 PCl_3$
$\rm 2 PCl_3 + 3 CaF_2 \rightarrow 2 PF_3 + 3 CaCl_2$
$\rm PCl_3 + PBr_3 \rightleftharpoons PCl_2Br + PClBr_2$
A summary of the physical properties of the phosphorous trihalides is given in Table $1$. All the compounds have a pyramidal structure in the vapor phase and in solution.
Table $1$: Selected physical properties of the phosphorous trihalides.
Compound Mp (°C) Bp (°C) P-X (Å) X-P-X (°)
PF3 -151.5 -101.8 1.56 96.3
PCl3 -93.6 76.1 2.04 100
PBr3 -41.5 173.2 2.22 101
PI3 61.2 200 (dec.) 2.43 102
The phosphorous trihalides hydrolyze to phosphoric acid, (8.5.4), and undergo alcoholysis to form the trialkyl phosphite derivative, (8.5.5). Phosphorous trifluoride is only slowly hydrolyzed by water, but reacts readily with alakaline solutions. In contrast, phosphorus triiodide is an unstable red solid that reacts violently with water. Phosphorous trichloride in particular is an excellent synthon for most trialkyl phosphines, (8.5.6).
$\rm PCl_3 + 3 H_2O \rightarrow H_3PO_3 + 3 HCl$
$\rm PCl_3 + 3 HOR \rightarrow P(OR)_3 + 3 HCl$
$\rm PCl_3 + 3 RMgBr \rightarrow PR_3 + 3 MgBrCl$
As with other P(III) compounds such as trialkyl phosphines, phosphorous trihalides can be oxidized to the analogous phosphene oxide, e.g., (8.5.7).
$\rm 2 PCl_3 + O_2 \rightarrow \text{O=P}Cl_3$
Phosphorous trihalides form Lewis acid-base complexes with main group metals, but the bonding to dn (n ≠ 0) transition metals occurs in the same manner to that of trialkyl phosphine, i.e., with dπ-pπ back donation. In fact one of the first examples of complexation of phosphorous to a low oxidation metal was the formation of PF3 complexes with Fe-porphyrin.
Phosphorous pentahalides
The reaction of white phosphorous with excess halogen yields the pentahalides, (8.5.8). However, the pentafluoride is best prepared by halide exchange, (8.5.9).
$\rm P_4 + 10 Cl_2 \rightarrow 4 PCl_5$
$\rm 2 PCl_5 + 5 CaF_2 \rightarrow 2 PF_3 + 5 CaCl_2$
The pentaiodide is unknown, however, selected physical properties are given for the others in Table $2$.
Table $2$: Selected physical properties of phosphorous pentahalides.
Compound Mp (°C) Bp (°C)
PF5 -93.78 -84.5
PCl5 166.8 160 (subl.)
PBr5 100 106.0
The structures of the phosphorous pentahalides are all trigonal bipyramidal in the gas phase (Figure $1$). Phosphorous pentafluoride maintains the trigonal bipyramidal structure in solid state, but the chloride and bromide are ionic solids, [PCl4]+[PCl6]- and [PBr4]Br (Figure $2$), respectively. The tetrahedral [PCl4]+ ion is also formed with the reaction of PCl5 with other chloride ion acceptors, (8.5.10) and (8.5.11).
$\rm PCl_5 + TiCl_4 \rightarrow [PCl_4]^+[Ti_2Cl_9]^-$
$\rm PCl_5 + NbCl_5 \rightarrow [PCl_4]^+[NbCl_6]^-$
All of the pentahalides undergo thermal decomposition, (8.5.12).
$\rm PX_5 \rightleftharpoons PX_3 + X_2$
Careful hydrolysis of the pentahalides yields the oxide of the appropriate trihalide, (8.5.13), while excess hydrolysis yields phosphoric acid, (8.5.14).
$\rm PX_5 + H_2O \rightarrow \text{O=}PX_3 + HX$
$\rm PX_5 + 4 H_2O \rightarrow H_3PO_4 + HX$ | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/08%3A_Group_15_-_The_Pnictogens/8.05%3A_Halides_of_Phosphorous.txt |
Thumbnail: A sample of sulfur a member of the oxygen group of elements. (Public Domain; Ben Mills).
09: Group 16
The elements
The Group 16 elements have a particular name chalcogenes. Table 1 lists the derivation of the names of the halogens.
Table $1$: Derivation of the names of each of the Group 16(VI) elements.
Element Symbol Name
Oxygen O Greek oxys (sharp, from the taste of acids) and genēs (producer)
Sulfur (sulphur) S From the Latin sulphurium
Selenium Se Greek selene meaning Moon
Tellurium Te Latin tellus meaning earth
Polonium Po Named after Poland, Latin Polonia
Note
In Latin, the word is variously written sulpur, sulphur, and sulfur. It is an original Latin name and not a classical Greek loan, so the ph variant does not denote the Greek letter φ. Sulfur in Greek is thion, whence comes the prefix thio- to donate a sulfur derivative, e.g., a thioketone, R2C=S. The simplification of the Latin words p or ph to an f appears to have taken place towards the end of the classical period. The element has traditionally been spelled sulphur in the United Kingdom, India, Malaysia, South Africa, Australia, Ireland, and Canada, but sulfur in the United States. IUPAC adopted the spelling “sulfur” in 1990, as did the Royal Society of Chemistry Nomenclature Committee in 1992.
Discovery
Oxygen
The 2nd century BC Greek writer, Philo of Byzantium, observed that inverting a jar over a burning candle and surrounding the jar’s neck with water resulted in some water rising into the neck. He incorrectly ascribed this to the idea that part of the air in the vessel were converted into the element fire and thus were able to escape through pores in the glass. Much later Leonardo da Vinci (Figure $1$) suggested that this effect was actually due to a portion of air being consumed during combustion.
By the late 17th century, Robert Boyle (Figure $2$) showed that air is necessary for combustion. His work was expanded by English chemist John Mayow (Figure $3$) by showing that fire requires only a part of air that he called spiritus nitroaereus or just nitroaereus.
The reactive nature of nitroaereus was implied by Mayow from his observation that antimony (Sb) increased in weight when heated in air. He also suggested that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body; both concepts that were proven to be correct.
Robert Hooke (Figure $4$), Ole Borch (Figure $5$), Mikhail Lomonosov (Figure $6$), and Pierre Bayen (Figure $7$) all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as an element, probably since the prevalence at that time of the phlogiston, and their attempts to fit their experimental observations to that theory.
Note
The phlogiston theory was postulated in 1667 by the German alchemist J. J. Becher, and modified in 1731 by the chemist Georg Ernst Stahl. Phlogiston theory stated that all combustible materials were made of two parts. One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated component was thought to be its true form, or calx. Highly combustible materials that leave little residue (e.g., wood) were thought to mostly comprise of phlogiston, while non-combustible substances that corrode (e.g., iron) contained very little phlogiston. Air did not play a role in phlogiston theory, instead, it was based on observations of what happens when something burns, that most common objects appear to become lighter and seem to lose something in the process. However, one observation that overturned phlogiston theory was that metals, gain weight in rusting when they were supposedly losing phlogiston!
Oxygen was first discovered by Carl Wilhelm Scheele (Figure $9$ ) by heating mercuric oxide (HgO). Scheele called the gas fire air because it was the only known supporter of combustion. He wrote an account of this discovery in a manuscript (Treatise on Air and Fire) submitted in 1775. Unfortunately for Scheele his work was not published until 1777. In August 1774, an experiment conducted by Joseph Priestley (Figure $9$) sunlight on mercuric oxide (HgO) inside a glass tube, which liberated a gas he named dephlogisticated air. Priestley noted that candles burned brighter in this gas. He even went as far as breathing the gas himself, after which he wrote: "The feeling of it to my lungs was not sensibly different from that of common air, but I fancied that my breast felt peculiarly light and easy for some time afterwards." Priestley published his findings in 1775. Because he published his findings first, Priestley is usually given credit for the discovery of what became known as oxygen.
Interestingly, Lavoisier (Figure $10$) claimed to have discovered this new substance independently. However, Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Furthermore, Scheele also posted a letter to Lavoisier on September 30, 1774 that described his own discovery. Lavoisier never acknowledged receiving it, however, a copy of the letter was found in Scheele's belongings after his death.
Raoul Pictet (figure 11) showed that by the evaporation of liquid sulfur dioxide (SO2), carbon dioxide could be liquefied, which in turn was evaporated to cool oxygen gas enough to liquefy it. Pictet reported his results on December 22, 1877. Two days later, Louis Cailletet (Figure $12$) announced his own method of liquefying oxygen. In both cases only a few drops could be produced, making analysis difficult. In 1891 James Dewar (Figure $13$) was able to produce enough liquid oxygen to study. However, it was the process developed independently by Carl von Linde (Figure $14$) and William Hampson (1854 - 1926).
Sulfur
Sulfur was known in ancient times and is referred to in the Bible. English translations of the Bible commonly referred to burning sulfur as brimstone, giving rise to the name of fire-and-brimstone sermons, in which listeners are reminded of the fate of eternal damnation that await the unbelieving and unrepentant. It is from this part of the Bible that Hell is implied to smell of sulfur (likely due to its association with volcanic activity). Sulfur ointments were used in ancient Egypt, while it was used for fumigation in Greece. A natural form of sulfur known as shiliuhuang was known in China since the 6th century BC. However, it was not until 1777 that Lavoisier (Figure $10$) convinced the scientific community that sulfur was an element and not a compound.
Selenium
The element was discovered in 1817 by Berzelius (Figure $15$), who found the element associated with tellurium. It was discovered as a byproduct of sulfuric acid production.
Tellurium
Tellurium was discovered in the 18th century in gold ore from the mines in Zlatna, Transylvania. In 1782 Müller von Reichenstein (Figure $16$), the Hungarian chief inspector of mines in Transylvania, concluded that the ore was bismuth sulfide. However, the following year, he reported that this was erroneous and that the ore contained mostly gold and an unknown metal very similar to antimony. After three years of work Müller determined the specific gravity of the mineral and noted the radish-like smell of the white smoke evolved when the new metal was heated. Nevertheless, he was not able to identify this metal and gave it the names aurum paradoxium and metallum problematicum, as it did not show the properties predicted for the expected antimony.
In 1789 Kitaibel (Figure $17$) also discovered the element independently in an ore from Deutsch-Pilsen which had been regarded as argentiferous molybdenite, but later he gave the credit to Müller. In 1798, the name was chosen by Klaproth (Figure $18$) who earlier isolated it from the mineral calaverite.
Polonium
Temporarily called radium F, polonium was discovered by Marie Curie and her husband Pierre Curie (Figure $19$) in 1898, but it was later named after Marie Curie's native land of Poland. At the time Poland was not an independent country, but partitioned under Russian, Prussian, and Austrian. It was Curie's hope that naming the element after her native land would publicize its lack of independence. Polonium was the first chemical element named to highlight a political controversy.
Abundance
The abundance of the chalcogenes is given in Table $2$.
Table $2$: Abundance of Group 16 elements.
Element Terrestrial abundance (ppm)
O 47 x 104 (Earth’s crust), constituent of water, 21 x 104 (atmosphere)
S 260 (Earth’s crust), 870 (sea water), 10-3 (atmosphere)
Se 0.05 (Earth’s crust), 5 (soil), 0.2 x 10-3 (sea water)
Te 5 x 10-3 (Earth’s crust), 0.03 (soil), 0.15 x 10-6 (sea water)
Po Trace (Earth’s crust)
Isotopes
The naturally abundant isotopes of the Group 16 elements are listed in Table $3$. All of the isotopes of polonium are radioactive.
Table $3$: Abundance of the non-synthetic isotopes of the Group 16 elements.
Isotope Natural abundance (%)
Oxygen-16 99.76
Oxygen-17 0.039
Oxygen-18 0.201
Sulfur-32 95.02
Sulfur-33 0.75
Sulfur-34 4.21
Sulfur-36 0.02
Selenium-74 0.87
Selenium-76 9.36
Selenium-77 7.63
Selenium-78 23.78
Selenium-80 49.61
Tellurium-120 0.09
Tellurium-122 2.55
Tellurium-123 0.89
Tellurium-124 4.74
Tellurium-125 7.07
Tellurium-126 18.84
Tellurium-128 31.74
Tellurium-130 34.08
There are 38 known nuclear isomers of tellurium with atomic masses that range from 105 to 142. Tellurium is the lightest element known to undergo alpha decay, with isotopes 106Te to 110Te being able to undergo this mode of decay.
Cigarettes: it is not only the smoke that kills, but also the radioactivity
Ever since the early 1960s, the presence of polonium-210 in tobacco smoke has been known. The world's biggest tobacco firms spent over 40 years trying to find ways to remove the polonium-210 without success: even to this day. However, they also never published the results, keeping the facts of the radioactive hazards from the consumer.
Radioactive polonium-210 is contained in phosphate fertilizers and is absorbed by the roots of plants (such as tobacco) and stored in its tissues. Tobacco plants fertilized by rock phosphates contain polonium-210, which emits alpha radiation estimated to cause about 11,700 lung cancer deaths annually worldwide.
Industrial production of the elements
Sulfur
Elemental sulfur is found near hot springs and volcanic regions in many parts of the world. Volcanic deposits are mined in Indonesia, Chile, and Japan. Significant deposits of sulfur also exist in salt domes along the coast of the Gulf of Mexico, and in eastern Europe and western Asia. The sulfur in these deposits is believed to come from the action of anaerobic bacteria on sulfate minerals. However, fossil-based sulfur deposits from salt domes are the basis for commercial production in the United States, Poland, Russia, Turkmenistan, and Ukraine. Sulfur is mainly extracted from natural sources by two processes: the Sicilian process and the Frasch process.
Sicilian process
First used in Sicily from where it takes its name, the Sicilian process was used in ancient times to get sulfur from rocks present in volcanic regions. The sulfur deposits are piled and stacked in brick kilns built on sloping hillsides, and with airspaces between them (Figure $20$). Then powdered sulfur is put on top of the sulfur deposit and ignited. As the sulfur burns, the heat melts the sulfur deposits, causing the molten sulfur to flow down the sloping hillside. The molten sulfur can then be collected in wooden buckets. The sulfur produced by the Sicilian process must be purified by distillation.
Frasch process
In 1867, sulfur was discovered in the caprock of a salt dome in Louisiana; however, it was beneath quicksand, which prevented mining. In 1894 Herman Frasch (Figure $21$), devised a method of sulfur removal using pipes to bypass the quicksand. The process proved successful, but the high cost of fuel needed to heat the water made the process uneconomic until the 1901 discovery of the Spindletop oil field in Texas (Figure $22$) provided cheap fuel oil to the region.
In the Frasch process three concentric pipes to extract sulfur at high purity directly out of the ground (Figure $23$). Superheated steam (160 °C) is pumped down the outermost pipe, which melts the sulfur. Hot compressed air is pumped down the innermost pipe, which serves to create foam and pressure. The resulting molten sulfur foam is then expelled through the middle pipe. The Frasch process produces sulfur with 99.5% purity, which needs no further purification.
Most of the world's sulfur was obtained using the Frasch process until the late 20th century, when sulfur recovered from petroleum sources (recovered sulfur) became more commonplace.
Selenium
Elemental selenium is a rare mineral, and most elemental selenium comes as a byproduct of refining copper or producing sulfuric acid. Isolation of selenium begins by oxidation with sodium carbonate to produce selenium dioxide. The selenium dioxide is then mixed with water and the solution is acidified to form selenous acid (oxidation step). Selenous acid is bubbled with sulfur dioxide (reduction step) to give elemental selenium.
Elemental selenium produced by chemical reactions appears as the amorphous red form. When the red form is rapidly melted, it forms the black, vitreous form. The most thermodynamically stable and dense form of selenium is the electrically conductive gray (trigonal) form, which is composed of long helical chains of selenium atoms (Figure $24$). The conductivity of this form is notably light sensitive. Selenium also exists in three different deep-red crystalline monoclinic forms, which is composed of Se8 molecules, similar to many allotropes of sulfur.
Tellurium
The principal source of tellurium is from anode sludges produced during the electrolytic refining of copper. Treatment of 500 tons of copper ore typically yields 1 lb (0.45 kg) of tellurium. The anode sludges contain the selenides and tellurides of the noble metals in compounds with the formula M2Se or M2Te (M = Cu, Ag, Au). At temperatures of 500 °C the anode sludges are roasted with sodium carbonate (Na2CO3) under air. The metals are reduced to the metals, while the tellurium is converted to sodium tellurite, (9.1.1).
$\rm M_2Te + O_2 + Na_2CO_3 \rightarrow Na_2TeO_3 + 2 M + CO_2$
Tellurites can be extracted from the mixture with water and are normally present as hydrotellurites HTeO3- in solution. Selenates are also formed during this process, but they can be separated by adding sulfuric acid. The hydrotellurites are converted into the insoluble tellurium dioxide while the selenites stay in solution, (9.1.2).
$\rm HTeO_3^- + OH^- + H_2SO_4 \rightarrow TeO_2 + 2 SO_4^{2-} + 2 H_2O$
The reduction to the metal is done either by electrolysis or by reacting the tellurium dioxide with sulfur dioxide in sulfuric acid, (9.1.3).
$\rm TeO_2 + 2 SO_2 + 2 H_2O \rightarrow Te + SO_4^{2-} + 4 H^+$
Physical properties
The physical properties of the Group 16 elements encompasses a gas (O2), a non-metallic solid (S2), and metals (Se, Te, Po), Table $4$.
Table $4$: Selected physical properties of the Group 16 elements.
Element Mp (°C) Bp (°C) Density (g/cm3)
O -218.79 -182.95 1.429 g/L
S 115.21 444.6 1.819
Se 221 685 4.81 (gray), 4.39 (alpha), 4.28 (vitreous)
Te 449.51 988 6.24 (solid), 5.70 (liquid)
Po 254 962 9.196 (alpha), 9.398 (beta)
Vapor phase
The lighter Group 16 elements form X2 dimers in the vapor phase. Sulfur also forms higher allotropes in the vapor phase (e.g., S8 and S6), while selenium and tellurium forms atomic vapor in addition to the dimmers. Unlike dihydrogen, however, the bonding is associated with the molecular orbital combination of the two π-orbitals (Figure $25$). All of the dimeric X2 molecules are paramagnetic.
Solid state
While sulfur forms over 30 allotropes, the common form of sulfur is cyclooctasulfur (S8) has three main allotropes: Sα, Sβ, and Sγ. The orthorhombic form (Sα) is more stable up to 95 °C, while the β-form is the thermodynamic form. The lone pairs of electrons make the S-S-S bend (108°), resulting in S8 having the shape of a crown (Figure $26$). When sulfur melts the S8 molecules break up. When suddenly cooled, long chain molecules are formed in the plastic sulfur that, behave as rubber. Plastic sulfur transform into rhombic sulfur over time.
Elemental selenium produced in chemical reactions appears as the amorphous red form. When rapidly melted, it forms the black, vitreous form, which is usually sold industrially as beads. The most thermodynamically stable and dense form of selenium is the electrically conductive gray (trigonal) form, which is composed of long helical chains of selenium atoms (Figure $24$). The conductivity of gray selenium is light sensitive and is hence used in photocopiers. Selenium also exists in three different deep-red crystalline monoclinic forms, which are composed of Se8 molecules, similar to many allotropes of sulfur. Unlike sulfur, however, selenium does not undergo the changes in viscosity when heated.
Tellurium is a crystalline metal with a triginal structure (a = 4.4572 Å, c = 5.929 Å). Polonium has a simple cubic structure in it’s a form (a = 3.359 Å)
Compounds of the Group 16 elements.
The chemistry of the Group 16 elements is dominated by the stability of the -2 oxidation state and the noble gas configuration of the X2- anion.
Oxidation state
The electronegativity of oxygen (3.5) results in it having predominantly -2 oxidation state, however, sulfur, selenium and tellurium all for compounds with higher oxidation states, especially with oxygen (Table $5$).
Table $5$: Examples of oxidation states in compounds of the Group 16 elements.
Element -2 -1 +4 +6
O Na2O, H2O H2O2 - -
S H2S H2S2 SO2 H2SO4, SO3
Se H2Se H2Se2 SeO2 SeF4
Te H2Te tBu2Te2 TeO2 Te(OH)6
Catenation
Catenation is the ability of a chemical element to form a long chain-like structure via a series of covalent bonds. Although oxygen shows this property only in the existence of ozone, sulfur is second only to carbon in exhibiting this mode of combination; the chalcogens beyond sulfur show it to diminishing degrees, polonium having no tendency to catenate.
When aqueous metal sulfide salts are heated with elemental sulfur a range of polysulfide ions are formed, (9.1.4). When alkali polysulfides dissolve in polar solvents (e.g., DMF or DMSO) a deep blue solution is formed. The absorption (λmax = 610 nm) is associated with the radical anion, S3-. While, polyselenides and polytellurides are less common, the Se32- and Te32- ions are known.
$\rm{ S^{2-} +} n \rm S \rightarrow S_{(n+1)}^{2-}$
The term polysulfide often refers to a class of polymers with alternating chains of several sulfur atoms and hydrocarbon substituents. The general formula is R2Sn, where n ranges from 2 – 10. For the selenium and tellurium analogs the extent of catenation is far more limited.
Bibliography
• D. Barisic, S. Lulic, and P. Miletic, Water Research, 1992, 26, 607.
• M. Davies, History of Science, 1989, 22, 63.
• T. F. Kelley, Science, 1965, 149, 537.
• M. B. Power, J. W. Ziller, A. N. Tyler, and A. R. Barron, Organometallics., 1992, 11, 1055. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/09%3A_Group_16/9.01%3A_The_Group_16_Elements-_The_Chalcogens.txt |
Ozone (O3) is an allotrope of oxygen that is much less stable than the diatomic molecule (O2). Ground-level ozone is an air pollutant with harmful effects on the respiratory system, hile the ozone layer in the upper atmosphere filters potentially damaging ultraviolet light from reaching the Earth's surface.
The structure of ozone is bent, with C2v symmetry, similar to water (Figure $1$a). The central oxygen has sp2 hybridization with one lone pair. As a consequence of the bent structure, and the resonance hybridization (Figure $1$b) ozone is a polar molecule (dipole moment = 0.5337 D).
Ozone is made by the exposure of oxygen (O2) to an electric discharge. Ozone has a characteristic smell can be commonly smelled after a lightening strike; in fact the name ozone comes from the Greek ozein meaning to smell. In the laboratory, ozone can also be produced by electrolysis using graphite rod cathode, a platinum wire anode, and sulfuric acid (3 M) electrolyte. The half cell reactions are as follows:
$\rm 3 H_2O \rightarrow O_3 + 6 H^+ + 6e^- \space\space\space\space\space\space\space (\Delta E_o = -1.53V)$
$\rm 6 H^+ + 6 e^- \rightarrow 3 H_2 \space\space\space\space\space\space\space (\Delta E_o = 0 V)$
$\rm 2 H_2O \rightarrow O_2 + 4 H^+ + 4 e^- \space\space\space\space\space\space\space (\Delta E_o = -1.23 V)$
Ozone is also produced through photolysis of oxygen, (Equation \ref{9.2.4} and \ref{9.2.5}), both in the laboratory and the atmosphere.
$\rm O_2 \xrightarrow{h\nu} 2 O\cdot \label{9.2.4}$
$\rm O\cdot + O_2 \rightarrow O_3 \label{9.2.5}$
Ozone is a very strong oxidizing agent, and will readily oxidize a range of materials, e.g., Equations \ref{9.2.6} and \ref{9.2.7}. It will also oxidize metals (except gold, platinum, and iridium) to their highest oxidation state, e.g., Equation \ref{9.2.8}.
$\rm O_3 + CO \rightarrow CO_2 + O_2 \label{9.2.6}$
$\rm O_3 + 2 I^- + H_2O \rightarrow O_2 + 2OH^- + I_2 \label{9.2.7}$
$\rm 2 Cu^+_{(aq)} + 2 H_3O^+_{(aq)} + O_{2(g)} \rightarrow 2 Cu^{2+}_{(aq)} + 3 H_2O_{(l)} + O_{2(g)} \label{9.2.8}$
Metal ozanides, which contain the ozonide anion (O3-) are explosive and must be stored at cryogenic temperatures. Ozonides for all the alkali metals are known. KO3, RbO3, and CsO3 can be prepared from their respective superoxides.
$\rm KO_2 + O_3 \rightarrow KO_3 + O_2$
Ozone as a modulator of life on Earth
The Earth’s atmosphere acts as a source of O2 and a repository of CO2, but its also acts as a shield for life. First, nearly all meteorites burn up on entry because of the high temperatures generated by the friction of the atmosphere. Second, the atmosphere acts as a shield for high energy UV radiation.
Although UV radiation converts 7-dehydrocholesterol into vitamin D3 in the skin (Figure $2$), and is therefore useful, high energy UV destroys living cells. In fact the darkening we call a suntan is actually the body’s mechanism for preventing further UV damage. Sun burn and skin cancer are caused by relatively weak UV light that reaches the Earth’s surface, without the atmosphere we would be exposed to high energy UV that would be a hazard to all life on Earth. The Earth’s “sun screen” is ozone (O3). And without ozone in the upper atmosphere there would be no life on Earth.
The ozone layer is located in the lower portion of the stratosphere from approximately 10 km to 50 km above Earth, though the thickness varies seasonally and geographically. This layer contains over 91% of the ozone in Earth’s atmosphere and absorbs 93-99% of the sun's high frequency ultraviolet light. The ozone decomposes photolytically to O2 and molecular oxygen, (9.2.10), and it is this reaction that accounts for the UV protection of the atmosphere. Ozone is naturally regenerated by the exothermic reaction of the molecular oxygen with O2, (9.2.11).
$\rm O_3 + h\nu \rightarrow O_2 + O$
$\rm O + O_2 \rightarrow O_3$
The balance between ozone formation and destruction is thus an important mechanism for the protection of living organisms on the planet. While the ozone layer had been relatively constant on Earth for millions of years, the last 70 have seen a dramatic change including the increase in the polar hole in the ozone layer. The ozone hole is defined geographically as the area where the total ozone concentration is less than 220 Dobson Units.
One Dobson unit refers to a layer of ozone that would be 10 µm (micrometre) that is 1 x 10-5 m thick under standard temperature (25 °C) and pressure (1 atmosphere).
The ozone hole has steadily grown in size and length of existence over the past two and half decades. At present the size of ozone hole over Antarctica is estimated to be about 30 million sq.km (Figure $3$). | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/09%3A_Group_16/9.02%3A_Ozone.txt |
Despite the greatest industrial complex of the Roman Empire being the imperial grain mill at Barbegal near Arles in what is now southern France (Figure \(1\)), and the knowledge of gearing used for the mill, the waterwheel (the power source of the Barbegal’s power) was little used in the ancient world. This was probably due to the high slave population obviating the need for labor saving devices. However, it may also be that because the Roman Empire was very centralized, they could provide flour on a large scale from a few highly mechanized locations. Upon the fall of the Roman Empire, the knowledge of water power would have been lost were it not for the writing departments of churches and monasteries that continued to operate through the subsequent Dark Ages.
The Barbegal mill is probably the first example of industrial mass production. It consisted of eight pairs of waterwheels positioned on a 65-foot slope (Figure \(2\)). The wheels were turned by water that fell from a reservoir, which in turn was fed by a magnificent aqueduct. The sixteen wheels each powered two grindstones using a set of gear that allowed the horizontal shaft from the wheel to turn a vertical shaft on which the grindstones were positioned. Grain for the mill was imported from as far as Egypt, and the flour production was eight times than that required for the local population of 10,000, resulting in an export business. Unfortunately, upon the fall of Rome the technology of waterpower almost ceased since the small city-states set up had no need for industrial complexes.
There are two general designs of waterwheel. The first is powered by water falling from above the wheel and is called an overshot wheel (Figure \(3\)). The alternative design, the undershot wheel, relied on water flowing on a river or pond such that the current moved the paddles at the bottom of the wheel (Figure \(4\)).
Despite its fall from use, the waterwheel was not forgotten. Due to the writings of 14th century BC engineers, such as Vitruvius (Figure \(5\)), whose texts had been preserved in the libraries in churches and monasteries across Europe, the waterwheel powered mills made a resurgence between the fifth and the tenth century. In most cases mills were owned by the church, since they had the knowledge (from ancient texts) to construct the mills and, possibly just as important, the literacy to develop the accounting system for their profitable use. The church would lease mills to farmers on a time usage, and they would take payment in flour. The owners of waterwheels and their mills were the first barons of industry post the Dark Ages. In fact, the fact that the Saxon word for an aristocrat is Lord, which means loaf giver, suggests the importance of the waterwheel. By the end of the tenth century the waterwheel was in widespread use across Europe. The Domesday Book (the nationwide census carried out by the Normans after the invasion of England) listed nearly six thousand grain mills in 1089.
Despite the large number of waterwheels, the gearing used up until the ninth century was little different to that described by Vitruvius and used at Barbegal. Then around 890 AD the monastery of St Gall a new device was attached to the waterwheel. Instead of gearing to transfer power from horizontal to vertical rotation, a piece of wood was set into the shaft driven by the waterwheel (Figure \(6\)). What had been created was a cam, since as the shaft turns the protruding piece of wood its anything in its way. For example, the first recorded use, by the monks at St Gall, was to crush malt for beer. However, the cam could be made to trip a hammer with every rotation (pounding), or to act on a crank to turn a rotary motion into a horizontal back-and-forward motion (cutting), or to push down a level and activate a suction pump (raising water from a well), or operate a bellows (for a metal forge). The range of motions meant that waterpower could now be used for a wide range of industries. By the end of the tenth century there were waterwheels powering forge hammers, oil and silk mills, sugar cane crusher, tanning mills pounding leather, grinding stones, ore crushing mills (Figure \(7\)), and as fulling mills for the rapidly expanding trans-European textile industry.
As a consequence of the waterwheel and the cam, the period between the tenth and fourteenth centuries has come to be known as the Medieval Industrial Revolution. It is interesting to note that water played a key role in the driving force during the next Industrial Revolution four hundred years later – the steam engine. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/09%3A_Group_16/9.03%3A_Water-_The_Fuel_for_the_Medieval_Industrial_Revolution.txt |
Hydrogen peroxide (H2O2) is a very pale blue liquid but appears colorless in dilute solution. It is prepared by the oxidation of anthraquinol (shown below). The hydrogen peroxide is extracted with water from the anthraquinone solution and the 20 - 40% solution is purified by solvent extraction. An alternative process involves the oxidation of isopropanol in either the vapor or liquid phase at 100 °C and ca. 15 atm, (9.4.1). The products are separated by fractional distillation.
$\rm Me_2C(H)OH + O_2 \rightarrow Me_2C(OOH)OH \rightarrow Me_2\text{C=O} + H_2O_2$
In the gas phase H2O2 adopts a gauche conformation (Figure $1$), but there is only a low barrier to rotation about the O-O bond.
Hydrogen peroxide is a liquid at standard temperature and pressure (25 °C, 1 atm) due to the presence of strong hydrogen bonding similar to found in water. In fact, the liquid range for H2O2 (Mp = -0.43 °C, Bp = 150.2 °C) is actually broader than water, and it is slightly more viscous than water. Hydrogen peroxide has a density of 1.44 g/cm3, and is 106 times less basic that water.
As with water, H2O2 is a good solvent because of its polar nature and broad liquid temperature range, however, it is dangerous in its pure state due to its facile (ΔH = -99 kJ/mol) auto decomposition, (9.4.2), as well as its strong oxidizing nature.
$\rm 2 H_2O_2 \rightarrow 2 H_2O + O_2$
Hydrogen peroxide is usually sold as 3 - 12% solution for home use; however, laboratory and certain industrial applications require 30% solutions.
Note
Hydrogen peroxide should be stored in a cool, dry, well-ventilated area and away from any flammable or combustible substances. It should be stored in a container composed of non-reactive materials such as stainless steel or glass (other materials including some plastics and aluminum alloys may also be suitable). Because it breaks down quickly when exposed to light, it should be stored in an opaque container, and pharmaceutical formulations typically come in brown bottles that filter out light.
Aqueous solution are weakly acidic (K = 1.5 x 10-12), (9.4.3). However, there is no exchange of oxygen atoms between H2O2 and H2O in the liquid phase.
$\rm H_2O_2 + H_2O \rightleftharpoons HO_2^- + H_3O^+$
As expected hydrogen peroxide is a strong oxidizing agent, (9.4.4), however, it can also act as a reducing agent, (9.4.5).
$\rm H_2O_3 + 2 HI \rightarrow I_2 + 2 H_2O$
$\rm 5 H_2O_2 + 2 MnO_4 + 6 H^+ \rightarrow 2 Mn^{2+} + 8 H_2O + 5 O_2 \uparrow$
9.05: Hydrogen Peroxide Providing a Lift for 007
In the pre-credit sequence of the 1965 film Thunderball, James Bond 007 (played by Sean Connery) uses a Jetpack to escape from two gunmen after killing Jacques Bouvar, SPECTRE Agent No. 6 (Figure \(1\)). The Jetpack was also used in the Thunderball posters, being the "Look Up" part of the "Look Up! Look Down! Look Out!" tagline (Figure \(2\)). The Jetpack returned in 2002 in Die Another Day (in which Pierce Brosnan played Bond) in the Q scene that showcased many other classic gadgets from previous Bond films.
The Jetpack was actually a real fully functional device: the Bell Rocket Belt. It was designed for use in the US Army, but was rejected because of its short flying time of 21-22 seconds. Powered by hydrogen peroxide (H2O2), it could fly about 250 m and reach a maximum altitude of 18 m, going 55 km/h. Despite its impracticality in the real world, the Jetpack made a spectacular debut in Thunderball. Although Sean Connery is seen in close-up during the takeoff and landings (Figure \(3\)), the main flight was actually piloted by Gordon Yeager (Fugre 4) and Bill Suitor (Figure \(5\)).
The Bell Rocket Belt is a low-power rocket propulsion device that allows an individual to safely travel or leap over small distances. All subsequent rocket packs were based on the construction design, developed in 1960-1969 by Wendell Moore. Moore's pack has two major parts:
1. Rigid glass-plastic corset (Figure \(6\)a), strapped to the pilot (Figure \(6\)b). The corset has a tubular metallic frame on the back, on which are fixed three gas cylinders: two with liquid hydrogen peroxide (Figure \(6\)c), and one with compressed nitrogen (Figure \(6\)d). When the pilot is on the ground, the corset distributes the weight of the pack to the pilot's back.
2. The rocket engine, able to move on a ball and socket joint (Figure \(6\)e) in the upper part of the corset. The rocket engine consists of a gas generator (Figure \(6\)f) and two pipes (Figure \(6\)g) rigidly connected with it, which end with jet nozzles with controlled tips (Figure \(6\)h). The engine is rigidly connected to two levers, which are passed under the pilot's hands. Using these levers the pilot inclines the engine forward or back and to the sides. On the right lever is the thrust control turning handle (Figure \(6\)i), connected with a cable to the regulator valve (Figure \(6\)j) to supply fuel to the engine. On the left lever is the steering handle, which controlled the tips of jet nozzles.
The operating principle of the Jetpack is shown in Figure. The hydrogen peroxide cylinders and compressed nitrogen cylinder are each at a pressure of ca. 40 atm or 4 MPa). To operate the pilot turns the engine thrust control handle, and opens the regulator valve (3 in Figure \(7\)). Compressed nitrogen (1 in Figure \(7\)) displaces liquid hydrogen peroxide (2 in Figure \(7\)), which enters the gas generator (4 in Figure \(7\)). In the gas generator, the hydrogen peroxide contacts the catalyst and is decomposed. The catalyst consists of thin silver plates, covered with a layer of samarium nitrate. The resulting hot high-pressure mixture of steam and gas enters two pipes, which emerge from the gas generator. These pipes are covered with a layer of heat insulator to reduce loss of heat. The hot gas enters the jet nozzles, where first they are accelerated, and then expand, acquiring supersonic speed and creating reactive thrust. The whole construction is simple and reliable; the rocket engine has no moving parts. The pack has two levers, rigidly connected to the engine installation. Pressing on these levers, the pilot deflects the nozzles back, and the pack flies forward. Accordingly, raising this lever makes the pack move back. It is possible to lean the engine installation to the sides (because of the ball and socket joint) to fly sideways. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/09%3A_Group_16/9.04%3A_Hydrogen_Peroxide.txt |
Size
Table \(1\) summarizes the comparative sizes of oxygen and sulfur.
Table \(1\): Comparison of physical characteristics for oxygen and sulfur.
Element Atomic radius (Å) Covalent radius (Å) Ionic radius (Å) van der Waal radius (Å)
Oxygen 0.48 0.66 1.40 1.52
Sulfur 0.88 1.05 1.84 1.80
Electronegativity
Sulfur is less electronegative than oxygen (2.4 and 3.5, respectively) and as a consequence bonds to sulfur are less polar than the corresponding bonds to oxygen. One significant result in that with a less polar S-H bond the subsequent hydrogen bonding is weaker than observed with O-H analogs. A further consequence of the lower electronegativity is that the S-O bond is polar.
Bonds formed
Sulfur forms a range of bonding types. As with oxygen the -2 oxidation state prevalent. For example, sulfur forms analogs of ethers, i.e., thioethers R-S-R. However, unlike oxygen, sulfur can form more than two covalent (non-dative) bonds, i.e., in compounds such as SF4 and SF6.
Such hypervalent compounds were originally thought be due to the inclusion of low energy d orbitals in hybrids (e.g., sp3d2 for SF6); however, a better picture involves a combination of s and p ortbitals in bonding (Figure \(1\)). Any involvement of the d orbitals is limited to the polarization of the p orbitals rather than direct hydridization. In this regard SF6 represents the archetypal hypervalent molecule. Finally, sulfur can form multiple bonds, e.g., Me2S=O.
Catenation
Catenation is defined as the ability of a chemical element to form a long chain-like structure via a series of covalent bonds. Oxygen’s extent of catenation is limited to ozone (O3) and peroxides (e.g., R-O-O-R). In contrast, the chemistry of sulfur is rich in the formation of multiple S-S bonds.
While elemental sulfur exists as a diatomic molecule (i.e., S2) in the gas phase at high temperatures, sulfur vapor consists of a mixture of oligomers (S3 to S8) as a temperature dependant equilibrium. In the solid state the formation of Sn dominates, and sulfur exists as a range of polymorphs in which extended S-S bonding occurs in either rings of 6 to 20 atoms (e.g., Figure \(2\)) or chains (catenasulfur).
The higher level of catenation for sulfur is due to the greater strength of a S-S bond (226 kJ/mol) as compared to the O-O bond (142 kJ/mol). In general the homoleptic bond strength is expected to decrease going down a period of the Periodic Table. The reason for the unexpected weakness of the O-O bond is that the electronegative oxygen atoms repel each other and thus weaken the bond.
9.07: Chalconide Hydrides
Dihydrides
The hydrides of sulfur, selenium and tellurium are all extremely toxic gases with repulsive smells. Hydrogen sulfide (H2S) is very toxic, in fact it is more than 5x as toxic as HCN (Table $1$). Hydrogen sulfide is considered a broad-spectrum poison, meaning that it can poison several different systems in the body, although the nervous system is most affected. It forms a complex bond with iron in the mitochondrial cytochrome enzymes, thereby blocking oxygen from binding and stopping cellular respiration. Exposure to low concentrations can result in eye irritation, a sore throat and cough, nausea, shortness of breath, and fluid in the lungs. Long-term, low-level exposure may result in fatigue, loss of appetite, headaches, irritability, poor memory, and dizziness.
Table $1$: Toxicity levels for hydrogen sulfide.
Concentration (ppm) Biological effect
0.00047 Threshold.
10–20 Borderline concentration for eye irritation.
50–100 Eye damage.
100–150 Olfactory nerve is paralyzed and the sense of smell disappears, often together with awareness of danger.
320–530 Pulmonary edema with the possibility of death.
530–1000 Stimulation of the central nervous system and rapid breathing, leading to loss of breathing.
800 Lethal concentration for 50% of humans for 5 minutes exposure (LC50).
+1000 immediate collapse with loss of breathing, even after inhalation of a single breath.
Each of the hydrides is prepared by the reaction of acid on a metal chalcogenide, e.g., (9.7.1) and (9.7.2). The unstable H2Po has been prepared by the reaction of HCl on Po metal.
$\rm Fe + S \rightarrow FeS$
$\rm FeS + 2 HCl \rightarrow H_2S \uparrow + FeCl_2$
The thermal stability and bond strength of the dihydrides follows the trend:
$\rm H_2S > H_2Se > H_2Te > H_2Po$
While H2Se is thermodynamically stable to 280 °C, H2Te and H2Po are thermodynamically unstable.
All the dihydrides behave as weak acids in water. Thus, dissolution of H2S is water results in the formation of the conjugate bases, (9.7.4) and (9.7.5), with dissociation constants of 10-7 and 10-17, respectively.
$\rm H_2S + H_2O \rightleftharpoons H_3O^+ + SH^-$
$\rm SH^- + H_2O \rightleftharpoons H_3O^+ + S^{2-}$
Sulfanes
The propensity of sulfur for catenation means that while the hydrides of oxygen are limited to water (H2O) and hydrogen peroxide (H2O2), the compounds H2Sn where n = 2 - 6 may all be isolated. Higher homologs are also known, but only as mixtures. All of the sulfanes are yellow liquids whose viscosity increases with increased chain length.
A mixture of lower sulfanes is prepared by the reaction of sodium sulfides (Na2Sn) with HCl, (9.7.6). From this mixture the compounds H2Sn where n = 2 - 5 are purified by fractional distillation. However, higher sulfanes are made by the reaction of either H2S or H2S2 with sulfur chlorides, (9.7.7) and (9.7.8).
$\rm Na_2S_2 + 2 HCl \rightarrow 2 NaCl + H_2S_n$
$\rm 2 H_2S + S_nCl_2 \rightarrow 2 HCl + H_2S_{n+2}$
$\rm 2 H_2S_2 + S+nCl_2 \rightarrow 2 HCl + H_2S_{n+4}$ | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/09%3A_Group_16/9.06%3A_Comparison_of_Sulfur_to_Oxygen.txt |
Note
Alternative spellings of sulfurous and sulfuric acids are based upon the traditional UK spelling of sulphur, i.e., sulphurous and sulphuric acid.
Sulfur dioxide and sulfurous acid solutions
The combustion of sulfur results in the formation of gaseous sulfur dioxide, (9.8.1).
$\rm S + O_2 \rightarrow SO_2$
The bent structure of SO2 is shown in Figure $1$, and as a consequence of the sp2 hybridization the molecule is polar.
The modest boiling temperature of SO2 (-10 °C) means that it is readily liquefied and easily kept as a liquid at room temperature under a slight pressure. The liquid is associated by dipole-dipole attractions due to the polar nature of SO2. Liquid SO2 is a good solvent due to the polarity of the molecule; as a consequence it readily solubalizes polar compounds and salts. It is also convenient since it is easy to remove from reaction products by evaporation.
Sulfur dioxide is soluble in water forming aqueous solutions where most of the SO2 is maintained as a hydrogen-bonded hydrate, in a similar manner to that observed for aqueous solutions of carbon dioxide. At equilibrium in neutral water (no added base) a small fraction reacts, to give a mixture of bisulfite (HSO3-, Figure $2$a) and sulfite (SO32-, Figure $2$b), (9.8.2). The free acid does not to exist.
$\rm SO_{2(aq)} + H_2O \rightleftharpoons \underset{\underset{H-SO_3^-}{\downarrow}}{HOSO_{2(aq)}^-} \rightleftharpoons SO_{3(aq)}^{2-}$
Bisulfite undergoes a further equilibrium, (9.8.3), to form disulfite, whose structure is shown in Figure $2$c.
$\rm 2 HSO_3^- \rightleftharpoons S_2O_5^{2-} + H_2O$
Salts of these anions are known, and complexes of the sulfite ion are known (Figure $3$), while SO2 itself can act as a ligand to heavy metals.
The bisulfite ion has strong reducing properties, e.g., (9.8.4) and (9.8.5).
$\rm 2 Fe^{3+} + SO_3^{2-} + 2 OH^- \rightarrow 2 Fe^{2+} + H_2O + SO_4^{2-}$
$\rm 2 MnO_4^+ + 5 SO_3^{2-} + 6 H^+ \rightarrow 2 Mn^{2+} + 3 H_2O + 5 SO_4^{2-}$
Bisulfite is also reduced by zinc in the presence of additional SO2, (9.8.6), to form the highly reducing dithionite anion (Figure $2$d). Reaction of bisulfite with elemental sulfur yields the thiosulfate anion (Figure $2$e), (9.8.7)f.
$\rm SO_3^{2-} + SO_2 \xrightarrow{Zn} S_2O_4^{2-}$
$\rm SO_3^{2-} + S \rightarrow S_2O_3^{2-}$
Sulfur trioxide and sulfuric acid
Oxidation of sulfur dioxide in the presence of a catalyst (e.g., platinum) yields sulfur trioxide, (9.8.8), which may be condensed to a liquid at room temperature (Bp = 45 °C).
$\rm 2 SO_2 + O_2 \xrightarrow{Pt} 2 SO_3$
Liquid SO3 exists as a mixture of monomer and trimers (Figure $4$a and 4b), while as a solid (Mp = 16.9 °C) it forms polymers (Figure $4$c).
The reaction of SO3 with water results in the formation of sulfuric acid, H2SO4, as a viscous, hydrogen bonded liquid. Sulfuric acid is a strong protic acid, which in dilute solutions (in water) reacts as a dibasic acid, (9.8.9), forming bisulfate (HSO4-) and sulfate (SO42-) anions. A large number of salts are known for both anions. In addition, sulfate is known to act as a monodentate or bidentate ligand in coordination complexes.
$\rm H_2SO_4 \xrightleftharpoons[+ H^+]{- H^+} HSO_4^- \xrightleftharpoons[+ H^+]{- H^+} SO_4^{2-}$
The dissolution of SO3 in concentrated sulfuric acid yields very corrosive, fuming sulfuric acid, which contains some pyrosulfuric acid, (9.8.10).
$\rm H_2SO_4 + SO_3 \rightleftharpoons H_2S_2O_7$
WARNING
The corrosive properties of sulfuric acid are accentuated by its highly exothermic reaction with water. Burns from sulfuric acid are potentially more serious than those of comparable strong acids (e.g., hydrochloric acid), as there is additional tissue damage due to dehydration and particularly secondary thermal damage due to the heat liberated by the reaction with water.
Sulfur as a source of atmospheric pollution and acid rain
Sulfur dioxide is formed as a pollutant during the combustion of sulfur containing fuels, in particular coal. While the emission of SO2 itself leads to concerns it is its conversion to sulfuric acid in the form of acid rain that has been of concern for several decades. The pathway for the formation of sulfuric acid in the atmosphere is depandant on whether the reaction occurs in dry atmosphere or in clouds and rain.
Gaseous reactions in a dry atmosphere
In the dry atmosphere, gaseous sulfur dioxide reacts with the hydroxide radical (formed by the photochemical decomposition of ozone, (9.8.11) and (9.8.12), in the presence of a non-reactive gas molecule such as nitrogen, (9.8.12). The sulfurous acid, thus formed reacts with oxygen to generate sulfur trioxide, (9.8.13), which reacts with water to form sulfuric acid, (9.8.14).
$\rm O_3 + h\nu \rightarrow O* + O_2$
$\rm O* + H_2O \rightarrow 2 HO\cdot$
$\rm HSO_3 + O_2 \rightarrow HO_2 + SO_3$
$\rm SO_3 + H_2O \rightarrow H_2SO_4$
Measurements indicate that the conversion rate of SO2 to H2SO4 is 4% per hour on a clear sunny day, but the rate is slower during the winter.
Liquid phase reactions in clouds and rain
In the liquid phase SO2 reacts directly with water, (9.8.15). The bisulfite (HSO3-) is oxidized by hydrogen peroxide forming a forming bisulfate (HSO4-) solution, (9.8.16).
$\rm 2 SO_2 + 2 H_2O \rightarrow SO_3^{2-} + HSO_3^- + 3 H^+$
$\rm HSO_3^- + H_2O_2 \rightarrow HSO_4^- + H_2O$
Water soluble hydrogen peroxide is formed by the oxidation of water, (9.8.17).
$\rm HO_2 + HO_2 \rightarrow H_2O_2 + O_2$
The HO2 radical is formed by the photolysis of organic carbonyl compounds, e.g., formaldehyde in (9.8.18) and (9.8.19).
$\rm H_2\text{C=O} + h\nu \rightarrow H\cdot + HCO\cdot$
$\rm HCO + O_2 \rightarrow HO_2 + CO$
The conversion rate is independent of pH is very fast: almost 100% per hour in summer. However, the conversion is limited by the supply of hydrogen peroxide, which is often present in much lower levels than SO2. Thus, a reduction in sulfur dioxide emissions does not always correlate with a reduction of wet acid deposition. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/09%3A_Group_16/9.08%3A_Oxides_and_Oxyacids_of_Sulfur.txt |
Sulfur hexafluoride
Sulfur hexafluoride (SF6) is a gas at standard temperature and pressure (25 °C, 1 atm). The most common synthesis involves the direct reaction of sulfur with fluorine yields SF6.
$\rm S + 3 F_2 \rightarrow SF_6$
It should be noted that while SF6 is highly stable, SCl6 is not formed. The explanation of this difference may be explained by a consideration of the Born-Haber cycle shown in Figure $1$. A similar cycle may be calculated for SCl6; however, a combination of a higher dissociation energy for Cl2 and a lower S-Cl bond energy (Table $1$) provide the rational for why SCl6 is not formed.
Table $1$: Comparison of diatomic bond dissociation and S-X bond energy for the fluorine and chlorine analogs.
Bond dissociation energy kJ/mol Bond energy kJ/mol
D(F-F) 158 E(S-F) 362
D(Cl-Cl) 262 E(S-Cl) 235
The S-F bond length (1.56 Å) is very short and consistent with π-bonding in addition to σ-bonding. Like SiF62-, SF6 is an example of a hypervalent molecule (Figure $2$).
Sulfur hexafluoride is an unreactive, non toxic compound. Its inert nature provides one of its applications, as a spark suppressor. The hexafluoride is generally resistant to chemical attack, e.g., no reaction is observed with potassium hydroxide (KOH) at 500 °C. The low reactivity is due to SF6 being kinetically inert due to:
• Coordination saturation precluding associative reactions with nucleophiles.
• Strong S-F bond (360 kJ/mol) limiting dissociative reactions.
Thermodynamically SF6 should react with water (ΔH = -460 kJ/mol), but the rate factors are too great. Sulfur hexafluoride can be reduced with sodium in liquid ammonia, (9.9.2), or with LiAlH4. In each of these reactions the mechanism involves the formation of a radical, (9.9.3). The reaction with sulfur trioxide yields SO2F2, (9.9.4), however, the reactions with carbon or CS2 only occur at elevated temperatures (500 °C) and pressure (4000 atm).
$\rm SF_6 + 8 Na \rightarrow Na_2S + 6 NaF$
$\rm SF_6 + e^- \rightarrow SF_6^-$
$\rm SF_6 + 2 SO_3 \xrightarrow{\text{250 °C}} 3 SO_2F_2$
Sulfur monochloride pentafluoride
Although the hexachloride is unknown, it is possible to isolate the monochloride derivative (SF5Cl) by the oxidative addition of Cl-F across SF4.
$\rm SF_4 + ClF \rightarrow SF_5Cl$
Sulfur monochloride pentafluoride is a gas (Bp = -21 °C), but unlike SF6 it is fairly reactive due to the polarization of the S-Cl bond (Figure $3$), and as a consequence it reacts with water, (9.9.6).
$\rm SF_5Cl + 3 H_2 \rightarrow SO_3 + 5 HF + HCl$
Sulfur pentafluoride
Although SF5 does not exist as a stable molecule, the gaseous dimmer S2F10 (Bp = 29 °C) may be isolated from the photochemical hydrogen reduction of SF5Cl, (9.9.7).
$\rm 2 SF_5Cl + H_2 \xrightarrow{h\nu} S_2F_{10} + 2 HCl$
While the sulfur is octahedral in S2F10 (Figure $4$a) the S-S bond is weak and long (2.21 Å versus an expected 2.08 Å for a single S-S bond). Despite the apparently weak S-S bond, S2F10 shows almost no reactivity at room temperature; however, the S-S bond undergoes homoleptic cleavage at high temperatures. The resultant SF5. radicals disproportionate to give highly reactive fluoride radicals, (9.9.8), which is the source of the highly oxidative properties of S2F10.
$\rm 2 SF_5 \rightarrow 2 SF_4 + 2 F\cdot$
The SF5. fragment is stabilized by the addition of an alkyl radical, and thus, there are a large number of RSF5 derivatives known. Unlike, the chloride analog, these are very stable.
Sulfur tetrafluoride
Sulfur tetrafluoride (SF4) is prepared from sulfur dichloride and sodium fluoride in acetonitrile solution at 70 - 80 °C.
$\rm 3 SCl_2 + 4 NaF \rightarrow SF_4 + S_2Cl_2 + 4 NaCl$
The structure of SF4 (and its substituted derivatives RSF3) is based upon a trigonal bipyramidal structure with one of the equatorial sites being occupied by a lone pair (Figure $4$b). Unlike the hexafluoride, sulfur tetrachloride is a highly reactive compound. It hydrolyzes readily, (9.9.10), and is a useful fluorinating agent (Figure $5$).
$\rm SF_4 + 2 H_2O \rightarrow SO_2 + 4 HF$
Sulfur chlorides
The chlorination of molten sulfur yields the fowl smelling disulfur dichloride (S2Cl2). If the reaction is carried out with a catalyst such as FeCl3, SnI4 or I2, an equilibrium mixture containing sulfur dichloride (SCl2) is formed. However, the dichloride dissociates readily, (9.9.11), although it can be isolated as a dark red liquid if it distilled in the presence of PCl5. The reaction of chlorine at -80 °C with SCl2 or S2Cl2 allows for the formation of SCl4 as a yellow crystalline compound which dissociates above -31 °C. Sulfur chlorides are readily hydrolyzed. Sulfur chlorides are used to dissolve sulfur (giving species up to S100Cl2) for the vulcanization of rubber.
$\rm 2 SCl_2 \rightleftharpoons S_2Cl_2 + Cl_2$
In the vapor phase S2Cl2 has C2 symmetry (Figure $6$a) while that of SCl2 has C2v symmetry (Figure $6$b). | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/09%3A_Group_16/9.09%3A_Sulfur_Halides.txt |
Thumbnail: Chlorine gas in an ampoule. (CC-BY-SA; W. Oelen (http://woelen.homescience.net/science/index.html))
10: The Halogens
The Group 17 elements have a particular name: the halogens meaning born of salt. This is due to the formation of salts when they form compounds with a metal. Table $1$ lists the derivation of the names of the halogens.
Table $1$: Derivation of the names of each of the halogens.
Element Symbol Name
Fluorine F Latin fluere meaning to flow
Chlorine Cl Greek khlôros meaning pale green
Bromine Br Greek brómos meaning stench
Iodine I Greek odes meaning violet or purple
Astatine At Greek astatos, meaning unstable
Discovery
Fluorine
The mineral fluorspar (also known as fluorite) consists mainly of calcium fluoride and was described in 1530 by Agricola (Figure $1$) for its use as a flux. Fluxes are used to promote the fusion of metals or minerals, and it was from this use that fluorine derived its name. In 1670 Heinrich Schwanhard found that when he mixed fluorspar with an acid the fumes produced (hydrogen fluoride) etched the glasses he was wearing. Despite many researchers investigating the chemistry of hydrogen fluoride (HF) the elemental form of fluorine was not isolated until 1886 when Henri Moissan (Figure $2$) studied the electrolysis of a solution of potassium hydrogen difluoride (KHF2) in liquid hydrogen fluoride (HF). The mixture was needed because hydrogen fluoride is a non-conductor. The device was built with platinum/iridium electrodes in a platinum holder and the apparatus was cooled to −50 °C.
The generation of elemental fluorine from hydrofluoric acid proved to be exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. The victims became known as fluorine martyrs.
Chlorine
Archaeological evidence has shown that sodium chloride (known as table salt) has been used as early as 3000 BC and brine (the saturated water solution) as early as 6000 BC. It is thought that hydrochloric acid was probably known to alchemist Jābir ibn Hayyān (Figure $3$) around 800 AD, while aqua regia (a mixture of nitric acid and hydrochloric acid) began to be used to dissolve gold sometime before 1400 AD. Upon dissolving gold in aqua regia, chlorine gas is released along with other nauseating and irritating gases.
Chlorine was first prepared and studied in 1774 by Carl Wilhelm Scheele (Figure $4$), and therefore he is credited for its discovery despite his failing to establish chlorine as an element, mistakenly thinking that it was the oxide obtained from the hydrochloric acid. Regardless of what he believed, Scheele did isolate chlorine by reacting MnO2 (as the mineral pyrolusite) with HCl, (10.1.1).
$\rm 4 HCl + MnO_2 \rightarrow MnCl_2 + 2 H_2O + Cl_2$
Bromine
Bromine was discovered independently by two chemists Antoine Balard (Figure $5$) in 1825 and Carl Jacob Löwig (Figure $6$) in 1826.
Balard found bromide salts in the ash of sea weed from the salt marshes of Montpellier. The sea weed was used to produce iodine, but also contained bromine. Balard distilled the bromine from a solution of seaweed ash saturated with chlorine. The properties of the resulting substance resembled that of an intermediate of chlorine and iodine; with those results he tried to prove that the substance was iodine monochloride (ICl), but after failing to do so he was sure that he had found a new element and named it muride, derived from the Latin word muria for brine.
In contrast, Löwig isolated bromine from a mineral water spring in Bad Kreuznach. Löwig used a solution of the mineral salt saturated with chlorine and extracted the bromine with Et2O. After evaporation a brown liquid remained. Unfortunately, the publication of his results were delayed and Balard published first.
Iodine
Iodine was discovered by Bernard Courtois (Figure $7$) in 1811 when he was destroying the waste material from the production of saltpeter (KNO3) during gunpowder production. Saltpeter produced from French niter beds required sodium carbonate (Na2CO3), which could be isolated from seaweed washed up on the coasts of Normandy and Brittany. To isolate the sodium carbonate, seaweed was burned and the ash then washed with water; the remaining waste was destroyed by the addition of sulfuric acid (H2SO4). After adding too much acid, Courtois observed a cloud of purple vapor that crystallized on cold surfaces making dark crystals. Courtois suspected that this was a new element but lacked the money to pursue his observations. In supplying samples to his friends, Charles Desormes and Nicolas Clément, he hoped his research was to continue. On 29 November 1813, Dersormes and Clément made public Courtois’s discovery, describing the substance to a meeting of the Imperial Institute of France.
Astatine
The existence of eka-iodine had been predicted by Mendeleev (Figure $8$), but astatine was first synthesized in 1940 by Corson (Figure $9$), MacKenzie (Figure $10$), and Segrè (Figure $11$) at the University of California, Berkeley by bombarding bismuth with alpha particles.
Abundance
The abundance of the halogens is given in Table $2$.
Table $2$: Abundance of the halogens.
Element Terrestrial abundance (ppm)
F 950 (Earth’s crust), 330 (soil), 1.3 (sea water), 6 x 10-4 (atmosphere)
Cl 130 (Earth’s crust), 50 – 2000 (soil), 1.8 x 104 (sea water)
Br 0.4 (Earth’s crust), 5 – 40 (soil), 65 (sea water)
I 0.14 (Earth’s crust), 3 (soil), 0.06 (sea water), 60 x 10-3 (atmosphere)
At Trace in some minerals
Isotopes
The naturally abundant isotopes of the halogens are listed in Table $3$. All 33 isotopes of astatine are radioactive.
Table $3$: Abundance of the major isotopes of the halogens.
Isotope Natural abundance (%)
Fluorine-19 100
Chlorine-35 75.77
Chlorine-36 trace
Chlorine-37 24.23
Bromine-79 50.69
Bromine-81 49.31
Iodine-127 100%
While 19F is the only naturally abundant isotope of fluorine, the synthetic isotope, 18F, has half life of about 110 minutes, and is commercially an important source of positrons for positron emission tomography (PET). PET is a nuclear medicine imaging technique that produces a 3-D image of processes within the body. The system detects pairs of γ-rays emitted indirectly by a positron-emitting radionuclide (tracer), which is introduced into the body on a biologically active molecule.
Trace amounts of radioactive 36Cl exist in the environment at about 7 x 10-11%. 36Cl is produced in the atmosphere by the interaction of cosmic rays with 36Ar. In the ground 36Cl is generated through neutron capture by 35Cl or muon (an elemental particle similar to an electron) capture by 40Ca. 36Cl decays with a half-life of 308,000 years making it suitable for geologic dating in the range of 60,000 to 1 million years. However, due to the large amounts of 36Cl produced by irradiation of seawater during atmospheric detonations of nuclear weapons between 1952 and 1958, it is also sued as an event marker for 1950s water in soil and ground water.
Iodine has 37 isotopes of iodine, but only one, 127I, is stable. Of the radioactive isotopes, 129I (half-life 15.7 million years) is used for radiometric dating of the first 85 million years of solar system evolution. 129I is also a product of uranium and plutonium fission, and as a consequence of nuclear fuel reprocessing and atmospheric nuclear weapons tests, the natural signal has been swamped. As a consequence it can now be used as a tracer of nuclear waste dispersion into the environment. 129I was used in rainwater studies to track fission products following the Chernobyl disaster.
Due to preferential uptake of iodine by the thyroid, isotopes with short half lives such as 131I can be used for thyroid ablation, a procedure in which radioactive iodine is administered intravenously or orally following a diagnostic scan. The lower energy isotopes 123I and 125I are used as tracers to evaluate the anatomic and physiologic function of the thyroid.
Industrial production.
Industrial production of fluorine involves the electrolysis of hydrogen fluoride (HF) in the presence of potassium fluoride (KF) during which fluorine gas is formed at the anode and hydrogen gas is formed at the cathode (Figure $12$). The potassium fluoride (KF) is converted to potassium bifluoride (KHF2), (10.1.2), which is the electrolyte and intermediate to the fluorine and hydrogen, (10.1.3).
$\rm HF + KF \rightarrow KHF_2$
$\rm 2 KHF_2 \rightarrow 2 KF + 2 F_2 + H_2$
The HF is formed as a byproduct of the production of phosphoric acid, since phosphate-containing minerals contain significant amounts of calcium fluorides, which upon treatment with sulfuric acid release hydrogen fluoride, (10.1.4).
$\rm CaF_2 + H_2SO_4 \rightarrow 2 HF + CaSO_4$
Chlorine is generally manufactured by electrolysis of a sodium chloride solution (brine). The production of chlorine results in the co-products caustic soda (sodium hydroxide, NaOH) and hydrogen gas (H2). Chlorine can also be produced by the electrolysis of a solution of potassium chloride, in which case the co-products are hydrogen and caustic potash (potassium hydroxide). There are three industrial methods for the extraction of chlorine by electrolysis of chloride solutions, all proceeding by the same reaction at the cathode, (10.1.5), and anode, (10.1.6), which lead to the overall reaction, (10.1.7), where M = Na or K.
$\rm 2 H^+_{(aq)} + 2 e^- \rightarrow H_{2(g)}$
$\rm 2 Cl^-_{(aq)} \rightarrow Cl_{2(g)} + 2 e^-$
$\rm 2 MCl + 2 H_2O \rightarrow Cl_2 + H_2 + 2 MOH$
Bromine exists exclusively as bromide salts in the Earth’s crust, however, due to leaching, bromide salts have accumulated in sea water, but at a lower concentration than chloride. The majority of bromine is isolated from bromine-rich brines, which are treated with chlorine gas, flushing through with air. In this treatment, bromide anions are oxidized to bromine by the chlorine gas, (10.1.8).
$\rm 2 Br^- + Cl_2 \rightarrow 2 Cl^- + Br_2$
Two major sources of iodine are used for commercial production: the caliche (a hardened sedimentary deposit of calcium carbonate found in Chile) and the iodine containing brines of gas and oil fields in Japan and the United States. The caliche contains sodium nitrate (NaNO3); in which traces of sodium iodate (NaIO3) and sodium iodide (NaI) are found. During the production of sodium nitrate the sodium iodate and iodide is extracted. Iodine sourced from brine involves the acidification with sulfuric acid to form hydrogen iodide (HI), which is then oxidized to iodine with chlorine, (10.1.9). The aqueous iodine solution is concentrated by passing air through the solution causing the iodine to evaporate. The iodine solution is then re-reduced with sulfur dioxide, (10.1.10). The dry hydrogen iodide (HI) is reacted with chlorine to precipitate the iodine, (10.1.11).
$\rm 2 HI_{(aq)} + Cl_2 \rightarrow I_{2(aq)} + 2 HCl_{(aq)}$
$\rm I_2 + 2 H_2O + SO_2 \rightarrow 2 HI + H_2SO_4$
$\rm 2 HI + Cl_2 \rightarrow I_2 \downarrow + 2 HCl$
Physical properties
The physical properties of the halogens (Table $4$) encompasses gases (F2 and Cl2), a liquid (Br2), a non-metallic solid (I2), and a metallic metal (At).
Table $4$: Selected physical properties of the halogens.
Element Mp (°C) Bp (°C) Density (g/cm3)
F -219.62 -188.12 1.7 x 10-3 @ 0 °C, 101 kPa
Cl -101.5 -34.04 3.2 x 10-3 @ 0 °C, 101 kPa
Br -7.2 58.8 3.1028 (liquid)
I 113.7 184.3 4.933
At 302 337 ca. 7
Reactivity
All the halogens are highly reactive and are as a consequence of the stability of the X- ion are strong oxidizing agents (Table $5$).
Table $5$: Electrochemical reduction potential for halogens.
Reduction Reduction potential (V)
F2 + 2 e- → 2 F- 2.87
Cl2 + 2 e- → 2 Cl- 1.36
Br2 + 2 e- → 2 Br- 1.07
I2 + 2 e- → 2 I- 0.53
WARNING
Elemental fluorine (fluorine gas) is a highly toxic, corrosive oxidant, which can cause ignition of organic material. Fluorine gas has a characteristic pungent odor that is detectable in concentrations as low as 20 ppb. As it is so reactive, all materials of construction must be carefully selected and metal surfaces must be passivated. In high concentrations, soluble fluoride salts are also toxic and skin or eye contact with high concentrations of many fluoride salts is dangerous.
Use of chlorine as a weapon
Chlorine gas, also known as bertholite, was first used as a weapon in World War I by Germany on April 22, 1915 in the Second Battle of Ypres. At around 5:00 pm on April 22, 1915, the German Army released one hundred and sixty eight tons of chlorine gas over a 4 mile front against French and colonial Moroccan and Algerian troops of the French 45th and 78th divisions (Figure $13$). The attack involved a massive logistical effort, as German troops hauled 5730 cylinders of chlorine gas, weighing ninety pounds each, to the front by hand. The German soldiers also opened the cylinders by hand, relying on the prevailing winds to carry the gas towards enemy lines. Because of this method of dispersal, a large number of German soldiers were injured or killed in the process of carrying out the attack. Approximately 6,000 French and colonial troops died within ten minutes at Ypres, primarily from asphyxiation and subsequent tissue damage in the lungs. Many more were blinded. Chlorine gas forms hydrochloric acid when combined with water, destroying moist tissues such as lungs and eyes. The chlorine gas, being denser than air, quickly filled the trenches, forcing the troops to climb out into heavy enemy fire.
As described by the soldiers it had a distinctive smell of a mixture between pepper and pineapple. It also tasted metallic and stung the back of the throat and chest. The damage done by chlorine gas can be prevented by a gas mask, or other filtration method, making the fatalities from a chlorine gas attack much lower than those of other chemical weapons. The use as a weapon was pioneered by Fritz Haber (Figure $14$) of the Kaiser Wilhelm Institute in Berlin, in collaboration with the German chemical conglomerate IG Farben, who developed methods for discharging chlorine gas against an entrenched enemy. It is alleged that Haber's role in the use of chlorine as a deadly weapon drove his wife, Clara Immerwahr, to suicide. After its first use, chlorine was used by both sides as a chemical weapon (Figure $15$), but it was soon replaced by the more deadly gases phosgene and mustard gas.
Vapor phase
All the halogens form X2 dimers in the vapor phase in an analogous manner to hydrogen. Unlike dihydrogen, however, the bonding is associated with the molecular orbital combination of the two p-orbitals (Figure $16$). The bond lengths and energies are given in Table $6$.
Table $6$: Bond lengths and energies for halogens.
Element Bond length (Å) Energy (kJ/mol)
F2 1.42 158
Cl2 1.99 243
Br2 2.29 193
I2 2.66 151
Solid state
Iodine crystallizes in the orthorhombic space group Cmca (Figure $17$). In the solid state, I2 molecules still contain short I-I bond (2.70 Å).
Compounds of the halogens.
The chemistry of the halogens is dominated by the stability of the -1 oxidation state and the noble gas configuration of the X- anion.
Oxidation state
The use of oxidation state for fluorine is almost meaningless since as the most electronegative element, fluorine exists in the oxidation state of -1 in all its compounds, except elemental fluorine (F2) where the oxidation state is zero by definition. Despite the general acceptance that the halogen elements form the associated halide anion (X-), compounds with oxidation states of +1, +3, +4, +5, and +7 are common for chlorine, bromine, and iodine (Table $7$).
Table $7$: Examples of multiple oxidation states for the halogens.
Element -1 +1 +3 +4 +5 +7
Cl HCl ClF ClF3, HClO2 ClO2 ClF5, ClO3- HClO4
Br HBr BrCl BrF3 Br2O4 BrF5, BrO3- BrO4-
I HI ICl IF3, ICl3 I2O4 IO3- IO4-
Bibliography
• G. Agricola, De Re Metallica, Dover Publications, UK (1950)
• K. Christe, Inorg. Chem., 1986, 25, 3721. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/10%3A_The_Halogens/10.01%3A_The_Group_17_Elements-_The_Halogens.txt |
Elemental fluorine (F2) is the most reactive element. Fluorine combines directly with all other elements, except nitrogen and the lighter noble gases. It also reacts with many compounds forming fluorides, and many organic compounds inflame and burn in the gas. The highly reactive nature is due to the weak F-F bond (thermodynamically unstable), which provides a low activation energy to reactions (kinetically unstable). The ΔG for reactions is often large due to the strength of the resulting X-F bonds. The weak F-F bond (158 kJ/mol) is due to the small size (0.5 Å) and high nuclear charge of fluorine that result in a small overlap of the bonding orbitals and a repulsion between the non-bonding orbitals (lone pairs) on the two fluorine atoms.
Ionic salts
The ease of formation of F- anion is due to the high electron affinity of fluorine (-322 kJ/mol). Since the fluoride ion is small (1.33 Å) and the least polarizable anion (i.e., hard) it is stable in ionic lattices with metal cations in a high oxidation state (high charge), e.g., MnF4 and CrF5. In general the highest oxidation states for any metal are found with the fluoride salts. The large ionization energies needed to produce the cations are recovered by the high lattice energies.
Covalent compounds
The high electronegativity of fluorine means that it forms a single electron pair bond polar bond with a high ionic character. The polar nature of the bond means that there is a large inductive effect within a molecule. For example, perfluoroethanol (CF3CF2OH) has an acidity comparable to acetic acid.
The high strength of X-F bonds (Table $1$) is also due to the high ionic character (up to 50%) that results in a high activation energy for bond breaking. In contrast, the low polarizability of the fluorine means that the inter-molecular van der Waals bonds are very weak. Thus, even with very high molecular weights the boiling point can be very low, e.g, WF6, Bp = 17 °C, Mw = 297.84 g/mol.
Table $1$: Typical bond energies for X-F bonds.
Bond Bond energy (kJ/mol)
C-F 486
N-F 272
P-F 490
A wide range of fluoride complexes may be prepared from both metal (FeF63-, RuF6-, PtF62-, and SnF62-) and non-metal (BF4-, SiF62-, and PF6-) fluorides. While many fluorides are salts, when the metal is in its higher oxidation states (e.g., OsF6 and WF6), the formation of an ionic lattice with the appropriate cation (i.e., Os6+ and W6+ respectively) is energetically unfavorable.
Hydrogen fluoride
Hydrogen fluoride (HF) is converted to highly corrosive hydrofluoric acid upon contact with moisture. Pure hydrogen fluoride must be handled in metal or polythene vessels, while aqueous solutions will readily etch and dissolve standard laboratory glassware requiring the use of fluorinated polymer (e.g., Teflon) containers.
Hydrogen fluoride is synthesized by the reaction of a fluoride salt with a concentrated acid, (10.2.1). The HF vapor may be condensed, and then subsequently purified by distillation.
$\rm CaF_2 + H_2SO_4 \rightarrow CaSO_4 + 2 HF \uparrow$
The H-F bonding in hydrogen fluoride involves an electron pair bond with a high degree of ionic character. This results in a very polar H-F bond and a large dipole moment (1.86 D).
In the vapor phase, hydrogen fluoride is monomeric above 80 °C, but at lower temperatures it associates into oligomers and small polymers, e.g., cyclic (HF)6, as a consequence of strong intermolecular hydrogen bonds. As a pure liquid (Mp = -83 °C, Bp = 19.5 °C) hydrogen fluoride is extensively associated by strong hydrogen bonding to form zig-zag polymers (Figure $1$).
Hydrogen fluoride has a high dielectric constant (84.2) and as such is a good solvent for polar molecules. However, it is not a good solvent for salts (even fluorides) because it does not solvate cations too well. Despite this, it is useful as a solvent because it is non-oxidizing and easy to evaporate off products.
In a similar manner to water, hydrogen fluoride self-ionizes, (10.2.2). Salts of H2F+ are known and the F- anion is further solvated by the HF to form a series of salts, (10.2.3). The complex anion HF2- is also formed in aqueous solutions of hydrogen fluoride (pK = 0.7).
$\rm 2 HF \rightleftharpoons H_2F^+ + F^-$
$\rm F^- + \text{n HF} \rightleftharpoons HF_2^- + H_2F_3^- + H_3F_4^- \text{ etc.}$
Hydrogen fluoride is actually a weak acid in aqueous solution with a low pK = 3.5. In fact, HF is a weaker acid that the other halogen analogs:
$\rm HF<HCl<HBr<HI$
This trend is despite the fluorine being more electronegative than the other halogens, but is consistent with the strength of the H-F bond (568 kJ/mol).
Hydrogen fluoride is used as a non-oxidizing acid for the hydrolysis of proteins and acid catalyzed condensation reaction. The stability of its salt (HF2-) allow for the study of very strong acids, (10.2.5) and (10.2.6).
$\rm 2 HF + R_2\text{C=O} \rightarrow R_2\text{C=OH}^+ + HF^+_2$
$\rm 2 HF + H_2O \rightarrow H_3O^+ + HF_2^-$
The acidity of HF may be increased sufficiently by the addition of a fluoride acceptor (e.g., SbF5) to facilitate the reaction with a weak base such as benzene, (10.2.7).
$\rm C_6H_6 + HF + SbF_5 \rightarrow C_6H_7^+ + SbF_6^-$
Finally, hydrogen fluoride can be used in the synthesis of other fluorine-containing compounds:
$\rm 6 HF + KCl + PCl_5 \rightarrow K[PF_6] + 6 HCl \uparrow$
Organic fluorine compounds
Organic compounds in which some or all of the hydrogen atoms are replaced by fluorine have unique (and often important) properties. The high stability of fluorocarbon compounds is a consequence of the C-F bond energy (486 kJ/mol) in comparison with that of C-H (415 kJ/mol); however, while kinetically stable, fluorocarbons are not necessarily particularly thermodynamically stable.
Replacement of hydrogen with fluorine results in an increased density; since the small size of fluorine means that the minimal distortion or structural change occurs as a result of the substitution. As with metal salts, the weak inter-molecular forces means that completely fluorinated organic compounds have low boiling points. One attribute of the low inter-molecular forces is the low coefficient of friction for fluoropolymers such as polytetrafluoroethylene, commonly known as Teflon (Figure $2$).
Synthetic routes to fluorocarbon compounds
The simplest route to a fluorocarbon compounds involves the direct replacement of another halogen by a metal fluoride, (10.2.9). The driving force for this reaction depends on the free-energy difference of MF and MCl, which is related to the difference in lattice energies. Thus, the larger the metal cation, the more favored the reaction. In this regard, AgF and CsF are the most effective fluorination agent.
$\rm \text{R-Cl} + MF \rightleftharpoons \text{R-F} + MCl$
Anhydrous hydrogen fluoride (HF) reacts with chlorocarbon compounds in the presence of a catalyst such as SbCl5 or CrF4, (10.2.10). However, elevated temperatures (50 – 150 °C) and high pressures (50 – 500 psi) are required.
$\rm 2 CCl_4 + 3 HF \rightarrow CCl_2F_2 + CCl_3F + 3 HCl$
The direct replacement of hydrogen with fluorine is possible if the reaction is carried out under dilute conditions in the presence of a catalyst, (10.2.11).
$\rm C_6H_6 + 9 F_2 \xrightarrow{Cu} C_6F_{12} + 6 HF$
Sulfur tetrafluoride (SF4) is a particularly selective fluorination agent. It can be used to convert ketones to difluoro compounds, (10.2.12).
$\rm 2 R_2\text{C=O} + SF_4 \rightarrow R_2CF_2 + SO_2$ | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/10%3A_The_Halogens/10.02%3A_Compounds_of_Fluorine.txt |
Comparison to fluorine
To appreciate the chemistry of chlorine in comparison to that of fluorine it is necessary to appreciate the differences and trends between the elements. As may be seen in Table $1$, chloride is significantly larger than fluorine. In addition while chlorine is an electronegative element its electronegativity is significantly less than that of fluorine, resulting in less polar bonding.
Table $1$: Comparison of physical characteristics for fluorine and chlorine.
Element Ionic radius (Å) Covalent radius (Å) van der Waal radius (Å) Electronegativity
Fluorine 1.33 0.64 1.47 -4.1
Chlorine 1.81 0.99 1.75 -2.9
The X-Cl bond is an electron pair covalent bond with a highly polar nature. In this regard, chlorine is similar to fluorine. However, there are two key features with regard to chlorine’s bonding that differentiates it from fluorine.
1. Unlike fluorine, chlorine can form multiple covalent bonds, e.g., ClO4- and ClF3.
2. Unlike fluorine, chlorine can form π-bonds with oxygen, i.e., Cl=O.
The chloride ion (Cl-) forms salts with ionic lattices (e.g., NaCl) but also forms a wide range of complexes, e.g., [Fe(H2O)5Cl]2+ and [RhCl6]3-. Chloride also acts as a bridging ligand in which one, two or three chlorides can bridge two metal centers (Figure $1$).
Chloride (and bromide) bridges are usually bent, whereas fluoride bridges can be either linear or bent. As an example, BeF2 and BeCl2 are isostructural, consisting of infinite chains with bent bridges (Figure $2$). In contrast, transition metal pentahalides show different structures depending on the identity of the halide. This, TaCl5 dimerizes with bent bridges (Figure $3$a), while TaF5 forms a cyclic tetramer with linear fluoride bridges (Figure $3$b).
The halide bridge
The bridging halide bonds can be described by both Lewis and molecular orbital (MO) theory. In a simple picture, the lone pair of a terminal halide can be thought to act as a Lewis base donor ligand to the second Lewis acidic metal center. Indeed some bridging halides are asymmetric consistent with this view; however, the symmetrical ones can be described by a resonance form. From a molecular orbital point of view, the bridging halide is represented by a combination of two metal centered orbitals with two halogen orbitals.
Hydrogen chloride
Hydrogen chloride (HCl) is prepared by the reaction of concentrated sulfuric acid (H2SO4) with either NaCl or concentrated HCl solution.
Hydrogen chloride is a polar molecule with a dipole of 1.08 D. However, the lower polarity as compared to that of hydrogen fluoride (1.91 D) is consistent with the physical and chemical properties. Hydrogen chloride is a gas at room temperature (Mp = -114.25 °C, Bp = -85.09 °C), and its low boiling point is consistent with weak hydrogen bonding in the liquid state. While self-ionization, (10.3.1), is very small, liquid HCl dissolves some inorganic compounds to give conducting solutions, (10.3.2).
$\rm 3 HCl \rightleftharpoons H_2Cl^+ + HCl_2^-$
$\rm R_3N + 2 HCl \rightleftharpoons R_3NH^+ + HCl_2^-$
Hydrogen chloride is soluble (and reacts) in water, (10.3.3). The pKa of the reaction (-7.0) is larger than observed for fluorine (3.2) and as such HCl is a stronger acid than HF.
$\rm HCl + H_2O \rightleftharpoons H_3O^+ + Cl^-$
Oxides of Chlorine
Chlorine forms a series of oxides (Table $2$) in which the chlorine has the formal oxidation states +1, +4, +6, and +7. The physical properties of the oxides are summarized in Table $2$. While, the oxides of chlorine are not very stable (in fact several are shock sensitive and are prone to explode) the conjugate oxyacids are stable.
Table $2$: Physical properties of the oxides of chlorine.
Compound Mp (°C) Bp (°C)
Cl2O -116 4
ClO2 -5.9 10
Cl2O4 -117 44.5
Cl2O6 3.5 unstable
Cl2O7 -91.5 82
Dichlorine monoxide (Cl2O, Figure $4$a) is a yellowish-red gas that is prepared by the reaction of chlorine with mercury oxide, (10.3.4), or with a solution of chlorine in CCl4.
$\rm 2 Cl_2 + 2 HgO \rightarrow HgCl_2 \cdot HgO + Cl_2O$
When heated or subject to a spark, Cl2O explodes to Cl2 and O2. Dichlorine monoxide reacts with water to form an orange-yellow solution of hypochlorous acid, (10.3.5).
$\rm H_2O_{(g)} + Cl_2O_{(g)} \rightleftharpoons 2 HOCl_{(g)}$
Chlorine dioxide (ClO2) is a yellowish gas at room temperature and is commonly used in industry as an oxidizing agent. The best synthesis of ClO2 involves the reduction of potassium chlorate (KClO3) by oxalic acid at 90 °C, since the CO2 formed acts as a diluent for the highly explosive ClO2. On an industrial scale ClO2 is made by the exothermic reaction of sodium chlorate with SO2 in sulfuric acid, (10.3.6). The photolysis of ClO2 yields a dark brown solid with the formula Cl2O3; however, its facile explosive decomposition precludes study.
$\rm 2 NaClO_3 + SO_2 + H_2SO_4 \rightarrow 2 ClO_2 + 2 NaHSO_4$
The structure of ClO2 (Figure $4$b) is equivalent to SO2 with one extra electron, resulting in a paramagnetic unpaired electron species. Unusually, despite the unpaired electron configuration, ClO2 shows no tendency to dimerize. This is unlike the analogous NO2 molecule.
Dichlorine tetraoxide (Cl2O4) is commonly called chlorine perchlorate as a consequence of its structure (Figure $4$c). Dichlorine hexaoxide (Cl2O6) is an unstable red oil that has the ionic structure in the solid state: [ClO2]+[ClO4]-.
Dichlorine heptoxide (Cl2O7) is a relatively stable oil, that is prepared by the dehydration of perchloric acid at -10 °C, (10.3.7), followed by vacuum distillation. The structure of Cl2O7 (Figure $4$d) has been determined by gas phase electron diffraction.
$\rm HClO_4 \xrightarrow[-H_2O]{+ P_4O_5} Cl_2O_7$
The reaction of Cl2O7 with alcohols and amines yields alkyl perchlorates (ROClO3) and amine perchlorates (R2NClO3), respectively.
Fluorides of chlorine
Given the isolobal relationship between the halogens it is not surprising that the mixed dihalogens can be prepared, e.g., ClF, ICl, and BrCl. Chlorine fluoride is a highly reactive gas (Bp = -100.1 °C) that is a powerful fluorinating agent, and is prepared by the oxidation of chlorine by chlorine trifluoride, (10.3.8).
$\rm Cl_2 + ClF_3 \rightarrow 3 ClF$
The higher electronegativity of fluorine as compared to chlorine (Table $1$), and the ability of chlorine to form more than one bond, means that higher fluorides of chlorine are also known, i.e., ClF3 and ClF5. Chlorine trifluoride (CF3, Bp = 11.75 °C) is a useful fluorinating agent, that is prepared by the high temperature reaction of elemental chlorine and fluorine, is a useful fluorinating age. The gaseous pentafluoride (ClF5, Bp = -31.1 °C) is prepared by the reaction of potassium chloride with fluorine, (10.3.10).
$\rm Cl_2 + 3 F_2 \xrightarrow{200 °C} 2 ClF_3$
$\rm KCl + 3 F_2 \xrightarrow{200 °C} ClF_5 + KF$
The structure of ClF3 is T-shaped with two lone pairs on chlorine (Figure $5$a), while that of ClF5 is square pyramidal with a single lone pair on chlorine (Figure $5$b).
In general the halogen fluorides are very reactive; explosive reactions occur with organic compounds. They are all powerful fluorinating agents when diluted with nitrogen, and the order of reactivity follows:
$\rm ClF_3 > BrF_3 > BrF_5 > IF_7 > ClF > IF_5 > BrF$
Like most halogen fluorides, ClF, ClF3 and ClF5 all react with strong bases (e.g., alkali metal fluorides) to form anions, (10.3.12) and (10.3.13), and strong acids (e.g., AsF5 and SbF5) to form cations, (10.3.14), (10.3.15), and (10.3.16).
$\rm ClF + CsF \rightarro Cs^+ + ClF_2^-$
$\rm ClF_3 + CsF \rightarrow Cs^+ + ClF_4^-$
$\rm 2 ClF + AsF_5 \rightarrow FCl_2^+ + AsF_6^-$
$\rm ClF_3 + AsF_5 \rightarrow ClF_2^+ + AsF_6^-$
$\rm ClF_5 + SbF_5 \rightarrow ClF_4^+ + SbF_6^-$ | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/10%3A_The_Halogens/10.03%3A_Compounds_of_Chlorine.txt |
Table $1$ lists the various oxyacids of chlorine. The relative strengths increase with the number of oxygen atoms since the more there are, the greater is the extent to which the negative charge on the resulting anion can be delocalized.
Table $1$: Relative acidity of oxyacids of chlorine.
Oxyacid Formula pKa
Hypochlorous HOCl 7.5
Chlorous HClO2 1.9
Chloric HClO3 -2
Perchloric HClO4 -10
Hypochlorous acid
Hypochlorous acid (HOCl) can be made pure in the gas phase, (10.4.1), while strong acid solutions can be made from Cl2O. In contrast, dilute aqueous solutions are obtained with a suspension of mercury oxide to remove the chloride, (10.4.2).
$\rm H_2O_{(g)} + Cl_2O_{(g)} \rightleftharpoons 2 HOCl_{(g)}$
$\rm 2 Cl_2 + 2 HgO_{(s)} + H_2O \rightarrow 2 HOCl + HgO \cdot HgCl_2 \downarrow$
Solutions of the anion, OCl-, are obtained by electrolysis of brine solution; allowing the products to mix at low temperature, (10.4.3).
$\rm Cl_2 + 2 OH^- \rightarrow ClO^- + Cl^- + H_2O$
The anion (hypochlorite) is a good oxidant, (10.4.4) and (10.4.5), but can undergo disproportionation, (10.4.6); slowly at 25 °C, but fast above 80 °C.
$\rm ClO^- + NH_3 \rightarrow NH_2Cl + OH^-$
$\rm ClO^- + 2 I^- + 2 H^+ \rightarrow I_2 + Cl^- + H_2O$
$\rm 3 ClO^- \rightarrow 2 Cl^- + ClO_3^-$
Chlorous Acid
Chlorous acid (HOClO) is prepared by the reaction of ClO2 with base, (10.4.7), followed by the precipitation of the ClO2- salt with barium chloride, (10.4.8). The barium salt is dried and then reacted with a calculated amount of H2SO4, (10.4.9).
$\rm 2 ClO_2 + 2 OH^- \rightarrow ClO_2^- + ClO_3^- + H_2O$
$\rm 2 ClO_2^- + BaCl_2 \rightarrow 2 Cl^- + Ba(ClO_2)_2 \downarrow$
$\rm Ba(ClO_2)_2 + H_2SO_4 \rightarrow Ba(SO_4) + 2 HClO_2$
The pure acid is unknown since it is too unstable, however, salts can be prepared directly, e.g., (10.4.10).
$\rm 2 ClO_2 + Na_2O_2 \rightarrow 2 NaClO_2 + O_2$
The anion (ClO2-) is stable in alkaline solutions but in acid solutions decomposition occurs, (10.4.11).
$\rm 5 HClO_2 \rightarrow 4 ClO_2 + Cl^- + H^+ + 2 H_2O$
As with hypochlorite, the chlorite anion is a strong oxidant, (10.4.12).
$\rm ClO_2^- + 4 I^- + 4 H^+ \rightarrow 2 I_2 + 2 H_2O + Cl^-$
Chloric Acid
The chloric anion (ClO3-) is made from the reaction of chlorine gas with hot alkali (80 °C) or by the electrolysis of hot NaCl solution.
$\rm 3 Cl_2 + 6 OH^- \rightarrow ClO_3^- + 5 Cl^- + 3 H_2O$
To obtain a solution of the acid, ClO3- is precipitated as the barium salt, (10.4.14), which is removed, dried, and suspended in water and treated with a calculated amount of H2SO4, (10.4.15). The free acid cannot be isolated and a maximum concentration of only 40% can be obtained in water.
$\rm 2 ClO_3^- + BaCl_2 \rightarrow 2 Cl^- + Ba(ClO_3)_2 \downarrow$
$\rm Ba(ClO_3)_2 + H_2SO_4 \rightarrow Ba(SO_4) + 2 HClO_3$
The ClO3- anion is pyramidal both in solid salts and in solution, and many salts are known; however, those with organic cations are explosive. The anion is a strong oxidizing agent, (10.4.16) and (10.4.17), and it disproportionates slowly in solution, (10.4.18).
$\rm ClO_3^- + 6 I^- + 6 H^+ \rightarrow 3 I_2 + 3 H_2O + Cl^-$
$\rm ClO_3^- + 3 NO_2^- \rightarrow 3 No_3^- + Cl^-$
$\rm 4 ClO_3^- \rightarrow 3 ClO_4^- + Cl^-$
Perchloric Acid
The perchlorate anion (ClO4-) is best made by electrolytic oxidation of chlorate in aqueous solution, (10.4.19). Fractional distillation can concentrate the solution to 72.5% which is a constant boiling mixture. This concentration is moderately safe to use, however, 100% perchloric acid may be obtained by dehydration with H2SO4.
$\rm ClO_3^- + H_2O \rightarrow ClO_4^- + 2 H^+ + 2 e^-$
WARNING
Perchloric acid is a very dangerous liquid that will explode if traces of metal ions are present. It is also a very strong oxidizing agent that will convert organic compounds to CO2 and H2O.
Perchloric acid is a very strong acid that is fully ionized in aqueous solution, such that the salt [H3O][ClO4] can be isolated. Many other perchlorate salts are known, but those with organic cations are explosive. Perchlorate salts of metals are often used when studying complex formation in aqueous solution, because ClO4- is a very weak ligand (PF6- is better) and unlikely to form complexes itself. However, perchlorate does complex with +3 and +4 cations.
10.05: Bromine Trifluoride as a Solvent
WARNING
Bromine trifluoride is a toxic, colorless, and corrosive liquid with a pungent choking smell that is soluble in sulfuric acid but explodes on contact with water and organic compounds. Vapors severely irritate and may burn the eyes, skin, and respiratory system. The liquid burns all human tissue and causes severe damage.
Bromine trifluoride (BrF3) has a liquid range similar to water (Mp = 8.8 °C and Bp = 127 °C), and like water it auto ionizes, (10.5.1).
$\rm BrF_3 \rightarrow BrF_2^+ + BrF_4^-$
The products, like those of water’s self-ionization, are an acid (BrF2+) and a base (BrF4-). However, unlike water, BrF3 reacts with fluoride acids and bases not proton acids and bases. Thus, in BrF3 a base is a salt that provides F-, i.e., potassium fluoride (KF) is a base in BrF3 solution in the same manner as potassium hydroxide (KOH) is a base in water. The product from the reaction of a fluoride donor salt with BrF3 is the formation of the conjugate base, BrF4-, (10.5.2).
$\rm AgF + BrF_3 \rightarrow Ag^+ + BrF_4^-$
Other examples of this type of reaction include:
$\rm KF + BrF_3 \rightarrow K^+ + BrF_4^-$
$\rm NOF + BrF_3 \rightarrow NO^+ + BrF_4^-$
By analogy, an acid in BrF3 solution is a compound that acts as a fluoride (F-) acceptor, i.e., a Lewis acid, (10.5.5).
$\rm SbF_5 + BrF_3 \rightarrow SbF_6^- + BrF_2^+$
Exercise $1$
What are the products from the reaction of HF with BrF3?
Answer
$\rm HF + BrF_3 \rightarrow HF_2^- + BrF_2^+$
Bromine trifluoride as a fluorinating agent
Bromine trifluoride is a strong fluorinating agent that is able to convert a metal (e.g., vanadium) to its associated fluoride compound, (i.e., VF5). A wide range of salts and oxides may be converted to fluorides with the metal in a high oxidation state. However, it should be noted that BeO, MgO, and Al2O3 form oxo fluorides rather than the fluoride.
The reaction of silver with BrF3 yields the monofluoride, while the same reaction with gold yields the trifluoride, Eq. If the reactions are combined in BrF3 solution a mixed metal fluoride salt is formed.
A similar reaction occurs with NOCl and V2O5.
Exercise $2$
What are the products from the reaction of BrF3 with (a) Sb2O5, (b) KCl, and (c) a mixture of Sb2O5 and KCl?
Answer
(a) SbF5, (b) KF, and (c) K[SbF6].
Bibliography
• J. H. Simons, Inorg. Synth., 1950, 3, 184. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/10%3A_The_Halogens/10.04%3A_Oxyacids_of_Chlorine.txt |
Thumbnail: Vial of glowing ultrapure neon. (CC SA; Jurii via http://images-of-elements.com/neon.php).
11: Group 18 - The Noble Gases
The elements
The Group 18 elements have a particular name Noble gases. Noble gas is translated from the German noun Edelgas, first used in 1898 by Hugo Erdmann (1862 - 1910) to indicate their extremely low level of reactivity. The noble gases were often also called the inert gases, however, since noble gas compounds are now known this name is no longer used. Table \(1\) lists the derivation of the names of the Noble gases.
Table \(1\): Derivation of the names of each of the Group 18(VIII) elements.
Element Symbol Name
Helium He Greek helios meaning the Sun
Neon Ne From the Greek meaning new one
Argon Ar From the Greek meaning inactive
Krypton Kr From the Greek kryptos meaning the hidden one
Xenon Xe From the Greek xenos], meaning foreigner, stranger, or guest
Radon Rn From its radioactive nature
Discovery
Helium
The first evidence of helium was the observation by astronomer Pierre Janssen (Figure \(1\)) on August 18, 1868 as a bright yellow line with a wavelength of 587.49 nm in the spectrum of the chromosphere of the Sun. On October 20 of the same year, English astronomer Norman Lockyer (Figure \(2\)) observed a yellow line in the solar spectrum, which he named the D3 Fraunhofer line because it was near the known D1 and D2 lines of sodium. He concluded that it was caused by an element in the Sun unknown on Earth. Lockyer and Edward Frankland (Figure \(3\)) named the element with the Greek word for the Sun, helios.
On March 26, 1895 British chemist Sir William Ramsay (Figure \(4\)) isolated helium on Earth by treating the mineral cleveite (a radioactive mineral containing uranium and found in Norway) with mineral acids.
Neon
Neon was discovered in 1898 by Sir William Ramsay (Figure \(4\)) and Morris Travers (Figure \(5\)). When Ramsay chilled a sample of air until it became a liquid, then warmed the liquid and captured the gases as they boiled off. After nitrogen, oxygen, and argon, the three gases that boiled off were krypton, xenon, and neon.
Argon
In 1785 Henry Cavendish (Figure \(6\)) suspected that argon was present in air but it was not isolated until 1894 by Lord Rayleigh (Figure \(7\)) and Sir William Ramsay (Figure \(4\)) in an experiment in which they removed all of the oxygen, carbon dioxide, water and nitrogen from a sample of clean air.
Krypton
Krypton was discovered in 1898 by Sir William Ramsay (Figure \(4\)) and Morris Travers (Figure \(5\)) in residue left from evaporating nearly all components of liquid air.
Note
In 1960, an international agreement defined the meter (m) in terms of wavelength of light emitted by the 86Kr isotope (wavelength of 605.78 nm). This agreement replaced the standard meter located in Paris, which was a metal bar made of a Pt-Ir alloy, and was itself replaced by a definition based on the speed of light, a fundamental physical constant. In October 1983, the Bureau International des Poids et Mesures defined the meter as the distance that light travels in a vacuum during 1/299,792,458 s.
Xenon
Xenon was discovered by William Ramsay (Figure \(4\)) and Morris Travers (Figure \(5\)) on July 12, 1898, shortly after their discovery of krypton and neon.
Radon
Radon was the fifth radioactive element to be discovered after uranium, thorium, radium and polonium. Discovered in 1900 by Friedrich Dorn (Figure \(8\)) after he noticed that radium compounds emanate a radioactive gas that he named Radium Emanation (Ra Em). Prior to these experiments, in 1899, Pierre and Marie Curie (Figure \(9\)) observed that the gas emitted by radium remained radioactive for a month. Later that year, Ernest Rutherford (Figure \(10\)) noticed variations when trying to measure radiation from thorium oxide. In 1901, he demonstrated that the emanations are radioactive, but credited the Curies for the discovery of the element.
Abundance
The abundance of the Noble gases is given in Table \(2\).
Table \(2\): Abundance of Group 18 elements.
Element Terrestrial abundance (ppm)
He 8 x 10-3 (Earth’s crust), 4 x 106 (sea water), 5 (atmosphere)
Ne 70 x 10-3 (Earth’s crust), 0.2 (sea water), 18 (atmosphere)
Ar 1.2 (Earth’s crust), 0.45 (sea water), 0.93 x 104 (atmosphere)
Kr 10 x 10-6 (Earth’s crust), 80 x 10-6 (sea water), 1 (atmosphere)
Xe 2 x 10-6 (Earth’s crust), 100 x 10-6 (sea water), 90 x 10-3 (atmosphere)
Isotopes
The naturally abundant isotopes of the Group 18 elements are listed in Table \(3\). All of the isotopes of radon are radioactive.
Table \(3\): Abundance of the non-synthetic isotopes of the Group 18 elements.
Isotope Natural abundance (%)
Helium-3 0.000137
Helium-4 99.999863
Neon-20 90.48
Neon-21 0.27
Neon-22 9.25
Argon-36 0.337
Argon-86 0.063
Argon-40 99.600
Krypton-78 0.35
Krypton-80 2.25
Krypton-81 trace
Krypton-82 11.6
Krypton-83 11.5
Krypton-84 57
Krypton-86 17.3
Xenon-124 0.095
Xenon-126 0.089
Xenon-128 1.91
Xenon-129 26.4
Xenon-130 4.07
Xenon-131 21.2
Xenon-132 26.9
Xenon-134 10.4
Xenon-136 8.86
Radon-222 trace
Unlike most elements, helium's isotopic abundance varies greatly by origin, due to the different formation processes. The most common isotope, 4He, is produced on Earth by a decay of heavier radioactive elements. It was also formed in enormous quantities during the Big Bang.
Naturally occurring 40K with a half-life of 1.25 × 109 years, decays to stable 40Ar (11.2%) by electron capture and positron emission, and also to stable 40Ca (88.8%) via beta decay. These properties and ratios are used to determine the age of rocks.
With a half-life of 230,000 years 81Kr is used for dating 50,000 - 800,000 year old groundwater. 85Kr is an inert radioactive noble gas with a half-life of 10.76 years. It is produced in nuclear bomb testing and nuclear reactors. 85Kr is released during the reprocessing of fuel rods from nuclear reactors.
Industrial production of the elements
Helium is extracted by fractional distillation from natural gas, which contains up to 7% helium. Since helium has a lower boiling point than any other element, low temperature and high pressure are used to liquefy nearly all the other gases. The resulting helium gas is purified by successive exposures to lowering temperatures. A final purification step with activated charcoal results in 99.995% pure Grade-A helium.
Argon is produced industrially by the fractional distillation of liquid air, a process that separates liquid nitrogen, which boils at 77.3 K, from argon, which boils at 87.3 K and oxygen, which boils at 90.2 K. Xenon is obtained commercially as a byproduct of the separation of air into oxygen and nitrogen.
Physical properties
The physical properties of the Group 18 elements are given in Table \(4\).
Table \(4\): Selected physical properties of the Group 18 elements.
Element Mp (°C) Bp (°C)
He -272.20 -268.93
Ne -248.59 -246.08
Ar -189.35 -185.85
Kr -157.36 -153.22
Xe -111.7 -108.12
Rn -71.15 -61.85
All of the Noble gases show characteristic spectral lines (Figure \(11\) – Figure \(15\)).
Compounds of the Group 18 elements.
Only a few hundred noble gas compounds have been formed. Neutral compounds of helium and neon have not been formed, while xenon, krypton, and argon have shown only minor reactivity. The reactivity follows the order:
Xenon compounds are the most numerous of the noble gas compounds. Oxidation states of +2, +4, +6, and +8 with electronegative elements, e.g., XeF2, XeF4, XeF6, XeO4, and Na4XeO6. Compounds of xenon bound to boron, hydrogen, bromine, iodine, beryllium, sulphur, titanium, copper, and silver have also been observed but only at low temperatures in noble gas matrices, or in supersonic noble gas jets.
Although radon is more reactive than xenon it should form chemical bonds more easily than xenon, however, due to the high radioactivity and short half-life of radon isotopes, only a few fluorides and oxides of radon have been formed.
Krypton is less reactive than xenon, and oxidation states are generally limited to +2, KrF2. Compounds in which krypton forms a bond to nitrogen and oxygen are only stable below -60 °C and -90 °C, respectively. Krypton atoms chemically bound to other nonmetals (hydrogen, chlorine, carbon) as well as some late transition metals (copper, silver, gold), but only at low temperatures in noble gas matrices, or in supersonic noble gas jets. Similar conditions were used to obtain the first compounds of argon.
Noble gases also form non-covalent compounds, for example clathrates that consist of an atom trapped within cavities of crystal lattices of organic and inorganic compounds. Noble gases can form endohedral fullerene compounds, in which the noble gas atom is trapped inside a fullerene molecule (Figure \(16\)).
Bibliography
• L. Pauling, J. Am. Chem. Soc., 1933, 55, 1895.
• M. Saunders, H. A. Jiménez-Vázquez, R. J. Cross, and R. J. Poreda, Science, 1993, 259, 1428. | textbooks/chem/Inorganic_Chemistry/Chemistry_of_the_Main_Group_Elements_(Barron)/11%3A_Group_18_-_The_Noble_Gases/11.01%3A_The_Group_18_Elements-_The_Noble_Gases.txt |
This chapter introduces some history and context about the field of Inorganic Chemistry.
• 1.1: What is Inorganic Chemistry?
A generally-accepted definition of Inorganic Chemistry is the study of non-carbon molecules, or all the elements on the periodic table except carbon. But, this definition is not completely correct because the field of Inorganic Chemistry also includes organometallic compounds and the study of some carbon-based molecules that have properties that are familiar to metals (like conduction of electricity). This makes the field of inorganic chemistry very broad, and practically limitless.
• 1.2: Inorganic vs Organic Chemistry
The division between the fields of Inorganic and Organic chemistry has become blurred. For example, let's look at one of the major classes of catalysts used for organic synthesis reactions; organometalic catalysts. Organometallic catalysts like these, and all organometallic compounds, contain metals that are bonded to carbon or carbon-containing molecules. So, are they "inorganic" because they contain metals, or "organic" because they contain carbon?
• 1.3: History of Inorganic Chemistry
• 1.4: Perspectives
• 1.5: Practice problems
01: Introduction to Inorganic Chemistry
Where did the name "Inorganic Chemistry" come from? Well, the term "Organic Chemistry" literally means the chemistry of life. Organic chemistry is the study of carbon-based molecules because the first molecules that were isolated from living organisms contained carbon. On the other hand, minerals and other non-living things seemed to be made of other elements. For some time in our history, scientists believed that the chemical difference between living and non-living things was carbon. So, if "organic" molecules are the molecules of life, then is "inorganic chemistry" the "chemistry of death"? Almost? "Inorganic" chemistry historically meant the chemistry of "non-living" things; and these were non-carbon based molecules and ions.
The names "organic" and "inorganic" come from science history, and still today a generally-accepted definition of Inorganic Chemistry is the study of non-carbon molecules, or all the elements on the periodic table except carbon (Figure \(1\). But, this definition is not completely correct because the field of Inorganic Chemistry also includes organometallic compounds and the study of some carbon-based molecules that have properties that are familiar to metals (like conduction of electricity). This makes the field of inorganic chemistry very broad, and practically limitless. A great way to understand the breadth of the field is to take a look at the abstracts in the latest article of Inorganic Chemistry. Or, check out the 20 most-read articles from this past year using the links below.
Practice
What are the Sub-Fields of Inorganic Chemistry?
To appreciate the breadth of Inorganic Chemistry, go to the most recent issue of Inorganic Chemistry and look at the titles and visual abstracts. Identify at least 4 sub-fields of Inorganic Chemistry.
Answer
There are a lot of correct answers here! The point here is that you notice that Inorganic Chemistry is a very broad field. It has something for almost everyone because many other fields overlap with Inorganic Chemistry. You might notice that some of the sub-fields you identified are also interdisciplinary fields between inorganic chemistry and another discipline. For a list of some of the subfields of Inorganic Chemistry, check this Wikipedia article.
1.02: Inorganic vs Organic Chemistry
The division between the fields of Inorganic and Organic chemistry has become blurred. For example, let's look at one of the major classes of catalysts used for organic synthesis reactions: organometalic catalysts (Figure $1$). Organometallic catalysts like these, and all organometallic compounds, contain metals that are bonded to carbon or carbon-containing molecules. So, are they "inorganic" because they contain metals, or "organic" because they contain carbon? These illustrate that clear divisions between organic and inorganic chemistry do not exist. Further, metal ions are common in biology and so the idea that metals are "inorganic" and thus classed as "non-living or non-biological" is incorrect. A canonical example is the organometallic catalyst adenosylcobalbumin, which is an important biological cofactor containing a cobalt (Co) ion (Figure $1$, right) and a cobalt-carbon bond.
Some of the subfields of Inorganic Chemistry focus on electrical conductivity of inorganic materials (i.e., conduction, superconduction, and semiconduction) and on the study of optical and electronic properties of inorganic nanomaterials. Electrical conductivity is a canonical property of metals, but carbon-based materials also demonstrate electrical conductivity. For example, carbon nanotubes conduct electricity through their extended conjugated $\pi$ systems. Fullerenes, of which the most famous is Buckminsterfullerene, or Buckeyball (C60), demonstrate interesting properties that are similar to nanoparticles, and when combined with metals and crystallized can demonstrate superconductivity.
Although carbon nanotubes and fullerenes are allotropes of carbon, their material properties are somewhat foreign to many organic chemists, who traditionally have focused on smaller organic molecules having very different properties. However, these properties are familiar to inorganic chemists. Thus, inorganic chemists have embraced these molecules as "inorganic" due to the fact that they behave more like inorganic materials than smaller organic molecules. This class of carbon-based molecules serves as another example of molecules that are not perfectly matched to the traditional definitions of "organic" and "inorganic" chemistry. Certainly, the future will hold more and more examples of molecules that do not fit into the traditional disciplines of chemistry. | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/01%3A_Introduction_to_Inorganic_Chemistry/1.01%3A_What_is_Inorganic_Chemistry.txt |
Metals serve an essential role in many aspects of human civilization and have defined ages of human history. The period of time from about 3300 BC to 1200 BC is often referred to as the Bronze Age. During this period our ancestors first started using metal and learned to mix various elements with copper to make a strong alloy, called bronze. This age yielded significant advancement in the crafting of sharper knives and stronger weapons out of metal instead of rock, wood and bone. Around 1200 BC the human race found an even harder metal and discovered a much stronger alloy called steel. This period is known as the Iron Age. More recently, periods of time known as Gold Rushes have caused huge changes in population distributions and wealth in some countries. Metal has obvious importance in our modern way of life. Today, iron and steel are used for making buildings, machines, automobiles, jewelry, cooking pots, tools, weapons, vehicles, electronics, surgical instruments and symbolic structures like the Eiffel Tower and the Statue of Liberty. Gold, silver, and copper still serve as currency for trade and exchange of goods and services.
The existence of chemistry as a field of study owes much to the fact that gold was a valuable commodity throughout our history. In both the ancient Egyptian society and during the Roman Empire, gold mines were the property of the state, not an individual or group. So there were few ways for most people to legally get any gold for themselves. The Alchemists were a varied group of scholars and charlatans who aimed to solve this problem by creating the Philosopher's Stone (which caused the transmutation of lead into gold). Three major streams of alchemy are known: Chinese, Indian, and European, with all three streams having some factors in common. Techniques developed in the European stream ultimately influenced the development of the science of chemistry.
Although alchemists were never successful in changing lead into gold, they made several contributions to modern-day chemistry. Strong acids and bases were discovered, including nitric acid ($ce{HNO3}$), sulfuric acid ($\ce{H2SO4}$), and hydrochloric acid ($\ce{HCl}$), as well as sodium hydroxide ($\ce{NaOH}$). Glassware for running chemical reactions was developed, as were methods for distillation, crystallization, and sublimation. Alchemy helped improve the study of metallurgy and the extraction of metals from ores. More systematic approaches to research were developed that allowed the discovery of atoms and laid the groundwork for development of the periodic table. For more about the History of Chemistry in general, try the LibreText page A Brief History of Chemistry.
Inorganic compounds have been known and used since antiquity; probably the oldest is the deep blue pigment called Prussian blue ($\ce{Fe4[Fe(CN)6]3}$). However, the chemical nature of these substances was unknown until the late nineteenth and early twentieth century when the modern field of coordination chemistry emerged. Much of what we know about inorganic chemistry is based largely on the work of and debates between Alfred Werner (1866–1919; Nobel Prize in Chemistry in 1913) and Sophus Mads Jørgensen (1837 –1914). After Werner succeeded in these debates, the field of inorganic chemistry declined in popularity until the mid-twentieth century, when the second world war stimulated renewed interest. During the post-war era, several important discoveries and theories were developed. For example, important theories of bonding in coordination compounds were developed.
Soon after World War II, Crystal Field Theory (CFT) and Ligand Field Theory (LFT) were developed. These are two critical and complimentary theories that provide explanations of spectroscopic, chemical, and structural properties of inorganic coordination compounds, CFT being more simple, and LFT more accurate. In the 1950's, organometallic catalysts were discovered that catalyzed important organic reactions and the Haber-Bosch Process was discovered. The Haber-Bosch Process is catalyzed by an inorganic oxide catalyst and is one of the world's most important industrial reactions. It provides for the synthesis of ammonia directly from elemental nitrogen, N2, and hydrogen, H2.
$\ce{N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3 (g)} \label{eq1}$
Since its development in the early twentieth century, it has led to the production of an enormous quantity of fertilizer, vastly increasing global food production. As a result, it is estimated that a significant fraction of the nitrogen content in the typical human body is ultimately derived from this process. Yet while the reaction must be run at high temperatures and pressures in the industrial setting, the nitrogenase enzyme in the roots of plants can carry out this reaction at the mild conditions within the soil. Intense investigations were then aimed at improving inorganic catalysts through understanding the metal cofactors in enzymes. The link between the Haber-Bosche industrial process and the nitrogenase enzyme was an early bridge between the fields of organometallic chemistry and biochemistry.
https://chem.libretexts.org/Bookshel...tion_Compounds
Contributions
Part of this page is adapted from K.Haas Ph.D. dissertation
Part of this page is adapted from the follwing LibreTexts pages
1.04: Perspectives
The field of inorganic chemistry stands on the shoulders of giants (or in this case, many talented and dedicated scientists). If you are lucky enough to have journal access, peruse the titles and pages of the first issues of Inorganic Chemistry. You will find some lively debate about the puzzles of determining inorganic structure and reactivity. The scientists that made early discoveries in this field practiced elegant science.
The inorganic chemists who built our field are diverse. They are the ones who discovered new elements, developed new materials, characterized the fundamental principles of reactions, and developed the theories that we use to understand and predict metal structure and reactivity. And, there are people who are still working on these things today.
To gain more perspective on Inorganic Chemistry, compare the topics in the very first issue (Feb 1, 1962) to the most recent issue or the current most popular articles using the links below.
1.05: Practice problems
These questions are designed to check your understanding of the reading in this chapter.
1. What is the historic definition of inorganic chemistry? What are some problems that arise from this definition?
2. Name three sub-fields of inorganic chemistry.
3. What is the class of inorganic compounds that have bonds between metals and carbon? | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/01%3A_Introduction_to_Inorganic_Chemistry/1.03%3A_History_of_Inorganic_Chemistry.txt |
Notes about this page
The page below is a brief overview on the history of atomic theory. It contains lots of video media.
Alternatives pages on the history of atomic theory are:
Skills to Develop
By the end of this section, you will be able to:
• State the postulates of Dalton’s atomic theory
• Use postulates of Dalton’s atomic theory to explain the laws of definite and multiple proportions
• Outline milestones in the development of modern atomic theory
• Summarize and interpret the results of the experiments of Thomson, Millikan, and Rutherford
A Video Introduction to Atomic Theory through the Nineteenth Century From Crash Course Chemistry
Video $1$: Lavoisier's discovery of The Law of Conservation of Matter led to the Laws of Definite and Multiple Proportions and eventually Dalton's Atomic Theory.
Atomic Theory through the Nineteenth Century
The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers, the scientists of their day. In the fifth century BC, Leucippus and Democritus argued that all matter was composed of small, finite particles that they called atomos, a term derived from the Greek word for “indivisible.” They thought of atoms as moving particles that differed in shape and size, and which could join together. Later, Aristotle and others came to the conclusion that matter consisted of various combinations of the four “elements”—fire, earth, air, and water—and could be infinitely divided. Interestingly, these philosophers thought about atoms and “elements” as philosophical concepts, but apparently never considered performing experiments to test their ideas.
The Aristotelian view of the composition of matter held sway for over two thousand years, until English schoolteacher John Dalton helped to revolutionize chemistry with his hypothesis that the behavior of matter could be explained using an atomic theory. First published in 1807, many of Dalton’s hypotheses about the microscopic features of matter are still valid in modern atomic theory. Here are the postulates of Dalton’s atomic theory.
1. Matter is composed of exceedingly small particles called atoms. An atom is the smallest unit of an element that can participate in a chemical change.
2. An element consists of only one type of atom, which has a mass that is characteristic of the element and is the same for all atoms of that element (Figure $1$). A macroscopic sample of an element contains an incredibly large number of atoms, all of which have identical chemical properties.
3. Atoms of one element differ in properties from atoms of all other elements.
4. A compound consists of atoms of two or more elements combined in a small, whole-number ratio. In a given compound, the numbers of atoms of each of its elements are always present in the same ratio (Figure $2$).
5. Atoms are neither created nor destroyed during a chemical change, but are instead rearranged to yield substances that are different from those present before the change (Figure $3$).
Figure $1$: A pre-1982 copper penny (left) contains approximately 3 $\times$ 1022 copper atoms (several dozen are represented as brown spheres at the right), each of which has the same chemical properties. (credit: modification of work by “slgckgc”/Flickr)
Figure $2$: Copper(II) oxide, a powdery, black compound, results from the combination of two types of atoms—copper (brown spheres) and oxygen (red spheres)—in a 1:1 ratio. (credit: modification of work by “Chemicalinterest”/Wikimedia Commons)
Figure $3$: When the elements copper (a shiny, red-brown solid, shown here as brown spheres) and oxygen (a clear and colorless gas, shown here as red spheres) react, their atoms rearrange to form a compound containing copper and oxygen (a powdery, black solid). (credit copper: modification of work by http://images-of-elements.com/copper.php).
Dalton’s atomic theory provides a microscopic explanation of the many macroscopic properties of matter that you’ve learned about. For example, if an element such as copper consists of only one kind of atom, then it cannot be broken down into simpler substances, that is, into substances composed of fewer types of atoms. And if atoms are neither created nor destroyed during a chemical change, then the total mass of matter present when matter changes from one type to another will remain constant (the law of conservation of matter (or mass)).
Want to learn more about the Law of Conservation of Mass?
Video $2$: "We are made of star stuff" - Carl Sagan.
Example $1$: Testing Dalton’s Atomic Theory
In the following drawing, the green spheres represent atoms of a certain element. The purple spheres represent atoms of another element. If the spheres touch, they are part of a single unit of a compound. Does the following chemical change represented by these symbols violate any of the ideas of Dalton’s atomic theory? If so, which one?
Solution
The starting materials consist of two green spheres and two purple spheres. The products consist of only one green sphere and one purple sphere. This violates Dalton’s postulate that atoms are neither created nor destroyed during a chemical change, but are merely redistributed. (In this case, atoms appear to have been destroyed.)
Exercise $1$
In the following drawing, the green spheres represent atoms of a certain element. The purple spheres represent atoms of another element. If the spheres touch, they are part of a single unit of a compound. Does the following chemical change represented by these symbols violate any of the ideas of Dalton’s atomic theory? If so, which one
Answer
The starting materials consist of four green spheres and two purple spheres. The products consist of four green spheres and two purple spheres. This does not violate any of Dalton’s postulates: Atoms are neither created nor destroyed, but are redistributed in small, whole-number ratios.
Dalton knew of the experiments of French chemist Joseph Proust, who demonstrated that all samples of a pure compound contain the same elements in the same proportion by mass. This statement is known as the law of definite proportions or the law of constant composition. The suggestion that the numbers of atoms of the elements in a given compound always exist in the same ratio is consistent with these observations. For example, when different samples of isooctane (a component of gasoline and one of the standards used in the octane rating system) are analyzed, they are found to have a carbon-to-hydrogen mass ratio of 5.33:1, as shown in Table $1$.
Table $1$: Constant Composition of Isooctane
Sample Carbon Hydrogen Mass Ratio
A 14.82 g 2.78 g $\mathrm{\dfrac{14.82\: g\: carbon}{2.78\: g\: hydrogen}=\dfrac{5.33\: g\: carbon}{1.00\: g\: hydrogen}}$
B 22.33 g 4.19 g $\mathrm{\dfrac{22.33\: g\: carbon}{4.19\: g\: hydrogen}=\dfrac{5.33\: g\: carbon}{1.00\: g\: hydrogen}}$
C 19.40 g 3.64 g $\mathrm{\dfrac{19.40\: g\: carbon}{3.63\: g\: hydrogen}=\dfrac{5.33\: g\: carbon}{1.00\: g\: hydrogen}}$
It is worth noting that although all samples of a particular compound have the same mass ratio, the converse is not true in general. That is, samples that have the same mass ratio are not necessarily the same substance. For example, there are many compounds other than isooctane that also have a carbon-to-hydrogen mass ratio of 5.33:1.00.
Dalton also used data from Proust, as well as results from his own experiments, to formulate another interesting law. The law of multiple proportions states that when two elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small, whole numbers. For example, copper and chlorine can form a green, crystalline solid with a mass ratio of 0.558 g chlorine to 1 g copper, as well as a brown crystalline solid with a mass ratio of 1.116 g chlorine to 1 g copper. These ratios by themselves may not seem particularly interesting or informative; however, if we take a ratio of these ratios, we obtain a useful and possibly surprising result: a small, whole-number ratio.
$\mathrm{\dfrac{\dfrac{1.116\: g\: Cl}{1\: g\: Cu}}{\dfrac{0.558\: g\: Cl}{1\: g\: Cu}}=\dfrac{2}{1}}$
This 2-to-1 ratio means that the brown compound has twice the amount of chlorine per amount of copper as the green compound.
This can be explained by atomic theory if the copper-to-chlorine ratio in the brown compound is 1 copper atom to 2 chlorine atoms, and the ratio in the green compound is 1 copper atom to 1 chlorine atom. The ratio of chlorine atoms (and thus the ratio of their masses) is therefore 2 to 1 (Figure $4$).
Figure $4$: Compared to the copper chlorine compound in (a), where copper is represented by brown spheres and chlorine by green spheres, the copper chlorine compound in (b) has twice as many chlorine atoms per copper atom. (credit a: modification of work by “Benjah-bmm27”/Wikimedia Commons; credit b: modification of work by “Walkerma”/Wikimedia Commons)
Example $2$: Laws of Definite and Multiple Proportions
A sample of compound A (a clear, colorless gas) is analyzed and found to contain 4.27 g carbon and 5.69 g oxygen. A sample of compound B (also a clear, colorless gas) is analyzed and found to contain 5.19 g carbon and 13.84 g oxygen. Are these data an example of the law of definite proportions, the law of multiple proportions, or neither? What do these data tell you about substances A and B?
Solution
In compound A, the mass ratio of carbon to oxygen is:
$\mathrm{\dfrac{1.33\: g\: O}{1\: g\: C}} \nonumber$
In compound B, the mass ratio of carbon to oxygen is:
$\mathrm{\dfrac{2.67\: g\: O}{1\: g\: C}} \nonumber$
The ratio of these ratios is:
$\mathrm{\dfrac{\dfrac{1.33\: g\: O}{1\: g\: C}}{\dfrac{2.67\: g\: O}{1\: g\: C}}=\dfrac{1}{2}} \nonumber$
This supports the law of multiple proportions. This means that A and B are different compounds, with A having one-half as much carbon per amount of oxygen (or twice as much oxygen per amount of carbon) as B. A possible pair of compounds that would fit this relationship would be A = CO2 and B = CO.
Exercise $2$
A sample of compound X (a clear, colorless, combustible liquid with a noticeable odor) is analyzed and found to contain 14.13 g carbon and 2.96 g hydrogen. A sample of compound Y (a clear, colorless, combustible liquid with a noticeable odor that is slightly different from X’s odor) is analyzed and found to contain 19.91 g carbon and 3.34 g hydrogen. Are these data an example of the law of definite proportions, the law of multiple proportions, or neither? What do these data tell you about substances X and Y?
Answer
In compound X, the mass ratio of carbon to hydrogen is $\mathrm{\dfrac{14.13\: g\: C}{2.96\: g\: H}}$.
In compound Y, the mass ratio of carbon to oxygen is $\mathrm{\dfrac{19.91\: g\: C}{3.34\: g\: H}}$.
The ratio of these ratios is
$\mathrm{\dfrac{\dfrac{14.13\: g\: C}{2.96\: g\: H}}{\dfrac{19.91\: g\: C}{3.34\: g\: H}}=\dfrac{4.77\: g\: C/g\: H}{5.96\: g\: C/g\: H}=0.800=\dfrac{4}{5}}. \nonumber$
This small, whole-number ratio supports the law of multiple proportions. This means that X and Y are different compounds.
In the two centuries since Dalton developed his ideas, scientists have made significant progress in furthering our understanding of atomic theory. Much of this came from the results of several seminal experiments that revealed the details of the internal structure of atoms. Here, we will discuss some of those key developments, with an emphasis on application of the scientific method, as well as understanding how the experimental evidence was analyzed. While the historical persons and dates behind these experiments can be quite interesting, it is most important to understand the concepts resulting from their work.
Atomic Theory after the Nineteenth Century
If matter were composed of atoms, what were atoms composed of? Were they the smallest particles, or was there something smaller? In the late 1800s, a number of scientists interested in questions like these investigated the electrical discharges that could be produced in low-pressure gases, with the most significant discovery made by English physicist J. J. Thomson using a cathode ray tube. This apparatus consisted of a sealed glass tube from which almost all the air had been removed; the tube contained two metal electrodes. When high voltage was applied across the electrodes, a visible beam called a cathode ray appeared between them. This beam was deflected toward the positive charge and away from the negative charge, and was produced in the same way with identical properties when different metals were used for the electrodes. In similar experiments, the ray was simultaneously deflected by an applied magnetic field, and measurements of the extent of deflection and the magnetic field strength allowed Thomson to calculate the charge-to-mass ratio of the cathode ray particles. The results of these measurements indicated that these particles were much lighter than atoms (Figure $1$).
Figure $5$: (a) J. J. Thomson produced a visible beam in a cathode ray tube. (b) This is an early cathode ray tube, invented in 1897 by Ferdinand Braun. (c) In the cathode ray, the beam (shown in yellow) comes from the cathode and is accelerated past the anode toward a fluorescent scale at the end of the tube. Simultaneous deflections by applied electric and magnetic fields permitted Thomson to calculate the mass-to-charge ratio of the particles composing the cathode ray. (credit a: modification of work by Nobel Foundation; credit b: modification of work by Eugen Nesper; credit c: modification of work by “Kurzon”/Wikimedia Commons).
Based on his observations, here is what Thomson proposed and why: The particles are attracted by positive (+) charges and repelled by negative (−) charges, so they must be negatively charged (like charges repel and unlike charges attract); they are less massive than atoms and indistinguishable, regardless of the source material, so they must be fundamental, subatomic constituents of all atoms. Although controversial at the time, Thomson’s idea was gradually accepted, and his cathode ray particle is what we now call an electron, a negatively charged, subatomic particle with a mass more than one thousand-times less that of an atom. The term “electron” was coined in 1891 by Irish physicist George Stoney, from “electric ion.”
In 1909, more information about the electron was uncovered by American physicist Robert A. Millikan via his “oil drop” experiments. Millikan created microscopic oil droplets, which could be electrically charged by friction as they formed or by using X-rays. These droplets initially fell due to gravity, but their downward progress could be slowed or even reversed by an electric field lower in the apparatus. By adjusting the electric field strength and making careful measurements and appropriate calculations, Millikan was able to determine the charge on individual drops (Figure $2$).
Figure $6$: Millikan’s experiment measured the charge of individual oil drops. The tabulated data are examples of a few possible values.
Looking at the charge data that Millikan gathered, you may have recognized that the charge of an oil droplet is always a multiple of a specific charge, 1.6 $\times$ 10−19 C. Millikan concluded that this value must therefore be a fundamental charge—the charge of a single electron—with his measured charges due to an excess of one electron (1 times 1.6 $\times$ 10−19 C), two electrons (2 times 1.6 $\times$ 10−19 C), three electrons (3 times 1.6 $\times$ 10−19 C), and so on, on a given oil droplet. Since the charge of an electron was now known due to Millikan’s research, and the charge-to-mass ratio was already known due to Thomson’s research (1.759 $\times$ 1011 C/kg), it only required a simple calculation to determine the mass of the electron as well.
$\mathrm{Mass\: of\: electron=1.602\times 10^{-19}\:\cancel{C}\times \dfrac{1\: kg}{1.759\times 10^{11}\:\cancel{C}}=9.107\times 10^{-31}\:kg} \tag{2.3.1}$
Scientists had now established that the atom was not indivisible as Dalton had believed, and due to the work of Thomson, Millikan, and others, the charge and mass of the negative, subatomic particles—the electrons—were known. However, the positively charged part of an atom was not yet well understood. In 1904, Thomson proposed the “plum pudding” model of atoms, which described a positively charged mass with an equal amount of negative charge in the form of electrons embedded in it, since all atoms are electrically neutral. A competing model had been proposed in 1903 by Hantaro Nagaoka, who postulated a Saturn-like atom, consisting of a positively charged sphere surrounded by a halo of electrons (Figure $3$).
Figure $7$: (a) Thomson suggested that atoms resembled plum pudding, an English dessert consisting of moist cake with embedded raisins (“plums”). (b) Nagaoka proposed that atoms resembled the planet Saturn, with a ring of electrons surrounding a positive “planet.” (credit a: modification of work by “Man vyi”/Wikimedia Commons; credit b: modification of work by “NASA”/Wikimedia Commons).
The next major development in understanding the atom came from Ernest Rutherford, a physicist from New Zealand who largely spent his scientific career in Canada and England. He performed a series of experiments using a beam of high-speed, positively charged alpha particles (α particles) that were produced by the radioactive decay of radium; α particles consist of two protons and two neutrons (you will learn more about radioactive decay in the chapter on nuclear chemistry). Rutherford and his colleagues Hans Geiger (later famous for the Geiger counter) and Ernest Marsden aimed a beam of α particles, the source of which was embedded in a lead block to absorb most of the radiation, at a very thin piece of gold foil and examined the resultant scattering of the α particles using a luminescent screen that glowed briefly where hit by an α particle.
What did they discover? Most particles passed right through the foil without being deflected at all. However, some were diverted slightly, and a very small number were deflected almost straight back toward the source (Figure $4$). Rutherford described finding these results: “It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you”1 (p. 68).
Figure $8$: Geiger and Rutherford fired α particles at a piece of gold foil and detected where those particles went, as shown in this schematic diagram of their experiment. Most of the particles passed straight through the foil, but a few were deflected slightly and a very small number were significantly deflected.
Here is what Rutherford deduced: Because most of the fast-moving α particles passed through the gold atoms undeflected, they must have traveled through essentially empty space inside the atom. Alpha particles are positively charged, so deflections arose when they encountered another positive charge (like charges repel each other). Since like charges repel one another, the few positively charged α particles that changed paths abruptly must have hit, or closely approached, another body that also had a highly concentrated, positive charge. Since the deflections occurred a small fraction of the time, this charge only occupied a small amount of the space in the gold foil. Analyzing a series of such experiments in detail, Rutherford drew two conclusions:
1. The volume occupied by an atom must consist of a large amount of empty space.
2. A small, relatively heavy, positively charged body, the nucleus, must be at the center of each atom.
This analysis led Rutherford to propose a model in which an atom consists of a very small, positively charged nucleus, in which most of the mass of the atom is concentrated, surrounded by the negatively charged electrons, so that the atom is electrically neutral (Figure $5$).
Figure $9$: The α particles are deflected only when they collide with or pass close to the much heavier, positively charged gold nucleus. Because the nucleus is very small compared to the size of an atom, very few α particles are deflected. Most pass through the relatively large region occupied by electrons, which are too light to deflect the rapidly moving particles.
After many more experiments, Rutherford also discovered that the nuclei of other elements contain the hydrogen nucleus as a “building block,” and he named this more fundamental particle the proton, the positively charged, subatomic particle found in the nucleus. With one addition, which you will learn next, this nuclear model of the atom, proposed over a century ago, is still used today.
Another important finding was the discovery of isotopes. During the early 1900s, scientists identified several substances that appeared to be new elements, isolating them from radioactive ores. For example, a “new element” produced by the radioactive decay of thorium was initially given the name mesothorium. However, a more detailed analysis showed that mesothorium was chemically identical to radium (another decay product), despite having a different atomic mass. This result, along with similar findings for other elements, led the English chemist Frederick Soddy to realize that an element could have types of atoms with different masses that were chemically indistinguishable. These different types are called isotopes—atoms of the same element that differ in mass. Soddy was awarded the Nobel Prize in Chemistry in 1921 for this discovery.
One puzzle remained: The nucleus was known to contain almost all of the mass of an atom, with the number of protons only providing half, or less, of that mass. Different proposals were made to explain what constituted the remaining mass, including the existence of neutral particles in the nucleus. As you might expect, detecting uncharged particles is very challenging, and it was not until 1932 that James Chadwick found evidence of neutrons, uncharged, subatomic particles with a mass approximately the same as that of protons. The existence of the neutron also explained isotopes: They differ in mass because they have different numbers of neutrons, but they are chemically identical because they have the same number of protons. This will be explained in more detail later in this unit.
Video $2$: An Introduction to Subatomic Particles
Summary
Video $3$: A summary of discoveries in atomic theory.
Video $4$: A different summary of discoveries in atomic theory.
The ancient Greeks proposed that matter consists of extremely small particles called atoms. Dalton postulated that each element has a characteristic type of atom that differs in properties from atoms of all other elements, and that atoms of different elements can combine in fixed, small, whole-number ratios to form compounds. Samples of a particular compound all have the same elemental proportions by mass. When two elements form different compounds, a given mass of one element will combine with masses of the other element in a small, whole-number ratio. During any chemical change, atoms are neither created nor destroyed.
Although no one has actually seen the inside of an atom, experiments have demonstrated much about atomic structure. Thomson’s cathode ray tube showed that atoms contain small, negatively charged particles called electrons. Millikan discovered that there is a fundamental electric charge—the charge of an electron. Rutherford’s gold foil experiment showed that atoms have a small, dense, positively charged nucleus; the positively charged particles within the nucleus are called protons. Chadwick discovered that the nucleus also contains neutral particles called neutrons. Soddy demonstrated that atoms of the same element can differ in mass; these are called isotopes.
Footnotes
1. Ernest Rutherford, “The Development of the Theory of Atomic Structure,” ed. J. A. Ratcliffe, in Background to Modern Science, eds. Joseph Needham and Walter Pagel, (Cambridge, UK: Cambridge University Press, 1938), 61–74. Accessed September 22, 2014, https://ia600508.us.archive.org/3/it...e032734mbp.pdf.
Glossary
Dalton’s atomic theory
set of postulates that established the fundamental properties of atoms
law of constant composition
(also, law of definite proportions) all samples of a pure compound contain the same elements in the same proportions by mass
law of multiple proportions
when two elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small whole numbers
law of definite proportions
(also, law of constant composition) all samples of a pure compound contain the same elements in the same proportions by mass
alpha particle (α particle)
positively charged particle consisting of two protons and two neutrons
electron
negatively charged, subatomic particle of relatively low mass located outside the nucleus
isotopes
atoms that contain the same number of protons but different numbers of neutrons
neutron
uncharged, subatomic particle located in the nucleus
proton
positively charged, subatomic particle located in the nucleus
nucleus
massive, positively charged center of an atom made up of protons and neutrons
Contributors
• Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors. Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/[email protected]).
• Adelaide Clark, Oregon Institute of Technology
• Crash Course Chemistry: Crash Course is a division of Complexly and videos are free to stream for educational purposes.
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2.01: Historical Development of Atomic Theory
Learning Objectives
• To become familiar with the history of the periodic table.
The modern periodic table has evolved through a long history of attempts by chemists to arrange the elements according to their reactivity and other properties as an aid in predicting chemical behavior. Now that we have arranged the table according to electronic structure, it makes sense to go back and look at earlier efforts in the light of what we know about electronic structure.
One of the first to suggest such an arrangement was the German chemist Johannes Dobereiner (1780–1849), who noticed that many of the known elements could be grouped in triads, sets of three elements that have similar properties—for example, chlorine, bromine, and iodine; or copper, silver, and gold. Dobereiner proposed that all elements could be grouped in such triads, but subsequent attempts to expand his concept were unsuccessful. We now know that portions of the periodic table—the d block in particular—contain triads of elements with substantial similarities. The middle three members of most of the other columns, such as sulfur, selenium, and tellurium in group 16 or aluminum, gallium, and indium in group 13, also have remarkably similar chemistry.
By the mid-19th century, the atomic masses of many of the elements had been determined. The English chemist John Newlands (1838–1898), hypothesizing that the chemistry of the elements might be related to their masses, arranged the known elements in order of increasing atomic mass and discovered that every seventh element had similar properties (Figure 3.4.1 ). Newlands therefore suggested that the elements could be classified into octaves: a group of seven elements (not counting the noble gases, which were unknown at the time) that correspond to the horizontal rows in the main group elements. Unfortunately, Newlands’s “law of octaves” did not seem to work for elements heavier than calcium, and his idea was publicly ridiculed. At one scientific meeting, Newlands was asked why he didn’t arrange the elements in alphabetical order instead of by atomic mass, since that would make just as much sense! Actually, Newlands was on the right track—with only a few exceptions, atomic mass does increase with atomic number, and similar properties occur every time a set of ns2np6 subshells is filled. Despite the fact that Newlands’s table had no logical place for the d-block elements, he was honored for his idea by the Royal Society of London in 1887.
John Newlands (1838–1898)
Newlands noticed that elemental properties repeated every seventh (or multiple of seven) element, as musical notes repeat every eighth note.
The periodic table achieved its modern form through the work of the German chemist Julius Lothar Meyer (1830–1895) and the Russian chemist Dimitri Mendeleev (1834–1907), both of whom focused on the relationships between atomic mass and various physical and chemical properties. In 1869, they independently proposed essentially identical arrangements of the elements. Meyer aligned the elements in his table according to periodic variations in simple atomic properties, such as “atomic volume” (Figure 3.4.2 ), which he obtained by dividing the atomic mass (molar mass) in grams per mole by the density of the element in grams per cubic centimeter. This property is equivalent to what is today defined as molar volume, the molar mass of an element divided by its density (measured in cubic centimeters per mole):
$\dfrac{molar\; mass\left ( \cancel{g}/mol \right )}{density\left ( \cancel{g}/cm^{3} \right )}=molar\; volume\left ( cm^{3}/mol \right ) \tag{3.4.1}$
As shown in Figure 3.4.2 , the alkali metals have the highest molar volumes of the solid elements. In Meyer’s plot of atomic volume versus atomic mass, the nonmetals occur on the rising portion of the graph, and metals occur at the peaks, in the valleys, and on the down slopes.
Dimitri Mendeleev (1834–1907)
When his family’s glass factory was destroyed by fire, Mendeleev moved to St. Petersburg, Russia, to study science. He became ill and was not expected to recover, but he finished his PhD with the help of his professors and fellow students. In addition to the periodic table, another of Mendeleev’s contributions to science was an outstanding textbook, The Principles of Chemistry, which was used for many years.
Mendeleev’s Periodic Table
Mendeleev, who first published his periodic table in 1869 (Figure 3.4.3 ), is usually credited with the origin of the modern periodic table. The key difference between his arrangement of the elements and that of Meyer and others is that Mendeleev did not assume that all the elements had been discovered (actually, only about two-thirds of the naturally occurring elements were known at the time). Instead, he deliberately left blanks in his table at atomic masses 44, 68, 72, and 100, in the expectation that elements with those atomic masses would be discovered. Those blanks correspond to the elements we now know as scandium, gallium, germanium, and technetium.
The groups in Mendeleev's table are determined by how many oxygen or hydrogen atoms are needed to form compounds with each element. For example, in Group I, two atoms of hydrogen (H), lithium (Li), sodium (Na), and potassium (K) form compounds with one atom of oxygen. In Group VII, one atom of fluorine (F), chlorine (Cl), and bromine (Br), reacts with one atom of hydrogen. Notice how this approach has trouble with the transition metals. Until roughly 1960, a rectangular table developed from Mendeleev's table and based on reactivity was standard at the front of chemistry lecture halls.
The most convincing evidence in support of Mendeleev’s arrangement of the elements was the discovery of two previously unknown elements whose properties closely corresponded with his predictions (Table 3.4.1 ). Two of the blanks Mendeleev had left in his original table were below aluminum and silicon, awaiting the discovery of two as-yet-unknown elements, eka-aluminum and eka-silicon (from the Sanskrit eka, meaning “one,” as in “one beyond aluminum”). The observed properties of gallium and germanium matched those of eka-aluminum and eka-silicon so well that once they were discovered, Mendeleev’s periodic table rapidly gained acceptance.
Table 3.4.1 Comparison of the properties predicted by Mendeleev in 1869 for eka-aluminum and eka-silicon with the properties of gallium (discovered in 1875) and germanium (discovered in 1886)
Property eka-Aluminum (predicted) Gallium (observed) eka-Silicon (predicted) Germanium (observed)
atomic mass 68 69.723 72 72.64
element metal metal dirty-gray metal gray-white metal
low mp* mp = 29.8°C high mp mp = 938°C
d = 5.9 g/cm3 d = 5.91 g/cm3 d = 5.5 g/cm3 d = 5.323 g/cm3
oxide E2O3 Ga2O3 EO2 GeO2
d = 5.5 g/cm3 d = 6.0 g/cm3 d = 4.7 g/cm3 d = 4.25 g/cm3
chloride ECl3 GaCl3 ECl4 GeCl4
volatile
mp = 78°C
bp* = 201°C
bp < 100°C bp = 87°C
*mp = melting point; bp = boiling point.
When the chemical properties of an element suggested that it might have been assigned the wrong place in earlier tables, Mendeleev carefully reexamined its atomic mass. He discovered, for example, that the atomic masses previously reported for beryllium, indium, and uranium were incorrect. The atomic mass of indium had originally been reported as 75.6, based on an assumed stoichiometry of InO for its oxide. If this atomic mass were correct, then indium would have to be placed in the middle of the nonmetals, between arsenic (atomic mass 75) and selenium (atomic mass 78). Because elemental indium is a silvery-white metal, however, Mendeleev postulated that the stoichiometry of its oxide was really In2O3 rather than InO. This would mean that indium’s atomic mass was actually 113, placing the element between two other metals, cadmium and tin.
One group of elements absent from Mendeleev’s table was the noble gases, all of which were discovered more than 20 years later, between 1894 and 1898, by Sir William Ramsay (1852–1916; Nobel Prize in Chemistry 1904). Initially, Ramsay did not know where to place these elements in the periodic table. Argon, the first to be discovered, had an atomic mass of 40. This was greater than chlorine’s and comparable to that of potassium; so Ramsay, using the same kind of reasoning as Mendeleev, decided to place the noble gases between the halogens and the alkali metals.
The Role of the Atomic Number in the Periodic Table
Despite its usefulness, Mendeleev’s periodic table was based entirely on empirical observation supported by very little understanding. It was not until 1913, when a young British physicist, H. G. J. Moseley (1887–1915), while analyzing the frequencies of x-rays emitted by the elements, discovered that the underlying foundation of the order of the elements was the atomic number, not the atomic mass. Moseley hypothesized that the placement of each element in his series corresponded to its atomic number Z, which is the number of positive charges (protons) in its nucleus. Argon, for example, although having an atomic mass greater than that of potassium (39.9 amu versus 39.1 amu, respectively), was placed before potassium in the periodic table. While analyzing the frequencies of the emitted x-rays, Moseley noticed that the atomic number of argon is 18, whereas that of potassium is 19, which indicated that they were indeed placed correctly. Moseley also noticed three gaps in his table of x-ray frequencies, so he predicted the existence of three unknown elements: technetium (Z = 43), discovered in 1937; promethium (Z = 61), discovered in 1945; and rhenium (Z = 75), discovered in 1925.
H. G. J. Moseley (1887–1915)
Moseley left his research work at the University of Oxford to join the British army as a telecommunications officer during World War I. He was killed during the Battle of Gallipoli in Turkey.
Example 3.4.1
Before its discovery in 1999, some theoreticians believed that an element with a Z of 114 existed in nature. Use Mendeleev’s reasoning to name element 114 as eka-______; then identify the known element whose chemistry you predict would be most similar to that of element 114.
Given: atomic number
Asked for: name using prefix eka-
Strategy:
A Using the periodic table locate the n = 7 row. Identify the location of the unknown element with Z = 114; then identify the known element that is directly above this location.
B Name the unknown element by using the prefix eka- before the name of the known element.
Solution:
A The n = 7 row can be filled in by assuming the existence of elements with atomic numbers greater than 112, which is underneath mercury (Hg). Counting three boxes to the right gives element 114, which lies directly below lead (Pb). B If Mendeleev were alive today, he would call element 114 eka-lead.
Exercise
Use Mendeleev’s reasoning to name element 112 as eka-______; then identify the known element whose chemistry you predict would be most similar to that of element 112.
Answer: eka-mercury
Summary
The periodic table arranges the elements according to their electron configurations, such that elements in the same column have the same valence electron configurations. Periodic variations in size and chemical properties are important factors in dictating the types of chemical reactions the elements undergo and the kinds of chemical compounds they form. The modern periodic table was based on empirical correlations of properties such as atomic mass; early models using limited data noted the existence of triads and octaves of elements with similar properties. The periodic table achieved its current form through the work of Dimitri Mendeleev and Julius Lothar Meyer, who both focused on the relationship between atomic mass and chemical properties. Meyer arranged the elements by their atomic volume, which today is equivalent to the molar volume, defined as molar mass divided by molar density. The correlation with the electronic structure of atoms was made when H. G. J. Moseley showed that the periodic arrangement of the elements was determined by atomic number, not atomic mass.
Key Takeaways
• The elements in the periodic table are arranged according to their properties, and the periodic table serves as an aid in predicting chemical behavior.
Conceptual Problems
1. Johannes Dobereiner is credited with developing the concept of chemical triads. Which of the group 15 elements would you expect to compose a triad? Would you expect B, Al, and Ga to act as a triad? Justify your answers.
2. Despite the fact that Dobereiner, Newlands, Meyer, and Mendeleev all contributed to the development of the modern periodic table, Mendeleev is credited with its origin. Why was Mendeleev’s periodic table accepted so rapidly?
3. How did Moseley’s contribution to the development of the periodic table explain the location of the noble gases?
4. The eka- naming scheme devised by Mendeleev was used to describe undiscovered elements.
1. Use this naming method to predict the atomic number of eka-mercury, eka-astatine, eka-thallium, and eka-hafnium.
2. Using the eka-prefix, identify the elements with these atomic numbers: 79, 40, 51, 117, and 121.
Numerical Problem
1. Based on the data given, complete the table.
Species Molar Mass (g/mol) Density (g/cm3) Molar Volume (cm3/mol)
A 40.078 25.85
B 39.09 0.856
C 32.065 16.35
D 1.823 16.98
E 26.98 9.992
F 22.98 0.968
Plot molar volume versus molar mass for these substances. According to Meyer, which would be considered metals and which would be considered nonmetals?
Answer
1. Species Molar Mass (g/mol) Density (g/cm3) Molar Volume (cm3/mol)
A 40.078 1.550 25.85
B 39.09 0.856 45.67
C 32.065 1.961 16.35
D 30.95 1.823 16.98
E 26.98 2.700 9.992
F 22.98 0.968 23.7
Meyer found that the alkali metals had the highest molar volumes, and that molar volumes decreased steadily with increasing atomic mass, then leveled off, and finally rose again. The elements located on the rising portion of a plot of molar volume versus molar mass were typically nonmetals. If we look at the plot of the data in the table, we can immediately identify those elements with the largest molar volumes (A, B, F) as metals located on the left side of the periodic table. The element with the smallest molar volume (E) is aluminum. The plot shows that the subsequent elements (C, D) have molar volumes that are larger than that of E, but smaller than those of A and B. Thus, C and D are most likely to be nonmetals (which is the case: C = sulfur, D = phosphorus).
Contributors and Attributions
• Anonymous
Modified by Joshua Halpern
Video from TED-Ed Lou Serocp | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/02%3A_Atomic_Structure/2.01%3A_Historical_Development_of_Atomic_Theory/2.1.01%3A_The_Periodic_Table.txt |
Hydrogen Absorption and Emission Spectra
When a high-voltage electrical discharge is passed through a sample of hydrogen gas (H2) at low pressure, the result is individual isolated hydrogen atoms that emit a red light. Unlike blackbody radiation, the color of the light emitted by the hydrogen atoms does not depend greatly on the temperature of the gas in the tube. When the emitted light is passed through a prism, only a few narrow lines of particular wavelengths, called a line spectrum, are observed rather than a continuous range of wavelengths (Figure $1$). The light emitted by hydrogen atoms is red because, of its four characteristic lines, the most intense line in its spectrum is in the red portion of the visible spectrum, at 656 nm.
The Balmer Series
In 1885, a Swiss mathematics teacher, Johann Balmer (1825–1898), showed that the frequencies of the lines observed in the visible region of the hydrogen line spectrum fit a simple equation that can be expressed as follows:
$u=constant\; \left ( \dfrac{1}{2^{2}}-\dfrac{1}{n^{^{2}}} \right ) \label{6.3.1}$
where n = 3, 4, 5, 6. As a result, these lines are known as the Balmer series. The Swedish physicist Johannes Rydberg (1854–1919) subsequently restated and expanded Balmer’s result in the Rydberg equation:
$\dfrac{1}{\lambda }=R_H\; \left ( \dfrac{1}{n_l^{2}}-\dfrac{1}{n_h^{2}} \right ) \label{6.3.2}$ ,
where $n_l$ and $n_h$ are positive integers, $n_h > n_l$, and $R_H$, the Rydberg constant, has a value of 1.09737 × 107 m−1. Like Balmer’s equation, Rydberg’s simple equation described the wavelengths of the visible lines in the emission spectrum of hydrogen (with $n_l = 2, n_h = 3, 4, 5,…$). More importantly, Rydberg’s equation also predicted the wavelengths of other series of lines that would be observed in the emission spectrum of hydrogen: one in the ultraviolet ($n_l = 1, n_h = 2, 3, 4,…$) and one in the infrared ($n_l = 3, n_h = 4, 5, 6$).
Other Series
The results given by Balmer and Rydberg for the spectrum in the visible region of the electromagnetic spectrum start with $n_h = 3$, and $n_l=2$.
Is there a different series with the following formula (e.g., $n_l=1$)?
$\dfrac{1}{\lambda} = R_{\textrm H} \left(\dfrac{1}{1^2} - \dfrac{1}{n^2} \right ) \label{1.5.2}$
The values for $n_h$ and wavenumber $u$ for this series would be:
Table $1$: The Lyman Series of Hydrogen Emission Lines ($n_l=1$)
$n_h$ 2 3 4 5
$\lambda$ (nm) 121 102 97 94
$\widetilde{ u}$ (cm-1) 82,2291 97,530 102,864 105,332
Do you know what region of the electromagnetic radiation these lines are in? Of course, these lines are in the UV region, and they are not visible, but they are detected by instruments; these lines form a Lyman series. The existences of the Lyman series and Balmer's series suggest the existence of more series. For example, the series with $n_2^2 = 3$ and $n_1^2$ = 4, 5, 6, 7, ... is called the Pashen series.
Multiple series
The spectral lines are grouped into series according to $n_1$ values. Lines are named sequentially starting from the longest wavelength/lowest frequency of the series using Greek letters within each series. For example, the ($n_1=1/n_2=2$) line is called "Lyman-alpha" (Ly-α), while the ($n_1=3/n_2=7$) line is called "Paschen-delta" (Pa-δ). The first six series have specific names:
• Lyman series with $n_1 = 1$
• Balmer series with $n_1 = 2$
• Paschen series (or Bohr series) with $n_1 = 3$
• Brackett series with $n_1 = 4$
• Pfund series with $n_1 = 5$
• Humphreys series with $n_1 = 6$
The spectral series of hydrogen based on the Rydberg Equation (on a logarithmic scale).
Example $1$: The Lyman Series
The so-called Lyman series of lines in the emission spectrum of hydrogen corresponds to transitions from various excited states to the n = 1 orbit. Calculate the wavelength of the lowest-energy line in the Lyman series to three significant figures. In what region of the electromagnetic spectrum does it occur?
Given: lowest-energy orbit in the Lyman series
Asked for: wavelength of the lowest-energy Lyman line and corresponding region of the spectrum
Strategy:
1. Substitute the appropriate values into Equation $Ref{2.1.2.2}$ (the Rydberg equation) and solve for $\lambda$.
2. Locate the region of the electromagnetic spectrum corresponding to the calculated wavelength.
Solution:
We can use the Rydberg equation (Equation $Ref{2.1.2.2}$ to calculate the wavelength:
$\dfrac{1}{\lambda }=R_H \left ( \dfrac{1}{n_l^{2}} - \dfrac{1}{n_h^{2}}\right ) \nonumber$
A For the Lyman series, $n_1 = 1$.
\begin{align*} \dfrac{1}{\lambda } &=R_H \left ( \dfrac{1}{n_l^{2}} - \dfrac{1}{n_h^{2}}\right ) \[4pt] &=1.097 \times m^{-1}\left ( \dfrac{1}{1}-\dfrac{1}{4} \right )\[4pt] &= 8.228 \times 10^{6}\; m^{-1} \end{align*} \nonumber
Spectroscopists often talk about energy and frequency as equivalent. The cm-1 unit (wavenumbers) is particularly convenient. We can convert the answer in part A to cm-1.
\begin{align*} \widetilde{ u} &=\dfrac{1}{\lambda } \[4pt] &= 8.228\times 10^{6}\cancel{m^{-1}}\left (\dfrac{\cancel{m}}{100\;cm} \right ) \[4pt] &= 82,280\: cm^{-1} \end{align*} \nonumber
and
$\lambda = 1.215 \times 10^{−7}\; m = 122\; nm \nonumber$
This emission line is called Lyman alpha and is the strongest atomic emission line from the sun; it drives the chemistry of the upper atmosphere of all the planets producing ions by stripping electrons from atoms and molecules. It is completely absorbed by oxygen in the upper stratosphere, dissociating O2 molecules to O atoms, which react with other O2 molecules to form stratospheric ozone (O3).
B This wavelength is in the ultraviolet region of the spectrum.
Exercise $1$: The Pfund Series
The Pfund series of lines in the emission spectrum of hydrogen corresponds to transitions from higher excited states to $n_1 = 5$. Calculate the wavelength of the second line in the Pfund series to three significant figures. In which region of the spectrum does it lie?
Answer
$4.65 \times 10^3\, nm$; infrared
The above discussion presents only a phenomenological description of hydrogen emission lines. Balmer and Rydberg could predict where emission lines could occur, but they could not explain why lines followed this pattern using any useful theory. For an explanation of why atoms demonstrate discrete emission spectra, quantum theories were developed.
The Bohr Model
In 1913, a Danish physicist, Niels Bohr (1885–1962; Nobel Prize in Physics, 1922), proposed a theoretical model for the hydrogen atom that explained its emission spectrum. Bohr’s model required only one assumption: The electron moves around the nucleus in circular orbits that can have only certain allowed radii. Rutherford’s earlier model of the atom had also assumed that electrons moved in circular orbits around the nucleus and that the atom was held together by the electrostatic attraction between the positively charged nucleus and the negatively charged electron. Although we now know that the assumption of circular orbits was incorrect, Bohr’s insight also posited that the electron could occupy only certain regions of space.
Using classical physics, Niels Bohr showed that the energy of an electron in a particular orbit is given by
$E_{n}=\dfrac{-R_Hhc}{n^{2}} \label{2.1.2.4}$
where $R_H$ is the Rydberg constant, h is Planck’s constant, c is the speed of light, and n is a positive integer corresponding to the number assigned to the orbit, with n = 1 corresponding to the orbit closest to the nucleus. In this model n = ∞ corresponds to the level where the energy holding the electron and the nucleus together is zero. In that level, the electron is unbound from the nucleus and the atom has been separated into a negatively charged (the electron) and a positively charged (the nucleus) ion. In this state the radius of the orbit is also infinite. The atom has been ionized.
Niels Bohr (1885–1962)
During the Nazi occupation of Denmark in World War II, Bohr escaped to the United States, where he became associated with the Atomic Energy Project.
In his final years, he devoted himself to the peaceful application of atomic physics and to resolving political problems arising from the development of atomic weapons.
As n decreases, the energy holding the electron and the nucleus together becomes increasingly negative; the radius of the orbit shrinks and more energy is needed to ionize the atom. The orbit with n = 1 is the lowest lying and most tightly bound. The negative sign in Equation $\ref{6.3.3}$ indicates that the electron-nucleus pair is more tightly bound (i.e. at a lower potential energy) when they are near each other than when they are far apart. Because a hydrogen atom with its one electron in this orbit has the lowest possible energy, this is the ground state (the most stable arrangement of electrons for an element or a compound) for a hydrogen atom. As n increases, the radius of the orbit increases; the electron is farther from the proton, which results in a less stable arrangement with higher potential energy (Figure $\PageIndex{2a}$). A hydrogen atom with an electron in an orbit with n > 1 is therefore in an excited state, defined as any arrangement of electrons that is higher in energy than the ground state. When an atom in an excited state undergoes a transition to the ground state in a process called decay, it loses energy by emitting a photon whose energy corresponds to the difference in energy between the two states (Figure $1$).
So the difference in energy (ΔE) between any two orbits or energy levels is given by $\Delta E=E_{n_l}-E_{n_h}$, where $n_l$ is the final orbit and $n_h$ the initial orbit. Substituting from Bohr’s equation (Equation 2.1.2.4) for each energy value gives
$\Delta E=E_{final}-E_{initial}=-\dfrac{R_H hc}{n_h^{2}}-\left ( -\dfrac{R_H hc}{n_l^{2}} \right )=-R_H hc\left ( \dfrac{1}{n_h^{2}} - \dfrac{1}{n_l^{2}}\right ) \label{6.3.4}$
If nh > nl, the transition is from a higher energy state (larger-radius orbit) to a lower energy state (smaller-radius orbit), as shown by the dashed arrow in part (a) in Figure $3$. Substituting hc/λ for ΔE gives
$\Delta E = \dfrac{hc}{\lambda }=-R_H hc\left ( \dfrac{1}{n_h^{2}} - \dfrac{1}{n_l^{2}}\right ) \label{6.3.5}$
Canceling hc on both sides gives
$\dfrac{1}{\lambda }=-R_H \left ( \dfrac{1}{n_h^{2}} - \dfrac{1}{n_l^{2}}\right ) \label{6.3.6}$
Except for the negative sign, this is the same equation that Rydberg obtained experimentally. The negative sign in Equations $\ref{6.3.5}$ and $\ref{6.3.6}$ indicates that energy is released as the electron moves from orbit $n_h$ to orbit $n_l$ because orbit $n_h$ is at a higher energy than orbit $n_l$. Bohr calculated the value of $R_H$ from fundamental constants such as the charge and mass of the electron and Planck's constant and obtained a value of 1.0974 × 107 m−1, the same number Rydberg had obtained by analyzing the emission spectra.
We can now understand the physical basis for the Balmer series of lines in the emission spectrum of hydrogen ($\PageIndex{3b}$); the lines in this series correspond to transitions from higher-energy orbits (n > 2) to the second orbit (n = 2). Thus, the hydrogen atoms in the sample have absorbed energy from the electrical discharge and decayed from a higher-energy excited state (n > 2) to a lower-energy state (n = 2) by emitting a photon of electromagnetic radiation whose energy corresponds exactly to the difference in energy between the two states (Figure $\PageIndex{3a}$). The n = 3 to n = 2 transition gives rise to the line at 656 nm (red), the n = 4 to n = 2 transition to the line at 486 nm (green), the n = 5 to n = 2 transition to the line at 434 nm (blue), and the n = 6 to n = 2 transition to the line at 410 nm (violet). Because a sample of hydrogen contains a large number of atoms, the intensity of the various lines in a line spectrum depends on the number of atoms in each excited state. At the temperature in the gas discharge tube, more atoms are in the n = 3 than the n ≥ 4 levels. Consequently, the n = 3 to n = 2 transition is the most intense line, producing the characteristic red color of a hydrogen discharge (Figure $\PageIndex{1a}$). Other families of lines are produced by transitions from excited states with n > 1 to the orbit with n = 1 or to orbits with n ≥ 3. These transitions are shown schematically in Figure $4$ :
Using Atoms to Time
In contemporary applications, electron transitions are used in timekeeping that needs to be exact. Telecommunications systems, such as cell phones, depend on timing signals that are accurate within a millionth of a second per day; the same goes for the devices that control the US power grid. Global positioning system (GPS) signals must be accurate within a billionth of a second per day, which is equivalent to gaining or losing no more than one second in 1,400,000 years. Quantifying time requires finding an event with an interval that repeats on a regular basis.
To achieve the accuracy required for modern purposes, physicists have turned to the atom. The current standard used to calibrate clocks is the cesium atom. Supercooled cesium atoms are placed in a vacuum chamber and bombarded with microwaves whose frequencies are carefully controlled. When the frequency is exactly right, the atoms absorb enough energy to undergo an electronic transition to a higher-energy state. Decay to a lower-energy state emits radiation. The microwave frequency is continually adjusted, serving as the clock’s pendulum.
In 1967, the second was defined as the duration of 9,192,631,770 oscillations of the resonant frequency of a cesium atom, called the cesium clock. Research is currently under way to develop the next generation of atomic clocks that promise to be even more accurate. Such devices would allow scientists to monitor vanishingly faint electromagnetic signals produced by nerve pathways in the brain and allow geologists to measure variations in gravitational fields, which cause fluctuations in time, that would aid in the discovery of oil or minerals.
Example $1$: The Lyman Series
The so-called Lyman series of lines in the emission spectrum of hydrogen corresponds to transitions from various excited states to the n = 1 orbit. Calculate the wavelength of the lowest-energy line in the Lyman series to three significant figures. In what region of the electromagnetic spectrum does it occur?
Given: lowest-energy orbit in the Lyman series
Asked for: wavelength of the lowest-energy Lyman line and corresponding region of the spectrum
Strategy:
1. Substitute the appropriate values into Equation 2.1.2.4 (the Rydberg equation) and solve for $\lambda$.
2. Use Figure 2.2.1 to locate the region of the electromagnetic spectrum corresponding to the calculated wavelength.
Solution:
We can use the Rydberg equation to calculate the wavelength:
$\dfrac{1}{\lambda }=-R_H \left ( \dfrac{1}{n_h^{2}} - \dfrac{1}{n_l^{2}}\right ) \nonumber$
A For the Lyman series, /(n_l = 1\). The lowest-energy line is due to a transition from the n = 2 to n = 1 orbit because they are the closest in energy.
$\dfrac{1}{\lambda }=-R_H \left ( \dfrac{1}{n_h^{2}} - \dfrac{1}{n_l^{2}}\right )=1.097\times m^{-1}\left ( \dfrac{1}{1}-\dfrac{1}{4} \right )=8.228 \times 10^{6}\; m^{-1} \nonumber$
It turns out that spectroscopists (the people who study spectroscopy) use cm-1 rather than m-1 as a common unit. Wavelength is inversely proportional to energy but frequency is directly proportional as shown by Planck's formula, E=h$u$.
Spectroscopists often talk about energy and frequency as equivalent. The cm-1 unit is particularly convenient. The infrared range is roughly 200 - 5,000 cm-1, the visible from 11,000 to 25,000 cm-1 and the UV between 25,000 and 100,000 cm-1. The units of cm-1 are called wavenumbers, although people often refer to the unit as inverse centimeters. We can convert the answer in part A to cm-1.
$\widetilde{ u} =\dfrac{1}{\lambda }=8.228\times 10^{6}\cancel{m^{-1}}\left (\dfrac{\cancel{m}}{100\;cm} \right )=82,280\: cm^{-1} \nonumber$
and
$\lambda = 1.215 \times 10^{−7}\; m = 122\; nm \nonumber$
This emission line is called Lyman alpha. As the strongest atomic emission line from the sun, it drives the chemistry of the upper atmosphere of all the planets, producing ions by stripping electrons from atoms and molecules. It is completely absorbed by oxygen in the upper stratosphere, dissociating O2 molecules to O atoms which react with other O2 molecules to form stratospheric ozone.
B This wavelength is in the ultraviolet region of the spectrum.
Exercise $1$: The Pfund Series
The Pfund series of lines in the emission spectrum of hydrogen corresponds to transitions from higher excited states to the n = 5 orbit. Calculate the wavelength of the second line in the Pfund series to three significant figures. In which region of the spectrum does it lie?
Answer
4.65 × 103 nm; infrared
Bohr’s model of the hydrogen atom gave an exact explanation for its observed emission spectrum. The following are his key contributions to our understanding of atomic structure:
• Electrons can occupy only certain regions of space, called orbits.
• Orbits closer to the nucleus are lower in energy.
• Electrons can move from one orbit to another by absorbing or emitting energy, giving rise to characteristic spectra.
Unfortunately, Bohr could not explain why the electron should be restricted to particular orbits. Also, despite a great deal of conjecturing, such as assuming that orbits could be ellipses rather than circles, his model could not quantitatively explain the emission spectra of any element other than hydrogen (Figure $5$). In fact, Bohr’s model worked only for species that contained just one electron: H, He+, Li2+, and so forth. Scientists needed a fundamental change in their way of thinking about the electronic structure of atoms to advance beyond the Bohr model.
The Energy States of the Hydrogen Atom
Thus far, we have explicitly considered only the emission of light by atoms in excited states, which produces an emission spectrum (a spectrum produced by the emission of light by atoms in excited states). The converse, absorption of light by ground-state atoms to produce an excited state, can also occur, producing an absorption spectrum (a spectrum produced by the absorption of light by ground-state atoms).
When an atom emits light, it decays to a lower energy state; when an atom absorbs light, it is excited to a higher energy state.
If white light is passed through a sample of hydrogen, hydrogen atoms absorb energy as an electron is excited to higher energy levels (orbits with n ≥ 2). If the light that emerges is passed through a prism, it forms a continuous spectrum with black lines (corresponding to no light passing through the sample) at 656, 468, 434, and 410 nm. These wavelengths correspond to the n = 2 to n = 3, n = 2 to n = 4, n = 2 to n = 5, and n = 2 to n = 6 transitions. Any given element therefore has both a characteristic emission spectrum and a characteristic absorption spectrum, which are essentially complementary images.
Emission and absorption spectra form the basis of spectroscopy, which uses spectra to provide information about the structure and the composition of a substance or an object. In particular, astronomers use emission and absorption spectra to determine the composition of stars and interstellar matter. As an example, consider the spectrum of sunlight shown in Figure $7$. Since the sun is very hot, it emits light in the form of a continuous emission spectrum. Superimposed on the spectrum, however, is a series of dark lines primarily resulting from the absorption of specific frequencies of light by cooler atoms in the outer atmosphere of the sun. By comparing these lines with the spectra of elements measured on Earth, we now know that the sun contains large amounts of hydrogen, iron, and carbon, along with smaller amounts of other elements. During the solar eclipse of 1868, the French astronomer Pierre Janssen (1824–1907) observed a set of lines that did not match those of any known element. He suggested that they were due to the presence of a new element, which he named helium, from the Greek helios, meaning “sun.” Helium (He) was finally discovered in uranium ores on Earth in 1895. Alpha particles are helium nuclei. Alpha particles emitted by radioactive uranium nuclei pick up electrons from the rocks to form helium atoms.
The familiar red color of neon signs used in advertising is due to the emission spectrum of neon shown in part (b) in Figure $5$. Similarly, the blue and yellow colors of certain street lights are caused, respectively, by mercury and sodium discharges. In all these cases, an electrical discharge excites neutral atoms to a higher energy state, and light is emitted when the atoms decay to the ground state. In the case of mercury, most of the emission lines are below 450 nm, which produces a blue light (part (c) in Figure $5$). In the case of sodium, the most intense emission lines are at 589 nm, which produces an intense yellow light.
Summary
There is an intimate connection between the atomic structure of an atom and its spectral characteristics. Atoms of individual elements emit light at only specific wavelengths, producing a line spectrum rather than the continuous spectrum of all wavelengths produced by a hot object. Niels Bohr explained the line spectrum of the hydrogen atom by assuming that the electron moved in circular orbits and that orbits with only certain radii were allowed. Lines in the spectrum were due to transitions in which an electron moved from a higher-energy orbit with a larger radius to a lower-energy orbit with smaller radius. The orbit closest to the nucleus represented the ground state of the atom and was most stable; orbits farther away were higher-energy excited states. Transitions from an excited state to a lower-energy state resulted in the emission of light with only a limited number of wavelengths. Atoms can also absorb light of certain energies, resulting in a transition from the ground state or a lower-energy excited state to a higher-energy excited state. This produces an absorption spectrum, which has dark lines in the same position as the bright lines in the emission spectrum of an element. Bohr’s model revolutionized the understanding of the atom but could not explain the spectra of atoms heavier than hydrogen. | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/02%3A_Atomic_Structure/2.01%3A_Historical_Development_of_Atomic_Theory/2.1.02%3A_Discovery_of_Subatomic_Particles_and_the_Bohr_Atom.txt |
Considering the failures of the Bohr model, Erwin Schrödinger and Werner Heisenberg proposed a major change in the paradigm regarding the electron. In several breakthrough papers (1925-1927) they attributed wave properties to electrons and they each received Nobel prizes for developing the theories of Wave Mechanics (or the "New Quantum Mechanics"). This approach treated electrons as having "dual" nature: possessing properties of both waves and particles.
Schrödinger's Equation describes the behavior of the electron (in a hydrogen atom) in three dimensions. It is a mathematical equation that defines the electron’s position, mass, total energy, and potential energy. The simplest form of the Schrödinger Equation is as follows:
$\hat{H}\psi = E\psi \nonumber$
where $\hat{H}$ is the Hamiltonian operator, $E$ is the energy of the electron, and $\psi$ is the wavefunction.
The Hamiltonian, $\hat{H}$
The Hamiltonian operator is like a set of instructions that tells us what to do with the function that follows it. A Hamiltonian operator is a function over three-dimensional space that corresponds to the sum of kinetic energies and potential energies of the particles in a system, one electron and its nucleus in this case. The Hamiltonian operator for a one-electron system is:
$\hat{H}=\dfrac{-h{^2}}{8\pi{^2}m_e}\left(\dfrac{\partial{^2}}{\partial{x^2}}+\dfrac{\partial{^2}}{\partial{y^2}}+\dfrac{\partial{^2}}{\partial{z^2}}\right)-\dfrac{Ze^2}{4\pi{}\epsilon_0{r}}, \nonumber$
where $h$ is Planck's constant, $m_e$ is the mass of the electron, $e$ is the charge of the electron, $r$ is the distance from the nucleus ($r=\sqrt{x^2+y^2+z^2}$), $Z$ is the charge of the nucleus, and $4\pi{}\epsilon_0$ is the permittivity of a vacuum.
Kinetic Energy
The first part of the Hamiltonian written above, $\dfrac{-h{^2}}{8\pi{^2}m_e}\left(\dfrac{\partial{^2}}{\partial{x^2}}+\dfrac{\partial{^2}}{\partial{y^2}}+\dfrac{\partial{^2}}{\partial{z^2}}\right)$ describes the kinetic energy of the electron. This is the energy due to motion of the electron.
Potential energy
The second part written above, $\dfrac{-Ze^2}{4\pi{}\epsilon_0{r}}$, describes the potential energy of the electron, and is commonly written as $V(r)$ or $V(x,y,z)$.
$V(x,y,x) = \dfrac{-Ze^2}{4\pi{}\epsilon_0{r}} = \dfrac{-Ze^2}{4\pi{}\epsilon_0{\sqrt{x^2+y^2+z^2}}} \nonumber$
The potential energy depends on the attractive electrostatic force between the electron and the nucleus. You might notice that this attraction is essentially the same as the electrostatic force defined by Coulomb's law. And, just as in Coulomb's law, when two opposite charges are attracted to one another, the potential energy of the force is negative. Thus, when an electron is close to the nucleus, the potential energy is a large negative number corresponding to a strong attractive force. When an electron is farther from the nucleus, the potential energy is still negative but with a smaller magnitude, corresponding to a weaker attractive force. If the electron is very far from the nucleus ($r = \infty$) then the attractive force, and the potential energy, is zero.
The Wavefunction, $\psi$
In simple terms, the wavefunction ($\psi$) of an electron describes the electron's position in space, relative to the nucleus. The square of the $\psi$ describes an atomic orbital. We can't define the position too exactly because we would violate the Heisenberg Uncertainty principle, but we can define its wave. A simple example of a $\psi$ is described in the next section: Particle in a Box. Here, we will describe the $\psi$ in general terms. Generally, in a one-electron atom, the electron $\psi$ is defined by the wave's distance from the nucleus and its angle with respect to the x, y, and z axes of the atom's Cartesian coordinates (the nucleus is at the origin). The general form of the ($\psi$) for an electron in a hydrogen atom can be written as follows:
$\psi_{n,l,m_l} = R_{n,l}(r) + Y_{l,m_l}(\theta,\phi) \nonumber$
Quantum Numbers define $\psi$
The ($\psi$) is defined by three of the quantum numbers: $n$, $l$, and $m_l$. These quantum numbers will be discussed more in a later section (2.2.2) The radial variation, $R$, depends on the electron's distance from the nucleus. The quantum numbers $n$ (energy level) and $l$ (orbital type) define $R$. Since $n$ must be an integer, there are only certain allowed values for the solution to the wavefunction.
The angular variation, $Y$, depends on the angle with respect to the x, y, and z coordinates, and depends on the quantum numbers $l$ (the orbital type) and $m_l$ (the angular momentum, or the specific orbital). For example $p_x$ lies along the $x$ axis, while $p_y$ points in a different direction in space.
For review, a list of the quantum numbers, their values, and meanings are in the table below.
SYMBOL NAME VALUES MEANING
$n$ principal $1,2,3...$(any integer) energy level, shell
$l$ angular momentum $0 \rightarrow n-1$
subshell, $0=s, 1=p, 2=d, 3=f...$
this is the angular dependence of the orbital, shape of the orbital
*letters have historical meaning, sharp, principle, diffuse, fundamental
$m_l$ magnetic $+l \rightarrow -l$ orientation of angular momentum in space, orbital
$m_s$ spin $+\frac{1}{2}, -\frac{1}{2}$ the imaginary property we call "spin", up or down
Some important considerations and limitations
Although it might seem like there could be any value of x, y, and z for the Hamiltonian, these values are limited by the allowed positions of electrons according to $\psi$, which is limited by integer values of $n$. In other words the allowed solutions are quantized. However, there are an infinite number of values for $n$ from $n=1\rightarrow\infty$, so there are also infinite solutions to the Schrödinger equation.
The $\psi$ describes the wave properties of an electron. The probability of finding the electron somewhere in space is the square of the wavefunction ($\psi^2$ or $\psi \psi^*$). In other words, $\psi^2$ describes the shape and size of an electron's orbital (the shapes you already know).
There are some requirements for a physically realistic and meaningful solution for $\psi$, and thus $\psi^2$.
1. There is only one possible value for $\psi$ for any set of the three quantum numbers $n, l, m_l$.
2. The $\psi$ approaches zero as $r$ approaches infinity, and so $\psi^2$ also approaches zero as $r\rightarrow\infty$.
3. The wavefunction must be normalized. In other words, the total probability of finding the electron in all of space must be 1. $\int_{\text {all space}} \psi_{A} \psi_{A}^{*} d \tau=1 \nonumber$
4. Any two orbitals must not occupy the same space. In other words, any two orbitals in an atom are orthogonal. If $\psi_{A}$ and $\psi_{B}$ are wavefunctions for different orbitals in the same atom, $\int_{\text {all space}} \psi_{A} \psi_{B}^{*} d \tau=0 \nonumber$
5. The probability of finding the electron anywhere in infinite space must be defined. This means that the wave functions and their first derivatives must be continuous (i.e. not change abruptly from one point to the next).
2.02: The Schrodinger equation particle in a box and atomic wavefunctions
The particle in the box is a model that can illustrate how a wave equation works. Although it does not represent a real situation, we can limit our model to just one dimension (the x-dimension, for instance) such that the Schrödinger equation becomes significantly simplified. Despite being unrealistic, this simplification is quite useful for gaining an understanding of the Schrödinger equation.
The model of a particle in a box
The particle in the box is a hypothetical situation with a particle trapped in a one-dimensional “box”. Let’s not get hung-up on the fact that the common object called a “box” is typically an object with three dimensions instead of just one dimension. And let’s also not get hung up on the word “particle”. This is a particle that has properties of a wave…so it is unlike the macroscopic particle that you’re probably imagining. This “box” is more like a line, or an x-axis; it is just a one-dimensional space in which a particle-wave is trapped.
The particle-wave can only exist inside the walls (where $0<x<a$), along the x axis in the figure shown above. In terms of energy, the potential energy is zero ($V=0$) because the particle is in an energetically-favorable position here. On the other hand, outside the box, the particle cannot exist and the potential energy is infinitely large ($V=\infty$) outside the walls (where $x<0$ or $x>a$). This means that it is infinitely unfavorable for the particle-wave to exist outside the box, and so it never does. The particle-wave is trapped between the walls, along the 1-dimensional $x$ axis, and there are no forces acting on the particle-wave inside this “box”.
The Schrödinger wave equation for a particle in a box
The particle in a box model lets us consider a simple version of the Schrödinger equation. Before we simplify, let's take another look at the full Hamiltonian for a particle-wave in three dimensions (see equation 2.2.2) and the simplest form of the Schrödinger equation (see equation 2.2.1). Both of these equations are described in the previous section and are written below for convenience.
The Schrödinger Equation (from equation 2.2.1): $\hat{H}\psi = E\psi$
The Hamiltonian in three dimensions (from equation 2.2.2):
$\hat{H}=\dfrac{-h{^2}}{8\pi{^2}m_e}\left(\dfrac{\partial{^2}}{\partial{x^2}}+\dfrac{\partial{^2}}{\partial{y^2}}+\dfrac{\partial{^2}}{\partial{z^2}}\right)-\dfrac{Ze^2}{4\pi{}\epsilon_0{r}}\nonumber$
In the particle in a box model, there is only one dimension, $x$. Because the $y$ and $z$ values are zero, we can drop $y$ and $z$ out of our Hamiltonian equation. And, since $V=0$ inside the box, we can drop the whole part of the Hamiltonian equation that describes the potential energy ($\frac{-Ze^2}{4\pi{}\epsilon_0{r}}$). This leaves us with a much simplified Hamiltonian operator, which we can then use to write the Schrödinger wave equation for a particle moving in one dimension. One last thing we'll do is use a more general value $m$ for the mass of the particle, rather than $m_e$ that specifically represents the mass of an electron. The Schrödinger equation for a particle-wave in a one-dimensional box is:
$\dfrac{-h{^2}}{8\pi{^2}m}\left(\dfrac{\partial{^2}\psi(x)}{\partial{x^2}}\right) = E\psi \label{1DSchr}$
Exercise $1$
Follow the steps that were described in the last paragraph to find Equation \ref{1DSchr} from equation 2.2.1 and 2.2.2. In other words, derive the Schrödinger Equation for the particle in a box given the three-dimensional Hamiltonian and Schrödinger Equations.
Answer
Here is the Hamitlonian Equation for in three dimensions.
$\hat{H}=\textcolor{red}{\dfrac{-h{^2}}{8\pi{^2}m}\left(\dfrac{\partial{^2}}{\partial{x^2}}+\dfrac{\partial{^2}}{\partial{y^2}}+\dfrac{\partial{^2}}{\partial{z^2}}\right)}-\textcolor{blue}{\dfrac{Ze^2}{4\pi{}\epsilon_0{r}}}$
There are two parts to this equation: the kinetic energy contribution and the potential energy contribution. While the first part of this expression, $\textcolor{red}{\dfrac{-h{^2}}{8\pi{^2}m}\left(\dfrac{\partial{^2}}{\partial{x^2}}+\dfrac{\partial{^2}}{\partial{y^2}}+\dfrac{\partial{^2}}{\partial{z^2}}\right)}$, is the kinetic energy contribution, the second part of the expression, $\textcolor{blue}{-\dfrac{Ze^2}{4\pi{}\epsilon_0{r}}}$, is the potential energy ($V$) contribution.
(1) Simplify the expression for kinetic energy
Your particle-wave is moving in only one dimension. We can arbitrarily choose the dimension $x$. This means that $x$ is a variable and will have some value from $0$ to $a$. On the other hand, the particle does not move in the other two dimensions, $y$ and $z$. You can assign the position of the particle along $x$ and $y$ axes as zero or you can just understand that they do not exist in any significant capacity in our 1-D space. Either way, you can just drop them out of the kinetic energy part of the Hamiltonian, as so:
When $y=z=0$ and the particle-wave only exists/moves in the $x$ dimension, $\textcolor{red}{\dfrac{-h{^2}}{8\pi{^2}m}\left(\dfrac{\partial{^2}}{\partial{x^2}}+\dfrac{\partial{^2}}{\partial{y^2}}+\dfrac{\partial{^2}}{\partial{z^2}}\right)}$ = $\textcolor{red}{\dfrac{-h{^2}}{8\pi{^2}m}\left(\dfrac{\partial{^2}}{\partial{x^2}}\right)}$
And the overall Hamiltonian is simplified:
$\hat{H}=\textcolor{red}{\dfrac{-h{^2}}{8\pi{^2}m}\left(\dfrac{\partial{^2}}{\partial{x^2}}\right)}-\textcolor{blue}{\dfrac{Ze^2}{4\pi{}\epsilon_0{r}}}$
(2) Simplify the expression for potential energy ($V$).
The particle's potential energy inside the box is zero ($V=0$). Therefore, $\textcolor{blue}{V=-\dfrac{Ze^2}{4\pi{}\epsilon_0{r}}}=0$. We can simply drop the potential energy portion from the Hamiltonian because its value is zero.
The simplified Hamiltonian is now:
$\hat{H}=\textcolor{red}{\dfrac{-h^2}{8\pi^2m}\left(\dfrac{\partial{^2}}{\partial{x^2}}\right)}$
(3) Substitute the simplified Hamiltonian into the Schrödinger Equation.
The Schrödinger is: $\textcolor{red}{\hat{H}}\psi = E\psi$
When the simplified Hamiltonian is substituted into this equation, the result is the 1-D Schrödinger Equation \ref{1DSchr}.
$\textcolor{red}{\left(\dfrac{-h{^2}}{8\pi{^2}m}\left(\dfrac{\partial{^2}\psi(x)}{\partial{x^2}}\right)\right)} = E\psi$
The wavefunction of a one-dimensional wave
Although we've simplified the Schrödinger Equation by considering the particle in the box, Equation \ref{1DSchr} may still look mysterious to you. But it's simpler than you might realize. Let's unpack it! Here we'll go through the steps of deriving the 1-dimensional wavefunction for the particle in a box. We won't try to derive a three-dimensional wavefunction for a "real" electron, but when you understand how to find the 1-dimensional wavefunction, and what it tells us, then you can conceptually extend this to understand what the wavefunction tells us about electrons in three dimensions.
Let's dissect Equation \ref{1DSchr}. Here it is again for reference:
$\textcolor{green}{\dfrac{-h{^2}}{8\pi{^2}m_e}}\left(\dfrac{\partial{^2}\psi(x)}{\partial{x^2}}\right) = \textcolor{green}{E}\psi\nonumber$
The expression $\textcolor{green}{\frac{-h^2}{8\pi^2m_e}}$ is just a negative constant. This negative constant is multiplied by the second derivative of the wavefunction in the expression above. And $\textcolor{green}{E}$ is a constant multiplied by that same function. We can rearrange this equation as follows:
$\dfrac{\partial{^2}\psi(x)}{\partial{x^2}} = \textcolor{green}{\dfrac{-8\pi{^2}m_eE}{h{^2}}}\psi = -\textcolor{green}{C}\psi \label{Schrx}$
where $C$ is just the constant; $\textcolor{green}{C}=\textcolor{green}{\dfrac{8\pi{^2}m_eE}{h{^2}}}$
To an expert in the mathematical field of Differential Equations, this expression \ref{1DSchr} follows a familiar pattern that makes it easy to translate. If you're not an expert in differential equations, that is OK; you'll just have to bear with this while things get a little "hand-wavy" and try Exercise $2$ to prove it to yourself that the following "hand waviness" is true. An expert in differential equations could tell you that when the second derivative of a function is equal to a negative constant times that same function, the function can be written in terms of $sin$ and $cos$, as shown in Equation \ref{wf1} (that was the hand-wavy part...see, not so bad):
$\psi=A\sin(rx) + B\cos(sx) \label{wf1}$
$A, B, r,$ and $s$ are constants in Equation \ref{wf1}. This expression is useful because it is in a form that we could plot with a graphing program (or calculator). However, to do so we must know the constants $A, B, r,$ and $s$.
What are $r$ and $s$?
If we substitute expression \ref{wf1} into Equation \ref{1DSchr} and solve for r and s, we find the expressions below.
$r=s=\sqrt{2m_eE}\left(\frac{2\pi}{h}\right) \label{rs}$
This shows us that r and s are equal and constant; we know the values for $\pi, h,$ and $m_e$, and we can solve for the constant $E$.
Exercise $2$: Prove that Equation \ref{wf1} is true
Show that Equation \ref{wf1} is true in the case that Equation \ref{Schrx} ( $\frac{\partial{^2}\psi(x)}{\partial{x^2}}=-C\psi$) is true. Also consider Equation \ref{rs} ($r,s$ are constants and $r=s$) and explain why the constant $r$ must equal $s$.
To do this problem, you'll also need to know the basic differentiation rules of $\sin$ and $\cos$ functions, given below for convenience.
Differentiation Rules for $sin$ and $cos$:
$\frac{d}{dx}\big(\sin(rx)\big)=r\cos(rx)$, where r is a constant.
$\frac{d}{dx}\big(\cos(rx)\big)=-r\sin(rx)$, where r is a constant.
Answer
You can show that \ref{wf1} is true by substituting \ref{wf1} back into Equation \ref{Schrx}; in other words take the second derivative of the function, $\psi=A\sin(rx) + B\cos(sx)$ to show that $\frac{\partial{^2}\psi(x)}{\partial{x^2}}=-C\psi$. $C, r,$ and $s$ are arbitrary constants, and $r=s$. The actual value of these constants is irrelevant for this problem. Here is one way to approach the problem:
(1) Substitute $\textcolor{red}{\psi=A\sin(rx) + B\cos(sx)}$ into $\textcolor{blue}{\frac{\partial{^2}\psi(x)}{\partial{x^2}}=-C\psi}$.
$\textcolor{blue}{\frac{\partial{^2}}{\partial{x^2}}}\left(\textcolor{red}{A\sin(rx) + B\cos(sx)}\right)= \textcolor{blue}{-C}\textcolor{red}{\psi}$
(2) Take the first and second derivative of $\textcolor{red}{A\sin(rx) + B\cos(sx)}$.
The result of taking the first derivative...
$\textcolor{blue}{\frac{\partial{}}{\partial{x}}}\left(\textcolor{green}{r}\textcolor{red}{A\cos(rx) - \textcolor{green}{s}B\sin(sx)}\right)= \textcolor{blue}{-C}\textcolor{red}{\psi}$
...and the second derivative...
$\textcolor{green}{-r^2}\textcolor{red}{A\sin(rx)} - \textcolor{green}{s^2}\textcolor{red}{B\cos(sx)} = \textcolor{blue}{-C}\textcolor{red}{\psi}$
This is where we see that the only way $\psi=A\sin(rx) + B\cos(sx)$ is if we can factor the values $r^2$ and $s^2$ out of the left side. We have a hint that $r=s$ from the discussion above and Equation \ref{rs}. Now we see that this is a necessary condition for the expression $\textcolor{green}{-r^2}\textcolor{red}{A\sin(rx)} - \textcolor{green}{s^2}\textcolor{red}{B\cos(sx)} = \textcolor{blue}{-C}\textcolor{red}{\psi}$ to be true.
(3) Simplify
If $r=s$, then $r^2=s^2$ and we can replace s with r in the equation above, as so...
$\textcolor{green}{-r^2}\textcolor{red}{A\sin(rx)} - \textcolor{green}{r^2}\textcolor{red}{B\cos(sx)} = \textcolor{blue}{-C}\textcolor{red}{\psi}$
Now, we can factor the constant, $-r^2$ out of the left side of this expression,
$\textcolor{green}{-r^2}\left(\textcolor{red}{A\sin(rx) + B\cos(sx)}\right) = \textcolor{blue}{-C}\textcolor{red}{\psi}$
Now we see that $\psi=A\sin(rx) + B\cos(sx)$ when $-C=-r^2$. In other words, we have shown that $\psi=A\sin(rx) + B\cos(sx)$ is true.
Why is this a useful exercise?
Although it is difficult to explain the derivation of the expression $\psi=A\sin(rx) + B\cos(sx)$ without differential equations and only knowing that $\frac{\partial{^2}\psi(x)}{\partial{x^2}}=-C\psi$, to show this is true is as simple as solving this exercise. This exercise is here to help you review your calculus and to show you that the "hand waving" in the text above is not magic, but it is coming from math that you have seen in your calculus courses. Math is crucial for explaining the nature of the universe! Eat your vegetables and practice your math.
Exercise \ref{rs}
*Complete Exercise $2$ before attempting this one.
Derive the expression for r below (from Equation \ref{rs}) using the 1-D wavefunction (\ref{wf1}: $\psi=A\sin(rx) + B\cos(sx)$) and the 1-D Schrödinger equation (\ref{1DSchr}: $\left(\frac{-h{^2}}{8\pi{^2}m_e}\left(\frac{\partial{^2}\psi(x)}{\partial{x^2}}\right) = E\psi\right)$).
$r=\dfrac{2\pi}{h}\sqrt{2m_eE}$
Answer
The steps below could be carried out in a different sequence:
(1) Rearrange
Let's move all constants to the right side of \ref{1DSchr} so that we arrive at the expression that is shown in part of Equation \ref{Schrx} .
$\dfrac{\partial{^2}\psi(x)}{\partial{x^2}} = \textcolor{green}{\dfrac{-8\pi{^2}m_eE}{h{^2}}}\psi$
(2) Take the second derivative of $\psi$, which we found in Exercise $2$.
$\dfrac{\partial{^2}\psi(x)}{\partial{x^2}} =\textcolor{green}{-r^2}\left(\textcolor{red}{A\sin(rx) + B\cos(sx)}\right) = \textcolor{green}{\dfrac{-8\pi{^2}m_eE}{h{^2}}}\textcolor{red}{\psi}$
And since $\textcolor{red}{\psi=A\sin(rx) + B\cos(sx)}$, we can simplify this by dividing both sides by $\psi$ (or by $A\sin(rx) + B\cos(sx)$).
**recall that $s=r$. So, you could have substituted $r$ for $s$ in the expression above, and you'd still be on the right track.
$\textcolor{green}{-r^2}= \textcolor{green}{\dfrac{-8\pi{^2}m_eE}{h{^2}}}$
(3) Solve for r (*we already showed $r=s$ in the Exercise $2$).
$\textcolor{green}{r} = \sqrt{\dfrac{8\pi{^2}m_eE}{h{^2}}} = \dfrac{2\pi}{h}\sqrt{2m_eE}$
What are $A$ and $B$?
We can find the possible values of $A$ and $B$ by looking at the possible extremes for the value of $x$. In our model in Figure $1$, we can see that x can have values from $0\rightarrow a$. The two extremes, where $x=0$ and $x=a$, lie at the walls of the box. The electron cannot exist beyond these values because it is trapped inside the walls, so its wavefunction must be zero at $x=0$ and $x=a$.
The case where $x=0$ can help us find the value of $B$. We just said that the wavefunction must be zero where $x=0$. This means:
$\psi_{x=0} = A\sin(rx) + B\cos(sx) = A\sin(0) + B\cos(0) = 0 \nonumber$
Because $A\sin(0)=0$, the expression can be simplified to:
$\psi_{x=0} = B\cos(0) = B(1) = 0 \nonumber$
Now we can see that there is only one possible value for the constant, $B$, that would allow this expression to be true, $B=0$. And since $B$ is constant then it will always be zero despite the value of $x$. This simplifies our 1-D wavefunction:
$\psi=A\sin(rx) \label{wf2}$
The case where $x=a$ can help us find the value for $A$. Since the wavefunction must be zero where $x=a$:
$\psi_{x=a} = A\sin(rx) = A\sin(ra) = 0 \nonumber$
In this case, $A$ cannot be zero because if both A and B are zero, the wavefunction is zero at all values of $x$, and so our wavefunction would not exist inside the box. So, let's assume A is not zero. Then, it must be the case that
$\sin(ra)=0 \nonumber$
and then because of the way that $\sin$ functions work, the quantity $ra$ must be an integer value of $\pi$ if $\sin(ra)=0$.
$ra = \pm n\pi \nonumber$
where $n\ = 1, 2, 3...$ any non-zero integer. We can ignore negative values here since both positive and negative values of $r$ will give the same value for $\sin(ra)$. We can then solve for r:
$r=\frac{n\pi}{a} \label{r}$
and substitute the value $r=\dfrac{n\pi}{a}$ into Equation \ref{wf2} to get:
$\psi = A\sin\left(\frac{n\pi{x}}{a}\right) \label{wf3}$
Recall that the square of the wavefunction ($\psi^2$) gives the probability of finding the electron anywhere in space. In our model, the probability of finding the particle-wave inside the box is unity (it is 1). Mathematically, this is the normalizing requirement (expressed as $\int \psi_A \psi_A^* d\tau=1$), which can be solved to find the value of A.
$A = \sqrt{\frac{2}{a}} \nonumber$
Now, we can substitute the value of $A$ into expression \ref{wf3}, and we have a wavefunction that can be visualized using any graphing program/calculator!
$\psi=\sqrt{\frac{2}{a}}\left(\sin\left(\frac{n\pi{x}}{a}\right)\right) \label{wf4}$
What is the energy (E) of a particle in the box?
So, now you can see how a 1-D wavefunction can be solved. But how do we find the E of the electron from this? Well, in order to not distract you earlier, we skipped over a very simple way to do this. The E falls right out of the equations you just saw. We can set the two expressions we found above for the constant $r$, (equations \ref{rs} and \ref{r}) equal to one another and then solve for E:
$r=\dfrac{n\pi}{a}=\sqrt{2mE}\left(\dfrac{2\pi}{h}\right) \nonumber$
$E=\frac{n^2h^2}{8ma^2} \label{E}$
Expression \ref{E} can be used to calculate the energy of a particle in a one-dimensional box of length $a$, given its integer energy level, $n$. Here we can see that the energy is quantized because $n$ is an integer ($n=1,2,3...$). In other words, $n$ is a quantum number.
Making sense of the particle in a box: plotting the 1-D $\psi$ and $\psi^2$.
No matter whether you're dealing with a 1-dimensional or 3-dimensional wave equation, the wavefunction itself ($\psi$) describes the particle's wave properties. This $\psi$ doesn't have actual physical meaning, so it's hard to imagine what it "looks like" other than just plotting its function. However, the probability of finding the particle-wave at any specific position along the x-axis between $x=0$ and $x=a$ is more physically meaningful. The probability of finding the particle is proportional to the square of the wave function, which is represented by either $\psi^2$ or, sometimes, $\psi\psi^*$. The plots of the functions for $\psi$ and $\psi^2$ for the first three possible values of $n$ are shown below in Figure $2$. You could create plots similar to these simply by plotting the function shown in Equation \ref{wf4}. To plot, you just need to assign a value of $n$, and it is convenient to assign the length of the box as $a=1$. The plot generated would have the general ratio of $\frac{x}{a}$ on the x axis, and thus would be relevant for any length box.
The graphs above represent "solutions" to the wave function. For example, the solution $n=1$ gives the plot shown in Figure $2$ B. The solution of $n=2$ is plotted in Figure $2$ C, and so on. These values for $n$ are some of the possible solutions to the 1-dimensional $\psi$, and they yield descriptions of the wave behavior and probability of finding a particle in 1-dimensional space.
How does this apply to atoms?
The 1-dimensional particle in a box does not represent a real situation; but rather, it is a simple model that we can use to understand a more complex system, like a electron orbiting the nucleus in three dimensions. It's useful to recognize the analogies here that might represent something familiar in a real situation.
$x=0$ is analogous to the nucleus in an atom. Here in the particle in a box model, $\psi=0$ and there is zero probability of finding an electron. When this is extended to an atom, this position is analogous to the nucleus (at the origin of a coordinate system where $x=y=z=0$. A 3-dimensional $\psi$ is also zero at the nucleus and the electron cannot exist there; also, $\psi$ approaches zero as it approaches this position in both the particle in a box model and in a more realistic case of an electron in an atom.
$x=a$ is analogous to a boundary surface far from the nucleus. In the particle in a box model, the 1-dimensional $\psi$ is zero at $x=a$. This is analogous to an electron's $\psi$ approaching zero as it gets further from the nucleus. There will be a more in depth description of boundary surface in the next section (2.2.2). You already know this as the outermost surface of an electron orbital. One minor but notable difference between the 1-dimensional particle in a box and the case of a real atom is that the $\psi$ in three dimensions approaches but never quite reaches zero as distance from the nucleus increases, while $\psi=0$ for $x=a$ in the more simple case of a 1-dimensional particle in a box.
A change in sign of the $\psi$ (and where both $psi=0$ and $\psi^2=0$) is a node. In both the 1-dimensional particle in a box model and in the more realistic 3-dimensional case, a node is found where the $\psi$ changes sign. At this point, $\psi=\psi^2=0$; in other words there is zero probability of finding the particle or electron at these points. It is easy to spot these points in the plots above because the wave function crosses the x-axis and the $\psi^2$ meets zero.
Exercise $4$
Identify the points on plots B, C, and D in Figure $2$ that are nodes. How many nodes are there for $n=1, n=2,$ and $n=3$?
Answer
The nodes are annotated with red circles on the figure below. Panel B ($n=1$) has zero nodes. There are two points where $\psi=0$ in the case of $n=1$, but these points are at the walls of the box and they are not nodes. The walls of the box are analogous to the nucleus and the boundary surface of an electron orbital. Panel C ($n=2$) has one node where the $\psi$ crosses zero. Panel D ($n=3$) shows two nodes.
Exercise $5$
Use a graphing program to plot the 1-dimensional $\psi$ and $\psi^2$ for $n=1,2,3,4$. How many nodes are there for $n=4$? Is this expected? Predict how many nodes we should expect for $n=5$.
You can use any plotting program to do this, and if you aren't familiar with any, try this one: Desmos
Answer
You can use any plotting program to do this. This is the function you should plot (it is Equation \ref{wf4}):
$\psi=\sqrt{\frac{2}{a}}\left(\sin\left(\frac{n\pi{x}}{a}\right)\right)$
For n=1: Assign a value of $n=1$, as stated in the problem. In the text above, it also states that it is convenient to assign a value of $a=1$. Assigning these values results in the following function:
$\psi=\sqrt{\frac{2}{1}}\left(\sin\left(\frac{1\pi{x}}{1}\right)\right)$ and simplifying leads to $\psi=\sqrt{2}\left(\sin\left(\pi{x}\right)\right)$
For n=2,3,4 you would repeat the process above. You will get the following functions:
For n=2: $\psi=\sqrt{2}\left(\sin\left(2\pi{x}\right)\right)$
For n=3: $\psi=\sqrt{2}\left(\sin\left(3\pi{x}\right)\right)$
For n=4: $\psi=\sqrt{2}\left(\sin\left(4\pi{x}\right)\right)$
In a program like Desmos, you need to input the function correctly. It's useful to know the code for at least one graphing program. For Desmos, and most others, you can find help online. For example, the correct input for Desmos for the n=4 case is sqrt(2)*(sin(4*pi*x)). But, you have to type it in because copy/paste doesn't work. You'll also want to display your graph only from $x=0$ to $x=1$ by hitting the settings button (a little wrench in the upper right corner) and changing the x scale. Go ahead and change the y scale too, to $y=-2$ to $y=2$. If you do this, you should get something that looks like this:
To graph the square of each wavefunction, you'd just square the functions that we just plotted:
In both the case of $\psi$ and $\psi^2$, the $n=4$ function shows three nodes (see the bolded purple line, where nodes are tiny grey spots on the axis). This follows a pattern that you might have noticed with the cases of $n=1,2,$ and $3$ where the number of nodes is $n-1$ (nodes = $n-1$). We should expect three nodes for $n=4$ based on this pattern, and four nodes for $n=5$.
In the next section, we will extend these ideas qualitatively to three dimensions. While there is only one quantum number in one dimension, there are three quantum numbers in three dimensions that (when combined) give discrete descriptions of an electron in three-dimensional space (plus a fourth quantum number that explains other properties of the electron). | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/02%3A_Atomic_Structure/2.02%3A_The_Schrodinger_equation_particle_in_a_box_and_atomic_wavefunctions/2.2.01%3A_Particle_in_a_Box.txt |
The one-dimensional particle in a box model from the previous section shows us how a wavefunction works in one dimension (the x- dimension). In one dimension, the wavefunction requires only one quantum number, $n$. A full explanation of the three-dimensional wavefunction for an electron is outside of the scope of this course. Instead, we will focus on gaining a conceptual understanding by extending knowledge of the one-dimensional wavefunction to a three-dimensional space.
Quantum numbers
Extending the wavefunction to three dimensions requires a total of three quantum numbers. In addition to $n$, which we saw in the one-dimensional case, we also need $l$ and $m_l$ to express the wavefunction in three dimensions. A complete solution to the Schrödinger equation, both the three-dimensional wavefunction and energy, includes a set of three quantum numbers ($n, l, m_l$). The wavefunction describes what we know as an atomic orbital; it defines the region in space where the electron is located. Additionally, there is a fourth quantum number, $m_s$. The $m_s$ quantum number accounts for the observed interaction of electrons with an applied magnetic field; it is an additional postulate that is not part of the wavefunction. These four quantum numbers are described below.
The quantum number, $n$: This is the principle quantum number. This number represents the shell, including both the overall energy of the electron in that shell and the size of that shell. An allowed value for $n$ is any non-zero, positive integer (1, 2, 3, 4, ... etc are allowed, but 4.1 is not allowed).
The quantum number, $l$: This is the angular momentum quantum number that corresponds to the subshell and its shape. It represents the angular dependence of the subshell, or the "shape" of the orbitals within a subshell. The allowed values of $l$ depend on $n$. The allowed values of $l$ for an electron in shell $n$ are integer values between $0$ to $n-1$, or $l = 0\rightarrow n-1$. These values correspond to the orbital shape where $l=0$ is an s-orbital, $l=1$ is a p-orbital, $l=2$ is a d-orbital, and $l=3$ is an f-orbital.
The quantum number $m_l$: This is the magnetic quantum number. Its possible values give the number of orbitals within a subshell and its specific value gives the orbital's orientation in space. The allowed values of $m_l$ depend on the value of $l$. The value of $m_l$ is allowed to be any positive or negative integer between $+l$ and $-l$. In other terms, $m_l=+l \rightarrow -l$. For example, if the electron is in a 3-p-orbital, then $n=3, l=1$, and the possible values of $m_l$ are $-1, 0,$ and $+1$. Since there are three possible values of $m_l$ there are three orbitals in the $p$ subshell. The specific $m_l$ value defines in which of the three possible p-orbitals ($p_x, p_y,$ or $p_z$) the electron exists. In the case of the $s$ subshell, there is only one value, $m_l=0$ because $l=0$. The one value corresponds to the fact that there is only one $s$ orbital in any shell.
The quantum number $m_s$: This quantum number accounts for the electron's "spin". In short, electrons interact with magnetic fields in a way that is similar to how a tiny bar magnet would interact with a magnetic field. The allowed values for $m_s$ are $+\frac{1}{2}$ and $-\frac{1}{2}$.
Table $1$. Quantum Numbers
SYMBOL NAME VALUES MEANING
$n$ principal $1,2,3...$(any integer) energy level, shell
$l$ angular momentum $0 \rightarrow n-1$
subshell, $0=s, 1=p, 2=d, 3=f...$
this is the angular dependence of the orbital, shape of the orbital
*letters have historical meaning, sharp, principle, diffuse, fundamental
$m_l$ magnetic $+l \rightarrow -l$ orientation of angular momentum in space, orbital
$m_s$ spin $+\frac{1}{2}, -\frac{1}{2}$ the imaginary property we call "spin", up or down
1-Electron Wavefunctions (Atomic orbitals)
One simplified representation of the three-dimensional wavefunction is shown below. This representation breaks the wavefunction into two parts: the radial contribution ($\textcolor{blue}{R_{n,l}(r)}$) and the angular contribution ($\textcolor{red}{Y_{l,m_l}(\theta,\phi)}$).
$\psi_{(n,l,m_l)}= \textcolor{blue}{R_{n,l}(r)} \times \textcolor{red}{Y_{l,m_l}(\theta,\phi)} \label{WF1}$
Allowed Energies
From the wavefunction, we can find the allowed energies of an electron in an atom. We have not given you the more complex form of the wavefunction where you'd need to derive \ref{E} from \ref{WF1} because it is outside the scope of this course. Instead, consider the process that we used to derive the energy of a particle in a one-dimensional box in the previous section. A similar process can be used to find the energy of an electron in three dimensions, shown in \ref{E}.
$E_n = -\frac{hcRZ^2}{n^2}=-(13.607\thinspace eV) \left(\frac{Z}{n}\right)^2 \label{E}$,
where $n$ is the principle quantum number, $h$ is Planck's constant, $c$ is the speed of light, $R$ is the Rydberg constant, and $Z$ is the charge of the nucleus. It is useful to know that the value of $hcR = 13.6 eV$. This is also the value of the ionization energy of an electron in a hydrogen atom. This equation only works for hydrogenic atoms (atoms or ions that are "like" hydrogen in that they have only one electron).
Exercise $1$
The energy that it takes to eject a ground-state electron from a hydrogen atom (its ionization energy, IE) is measured to be approximately 13.6 eV, while the IE energy of a He+ ion is four times greater at 54.4 eV. These values can be predicted using both Equation \ref{E} and the Rydberg equation.
1. Using Equation \ref{E} for allowed energies of electrons in a hydrogen atom, derive the energy of the ground state electron in H. Then repeat this process for a He+ ion. What do the values indicate about the relative attraction between an electron and the nuclei of these two hydrogenic atoms?
2. Show that the IE of H's 1s electron can also be predicted by the Rydberg formula (equation 2.1.2.2).
Answer (a)
Use Equation \ref{E} . For a hydrogen atom with one proton, $Z=1$, and the ground state energy level is the lowest energy level, $n=1$. Therefore, the energy of an electron in the ground state of a H atom is:
$E_n = -\frac{hcRZ^2}{n^2} = -13.607\thinspace eV(\frac{1}{1})^2 = -13.607 \thinspace eV$, or approximately $-13.6 \thinspace eV$. Thus, the IE (energy necessary to remove that electron) is +13.6 eV.
We can use the same process to find the ground state energy of an electron in He+. This time, $Z=2$ because helium has two protons in its nucleus, and as before, the ground state is $n=1$.
$E_n = -\frac{hcRZ^2}{n^2} = -13.607\thinspace eV(\frac{2}{1})^2 = -54.428 eV$, or approximately $-54.4 \thinspace eV$. Thus, the IE is +54.4 eV.
These ground-state energies are the same as the energy required to remove the ground state electron from H and He+. In other words, these are the predicted ground state IE values for one electron in H and He+. The measured IEs for these two species are in fact 13.6 eV and 54.4 eV, respectively.
The He+ ion has a more negative value for its ground state energy with a magnitude four times greater than that of the ground state H electron. This means that the He+ electron is more strongly attracted to the (+2) nucleus than the H electron is attracted to the (+1) nucleus, and more energy is required to eject the He+ electron than the H electron. It takes more energy to remove the electron from He+ than from H. This makes sense in light of Coulomb's law: The attractive force between the H nucleus (+1) and an electron (-1) would be weaker than that of a He nucleus (+2) and an electron (-1) because the positive charge of the He nucleus has a greater magnitude.
Answer (b)
The Rydberg equation is: $\dfrac{1}{\lambda }= E (cm^{-1})= R_H\; \left ( \dfrac{1}{n_l^{2}}-\dfrac{1}{n_h^{2}} \right )$ where $n_l$ and $n_h$ are positive integers, $n_h > n_l$, and the Rydberg constant ($R_H$), has a value of 1.09737 × 107m−1. We should set $n_l=1$ to represent the initial ground state, and $n_h=\infty$ to represent the removal of the electron from the atom.
Step 1.
Calculate Energy (in m-1) using the Rydberg equation:
$E = R_H\left( \dfrac{1}{n_l^{2}}-\dfrac{1}{n_h^{2}} \right) = 1.09737\times10^7 m^{-1} \left( \dfrac{1}{1^{2}}-\dfrac{1}{\infty^{2}} \right ) = 1.09737\times10^7 m^{-1} \left( 1-0 \right ) = 1.09737 \times 10^7 m^{-1}$
Ultimately, we want a value of energy in units of eV so we can compare it to the answers in part (a), but our value here is in units of inverse meters (m-1). A quick internet search can tell you the conversion between inverse centimeters (cm-1) and eV is $1 eV = 8065.6 cm^{-1}$.
Step 2.
Convert from m-1 to cm-1 to eV so that we can compare to the answer in part (a):
$1.09737\times 10^7 \frac{1}{m} \times \dfrac{1m}{100cm} \times \dfrac{1eV}{8065.6 \frac{1}{cm}} = 13.6056 eV,$ or approximately 13.6 eV.
Exercise $2$
Which of the following are hydrogenic atoms? H, H+, H-, He, He2+, He+, He-, Li, Li3+, Li2+, Li+, Li- .
Answer
Hydrogenic atoms are atoms or ions that contain only one electron. H, He+, and Li2+ are hydrogenic atoms/ions that contain only one electron.
Radial and Angular contributions to the wavefunction:
Radial contribution, $\textcolor{blue}{R_{n,l}(r)}$
The radial part of the wavefunction, $\textcolor{blue}{R_{n,l}(r)}$ gives the radial variation of $\psi$. In other words, $\textcolor{blue}{R_{n,l}(r)}$ defines how the wavefunction depends on the distance of the electron from the nucleus (the radius). The $\textcolor{blue}{R_{n,l}(r)}$ parts of the wavefunction for a hydrogenic atom are listed in Table $2$, and they are plotted in Figure $1$. Notice that the $R_{n,l}(r)$ of all s-orbitals (solid lines) reaches a maximum at $r=0$. This is unique to the s-orbitals' $R_{n,l}(r)$. The $R_{n,l}(r)$ of p- and d-orbitals orbitals approaches zero as r approaches zero. This has important consequences for how closely an electron in these orbitals can approach the nucleus.
Table $2$. Radial wavefunctions ($R(r)$) for the first three shells of a hydrogen atom. $Z$ is the nuclear charge, and $a_0 = 52.9 pm = 0.529 Å$ is the Bohr radius (the radius of a hydrogen 1s orbital). For the H atom, $Z=1$ (the nuclear charge of hydrogen). Plotting of the functions can be simplified if we set $a_0=1$ and plot the functions with respect to $r/a_0$, which was done to find the plots shown in Figure $1$.
Orbital $n$ $l$ $\textcolor{blue}{R_{n,l}(r)}$
1s 1 0 $2\left[\frac{Z}{a_0}\right]^{3/2}e^{-Zr/a_0}$
2s 2 0 $2\left[\frac{Z}{2a_0}\right]^{3/2}\left(2-\frac{Zr}{a_0}\right)e^{-Zr/2a_0}$
2p 2 1 $\frac{1}{\sqrt{3}}\left[\frac{Z}{3a_0}\right]^{3/2}\left(\frac{Zr}{a_0}\right)e^{-Zr/2a_0}$
3s 3 0 $\frac{2}{27}\left[\frac{Z}{3a_0}\right]^{3/2}\left(27-18\frac{Zr}{a_0}-2(\frac{Zr}{a_0})^2\right)e^{-Zr/3a_0}$
3p 3 1 $\frac{1}{81\sqrt{3}}\left[\frac{2Z}{a_0}\right]^{3/2}\left(6-\frac{Zr}{a_0}\right)(\frac{Zr}{a_0})e^{-Zr/3a_0}$
3d 3 2 $\frac{1}{81\sqrt{15}}\left[\frac{2Z}{a_0}\right]^{3/2}\left(\frac{Zr}{a_0}\right)^2e^{-Zr/3a_0}$
Radial probability function, $\textcolor{black}{4 \pi r^2}\textcolor{blue}{(R_{n,l}(r))^2}$
Just like in the particle-in-a-box model, the square of the wavefunction is proportional to the probability of finding a particle (electron) at some point in space. The square of the radial part of the wavefunction is called the radial distribution function $4 \pi r^2\textcolor{blue}{(R_{n,l}(r))^2}$, and it describes the probability of locating the electron at some distance $r$ away from the nucleus. When we normalize the probability functions by dividing the function by its integral over all space, we get the plots shown in Figure $2$. The normalized probability functions are compared to the original radial part of the wavefunctions in Figure $3$. The most probable distance for finding an electron is shown by the maximum value of the function. For an electron in the 1s orbital of H, the most probable distance from the nucleus occurs at $r=1a_0$. This is the Bohr Radius, and it has a value of $a_0 = 52.9 pm = 0.529 Å$. It is convenient to plot the functions of the hydrogen atomic orbitals relative to the size of its smallest orbital, the 1s orbital; this is the reason we plot $R_{n,l}(r)$ and $4 \pi r^2(R_{n,l}(r))^2$ relative to $\frac{r}{a_0}$.
Radial nodes
From the discussion of the 1-dimensional particle in a box, we learned that nodes exist where $\psi=0$. In the case of the 3-dimensional wavefunction, there are two different types of nodes: radial nodes and angular nodes. Radial nodes occur where the radial part of the wavefunction is zero ($R(x)=0$). These are easy to find by plotting the radial part of the wavefunction and finding where the radial part of the wavefunction (and the radial probability function) is zero (where $\textcolor{blue}{R_{n,l}(r)}=0$) and where $4 \pi r^2\textcolor{blue}{(R_{n,l}(r))^2}=0$. These nodes are spherical in shape and depend on the energy level and subshell (the values of $n$ and $l$). The number of radial nodes is $n-l-1$. An example of a radial node is the single node that occurs in the $2s$ orbital ($2-0-1=1$ node). In contrast, the 1s orbital has zero radial nodes ($1-0-1=0$ nodes). Where there is a node, there is zero probability of finding an electron.
Boundary surfaces
A true node occurs where the probability of finding an electron is zero. Nodes occur where $\psi=\psi^2=0$. Far from the nucleus, the probability of finding the electron rapidly approaches zero, but is never exactly zero. What this means is that it is rather improbable to find the electron at far distances from the nucleus, but it's not impossible. The rapid fall of probability creates a boundary surface rather than a node. The boundary surface represents the area around the nucleus where the electron exists most of the time.
Exercise $1$
How many radial nodes are in the s, p, d, and f orbitals in the first four shells? $n=1,2,3,4$?
Answer
You can approach this answer by using the mathematical relationship between the number of radial nodes and the values of $n$ and $l$: number of radial nodes is $n-1-l$. You could also notice the pattern that the first orbital of any type (1s, 2p, 3d...) has zero radial nodes, the second orbital of a type (2s, 3p, 4d...) has one radial node, the third orbital of a type has two radial nodes (3s, 4p, 5d...)...etc.
A full list of all of the orbitals in the first four shells and their number of radial nodes is below. It's a good idea to prove to yourself that the numbers given below are consistent with the plots of the orbitals' wavefunctions for those orbitals in the first three shells:
s-orbitals:
1s: $n-1-l = 1-1-0 = 0$, zero radial nodes
2s: $n-1-l = 2-1-0 = 1$, one radial node
3s: $n-1-l = 3-1-0 = 2$, two radial nodes
4s: $n-1-l = 4-1-0 = 3$, three radial nodes
p-orbitals
2p: $n-1-l = 2-1-1 = 0$, zero radial nodes
3p: $n-1-l = 3-1-1 = 1$, one radial node
4p: $n-1-l = 4-1-1 = 2$, two radial nodes
d-orbitals
3d: $n-1-l = 3-1-2 = 0$, zero radial nodes
4d: $n-1-l = 4-1-2 = 1$, one radial nodes
f-orbital
4f: $n-1-l = 4-1-3 = 0$, zero radial nodes
Exercise $1$
Draw a rough plot of the following:
1. The radial wavefunction, $R_{n,l}(r)$ and the probability density function, $4 \pi r^2(R_{n,l}(r))^2$, for the 1s orbital.
2. The radial wavefunction, $R_{n,l}(r)$ and the probability density function, $4 \pi r^2(R_{n,l}(r))^2$, for the 4s orbital.
Answer, 1s
The 1s orbital's $R_{n,l}(r)$ reaches a maximum near the origin and approaches zero far from the nucleus. The $4 \pi r^2(R_{n,l}(r))^2$ approaches zero at the origin, has a maximum intensity at the Bohr radius, and approaches zero far from the nucleus (its boundary surface).
.
Answer, 4s
Like all s orbitals, the 4s orbital's $R_{n,l}(r)$ reaches a maximum near X=0, but it would be less intense than the function of the 1s, 2s, and 3s orbitals. It should have three radial nodes and so its wavefunction would change sign three times. This would give four regions of electron density in the probability function while the function would approach zero (boundary surface) far from the nucleus.
Angular contribution, $\textcolor{red}{Y_{l,m_l}(\theta,\phi)}$, and angular probability function, $\textcolor{red}{(Y_{l,m_l}(\theta,\phi))^2}$
The angular contribution to the wavefunction, $\textcolor{red}{Y_{l,m_l}(\theta,\phi)}$, describes the wavefunction's shape, or the angle with respect to a coordinate system. To describe the direction in space, we use spherical coordinates that tell us distance and orientation in 3-dimensional space. There are three spherical coordinates: $r, \phi,$ and $\theta$. $r$ is the radius, or the actual distance from the origin. $\phi$ and $\theta$ are angles. $\phi$ is measured from the positive x axis in the xy plane and may be between 0 and $2\pi$. $\theta$ is measured from the positive z axis towards the xy plane and may be between 0 and $\pi$.
$\textcolor{red}{Y_{l,m_l}(\theta,\phi)}=\left ( \dfrac{1}{4\pi} \right )^{1/2}y\left (\theta,\phi \right ) \nonumber$
$\textcolor{red}{Y_{l,m_l}(\theta,\phi)}$, is slightly more difficult to describe than the radial contribution was. This is partly because $\textcolor{red}{Y_{l,m_l}(\theta,\phi)}$ contains imaginary numbers, which have no real, physical meaning. However, the angular part of the wavefunction becomes more "real" when you square it to get angular probability density, a more tangible concept described as the shapes of orbitals.
The values of $\theta$, $\phi$, and $y(\theta,\phi)$ for orbitals in the hydrogen atom are listed in Table $3$. But $Y_{l,m_l}(\theta,\phi)$ is only a mathematical function and has no real physical meaning. The square of the radial wavefunction, $Y_{l,m_l}(\theta,\phi)^2$, gives the probability of finding the electron at a point in space on a ray described by $(\phi, \theta)$. Thus $Y_{l,m_l}(\theta,\phi)^2$ describes the shape of the orbital.
Table $3$. Components of the angular wavefunction and the resulting orbital shapes. The plots in the two right hand columns were created using Desmos and Mathematica: specifically "Spherical Harmonics" from the Wolfram Demonstrations Project, http://demonstrations.wolfram.com/SphericalHarmonics/ and "Visualizing Atomic Orbitals" demonstrations.wolfram.com/Vi...tomicOrbitals/.
subshell $m_l$ (orbital) $\theta$ $\phi$ Plots of $\theta^2$ orbital shapes
$l=0$, s $m_l=0$ $\frac{1}{\sqrt{2}}$ $\frac{1}{\sqrt{2 \pi}}$
$l=1$, p $m_l=0$ $\frac{\sqrt{6}}{2} \cos \theta$ $\frac{1}{\sqrt{2 \pi}}$
$l=1$, p $m_l=+1$ $\frac{\sqrt{3}}{2} \sin \theta$ $\frac{1}{\sqrt{2 \pi}} e^{i \phi}$
$l=1$, p $m_l=-1$ $\frac{\sqrt{3}}{2} \sin \theta$ $\frac{1}{\sqrt{2 \pi}} e^{-i \phi}$
$l=2$, d 0 $\frac{\sqrt{10}}{4}\left(3 \cos ^{2} \theta-1\right)$ $\frac{1}{\sqrt{2 \pi}}$
$l=2$, d $m_l=+1$ $\frac{\sqrt{15}}{2} \sin \theta \cos \theta$ $\frac{1}{\sqrt{2 \pi}} e^{i \phi}$
$l=2$, d $m_l=-1$ $\frac{\sqrt{15}}{2} \sin \theta \cos \theta$ $\frac{1}{\sqrt{2 \pi}} e^{-i \phi}$
$l=2$, d $m_l=+2$ $\frac{\sqrt{15}}{4} \sin ^{2} \theta$ $\frac{1}{\sqrt{2 \pi}} e^{i 2 \phi}$
$l=2$, d $m_l=-2$ $\frac{\sqrt{15}}{4} \sin ^{2} \theta$ $\frac{1}{\sqrt{2 \pi}} e^{-i 2 \phi}$
Angular nodes
Angular nodes exist where $\textcolor{red}{(Y_{l,m_l}(\theta,\phi))^2}=0$. These nodes are planar in shape, and they depend on the value of $l$. The number of angular nodes in any orbital is equal to $l$. This means that s-orbitals ($l=0$) have zero angular nodes, p-orbitals ($l=1$) have one angular node, d-orbitals ($l=2$) have two angular nodes, and so on. Planar nodes can be flat planes (like the nodes in all p orbitals) or they can have a conical shape, like the two angular nodes in the $d_{Z^2}$ orbital. Angular nodes in some p and d orbitals are shown in Figure $4$.
Atomic Orbitals
Atomic orbitals result from a combination of both the radial and angular contributions of the wavefunction. Atomic orbitals can have both angular nodes and radial nodes, depending on the values of $n$ and $l$.
The chart below compares the radial variation, angular variation, and their combinations (orbitals).
Orbital Radial Probability Radial Nodes Angular Probability Angular Nodes
Combination
(Orbital)
Total Nodes
1s 0 0 0
2s 1 0 1
2p 0 1 1
3s 2 0 2
3p 1 1 2
3d 0 2 2 | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/02%3A_Atomic_Structure/2.02%3A_The_Schrodinger_equation_particle_in_a_box_and_atomic_wavefunctions/2.2.02%3A_Quantum_Numbers_and_Atomic_Wave_Functions.txt |
Introduction
The Aufbau Principle (also called the building-up principle or the Aufbau rule) states that, in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy level before occupying higher-energy levels. In general, an electron will occupy an atomic orbital with the lowest value of $n, l, m_l$, in that order of priority. The value of $m_s$ for an unpaired electron is conventionally assigned a value of $+\frac{1}{2}$. Each electron in an atom/ion must have a unique set of values for all four quantum numbers.
The Ground State of Hydrogenic Atoms/Ions
In a hydrogenic atom/ion, there is only one electron. In this case, the only factor determining energy is the value of $n$. The ground state will always be the 1s orbital ($n=1, l=0, m_l=0, m_s=+\frac{1}{2}$).
The Ground State of Multi-electron Atoms/Ions
In atoms/ions with two or more electrons, the ground state electron configuration (1) minimizes the total energy of the electrons, (2) obeys the Pauli exclusion principle, (3) obeys Hunds rule of maximum multiplicity, and (4) considers the exchange interaction. These "rules" are described below.
(1) Electrons will occupy the lowest energy orbitals in order to minimize the total energy. The two quantum numbers that are related to energy in multi-electron atoms are $n$, and $l$. Thus, orbitals with the lowest values of $n$ and $l$ will fill first.
(2) Hund's rule of maximum multiplicity states that for a given electron configuration, the lowest energy arrangement of electrons in degenerate orbitals is the one with the greatest "multiplicity," where multiplicity is the number of unpaired electrons (n) plus 1 (multiplicity = n + 1). This rule is used to predict the ground state of an atom or molecule with one or more open electronic shells.
Hund's rule is based on empirical observation of atomic spectra, and it is a consequence of the energy required to pair two electrons in the same orbital. This energy of repulsion between two electrons in the same orbitals is a Coulombic energy of repulsion, $\Pi_c$, caused by two electrons with like charge sharing the same area of space (an orbital). When more than one electron occupies a set of degenerate orbitals, the most favorable arrangement is one where the number of paired electrons is minimized.
A simplified definition of Hund's rule is that the lowest energy arrangement is the one with the greatest number of unpaired electrons. This implies that if two or more orbitals of equal energy are available, electrons will occupy them singly before filling them in pairs. Examples of ground state arrangements of electrons in three degenerate p-orbitals is given in the figure shown here.
Figure $2$. Hund's rule is that spin multiplicity must be maximized in the ground state. (CC-BY-NC-SA; Libretexts)
(3) The Pauli exclusion principle states that it is impossible for two electrons of a multi-electron atom to have the same set of values for all four quantum numbers. Two electrons in different orbitals will have a different set of $n, l$, and $m_l$ values. When two electrons reside in the same orbital, they posses the same $n, l$, and $m_l$ values, therefore their ms must be different. Thus, two electrons in the same orbital must have opposite half-integer spin projections of $+\frac{1}{2}$ and $-\frac{1}{2}$. Figure $3$. Pauli's principle tells us that paired electrons must have opposite spin. CC-BY-NC-SA; Libretexts)
(4) The exchange interaction is sometimes called the exchange energy or exchange force. However, it is not a true energy or force. Rather, it is a quantum mechanical effect that takes place between identical particles. The exchange interaction results in a ground state electron configuration with unpaired electrons all being of the same spin. Unpaired electrons are conventionally written in the "spin up" direction. Figure $4$. Unpaired electrons should have identical spins due to the exchange interaction. CC-BY-NC-SA; Libretexts)
Trends in Expected Electron Configuration
The four "rules" above can be used as guidelines to predict the ground state electron configuration of atoms, the filling of subshells, and the configuration of electrons in degenerate orbitals. However, the utility of these guidelines for predicting actual electron configurations requires more nuanced knowledge of the relative energy levels of orbitals.
Generally, orbital energy levels directly correspond to their shell number. Additionally, orbitals within a shell generally follow the energetic trend where s<p<d<f. Although these general trends for relative orbital energy levels hold true for most of the main block elements (the s - and p-blocks), there are important exceptions in the orbital energy levels of transition metal atoms and ions of the d- and f-blocks.
Elements that violate general trends in electron configuration are outlined with a dark line in Figure $5$. All of the exceptions are within the d- and f- blocks, and the violations are caused by an unexpected order of the orbital energy levels.
In the next section you will learn why the orbital energy levels correlate with shell number and why subshells within a shell usually follow the trend that s<p<d<f. You will also learn why there are occasional exceptions to this trend and how these exceptions influence elemental properties.
References
1. Miessler, Gary L., and Donald A. Tarr. Inorganic Chemistry. Upper Saddle River, NJ: Pearson Prentice Hall, 2010. Print.
2. Brown, Ian David. The Chemical Bond in Inorganic Chemistry the Bond Valence Model. Oxford: Oxford UP, 2006. Print. | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/02%3A_Atomic_Structure/2.02%3A_The_Schrodinger_equation_particle_in_a_box_and_atomic_wavefunctions/2.2.03%3A_Aufbau_Principle.txt |
Introduction
Coulomb's Law is from classical physics; it tells us that particles with opposite electrostatic charge are attracted to each other, and the larger the charge on either particle or the closer the distance between them, the stronger the attraction. Coulomb's law explains why atomic size decreases as the charge on the nucleus increases, but it can't explain the nuances and variations in size as we go across the periodic table. Coulomb's Law also explains why electrons in different shells (n), at different distances from the nucleus, have different energies. But on its own, Coulomb's law doesn't quite explain why electron subshells within a shell (like 2s vs. 2p) would have different energies. To explain these things, we need to consider how both electron shielding and penetration result in variations in effective nuclear charge (Z*) that depend on shell and subshell.
Effective Nuclear Charge (Z*)
Coulombs' law works well for predicting the energy of an electron in a hydrogen atom (because H has only one electron). It also works for hydrogen-like atoms: any nucleus with exactly one electron (a He+ ion, for example, has one electron). However, Coulomb's law is insufficient for predicting the energies of electrons in multi-electron atoms and ions.
Electrons within a multi-electron atom interact with the nucleus and with all other electrons. Each electron in a multi-electron atom experiences both attraction to the nucleus and repulsion from interactions with other electrons. The presence of multiple electrons decreases the nuclear attraction to some extent. Each electron in a multi-electron atom experiences a different magnitude of (and attraction to) the nuclear charge depending on what specific shell and subshell the electron occupies. The amount of positive nuclear charge experienced by any individual electron is the effective nuclear charge (Z*).
For example, in lithium (Li), none of the three electrons "feel" the full +3 charge from the nucleus (see Figure $1$). Rather, each electron "feels" a Z* that is less than the actual Z and that depends on the electron's orbital. The actual nuclear charge in Li is $Z=+3$; the 1s electrons experience a $Z^* =+2.69$, and the 2s electron experiences a $Z^* = +1.28$. In general, core electrons (or the electrons closest to the nucleus), "feel" a Z* that is close to, but less than, Z. On the other hand, outer valence electrons experience a Z* that is much less than Z.
In summary:
• Core electrons: $Z^* \lessapprox Z$
• Valence electrons: $Z^* \ll Z$
Shielding:
Shielding is the reduction of true nuclear charge (Z) to the effective nuclear charge (Z*) by other electrons in a multi-electron atom or ion. Shielding occurs in all atoms and ions that have more than one electron. H is the only atom in which shielding does not occur.
Explanation of shielding: Electrons in a multi-electron atom interact with the nucleus and all other electrons in the atom. To describe shielding, we can use a simplified model of the atom: we will choose an electron-of-interest in a multi-electron atom and treat all the "other" electrons as a group of spherically-distributed negative charge. Classical electrostatics allows us to treat spherical distribution of charge as a point of charge at the center of the distribution. Thus, we consider all the "other" electrons in our atom as a point of negative charge in the center of the atom. While the positive charge of the nucleus provides an attractive force toward our electron, the negative charge distribution at the center of the atom would provide a repulsive force. The attractive and repulsive forces would partially cancel each other; but since there are less "other" electrons than there are protons in our atom, the nuclear charge is never completely canceled. The "other" electrons partially block, or shield, part of the nuclear charge so that our electron-of-interest experiences a partially-reduced nuclear charge, the Z*.
In reality, there is not a one-for-one "canceling" of the nuclear charge by each electron. Partly due to penetration, no single electron can completely shield a full unit of positive charge. Core electrons shield valence electrons, but valence electrons have little effect on the Z* of core electrons. The ability to shield, and be shielded by, other electrons strongly depends the electron orbital's average distance from the nucleus and its penetration; thus shielding depends on both shell ($n$) and subshell ($l$).
Shielding depends on electron penetration
Coulomb's law shows us that distance of an electron from its nucleus is important in determining the electron's energy (its attraction to the nucleus). The shell number ($n$) determines approximately how far an electron is from the nucleus on average. Thus, all orbitals in the same shell (s,p,d) have similar sizes and similar average distance of their electrons from the nucleus. But there is another distance-related factor that plays a critical role in determining orbital energy levels: penetration. Penetration describes the ability of an electron in a given subshell to penetrate into other shells and subshells to get close to the nucleus. Penetration is the extent to which an electron can approach the nucleus. Penetration depends on both the shell ($n$) and subshell ($l$).
The penetration of individual orbitals can be visualized using the radial probability functions. For example, Figure $2$ below shows plots of the radial probability function of the 1s, 2s, and 2p orbitals. From these plots, we can see that the 1s orbital is able to approach the closest to the nucleus; thus it is the most penetrating. While the 2s and 2p have most of their probability at a farther distance from the nucleus (compared to 1s), the 2s orbital and the 2p orbital have different extents of penetration. Notice that the 2s orbital is able to penetrate the 1s orbital because of the central 2s lobe. The 2p orbital penetrates somewhat into the 1s, but it cannot approach the nucleus as closely as the 2s orbital can. While the 2s orbital penetrates more than 2p (the 2s orbital can approach closer to the nucleus), the 2p is slightly closer on average than 2s. The order of Z* in 2s and 2p subshells depends on which factor (average distance or penetration) is more important. In the first two rows of the periodic table, penetration is the dominant factor that results in 2s having a lower energy than 2p (see Figure $4$ for values).
An electron orbital's penetration affects its ability to shield other electrons and affects the extent to which it is shielded by other electrons. In general, electron orbitals that have greater penetration experience stronger attraction to the nucleus and less shielding by other electrons; these electrons thus experience a larger Z*. Electrons in orbitals that have greater penetration also shield other electrons to a greater extent.
Within the same shell value (n), the penetrating power of an electron follows this trend in subshells (ml):
s>p>d>f
Exercises
Exercises
1. Compare the 2s and 2p orbitals:
(a) Which is closer to the nucleus on average?
(b) Which is more penetrating?
(c) Which orbital experiences a stronger Z* and is thus lower in energy (consider your experience, but also inspect Figures 1.1.2.3 and 1.1.2.4 from the previous page)? Please explain.
2. Peruse the Hyperphysics page (click) that shows radial probability functions of several orbitals (click around on various orbitals). Compare the 2p and 3s orbitals:
(a) Which is farther from the nucleus on average?
(b) Which is more penetrating?
(c) Which orbital is lower in energy?
3. Which atom, Li, or N, has a stronger valence Z*? Explain why.
4. Explain why 2s and 2p subshells are completely degenerate in a hydrogen atom.
5. Which atom has a smaller radius: Be or F? Explain.
6. Which electrons shield others more effectively: 3p or 3d?
7. Use the clues given in the figures below to label the radial distribution functions shown.
8. Examine the plot below. Notice that the probability density plots for the 3s, 3p, and 3d subshells are highlighted.
(a) For which of these three functions is the highest probability density at the smallest $r$ value (which is closest to the nucleus on average)? Is this the same subshell that penetrates most?
(b) Use this example to describe how penetration and shielding result in a splitting of subshell energy level in multi-electron atoms.
9. Explain why the ground state (most energetically favorable) electron configuration of Be is 1s22s2 rather than alternative configurations like 1s22s12p1 or 1s22p2.
Answer 1
(a) The 2s orbital is closer to the nucleus on average.
(b) The 2s orbital is more penetrating than 2p.
(c) You might "know" that the 2s orbital is lower in energy than 2p because 2s fills first. But a close inspection of Figures 1.1.2.3 and 1.1.2.4 from the previous page indicates that while the 2s and 2p elements are degenerate in Ne (element 10), for elements with atomic number 11 and greater 2p has a higher Z* than 2s! This example illustrates that both average distance and penetration are factors in determining Z*, and the factor that is more important may change as we increase in atomic number.
Answer 2
(a) The 3s orbital reaches farther away from the nucleus and is on average farther from the nucleus than 2p.
(b) The 3s orbital is more penetrating than 2p, even though 3s is farther on average!
(c) The 2p orbital is lower in energy than 3s; this is because 2p is still significantly closer to the nucleus on average and experiences a stronger Z*. (Penetration is not the only consideration!)
Answer 3
A nitrogen atom has a stronger effective nuclear charge (Z*) than lithium due to its greater number of protons; even though N also has more electrons that would shield the nuclear charge, each electron only partially shields each proton. This means that atoms with greater atomic number always have greater Z* for any given electron.
Answer 4
The hydrogen atom has only one electron; thus there is no shielding to consider. When there are not other electrons to shield the nucleus, penetration and shielding are irrelevant, and subshells within a shell are degenerate.
Answer 5
Fluorine has a smaller radius than beryllium because F has a greater valence Z* and therefore pulls the valence electrons closer to the nucleus and provides a smaller atomic radius.
Answer 6
3p shields better than 3d because p orbitals penetrate more than d orbitals within the same shell.
Answer 7
Answer 8
3d is closest on average, but 3s penetrates most. The three subshells of $n=3$ differ in their average distance and in their ability to penetrate; these factors result in differences in the Z* experienced by electrons in each orbital. We would expect 3s to be lowest in energy followed by 3p and then 3d.
Answer 9
This question is asking why the 2s orbital fills in Be before 2p is occupied. This is a multi-electron atom, therefore core electrons shield the 2s and 2p orbitals to different extents. In Be we expect the 2s orbital to fill before 2p because 2s penetrates more and experiences a higher Z*.
Slater's rules for estimating Z*
The Z* can be estimated using a number of different methods; probably the best known and most commonly used method is known as Slater's Rules. Slater developed a set of rules to estimate Z* based on how many other electrons exist in the atom and on the orbital location of the electron-of-interest. These two factors are important determinants in shielding, and they are used to calculate a shielding constant (S) used in Slater's formula:
$Z*=Z-S \nonumber$
where Z is the actual nuclear charge (the atomic number) and Z* is the effective nuclear charge.
In the calculation of S, it is assumed that electrons closer to the nucleus than our electron-of-interest cancel some of the nuclear charge; those farther from the nucleus have no effect. To calculate S, all the relevant orbitals in an atom are written out in order of increasing energy, separating them into in "groups". Each change in shell number is a new group; s and p subshells are in the same group but d and f orbitals are their own group. You write out all the orbitals using parentheses until you get to the group of the electron-of-interest, like this:
(1s)(2s,2p)(3s,3p)(3d)(4s,4p)(4d)(4f)(5s,5p) etc.
**Critical: The orbitals must be written in order of increasing energy!
1. Electrons in the same Group(): Each other electron (not counting the electron-of-interest) in the same group () as the chosen electron, contributes 0.35 to S.
Conceptually, this means electrons in the same group shield each other 35%.
2. Electrons in Groups() to the left:
• If the electron-of-interest is in a d or f subshell, every electron in groups () to the left contributes 1.00 to S.
Conceptually, this means that d and f electrons are shielded 100% by all electrons in the same shell with a smaller value of $l$, as well as all electrons in lower shells ($n$).
• If the electron-of-interest is in an s or p subshell, all electrons in the next lower shell (n - 1) contribute 0.85 to S. And all the electrons in even lower shells contribute 1.00 to S.
Conceptually, this means that s and p electrons are shielded 85% by the electrons one shell lower, and 100% by all electrons in shells n - 2 or lower.
3. 1s electrons: S of a 1s electron is just S=0.3, no matter the element.
A video explaining how to use Slater's Rules
Example $1$: Fluorine, Neon, and Sodium
What is the Z* experienced by the valence electrons in the three isoelectronic species: fluorine anion (F-), neutral neon atom (Ne), and sodium cation (Na+)?
Solution
Each species has 10 electrons, and the number of core electrons is 2 (10 total electrons - 8 valence), but the effective nuclear charge varies because each has a different atomic number (Z). The approximate Z* can be found with Slater's Rules. For all of these species, we would calculate the same sigma value:
Calculating $S$: (1s)(2s,2p), $S = 2(0.85) + 7(0.35) = 1.7 + 2.45 = 4.15$
Fluorine anion: $Z*= 9 - S = 9 - 4.15 = 4.85$
Neon atom: $Z*= 10 - S = 10 - 4.15 = 5.85$
Sodium Cation: $Z*= 11 - S = 11 - 4.15 = 6.85$
So, the sodium cation has the greatest effective nuclear charge.
Exercise $1$
Calculate Z* for a 3d-electron in a zinc (Zn) atom.
Answer
Write out the relevant orbitals: (1s)(2s,2p)(3s,3p)(3d) (4s)
Notice that although 4s is fully occupied, we don't include it because in Zn, 4s is higher in energy than 3d. Therefore, 4s is to the right of the d electrons we are considering. The electron-of-interest is in 3d, so the other nine electrons in 3d each contribute 0.35 to the value of S. The other 18 electrons each contribute 1 to the value of S.
$S=18(1)+9(0.35)=21.15$
$Z*=30-21.15=8.85$
So, although the nuclear charge of Zn is 30, the 3d electrons only experience a $Z* \approx 8.85$!
"Best" values for Z*
Slater's rules are a set of simple rules for predicting $S$ and Z* based on empirical evidence from quantum mechanical calculations. In other words, the Z* calculated from Slater's rules are approximate values. The values considered to be the most accurate are derived from quantum mechanical calculations directly. You can find these values in a nice chart on the Wikipedia article of Effective Nuclear Charge. The chart is recreated here in Figure $3$ for convenience:
Z* modulates attraction
When valence electrons experience less nuclear charge than core electrons, different electrons experience different magnitudes of attraction to the nucleus. A modified form of Coulomb's Law is written below, where $e$ is the charge of an electron, Z* is the effective nuclear charge experienced by that electron, and $r$ is the radius (distance of the electron from the nucleus).
$F_{eff}=k \dfrac{Z*e^2}{r^2} \nonumber$
This formula suggests that if we can estimate Z*, then we can predict the attractive force experienced by, and the energy of, an electron in a multi-electron atom (ex. Li).
The attraction of the nucleus to valence electrons determines the atomic or ionic size, ionization energy, electron affinity, and electronegativity. The stronger the attraction, and the stronger Z*, the closer the electrons are pulled toward the nucleus. This in turn results in a smaller size, higher ionization energy, higher electron affinity, and stronger electronegativity.
General Periodic Trends in Z*
Close inspection of Figure $3$ and analysis of Slater's rules indicate that there are some predictable trends in Z*. The data from Figure $3$ is plotted below in Figure $4$ to provide a visual aid to the discussion below.
Trends in Z* for electrons in a specific shell and subshell
The Z* for electrons in a given shell and subshell generally increases as atomic number increases; this trend holds true going across the periodic table and down the periodic table. Convince yourself that this is true for any subshell by examining Figure $4$. (CC-BY-NC-SA; Kathryn Haas)
Do you notice any exceptions to this general trend?
Inspection of Figure $4$ should confirm for you that the Z* increases as Z increases for electrons in any subshell (like the 1s subshell for example, which is plotted above as a red line with square points). You can see this trend as the positive slope in each series. There is one obvious exception in Period 5 in elements 39 (Y) to 41 (Nb; the Z* of 4s actually decreases across these three elements as atomic number increases. There is also an exception between Y and Zr in the 3d subshell, and between Tc and Ru in the 5s subshell.
For valence electrons:
It is useful to understand trends in valence Z* because the valence Z* determines atomic/ionic properties and chemical reactivity. The trends in the valence Z* are not simple because as atomic number increases, the valence shell and/or subshell also changes. The valence Z* is indicated in Figure $4$ as a black line with open circles.
Down the table: As we go down a column of the periodic table, the valence Z* increases. This is a simple trend because type of subshell is consistent and there is an increase only in shell and in atomic number, Z. This trend is best illustrated by inspection of Figure $3$.
Across the table: the trend depends on shell and subshell, but generally Z* increases across a period.
Periods 1-3 (s and p only): As we go across the table in periods 1-3, the shell stays constant as Z increases and the subshell changes from s to p. In these periods, there is a gradual increase in valence Z* as we move across any of the first three periods.
Periods 4 and 5 (s, p, and d): Now we have some more complex trends because valence subshell and shell are changing as we increase in atomic number. Notice that the valence Z* generally increases going across a period as long as the subshell isn't changing; the exception is within the 4d subshell (elements 39-44 or Y-Ru). In general, going from an $(n)s$ subshell to an $(n-1)d$ subshell, there a relatively large increase in valence Z*. And in going from an $(n-1$d\) subshell to an $(n)p$ subshell, there is a relatively large decrease in Z*.
From one period to another: From Figure $4$, we can see that as we increase Z by one proton, going from one period to the next, there is a relatively large decrease in Z* (from Ne to Na, for example). This is because as Z increases by a small interval, the shell number increases, and so the electrons in the valence shell are much farther from the nucleus and are more shielded by all the electrons in the lower shell numbers.
Exercises
Exercise $2$
1. Compare trends in Z* and atomic size. Explain how and why atomic size depends on Z*.
2. Compare trends in Z* and ionization energy. Explain how and why ionization energy depends on Z*.
Answer
1. On the periodic table, atomic radius generally decreases across the periods (left to right) and increases down the groups. As atomic number increases across the periodic table, nuclear charge (Z) increases and Z* increases. In turn, the atomic radius decreases because the higher nuclear charge (and thus higher Z*) pulls electrons closer to the nucleus. Atomic radius increases down the periodic table because the shell number increases. Despite an increase in Z* going down the periodic table, larger atomic radii result from electrons occupying higher shells.
2. Ionization energies (IE) are inversely related to atomic radius; IE increases across the periods and decreases down the groups. Since the nucleus holds valence electrons more strongly (due to higher Z*) across the periods, IE increases because valence electrons are harder to remove. Down the periodic table, larger atomic radii cause electrons in valence orbitals to be shielded by core electrons. Recall that shielding reduces the nuclear charge available to electrons in higher orbital levels, resulting in a lower Z*. With more shielding and lower Z*, the valence electrons are held less tightly by the nucleus such that ionization energy decreases (i.e., valence electrons are easier to remove). | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/02%3A_Atomic_Structure/2.02%3A_The_Schrodinger_equation_particle_in_a_box_and_atomic_wavefunctions/2.2.04%3A_Shielding.txt |
General periodic trends are specific patterns that are present within the periodic table; these are patterns in properties like electronegativity, ionization energy, electron affinity, atomic radius, melting point, and metallic character. General periodic trends provide chemists with an invaluable tool to quickly predict an element's properties. These trends exist because of the similar atomic structure of the elements within their respective group families or periods, and because of the periodic nature of the elements. Some of the general periodic trends are described in this section.
2.03: Periodic Properties of Atoms
Ionization energy (IE) is the energy required to remove an electron from a neutral atom or cation in its gaseous phase. IE is also known as ionization potential.
$A^{n+}_{(g)} \longrightarrow A^{(n+1)+}_{(g)} + e^- \hspace{1cm} IE = \Delta U \nonumber$
Conceptually, ionization energy is the affinity of an element for its outermost electron (an electron it already has in its valence shell).
1st, 2nd, and 3rd Ionization Energies
The symbol $I_1$ stands for the first ionization energy (energy required to take away an electron from a neutral atom, where $n=0$). The symbol $I_2$ stands for the second ionization energy (energy required to take away an electron from an atom with a +1 charge, $n=2$.)
• First Ionization Energy, $I_1$ (general element, A): $A_{(g)} \rightarrow A^{1+}_{(g)} + e^- \nonumber$
• Second Ionization Energy, $I_2$ (general element, A): $A^{1+}_{(g)} \rightarrow A^{2+}_{(g)} + e^- \nonumber$
• Third Ionization Energy, $I_3$ (general element, A): $A^{2+}_{(g)} \rightarrow A^{3+}_{(g)} + e^- \nonumber$
Each succeeding ionization energy is larger than the preceding energy. This means that $I_1 < I_2 < I_3 < ... < I_n$ will always be true. For example, ionization energy increases as succeeding electrons are taken away from $\ce{Mg}$.
$Mg \,_{(g)} \rightarrow Mg^+\,_{(g)} + e^- \;\;\; I_1= 738\, kJ/mol \nonumber$
$Mg^+ \,_{(g)} \rightarrow Mg^{2+}\,_{(g)} + e^- \;\;\; I_2= 1451\, kJ/mol \nonumber$
Ionization energy is correlated with the strength of attraction between the positively-charged nucleus and the negatively-charged valence electrons. The higher the ionization energy, the stronger the attractive force between nucleus and valence electrons, and the more energy is required to remove a valence electron. The lower the ionization energy, the weaker the attractive force between nucleus and valence electrons, and the less energy required to remove a valence electron.
General periodic trends in electron affinity
In general, ionization energies increase from left to right and decrease down a group; however, there are variations in these trends that would be expected from the effects of penetration and shielding. The trends in first ionization energy are shown in Figure $1$ and are summarized below.
• Across a period: As Z* increases across a period, the ionization energy of the elements generally increases from left to right. However there are breaks or variation in the trends in the following cases:
• IE is especially low when removal of an electron creates a newly empty p subshell (examples include $I_1$ of B, Al, Sc)
• IE energy is especially low where removal of an electron results in a half-filled p or d subshell (examples include $I_1$ of O, S)
• IE increases more gradually across the d- and f-subshells compared to s- and p- subshells. This is because d- and f- electrons are weakly penetrating and experience especially low Z*.
• From one period to the next: There is an especially large decrease in IE with the start of every new period (from He to Li or from Ne to Na for example). This is consistent with the idea that IE is especially low when removal of an electron creates a newly empty s-subshell.
• Noble gases: The noble gases posses very high ionization energies. Note that helium has the highest ionization energy of all the elements.
• Down a group: Although Z* increases going down a group, there is no reliable trend in IE going down any group; in some cases IE increases going down a group, while in other cases IE decreases going down a group.
Plots of the $I_1, I_2,$ and $I_3$ of elements from hydrogen to krypton (first four periods) are shown in Figure $2$. Notice that $I_3>I_2>I_1$. Also notice that trends mentioned above for $I_1$ hold true for subsequent ionizations when electron configurations are considered! | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/02%3A_Atomic_Structure/2.03%3A_Periodic_Properties_of_Atoms/2.3.01%3A_Ionization_energy.txt |
Definitions of Electron Affinity
According to IUPAC, there are two different, but equivalent, definitions of electron affinity (EA).1
Definition: Electron Affinity defined as removal of an electron
Electron affinity can be defined as the energy required when an electron is removed from a gaseous anion. The reaction as shown in equation $\ref{EA1}$ is endothermic (positive $\Delta U$) for elements except noble gases and alkaline earth metals. Under this definition, the more positive the EA value, the higher an atom's affinity for electrons.
$A^{-}_{(g)} \longrightarrow A_{(g)} + e^- \hspace{1cm} EA = \Delta U \label{EA1}$
The reaction shown in Equation $\ref{EA1}$ is similar those that define ionization energy. For this reason, the EA is also described as the zeroth ionization energy.
Definition: Electron Affinity defined as addition of an electron
An alternate and more common definition is the microscopic reverse of Equation $\ref{EA1}$. This more common definition states that electron affinity is the energy released when an electron is added to a gaseous atom, as shown in Equation $\ref{EA2}$. The reaction as shown in equation $\ref{EA2}$ is exothermic (negative $\Delta U$) for elements except noble gases and alkaline earth metals. The more negative this EA value, the higher an atom's affinity for electrons.
$A_{(g)} + e^- \longrightarrow A^{-}_{(g)} \hspace{1cm} EA = \Delta U \label{EA2}$
Conceptually, this second definition is quite similar to the concept of electronegativity; but unlike electronegativity, EA is a well-defined quantitative measurement.
Trends in Electron Affinity
For this discussion, we will use the definition of EA that is consistent with it being a zeroth ionization energy: a more positive (larger) value means that the EA is higher (meaning stronger affinity toward an electron).
• Across a period: Similar to ionization energy, EA generally increases across a row of the periodic table; this observation is consistent with the increase in effective nuclear charge (Z*) from left to right across a period. However, there are variations across a period that are similar to variations in ionization energy and that can be explained by shielding, penetration, and electron configuration.
• Down a group: Like the case of ionization energy trends, EA does not consistently decrease going down a column of the periodic table despite the fact that $Z^*$ increases down a group.
The trend in EA follows a zig-zag pattern similar to the one seen with ionization energies, except that it is displaced by one unit from the trend in $I_1$, two units from $I_2$, and so on. For example, EA peaks at F, while $I_1$ peaks at Ne, $I_2$ peaks at Na, and $I_3$ peaks at Mg. A plot of EA for the first 13 elements is shown overlaid on plots of $I_1, I_2$ and $I_3$ in Figure $1$., where the shifts in the peaks and valleys within each zig-zag trend are indicated.
Sources
1. IUPAC. Compendium of Chemical Terminology, 2nd ed. (the "Gold Book"). Compiled by A. D. McNaught and A. Wilkinson. Blackwell Scientific Publications, Oxford (1997). Online version (2019-) created by S. J. Chalk. ISBN 0-9678550-9-8. doi.org/10.1351/goldbook.
2. Electron Affinity (data page), Wikipedia. en.Wikipedia.org/wiki/Electron_affinity_(data_page) Accessed 12/3/19.
2.3.03: Covalent and Ionic Radii
Measurement of Radius
There are several methods that can be used to determine radii of atoms and ions:
• Nonpolar atomic radii: The radius of an atom is derived from the bond lengths within nonpolar molecules; one-half the distance between the nuclei of two atoms within a covalent bond.
• van der Waals radius: The radius of an atom is determined by collision with other atoms.
• Crystal radii: The atomic or ionic radius is determined using electron density maps from X-ray data.
The measurement of atomic or ionic size will depend on a number of factors, including the covalent character of bonding in any particular molecule, coordination number, physical state (liquid, solid, gas), the identity of nearby atoms/ions, variation in crystal structure, and distortions within regular crystal structures. You should keep in mind that the size of an atom or ion is a "fuzzy" measure, and the radius under a different set of conditions will probably change slightly.
Regardless, measured atomic and ionic radii reveal obvious trends across the periodic table and between atoms and ions. The relative atomic sizes shown in Figure \(1\) were derived from crystallographic data.1
Trends in Atomic Radius
Atomic size generally decreases gradually from left to right across a period of elements. As nuclear charge (Z) increases, we expect the effective nuclear charge (Z*) of the valence electrons to also increase. Increasing Z pulls electrons closer to the nucleus. However, with each additional unit of Z, there is also an additional electron. The change in size is a balance of a compression caused by increasing Z and an expansion in the number of electrons. As a result, the atomic radius decreases gradually across a period.
Atomic size generally increases going down a group. As valence electrons occupy higher level shells due to the increasing quantum number (n), size increases despite the fact that Z and Z* are increasing going down the group.
Trends in Ionic Radius
Trends in ionic radius follow general trends in atomic radius for ions of the same charge. Ionic radius varies with the charge of the ion (and number of electrons) and the electron configuration (e.g. high spin or low spin).
Cations
Compared to their atoms, cations have the same Z but fewer electrons. Removal of electrons from an atom to form a cation results in a significant increase in effective nuclear charge, resulting in all other electrons being more strongly attracted to the nucleus, and having a lower energy level. The result is a contraction in size from the atom to cation. Figure \(2\) visually illustrates the relative size of atoms and some cations of the first four periods; the data is available in tabular format in Figure \(3\).
Anions
Compared to their atoms, anions have the same Z but more electrons. Addition of electrons to an atom to form an anion results in a decrease in effective nuclear charge, which corresponds to a decrease in attractive force between the nucleus and electrons. Lower attractive force leads to expansion, where the size of the atom becomes larger in the formation of an anion. Figure \(2\) visually illustrates the relative size of atoms and some anions of the first four periods, while the data is available in tabular format in Figure \(3\).
2.04: Problems (do we want this here)
Example Exercises
The following series of problems reviews general understanding of the aforementioned material.
Exercises
1. Based on the periodic trends for ionization energy, which element has the highest ionization energy?
2. Which has a larger atomic radius: nitrogen or oxygen?
3. Which element is more electronegative, sulfur (S) or selenium (Se)?
4. Why is the electronegativity value of most noble gases zero?
5. Rewrite the following list in order of decreasing electron affinity: fluorine (F), phosphorous (P), sulfur (S), boron (B).
6. Which of these elements has a smaller atomic radius than sulfur (S): O, Cl, Ca, Li
Answer 1
Helium (He).
Answer 2
Atomic radius increases from right to left on the periodic table. Therefore, nitrogen is larger than oxygen.
Answer 3
Sulfur (S). Note that sulfur and selenium share the same column. Electronegativity increases up a column. This indicates that sulfur is more electronegative than selenium.
Answer 4
Because of their full valence electron shell, the noble gases are extremely stable and do not readily lose or gain electrons.
Answer 5
Fluorine (F)>Sulfur (S)>Phosphorous (P)>Boron (B). Explanation: Electron affinity generally increases from left to right and from bottom to top.
Answer 6
Oxygen (O) is the only element in the list with a smaller atomic radius than S. Periodic trends indicate that atomic radius increases down a group (from top to bottom) and from left to right across a period. | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/02%3A_Atomic_Structure/2.03%3A_Periodic_Properties_of_Atoms/2.3.02%3A_Electron_Affinity.txt |
• 3.1: Lewis Electron-Dot Diagrams
The bonding between atoms in a molecule can be topically modeled though Lewis electron dot diagrams. Creating Lewis diagrams is rather simple and requires only a few steps and some accounting of the valence electrons on each atom. Valence electrons are represented as dots. When two electrons are paired (lone pairs), they are represented by two adjacent dots located on an atom, and when two paired electrons are shared between atoms (bonds), they are shown as lines.
• 3.2: Valence Shell Electron-Pair Repulsion
The Valence Shell Electron Repulsion (VSEPR) model can predict the structure of most molecules and polyatomic ions in which the central atom is a nonmetal; it also works for some structures in which the central atom is a metal. VSEPR builds on Lewis electron dot structures and together can predict the geometry of each atom in a molecule. The main idea of VSEPR theory is that pairs of electrons (in bonds and in lone pairs) repel each other.
• 3.3: Molecular Polarity
Dipole moments occur when there is a separation of charge. They can occur between two ions in an ionic bond or between atoms in a covalent bond; dipole moments arise from differences in electronegativity. The larger the difference in electronegativity, the larger the dipole moment. The distance between the charge separation is also a deciding factor into the size of the dipole moment. The dipole moment is a measure of the polarity of the molecule.
• 3.4: Hydrogen Bonding
A hydrogen bond is an intermolecular force (IMF) that forms a special type of dipole-dipole attraction when a hydrogen atom bonded to a strongly electronegative atom exists in the vicinity of another electronegative atom with a lone pair of electrons. Hydrogen bonds are are generally stronger than ordinary dipole-dipole and dispersion forces, but weaker than true covalent and ionic bonds.
03: Simple Bonding Theory
In 1916, Gilbert Lewis Newton introduced a simple way to show the bonding between atoms in a molecule though Lewis electron dot diagrams. Creating Lewis diagrams is rather simple and requires only a few steps and some accounting of the valence electrons on each atom. Valence electrons are represented as dots. When two electrons are paired (lone pairs), they are represented by two adjacent dots located on an atom, and when two paired electrons are shared between atoms (bonds), they are shown as lines. For example, below are the electron dot structures of atoms and the Lewis electron dot structures of the molecules.
These diagrams are helpful because they allow us to show how atoms are connected, and when coupled with Valence Shell Electron Repulsion Theory (VSEPR), we can use Lewis structures to predict the shape of the molecule.
The drawing of Lewis electron-dot structures is guided largely by the octet rule: that atoms form bonds to achieve eight electrons in their valence shell. For many elements, a full valence shell has an electron configuration of \(s^2p^6\), or eight electrons. A common exception to this rule is the first row elements, H and He. These two elements have \(n=1\) as their valence shell, and so they have only two electrons in a full valence shell (\(1s^2\) electron configuration). Although H and He are exceptions to the "octet rule", they still form bonds to achieve a full valence shell. It may be better to think of this as the "full valence" rule of bonding. We will see many more violations to the "octet" rule as we progress through this course. In the case of metals and metalloids, breaking of the rules is particularly common. (CC-BY-NC-SA; Kathryn Haas)
Common violations of the octet rule
• Less than eight electrons (hypovalency): H and He are examples of elements that cannot have more than two electrons in their full valence shell. Additionally, there are cases where a valid Lewis structure contains atoms with hypovalency: partially-filled valence shells. An example of this is a carbocation, a positively charged carbon with only six electrons in its valence shell. Carbocations are important intermediates in organic chemistry, and they are highly reactive (unstable) Lewis acids (electrophiles). Another example of hypovalency is the Lewis structure for BH3, which shows boron with three bonds, and only six electrons in its valence shell. Boron hydride is a strong Lewis Acid. (Side note: the actual chemical form of BH3 is not well-predicted by the Lewis structure and we'll see more about this in Section 3.1.4.)
• More than eight electrons (hypervalency): This is a case where an element has more than eight electrons in its valence shell. It is common for larger atoms (\(n\geq3\)), and it is discussed further in Section 3.1.2.
Pitfall alert! There are different rules for counting electrons depending on the purpose of the counting. These are the rules for counting for "octets". The rules for counting for calculation of an atom's formal charge are !!different!! and are described in Section 3.1.1. When an atom is part of a molecule, all electrons that are associated with the atom are counted as contributing to the atom's valence. This includes electrons that are lone pairs on the atoms, and all electrons that are shared in bonds. If four electrons are shared between two atoms, it is a double bond. If six electrons are shared between atoms, it is a triple bond.
Electrons in valence (octet) = total unbonded electrons on the atom + total bonded electrons (2 electrons per bond)
Even if you're an old pro at drawing Lewis structures, it's a good idea to polish up. Please complete the practice exercises below. You should get out an actual piece of paper (I know...just do it. It's good for you.) and a writing tool and try to complete each problem before checking the answer.
Exercise \(1\)
Draw the Lewis structures for H2O, CO2, and N2.
Answer
Exercise \(2\)
Draw the Lewis structures for H2, BH3 and BF3.
Answer
These three examples include atoms that have less than eight electrons in their valence shell. In the case of H, it is satisfied with only two electrons in its valence, as was discussed earlier in this section. The case of BH3 was also discussed above. In the Lewis structure of BH3, the boron can only have six electrons in its octet and it is neutral in charge. The boron is electron deficient even though it has neutral charge.
The case of BF3 deserves a discussion: If you are unfamiliar with resonance and formal charge, see Sections 3.1.1 (Resonance) and 3.1.3 (Formal Charge) first and come back to this afterward. You might have drawn the BF3 structure similar to the one that is drawn for BH3, where boron has three bonds and only six electrons in its valence shell. If you did that, then you are correct; but if you only gave this structure, then your answer is not complete. There are three other correct ways to draw the structure. Since F has lone pairs of electrons, other valid Lewis structures would each have a double bond to one of the fluorines (three total). All of these structures are called resonance structures, and based on the four of them, we could predict that BF3 would have B-F bonds that have some double-bond character. If fact, this is the case for BF3; it has bond lengths that are shorter than single bonds, but longer than double bonds. Read more about it on its Wikipedia page here (click).
3.01: Lewis Electron-Dot Diagrams
Learning Objectives
• To understand the concept of resonance.
Resonance structures are a set of two or more Lewis Structures that collectively describe the electronic bonding of a single polyatomic species including fractional bonds and fractional charges. Resonance structures are capable of describing delocalized electrons that cannot be expressed by a single Lewis formula with an integral number of covalent bonds.
Sometimes one Lewis Structure is not Enough
Sometimes, even when formal charges are considered, the bonding in some molecules or ions cannot be described by a single Lewis structure. Resonance is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by a single Lewis formula. A molecule or ion with such delocalized electrons is represented by several contributing structures (also called resonance structures or canonical forms). Such is the case for ozone (\(\ce{O3}\)), an allotrope of oxygen with a V-shaped structure and an O–O–O angle of 117.5°.
Ozone (\(O_3\))
1. We know that ozone has a V-shaped structure, so one O atom is central:
2. Each O atom has 6 valence electrons, for a total of 18 valence electrons.
3. Assigning one bonding pair of electrons to each oxygen–oxygen bond gives
with 14 electrons left over.
4. If we place three lone pairs of electrons on each terminal oxygen, we obtain
and have 2 electrons left over.
5. At this point, both terminal oxygen atoms have octets of electrons. We therefore place the last 2 electrons on the central atom:
6. The central oxygen has only 6 electrons. We must convert one lone pair on a terminal oxygen atom to a bonding pair of electrons—but which one? Depending on which one we choose, we obtain either
Which is correct? In fact, neither is correct. Both predict one O–O single bond and one O=O double bond. As you will learn, if the bonds were of different types (one single and one double, for example), they would have different lengths. It turns out, however, that both O–O bond distances are identical, 127.2 pm, which is shorter than a typical O–O single bond (148 pm) and longer than the O=O double bond in O2 (120.7 pm).
Equivalent Lewis dot structures, such as those of ozone, are called resonance structures. The position of the atoms is the same in the various resonance structures of a compound, but the position of the electrons is different. Double-headed arrows link the different resonance structures of a compound:
The double-headed arrow indicates that the actual electronic structure is an average of those shown, not that the molecule oscillates between the two structures.
When it is possible to write more than one equivalent resonance structure for a molecule or ion, the actual structure is the average of the resonance structures.
The Carbonate (\(CO_3^{2−} \)) Ion
Like ozone, the electronic structure of the carbonate ion cannot be described by a single Lewis electron structure. Unlike O3, though, the actual structure of CO32 is an average of three resonance structures.
1. Because carbon is the least electronegative element, we place it in the central position:
2. Carbon has 4 valence electrons, each oxygen has 6 valence electrons, and there are 2 more for the −2 charge. This gives 4 + (3 × 6) + 2 = 24 valence electrons.
3. Six electrons are used to form three bonding pairs between the oxygen atoms and the carbon:
4. We divide the remaining 18 electrons equally among the three oxygen atoms by placing three lone pairs on each and indicating the −2 charge:
5. No electrons are left for the central atom.
6. At this point, the carbon atom has only 6 valence electrons, so we must take one lone pair from an oxygen and use it to form a carbon–oxygen double bond. In this case, however, there are three possible choices:
As with ozone, none of these structures describes the bonding exactly. Each predicts one carbon–oxygen double bond and two carbon–oxygen single bonds, but experimentally all C–O bond lengths are identical. We can write resonance structures (in this case, three of them) for the carbonate ion:
The actual structure is an average of these three resonance structures.
The Nitrate (\(NO_3^-\)) ion
1. Count up the valence electrons: (1*5) + (3*6) + 1(ion) = 24 electrons
2. Draw the bond connectivities:
3. Add octet electrons to the atoms bonded to the center atom:
4. Place any leftover electrons (24-24 = 0) on the center atom:
5. Does the central atom have an octet?
• NO, it has 6 electrons
• Add a multiple bond (first try a double bond) to see if the central atom can achieve an octet:
6. Does the central atom have an octet?
• YES
• Are there possible resonance structures? YES
Note: We would expect that the bond lengths in the \(\ce{NO_3^{-}}\) ion to be somewhat shorter than a single bond.
Example \(1\): Benzene
Benzene is a common organic solvent that was previously used in gasoline; it is no longer used for this purpose, however, because it is now known to be a carcinogen. The benzene molecule (\(\ce{C6H6}\)) consists of a regular hexagon of carbon atoms, each of which is also bonded to a hydrogen atom. Use resonance structures to describe the bonding in benzene.
Given: molecular formula and molecular geometry
Asked for: resonance structures
Strategy:
1. Draw a structure for benzene illustrating the bonded atoms. Then calculate the number of valence electrons used in this drawing.
2. Subtract this number from the total number of valence electrons in benzene and then locate the remaining electrons such that each atom in the structure reaches an octet.
3. Draw the resonance structures for benzene.
Solution:
A Each hydrogen atom contributes 1 valence electron, and each carbon atom contributes 4 valence electrons, for a total of (6 × 1) + (6 × 4) = 30 valence electrons. If we place a single bonding electron pair between each pair of carbon atoms and between each carbon and a hydrogen atom, we obtain the following:
Each carbon atom in this structure has only 6 electrons and has a formal charge of +1, but we have used only 24 of the 30 valence electrons.
B If the 6 remaining electrons are uniformly distributed pairwise on alternate carbon atoms, we obtain the following:
Three carbon atoms now have an octet configuration and a formal charge of −1, while three carbon atoms have only 6 electrons and a formal charge of +1. We can convert each lone pair to a bonding electron pair, which gives each atom an octet of electrons and a formal charge of 0, by making three C=C double bonds.
C There are, however, two ways to do this:
Each structure has alternating double and single bonds, but experimentation shows that each carbon–carbon bond in benzene is identical, with bond lengths (139.9 pm) intermediate between those typically found for a C–C single bond (154 pm) and a C=C double bond (134 pm). We can describe the bonding in benzene using the two resonance structures, but the actual electronic structure is an average of the two. The existence of multiple resonance structures for aromatic hydrocarbons like benzene is often indicated by drawing either a circle or dashed lines inside the hexagon:
Exercise \(1\): Nitrite Ion
The sodium salt of nitrite is used to relieve muscle spasms. Draw two resonance structures for the nitrite ion (NO2).
Answer
Resonance structures are particularly common in oxoanions of the p-block elements, such as sulfate and phosphate, and in aromatic hydrocarbons, such as benzene and naphthalene.
Warning
If several reasonable resonance forms for a molecule exists, the "actual electronic structure" of the molecule will probably be intermediate between all the forms that you can draw. The classic example is benzene in Example \(1\). One would expect the double bonds to be shorter than the single bonds, but if one overlays the two structures, you see that one structure has a single bond where the other structure has a double bond. The best measurements that we can make of benzene do not show two bond lengths - instead, they show that the bond length is intermediate between the two resonance structures.
Resonance structures is a mechanism that allows us to use all of the possible resonance structures to try to predict what the actual form of the molecule would be. Single bonds, double bonds, triple bonds, +1 charges, -1 charges, these are our limitations in explaining the structures, and the true forms can be in between - a carbon-carbon bond could be mostly single bond with a little bit of double bond character and a partial negative charge, for example.
Summary
Some molecules have two or more chemically equivalent Lewis electron structures, called resonance structures. Resonance is a mental exercise and method within the Valence Bond Theory of bonding that describes the delocalization of electrons within molecules. These structures are written with a double-headed arrow between them, indicating that none of the Lewis structures accurately describes the bonding but that the actual structure is an average of the individual resonance structures. Resonance structures are used when one Lewis structure for a single molecule cannot fully describe the bonding that takes place between neighboring atoms relative to the empirical data for the actual bond lengths between those atoms. The net sum of valid resonance structures is defined as a resonance hybrid, which represents the overall delocalization of electrons within the molecule. A molecule that has several resonance structures is more stable than one with fewer. Some resonance structures are more favorable than others. | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/03%3A_Simple_Bonding_Theory/3.01%3A_Lewis_Electron-Dot_Diagrams/3.1.01%3A_Resonance.txt |
The octet rule applies well to atoms in the second row of the periodic table, where a full valence shell includes eight electrons with an electron configuration of $s^2p^6$. Even elements in the third and fourth row are known to follow this rule sometimes, but not always. In larger atoms, where $n\geq3$ the valence shell contains additional subshells: the $d, f, g...$ subshells. Therefore, atoms with $n\geq3$ can have higher valence shell counts by "expanding" into these additional subshells. When atoms contain more than eight electrons in their valence shell, they are said to be hypervalent. Hypervalency allows atoms with $n\geq3$ to break the octet rule by having more than eight electrons. This also means they can have five or more bonds; something that is nearly unheard of for atoms with $n\leq2$. Complete the exercises below to see examples of molecules containing hypervalent atoms.
Exercise $1$
Draw the Lewis structures for sulfur hexafluoride ($\ce{SF6}$).
Answer
Each fluorine atom has one valence electron and will make one bond each. The sulfur has six valence electrons, and must make six bonds to form a molecule with the six fluorine atoms. The molecular structure has six bonds to sulfur, with twelve valence electrons. The sulfur is hypervalent.
Exercise $1$
Draw the Lewis structure for chlorine trifluoride ($\ce{ClF3}$.
Answer
All atoms are halogens and each has seven valence electrons. Chlorine is capable of hypervalency because it is in the third row of the periodic table; however fluorine cannot have more than eight valence electrons in its valence because it is in the second row. The structure has the three fluorine atoms bonded to a central chlorine atom. The chlorine has a valence of ten electrons due to its three bonds and two lone pairs.
.
Is hypervalency real? Not exactly. Hypervalency is a concept associated with hybrid orbital theory and Lewis theory. It's useful for some simple things, like predicting how atoms are connected and predicting molecular shape. But the idea that the d-orbitals are involved in bonding isn't accurate according to wave mechanics.
For main group molecules, chemists (like Pauling) thought a long time ago that hypervalence is due to expanded s2p6 octets. The consensus is now clear that d orbitals are NOT involved in bonding in molecules like SF6 any more than they are in SF4 and SF2. In all three cases, there is a small and roughly identical participation of d-orbitals in the wavefunctions. This has been established in both MO and VB theory. However, using hybrid orbitals with d-orbital contributions equips us with a language which can pragmatically describe the geometries of highly coordinated substances.
While hybrid orbitals are a powerful tool to describe the geometries and shape of molecules and metal complexes, in "real" molecules their significance may be debated. Often a more realistic molecular orbitals approach is needed. However, from an epistemologically simple point of view, bonding theories can only be judged by their predictions. To the extent that hybridization can explain the shapes of PF5 and SF6, valence bond theory is a perfectly good theory. To the extent that if you write out the valence bond wavefunction using hybridized orbitals and calculate energies and other properties à la Pauling (i.e., ionization energy and electron affinities) and find them to be off from experimental results (by tens of kcals/mol), then valence bond theory is not accurate.
Bonding theories can only be judged by their predictions.
A simple explanation that can be given is that molecular wavefunctions constructed out of hybridized atomic orbitals are accurate enough to predict some things, but not others. Predictions of any theory must be compared with empirical evidence to assess when they work and when they fail. When a theory gives the wrong answer, at least one assumption must not hold. In this case, the valence bond wavefunction is not accurate enough to capture some important features of a system's electronic structure. It may not be the most intellectually satisfying answer, but to say more would result in a much more complicated answer and certainly far beyond the level reasonably expected from general chemistry discussions.
3.1.03: Formal Charge
The formal charge of an atom in a molecule is the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms. Another way of saying this is that formal charge results when we take the number of valence electrons of a neutral atom, subtract the nonbonding electrons, and then subtract the number of bonds connected to that atom in the Lewis structure.
Calculating Formal Charges
We calculate the formal charge of an atom in a molecule or polyatomic ions as follows:
Formal Charge = (valence electrons of the "free" element) - (unshared electrons) - (bonds).
We can double-check formal charge calculations by determining the sum of the formal charges for the whole structure. The sum of the formal charges of all atoms in a molecule must be zero; the sum of the formal charges in an ion should equal the charge of the ion.
We must remember that the formal charge calculated for an atom is not the actual charge of the atom in the molecule. Formal charge is only a useful bookkeeping procedure; it does not indicate the presence of actual charges.
Calculating Formal Charge from Lewis Structures
Assign formal charges to each atom in the interhalogen ion \(\ce{ICl4-}\).
Solution
We divide the bonding electron pairs equally for all \(\ce{I–Cl}\) bonds:
We assign lone pairs of electrons to their atoms. Each \(\ce{Cl}\) atom now has seven electrons assigned to it, and the I atom has eight.
Subtract this number from the number of valence electrons for the neutral atom:
• I: 7 – 8 = –1
• Cl: 7 – 7 = 0
The sum of the formal charges of all the atoms equals –1, which is identical to the charge of the ion (–1).
Exercise \(1\)
Calculate the formal charge for each atom in the carbon monoxide molecule:
Answer
C −1, O +1
Example: Calculating Formal Charge from Lewis Structures
Assign formal charges to each atom in the interhalogen molecule \(\ce{BrCl3}\).
Solution
Assign one of the electrons in each Br–Cl bond to the Br atom and one to the Cl atom in that bond:
Assign the lone pairs to their atom. Now each Cl atom has seven electrons and the Br atom has seven electrons.
Subtract this number from the number of valence electrons for the neutral atom. This gives the formal charge:
• Br: 7 – 7 = 0
• Cl: 7 – 7 = 0
All atoms in \(\ce{BrCl3}\) have a formal charge of zero, and the sum of the formal charges totals zero, as it must in a neutral molecule.
Exercise \(2\)
Determine the formal charge for each atom in \(\ce{NCl3}\).
Answer
N: 0; all three Cl atoms: 0 | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/03%3A_Simple_Bonding_Theory/3.01%3A_Lewis_Electron-Dot_Diagrams/3.1.02%3A_Breaking_the_octet_rule_with_higher_electron_counts_%28hypervalent_atoms%29.txt |
Two notable cases where Lewis theory fails to predict structure are the cases of beryllium (\(\ce{Be}\)) and boron (\(\ce{B}\)). These two atoms are in period 2 (\(n=2\)) of the periodic table and their atoms have the valence electron configurations of \(2s^2\) and \(2s^22p^1\), respectively.
Beryllium
The Lewis electron dot structures are shown below for \(\ce{BeX2}\), where \(\ce{X}\) is one of the halogens, \(\ce{F}\) or \(\ce{Cl}\).
Each of the structures above would predict a linear geometry for the \(\ce{BeX2}\) molecule. Together the three resonance structures suggest partial double-bond character in the Be-X bond, which results in an intermediate bond length between a single and double bond.
There are issues with each of these resonance structures. The structure on the left would predict only four electrons around \(\ce{Be}\); thus, the atom does not fulfill the octet rule. The structure on the left suggests multiple bonds for the halogen (\(\ce{X}\)) and high separation of charge with formal charge on each atom. The structure in the middle is a mix of these problems. None of these situations is ideal according to Lewis theory. Further, experimental data is not consistent with any of these structures or their resonance hybrid (except in the case of \(\ce{BeCl2}\) at very high temperatures).
It turns out that the monomer of \(\ce{BeX2}\) does exist, but only at very high temperatures and low pressures. Even under extreme conditions, the monomer is not particularly stable due to the electron deficiency around \(\ce{Be}\).
BeF2
At ambient temperature and pressure, BeF2 is a solid that looks similar to quartz (Figure \(1\)) The Be is four-coordinate with tetrahedral geometry; each F is two-coordinate and the Be-F bond length is 1.54 Å. This structure is possible due to an extended 3-dimensional network in the solid where adjacent BeF2 units are bonded to one another, as shown in Figure \(1\).
In the liquid phase, \(\ce{BeF2}\) has a fluctuating tetrahedral structure where Be and F ions exchange. The vapor phase is reached at temperatures higher than 1000 °C (at ~ 1 atm). In the vapor phase, \(\ce{BeF2}\) exists as a monomer with linear geometry and a bond length of 1.43 Å, consistent with a double bond between \(\ce{Be}\) and \(\ce{F}\).
BeCl2
At ambient temperature and pressure, \(\ce{BeCl2}\) is a solid. As in \(\ce{BeF2}\) described above, \(\ce{BeCl2}\) has four-coordinate, tetrahedral Be and two-coordinate Cl. In contrast to BeF2, solid BeCl2 is a 1-dimensional polymer consisting of edge-shared tetrahedral.
In the gas phase, BeCl2 exists as a dimer with two chlorine atoms bridging two Be atoms. In the dimer, the Be atoms are 3-coordinate. Bridging Cl atoms are two-coordinate, while terminal Cl atoms are one-coordinate. At higher temperatures in the vapor phase, the linear monomer also exists.
Boron (\(2s^22p^1\))
Prediction based on Lewis structures:
Lewis structures of \(\ce{BH3}\) and \(\ce{BF3}\) were described in Exercise 3.1.2, and are drawn again below for convenience.
Boron trihalides
Boron trihalides, like \(\ce{BF3}\), have properties that are largely predicted by Lewis structures and VSEPR theory. The Lewis structure for \(\ce{BF3}\) includes several resonance structures. The structure with only single bonds is the most common representation for this molecule because the charge separation shown in the other structures is considered to be unfavorable. The highly polarized B-F bond has a dipole moment that lies opposite to the indicated formal charges shown in the resonance structures with double bonds between boron and fluorine.
The resonance hybrid of \(\ce{BF3}\) predicts partial double bond character between boron and fluorine, thus a bond length shorter than a single bond. Using the Lewis structures and VSEPR theory, we would predict a trigonal planar geometry around boron. In fact, the actual structure of \(\ce{BF3}\) is a monomer with trigonal planar geometry and with bond length that is shorter than a single bond. The case is similar to structures of other boron trihalides as well.
Boron trihalides are electron deficient at the boron center and react readily with Lewis bases. In other words they are strong Lewis acids (electrophiles).
Boron trihydride (BH3 is really B2H6)
The properties of boron trihydride (\(\ce{BH3}\)) are not predicted by the simple predictions made through Lewis structures and VSEPR. The monomer, BH3, is not stable, but when dissolved in the presence of a Lewis base, BH3 can form a stable acid-base adduct. In its pure form, the compound actually exists as a dimeric gas with a molecular unit of B2H6 (try drawing a valid Lewis structure for that!). Its unexpected structure includes two H's that bridge the two boron atoms in 3-center-2-electron bonds. You can read more about B2H6 on the Wikipedia page for Diborane(6). The unique bonding observed for boron is described more in Chapter 8 and Chapter 15. | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/03%3A_Simple_Bonding_Theory/3.01%3A_Lewis_Electron-Dot_Diagrams/3.1.04%3A_Lewis_Fails_to_Predict_Unusual_Cases_-_Boron_and_Beryllium.txt |
Introduction to VSEPR
The Valence Shell Electron Repulsion (VSEPR) model can predict the structure of most molecules and polyatomic ions in which the central atom is a nonmetal; it also works for some structures in which the central atom is a metal. VSEPR builds on Lewis electron dot structures (discussed in Section 3.1); Lewis structures alone predict only connectivity while the Lewis structure and VSEPR together can predict the geometry of each atom in a molecule. The main idea of VSEPR theory is that pairs of electrons (in bonds and in lone pairs) repel each other. The pairs of electrons (in bonds and in lone pairs) are called "groups". Because electrons repel each other electrostatically, the most stable arrangement of electron groups (i.e., the one with the lowest energy) is the one that minimizes repulsion. Groups are positioned around the central atom in a way that produces the molecular structure with the lowest energy. In other words, the repulsion between groups around an atom favors a geometry in which the groups are as far apart from each other as possible. Although VSEPR is simplistic because it does not account for the subtleties of orbital interactions that influence molecular shapes, it accurately predicts the three-dimensional structures of a large number of compounds.
We can use the VSEPR model to predict the geometry around the atoms in a polyatomic molecule or ion by focusing on the number of electron pairs (groups) around a central atom of interest. Groups include bonded and unbonded electrons; a single bond, a double bond, a triple bond, a lone pair of electrons, or even a single unpaired electron each count as one group. The molecule or polyatomic ion is given an AXmEn designation, where A is the central atom, X is a bonded atom, E is a nonbonding valence electron group (usually a lone pair of electrons), and m and n are integers. The number of groups is equal to the sum of m and n. Using this information, we can describe the molecular geometry around a central atom, that is, the arrangement of the bonded atoms in a molecule or polyatomic ion. The geometries that are predicted from VSEPR when a central atom has only bonded groups (n = 0) are listed below in Table $1$. The cases where lone pairs contribute to the total groups (n $\geq$ 1) are discussed in the next section about lone pair repulsion.
Table $1$. Geometries predicted using VSEPR theory (bonded groups only).
Groups around central atom
(m + n)
Geometry Name Geometry Sketch Predicted bond Angle Example
2 linear 180°
3 trigonal plane 120°
4 tetrahedron 109.5°
5 trigonal bipyramid 90° and 120°
6 octahedron 90°
7 pentagonal bipyramid 90° and 72°
8 square antiprism 70.5°, 99.6° and 109.5°
Practice
VSEPR to predict Molecular Geometry
You can follow these four steps to predict the geometry around an atom using VSEPR:
1. Draw the Lewis electron structure of the molecule or polyatomic ion.
2. For the central atom of interest, assign the AXmEn designation and the total number of groups (m+n).
3. Determine the electron group arrangement around the central atom that minimizes repulsions.
4. Describe the molecular geometry.
Use the procedure above to complete the exercises below
Exercise $1$
Predict the geometry around the central atom in BeH2 and CO2.
Answer BeH2
1. The central atom, beryllium, contributes two valence electrons, and each hydrogen atom contributes one. The Lewis electron structure is
2. There are two groups around the central atom, and both groups are single bonds. Thus BeH2 is designated as AX2.
3. We see from Table $1$ that the arrangement that minimizes repulsions places the groups 180° apart.
4. From Table $1$ we see that with two bonding pairs, the molecular geometry that minimizes repulsions in BeH2 is linear.
Answer CO2
1. The central atom, carbon, contributes four valence electrons, and each oxygen atom contributes six. The Lewis electron structure is
2. The carbon atom forms two double bonds. Each double bond is counted as one group, so there are two groups around the central atom. Once again, both groups around the central atom are bonds, so CO2 is designated as AX2.
3. Like BeH2, the arrangement that minimizes repulsions places the groups 180° apart.
4. VSEPR only recognizes groups around the central atom (the carbon). Thus the lone pairs on the oxygen atoms do not influence the molecular geometry. With two bonded groups on the central atom and no lone pairs, the molecular geometry of CO2 is linear (Table $1$). The structure of $\ce{CO2}$ is shown in Table $1$.
Exercise $2$
Predict the geometry around the central atom in BCl3 and CO32-.
Answer BCl3
1. The central atom, boron, contributes three valence electrons, and each chlorine atom contributes seven valence electrons. The Lewis electron structure is
2. There are three groups around the central atom and all are single bonds. The structure is designated as AX3.
3. To minimize repulsions, the groups are placed 120° apart (Table $1$).
4. From Table $1$ we see that with three bonding pairs around the central atom, the molecular geometry of BCl3 is trigonal planar.
Answer CO32-
1. The central atom, carbon, has four valence electrons, and each oxygen atom has six valence electrons. The Lewis electron structure of one of three resonance forms is represented as
2. The structure of CO32− is a resonance hybrid. It has three identical bonds, each with a bond order of $1 \frac{1}{3}$. All electron groups are bonds. With three bonding groups around the central atom, the structure is designated as AX3.
3. We minimize repulsions by placing the three groups 120° apart (Table $1$).
4. We see from Table $1$ that the molecular geometry of CO32− is trigonal planar with bond angles of 120°.
In our next example we encounter the effects of lone pairs and multiple bonds on molecular geometry for the first time.
Exercise $3$
Predict the geometry around the central atom in CH4, PCl5 and SF6.
Answer CH4
1. The central atom, carbon, contributes four valence electrons, and each hydrogen atom has one valence electron, so the full Lewis electron structure is
2. There are four electron groups around the central atom. All electron groups are bonding pairs, so the structure is designated as AX4.
3. As shown in Table $1$, repulsions are minimized by placing the groups in the corners of a tetrahedron with bond angles of 109.5°.
4. With four bonding pairs, the molecular geometry of methane is tetrahedral (Table $1$).
Answer PCl5
1. Phosphorus has five valence electrons and each chlorine has seven valence electrons, so the Lewis electron structure of PCl5 is
2. There are five bonding groups around phosphorus, the central atom. All electron groups are bonds, so the structure is designated as AX5.
3. The structure that minimizes repulsions is a trigonal bipyramid, which consists of two trigonal pyramids that share a base (Table $1$).
4. The molecular geometry of PCl5 is trigonal bipyramidal, as shown below. The molecule has three atoms in a plane in equatorial positions and two atoms above and below the plane in axial positions. The three equatorial positions are separated by 120° from one another, and the two axial positions are at 90° to the equatorial plane. The axial and equatorial positions are not chemically equivalent.
Answer SF6
1. The central atom, sulfur, contributes six valence electrons, and each fluorine atom has seven valence electrons, so the Lewis electron structure is
With an expanded valence, this species is an exception to the octet rule.
2. There are six electron groups around the central atom, each a bonding pair. We see from Figure $2$ that the geometry that minimizes repulsions is octahedral.
3. With only bonding pairs, SF6 is designated as AX6. All positions are chemically equivalent, so all electronic interactions are equivalent.
4. There are six nuclei, so the molecular geometry of SF6 is octahedral.
3.02: Valence Shell Electron-Pair Repulsion
In the previous section, we saw how to use VSEPR to predict the geometry around a central atom based on the number of groups attached to a central atom. However, our previous discussion was limited to the simple cases where all of the groups were bonded groups (i.e., in the designation AXmEn , n=0). When all of the groups are bonds, the geometries can be predicted using information in Table 3.2.1 in the previous section. Now we will consider cases where one or more of these groups are lone pairs.
Lone pairs have stronger repulsive forces than bonded groups.
When one or more of the groups is a lone pair of electrons (non-bonded electrons), the experimentally-observed geometry around an atom is slightly different than in the case where all groups are bonds. The actual bond angles are similar, but not exactly the same, as those predicted based on the total number of groups (the "parent" geometry). When there is a mixture of group types (lone pairs (E) and bonded groups (X)) there are three different types of angles to consider: bond angles between two bonded atoms (X-X angles), angles between a bonded atom and a lone pair (X-E angles), and angles between two lone pairs (E-E angles). Empirical evidence shows the following trend in the degree of bond angles around atoms with a mixture of group types:
Trend in bond angles:
E-E >X-E >X-X
Using empirical evidence as a guide, we can predict that lone pairs repel other electron groups more strongly than bonded pairs. The molecular geometry of molecules with lone pairs of electrons is better predicted when we consider that electronic repulsion created by lone pairs is stronger than the repulsion from bonded groups. It is difficult to predict the exact bond angle based on this principle, but we can predict approximate angles, as described and summarized below in Table \(1\).
Table \(1\): Predictions of molecular geometry and bond angles around atoms with a mixture of bonded (X) and unbonded (E) electron groups.
Electron Groups (m + n)
2
(steric number = 2)
3
(steric number = 3)
4
(steric number = 4)
5
(steric number = 5)
6
(steric number = 6)
Parent Geometry
(0 lone pairs)
AXm
AX2, linear
180°
AX3, trig. plane
120°
AX4, tetrahedron
109.5°
AX5, trig. bipyramid
90°, 120°
AX6, octahedron
90°
1 lone pair
AXmE1
AX2E1,
bent
<120°
AX3E1,
trig. pyramid
<109.5°
AX4E1, see-saw
<90°, <120°
AX5E1,
square pyramid
90°, <90°
2 lone pairs
AXmE2
AX2E2, bent
<109.5°
AX3E2, T-shape
<90°
AX4E2,
square plane
90°
3 lone pairs
AXmE3
AX2E3, linear
180°
Table \(1\) summarizes the geometries and bond angles predicted for nearest-neighbor bonded groups on central atoms with a mixture of lone pairs and bonded groups. The table does not cover all possible situations; it only includes cases where there are two bonded groups in which an X-X angle is measurable between nearest-neighbors. A more detailed description of some selected cases is given below.
Two Electron Groups (m + n = 2)
(Steric number = 2) In the case that there are only two electron groups around a central atom, those groups will lie 180° from one another. This results in a linear molecular geometry with 180° bond angles.
Three Electron Groups (m + n = 3)
(Steric number = 3) In the case that there are three electron groups around a central atom, those groups will lie approximately 120° from one another in space. This results in an electronic geometry that is approximately trigonal planar. There are two different molecular geometries that are possible in this category:
• When all of the electron groups are bonds (m = 3 or AX3), the molecular geometry is a trigonal plane with 120° bond angles.
• When there is one lone pair (m=2, n=1 or AX2E1), the molecular geometry is bent with a bond angle that is slightly less than 120°.
AX2E Molecules: Example SO2
1. The central atom, sulfur, has 6 valence electrons, as does each oxygen atom. With 18 valence electrons, the Lewis electron structure is shown below.
2. There are three electron groups around the central atom: two double bonds and one lone pair. We initially place the groups in a trigonal planar arrangement to minimize repulsions (Table \(1\)).
3. With two bonding pairs and one lone pair, the structure is designated as AX2E. This designation has a total of three electron pairs, two X and one E. The lone pair occupies more space around the central atom than a bonding pair (even double bonds!). Bonding pairs and lone pairs repel each other electrostatically in the order BP–BP < LP–BP < LP–LP. In SO2, we have one BP–BP interaction and two LP–BP interactions.
4. The molecular geometry is described only by the positions of the nuclei, not by the positions of the lone pairs. Thus, with two nuclei and one lone pair the shape is bent, or V shaped, which can be viewed as a trigonal planar arrangement with a missing vertex. The O-S-O bond angle is expected to be less than 120° because of the extra space taken up by the lone pair.
Four Electron Groups (m + n = 4)
(Steric number = 4) In the case that there are four electron groups around a central atom, those groups will lie approximately 109.5° from one another in space. This results in an electronic geometry that is approximately tetrahedral. There are three different molecular geometries that are possible in this category:
• When all electron groups are bonds (m=4 or AX4), the molecular geometry is a tetrahedron with bond angles of 109.5°.
• When there is one lone pair (m=3, n=1 or AX3E1), the molecular geometry is a trigonal pyramid with bond angles of slightly less than 109.5°.
• When there are two lone pairs (m=2, n=2 or AX2E2), the molecular geometry is bent with bond angles of slightly less than 109.5°.
AX3E Molecules: Example NH3
One of the limitations of Lewis structures is that they depict molecules and ions in only two dimensions. With four electron groups, we must learn to show molecules and ions in three dimensions.
1. In ammonia, the central atom, nitrogen, has five valence electrons and each hydrogen donates one valence electron, producing the Lewis electron structure
2. There are four electron groups around nitrogen, three bonding pairs and one lone pair. Repulsions are minimized by directing each hydrogen atom and the lone pair to the corners of a tetrahedron.
3. With three bonding pairs and one lone pair, the structure is designated as AX3E. This designation has a total of four electron pairs, three X and one E. We expect the LP–BP interactions to cause the bonding pair angles to deviate significantly from the angles of a perfect tetrahedron.
The Difference in the Space Occupied by a Lone Pair of Electrons and by a Bonding Pair
As with SO2, this composite model of electron distribution and negative electrostatic potential in ammonia shows that a lone pair of electrons occupies a larger region of space around the nitrogen atom than does a bonding pair of electrons that is shared with a hydrogen atom.
4. There are three nuclei and one lone pair, so the molecular geometry is trigonal pyramidal. In essence, this is a tetrahedron with a vertex missing. However, the H–N–H bond angles are less than the ideal angle of 109.5° because of LP–BP repulsion. The bond angles in ammonia are 106.6°.
AX2E2Molecules: Example H2O
1. Oxygen has six valence electrons and each hydrogen has one valence electron, producing the Lewis electron structure
2. There are four groups around the central oxygen atom, two bonding pairs and two lone pairs. Repulsions are minimized by directing the bonding pairs and the lone pairs to the corners of a tetrahedron.
3. With two bonding pairs and two lone pairs, the structure is designated as AX2E2 with a total of four electron pairs. Due to LP–LP, LP–BP, and BP–BP interactions, we expect a significant deviation from idealized tetrahedral angles.
4. With two hydrogen atoms and two lone pairs of electrons, the structure has significant lone pair interactions. There are two nuclei about the central atom, so the molecular shape is bent, or V shaped, with an H–O–H angle that is even less than the H–N–H angles in NH3, as we would expect because of the presence of two lone pairs of electrons on the central atom rather than one. This molecular shape is essentially a tetrahedron with two missing vertices.
Five Electron Groups (m + n = 5)
(Steric number = 5) In the case that there are five electron groups around a central atom, there are two different types of positions around the central atom: equatorial positions and axial positions. The three equatorial ligands are 120° from one another and are 90° from each of the two axial ligands. The axial positions have three adjacent groups oriented 90° away in space. Axial groups are thus more crowded than the equatorial positions with only two adjacent groups at 90°. The crowding of axial positions results in slight differences in bond distances; crowded axial groups have longer bonds than the less crowded equatorial groups. Lone pairs of electrons generally prefer to occupy equatorial positions rather than axial positions. The justification for this preference, according to VSEPR theory, is that the lone electron pairs are more repulsive than bonding electron pairs, and thus the lone pairs prefer the less crowded equatorial positions.
The arrangement of five groups around a central atom results in a trigonal bipyramidal electronic geometry. There are four different molecular geometries that are possible in this category, depending upon the number of bonded groups and lone pairs of electrons:
• When all electron groups are bonds (m=5 or AX5), the molecular geometry is a trigonal bipyramid with bond angles of 120° and 90° between adjacent ligands.
• When there is one lone pair (m=4, n=1 or AX4E1), the lone pair occupies one of the equatorial positions. The molecular geometry is called a see saw with bond angles of slightly less than 120° and slightly less than 90°.
• When there are two lone pairs (m=3, n=2 or AX3E2), each lone pair occupies one of the three equatorial positions. The molecular geometry is T-shaped with bond angles of slightly less than 120° and slightly less than 90°.
• When there are three lone pairs (m=1, n=3 or AX3E2), the lone pairs occupy the three equatorial positions. The molecular geometry is linear with bond angles of 180°.
AX4EMolecules: SF4
1. The sulfur atom has six valence electrons and each fluorine has seven valence electrons, so the Lewis electron structure is
With an expanded valence, this species is an exception to the octet rule.
2. There are five groups around sulfur, four bonding pairs and one lone pair. With five electron groups, the lowest energy arrangement is a trigonal bipyramid.
3. We designate SF4 as AX4E; it has a total of five electron pairs. However, because the axial and equatorial positions are not chemically equivalent, where do we place the lone pair? If we place the lone pair in the axial position, we have three LP–BP repulsions at 90°. If we place it in the equatorial position, we have two 90° LP–BP repulsions at 90°. With fewer 90° LP–BP repulsions, we can predict that the structure with the lone pair of electrons in the equatorial position is more stable than the one with the lone pair in the axial position. We also expect a deviation from ideal geometry because a lone pair of electrons occupies more space than a bonding pair.
Illustration of the Area Shared by Two Electron Pairs versus the Angle between Them
At 90°, the two electron pairs share a relatively large region of space, which leads to strong repulsive electron–electron interactions.
4. With four nuclei and one lone pair of electrons, the molecular structure is based on a trigonal bipyramid with a missing equatorial vertex; it is described as a seesaw. The Faxial–S–Faxial angle is 173° rather than 180° because of the lone pair of electrons in the equatorial plane.
AX3E2Molecules: BrF3
1. The bromine atom has seven valence electrons, and each fluorine has seven valence electrons, so the Lewis electron structure is
Once again, we have a compound that is an exception to the octet rule.
2. There are five groups around the central atom, three bonding pairs and two lone pairs. We again direct the groups toward the vertices of a trigonal bipyramid.
3. With three bonding pairs and two lone pairs, the structural designation is AX3E2 with a total of five electron pairs. Because the axial and equatorial positions are not equivalent, we must decide how to arrange the groups to minimize repulsions. If we place both lone pairs in the axial positions, we have six LP–BP repulsions at 90°. If both are in the equatorial positions, we have four LP–BP repulsions at 90°. If one lone pair is axial and the other equatorial, we have one LP–LP repulsion at 90° and three LP–BP repulsions at 90°:
Structure (c) can be eliminated because it has a LP–LP interaction at 90°. Structure (b), with fewer LP–BP repulsions at 90° than (a), is lower in energy. However, we predict a deviation in bond angles because of the presence of the two lone pairs of electrons.
4. The three nuclei in BrF3 determine its molecular structure, which is described as T shaped. This is essentially a trigonal bipyramid that is missing two equatorial vertices. The Faxial–Br–Faxial angle is 172°, less than 180° because of LP–BP repulsions.
Because lone pairs occupy more space around the central atom than bonding pairs, electrostatic repulsions are more important for lone pairs than for bonding pairs.
AX2E3Molecules: I3−
1. Each iodine atom contributes seven electrons and the negative charge one, so the Lewis electron structure is
2. There are five electron groups about the central atom in I3, two bonding pairs and three lone pairs. To minimize repulsions, the groups are directed to the corners of a trigonal bipyramid.
3. With two bonding pairs and three lone pairs, I3 has a total of five electron pairs and is designated as AX2E3. We must now decide how to arrange the lone pairs of electrons in a trigonal bipyramid in a way that minimizes repulsions. Placing them in the equatorial positions eliminates 90° LP–LP repulsions and minimizes the number of 90° LP–BP repulsions.
The three lone pairs of electrons have equivalent interactions with the three iodine atoms, so we do not expect any deviations in bonding angles.
4. With three nuclei and three lone pairs of electrons, the molecular geometry of I3 is linear. This can be described as a trigonal bipyramid with three equatorial vertices missing. The ion has an I–I–I angle of 180°, as expected.
Six Electron Groups (m + n = 6)
(Steric number = 6) In the case that there are six electron groups around a central atom, the nearest groups will lie approximately 90° from one another in space. This results in an electronic geometry that is approximately octahedral. There are three relevant molecular geometries in this category:
• When all electron groups are bonds (m=6 or AX6), the molecular geometry is an octahedron with bond angles of 90° between adjacent bonds.
• When there is one lone pair (m=5, n=1 or AX5E1) we now distinguish between the axial and equitorial positions; the lone pair is considered to be in one of the axial positions, while the bond directly opposite of the lone pair is the axial bond. The molecular geometry is a square pyramid with bond angles of 90° between adjacent equatorial bonds and slightly less than 90° between the axial bond and equatorial groups.
• When there are two lone pairs (m=4, n=2 or AX4E2), the lone pairs are opposite of one another and each occupy an axial position. The molecular geometry is square planar with bond angles of 90°.
AX5EMolecules: BrF5
1. The central atom, bromine, has seven valence electrons, as does each fluorine, so the Lewis electron structure is
With its expanded valence, this species is an exception to the octet rule.
2. There are six electron groups around the Br, five bonding pairs and one lone pair. Placing five F atoms around Br while minimizing BP–BP and LP–BP repulsions gives the following structure:
3. With five bonding pairs and one lone pair, BrF5 is designated as AX5E; it has a total of six electron pairs. The BrF5 structure has four fluorine atoms in a plane in an equatorial position and one fluorine atom and the lone pair of electrons in the axial positions. We expect all Faxial–Br–Fequatorial angles to be less than 90° because of the lone pair of electrons, which occupies more space than the bonding electron pairs.
4. With five nuclei surrounding the central atom, the molecular structure is based on an octahedron with a vertex missing. This molecular structure is square pyramidal. The Faxial–B–Fequatorial angles are 85.1°, less than 90° because of LP–BP repulsions.
AX4E2Molecules: ICl4−
1. The central atom, iodine, contributes seven electrons. Each chlorine contributes seven, and there is a single negative charge. The Lewis electron structure is
2. There are six electron groups around the central atom, four bonding pairs and two lone pairs. The structure that minimizes LP–LP, LP–BP, and BP–BP repulsions is
3. ICl4 is designated as AX4E2 and has a total of six electron pairs. Although there are lone pairs of electrons, with four bonding electron pairs in the equatorial plane and the lone pairs of electrons in the axial positions, all LP–BP repulsions are the same. Therefore, we do not expect any deviation in the Cl–I–Cl bond angles.
4. With five nuclei, the ICl4− ion forms a molecular structure that is square planar, an octahedron with two opposite vertices missing.
Summary
The arrangement of bonded atoms in a molecule or polyatomic ion is crucial to understanding the chemistry of a molecule, but Lewis electron structures give no information about molecular geometry. The valence-shell electron-pair repulsion (VSEPR) model allows us to predict which of the possible structures is actually observed in most cases. VSEPR is based on the assumption that pairs of electrons occupy space, and the lowest-energy structure is the one that minimizes repulsions between electron pairs. In the VSEPR model, the molecule or polyatomic ion is given an AXmEn designation, where A is the central atom, X is a bonded atom, E is a nonbonding valence electron group (usually a lone pair of electrons), and m and n are integers. Each group around the central atom is designated as a bonding pair (BP) or lone (nonbonding) pair (LP). From the BP and LP interactions we can predict both the relative positions of the atoms and the angles between the bonds, called the bond angles. From this we can describe the molecular geometry. The VSEPR model can be used to predict the shapes of many molecules and polyatomic ions, but it gives no information about bond lengths and the presence of multiple bonds. A combination of VSEPR and a bonding model, such as Lewis electron structures, is necessary to understand the presence of multiple bonds.
The relationship between the number of electron groups around a central atom, the number of lone pairs of electrons, and the molecular geometry is summarized in Table \(1\).
Example Excercises
Example \(1\)
Using the VSEPR model, predict the molecular geometry of each molecule or ion.
1. PF5 (phosphorus pentafluoride, a catalyst used in certain organic reactions)
2. H3O+ (hydronium ion)
Given: two chemical species
Asked for: molecular geometry
Strategy:
1. Draw the Lewis electron structure of the molecule or polyatomic ion.
2. Determine the electron group arrangement around the central atom that minimizes repulsions.
3. Assign an AXmEn designation; then identify the LP–LP, LP–BP, or BP–BP interactions and predict deviations in bond angles.
4. Describe the molecular geometry.
Solution:
1. A The central atom, P, has five valence electrons and each fluorine has seven valence electrons, so the Lewis structure of PF5 is
B There are five bonding groups about phosphorus. The structure that minimizes repulsions is a trigonal bipyramid.
C All electron groups are bonding pairs, so PF5 is designated as AX5. Notice that this gives a total of five electron pairs. With no lone pair repulsions, we do not expect any bond angles to deviate from the ideal.
D The PF5 molecule has five nuclei and no lone pairs of electrons, so its molecular geometry is trigonal bipyramidal.
• A The central atom, O, has six valence electrons, and each H atom contributes one valence electron. Subtracting one electron for the positive charge gives a total of eight valence electrons, so the Lewis electron structure is
B There are four electron groups around oxygen, three bonding pairs and one lone pair. Like NH3, repulsions are minimized by directing each hydrogen atom and the lone pair to the corners of a tetrahedron.
C With three bonding pairs and one lone pair, the structure is designated as AX3E and has a total of four electron pairs (three X and one E). We expect the LP–BP interactions to cause the bonding pair angles to deviate significantly from the angles of a perfect tetrahedron.
D There are three nuclei and one lone pair, so the molecular geometry is trigonal pyramidal, in essence a tetrahedron missing a vertex. However, the H–O–H bond angles are less than the ideal angle of 109.5° because of LP–BP repulsions:
Exercise \(1\)
Using the VSEPR model, predict the molecular geometry of each molecule or ion.
1. XeO3
2. PF6
3. NO2+
Answer a
trigonal pyramidal
Answer b
octahedral
Answer c
linear
Example \(2\)
Predict the molecular geometry of each molecule.
1. XeF2
2. SnCl2
Given: two chemical compounds
Asked for: molecular geometry
Strategy:
Use the strategy given in Example\(1\).
Solution:
1. A Xenon contributes eight electrons and each fluorine seven valence electrons, so the Lewis electron structure is
B There are five electron groups around the central atom, two bonding pairs and three lone pairs. Repulsions are minimized by placing the groups in the corners of a trigonal bipyramid.
C From B, XeF2 is designated as AX2E3 and has a total of five electron pairs (two X and three E). With three lone pairs about the central atom, we can arrange the two F atoms in three possible ways: both F atoms can be axial, one can be axial and one equatorial, or both can be equatorial:
The structure with the lowest energy is the one that minimizes LP–LP repulsions. Both (b) and (c) have two 90° LP–LP interactions, whereas structure (a) has none. Thus both F atoms are in the axial positions, like the two iodine atoms around the central iodine in I3. All LP–BP interactions are equivalent, so we do not expect a deviation from an ideal 180° in the F–Xe–F bond angle.
D With two nuclei about the central atom, the molecular geometry of XeF2 is linear. It is a trigonal bipyramid with three missing equatorial vertices.
• A The tin atom donates 4 valence electrons and each chlorine atom donates 7 valence electrons. With 18 valence electrons, the Lewis electron structure is
B There are three electron groups around the central atom, two bonding groups and one lone pair of electrons. To minimize repulsions the three groups are initially placed at 120° angles from each other.
C From B we designate SnCl2 as AX2E. It has a total of three electron pairs, two X and one E. Because the lone pair of electrons occupies more space than the bonding pairs, we expect a decrease in the Cl–Sn–Cl bond angle due to increased LP–BP repulsions.
D With two nuclei around the central atom and one lone pair of electrons, the molecular geometry of SnCl2 is bent, like SO2, but with a Cl–Sn–Cl bond angle of 95°. The molecular geometry can be described as a trigonal planar arrangement with one vertex missing.
Exercise \(2\)
Predict the molecular geometry of each molecule.
1. SO3
2. XeF4
Answer a
trigonal planar
Answer b
square planar | textbooks/chem/Inorganic_Chemistry/Inorganic_Chemistry_(LibreTexts)/03%3A_Simple_Bonding_Theory/3.02%3A_Valence_Shell_Electron-Pair_Repulsion/3.2.01%3A_Lone_Pair_Repulsion.txt |
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