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Structure maps, which plot structures against properties such electronegativity, are more consistent than radius ratio rules in correctly predicting coordination numbers and crystal structures. One of the early examples of this approach was published by Mooser and Pearson in 1959.[3]
A Mooser-Pearson diagram maps crystal structures according to the average principal quantum numbers of the atoms and their electronegativity difference. The basic ideas behind such a plot are:
• The greater the electronegativity difference, the more ionic is the compound. Higher ionicity results in higher coordination numbers because anions like to surround cations (and vice versa).
• Higher principal quantum numbers result in less s-p hybridization, less directional bonding, and therefore higher coordination number. We saw this trend before with the structures of elements in group IV: descending the group the coordination number increases progressively from 3-4 (carbon) to 12 (Pb).
The lines in the Mooser-Pearson diagram separate MX compounds with CsCl, NaCl, and tetrahedral (wurtzite and zincblende) structures. Note that wurtzite has higher ionicity than zincblende in the plot, consistent with our discussion of the "boat" and "chair" ring structures in Chapter 8. Diamorphic compounds tend to fall on the boundaries. On the whole, the Mooser-Pearson diagram makes far fewer errors in predicting structures than the radius ratio rule. There are similar diagrams for MX2 structures, in which the order of ionicity is CaF2 (8:4 coordination) > rutile (6:3) > silica structures (4:2).
9.03: Energetics of Crystalline Solids- The Ionic Model
Many ionic compounds have simple structures. Because the forces holding the atoms together are primarily electrostatic, we can calculate the cohesive energy of the crystal lattice with good accuracy. Interesting questions to ask about these lattice energy calculations are:
• How accurate are lattice energy calculations?
• What do they teach us about the chemical bonds in ionic crystals?
• Can we use lattice energies to predict properties such as solubility, stability, and reactivity?
• Can we use lattice energies to predict the crystal structures of ionic compounds?
Let's start by looking at the forces that hold ionic lattices together. There are mainly two kinds of force that determine the energy of an ionic bond.
The NaCl crystal structure is the archetype for calculating lattice energies and computing enthalpies of formation from Born-Haber cycles.
1) Electrostatic Force of attraction and repulsion (Coulomb's Law): Two ions with charges z+ and z-, separated by a distance r, experience a force F:
$\mathbf{F} = -\frac{e^{2}}{4\pi \varepsilon_{0}} \frac{z_{+}z_{-}}{r^{2}}$
where
e = 1.6022×10−19 C
4 π ε0 = 1.112×10−10 C²/(J m)
This force is attractive for ions of opposite charge.
The electrostatic potential energy, Eelec, is then given by
$\mathbf{E}_{elec} \int_{\infty}^{r} F(r)dr = \frac{e^{2}}{4\pi \varepsilon_{0}} \frac{z_{+}z_{-}}{r}$
The sign of Eelec is negative for the attractive interaction between a cation and an anion. That is, the closer oppositely charged ions approach each other, the lower the potential energy.
2) Closed-shell repulsion. When electrons in the closed shells of one ion overlap with those of another ion, there is a repulsive force comes from the Pauli exclusion principle. A third electron cannot enter an orbital that already contains two electrons. This force is short range, and is typically modeled as falling off exponentially or with a high power of the distance r between atoms. For example, in the Born approximation, B is a constant and ρ is a number with units of length, which is usually empirically determined from compressibility data. A typical value of ρ is 0.345 Å.
$\mathbf{E}_{repulsion} = Bexp(\frac{-r}{\rho})$
The energy of the ionic bond between two atoms is then calculated as the combination of net electrostatic and the closed-shell repulsion energies, as shown in the figure at the right. Note that for the moment we are ignoring the attractive van der Waals energy between ions, which we will explain below. For a pair of ions, the equilibrium distance between ions is determined by the minimum in the total energy curve. At this distance, the net force on each ion is zero.
Electrostatic energy of a crystal lattice
We can use these equations to calculate the lattice energy of a crystal by summing up the interactions between all pairs of ions. Because the closed-shell repulsion force is short range, this term is typically calculated only for interactions between neighboring ions. However, the Coulomb force is long range, and must be calculated over the entire crystal. This problem was first solved in 1918 by Erwin Madelung, a German physicist.[4]
Consider an ion in the NaCl structure labeled "O" in the diagram at the right. We can see that the nearest neighbor interactions (+ -) with ions labeled "1" are attractive, the next nearest neighbor interactions (- - and + +) are repulsive, and so on. In the NaCl structure, counting from the ion in the center of the unit cell, there are 6 nearest neighbors (on the faces of the cube), 12 next nearest neighbors (on the edges of the cube), 8 in the next shell (at the vertices of the cube), and so on. Their distances from ion "0" increase progressively: ro, √2 ro, √3 ro, and so on, where ro is the nearest neighbor distance.
We can now write the electrostatic energy at ion "O" as:
$\mathbf{E}_{elec} = -6\frac{e^{2}}{4\pi \varepsilon_{0}} \frac{z_{+}z_{-}}{r_{o}} + 12 \frac{e^{2}}{4 \pi \varepsilon_{0}} \frac{z_{+}z_{-}}{\sqrt{2}r_{o}} -8 \frac{e^{2}}{4 \pi \varepsilon_{0}} \frac{z_{+}z_{-}}{\sqrt{3}r_{o}} + \dots$
Factoring out constants and the nearest-neighbor bond distance ro we obtain:
$\mathbf{E}_{elec} = \frac{e^{2}}{4 \pi \varepsilon_{0}} \frac{z_{+}z_{-}}{r_{o}}(6 - \frac{12}{\sqrt{2}} + \frac{8}{\sqrt{3}} - \frac{6}{\sqrt{4}} + \dots)$
Where the sum in parentheses, which is unitless, slowly converges to a value of A = 1.74756. Generalizing this formula for any three-dimensional ionic crystal we get a function:
$\mathbf{E}_{elec} = \frac{e^{2}}{4 \pi \varepsilon_{0}} \frac{z_{+}z_{-}}{r_{o}} NA$
where N is Avogadro's number (because we are calculating energy per mole of ions) and A is called the Madelung constant. The Madelung constant depends only on the geometrical arrangement of the ions and so it varies between different types of crystal structures, but within a given structure type it does not change. Thus MgO and NaCl have the same Madelung constant because they both have the NaCl structure
The Madelung constant is calculated by summing up electrostatic interactions with ion labeled 0 in the expanding spheres method. Each number designates the order in which it is summed. For example, ions labeled 1 represent the six nearest neighbors (attractive interaction), ions labeled 2 are the 12 next nearest neighbors (repulsive interaction) and so on. Note that if the sum is carried out over shells 1-2-3..., it converges very slowly, but there are mathematical methods for summing it which give a rapidly converging series.
The table below lists Madelung constants for some common structures. The reduced Madelung constant is obtained by normalizing the values to the number of ions in the formula unit. It can be seen from the table that the reduced Madelung constants are quite similar for different structures. This makes it hard to determine on the basis of electrostatic energy calculations which structure will be most stable for a given compound. It is interesting to note that the trend in reduced Madelung constants roughly follows the trend in ionicity (cf. the Mooser-Pearson diagram in Section 9.2). For example, wurtzite has a slightly higher Madelung constant than zincblende, consistent with our earlier conclusion that the wurtzite structure is favored by more polar compounds.
Structure Madelung Constant, A Reduced Madelung Constant, 2A/n
NaCl (halite) 1.7476 1.7476
CsCl 1.7627 1.7627
ZnS (zincblende) 1.6381 1.6381
ZnS (wurtzite) 1.6413 1.6413
CaF2 (fluorite) 2.5194 1.6796
TiO2 (rutile) 2.4080 1.6053
Al2O3 (corundum) 4.172 1.6688
Total lattice energy of a crystal
Having in hand a formula for the electrostatic energy, we can now add in the closed-shell repulsion term to obtain an equation that gives us the total lattice energy.
$E_{L} = \frac{e^{2}}{4\pi \varepsilon_{0}} \frac{z_{+}z_{-}}{r_{o}} NA + NBexp(\frac{-r}{\rho})$
At the equilibrium bond distance, the forces on all the ions are zero, and we can use this fact to eliminate the constant B:
$[\frac{dE}{dr}]_{r=r_{o}} = 0$
Expressed this way, EL is a negative number (because z+ and z- have opposite signs). It represents the energy change for forming one mole of solid salt from one mole of the gaseous ions, separated initially at an infinite distance.
Lithium fluoride (shown here as a large single crystal in a beaker of water) is the only alkali halide that is not freely soluble in water. The lattice energy of LiF is the most negative of the alkali fluorides because Li+ and F- are both small ions and EL is proportional to 1/r0. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/09%3A_Ionic_and_Covalent_Solids_-_Energetics/9.02%3A_Structure_Maps.txt |
Now that we have an equation for the lattice energy of an ionic crystal, we can ask the question of how accurate it is. Remember, we made several approximations in arriving at this formula. We assumed that the lattice was completely ionic, we ignored the van der Waals attractive energy of the ions, and we assumed that there was no covalent contribution to the bonding.
Let's consider the lattice energy of table salt (NaCl)
$\ce{Na^{+}_{(g)} + Cl^{-}_{(g)} -> NaCl_{(s)}}$
To calculate the lattice energy, we lump together the physical constants:
$E_{L} (\frac{kJ}{mol}) = (1389.3) \frac{Aq_{1}q_{2}}{r_{o}}(1-\frac{.345}{r_{o}})$
where ro is expressed in Å. Now we can calculate the lattice energy for NaCl using ro = 2.814 Å, as:
$E_{L} = -(1389.3) \frac{1.7476}{2.814} (1 -\frac{.345}{2.814}) = -766.5 \frac{kJ}{mol}$
We can alternatively construct a Born-Haber cycle for the formation of NaCl from the elements and calculate the lattice energy as the "missing" term in the cycle.
S= Sublimation energy of Na(s)
IP= Ionization potential of Na(g)
D= Bond dissocation energy of Cl2(g)
EA= Electron affinity of Cl(g)
EL=Lattice energy of NaCl
R= Gas constant
T= Absolute temperature
From Hess' Law: $\Delta H_{f} = s + \frac{1}{2}D + IP + EA + E_{L} -2RT = \mathbf{-396 \frac{kJ}{mol}}$
Here we have to subtract 2RT to convert our cycle of energies to a cycle of enthalpies, because we are compressing two moles of gas in making NaCl(s) and PΔV = ΔnRT, where Δn = -2.
Experimentally ΔHf for NaCl is -411 kJ/mol
Because all the other numbers in the cycle are known accurately, the error in our calculation is only about 15 kJ (about 2% of EL). The result is promising because we neglected the van der Waals term.
But....how did we get away with neglecting the van der Waals term?
This is because we used energy minimization to obtain the repulsion energy in the Born-Mayer equation. If we underestimate the attractive energy of the crystal lattice, the energy minimization criterion ensures that the repulsion energy is underestimated as well. The two errors partially compensate, so the overall error in the calculation is small.
We can do better by explicitly including the short-range van der Waals attractive energy between ions. The table below shows results of more detailed lattice energy calculations for ionic fluorides in which the van der Waals term is explicitly included. The errors in this case are only about 1% of EL.
Compound Calculated Lattice Energy (kJ/mol) Experimental EL from Born-Haber Cycle
MgF2 (rutile structure) -2,920 -2,908
CaF2 (fluorite structure) -2,586 -2,611
BaF2 (fluorite structure) -2,326 -2,368
Silver Halides
It is interesting to repeat this exercise for the silver halides, which have either the NaCl structure (AgF, AgCl, AgBr) or zincblende structure (AgI).
Silver Halide Calculated Cycle Difference (kJ/mol)
AgF -920 -954 34
AgCl -833 -908 75
AgBr -816 -900 84
AgI -778 -883 105
Looking at the table, we see that the error is small for AgF and becomes progressively larger for the heavier silver halides. However we are still obtaining answers within about 12% error even for AgI. Should we interpret the good agreement with values calculated from the ionic model to mean that these compounds are ionic? Clearly, this description is inappropriate for AgI, where the electronegativity difference Δχ is only 0.6 (compare this value to 0.4 for a C-H bond, which we typically view as non-polar).
A drop of siver nitrate solution, when added to a dilute hydrochloric acid solution, results in the immediate formation of a white silver chloride precipitate. This reaction is used as a qualitative test for the presence of halide ions in solutions. The covalent bonding contribution to the lattice energies of AgCl, AgBr, and AgI makes these salts sparingly soluble in water.
Again, we can interpret the fortuitous agreement between the calculated and experimentally obtained energies in terms of compensating errors. Our lattice energy calculation overestimates the ionic contribution in the case of the heavier silver halides, but underestimates the covalent contribution. Of these compounds, only AgF is soluble in water and should be thought of as an ionic compound. The others are progressively more insoluble in water (Ksp is 10-10, 10-13, and 10-16 for AgCl, AgBr, and AgI), reflecting increasing covalency as Δχ decreases.
The moral of the story is that simple lattice energy calculations based on the ionic model work well, but they do not necessarily imply that the compounds are ionic! | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/09%3A_Ionic_and_Covalent_Solids_-_Energetics/9.04%3A_Born-Haber_Cycles_for_NaCl_and_Silver_Halides.txt |
From the discussion above, it is clear that the lattice energy, EL, of an ionic crystal can be calculated with reasonable accuracy if the structure is known. But how can we calculate EL for a new or hypothetical compound of unknown structure? Recall that the reduced Madelung constant is about the same for different crystal structures. Russian chemist A. F. Kapustinskii recognized this fact and devised a formula that allows one to calculate EL for any compound if we know the univalent radii of the constituent ions.[5]
The Madelung constant, A, is proportional to the number of ions (n) in formula unit, so dividing by the n gives similar values as shown in the table below:
A/n ~ invariant
Structure A/n
NaCl 0.874
CsCl 0.882
Rutile 0.803
Fluorite 0.800
Kapustinskii noticed that the difference in ionic radii between M+ and M2+ (the monovalent vs. divalent radius) largely compensates for the differences in A/n between monovalent (NaCl, CsCl) and divalent (rutile, CaF2) structures. He thus arrived at a lattice energy formula using an average Madelung constant, corrected to monovalent radii. In the Kapustinskii formula, the lattice energy (kJ/mol) is given by:
$E_{L} = \frac{1213.8z_{+}z_{-}n}{r_{+}+r_{-}} (1 - \frac{0.345}{r_{+}+r_{-}})$
Here the sum of the monovalent radii is used in place of ro, the bond distance in the Born-Mayer equation. The beauty of this formula is that it requires no knowledge of the structure of the compound. Therefore it can be used, in combination with Born-Haber cycles, to predict the stability of unknown compounds. As we show below, this is a broadly useful tool in guiding syntheses and predicting the reactivity of inorganic solids.
9.06: Discovery of Noble Gas Compounds
In 1962 at the University of British Columbia, Neil Bartlett was working with the powerful oxidizer PtF6 and, because of an accidental leak in his vacuum line, noticed the compound’s reaction with O2 to generate a solid with formula "PtF6O2." The formula suggested Pt in the +10 oxidation state, which was clearly unreasonable because PtF6 was known to be a more powerful oxidizer than either molecular fluorine (F2) or molecular oxygen (O2). Bartlett noticed that the X-ray powder diffraction pattern of the compound was similar to that of Cs+AsF6-, a salt with the CsCl structure in which octahedral AsF6- ions occupy the chloride ion sites. This led Bartlett to propose a formulation of O2+PtF6- for his new compound.[6] Magnetic susceptibility data subsequently confirmed the presence of the paramagnetic O2+ cation, which (see Chapter 2) has a bond order of 2.5. This formulation implies that PtF6 was a strong enough oxidizing agent to oxidize molecular oxygen.
But just how strong an oxidizer is PtF6? Its electron affinity could be estimated by using a Born-Haber cycle, filling in the lattice energy of O2+PtF6- by means of Kapustinskii's formula:
The electron affinity (EA) for PtF6 can be calculated as EA = -159 - 1167 + 571 = -751 kJ/mol. To put it in perspective, this is 417 kJ/mol more exothermic than the electron affinity of atomic fluorine (334 kJ). PtF6 was by far the strongest oxidizer that had ever been made!
Bartlett recognized that Xe has ionization energy of +1170 kJ, which is very close to the ionization energy of O2. Since Xe+ should be about the same size as O2+, the lattice energy should be about the same with Xe+ in the cation site of the O2+PtF6- structure. Since all of the other terms in the Born-Haber cycle for the reaction of Xe with PtF6 are the same, Bartlett concluded that Xe+PtF6-, like O2+PtF6-, should be a stable compound. He purchased a lecture bottle of xenon gas and reacted the two compounds, producing an orange solid.[7] While the product initially formed in the reaction may in fact be Xe+PtF6-, the Xe+ free radical is a powerful Lewis acid and reacts further with excess PtF6. The ultimate product of the reaction is formulated [XeF+][Pt2F11-], a salt which contains Xe in the +2 oxidation state and Pt in the +5 oxidation state. This was an important discovery because it shattered the dogmatic notion, which derived from the octet rule, that elements in group VIII could not form bonds with other elements. The name of this group was changed from the "inert gases" to the "noble gases." Subsequently, many compounds of Xe and a few of Kr and even Ar (which is much harder to oxidize) were synthesized and characterized. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/09%3A_Ionic_and_Covalent_Solids_-_Energetics/9.05%3A_Kapustinskii_Equation.txt |
Lattice energies, in addition to guiding the discovery of unknown compounds, are useful in explaining the absence (i.e., the thermodynamic instability) of non-existent compounds.[8] For example, CuF and AuF are unknown compounds, whereas CuF2, AuF3, and AuF5 are stable. In contrast AgF is a known, stable compound.
From the Born Haber cycle for CuF, the compound should be marginally stable (ΔHfo = -140 kJ/mol) with respect to the elements. Why then is CuF unknown?
To gain insight into this question, we first construct a Born-Haber cycle for the formation of CuF2 from the elements. This compound is stable with respect to the elements by -368 kJ/mol.
Combining the two cycles we see that the disproportionation of CuF to Cu and CuF2 is spontaneous. From similar cycles, we can also predict that the reaction 3AuF → AuF3 + 2Au should be spontaneous.
Why is the lowest oxidation state unstable for these fluorides? The key point is that the large difference in EL values (2908-972=1926 kJ in the case of copper fluorides) drives their disproportionation reactions. Note that when we use the Kapustinskii equation, we calculate that EL for CuF2 is approximately three times that of CuF. We use the same univalent radii in both calculations, but Cu has a 2+ charge in CuF2 (doubling the lattice energy relative to CuF), and contains 3/2 as many ions. The product z+z-n is thus three times larger for CuF2. The difference in EL values will thus increase as EL for the monovalent salt increases. We know that fluorides, having a small anion radius, will give larger EL values than iodides, which have larger anions. Thus the disproportionation reaction becomes more favorable for CuF than it is for CuI.
The stability of the lower vs. higher oxidation state thus depends on the size of the anion. For example, in fluorides, CuF is unstable but CuF2 is stable. However, in iodides, CuI is stable whereas CuI2 is unstable. From this we can develop a broad conclusion: small anions (O,F) tend to stabilize higher oxidation states, whereas large anions (S, Br, I...) stabilize lower oxidation states. Note that this trend has to do with the size and not with the electronegativity of the anion. Coincidentally, F and O are electronegative elements, but it is really their small size that has consequences for the lattice energy and their stabilization of higher oxidation states.
Cuprous iodide (CuI) is a crystalline compound used in organic synthesis and cloud seeding. This compound can be made in the laboratory by reacting soluble Cu2+ salts with a solution of sodium or potassium iodide. Because CuI2 is thermodynamicaly unstable, the reaction liberates I2 and a CuI precipitate forms.
Remember that the hard-soft acid-base rules could be interpreted in terms of the dominance of ionic vs. covalent interactions. Here we have put the hard-hard interaction in quantitative terms, based on (electrostatic) lattice energies.
Ag appears to buck the periodic trend. Why is AgF stable? This is because the second IP is very high (2071 kJ vs. 1958 kJ for Cu, 1979 for Au). Thus both AgF and AgF2 are known fluorides of Ag. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/09%3A_Ionic_and_Covalent_Solids_-_Energetics/9.07%3A_Stabilization_of_High_and_Low_Oxidation_States.txt |
Another interesting consequence of lattice energies involves the formation of certain salts containing Na- and e- anions. These compounds are known as alkalides and electrides, respectively.[9] Most of these compounds have been discovered by Prof. James Dye at Michigan State University.
The alkali metals have one electron in their valence shell. For example, the electronic configurations of Na and K are [Ar]3s1 and [Kr]4s1, respectively. Although we are accustomed to seeing these very electropositive elements give up their electrons when they make compounds with electronegative elements, they can also gain an electron to achieve a [noble gas]ns2 configuration. This is possible with strong electron donors such as alkali metals, especially when the cation that is formed is stabilized by coordination to a crown ether. Typically, these compounds are synthesized by combining the alkali metal and the appropriate crown ether in liquid ammonia, and then evaporating the ammonia.
Cavities and channels in an electride
Electride salts are formed under similar conditions, except in this case the anion is simply an electron that exists in an anion "cavity" in the crystal. The crystal structures are clearly salt-like, with the cations (alkali cations stabilized by crown or cryptand ligands) alternating in the structure with electrons.
Alkalide Salt
\[\ce{2Na ->[NH_{3}] Na^{+}(NH3)+ e^{-}(NH3) ->[18C6] [Na(18C6)]^{+} Na^{-}}\]
Electride Salt
\[\ce{2Cs ->[NH_{3}][18C6] [Cs(18C6)]^{+} e^{-}}\]
Complexing Na+ (K+, Rb+, Cs+) with crown ether ligands stabilizes the M+ form of the metal ("salt" form). Because the metal cation with its ligand shell is rather large, the lattice energy of these salts is rather low.
18-Crown-6, a crown ether that strongly complexes Na+ cations
Solutions of electride salts are powerful reducing agents, as demonstrated by their use in the Birch reduction, in which aromatic compounds are hydrogenated to produce dienes. Electrides are also useful for reducing metal ions to metals. Evaporation of blue electride solutions in anhydrous ammonia affords a mirror of Na. Such solutions slowly lose their color as the electrons reduce ammonia to the amide anion:
\(\ce{[Na(NH3)6]^{+}e^{-} + NH3 -> NaNH2 + H2}\)
It is interesting to consider, in the context of lattice energies and Born-Haber cycles, what might happen without these ligands present. That is, we can ask the question of whether sodium metal would prefer to exist in the metallic form as Na(s), or to form the sodide salt Na+Na-, or the electride salt Na+ e-. Of course, we already know the answer to this question. Elemental sodium is clearly a metal (it is shiny, conducts electricity, and has a bcc crystal structure) and has never been observed in either of the "salt" forms. But how far away are these forms energetically?
We can calculate the energetics by assuming that the Na- ion is about the same size as Br-, and that the e- anion is about the same size as Cl-. Then the lattice energies in the cycles become the same as those of NaBr and NaCl:
The rather surprising result from these calculations is that sodium would be marginally stable as a sodide salt and very stable (by 161 kJ) as an electride. All the terms in these simple cycles are known precisely. But we must be doing something wrong here, because Na(s) is clearly metallic.
The key, subtle point here is that Born-Haber cycles consider only the potential energy (rather than the total energy) of the substances in the cycle. Normally we can ignore the kinetic energy part of the total energy, but in this case we cannot because of the quantum mechanical effect of resonance. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/09%3A_Ionic_and_Covalent_Solids_-_Energetics/9.08%3A_Alkalides_and_Electrides.txt |
The solution to the problem above must consider the quantum mechanical nature of the electron. The valence electrons in Na metal are in orbitals that are delocalized over the entire crystal. However in the Na+ e- "salt" form, the electrons are localized on specific anion sites. This localization imparts an additional kinetic energy (via the "particle in a box" effect) that adds to the total energy. From the quantum mechanical result for a particle in a one-dimensional box, we obtain
$KE= \frac{h^{2}n^{2}}{8mL^{2}}$
where:
• h = Plank's constant = 6.626 x 10-34 J s
• n = energy level, assumed to be the lowest, n = 1
• m = electron mass = 9.109 x 10-31 kg
• L = size of the box
If we approximate the size of the electron "box" as 3 Å (3 x 10-10 m), we obtain:
$KE= \frac{(6.626 \cdot 10^{-34})^{2}(6.022 \cdot 10^{23})}{8 \cdot (9.1 \cdot 10^{-31})(3 \cdot 10^{-10})^{2}} = 4.04 \cdot 10^{5} \frac{J}{mol} = + 404 \frac{kJ}{mol}$
This extra kinetic energy makes the Na+e- "salt" unstable relative to the electron-delocalized metal.
Sodium owes its metallic properties to the resonance stabilization of its delocalized valence electrons.
The calculation is not very accurate because the electron kinetic energy is not zero in the metal, and because the "box" size is not so well defined. However, it does illustrate that electron delocalization has a substantial effect in thermodynamically stabilizing metals. The situation is entirely analogous to the stabilization of aromatic molecules by electron delocalization. In molecules like benzene, resonance is also a quantum-mechanical kinetic energy effect. In general, the resonance stabilization energy is significantly larger in metals than it is in π-delocalized organic molecules. For example, the resonance energy of the six π-electrons in benzene is approximately 151 kJ/mol, less than half of the value we have calculated (per electron) in sodium metal.
Interestingly, several of the alkali metals (and other metallic elements) transform at ultrahigh pressures to optically transparent, insulating phases in which the valence electrons are localized.[10][11] These high pressure electride phases defy the general rule (which we will encounter in Chapter 10) that insulators transform to metals at sufficiently high density. They illustrate how the stable structure of an element can depend on the trade-off between the lattice energy of an ionic electride structure and the resonance energy of an electron-delocalized metal.
Crystal structure of Ca2N, with 2D electron layers shown schematically in blue.
In addition to the salt-like electrides that are formed by complexing alkali metals with crown ether ligands, there are a number of recently discovered solid state nitrides and oxides, such as Ca2N and Ca24Al28O64, that are more properly formulated as electrides, i.e. as [Ca2N+](e-)[12] and [Ca24Al28O64]4+(e)4.[13] In these compounds the Ca, Al, O, and N atoms have their ordinary octet oxidation states (+2, +3, -2, and -3, respectively), and electrons act as anions, filling in cage-like voids or layers in the crystal. For example, Ca2N adopts the anti-CdCl2 structure, as shown at the right, with void spaces between layers that are occupied by a 2D gas of electrons. Like other layered materials, Ca2N can be easily delaminated into thin nanosheets while retaining its structure and properties as an electride.[14] These compounds are powerful reducing agents and also have interesting activity as catalysts.[15]
9.10: Prelude to Ionic and Covalent Solids - Energetics
In Chapter 8, we learned all about crystal structures of ionic compounds. A good question to ask is, what makes a compound choose a particular structure? In addressing this question, we will learn about the forces that hold crystals together and the relative energies of different structures. This will in turn help us understand in a more quantitative way some of the heuristic concepts we have learned about in earlier chapters, such as hard-soft acid-base theory.
At ordinary pressures, the CsCl structure is adopted by only three of the alkali halides, CsCl, CsBr, and CsI. Under high pressure however, other alkali halides transform from the NaCl to the CsCl structure because of the higher Madelung constant of CsCl. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/09%3A_Ionic_and_Covalent_Solids_-_Energetics/9.09%3A_Resonance_Energy_of_Metals.txt |
The alkali oxides, made by reacting alkali metals (Li, Na, K, Rb, Cs) with oxygen, show an unusual trend. When lithium reacts with oxygen we obtain the binary oxide Li2O, as expected from combining an element in group I with one in group VI. Curiously, the oxide that forms most readily when sodium metal is oxidized is not Na2O, but is instead the peroxide Na2O2, which we can formulate as (Na+)2(O22-). With potassium, rubidium, and cesium we obtain the superoxides MO2, which contain the superoxide radical anion (O2-.) and should be formulated as (M+)(O2-). While it is possible to make Na2O, K2O, Rb2O, and Cs2O by reaction of the appropriate metal nitrate (MNO3) with elemental alkali metal M,[16] it is curious that these "normal valent" compounds do not form by direct reaction of the metal with oxygen.
Sodium metal is oxided in air to sodium peroxide, Na2O2
Because the alkali metals are all very electropositive (χ = 0.8-1.0), and oxygen is very electronegative (χ = 3.5), we expect all the compounds we make by combining them to be reliably ionic. Consistent with this picture we find that Li2O (along with Na2O, K2O, and Rb2O) adopts the antifluorite structure (8:4 coordination - see problem 8.8.2), which we expect to find with relatively ionic M2X compounds. Strangely however, Cs2O crystallizes in the anti-CdCl2 structure. This is odd because CdCl2 has a layered structure that we normally associate with polar covalent MX2 compounds (see section 8.4). In Cs2O, six Cs+ cations surround each O2- anion in an octahedron. Each Cs+ is coordinated to three O2- ions, and the Cs+ ions contact each other across a van der Waals gap. The juxtaposition of Cs+ ions near each other is clearly electrostatically unfavorable, so why does Cs2O prefer the anti-CdCl2 structure to antifluorite?
The answer has to do with the crowding of alkali ions around oxygen, as illustrated for K2O at the right. Because eight large K+ ions surround each O2- ion in the structure, the cations are essentially in contact. Indeed, the metal-oxygen bonds are "stretched" in Na2O, K2O, and Rb2O relative to M-O bonds with the same bond order in other structures.[17] The situation is so extreme for Cs2O that it finds an (electrostatically unfavorable) structure in which the coordination is lowered to 6:3. This packing problem is relieved somewhat in the peroxides, where the coordination is still 8:4 but the anion is larger, and especially in the superoxides where the cation:anion ratio is 1:1 and the coordination is 6:6. Thus the larger alkali ions (K+, Rb+, Cs+) tend to form superoxides.
Space-filling models of the crystal structures of K2O (top) and Cs2O (bottom). Oxygen atoms are red, potassium ions are blue, and cesium ions are magenta
Another way that we can rationalize this trend is through the energetics of forming the oxides, peroxides, and superoxides.
Let's calculate the enthalpy change (per mole of metal) for forming a metal oxide M2O from the metal and oxygen:
$\ce{M_{(s)} + \frac{1}{4} O2 _{(g)} -> \frac{1}{2} M2O_{(s)}}$
We can use Hess' law to write this as a sum of reactions:
Reaction ΔH
$\ce{M_{(s)}-> M_{(g)} -> M^{+}_{(g)}}$ $\Delta_{s} + IE = \Delta_{f, M^{+} (g)}$
$\ce{\frac{1}{4} O2_{(g)} -> \frac{1}{2}O_{(g)} -> \frac{1}{2} O^{2-}}$ $\frac{1}{4} \Delta H_{d} + \frac{1}{2} EA_{1} + \frac{1}{2}EA_{2} = \frac{1}{2} \Delta H_{f, O^{2-} (g)}$
$\ce{M^{+}_{(g)} + \frac{1}{2}O^{2-}_{(g)} -> \frac{1}{2}M2O_{(s)}}$ $\frac{1}{2} E_{L, M_{2}O} - \frac{3}{2}RT$
Overall:
$\ce{M_{(s)} + \frac{1}{4}O2_{(g)} -> \frac{1}{2}M2O_{(s)}}$ $\Delta H_{f, M^{+}(g)} + \frac{1}{2} \Delta H_{f, O^{2-} (g)} + \frac{1}{2} E_{L, M_{2}O} -\frac{3}{2} RT$
To get the enthalpy change for the overall reaction (the heat of formation of 1/2 mole of M2O) we will need the heats of formation of M+(g) and O2-(g), which are available from tabulated values, and EL, which we can calculate from Kapustinskii's equation.
Similarly, we can write for the formation of the alkali peroxides:
$\ce{M_{(s)} + \frac{1}{2}O2_{(g)} -> \frac{1}{2}M2O2_{(s)}}$
Reaction
Reaction ΔH
$\ce{M_{(s)} -> M_{(g)} -> M^{+}_{(g)}}$ $\Delta H _{(f, M^{+} (g)}$
$\ce{\frac{1}{2} O2_{(g)} -> \frac{1}{2}O2^{2-}_{(g)}}$ $\frac{1}{2} \Delta H_{f, O_{2}^{2-} (g)}$
$\ce{M^{+}_{(g)} + \frac{1}{2} O2^{2-}_{(g)} -> \frac{1}{2} M2O2_{(s)}}$ $\frac{1}{2} E_{L, M_{2}O_{2}} - \frac{3}{2} RT$
Overall:
$\ce{M_{(s)} + \frac{1}{2}O2_{(g)} ->\frac{1}{2}M2O2_{(s)}}$ $\Delta H_{f, M^{+} (g)} + \frac{1}{2} \Delta H_{f, O_{2}^{2-} (g)} + \frac{1}{2} E_{L, M_{2}O_{2}} - \frac{3}{2} RT$
and for the superoxides:
$\ce{M_{(s)} + O2_{(g)} -> MO2_{(s)}}$
Reaction ΔH
$\ce{M_{(s)} -> M_{(g)} -> M^{+}_{(g)}}$ $\Delta H_{f, M^{+} (g)}$
$\ce{O2_{(g)} -> O2^{-}_{(g)}}$ $\Delta H_{f, O_{2}^{-} (g)}$
$\ce{M^{+}_{(g)} + O2^{-}_{(g)} -> MO2_{(s)}}$ $E_{L, MO_{2}} - 2RT$
Overall:
$\ce{M_{(s)} + O2_{(g)} -> MO2_{(s)} }$ $\Delta H_{f, M^{+} (g)} + \Delta H_{(f, O_{2}^{-}(g)} + E_{L, MO_{2}} - 2RT$
For the gaseous anions and cations, we have the following heats of formation and ionic radii (CN=6):
Ion ΔHf, kJ ionic radius, Å
Li+ 678 0.76
Na+ 602 1.02
K+ 506 1.38
Rb+ 485 1.52
Cs+ 473 1.67
O2- 500 1.20
O22- 519 1.59
O2- -88 1.49
Now using Kapustinskii's equation, we can calculate the lattice energies for each compound; these have been converted to lattice enthalpies by subtracting 2 RT or 3 RT as appropriate:
$\mathbf{E}_{L} = \frac{1213.8z_{+}z_{-}n}{r_{+}+r_{-}}(1 - \frac{0.345}{r_{+}+r_{-}})$
Metal ΔHL,M2O ΔHL,M2O2 ΔHL,MO2
Li -3,065 kJ -2,651 kJ -918 kJ
Na -2,776 -2,433 -838
K -2,454 -2,178 -751
Rb -2,345 -2,090 -721
Cs -2,241 -2,007 -678
As expected, the lattice energies for M2O and M2O2 are comparable, the latter being somewhat smaller in magnitude because of the larger size of the O22- anion. The lattice energies of the superoxides, MO2, are about 1/3 those of the corresponding peroxides because both the anion and cation are singly charged, and there are only two ions per formula unit.
Now, putting it all together, we can use the lattice energies and heats of formation of the individual ions to compare the heats of formation (per mole of metal) of each of the oxides:
Metal 1/2 ΔHf,M2O 1/2 ΔHf,M2O2 ΔHf,MO2
Li -404 kJ -388 kJ -328 kJ
Na -338 -354 -324
K -271 -321 -328
Rb -241 -300 -324
Cs -53 -70 -81
We can see that for Li, the formation of Li2O is favored over Li2O2 or LiO2 because of the very favorable lattice energy of Li2O. As the lattice energy becomes less negative with increasing cation size, the peroxide becomes the most stable at Na. For the heavier alkalis, M2O becomes quite unstable and the superoxides MO2 are the most stable. This is consistent with our observations of the chemistry of the group I oxides.
Metal-air batteries
The alkali oxides are quite interesting in the context of metal-air batteries because of their potential for extremely high energy storage on a mass basis. Such batteries have alkali metal (typically Li) or Zn anodes and utilize oxygen from the air at the cathode. Although lithium is the lightest and therefore the most energy-dense alkali metal, there are materials problems associated with the formation of Li dendrites when the battery is recharged, and also with the slow kinetics of the four-electron interconversion between O2(g) and 2 O2- at the cathode. For this reason, superoxide batteries are currently being studied as alternatives. The one-electron cathode reaction O2 + e- = O2- is kinetically fast, and potassium[18] and sodium[19] represent potentially viable alternatives to lithium for the anode of these air-breathing batteries. Recently, it has been shown that LiO2 can be kinetically stabilized by template growth on iridium nanoparticles, potentially opening the door to very high energy density lithium-air batteries.[20]
Schematic of an air-breathing lithium battery | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/09%3A_Ionic_and_Covalent_Solids_-_Energetics/9.11%3A_The_Strange_Case_of_the_Alkali_Oxides.txt |
Lattice energies can also help predict compound solubilities. Let's consider a Born-Haber cycle for dissolving a salt in water. We can imagine this as the sum of two processes: (1) the vaporization of the salt to produce gaseous ions, characterized by the lattice enthalpy, and (2) the hydration of those ions to produce the solution. The enthalpy change for the overall process is the sum of those two steps. We know that the entropy change for dissolution of a solid is positive, so the solubility depends on the enthalpy change for the overall process.
Here we need to consider the trends in both the lattice energy EL and the hydration energy EH. The lattice energy depends on the sum of the anion and cation radii (r+ + r-), whereas the hydration energy has separate anion and cation terms. Generally the solvation of small ions (typically cations) dominates the hydration energy because of the 1/r2 dependence.
$E_{L} \alpha \frac{1}{r_{+}+ r_{-}}$
$E_{H} \alpha \frac{1}{r_{+}^{2}} + \frac{1}{r_{-}^{2}}$
For salts that contain large anions, EL doesn't change much as r+ changes. That is because the anion dominates the r+ + r- term in the denominator of the formula for EL. On the other hand, EH changes substantially with r+, especially for small cations.
As a result, sulfate salts of small divalent cations, such as MgSO4 (epsom salts), are soluble, whereas the lower hydration energy of Ba2+ in BaSO4 makes that salt insoluble (Ksp = 10-10).
Left: EL diagram for sulfate salts. The large SO42- ion is size-mismatched to small cations such as Mg2+, which have large hydration energies, resulting soluble salts. With larger cations such as Ba2+, which have lower EH, the lattice energy exceeds the solvation enthalpy and the salts are insoluble.. Right: In the case of small anions such as F- and OH-, the lattice energy dominates with small cations such as transition metal ions (TMn+), Mg2+, and Li+. Anion-cation size mismatch occurs with larger cations, such as Cs+ and Ba2+, which make soluble fluoride salts.
For small anions, EL is more sensitive to r+, whereas EH does not depend on r+ as strongly. For fluorides and hydroxides, LiF is slightly soluble whereas CsF is very soluble, and Mg(OH)2 is insoluble whereas Ba(OH)2 is very soluble.
Putting both trends together, we see that low solubility is most often encountered when the anion and cation match well in their sizes, especially when one or both are multiply charged.
Space-filling models showing the van der Waals surfaces of Ba2+ and SO42-. The similarity in size of the two ions contributes to the low solubility of BaSO4 in water.
Combining all our conclusions about solubility, we note the following trends:
1) Increasing size mismatch between the anion and cation leads to greater solubility, so CsF and LiI are the most soluble alkali halides.
2) Increasing covalency leads to lower solubility in the salts (due to larger EL. For example, AgF, AgCl, AgBr, and AgI exhibit progressively lower solubility because of increasing covalency.
$\ce{AgF > AgCl > AgBr > AgI}$
3) Increasing the charge on the anion lowers the solubility because the increase in EL is large relative to the increase in EH.
4) Small, polyvalent cations (having large EH) make soluble salts with large, univalent anions such as I-, NO3-, ClO4-, PF6-, and acetate.
Examples: Salts of transition metal and lanthanide ions
• Ln3+: Nitrate salts are soluble, but oxides and hydroxides are insoluble.
• Fe3+: Perchlorate is soluble, but sulfate is insoluble.
5) Multiple charged anions such as O2-, S2-, PO43-, and SO42- make insoluble salts with most M2+, M3+, and M4+ metals. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/09%3A_Ionic_and_Covalent_Solids_-_Energetics/9.12%3A_Lattice_Energies_and_Solubility.txt |
• Explain why lattice energy calculations are very accurate for NaCl and CaCl2, but less accurate (by about 10%) for AgCl and PbCl2. Does the Born-Mayer equation under- or overestimate the latter values?
• Fluorine is more electronegative than oxygen. However, for many transition metals, we can make higher oxidation states in oxides than we can in fluorides. For example, Mn(IV) is stable in an oxide (MnO2), but MnF4 is unstable relative to MnF3 and fluorine.[21] Can you explain this in terms of lattice energies?
9.14: Problems
1. Use lattice energies to explain why MgSO4 decomposes to magnesium oxide and SO3 at a much lower temperature than does BaSO4.
2. Solid MgO might be formulated as Mg+O- or Mg2+O2-. Use the thermochemical data below (some of which are irrelevant) and Kapustinskii's formula to determine which is more stable. The lattice constant for MgO (NaCl structure) is 4.213 Å. While the idea of an O- ion might seem strange, note that the second electron affinity of O and the second ionization potential of Mg (in the table below) are both quite endothermic.
Reaction ∆Ho, kJ/mol
Mg(s) = Mg(g) 148
Mg(g) = Mg+(g) + e- 739
Mg+(g) = Mg2+(g) + e- 1,452
O2(g) = 2 O(g) 498
O(g) + e- = O-(g) -141
O-(g) + e- = O2-(g) 790
3. From the heat of formation of solid NH4Cl (-315 kJ/mol) and gaseous NH3 (-46), the bond dissociation energies of H2 (436) and Cl2 (244), the ionization potential of atomic hydrogen (1,311), and the electron affinity of atomic chlorine (-349), calculate the gas-phase proton affinity of NH3. The lattice energy of NH4Cl may be estimated from Kapustinskii's formula using rN-Cl = 3.50 Å.
4. Bottles of aqueous ammonia are often labeled “ammonium hydroxide.” We will test this idea by using a lattice energy calculation to determine whether the salt NH4+OH- can exist.
The heats of formation of gaseous OH- and H2O are respectively -141 and -242 kJ/mol. Assuming that NH4+ is about the same size as Rb+, and OH- about the same size as F-, using Kapustinskii's formula, ionic radii, and the NH3 proton affinity calculated in problem 3, determine whether NH4+OH- should be a stable salt relative to NH3 and H2O. At what temperature should NH4+Cl- be unstable relative to NH3 and HCl, if ΔHfo for HCl is -92 kJ/mol and ΔSo (\(\ce{NH4Cl -> NH3 + HCl}\)) = 280 J/mol K?
5. Lithium metal burns in nitrogen to make the nitride Li3N. The heavier alkali metals (K, Rb, Cs) can form stable azides (MN3), but not M3N nitrides. Explain why this is so.
6. (a) Do you expect BaSO4 or MgSO4 to be more soluble in water? (b) Is LiF more soluble than LiClO4? Explain.
7. Which polymorph of ZnS (zincblende or wurzite) would you expect to be more stable on the basis of electrostatic energy?
8. Arsenic contamination of ground water is a serious problem in Bangladesh, Chile, Argentina, and other parts of the world including the western United States. Arsenic poisoning been widespread in the Ganges river delta, where tube wells bring contaminated water up from 20-100 meters below the surface. One simple treatment that has been proposed is to precipitate the arsenic by aeration of the well water, which also contains high concentrations of Fe2+. Referring to the Pourbaix diagram of arsenic below and the Pourbaix diagram of iron in Chapter 4, identify the iron and arsenic species that are present in aerated water at neutral pH. What insoluble compound precipitates to lower the concentration of arsenic? (Hint: which compound would have the largest lattice energy?)
9.15: References
1. Pauling, L. (1929). "The principles determining the structure of complex ionic crystals". J. Am. Chem. Soc. 51 (4): 1010–1026. doi:10.1021/ja01379a006.
2. K.-I. Kobayashi, T. Kimura, H. Sawada, K. Terakura, and Y. Tokura, Room-temperature magnetoresistance in an oxide material with an ordered double-perovskite structure, Nature (1998) 395, 677-680. DOI:10.1038/27167
3. E. Mooser and W. B. Pearson, On the Crystal Chemistry of Normal Valence Compounds, Acta. Cryst. 12, 1015 (1959).
4. Madelung E (1918). "Das elektrische Feld in Systemen von regelmäßig angeordneten Punktladungen". Phys. Zs. XIX: 524–533.
5. A. F. Kapustinskii: Lattice energy of ionic crystals, Quart. Rev. Chem. Soc. Nr. 10, 1956, pp. 283–294. DOI|10.1039/QR9561000283
6. Neil Bartlett and D. H. Lohmann (March 1962). "Dioxygenyl hexafluoroplatinate (V), O2+[PtF6]". Proceedings of the Chemical Society (London: Chemical Society) (3): 115. doi:10.1039/PS9620000097.
7. Bartlett, N. (June 1962). "Xenon hexafluoroplatinate (V) Xe+[PtF6]". Proceedings of the Chemical Society (London: Chemical Society) (6): 218. doi:10.1039/PS9620000197.
8. W. E. Dasent, Non-Existent Compounds, J. Chem. Educ., 1963, 40, p 130, DOI: 10.1021/ed040p130
9. Dye, J. L. (2003). "Electrons as Anions". Science 301 (5633): 607–608. doi:10.1126/science.1088103. PMID 12893933.
10. Ma, Y.; Eremets, M.; Oganov, A. R.; Xie, Y.; Trojan, I.; Medvedev, S.; Lyakhov, A. O.; Valle, M.; Prakapenka, V., "Transparent Dense Sodium," Nature 2009, 458, 182–183. doi:10.1038/nature07786
11. M.-S. Miao and R. Hoffmann, "High Pressure Electrides: A Predictive Chemical and Physical Theory," Acc. Chem. Res. 2014, 47, 1311–1317. DOI: 10.1021/ar4002922
12. K. Lee, et al., "Dicalcium nitride as a two-dimensional electride with an anionic electron layer," Nature, 2013, 494, 336–340. DOI:10.1038/nature11812
13. S. Matsuishi, et al., "High-Density Electron Anions in a Nanoporous Single Crystal: [Ca24Al28O64]4+(e)4," Science, 2003, 301, 626-629. DOI: 10.1126/science.1083842
14. D. L. Druffel et al., "Experimental Demonstration of an Electride as a 2D Material," J. Am. Chem. Soc., 2016, 138, 16089–16094. DOI: 10.1021/jacs.6b10114
15. Y. Inoue et al., "Highly Dispersed Ru on Electride [Ca24Al28O64]4+(e)4 as a Catalyst for Ammonia Synthesis," ACS Catal., 2014, 4, 674–680. DOI: 10.1021/cs401044a
16. Holleman, A.F.; Wiberg, E., eds (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 978-0-12-352651-9.
17. N. K. McGuire and M. O'Keeffe, "Bond lengths in alkali metal oxides," J. Solid State Chem. 1984, 54, 49-53. DOI:10.1016/0022-4596(84)90129-4
18. X. Ren and Y. Wu, "A low-overpotential potassium–oxygen battery based on potassium superoxide," J. Am. Chem. Soc. 2013, 135, 2923–2926. DOI:10.1021/ja312059q
19. P. Hartmann, et al., "A rechargeable room-temperature sodium superoxide (NaO2) battery," Nature Materials 2012, 12, 228–232. DOI:10.1038/nmat3486
20. Lu et al., "A lithium–oxygen battery based on lithium superoxide," Nature 2016, 529, 377-382. DOI:10.1038/nature16484
21. K. O. Christe, "Chemical synthesis of elemental fluorine," Inorg. Chem. 1986, 25,3721–3722. DOI: 10.1021/ic00241a001 | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/09%3A_Ionic_and_Covalent_Solids_-_Energetics/9.13%3A_Discussion_Questions.txt |
Learning Objectives
• Explain the physical basis of the Hubbard and Mott models of metal-insulator transitions.
• Understand why good superconductors derive from bad metals.
• Know the structures and the periodic trends in band gaps and colors of semiconductors.
• Obtain the band gap of an intrinsic semiconductor from the temperature dependence of the conductivity.
• Predict the doping type when impurities or defects are introduced into a semiconductor.
• Correlate the band picture and Fermi level with n- or p-type doping.
• Understand the physical principles of operation of diodes, LEDs, solar cells, and FETs.
• Explain the differences in structures and electronic properties of crystalline and amorphous semiconductors.
The band model (like MO theory) is based on a one-electron model. This was an approximation we made at the very beginning of our discussion of MO theory: we used hydrogen-like (one-electron) solutions to the Schrödinger equation to give us the shapes of s, p, d, and f atomic orbitals. In a one-electron atom, these orbitals are degenerate within a given shell, and the energy differences between, e.g., 2s and 2p orbitals arise only when we consider the energy of an electron in the field of other electrons in the atom. Moving from atoms to molecules, we made linear combinations to generate one-electron molecular orbitals (and, in solids, one-electron energy bands). But as in multi-electron atoms, life is not so simple for real molecules and solids that contain many electrons. Electrons repel each other and so their movement in molecules and in solids is correlated.
• 10.1: Prelude to Electronic Properties of Materials - Superconductors and Semiconductors
Correlated electron effects give rise to metal-insulator transitions that are driven by small changes in temperature, pressure, or composition, as well as to superconductivity - the passage of current with zero resistance at low temperatures. In this chapter we will develop some simple models to understand these interesting and important electronic properties of solids.
• 10.2: Metal-Insulator Transitions
Under experimentally accessible temperatures and pressures, Si and Ge are always semiconducting (i.e., insulating), and Pb is always metallic. Why is Sn different? The reason has to do with orbital overlap. Theory tells us in fact that any (and all) insulators should become metallic at high enough pressure, or more to the point, at high enough density. For most insulators, however, the pressures required are far beyond those that we can achieve in the laboratory.
• 10.3: Superconductors
The phenomenon of superconductivity, first discovered in Hg metal in 1911 by Onnes, continues to be only partially understood. It is of great interest to physicists as a macroscopic quantum phenomenon, and to chemists and materials scientists who try to make better superconductors (especially those that superconduct at higher temperatures) and devices derived from them, such as superconducting quantum interference devices (SQUIDs), which are extremely sensitive magnetometers.
• 10.4: Periodic Trends- Metals, Semiconductors, and Insulators
• 10.5: Semiconductors- Band Gaps, Colors, Conductivity and Doping
There are a number of places where we find semiconductors in the periodic table.
• 10.6: Semiconductor p-n Junctions
• 10.7: Diodes, LEDs and Solar Cells
Diodes are semiconductor devices that allow current to flow in only one direction. Diodes act as rectifiers in electronic circuits, and also as efficient light emitters (in LEDs) and solar cells (in photovoltaics). The basic structure of a diode is a junction between a p-type and an n-type semiconductor, called a p-n junction. Typically, diodes are made from a single semiconductor crystal into which p- and n-dopants are introduced.
• 10.8: Amorphous Semiconductors
Amorphous semiconductors are disordered or glassy forms of crystalline semiconductor materials with network structures involving primarily covalent bonding. Crystalline silicon, which has the diamond structure, is an ordered arrangement of fused six-membered silicon rings and the local bonding environment of the silicon atoms is tetrahedral. The silicon atoms in amorphous silicon (a-Si) are also predominantly tetrahedrally coordinated, but there is no long-range order in the structure.
• 10.9: Discussion Questions
• 10.10: Problems
• 10.11: References
10: Electronic Properties of Materials - Superconductors and Semiconductors
In Chapter 6 we developed an energy band picture for metals, starting from atomic orbitals and building up the molecular orbitals of the solid metallic crystal. This treatment gave us a useful picture of how electrons behave in metals, moving at very fast speed between scattering events, and migrating in an electric field at a slow drift velocity. It also taught us that a metal is something with a partially filled band, meaning that the Fermi level cuts through one of its bands of orbitals. An insulator or a semiconductor has a similar band picture, except that the bands are either completely full or completely empty. In this case the Fermi level lies in the gap between fully occupied and unoccupied bands. We will see in this chapter that the properties of semiconductors (along with their useful electronic applications) depend on the addition of small amounts of impurities ("dopants") that change the position of the Fermi level, resulting in conduction by electrons or "holes."
Modern integrated circuits contain billions of nanoscale transistors and diodes that are essential for logic and memory functions. Both kinds of devices rely on junctions between crystalline silicon regions that contain a few parts per million of boron or phosphorus impurities.
While the band picture works well for most crystalline materials, it does not tell us the whole story of conduction in solids. That is because the band model (like MO theory) is based on a one-electron model. This was an approximation we made at the very beginning of our discussion of MO theory: we used hydrogen-like (one-electron) solutions to the Schrödinger equation to give us the shapes of s, p, d, and f atomic orbitals. In a one-electron atom, these orbitals are degenerate within a given shell, and the energy differences between, e.g., 2s and 2p orbitals arise only when we consider the energy of an electron in the field of other electrons in the atom. Moving from atoms to molecules, we made linear combinations to generate one-electron molecular orbitals (and, in solids, one-electron energy bands). But as in multi-electron atoms, life is not so simple for real molecules and solids that contain many electrons. Electrons repel each other and so their movement in molecules and in solids is correlated. While this effect is weak in a "good" metal such as sodium - where the wavefunctions are highly delocalized - it can be quite important in other materials such as transition metal oxides. Correlated electron effects give rise to metal-insulator transitions that are driven by small changes in temperature, pressure, or composition, as well as to superconductivity - the passage of current with zero resistance at low temperatures. In this chapter we will develop some simple models to understand these interesting and important electronic properties of solids. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/10%3A_Electronic_Properties_of_Materials_-_Superconductors_and_Semiconductors/10.01%3A_Prelude_to_Electronic_Properties_of_Materials_-_Superconduct.txt |
In Chapter 6 we learned that metals and insulators not only have different electrical properties but also have very different crystal structures. Metals tend to have high coordination numbers (typically 8 or 12) whereas insulators have low coordination numbers that can be rationalized as "octet" bonding arrangements. For example, in crystalline Si or Ge (diamond structure), each atom has four nearest neighbors. There are two electrons per bond, and thus each atom has eight electrons in its valence shell. Sn, the element below Ge, exists in two different forms, one (gray tin) with the diamond structure that is a brittle narrow-gap semiconductor, and the other (white tin) with a body-centered tetragonal structure that is a malleable metal. These two forms are very close in energy, and in fact metallic white tin transforms to the brittle semiconducting gray form at low temperature. Extremely cold weather in 18th century Europe caused many tin organ pipes to break and eventually turn to dust. This transformation has been called tin blight, tin disease, tin pest or tin leprosy. The dust is actually grey tin, which lacks the malleability of its metallic cousin white tin.
Under experimentally accessible temperatures and pressures, Si and Ge are always semiconducting (i.e., insulating), and Pb is always metallic. Why is Sn different? The reason has to do with orbital overlap. Theory tells us in fact that any (and all) insulators should become metallic at high enough pressure, or more to the point, at high enough density. For most insulators, however, the pressures required are far beyond those that we can achieve in the laboratory.
How can we rationalize the transition of insulators to the metallic state? Indeed, how can we understand the existence of insulators at all?
The Hubbard Model
Let's consider a chain of a large number (N) of atoms as we did in Chapter 6. For convenience, we can say that these are atoms such as H, Na, or Cs that have one valence electron. The simple band model we developed earlier suggests that the chain should be metallic, because N atoms combine to make N orbitals, and the N valence electrons only fill the band of orbitals halfway. But this conclusion doesn't depend on the density, which creates a paradox. If atoms in the chain are very far apart, we suspect that the electrons should localize on the atoms.
Electron hopping in a 1-D chain of atoms.
A solution to this problem was proposed by J. Hubbard in 1963.[1] Hubbard considered the energy required to transfer an electron from an atom to its nearest neighbor, as shown in the picture at the right. Because each atom already has one electron (with random spin), moving an electron over by one atom requires overcoming the energy of electron-electron repulsion to make a cation-anion pair. For well-separated atoms this energy (U) is given by:
$U = IP -Ea -\frac{e^{2}}{4\pi \varepsilon_{0} d}$
where IP and EA are the ionization energy and electron affinity, ε0 is the permittivity of free space, and the last term in the equation represents the coulombic attraction between the cation and the anion. For atoms such as alkali metals, U is on the order of 3–5 eV, which is much larger than the thermal energy kT. Thus we expect there to be very few anion-cation pairs at room temperature, and the chain of atoms should be insulating.
Energy vs. DOS for the chain of atoms as the density and degree of orbital overlap between atoms increases. Increasing overlap broadens the neutral atom and anion-cation states into bands, each of which has a bandwidth Δ. A transition to the metallic state occurs abruptly when Δ exceeds the Hubbard gap U.
What happens when we squeeze the atoms together? In the Hubbard model, as the distance between atoms decreases, the energies of both the neutral atom states and the anion-cation states broaden into bands, each of which has a band width Δ. The lower band can accommodate exactly N electrons (not 2N as in the MO picture we developed earlier) because each orbital can only take one electron without spin-pairing. Thus for small Δ the lower band is full and the upper band is empty. However, as we continue to compress the chain, the orbital overlap becomes so strong that Δ ≈ U. At this point, the bands overlap and some of the electrons fill the anion-cation states. The chain then becomes conducting and the material is metallic.
Some materials, such as Sn and VO2, happen to have just the right degree of orbital overlap to make the Hubbard transition occur by changing the temperature or pressure. Such materials can be very useful for electrical switching, as illustrated at the right for rutile structure VO2. Most materials are far away from the transition, either on the metallic or insulating side. An interesting periodic trend that illustrates this concept can be seen among the transition metal monoxides, MO (M = Ti, V, Cr, Mn, Fe, Co, Ni), all of which have the NaCl structure. TiO and VO are metallic, because the 3d orbitals have significant overlap in the structure. However, CrO, MnO, FeO, CoO, and NiO are all insulators, because the 3d orbitals contract (and therefore Δ < U) going across the transition metal series. In contrast, the analogous sulfides (TiS, VS,....NiS) are all metallic. The sulfides have the NiAs structure, in which all the metal atoms are eclipsed along the stacking axis (the hexagonal c-axis). The short metal-metal distances along that axis result in strong orbital overlap, making Δ > U.
Vanadium dioxide has the rutile structure, and in its undistorted form it is metallic, with one valence electron per V atom. Distortion of the lattice makes pairs of V atoms, resulting in an electronically insulating state. The metal-insulator transition can be driven reversibly by changing the temperature, pressure, or orbital occupancy. Electrical switching of this transition in VO2 is being studied for applications in high performance thin film transistors[2]
The Mott Model
A simpler, less atomistic model of the metal-insulator transition was formulated by Neville Mott.[3] The Mott model considers the behavior of an electron in a material as a function of the density of all the other valence electrons. We know for a one-electron hydrogen-like atom (H, Na, Cs, etc.) the Schrödinger equation contains a potential energy term:
$V(r) = -(\frac{e^{2}}{4 \pi \varepsilon_{0} r})$
This potential energy function gives rise to familiar ladder of allowed energy levels in the hydrogen atom. However, in a metal, this Coulomb potential must be modified to include the screening of nuclear charge by the other electrons in the solid. In this case there is a screened Coulomb potential:
$V(r) = -(\frac{e^{2}}{4 \pi \varepsilon_{0}r}) exp(-qr)$
where q, which is the inverse of the screening length, is given by:
$q^{2} = 4m_{e}^{2}(\frac{3n}{\pi})^{\frac{1}{3}}(\frac{2\pi}{h})$
Here n is the density of atoms (or valence electrons), me is the electron mass, and h is Planck's constant. At distances much larger than the screening length q-1, the electron no longer "feels" the charge on the nucleus. Mott showed that there is a critical density of electrons nc above which the valence electrons are no longer bound by individual nuclei and are free to roam the crystal. This critical density marks the transition to the metallic state, and is given by the Mott criterion:
$\mathbf{n_{c}^{\frac{1}{3}}a_{H} \approx 0.26}$
In this equation, aH is an effective Bohr radius for the valence electrons in the low-density limit, e.g. the average orbital radius of electrons in the 6s shell of a Cs atom when computing the value for Cs metal.
Solutions of lithium metal in liquid ammonia at low (top, ionic conductor) and higher (bottom, metal) Li concentration. From a video by Joshua Judkins
The important concept from the Mott model is that the metal-insulator transition depends very strongly on the density of valence electrons. This is consistent with the orbital overlap model of Hubbard, but also more general in the sense that it does not depend on a periodic structure of atoms. The Mott model is thus applicable to such diverse systems as metal atoms dissolved in liquid ammonia, metal atoms trapped in frozen gas matrices, and dopants in semiconductors.[4] In some systems, it is possible to continuously tune the density of valence electrons with rather striking results. For example, dissolving alkali metal (Li, Na, ...) in liquid ammonia (bp -33 oC) produces a blue liquid. The solvated alkali cations and negatively charged electrons impart ionic conductivity (as in a salt solution) but not electronic conductivity to the blue liquid ammonia solution. But as the concentration of electrons increases, a reflective, bronze-colored liquid phase forms that floats over the blue phase. This bronze phase is metallic and highly conducting. Eventually, with enough alkali metal added, the entire liquid is converted to the electronically conducting bronze phase.
The electrical switching of VO2 between insulating and metallic phases (see above) can also be rationalized in terms of the Mott transition. Adding more electron density (by chemical or electric field doping) increases the concentration of valence electrons, driving the phase transition to the metallic state.
Thermodynamics and phase transitions
Thermodynamically, the metal-insulator transition is a first-order phase transition. In such a transition, the structure and properties change abruptly (think of the breakfast-to-lunch transition at McDonalds - there is just no way to get pancakes after, or hamburgers before 10:30 AM![5]). Thus in the case of Sn metal, the changes in structure (from four- to eight-coordination) and in electronic conductivity (insulator to metal) occur simultaneously. As in other first order phase transitions such as ice to water to steam, there is a latent heat associated with the transition and a discontinuity in derivative properties such as the heat capacity.
A typical phase diagram for a metal-insulator transition is shown at the right for V2O3. The octahedrally coordinated V3+ ion has a d2 electron count, so there are two unpaired spins per atom, and at low temperature the spins in the lattice order antiferromagnetically. As we learned in Chapter 8, above the Néel temperature an antiferromagnet becomes a paramagnet, which is also a Mott insulator. Increasing the pressure, or doping with electrons (e.g., by substituting some d3 Cr3+ for V3+) pushes the electron density over the Mott transition, the spins pair, and the solid becomes metallic. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/10%3A_Electronic_Properties_of_Materials_-_Superconductors_and_Semiconductors/10.02%3A_Metal-Insulator_Transitions.txt |
Superconductivity refers to the flow of electrical current in a material with zero resistance. Such materials are very important for use in electromagnets, e.g., in magnetic resonance imaging (MRI) and nuclear magnetic resonance (NMR) machines, because once the current starts flowing in the coils of these magnets it doesn't stop. Magnetic levitation using superconductors - which, below a critical field strength, are perfect diamagnets that are not penetrated by magnetic flux lines - is also potentially relevant to future technologies such as magnetically levitated trains.
The phenomenon of superconductivity, first discovered in Hg metal in 1911 by Onnes, continues to be only partially understood. It is of great interest to physicists as a macroscopic quantum phenomenon, and to chemists and materials scientists who try to make better superconductors (especially those that superconduct at higher temperatures) and devices derived from them, such as superconducting quantum interference devices (SQUIDs), which are extremely sensitive magnetometers.
A magnet levitating above a high-temperature superconductor, cooled with liquid nitrogen. Persistent electric current flows on the surface of the superconductor, excluding the magnetic field of the magnet. This current effectively forms an electromagnet that repels the magnet.
Spin pairing and zero resistance
The transition from the metallic to the superconducting state is related to the quantum phenomena of Bose-Einstein condensation and superfluidity. Individual electrons have spin = 1/2, and as such are fermions (particles with half-integer spin). Because of the Pauli exclusion principle, no more than two fermions can occupy the same quantum state (such as an orbital in a molecule or a solid). The familiar consequence of this rule is the aufbau filling of orbitals with spin-paired electrons in each energy level. In contrast, particles with integer spins - which are called bosons - do not have this restriction, and any number of bosons can occupy the same quantized energy level.
Superconductivity occurs when electrons spin pair into so-called Cooper pairs, which can travel through the lattice together. The electrons in a Cooper pair, although spin-paired, have a long-distance relationship: the spatial extent of a Cooper pair is a few nanometers in cuprate superconductors, and up to one micron in low Tc superconductors such as aluminum. Because its overall spin angular momentum is zero, a Cooper pair is a boson. When the temperature is low enough, the Cooper pairs "condense" into the lowest energy level. The second lowest energy level - which is typically a few meV above the ground state - is not accessible to them as long as the energy gap is larger than the thermal energy, kT. The scattering of electrons by the lattice then becomes forbidden by energy conservation because scattering dissipates energy, and the Cooper pairs cannot change their energy state. Thus the resistance (which arises from scattering, as we learned in Ch. 6) drops abruptly to zero below Tc. However, the Cooper pairs can be broken apart when they move fast, and thus superconductors turn back into normal metals (even below Tc) above some critical current density jc. This phenomenon is also related to the critical magnetic field, Hc, that quenches superconductivity.
A trampoline for electrons
What causes electrons, which repel each other because of their negative charge, to pair up and travel together in superconductors? The mechanism - which must involve some kind of attractive interaction between electrons - is well understood for "conventional" superconductors which have relatively low transition temperatures, but is not yet known with certainty for high temperature oxide superconductors. In conventional, or BCS superconductors, the spin pairing is mediated by the lattice as shown in the figure at the left. A strong electron-lattice interaction causes a distortion in the lattice as an electron moves through. This elastic deformation is felt as an attractive force by a second electron moving in the opposite direction. This can be thought of as analogous to the interaction of two people jumping on a trampoline. The weight of the first person on the trampoline creates a "well" that attracts the second one, and they tend to move together (even if they don't like each other). Strange as this interaction seems, it is supported experimentally by isotope effects on Tc and by quantitative predictions of Tc values in conventional superconductors.
Bad metals make good superconductors. All superconductors are "normal" metals - with finite electrical resistance - above their critical transition temperature, Tc. If you ask where in the periodic table one might look for superconductors, the answer is surprising. The most conductive metals (Ag, Au, Cu, Cs, etc.) make the worst superconductors, i.e., they have the lowest superconducting transition temperatures, in many cases below 0.01 K. Conversely "bad" metals, such as niobium alloys, certain copper oxides, KxBa1-xBiO3, MgB2, FeSe, and alkali salts of C60n- anions, can have relatively high transition temperatures.
Timeline of superconducting materials, showing Tc vs. year of discovery.
We observe that most good superconductors appear in composition space very near a metal-insulator transition. In terms of our microscopic picture, orbital overlap in superconductors is poor, just barely enough to make them act as metals (Δ ≈ U) above Tc. In the normal state, superconductors with high Tc - which can be as high as 150 K - are typically "bad" metals. An important characteristic of such metals is that the mean free path of electrons (in the normal state, above Tc) is on the order of the lattice spacing, i.e., only a few Å. In contrast, we learned in Ch. 6 that good metals such as Au, Ag, and Cu have electron mean free paths that are two orders of magnitude longer (ca. 40 nm). In a bad metal, the electron "feels" the lattice rather strongly, whereas in a good metal, the electrons are insensitive to small changes in the distance between metals atoms.
What does the band picture look like for a bad metal? The key point is that because orbital overlap is poor, the metal has a high density of states at the Fermi level. This is a universal property of high temperature superconductors and provides a clue of where to look for new and improved superconducting materials. Recall that transition elements in the middle of the 3d series (Cr, Fe, Co, Ni) were magnetic because of poor orbital overlap and weak d-d bonding. The elements below these - especially Nb, Ta, and W - have just barely enough d-d orbital overlap to be on the metallic side of the metal-insulator transition and to be "bad" metals. Carbides and nitrides of these elements are typically superconducting, with the carbon and nitrogen atoms serving to adjust the valence electron density, as illustrated in the table below.
Generic E vs. DOS for a bad metal.
Compound NbC Mo2N TaC VN NbN TaN Nb3Ge
Tc (K) 11.1 5.0 9.7 7.5 15.2 17.8 22.3
High Tc superconductors
In addition to having weak orbital overlap in the metallic state - which results in a high DOS at EF - high temperature superconductors also typically contain elements in mixed oxidation states (for example, Cu2+/3+ or Bi3+/5+) that are close in energy to the O2-/O- couple in the lattice. At ambient pressure, cuprate superconductors have the highest known Tc values, ranging between about 35 and 150 K. The crystal structures of these materials are almost all variants of the perovskite lattice, as shown at the right for the 1-2-3 superconductor YBa2Cu3O7. An ideal perovskite lattice would have formula ABO3 = A3B3O9. In YBa2Cu3O7, Y and Ba occupy the A cation sites, Cu occupies the B sites, and two of the nine O atoms are missing.
The YBa2Cu3O7 lattice consists of mixed-valent copper(II/III) oxide sheets capped by oxygen atoms to form CuO5 square pyramids. These sheets encapsulate the Y3+ cations. Copper(II) oxide ribbons that share the apical oxygen atoms of the square pyramids run in one direction through the structure. In YBa2Cu3O7 and related materials, one component of the structure (here the Cu-O ribbons) acts as a charge reservoir to control the doping of the planar CuO2 sheets, which are the elements of the structure that carry the supercurrent. Cuprate superconductors with Bi, Tl, or Hg-containing charge reservoir layers and multiple, eclipsed CuO2 sheets in the unit cell tend to have the highest Tc values.
Crystal structure of YBa2Cu3O7 (YBCO), the first superconductor with Tc above the boiling point of liquid nitrogen.
The connection between the metal-insulator transtion and superconductivity is nicely illustrated in the phase diagram of La2-xSrxCuO4, the first cuprate superconductor, which was discovered in 1986 by Georg Bednorz and K. Alex Müller. This compound has a rather simple structure in which rocksalt La(Sr)O layers are intergrown with perovskite La(Sr)CuO3 layers. Undoped La2CuO4 contains only Cu2+ ions and is an antiferromagnetic insulator. As a small amount of Sr2+ is substituted for La3+, some of the Cu2+ is oxidized to Cu3+, and the lattice is doped with holes. As the doping level increases, the antiferromagnetic phase undergoes a first-order phase transition to a "bad" metal, and at slightly higher doping density the superconducting phase appears. The proximity of the superconducting phase to the metal-insulator transition is a hallmark of cuprate superconductors. A maximum Tc of 35K is observed at x = 0.15. Doping at higher levels moves the Fermi level beyond the point of highest DOS in the d-band of Cu and the superconducting phase then gradually disappears. It is interesting to compare this phase diagram with that of V2O3 (above), which also undergoes an antiferromagnetic insulator to "bad" metal transition as it is doped.
Crystal structure and phase diagram of the cuprate superconductor La2-xSrxCuO4. (LSCO)
10.04: Periodic Trends- Metals Semiconductors and Insulators
As we consider periodic trends in the electronic properties of materials, it is important to review some of the key bonding trends we have learned in earlier chapters:
• Going down the periodic table, atoms in solids tend to adopt structures with higher coordination numbers.
• The second row of the periodic table is special, with strong s-p hybridization and π-bonding between atoms.
• Electrons in higher quantum shells are less strongly bound, so the energy difference between bonding and antibonding orbitals becomes smaller for heavier atoms.
We also know that most of the elements in the periodic table are metals, but the elements in the top right corner are insulating under ordinary conditions (1 atm. pressure) and tend to obey the octet rule in their compounds.
At the transition between metals and non-metals in the periodic table we encounter a crossover in electronic properties, as well as in other properties such as the acidity of the oxides (see Ch. 3). The group of elements at the border is loosely referred to as the metalloids. Several of these elements (such as C, Sn, and As) can exist as different allotropes that can be metals, insulators, or something in between.
A more rigorous delineation of the electronic properties of these elements (and of many compounds) can be made by considering their band structures and the temperature dependence of the electronic conductivity. As we have previously discussed, metals have partially filled energy bands, meaning that the Fermi level intersects a partially filled band. With increasing temperature, metals become poorer conductors because lattice vibrations (which are called phonons in the physics literature) scatter the mobile valence electrons. In contrast, semiconductors and insulators, which have filled and empty bands, become better conductors at higher temperature, since some electrons are thermally excited to the lowest empty band. The distinction between insulators and semiconductors is arbitrary, and from the point of view of metal-insulator transitions, all semiconductors are insulators. We typically call an insulator a semiconductor if its band gap (Egap) is less than about 3 eV. A semimetal is a material that has a band gap near zero, examples being single sheets of sp2-bonded carbon (graphene) and elemental Bi. Like a narrow gap semiconductor, a semimetal has higher conductivity at higher temperature. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/10%3A_Electronic_Properties_of_Materials_-_Superconductors_and_Semiconductors/10.03%3A_Superconductors.txt |
Semiconductors, as we noted above, are somewhat arbitrarily defined as insulators with band gap energy < 3.0 eV (~290 kJ/mol). This cutoff is chosen because, as we will see, the conductivity of undoped semiconductors drops off exponentially with the band gap energy and at 3.0 eV it is very low. Also, materials with wider band gaps (e.g. SrTiO3, Egap = 3.2 eV) do not absorb light in the visible part of the spectrum
There are a number of places where we find semiconductors in the periodic table:
• Early transition metal oxides and nitrides, especially those with d0 electron counts such as TiO2, TaON, and WO3
• Oxides of later 3d elements such as Fe2O3, NiO, and Cu2O
• Layered transition metal chalcogenides with d0, d2 and d6 electron counts including TiS2, ZrS2, MoS2, WSe2, and PtS2
• d10 copper and sliver halides, e.g., CuI, AgBr, and AgI
• Zincblende- and wurtzite-structure compounds of the p-block elements, especially those that are isoelectronic with Si or Ge, such as GaAs and CdTe. While these are most common, there are other p-block semiconductors that are not isoelectronic and have different structures, including GaS, PbS, and Se.
A 2" wafer cut from a GaAs single crystal. GaAs, like many p-block semiconductors, has the zincblende structure.
The p-block octet semiconductors are by far the most studied and important for technological applications, and are the ones that we will discuss in detail.
Zincblende- and wurtzite-structure semiconductors have 8 valence electrons per 2 atoms. These combinations include 4-4 (Si, Ge, SiC,…), 3-5 (GaAs, AlSb, InP,…), 2-6 (CdSe, HgTe, ZnO,…), and 1-7 (AgCl, CuBr,…) semiconductors. Other variations that add up to an octet configuration are also possible, such as CuIInIIISe2, which has the chalcopyrite structure, shown at the right.
The chalcopyrite structure is adopted by ABX2 octet semiconductors such as CuIInIIISe2 and CdIISnIVP2. The unit cell is doubled relative to the parent zincblende structure because of the ordered arrangement of cations. Each anion (yellow) is coordinated by two cations of each type (blue and red).
How does the band gap energy vary with composition? There are two important trends
(1) Going down a group in the periodic table, the gap decreases:
C (diamond) > Si > Ge > α-Sn
Egap (eV): 5.4 1.1 0.7 0.0
This trend can be understood by recalling that Egap is related to the energy splitting between bonding and antibonding orbitals. This difference decreases (and bonds become weaker) as the principal quantum number increases.
(2) For isoelectronic compounds, increasing ionicity results in a larger band gap.
Ge < GaAs < ZnSe
0.7 1.4 2.8 eV
Sn < InSb < CdTe < AgI
0.0 0.2 1.6 2.8 eV
This trend can also be understood from a simple MO picture, as we discussed in Ch. 2. As the electronegativity difference Δχ increases, so does the energy difference between bonding and antibonding orbitals.
The band gap is a very important property of a semiconductor because it determines its color and conductivity. Many of the applications of semiconductors are related to band gaps:
• Narrow gap materials (HgxCd1-xTe, VO2, InSb, Bi2Te3) are used as infrared photodetectors and thermoelectrics (which convert heat to electricity).
• Wider gap materials (Si, GaAs, GaP, GaN, CdTe, CuInxGa1-xSe2) are used in electronics, light-emitting diodes, and solar cells.
Color wheel showing the colors and wavelengths of emitted light.
Semiconductor solid solutions such as GaAs1-xPx have band gaps that are intermediate between the end member compounds, in this case GaAs and GaP (both zincblende structure). Often, there is a linear relation between composition and band gap, which is referred to as Vegard's Law. This "law" is often violated in real materials, but nevertheless offers useful guidance for designing materials with specific band gaps. For example, red and orange light-emitting diodes (LED's) are made from solid solutions with compositions of GaP0.40As0.60 and GaP0.65As0.35, respectively. Increasing the mole fraction of the lighter element (P) results in a larger band gap, and thus a higher energy of emitted photons.
Colors of semiconductors
The color of absorbed and emitted light both depend on the band gap of the semiconductor. Visible light covers the range of approximately 390-700 nm, or 1.8-3.1 eV. The color of emitted light from an LED or semiconductor laser corresponds to the band gap energy and can be read off the color wheel shown at the right.
Fe2O3 powder is reddish orange because of its 2.2 eV band gap
The color of absorbed light includes the band gap energy, but also all colors of higher energy (shorter wavelength), because electrons can be excited from the valence band to a range of energies in the conduction band. Thus semiconductors with band gaps in the infrared (e.g., Si, 1.1 eV and GaAs, 1.4 eV) appear black because they absorb all colors of visible light. Wide band gap semiconductors such as TiO2 (3.0 eV) are white because they absorb only in the UV. Fe2O3 has a band gap of 2.2 eV and thus absorbs light with λ < 560 nm. It thus appears reddish-orange (the colors of light reflected from Fe2O3) because it absorbs green, blue, and violet light. Similarly, CdS (Egap = 2.6 eV) is yellow because it absorbs blue and violet light.
Electrons and holes in semiconductors
Pure (undoped) semiconductors can conduct electricity when electrons are promoted, either by heat or light, from the valence band to the conduction band. The promotion of an electron (e-) leaves behind a hole (h+) in the valence band. The hole, which is the absence of an electron in a bonding orbital, is also a mobile charge carrier, but with a positive charge. The motion of holes in the lattice can be pictured as analogous to the movement of an empty seat in a crowded theater. An empty seat in the middle of a row can move to the end of the row (to accommodate a person arriving late to the movie) if everyone moves over by one seat. Because the movement of the hole is in the opposite direction of electron movement, it acts as a positive charge carrier in an electric field.
The opposite process of excitation, which creates an electron-hole pair, is their recombination. When a conduction band electron drops down to recombine with a valence band hole, both are annihilated and energy is released. This release of energy is responsible for the emission of light in LEDs.
An electron-hole pair is created by adding heat or light energy E > Egap to a semiconductor (blue arrow). The electron-hole pair recombines to release energy equal to Egap (red arrow).
At equilibrium, the creation and annihilation of electron-hole pairs proceed at equal rates. This dynamic equilibrium is analogous to the dissociation-association equilibrium of H+ and OH- ions in water. We can write a mass action expression:
$n \times p = K_{eq} = n_{i}^{2}$
where n and p represent the number density of electrons and holes, respectively, in units of cm-3. The intrinsic carrier concentration, ni, is equal to the number density of electrons or holes in an undoped semiconductor, where n = p = ni.
Note the similarity to the equation for water autodissociation:
$[H^{+}][OH^{-}] = K_{w}$
By analogy, we will see that when we increase n (e.g., by doping), p will decrease, and vice-versa, but their product will remain constant at a given temperature.
Temperature dependence of the carrier concentration. Using the equations $K_{eq} = e^{(\frac{- \Delta G^{o}}{RT})}$ and $\Delta G^{o} = \Delta H^{o} - T \Delta S^{o}$, we can write:
$n \times p = n_{i}^{2} = e^{(\frac{\Delta S^{o}} {R})} e^{(\frac{- \Delta H^{o}}{RT})}$
The entropy change for creating electron hole pairs is given by:
$\Delta S^{o} = R ln (N_{V}) + R ln (N_{V}) = R ln (N_{C}N_{V})$
where NV and NC are the effective density of states in the valence and conduction bands, respectively.
and thus we obtain
$n_{i}^{2} = N_{C}N_{V} e^{({- \Delta H^{o}}{RT})}$
Since the volume change is negligible, $\Delta H^{o} \approx \Delta E^{o}$, and therefore $\frac {\Delta H^{o}}{R} \approx \frac{E_{gap}}{k}$, from which we obtain
$n_{i}^{2} = N_{C}N_{V} e^{(\frac{-E_{gap}}{kT})}$
and finally
$\mathbf{n= p = n_{i} = (N_{C}N_{V})^{\frac{1}{2}} e^{(\frac{-E_{gap}}{2kT})}}$
For pure Si (Egap = 1.1 eV) with N ≈ 1022/cm3, we can calculate from this equation a carrier density ni of approximately 1010/cm3 at 300 K. This is about 12 orders of magnitude lower than the valence electron density of Al, the element just to the left of Si in the periodic table. Thus we expect the conductivity of pure semiconductors to be many orders of magnitude lower than those of metals.
Conductivity of intrinsic semiconductors
The conductivity (σ) is the product of the number density of carriers (n or p), their charge (e), and their mobility (µ). Recall from Chapter 6 that µ is the ratio of the carrier drift velocity to the electric field and has units of cm2/Volt-second. Typically electrons and holes have somewhat different mobilities (µe and µh, respectively) so the conductivity is given by:
$\sigma = ne\mu_{e} + pe\mu_{h}$
For either type of charge carrier, we recall from Ch. 6 that the mobility μ is given by:
$\mu = \frac{v_{drift}}{E} = \frac{e\tau}{m}$
where e is the fundamental unit of charge, τ is the scattering time, and m is the effective mass of the charge carrier.
Taking an average of the electron and hole mobilities, and using n = p, we obtain
$\mathbf{\sigma= \sigma_{o} e^{(\frac{-E_{gap}}{2kT})}}, \: where \: \sigma_{o} = 2(N_{C}N_{V})^{\frac{1}{2}}e\mu$
By measuring the conductivity as a function of temperature, it is possible to obtain the activation energy for conduction, which is Egap/2. This kind of plot, which resembles an Arrhenius plot, is shown at the right for three different undoped semiconductors. The slope of the line in each case is -Egap/2k.
Plots of ln(σ) vs. inverse temperature for intrinsic semiconductors Ge (Egap = 0.7 eV), Si (1.1 eV) and GaAs (1.4 eV). The slope of the line is -Egap/2k.
Doping of semiconductors. Almost all applications of semiconductors involve controlled doping, which is the substitution of impurity atoms, into the lattice. Very small amounts of dopants (in the parts-per-million range) dramatically affect the conductivity of semiconductors. For this reason, very pure semiconductor materials that are carefully doped - both in terms of the concentration and spatial distribution of impurity atoms - are needed.
n- and p-type doping. In crystalline Si, each atom has four valence electrons and makes four bonds to its neighbors. This is exactly the right number of electrons to completely fill the valence band of the semiconductor. Introducing a phosphorus atom into the lattice (the positively charged atom in the figure at the right) adds an extra electron, because P has five valence electrons and only needs four to make bonds to its neighbors. The extra electron, at low temperature, is bound to the phosphorus atom in a hydrogen-like molecular orbital that is much larger than the 3s orbital of an isolated P atom because of the high dielectric constant of the semiconductor. In silicon, this "expanded" Bohr radius is about 42 Å, i.e., 80 times larger than in the hydrogen atom. The energy needed to ionize this electron – to allow it to move freely in the lattice - is only about 40–50 meV, which is not much larger the thermal energy (26 meV) at room temperature. Therefore the Fermi level lies just below the conduction band edge, and a large fraction of these extra electrons are promoted to the conduction band at room temperature, leaving behind fixed positive charges on the P atom sites. The crystal is n-doped, meaning that the majority carrier (electron) is negatively charged.
Alternatively, boron can be substituted for silicon in the lattice, resulting in p-type doping, in which the majority carrier (hole) is positively charged. Boron has only three valence electrons, and "borrows" one from the Si lattice, creating a positively charged hole that exists in a large hydrogen-like orbital around the B atom. This hole can become delocalized by promoting an electron from the valence band to fill the localized hole state. Again, this process requires only 40–50 meV, and so at room temperature a large fraction of the holes introduced by boron doping exist in delocalized valence band states. The Fermi level (the electron energy level that has a 50% probability of occupancy at zero temperature) lies just above the valence band edge in a p-type semiconductor.
n- and p-type doping of semiconductors involves substitution of electron donor atoms (light orange) or acceptor atoms (blue) into the lattice. These substitutions introduce extra electrons or holes, respectively, which are easily ionized by thermal energy to become free carriers. The Fermi level of a doped semiconductor is a few tens of mV below the conduction band (n-type) or above the valence band (p-type).
As noted above, the doping of semiconductors dramatically changes their conductivity. For example, the intrinsic carrier concentration in Si at 300 K is about 1010 cm-3. The mass action equilibrium for electrons and holes also applies to doped semiconductors, so we can write:
$n \times p = n_{i}^{2} = 10^{20} cm^{-6} \: at \: 300K$
If we substitute P for Si at the level of one part-per-million, the concentration of electrons is about 1016 cm-3, since there are approximately 1022 Si atoms/cm3 in the crystal. According to the mass action equation, if n = 1016, then p = 104 cm-3. There are three consequences of this calculation:
• The density of carriers in the doped semiconductor (1016 cm-3) is much higher than in the undoped material (~1010 cm-3), so the conductivity is also many orders of magnitude higher.
• The activation energy for conduction is only 40–50 meV, so the conductivity does not change much with temperature (unlike in the intrinsic semiconductor)
• The minority carriers (in this case holes) do not contribute to the conductivity, because their concentration is so much lower than that of the majority carrier (electrons).
Similarly, for p-type materials, the conductivity is dominated by holes, and is also much higher than that of the intrinsic semiconductor.
Chemistry of semiconductor doping. Sometimes it is not immediately obvious what kind of doping (n- or p-type) is induced by "messing up" a semiconductor crystal lattice. In addition to substitution of impurity atoms on normal lattice sites (the examples given above for Si), it is also possible to dope with vacancies - missing atoms - and with interstitials - extra atoms on sites that are not ordinarily occupied. Some simple rules are as follows:
• For substitutions, adding an atom to the right in the periodic table results in n-type doping, and an atom to the left in p-type doping.
For example, when TiO2 is doped with Nb on some of the Ti sites, or with F on O sites, the result is n-type doping. In both cases, the impurity atom has one more valence electron than the atom for which it was substituted. Similarly, substituting a small amount of Zn for Ga in GaAs, or a small amount of Li for Ni in NiO, results in p-type doping.
• Anion vacancies result in n-type doping, and cation vacancies in p-type doping.
Examples are anion vacancies in CdS1-x and WO3-x, both of which give n-type semiconductors, and copper vacancies in Cu1-xO, which gives a p-type semiconductor.
• Interstitial cations (e.g. Li) donate electrons to the lattice resulting in n-type doping. Interstitial anions are rather rare but would result in p-type doping.
Sometimes, there can be both p- and n-type dopants in the same crystal, for example B and P impurities in a Si lattice, or cation and anion vacancies in a metal oxide lattice. In this case, the two kinds of doping compensate each other, and the doping type is determined by the one that is in higher concentration. A dopant can also be present on more than one site. For example, Si can occupy both the Ga and As sites in GaAs, and the two substitutions compensate each other. Si has a slight preference for the Ga site, however, resulting in n-type doping. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/10%3A_Electronic_Properties_of_Materials_-_Superconductors_and_Semiconductors/10.05%3A_Semiconductors-_Band_Gaps_Colors_Conductivity_and_Doping.txt |
Semiconductor p-n junctions are important in many kinds of electronic devices, including diodes, transistors, light-emitting diodes, and photovoltaic cells. To understand the operation of these devices, we first need to look at what happens to electrons and holes when we bring p-type and n-type semiconductors together. At the junction between the two materials, mobile electrons and holes annihilate each other, leaving behind the fixed + and - charges of the electron donor and electron acceptor dopants, respectively. For example, on the n-side of a silicon p-n junction, the positively charged dopants are P+ ions and on the p-side the negatively charged dopants are B-. The presence of these uncompensated electrical charges creates an electric field, the built-in field of the p-n junction. The region that contains these charges (and a very low density of mobile electrons or holes) is called the depletion region.
The electric field, which is created in the depletion region by electron-hole recombination, repels both the electrons (on the n-side) and holes (on the p-side) away from the junction. The concentration gradient of electrons and holes, however, tends to move them in the opposite direction by diffusion. At equilibrium, the flux of mobile carriers is zero because the field-driven migration flux is equal and opposite to the concentration-driven diffusion flux.
When p-type and n-type semiconductors are joined, electrons and holes are annihilated at the interface, leaving a depletion region that contains positively and negatively charged donor and acceptor atoms, respectively. At equilibrium, the Fermi level (EF) is uniform throughout the junction. EF lies just above the valence band on the p-type side of the junction and just below the conduction band on the n-type side.
The width of the depletion layer depends on the screening length in the semiconductor, which in turn depends on the dopant density. At high doping levels, the depletion layer is narrow (tens of nanometers across), whereas at low doping density it can be as thick as 1 µm. The depletion region is the only place where the electric field is nonzero, and the only place where the bands bend. Elsewhere in the semiconductor the field is zero and the bands are flat.
In the middle of p-n junction, the Fermi level energy, EF, is halfway between the valence band, VB, and the conduction band, CB, and the semiconductor is intrinsic (n = p = ni)
10.07: Diodes LEDs and Solar Cells
Diodes are semiconductor devices that allow current to flow in only one direction. Diodes act as rectifiers in electronic circuits, and also as efficient light emitters (in LEDs) and solar cells (in photovoltaics). The basic structure of a diode is a junction between a p-type and an n-type semiconductor, called a p-n junction. Typically, diodes are made from a single semiconductor crystal into which p- and n- dopants are introduced.
Closeup of a diode, showing the square-shaped semiconductor crystal (black object on left) (John Maushammer, Wikipedia, CC-BY-SA)
If the n-side of a diode is biased at positive potential and the p-side is biased negative, electrons are drawn to the n-side and holes to the p-side. This reinforces the built in potential of the p-n junction, the width of the depletion layer increases, and very little current flows. This polarization direction is referred to as "back bias." If the diode is biased the other way, carriers are driven into the junction where they recombine. The electric field is diminished, the bands are flattened, and current flows easily since the applied bias lowers the built-in potential. This is called "forward bias."
Electrons (red) and holes (white) in a forward-biased diode. (S-kei. Wikipedia, CC-BY-SA)
The figure on the left illustrates a forward-biased diode, through which current flows easily. As electrons and holes are driven into the junction (black arrows in lower left figure), they recombine (downward blue arrows), producing light and/or heat. The Fermi level in the diode is indicated as the dotted line. There is a drop in the Fermi level (equal to the applied bias) across the depletion layer. The corresponding diode i-V curve is shown on the right. The current rises exponentially with applied voltage in the forward bias direction, and there is very little leakage current under reverse bias. At very high reverse bias (typically tens of volts) diodes undergo avalanche breakdown and a large reverse current flows.
Diode i-V curve
A light-emitting diode or LED is a kind of diode that converts some of the energy of electron-hole recombination into light. This radiative recombination process always occurs in competition with non-radiative recombination, in which the energy is simply converted to heat. When light is emitted from an LED, the photon energy is equal to the bandgap energy. Because of this, LED lights have pure colors and narrow emission spectra relative to other light sources, such as incandescent and fluorescent lights. LED lights are energy-efficient and thus are typically cool to the touch.
Light-emitting diode (LED). (S-kei. Wikipedia, CC-BY-SA)
Direct-gap semiconductors such as GaAs and GaP have efficient luminescence and are also good light absorbers. In direct gap semiconductors, there is no momentum change involved in electron-hole creation or recombination. That is, the electrons and holes originate at the same value of the momentum wavevector k, which we encountered in Ch. 6. k is related to the momentum (also a vector quantity) by p = hk/2π. In a direct-gap semiconductor, the top of the valence band and the bottom of the conduction band most typically both occur at k = 0. Since the momentum of the photon is close to zero, photon absorption and emission are strongly allowed (and thus kinetically fast). Polar semiconductors such as GaAs, GaN, and CdSe are typically direct-gap materials. Indirect-gap semiconductors such as Si and Ge absorb and emit light very weakly because the valence band maximum and conduction band minimum do not occur at the same point in k-space. This means that a lattice vibration (a phonon) must also be created or annihilated in order to conserve momentum. Since this "three body" (electron, hole, phonon) process has low probability, the radiative recombination of electrons and holes is slow relative to non-radiative decay - the thermalization of electron-hole energy as lattice vibrations - in indirect-gap semiconductors. The momentum selection rule thus prevents light absorption/emission and there are no pure Si LEDs or Si-based lasers.
Prof. Shuji Nakamura holding a blue LED.
While red, orange, yellow, and green LEDs can be fabricated relatively easily from AlP-GaAs solid solutions, it was initially very difficult to fabricate blue LEDs because the best direct gap semiconductor with a bandgap in the right energy range is a nitride, GaN, which is difficult to make and to dope p-type. Working at Nichia Corporation in Japan, Shuji Nakamura succeeded in developing a manufacturable process for p-GaN, which is the basis of the blue LED. Because of the importance of this work in the development of information storage (Blu-Ray technology) and full-spectrum, energy-efficient LED lighting, Nakamura shared the 2014 Nobel Prize in Physics with Isamu Asaki and Hiroshi Amano, both of whom had made earlier contributions to the development of GaN diodes.
A Solar cell, or photovoltaic cell, converts light absorbed in a p-n junction directly to electricity by the photovoltaic effect. Photovoltaics is the field of technology and research related to the development of solar cells for conversion of solar energy to electricity. Sometimes the term solar cell is reserved for devices intended specifically to capture energy from sunlight, whereas the term photovoltaic cell is used when the light source is unspecified.
Photovoltaic effect in a semiconductor p-n junction. (S-kei. Wikipedia, CC-BY-SA)
Photocurrent in p-n junction solar cells flows in the diode reverse bias direction. In the dark, the solar cell simply acts as a diode. In the light, the photocurrent can be thought of as a constant current source, which is added to the i-V characteristic of the diode. The relationship between the dark and light current in a photovoltaic cell is shown in the diagram at the left.
Current-voltage characteristic of a solar cell in the dark and under illumination with band gap light. The short-circuit photocurrent is indicated as isc, and the open-circuit photovoltage is Vphoto. The maximum power generated by the solar cell is determined by the area of the orange box.
The built-in electric field of the p-n junction separates e- h+ pairs that are formed by absorption of bandgap light in the depletion region. The electrons flow downhill, towards the n-type side of the junction, the holes flow uphill towards the p-side. If hν ≥ Egap, light can be absorbed by promoting an electron from the valence band to the conduction band. Any excess energy is rapidly thermalized. Light with hν > Eg thus can store only Eg worth of energy in an e- h+ pair. If light is absorbed outside of depletion region, i.e., on the n- or p-side of the junction where there is no electric field, minority carriers must diffuse into the junction in order to be collected. This process occurs in competition with electron-hole recombination. Because impurity atoms and lattice defects make efficient recombination centers, semiconductors used in solar cells (especially indirect-gap materials such as Si, which must be relatively thick in order to absorb most of the solar spectrum) must be very pure. Most of the cost of silicon solar cells is associated with the process of purifying elemental silicon and growing large single crystals from the melt.
In the photodiode i-V curve above, Vphoto is typically only about 70% of the bandgap energy Egap. The photocurrent is limited by the photon flux, the recombination rate, and the re-emission of absorbed light.[6] The area of the orange rectangle indicates the power generated by the solar cell, which can be calculated as P = i x V. In good single crystal or polycrystalline solar cells made of Si, GaAs, CdTe, CuInxGa1-xSe2, or (CH3NH3)PbI3 the quantum yield (the ratio of short circuit photocurrent to photon flux) is close to unity.
The equivalent circuit of a p-n junction solar cell, which results in the "light" i-V curve shown in the figure above. The solar cell is effectively a diode with a reverse-bias current source provided by light-generated electrons and holes. The shunt resistance (Rsh) in the equivalent circuit represents parasitic electron-hole recombination. A high shunt resistance (low recombination rate) and low series resistance (Rs) are needed for high solar cell efficiency.
Solar cells have many current applications. Individual cells are used for powering small devices such as electronic calculators. Photovoltaic arrays generate a form of renewable electricity, particularly useful in situations where electrical power from the grid is unavailable such as in remote area power systems, Earth-orbiting satellites and space probes, remote radio-telephones and water pumping applications. Photovoltaic electricity is also increasingly deployed in grid-tied electrical systems.
The cost of installed photovoltaics (calculated on a per-watt basis) has dropped over the past decade at a rate of about 13% per year, and has already reached grid parity in Germany and a number of other countries.[7] Photovoltaic grid parity is anticipated in U.S. power markets in the 2020 timeframe.[8] A major driver in the progressively lower cost of photovoltaic power is the steadily increasing efficiency of solar cells, which is shown in the graphic at the right. Higher efficiency solar cells require less area to deliver the same amount of power, and this lowers the "balance of system" costs such as wiring, roof mounting, etc., which scale as the area of the solar panels. Progress towards higher efficiency reflects improved processes for making photovoltaic materials such as silicon and gallium arsenide, as well as the discovery of new materials. Silicon solar cells are a mature technology, so they are now in the flat part of the learning curve and are approaching their maximum theoretical efficiencies. Newer technologies such as organic photovoltaics, quantum dot solar cells, and lead halide perovskite cells are still in the rising part of the learning curve.
Reported timeline of solar cell energy conversion efficiencies since 1976 (National Renewable Energy Laboratory)
A field effect transistor (FET) is a transistor that uses an electric field to control the width of a conducting channel and thus the current in a semiconductor material. It is classified as unipolar transistor, in contrast to bipolar transistors.
Field effect transistors function as current amplifiers. The typical structure of Si-based FETs is one in which two n-type regions (the source and the drain) are separated by a p-type region. An oxide insulator over the p-type region separates a metal gate lead from the semiconductor. This structure is called a metal-oxide-semiconductor FET (or MOSFET). When voltage is applied between source and drain, current cannot flow because either the n-p or the p-n junction is back-biased. When a positive potential is applied to the gate, however, electrons are driven towards the gate, and locally the semiconductor is "inverted" to n-type. Then the current flows easily between the n-type source and drain through the n-channel. The current flow between the source and drain is many times larger than the current through the gate, and thus the FET can act as an amplifier. Current flow can also represent a logical "1," so FETs are also used in digital logic.
Cross section of an n-type MOSFET
In electronic devices such as microprocessors, field-effect transistors are kept in the off-state most of the time in order to minimize background current and power consumption. The FET shown above, which has n-type source and drain regions, is called an NMOS transistor. In a PMOS transistor, the source and drain regions are p-type and the gate is n-type. In CMOS (complementary metal-oxide semiconductor) integrated circuits, both NMOS and PMOS transistors are used. CMOS circuits are constructed in such a way that all PMOS transistors must have either an input from the voltage source or from another PMOS transistor. Similarly, all NMOS transistors must have either an input from ground or from another NMOS transistor. This arrangement results in low static power consumption.
Transistors are most useful in the range of gate voltage (indicated by the red circle in the figure at the left) where the source-drain current changes rapidly. In this region it is possible to effect a large change in current between source and drain when a small signal is applied to the gate. An important figure of merit for FETs is the subthreshold slope, which is the slope a plot of log(current) vs. Vgate. An ideal subthreshold slope is one decade of current per 60 mV of gate bias. Typically, a decade change in source-drain current can be achieved with a change in gate voltage of ~70 mV. The performance of FETs as switches and amplifiers is limited by the subthreshold slope, which in turn is limited by the capacitance of the gate. It is desirable to have a very high gate capacitance, which requires a thin insulating oxide, but also to have a small leakage current, which requires a thick oxide. A current challenge in the semiconductor industry is to continue to scale FETs to even smaller nanoscale dimensions while maintaining acceptable values of these parameters. This is being done by developing new gate insulator materials that have higher dielectric constants than silicon oxide and do not undergo redox reactions with silicon or with metal gate leads. Only a few known materials (such as hafnium oxynitride and hafnium silicates) currently meet these stringent requirements. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/10%3A_Electronic_Properties_of_Materials_-_Superconductors_and_Semiconductors/10.06%3A_Semiconductor_p-n_Junctions.txt |
Amorphous semiconductors are disordered or glassy forms of crystalline semiconductor materials. Like non-conducting glasses, they are network structures with primarily covalent bonding. Crystalline silicon, which has the diamond structure, is an ordered arrangement of fused six-membered silicon rings, all in the "chair" conformation, as we saw in Ch. 8. The local bonding environment of the silicon atoms is tetrahedral. The silicon atoms in amorphous silicon (a-Si) are also predominantly tetrahedrally coordinated, but there is no long-range order in the structure. In addition to six-membered rings, there are five- and seven-membered rings, as well as some "dangling bond" sites in which Si atoms have only three nearest neighbors.
Schematic illustration of the structures of crystalline silicon (left), amorphous silicon (middle), and amorphous hydrogenated silicon (right)
Two of the most widely studied amorphous semiconductors are a-Si and amorphous selenium, a-Se. Si and Se can both be made in glassy form, usually by sputtering or evaporation at relatively low temperature. In a-Se, as in a-Si, locally, most of the atoms have their "normal" valence, but there are many defects and irregularities in the structure. Dangling bonds in amorphous semiconductors have orbital energies in the middle of gap, and electrons in these states are effectively non-bonding. Because these dangling bond sites are far apart from each other, there is little orbital overlap between them, and they also exist over a range of energies. Electrons in these mid-gap states are therefore localized, a phenomenon known as Anderson localizaton. Amorphous Si is insulating because electrons the Fermi level (in the middle of the gap) are not mobile in the lattice. These localized states create a mobility gap, and only electrons in states that are strongly bonding or antibonding are delocalized. Therefore, unmodified a-Si is not very useful as a semiconductor. However, by hydrogenating the material as it is formed (typically in a plasma of H atoms), the under-coordinated Si atoms are bonded to hydrogen atoms. This generates filled bonding and empty antibonding orbitals, the energies of which are outside the mobility gap. Hydrogenation thus lowers the density of states in the mobility gap. Hydrogenated amorphous silicon (a-Si:H) is insulating in the dark, but is a good photoconductor because light absorption creates electrons and holes in mobile states that are outside the mobility gap.
Energy vs. DOS for an amorphous semiconductor. Disorder and dangling bonds result in localized mid-gap states.
The photoconductivity of amorphous Se is exploited in xerography. A conductive drum coated with a-Se, which is insulating, is charged with static electricity by corona discharge from a wire. When the drum is exposed to a pattern of light and dark (the image to be duplicated), the illuminated a-Se areas become conductive and the static charge is dissipated from those parts of the drum. Carbon-containing toner particles then adhere via static charge to the areas that were not exposed to light, and are transferred and bonded to paper to make the copy. The speed of the process and the high resolution of pattern transfer depend on the very low conductivity of a-Se in the dark and its high conductivity under illumination.
Charging of amorphous Se and pattern transfer in the xerographic cycle.
Amorphous hydrogenated Si is used in inexpensive thin film solar cells. The mobility gap is about 1.7 eV, which is larger than the bandgap crystalline of Si (1.1 eV). a-Si:H is a direct-gap material, and therefore thin films are good light absorbers. a-Si:H solar cells can be vapor-deposited in large-area sheets. p+Si-a-Si:H-n+Si cells have around 10% power conversion efficiency. However amorphous Si solar cells gradually lose efficiency as they are exposed to light. The mechanism of this efficiency loss, called the Staebler-Wronski effect,[9], involves photogenerated electron-hole pairs which have sufficient energy to cause chemical changes in the material. While the exact mechanism is still unclear, it has been proposed that the energy of electron-hole recombination breaks a weak Si-Si bond, and that one of the resulting dangling bonds abstracts a H atom, leaving a passivated Si-H center and a permanent dangling bond. The effect is minimized by hydrogenating a-Si and can be partially reversed by annealing.
A calculator that runs on solar and battery power
Thin layers of amorphous silicon are used in conjunction with crystalline silicon in heterojunction intrinsic thin-layer (HIT) solar cells.[10] Because the mobility gap of a-Si is wider than the bandgap of c-Si, there is a potential energy barrier at the amorphous-crystalline interface that reflects electrons and holes away from that interface. At the p+ contact, only holes can tunnel through the barrier, whereas only electrons can tunnel through the barrier to the n+ contact. The passivation of surface defects that are sites of electron-hole recombination prevents a major loss mechanism in solar cells, increasing both the photovoltage and the photocurrent relative to conventional c-Si p-n junction cells. Panasonic and Sanyo have announced the production of HIT cells with power conversion efficiencies as high as 23%.
Layered structure of a HIT solar cell. The layers are not drawn to scale. A thick crystalline n-silicon layer is the light absorber, and photogenerated holes, which are the minority carriers, are reflected away from the aluminum back contact by the thin intrinsic a-Si layer there. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/10%3A_Electronic_Properties_of_Materials_-_Superconductors_and_Semiconductors/10.08%3A_Amorphous_Semiconductors.txt |
• How is magnetic ordering in the 3d transition metals (Fe, Co, Ni) and the absence of magnetism in the elements just below them (Ru, Ir, Pd) related to the metal-insulator transition?
• Why are good metals bad superconductors and vice-versa?
• Discuss why semiconducting oxides of early transition metals such as TiO2 and Nb2O5 can be doped n-type but not p-type. Conversely, semiconducting late transition metal oxides such as NiO and Cu2O can be doped p-type but not n-type.
10.10: Problems
1. The structure of a high temperature superconductor containing barium, europium, copper, and oxygen is shown below. What are the coordination environments of Cu in this compound? This structure is actually closely related to perovskite, ABO3. Explain the relationship between this structure and the ideal perovskite structure.
2. VO2 can exist in insulating or metallic form, depending on temperature and pressure. Which form would be stabilized by increasing the temperature? Explain your answer.
3. Explain briefly how and why the bandgaps for octet p-block semiconductors vary (1) with the average principal quantum number, and (2) with the electronegativity difference between anion and cation.
4. Indicate the type of conduction (n or p) in the following: (a) Se-doped GaAs, (b) InAs1-x, where x << 1, (c) Li0.05Ni0.95O, (d) LixWO3, where x << 1.
5. The structure of copper indium selenide, a semiconductor used in thin film solar cells, is shown below in sections.
(a) What is the stoichiometry of the compound?
(b) What kind of doping (n or p) will occur if a small amount of iodine is substituted for selenium?
(c) What kind of doping (n or p) will occur if a small fraction of the indium sites are occupied by copper atoms?
6. Using 1 eV = 1240 nm, predict the colors of anatase TiO2 (Eg = 3.1 eV), SiC (2.0 eV), ZnSnP2 (1.7 eV), ZnGeP2 (1.9 eV), and InP (1.27 eV).
7. The conductivity of a certain intrinsic (undoped) semiconductor increases by a factor of two when the temperature is raised from 300 to 330 K. What is the bandgap (in eV)? R = 8.314 J/mol-K, 1 eV/atom = 96.52 kJ/mole.
8. Pure Ge is much more conductive than pure Si. Given their bandgaps (0.74 and 1.15 eV, respectively), estimate the ratio of their conductivities at room temperature.
9. The figure below illustrates the trends in conductivity vs. inverse temperature for Si, Ge, and As-doped Ge. Identify lines (i), (ii), and (iii) with the appropriate materials. Explain why the slope of line (i) is close to zero.
10. Sketch a silicon p-n junction, showing the depletion region, band bending, and the Fermi level in the absence of light or applied potential. In the dark, the p-n junction acts as a rectifier. (a) Which way do electrons and holes flow most easily in the dark? (b) Does the built in electric field increase or decrease under forward bias? (c) In the light, the junction acts as a photodiode. In this case, under short circuit conditions, do electrons flow in the same direction or in the opposite direction as in (a)? Explain.
10.11: References
1. Hubbard, J. (1963). "Electron Correlations in Narrow Energy Bands". Proceedings of the Royal Society of London 276 (1365): 238–257. doi:10.1098/rspa.1963.0204. Bibcode: 1963RSPSA.276..238H.
2. T. Mizokawa, Metal-insulator transitions: orbital control, Nature Physics 9, 612–613 (2013), doi: doi:10.1038/nphys2769
3. N. F. Mott, (1961) "The Transition to the Metallic State," Phil. Mag. 6, 287. DOI: 10.1080/14786436108243318.
4. P. P. Edwards and M. J. Sienko (1982) "The Transition to the Metallic State," Acc. Chem. Res. 15, 87-93. DOI: 10.1021/ar00075a004
5. Recent data have disproven this assertion; MacDonalds has finally responded to public opinion and is offering breakfast after 10:30 AM. But the laws of thermodynamics remain immutable and eternal
6. E. Yablonovitch, O. Miller, and S. Kurtz, "Strong Internal and External Luminescence as Solar Cells Approach the Shockley–Queisser Limit," IEEE Journal of Photovoltaics, vol. 2, no. 3, pp. 303-311, July 2012.
7. "Recent facts about photovoltaics in Germany". Fraunhofer ISE. 7 January 2015. Retrieved 17 February 2015.
8. Energy Information Administration, (November 2010). Levelized Cost of New Generation Resources in the Annual Energy Outlook 2011.
9. Staebler, D. L. and Wronski, C. R. Optically induced conductivity changes in discharge-produced hydrogenated amorphous silicon. J. Appl. Physics. 51(6), June 1980.
10. Mishima, T., Taguchi, M., Sakata, H., Maruyama, E., 2011. Development status of high efficiency HIT solar cells. Sol. Energy Mater. Sol. Cell. 95, 18–21. doi:10.1016/j.solmat.2010.04.030 | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/10%3A_Electronic_Properties_of_Materials_-_Superconductors_and_Semiconductors/10.09%3A_Discussion_Questions.txt |
Learning Objectives
• Understand the physical basis of mesoscopic behavior in nanoscale semiconducting and magnetic particles.
• Describe how the particle-in-a-box equation applies to electrons in quantum wells.
• Use the Brus formula to calculate the band gap energy of nanoscale semiconductor particles.
• Use the surface energy concept to calculate changes in the melting point and vapor pressure of nanoparticles.
• Describe the methods used to make semiconducting and metal nanocrystals of uniform size.
• Explain the origin of the localized surface plasmon resonance effect in metal nanoparticles.
• Describe the emerging analytical and biomedical applications of metal nanoparticles.
Nanomaterials describe materials of which a single unit is sized (in at least one dimension) between 1 and 1000 nanometers, but is usually 1—100 nm. Nanomaterials research takes a materials science-based approach to nanotechnology, leveraging advances in materials metrology and synthesis which have been developed in support of microfabrication research. Materials with structure at the nanoscale often have unique optical, electronic, or mechanical properties. In this chapter we will learn about the basic science of nanomaterials, i.e., what it is about their size that makes them different.
11: Basic Science of Nanomaterials
What do stained glass, sunscreen, magnetic hard drives, heterogeneous catalysts, consumer electronics, stain-resistant clothing, self-cleaning glass, and medical diagnostics all have in common? All of them derive some special property and utility from nanoscale materials: ordinary elements and inorganic compounds such as gold, silver, TiO2, chromium, SiO2, and silicon that acquire different properties when their characteristic dimensions are somewhere between 1 and 100 nm. In this chapter we will learn about the basic science of nanomaterials, i.e., what it is about their size that makes them different.
11.02: Physics and Length Scales- Cavity Laser Coulomb Blockade Nanoscale Magnets
The special properties of nanomaterials do not derive from different laws of physics, which are the same for objects large and small. For example, Newton's second law ($F = ma$), Coulomb's law ($E= \frac{q_{1}q_{2}}{4 \pi \varepsilon_{0}r}$), and the laws of energy and momentum conservation are the same for buckyballs (C60) and full-size soccer balls. Nevertheless, the physics of electrons, atoms, and photons naturally produce characteristic length scales, some of which we have already seen. For example, in Chapter 6 we discovered that the mean free path of an electron in a good metal is about 40 nm. In Chapter 10, we learned that the Bohr radius of an electron or hole in doped Si is about 4 nm, and that the coherence length of Cooper pairs in semiconductors is somewhere between a few nm and 1 µm. When objects become small relative to these characteristic lengths, their physical properties change in interesting ways. Materials that exist at the relevant length scale are called mesoscopic (meso = "between," scopic = "size") meaning that they cross over from one kind of behavior - the bulk behavior of large objects - to another. This length scale is different for different kinds of properties, but for many it happens between 1 and 100 nm. We illustrate this point with a few examples.
The cavity laser
A vertical cavity surface-emitting laser (or VCSEL) is a semiconductor-based device that emits light in the vertical direction relative to the plane of the chip. These devices are being developed and used for high power applications such as laser surgery, infrared illumination for military surveillance, and laser cutting tools. The basic design of a VCSEL is shown at the right. The VCSEL is basically a light-emitting diode, incorporating p-type and n-type regions of the III-V semiconductor (Al,Ga)As. However, it has two special features. First, the refractive index of the semiconductor is modulated above and below the junction to make Bragg mirrors. These mirrors reflect light emitted in the junction, so that the photon density becomes very high there, a necessary condition for stimulated emission and lasing. The Bragg mirror stack is asymmetric (thinner on the bottom), so that light can escape from the junction in one direction only. Second, the junction itself consists of a thin "quantum well" layer of (In,Ga)As, a III-V semiconductor with a smaller band gap than the surrounding (Al,Ga)As layers.
A VCSEL device structure. This is a bottom-emitting multiple-quantum-well VCSEL.
The quantum well structure, the band diagram of which is illustrated at the left, is the nanoscale part of the laser. The energies of the conduction and valence bands of (Al,Ga)As flank those of the thin (In,Ga)As layer. Therefore, electrons and holes injected into that layer cannot escape: electrons in (In,Ga)As do not have enough energy to climb "up" to the conduction band of (Al,Ga)As and holes cannot climb "down." Electrons confined to such a small well behave as a particle-in-a-box (as we learned in the context of electrides in Chapter 9).
The electron has a kinetic energy defined by the equation:
$KE= \frac{h^{2}n^{2}}{8mL^{2}}$
We can calculate the energy difference between the lowest (n=1) and next lowest (n=2) levels, which is inversely proportional to the square of the thickness (L) of the (In,Ga)As layer. In this calculation we need to use the effective mass of the electron in (In,Ga)As, which is about 7% of the electron rest mass[1]. With an 8 nm thick layer, this energy is:
$E = \frac{(2^{2}-1^{2})(6.626 \cdot 10^{-34} Js)^{2}}{8 \cdot 0.07 \cdot (9.1 \cdot 10^{-31} kg)(8 \cdot 10^{-9} m)^{2}} = 4.0 \cdot 10^{-20} J = 0.25eV$
The VCSEL cavity will thus have a resonant energy of 0.25 eV and emit photons at this energy in the infrared (λ ≈ 5000 nm). Note that because of the inverse square dependence of the cavity energy on layer thickness, lasers based on this design can only function at nanoscale dimensions. When the cavity is three times thicker, its resonant energy becomes comparable to the thermal energy at room temperature (kT = 0.026 eV), and the lasing effect is thermally "washed out."
Coulomb blockade
A capacitor is a (macroscopic) device that stores electrical charge. The basic structure of a capacitor is shown at the right. When a voltage is applied to such a device, it develops a charge (± Q) on the two plates that is proportional to the voltage:
$C= \frac{Q}{V}$
Basic design of a parallel plate capacitor of area A and with a dielectric thickness of d
The magnitude of the capacitance C is determined by the permittivity ε and the dimensions of the dielectric layer, A and d.
$C= \frac{\varepsilon A}{d}$
We can also calculate the work done in charging the capacitor up (i.e., the energy stored by charging the capacitor) by integrating the voltage times the charge:
$E = int^{Q}_{0} V(q)dq = \int^{Q}_{0} \frac{q}{C} dq = \frac{1}{2} \frac{Q^{2}}{C} = \frac{1}{2}CV^{2} = \frac{1}{2}VQ$
Now it is interesting to ask, what happens to a capacitor when we make it very small? This is of particular interest in a device called a single electron transistor, a schematic diagram of which is shown at the right. The metallic gate lead is separated from a "quantum dot," which can be a metal or semiconductor particle, by a thin dielectric layer. This metal-dielectric-metal sandwich acts as a capacitor, and from the equation above, the energy needed to charge it by a single electron (Q = e) is:
$E= \frac{e^{2}}{2C}$
where e is the charge of the electron, 1.602 x 10-19 Coulomb. If the gate width and lateral dimensions are very small - say 2 nm as is readily achievable in self-assembled Coulomb blockade devices[2][3] - then for a typical insulating dielectric, a voltage of about 200 mV is needed to charge the quantum dot by a single electron. Again, this effect is unique to the nanoscale, because a 10 times larger device area would make the single-electron charging voltage about 20 mV, which is smaller than the thermal energy kT (26 meV). Thus for devices larger than about 5-6 nm, individual electron charging events are washed out at room temperature by thermal fluctuations.
Schematic of a single-electron transistor.
How can a nanoscale capacitor like this act as a transistor, which functions as a switch in an electrical circuit? The effect comes from the mutual repulsion of electrons. An electron on the quantum dot repels any other electron that would be forced onto it by applying a small voltage between the source and the drain. Hence the conductance of the quantum dot is very low at a gate bias of zero volts, or at any gate bias (200, 400, 600 mV...) that places an integer number (1, 2, 3,...) of electrons on the dot. But halfway in between these voltages (e.g., at 100, 300, 500 mV) the energy is the same whether there are n or n+1 electrons on the dot. This means that electrons can hop on and off without changing their energy, i.e., they can tunnel through the dot from source to drain. This effect gives peaks in the conductance of the dot at regular steps in the gate voltage. In effect, the gate can act as a switch, as in a conventional field-effect transistor. Single-electron transistors are being researched as ultra-sensitive electrometers and single-molecule chemical sensors, since a tiny change in the electrostatic environment of the dot can switch the device on or off.
Nanoscale magnets
Ferro- and ferrimagnetic materials such as iron and chromium oxide are used for digitial storage of information in hard disks. The individual memory bits, which can be oriented perpendicular or parallel to the plane of the disk as shown at the right, store a logical "0" or "1" depending on the orientation of their magnetic dipole. To be useful, this information must be non-volatile, i.e., the magnetic bit must retain its polarization in the absence of an applied field from the read/write head.
Longitudinal and perpendicular recording, two types of writing heads on a hard disk.
The storage density of such magnetic memories is impressive. A 2.5" hard drive can now store 1 TB of information, using rod-shaped magnetic grains that are approximately 0.5 µm long. We now have good synthetic methods for making these same materials as crystals with dimensions of only a few nanometers. Why aren't those nanocrystals used to make even more dense memory disks?
The reason is that the energy needed to flip the magnetization (i.e., to turn a "0" into a "1" and vice-versa) is strongly size-dependent. For a ferro- or ferrimagnet this energy is equal to Mr3, where M is the magnetic energy per unit volume and r is the characteristic dimension (e.g., the length of the edge of a cube, or the diameter of a sphere) of the magnetic grain. For typical materials such as iron, this energy becomes comparable to kT when r is about 3-5 nm. Such small particles are superparamagnetic, meaning that they still have a large magnetic moment because of the ordering of their spins, but they do not retain a permanent polarization in the absence of an applied magnetic field. Superparamagnetic particles are thus not useful for magnetic memories, but they are interesting and practical in other ways, for example in ferrofluids, magnetic resonance imaging (MRI), and some emerging medical diagnostic and therapeutic applications.
A ferrofluid containing superparamagnetic Fe3O4 nanoparticles, which are coated with oleic acid and suspended in oil, in the field of a strong permanent magnet.
In these three illustrative examples (involving light emission, electronic conduction, and magnetic behavior), the transition to new properties involves a crossover in which the characteristic energy of the system is comparable to the thermal energy kT. It just so happens that for many physical phenomena, this crossover occurs on the length scale of nanometers. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/11%3A_Basic_Science_of_Nanomaterials/11.01%3A_Prelude_to_Basic_Science_of_Nanomaterials.txt |
One of the most fascinating and well-studied mesoscopic effects occurs with semiconductor particles of various shapes when one or more of their dimensions is in the range of a few nanometers. These so-called "quantum dots" (0D), "quantum rods" (1D) and "nanosheets" (2D) acquire striking new electronic and optical properties.
Atomic resolution image of a CdSe nanoparticle.
The synthesis of semiconductor quantum dots, which is discussed in more detail below, is sufficiently well controlled to give essentially perfect crystals of a few thousand atoms, something that is statistically impossible in a macroscopic crystal. An image of such a crystal is shown at the right. In the synthesis, these crystals can be capped with an epitaxial layer of ligands or by a shell of a wider bandgap semiconductor (such as ZnS in the case of CdSe), so that the inner core is effectively a quantum well. Electron-hole pairs formed by excitation of the core semiconductor are confined there and cannot reach the external surface of the particle, where they might otherwise be trapped and recombine thermally. Consequently, the quantum yield for bandgap emission of semiconductor quantum dots is typically high, giving rise to the bright emission colors shown at the right for CdSe particles of different sizes. Because of their strong and narrow emission bands, quantum dots are of interest as luminescent tags for biological imaging applications, and also as light absorbers and emitters in solar cells and LEDs.
Emission colors of CdSe nanoparticles of different sizes. Smaller particles emit blue light because the exciton energy increases as the size decreases.
The size-dependence of the emission color comes primarily from a particle-in-a-box effect. The electron and hole that are created when the quantum dot absorbs light are bound together as an exciton by the confines of the "box". Louis Brus used first-order perturbation theory to determine that the bandgap of a semiconductor quantum dot is given approximately by:[4]
$E_{gap} = E_{gap, bulk} + \frac{h^{2}}{8 \mu R^{2}} - \frac{1.8e^{2}}{4R \pi \varepsilon \varepsilon_{0}} + \dots$
where R is the particle radius, µ is the electron-hole reduced mass (1/µ = 1/me* + 1/mh*), me* and mh* are the electron and hole effective masses, and ε is the dielectric constant of the semiconductor. In this equation, the first term after the bulk bandgap is the kinetic energy due to confinement of the exciton, and the second term represents the electrostatic attractive energy between the confined electron and hole. Because the energy is a function of R2, it can be widely tuned across the visible spectrum by changing the size of the quantum dot.
11.04: Synthesis of Semiconductor Nanocrystals
Early work on the quantum size effect in semiconductor nanoparticles used simple metathesis reactions in the synthesis. For example, CdSe and PbS can be precipitated at ambient temperature by the reactions:
\(\ce{CdCl2_{(aq)} + H2Se_{(g)} = CdSe_{(s)} + 2HCl_{(aq)}}\)
\(\ce{Pb(NO3)2 + H2S_{(g)} = PbS_{(s)} + 2HNO3_{(aq)}}\)
Complete atomistic model of a 5 nm diameter colloidal lead sulfide nanoparticle with surface passivation by oleic acid, oleyl and hydroxyl groups.
The growth of the particles was restricted by carrying out these reactions in different matrices, such as in polymer films or the silicate cages of zeolites, and capping ligands were also sometimes used to limit particle growth. While these reactions did produce nanoparticles, in general a broad distribution of particle sizes was obtained. The particles were also unstable to Ostwald ripening - in which large particles grow at the expense of smaller ones in order to minimize the total surface energy - because of the reversibility of the acid-forming synthetic reactions in aqueous media. The lack of good samples prevented detailed studies and the development of applications for semiconductor quantum dots.
A very important development in nanoparticle synthesis came in the early 1990's, when Murray, Norris, and Bawendi introduced the first non-aqueous, controlled growth process for II-VI semiconductor quantum dots.[5] The keys to this synthesis were (1) to use non-aqueous solvents and capping ligands to stabilize the products against ripening, (2) to carry out the reaction at high temperature to ensure good crystallinity, and (3) to separate the steps of particle nucleation and growth, and thereby obtain particles of uniform size. This procedure is illustrated below:
The synthesis is carried out in a coordinating, high boiling solvent that is a mixture of trioctylphosphine (TOP) and trioctylphosphine oxide (TOPO). In early experiments, organometallic cadmium compounds such as diethylcadmium were used as the metal source, but it was later found that these highly toxic and pyrophoric compounds could be replaced by CdO. At the start of the reaction, a selenium source, typically bis(trimethylsilyl)selenium, [(CH3)3Si]2Se, dissolved in TOP is rapidly injected into the hot (350 °C) reaction mixture. The reaction causes a rapid burst of nanoparticle nucleation, but the temperature also drops as the cold solvent is injected and so the nucleation event ends quickly. The cooled solution now contains nanocrystal seeds. It is supersaturated in TOPO-Cd and TOP-Se, but particle growth proceeds slowly until the solution is heated again to the growth temperature, ca. 250°C. Particle growth and size-focusing occurs because small particles require less added material to grow by an amount ΔR than larger particles. This is because the volume of an added shell around a spherical seed is 4πR2ΔR, so for larger R, ΔR is smaller. Very narrow particle size distributions can be obtained under conditions of high supersaturation, where the rate of nanoparticle growth is fast relative to particle dissolution and Ostwald ripening.[6] The size distribution can then be narrowed further by adding a non-solvent such as hexane to the cooled reaction mixture. The largest particles precipitate first, followed by smaller particles. Because the nanoparticles are capped with a ligand shell of TOP, they can be re-suspended in organic solvents once they are size-separated.
The high temperature synthesis of semiconductor quantum dots has been applied to a broad variety of materials including II-VI, III-V, and IV-VI semiconductors. Monodisperse nanoparticles of controlled shapes can be made by variants of this method. For example, it is possible to adjust the conditions so that CdSe nucleates in the zincblende polymorph as tetrahedrally shaped seeds, and then grow polar wurtzite "arms" onto each triangular face, resulting in nanocrystal tetrapods. Numerous other nanocrystal shapes such as rods, arrowheads, rice (tapered rods), and polar structures such as Janus rods can be made by variants of this technique. These shape-control strategies often involve the use of ligands that adsorb specifically to certain crystal faces and inhibit their growth. For example, hexylphosphonic acid ligands adsorb selectively to Cd-rich crystal faces and thus lead to the growth of prismatic wurtzite-phase CdSe nanocrytals.
A second widely used method of semiconductor nanocrystal synthesis involves growth from molecular precursors and molten metal droplets, as shown in the figure at the right. The vapor-liquid-solid (VLS) and related solution-liquid-solid (SLS) growth processes rely on the fact that semiconductors such as Si, Ge, GaAs, InP, and others are soluble at high temperatures in liquid metals such as Au and Cu. The catalytic reduction of a molecule such as silicon tetrachloride (SiCl4) at the surface of a gold nanocrystal liberates HCl gas and creates a solid solution of Si in Au. The presence of Si lowers the melting point of Au, and as more SiCl4 reacts, a liquid eutectic droplet of the Si-Au alloy is formed. When this droplet becomes supersaturated in Si, a silicon nanocrystal nucleates and grows. This reaction can be performed on the surface of a macroscopic Si crystal, in which case nanocrystal "whiskers" grow from the surface, typically as single crystals and with an epitaxial orientation that is determined by the Si crystal face of the substrate. The diameter of the whiskers is controlled by the radius of the Au drops, which can be as small as a few nanometers and as large as several microns. Using this technique, "forests" of nanowires or microwires can be grown. Because the composition of the nanowire depends on the precursor being fed to the Au droplets, it is possible to make "totem pole" structures with varied compositions along the nanowire axis. Semiconductor shells can also be grown around the wires by chemical vapor deposition (CVD). The VLS process can also be adapted to complex compositions for which no molecular precursor is available by using a laser to ablate the semiconductor from a solid target.
Epitaxial growth of silicon nanowires catalyzed by gold nanoparticles in the vapor-liquid-solid (VLS) process.
Semiconductor nanowires made by this method are the basis of extremely sensitive biosensors, in which a molecular binding event anywhere along the wire strongly affects its electrical conductivity.[7] Nanowire and microwire arrays are also being studied as solar cell and lithium battery materials, as well as nanoscale electronic and optoelectronic devices. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/11%3A_Basic_Science_of_Nanomaterials/11.03%3A_Semiconductor_Quantum_Dots.txt |
Nanoparticles have a substantial fraction of their atoms on the surface, as shown in the plot at the right.[8] This high surface area to volume ratio is an important factor in many of the physical properties of nanoparticles, such as their melting point and vapor pressure, and also in their reactivity. Heterogeneous catalysts, for example, are often based on nanoparticles because the catalytically reactive atoms are those that are on the surface of the particle.
The surface energy is always positive. A key quantity that is connected with the chemistry of all surfaces is the surface energy. This is the (thermodynamically unfavorable) energy of making "dangling bonds" at the surface. Atoms at the surface are under-coordinated, and because breaking bonds costs energy, surface atoms always have higher energy than atoms in the bulk. This happens regardless of whether the bonding is covalent (as in a metal), ionic (in a salt), or non-covalent (in a liquid such as water). We see this effect, for example, in water droplets that bead up on a waxy surface. The droplet contracts into a sphere (against the force of gravity that works to flatten it) in order to minimize the number of dangling hydrogen bonds at the surface.
In the case of metal or semiconductor particles, strong covalent bonds are broken at the surface. For example, a gold atom in bulk face-centered cubic Au has 12 nearest neighbors, but a gold atom on the (111) surface of the crystal (the most dense crystal plane of gold) has six nearest neighbors in-plane and three underneath, for a total of 9. We might expect the surface energy of this crystal face to be a little less (because the remaining bonds will become slightly stronger) than 3/12 = 1/4 of the bonding energy of bulk Au, and this is in fact a fairly good rule of thumb for many materials. When translated into energy per unit area, the surface energy of metals and inorganic salts is usually in the range of 1-2 J/m2.
Example: The sublimation energy of bulk gold is 334 kJ/mol, and the surface energy is 1.5 J/m2. What percentage of the bulk bonding energy is lost by atoms at the (111) surface of a gold crystal?
To solve this problem, we need to know the surface area per Au atom. Gold has the face-centered cubic structure, and the unit cell edge length is 4.08 Å. From this we can determine that the Au-Au distance is 4.08/1.414 = 2.88 Å. In a hexagonal array of gold atoms with this interatomic spacing, the surface area per atom is (2.88 Å)2 x 0.866 = 7.2 Å2. Multiplying by Avogadro's number, we find that the area per mole of Au surface atoms on the (111) crystal face is 4.3 x 104 m2.
Now we multiply this area by the surface energy:
$(4.3 \times 10^{4} \frac{m^{2}}{mol}) \times (1.5 \frac{J}{mol}) \times (\frac{1kJ}{1000J}) = \mathbf{+65 kJ}$ per mole of Au surface atoms.
$\frac{65 kJ}{334kJ} \times 100 \% = \mathbf {19 \%}$ of the bulk bonding energy is lost by atoms at the surface.
It is clear from this example that the surface energy of nanoparticles can have a major effect on their physical properties, since a large fraction of the atoms in a nanoparticle are on the surface. A good example of this is the dramatic depression of the melting point.
Solid nanocrystals are in general faceted, whereas liquid droplets made by melting nanocrystals adopt a spherical shape to minimize the surface area.
A faceted solid nanocrystal melts into a spherical droplet in order to minimize its surface energy.
Let's consider melting a silver nanocrystal that is 2 nm in diameter, meaning that about 1/2 of the atoms are on the surface. The spherical liquid droplet has lower surface area than the faceted crystal. For example, a cube has 1.24 times the surface area of a sphere of the same volume. If the decrease in surface area is about 20% upon melting, and the surface energy is about 1/4 of the bulk bonding energy of the atoms, then for a 2 nm diameter nanocrystal we would estimate that:
$\Delta H^{o}_{fusion} \approx \Delta H^{o}_{fusion, \: bulk} - (\frac{1}{2})(0.25)(\frac{1}{4}) \Delta H^{o}_{vap} = \Delta H^{o}_{fusion, \: bulk} - 0.025 \Delta H^{o}_{vap, \: bulk}$
where ΔH°vap, bulk, the heat of vaporization, is the total bonding energy of atoms in a bulk crystal. For silver, ΔH°fusion, bulk = 11.3 kJ/mol and ΔH°vap, bulk = 250 kJ/mol. From this we can calculate the heat of fusion of a 2 nm Ag nanocrystal:
$\Delta H^{o}_{fusion} \approx 11.3 - (0.025)(25) = 5.1 \frac{kJ}{mol}$
The melting point of bulk silver is 962 °C = 1235 K. Assuming that the entropy of fusion is the same in the bulk and in the nanocrystal, the melting point of the nanocrystal should be $1234K \times (\frac{5.1}{11.3}) = 557K = 284^{\circ} C$, a drop of almost 700 degrees from the bulk value. Experimentally, the melting point of a 2 nm diameter silver nanocrystal drops about 800 degrees below that of the bulk, to 127 °C.[9] This is a whopping big change in the melting point, which is in reasonable agreement with our rough estimate.
The same effect - energetic destabilization of the surface atoms relative to bulk atoms - results in the lower boiling point, higher vapor pressure, higher solubility, and higher reactivity of nanocrystals relative to microcrystals (or larger crystals) of the same material. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/11%3A_Basic_Science_of_Nanomaterials/11.05%3A_Surface_Energy.txt |
Nanoscale metal particles have been the subject of intense research over the past 20 years, especially because of their unusual optical, magnetic, and catalytic properties. The synthesis of metal nanocrystals in various shapes has become increasingly sophisticated and rational, like the synthesis of semiconductor nanocrystals described above. By controlling the separate phases of nucleation and growth, and by using ligands that cap specific crystal faces during growth, it is possible to make metal nanocrystals of uniform size in a variety of interesting and useful shapes including cubes, truncated cubes, octahedra, triangular prisms, and high aspect ratio rods. By exploiting displacement reactions that replace one metal with another, complex hollow shapes such as nanocages (as shown at the left) can be made starting with other shapes. In this case, solid silver nanocubes are transformed to gold nanocages.
Five-fold twinning in a gold nanoparticle
The interesting optical properties of nanocrystalline Au, Ag, Cu, and a number of other metals, derive from the collective oscillation of their valence electrons, a phenomenon known as plasmon resonance. Remember that in these metals, the electron mean free path is long (about 100 times larger than the size of the atoms), so the valence electrons feel only the average positive charge of the atomic cores as they zoom around the crystal. Light impinging on the metal acts as an oscillating electric field, pushing and pulling on the valence electrons at the characteristic frequency of the light wave. The situation is very much like a pendulum or a weight on a spring. The electrons, pushed away from their equilibrium positions, feel a restoring force that is proportional to their displacement. Their motion can be described by Hooke's Law:
$F = kx$
where the spring constant k determines the "stiffness" of the spring. In the case of the plasmon resonance, k is proportional to the number density of valence electrons n, and the square of the electronic charge e:
$k = \frac{ne^{2}}{\varepsilon_{0}}$
The resonant frequency of the plasma oscillation is given by:
$\omega_{p} = (\frac{k}{m})^{\frac{1}{2}} = (\frac{ne^{2}}{m_{e}\varepsilon_{0}}^{\frac{1}{2}}$
where me is the electron mass. For most metals, the plasmon resonance is in the ultraviolet part of the spectrum, but for a few metals like Au, Ag, and Cu it is in the visible.
The Lycurgus cup (4th century Roman glass) derives its unique coloration from noble metal nanoparticles. The cup is red in transmitted light and green in scattered (reflected) light.
For metal particles that are much smaller than the wavelength of light, this effect is called the localized surface plasmon resonance, or LSPR. There are three important consequences of the LSPR effect:
• The local electric field of the incoming light wave is greatly enhanced at the particle surface. This gives rise to huge enhancement factors in optical processes such as Raman scattering and fluorescence. Thus, certain analytical spectroscopic techniques are greatly enhanced by LSPR.
• Near the plasmon resonance frequency, metal nanocrystals absorb and scatter light very strongly. This makes them brightly reflective, and the strong light absorption can be exploited for light-induced local heating. These properties are being applied in medical diagnostics and therapy, e.g., for detection and photothermal destruction of cancer cells. By adjusting the size and shape of the gold nanoparticles, which are more stable than Ag and Cu in biological media, the plasmon frequency can be tuned to the tissue-transparent near-IR region of the spectrum between 700 and 900 nm. Small quantities of plasmonic Ag and Au particles also make brightly colored and strongly scattering pigments, e.g. in stained glass as shown above at the right.
• The plasmon frequency is sensitive to the refractive index of the particle's surroundings, i.e., its chemical environment. This makes metal nanoparticles of special interest for sensing and biosensing applications.
The colors of plasmonic gold nanoparticles depend on their size and shape.
Theory of light scattering and absorption by metal nanoparticles
The valence electrons in metal nanoparticles oscillate in the electric field of a light wave. While the nature of these oscillations is somewhat complex in metal particles that are non-spherical, the theory for spherical particles is relatively simple and in fact was worked out over 100 years ago by German physicist Gustav Mie.
Mie considered the interaction of a spherical particle with a uniform electric field, E, oscillating at angular frequency ω (= 2π f). This is a good approximation when the particle diameter is much smaller than the wavelength of light, as shown on the left. The particle is embedded in a uniform, insulating material (e.g. a solvent) that has a dielectric constant εdiel. For insulators, εdiel is a positive real number.
The dielectric constant ε of a metal is actually a complex number:
$\varepsilon = \varepsilon^{'} + i \varepsilon^{"}$
Here the real part, ε', is related to the refraction of light, and the imaginary part, ε", is related to light absorption. Both ε' and ε" are dependent on the frequency of the light. For metals near the plasmon resonance frequency, ε' is typically a negative number.
The cross-section for absorption of the light wave by the particle is:
$\sigma_{absorption} = \frac{9 \omega}{c} \epsilon^{\frac{3}{2}}_{diel} V \frac{\epsilon^{"}_{metal}}{(\epsilon^{'}_{metal} + 2 \epsilon_{diel})^{2} + (\epsilon^{"}_{metal})^{2}}$
and the cross-section for scattering is:
$\sigma_{scattering} = \frac{3}{2\pi} (\frac{\omega^{4}}{c}) \epsilon^{2}_{diel} V^{2} \frac{(\epsilon^{'}_{metal} - \epsilon_{diel})^{2} + (\epsilon^{"}_{metal})^{2}}{(\epsilon^{'}_{metal} + 2 \epsilon_{diel})^{2} + (\epsilon^{"}_{metal})^{2}}$
The sum of these two is the cross-section for extinction:
$\sigma_{extinction} = \sigma_{absorption} + \sigma_{scattering}$
These cross-sections become large when the (ε'metal + 2εdiel) term in the denominator becomes small. This occurs when
$\epsilon^{'}_{metal} \simeq -2\epsilon_{diel}$
For 15 nm diameter gold nanoparticles in water, this happens at about 580 nm, resulting in the characteristic wine-red color of colloidal gold solutions. Changing the solution environment (e.g., by adsorbing a molecule onto the gold surface) changes εdiel and thus alters the color slightly.
It is important to note that the cross-section for scattering is proportional to the square of the volume of the particle, V2, whereas the absorption is proportional to V. This means that very small gold particles (< 5 nm) are strongly absorbing but not strongly scattering. Larger particles (>30 nm) scatter light very strongly. Depending on the application, therefore, we choose larger or smaller particles.
One of the key complementary properties of noble metal nanoparticles that is important to their use in biomedicine is the ease with which they can be covalently conjugated with polymers or small molecules, typically via thiol or amine bonds at their surface. This imparts biological recognition properties to the particles that enables them to bind to specific biomolecular targets. The figure at the left illustrates some of the functionality that can be imparted to nanoparticles through surface functionalization.
Functionalization of gold nanoparticles with thiol-terminated single-stranded DNA was the basis of one of the first nanoparticle sensors, developed by the Mirkin group at Northwestern University. DNA-coated nanoparticles have the characterstic wine-red plasmonic color of spherical nano-gold. However, when these particles are linked together by a complementary DNA strand, the resonance frequency shifs, resulting in a blue color. This color change, illustrated in the figure at the right, provides a "litmus test" for the presence of the target DNA sequence.[10] "Melting" of the DNA - heating it to the temperature at which double stranded DNA dissociates to make single strands - reverses the color change. The DNA hybridization/melting transition is highly cooperative because of the aggregation of many gold particles, so the transition temperature is very sharp. With proper temperature control, the color change can be sensitive to a single base mismatch in the target DNA that is detected by this method.
Aggregation of Au nanoparticles (in this case by adding salt to a colloidal solution) causes a color change from wine red (left) to blue (right). Photo credit: George Lisensky, Beloit College
Subsequent research has developed sophisticated diagnostic and therapeutic ("theranostic") applications for these spherical nucleic acid[11] particles. These particles easily penetrate cell membranes and can report on the chemistry happening inside living cells.
General schematic of nanoflare-based detection.
An important property of gold nanoparticles in these applications is their ability to quench the fluorescence of reporter molecules that are near their surface. Nucleic acid strands that contain a hairpin loop can position fluorescent molecules near the gold surface, where their fluorescence is turned off by nanoparticle quenching. Hybridization of these sequences to target RNA or DNA causes the fluorescence to turn on by moving the fluorescent molecule away from the nanoparticle surface. These so called "nanoflares" can thus signal the up- or down-regulation of specific genes inside cells. Nanoflares are the basis of the Verigene System, developed and commercialized by Nanosphere, Inc. to detect markers for infectious diseases and cancers. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/11%3A_Basic_Science_of_Nanomaterials/11.06%3A_Nanoscale_Metal_Particles.txt |
PowerPoint slides for this section
11.08: Discussion Questions
• Explain how separating the nucleation and growth steps leads to nanoparticles of uniform size.
• A recent paper by Delia Milliron and coworkers (Nature, 2013, 500, 323–326, doi:10.1038/nature12398) describes plasmonic indium-tin-oxide (ITO) nanoparticles that can control the infrared transparency of windows. Explain how the plasmon resonance wavelength of the nanoparticles can be tuned by electrochemical doping, and how the invention could save some of the energy used to heat and cool buildings.
11.09: Problems
1. Consider a spherical gold nanoparticle that contains 500 atoms. If the diameter of an atom is approximately 3 Å, what fraction of the gold atoms in the particle are on the surface?
2. Now consider small droplets of mercury that contains 500 atoms. Mercury atoms are also about 3 Å in diameter. The heat of vaporization of bulk mercury is 64.0 kJ/mol, and the vapor pressure of mercury is 0.00185 torr = 2.43 x 10-6 atm. The surface tension of mercury (γHg) is 0.518 N/m, and the surface excess energy can be calculated as γHgA, where A is the surface area. Using this information and the Clausius-Clapyron equation (P = const•exp(-ΔHvap/RT)), calculate the vapor pressure of these small droplets of mercury.
3. James Heath and coworkers (Phys. Rev. Lett. 1995, 75, 3466) have observed Ostwald ripening in thin films of gold nanoparticles at room temperature. Starting with an uneven distribution of particle sizes, they find that the large particles grow at the expense of smaller ones. Can you explain this observation, based on your answers to problems (1) and (2)?
4. The bandgap of bulk germanium is 0.67 eV. What bandgap would you expect for a 4 nm diameter Ge nanocrystal? Use the Brus formula,
$\Delta E_{gap} \approx \frac{h^{2}}{8 \mu R^{2}} - \frac{1.8 e^{2}}{4R \pi \varepsilon \varepsilon_{0}} + \dots$
where R is the particle radius, εGe = 16.2, h = 6.6 10-34 J s, 1 eV = 1.6 10-19 J, 1 J = 1 kg m2/s2. and e2/4πε0 = 1.44 10-9 eV m. Assume that the electron-hole reduced mass μ is approximately 40% of the free electron mass, me = 9.1 10-31 kg.
5. Grecian Formula (a hair coloring product) until recently contained lead acetate, which reacts with the cysteine in hair to make PbS. In bulk form, PbS is a semiconductor with a band gap of 0.3 eV (1 eV = 1240 nm). The particles are initially very small but grow as more Grecian Formula is applied and reacts with cysteine. As the particles grow, they change progressively from colorless to yellow to black. Explain why the particles are initially colorless and why their color changes. (Grecian formula now uses Bi, which is less toxic than Pb, and works by the same mechanism)
11.10: References
1. C.T. Liu, S. Y. Lin, D. C. Tsui, H. Lee, and D. Ackley, Appl. Phys. Lett. 1988, 53, 2510. DOI: 10.1063/1.100409.
2. D. L. Feldheim, K. C. Grabar, M. J. Natan, and T. E. Mallouk, "Electron Transfer in Self-Assembled Inorganic Polyelectrolyte/Metal Nanoparticle Heterostructures," J. Am. Chem. Soc., 118, 7640-1 (1996)
3. S. Chen, R. W. Murray, and S. W. Feldberg, "Quantized Capacitance Charging of Monolayer-Protected Au Clusters," J. Phys. Chem. B 1998, 102, 9898-9907.
4. Brus, Louis E. (1984). "Electron–electron and electron‐hole interactions in small semiconductor crystallites: The size dependence of the lowest excited electronic state". J. Chem. Phys. 80, 4403. DOI: 10.1063/1.447218.
5. C. B. Murray, D. J. Norris, and M. G. Bawendi, "Synthesis and characterization of nearly monodisperse CdE (E = sulfur, selenium, tellurium) semiconductor nanocrystallites," J. Am. Chem. Soc. 1993, 115, 8706–8715. DOI: 10.1021/ja00072a025.
6. Y. Yin and A. P. Alivisatos, "Colloidal nanocrystal synthesis and the organic–inorganic interface," Nature 2005, 437, 664-670. DOI: 10.1038/nature04165
7. G. Zheng, F. Patolsky, Y. Cui, W. U. Wang and C. M. Lieber, "Multiplexed electrical detection of cancer markers with nanowire sensor arrays," Nature Biotechnol. 2005, 23, 1294 - 1301. DOi:10.1038/nbt1138.
8. K. J. Klabunde, J. Stark, O. Koper, C. Mohs, D. G. Park, S. Decker, Y. Jiang, I. Lagadic, and D. Zhang, "Nanocrystals as Stoichiometric Reagents with Unique Surface Chemistry," J. Phys. Chem. 1996, 100, 12142–12153. DOI: 10.1021/jp960224x.
9. S. A. Little, T. Begou, R. W. Collins, and S. Marsillac, Appl. Phys. Lett. 2012, 100, 051107. DOI: 10.1063/1.3681367
10. Mirkin, C. A., et al., A DNA-based method for rationally assembling nanoparticles into macroscopic materials. Nature 1996, 382 (6592), 607-609.
11. Cutler, J. I., et al., Spherical Nucleic Acids. J Am Chem Soc 2012, 134 (3), 1376-1391. | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/11%3A_Basic_Science_of_Nanomaterials/11.07%3A_Applications_of_Nanomaterials.txt |
Powerpoint slides to accompany the chapters of this book will be provided to instructors upon request ([email protected])
Additional resources:
12: Resources for Students and Teachers
IONIC VIPEr is a cyber-interface that facilitates collaborative development of learning materials and their dissemination to the wider inorganic community. This website, VIPEr (Virtual Inorganic Pedagogical Electronic Resource), serves both as a repository and as a user-friendly platform for social networking tools that facilitate virtual collaboration and community building. The VIPEr community seeks to develop and disseminate best practices for teaching inorganic chemistry.
IONIC VIPEr website
12.02: Beloit College University of Wisconsin Video Lab Manual
The video lab manual provides a broad range of videos of laboratory experiments in inorganic nanoscience and materials chemistry. This is an especially useful resource for teachers who would like to develop laboratory experiments for their classes in inorganic chemistry.
Video Lab Manual website
12.03: Atomic and Molecular Orbitals (University of Liverpool)
The University of Liverpool website shows animations of atomic orbitals, selected molecular orbital diagrams, and the VSEPR shapes of selected molecules.
Structure and Bonding website
12.04: Interactive 3D Crystal Structures (University of Liverpool)
The University of Liverpool website also provides a menu of inorganic structures (both molecules and extended solids) that can be visualized and manipulated in 3D. This is a valuable tool that supplements the descriptions of structures in Chapters 5-8 of this book.
3D Crystal Structure website
12.05: Appendix 1- Periodic Tables
Interactive Periodic Table (Wikipedia)
-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-.-
Mendeleev's periodic table (1869)
18-column periodic table
A 32-column periodic table with Sc, Y, Lu and Lr in group 3
12.06: Appendix 2- Selected Thermodynamic Values
Selected Thermodynamic Values (at 298.15 K)
Substance ΔHf° (kJ/mol) S° (J/K·mol) ΔGf° (kJ/mol)
Aluminum
Al(s) 0 28.3 0
AlCl3(s) -704.2 110.67 -628.8
Al2O3(s) -1675.7 50.92 -1582.3
Barium
BaCl2(s) -858.6 123.68 -810.4
BaCl2 • 2 H2O (s) -1460.1 203 -1296.5
BaO(s) -553.5 70.42 -525.1
Ba(OH)2 • 8 H2O (s) -3342 427 -2793
BaSO4(s) -1473.2 132.2 -1362.2
Beryllium
Be(s) 0 9.5 0
Be(OH)2(s) -902.5 51.9 -815
Bromine
Br(g) 111.884 175.022 82.396
Br2(liq) 0 152.2 0
Br2(g) 30.907 245.463 3.11
BrF3(g) -255.6 292.53 -229.43
HBr(g) -36.4 198.695 -53.45
Calcium
Ca(s) 0 41.42 0
Ca(g) 178.2 158.884 144.3
Ca2+(g) 1925.9
CaC2(s) -59.8 69.96 -64.9
CaCO3 (s; calcite) -1206.92 92.9 -1128.79
CaCl2(s) -795.8 104.6 -748.1
CaF2(s) -1219.6 68.87 -1167.3
CaH2(s) -186.2 42 -147.2
CaO(s) -635.09 39.75 -604.03
CaS(s) -482.4 56.5 -477.4
Ca(OH)2(s) -986.09 83.39 -898.49
Ca(OH)2(aq) -1002.82 -74.5 -868.07
CaSO4(s) -1434.11 106.7 -1321.79
Carbon
C(s, graphite) 0 5.74 0
C(s, diamond) 1.895 2.377 2.9
C(g) 716.682 158.096 671.257
CCl4(liq) -135.44 216.4 -65.21
CCl4(g) -102.9 309.85 -60.59
CHCl3(liq) -134.47 201.7 -73.66
CHCl3(g) -103.14 295.71 -70.34
CH4 (g, methane) -74.81 186.264 -50.72
C2H2 (g, ethyne) 226.73 200.94 209.2
C2H4 (g,ethene) 52.26 219.56 68.15
C2H6 (g, ethane) -84.68 229.6 -32.82
C3H8 (g, propane) -103.8 269.9 -23.49
C4H10 (g, butane) -888.0
C6H6 (liq, benzene) 49.03 172.8 124.5
C6H14(liq) -198.782 296.018 -4.035
C8H18(liq) -249.952 361.205 6.707
CH3OH(liq, methanol) -238.66 126.8 -166.27
CH3OH(g, methanol) -200.66 239.81 -161.96
C2H5OH(liq, ethanol) -277.69 160.7 -174.78
C2H5OH(g, ethanol) -235.1 282.7 -168.49
CH3COOH(liq) -276.981 160.666 -173.991
CO(NH2)2(s, urea) -333.5 104.6 -197.4
CO(g) -110.525 197.674 -137.168
CO2(g) -393.509 213.74 -394.359
CS2(g) 117.36 237.84 67.12
COCl2(g) -218.8 283.53 -204.6
Cesium
Cs(s) 0 85.23 0
Cs+(g) 457.964
CsCl(s) -443.04 101.17 -414.53
Chlorine
Cl(g) 121.679 165.198 105.68
Cl-(g) -233.13
Cl2(g) 0 223.066 0
HCl(g) -92.307 186.908 -95.299
HCl(aq) -167.159 56.5 -131.228
Chromium
Cr(s) 0 23.77 0
Cr2O3(s) -1139.7 81.2 -1058.1
CrCl3(s) -556.5 123 -486.1
Copper
Cu(s) 0 33.15 0
CuO(s) -157.3 42.63 -129.7
CuCl2(s) -220.1 108.07 -175.7
Fluorine
F2(g) 0 202.78 0
F(g) 78.99 158.754 61.91
F-(g) -255.39
F-(aq) -332.63 -13.8 -278.79
HF(g) -271.1 173.779 -273.2
HF(aq) -332.63 -13.8 -278.79
Hydrogen
H2(g) 0 130.684 0
H(g) 217.965 114.713 203.247
H+(g) 1536.202
H2O(liq) -285.83 69.91 -237.129
H2O(g) -241.818 188.825 -228.572
H2O2(liq) -187.78 109.6 -120.35
Iodine
I2(s) 0 116.135 0
I2(g) 62.438 260.69 19.327
I(g) 106.838 180.791 70.25
I-(g) -197
ICl(g) 17.78 247.551 -5.46
Iron
Fe(s) 0 27.78 0
FeO(s) -272
Fe2O3(s, hematite) -824.2 87.4 -742.2
Fe3O4(s, magnetite) -1118.4 146.4 -1015.4
FeCl2(s) -341.79 117.95 -302.3
FeCl3(s) -399.49 142.3 -344
FeS2(s, pyrite) -178.2 52.93 -166.9
Fe(CO)5(liq) -774 338.1 -705.3
Lead
Pb(s) 0 64.81 0
PbCl2(s) -359.41 136 -314.1
PbO(s, yellow) -217.32 68.7 -187.89
PbS(s) -100.4 91.2 -98.7
Lithium
Li(s) 0 29.12 0
Li+(g) 685.783
LiOH(s) -484.93 42.8 -438.95
LiOH(aq) -508.48 2.8 -450.58
LiCl(s) -408.701 59.33 -384.37
Magnesium
Mg(s) 0 32.68 0
MgCl2(g) -641.32 89.62 -591.79
MgO(s) -601.7 26.94 -569.43
Mg(OH)2(s) -924.54 63.18 -833.51
MgS(s) -346 50.33 -341.8
Mercury
Hg(liq) 0 29.87 0
HgCl2(s) -224.3 146 -178.6
HgO(s, red) -90.83 70.29 -58.539
HgS(s, red) -58.2 82.4 -50.6
Nickel
Ni(s) 0 29.87 0
NiO(s) -239.7 37.99 -211.7
NiCl2(s) -305.332 97.65 -259.032
Nitrogen
N2(g) 0 191.61 0
N(g) 472.704 153.298 455.563
NH3(g) -46.11 192.45 -16.45
N2H4(liq) 50.63 121.21 149.34
NH4Cl(s) -314.43 94.6 -202.87
NH4Cl(aq) -299.66 169.9 -210.52
NH4NO3(s) -365.56 151.08 -183.87
NH4NO3(aq) -339.87 259.8 -190.56
NO(g) 90.25 210.76 86.55
NO2(g) 33.18 240.06 51.31
N2O(g) 82.05 219.85 104.2
N2O4(g) 9.16 304.29 97.89
NOCl(g) 51.71 261.69 66.08
HNO3(liq) -174.1 155.6 -80.71
HNO3(g) -135.06 266.38 -74.72
HNO3(aq) -207.36 146.4 -111.25
Oxygen
O2(g) 0 205.138 0
O(g) 249.17 161.055 231.731
O3(g) 142.7 238.93 163.2
Phosphorus
P4(s, white) 0 164.36 0
P4(s, red) -70.4 91.2 -48.4
P(g) 314.64 163.193 278.25
PH3(g) 5.4 310.23 13.4
PCl3(g) -287 311.78 -267.8
P4O10(s) -2984 228.86 -2697.7
H3PO4(s) -1279 110.5 -1119.1
Potassium
K(s) 0 64.18 0
KCl(s) -436.747 82.59 -409.14
KClO3(s) -397.73 143.1 -296.25
KI(s) -327.9 106.32 -324.892
KOH(s) -424.764 78.9 -379.08
KOH(aq) -482.37 91.6 -440.5
Silicon
Si(s) 0 18.83 0
SiBr4(liq) -457.3 277.8 -443.8
SiC(s) -65.3 16.61 -62.8
SiCl4(g) -657.01 330.73 -616.98
SiH4(g) 34.3 204.62 56.9
SiF4(g) -1614.94 282.49 -1572.65
SiO2(s, quartz) -910.94 41.84 -856.64
Silver
Ag(s) 0 42.55 0
Ag2O(s) -31.05 121.3 -11.2
AgCl(s) -127.068 96.2 -109.789
AgNO3(s) -124.39 140.92 -33.41
Sodium
Na(s) 0 51.21 0
Na(g) 107.32 153.712 76.761
Na+(g) 609.358
NaBr(s) -361.062 86.82 -348.983
NaCl(s) -411.153 72.13 -384.138
NaCl(g) -176.65 229.81 -196.66
NaCl(aq) -407.27 115.5 -393.133
NaOH(s) -425.609 64.455 -379.484
NaOH(aq) -470.114 48.1 -419.15
Na2CO3(s) -1130.68 134.98 -1044.44
Sulfur
S(s, rhombic) 0 31.8 0
S(g) 278.805 167.821 238.25
S2Cl2(g) -18.4 331.5 -31.8
SF6(g) 1209 291.82 -1105.3
H2S(g) -20.63 205.79 -33.56
SO2(g) -296.83 248.22 -300.194
SO3(g) -395.72 256.76 -371.06
SOCl2(g) -212.5 309.77 -198.3
H2SO4(liq) -813.989 156.904 -690.003
H2SO4(aq) -909.27 20.1 -744.53
Tin
Sn(s, white) 0 51.55 0
Sn(s, gray) -2.09 44.14 0.13
SnCl4(liq) -511.3 248.6 -440.1
SnCl4(g) -471.5 365.8 -432.2
SnO2(s) -580.7 52.3 -519.6
Titanium
Ti(s) 0 30.63 0
TiCl4(liq) -804.2 252.34 -737.2
TiCl4(g) -763.2 354.9 -726.7
TiO2(s) -939.7 49.92 -884.5
Zinc
Zn(s) 0 41.63 0
ZnCl2(s) -415.05 111.46 -369.398
ZnO(s) -348.28 43.64 -318.3
ZnS(s, sphalerite) -205.98 57.7 -201.29
Aqueous Ions and Molecules
Ca2+(aq) -542.96 -55.2 -553.04
CO32-(aq) -676.26 -53.1 -528.1
CO2(aq) -413.8 117.6 -386.0
Cl-(aq) -167.16 56.5 -131.26
H+(aq) 0 0
HCO2-(aq) -410 91.6 -335
HCO2H(aq) -410 164 -356
HCO3-(aq) -691.11 95 -587.06
H2CO3(aq) -698.7 191 -623.42
NH3(aq) -80.29 111 -26.6
OH-(aq) -229.94 -10.54 -157.3
Ag+(aq) 105.58 72.68 77.124
12.07: Appendix 3- Bond Enthalpies
Average single bond enthalpies (kJ/mol)
H-H 436 C-H 413 N-H 391 O-H 483 F-F 155
H-F 567 C-C 348 N-N 163 O-O 146
H-Cl 431 C-N 293 N-O 201 O-F 190 Cl-F 253
H-Br 366 C-O 358 N-F 272 O-Cl 203 Cl-Cl 242
H-I 299 C-F 485 N-Cl 200 O-I 234
C-Cl 328 N-Br 243 Br-F 237
C-Br 276 S-H 339 Br-Cl 218
C-I 240 P-H 322 S-F 327 Br-Br 193
C-S 259 P-F 490 S-Cl 253
S-Br 218 I-Cl 208
Si-H 323 S-S 266 I-Br 175
Si-Si 226 I-I 151
Si-C 301
Si-O 368
Si-F 565
Si-Cl 464
Average double bond enthalpies (kJ/mol)
C=C 614 N=N 418 O=O 495
C=N 615 N=O 607 S=O 523
C=O 799 S=S 418
Average triple bond enthalpies (kJ/mol)
C≡C 839 C≡N 891 N≡N 941 C≡O 1072
13.01: Prelude to Metals and Alloys - Mechanical Properties
Thumbnail: Microstructure of rolled and annealed brass (400× magnification). Public Domain via Wikipedia
13: Metals and Alloys - Mechanical Properties
How much do the mechanical properties of metals and alloys vary with processing? The answer is, a great deal. Consider the following hypothetical situation: Upon graduation, you go to work as an engineer for Boeing. Your job is to work with aluminum companies to help them produce high strength alloys. Why? A large jet airplane weighs a total of 500 tons. Of that total, 50 tons is cargo, 150 tons is the plane structure, and the remainder is fuel. If you can triple the strength of the materials in the structure (aluminum), you can reduce the mass of the structure to 50 tons and increase the cargo to 150 tons. Look at what has been done already:
Material Tensile strength yield (psi)
pure (99.45%) annealed Al 4 x 103
pure (99.45%) cold drawn Al 24 x 103
Al alloy - precipitated, hardened 50 x 103
By chemical and physical manipulation we have already increased the yield strength 12 times over annealed Al. Yet the yield strength of a "perfect" single crystal of pure Al is ca. 106 psi. We still have 3 orders of magnitude to go. This just shows that there will still be plenty to do on this project between now and graduation! | textbooks/chem/Inorganic_Chemistry/Book%3A_Introduction_to_Inorganic_Chemistry_(Wikibook)/12%3A_Resources_for_Students_and_Teachers/12.01%3A_VIPEr-_Virtual_Inorganic_Pedagogical_Electronic_Resource-_A_Community_for_Teachers_and_Students_.txt |
Living organisms store and transport transition metals both to provide appropriate concentrations of them for use in metalloproteins or cofactors and to protect themselves against the toxic effects of metal excesses; metalloproteins and metal cofactors are found in plants, animals, and microorganisms. The normal concentration range for each metal in biological systems is narrow, with both deficiencies and excesses causing pathological changes. In multicellular organisms, composed of a variety of specialized cell types, the storage of transition metals and the synthesis of the transporter molecules are not carried out by all types of cells, but rather by specific cells that specialize in these tasks. The form of the metals is always ionic, but the oxidation state can vary, depending on biological needs. Transition metals for which biological storage and transport are significant are, in order of decreasing abundance in living organisms: iron, zinc, copper, molybdenum, cobalt, chromium, vanadium, and nickel. Although zinc is not strictly a transition metal, it shares many bioinorganic properties with transition metals and is considered with them in this chapter. Knowledge of iron storage and transport is more complete than for any other metal in the group.
III. Summary
Transition metals (Fe, Cu, Mo, Cr, Co, Mn, V) play key roles in such biological processes as cell division (Fe, Co), respiration (Fe, Cu), nitrogen fixation (Fe, Mo, V), and photosynthesis (Mn, Fe). Zn participates in many hydrolytic reactions and in the control of gene activity by proteins with "zinc fingers." Among transition metals, Fe predominates in terrestial abundance; since Fe is involved in a vast number of biologically important reactions, its storage and transport have been studied extensively. Two types of Fe carriers are known: specific proteins and low-molecular-weight complexes. In higher animals, the transport protein transferrin binds two Fe atoms with high affinity; in microorganisms, iron is transported into cells complexed with catecholates or hydroxamates called siderophores; and in plants, small molecules such as citrate, and possibly plant siderophores, carry Fe. Iron complexes enter cells through complicated paths involving specific membrane sites (receptor proteins). A problem yet to be solved is the form of iron transported in the cell after release from transferrin or siderophores but before incorporation into Fe-proteins.
Iron is stored in the protein ferritin. The protein coat of ferritin is a hollow sphere of 24 polypeptide chains through which Fe2+ passes, is oxidized, and mineralizes inside in various forms of hydrated Fe2O3. Control of the formation and dissolution of the mineral core by the protein and control of protein synthesis by Fe are subjects of current study.
Biomineralization occurs in the ocean (e.g., Ca in shells, Si in coral reefs) and on land in both plants (e.g., Si in grasses) and animals (e.g., Ca in bone, Fe in ferritin, Fe in magnetic particles). Specific organic surfaces or matrices of protein and/or lipid allow living organisms to produce minerals of defined shape and composition, often in thermodynamically unstable states.
Contributors and Attributions
• Elizabeth C. Theil (North Carolina State University, Department of Biochemistry)
• Kenneth N. Raymond (University of California at Berkeley, Department of Chemistry)
01: Transition-Metal Storage Transport and Biomineralization
The transition metals and zinc are among the least abundant metal ions in the sea water from which contemporary organisms are thought to have evolved (Table 1.1).1-5 For many of the metals, the concentration in human blood plasma greatly exceeds that in sea water. Such data indicate the importance of mechanisms for accumulation, storage, and transport of transition metals and zinc in living organisms.
Table 1.1: Concentrations of transition metals and zinc in sea water and human plasma.a*Data from References 1 - 5 and 12
Element Sea Water (M) x 108 Human Plasma (M) x 108
Fe 0.005 - 2 2230
Zn 8.0 1720
Cu 1.0 1650
Mo 10.0 1000
Co 0.7 0.0025
Cr 0.4 5.5
V 4.0 17.7
Mn 0.7 10.9
Ni 0.5 4.4
The metals are generally found either bound directly to proteins or in cofactors such as porphyrins or cobalamins, or in clusters that are in turn bound by the protein; the ligands are usually O, N, S, or C. Proteins with which transition metals and zinc are most commonly associated catalyze the intramolecular or intermolecular rearrangement of electrons. Although the redox properties of the metals are important in many of the reactions, in others the metal appears to contribute to the structure of the active state, e.g., zinc in the Cu-Zn dismutases and some of the iron in the photosynthetic reaction center. Sometimes equivalent reactions are catalyzed by proteins with different metal centers; the metal binding sites and proteins have evolved separately for each type of metal center.
Iron is the most common transition metal in biology.6,7 Its use has created a dependence that has survived the appearance of dioxygen in the atmosphere ca. 2.5 billion years ago, and the concomitant conversion of ferrous ion to ferric ion and insoluble rust (Figure 1.1 See color plate section, page C-1.). All plants, animals, and bacteria use iron, except for a lactobacillus that appears to maintain high concentrations of manganese instead of iron. The processes and reactions in which iron participates are crucial to the survival of terrestrial organisms, and include ribonucleotide reduction (DNA synthesis), energy production (respiration), energy conversion (photosynthesis), nitrogen reduction, oxygen transport (respiration, muscle contraction), and oxygenation (e.g., steroid synthesis, solubilization and detoxification of aromatic compounds). Among the transition metals used in living organisms, iron is the most abundant in the environment. Whether this fact alone explains the biological predominance of iron or whether specific features of iron chemistry contribute is not clear.
Many of the other transition metals participate in reactions equivalent to those involving iron, and can sometimes substitute for iron, albeit less effectively, in natural Fe-proteins. Additional biological reactions are unique to nonferrous transition metals.
Zinc is relatively abundant in biological materials.8,9 The major location of zinc in the body is metallothionein, which also binds copper, chromium, mercury, and other metals. Among the other well-characterized zinc proteins are the Cu-Zn superoxide dismutases (other forms have Fe or Mn), carbonic anhydrase (an abundant protein in red blood cells responsible for maintaining the pH of the blood), alcohol dehydrogenase, and a variety of hydrolases involved in the metabolism of sugars, proteins, and nucleic acids. Zinc is a common element in nucleic-acid polymerases and transcription factors, where its role is considered to be structural rather than catalytic. Interestingly, zinc enhances the stereoselectivity of the polymerization of nucleotides under reaction conditions designed to simulate the environment for prebiotic reactions. Recently a group of nucleic-acid binding proteins, with a repeated sequence containing the amino acids cysteine and histidine, were shown to bind as many as eleven zinc atoms necessary for protein function (transcribing DNA to RNA).10 Zinc plays a structural role, forming the peptide into multiple domains or "zinc fingers" by means of coordination to cysteine and histidine (Figure 1.2A See color plate section, page C-l.). A survey of the sequences of many nucleic-acid binding proteins shows that many of them have the common motif required to form zinc fingers. Other zinc-finger proteins called steroid receptors bind both steroids such as progesterone and the progesterone gene DNA (Chapter 8). Much of the zinc in animals and plants has no known function, but it may be maintaining the structures of proteins that activate and deactivate genes.11
Copper and iron proteins participate in many of the same biological reactions:
1. reversible binding of dioxygen, e.g., hemocyanin (Cu), hemerythrin (Fe), and hemoglobin (Fe);
2. activation of dioxygen, e.g., dopamine hydroxylase (Cu) (important in the synthesis of the hormone epinephrine), tyrosinases (Cu), and catechol dioxygenases (Fe);
3. electron transfer, e.g., plastocyanins (Cu), ferredoxins, and c-type cytochromes (Fe);
4. dismutation of superoxide by Cu or Fe as the redox-active metal (superoxide dismutases).
The two metal ions also function in concert in proteins such as cytochrome oxidase, which catalyzes the transfer of four electrons to dioxygen to form water during respiration. Whether any types of biological reactions are unique to copper proteins is not clear. However, use of stored iron is reduced by copper deficiency, which suggests that iron metabolism may depend on copper proteins, such as the serum protein ceruloplasmin, which can function as a ferroxidase, and the cellular protein ascorbic acid oxidase, which also is a ferrireductase.
Cobalt is found in vitamin B12 , its only apparent biological site.12 The vitamin is a cyano complex, but a methyl or methylene group replaces CN in native enzymes. Vitamin-B12 deficiency causes the severe disease of pernicious anemia in humans, which indicates the critical role of cobalt. The most common type of reaction in which cobalamin enzymes participate results in the reciprocal exchange of hydrogen atoms if they are on adjacent carbon atoms, yet not with hydrogen in solvent water:
(An important exception is the ribonucleotide reductase from some bacteria and lower plants, which converts ribonucleotides to the DNA precursors, deoxyribonucleotides, a reaction in which a sugar -OH is replaced by -H. Note that ribonucleotide reductases catalyzing the same reaction in higher organisms and viruses are proteins with an oxo-bridged dimeric iron center.) The cobalt in vitamin B12 is coordinated to five N atoms, four contributed by a tetrapyrrole (corrin); the sixth ligand is C, provided either by C5 of deoxyadenosine in enzymes such as methylmalonyl-CoA mutase (fatty acid metabolism) or by a methyl group in the enzyme that synthesizes the amino acid methionine in bacteria.
Nickel is a component of a hydrolase (urease), of hydrogenase, of CO dehydrogenase, and of S-methyl CoM reductase, which catalyzes the terminal step in methane production by methanogenic bacteria. All the Ni-proteins known to date are from plants or bacteria.13,14 However, about 50 years elapsed between the crystallization of jack-bean urease in 1925 and the identification of the nickel component in the plant protein. Thus it is premature to exclude the possibility of Ni-proteins in animals. Despite the small number of characterized Ni-proteins, it is clear that many different environments exist, from apparently direct coordination to protein ligands (urease) to the tetrapyrrole F430 in methylreductase and the multiple metal sites of Ni and Fe-S in a hydrogenase from the bacterium Desulfovibrio gigas. Specific environments for nickel are also indicated for nucleic acids (or nucleic acid-binding proteins), since nickel activates the gene for hydrogenase.15
Manganese plays a critical role in oxygen evolution catalyzed by the proteins of the photosynthetic reaction center. The superoxide dismutase of bacteria and mitochondria, as well as pyruvate carboxylase in mammals, are also manganese proteins.16,17 How the multiple manganese atoms of the photosynthetic reaction center participate in the removal of four electrons and protons from water is the subject of intense investigation by spectroscopists, synthetic inorganic chemists, and molecular biologists.17
Vanadium and chromium have several features in common, from a bioinorganic viewpoint.18a First, both metals are present in only small amounts in most organisms. Second, the biological roles of each remain largely unknown.18 Finally, each has served as a probe to characterize the sites of other metals, such as iron and zinc. Vanadium is required for normal health, and could act in vivo either as a metal cation or as a phosphate analogue, depending on the oxidation state, V(lV) or V(V), respectively. Vanadium in a sea squirt (tunicate), a primitive vertebrate (Figure 1.2B), is concentrated in blood cells, apparently as the major cellular transition metal, but whether it participates in the transport of dioxygen (as iron and copper do) is not known. In proteins, vanadium is a cofactor in an algal bromoperoxidase and in certain prokaryotic nitrogenases. Chromium imbalance affects sugar metabolism and has been associated with the glucose tolerance factor in animals. But little is known about the structure of the factor or of any other specific chromium complexes from plants, animals, or bacteria.
Molybdenum proteins catalyze the reduction of nitrogen and nitrate, as well as the oxidation of aldehydes, purines, and sulfite.19 Few Mo-proteins are known compared to those involving other transition metals. Nitrogenases, which also contain iron, have been the focus of intense investigations by bioinorganic chemists and biologists; the iron is found in a cluster with molybdenum (the iron-molybdenum cofactor, or FeMoCo) and in an iron-sulfur center (Chapter 7). Interestingly, certain bacteria (Azotobacter) have alternative nitrogenases, which are produced when molybdenum is deficient and which contain vanadium and iron or only iron. All other known Mo-proteins are also Fe-proteins with iron centers, such as tetrapyrroles (heme and chlorins), Fe-sulfur clusters, and, apparently, non-heme/non-sulfur iron. Some Mo-proteins contain additional cofactors such as the Havins, e.g., in xanthine oxidase and aldehyde oxidase. The number of redox centers in some Mo-proteins exceeds the number of electrons transferred; reasons for this are unknown currently. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/01%3A_Transition-Metal_Storage_Transport_and_Biomineralization/1.01%3A_Biological_Significance_of_Iron_Zinc_Copper_Molybdenum_Cobalt_Chromium_Vanadium_and_.txt |
Iron
Three properties of iron can account for its extensive use in terrestrial biological reactions:
1. facile redox reactions of iron ions;
2. an extensive repertoire of redox potentials available by ligand substitution or modification (Table 1.4);
3. abundance and availability (Table 1.1) under conditions apparently extant when terrestrial life began (see Section I.B.).
Ferrous ion appears to have been the environmentally stable form during prebiotic times. The combination of the reactivity of ferrous ion and the relatively large amounts of iron used by cells may have necessitated the storage of ferrous ion; recent results suggest that ferrous ion may be stabilized inside ferritin long enough to be used in some types of cells. As primitive organisms began to proliferate, the successful photosynthetic cells, which trapped solar energy by reducing CO2 to make carbohydrates (CH2O)n and produce O2, exhausted from the environment the reductants from H2 or H2S or NH3. The ability of primitive organisms to switch to the use of H2O as a reductant, with the concomitant production of dioxygen, probably produced the worst case of environmental pollution in terrestrial history. As a result, the composition of the atmosphere, the course of biological evolution, and the oxidation state of environmental iron all changed profoundly. Paleogeologists and meteorologists estimate that there was a lag of about 200 - 300 million years between the first dioxygen production and the appearance of significant dioxygen concentrations in the atmosphere, because the dioxygen produced at first was consumed by the oxidation of ferrous ions in the oceans. The transition in the atmosphere, which occurred about 2.5 billion years ago, caused the bioavailability of iron to plummet and the need for iron storage to increase. Comparison of the solubility of Fe3+ at physiological conditions (about 10-18 M) to the iron content of cells (equivalent to 10-5 to 10-8 M) emphasizes the difficulty of acquiring sufficient iron.
Iron is stored mainly in the ferritins, a family* of proteins composed of a protein coat and an iron core of hydrous ferric oxide [Fe2O3(H2O)n] with various amounts of phosphate.6,7 As many as 4,500 iron atoms can be reversibly stored inside the protein coat in a complex that is soluble; iron concentrations equivalent to 0.25 M [about 1016-fold more concentrated than Fe(III) ions] can be easily achieved in vitro (Figure 1.1). Ferritin is found in animals, plants, and even in bacteria; the role of the stored iron varies, and includes intracellular use for Fe-proteins or mineralization, long-term iron storage for other cells, and detoxification of excess iron. Iron regulates the synthesis of ferritin, with large amounts of ferritin associated with iron excess, small or undetectable amounts associated with iron deficiency. [Interestingly, the template (mRNA) for ferritin synthesis is itself stored in cells and is recruited by intracellular iron or a derivative for efficient translation into protein.31 Iron does not appear to interact directly with ferritin mRNA nor with a ferritin mRNA-specific regulatory (binding) protein; however, the specific, mRNA regulatory (binding) protein has sequence homology to aconitase, and formation of an iron-sulfate cluster prevents RNA binding.] Because iron itself determines in part the amount of ferritin in an organism, the environmental concentration of iron needs to be considered before one can conclude that an organism or cell does not have ferritin.
Ferritin is thought to be the precursor of several forms of iron in living organisms, including hemosiderin, a form of storage iron found mainly in animals. The iron in hemosiderin is in a form very similar to that in ferritin, but the complex with protein is insoluble, and is usually located within an intracellular membrane (lysosomes). Magnetite (Fe3O4) is another form of biological iron derived, apparently, from the iron in ferritin. Magnetite plays a role in the behavior of magnetic bacteria, bees, and homing pigeons (see Section II.C).
The structure of ferritin is the most complete paradigm for bioinorganic chemistry because of three features: the protein coat, the iron-protein interface, and the iron core.6,7
* A family of proteins is a group of related but distinct proteins produced in a single organism and usually encoded by multiple, related genes.
Protein Coat
Twenty-four peptide chains (with about 175 amino acids each), folded into ellipsoids, pack to form the protein coat,* which is a hollow sphere about 100 Å in diameter; the organic surface is about 10 Å thick (Figure 1.9). Channels which occur in the protein coat at the trimer interfaces may be involved in the movement of iron in and out of the protein.62,63,65 Since the protein coat is stable with or without iron, the center of the hollow sphere may be filled with solvent, with Fe2O3 • H2O, or, more commonly, with both small aggregates of iron and solvent. Very similar amino-acid sequences are found in ferritin from animals and plants. Sorting out which amino acids are needed to form the shape of the protein coat and the ligands for iron core formation requires the continued dedication of bioinorganic chemists; identification of tyrosine as an Fe(III)-ligand adds a new perspective.64
* Some ferritin subunits, notably in ferritin from bacteria, bind heme in a ratio of less than one heme per two subunits. A possible role of such heme in the oxidation and reduction of iron in the core is being investigated.
Iron-Protein Interface
Formation of the iron core appears to be initiated at an Fe-protein interface where Fe(II)-O-Fe(Ill) dimers and small clusters of Fe(Ill) atoms have been detected attached to the protein and bridged to each other by oxo/hydroxo bridges. Evidence for multiple nucleation sites has been obtained from electron microscopy of individual ferritin molecules (multiple core crystallites were observed) and by measuring the stoichiometry of binding of metal ions, which compete with binding of monoatomic iron, e.g., VO(IV) and Tb(III) (about eight sites per molecule). EXAFS (Extended X-ray Absorption Fine Structure) and Mössbauer spectroscopies suggest coordination of Fe to the protein by carboxyl groups from glutamic (Glu) and aspartic (Asp) acids. Although groups of Glu or Asp are conserved in all animal and plant ferritins, the ones that bind iron are not known. Tyrosine is an Fe(III)-ligand conserved in rapid mineralizing ferritins identified by Uv-vis and resonance Raman spectroscopy.64
Iron Core
Only a small fraction of the iron atoms in ferritin bind directly to the protein. The core contains the bulk of the iron in a polynuclear aggregate with properties similar to ferrihydrite, a mineral found in nature and formed experimentally by heating neutral aqueous solutions of Fe(III)(NO3)3. X-ray diffraction data from ferritin cores are best fit by a model with hexagonal close-packed layers of oxygen that are interrupted by irregularly incomplete layers of octahedrally coordinated Fe(III) atoms. The octahedral coordination is confirmed by Mössbauer spectroscopy and by EXAFS, which also shows that the average Fe(III) atom is surrounded by six oxygen atoms at a distance of 1.95 Å and six iron atoms at distances of 3.0 to 3.3 Å.
Until recently, all ferritin cores were thought to be microcrystalline and to be the same. However, x-ray absorption spectroscopy, Mössbauer spectroscopy, and high-resolution electron microscopy of ferritin from different sources have revealed variations in the degree of structural and magnetic ordering and/or the level of hydration. Structural differences in the iron core have been associated with variations in the anions present, e.g., phosphate29 or sulfate, and with the electrochemical properties of iron. Anion concentrations in turn could reflect both the solvent composition and the properties of the protein coat. To understand iron storage, we need to define in more detail the relationship of the ferritin protein coat and the environment to the redox properties of iron in the ferritin core.
Experimental studies of ferritin formation show that Fe(II) and dioxygen are needed, at least in the early stages of core formation. Oxidation to Fe(III) and hydrolysis produce one electron and an average of 2.5 protons for iron atoms incorporated into the ferritin iron core. Thus, formation of a full iron core of 4,500 iron atoms would produce a total of 4,500 electrons and 11,250 protons. After core formation by such a mechanism inside the protein coat, the pH would drop to 0.4 if all the protons were retained. It is known that protons are released and electrons are transferred to dioxygen. However, the relative rates of proton release, oxo-bridge formation, and electron transfer have not been studied in detail. Moreover, recent data indicate migration of iron atoms during the early stages of core formation and the possible persistence of Fe2+ for periods of time up to 24 hours. When large numbers of Fe(II) atoms are added, the protein coat appears to stabilize the encapsulated Fe(III).34a,b Formation of the iron core of ferritin has analogies to surface corrosion, in which electrochemical gradients are known to occur. Whether such gradients occur during ferritin formation and how different protein coats might influence proton release or alter the structure of the core are subjects only beginning to be examined.
Zinc, Copper, Vanadium, Chromium, Molybdenum, Cobalt, Nickel, and Manganese
Ions of nonferrous transition metals require a much less complex biological storage system, because the solubilities are much higher (≥ 10-8 M) than those for Fe3+. As a result, the storage of nonferrous transition metals is less obvious, and information is more limited. In addition, investigations are more difficult than for iron, because the amounts in biological systems are so small. Essentially nothing is known yet about the storage of vanadium, chromium, molybdenum, cobalt, nickel, and manganese, with the possible exception of accumulations of vanadium in the blood cells of tunicates.
Zinc and copper, which are used in the highest concentrations of any of the non-ferrous transition metals, are specifically bound by the protein metallothionein35,36 (see Figure 1.10). Like the ferritins, the metallothioneins are a family of proteins, widespread in nature and regulated by the metals they bind. In contrast to ferritin, the amounts of metal stored in metallothioneins are smaller (up to twelve atoms per molecule), the amount of protein in cells is less, and the template (mRNA) is not stored. Because the cellular concentrations of the metallothioneins are relatively low and the amount of metal needed is relatively small, it has been difficult to study the biological fate of copper and zinc in living organisms, and to discover the natural role of metallothioneins. However, the regulation of metallothionein synthesis by metals, hormones, and growth factors attests to the biological importance of the proteins. The unusual metal environments of metallothioneins have attracted the attention of bioinorganic chemists.
Metallothioneins, especially in higher animals, are small proteins35,36 rich in cysteine (20 per molecule) and devoid of the aromatic amino acids phenylalanine and tyrosine. The cysteine residues are distributed throughout the peptide chain. However, in the native form of the protein (Figure 1.10), the peptide chains fold to produce two clusters of -SH, which bind either three or four atoms of zinc, cadmium, cobalt, mercury, lead, or nickel. Copper binding is distinct from zinc, with 12 sites per molecule.
In summary, iron is stored in iron cores of a complicated protein. Ferritin, composed of a hollow protein coat, iron-protein interface, and an inorganic core, overcomes the problems of redox and hydrolysis by directing the formation of the quasi-stable mineral hydrous ferric oxide inside the protein coat. The outer surface of the protein is generally hydrophilic, making the complex highly soluble; equivalent concentrations of iron are ≤ 0.25 M. By contrast to iron, storage of zinc, copper, chromium, manganese, vanadium, and molybdenum is relatively simple, because solubility is high and abundance is lower. Little is known about the molecules that store these metals, with the possible exception of metallothionein, which binds small clusters of zinc or copper. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/01%3A_Transition-Metal_Storage_Transport_and_Biomineralization/1.02%3A_Biological_Systems_of_Metal_Storage.txt |
Iron
Iron is the most abundant transition element in the Earth's crust and, in general, in all life forms. An outline of the distribution of iron in the Earth's crust20,21 is shown in Table 1.2. As can be seen, approximately one-third of the Earth's mass is estimated to be iron. Of course, only the Earth's crust is relevant for life forms, but even there it is the most abundant transition element. Its concentration is relatively high in most crustal rocks (lowest in limestone, which is more or less pure calcium carbonate). In the oceans, which constitute 70 percent of the Earth's surface, the concentration of iron is low but increases with depth, since this iron exists as suspended particulate matter rather than as a soluble species. Iron is a limiting factor in plankton growth, and the rich fisheries associated with strong up-welling of ocean depths result at least in part from the biological growth allowed by these iron supplies. Properties that dominate the transport behavior of most transition metal ions are: (l) redox chemistry, (2) hydrolysis, and (3) the solubility of the metal ions in various complexes, particularly the hydroxides.
Table 1.2 - Iron: its terrestrial distribution.a
a) Data from References 1a and 20
One third of Earth's mass, most abundant element by weight
• Distribution in crustal rocks (weight %):
• igneous 5.6
• shale 4.7
• sandstone 1.0
• limestone 0.4
• Ocean (70% of Earth's surface):
• 0.003 - 0.1 ppb, increasing with depth; limiting factor in plankton growth
• Rivers:
• 0.07 - 7 ppm
• Ksp for Fe(OH)3 is approximately 10-39, hence at pH 7 [Fe3+] 10-18 M
As an example of the effects of solubility, consider the enormous variation in the concentration of iron in rivers, depending on whether the water is from a clear mountain stream running over rock or a muddy river carrying large amounts of sediment. However, the amount of dissolved iron in the form of free ferric ion or its hydrolysis products, whatever the source of water, is extremely low. As can be seen from the solubility of hydrated Fe(III) (Ks ~ 10-18 M) (Table 1.2), the concentration of free ferric ion is extraordinarily low at neutral pH; so significant concentrations of soluble iron species can be attained only by strong complex formation.
One example of the versatility of iron as a function of its environment is how the ligand field can strongly alter the structural and ligand exchange properties of the metal ion (Figure 1.3). The ligand field can also alter the redox properties. For high-spin ferric ion, as found in the aquo complex or in many other complexes (including the class of microbial iron-transport agents called siderophores, to be discussed later), the coordination geometry is octahedral or pseudo-octahedral. In the relatively weak ligand field (high-spin ground state), the complex is highly labile. In a strong ligand field, such as an axially ligated porphyrin complex of ferric ion, or the simple example of the ferrocyanide anion, the low-spin complex is exchange-inert. Similarly, the high-spin octahedral ferrous complexes are exchange-labile, but the corresponding axially ligated porphyrin complexes, or the ferrocyanide complexes, are spin-paired (diamagnetic) and ligand exchange-inert. Large, bulky ligands or constrained ligands, such as those provided by metalloprotein and enzyme sites, can cause a tetrahedral environment, in which both ferrous ion and ferric ion form high-spin complexes.
The distribution of specific iron complexes in living organisms depends strongly on function. For example, although there are many different iron complexes in the average human, the relative amounts of each type differ more than 650-fold (Table 1.3). The total amount of iron in humans is quite large, averaging more than three and up to five grams for a healthy adult. Most of the iron is present as hemoglobin, the plasma oxygen-transport protein, where the function of the iron is to deliver oxygen for respiration. A much smaller amount of iron is present in myoglobin, a muscle oxygen-storage protein. For transport, the most important of these iron-containing proteins is transferrin, the plasma iron-transport protein that transfers iron from storage sites in the body to locations where cells synthesizing iron proteins reside; the major consumers of iron in vertebrates are the red blood cells. However, at any given time relatively little of the iron in the body is present in transferrin, in much the same way that at any given time in a large city only a small fraction of the population will be found in buses or taxis. Other examples of iron-containing proteins and their functions are included in Table 1.3 for comparison.
Table 1.3 - Average human Fe distribution
Protein Function Oxidation State of Fe Amount of Fe (g) Percent of Total
Hemoglobin Plasma O2 transport 2 2.6 65
Myoglobin Muscle O2 storage 2 0.13 6
Transferrin Plasma Fe transport 3 0.007 0.2
Ferritin Cell Fe storage 3 0.52 13
Hemosiderin Cell Fe storage 3 0.48 12
Catalase H2O2 metabolism 2 0.004 0.1
Cytochrome c Electron transport $\frac{2}{3}$ 0.004 0.1
Other Oxidases, other enzymes, etc. 0.14 3.6
An example of different iron-coordination environments, which alter the chemical properties of iron, is the difference in the redox potentials of hydrated Fe3+ and the electron-transport protein cytochrome c (Table 1.4). The coordination environment of iron in cytochrome c is illustrated in Figure 1.4. For example, the standard reduction potential for ferric ion in acid solution is 0.77 volts; so here ferric ion is quite a good oxidant. In contrast, cytochrome c has a redox potential of 0.25 volts. A wide range of redox potentials for iron is achieved in biology by subtle differences in protein structure, as listed in Table 1.4. Notice the large difference in the potential of cytochrome c and rubredoxin (Figure 1.5), 0.25 volts vs. -0.06 volts, respectively. In polynuclear ferredoxins, in which each iron is tetrahedrally coordinated by sulfur, reduction potentials are near -0.4 volts. Thus, the entire range of redox potentials, as illustrated in Table 1.4, is more than one volt.
Table 1.4 - Fe redox potentials
Complex Coord. no., type Fe3+/Fe2+ Eo (mV)
Fe(OH2)63+ 6, aquo complex 770
Cytochrome a3 6, heme 390
HIPIP 4, Fe4S4(SR)4- 350
Cytrochrome c 6, heme 250
Rubredoxin 4, Fe(SR)4 -60
Ferredoxins 4, Fe4S4(SR)42- -400
Zinc, Copper, Vanadium, Chromium, Molybdenum, and Cobalt
The chemical properties of the other essential transition elements simplify their transport properties. For zinc there is only the +2 oxidation state, and the hydrolysis of this ion is not a limiting feature of its solubility or transport. Zinc is an essential element for both animals and plants.8,9,20,21 In general, metal ion uptake into the roots of plants is an extremely complex phenomenon. A cross-sectional diagram of a root is shown in Figure 1.6. It is said that both diffusion and mass flow of the soil solution are of significance in the movement of metal ions to roots. Chelation and surface adsorption, which are pH dependent, also affect the availability of nutrient metal ions. Acid soil conditions in general retard uptake of essential divalent metal ions but increase the availability (sometimes with toxic results) of manganese, iron, and aluminum, all of which are normally of very limited availability because of hydrolysis of the trivalent ions.
Vanadium is often taken up as vanadate, in a pathway parallel to phosphate.18 However, its oxidation state within organisms seems to be highly variable. Unusually high concentrations of vanadium occur in certain ascidians (the specific transport behavior of which will be dealt with later). The workers who first characterized the vanadium-containing compound of the tunicate, Ascidia nigra, coined the name tunichrome.22 The characterization of the compound as a dicatecholate has been reported.23
Quite a different chemical environment is found in the vanadium-containing material isolated from the mushroom Amanita muscaria. Bayer and Kneifel, who named and first described amavadine,24 also suggested the structure shown in Figure 1.7.25 Recently the preparation, proof of ligand structure, and (by implication) proof of the complex structure shown in Figure 1.7 have been established.26 Although the exact role of the vanadium complex in the mushroom remains unclear, the fact that it is a vanadyl complex is now certain, although it may take a different oxidation state in vivo.
The role of chromium in biology remains even more mysterious. In human beings the isolation of "glucose tolerance factor" and the discovery that it contains chromium goes back some time. This has been well reviewed by Mertz, who has played a major role in discovering what is known about this elusive and apparently quite labile compound.27 It is well established that chromium is taken up as chromic ion, predominantly via foodstuffs, such as unrefined sugar, which presumably contain complexes of chromium, perhaps involving sugar hydroxyl groups. Although generally little chromium is taken up when it is administered as inorganic salts, such as chromic chloride, glucose tolerance in many adults and elderly people has been reported to be improved after supplementation with 150 - 250 mg of chromium per day in the form of chromic chloride. Similar results have been found in malnourished children in some studies in Third World countries. Studies using radioactively labeled chromium have shown that, although inorganic salts of chromium are relatively unavailable to mammals, brewer's yeast can convert the chromium into a usable form; so brewer's yeast is today the principal source in the isolation of glucose tolerance factor and has been used as a diet supplement.
Although chromium is essential in milligram amounts for human beings as the trivalent ion, as chromate it is quite toxic and a recognized carcinogen.30 The uptake-reduction model for chromate carcinogenicity as suggested by Connett and Wetterhahn is shown in Figure 1.8. Chromate is mutagenic in bacterial and mammalian cell systems, and it has been hypothesized that the difference between chromium in the +6 and +3 oxidation states is explained by the "uptake-reduction" model. Chromium(III), like the ferric ion discussed above, is readily hydrolyzed at neutral pH and extremely insoluble. Unlike Fe3+, it undergoes extremely slow ligand exchange. For both reasons, transport of chromium(III) into cells can be expected to be extremely slow unless it is present as specific complexes; for example, chromium(III) transport into bacterial cells has been reported to be rapid when iron is replaced by chromium in the siderophore iron-uptake mediators. However, chromate readily crosses cell membranes and enters cells, much as sulfate does. Because of its high oxidizing power, chromate can undergo reduction inside organelles to give chromium(m), which binds to small molecules, protein, and DNA, damaging these cellular components.
In marked contrast to its congener, molybdenum is very different from chromium in both its role in biology and its transport behavior, again because of fundamental differences in oxidation and coordination chemistry properties. In contrast to chromium, the higher oxidation states of molybdenum dominate its chemistry, and molybdate is a relatively poor oxidant. Molybdenum is an essential element in many enzymes, including xanthine oxidase, aldehyde reductase, and nitrate reductase.19 The range of oxidation states and coordination geometries of molybdenum makes its bioinorganic chemistry particularly interesting and challenging.
The chemistry of iron storage and transport is dominated by high concentrations, redox chemistry (and production of toxic-acting oxygen species), hydrolysis (pKa is about 3, far below physiological pH), and insolubility. High-affinity chelators or proteins are required for transport of iron and high-capacity sequestering protein for storage. By comparison to iron, storage and transport of the other metals are simple. Zinc, copper, vanadium, chromium, manganese, and molybdenum appear to be transported as simple salts or loosely bound protein complexes. In vanadium or molybdenum, the stable anion, vanadate or molybdate, appears to dominate transport. Little is known about biological storage of any metal except iron, which is stored in ferritin. However, zinc and copper are bound to metallothionein in a form that may participate in storage. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/01%3A_Transition-Metal_Storage_Transport_and_Biomineralization/1.03%3A_Chemical_Properties_Relative_to_Storage_and_Transport.txt |
Many structures formed by living organisms are minerals. Examples include apatite [Ca2(OH)PO4] in bone and teeth, calcite or aragonite (CaCO3) in the shells of marine organisms and in the otoconia (gravity device) of the mammalian ear, silica (SiO2) in grasses and in the shells of small invertebrates such as radiolara, and iron oxides, such as magnetite (Fe3O4) in birds and bacteria (navigational devices) and ferrihydrite FeO(OH) in ferritin of mammals, plants, and bacteria. Biomineralization is the formation of such minerals by the influence of organic macromolecules, e.g., proteins, carbohydrates, and lipids, on the precipitation of amorphous phases, on the initiation of nucleation, on the growth of crystalline phases, and on the volume of the inorganic material.
Iron oxides, as one of the best-studied classes of biominerals containing transition metals, provide good examples for discussion. One of the most remarkable recent characterizations of such processes is the continual deposition of single-crystal ferric oxide in the teeth of chiton.48 Teeth of chiton form on what is essentially a continually moving belt, in which new teeth are being grown and moved forward to replace mature teeth that have been abraded. However, the study of the mechanisms of biomineralization in general is relatively recent; a great deal of the information currently available, whether about iron in ferritin or about calcium in bone, is somewhat descriptive.
Three different forms of biological iron oxides appear to have distinct relationships to the proteins, lipids, or carbohydrates associated with their formation and with the degree of crystallinity.49 Magnetite, on the one hand, often forms almost perfect crystals inside lipid vesicles of magneto-bacteria.50 Ferrihydrite, on the other hand, exists as large single crystals, or collections of small crystals, inside the protein coat of ferritin; however, iron oxides in some ferritins that have large amounts of phosphate are very disordered. Finally, goethite [α-FeO(OH)] and lepidocrocite [γ-FeO(OH)] form as small single crystals in a complex matrix of carbohydrate and protein in the teeth of some shellfish (limpets and chitons); magnetite is also found in the lepidocrocite-containing teeth. The differences in the iron-oxide structures reflect differences in some or all of the following conditions during formation of the mineral: nature of co-precipitating ions, organic substrates or organic boundaries, surface defects, inhibitors, pH, and temperature. Magnetite can form in both lipid and protein/carbohydrate environments, and can sometimes be derived from amorphous or semicrystalline ferrihydrite-like material (ferritin). However, the precise relationship between the structure of the organic phase and that of the inorganic phase has yet to be discovered. When the goal of understanding how the shape and structure of biominerals is achieved, both intellectual satisfaction and practical commercial and medical information will be provided.
Synthetic iron complexes have provided models for two stages of ferritin iron storage and biomineralization:51-59
1. the early stages, when small numbers of clustered iron atoms are bound to the ferritin protein coat, and
2. the final stages, where the bulk iron is a mineral with relatively few contacts to the protein coat. In addition, models have begun to be examined for the microenvironment inside the protein coat.54
Among the models for the early or nucleation stage of iron-core formation are the binuclear Fe(III) complexes with [Fe2O(O2CR2)]2+ cores;55,56 the three other Fe(III) ligands are N. The μ-oxo complexes, which are particularly accurate models for the binuclear iron centers in hemerythrin, purple acid phosphatases, and, possibly, ribonucleotide reductases, may also serve as models for ferritin, since an apparently transient Fe(II)-O-Fe(III) complex was detected during the reconstitution of ferritin from protein coats and Fe(II). The facile exchange of (O2CR) for (O2PR) in the binuclear complex is particularly significant as a model for ferritin, because the structure of ferritin cores varies with the phosphate content. An asymmetric trinuclear (Fe3O)7+ complex57 and an (FeO)11 complex (Figure 1.21) have been prepared; these appear to serve as models for later stages of core nucleation (or growth).59
Models for the full iron core of ferritin include ferrihydrite, which matches the ordered regions of ferritin cores that have little phosphate; however, the site vacancies in the lattice structure of ferrihydrite [FeO(OH)] appear to be more regular than in crystalline regions of ferritin cores. A polynuclear complex of iron and microbial dextran (α-1,4-D-glucose)n has spectroscopic (Mössbauer, EXAFS) properties very similar to those of mammalian ferritin, presumably because the organic ligands are similar to those of the protein (-OH, -COOH). In contrast, a polynuclear complex of iron and mammalian chondroitin sulfate (α-1,4-[α-1,3-D-glucuronic acid-N-acetyl-D-galactosamine-4-sulfate]n) contains two types of domains: one like mammalian ferritin [FeO(OH)] and one like hematite (α-Fe2O3), which was apparently nucleated by the sulfate, emphasizing the importance of anions in the structure of iron cores.60 Finally, a model for iron cores high in phosphate, such as those from bacteria, is Fe-ATP (4:1), in which the phosphate is distributed throughout the polynuclear iron complex, providing an average of 1 or 2 of the 6 oxygen ligands for iron.61
The microenvironment inside the protein coat of ferritin has recently been modeled by encapsulating ferrous ion inside phosphatidylcholine vesicles and studying the oxidation of iron as the pH is raised. The efficacy of such a model is indicated by the observation of relatively stable mixtures of Fe(II)/Fe(III) inside the vesicles, as have also been observed in ferritin reconstituted experimentally from protein coats and ferrous ion.43,54
Models for iron in ferritin must address both the features of traditional metalprotein interactions and the bulk properties of materials. Although such modeling may be more difficult than other types of bioinorganic modeling, the difficulties are balanced by the availability of vast amounts of information on Fe-protein interactions, corrosion, and mineralization. Furthermore, powerful tools such as x-ray absorption, Mössbauer and solid state NMR spectroscopy, scanning electron and proton microscopy, and transmission electron microscopy reduce the number of problems encountered in modeling the ferritin ion core.
Construction of models for biomineralization is clearly an extension of modeling for the bulk phase of iron in ferritin, since the major differences between the iron core of ferritin and that of other iron-biominerals are the size of the final structure, the generally higher degree of crystallinity, and, at this time, the more poorly defined organic phases. A model for magnetite formation has been provided by studying the coulometric reduction of half the Fe3+ atoms in the iron core of ferritin itself. Although the conditions for producing magnetite have .yet to be discovered, the unexpected observation of retention of the Fe2+ by the protein coat has provided lessons for understanding the iron core of ferritin. Phosphatidyl choline vesicles encapsulating Fe2+ appear to serve as models for both ferritin and magnetite; only further investigation will allow us to understand the unique features that convert Fe2+ to [FeO(OH)], on the one hand, and Fe3O4, on the other. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/01%3A_Transition-Metal_Storage_Transport_and_Biomineralization/1.04%3A_Iron_Biomineralization.txt |
The storage of iron in humans and other mammals has been dealt with in the previous section. Only a small fraction of the body's inventory of iron is in transit at any moment. The transport of iron from storage sites in cellular ferritin or hemosiderin occurs via the serum-transport protein transferrin. The transferrins are a class of proteins that are bilobal, with each lobe reversibly (and essentially independently) binding ferric ion.37-39 This complexation of the metal cation occurs via prior complexation of a synergistic anion that in vivo is bicarbonate (or carbonate). Serum transferrin is a monomeric glycoprotein of molecular weight 80 kDa. The crystal structure of the related protein, lactoferrin,39 has been reported, and recently the structure of a mammalian transferrin40 has been deduced.
Ferritin is apparently a very ancient protein and is found in higher animals, plants, and even microbes; in plants and animals a common ferritin progenitor is indicated by sequence conservation.41 In contrast, transferrin has been in existence only relatively recently, since it is only found ia the phylum Chordata. Although the two iron-binding sites of transferrin are sufficiently different to be distinguishable by kinetic and a few other studies, their coordination environments have been known for some time to be quite similar. This was first discovered by various spectroscopies, and most recently was confirmed by crystal-structure analysis, which shows that the environment involves two phenolate oxygens from tyrosine, two oxygens from the synergistic, bidentate bicarbonate anion, nitrogen from histidine, and (a surprise at the time of crystal-structure analysis) an oxygen from a carboxylate group of an aspartate.39
The transferrins are all glycoproteins, and human serum transferrin contains about 6 percent carbohydrate. These carbohydrate groups are linked to the protein, and apparently strongly affect the recognition and conformation of the native protein.
Although transferrins have a high molecular weight and bind only two iron atoms, transferrin is relatively efficient, because it is used in many cycles of iron transport in its interaction with the tissues to which it delivers iron. Transferrin releases iron in vivo by binding to the cell surface and forming a vesicle inside the cell (endosome) containing a piece of the membrane with transferrin and iron still complexed. The release of the iron from transferrin occurs in the relatively low pH of the endosome, and apoprotein is returned to the outside of the cell for delivery of another pair of iron atoms. This process in active reticulocytes (immature red blood cells active in iron uptake) can turn over roughly a million atoms of iron per cell per minute.38 A schematic structure of the protein, deduced from crystal-structure analysis, is shown in Figure 1.11.
Transferrin is an ellipsoidal protein with two subdomains or lobes, each of which binds iron. The two halves of each subunit are more or less identical, and are connected by a relatively small hinge. In human lactoferrin, the coordination site of the iron is the same as the closely related serotransferrin site. A major question that remains about the mechanism of iron binding and release is how the protein structure changes in the intracellular compartment of low pH to release the iron when it forms a specific complex with cell receptors (transferrin binding proteins) and whether the receptor protein is active or passive in the process. Recent studies suggest that the cell binding site for transferrin (a membrane, glycoprotein called the transferrin receptor) itself influences the stability of the iron-transferrin complex. The path of iron from the endosome to Feproteins has not been established; and the form of transported intracellular iron is not known.
Another major type of biological iron transport occurs at the biological opposite of the higher organisms. Although almost all microorganisms have iron as an essential element, bacteria, fungi, and other microorganisms (unlike humans and other higher organisms) cannot afford to make high-molecular-weight protein-complexing agents for this essential element when those complexing agents would be operating extracellularly and hence most of the time would be lost to the organism. As described earlier, the first life forms on the surface of the Earth grew in a reducing atmosphere, in which the iron was substantially more available because it was present as ferrous-containing compounds. In contrast to the profoundly insoluble ferric hydroxide, ferrous hydroxide is relatively soluble at near neutral pH. It has been proposed that this availability of iron in the ferrous state was one of the factors that led to its early incorporation in so many metabolic processes of the earliest chemistry of life.6,38 In an oxidizing environment, microorganisms were forced to deal with the insolubility of ferric hydroxide and hence when facing iron deficiency secrete high-affinity iron-binding compounds called siderophores (from the Greek for iron carrier). More than 200 naturally occurring siderophores have been isolated and characterized to date.42
Most siderophore-mediated iron-uptake studies in microorganisms have been performed by using cells obtained under iron-deficient aerobic growth conditions. However, uptake studies in E. coli grown under anaerobic conditions have also established the presence of siderophore-specific mechanisms. In both cases, uptake of the siderophore-iron complex is both a receptor- and an energy-dependent process. In some studies the dependence of siderophore uptake rates on the concentration of the iron-siderophore complex has been found to conform to kinetics characteristic of protein catalysts, i.e., Michaelis-Menten kinetics. For example, saturable processes with very low apparent dissociation constants of under one micromolar (1 μM) have been observed for ferric-enterobactin transport in E. coli (a bacterium), as shown in Figure 1.12.
Similarly, in a very different microorganism, the yeast Rhodoturala pilimanae, Michaelis-Menten kinetics were seen again with a dissociation constant of approximately 6 μM for the ferric complex of rhodotoroulic acid; diagrams of some representative siderophores are shown in Figure 1.13. The siderophore used by the fungus Neurospora crassa was found to have a dissociation constant of about 5 μM and, again, saturable uptake kinetics.
Although the behavior just described seems relatively simple, transport mechanisms in living cells probably have several more kinetically distinct steps than those assumed for the simple enzyme-substrate reactions underlying the Michaelis-Menten mechanism. For example, as ferric enterobactin is accumulated in E. coli, it has to pass through the outer membrane, the periplasm, and the cytoplasm membrane, and is probably subjected to reduction of the metal in a low-pH compartment or to ligand destruction.
A sketch of a cell of E. coli and some aspects of its transport behavior are shown in Figure 1.14. Enterobactin-mediated iron uptake in E. coli is one of the best-characterized of the siderophore-mediated iron-uptake processes in microorganisms, and can be studied as a model. After this very potent iron-sequestering agent complexes iron, the ferric-enterobactin complex interacts with a specific receptor in the outer cell membrane (Figure 1.14), and the complex is taken into the cell by active transport. The ferric complexes of some synthetic analogs of enterobactin can act as growth agents in supplying iron to E. coli. Such a feature could be used to discover which parts of the molecule are involved in the sites of structural recognition of the ferric-enterobactin complex. Earlier results suggested that the metal-binding part of the molecule is recognized by the receptor, whereas the ligand platform (the triserine lactone ring; see Figure 1.13) is not specifically recognized.
To find out which domains of enterobactin are required for iron uptake and recognition, rhodium complexes were prepared with various domains of enterobactin (Figure 1.15) as ligands to use as competitors for ferric enterobactin.44 The goal was to find out if the amide groups (labeled Domain II in Figure 1.15), which linked the metal-binding catechol groups (Domain III, Figure 1.15) to the central ligand backbone (Domain I, Figure 1.15), are necessary for recognition by the receptor protein. In addition, synthetic ligands were prepared that differed from enterobactin by small changes at or near the catecholate ring. Finally, various labile trivalent metal cations, analogous to iron, were studied to see how varying the central metal ion would affect the ability of metal enterobactin complexes to inhibit competitively the uptake of ferric enterobactin by the organism. For example, if rhodium MECAM (Figure 1.16) is recognized by the receptor for ferric enterobactin on living microbial cells, a large excess of rhodium MECAM will block the uptake of radioactive iron added as ferric enterobactin. In fact, the rhodium complex completely inhibited ferric-enterobactin uptake, proving that Domain I is not required for recognition of ferric enterobactin.
However, if only Domain III is important in recognition, it would be expected that the simple tris(catecholato)-rhodium(III) complex would be an equally good inhibitor. In fact, even at concentrations in which the rhodium-catechol complex was in very large excess, no inhibition of iron uptake was observed, suggesting that Domain II is important in the recognition process.
The role of Domain II in the recognition process was probed by using a rhodium dimethyl amide of 2,3-dihydroxybenzene (DMB) as a catechol ligand, with one more carbonyl ligand than in the tris(catecholato)-rhodium(III) complex. Remarkably, this molecule shows substantially the same inhibition of enterobactin-mediated iron uptake in E. coli as does rhodium MECAM itself. Thus, in addition to the iron-catechol portion of the molecule, the carbonyl groups (Domain II) adjacent to the catechol-binding subunits of enterobactin and synthetic analogs are required for recognition by the ferric-enterobactin receptor. In contrast, when a methyl group was attached to the "top" of the rhodium MECAM complex, essentially no recognition occurred.
In summary, although the structure of the outer-membrane protein receptor of E. coli is not yet known, the composite of the results just described gives a sketch of what the ferric-enterobactin binding site must look like: a relatively rigid pocket for receiving the ferric-catecholate portion of the complex, and proton donor groups around this pocket positioned to hydrogen bond to the carbonyl oxygens of the ferric amide groups. The mechanisms of iron release from enterobactin, though followed phenomenologically, are still not known in detail. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/01%3A_Transition-Metal_Storage_Transport_and_Biomineralization/1.05%3A_Transport_of_Iron.txt |
As described in an earlier section, transport problems posed by the six elements listed in the heading are somewhat simpler (with the exception of chromium) than those for iron. One very interesting recent development has been the characterization of sequestering agents produced by plants which complex a number of metal ions, not just ferric ions. A key compound, now well-characterized, is mugeneic acid (Figure 1.17).45 The structural and chemical similarities of mugeneic acid to ethylenediaminetetraacetic acid (EDTA) have been noted. Like EDTA, mugeneic acid forms an extremely strong complex with ferric ion, but also forms quite strong complexes with copper, zinc, and other transition-metal ions. The structure of the cobalt complex (almost certainly essentially identical with that of the iron complex) is shown in Figure 1.18. Like the siderophores produced by microorganisms, the coordination environment accommodated by mugeneic acid is essentially octahedral. Although the coordination properties of this ligand are well laid out, and it has been shown that divalent metal cations, such as copper, competitively inhibit iron uptake by this ligand, the detailed process of metal-ion delivery by mugeneic acid and related compounds has not been elucidated.
As noted in an earlier section, the biochemistry of vanadium potentially involves four oxidation states that are relatively stable in aqueous solution. These are V2+, V3+, VO2+ , and VO2+ (the oxidation states 2, 3, 4, and 5, respectively). Since even without added sequestering agents, V2+ slowly reduces water to hydrogen gas, it presumably has no biological significance. Examples of the remaining three oxidation states of vanadium have all been reported in various living systems. One of the most extensively investigated examples of transition-metal-ion accumulation in living organisms is the concentration of vanadium in sea squirts (tunicates), which is reported to be variable; many species have vanadium levels that are not exceptionally high. Others such as Ascidia nigra show exceptionally high vanadium concentrations.46
In addition to showing a remarkable concentration of a relatively exotic transition-metal ion, tunicates are a good laboratory model for uptake experiments, since they are relatively simple organisms. They possess a circulation system with a one-chambered heart, and a digestive system that is essentially a pump and an inlet and outlet valve connected by a digestive tract. The organism can absorb dissolved vanadium directly from sea water as it passes through the animal. The influx of vanadate into the blood cells of A. nigra has been studied by means of radioisotopes. The corresponding influx of phosphate, sulfate, and chromate (and the inhibition of vanadate uptake by these structurally similar oxoanions) has been measured. In the absence of inhibitors, the influx of vanadate is relatively rapid (a half-life on the order of a minute near 0 ºC) and the uptake process shows saturation behavior as the vanadate concentration is increased. The uptake process (in contrast to iron delivery in microorganisms, for example, and to many other uptake processes in microorganisms or higher animals) is not energy-dependent. Neither inhibitors of glycolysis nor decouplers of respiration-dependent energy processes show any significant effect on the rate of vanadate influx.
Phosphate, which is also readily taken up by the cells, is an inhibitor of vanadate influx. Neither sulfate nor chromate is taken up significantly, nor do they act as significant inhibitors for the vanadate uptake. Agents that inhibit transport of anions, in contrast, were found to inhibit uptake of vanadate into the organism. These results have led to the model proposed in Figure 1.19:
1. vanadate enters the cell through anionic channels; this process eliminates positively charged metal ion or metal-ion complexes present in sea water;
2. vanadate is reduced to vanadium(III); since the product is a cation, and so cannot be transported through the anionic channels by which vanadate entered the cell, the vanadium(III) is trapped inside the cell-the net result is an accumulation of vanadium. [It has been proposed that the tunichrome could act either as a reducing agent (as the complex) or (as the ligand) to stabilize the general vanadium(III); however, this seems inconsistent with its electrochemical properties (see below).]
Synthetic models of tunichrome b-1 (Figure 1.20) have been prepared. Tunichrome is a derivative of pyrogallol whose structure precludes the formation of an octahedral complex of vanadium as a simple 1:1 metal:ligand complex. The close analogue, described as 3,4,5-TRENPAMH9, also cannot form a simple octahedral 1:1 complex. In contrast, the synthetic ligands TRENCAM and 2,3,4-TRENPAM can form pseudo-octahedral complexes. The structure of the vanadium TRENCAM complex shows that it is indeed a simple pseudo-octahedral tris-catechol complex.47 The electrochemical behavior of these complexes is similar, with vanadium(IV/III) potentials of about -0.5 to -0.6 volts versus NHE. These results indicate that tunichrome b-1 complexes of vanadium(IV/III) would show similar differences in their redox couples at high pH. At neutral pH, in the presence of excess pyrogallol groups, vanadium(IV) can be expected to form the intensely colored tris-catechol species. However, comparison of the EPR properties reported for vanadium-tunichrome preparations with model vanadium(lV)-complexes would indicate predominantly bis(catechol) vanadyl coordination. In any case, the vanadium(III) complexes must remain very highly reducing. It has been pointed out that the standard potential of pyrogallol is 0.79 V and decreases 60 mV per pH unit (up to about pH 9), so that at pH 7 the potential is about 0.4 V. The potentials of the vanadium couples for the tunichrome analogs are about -0.4 V. It has been concluded, therefore, that tunichrome or similar ligands cannot reduce the vanadium(IV) complex; so the highly reducing vanadium(III) complex of tunichrome must be generated in some other way.47
Although a detailed presentation of examples of the known transport properties of essential transition-metal ions into various biological systems could be the subject of a large book, the examples that we have given show how the underlying inorganic chemistry of the elements is used in the biological transport systems that are specific for them. The regulation of metal-ion concentrations, including their specific concentration when necessary from relatively low concentrations of surrounding solution, is probably one of the first biochemical problems that was solved in the course of the evolution of life.
Iron is transported in forms in which it is tightly complexed to small chelators called siderophores (microorganisms) or to proteins called transferrins (animals) or to citrate or mugeneic acid (plants). The problem of how the iron is released in a controlled fashion is largely unresolved. The process of mineral formation, called biomineralization, is a subject of active investigation. Vanadium and molybdenum are transported as stable anions. Zinc and copper appear to be transported loosely associated with peptides or proteins (plants) and possibly mugeneic acid in plants. Much remains to be learned about the biological transport of nonferrous metal ions. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/01%3A_Transition-Metal_Storage_Transport_and_Biomineralization/1.06%3A_Transport_of_Zinc_Copper_Vanadium_Chromium_Molybdenum_and_Cobalt.txt |
I. Introduction
This chapter deals with metalloenzymes wherein the metal acts mainly as a Lewis acid; i.e., the metal does not change its oxidation state nor, generally, its protein ligands. Changes in the coordination sphere may occur on the side exposed to solvent. The substrate interacts with protein residues inside the active cavity and/or with the metal ion in order to be activated, so that the reaction can occur. Under these circumstances the catalyzed reactions involve, as central steps with often complex reaction pathways, the following bond-breaking and/or formation processes:
$\tag{2.1}$
Peptide Hydrolysis
$\tag{2.2}$
Carboxylic Ester Hydrolysis
$\tag{2.3}$
Phosphoric Ester Hydrolysis
$\tag{2.4}$
Nucleophilic Addition of OH- and H-
Scheme (2.3) also pertains to the reactions which need ATP hydrolysis to promote endoenergetic reactions. We will also briefly deal with coenzyme B12; this is a cobalt(III) complex that, by interacting with a number of proteins, produces an R-CH2 radical by homolytic breaking of the Co-C bond as follows:
$\tag{2.5}$
After an R-CH2 radical is formed, it initiates a radical reaction. This is the only system we treat in which the oxidation state changes.
VI. Perspectives
Although a great deal is known about the biophysical characteristics of the various enzyme derivatives mentioned in this chapter, we are still far from a clear understanding of their mechanisms of action, especially if we take into consideration the role of each amino-acid residue inside the active-site cavity. Although we can successfully discuss why certain metal ions are used in certain biological reactions, we still do not know why nickel(II), for example, is involved in the enzymatic hydrolysis of urea.199,200 If we are content with the explanations given in Sections III.A or V.D, we would need model compounds that are good catalysts and perform the job in several steps. This latter requirement would make the various models much more interesting, and would represent a new objective in the investigation of the structure-function relationship of catalytically active molecules. Indeed, the synthesis of large polypeptides may in principle provide such models. In this respect we need to know more about protein folding, for which emerging techniques like protein computer graphics and molecular dynamics are very promising.
Chemical modifications of proteins like the alkylation of carboxylate124,201 or histidine202 residues have been performed for a long time. A newer approach toward modeling the function of a protein, and understanding the role of the active site, involves cleaving part of a naturally occurring protein through enzymatic or chemical procedures, and then replacing it with a synthetic polypeptide. The use of modem techniques of molecular genetics has allowed site-directed mutagenesis to become in principle a very powerful technique for changing a single residue in a cavity. Site-directed mutagenesis is a very popular approach, and its principal limitation with respect to the synthetic polypeptide route is that only natural amino acids can be used (aside from the technical difficulties in both approaches). Small quantities of site-directed mutants have been obtained for CPA125-127 and AP,203 whereas the expression of CA204,205 is now satisfactory.
Predictions of the changes in structure needed to affect the reaction pathway can nowadays be made with the aid of computers. The occurrence of the predicted change can be checked through x-ray analysis and NMR. The latter spectroscopy is today well-recognized as being able to provide structural information on small ($\leq$20 kDa) proteins through 2- or 3-dimensional techniques.206-208 These techniques are increasingly being applied to paramagnetic metalloproteins such as many of those discussed here.208,209 The advantage of handling a paramagnetic metalloprotein is that we can analyze signals shifted far away from their diamagnetic positions, which correspond to protons close to the metal ion,69 even for larger proteins. It is possible to monitor the distances between two or more protons under various conditions, such as after the addition of inhibitors or pseudosubstrates, chemical modification, or substitution of a specific amino acid.
Contributors and Attributions
• Ivano Bertini (University of Florence, Department of Chemistry)
• Claudio Luchinat (University of Bologna, Institute of Agricultural Chemistry)
02: The Reaction Pathways of Zinc Enzymes and Related Biological Catalysts
Introduction
Carbon-dioxide (CO2) hydration and its mechanism in living systems are of fundamental importance for bioinorganic chemistry. In 1932 the existence of an enzyme catalyzing CO2 hydration in red blood cells was established,31 The enzyme was named carbonic anhydrase (abbreviated CA). In 1939 the enzyme was recognized to contain zinc (Zn).32 Because CO2 is either the starting point for photosynthesis or the endpoint of substrate oxidation, carbonic anhydrases are now known to be ubiquitous, occurring in animals, plants, bacteria, and fungi. Different enzymes from different sources, catalyzing the same reaction and usually having homologous structures, are termed isoenzymes. Thus far, a total of 7 distinct classes of CAs have been identified based upon organism: alpha, beta, gamma, delta, zeta, eta, and theta. Each class may contain multiple isoenzymes. Sometimes the same organism has more than one isoenzyme for a particular function, as is true for human carbonic anhydrase. Humans have 15 CAs that belong to the alpha class; these isozymes vary by location in the body and by catalytic activity. CA is a classic example of a hydrolytic enzyme, one that catalyzes addition or removal of water to a substrate molecule. More specifically, CA catalyzes the reversible conversion of carbon dioxide (CO2) to bicarbonate (HCO3-), also referred to as carbonic acid.
Although hydration of CO2 is spontaneous in water at pH 7, the reaction is kinetically slow (k = 10-1 s-1), too slow to convert all CO2 produced in respiration. Only above pH 9 does the uncatalyzed reaction become fast, owing to direct attack of OH-, which is a much better nucleophile than H2O (k = 104 M-1 s-1, where M-1 refers to the OH- concentration). The figure below compares nucleophilic attack of water versus hydroxide (OH-) on CO2.
Between H2O and OH-, formation of HCO3- occurs faster when OH- is the nucleophile. A faster reaction at higher pH, when more OH- ions are present, suggests OH- is involved in the rate determining step. However, realistically, the pH of human blood cannot be changed to speed up hydration of CO2. Instead, humans use carbonic anhydrase to catalyse the reaction. When CA is present, the reaction is sped up to a rate of k = 106 s-1.
The ubiquity of CA in different organisms reflects the importance of these enzymes in sustaining life. The speed of CA-catalyzed CO2 hydration is essential to meet the needs of living cells. Some physiological CA functions include pH regulation, electrolyte secretion, ion transport, and CO2 homeostasis. In the digestive tract, CAs plays a role in the secretion of acid and keep saliva neutral by modulating pH.G,H Among these functions, CA most notably plays a role in transport of CO2 and HCO3- related to respiration, the process of atmospheric oxygen and carbon dioxide exchange that occurs when humans inhale oxygen and exhale carbon dioxide. With low solubility, CO2 must be converted to a more soluble form, HCO3-, for transport throughout the body. Bicarbonate ion (HCO3-) eventually reaches the lungs, gets converted back to CO2, and exits the body through exhalation.C
Medical research revolving around CA focuses on the Zn-containing active site as a therapeutic target for various disease treatments; both CA inhibitors and activators are incorporated in drug design. CA inhibitors are used as treatment for epilepsy, ulcers, cancer, obesity, and other neurological disorders. In the eye, CA produces hydrogen ions that maintain optic pressure. However, too much pressure in the eye can damage the optic nerve and cause glaucoma. CA activity can create a concentration gradient that drives the transport of water to the optical nerve. When too much water is around the optical nerve, pressure around the nerve increases causing damage. Inhibition of CA has become a key treatment of glaucoma.
Beyond pharmaceutical applications, CA has also been investigated for its utility in carbon capture and carbon sensor applications. Carbon capture and storage happens as CAs convert CO2 to bicarbonate. Increased availability of bicarbonate in the presence of calcium ions (Ca2+) causes precipitation of calcite (CaCO3). This process, called bio-mineralization, may be a viable mode of long term storage of CO2 in calcite to mitigate CO2 emissions.
Details about the structure and function of CA’s Zn-containing active site have been elucidated over 80 years of research. The current article delves into the metalloenzyme active site using bioinorganic concepts.
Old Text Below --------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------
Carbon-dioxide hydration and its mechanism in living systems are of fundamental importance for bioinorganic chemistry. In 1932 the existence of an enzyme catalyzing CO2 hydration in red blood cells was established,31 The enzyme was named carbonic anhydrase (abbreviated CA). In 1939 the enzyme was recognized to contain zinc.32 Because CO2 is either the starting point for photosynthesis or the endpoint of substrate oxidation, carbonic anhydrases are now known to be ubiquitous, occurring in animals, plants, and several bacteria. Different enzymes from different sources, catalyzing the same reaction and usually having homologous structures, are termed isoenzymes. Sometimes the same organism has more than one isoenzyme for a particular function, as is true for human carbonic anhydrase.
CO2 gas is relatively soluble in water (3 x 10-2 M at room temperature under pCO2 = 1 atm), equilibrating with hydrogen carbonate at pKa 6.1:
$CO_{2} + H_{2}O \rightleftharpoons HCO_{3}^{-} + H^{-} \tag{2.6}$
The uncatalyzed reaction is kinetically slow around physiological pH (k $\simeq$ 10-1 s-1), whereas, in the presence of the most efficient isoenzyme of CA, the maximal CO2 turnover number (i.e., the number of substrate molecules transformed per unit time by each molecule of enzyme)33 is $\simeq$ 106 s-1. The uncatalyzed attack by water on CO2 may be facilitated by two hydrogen-bonded water molecules, one of which activates the carbon by means of a hydrogen bond to a terminal CO2 oxygen, the other of which binds the carbon atom via oxygen: 34,35
$\tag{2.7}$
Only above pH 9 does the uncatalyzed reaction become fast, owing to direct attack of OH-, which is a much better nucleophile than H2O (k $\simeq$ 104 M-1s-1, where M-1 refers to the OH- concentration):
$CO_{2} + OH^{-} \rightleftharpoons HCO_{3}^{-} \tag{2.8}$
On the other hand, the rate constant in the presence of the enzyme, called kcat, is pH-independent above pH 8 in every CA isoenzyme (Figure 2.2).33,36
In vitro, carbonic anhydrase is quite versatile, catalyzing several reactions that involve both OH- and H+, such as the hydrolysis of esters and the hydration of aldehydes. The various isoenzymes have been characterized to different degrees of sophistication. High-activity forms are labeled II (kcat $\simeq$ 106 s-1 at 25 °C); low-activity forms I (kcat $\simeq$ 105 s-1), and the very-low-activity forms III (kcat $\simeq$ 103 s-1).37 X-ray structural information at nominal 2 Å resolution is available for HCA I38 and HCA II,39 where H indicates human. The structure of HCA II has been refined recently.40 High-resolution structures of mutants and of their substrate and inhibitor derivatives are being reported.211 All isoenzymes are single-chain polypeptides, with M.W. about 30 kDa and one zinc ion per molecule. They have the shape of a rugby ball with a crevice 16 Å deep running through the south pole (Figure 2.3 See color plate section, page C-2.). At the bottom of the crevice, the zinc ion is anchored to the protein by three histidine nitrogen atoms and is exposed to solvent. Two histidines (His-94 and His-96, HCA I numbering) are bound to zinc via their N$\epsilon$2 atoms, whereas one (His-119) is bound via its N$\delta$1 atom (Figure 2.4). It is quite general that histidines bind zinc equally well by either of the two histidine nitrogens, the preference being probably dictated by the steric constraints imposed by the protein folding. The three histidine NH protons are all engaged in H-bonding (Figure 2.4). Histidine-119 is involved in H-bonding with a glutamate residue. As mentioned, this could be a way of controlling the basicity of the metal ligands. A solvent molecule bound to zinc is involved in an H-bond with Thr-199, which in turn is H-bonded to Glu-106. This H-bonding network is important for understanding the subtle structural changes that occur with pH changes; these could, in principle, account for the pH-dependent properties. Although the structure of crystals grown at pH 8 in sulfate-containing buffer gives some indication of a single solvent molecule bound to zinc (Figures 2.3 and 2.5 See color plate section, pages C2, C3.), theoretical studies indicate that two water molecules can be at bonding distances.42 Such a finding is consistent with spectroscopic studies on other derivatives and with the concept that attachment and detachment of substrates occur through five coordination.
Just as is true for every zinc enzyme in which zinc is at the catalytic site, activity is lost if the metal is removed, and is restored by zinc uptake. The tertiary structure of carbonic anhydrase is maintained in the absence of zinc; even the denatured apoprotein can refold spontaneously from a random coil to a native-like conformation. Although such a process is accelerated by zinc,43,44 the presence of the metal does not seem to be an absolute requirement for the correct folding of CA, whereas it is an absolute requirement for several other metalloproteins.23 ,29,30
Anions are attracted in the metal cavity by the positive Zn(N3OH2)2+ moiety, and are believed to bind to zinc in carbonic anhydrase very effectively; so their use should be avoided as much as possible if the goal is to study the enzyme as it is. When the protein is dialyzed against freshly doubly distilled or carefully deionized water under an inert atmosphere, the pH of the sample approaches the isoelectric point, which is below 6 for HCA I and bovine (BCA II) enzymes. The pH can then be adjusted by appropriate additions of NaOH. All the measurements reported in the literature performed in acetate, phosphate, imidazole, or tris sulfate buffers are affected by the interference of the anion with the metal ion. However, buffer species containing large anions like Hepes (4[(2-hydroxyethyl)-l-piperazinyl]ethanesulfonic acid) can be used,45 since these anions do not enter the cavity.
There are many indications that zinc in the high-pH form of CA is four-coordinate with an OH group in the fourth coordination site. At low pH the enzyme exists in a form that contains coordinated water; the coordination number can be four (one water molecule) or five (two water molecules). Of course, the occurrence of the low-pH species depends on the pKa's of the complex acid-base equilibria. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.01%3A_About_Carbonic_Anhydrase.txt |
It is convenient to discuss the cobalt-substituted carbonic anhydrase enzyme, since its electronic spectra are markedly pH-dependent and easy to measure (Figures 2.7 and 2.8).56,57 The spectra are well-shaped, and a sharp absorption at 640 nm is present at high pH and absent at low pH. Whereas CoHCA I is almost entirely in the low-pH form at pH 5.7, this is not true for the CoBCA II isoenzyme. The acid-base equilibrium for Co-substituted carbonic anhydrase (deprotonation of the metal-coordinated water) involves three species:
* An isosbestic point is a value of frequency where the two species in an A $\rightleftharpoons$ B equilibrium have the same absorption. As a consequence, all mixtures of A and B also show the same absorption at that frequency, and all the spectra along, e.g., a pH titration from A to B, plotted one on top of the other, cross at the isosbestic point. The presence of isosbestic points thus indicates the presence of only two species in equilibrium.
These kinds of isoenzymes contain at least another histidine in the cavity, which represents another acidic group, with a pKa of about 6.5 in its free state. The interaction between such an acidic group and metal-coordinated water, for example, via a network of hydrogen bonds, provides a physical picture that can account for the observed experimental data.49 Two apparent acid dissociation constants Ka can be obtained from the fitting of the curves of Figure 2.8. They are called apparent, because they do not represent actual acid dissociations at the microscopic level. When there are two acidic groups interacting with each other, the system must be described in terms of four constants, also called microconstants, because the dissociation of each of the two groups is described by two different pKa's, depending on the ionization state of the other group (Figure 2.9); so the two apparent constants can be expressed in terms of four microconstants describing two interacting acidic groups.
It is again a general feature of these systems that the four microconstants can be obtained only by making some assumptions. In one analysis the molar absorbances of species (1) and (3), and of species (2) and (4), were assumed to be equal.58 In other words, it is assumed that the changes in the electronic spectra of cobalt(II) (Figure 2.7) are due entirely to the ionization of the coordinated water, not at all to the ionization state of the other group. This assumption accounts for the observation of approximate isosbestic points, even though there is an equilibrium between more than two species. With this assumption the four microconstants could be obtained (Table 2.5).
Table 2.5 - Values of microconstants associated with acid-base equilibriaa in cobalt(II)-substituted carbonic anhydrases.58
a) As defined in Figure 2.9
pKa1 pKa2 pKa3 pKa4
CoHCA I 7.14 7.21 8.45 8.38
CoHCA II 5.95 562 6.62 6.95
CoBCA II 6.12 6.28 7.75 7.59
Recall that the activity and spectroscopic profiles follow one another (see Figure 2.2 and Section IV.B). Furthermore, similar microconstant values had been obtained on ZnHCA II by analyzing the pH dependence of the maximum velocity of the hydration reaction, Vmax, assuming that the two hydroxo-containing species had the same activity.49 The present analysis implies that species (2) and (3) of Figure 2.9 are distinguishable, although their interconversion may be fast.
Metal coordination lowers the pKa of coordinated water. Factors affecting the acidity of the coordinated water are many, and their effects are probably overlapping, making the analysis quite complex (see also Section III.A). Nonetheless, the following factors probably contribute to the lowering of the pKa:
1. the charge of the chromophore, which in this case is 2+, although it may be somewhat lowered by the H-bonding between a coordinated histidine and a negative glutamate residue;
2. the coordination number (which is four), since a higher value leads to a larger electron density on the metal ion ligands;
3. the presence of other acidic groups with which the coordinated water interacts;
4. the presence of positively charged residues inside the metal binding cavity that favors the removal of a proton from the cavity.
This last factor is presumably operating for CA III, which contains several arginine residues in the cavity; the same factor may also induce changes in the microscopic properties of the solvent inside the active cavity. These considerations account for the observation that most model complexes have a significantly higher pKa value than the protein itself. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.02%3A_Acid-base_Equilibria.txt |
All the above structural and kinetic information obtained under a variety of conditions with different metal ions can be used to propose a catalytic cycle for carbonic anhydrase (Figure 2.21), As shown by studies on the pH-dependent properties of native and metal-substituted CAs, both type-I and type-II proteins have two acidic groups, the zinc-coordinated water and a free histidine. At physiological pH the enzyme is essentially in the Zn—OH form (step A in Figure 2.21). A Zn—OH moiety is a relatively good nucleophile, poised for nucleophilic attack on carbon dioxide. It is possible that the hydrogen bond with Thr-199, which seems to be consistent with an sp3 oxygen, orients the OH for attack at the substrate CO2 . Studies of the copper derivative indicated that the concentration of CO2 in the cavity is higher than in bulk solution (step B).
Molecular dynamics calculations have shown that there are either three96 or two97 potential wells for CO2 in the hydrophobic pocket. It was shown98 that when Val-143 is replaced by the much larger Phe, the activity decreases by a factor of 103. Apparently the large Phe residue does not leave space within the cavity to accomodate CO2.
It would also be nice if the enzyme were able to activate CO2. There is no evidence that it does, even though the positive charge around zinc and the NH of Thr-199 would represent two electrostatic attraction points that could activate CO2. It is well-known that CO2's interactions with positive charges activate the carbon for nucleophilic attack.99,100 The positioning of CO2 between zinc and the peptide NH of Thr-199 would be ideal for the OH attack. Merz97 locates it as shown in Figure 2.22.
It was believed that, once bicarbonate is formed (C), the proton has to transfer to a terminal oxygen atom, either via an intermediate in which bicarbonate is bidentate (D) or via a hydrogen-bond network (E). Indeed, in model compounds one would expect HCO3- to bind through a nonprotonated oxygen. However, the possibility of restoring the hydrogen bond with Thr-199 as in sulfonamide adducts could justify the presence of the hydrogen on the coordinating oxygen.214 The bicarbonate derivative is presumably in equilibrium between four- and five-coordinate species (F), as shown by the electronic spectra of the cobalt derivative.59 The five-coordinate species provides a low barrier for the substrate detachment step via an associative mechanism involving coordination of a water molecule (G). A possible five-coordinate species would contain bicarbonate in the B site and water in the C site (Figure 2.12). It is reasonable that the measured Km for the reaction of bicarbonate dehydration is the thermodynamic dissociation constant of the M—HCO3- species. Anionic or neutral inhibitors are competitive with bicarbonate because they tend to bind at the same site. At this stage the second substrate, which is H+, has to be released (H). It is reasonable that the water proton transfers to a group inside the cavity, e.g. , the free histidine mentioned above, and subsequently to the solvent. In the absence of buffers the latter step is rate-limiting for the high-activity isoenzymes, since the diffusion rate cannot exceed the product of the concentration times the diffusion coefficient, i.e., 10-7 M x 1011 M-1s-1. Such a limit is then 104 s-1, whereas the turnover rate is 106 s-1. The presence of buffer can assist in proton transfer at this stage, in such a way that the rate-limiting step becomes the internal proton transfer. The release of H+ from the Zn—OH2 moiety is also the rate-limiting step for the low-activity CA III, as nicely shown by the electronic spectra of CoCA III. These spectra change from the basic form at the beginning of the reaction to the acidic form upon CO2 addition (Figure 2.23).101 After the interconversion of CO2 into bicarbonate, there is an accumulation of the CoOH2 species, the deprotonation of which is slower than the release of HCO3-. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.03%3A_Catalytic_Mechanism.txt |
It is quite relevant to know whether a water molecule is coordinated to the metal ion in a metalloenzyme, and whether it is still coordinated in the presence of substrates and inhibitors. The presence or absence of H2O coordinated to a paramagnetic center can in principle be monitored by solvent water 1H NMR,69 by exploiting the occurrence of a magnetic interaction between the magnetic moments of the unpaired electrons and the nuclear magnetic moments of the water protons. When this interaction fluctuates with time, it causes a shortening of the water-proton relaxation times.*
* The nuclear longitudinal relaxation time, T1 can be defined as the rate constant by which the populations of the MI = $\frac{1}{2}$ and MI = $-\frac{1}{2}$ (for protons) levels reach their equilibrium value after an external perturbation (e.g., a radiofrequency pulse in an NMR experiment). The transverse relaxation time, T2, can be defined as the average lifetime of a hydrogen nucleus in a given spin state. The NMR linewidth is inversely proportional to T2. The relation T2 ≤ T1, always holds.
The longitudinal relaxation rate values, T1-1, of all the solvent water protons increase when even a single water molecule interacts with a paramagnetic center, provided that this bound water exchanges rapidly with free water molecules. To obtain the necessary experimental data, a methodology has been developed based on the measurement of water 1H T1-1 values at various magnetic fields (Nuclear Magnetic Relaxation Dispersion, NMRD).69-71 The experimental data contain information on the correlation time, i.e., the time constant for the dynamic process that causes the proton-unpaired electron interaction to fluctuate with time; furthermore, under certain conditions, they may provide quantitative information on the number of interacting protons and their distance to the metal. The enhancement of T1-1, called T1p-1, is caused by the paramagnetic effect on bound water molecules and by the exchange time $\tau$m, according to the relationship
$(T_{1p})^{-1} = f_{M}(T_{1M} + \tau_{M})^{-1} \tag{2.11}$
where fM is the molar fraction of bound water and T1M is the relaxation time of a bound water proton. Therefore we measure the water 1H T1-1, subtract the diamagnetic effect (i.e., the water-proton relaxation rate measured in a solution of a diamagnetic analogue), obtain T1p-1, then check that $\tau$m is negligible with respect to T1M For high-spin cobalt(II), T1M is of the order of 10-3 s, whereas $\tau$m is about 10-5 s. Then the experimental T1p can be safely related to T1M. It is now important, in order to proceed with the analysis, to define the correlation time for the interaction between proton nuclei and unpaired electrons, $\tau$c. Its definition is important in order to obtain a physical picture of the system, and to quantitatively analyze the obtained T1M values.69 $\tau$c is defined by
$\tau_{c}^{-1} = \tau_{r}^{-1} + \tau_{s}^{-1} + \tau_{m}^{-1} \tag{2.12}$
where (\tau\)r is the rotational correlation time, (\tau\)s is the electronic relaxation time, and (\tau\)mhas been previously defined. (\tau\)r depends on the size of the molecule, which can be calculated rigorously if the molecule is spherical, or approximately if it is not. The appropriate expression is
$\tau_{r} = \frac{4\pi \eta a^{3}}{3k_{B}T} \tag{2.13}$
where $\eta$ is the microviscosity of the solution, a is the radius (or approximate radius) of the molecule, kB is the Boltzmann constant, and T is the absolute temperature. For CA, $\tau$r can be safely calculated to be $\simeq$10-8 s at room temperature. Since the correlation time $\tau$c in high-spin cobalt proteins varies between 10-11 and 10-12 s, it must therefore be determined by the electronic relaxation time.
Water 1H NMRD profiles are often analyzed by using the classical dipolar interaction approach, as first described by Solomon:72
$T_{1M}^{-1} = \frac{2}{15} \left(\frac{\mu_{0}}{4 \pi}\right)^{2} \frac{\gamma_{I}^{2} g_{e}^{2} \mu_{B}^{2} S(S+1)}{r^{6}} \left(\frac{7 \tau_{c}}{1 + \omega_{S}^{2} \tau_{c}^{2}} + \frac{3 \tau_{c}}{1 + \omega_{I}^{2} \tau_{c}^{2}} \right) \tag{2.14}$
where $\mu$0 is the permeability of vacuum, $\gamma$I is the nuclear magnetogyric ratio, ge is the electron g-factor, S is the electron spin quantum number, r is the electron-nucleus distance, and $\omega$S and $\omega$I are the electron and nuclear Larmor frequencies, respectively. This equation describes the dipolar interaction between the magnetic moment of nucleus $I(\hbar \gamma_{I} \sqrt{I(I+1)})$ and the magnetic moment of the electrons $S(g_{e} \mu_{B} \sqrt{S(S+1)})$as a function of the correlation time ($\tau$c) and of the magnetic field (expressed as (\omega\)I and $\omega$S). Neglect of the zero-field splitting of the S = $\frac{3}{2}$ manifold may introduce an error in the quantitative estimates within a factor of two. 73
Fitting of the data for pseudotetrahedral complexes shows that they have $\tau$s of 10-11 s, whereas five-coordinate complexes have a shorter $\tau$s, on the order of 10-12 s. The latter derivatives also have exchangeable protons that could correspond to a water molecule in the coordination sphere, whereas the former do not.25 The $\tau$s values are thus proposed as indicators of the coordination number in low-symmetry, four- and five-coordinate cobalt complexes. The shorter electronic relaxation times are related to low-lying excited states, which, independently of the particular mechanism, favor electron relaxation.74
Short electronic relaxation times in paramagnetic compounds cause only minor broadening of 1H NMR lines, whereas the isotropic shifts (i.e., the shifts due to the presence of unpaired electron(s), usually very large) are independent of the value of the electronic relaxation times. For cobalt-substituted carbonic anhydrase, the 1H NMR spectra have been recorded for several derivatives, and the proton signals of histidines coordinated to the metal were found to be shifted well outside the diamagnetic region (Figure 2.14).75 Five-coordinate species give sharper signals than four-coordinate ones. The spectra in D2O for both kinds of derivatives show three fewer isotropically shifted signals than in H2O. These signals are assigned to histidine NH protons, which are replaced by deuterons in D2O. Five-coordinate species provide 1H NMR spectra with many signals slightly shifted from the diamagnetic position. It is believed that such complexes have relatively large magnetic anisotropy, which, summed up to the external magnetic field, provides further differentiation in shifts of the protons. Such shift contributions are called pseudocontact shifts. These shifts depend on the third power of the distance from the metal and on the position of the proton with respect to the molecular axes. These signals belong to protons of noncoordinated residues from 5 to 10 Å from the metal. Their assignment in principle provides further information on the structure in the vicinity of the metal ion. The 1H NMR spectra of cobalt(II) enzymes thus afford a powerful method for monitoring structure and reactivity of the metal-bound residues. This is one task for future investigations of the enzyme. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.04%3A_Coordinated_Water_and_NMR.txt |
The binding of inhibitors is also pH-dependent. It is possible, however, to obtain fully inhibited systems by adjusting the inhibitor concentration and pH. In this manner the so-called limit spectra of CoCA derivatives are obtained. Many systems have been characterized, providing a variety of spectral characteristics59 (Figure 2.10).
The differences in molar absorbance are larger than expected for changing only one coordinated atom. A rationalization of the experimental data came by applying a criterion, first suggested by Gray,60 according to which four-coordinate species have larger maximal absorption than five-coordinate species. This property theoretically arises from greater mixing of p and d metal orbitals in the four-coordinate case, which makes the d-d transitions partially allowed, neglecting other factors such as the covalency of the coordination bond, nephelauxetic effects,* or vicinity of charge transfer bands. Subsequent extension of the measurements to the near-infrared region was instructive:59 the low-intensity spectra exhibited a weak absorption between 13,000 and 15,000 cm-1. The latter band was assigned to the highest in energy of the F → F transitions, which increases in energy with the coordination number. $\dagger$ Therefore both the low intensity of the bands ($\epsilon$max < 200 M-1cm-1) and the presence of the F → F transition at high energy were taken as evidence for five coordination. Spectra showing high maximal absorption ($\epsilon$max > 300 M-1cm-1) were assigned as arising from four-coordinate species. The corresponding chromophores are CoN3In(OH2) and CoN3In, where In denotes inhibitor. Intermediate maximal absorptions may indicate an equilibrium between four- and five-coordinate species. In Table 2.6 some inhibitors are classified according to their behavior. Bicarbonate, which is a substrate of the enzyme, gives rise to an equilibrium between four- and five-coordinate species.48,59
Table 2.6 - Classification of inhibitors of bovine carbonic anhydrase II according to the electronic spectral properties of the adducts with cobalt(lI) derivatives. a 48,59
*Donor sets in parentheses
Four-coordinate Equilibria between four- and five-coordinate species Five-coordinate
Sulphonamides (N4) Bicarbonate (N3O—N3O2) Carboxylates (N3O2)
Cyanide (N3C) Chloride (N3Cl—N3OCl) Thiocyanate (N4O)
Cyanate (N4) Bromide (N3Br—N3OBr) Nitrate (N3O2)
Aniline (N4) Azide (N4—N4O) Iodide (N3OI)
Phenol (N3O)
Chlorate (N3O)
* Nephelauxetic (literally, cloud-expanding) effects are due to partial donation of electrons by the ligand to the metal, and are stronger for less electronegative and more reducing ligands
$\dagger$ By F → F transition we mean here a transition between two electronic states originating from the same F term (the ground term) in the free ion and split by the ligand field; the stronger the ligand field, the larger the splitting. For high-spin cobalt(Il), the free-ion ground state 4F (quartet F) is split in octahedral symmetry into 4T2g , 4T1g , and 4A2g states, the 4T2g lying lowest; in lower symmetries the T states are further split. The highest F → F transition is, therefore, that from the ground state 4T2g , or the lowest of its substates in low symmetry, to the 4A2g state. For the same type of ligands, e.g., nitrogens or oxygens, the ligand field strength, and therefore the energy of the F → F transition, increases with the number of ligands.
The differences in the electronic spectra outlined above also have been detected in both CD and MCD spectra. In the latter, pseudotetrahedral species give a sizably positive band in the high-energy region, whereas five-coordinate species show a much weaker positive band and six-coordinate complexes have only weak negative bands (Figure 2.11).21,61 This additional empirical criterion may be helpful in assigning the coordination number. A further criterion is based on how much of the splitting of the S = $\frac{3}{2}$ ground state is caused by spin-orbit coupling (zero-field splitting). This splitting can be indirectly measured from the temperature dependence of the electronic relaxation times of the cobalt complexes, in turn estimated from their ability to saturate the EPR lines of the complexes at low temperatures.62 There are theoretical reasons to predict that the above splitting increases in the order four coordination < five coordination < six coordination.63
Three binding sites have been identified in the cavity of CA40,64-66 (Figure 2.12). The OH- binding site, which provides a tetrahedral structure around the metal ion, is called the A site. The hydrogen interacts via hydrogen bonding with the oxygen of Thr-199. Thr-199 and Thr-200, together with their protein backbone, identify a hydrophilic region that probably plays a fundamental role in the energetic balance of ligand binding. On the back of the cavity there is a hydrophobic region formed by Val-143, Leu-198, and Trp-209. Although this cavity is hydrophobic, the x-ray structure shows evidence of a water molecule, H-bonded to the coordinated water. Ligands with a hydrophobic end could easily be located in this binding position, which is called B. The coordinated water molecule would change its position in order to make reasonable angles between coordinated groups. The new position is labeled C. The x-ray structure of the thiocyanate derivative of HCA II40,64 illustrates the latter case (see Figure 2.5). The NCS- ion is in van der Waals contact with Val-143, Leu-198, and Trp- 209. The water interacts with the hydroxyl group of Thr-199. The geometry of the five-coordinate derivative can be roughly described as a distorted square pyramid with His-94 in the apical position (Figure 2.13A). This could be a typical structure for those derivatives that have spectra typical of five-coordinate adducts, like the carboxylate derivatives.
In aromatic sulfonamide (Ar—SO2—NH2) derivatives, which probably bind as anions (see Section IV.C.4), the NH- group binds zinc in the A position,64--66 giving rise to an H-bond with Thr-199. The oxygens do not interact with the metal; one of them sits in the hydrophobic pocket. The chromophore around zinc is pseudotetrahedral (Figure 2.13B). The energy involved in the coordination includes the coordination bond, the hydrophobic interactions of the aromatic sulfonamide ring, and the maintainance of the Zn-X-H-Thr-199 hydrogen bonding (X=N,O). It is interesting to note that cyanate, according to spectroscopic studies,48,59 gives rise to tetrahedral derivatives, probably because the terminal oxygen can enter into H-bonds with the hydrophilic region of the cavity. 13C NMR data on N13CO- interacting with CoBCA indicate that the anion interacts directly with the metal ion.67 We do not have direct information on where it binds.212
The fine balance between hydrophobic and hydrophilic interactions, as well as major steric requirements, play important roles in the binding of inhibitors. Cyanide is the only ligand that may bind in a 2:1 ratio.68 It is likely that the bis-cyanide adduct has the same arrangement as the NCS-—H2O derivative. The spin state of the bis-cyanide adduct is S = $\frac{1}{2}$.68 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.05%3A_Coordination_Geometries.txt |
Hydrolysis of carboxylic and phosphoric esters is also a slow process at neutral pH, and is catalyzed by acids and bases by mechanisms similar to those involved in amide and peptide hydrolysis. Metal ions are also good catalysts of both carboxylic and phosphoric ester hydrolysis, typically with rate increases much higher than those observed for hydrolysis of amides or peptides (Table 2.8). The ability of metal ions to coordinate to the carbonyl oxygen—which is higher in amides than in esters—is inversely correlated with their catalytic properties, perhaps because the main role of the metal ion is not in polarizing the carbonyl group, but in providing a metal-coordinated hydroxide as the attacking nucleophile.108 For the hydrolysis of phosphate esters, it is difficult to draw conclusions based on experience with carboxylic esters, because, although the coordinating ability of the phosphoric oxygen may be higher, thus favoring the polarizing role of the metal, the nucleophilic attack is also likely to be easier, because the energy of the trigonal bipyramidal intermediate is probably rather low. Base-catalyzed hydrolysis of phosphate esters occurs with inversion of configuration, and this supports the existence of a trigonal bipyramidal intermediate.140 The metal acts both as activator of substrate through binding and as Lewis acid to provide the OH moiety for the nucleophilic attack:
\(\tag{2.20}\)
As with peptide hydrolysis, several enzyme systems exist that catalyze carboxylic and phosphoric ester hydrolysis without the need for a metal ion. They generally involve a serine residue as the nucleophile; in turn, serine may be activated by hydrogen-bond formation—or even proton abstraction—by other acid-base groups in the active site. The reaction proceeds to form an acyl- or phosphoryl-enzyme intermediate, which is then hydrolyzed with readdition of a proton to the serine oxygen. Mechanisms of this type have been proposed for chymotrypsin.141 In glucose-6-phosphatase the nuc1eophile has been proposed to be a histidine residue.142
Again by analogy with peptide hydrolysis, metalloenzymes catalyzing ester hydrolysis may take advantage of additional chemical features provided by amino-acid residues present in the active-site cavity. This situation occurs with carboxypeptidase,143 which shows esterase activity in vitro. Although the rate-limiting steps for carboxylic esters and peptides may differ, several features, such as the pH dependences of kcat and Km and the presence of two spectroscopically observable intermediates, point to substantially similar mechanisms. On the other hand, carboxylic ester hydrolysis catalyzed by carbonic anhydrase seems to rely on fewer additional features of the active-site cavity, perhaps only on the presence of a metal-coordinated hydroxide that can perform the nucleophilic attack on the carbonyl carbon atom.47
Metalloenzyme-catalyzed phosphoric ester hydrolysis can be illustrated by alkaline phosphatase, by far the most-investigated enzyme of this class. The protein is a dimer of 94 kDa containing two zinc(II) and one magnesium(II) ions per monomer, and catalyzes, rather unspecifically, the hydrolysis of a variety of phosphate monoesters as well as transphosphorylation reactions. The x-ray structure at 2.8 Å resolution obtained on a derivative in which all the native metal ions were replaced by cadmium(II) reveals three metals in each subunit, all located in a single binding region (Figure 2.32). In the native enzyme M1 and M2 sites are occupied by zinc and M3 by magnesium.144 M1 was first reported to be coordinated to three histidine residues (His-331, 372, and 412 in Figure 2.32). Further refinement indicated that Asp-327 could be a ligand to M1, in the place of His-372.145 1H NMR spectroscopy of the enzyme with cobalt substituted in the M1 site shows that there are three exchangeable protons sensing the paramagnetic metal ion.146 They could come from three histidine NHs, or from two histidine NHs and another group containing the exchangeable proton very close to the metal ion, like an arginine. Protein ligands to M2 are Asp-369, His-370, and Asp-51, the latter probably bridging M2 to M3 with the other carboxyl oxygen. M3 is coordinated, in addition, to Asp-51, to Asp-153, to Thr- 155, and to Glu-322. Several spectroscopic pieces of evidence on the native and metal-substituted derivatives indicate that M1 is five-coordinate, but M2 and M3 are six-coordinate, probably with water molecules completing the coordination spheres.28
M1 is essential for activity, but full catalytic efficiency is reached only when all metal ions are present. These data suggest that maximum activity is the result of fine-tuning several chemical properties of the active site as a whole, including the nature of the M1 metal, which can be only zinc or cobalt (Table 2.4).
A further key feature of the active site is the presence of a serine residue (Ser-102), the oxygen atom of which is close to the M1 - M2 pair (especially to M2), although not at direct binding distance according to the crystal structure. There is ample and direct evidence that Ser-102 is reversibly phosphorylated during the course of the catalytic reaction, and that M1 is able to coordinate a phosphate ion.28
Another crucial piece of information obtained by physico-chemical techniques is that the lability of the phosphoseryl intermediate and the catalytic activity increase with pH, depending on the state of ionization of an active-site group, which is most likely a water molecule coordinated to M1.147 Thus the active form of the enzyme is again a metal-hydroxide species. Furthermore, an inactive derivative with copper ions in the M1 and M2 sites shows evidence of magnetic coupling between the metal ions, of the magnitude expected if the two metals shared a common donor atom.148 Likely candidates are a bridging hydroxide ion or Ser-102, which thus might be somewhat mobile relative to the position occupied in the x-ray structure, and demonstrate its potential ability to be activated for the nucleophilic attack by coordination to a metal ion. Such a mechanism would be an "inorganic" version of the type of activation postulated for chymotrypsin and other hydroIases.
A possible mechanism for alkaline phosphatase-catalyzed phosphoric ester hydrolysis could involve the following steps (Figure 2.33):
1. Binding of the phosphate group to M1—in the place of a water molecule—by one of the nonprotonated oxygens, and subsequent activation of the phosphorus atom for nucleophilic attack. The binding of the substrate may be strengthened by interaction with the positively charged Arg-166 residue149 (not shown). The steric alteration in the active site could cause movement of Ser-102 toward M2, with deprotonation upon binding.
2. Nucleophilic attack on phosphorus by the coordinated serine alkoxide, cleaving the ester bond and liberating the alcohol product.
3. Formation of the phosphoseryl intermediate with cleavage of the M1- phosphate bond, decreasing the pKa of the second coordinated water molecule, the proton of which could be taken up by the leaving alcohol.
4. Attack by the metal-coordinated hydroxide on the phosphoryl derivative, possibly with M2 again polarizing the seryl oxygen, yielding a free phosphate ion coordinated to M1. A further water molecule could aid in the liberation of phosphate via an associative mechanism.
In the presence of alcohols, alkaline phosphatase displays transphosphorylation activity, i.e., hydrolysis of the starting ester and esterification of the phosphate group with a different alcohol. This ability is easily understood if one keeps in mind that the reaction depicted above is reversible, and that a different alcohol may be involved in the formation of the ester bond. Most group-transfer reactions catalyzed by metalloenzymes are likely to proceed through the same elementary steps proposed for hydrolytic reactions. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.06%3A_Ester_Hydrolysis_and_Phosphoryl_Transfer.txt |
Group Transfer Enzymes
The phosphodiester bond in ATP and in related molecules is a high-energy bond whose hydrolysis liberates a large quantity of energy:
$ATP + H_{2}O \rightleftharpoons ADP + P_{i} + 30-50\; kJ\; mol^{-1} \tag{2.25}$
In many systems, typically the ATPases, the terminal phosphoryl group is transferred to another acidic group of the enzyme, e.g., a carboxylate group, to form another high-energy bond whose energy of hydrolysis is needed later for some endoenergetic transformation. Therefore the first step of the reaction is the phosphoryl transfer to a group of the enzyme.
Kinases, a subset of the class of transferases, constitute a large group of enzymes that phosphorylate organic substrates:
$ATP_{4-} + HX \rightleftharpoons ADP^{3-} + PX^{2-} + H^{+} \tag{2.26}$
In some kinases, such as nucleoside diphosphate kinase,168,169 an intermediate step is the phosphoryl transfer to a group belonging to the enzyme, as happens in ATPase and as was discussed in detail for alkaline phosphatase (Section V.B). In other kinases the phosphoryl transfer occurs directly from the donor to the acceptor in a ternary complex of the enzyme with the two substrates.170 Often metal ions like magnesium or manganese are needed. These ions interact with the terminal oxygen of the ATP molecule, thus facilitating the nucleophilic attack by the acceptor. The metal ion is often associated with the enzyme. For mechanistic schemes, see the proposed mechanism of action of alkaline phosphatase, especially when a phosphoryl enzyme intermediate is involved.
The B12-dependent Enzymes
There are many enzymes that need a cobalt complex as cofactor in order to carry out vicinal 1,2 interchange:
$\tag{2.27}$
or
$\tag{2.28}$
For the former type of reactions, X can be a group containing either C or N. Typical reactions171 include insertion of a secondary methyl group into a main chain
$\tag{2.29}$
isomerization of an amino group from a primary to a secondary carbon
$\tag{2.30}$
and deamination reactions
$\tag{2.31}$
A list of coenzyme-B12-dependent enzymes is given in Table 2.10.
Table 2.10 - Some coenzyme-B12-dependent enzymes.
MethylmalonylCoA mutase
Glutamate mutase
$\alpha$-Methylene-glutarate mutase
Dioldehydrase
Glyceroldehydrase
Ethanoldeaminase
L-$\beta$-lysine mutase
D-$\alpha$-Iysine mutase
Ribonucleotide reductase
Methionine synthetase
Methane synthetase
Methyl transferase
Acetate synthetase
In coenzyme-B12, cobalt is bound to a tetraazamacrocyclic ligand172 (Figure 2.40). The cobalt atom lies approximately in the plane of the corrin ligand (shown in bold). Note that rings A and D are directly linked. The conjugation therefore extends over only 13 atoms, excluding the cobalt, and involves 14 $\pi$ electrons. Complexes that possess the $\alpha$-D-ribofuranose-3-phosphate and the terminal 5,6-dimethylbenzimidazole as an axial ligand are called cobalamins. The name cobamides applies to complexes that lack or have different heterocyclic groups. Finally, the upper or $\beta$ position is occupied by another ligand, which may be water, OH-, CN-, an alkyl group, etc. The cyano derivative (ii) is vitamin B12. 5'-deoxyadenosylcobalamin (i) is called coenzyme B12. The cobalt atom in these complexes is a diamagnetic cobalt(III) system (d6).
The aquo complex has a pKa of 7.8, which compares with that of 5.7 for the aquopentaamminecobalt(III) complex at 298 K.173 The difference has been mainly ascribed to the difference in solvation of the two complexes, although the corrin ligand bears a negative charge, which reduces the positive charge and therefore the Lewis acidity of the metal ion. The standard reduction potential between pH 2.9 and 7.8 is -0.04 V vs. SCE, featuring the conversion from aquocobalamin with bound benzimidazole (base on) to base-on cob(II)alamin.174 The potential decreases with pH above pH 7.8 down to -0.3 V. The reduced form is five-coordinate, without the water molecule above pH 2.9, and lowspin.175,176 The system can be further reduced at a potential of -0.85 V to obtain cob(I)alamin, in which the metal ion is four-coordinate and low-spin (d8). The standard reduction potential for the hexaaquocobalt(III) complex is 1.95 V, which is lowered to 0.10 for the hexaammine complex, to -0.13 for the tris-ethylenediamine complex, and to -0.80 for the hexacyanocobaltate(III) ion.173 After reduction to cobalt(II), the model complexes are reduced to the metal.
The electronic spectrum of the metal-free corrin resembles that of metal derivatives; it seems therefore that the bands are essentially $\pi$-$\pi$* transitions modified by the central atom and by the axial ligands.177,178 The cob(III)alamins are red, whereas the cob(II)alamins, which are brown, show an additional band at 600 nm.179 The latter have an EPR spectrum typical of an unpaired electron in the dz2 orbital with some 4s mixing: the cob(II)alamin at pH 7 has g$\parallel$ = 2.004, g$\perp$ = 2.32, A$\parallel$(Co) = 0.0100, A$\perp$ (Co) = 0.0027 cm-1, and A$\parallel$ (N) = 0.00173 cm-1.
Both cobalt(I) and cobalt(II)-containing cobalamins readily react with alkyl derivatives to give alkylcob(III)alamins:
$Cob(II)alamin + CH_{3}I \rightarrow CH_{3}-Cob(III)alamin + \frac{1}{2}I_{2} \tag{2.32}$
$Cob(I)alamin + CH_{3}I \rightarrow CH_{3}-Cob(III)alamin + I^{-} \tag{2.33}$
These can formally be regarded as complexes of cobalt(III) with a carbanion. These are rare examples of naturally occurring organometallic compounds. The Co—C bond in alkylcobalamins is relatively weak (bond dissociation energy = 100 kJ mol-1, though higher values are reported in the literature182,183) and can be broken thermally (by heating the complex above 100 °C)182-184 or photochemically, even in daylight exposure.180,181 The energy of the Co—C bond is about 17 kJ mol-1 greater when the transaxial base is absent.184
The cobalamin coenzyme is bound by the apoenzyme with no significant change in the absorption spectrum.185 This suggests that no major change occurs in the coordination of cobalt(III). The first step of the reaction involves homolytic fission of the Co—C bond:182-184, 186-188
$B-CO^{III}-R \rightarrow B-CO^{II.}+R\cdotp \tag{2.34}$
where Band R are the ligands at the $\alpha$ and $\beta$ apical positions. The 5'-deoxyadenosyl radical probably reacts with the substrate, generically indicated as SubH, to give the Sub$\cdotp$ radical and RH. Then the rearrangement reaction proceeds along a not-well-established pathway. It is the protein-substrate binding that controls the subsequent chemistry. In the absence of protein the Co—C bond is kinetically stable; in the presence of protein and substrate the rate of labilization of the Co—C bond increases by a factor of 1011-1012.182-185 By generating the radical in the coenzyme without the protein by means of photolysis or thermolysis, we enable the coenzyme to catalyze some rearrangement reactions without the protein. It may therefore be that the protein plays a major role in inducing the homolytic fission, but a relatively minor role in the subsequent steps, perhaps confined to preventing the various species from diffusing away from each other.
Studies on protein-free corrinoids and model complexes have shown that increasing the steric bulkiness around the coordinated C$\alpha$ atom can cause a dramatic labilization of the Co—C bond.189 The protein-coenzyme adduct might contain the coenzyme in a resting state and the protein in a strained state; the substrate would then switch the system into a strained coenzyme and a relaxed enzyme with little thermodynamic barrier. The strained form of the coenzyme is then in labile equilibrium with base-on cobalt(II) and the free radical.190 This hypothesis, that conformational changes in cobalamin can switch chemical reactions on and off, is closely analogous with the known aspects of hemoglobin function.
It has been suggested that the radical formation in the coenzyme is triggered by a steric perturbation involving an enzyme-induced conformational distortion of the corrin ring toward the deoxyadenosyl group, thereby weakening the cobalt-carbon bond.187,190-194 Structural studies of different corrinoid complexes reveal highly puckered and variable conformations of the corrin ring, attesting to its flexibility.195 For the dimethylglyoxime models, it has been shown that increasing the size of the axial ligand B does induce Co—C bond lengthening and weakening because of conformational distortion of the equatorial ligand away from B and toward the R group.196 It has been proposed that the flexibility of the corrin ligand is the reason why Nature does not use the porphyrin ligand in vitamin B12.197 In an alternative explanation, the weakening of the Co—C bond would be an electronic effect associated with the labilization of the Co—N bond.198 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.07%3A_Group_Transfer_and_Vitamin_B-12.txt |
With zinc enzymes, metallosubstitution is a convenient tool for monitoring the protein and its function by means of spectroscopic techniques. Furthermore, it is interesting to learn how reactivity depends on the nature of the metal ion and its coordination properties, because much of it depends on the protein structure, which seemingly remains constant. As discussed, zinc enzymes can be studied by replacing zinc with other spectroscopically useful metal ions, whose activities have been checked, and by transferring the information obtained to the native enzyme. The strategy of metal substitution is not limited to zinc enzymes, since it has been used for magnesium-activated enzymes and, occasionally, other metalloenzymes as well.
By dialyzing a protein solution against chelating agents, such as EDTA, 1,10-phenanthroline, or 2,6-dipicolinic acid at moderately acidic pH, or by reversibly unfolding the protein with denaturing agents (as has been done with alkaline phosphatase), one can cause zinc proteins to release their metal ions, giving rise to the corresponding but inactive apoprotein. Sometimes (e.g., by using alcohol dehydrogenase) dialysis against chelating agents can be applied to a suspension of protein microcrystals. In this way the chelating agent is still able to reach and remove the active site metal by slowly diffusing in the crystals through the hydration water, while the apoprotein is maintained in the native conformation by the crystal packing forces and denaturation is avoided. After the chelating agent is dialyzed out, often against a high-salt (e.g., CIO4-) buffer to reduce nonspecific binding, a new metalloprotein can obtained by addition of the appropriate metal salt.23
Cobalt(II)-substituted zinc proteins often show about as much activity as the native zinc enzymes (Table 2.4). This is a general characteristic of the cobalt-substituted zinc enzymes,24 since the coordination chemistry of cobalt(II) is very similar to that of zinc(II). The two ions also show virtually identical ionic radii. Cobalt(II) derivatives generally display useful electronic spectra. High-spin cobalt(II) ions are paramagnetic, containing three unpaired electrons (S = $\frac{3}{2}$); thus they can also give rise to EPR spectra. The electronic relaxation times, i.e., the average lifetimes of the unpaired electrons in a given spin state of the S manifold ( $- \frac{3}{2}$, $- \frac{1}{2}$, $\frac{1}{2}$, $\frac{3}{2}$), are very short (10-11 to 10-12 s) at room temperature. In order to detect EPR spectra, the sample temperature is usually decreased, often down to liquid helium temperature, to increase the electronic relaxation times and sharpen the EPR linewidths. On the other hand, as the paramagnetic broadening of the NMR lines in such systems is inversely proportional to the electronic relaxation times (see Section IV.C.3), room-temperature 1H NMR spectra of cobalt(II) complexes can be easily detected, even in the absence of chemical exchange. Therefore cobalt(II) is an exceptional probe to monitor the structure and reactivity of zinc enzymes. Of course, the transfer of information from the artificial to the native enzyme must be done with caution. However, if we can understand the functioning of the cobalt enzyme, we then have a reference frame by which to understand the kinetic properties of the native enzyme. The spectroscopic properties of cobalt(Il) in cobalt-substituted proteins have been reviewed.25
Copper(Il)-substituted zinc proteins are generally inactive with respect to the natural and most artificial substrates (Table 2.4). In model compounds copper(Il) is often principally four-coordinate, with at most two more ligands present at metal-ligand distances that are longer than normal coordination bonds. As a consequence, the ability of zinc to switch between four- and five-coordinate species without any appreciable barrier and with usual metal-donor distances is not mimicked by copper. Furthermore, binding at the four principal coordination positions is generally stronger for copper than for zinc. It follows that substrates may have slow detachment kinetics. These properties are unfavorable for catalysis.
Copper(II) can be easily and meaningfully studied by means of electronic spectroscopy. Moreover, the EPR spectra can be recorded even at room temperature because of the long electronic relaxation times, which are of the order of 10-9 s. Because a protein is a macromolecule, it rotates slowly, and the EPR spectra in solution at room temperature look like those of crystalline powders or frozen solutions (powder-like spectra). ENDOR spectra are also easily obtained for copper proteins at low temperatures, because at low temperature the electronic relaxation times are even longer, and saturation of the EPR lines (which is a requirement to obtain ENDOR spectra) is easy to accomplish. The long electronic relaxation times make the broadening effects on the NMR lines of nuclei sensing the metal ion too severe; so these lines, unlike those of cobalt(Il) complexes, generally escape detection. However, if the nucleus under investigation is in fast exchange between a free species in large excess and a bound species, the line may be observed, because the broadening effects are scaled down by a factor equal to the molar fraction of bound species. The nuclear relaxation parameters contain precious structural and/or dynamic information (see Section IV.C.3). The spectroscopic properties of copper(II) in proteins have been extensively reviewed.26,27
Cadmium-substituted zinc proteins may also be active (Table 2.4), although usually at higher pH. This observation is readily explained in terms of the pKa of a coordinated water, which is expected to be higher than that of analogous zinc complexes because the cadmium ion is larger and polarizes the Cd-OH2 bond less.
113Cd and 111Cd are nuclei with relatively high sensitivity for NMR spectroscopic study. The 113Cd chemical shift spans from -200 to 800 ppm relative to CdSO4 in H2O, depending on the number and nature of donor atoms.24,28 Sulfur donor atoms cause larger downfield shifts than oxygens or nitrogens, and the downfield shift increases with decreasing number of donor atoms. Therefore, 113Cd NMR probes have been used extensively to study zinc enzymes, metal-storage proteins like thioneins, and other proteins with cysteine ligands, and chemical shifts in various cadmium proteins, together with the proposed ligand donor set, have been obtained (Figure 2.1).
Manganese(II)-containing proteins give rise to detectable EPR signals; however, their interpretation in terms of structure and dynamics is not always informative. The electronic relaxation times of Mn2+ are the longest among metal ions, of the order of 10-8 s at room temperature and at the magnetic fields of interest. This property and the large S = $\frac{5}{2}$ value account for a large NMR linewidth, even larger than in copper(II) systems. Manganese(II)- and nickel(II)- substituted zinc proteins have often been reported to have fractional activity (Table 2.4).24 Several efforts have been devoted to Mn(II) derivatives, especially by studying the NMR signals of nuclei in molecules that exchange rapidly with the metalloprotein.
Finally, several other metal-substituted zinc metalloprotein derivatives have been prepared, including those of VO2+ , Fe(II), Co(III), Pt(II), and HgCl2. Although these systems add little directly to our understanding of the relationship between structure and function of the enzymes, nonetheless they represent new bioinorganic compounds and are of interest in themselves, or can add information on the coordinating capabilities, and reactivity in general, of the residues present in the active cavity.
Under the heading zinc enzymes there are several enzymes in which zinc is essential for the biological function, but is not present in the catalytic site. Among the most-studied enzymes, zinc has a structural role in superoxide dismutase, where the ligands are three histidines and one aspartate. In alcohol dehydrogenase there is a zinc ion that has a structural role, besides the catalytically active one. The former zinc has four cysteine ligands. Cysteine ligands are also present in zinc thioneins, which are zinc-storage proteins. The recently discovered class of genetic factors containing "zinc fingers" are zinc proteins in which the metal has an essentially structural role.29 Such a role may consist of lowering the folding enthalpy of a protein to induce an active conformation or to stabilize a particular quaternary structure.
Zinc may also have a regulatory role; i.e., it does not participate in the various catalytic steps, but its presence increases the catalytic rate. This is a rather loose but common definition. Typically, zinc in the B site of alkaline phosphatase (Section V.B) has such a role, and the ligands are histidines, aspartates, and water molecules.
The enzymes in which zinc plays a structural or regulatory role will not be further discussed here, because they do not participate in the catalytic mechanisms; see the broader review articles.23 ,29,30 Rather, we will describe in some detail the enzyme carbonic anhydrase, in order to show how researchers have investigated such complicated systems as enzymes. We will discover as we look at the details of the structures and mechanisms of enzymes that there are large differences between reactivities in solution and in enzymatic cavities. The fundamental properties underlying these differences are still not fully understood.
2.09: Model Chemistry
Some efforts have been reported in the literature to simulate the activity of CA and therefore to obtain further information on the mechanism. The pKa of Zn—OH2 moieties in various complexes has been studied as discussed in Section III.A. The electronic spectra of some cobalt analogues have been found to be similar. One such example is shown in Figure 2.24; the complex Co(TPyMA)OH22+ (Table 2.3) 10 provides a five-coordinate adduct with a weakly bound axial nitrogen (Figure 2.25A).
The interconversion between Co(TPyMA)OH22+ and Co(TPyMA)OH+ was studied by electronic spectroscopy (Figure 2.24). Despite the difference in the number of coordinated nitrogens, the difference between the high- and low-pH forms resembles that of the cobalt enzyme (cf. Figure 2.7).10
Table 2.3 shows that only one compound, with zinc(II) as the metal ion, seems to have three nitrogens and a water, whereas all the other models have a higher coordination number15-17 The simple [CoIlI(NH3)5OH]2+ complex has been shown to accelerate the formation of bicarbonate (k = 2 x 102 M-1s-1), but, of course, bicarbonate remains coordinated to the metal because of the kinetic inertness of cobalt(III).102,103 Some relatively ill-defined systems have been reported to have some kind of activity. The ligand shown in Figure 2.25B, with zinc(II) as the metal ion in H2O, accelerates the attainment of the equilibrium104
$CO_{2} + H_{2}O \xrightleftharpoons[k_{-1}]{k_{1}} HCO_{3}^{-} + H^{+} \tag{2.17}$
with kobs = k1 + k-1 $\simeq$ 103 M-1s-1. The system in Figure 2.25C, with Zn2+ and excess imidazole, promotes CO2 hydration, though not the back reaction.105 The cobalt(II) analogue shows no activity.106
It can be concluded that the M—OH group can indeed be involved in one step of the enzymatic pathway. The sophistication of the whole enzymatic function has not yet been fully achieved with the present generation of models, even though the functionalization of both hydrophilic and hydrophobic molecules like cyclodextrins (Figure 2.25C) has also been used.105 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.08%3A_Metal_Substitution.txt |
Nucleophilic addition of OH- ions as a step in enzymatic pathways is not restricted to hydrolytic processes; it often occurs in lyases, the class of enzymes catalyzing removal (or incorporation in the reverse reaction) of neutral molecules such as H2O—but also NH3, CO2, etc.—from a substrate. It is outside the scope of this section to review all other mechanisms involved in lyase reactions, especially because they are not reducible to common steps and because several of them do not require the presence of a metal ion. We restrict ourselves to H2O removal (or incorporation), a widespread feature of which seems to be the splitting of water into the constituents H+ and OH- ions at some step of the mechanism. As an example, the dehydration of 2-phospho-D-glycerate to phosphoenolpyruvate catalyzed by enolase, a Mg-activated enzyme,
$\tag{2.21}$
has been shown by kinetic isotope-effect studies150 to proceed via fast H+ removal from substrate followed by slow release of the product, and finally by release of OH-. The role of a metal ion like magnesium might be to activate the substrate by coordinating the phosphate group, rather than by providing a coordinated hydroxide for nucleophilic attack.
Other lyases, however, contain transition metal ions [often iron(II)], and their main role might well be that of lowering the pKa of water. None of them, however, is yet known well enough to allow a detailed discussion of the molecular mechanism. A striking exception is carbonic anhydrase, which has been so extensively and successfully studied that it is ideal as a case study (Section IV).
Hydride transfer is another elementary process encountered in many enzymatic reactions. Although hydride transfer implies a redox reaction, it also involves nucleophilic attack on substrate as in the foregoing examples. Unlike OH-, hydride ions do not exist in aqueous solutions as free ions. In biological systems hydride is always directly transferred from one organic moiety to another by simultaneous breakage and formation of covalent bonds. The activation energy for this process is much higher than, for example, that of H+ transfer via the formation of hydrogen bonds. Moreover, unlike hydrogen-bonded species, there is no intermediate in the process that can be stabilized by the catalyst. Instead, reacting species can be destabilized in order to lower the activation energy barrier. The role of the enzyme, and of the metal ion when present, is to provide binding sites for both substrates. The enzyme achieves this both geometrically, by allowing for proper orientation of the groups, and electronically, by providing energy to overcome the activation barrier.
These general concepts can be exemplified by liver alcohol dehydrogenases (LADH), dimeric zinc enzymes of 80 kDa that catalyze the following class of reactions using the NADH/NAD+ system as coenzyme (or, really, as cosubstrate):
$\tag{2.22}$
In particular, LADHs catalyze the reversible dehydrogenation of primary and secondary alcohols to aldehydes and ketones, respectively. Other enzymatic activities of LADHs are aldehyde dismutation and aldehyde oxidation.151 The physiological role, although surely related to the metabolism of the above species, is not definitely settled. Much effort is being devoted to understanding the mechanism of action of this class of enzymes, which have obvious implications for the social problem of alcoholism.
Each monomer unit of LADH contains two zinc ions: one coordinated to four cysteine sulfurs, the other coordinated to two cysteine sulfurs, one histidine nitrogen, and a water molecule. The former has no apparent role in catalysis; the latter is essential for catalytic activity. The x-ray structure of the metal-depleted enzyme from horse liver has been solved at 2.4 Å resolution, and that of the holoenzyme at 2.9 Å resolution (Figure 2.34 See color plate section, page C-6.). Many crystal structures are also available for binary complexes with substrates, pseudosubstrates, or coenzymes, as well as for ternary complexes with coenzyme and substrates.152 The very detailed picture emerging from such structural information has helped us understand how LADH functions. As will be evident from the following discussion, elucidation of this mechanism also reveals some important fundamental chemistry.
A key property of the enzyme, established by x-ray data, is the existence of two protein domains in each monomer that are relatively free to rotate relative to each other. The apo- and holo-enzymes exist in the so-called open form, whereas binding of NADH coenzyme induces rotation of one domain, resulting in the so-called closed form153,154 (Figures 2.34 and 2.35). Closure brings the catalytic zinc ion into an ideal position to bind the aldehyde substrate in such a way that the reactive CH2 group of the nicotinamide ring of NADH points toward the carbonyl carbon (Figure 2.35).
The main functions of the metal are thus to orient the substrate geometrically and to polarize the carbon-oxygen bond. Although the latter makes obvious chemical sense for the aldehyde reduction reaction, since polarization of the C=O bond facilitates nucleophilic attack of hydride at the carbonyl carbon, coordination of an alcohol to a metal is expected to decrease the alcohol's tendency to transfer hydride to NAD+, unless the hydroxyl proton is released upon coordination.155
$\tag{2.23}$
Formation of an alkoxide ion as an intermediate has often been questioned, because the pKa of the alcohol would have to be reduced by about 10 units upon coordination.156 The possibility that hydride transfer from alcohol to NAD+ and hydroxyl proton release could occur simultaneously is attractive, but careful experiments have shown that the two steps must be kinetically separate.157 We summarize here the key information that leads to a full, although circumstantial, rationalization of the chemical behavior of the enzyme.
1. The activity versus pH profiles156,158 are bell-shaped, with kcat increasing with a pKa below 7, reaching a plateau, and decreasing with a pKa above 11, and Km increasing with a pKa of about 9.
2. X-ray data show that the zinc ion is accessible to solvent in the open conformation, much less so in the closed conformation when the reduced coenzyme is bound, and inaccessible when the substrate is coordinated to the metal in the ternary complex, extruding all the water molecules from the active site.152 None of the complexes has a coordinated water molecule as a fifth ligand when substrates or inhibitors are bound to the metal. The metal ion is always four coordinate and pseudotetrahedral. Computer graphics reveal beyond any doubt that there is no room for a fifth ligand in the active site, at least in the closed form.
3. Many (although not all) spectroscopic data on metal-substituted derivatives and their binary and ternary complexes have also been interpreted as indicative of a four-coordinate metal.159 Even nickel(II) and copper(II), which have little tendency to adapt to a pseudotetrahedral ligand environment, do so in LADH, the electronic structure of the latter resembling that of blue proteins (Figure 2.36).160
1. The substrate binding site is actually "created" by the closure of the protein (Figure 2.34). The reactive species are thus trapped in an absolutely anhydrous environment. The chromophoric aldehyde DACA has been extensively used as an "indicator" of the polarity of the binding site. Large red shifts of the ligand $\pi$-$\pi$* transition upon binding indicate the polarity of the site to be much higher than in water; there is a further sizeable increase in polarity when NAD+ instead of NADH is bound in the ternary complex.163
$\tag{2.24}$
1. The electronic spectra of the cobalt-substituted derivative are characteristically different when different anions are bound to the metal (Figure 2.37).164 A catalytically competent ternary complex intermediate displays the electronic absorption pattern typical of anion adducts.166
1. From extended kinetics measurements a protonation scheme (Figure 2.38) has been proposed that accounts for the many pKa values observed under different conditions.167 This scheme again requires formation of a coordinated alkoxide intermediate, but has the advantage of rationalizing in a simple way a complex pattern. In essence, the only relevant acid-base group supplied by the enzyme is the metal-coordinated water, which has a pKa of 9.2 in the free enzyme (open form). Upon binding of NADH the pKa increases to 11.2. Since NADH dissociation is the last and rate-limiting step of the alcohol oxidation reaction, the decrease in kcat with this pKa is accounted for by a decrease in dissociation rate of NADH from the hydroxo form. On the other hand, the pKa of water is decreased to 7.6 upon binding of NAD+. These rather large changes in both directions are best explained by a marked sensitivity of the coordinated water molecule to the polarity of the environment, which, with the possible exception of the unligated form that has a more or less "regular" pKa value of 9.2, can be almost completely anhydrous and much different from that of bulk water. The nonpolar nicotinamide ring of NADH decreases the overall electrostatic interactions of the water molecule, whereas the positive charge of NAD+ drastically increases them. In this scheme, the association rates of both coenzymes are predicted to (and, in fact, do) decrease with a pKa of 9.2, the dissociation rate of NAD+ is predicted to (and does) decrease with a pKa of 7.6, and the dissociation rate of NADH is predicted to decrease with a pKa of 11.2 (and, indeed, it is pH-independent up to and above pH 10).
The decrease of kcat at low pH depends on an ionization that in turn depends on the substrate. This pKa must be that of the coordinated alcohol; at too Iowa pH, deprotonation of the coordinated alcohol becomes the rate-limiting step. The pKa values observed for this process range from 6.4 for ethanol to 4.3 for triftuoroethanol. What is surprising for aqueous-solution chemistry—that the pKa of a coordinated alcohol is lower than the pKa of a coordinated water molecule—can now be explained in terms of the different polarity of the two adducts in LADH. In the binary complex with NAD+ (pKa = 7.6), the water molecule is still free to interact through H-bonding with the solvent and partially dissipate the electrostatic charge. In the ternary complex with any alcohol, the R group may prevent access of the solvent to the cavity, decreasing the dielectric constant of the medium. As a consequence, the polarity of the environment is increased. It is interesting to speculate that Nature could have chosen a stronger Lewis acid than a zinc ion coordinated to two negatively charged residues to decrease the pKa of a coordinated alkoxide, but then the pKa of the coordinated water would have simultaneously undergone a parallel and possibly even stronger decrease. Instead, LADH provides a self-regulating environment that is tailored to decrease the pKa of a coordinated alcohol, once properly positioned, more than that of a coordinated water. The full catalytic cycle for the dehydrogenation reaction at pH around 7 can be summarized as follows (Figure 2.39):
1. NAD+ binds to the open, water-containing form of the enzyme with a maximal on-rate. The pKa of water is decreased to 7.6, but water is still mostly unionized.
2. A neutral alcohol molecule enters the crevice between the two domains, and coordinates the zinc ion by displacing the water molecule. The protein is still in the open form.
3. Domain rotation brings the protein into the closed form, excluding all the residual water molecules from the active site; the combined effect of the metal positive charge and of the unshielded positive charge of the nicotinamide ring lowers the pKa of the coordinated alcohol below 7. A proton is expelled from the cavity, possibly via a hydrogen-bond network of protein residues.
4. Direct hydride transfer takes place from the alcohol CH to the 4-position of the properly oriented nicotinamide ring. The resulting ternary complex is an NADH-aldehyde adduct. The polarity of the active site dramatically drops.
5. The aldehyde product leaves and is replaced by a neutral water molecule (its pKa now being 11.2). Additional water molecules can now enter the crevice, favoring the partial opening of the structure.
6. The loss of contacts between the two halves of the channel favors a complete opening and then the release of NADH, whose dissociation rate is maximal and pH-independent. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.10%3A_Nucleophilic_Addition_of_%28OH-%29_and_%28H-%29.txt |
At neutral pH the uncatalyzed hydrolysis of amides or peptides
$R-CO-NH-R' + H_{2}O \rightleftharpoons R-COO^{-}+R'-NH_{3}^{+} \tag{2.18}$
is a slow process, with rate constants as low as 10-11s-1. Peptide hydrolysis catalyzed by carboxypeptidase or thermolysin can attain kcat values of 104 s-1. Organic chemistry teaches us that amide hydrolysis is relatively efficiently catalyzed by acids and bases. The general mechanisms involve protonation of the carbonyl oxygen (or amide nitrogen), and addition of OH- (or of a general nucleophile) to the carbonyl carbon atom. Several organic and inorganic bases have been found to be reasonably efficient catalysts. On the other hand, transition metal aqua-ions or small metal-ion complexes also display catalytic efficiency (Table 2.8).107-114 A metal ion is a Lewis acid, capable of effectively polarizing the carbonyl bond by metal-oxygen coordination. Furthermore, the metal ion can coordinate a hydroxide group in such a way that there is a high OH- concentration at neutral or slightly alkaline pH. It is thus conceivable that a metalloenzyme may combine some or all of these features and provide a very efficient catalyst.
Table 2.8 - Rate constants for amide and ester hydrolysis catalyzed by acids, bases, or metal ions. * autohydrolysis
Compound Catalyst and Conditions Rate Constant Reference
Glycine amide
pH 9.35
Cu2+, pH 9.35
1.9 x 10-5 s-1
2.6 x 10-3 M-1s -1
107
107
[Co(en)2(glycine amide)]3+ pH 9.0a 2.6 x 10-4 s-1 108
D,L-phenylalanine ethylester
pH 7.3
Cu2+, pH 7.3
5.8 x 10-9 s-1
3.4 x 10-2 M-1s-1
109
109
(tn)2O3PO—C6H4NO2 OH- 5.1 M-1s-1 110
CoIII—(tn)2O3PO—C6H4NO2 a 7 x 10-5 s-1 110
Ethyl-$\beta$-phenylpropionate
H+
OH-
5 x 10-3 M-1 s--1
1.3 x 10- 6 M-1s-1
111
111
Adenosine triphosphate
pH 5.3
Cu2+, pH 5.3
5.6 x 10-6 s-1
1.1 x 10-2 M-1s-1
112
112
Glycine methylester
Co2+(1:1), pH 7.9
Cu2+(1:1), pH 7.3
1.6 x 10-2 s-1
4.2 x 10- 2 s-1
113
113
Glycine propylester
Co3+(1:1), pH 0
Co3+(1:1), pH 8.5
1.1 x 10-3 s-1
>1 x 10-2 s-1
114
114
Much experimental work has been done on mimicking ester and especially peptide hydrolysis with model coordination compounds. Most of the work carried out has involved108,110,115,116 cobalt(III), Although such an ion may not be the best conceivable model for zinc-promoted hydrolytic reactions (see Section IV.G), it has the great advantage of being substitutionally inert, thus removing mechanistic ambiguities due to equilibration among isomeric structures in the course of the reaction, Interesting amide hydrolysis reactions also have been described using complexes with other metal ions, such as copper(Il)117 and zinc(Il)118 itself. In recent years efforts have focused on the construction of bifunctional catalysts to better mimic or test the enzymatic function. For instance, phenolic and carboxylic groups can be placed within reach of Co(III)-chelated amides in peptidase models.116 The presence of the phenolic group clearly accelerates amide hydrolysis, but carboxyl groups are ineffective. This model chemistry is too simple to provide insights into the actual enzymatic mechanism, which must start with recognizing the substrate through several steps, orienting it, activating it, performing the reaction, and finally releasing the products. See the more specialized reviews dealing with nonenzymatic reactivity.119-121
From basic knowledge of the chemistry of hydrolytic reactions, the x-ray structures of carboxypeptidase A and a variety of its derivatives with inhibitors as substrate analogues, product analogues, and transition-state analogues have revealed several features of the active site that are potentially relevant for the catalytic mechanism (Figures 2.26-2.28 See color plate section, pages C4, C5.).122 The metal ion is coordinated to two histidine residues (His-69 and His-196), to a glutamate residue that acts as a bidentate ligand (Glu-72), and to a water molecule. The metal is thus solvent-accessible and, as such, can activate the deprotonation of a water molecule to form a hydroxide ion, or polarize the carbonyl oxygen of the substrate by coordinating it in the place of the solvent molecule, or both, if some flexibility of the coordination sphere is allowed. Another glutamic-acid residue (Glu-270) is in close proximity to the metal center. If the role of the metal were mainly to polarize the carbonyl carbon, Glu-270 in its deprotonated form could be positioned to perform a nucleophilic attack on the carbonyl carbon, yielding an anhydride intermediate. Alternatively, the metal could mainly serve to provide a coordinated hydroxide ion that, in turn, could attack the carbonyl carbon; here Glu-270 would help form ZnOH by transferring the proton to the carboxylate group:
On the opposite side of the cavity is a tyrosine residue that has been shown to be quite mobile and therefore able to approach the site where the catalytic events occur. The cavity has a hydrophobic pocket that can accommodate the residue, R, of nonpolar C-terminal amino acids of the peptide undergoing hydrolysis (Figures 2.26 and 2.28), thereby accounting for the higher efficiency with which hydrophobic C-terminal peptides are cleaved. Finally, an Asn and three Arg residues are distributed in the peptide-binding domain; Asn-144 and Arg-145 can interact via hydrogen bonds with the terminal carboxyl group. Arg-127 can hydrogen-bond the carbonyl oxygen of the substrate.
All these features have enabled detailed interpretation of many chemical and physico-chemical data at the molecular level. The essential data are as follows:
1. Metal substitution. Table 2.9 lists the divalent metals that have been substituted for zinc(II) in CPA, together with their relative peptidase (and esterase) activities.22 For some of them, the available x-ray data show123 that the active-site structure is essentially maintained. Even the copper derivative is slightly active. The apoenzyme is completely inactive, however.
Table 2.9 - Catalytic activities of metal-substituted carboxypeptidases.a 22 a) Activities are relative to the native enzyme, taken as 100%. b) Some activity toward both peptides and esters has recently been observed. 22
Peptidase Esterase
Apo 0 0
Cobalt 200 110
Nickel 50 40
Manganese 30 160
Cadmium 5 140
Mercury 0 90
Rhodium 0 70
Lead 0 60
Copper b b
1. Active-site modifications. Chemical modification and site-directed mutagenesis experiments suggest that Glu-270 is essential for catalysis.124,125 Tyr-248,126 Tyr-198,127 and one or more of the arginines124 are involved but not essential.
2. Kinetics. kcat/Km pH profiles are bell-shaped, characterized by an acid pKa, limb around 6 and an alkaline pKa limb around 9: kcat increases with the pKa of 6 and then levels off, and Km increases with a pKa of 9. Several lines of evidence suggest that the pKa $\approx$ 6 corresponds to the ionization of the Glu-270-coordinated H2O moiety:
$\tag{2.19}$
Site-directed mutagenesis has ruled out Tyr-248 as the group with the pKa of 9 in the rat enzyme.125,126 Unfortunately, in this enzyme the pKa of 9 is observed in kcat rather than Km; so the situation for the most studied bovine enzyme is still unclear. Tyr-248 favors substrate binding three to five times more than the mutagenized Phe-248 derivative.126 The three possible candidates for this pKa, are the coordinated water, Tyr-248, and the metal-coordinated His-196, whose ring NH is not hydrogen-bonded to any protein residue.128 The x-ray data at different pH values show a shortening of the Zn—O bond upon increasing pH.129 This favors the ZnOH hypothesis.
1. Anion binding. The metal binds anionic ligands only below pH 6, i.e., when Glu-270 is protonated, when Glu-270 is chemically130 or genetically125 modified, or when aromatic amino acids or related molecules are bound in the C-terminal binding domain (Arg-145 + hydrophobic pocket).131-134
2. Intermediates. An anhydride intermediate involving Glu-270 for a slowly hydrolyzed substrate may have been identified.135 Some other intermediates have been observed spectroscopically at subzero temperatures with the cobalt(II) derivative.22,136 Peptides bind in a fast step without altering the spectroscopic properties of cobalt(II), following which a metal adduct forms and accumulates.22 Thus, if an anhydride intermediate is formed, it is further along the catalytic path.
On the basis of these data, and many related experiments, a detailed mechanism can be formulated (Figure 2.29). The incoming peptide interacts with arginine residues through its terminal carboxylate group. The interaction could initially involve Arg-71 (not shown); then the peptide would smoothly slide to its final docking position at Arg-145, while the R residue, if hydrophobic, moves to the hydrophobic pocket (Figure 2.29B). The carbonyl oxygen forms a strong hydrogen bond with Arg-127. Additional stabilization could come from hydrogen bonding of Tyr-248 to the penultimate peptide NH. This adduct might be the first intermediate suggested by cryospectroscopy 22,136 (Figure 2.24).
At this point the metal-bound hydroxide, whose formation is assisted by Glu-270, could perform a nucleophilic attack on the carbonyl carbon activated by Arg-127 and possibly, but not necessarily, by a further electrostatic interaction of the carbonyl oxygen with the metal ion. The structure of the substrate analogue $\alpha$-R-$\beta$-phenylpropionate shows that the carbonyl binds in a bidentate fashion:
The system then evolves toward breaking of the C—N bond, caused by addition of a proton to the amino nitrogen. This proton could come from Glu-270, which thereby returns to the ionized state. The breaking of the peptide bond could be the rate-limiting step.22 The second proton required to transform the amino nitrogen into an NH3+ group could come from the coordinated carboxylic group of the substrate, which now bears one excess proton, again through Glu-270 (Figure 2.29D). The system shown in Figure 2.29D can, in fact, be seen as a ternary complex with a carboxylate ligand and an amino-acid zwitterion, bound synergistically.131-134 Finally, the metal moves back to regain a bidentate Glu-72 ligand, and the cleaved peptide leaves, while a further water molecule adds to the metal ion and shares its proton with the free carboxylate group of Glu-270.
Once the hydrolysis has been performed, the cleaved amino acid still interacts with Arg-145 and with the hydrophobic pocket, whereas the amino group interacts with Glu-270. The carboxylate group of the new terminal amino acid interacts with zinc. This picture, which is a reasonable subsequent step in the catalytic mechanism, finds support from the interaction of L- and D-phenylalanine with carboxypeptidase.131-134,138
This mechanism, essentially based on the recent proposal by Christianson and Lipscomb,137 underlines the role of the Zn—OH moiety in performing the nucleophilic attack much as carbonic anhydrase does. This mechanism can apply with slight changes to thermolysin139 and other proteases. Thermolysin cleaves peptidic bonds somewhere in the peptidic chain. The mechanism could be very similar, involving zinc bound to two histidines and Glu-166 (Figure 2.31).139 Glu-166 is monodentate. The role of Glu-270 in CPA is played by Glu-143 and the role of Arg-127 is played by His-231. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.11%3A_Peptide_Hydrolysis.txt |
The ease with which electronic spectra can be obtained provides a simple way of determining the affinity constants of inhibitors for the cobalt-substituted enzymes. An aliquot of enzyme is diluted in a spectrophotometric cell up to a fixed volume, and the spectrum is measured. Then the spectra are remeasured on samples containing the same amount of enzyme plus increasing amounts of inhibitor in the same cell volume. The pH is rigorously controlled. If solutions of enzyme and inhibitor have the same pH, the pH should be verified after the spectral measurements, in order to avoid contamination from the electrode salt medium. Both absolute values and pH dependences of affinity constants obtained from electronic spectra are the same as those obtained from inhibition measurements, where known, and are comparable to those obtained on the native enzyme.
Although affinity constant values reported in the literature were measured under different experimental conditions of, e.g., pH, buffer type, and buffer concentration, several pH-dependent trends are apparent. According to such dependences, three classes of inhibitors can be identified48 (Figure 2.15).
In the first class, the affinity constant, expressed as log K, decreases linearly with increasing pH. Anions that are weak Lewis bases (Cl-,N3-,CH3COO-, NO3-, etc.) behave in this manner, as do neutral ligands like CH3OH and aniline. An example is shown in Figure 2.16. A qualitative fit to such curves can be obtained using a single pKa. This behavior could be accounted for by assuming that the ligand binds only the low-pH form of the enzyme, in a simplified scheme in which only one pKa value determines the species distribution in CA. We know, however, that the picture is more complex.
Figure 2.16: pH dependence of the apparent affinity constants of nitrate for human I \(\blacksquare\), bovine II (▲), and human II (•) carbonic anhydrases. The curves are best-fit curves obtained assuming non-zero affinity of the anion for species I and 3 of Figure 2.9. The best-fit parameters are reported in Table 2.6. Points in parentheses for HCA I reflect possible binding of a second nitrate ion and have been excluded from the fit.57
If the species distribution calculated according to the scheme of Figure 2.9 is assumed to hold, and if it is assumed that only the two water-containing species (1) and (3) can be bound by the ligand, then actual affinity constants can be evaluated for both species (1) and (3)57 (see Table 2.7). Such constants are similar for the three isoenzymes, whereas the apparent affinity constants at pH 7, for example, mainly depend on the pKa's of the coordinated water according to the values of Table 2.5. Therefore, the low-activity species CA I has larger affinity for anions like nitrate (and bicarbonate) than do the high-activity forms at pH 7.
Table 2.7: Affinity constants of nitrate for species 1 and 3a of cobalt(II)-substituted carbonic anhydrases.57 * As defined in Figure 2.9
HCA I BCA II HCA II
log K1 3.74 ± 0.04 4.01 ± 0.02 4.34 ± 0.04
log K3 2.62 ± 0.06 2.56 ± 0.04 2.61 ± 0.05
A second type of behavior occurs for weak acids like HCN, H2S, and aromatic sulfonamides (ArSO2NH2).76,77 Assuming that the anions (conjugated bases) bind the low-pH species of the enzyme, the bell-shaped plot of log K versus pH (Figure 2.15) can be accounted for. In fact, at low pH, the inhibitors are in the protonated form, which is not suitable for metal binding. At high pH the concentration of the low-pH species of the enzyme decreases. The maximal apparent affinity is experimentally halfway between the pKa of the inhibitor and the "pKa" of the enzyme, treated as if it were only one. The same type of curve is also expected if the high-pH species of the enzyme binds the weak acid. Indeed, kinetic measurements seem to favor this hypothesis for sulfonamides.78
A third type of behavior obtains for inhibitors like imidazole and triazoles, which bind the enzyme with similar affinities over a large range of pH (Figure 2.15),79,80 because both the imidazolate anion and the neutral imidazole can bind to the aquo forms of the enzyme with essentially the same affinity,48,80,81 and the reaction of imidazole with the Zn—OH species cannot be distinguished thermodynamically from the reaction of imidazolate with the aquo forms:
It is possible that the noncoordinated nitrogen can interact with a group in the protein via a hydrogen bond. A candidate could be the NH group of His-200 in HCA I or the hydroxyl group of Thr-200 in HCA II. Indeed, only imidazole and triazoles, which have two nitrogens in 1,3-positions, seem to have this ability.213
In summary, from cobalt substitution we have learned:
1. the coordination geometry of the high- and low-pH forms by means of electronic spectroscopy;
2. the values of the pKa's from the pH dependence of the electronic spectra;
3. the four and five coordination of the various derivatives with exogenous ligands;
4. the affinity constants of exogenous ligands and their pH dependence;
5. a fingerprint in the 1H NMR spectra that can be used to monitor structural variations.
Most of these conclusions can be safely transferred to the native zinc enzyme, although minor differences can occur, for example, in the position of the equilibrium between four- and five-coordinate species. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.12%3A_pH_Dependence_of_Inhibitor_Binding.txt |
Zinc has a specific role in bioinorganic processes because of the peculiar properties of the coordination compounds of the zinc(II) ion.
(1) Zinc(II) can easily be four-, five-, or six-coordinate, without a marked preference for six coordination. The electronic configuration of zinc(II) is 3d10 with two electrons per orbital. In coordination compounds, there is no ligand-field stabilization energy, and the coordination number is determined by a balance between bonding energies and repulsions among the ligands. Tetrahedral four-coordinate complexes have shorter metal-donor distances than five-coordinate complexes, and the latter have shorter ones than six-coordinate complexes (Table 2.2), whereas the ligand repulsion increases in the same order.
Table 2.2 - Average zinc(II)-donor atom distances (A) for some common zinc(II) ligands in four-, five-, and six-coordinate complexes.5
Ligand Coordination
4
5
Number
6
H2O 2.00 2.08 2.10
R-COO- 1.95 2.02 2.07
Imidazole 2.02 2.08
Pyridine 2.06 2.12 2.11
R-NH2 2.06 2.15
R, R'NH 2.19 2.27
The repulsion can be both steric and electronic. In enzymes, zinc(II) usually has coordination numbers smaller than six, so that they have available binding sites in their coordination spheres. Substrate can in principle bind to zinc by substituting for a coordinated water or by increasing the coordination number. This behavior would be typical of Lewis acids, and, indeed, zinc is the most common Lewis acid in bioinorganic chemistry. Zinc could thus substitute for protons in the task of polarizing a substrate bond, e.g., the carbonyl C-O bond of peptides and esters, by accepting a substrate atom (oxygen) as a ligand. This has been shown to be possible in model systems. Relative to the proton, a metal ion with an available coordination position has the advantage of being a "superacid," 6 in the sense that it can exist at pH values where the H3O+ concentration is extremely low. Also, relative to the proton, the double positive charge partly compensates for the smaller electrophilicity due to the smaller charge density.
(2) As a catalyst, zinc in zinc enzymes is exposed to solvent, which for enzymes is almost always water. A coordinated water molecule exchanges rapidly, because ligands in zinc complexes are kinetically labile. This, again, can be accounted for by zinc's lack of preference for a given coordination number. A six-coordinate complex can experience ligand dissociation, giving rise to a five-coordinate complex with little energy loss and then little energetic barrier. On the other side, four-coordinate complexes can add a fifth ligand with little energetic barrier and then another ligand dissociates.7 The coordinated water has a pKa sizably lower than free water. Suitable models have been synthesized and characterized in which a solvent water molecule coordinated to various dipositive metal ions has pKa values as low as 7 (Table 2.3).
Table 2.3 - The pKa values of coordinated water in some metal complexes.
a) dacoda = 1,4-diaza-cyciooctane-1 ,4-diacetate.
b) TPyMA = tris(3,5-dimethyl-1-pyrazolylmethyl)amine.
c) TMC = 1,4,8,11-tetramethyl-1,4,8,11-tetraaza-cyclotetradecane.
d) CR = Schiff base between 2,6-diacetylpyridine and bis(3-aminopropyl)amine.
e) DMAM-PMHD = 1-[(6(dimethylamino)methyl)-2-pyridyl)methyl]hexahydro-1,4-diazepin-5-one.
f) C-PMHD = 1[(6-carboxy)-2-pyridyl)methyl]hexa-hydro-1,4-diazepin-5-one.
g) [12]aneN3 = 1,5,9-triaza-cyclododecane.
h) HP[12]aneN3 = 2-(2-hydroxyphenylate)-1,5,9-triaza-cyclododecane.
i)TImMP = tris(4,5-dimethyl-2-imidazolylmethyl)phosphinoxide.
Complex Note Donor Set pKa Reference
Ca(NO3)2(OH2)42+ O6 10.3 8
Cr(OH2)63+ O6 4.2 8
Cr(NH2)5OH23+ N5O 5.1 8
Mn(OH2)62+ O6 10.5 8
Fe(OH2)63+ O6 1.4 8
Co(OH2)62+ O6 9.8 8
Co(dacoda)OH22+ (a) N2O3 9.4 9
Co(TPyMA)OH22+ (b) N4O 9.0 10
Co(TMC)OH22+ (c) N4O 8.4 11
Co(CR)OH22+ (d) N4O 8.0 12
Co(NH3)5OH23+ N5O 6.2 8
Ni(OH2)62+ O6 10.0 8
Cu(OH2)62+ O6 7.3 8
Cu(DMAM-PMHD)OH22+ (e) N3O2 7.1 13
Cu(C-PMHD)OH2+ (f) N2O3 6.6 14
Zn(OH2)62+ O6 9.0 8
Zn(DMAM-PMHD)OH22+ (e) N3O 9.2 13
Zn(C-PMHD)OH2+ (f) N3O2 7.1 14
Zn(CR)OH22+ (d) N4O 8.7 12
ZN([12]aneN3)OH22+ (g) N3O 7.3 15, 16
Zn(HP[12]aneN3)OH2+ (h) N3O2 10.7 16
Zn(TImMP)OH22+ (i) N3O < 7 17
Co(TImMP)OH22+ (i) N3O 7.8 17
This is the result of the formation of the coordination bond. The oxygen atom donates two electrons to the metal ion and formally becomes positively charged:
Under these conditions a proton is easily released. The nucleophilicity of coordinated water is, of course, decreased with respect to free water, owing to the decreased electronic charge on the oxygen atom, but a significant concentration of M-OH species may exist in neutral solution. In turn, the coordinated hydroxide is a slightly poorer nucleophile than the free OH- ion, but better than water. On the basis of recent MO calculations,18 the order of nucleophilicity for solvent-derived species can be summarized as follows:
\[H_{3}O^{+} <H_{2}O-M^{2+} \simeq H_{2}O < HO-M^{+} \simeq HO^{-}\]
Therefore, at neutral or slightly alkaline pH, the small decrease in efficiency of coordinated vs. free hydroxide ions is more than compensated for by the higher concentration of reactive species available (i.e., HO-M+ vs. HO-). Another common role for zinc enzymes is thus to provide a binding site at which the substrate can be attacked by the metal-coordinated hydroxide:
The pKa of coordinated water in zinc complexes is controlled by the coordination number and by the total charge of the complex, in the sense that it decreases with decreasing coordination number and with increasing positive charge, because a zinc ion, bearing in effect a more positive charge, will have greater attraction for the oxygen lone pair, thus lowering the pKa. Charged ligands affect water pKa's more than does the number of ligands.18 The pKa in metalloproteins is further controlled by the presence of charged groups from protein side chains inside the cavity or by the binding of charged cofactors. The coordinated water may have a pKa as low as 6, as in carbonic anhydrase (see later). On the other hand, the pKa of the coordinated water is 7.6 in liver alcohol dehydrogenase (LADH) when NAD+ is bound, 9.2 in the coenzyme-free enzyme, and 11.2 in the presence of NADH (see Section V.C).
(3) As mentioned before, Zn complexes show facile four- to five-coordinate interconversion. The low barrier between these coordination geometries is quite important, because the substrate may add to the coordination sphere in order to replace the solvent or to be coordinated together with the solvent. If the interconversion between four- and five-coordination is fast, catalysis is also fast.
Thus, to summarize, zinc is a good Lewis acid, especially in complexes with lower coordination numbers; it lowers the pKa of coordinated water and is kinetically labile, and the interconversion among its four-, five-, and six-coordinate states is fast. All of these properties make zinc quite suitable for biological catalysis.19 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.13%3A_Selecting_Zinc.txt |
The CO2 $\rightleftharpoons$ HCO3- interconversion catalyzed by CA is extremely fast. The usual kinetic parameters describing an enzymatic reaction are the turnover number or kinetic constant for the reaction, kcat, and the Michaelis constant Km. In the simple catalytic scheme
$E+S \xrightleftharpoons[k_{-1}]{k_{1}} ES \xrightarrow{k_{2}} E+P,$
where E stands for enzyme, S for substrate, and P for product, Km-1 is given by k1/(k-1 + k2). If k2 is small, kcat = k2 and Km-1 = k1/k-1, the latter corresponding to the thermodynamic affinity constant of the substrate for the enzyme. The pH dependences46 of kcat and Km for CO2 hydration for the high- and low-activity isoenzymes have been determined (Figure 2.2).33,36 It appears that Km is pH-independent, whereas kcat increases with pH, reaching a plateau above pH 8. For bicarbonate dehydration (the reverse of Equation 2.6), H+ is a cosubstrate of the enzyme. The pH dependence of kcat/Km for HCO3- dehydration is also mainly due to kcat, which shows the same pH profile as that for CO2 if the experimental kinetic data are divided by the available concentration of the H+ cosubstrate.47,48 Further measurements have shown that the pH dependence of kcat reflects at least two ionizations if the measurements are performed in the absence of anions.49 The value of kcat reaches its maximum at alkaline pH only when buffer concentrations exceed 10-2 M.50 In other words, the exchange of the proton with the solvent is the rate-limiting step along the catalytic pathway if relatively high concentrations of proton acceptors and proton donors are not provided by a buffer system. This limit results from the high turnover of the enzyme, which functions at the limit imposed by the diffusion rate of the H+ cosubstrate. At high buffer concentration, kcat shows an isotope effect consistent with the occurrence of an internal proton transfer as the new rate-limiting step.51
Measurements of the catalyzed reaction performed at chemical equilibrium starting from mixtures of 12C-18O-labeled HCO3- and 13C-16O-labeled CO2 have shown the transient formation of 13C-18O-labeled species (both CO2 and HCO3-) before 18O-labeled water appears in solution.52 These experiments provided evidence that, at chemical equilibrium, an oxygen atom can pass from HCO3- to CO2 and vice versa several times before being released to water. Furthermore, maximal exchange rates are observed even in the absence of buffers.
Under chemical equilibrium conditions, 13C NMR spectroscopy is particularly useful in investigating substrate interconversion rates, since the rates pass from a slow-exchange regime in the absence of enzyme to fast exchange at sufficient enzyme concentration. In the absence of enzyme two 13C signals are observed, one for CO2 and the other for HCO3-. In the presence of enzyme only one averaged signal is observed (Figure 2.6). Starting from the slow exchange situation, in the absence of enzyme, the increase in linewidth ($\Delta$ $\nu$) of the substrate (A) and product (B) signals (caused by exchange broadening that is caused in turn by the presence of a small amount of catalyst) depends on the exchange rate and on the concentration of each species, according to the following relation:
$\Delta \nu_{A}[A] = \Delta \nu_{B}[B] = \tau_{esch}^{-1} \tag{2.9}$
Therefore, the exchange rate $\tau_{esch}^{-1}$ can be calculated.53 The appearance of the NMR spectrum for different $\tau_{esch}^{-1}$ values is illustrated in Figure 2.6 under the condition [A] = [B]. For the high-activity enzyme it was found that the maximal exchange rates are larger than the maximal turnover rates under steady-state conditions; the ratio between kexch of the high-activity (type II) and low-activity (type I) forms is 50, i.e., larger than the ratio in kcat49,54 This result is consistent with the idea that the rate-limiting step in the steady-state process is an intramolecular proton transfer in the presence of buffer for type II enzymes, whereas it may not be so for the type I enzymes. The exchange is pH-independent in the pH range 5.7 - 8, and does not show a proton-deuteron isotope effect. The apparent substrate binding (HCO3-) is weaker than steady-state Km values, indicating that these values are not true dissociation constants. Chloride is a competitive inhibitor of the exchange.49
A similar investigation was conducted for type I CoHCA at pH 6.3, where the concentrations of CO2 and HCO3- are equal.55 The two lines for the two substrates were found to have different linewidths but equal T 1 values. Measurements at two magnetic fields indicate that the line broadening of the HCO3- resonance is caused by substrate exchange and by a paramagnetic contribution due to bonding. The temperature dependence of the linewidth shows that the latter is determined by the dissociation rate. Such a value is only about 2.5 times larger than the overall CO2 $\rightleftharpoons$ HCO3- exchange-rate constant. Therefore the exchange rate between bound and free HCO3- is close to the threshold for the rate-limiting step. Such an exchange rate is related to the higher affinity of the substrate and anions in general for type I isoenzymes than for type II isoenzymes. This behavior can be accounted for in terms of the pKa of coordinated water (see Section C). | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.14%3A_Steady-State_and_Equilibrium_Kinetics_of_Carbonic_Anhydrase-Catalyzed_.txt |
The Groups to which Zinc(II) is Bound
Zinc(II) is an ion of borderline hardness and displays high affinity for nitrogen and oxygen donor atoms as well as for sulfur. It is therefore found to be bound to histidines, glutamates or aspartates, and cysteines. When zinc has a catalytic role, it is exposed to solvent, and generally one water molecule completes the coordination, in which case the dominating ligands are histidines. It has been noted20 recently that coordinated histidines are often hydrogen-bonded to carboxylates:
It is possible that the increase in free energy for the situation in which the hydrogen is covalently bound to the carboxylate oxygen and H-bonded to the histidine nitrogen is not large compared to kBT. Under these circumstances the protein could determine the degree of imidazolate character of the ligand and therefore affect the charge on the metal.
The binding of zinc(II) (like that of other metal ions) is often determined by entropic factors. Water molecules are released when zinc(II) enters its binding position, thus providing a large entropy increase. Most commonly zinc is bound to three or four protein ligands. Large entropy increases are not observed, however, when zinc(II) binds to small polypeptides like the recently discovered zinc fingers, for here the binding site is not preformed (see Section III.B), and zinc(II) must be present for the protein to fold properly into the biologically active conformation.
The Reactivity of Zinc(II) in Cavities
In the preceding section we discussed the properties of zinc(II) as an ion. These properties are, of course, important in understanding its role in biological catalysis, but it would be too simplistic to believe that reactivity can be understood solely on this basis. Catalysis occurs in cavities whose surfaces are constituted by protein residues. Catalytic zinc is bound to a water molecule, which often is H-bonded to other residues in the cavity and/or to other water molecules. The structure of the water molecules in the cavity cannot be the same as the structure of bulk water. Furthermore, the substrate interacts with the cavity residues through either hydrophilic (H-bonds or electric charges) or hydrophobic (London dispersive forces) interactions. As a result, the overall thermodynamics of the reaction pathway is quite different from that expected in bulk solutions. Examples of the importance of the above interactions will be given in this chapter.
The Investigation of Zinc Enzymes
Direct spectroscopic investigation of zinc enzymes is difficult, because zinc(II) is colorless and diamagnetic; so it cannot be studied by means of electronic or EPR spectroscopy. Its NMR-active isotope, 67Zn, the natural abundance of which is 4.11%, has a small magnetic moment, and cannot (with present techniques) be examined by means of NMR spectroscopy at concentrations as low as 10-3 M. The enzymes could be reconstituted with 67Zn. However, 67Zn has a nuclear quadrupolar moment, which provides efficient relaxation times, especially in slow-rotating proteins and low-symmetry chromophores, making the line very broad.210 Of course, 1H NMR can be useful for the investigation of the native enzymes. However, often the molecular weight is such that the proteins are too large for full signal assignment given the current state of the art. At the moment the major source of information comes from x-ray data. Once the structure is resolved, it is possible to obtain reliable structural information on various derivatives by the so-called Fourier difference map. The new structure is obtained by comparing the Fourier maps of the native and of the derivative under investigation. Many x-ray data are now available on carboxypeptidase (Section V.A) and alcohol dehydrogenase (Section V.C).
The zinc ion can be replaced by other ions, and sometimes the enzymatic activity is retained fully or partially (Table 2.4). These new systems have attracted the interest of researchers who want to learn about the role of the metal and of the residues in the cavity, and to characterize the new systems per se. Spectroscopic techniques can be appropriate for the new metal ions; so it is possible to quickly monitor properties of the new derivative that may be relevant for the investigation of the zinc enzyme.
Table 2.4 - Representative metal-substituted zinc enzymes. Percent activities with respect to the native zinc enzyme in parentheses.a
Enzyme Substituted Metals
Alcohol dehydrogenase Co(II)(70), Cu(II)(1), Cu(I)(8), Cd(II)(30), Ni(II)(12)
Superoxide dismutase Co(II)(90), Hg(II)(90), Cd(II)(70), Cu(II)(100)
Aspartate transcarbamylase Cd(II)(100), Mn(II)(100), Ni(II)(100)
Transcarboxylase Co(II)(100), Cu(II)(0)
RNA polymerase Co(II)(100)
Carboxypeptidase A Mn(II)(30), Fe(II)(30), Co(II)(200), Ni(II)(50), Cu(II)(0),b Cd(II)(5), Hg(II)(0), Co(III)(0), Rh(II)(0), Pb(II)(0)
Thermolysin Co(Il)(200), Mn(II)(10), Fe(II)(60), Mg(II)(2), Cr(II)(2), Ni(II)(2), Cu(II)(2), Mo(II)(2), Pb(II)(2), Cd(II)(2), Nd(III)(2), Pr(III)(2)
Alkaline phosphatase Co(II)(30), Cd(II)(1), Mn(II)(1), Ni(II)(0), Cu(II)(0), Hg(II)(0)
β-Lactamase II
Mn(II)(3), Co(II)(11), Ni(II)(0), Cu(II)(0), Cd(II)(11), Hg(II)(4)
Carbonic anhydrasec Cd(II)(2), Hg(II)(0), Cu(II)(0), Ni(II)(2), Co(II)(50), Co(III)(0), Mn(II)(18), V(IV)O2+(0)
Aldolase Mn(II)(15), Fe(II)(67), Co(II)(85), Ni(II)(11), Cu(II)(0), Cd(II)(0), Hg(II)(0)
Pyruvate carboxylase Co(II)(100)
Glyoxalase Mg(II)(50), Mn(II)(50), Co(II)(50)
a) Taken from Reference 21. b) Recent data indicate nonnegligible catalytic activity.22 c) BCA II, except the value for Cd(II) obtained with HCA II | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.15%3A_Strategies_for_the_Investigation_of_Zinc_Enzymes.txt |
Table 2.1 lists metalloenzymes that catalyze hydrolytic and related reactions. According to the above guidelines the hydrolysis of peptide bonds is catalyzed by enzymes called peptidases that belong to the class of hydrolases (according to the official enzyme classification). Two peptidases (carboxypeptidase and thermolysin) are known in great detail, because their structures have been elucidated by high-resolution x-ray crystallography. They share many features; e.g., their metal ions coordinate to the same kind of protein residues. A discussion of the possible mechanism of carboxypeptidase A will be given in Section V.A. Metallopeptidases are zinc enzymes: generally they are single polypeptide chains with molecular weights in the range 30 to 40 kDa. Metallohydrolases of carboxylic and phosphoric esters are also often zinc enzymes. Alkaline phosphatase will be described in Section V.B as a representative of this class. Magnesium is sometimes involved in hydrolytic reactions. This is common when phosphate groups are involved, probably because the affinity of Mg2+ for phosphate groups is high.1 However, hydrolytic reactions can be performed by other systems (not treated here) like urease, which contains nickel(II),2 or acid phosphatase, which contains two iron ions,3 or aconitase, which contains an Fe4S4 cluster.4
Table 2.1 - Representative metalloenzymes catalyzing hydrolytic and related reactions
Enzyme Metal(s) Function
Carboxypeptidase Zn2+ Hydrolysis of C-terminal peptide residues
Leucine aminopeptidases Zn2+ Hydrolysis of leucine N-terminal peptide residues
Dipeptidase Zn2+ Hydrolysis of dipeptides
Neutral protease Zn2+, Ca2+ Hydrolysis of peptides
Collagenase Zn2+ Hydrolysis of collagen
Phospholipase C Zn2+ Hydrolysis of phospholipids
β-Lactamase II Zn2+ Hydrolysis of
β-lactam ring
Thermolysin Zn2+, Ca2+ Hydrolysis of peptides
Alkaline phosphatase Zn2+, Mg2+ Hydrolysis of phosphate esters
Carbonic anhydrase Zn2+ Hydration of CO2
α-Amylase Ca2+, Zn2+ Hydrolysis of glucosides
Phospholipase A2 Ca2+ Hydrolysis of phospholipids
Inorganic pyrophosphatase Mg2+ Hydrolysis of pyrophosphate
ATPase Mg2+ Hydrolysis of ATP
Na+ - K+ - ATPase Na+, K+ Hydrolysis of ATP with transport of cations
Mg2+ - Ca2+ - ATPase Mg2+, Ca2+ Hydrolysis of ATP with transport of cations
Phosphatases Mg2+, Zn2+ Hydrolysis of phosphate esters
Creatine kinase M2+ Phosphorylation of creatine
Pyruvate kinase M+, M2+ Dephosphorylation of phosphoenolpyruvate
Phosphoglucomutase Mg2+ Phosphate transfer converting glucose-I-phosphate to glucose-6-phosphate
DNA polymerase Mg2+ ( Mn2+) Polymerization of DNA with formation of phosphate esters
Alcohol dehydrogenase Zn2+ Hydride transfer from alcohols to NAD+
Examples of enzymes catalyzing nucleophilic addition of OH- (other than hydrolysis) and H- are carbonic anhydrase and alcohol dehydrogenase. Both are zinc enzymes. In the official biochemical classification of enzymes, carbonic anhydrase belongs to the class of lyases. Lyases are enzymes that cleave C-C, C-O, C-N, or other bonds by elimination, leaving double bonds, or conversely add groups to double bonds. Carbonic anhydrase has a molecular weight around 30 kDa, and is among the most-studied metalloenzymes. It catalyzes the deceivingly simple CO2 hydration reaction. The subtleties of its biological function, unraveled by a combination of techniques, make it an ideal example for bioinorganic chemistry. Section IV is fully dedicated to this enzyme. Alcohol dehydrogenase is a 90-kDa enzyme that catalyzes the reversible transfer of a hydride ion from alcohols to NAD+. Although it is a redox enzyme (in fact, classified as an oxidoreductase) and not a hydrolytic one, it will illustrate a different use that Nature makes of zinc to catalyze nucleophilic attack at carbon (Section V.C). Finally, the enzymatic transfer of organic radicals by enzymes involving coenzyme B12 will be briefly considered. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.16%3A_The_Natural_Catalysts.txt |
From Copper Substitution
The coordination chemistry of CuCA is not yet fully understood, since the electronic spectra are not very pH-sensitive. Nevertheless, the affinity of anions is pH-dependent, as it is for CoCA.82 As could be anticipated from Section III.B, the affinity of anions, including HCO3-, is higher than that of CoCA. Water is usually present in the coordination sphere, along with the anion, as checked by water 1H NMRD.83,84 The steric requirements of the three histidines and of the cavity allow the anion and the water molecule to arrange in an essentially square pyramidal geometry (Figure 2.17).
This is consistent with the electronic and EPR spectra. In particular, the EPR spectra are all axial, with g-values decreasing from 2.31 in the nonligated enzyme to 2.24 in the various anion adducts.84 The water molecule would be in the C site or hydrophilic binding site, and the anion would be in the B site or hydrophobic pocket. His-94 would be in the apical position of the square pyramid. It has been shown by EPR spectroscopy that at low temperature two cyanide anions bind to copper. The donor atoms are two cyanide carbon and two histidine nitrogen atoms in the basal plane, and the third histidine nitrogen in the axial position.85 The hyperfine splitting is observed only with nuclei in the basal plane. It is observed both with 13C nuclei of 13C-enriched CN- and with the two 14N of two histidines. The second cyanide may thus displace the coordinated water (Figure 2.17). Oxalate and sulfonamides displace water from the coordination sphere.85,86 For the oxalate ion this may occur through bidentate behavior. Coordination to an oxygen of the sulfonamide cannot be ruled out, although the electronic and EPR spectra of the sulfonamide complex are more consistent with a pseudotetrahedral chromophore. The SO2 moiety would in any case point toward the B binding site. It is likely that sulfonamides bind as in ZnCA. Bicarbonate also shows less water relaxivity than other monodentate anions.83,84.86
13C NMR spectroscopy has been used to investigate the location of the two substrates, CO2 and HCO3 ,with respect to the metal ion in CuCA.86-88 As was pointed out in Section IV.B, the interconversion between the two species is slow on the NMR timescale in the absence of catalysts. Therefore, two signals are observed (Figures 2.6 and 2.18). In the presence of the catalytically active CoCA, only one signal is observed at suitable enzyme concentrations, and individual information on CO2 binding cannot be obtained.89,90 In the presence of inactive CuCA, two signals are again observed, which are broadened to different extents.
For the HCO3- signal the T2-1 values as estimated from the linewidth are much larger than T1-1. Since the equation for T2-1, analogous to Equation (2.14), would predict similar T1 and T2 values,69,72 a sizeable broadening due to chemical exchange must be present. Indeed, unlike T1p-1 (Equation 2.1l), T2p-1 may be a complicated function of the exchange time $\tau$M and of the isotropic shift, $\Delta \omega$M,
$T_{2p}^{-1} = \frac{f_{M}}{\tau_{M}} \frac{T_{2M}^{-2} + T_{2M}^{-1} \tau_{M}^{-1} + (\Delta \omega_{M})^{2}}{(T_{2M}^{-1} + \tau_{M}^{-1})^{2} + (\Delta \omega_{M})^{2}} \tag{2.15}$
In the slow-exchange region, i.e., when two separate signals are observed and the broadening is due to exchange, $T_{2p}^{-1} = f_{M} \tau_{M}^{-1}$. This region is characterized by a marked increase in linewidth with increasing temperature, as confirmed by measurements at 4 and 25 °C. Therefore, T2p gives a direct measure of $\tau$M.56 The 13C T1-1 values of HCO3- are consistent with bicarbonate bound to the metal. The Cu—C distance would be 2.5 Å if the unpaired electron were completely on the copper ion, as estimated by using Equation (2.1) and a value of $\tau$c = 2.1 x 10-9 s independently obtained from water 1H NMRD.83 This distance is much too short for a coordinated bicarbonate; however, electron delocalization on the bicarbonate ligand may account for such a short calculated distance; the possibility of a bidentate type of ligation cannot be discarded. The dissociation rate, which is very low, by itself accounts for the lack of activity of the derivative.
For CO2, a carbon-copper distance could be calculated if the affinity constants of the substrate for the protein were known. When the binding site, if any, starts being saturated, fast exchange with excess ligand (in this case, CO2) decreases the observed paramagnetic effect. From this behavior, the affinity constant may be estimated. For CO2 the paramagnetic effect remained constant up to 1 M CO2; i.e., the affinity constant is smaller than 1 M-1. This means that practically there is no affinity for copper; yet the paramagnetic effect is paradoxically high.88
Another picture comes by analyzing the NMR data in terms of a pure diffusive model.88 Here Hubbard's equation91 has been used:
$T_{2p}^{-1} = N_{M} \left(\dfrac{\mu_{0}}{4 \pi}\right)^{2} \frac{8 \pi}{225} \frac{\gamma_{I}^{2} n_{e}^{2} \mu_{B}^{2} S(S+1)}{d(D_{N}+D_{M})} \bigg(13f(\omega_{S}, \tau_{D})+3f(\omega_{S}, \tau_{D})\bigg) \tag{2.16}$
where
$f(\omega, \tau_{D}) = \frac{15}{2} I(u)$
$u = [\omega \tau_{D}]^{1/2}$
$I(u) = u^{-5} \{u^{2} - 2 + e^{-u}[(u^{2} - 2) \sin u + (u^{2} + 4u + 2) \cos u]\}$
d is the distance of closest approach, DN and DM are the diffusion coefficients of the molecules containing the nucleus under investigation, and $\tau$D = 2d2/(DN + DM). The experimental paramagnetic effect can be reproduced with a CO2 concentration inside the cavity much larger than the one in the bulk solution. This result indicates that substrate does not bind to a specific site, but probably binds in the hydrophobic region. Note that CO2 is more soluble in organic solvents than in water.
The effect of the cavity is to attract CO2 by interaction either with the metal ion or with a hydrophobic part of the cavity itself. But the affinity constant is in any case lower than expected from the Michaelis constant (see Section IV.B) measured under steady-state conditions, indicating that the latter does not represent the dissociation constant of the enzyme-CO2 system.
In summary, the main information concerning the catalytic cycle obtained from the copper derivative is the structural and kinetic characterization of both CO2 and HCO3- species when they are not interconverting but present within the cavity. In this way we have further proof that HCO3- is bound to the metal and that CO2 is attracted inside the cavity either by hydrophobic interactions or by the metal ion or both. The data obtained on the geometry around copper are consistent with those obtained on cobalt.
From Manganese and Cadmium Substitutions
Several studies have been performed on MnCA. Although CA is not the protein for which Mn(II) has been most extensively used as a paramagnetic probe to map substrates and inhibitors within the metal cavity, by measuring the T2M-1 values of protons of the inhibitor N-acetyl-sulfanilamide, and by assuming that dipolar contributions are dominant, researchers have mapped the orientation of the inhibitor inside the active cavity (Figure 2.19).92 This orientation is consistent with x-ray data on stronger binding sulfonamides.64-66 MnCA is not completely inactive. 13C NMR studies of the CO2 $\rightleftharpoons$ HCO3- interconversion at pH 8.5 showed that the interconversion rate is about 4 percent that of the native enzyme.93 The T1-1 and T2-1 values of H13CO3- suggest that bicarbonate might be bidentate in the central step of the catalytic cycle.93
Data from 113Cd studies that have been performed on CdBCA II and CdHCA I are consistent with the general picture presented here.94 The 113Cd chemical shifts are indeed consistent with a donor set of three nitrogens and two oxygens. The cadmium(II) derivative could thus be five-coordinate with two water molecules, in agreement with the expectation based on its ionic radius being larger than that of zinc(II). The 113Cd signal of CdBCA II in the presence of benzene-sulfonamide enriched in 15N is split into a doublet because of the nitrogen-cadmium coupling (Figure 2.20).95 This result provides direct evidence for metal-nitrogen bonding in sulfonamides, which has been confirmed by x-ray data.65 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/02%3A_The_Reaction_Pathways_of_Zinc_Enzymes_and_Related_Biological_Catalysts/2.17%3A_What_Do_We_Learn.txt |
Calcium, like many other "inorganic elements" in biological systems, has during the last decade become the subject of much attention both by scientists and by the general public.1 The presence and central role of calcium in mammalian bones and other mineralized tissues were recognized soon after its discovery as an element by Davy in 1808. Much later, the insight arrived that Ca2+ ions could play an important role in other tissues as well. Experiments of great historical influence were performed by the British physiologist Sidney Ringer a little over a century ago.2 He was interested in the effects of various cations on frog-heart muscle and somewhat serendipitously discovered that Ca2+ ions, everpresent in the tap water distributed in central London, in millimolar concentrations were necessary for muscle contraction and tissue survival.
Today it is widely recognized that Ca2+ ions are central to a complex intracellular messenger system that is mediating a wide range of biological processes: muscle contraction, secretion, glycolysis and gluconeogenesis, ion transport, cell division and growth (for definitions of terms in boldface, see Appendix A in Section IX). The detailed organization of this messenger system is presently the subject of considerable scientific activity, and some details are already known. One of the links in the system is a class of highly homologous Ca2+-binding proteins, to be discussed later on in this chapter, that undergo Ca2+-dependent conformational changes and respond to transitory increases in intracellular Ca2+-ion concentrations. A prerequisite for the proper function of the calcium messenger system in higher organisms is that the cytosolic Ca2+ concentration in a "resting" cell be kept very low, on the order of 100 to 200 nM. Transitory increases in the Ca2+ concentration that may result from hormonal action on a membrane receptor must rapidly be reduced. Several transport proteins, driven either by ATP hydrolysis or by gradients of some other ion like Na+, are involved in this activity.
Ca2+ ions are also known to play various roles outside cells. In the plant kingdom Ca2+ ions often form links between individual cells and are required for maintaining the rigidity of whole plants; some seaweeds are typical examples. In the blood plasma of mammals, in which the Ca2+ concentration exceeds the intracellular by a factor of about 104, Ca2+ ions are instrumental in joining certain proteins in the blood-clotting system with membrane surfaces of circulating cells. Many extracellular enzymes also contain Ca2+ ions, sometimes at the active site but most often at other locations. It is generally believed that Ca2+ ions confer on proteins an increased thermal stability, and indeed proteins in heat-tolerant microorganisms often hold many such ions.
Vertebrates require much calcium in their food; in the USA the recommended daily allowance (RDA) for adult humans is 800 mg, and most other countries have comparable recommendations. During gestation in mammals, calcium must be transported across the placenta into the fetus, in particular during those phases of pregnancy when bone formation is most rapid. Interestingly, there appear to be some parallels between intestinal and placental transport that will be discussed further below. The role of calcium in biominerals is a vast subject that we can treat only superficially in this chapter.
To provide a background to the more biologically oriented sections that follow, we begin with a brief recapitulation of some basic facts about calcium. Then we continue with an outline of calcium distribution in biological tissues and organelles, and of the methods that can be used to obtain this information. After this follows a brief section on Ca2+ transport, and an account of the mechanism of intracellular Ca2+ release as it is presently understood. A discussion of some selected Ca2+-binding proteins of general interest, both intracellular and extracellular, then follows. Before we conclude the chapter, we will summarize some recent observations on Ca2+-binding proteins in prokaryotes.
III. Calcium In Living Cells: Methods For Determining Concentrations And Spatial Distributions
1. Measurements of "Free" Calcium Concentrations
1. Ca2+-selective microelectrodes
2. Bioluminescence
3. Complexing Agents with Ca2+-dependent Light Absorption or Fluorescence
4. Complexing Agents with Ca2+-dependent NMR Spectra
3. Summary
Much of our present knowledge about the biological role of Ca2+ rests on detailed measurements of the concentration, distribution, and chemical nature of Ca2+ and its complexes. Concentrations of uncomplexed, or "free," Ca2+ can be measured by Ca2+-selective microelectrodes, bioluminescence and complexing agents with Ca2+-dependent light absorption, fluorescence, or NMR spectra. An outcome of such studies is that the "free" Ca2+ concentration in resting eukaryotic cells generally is very low, on the order of 100 to 200 nM. Total Ca2+ concentrations, uncomplexed and complexed, can be measured by a variety of physical techniques. Some techniques, like atomic absorption, are sensitive but give poor spatial resolution. Others involve the bombardment of the sample with electrons or charged atoms, and can yield spatial resolutions of the order of a few nm; however, there is a trade-off between detectability and resolution
IV. The Transport and Regulation of Ca2+ Ions in Higher Organisms
2. Intracellular Ca2+ Transport
1. The Ca2+-ATPases
2. The Na+/Ca2+ Exchanger of the Plasma Membrane
3. Mitochondrial Ca2+ Transport: Influx
4. Mitochondrial Ca2+ Transport: Efflux
5. Ca2+ Efflux from Non-mitochondrial Stores
6. Other Voltage-gated or Receptor-activated Ca2+ Channels
4. Summary
The fluxes of Ca2+ ions and their regulation in higher organisms, as well as in microorganisms, depend on several transport proteins in addition to vesicular and gated processes. An important class of transport proteins are the Ca2+-ATPases, which are particularly abundant in muscle cells. These proteins translocate Ca2+ ions against large activity (or concentration) gradients through the expenditure of ATP. Transport of Ca2+ ions against activity gradients across membranes may also be accomplished by coupled transport of other ions, like Na+, with a gradient in the opposite direction.
As a result of some external stimulus—the action of a hormone, for example—the "free" Ca2+-ion concentrations in the cytoplasm of many cell types may transiently increase several orders of magnitude. This increase largely results from the release of Ca2+ from intracellular stores (ER, SR) in response to the initial formation of a new type of messenger, 1,4,5-IP3. The activity of Ca2+-transport proteins eventually restores the Ca2+ concentration levels to resting levels. This sequence of events forms the basis for Ca2+'s role in the regulation of a wide variety of cellular activities (see Section V).
V. Molecular Aspects of Ca2+-regulated Intracellular Processes
8. Summary
Many different biological processes in eukaryotic cells are regulated by intracellular Ca2+ concentration levels. Examples of such processes are muscle contraction, transport processes, cell division and growth, enzyme activities, and metabolic processes. A link in this regulatory chain is a number of intracellular Ca2+ receptors with Ca2+ affinities such that their binding sites are largely unoccupied at resting Ca2+ concentration levels, but are occupied at Ca2+ levels reached as a result of some external stimulus. This class of Ca2+ receptors is often called the "calmodulin superfamily" and includes the well-known members troponin C (regulating muscle contraction in striated muscle) and calmodulin (playing an important role in the regulation of many cellular processes). Amino-acid sequence determinations as well as x-ray and 2D 1H NMR studies have revealed a strong homology between the regulatory Ca2+-binding proteins. The Ca2+-binding sites are located in a loop flanked by two helices, and the Ca2+ ions are ligated with approximately octahedral or pentagonal bipyramidal symmetry. The ligands are six or seven oxygen atoms that are furnished by side-chain carboxylate or hydroxyl groups, backbone carbonyls, and water molecules. Pairs of these Ca2+ sites, rather than individual sites, appear to be the functional unit, and a common consequence of their arrangement is cooperative Ca2+ binding. Ca2+ binding to the intracellular receptor proteins is accompanied by structural changes that expose hydrophobic patches on their surfaces, thereby enabling them to bind to their target proteins.
VI. Extracellular Ca2+-binding Proteins
2. Summary
In higher organisms, the Ca2+ concentration in extracellular fluids generally is considerably higher than the intracellular concentrations. In mammalian body fluids, the Ca2+ concentration is typically on the order of a few mM. The extracellular concentration levels are highly regulated and undergo only minor variations. A consequence of these high levels of Ca2+ in extracellular fluids is that the binding constant need be only 103 to 104 M-1 in order for a protein site to be highly occupied by Ca2+. Several extracellular enzymes and enzyme activators have one or more Ca2+ ions as integral parts of their structures. Some Ca2+ ions are bound at, or near, the active cleft and may take part in the enzymatic reactions (e.g., phospholipase A2, $\alpha$-amylase). In other molecules, for example, serine proteases like trypsin and chymotrypsin, the Ca2+ ion is not essential for enzymatic activity, and may play more of a structural role. Ca2+ ions are involved in the cascade of enzymatic events that results in blood clotting in mammals. Several of the proteins in this system contain two new amino acids, $\gamma$-carboxyglutamic acid (Gla) and $\beta$-hydroxyaspartic acid (Hya), which are strongly suspected to be involved as ligands in Ca2+ binding. In the presence of Ca2+ ions, prothrombin and other Gla-containing proteins will bind to cell membranes containing acidic phospholipids, in particular, the platelet membrane. It appears likely that Ca2+ ions form a link between the protein and the membrane surface.
VII. Calcium in Mineralized Tissues
Summary
Calcium is, along with iron, silicon, and the alkaline earth metals, an important constituent of mineralized biological tissues. Some Ca2+-based biominerals, like bone or mother-of-pearl, can be regarded as complex composites with microscopic crystallites embedded in a protein matrix. The formation of calcified biominerals is a highly regulated process, and human bone, for instance, is constantly being dissolved and rebuilt. When the rates of these two counteracting processes are not in balance, the result may be decalcification, or osteoporosis, which seriously reduces the strength of the bone.
VIII. Ca2+-binding Proteins in Microorganisms: The Search for a Prokaryotic Calmodulin
Summary
The role of Ca2+ ions in the regulation of biological activities of procaryotic organisms is still largely unsettled. Over the last decade, however, evidence has gradually accumulated that calcium ions are involved in diverse bacterial activities, such as chemotaxis and substrate transport, sporulation, initiation of DNA replication, phospholipid synthesis, and protein phosphorylation.168 An important landmark is the recent demonstration that the intracellular Ca2+ concentration in E. coli is tightly regulated to about 100 nM, a level similar to that typical of resting eukaryotic cells.169 Furthermore, increasing numbers of calcium-binding proteins, some of which also have putative EF-hand Ca2+ sites characteristic of the calmodulin superfamily of intracellular regulatory proteins, have been isolated in bacteria.168
IX. Appendixes
1. Definition of Biochemical Terms
Antiport A transport protein that carries two ions or molecules in opposite directions across a membrane.
Basal lateral membrane The membrane in intestinal epithelial cells that is located on the base of the cells, opposite the microvilli that face the intestinal lumen.
Cytosol The unstructured portion of the interior of a cell—the cell nucleus excluded—in which the organelles are bathed.
Electrogenic A biological process driven by electric field gradients.
Endocytosis The process by which eukaryotic cells take up solutes and/or particles by enclosure in a portion of the plasma membrane to (temporarily) form cytoplasmic vesicles.
Endoplasmic reticulum (ER) Sheets of folded membranes, within the cytoplasm of eukaryotic cells, that are the sites for protein synthesis and transport.
Epithelial cells Cells that form the surface layer of most, if not all, body cavities (blood vessels, intestine, urinary bladder, mouth, etc.).
Erythrocytes Red-blood corpuscles.
Eukaryotic cells Cells with a well-definied nucleus.
Exocytosis The process by which eukaryotic cells release packets of molecules (e.g., neurotransmitters) to the environment by fusing vesicles formed in the cytoplasm with the plasma membrane.
Gluconeogenesis Metabolic synthesis of glucose.
Glycolysis Metabolic degradation of glucose.
Hydropathy A measure of the relative hydrophobic or hydrophilic character of an amino acid or amino-acid side chain.
Lamina propria mucosae The layer of connective tissue underlying the epithelium of a mucous membrane.
Mitrochondrion A double-membrane organelle in eukaryotic cells that is the center for aerobic oxidation processes leading to the formation of energy-rich ATP.
Organelle A structurally distinct region of the cell that contains specific enzymes or other proteins that perform particular biological functions.
Osteoporosis Brittle-bond disease.
Phorbol esters Polycyclic organic molecules that act as analogues to diacylglycerol and therefore are strong activators of protein kinase C.
Prokaryotic cells Cells lacking a well-defined nucleus.
Sarcoplasmic reticulum The ER of muscle cells.
Trophoblasts The cells between the maternal and fetal circulation systems.
Tryptic digest Fragmentation of proteins as a result of treatment with the proteolytic enzyme trypsin.
Uniporter A transport protein that carries a particular ion or molecule in one direction across a membrane.
1. One-Letter Code for Amino-Acid Residues
A—alanine, C—cysteine, D—aspartate, E—glutamate, F—phenylalanine, G—glycine, H—histidine, I—isoleucine, K—Iysine, L—Ieucine, M—methionine, N—asparagine, P—proline, Q—glutamine, R—arginine, S—serine, T—threonine, V—valine, W—tryptophan, Y—tyrosine.
1. The Activity of a Transport Protein
This is usually described in terms of the classical Michaelis-Menten scheme:
$V (= transport\; rate) = V_{max} \cdotp \frac{[S]}{[S] + K_{m}},$
where [S] is the concentration of the solute to be transported and Km = (k-1+k2)/k1 is the Michaelis constant (dimension "concentration") for the reaction
$E+S \xrightleftharpoons[k_{-1}]{k_{1}} ES \xrightarrow{k_{2}} R \ldotp$
Approximated as the reciprocal ratio between on- and off-rate constants relevant to the solute-protein complex, 1/Km = k1/k-1 may be taken as a lower limit of the affinity of the protein for the solute.
Contributors and Attributions
• Sture Forsén (University of Lund, Chemical Centre, Physical Chemistry 2)
• Johan Kördel (University of Lund, Chemical Centre, Physical Chemistry 2)
03: Calcium in Biological Systems
Since Ca2+ ions evidently play an important role in regulating a variety of cellular responses in animals and higher organisms, one may ask whether this use of Ca2+ is a recent discovery of Nature, or if it was invented early in evolution. It now appears well-established that the key intracellular "Ca2+- receptor" protein calmodulin (CaM; see Section V.A) is present in all eukaryotic cells. Even in a unicellular eukaryote like common yeast (Saccharomyces cerevisiae), Ca2+ has an important regulatory role, and recently yeast CaM, as well as the single-gene encoding for it, was isolated.160
The amino-acid sequence of the yeast CaM (147 a.a.; Mr = 16.1 kDa) is 60 percent identical with the sequences of all other CaMs known. In fact, if generally accepted conservative amino-acid replacements are allowed, the homology increases to 80 percent or more, the most highly conserved portions being the four putative Ca2+-binding sites. Sites I and III match the EF-hand test sequence (see Figure 3.24) very well; in site a His occurs after the "z"- ligand instead of the archetypal Gly; and in site IV there is no amino acid between the residues that usually make up ligands "x" and "y." The effect of these alterations on the Ca2+ affinity of yeast CaM is not yet known.
That CaM is essential for the growth of yeast cells was shown by deletion or disruption of the gene. This constitutes, in fact, the first demonstration in any organism that CaM is an essential protein. (Deletions of genes in mammals are ethically questionable research procedures!)
In the biochemically less sophiscated (than eukaryotes) prokaryotic cells, a regulatory role of Ca2+ is not well-established. What is known is that calcium is massively accumulated during sporulation in many bacteria, for example, in strains of Bacillus, Streptomyces, and Myxococcus. In Myxococcus xanthus a development-specific protein called protein S assembles at the surface of myxospores in the presence of Ca2+. The DNA sequence of the gene that encodes this protein has been deciphered.161 The primary sequence of protein S (175 a.a., Mr = 19.2 kDa) turns out to closely resemble mammalian CaM. It has four internally homologous regions with putative Ca2+ sites. At least two of these are partly similar to the typical EF-hand, but uncharacteristically there are many more prolines in the M. xanthus protein than in bovine CaM (12 versus 2); so it is questionable if the bacterial protein really has the repeated helix-loop-helix structure found in mammalian CaM.162
One candidate for a prokaryotic CaM was reported by Leadlay et al.163 in Streptomyces erythreaus, the bacterium that produces the well-known antibiotic "erythromycin." The amino-acid sequence of a low-molecular-weight Ca2+- binding protein, as determined from the gene encoding it, revealed a high homology with mammalian CaM. The protein is made up of 177 amino acids (Mr = 20.1 kDa), and has four regions that are predicted to have the helix-loop-helix secondary structure typical of EF-hand proteins. The aligned sequences of the 12 residues in each of the four potential calcium-binding loops in the S. erythreaus protein are compared with those of human calmodulin in Table 3.6.
Table 3.6 - Aligned EF-hand sequences for the prokaryotic and human calmodulins
Ligands 1 3 5 7 9 12
S. erythraeus protein I D F D G N G A L E R A D
S. erythraeus protein II
G V G S D G S L T E E Q
S. erythraeus protein III D K N A D G Q I N A D E
S. erythraeus protein IV D T N G N G E L S L D E
Human calmodulin I D K D G D G T I T T K E
Human calmodulin II D A D G N G T I D F P E
Human calmodulin III D K D G N G Y I S A A E
Human calmodulin IV D I D G D G Q V N Y E E
The pattern of residues in the S. erythraeus protein is typical of an EF-hand at least in sites I, III, and IV. Site II is unusual in having Gly at both positions 1 and 3. 113Cd NMR studies show that the bacterial protein binds three metal ions strongly (K ≳ 105 M-1) with chemical shifts close to those expected for EF-hands, and 1H NMR studies show that it undergoes a Ca2+- dependent conformational change.164
Although the S. erythraeus protein has a homology with eukaryotic CaM, it has been pointed out that the protein has an even higher homology with a group of eukaryotic sarcoplasmic Ca2+-binding proteins165 (see Section V.D). The search for a prokaryotic CaM analogue continues, and the prospect of success has been improved after recent reports of a 21-amino-acid-Iong polypeptide from an E. coli heat-shock protein166 that shows the typical structural features of CaM-binding domains in other eukaryotic proteins.167 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.01%3A_%28Ca2%29-binding_Proteins_in_Microorganisms-_The_Search_for_a_Prokaryotic_Calmodulin.txt |
Basic Facts
Calcium was first recognized as an element in 1808 by Humphry Davy, and the name was given after the Latin for lime: calx. Several isotopes of calcium are known. The stable isotopes are, in order of decreasing natural abundance, 40Ca (96.94%), 44Ca (2.1%), 42Ca (0.64%), and 43Ca (0.145%). 43Ca is the only isotope with a nuclear spin (I = $\frac{7}{2}$) different from zero, which makes it amenable to NMR studies. 45Ca is a radioactive isotope of some importance ($\beta^{-}$ decay; 8.8 min half life).3 It has been used in studies of calcium localization and transport in biological systems.
Calcium constitutes about 3 percent by weight of the Earth's crust, mostly in the form of sedimentary rocks of biological origin dating back some three billion years. In sea water the total concentration of calcium ranges from 5 to 50 times higher than in fresh water, which, in turn, has a calcium concentration ten times that of rain water (see Table 3.1). This explains the pleasant feeling when ordinary soaps are used in rain water. The calcium concentration in ordinary tap water varies with location; calcium is usually added to water in distributing networks in order to prevent corrosion of iron pipes. Tap water with a calcium concentration above 1.5 mM is usually classified as "hard." Interestingly, the taste of beer seems related to the calcium concentration, and it is claimed that "good" beer should have a concentration higher than that of "hard" tap water.
In the body fluids of higher organisms the total calcium concentration is usually on the order of a few millimolar (see Table 3.1). In adult human serum, the concentration is observed to be, within narrow limits, 2.45 mM.
Table 3.1 - Ca2+ concentrations in fluids and tissues.6-9
Specimen Units are mM if not otherwise stated
Sea water 10
Fresh water 0.02 - 2
Rain water 0.002 - 0.02
"Hard" tap water 1.5
"Good" beer 4
Adult human serum 2.45 $\pm$ 0.05
Serum of other vertebrates 1.5 - 5
Nematote body fluids 6
Molluscan serum - marine 9 - 15
Molluscan serum - fresh water 1.5 - 7.8
Molluscan serum - land 3.3 - 12.3
Milk 70
Bone 0.8 - 1.0
Mitochondria from rat liver 0.8 $\pm$ 0.1 mmol/kg
Endoplasmatic reticulum 8 - 10 mmol/kg
Cytoplasm of a resting mammalian cell 0.0001
Cytoplasm of E. coli 0.0001
Essentials of Ca2+ Chemistry
Since the Ca2+ ion accomplishes its biological tasks in an environment with 1 to 3 mM Mg2+, it is of particular interest to compare the properties of these two ions in order to understand how a discrimination is made in biological systems. In addition, the coordination chemistry of Ca2+ is closely related to that of Mg2+ (as well as Cd2+), though there are several obvious differences. First of all, the ionic radius of a Ca2+ ion with a given coordination number (CN) is always higher than that of an Mg2+ or Cd2+ ion with the same CN. At CN = 6, the ionic radii of Ca2+, Cd2+, and Mg2+ are 1.00, 0.95, and 0.72 Å, respectively, whereas at CN = 8 they are 1.12, 1.10, and 0.89 Å, respectively.4
Ligand preferences of Ca2+ depend on the fact that it is a hard metal ion. Thus Ca2+ strongly prefers oxygen ligands over nitrogen or sulfur ligands; Ca2+••••N bonds are about 0.25—0.3 Å longer than Ca2+••••O bonds.5,6,10 Large differences in coordination number and geometry have been observed for Ca2+ complexes. In a study of 170 x-ray structures of Ca2+ complexes involving carboxylate groups,11 binding was found to be either (i) unidentate, in which the Ca2+ ion interacts with only one of the two carboxylate oxygens, (ii) bidentate, in which the Ca2+ ion is chelated by both carboxylate oxygens, or (iii) mixed ("$\alpha$-mode") in which the Ca2+ ion is chelated by one of the carboxylate oxygens and another ligand attached to the $\alpha$-carbon (see Figure 3.1). The Ca2+- oxygen distances span a range from 2.30 to 2.50 Å, with the average distance being 2.38 Å in the unidentate and 2.53 Å in the bidentate mode, respectively.11 Observed coordination numbers follow the order 8 > 7 > 6 > 9. By contrast, Mg2+ nearly always occupies the center of an octahedron of oxygen atoms (CN = 6) at a fixed Mg2+-oxygen distance of 2.05 ± 0.05 Å.
In Table 3.2, stability constants for the binding of Ca2+ and Mg2+ to various ligands are collected. We may note that selectivity of Ca 2+ over Mg2+ is not very great for simple carboxylate ligands, but that it tends to increase for large multidentate ligands, such as EDTA and in particular EGTA. The Ca2+ sites in many intracellular proteins with "EF-hand" binding sites (see Section V. C) bind Ca2+ about 104 times more strongly than Mg2+.
Table 3.2 - Ca2+ and Mg2+ (where available) stability constants (log K) for different organic and biochemical ligands. Most values are at ionic strength 0.1 and 25 °C.5,6,12-15
a) EGTA: ethylenebis(oxyethylenenitrilo)tetraacetate
b) EDTA: ethylenedinitrilotetraacetate
Ligand Ca2+ Mg2+
Acetate 0.5 0.5
Lactate 1.1 0.9
Malonate 1.5 2.1
Aspartate 1.6 2.4
Citrate 3.5 3.4
Nitrilotriacetate 6.4 5.5
EGTAa 10.9 5.3
EDTAb 10.6 8.8
Glycine (Gly) 1.4 3.4
$\gamma$-Carboxyglutamic acid (Gla) 1.3
Gly-Gly dipeptide 1.2
Gla-Gla dipeptide 3.2
Macrobicyclic amino cryptate [2.2.2] 4.5
Fluo-3 6.2 2.0
Fura-2 6.9 2.0
BAPTA 7.0 1.8
Quin-2 7.1 2.7
Phospholipase A2 3.6
Thrombin fragment 1 3.7 3.0
Trypsinogen 3.8
Chymotrypsinogen 3.9
Chymotrypsin 4.1
Calmodulin, N-terminal 4.5 3.3
Trypsin 4.6
Calmodulin, C-terminal 5.3
Protein kinase C $\sim$7
$\alpha$-Lactabumin $\sim$7
Rabbit skeletal muscle
Troponin C, Ca2+/Mg2+ sites 7.3 3.6
Carp parvalbumin $\sim$8.5 4.2
Bovine calbindin D9K 8.8 $\sim$4.3
Another difference in ligand-binding properties of Mg2+ and Ca2+ can be seen by comparing the rates of substitution of water molecules in the inner hydration sphere by simple ligands, according to
$M(H_{2}O)_{n}^{2+} + L \xrightarrow{k} ML(H_{2}O)_{n-1}^{2+} + H_{2}O$
This rate (log k, with k in s-1) has been determined to be 8.4 for Ca2+ and 5.2 for Mg2+.16
The formation of biominerals is a complex phenomenon. In order to obtain a feeling for the conditions under which inorganic solid phases in biological systems are stable, it is of some interest to look at solubility products. Solubility products, Kspo, have a meaning only if the composition of the solid phase is specified. For a solid compound with the general composition (A)k(B)l(C)m the solubility product is defined as
$K_{sp}^{o} = [A]^{k} [B]^{l} [C]^{m} \tag{3.1}$
where [A], [B], etc., denote activities of the respective species, usually ionic, in equilibrium with the solid. Activities are concentrations multiplied by an activity coefficient, $\gamma$, nearly always less than unity. Activity coefficients for ions in real solutions can be estimated from Debye-Hückel theory17 if the ionic strength of the solution is known. In human blood plasma, the ionic strength, I, is about 0.16, and the activity coefficient for Ca2+ at 37 °C is 0.34. In many discussions it may be sufficient to equate concentrations with activities.
The solid phase involved is essentially assumed to be an infinitely large, defect- and impurity-free crystal with a well-defined structure. Microscopic crystals have higher solubilities than large crystals, a well-known phenomenon that leads to "aging" of precipitates, in which larger crystals grow at the expense of smaller ones.
Many anionic species appearing in the solubility products may also be involved in protonation equilibria in solution, such as those of phosphoric acid: H2PO4- $\rightleftharpoons$ H+ + HPO42-; HPO42- $\rightleftharpoons$ PO43- + H+; etc. When the prospects for the formation of a solid phase under certain solution conditions are investigated, the activity, or concentration, of the particular anionic species specified in the solubility product must be known, not only "total phosphate" or "total calcium," etc. The data in Table 3.3 show that, at pH > 5, the most stable (i.e., insoluble) solid calcium phosphate is hydroxyapatite.
Table 3.3 - Solubility products, at pH 5 and 25 °C, for solid calcium phosphates
Solid Phase -log Kspo -log Kspo of corresponding Mg2+ compound where applicable
CaSO4 • 2H2O (sulfate, "gypsum") 5.1 < 1.0
Ca(OH)2 (hydroxide) 5.3 10.7
CaHPO4 • 2H2O (hydrogen phosphate) 6.6
CaCO3 (carbonate, "calcite," "aragonite") 8.5 7.5
CaC2O4 • H2O (oxalate, "whewellite") 10.5 5.0
$\beta$-Ca3(PO4)2 ($\beta$-phosphate) 29
Ca5(PO4)3OH (hydroxyapatite) 58 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.02%3A_Basic_Facts_About_Calcium-_Its_Compounds_and_Reactions.txt |
The formation of calcified tissue—shells, bone, and teeth—is a very complex process that is under strict regulatory control. Despite the obvious importance of this field, relatively little research has been directed toward elucidation of the underlying mechanisms, perhaps because the field spans a broad range of subjects, from inorganic solution and solid-state chemistry to cellular physiology.155
Historically, it was long held that formation of biological minerals such as bone was simply the nucleation and growth of calcium hydroxyapatite within an extracellular matrix of collagen. Many proteins other than collagen have now been discovered in appreciable quantities in bone and other biological minerals. It is also apparent that the pattern of calcification differs in shells, bone, teeth, and other mineralized tissues; so it is not likely that there is only one underlying mechanism. Considering the immensity of the subject, we will here only make a few brief comments, mainly about bone and teeth.
As was briefly mentioned earlier in this chapter, the inorganic matter of bone and teeth in many ways resembles apatite minerals (Ca5(OH)(PO4)3). Table 3.5 summarizes inorganic solid components of other biominerals.
Table 3.5 - A summary of the main inorganic solid component8 of the most-common biominerals in living systems.148
a) Most real biominerals are actually nonstoichiometric, and contain a number of additional cations (e. g., Mg2+) or anions (e.g., F-). In addition, the inorganic phase may be interpenetrated by a biopolymer.
Anion Formula Crystal Form Occurrence Main Function
Carbonate CaCO3
Calcite
Aragonite
Valerite
Sea corals, molluscs, and many animals and plants Exoskeleton; Ca-store; eye lens
Oxalate
Ca(COO)2•H2O
Ca(COO)2•2H2O
Whewellite
Weddelite
Insect eggs; vertebrate stones Deterrent; cytoskeleton; Ca store
Phosphate
(Ca)10(PO4)6(OH)2
(unit cell comp.)
Hydroxyapatite Bones; teeth; shells; intracellular in some bacteria Skeletal; Ca storage; pressure-transducer (piezo-electric)
Sulfate CaSO4•H2O Gypsum Jelly fish; plants Gravity device; S and Ca store
A detailed analysis156 shows that, apart from Ca2+ and PO43-, many other cations and anions occur in bone, e.g., Mg2+, Na+, K+, Sr2+, CO32-, F-, Cl-, and citrate. X-ray diffraction patterns and electron-microscope pictures of bone show that the inorganic phase is made up of many very small and imperfect crystals. By contrast, dental enamel is made up of much larger and uniform thin crystals. Although the solubility product of calcium hydroxyapatite (see Section II) is such that the equilibrium Ca2+ concentration should be in the low micromolar range, bone mineral appears to be in equilibrium with much higher Ca2+ concentrations (0.8-1.0 mM).157 This discussion brings us to the question of how the inorganic crystallites are formed. Obviously both Ca2+ and PO43- ions must be concentrated in cells or organelles bordering on the regions where mineralization is to take place. Fresh layers of bone matrix are formed by a continuously replenished layer of cells called osteoblasts (Figure 3.32A), which, in addition to apatite crystallites, also secrete collagen, and large specific proteins called osteonectin, osteocalcin (a Gla protein), proteoglycans, and phosphoproteins. In tissues undergoing rapid mineral deposition, the crystallites appear to be formed in vesicles that may have peeled off from the adjacent cell layers. These vesicles seem able to concentrate calcium and phosphate in a manner not well understood.
Bone, unlike diamond, is not forever. It can be remodeled and dissolved. A serious medical problem, which affects some women after menopause, is osteoporosis, i.e., the decalcification of bone. This loss of bone mass, which occurs with increasing age, makes bones more susceptible to breaking under stress. About 50 percent of American women, and 25 percent of American men, over 45 years of age are affected by osteoporosis.158 Whereas osteoblast cells handle bone formation, another type of cells, osteoclasts, can erode it (Figure 3.32B). These macrophage-like cells can form deep tunnels in a bone matrix, and the cavities left behind are rapidly invaded by other cells forming blood vessels and new layers of osteoblasts. The modus operandi of osteoclast cells is not well understood at present. They may secrete calcium-chelating organic anions, such as citrate, to assist in the solubilization of the bones, as well as extracellular proteases that degrade the organic part of the matrix. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.03%3A_Calcium_in_Mineralized_Tissues.txt |
Calmodulin is a small acidic protein (Mr ≈ 16,700), the amino-acid sequence of which has been remarkably preserved during evolution. Early on, an analysis of its amino-acid sequence indicated that it should have four Ca2+-binding sites, a deduction that proved to be correct. The three-dimensional x-ray structure of bovine brain calmodulin85 has been solved to a resolution of 2.2 Å. A space-filling model is shown in Figure 3.17. (See color plate section, page C-9.) The molecule has a dumbbell-like shape, with two globular domains connected by an eight-turn $\alpha$-helix—an unusual structural feature. In the crystal structure, there are no direct contacts between the two globular domains, each of which contains two Ca2+-binding sites. The Ca2+sites are all constructed in the same way: two $\alpha$-helices separated by a calcium-binding loop, 12 amino acids long, and wrapped around the Ca2+ ion. This structural arrangement is nearly identical with that first observed in the x-ray structure of carp parvalbumin, and is colloquially termed "the EF-hand."86 This structural unit is also observed in all available x-ray structures of proteins of the calmodulin superfamily (see Sections V.B and V.C). The Ca2+ ligands are all oxygen atoms, located approximately at the vertices of a pentagonal bipyramid.
The binding of Ca2+ and other cations to CaM has been extensively investigated.87 The first two Ca2+ ions are bound in a cooperative manner, with an average binding constant of about 2 x 105 M-1 in 150 mM KCl and 1 mM Mg2+. The third and fourth Ca2+ ions are bound with binding constants of about 3 x 104 M-1 under the same conditions. Spectroscopic evidence has shown that the first two Ca2+-ions are bound in the C-terminal domain. Mg2+ has been shown to bind primarily to the N-terminal domain (see Table 3.2).88
The rates of dissociation of Ca 2+ from the (Ca2+)4 CaM complex have been studied by both stopped-flow and NMR techniques.89,90 Fast and slow processes are observed, both corresponding to the release of two Ca2+ ions. At an ionic strength I = 0.1 and 25 °C, the rates for the two processes differ by a factor of 30 (see Table 3.4).
Table 3.4 - Rates of Ca2+-dissociation and -association of some enzymes and proenzymes.
koff [s-1] kon [M-1 s-1]
Macrobicyclic amino cryptate [2.2.2] 0.3 104
Phospholipase A2 1.1 x 103 4 x 106
sTroponin C: Ca2+ sites 300
sTroponin C: Ca2+-Mg2+ sites 5
Trypsin 3 1.1 x 105
Trypsinogen $\leq$ 10 6 x 104
Chymotrypsin 70 ~ 106
Chymotrypsinogen 350 2.8 x 105
Calmodulin: N-terminal 300-500 107
Calmodulin: C-terminal 10-20
A body of biophysical measurements, mostly made before the advent of x-ray structures, indicated that CaM is constructed from two largely independent domains.87 This conclusion emanated from studies of the two tryptic fragments, TR1C and TR2C. The major site of cleavage is between Lys-77 and Asp-78 of the central helix, and results in N-terminal and C-terminal fragments of nearly equal size. To a good approximation, the biophysical properties of the intact CaM molecule—NMR, UV and CD spectra, kinetic properties, thermochemical data, etc.—are the sum of the same properties of the fragments TR1C and TR2C. This means that we may assign the slow dissociation process, kofff, to the C-terminal domain, and the fast, k;ff' to the N-terminal domain of CaM. Combining binding constants and off-rates, we may calculate that the rates of Ca2+ binding to CaM are on the order of 107 M-1 s-1 at high ionic strength, and 108 M-1 s-1 or higher at low ionic strength. Recently the x-ray structure of the C-terminal fragment TR2C was solved, and indeed showed a structure nearly identical with C-terminal domain of intact CaM.91
The structural changes occurring in CaM as Ca2+ ions are bound are associated with pronounced changes in 1H NMR, UV, fluorescence, and CD spectra.87 The observed changes in CD and fluorescence spectra in the presence of Mg2+ are only about 20 to 25 percent of those induced by Ca2+. A comparison of the CD spectra of CaM and its tryptic fragments indicates that the structural changes induced by Ca2+ are substantially greater in the C-terminal than in the N-terminal half.92 By and large, few structural details of the conformation changes have as yet been obtained. However, one aspect of the Ca2+-induced conformation change is that hydrophobic sites, probably one on each domain of the molecule, become exposed. In the presence of excess Ca2+, CaM will bind to other hydrophobic molecules, e.g., phenyl-Sepharose, a variety of drugs, many small peptides, and—last but not least—its target proteins. This brings us to the question of how CaM recognizes and interacts with the latter. We may suspect that the hydrophobic sites on each domain are somehow involved, but the role played by the central helix is still not clear. To explain small-angle x-ray scattering data, the interconnecting helix needs to be kinked, bringing the intact globular domains closer.93
A putative CaM-binding segment (27 amino acids long) of myosin light-chain kinase (MLCK), an enzyme activated by CaM, has been identified.94 The interaction between the segment peptide ("M13") and CaM has been studied95 by CD spectroscopy and 1H NMR. From these studies it appears that a unique 1:1 complex is formed, and that secondary and tertiary structural changes occur not only in the peptide M13 but also in both halves of the CaM molecule. Further NMR studies 96,97 of the interaction between CaM and naturally occurring peptides (mellitin and mastoparan) that share some structural features of M13—clusters of basic residues, hydrophobic residues adjacent to the basic residues, and a predicted high $\alpha$-helical content—show very much the same results. Based on these results, a model, shown in Figure 3.18, for the interaction between CaM and M13 has been proposed. In this model the central helix is kinked at position81, allowing the two domains to wrap around the assumed $\alpha$-helical M13. Preliminary structure calculations of calcium-loaded CaM, based on NMR data, indicate that the central helix in solution indeed is kinked and very flexible,99 and comparisons100 of chemical shifts in calmodulin with and without M13 complexed supports the model in Figure 3.18. Recent structural studies using NMR spectroscopy and x-ray diffraction have essentially confirmed the general features of this model, although the orientation of the peptide is found to be reversed.173
In conclusion, two important features of the protein should be recognized.
1. The binding of Ca2+ to CaM (and to its complex with the target protein) is quite likely cooperative, meaning that the switch from inactive to active conformation may occur over a much more narrow Ca2+-concentration interval than otherwise.
2. The effective Ca2+ affinity will be different in the presence of the target proteins. To illustrate this second point, consider the standard free energies in the minimum scheme depicted in Figure 3.19. If the affinity of the Ca2+-calmodulin complex (CaM(Ca)4) for the target protein (P) is greater than that of Ca2+-free calmodulin (CaM)—i.e., |$\Delta$GIII| > |$\Delta$GII|—it follows that the Ca2+ affinity of the complex between P and CaM (P•CaM) must be higher than in CaM itself. This effect is also found experimentally in model systems.101 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.04%3A_Calmodulin.txt |
The Ca2+ concentration in extracellular fluids is usually orders of magnitude higher than intracellular concentrations. In mammalian body fluids, the "free" Ca2+ concentration is estimated to be 1.25 mM (total Ca2+ is ~2.45 mM) with only minor variations.140 We would thus expect that Ca2+ ions in extracellular fluids playa very different role from that inside cells. To ensure Ca2+ binding the macromolecular binding sites need have only a modest Ca2+ affinity (KBCa2+ ≈ 103 to 104 M-1), and since extracellular Ca2+ does not seem to have a signaling function, the rates of Ca2+ association or dissociation in protein-binding sites need not be very high.
One particularly important aspect of Ca2+ in mammals is its role in the blood coagulation system. Here we will meet a new type of amino acid, $\gamma$-carboxyglutamic acid ("Gla" )-see Figure 3.29, that seems to have been designed by Nature as a Ca2+ ligand with rather special functions. Gla-containing proteins are also encountered in some mineralized tissues. The formation of bone, teeth, and other calcified hard structures is an intriguingly complicated phenomenon that will be dealt with in Section VII. We start, however, with a brief discussion of the role of Ca2+ in some extracellular enzymes.
Ca2+-binding in Some Extracellular Enzymes
Several extracellular enzymes have one or more Ca2+ ions as integral parts of their structure. In a very few of them the Ca2+ ion is bound at or near the active cleft, and appears necessary for maintaining the catalytic activity (phospholipase A2 , $\alpha$-amylase, nucleases), whereas other enzymes show catalytic activity even in the absence of Ca2+ (trypsin and other serine proteases). In the latter proteins, the Ca2+ ion is usually ascribed a "structural" role, although its function may be rather more related to "dynamics" and so be more subtle and complex.
Trypsin has one Ca2+-binding site with four ligands (two side-chain and two backbone oxygens) donated by the protein (Glu-70, Asn-72 , Val-75, and Glu-80) and two ligating water molecules, making the site roughly octahedral.141 The binding constant of Ca2+ to trypsin and its inactive precursor "proenzyme," trypsinogen, has been measured (see Table 3.2). The binding constant is slightly smaller for the precursor, as is also true for chymotrypsin and chymotrypsinogen.142 The Ca2+ affinities of the serine proteases and their proenzymes are such that their Ca2+ sites will be largely occupied in extracellular fluids, but would be unoccupied inside a cell. It has been suggested that this phenomenon constitutes a safeguard against unwanted conversion of the proenzymes into the active enzymes as long as they still are inside the cells where they are synthesized.
The rates of Ca2+ dissociation of the above enzymes and proenzymes have been measured by 43Ca NMR and stopped-flow techniques,142 and are collected in Table 3.4. We note that the values of kon and koff are generally much smaller than in the intracellular regulatory EF-hand proteins discussed in Section VI. Whereas the latter have dynamic and equilibrium properties similar to those of flexible low-molecular- weight chelators such as EDTA and EGTA, the serine proteases are more similar to the more-rigid cryptates, such as the macrobicyclic amino cryptate [2.2.2] (see Tables 3.2 and 3.4).
As mentioned above, there are a few enzymes in which a Ca2+ ion is present in the active cleft and essential for activity. Pancreatic phospholipase A2 (Mr ≈ 14 kDa) is an enzyme of this type. The x-ray structure is known to high resolution, and a single Ca2+ ion is found to be surrounded by six ligands, four presented by the protein (Tyr-28, Glu-30, Glu-32, and Asp-49) and two water molecules.143 A mechanism for the action of phospholipase A2 has been proposed144 and is shown in Figure 3.30.
This mechanism is based on three high-resolution x-ray crystal structures of phospholipase A2 with and without transition-state analogues bound. The binding constant for Ca2+ together with the rate of dissociation found from variable-temperature 43Ca NMR studies145 can be used to calculate kon ≈ 4 x 106 M-1s-1, again lower than in EF-hand proteins. Recent 1H NMR studies indicate that the global structure of the lipase is very much the same in the Ca2+-free and the Ca2+-bound forms. Structural changes upon Ca2+ binding appear primarily located in the region of the binding site.112,146
The mammary glands produce, among other substances, a Ca2+-binding enzyme activator, $\alpha$-lactalbumin, that has about 40 percent sequence identity with lysozyme. This protein, which is involved in the conversion of glucose into lactose, is secreted in large quantities, and in human milk constitutes some 15 percent of total protein. The Ca2+-binding constant of bovine or human $\alpha$-lactalbumin is on the order of 107 M-1 under physiological conditions. In addition to Ca2+, the enzyme also binds Zn2+. It appears that Ca2+-ion binding affects enzymatic activity, and somehow controls the secretion process, but the biological role of metal-ion binding to a-lactalbumin needs to be studied further. The x-ray structure of a-lactalbumin from baboon milk (Mr ≈ 15 kDa) has been determined147 to a high resolution (~1.7 Å). The Ca2+-binding site has an interesting structure. The ion is surrounded by seven oxygen ligands, three from the carboxylate groups of aspartyl residues (82, 87, and 88), two carbonyl oxygens (79 and 84), and two water molecules. The spatial arrangement is that of a slightly distorted pentagonal bipyramid with the carbonyl oxygens at the apices, and the five ligands donated by the proteins are part of a tight "elbow"-like turn. The $\alpha$-lactalbumin site has a superficial structural similarity to an "EF-hand," although the enzyme presumably has no evolutionary relationship with the intracellular Ca2+-binding regulatory proteins.
Blood clotting proceeds in a complicated cascade of linked events involving many enzymes and proenzymes. About a decade ago it was shown that several of these proteins contained a previously unknown amino acid, $\gamma$-carboxyglutamic acid (Gla), and more recently yet another new amino acid, $\beta$-hydroxyaspartie acid (Hya), has been discovered (see Figure 3.29). The former is formed postribosomally by a vitamin-K-dependent process in the liver.148 Presently the most-studied Gla protein in the blood-clotting system is prothrombin (Mr ≈ 66 kDa). Ten Gla residues are clustered pairwise in the N-terminal region, essentially lining one edge of the molecule, forming a highly negatively charged region.149 A small (48 residues) proteolytic fragment (F1) that contains all ten Gla amino acids can be prepared. Prothrombin can bind about 10 Ca2+ ions, but F1 binds only 7. Binding studies to F1 show that the Ca2+ ions bind at three high-affinity cooperative sites and four noninteracting sites,150 and that this binding takes places in conjunction with a spectroscopically detectable conformational change (see Table 3.1).
In the presence of Ca2+ ions, prothrombin and other vitamin-K-dependent proteins in the blood-coagulation system will bind to cell membranes containing acidic phospholipids, in particular, the platelet membrane, which is rich in phosphatidylserine. A proposed model for the prothrombin-membrane interaction is shown in Figure 3.31.
It has long been known that calcium ions are involved in cell-to-cell and cell-to-extracellular matrix interactions, but the molecular details largely remain to be unraveled. In the late 1980s a large, adhesive, calcium-binding matrix glycoprotein (Mr ~ 420 kDa) named thrombospondin was characterized. This multifunctional adhesion molecule is composed of three polypeptide chains, each with 38 amino-acid-Iong repeats that are homologous with the calcium-binding helix-loop-helix sites of the calmodulin superfamily.152 Each thrombospondin molecule is reported to bind 12 calcium ions with an affinity of about 104 M-1, and the removal of calcium is accompanied by a conformational change.153,154 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.05%3A_Extracellular_Calcium_Ion-binding_Proteins.txt |
A "second" messenger is an entity that inside a cell mediates the action of some hormone at the plasma membrane, the hormone being considered the "first" messenger. The first such second messenger to be discovered—in fact, the very molecule that led to the formulation of the whole concept—was cyclic AMP.69 During the decade following the discovery of cAMP, it was gradually realized that intracellular release of Ca2+ ions also accompanied hormonal stimuli, and the Ca2+ ion slowly became regarded as a second messenger. This idea was first clearly enunciated by Rasmussen70 as early as 1970, and gained general acceptance when the ubiquitous intracellular Ca2+-binding protein calmodulin (see Section V.A) was discovered. In the mid-1970s this protein was shown to be a Ca2+-dependent regulator of a large number of Ca2+-dependent enzymes, transport proteins, etc., establishing a molecular basis for Ca2+ action in cells.
There were some puzzling facts, however. Although a transitory increase in intracellular Ca2+ concentration in response to the binding of a hormone or transmitter substance to a surface receptor could result from extracellular Ca2+ being released into the cytoplasm, there was compelling evidence for muscle cells that the main Ca2+ source was the sarcoplasmic reticulum (SR). This result led to the hypothesis of ''Ca2+-induced Ca2+ release," i.e., that upon stimulation of the cell, a small amount of Ca2+ entered into the cytoplasm and triggered the release of greater amounts of Ca2+ from the SR. For some cell types it could, however, be shown that transient increases in intracellular Ca2+ could occur even when extracellular Ca2+ was removed, although prolonged responses required the presence of extracellular Ca2+. Although some specialized cells have gated plasma-membrane Ca2+ channels, release of Ca2+ into the cytoplasm from intracellular stores appears to be of at least equal importance. Furthermore, there is now overwhelming evidence63,70-72 that intracellular Ca2+ is released in response to the formation of a new type of intracellular messenger: 1,4,5-IP3. Receptors for this messenger have recently been found in the membranes of intracellular organelles, and binding of 1,4,5-IP3 to these receptors results in the release of Ca2+ ions.73
1,4,5-IP3 is formed as a product in the hydrolysis of a special phospholipid present in the cell membrane: phosphatidyl-inositol-4,5-bisphosphate. This reaction, then, is the initial receptor-stimulated event. The newly formed 1,4,5-IP3 is assumed to diffuse into the cytoplasm, and eventually reach intracellular 1,4,5-IP3 receptors on the ER, thereby triggering the release of Ca2+. A simplified reaction scheme is shown in Figure 3.15.
A diacylglycerol (DG) is also formed in the hydrolysis step. DG can also act as an intracellular messenger, and stimulates the activity of a membrane-bound protein kinase, known as protein kinase C (PKC). As a result, PKC may phosphorylate certain key proteins and influence their activity. Protein kinase C is also activated by Ca2+ ions, a fact that illustrates Nature's knack in designing regulatory networks! 1,4,5-IP3 is either directly degraded in a series of enzymatic steps back to inositol, which is then used to resynthesize the phospholipid, or it may be further phosphorylated to inositol-1,3,4,5,-tetraphosphate (1,3,4,5-IP4), which may undergo dephosphorylation to form inositol-1,3,4-trisphosphate (1,3,4-IP3). The biological functions of the latter compounds are now being investigated.
The intracellular levels of Ca2+ are restored back to the normal low resting values (100 to 200 nM) via transport back into the SR, and/or into mitochondria, or out through the plasma membrane by the pumping mechanisms discussed in Section IV.B. As was briefly mentioned above, depriving a cell of extracellular Ca2+ will eventually make the cell incapable of prolonged responses to external stimuli. It appears that the intracellular Ca2+ stores may become depleted if not replenished. It has been suggested that the intracellular ER Ca2+ pool has a direct route of access to the extracellular pool, a route that is closed when the ER pool is full.74
In a sense, then, Ca2+ seems to have been downgraded by the inositolphosphates from a "second" to a "third" messenger; however, the pivotal role of Ca2+ as a regulator of cellular activities remains undisputed. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.06%3A_Inositol_Trisphosphate_and_the_Calcium_Ion_Messenger_System.txt |
In order to provide a better understanding of the role of Ca2+ as an almost universal regulator of cellular function, we need to take a brief look at the many ways by which Ca2+ ions can be transported in or out of eukaryotic cells. Although various transport pathways have been elucidated, the present picture is probably not complete, since the molecular structures and properties of the transport proteins are only partially known. The major pathways for Ca 2+ transport across cellular membranes involve three membrane systems: the plasma membrane, the inner mitochondrial membrane, and the membrane of the endoplasmic reticulum (ER) (or, in striated muscle cells, a specialized form of ER called the sarcoplasmic reticulum (SR): (Figure 3.9). Two of the membrane-bound transport systems are Ca2+-ATPases, since they derive their main energy from the hydrolysis of ATP (1 and 2 in Figure 3.9). Their properties do, however, differ in many other respects, as we will see.
The Ca2+-ATPases
The plasma membrane Ca2+-ATPase (PM Ca2+-ATPase) of erythrocytes—first recognized by Schatzmann in 196645—was isolated in pure form by Niggli et al. in 1979, using an affinity column with an ATP-ase binding protein, calmodulin (see Section V.A), coupled to the gel.46 Ca2+-ATPases purified from other types of plasma membranes appear to be very similar. The schematic structure of the erythrocyte membrane Ca2+-ATPase is presented in Figure 3.10.47 The sarcroplasmic reticulum in muscle cells is abundant in Ca2+- ATPase. It is estimated that this protein constitutes more than 80 percent of the integral membrane proteins, and covers a third of the surface area.
The sarcoplasmic reticulum Ca2+-ATPase (SR Ca2+-ATPase) was first purified by MacLennan in 1970.48 Presently it is the best characterized Ca2+-ATPase. A schematic model and a summary of some properties are given in Figure 3.11.49 Ten hydrophobic segments of about 20 amino-acid residues each are revealed by hydropathy plots, and these segments are assumed to span the membrane as $\alpha$-helices. (For the one-letter codes for amino acids, see Appendix B in Section IX.) The phosphorylation site has been identified as Asp-351, and the nucleotide binding domain is following the phosphorylation domain. The Ca2+-binding sites are located within the predicted trans-membrane domains (see Figure 3.11). This was shown through a series of site-directed mutations in which likely Ca2+-liganding residues like Asp, GIu, and Thr were mutated into residues lacking possible side-chain ligands (e.g., Asn, GIn, and Ala).50
The presently accepted reaction cycle involves two main alternative conformations, E1 and E2, the former with two high-affinity sites (Km ≲ 1 $\mu$M)4 on the cytoplasmic side, which in E2 are open to the luminal side with Km ~ 1 mM.49,51 The mechanism suggested for Ca2+ transport (Figure 3.12) has many features similar to that suggested by Williams for H+ translocation in the mitochondrial ATPase.52
It is instructive to consider briefly the thermodynamic limits of the transport. (The discussions about the thermodynamics behind Ca2+/Na+ transport pertain to Na+/K+ gradients in excitable tissues as well). Let us define an "inside" and an "outside" separated by a membrane, as shown in Figure 3.13, where [Ca2+] and $\psi$ denote activities and membrane potentials, respectively. The difference in electrochemical potential, $\Delta \mu$, across the membrane for a Ca2+ ion is given by
$\Delta \mu_{Ca^{2+}} = + RT ln\frac{[Ca^{2+}]_{o}}{[Ca^{2+}]_{i}} + 2F \Delta \psi \tag{3.4}$
where F is Faraday's constant, T the temperature, and R the gas constant. If we assume $\Delta \Psi$ = 0, which appears reasonable for the SR membrane according to experimental evidence, we may calculate the free-energy change, $\Delta$G, at 25 °C for transferring $\Delta$n moles of Ca2+ across the membrane. This becomes $\Delta$G = -$\Delta$n x $\Delta \mu$Ca2+ = $\Delta$n x 4.1 kcal/mol if [Ca2+]o/[Ca2+]j= 10-3 and $\Delta$G = $\Delta$n x 5.4 kcal/mol if [Ca2+]o/[Ca2+]j= 10-4. Under the pertinent cellular conditions, the free-energy change associated with ATP hydrolysis to ADP and Pi has been calculated by Tanford to be $\Delta$G = -13 to -14 kcal/mol.53 In the absence of a membrane potential, it is thus possible to transport two Ca2+ ions for every ATP molecule hydrolyzed against a concentration (or activity) gradient of 104 or more. This treatment says nothing, of course, about the molecular details of this transport. A more detailed model for the transport cycle has been proposed by Tanford.53
In the specialized cells of muscle tissue, the sarcoplasmic reticulum may contain much calcium, and if all were "free" Ca2+, the concentration could be as high as 30 mM.54 This value would cause an osmotic pressure difference across the membrane, as well as put a high demand on the SR Ca2+-ATPase. A lowering of the free Ca2+ concentration inside the SR would clearly be beneficial. In the presence of oxalate or phosphate ions in the external medium, calcium oxalate or phosphate may precipitate inside the sarcoplasmic reticulum, but under normal circumstances it appears that Ca2+ ions inside the SR are bound to a very acidic protein, calsequestrin.54 Each molecule (Mr ≈ 40 kDa) is able to bind 40 to 50 Ca2+ ions with an effective dissociation constant of about 1 mM (at I = 0.1). The protein has a low cation specificity and behaves in many respects like a negatively charged polyelectrolyte. It has been crystallized55 and we may soon have access to its x-ray structure.
The Na+/Ca2+ Exchanger of the Plasma Membrane
Presently available information on the Na+/Ca2+ exchanger has mainly been obtained from studies of the large cells of the giant squid axon and of plasma-membrane vesicles from various other tissues.56,57 In heart plasma-membrane vesicles, the exchanger has the following characteristics: Km = 1.5-5 $\mu$M for Ca2+ and ~20 nM for Na+; Vmax ≈ 20 nmol Ca2+/mg protein.58 The stoichiometry is at least 3:1 Na+/Ca2+. Very few molecular details of the exchanger are available at present. We may again briefly consider the thermodynamic framework for an Na+/Ca2+ exchanger (Figure 3. 14). The difference in electrochemical potential for Na and Ca2+ across the membrane is:
$\Delta \mu_{Ca^{2+}} = RT ln\frac{[Ca^{2+}]_{o}}{[Ca^{2+}]_{i}} + 2F \Delta \Psi, \tag{3.5}$
$\Delta \mu_{Na^{+}} = RT ln\frac{[Na^{+}]_{o}}{[Na^{+}]_{i}} + F \Delta \Psi \ldotp \tag{3.6}$
The free-energy change, $\Delta$GtCa2+, associated with a transfer of $\Delta$nCa2+ moles of Ca2+ from the inside to the outside is $\Delta$GtCa2+ = $\Delta$nCa2+ x $\Delta \mu$Ca2+, and the corresponding change associated with the movement of $\Delta$nNa+ moles of Na+ from the outside in is $\Delta$GtNa+ = - $\Delta$nNa+ x $\Delta \mu$Na+. If these free-energy changes are coupled via the exchanger, there will be a net flux of Ca2+ as long as the free-energy difference,
$\Delta \Delta G = \Delta G_{t}^{Ca^{2+}} - \Delta G_{t}^{Na^{+}} = \Delta n_{Ca^{2+}} \times \Delta \mu_{Ca^{2+}} - \Delta n_{Na^{+}} \times \Delta \mu_{Na^{+}}, \tag{3.7}$
is less than zero. We can write $\Delta \Delta$G for the transport of 1 mol Ca2+ as
$\Delta \Delta G = 2.303 RT \bigg[ log\frac{[Ca^{2+}]_{o}}{[Ca^{2+}]_{i}} - \Delta n_{Na^{+}} \times log\frac{[Na^{+}]_{o}}{[Na^{+}]_{i}} \bigg] + (2 - \Delta n_{Na^{+}}) \times F \Delta \Psi \ldotp \tag{3.8}$
Equating ion activities with concentrations, we note that in a typical mammalian cell [Na+]o ≈ 110-145 mM, and [Na+]i ≈ 7-15 mM, or [Na+]o/[Na+]i ≈ 10. In the absence of a membrane potential difference ($\Delta \Psi$ = 0), Equation (3.8) can thus be simplified to
$\Delta \Delta G = 2.3 RT \bigg[ log\frac{[Ca^{2+}]_{o}}{[Ca^{2+}]_{i}} - \Delta n_{Na^{+}} \bigg] \ldotp \tag{3.9}$
To pump one Ca2+ ion out of a cell against a concentration gradient of about 103 (1 $\mu$M → 1 mM) requires that at least 3 Na+ ions pass in the opposite direction, thus maintaining $\Delta \Delta$G < O. What then will be the effect of a membrane potential difference? Most animal cells, particularly excitable cells such as nerve and muscle cells, have resting potential differences, $\Delta \Psi$, over the plasma membrane of 30 to 90 mV (cytoplasm negative). For this value we find the change in free energy, $\Delta \Delta$G, for the transport of one mol Ca2+ to be
$\Delta \Delta G = 2.3 RT \bigg[ log\frac{[Ca^{2+}]_{o}}{[Ca^{2+}]_{i}} - \Delta n_{Na^{+}} \bigg] + (2 - \Delta n_{Na^{+}}) 0.1 F \ldotp \tag{3.10}$
Thus for $\Delta$nNa+ > 2, we have $\Delta \Delta$G < 0, and the transport of Ca2+ against a concentration gradient of about 103 will be promoted. This is another good reason for having a Na+/Ca2+ exchange stoichiometry of 3:1. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.07%3A_Intracellular_Calcium_Ion_Transport.txt |
Much of our present knowledge of the biological role of Ca2+ ions in the regulation and modulation of cellular activities rests on the development of analytical techniques in three different areas: our ability to measure the low concentration levels in the cytoplasm of resting cells, follow the concentration changes, both temporally and spatially, that may occur as a result of an external stimulus, and measure the distribution of Ca2+ in various compartments of a cell. The last decade has seen the emergence of many such new techniques, and the improvement of old ones, which has had a major impact on our understanding of the detailed molecular mechanisms and dynamics of the Ca2+ messenger system. In this section, we will survey some of the most important techniques and results obtained using these. Broadly speaking there are two main groups of experimental techniques: those that aim at measuring the concentration of "free" (or uncomplexed) Ca2+-ion concentrations (or activities), and those that measure total calcium.
Ca2+-selective Microelectrodes
Ion-selective electrodes can be made from a micropipette (external diameter 0.1-1$\mu$m) with an ion-selective membrane at the tip.18,19 For Ca2+ the membrane can be made of a polyvinyl chloride gel containing a suitable Ca2+-selective complexing agent soluble in the polymer gel. A commonly used complexing agent is "ETH 1001" (see Figure 3.2A). An additional "indifferent" reference electrode is needed. For measurements inside cells, the reference electrode can also be made from a micropipette filled with an electrolyte gel. Often the ion-selective and reference electrodes are connected in a double-barrelled combination microelectrode.21 The whole assembly can then be inserted, using a micromanipulator, into a single cell typically 30-50 $\mu$m across. The arrangement is depicted in Figure 3.2B. With proper care, Ca2+ microelectrodes can be used to measure Ca2+-ion concentrations down to 10-8 M.19,21 One limitation of the technique is that the response time is usually in seconds or even minutes, making rapid concentration transients difficult to follow.
Bioluminescence
Several living oganisms are able to emit light. The light-emitting system in the jellyfish (Aequorea) is a protein called aequorin (Mr $\simeq$ 20 kDa). The light is emitted when a high-energy state involving a prosthetic group (coelenterazine) returns to the ground state in a chemical reaction that is promoted by Ca2+ ions. At Ca2+ concentrations below ~0.3 $\mu M$ the emission is weak, but in the range 0.5-10 $\mu M$ the emission is a very steep function of the concentration (roughly as [Ca2+]2.5).18,19,22 The response to a Ca2+-concentration transient is rapid ($\tau$1/2 = 10 ms at room temperature), and the light emitted can be accurately measured even at very low light levels by means of image intensifiers and/or photon counting. For measurements of Ca2+ concentrations inside cells, aequorin has usually been introduced either through microinjection or through some other means. A novel idea, however, is to utilize recombinant aequorin reconstituted within the cells of interest, thus circumventing the often difficult injection step.174
Complexing Agents with Ca2+-dependent Light Absorption or Fluorescence
An important advance in the field of Ca2+-ion detennination was made by R. Y. Tsien, who in 1980 described23 the synthesis and spectroscopic properties of several new tetracarboxylate indicator dyes that had high affinity and reasonable selectivity for Ca2+. All these dye molecules have a high UV absorbance that is dependent on whether Ca2+ is bound or not; a few also show a Ca2+- dependent fluorescence. Tsien has also demonstrated that these anionic chelators can be taken up by cells as tetraesters, which, once inside the cells, are rapidly enzymatically hydrolyzed to give back the Ca2+-binding anionic forms. Fluorescent tetracarboxylate chelators with somewhat improved Ca2+ selectivity such as "BAPTA," "Quin-2," and "Fura-2" (Figure 3.3) were later described.24 These chelators are very suitable for measurement of Ca2+-ion concentrations in the range 1 $\mu$M to 10 nM in the presence of 1 mM Mg2+ and 100 mM Na+ and/or K+—i.e., conditions typically prevailing in animal cells. Recently a new set of chelators that are more suitable for measurements of calcium concentrations above 1 $\mu$M was presented.25 The most interesting of these is "Fluo-3," with a calcium-binding constant of 1.7 x 106.
Whereas the emission spectrum for Fura-2 (Figure 3.3B), which peaks at 505-510 nm, hardly shifts wavelength when Ca2+ is bound, the absorption spectrum shifts toward shorter wavelengths. In studies of free Ca2+ concentrations where internal referencing is necessary, for example, in studies of single cells, it is therefore advantageous to excite alternately at $\sim$350 and 385 nm, and to measure the ratio of fluorescence intensity at $\sim$510 nm.
The use of fluorescent chelators has recently permitted studies in single cells of rapid fluctuations or oscillations of free Ca2+ and the formation of Ca2+ concentration gradients. Using a fluorescence microscope coupled to a low-light-level television camera feeding a digital image processor, Tsien et al.26 have been able to reach a time resolution of about 1 s in single-cell studies. The results of some highly informative studies made using this instrument are shown in Figure 3.4. (See color plate section, page C-7.) The concentration of free Ca2+ is presented in pseudocolor, and the Fura-2 concentration inside cells is 50-200 $\mu$M, as indicated in the figures. We see a Ca2+ gradient diffusing through an entire sea-urchin egg (~120 $\mu$ across) in 30 s. The free Ca2+ concentration of the resting egg (~100 nM) is increased to about 2 $\mu$M as Ca2+ diffuses through the egg. The mechanism of propagation is believed to be a positive feedback loop with inositol trisphosphate releasing Ca2+ and vice versa (see Section V).
A pertinent question concerning the uses of intracellular Ca2+ chelators is whether or not the chelator significantly perturbs the cell. The chelator will obviously act as a Ca2+ buffer in addition to all other Ca2+-binding biomolecules in the cell. The buffer effect is probably not of any major consequence, since the cell may adjust to the new situation by an increase in total Ca2+, especially if the chelator concentration is in the $\mu$M range. The chelators could, however, interact with and inhibit intracellular enzymes or other molecules, an effect that could result in aberrant cellular behavior. It is not unlikely that BAPTA will bind to certain proteins.27
Complexing Agents with Ca2+-dependent NMR Spectra
A series of symmetrically substituted fluorine derivatives of BAPTA (see Figure 3.3A) has been synthesized.28,29 One of these chelators is 5F-BAPTA (Figure 3.5A), which has a binding constant for Ca2+, KBCa, of 1.4 x 106 M-1 and a 19F NMR chemical shift, $\delta$, that in the free ligand is different from that in the complex with Ca2+ ($\Delta \delta$Ca2+ ≈ 6 ppm). The rate of Ca 2+ dissociation, koff, is 5.7 x 102 s-1, which gives the rate of association, kon, as 8 x 108 M-1s-1 according to
$K_{B} = k_{on} / k_{off} \tag{3.2}$
This exchange rate means that we are approaching the slow exchange limit in 19F NMR, and in subsaturating concentrations of Ca2+ two 19F signals are seen (see Figure 3.5B).
Since the areas of the NMR signals from the bound (B) and free (F) forms of the ligand are proportional to their concentration, the free Ca2+ concentration is obtained simply as
$[Ca^{2+}]_{free} = \frac{B}{F} \cdotp \frac{1}{K_{B}}\ldotp \tag{3.3}$
An additional beneficial property of 5F-BAPTA and other fluorinated analogues of BAPTA is that they will also bind other metal ions with a 19F chemical shift of the complex that is characteristic of the metal ion.29 Under favorable circumstances, it is thus possible to measure simultaneously the concentrations of several cations.
For 5F-BAPTA the selectivity for Ca2+ over Mg2+ is very good (KBMg2+ ≈ 1 M-1). In applications of 5F-BAPTA to intracellular studies, the same protocol is used as with the parent compound and its fluorescent derivatives: some esterified derivative, e.g., the acetoxymethyl ester, is taken up by the cells and allowed to hydrolyze in the cytoplasm. The intracellular concentrations of 5F-BAPTA needed to get good 19F NMR signals depend on the density of cells in the sample tube and the number of spectra accumulated. With accumulation times on the order of ten minutes (thus precluding the observation of concentration transients shorter than this time), Ca2+ concentrations of the order of 1 $\mu$M have been studied in perfused rat hearts using 5F-BAPTA concentrations of about 20 $\mu$M.34 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.08%3A_Measurements_of_Free_Calcium_Concentrations.txt |
The measurement of total calcium in a biological sample can be made by any method sensitive only to the element and not to its particular chemical form. Atomic absorption spectroscopy is excellent as such a method. Obviously, the spatial resolution that can be obtained with this method is limited, and it is hard to imagine its application to elemental mapping of single cells. The techniques discussed in this subsection have been limited to those that permit a spatial resolution of at least 1 $\mu$m on samples usually prepared by sectioning the frozen biological specimens.
Electron Probe and Electron Energy-loss Techniques
When the electron beam in an electron microscope hits a thin sample, some atoms in the sample will be excited or ionized, and returning to their ground state will emit characteristic x-rays. The x-ray emission at different wavelengths may then be measured by a photon-energy-sensitive detector. This is the basis of electron probe x-ray microanalysis (EPMA). The electrons that pass through the sample, and that give the transmission image in electron microscopy, will suffer energy losses that depend on the nature (to some extent also, the chemical state) and distribution of different elements. The outcome of these phenomena forms the basis of electron energy-loss spectroscopy (EELS; see Figure 3.6).
The EPMA technique as applied to calcium has been improved by Somlyo in particular.30 Typically samples are rapidly frozen and sectioned at low temperatures (-130 °C) to preserve the in vivo localization of diffusible ions and molecules. Spatial resolutions of 10 nm or better have been attained on ≳100 nm thick freeze-dried cryosections. The minimal detectable concentration, which requires some signal averaging, is approximately 0.3 mmol Ca per kg dry specimen (i.e., 10 ppm). The calcium content of mitochondria and endoplasmic reticulum in rat liver cells has been studied by EPMA (see Table 3.1).8
The high calcium content of endoplasmic reticulum (ER) is consistent with the view that this organelle is the major source of intracellular Ca2+ released through the messenger inositol trisphosphate (see Section IV.C). Other EPMA studies have shown mitochondria to have a large capacity for massive calcium accumulation in cells where cytoplasmic Ca2+ concentrations have been abnormally high, for example, as a result of damage of the cell membrane.30
EELS is presently less well-developed than EPMA. Two of the major difficulties in the use of EELS for quantitative analysis of calcium and other elements are (i) large background, since it is a difference technique, and (ii) sensitivity to specimen thickness. The major advantage of EELS is that the spatial resolution is potentially much better than in EPMA, and can be 1 to 2.5 nm in favorable specimens.
Proton-induced X-ray Emission (PIXE)
A specimen exposed to a beam of high-energy (1 to 4 MeV) protons will also emit characteristic x-rays just as in EPMA. The advantage of using protons instead of electrons is that protons are more likely to collide with an atom, thus producing excited atoms emitting x-rays. The sensitivity in detecting a particular element is therefore much higher in PIXE than in EPMA or EELS. The PIXE technique, which was developed at the University of Lund, Sweden, in the late 1960s, was originally used mainly for studies of fairly large objects.9
In 1980 a group at Oxford University succeeded in focusing the proton beam to a diameter of 1 $\mu$m with sufficient energy (4 MeV) and beam intensity (100 pA/$\mu$m2) to allow elemental mapping at ppm concentrations.31 Similar beam performances (~0.5 $\mu$m diameter) are now also available at the University of Lund and other laboratories. Beam diameters of 0.1 $\mu$m are likely to be achieved in the near future. Like EPMA, the PIXE method allows the simultaneous observation of several elements in the same sample. The biological applications of the microbeam PIXE technique are limited, but it is clear that its potential is great. Some representative results obtained with the Oxford microbeam are shown in Figure 3.7. (See color plate section, page C-8.)
Ion Microscopy
Ion microscopy is another technique capable of detecting all elements at the ppm level. The basic idea is to expose a freeze-fixed, cryofractured, and freeze-dried sample, which has been put onto a conducting substrate in a vacuum chamber, to a beam of ions (e.g., D2+ or Ar+). These ions will remove the top two or three atomic layers of the sample surface by sputtering. A certain fraction of the removed atoms will leave as ions. This secondary ion beam is accelerated into a double-focusing mass spectrometer, where the ions are separated according to their mass-to-charge ratio. The ion optics are designed to preserve the spatial distribution of the emitted secondary ions, and an element image of the sample can thus be produced with a spatial resolution of ~0.5 $\mu$m.32 The ion-microscope technique can form images of a particular isotope of an element. In principle, then, one could perform isotope labeling or "isotope chase" studies and follow, say, the fate of isotope-enriched 43Ca externally applied to a cell. The ion-microscope technique has not yet come into widespread use, but the quality of element (or ion) images obtained on single cells is impressive.33 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.09%3A_Measurements_of_Total_Calcium_Concentrations.txt |
Influx
Mitochondria isolated from various types of animal cells—but, interestingly, not those from plant cells—can rapidly accumulate exogenous Ca2+.59 The transporter is located in the inner membrane and the driving force behind the Ca2+ transport appears to be merely the high potential difference across this membrane ($\Delta \Psi$ ≈ 150 to 180 mV, negative in the inner matrix). This potential difference is fairly closely maintained by the pumping out of H+ from the matrix by cell respiration. For the transport of 1 mol Ca2+ from the "outside" (= cytoplasm) to the "inside" (= inner mitchondrial matrix), we may deduce from Equation (3.4) that the free-energy change $\Delta$G may be written ($\Delta$nCa2+ = -1)
$\Delta G = - RT \cdotp ln\frac{[Ca^{2+}]_{o}}{[Ca^{2+}]_{i}} - 2F \Delta \Psi \ldotp \tag{3.11}$
From this analysis it may be inferred that the limiting Ca2+ concentration (or activity) ratio that can be achieved by this electrogenic pump (i.e., $\Delta$G = 0) is
$\dfrac{[Ca^{2+}]_{o}}{[Ca^{2+}]_{i}} = e^{\frac{-2 F \Delta \Psi}{RT}} \tag{3.12}$
With $\Delta \Psi$ = 150 mV, this ratio is calculated to be 8.4 x 10-6 at 25 °C. It is evident that, as long as the Ca2+ influx would not lower the membrane potential difference, the Ca2+ uniporter has a very high pumping potential. Measured values of the pumping rate, Vmax, are indeed high (>10 nmol/mg protein59) and probably limited only by the rate of electron transport and H+ extrusion in the mitochondria.
Mitochondria may accumulate large quantities of Ca2+, probably to maintain electroneutrality. To prevent the buildup of high concentrations of free Ca2+ (and of osmotic pressure), phosphate ions are also transported into the inner matrix, where an amorphous calcium phosphate—or possibly a phosphocitrate60—is formed. The equilibrium concentration of free Ca2+ in the mitochondrial matrix may as a result be comparatively low, on the order of 1 $\mu$M.
The molecular nature of the mitochondrial Ca2+ uniporter continues to be elusive, and needs to be studied further.
Efflux
Mitochondria, as well as SR, release Ca2+ ions by mechanisms other than "back leakage" through the pumps. In mitochondria from excitable cells, the efflux occurs mainly through an antiport, where 2 Na+ ions are transported inward for every Ca2+ ion departing for the cytosolic compartment.61 In other cells there is evidence for the dominance of a 2H+-Ca2+ antiport.59 In all likelihood the Ca2+ efflux is regulated, possibly by the redox state of pyridine nucleotides in the mitochondria. As with the Ca2+ uniporter, few details on the molecular nature of the antiporters are presently available.
Ca2+ Efflux from Non-mitrochondrial Stores
Release of Ca2+ from ER and SR presently appears to be the prime effect of the new intracellular messenger 1,4,5-triphosphoinositol (1,4,5-IP3) released into the cytoplasm as a result of an external hormonal stimulus (see Section IV.C). It seems that receptors for 1,4,5-IP3 have been established on ER, and that the binding of 1,4,5-IP3 causes a release of Ca2+ stored in this organelle.62,63,170,171 In addition to the receptor-controlled Ca2+ efflux, there may be other pathways for Ca2+ release, and Ca2+ mobilization may be regulated by other intracellular entities, the Ca2+ ions themselves included.
Other Voltage-gated or Receptor-activated Ca2+ Channels
In addition to the transport pathways already discussed, some cells seem to have Ca2+ channels in the plasma membrane that can be opened by the action of an agonist on a receptor or that are gated in response to changes in membrane potential.64 For example, Ca2+ channels can be opened by nicotinic cholinergic agonists65 or by the excitatory amino acid N-methyl-D-aspartate (NMDA).66 Endochrine cells and also some muscle and neuronal cells have voltage-sensitive Ca2+ channels.67,68 We will not discuss these further, but merely point to their existence. We finally note that during the last few years knowledge about the mechanisms of Ca2+ entry and release to and from extracellular and intracellular pools has increased dramatically, and we refer the reader to recent reviews of the field.175,176 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.10%3A_Mitochondrial_Calcium_Ion_Transport.txt |
So far we have mainly discussed the routes and means by which the concentration of Ca2+ ions in the cytoplasm can be transiently increased and brought back to resting levels. But changing the cytoplasmic Ca2+ concentration is not enough. In order to influence the cellular machinery, the Ca2+ ions must interact with different proteins, intracellular Ca2+ receptors if you like. These intracellular Ca2+-receptor proteins must have certain properties in order to function.
1. Their Ca2+-affinity must be such that their Ca2+-binding sites are essentially unoccupied at resting levels of free Ca2+ (~10-7 M) and occupied at levels reached upon stimulus (generally assumed to be 10-5 to 10-6 M). This means that the binding constants KBCa2+ should be ~106 M-1.
2. We should also remember that Ca2+ must exert its function in the presence of a number of other ions; in mammalian cells the intracellular concentration of "free" Mg2+ ions is around 1 mM, and that of K+ ions around 100 to 150 mM. The receptors must therefore have an adequate selectivity for Ca2+.
3. In response to Ca2+ binding, a Ca2+ receptor must undergo some kind of conformation change that either alters its interaction with other molecules or changes its activity if it is an enzyme.
4. Finally, there are kinetic considerations. In many cells a rapid response is essential, and therefore the receptors must be able to interact swiftly—within milliseconds—with incoming Ca2+ ions, and the ions must also be able to depart almost as rapidly.
A few proteins have been discovered that qualify as intracellular Ca2+ receptors. The best known of these is calmodulin (CaM), which appears to be present in all eukaryotic cells. Most of the cellular responses elicited by Ca2+ appear to result from interactions between the Ca2+-calmodulin complex and various other target enzymes and proteins.75 Another important Ca2+-receptor protein is troponin C (TnC), which occurs in muscle cells and is instrumental in mediating muscle contraction.76 These two types of proteins are highly homologous, as we shall see, and may be considered members of a superfamily of closely related intracellular Ca2+-binding proteins. This superfamily has been given the name "the calmodulin superfamily," and close to 200 distinct family members are presently known.77 Not all members of the superfamily may qualify as Ca2+ receptors; some like parvalbumins and calbindins (see Section IV.A) appear to have a role in intracellular transport and/or Ca2+-buffering. For others, such as the S-100 proteins78 found predominantly in brain tissue, and calcimedins,79 isolated from smooth muscle, the biological function is still unclear.
One Ca2+ receptor with enzymatic activity is protein kinase C. Its activity is markedly increased in the presence of Ca2+, and it has a high calcium-binding constant (see Table 3.2) in the presence of diacylglycerol or phorbol esters.80
During recent years, groups interested in the role of Ca2+ in secretion and in the control of membrane cytoskeleton have identified some intracellular Ca2+/phospholipid-binding proteins that appear to be distinct from the calmodulin superfamily; these include lipocortin, endonexin, calelectrin, p36, and calpactin.81-83 These membrane-binding proteins are collectively called annexins,84 and contain repeated domains distinct from EF-hands. The Ca2+ sites are very similar to that observed in phospholipase A2, as shown by the recently determined x-ray structure of annexin V.172 A condensed overview of the interaction of Ca2+ with intracellular proteins is shown in Figure 3.16. We will now go on to discuss the molecular properties of some of the proteins mentioned above, starting with calmodulin. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.11%3A_Molecular_Aspects_of_Calcium_Ion-Regulated_Intracellular_Processes_%28Part_1%29.txt |
Sarcoplasmic Calcium-binding Protein from Nereis diversicolor
The calmodulin superfamily of proteins also includes' sarcoplasmic Ca2+-binding proteins (SCPs) that can be found in both vertebrate and invertebrate muscle.129 The function of SCPs is not yet known, but their sequence homology with Ca2+-binding proteins of known tertiary structure suggests that they originally contained four helix-loop-helix Ca2+-binding domains. Ca2+ binding has been preserved in the first and third domains of all known SCPs, but only one, if any, of domains II and IV is functional. The three-dimensional crystal structure of an SCP from the sandworm Nereis diversicolor analyzed at 3.0 Å resolution130 can be seen in Figure 3.26. (See color plate section, page C-11.) The C-terminal half (domains III and IV) of the molecule contains two Ca2+- binding EF-hands (green and red in Figure 3.26) similar to calbindin D9k and the globular domains of troponin C and calmodulin. The N-terminal half is, on the contrary, markedly different from the normal helix-loop-helix geometry. Domain I binds Ca2+ with a novel helix-loop-helix conformation, whereas domain H lacks calcium-binding capacity. The two halves are packed closely together, and are not, as in troponin C or calmodulin, connected by a solvent-exposed $\alpha$-helix.
Membrane Cytoskeleton and Phospholipid Binding Proteins
It has long been suspected that Ca2+ ions are somehow involved in exocytosis. Recently several groups131 have isolated intracellular proteins that associate with membranes, and/or membrane cytoskeleton proteins, in a Ca2+-dependent manner, and that seem able to mediate vesicle fusion or aggregation at Ca2+ concentrations above 200 $\mu$M. These proteins—endonexin, calelectrin, p36, and pII—have stretches of consensus amino-acid sequences that are also found in a phospholipase A2 inhibitor protein, lipocortin.132 It appears that further studies of this new class of proteins, known as annexins, will lead to new insights into cell-signaling pathways. Multiple functions have been proposed for the annexins, but no cellular role has yet been defined.133 The first crystal structure of an annexin, human annexin V—which in vitro will form voltage-gated Ca2+ channels—has been determined recently.172 In annexin, the three Ca2+-binding sites are located on the side of the molecule that is involved in membrane binding.
Ca2+-dependent Proteases
An interesting Ca2+-activated intracellular protease, sometimes called calpain, was discovered during the last decade.134 The ending -pain refers to its relation with other proteolytic enzymes like papain. It may seem dangerous to have a proteolytic enzyme loose inside a cell, and it must have rather specialized functions and be under strict control. The complete primary structure of the calcium protease (Mr ≈ 80,000) in chicken tissues has recently been deduced from the nucleotide sequence of cloned DNA.135,136 The findings are quite unexpected.
The protein contains four distinct domains. The first and third domains have no clear sequence homologies with known protein sequences, but the second domain has a high homology with the proteolytic enzyme papain, and the fourth domain is highly homologous to calmodulin. This fourth domain thus has four EF-hand-type Ca2+-binding sites, although the third site has a somewhat unusual loop sequence. Here we apparently are faced with an unusual invention by Nature: by fusing the gene for a protease with that of the canonical Ca2+ receptor, she has created a molecule in which a regulatory protein is covalently linked to its target enzyme!
Protein Kinase C
Before we leave our brief survey of intracellular Ca2+-binding proteins, we must write a few lines about an important Ca2+-regulated kinase (a phosphorylating enzyme), i.e., protein kinase C (PKC). The activity of this enzyme, or rather family of enzymes,137 appears to be regulated by three factors: phospholipids, in particular phosphatidylserine; diacyl-glycerols, one of the products of inositol lipid breakdown; and Ca2+ ions. The high-activity form of PKC, which appears responsible for much of the phosphorylation activity of many cells, is presumably membrane-bound, whereas the low-activity form may be partly cytosolic (Figure 3.27). The schematic structure of rabbit PKC (Mr ≈ 77 kDa) according to Ohno et al.138 is shown in Figure 3.28. The Ca2+ site(s) are presumably in the regulatory domain. No typical "EF-hand" pattern has been found in the amino-acid sequence. A protein kinase that requires Ca2+ but not phospholipids nor calmodulin for activity has been purified from soybean. From the amino-acid sequence the protein appears to have a calmodulin-like Ca2+-binding domain, very much as in calpain.139 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.12%3A_Molecular_Aspects_of_Calcium_Ion-regulated_Intracellular_Processes_%28Part_2%29.txt |
A few intracellular Ca2+-binding proteins have been discovered that by sequence homology clearly belong to the CaM-TnC family with Ca2+ sites of the "EF-hand"-type, but that do not appear to exert a direct regulatory function. Parvalbumins (Mr ≈ 12 kDa), calbindin D9K (Mr ≈ 8.7 kDa) and calbindin D28K (Mr ≈ 28 kDa) belong to this group. Parvalbumin(s) exist in two main types, $\alpha$ and $\beta$, found in large quantities in the white muscle of fish, amphibia, and reptiles, but also in different mammalian tissues,116,117 including neurons of the central and peripheral nervous system. The molecule has two fairly strong Ca2+- binding sites (see Table 3.2). The x-ray structure of carp parvalbumin was solved in 1973 by Kretsinger et al.,118 and for a decade provided the basis for all discussions on intracellular Ca2+-binding proteins. The concept of the canonical "EF-hand" Ca2+-binding site originated from the parvalbumin work, and the name "EF" derives from the labeling of the two helices that flank the second of the two Ca2+ sites in parvalbumin, as shown in Figure 3.23.
If the first Ca2+ ligand in the approximately octahedral coordination sphere is given number 1 (or "x") the others come in the order 3( "y"), 5("z"), 7("-y"), 9("-x"), and 12("-z"). In the second site of parvalbumin, "-x" is actually a H2O molecule, but in the first site it is the carboxylate of a Glu. Studies118 of putative Ca2+-binding sites in other proteins with known primary sequences led to the generalized EF-hand structure—including residues in the flanking $\alpha$-helices—shown in Figure 3.24.
This sequence, with minor modifications, has been widely used in searching for "EF-hands" in libraries of amino-acid (or DNA) sequences of new proteins with unknown properties. In this way, calbindin D28k a protein with unknown function, initially discovered in chicken intestine, but later found also in brain, testes, and other tissue, has been shown to have four EF-hand sites.119
Recently two structures of carp parvalbumin, both with a resolution of 1.6 Å, were published.120 One of these structures is the native calcium-loaded form of the protein; the second is the structure of parvalbumin in which Ca2+ has been replaced by Cd2+. No significant differences are observed upon replacement of calcium by cadmium. 113Cd has a nuclear spin of I = $\frac{1}{2}$, making it much more amenable to NMR studies than the quadrupolar 43Ca (I = $\frac{7}{2}$), This study supports the use of 113Cd NMR as a tool for the study of calcium-binding proteins.121
The function of parvalbumin has long been assumed to be that of buffering Ca2+ in muscle cells, i.e., taking up Ca2+ ions released from Ca2+-troponin complexes, thereby ensuring that the cytoplasmic levels of free Ca2+ are always kept very low, even during short bursts of muscle activity,122 The widespread occurrence of parvalbumin in non-muscle tissue indicates that it probably has other roles as well.
Calbindin D9k (Mr ≈ 8.7 kDa) is another intracellular Ca2+-binding protein with unknown function. It was briefly mentioned in connection with Ca2+ uptake and transport in the intestine and placenta (Section IV.A). Like the avian calbindin D28k the D9k calbindin has been observed in many types of tissue. The homology between the D9k and D28k calbindins is much less than the name suggests; both their syntheses are, however, regulated by vitamin D. The x-ray structure of bovine calbindin D9k has been determined123 and refined to a resolution of 2.3 Å, and a three-dimensional solution structure of porcine calbindin D9k is also available.124 The average solution structure calculated from NMR data is shown in Figure 3.25 (See color plate section, page C-10.)
The protein has four main $\alpha$-helices and two Ca2+-binding loops (I and II). The interior of the molecule shows a loose clustering of several hydrophobic side chains; in particular, three phenylalanine rings come very close in space. The Ca2+-binding loops constitute the least-mobile parts of the molecule. The crystallographic temperature factors have pronounced minima in these regions, with the lowest overall B-factor observed in loop II. Both Ca2+ ions are roughly octahedrally coordinated with protein oxygen atoms. There are some striking differences between the two sites, however. Whereas the C-terminal site (II) has a general structure very similar to the archetypal "EF-hand," as observed in CaM, sTnC, and parvalbumin, the N-terminal site (I) has an extra amino-acid residue inserted between vertices x and y, and z and -y (see Figure 3.24). As a consequence, the peptide fold in site I is different from that in site II. Three carboxylate groups are ligands in site II, but in site I there is only one.
Despite this marked difference in charge and peptide fold, the Ca2+ affinity of both Ca2+ sites is remarkably similar, as has been shown in a study in which site-directed mutagenesis was combined with different biophysical measurements.37 Cooperative Ca2+ binding in the native calbindin D9k (the "wild type") was first demonstrated at low ionic strength by means of the values of the two stoichiometric Ca2+-binding constants, K1 and K2, which could be measured with good accuracy (K1 = 4.4 x 108 M-1 and K2 = 7.4 x 108 M-1). The effects of amino-acid substitutions in Ca2+ site I were primarily localized to this site, with virtually no effects on the structure or other biophysical properties pertinent to site II. The appearance of sequential Ca2+ binding in some of the calbindin mutants did allow the identification of 1H NMR resonances that respond primarily to binding of Ca2+ to either one of the sites. This result in tum permitted an estimate of the ratio between the site-binding constants (KA and KB) in the wild-type protein and in one of the mutant proteins (Tyr-13 → Phe). In this way the reseachers125 could assess, to within narrow limits, the free energy of interaction, $\Delta \Delta$G, between the two Ca2+ sites as 7.7 kJ/mol at low ionic strength and 4.6 kJ/mol in the presence of 0.15 M KCl. How this site-site interaction is transmitted on a molecular level is still unknown.
Through a combination of site-specific mutations and biophysical measurements, it has recently been demonstrated that carboxylate groups at the surface of the protein, but not directly ligated to the bound Ca2+ ions, have a profound effect on the Ca2+ affinity.126 Neutralization of the surface charges reduces affinity and increases the stability of the protein toward unfolding by urea.127
A surprising discovery about the structure of bovine calbindin D9k in solution has also been made recently.128 Detailed analysis of the 2D 1H NMR spectrum of wild-type calbindin has revealed that it exists as a 3:1 equilibrium mixture of two forms, corresponding to a trans and cis conformation around the Gly-42-Pro-43 peptide bond. The global fold appears essentially the same in the two forms, and structural differences are primarily located in the inter-domain loop in which Pro-43 is located. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.13%3A_Paravalbumin_and_Calbindins_%28D_9K%29_and_%28D_28K%29.txt |
All living organisms need calcium, which must be taken up from the environment. Thus, Ca2+ ions have to be distributed throughout the organism and made available where needed. In higher organisms, such as humans, the blood-plasma level of total calcium is kept constant (≈2.45 mM) within narrow limits, and there must be a mechanism for regulating this concentration. On a cellular level we have already seen in the preceding section that the basal cytoplasmic Ca2+ concentration, at least in eucaryotic cells, is very low, on the order of 100 nM. At the same time the concentrations of Ca2+ in certain organelles, such as endoplasmic (or sarcoplasmic) reticulum or mitochondria, may be considerably higher. If Ca2+ ions are to be useful as intracellular "messengers," as all present evidence has it, Ca2+ levels in the cytoplasm would have to be raised transitorily as a result of some stimulus. Ca2+ ions may enter the cytoplasm either from the extracellular pool or from the Ca2+-rich organelles inside the cell (or both). We could imagine Ca2+ channels being regulated by chemical signaling, perhaps by a hormone acting directly on the channel, or by a small molecule released intracellularly when a hormone is attached to a membrane-bound receptor. Some channels may be switched on by voltage gradients, and both these mechanisms may operate concurrently.
Increased intracellular Ca2+ levels must eventually be brought back to the basal levels, in some cells very quickly. The ions could be transported out of the cell or back into the Ca2+-rich organelles. This transport will be against an electrochemical potential gradient, and thus requires energy. There are many possibilities for different forms of Ca2+ transport and regulation in living systems, and we still know fairly little about the whole picture. Detailed studies are also complicated by the fact that, in higher organisms, cells are differentiated. Nature is multifarious, and what is valid for one type of cell may not be relevant for another. With these words of caution we will start out on a macroscopic level and continue on toward molecular levels.
Ca2+ Uptake and Secretion
The uptake of Ca2+ from food has mostly been studied in typical laboratory animals, such as rats, hamsters, chickens, and humans. In humans, uptake occurs in the small intestine, and transport is regulated by a metabolite of vitamin D, calcitriol (1,25-dihydroxy vitamin D3).34 The uptake process is not without loss; roughly 50 percent of the calcium content in an average diet is not absorbed. To maintain homeostasis and keep the calcium level in blood plasma constant, excess Ca2+ is excreted through the kidney. The main factor controlling this phenomenon in vertebrates is the level of the parathyroid hormone that acts on kidney (increases Ca 2+ resorption), on bone, and, indirectly, via stimulated production of calcitriol, on the intestinal tract (increases Ca2+ uptake). Calcium enters the cells from the outside world, i.e., the intestinal lumen, by traveling through the brush-border membrane of the intestinal epithelial cells, through the cytosolic interior of these cells, and into the body fluids through the basal lateral membranes of the same cells. The molecular events involved need to be studied further. Figure 3.8 outlines the Ca2+ transport processes known or thought to occur.
Transfer through the brush-border membrane is assumed to be "passive" although indirectly facilitated by calcitriol. The calcitriol effect may be due to synthesis of a carrier protein,35 but could also be an effect of altered membrane lipid composition. 36 The fate of Ca2+ ions, once inside the epithelial cell, is a much-debated subject. What appears clear is that the Ca2+ ions entering through the brush-border membrane do not cause an increase of the low cytosolic Ca2+ concentration. It is thus quite likely that the Ca2+ ions are carried through the cell but the means of transportation is unknown. One plausible carrier is the intracellular low-molecular-weight Ca2+-binding protein calbindin D9K (Mr ≈ 9 kDa) formerly known as ICaBP (see Section V.C).35 Its synthesis is induced by vitamin D, and it is mainly found in mammalian intestines. The porcine and bovine calbindin D9K has a Ca2+ binding constant of KB ≈ 3 x 108 M-1 in low ionic strength media37 and KB = 2 x 106 M-1 in the presence of 1 mM Mg2+ and 150 mM K+.38 The concentration of calbindin D9K in epithelial cells can reach millimolar levels,35 which could allow it to facilitate Ca2+ diffusion across the cytosol. This was first suggested by Williams, subsequently elaborated by Kretsinger et aI. in 1982,39 and later demonstrated in a model cell by Feher.40 The basic idea is that, although the diffusion rate of Ca2+ ions (~10-5 cm2 s-1) is higher than for the (Ca2+)2 calbindin complex (~0.2 x 10-5 cm2 s-1), the fact that the concentration of the latter complex may be about 103 times higher than that of free Ca2+ will result in an increased net calcium transport rate. Calbindin would, in fact, act very much like myoglobin in facilitating oxygen transport through muscle tissue.
Plausible as the above mechanism may seem, it may, however, not be the whole truth. An alternative mechanism is vesicular transport. In chicken intestine it has been shown that the only epithelial organelles that increased in Ca2+ content as a result of calcitriol treatment were the lysosomes.41 The result lends support to a transport mechanism involving Ca2+ uptake across the brush-border membrane by endocytic vesicles, fusion of these vesicles with lysosomes, and possibly also delivery of Ca2+ to the basal lateral membrane of the epithelial cell by exocytosis. This process would also explain the vitamin-D-induced alterations in brush-border-membrane lipid compositions as a consequences of preferential incorporation of certain types of lipids into the vesicles. Interestingly, the lysosomes in the chicken studies also contained high levels of calbindin D28k—a type of vitamin-D-induced Ca2+-binding protein found in avian intestines—making it conceivable that this protein acts as a "receptor" for Ca2+ at the brush-border membrane and upon Ca2+ binding could become internalized in endocytic vesicles.41
The basal lateral plasma membrane contains at least two types of Ca2+ pumps that also may play a role in Ca2+ uptake, one ATP-driven, one driven by a concurrent flow of Na+ ions into the cytoplasm (i.e., a Na+-Ca2+ antiport; see Figure 3.8). We discuss these types of transporting proteins in the next subsection.
There are some apparent analogies between intestinal Ca2+ transport and that occurring in the placenta. Transplacental movements of Ca2+ increase dramatically during the last trimester of gestation.42 In mammalian placental trophoblasts, high concentrations of calbindin D9K are found.43,44 The protein synthesis also in this tissue appears to be under calcitriol regulation. Ca2+ ions have to be supplied by mammalian females, not only to the fetus during pregnancy, but also to the newborn child through the mother's milk. The molecular details of Ca2+ transport in the mammalian glands have not been extensively studied. In milk, Ca2+ is bound mainly to micelles of casein, and the average Ca2+ content is reported to be 2.5 g/liter (see Table 3.1). | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.14%3A_The_Transport_and_Regulation_of_%28Ca2%29_Ions_in_Higher_Organisms.txt |
The contraction of striated muscle is triggered by Ca2+ ions. Muscle cells are highly specialized, and contain two types of filaments that may slide past each other in an energy-consuming process. One of the filaments, the thin filament, is built up by actin molecules (Mr ≈ 42 kDa) polymerized end-to-end in a double helix. In the grooves of this helix runs a long rod-like molecule, tropomyosin; and located on this molecule at every seventh actin, is a complex of three proteins, troponin. The three proteins in the troponin complex are troponin I (TnI), troponin T (TnT), and troponin C (TnC). A schematic picture of the organization of the thin filament is shown in Figure 3.20.
Troponin C is the Ca2 -binding subunit of troponin, and it is structurally highly homologous to calmodulin. Skeletal-muscle troponin C (sTnC; Mr ≈ 18 kDa) can bind four Ca2+ ions, but cardiac-muscle troponin C (cTnC) has one of the four calcium sites modified, so that it binds only three Ca2+ ions. The x-ray structures of sTnC from turkey and chicken skeletal muscle have been determined to resolutions of 2.8 and 3.0 Å, respectively.102,103 The structure of turkey sTnC is shown in Figure 3.21.
The similarity between the structures of CaM (Figure 3.17) and sTnC is obvious. In sTnC we again find two domains, each with two potential Ca2+ sites, separated by a 9-turn $\alpha$-helix. The crystals were grown in the presence of Ca2+ at a low pH (pH = 5), and only two Ca2+ ions are found in the C-terminal domain. The two Ca2+-binding sites in this domain have the same helix-loop-helix motif that is found in CaM, and they both conform to the archetypal EF-hand structure. The interhelix angles between helices E and F and between G and H are close to 110°. By contrast, the helices in the N-terminal domain, where no Ca2+ ions are bound, are closer to being antiparallel, with interhelix angles of 133° (helices A and B) and 151° (helices C and D).
Both sTnC and cTnC have two high-affinity Ca2+-binding sites (see Table 3.2) that also bind Mg2+ ions competitively, although with a much lower affinity. These two sites are usually called "the Ca2+-Mg2+ sites." 76,104 In sTnC there are also two (in cTnC, only one) Ca2+-binding sites of lower affinity (KBCa2+ ≈ 105 M-1) that bind Mg2+ weakly and therefore have been called "the Ca2+-specific sites." Since Ca2+ binding to the latter sites is assumed to be the crucial step in the contractile event, they are often referred to as "the regulatory sites" (see below). The existence of additional weak Mg2+ sites (KB ≈ 300 M-1) on sTnC, not in direct competition with Ca2+, has also been inferred.76,104,105 Spectroscopic studies have shown that the two strong Ca2+- Mg2+ sites are located in the C-terminal domain, and the weaker Ca2+-specific sites in the N-terminal domain of sTnC.106 This pattern is similar to that observed with CaM. NMR spectroscopic studies strongly suggest that binding of Ca2+ to both sTnC and cTnC is cooperative.107 In sTnC the C-terminal domain binds Mg2+ much more strongly than the N-terminal domain, by contrast to CaM, where the reverse is true.
The rates of dissociation of Ca2+ and Mg2+ from sTnC have been measured by both stopped-flow and 43Ca NMR techniques.76,108 As with CaM, the actual numbers depend on the solution conditions, ionic strength, presence of Mg2+, etc. (see Table 3.4). On the rate of Mg2+ dissociation from the Ca2+- Mg2+ sites, quite different results have been obtained by stopped-flow studies76 of fluorescence-labeled sTnC (koffMg2+ ≈ 8 s-1) and by 25Mg NMR (koffMg2+ $\simeq$ 800-1000 s-1).109 This apparent discrepancy seems to have been resolved by the observation that both binding and release of Mg2+ ions to the Ca2+-Mg2+ sites occur stepwise, with koffMg2+ < 20 s-1 for one of the ions, and koffMg2+ ≥ 800 s-1 for the other.110 The rates of dissociation of the Mg2+ ions are important, since under physiological conditions the Ca2+-Mg2+ sites of sTnC are likely to be predominantly occupied by Mg2+ ions, release of which determines the rate at which Ca2+ can enter into these sites.
Spectroscopic and biochemical data111 indicate that upon binding Ca2+, sTnC and cTnC undergo significant conformation changes. Comparisons of NMR spectroscopic changes on Ca2+ binding to intact sTnC, as well as to the two fragments produced by tryptic cleavage (essentially the N-terminal and C-terminal halves of the molecule, just as was the case with CaM), have shown that the conformation changes induced are mainly localized within the domain that is binding added ions.110,112 Thus the central $\alpha$-helix connecting the domains seems unable to propagate structural changes from one domain to the other. It has been suggested that the structural differences found in the x-ray structure of turkey sTnC between the C-terminal domain, which in the crystal contains two bound Ca2+ ions, and the N-terminal domain, in which no Ca2+ ions were found, may represent these conformational changes.113 This rather substantial conformational change is schematically depicted in Figure 3.22.
However, preliminary structure calculations114 of the calcium-saturated and calcium-free forms of calbindin D9k indicate that much more subtle conformational changes take place upon binding Ca2+ in calbindin D9k. Interestingly, 1H NMR spectroscopy has provided evidence for the concept that the structural change induced by Mg2+ binding to the C-terminal domain of sTnC must be very similar to that induced by Ca2+ ions. Another result obtained by 113Cd NMR studies108 is that the cadmium-loaded N-terminal domain of sTnC in solution undergoes a rapid interchange between two or more conformations, with an exchange rate on the order of 103-104 s-1.
Just as CaM exerts its biological function in complexes with other proteins, TnC participates in the three-protein troponin complex. It presently appears that TnC and TnI form a primary complex that is anchored by TnT to a binding site on tropomyosin.115 In the troponin complex the Ca2+ affinity is increased by a factor of about ten over that in isolated sTnC, both at the Ca2+-Mg2+ sites and at the Ca2+-specific sites. A similar increase in affinity is found for Mg2+. Given the amounts of "free" Mg2+ inside muscle cells (1 to 3 mM), it seems likely that the Ca2+-Mg2+ sites in the resting state of troponin are filled with Mg2+, so that a transitory release of Ca2+ leads primarily to rapid Ca2+ binding to the Ca2+-specific sites, and subsequently to conformation change and contraction. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/03%3A_Calcium_in_Biological_Systems/3.15%3A_Troponin_C.txt |
Contributors and Attributions
• Geoffrey B. Jameson (Georgetown University, Department of Chemistry)
• James A. Ibers (Northwestern University, Department of Chemistry)
04: Biological and Synthetic Dioxygen Carriers
Most organisms require molecular oxygen in order to survive. The dioxygen is used in a host of biochemical transformations, although most is consumed in the reaction
$O_{2} + 4 H^{+} + 4 e^{-} \rightarrow 2H_{2}O \tag{4.1}$
that is the terminal (or primary) step of oxidative phosphorylation (Chapters 5 and 6). For some small animals and for plants, where the surface-to-volume ratio is large, an adequate supply of dioxygen can be obtained from simple diffusion across cell membranes. The dioxygen may be extracted from air or water; for plants that produce dioxygen in photosynthesis, it is also available endogenously. For other organisms, particularly those with non-passive lifestyles, from scorpions to whales, diffusion does not supply sufficient dioxygen for respiration.
An elegant three-component system has evolved to transport dioxygen from regions of high abundance—water (at least if free of pollutant reductants) and air—to regions of relatively low abundance and high demand—the interior cells of the organism. This process is illustrated in Figure 4.1.1-3
The central component is a dioxygen-carrier protein. In the three chemically distinct carriers that have evolved and are found today, the dioxygen-binding site in the protein, that is, the so-called "active site," is a complex either of copper or of iron.4-6 For hemoglobins, the most widely distributed family of dioxygen carriers, the active site has long been known to consist of an iron porphyrin (heme) group embedded in the protein. Almost all hemoglobins share the basic structure illustrated in Figure 4.2.7-12 Hemocyanin13-15 and hemerythrin,16-18 the other two biological dioxygen carriers, feature pairs of copper atoms and iron atoms, respectively, at the active sites.* Some basic properties of these metalloproteins are summarized in Table 4.1.4-6
Table 4.1: General Features of Dioxygen-carrier Proteins
Metalloprotein Active Site of deoxy Color Change deoxy → oxy MW (Dalton) # Subunits Average MW Subunit (Dalton)
Hemoglobins
Vertebrate
Human A heme FeII purple → red 64,000 4 16,000
Invertebrate
Erythrocruroin (Lumbricus terrestris, earthworm) heme FeII purple → red up to 3.3 x 106 192 17,000
Chlorocruorn (Eudistylia vancouveri) chloroheme FeII purple → green 3.1 x 106 192 15,000
Hemocyanins
Mollusc (Helix pomatia-$\alpha$, edible snail) Cu1 . . . Cu1 colorless → blue -9 x 106 160 52,700
Arthropod (Cancer magister, crab) Cu1 . . . Cu1 colorless → blue -9 x 105 12 76,600
Hemerythrins (Phascolopsis syn. golfingia gouldii) FeII . . . FeII colorless → burgundy 108,000 8 13,500
* The use of the prefix hem- is confusing. In this context hem connotes blood. Thus, since hemocyanin and hemerythrin lack a heme group [an iron(II) porphyrin], they are nonheme metalloproteins.
The second component of the dioxygen-transport system facilitates the sequestration of dioxygen by the dioxygen-carrier protein. Specialized organs, such as lungs in air-breathing creatures or gills in fish, offer a very large surface area to the outside environment in order to facilitate diffusion. The third component is the delivery system. The oxygen carrier is dissolved or suspended in a fluid, called blood plasma or hemolymph, that is pumped throughout the animal by another specialized organ, the heart, through a network of tubes, the blood vessels. In many organisms an additional dioxygen-binding protein, which stores dioxygen, is located in tissues that are subject to sudden and high dioxygen demand, such as muscles. These dioxygen-storage proteins are prefixed myo- (from the Greek root mys for muscle). Thus for the dioxygen-transport protein hemerythrin there exists a chemically similar dioxygen-storage protein myohemerythrin. For the hemoglobin family the corresponding storage protein is called myoglobin. Interestingly, some organisms that use hemocyanin as the dioxygen-transport protein use myoglobin as the dioxygen-storage protein.
At the center of biological dioxygen transport are transition-metal complexes of iron or copper. To model such systems, chemists have prepared several synthetic oxygen carriers, especially of iron and cobalt porphyrins. In this chapter the structures and properties of biological and nonbiological oxygen carriers are described, with particular attention to the hemoglobin family. This family has been studied in more detail than any other group of proteins, and as a result a deeper understanding of the relationships among structure, properties, and biological function (i.e., physiology) exists. The central focus of this chapter is to delineate chemical features that determine the affinity of an active site, especially an iron porphyrin, for molecular oxygen. In order to develop this theme, macroscopic (thermodynamic and kinetic) factors associated with dioxygen binding and release are summarized first. The nonbiological chemistry of iron and copper in the presence of dioxygen is described briefly to elucidate the key role that the protein plays in supporting oxygen transport by preventing irreversible oxidation of the binding site or of its ligands. The macroscopic behavior of the biological systems is related to the microscopic picture that has been developed over the last 30 years from x-ray crystallographic studies and a miscellany of spectroscopic probes of the oxygen-binding site. Relationships between the geometry and charge distribution in the metal-dioxygen moiety and the nature of the interactions between this moiety and its surroundings are examined. Nonbiological dioxygen carriers have proved particularly useful in providing precise and accurate structural information as well as thermodynamic and kinetic data against which the corresponding data from biological oxygen carriers can be contrasted.
The bioinorganic chemistry of the hemoglobin family of oxygen binders is particularly amenable to study by means of small-molecule model systems: four of the five ligands that make up the active site are provided by a square-planar tetradentate ligand, the protoporphyrin IX dianion (Figure 4.2). One axial ligand in hemoglobin, imidazole from a histidine residue, is provided by the protein, and the remaining sixth coordination site is available for the exogenous ligand, e.g., dioxygen or carbon monoxide. Thus a model system that approximates the stereochemistry of the active site in hemoglobin may be assembled from an iron(II) porphyrin and a ligand, such as imidazole or pyridine. On the other hand, in hemocyanin and hemerythrin most of the ligands are supplied by the protein. Thus the assembly of a model system that provides appropriate ligands correctly disposed around the pair of metal atoms poses a major synthetic challenge, especially for hemocyanin, where details on the number, type, and arrangement of ligands have been difficult to establish. Many aspects of the physical, inorganic, and structural chemistry underlying biological oxygen transport and utilization (Chapter 5) have been clarified through model systems. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.01%3A_Biological_Dioxygen_Transport_Systems.txt |
As noted earlier, three solutions to the problem of dioxygen transport have evolved: hemoglobin (Hb), hemocyanin (Hc), and hemerythrin (Hr). Their remarkable distribution over plant and animal kingdoms is shown in Figure 4.8.15
The hemoglobins and myoglobins found in plants, snails, and vertebrates all appear to share a common, very ancient ancestor. There is some evidence now for a common ancestral hemocyanin.42c The appearance of hemerythrin in a few annelid worms is an evolutionary curiosity. These few words and the diagram will suffice to give some hints about how respiratory proteins evolved, a subject that is outside the scope of this book.
The Hemoglobin Family
Hemoglobins are the most evolutionarily diverse family of dioxygen carriers. They are found in some plants (e.g., leghemoglobin in the nitrogen-fixing nodules of legumes), many invertebrates (including some insect larvae), crustaceans, molluscs (especially bivalves and snails), almost all annelid worms, and in all vertebrates with one possible exception, the Antarctic fish Cyclostomata.
With few exceptions the monomeric and oligomeric hemoglobins all share a basically similar building block: a single heme group is embedded in a folded polypeptide with a molecular weight of about 16 kDa (see Figure 4.2), and is anchored to the protein by coordination of the iron center to an imidazole ligand from a histidine residue. Mammalian myoglobin is often taken as the archetypical myoglobin (see Table 4.1). Sperm whale, bovine, or equine myoglobin are specific examples; the muscle tissue from which they may be extracted is more available than that from Homo sapiens. The archetypical oligomeric hemoglobin that shows cooperative binding of O2 is the tetrameric hemoglobin A. It is readily available from the blood of human donors.* In some invertebrate hemoglobins, especially those of annelids, aggregates may contain as many as 192 binding sites, to give a molecular weight of about 3 x 106 Dalton. These and other high-molecular-weight hemoglobins of arthropods are often referred to as erythrocruorins (Er). In a few annelid worms, the otherwise ubiquitous heme b or protoheme is replaced by chloroheme (see Figure 4.2) to give chlorocruorins (Ch), which tum green upon oxygenation (chloros, Greek for green). Some organisms, for example the clam Scapharca equivalvis, feature a dimeric hemoglobin.
The only known anomalous hemoglobin is Hb Ascaris, which comes from a parasitic nematode found in the guts of pigs. It has a molecular weight of about 39 kDa per heme; this value is not a multiple of the myoglobin building block.31 Moreover, presumably in response to the low availability of O2 in pigs' guts, Hb Ascaris has an extraordinarily high affinity for dioxygen, in large part owing to an extremely slow rate of dioxygen release.32 Leghemoglobin is another carrier with a high affinity for dioxygen, in this case because of a high rate of O2 binding. Since O2 is a poison for the nitrogenase enzyme, yet the nodules also require dioxygen, diffusion of O2 is facilitated, but the concentration of free dioxygen in the vicinity of nitrogen-fixing sites is minimized.33
Kinetic and thermodynamic data for dioxygen binding and release from a variety of hemoglobins are summarized in Table 4.2.9,10,31,34-36 Notice that for the hemoglobin tetramer, which comprises two pairs of slightly dissimilar subunits, the $\alpha$ and $\beta$ chains bind O2 with significantly different affinities and rate constants, especially in the T state. Isolated chains behave like monomeric vertebrate hemoglobins, such as whale myoglobin, which have affinities close to those of R-state hemoglobin. The chlorocruorins have a low affinity compared to other erythrocruorins. Especially for proteins that bind O2 cooperatively, a range of values is specified, since affinities and rates are sensitive to pH, ionic strength, specific anions and cations (allosteric effectors), and laboratory. For example, as we noted above, the O2 affinity of hemoglobin A is sensitive to the concentration of 2,3-DPG and to pH (Bohr effect). Trout hemoglobin I is insensitive to these species, whereas a second component of trout blood, trout hemoglobin IV, is so sensitive to pH (Root effect) that at pH < 7 trout hemoglobin IV is only partially saturated at P(O2) = 160 Torr.4 Note that O2 affinities span five orders of magnitude. Since heme catabolism produces carbon monoxide, and since in some environments CO is readily available exogenously, selected data for CO binding are also presented.
* Blood from human donors is also a source for a variety of abnormal hemoglobins, the most famous of which is HbS, the hemoglobin giving rise to sickle-cell anemia, It was Pauling and coworkers30 who first found that HbS differs from HbA through the single substitution of valine for glutamic acid in each of two of the four subunits comprising Hb, Sickle-cell anemia was the first condition to be denoted a "molecular disease."
The Hemocyanin Family
Hemocyanins (Hc), the copper-containing dioxygen carriers, are distributed erratically in two large phyla, Mollusca (for example, octopi and snails) and Arthropoda (for example, lobsters and scorpions). The functional form of hemocyanin consists of large assemblies of subunits.14,15,37 In the mollusc family the subunit has a molecular weight of about 50 kDa and contains two copper atoms. From electron-microscopic observations, hemocyanin molecules are cylindrical assemblies about 190 or 380 Å long and 350 Å in diameter comprising 10 or 20 subunits, respectively, for a molecular weight as high as 9 x 106 Dalton. In the arthropod family, the subunit has a molecular weight of about 70 kDa with two copper atoms. Molecular aggregates are composed of 6, 12, 24, or 48 subunits. Upon oxygenation the colorless protein becomes blue (hence cyanin from cyanos, Greek for blue). Spectral changes upon oxygenation, oxygen affinities, kinetics of oxygen binding (Table 4.2),4.5,14,15,38 anion binding, and other chemical reactions show that the active site in the phylum Arthropoda and that in Mollusca, although both containing a pair of copper atoms, are not identical.4,14
No monomeric hemocyanins, analogous to myoglobin and myohemerythrin (next section), are known. For some hemocyanins the binding of dioxygen is highly cooperative, if calcium or magnesium ions are present, with Hill coefficients as high as n ~ 9. However, the free energy of interaction per subunit can be small in comparison with that for tetrameric hemoglobin; 0.9 to 2.5 kcal/mol compared to 3.0 kcal/mol. Allosteric effects, at least for a 24-subunit tarantula hemocyanin, can be separated into those within a dodecamer (12 subunits)—the major contributor to overall allostery—and those between dodecamers.39c This has been termed nested allostery. In contrast to the hemoglobin family, isolated chains have affinities typical of the T-state conformation for hemocyanin. The binding of CO, which binds to only one copper atom, is at best weakly cooperative.39
As alluded to above, the distribution of hemocyanins is striking, Among the molluscs exclusive use of hemocyanin as the respiratory protein occurs only with the cephalopods (squid, octopi, and cuttlefish), and in the arthropods only among the decapod (ten-footed) crustaceans (lobsters, shrimp, and crabs). The bivalve molluscs (for example, oysters and scallops) all use small dimeric or octameric hemoglobins. The edible gastropod (snail) Helix pomatia uses hemocyanin, whereas the apparently closely related fresh-water snail Planorbis uses a high-oligomer hemoglobin. Both use a myoglobin as the oxygen-storage protein. The structure of the active site has been extensively probed by EXAFS methods,40,41 and the x-ray crystal structure of a hexameric deoxyhemocyanin is known.42 Each copper atom is coordinated to three imidazole groups from histidine residues. The pinwheel arrangement of the six subunits, the domain structure of a single subunit, and the domain containing the active site are shown in Figure 4.9.
The Hemerythrin Family
The biological occurrence of hemerythrins (Hr in Figure 4.8), the third class of dioxygen carriers, is relatively rare, being restricted to the sipunculid family (nonsegmented worms), a few members of the annelid (segmented worm) family, a couple of brachiopods (shrimps), and a couple of priapulids. The oxygen-binding site contains, like hemocyanin, a pair of metal atoms, in this case, iron. Upon oxygenation the colorless protein becomes purple-red. Monomeric (myohemerythrin), trimeric, and octameric forms of hemerythrin are known; all appear to be based on a similar subunit of about 13.5 kDa. When hemerythrin is extracted from the organism, its oxygen binding is at best only weakly cooperative, with Hill coefficients in the range 1.1 to 2.1.18 In coelomic cells (the tissue between the inner membrane lining the digestive tract and the outer membrane of the worm—analogous to flesh in vertebrates), oxygen apparently binds with higher cooperativity (n ~ 2.5).43 Perchlorate ions have been observed to induce cooperativity: since CIO4- has no biological role, it appears that in protein purifications the biological allosteric effector is lost. No Bohr effect occurs. Dioxygen binding data are accumulated in Table 4.2.36,44
The structure of hemerythrin in a variety of derivatives (oxy, azido, met, and deoxy) is now well-characterized. With three bridging ligands, a distinctive cofacial bioctahedral stereochemistry is seen (Figure 4.10).45-48
Table 4.2 - Thermodynamics and kinetics of ligand binding to biological oxygen carriers (at 20-25 °C and buffered at pH 6.5-8.5).
Solubility of O2 in water: 1.86 x 106 M/Torr
Solubility of CO in water: 1.36 x 10-6 M/Torr
a) 10 mM Ca2+ added: necessary for cooperativity
b) CO binding at pH 9.6.
Carrier
P1/2(O2)
Torr
$\Delta$H
kcal/mol
Dioxygen
$\Delta$S
eu
binding
kon
$\mu$M-1s-1
koff
s-1
P1/2(O2)
Torr
Carbon
$\Delta$H
kcal/mol
Monoxide
$\Delta$S
eu
Binding
kon
$\mu$M-1s-1
koff
s-1
Hemoglobins
Hb Ascaris 0.0047 1.5 0.0041 0.063 0.21 0.018
Leg Hb 0.047 -18.9 156. 1. 0.00074 13.5 0.012
whale Mb 0.51 -14.9 14. 12. 0.018 -13.5 0.51 0.019
Whale Mb 140. 1600.
HbA isolated chains - $\alpha$ 0.74 -142 -21. 50. 28. 0.0025 4.0 0.013
HbA isolated chains - $\beta$ 0.42 -16.9 -29 60. 16. 0.0016 4.5 0.008
HbA R $\alpha$ chain 0.15-1.5 -18. -30. 29. 10. 0.001-0.004 3.2 0.005
HbA R $\beta$ chain 0.15-1.5 -18. -30. 100. 21. 0.001-0.004 9.8 0.009
HbA R $\alpha$E7His→Gly 220. 620. 19. 0.007
HbA R $\beta$E7His→Gly 100. 3. 5.0 0.0013
HbA T $\alpha$ chain 9-160 -12 -35 2.9 183. 0.10-2.8 0.099 0.09
HbA T $\beta$ chain 9-160 -12 -35 11.8 2500. 0.10-2.8 0.099 0.09
Chironimus Mb 0.40 300. 218. 0.0019 27. 0.095
Glycera Mb 5.2 190. 1800. 0.00089 27. 0.042
Aplysia Mb 2.7 -13.6 15. 70. 0.013 0.49 0.02
Spriographis chlorocruorin 16-78 -4.5
Hemocyaninsa
Molluscan Hc
Helix pomatia R 2.7 -11.5 -12.6 3.8 10. 10. -13.5 -24 0.66 70.
Helix pomatia T 55. -15.4 -31.1 1.3 300. CO binding noncooperative since not measureable
Levantina hierosohimia R 3.8 -7.5 -1.8
Levantina hierosohimia T 18. +3.1 +31.
Arthropod Hc
Panulirus interruptus Rb 1.0 31. 60. 720.
-6.0
CO
-2.7
binding
4.1
noncooperative
8100.
P. interruptus monomer 9.3 57. 100.
Leirus quinquestris R 1.7 -7.4 0.
Leirus quinquestris T 117. +3.1 +27.
Hemerythrins
Phascolopsis gouldii 2.0 -12.4 -18 7.4 56 not known to bind CO
Themiste zostericola 8-mer 6.0 7.5 82
T. zostericola monomer 2.2 78. 315. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.02%3A_Biological_Oxygen_Carriers.txt |
With thermodynamic background and general structural features relevant to ligand affinity enumerated, attention may now be turned to the detailed structural aspects of the active site and its surroundings. As was shown crudely in Figure 4.3, the ligand affinity of an iron porphyrin may be perturbed either by modulating the structure of the deoxy material or by modulating the structure and surroundings of the liganded material or both. The model systems provide the reference points against which the protein structures may be compared.
Structures Relevant to Deoxy Hemoglobins
The structure of the picket-fence porphyrin compound, Fe(PF)(2-MeIm), is shown in Figure 4.28.172 Minus the pickets, it is essentially a magnified view of the active site of deoxymyoglobin, shown in Figure 4.29.181 Some metrical details of these structures, of a very similar unsubstituted tetraphenylporphyrin,110 and of several other deoxyhemoglobins11c,182-185 are listed in Table 4.7. In general they are all similar, but important differences exist.
Table 4.7 - Metrical detalis of deoxyhemoglobins and their modelsa
a) See Figure 4.25 for definition of symbols.
b) From a difference refinement of CoHb vs. Hb, where the difference in metal-to-porphyrin-plane separation was 0.24(2) Å and the difference in M-NIm was 0.13(4) Å. Doming is similar to Hb.
Compound
Resol.
(Å)
Fe-Np
(Å)
Fe • • • Porp
(Å)
Doming
(Å)
Fe—NIm
(Å)
$\phi$
(deg)
Tilt
(deg)
Fe(PF)(2-Melm) 2.072(5) 0.43 0.03 2.095(6) 22.8 9.6
Fe(TPP)(2-Melm) 2.086(6) 0.40 0.13 2.161(5) 7.4 10.3
Mb 1.4 2.03(10) 0.42 0.08 2.22 19 11
Er • • • H2O 1.4 2.02 0.17 -0.06 2.25 7 3
HbA ($\alpha$ • • • H2O)
1.74 2.08(3) 0.40(5) 0.16(6) 2.16(6) 18(1) 12(2)
HbA ($\beta$ • • • H2O) 1.74 2.05(3) 0.36(5) 0.10(6) 2.09(6) 24(1) 11(2)
CoHbb 2.5 0.14(5) 0.13 2.24(6)
Co(TPP)(1-MeIm) 1.977(6) 0.13 0.01 2.157(3) 3.8 0
Co(TPP)(1,2-Me2Im) 1.985(3) 0.15 0.05 2.216(2) 10
In all structures, except deoxyerythrocruorin,183 the iron atom is displaced about 0.4 to 0.5 Å from the plane of the porphyrin toward the axial base. For deoxyerythrocruorin the displacement is less than half this, perhaps because the water molecule is weakly coordinated to the iron center.
An imidazole group from a histidine residue—the distal histidine E7 in position 7 on helix labeled E—hovers over the binding site for most vertebrate hemoglobins, except for genetically engineered mutants of human hemoglobin ($\beta$E7His → Gly), pathological mutant hemoglobins, such as hemoglobin Zürich ($\beta$E7His → Arg), and some others, such as elephant hemoglobin. Long believed to be noncoordinating, this distal histidine may, in fact, coordinate weakly to the Fe center at low temperature.159 In the $\alpha$ chains of human deoxyhemoglobin, hemoglobin A, a water molecule is found in the binding cavity.182 For many years the binding cavity has been referred to as the hydrophobic pocket—literally, water-hating. Although many hydrophobic groups, such as valine, leucine, isoleucine, and phenylalanine are positioned over the porphyrin, the immediate environment around the binding site is, in fact, polar, with the distal histidine and associated water molecules, as well as the heme group itself. As will be shown in the next section, the label "hydrophobic pocket" becomes more misleading when the interaction of coordinated ligands with distal groups is examined.
The orientation of the axial base, angle $\phi_{1}$, is similar for Fe(PF)(2-MeIm) and for several vertebrate deoxyhemoglobins. On the other hand, Fe(TPP)(2- Melm) and deoxyerythrocruorin have a similar eclipsed axial-base orientation. At least for five-coordinate species, where the iron center is substantially out of the porphyrin plane, orientation of the axial base does not invariably induce structural perturbations, e.g., doming, in the porphyrin skeleton.
The conformation of the protein chain is such that the proximal histidine in deoxyhemoglobin coordinates in a slightly tilted manner,182,186 comparable to the tilt that the sterically active 2-methyl substituent induces in the synthetic systems.172 Clearly, coordination of the histidine to the heme in a symmetric manner, as would be expected in the absence of the protein constraints, does not produce the conformation of lowest free energy for the whole molecule. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.03%3A_Detailed_Structures_of_Hemoglobins_and_Model_Systems.txt |
To the student the subject of biological and synthetic molecular-oxygen carriers offers unusual insights into how bioinorganic chemistry works and what its aims and uses are. First, consider another bioinorganic problem, that of the nature of the blue copper proteins. When bioinorganic chemists entered the scene, the nature, function, and structure of copper blues were largely unknown. To take a Cu(lI) solution, add a nitrogen base, and obtain a spectrum that "resembled" that of the proteins was not a contribution to the solution of the biological puzzle, although it was the activity of some bioinorganic chemists. The difficulty here was that too little was known about the protein system. But the challenge did not diminish once the structure of a blue copper protein was known, for that structure allowed definition of the active site, one that contained (in the oxidized form) a Cu(lI) center surrounded by two imidazole, one cysteine, and one methionine residue. Now the bioinorganic chemist was faced with the formidable (and still incompletely solved) synthetic problem of designing a tetradentate ligand that (i) would present two N atoms, one thiolate S atom, and one thioether SR2 group to a Cu(lI) center and (ii) would remain intact if the Cu(lI) center were reduced to Cu(I). If we could prepare such complexes, we would be in a position to examine in some detail the effects on physical properties, such as redox potentials or spectra, of chemical substitution. In other words, we could learn about structure-function relationships in the copper blues. But the risks involved included the possibility that the specially designed ligand, even if it could be synthesized, might not bind Cu(Il) or Cu(l) in the desired manner.
Contrast such a situation with that of the oxygen carriers. Hemoglobin208 and myoglobin209 were the first crystallized proteins to have their structures determined. Their functions were well-known. They had been studied by a wide variety of physical techniques, in part because their structures were known, and even before that because of their role in human health. The central tetradentate ligand of the heme group, namely, the porphyrin, was well-defined and much porphyrin chemistry was known. The structural puzzles that intrigued chemists and biologists were not answered in the initial, early structural studies of the proteins; for example, how is O2 bound and why is CO not bound more firmly? Model chemistry in this area looked as though it would be easy; after all, the metalloporphyrins were readily synthesized, and all one needed was an axial base, some spectroscopic equipment, perhaps some single crystals, and then the structure-function relationships in the biological oxygen carriers would be understood! Indeed, as often happens, the situation was more complicated than it appeared. The irreversible oxidation of iron porphyrins was a major stumbling block to simple modeling. This obstacle was overcome in solution studies through the use of low temperatures and aprotic solvents; some very useful measurements of O2 and CO binding were made on model systems in such solutions. But in order to isolate oxygen complexes so that they could be studied by diffraction methods, another approach, that of synthesizing elaborated porphyrins, such as those in Figure 4.23, was necessary. This task entailed difficult organic chemistry that ultimately led to successful models that proved to be stable under ambient conditions. From such models we have learned much about local stereochemistry and, through spectroscopic congruence, about the biological systems. In short, bioinorganic chemistry has made a major contribution to the understanding of biological molecular-oxygen carriers, primarily because knowledge of the biological systems was advanced, the systems "self-assemble," and the goals of the studies were well-defined.
The complementarity of the two approaches continues. There are several unanswered questions, including:
1. What is the structural basis for cooperativity? Indeed, is there a structural basis at all, or is the ~6 kcallmol that represents the effect spread over many interactions, so that there is no obvious structural effect to be modeled?
2. Can one design a model where hydrogen bonding to the bound O2 molecule can be demonstrated by diffraction experiments? How will the oxygen uptake properties depend on the strength of the hydrogen bond?
3. Can one design a "high-affinity" model system? What will this tell us about the largely ill-defined high-affinity systems that are found in Nature?
There remain many intriguing questions about biological molecular-oxygen carriers, questions that will be answered by complementary studies on the biological and model systems. To make and study such model systems is an example of the challenge and excitement of this aspect of bioinorganic chemistry. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.04%3A_Dioxygen_Carriers_and_Bioinorganic_Chemistry.txt |
Many parallels exist between the chemistry of FeII- and CoII-porphyrinato systems. Dioxygen binds to many CoII complexes to give mononuclear 1:1 Co:O2 complexes with a bent geometry
$\tag{4.37}$
and dinuclear 2:1 Co:O2 complexes,64,66 analogous to those described for FeII systems in Reactions (4.29a) and (4.29b). Indeed, these dinuclear systems were the first nonbiological oxygen carriers to be isolated. The geometry of the dioxygen moiety, spanning two metals, may be cis or trans:
$\tag{4.38}$
However, whereas these dinuclear cobalt species are invariably octahedral, dinuclear copper-peroxo species are tetrahedral or distorted square pyramidal.40,41
In the late 1960s, 1:1 Co:O2 species were first isolated by use of a combination of low temperatures and specific Schiff-base ligands.104 It was found that cobalt corrins, such as vitamin B12, also formed 1:1 dioxygen adducts,105 although this chemistry is not known to be utilized by living systems.103 Cobalt(Il) porphyrins also form 1:1 adducts but with low O2 affinity, especially in nonpolar, aprotic solvents. Thus hemoglobin and myoglobin may be reconstituted from a cobaltoheme with preservation not only of dioxygen-binding capabilities but also of cooperativity.106 The synthetic 1:1 Co:O2 complexes have proven to be very useful in increasing our understanding of factors that determine oxygen affinity for cobalt systems and by extrapolation for iron systems. Two important differences make CoII systems more accessible. First, in contrast to iron systems, the cleavage reaction (4.29c) and redimerization to a $\mu$-oxo species (Reaction 4.29d) do not occur (see Figure 4.15). Thus CoII complexes of O2 are stable in solution at room temperature without the need for protection illustrated in Figure 4.14. Second, for CoII-porphyrinato systems, the equilibrium constant for the addition of a second axial base, such as pyridine or 1-methylimidazole, is small. Thus the disproportionation to four-coordinate and six-coordinate species that occurs for corresponding FeII systems (Reaction 4.33) does not occur. This difference simplifies the interpretation of spectral changes that are used to obtain thermodynamic and kinetic parameters of which there are now voluminous examples.66
Moreover, the 1:1 Co-O2 complexes are paramagnetic. From the small 59Co hyperfine splitting, it is deduced that the single unpaired electron resides primarily on the dioxygen moiety.104a,105 From other experiments107 it is apparent that net transfer of electron density from the metal onto the dioxygen varies considerably, from about 0.1e- to about 0.8e-. For example, it is found for a given CoIl Schiff base, Co(bzacen), that the redox potential of the cobalt-Schiffbase center LCo, measured by cyclic voltammetry, E1/2,
$B_{2}LCo^{III} + e^{-} \rightleftharpoons B_{2}LCo^{II} \tag{4.39}$
$B = substituted\; pyridine$
is a linear function of log K(O2) as the axial base B is varied. The more easily the CoIl center may be oxidized, the higher is the O2 affinity,103 as illustrated in Figure 4.18A. The dioxygen affinity also increases as the basicity of the axial nitrogenous ligand increases.104a This effect is illustrated in Figure 4.18B. Because of differing steric requirements, dimethylformamide (DMF), substituted imidazole, and piperidine (pip) ligands do not fall on the correlation defined by the series of substituted pyridine species. Note the synergistic nature of dioxygen binding: in general, the more electron density that is pumped onto the metal by the axial base, the more electron density is available for donation into the $\pi$* orbitals of the dioxygen ligand. E1/2 and log K(O2) are also correlated, although more weakly, for a number of hemoglobins (Figure 4.18C).108 Here the porphyrin and axial base remain constant, but presumably the surroundings of the heme group and O2 binding site vary in a manner that is less well-defined than in the model systems of Figure 4.18A and B. Notwithstanding these various perturbations to the metal center, the O—O stretch occurs at about 1140 cm-1, placing all 1:1 cobalt and iron-dioxygen complexes of nitrogenous and other hard ligands into the superoxo class.*
Cobalt(II) porphyrins and their adducts with diamagnetic molecules invariably have spin S = $\frac{1}{2}$. (See Figure 4.16, but add one electron.) Thus the structural changes are less pronounced than for corresponding iron(II) systems.110,111 From the similarities in geometries and differences in electronic structures between cobalt- substituted and native hemoglobins and their models, many insights have been gained about the factors that determine oxygen affinity as well as how cooperativity might, or might not, work at the molecular level.112,113 The mechanism of cooperativity has also been probed by the substitution of other metalloporphyrins into the globin: for example, zinc porphyrins have been used for their excited triplet-state properties,114 manganese porphyrins for their EPR activity,115 and ruthenium porphyrins as a member of the iron triad.116
* Because the O—O stretch may be coupled with other ligand modes,109 its value should not be used to estimate superoxo character, although in a series of $\mu$-superoxo and $\mu$-peroxo complexes of carefully controlled stereochemistry, small changes in v(O—O) have been correlated with the pKa of the suite of ligands.66
4.06: General Aspects of the Chemistry of Copper
The chemistry of copper in biological systems is limited to oxidation states I and II. The CuI state has electronic configuration d10. Unless there are ligand bands or strong ligand-to-copper charge-transfer bands, diamagnetic CuI species are colorless. Complexes of CuII (d9) are often blue in color. The single unpaired electron makes CuII amenable to electron paramagnetic resonance (EPR) techniques, at least if the electron spins of CuII centers are independent of one another. In oxyhemocyanin the spins are so strongly coupled (-J > 600 cm-1) that at room temperature and below the system is effectively diamagnetic and the pair of CuII ions is EPR silent.14
In aqueous solutions the CuI ion is unstable with respect to disproportionation to Cu metal and CuII ion:62
$Cu^{II} + 2e^{-} \rightleftharpoons Cu \qquad E^{o} = 0.3402 V$
$Cu^{+} + e^{-} \rightleftharpoons Cu \qquad E^{o} = 0.522 V \tag{4.35}$
$2 Cu^{+} \rightleftharpoons Cu + Cu^{II-} \qquad E^{o} = 0.182 V$
The CuI state may be stabilized by ligands, especially sulfur-containing ones, or by immobilization as afforded by a protein matrix, or in nonaqueous solvents, such as acetonitrile, in the absence of dioxygen. Whereas CuI thiolate species are stable, CuII thiolate species usually are unstable with respect to the disproportionation:101
$2 Cu^{II}—SR \rightarrow 2 Cu^{I} + R—S—S—R \tag{4.36}$
Again, immobilization may give kinetic stability to CuII thiolate species, as occurs in the blue-copper family of electron-transport proteins.
Copper(l) complexes are often two-coordinate with a linear arrangement of ligands. Three-, four-, and possibly five-coordinate complexes are known.
In the presence of O2, nonbiological copper(l) [and iron(II)] complexes are often susceptible to ligand degradation, which may give the illusion of O2 binding.102 The mechanisms by which this reaction occurs remain essentially unknown. Iron-porphyrin systems are rather more robust. Nonetheless, there are now several well-characterized copper(l) systems that reversibly bind dioxygen,15b,103 at least at low temperature. One that has been structurally characterized features a dicopper(II)-peroxo moiety,103f while a second, with more properties in common with oxyhemocyanin, features a moiety.103g | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.05%3A_General_Aspects_of_the_Chemistry_of_Cobalt.txt |
Redox Chemistry of Free Molecular Dioxygen
Dioxygen has a rich redox chemistry that is not explicitly exploited in the oxygen carriers, but which is central to enzymes that coordinate and activate dioxygen for subsequent reaction with a substrate. On reduction of dioxygen by one electron, the superoxide anion radical O2-• is formed. Concomitant with a reduction in bond order from 2.0 to 1.5 is an increase in bond length from 1.21 to 1. 30 Å. A second reduction step produces the peroxide anion 022 -; the bond order is one, and the O—O separation is 1.49 Å. Each of these reduced species, O2-• and O22-, has a characteristic O—O stretching vibration in the infrared region. The free-energy changes and electrochemical potentials for the reduction of dioxygen at unit activity, pH = 1 (E°), are different from those at pH 7.0 (E°'), as shown in Figure 4.11.58,62 The values at pH 7.0 are more relevant to physiological conditions. Note that the superoxide anion may function as either an oxidant or a reductant.
Geometry and Electronic Structure of Coordinated Dioxygen
In coordinating to metals, dioxygen shows a great variety of geometries and two formal oxidation states. Many complexes have v(O—O) values in the range 740 to 930 cm-1, and, where known, an O—O separation in the range 1.40 to 1.50 Å. By analogy with the peroxide anion, these species are designated peroxo, O2II-. Similarly, the designation superoxo O2I- is applied to those complexes where v(O—O) values are in the range 1075 to 1200 cm-1, and the O—O separation is around 1.30 Å.63 Although such O—O separations and vibrations are consistent with coordinated peroxide or superoxide moieties, the net amount of charge transferred onto the dioxygen ligand from the metal and its other ligands is difficult to measure experimentally and is probably variable. Thus the oxidation state of the dioxygen ligand and that of the metal are best considered in a formal sense rather than literally—hence the use of the terminology O2I- to indicate oxidation state I- for the O2 moiety as a unit (not each O atom). Because of the high degree of covalency in the M—O bond, a more sensible comparison, at least for the peroxo class of compounds, is with organic peroxides, ROOH or ROOR. The clear separation of coordinated dioxygen into either the superoxo or the peroxo class is shown in Figure 4.12.63-66 Only those compounds for which both stretching frequencies (v(O—O)) and O—O separations (r(O—O)) are available are shown; for the purpose of the plot, non-coordinated anions and cations, replacement of ethylenendiamine by two ammonia ligands, and replacement of triphenylphosphine by alkylphenylphosphines are assumed not to perturb significantly v(O—O) or r(O—O).
At least seven different geometries have been observed for the coordination of dioxygen (Figure 4.13),63-66 only three or four of which are currently known to be biologically relevant—the superoxo (for oxyhemoglobin), the peroxo or(for oxyhemocyanin), and the hydroperoxo (for oxyhemerythrin).
The geometry is a function of the metal, its oxidation state, and its ligands. For the late transition metals of the cobalt and nickel triads, with soft $\pi$-acid ligands, such as phosphines and carbonyls, and with an initially low oxidation state of the metal, triangular coordination of a peroxo species with covalent M-O2 bonds is common.63 Concomitant with the formal reduction of dioxygen, the metal center undergoes a formal two-electron oxidation:
$\tag{4.26}$
In this example, where the metal has undergone, at least formally, a two-electron oxidation, the UV-visible properties of the metal-dioxygen complex tend to resemble those of bona fide Mn+II rather than Mn species.
Early transition metals (Ti, V, Cr triads) often coordinate several peroxo species, leaving the metal in formally a very high oxidation state (e.g., Cr(O2)43-, a CrV ion).63 The M—O2 links have more ionic character, with the $\eta$-peroxo groups acting as bidentate ligands. Titanium and molybdenum(II) porphyrins bind, respectively, one and two dioxygen molecules in this manner.65
With harder $\sigma$-donor ligand systems, such as those containing nitrogen and oxygen donors, and the metal center in a normal oxidation state, a formal one-electron reduction to an end-on coordinated superoxo species occurs with a bent bond. Metal-dioxygen species can also be formed by adding the superoxide anion to the oxidized species:64
$\tag{4.27}$
In the absence of steric constraints, dimerization to a (bridging) $\mu$-peroxo species frequently occurs, especially for cobalt-dioxygen complexes:
These dicobalt species (right-hand side of Equation 4.28) may be oxidized by one electron to give a $\mu$-superoxo moiety. A clear shortening of the O—O bond and concomitant increase in the value of v(O—O) are observed in several superoxo-peroxo pairs. These and other modes of O2 attachment are illustrated in Figure 4.13. Some geometries are represented by only one or two examples, and some geometries, for example, a linear M—O—O species, have never been observed.
In binding to metals, O2 effectively functions both as a $\pi$ acid, accepting into its $\pi$* orbitals electron density from the filled d orbitals of the metal, and as a $\sigma$ donor, donating electron density into an empty metal d orbital. Thus other $\sigma$ donor or $\pi$ acceptor ligands, such as nitric oxide (NO), alkyl isocyanides (R—NC), alkyl nitroso (R—NO), and carbon monoxide (CO), are often observed to bind to the same metal complexes that bind O2. The nature of the metal-dioxygen linkage in biological oxygen carriers and their models will be examined in more detail later. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.07%3A_General_Aspects_of_the_Chemistry_of_Dioxygen.txt |
Irreversible Oxidation
In the presence of dioxygen, iron(II) species are readily oxidized to iron(III) species. In the presence of water, iron(III) species frequently associate into $\mu$-oxodiiron(III) dimers. For iron(II)-porphyrin complexes this process may take only milliseconds at room temperature. The following mechanism was proposed in 1968 for the irreversible oxidation of iron(II)-porphyrinato species;67,68 subsequent work has largely confirmed it.69-71
$Fe^{II} + O_{2} \rightleftharpoons Fe^{III}—O_{2}^{I-} \tag{4.29a}$
$Fe^{III}—O_{2}^{I-} + Fe^{II} \rightleftharpoons Fe^{III}—O_{2}^{II-}—Fe^{III} \tag{4.29b}$
$Fe^{III}—O_{2}^{II-}—Fe^{III} \rightarrow 2Fe^{IV}=O \tag{4.29c}$
$Fe^{IV}=O + Fe^{II} \rightarrow Fe^{III}—O—Fe^{III} \tag{4.29d}$
In particular, the dimerization reaction (4.29b) may be rendered less favorable by low temperatures (< -40 °C) or by sterically preventing the bimolecular contact of an FeIlI-O2I- moiety with an FeII moiety. In the latter case, sterically bulky substituents on the equatorial ligand surround the coordinated O2 ligand and the other axial position, trans to the coordinated dioxygen ligand, is protected with a nitrogenous base, such as imidazole, or with additional bulky substituents on the equatorial ligand (Figure 4.14).72 The protein effectively provides such protection and thus plays a key role in preventing the bimolecular contact of two hemes. The first observation of reversible binding of dioxygen to an iron(II)-porphyrin in the absence of protein was made in 1958.73 In that pioneering study, a heme group was immobilized on a polymer support specially modified to contain imidazole functions. The structurally characterizable hemoglobin or myoglobin species was replaced by a noncrystalline structurally uncharacterized polymer.
Why does this irreversible oxidation not occur analogously for cobalt systems? Step (4.29c) involves cleavage of the O—O bond, which in H2O2 has a bond energy of 34.3 kcal/mol or in Na2O2 of 48.4 kcal/mol. By way of comparison, for O2 the bond energy is 117.2 and for HO2• it is 55.5 kcal/mol.64 A simple molecular orbital picture gives insight into why an FeIV=O species is stabilized relative to the analogous CoIV=O species.74 From Figure 4.15 we see that for metals with electronic configuration dn, where n $\leq$ 5, no electrons occupy the antibonding orbital $\pi$* for FeIII-OI-• or FeIV=O moieties. For CoIII (d6) the extra electron goes into the antibonding orbital $\pi$*. As predicted by the model, MnIII is observed indeed to behave like FeIII.
A second oxidation pathway does not require the bimolecular contact of two iron(II)-porphyrins. Coordinated dioxygen may be released not as O2, as in normal dioxygen transport, but, as noted in Section I.C, as a superoxide radical anion O2-• in a process called autoxidation:
$Fe^{III}—O_{2}^{I-} \rightleftharpoons Fe^{III} + O_{2} \tag{4.30}$
This process is assisted by the presence of other nucleophiles that are stronger than the superoxide anion, such as chloride, and by protons that stabilize the O2-• anion as HO2:
$Fe_{III} + Cl^{-} \rightleftharpoons Fe^{III}—Cl^{-} \tag{4.31}$
The formation of methemoglobin occurs in vivo, probably by the above mechanism, at the rate of ~ 3 percent of total hemoglobin per day.
If exogenous reductants are present, then further reduction of dioxygen can occur:
$2H^{+} + Fe^{III}—O_{2}^{I-} + e^{-} \rightarrow Fe^{V}=O + H_{2}O \tag{4.32}$
Such processes are important, for example, in the cytochrome P-450 system. With suitably small reductants, oxygenase activity also has been observed for hemoglobin A. This has led to the characterization of hemoglobin as a "frustrated oxidase."75 Note the formal similarity between this process (Equation 4.32) and the bimolecular irreversible oxidation of iron(II) porphyrins: the second Fe(II) complex in Reaction (4.29b) functions like the electron in Reaction (4.32).
Spectroscopy of the FeIII—O—FeIIImoiety
The end products of the irreversible bimolecular oxidation of FeII species contain the FeIlI—O—FeIII fragment. Given the facile formation of $\mu$-oxodiiron(III) species, it is not surprising that the Fe—O—Fe motif is incorporated into a variety of metalloproteins, including the oxygen-carrier hemerythrin (Figure 4.10),16-18 the hydrolase purple acid phosphatase,76 the oxidoreductases ribonucleotide reductase77 and methane monoxygenase,78 an iron-sulfur protein rubrerythrin,79a and the iron-transport protein ferritin.79b In ferritin higher-order oligomers are formed.
This $\mu$-oxodiiron(III) moiety has a distinctive fingerprint that has made it easy to identify this motif in proteins.80 Regardless of the number (4, 5, 6, or 7), geometry (tetrahedral, square pyramidal, tetragonally distorted octahedral, or pentagonal bipyramidal), and type of ligands (halide, RO-, RCOO-, aliphatic N, or aromatic N) around the iron center, and of the Fe—O—Fe angle, the magnetic susceptibility at room temperature lies in the range 1.5 to 2.0 Bohr magnetons per Felll—O—Fe lll group, equivalent to about one unpaired electron. 81,82 In other words, the high-spin (S = $\frac{5}{2}$) iron centers are strongly antiferromagnetically coupled. Other bridging groups, such as OH-, Cl-, carboxylate, alkoxide, or phenoxide, give very weak coupling.83-86
The asymmetric Fe—O stretch, vas(Fe—O), lies in the range 730 to 880 cm-1; in multiply bridged complexes this mode is weak in the infrared region. The symmetric vibration, vs(Fe—O), forbidden in the infrared region for linear, symmetric Fe—O—Fe groups, occurs in the range 360 to 545 cm-1. The symmetric mode is usually,87a but not always,87b observed by resonance Raman techniques upon irradiating on the low-energy side of the Fe—O chargetransfer band that occurs at about 350 nm.
Few dinuclear iron(II) complexes are known where the ligands approximately resemble those believed or known to occur in the family of $\mu$-oxodiiron(III) proteins.88 The dioxygen-binding process in hemerythrin has no close nonbiological analogue. Although spectroscopically similar to oxyhemerythrin, the unstable monomeric purple peroxo complex formed by the addition of hydrogen peroxide to basic aqueous FeIII(EDTA) solutions remains structurally uncharacterized.89,90
Oxidation and Spin State of Iron Porphyrins
Iron porphyrins, the active sites of the hemoglobin family, have a rich magnetochemistry.91 Iron porphyrins may be octahedral (two axial ligands), square pyramidal (one axial ligand), or square planar (no axial ligand). The metal d orbitals, now having partial porphyrin $\pi$* character, are split, as shown in Figure 4.16. The radius of the metal atom is much greater when it is high spin (S = 2 for FeII, S = $\frac{5}{2}$ for FeIII) than when it is low spin (S = 0 for FeII, S = $\frac{1}{2}$ for FeIII). This difference influences Fe—Nporph separations, porphyrin conformation, and the displacement of the iron center with respect to the porphyrin plane. For iron(II)-porphyrins, two strong-field axial ligands, such as a pair of imidazoles or an imidazole and carbon monoxide, lead to diamagnetic complexes (S = 0) with the six 3d electrons occupying those orbitals of approximate t2g symmetry. In a classic experiment in 1936, Pauling and Coryell proved that oxyhemoglobin and carbonmonoxyhemoglobin are diamagnetic.92*
* There was a considerable flurry of interest when an Italian group, using a SQUID (Superconducting Quantum Mechanical Interference Device), reported that at room temperature oxyhemoglobin was significantly paramagnetic.93 Not surprisingly, several theoretical papers followed that "proved" the existence of low-lying triplet and quintet excited states94-96 Subsequently, the residual paramagnetism was doubted97 and shown to arise from incomplete saturation of hemoglobin by O2; in other words, small amounts of deoxy hemoglobin remained98 Since oxygen affinity increases with decreased temperature, the concentration of paramagnetic impurity decreased with decreasing temperature.
No axial ligands at all may lead to a spin state of S = 1, with unpaired electrons in the dxy and dz2 orbitals. Five-coordinate iron(II)-porphyrinato complexes are commonly high spin, S = 2, although strong $\sigma$-donor $\pi$-acceptor ligands, such as phosphines, carbon monoxide, nitric oxide, and benzyl isocyanide,99 enforce a low-spin state. Five-coordinate iron(II)-porphyrinato complexes with aromatic nitrogenous axial ligands, such as pyridine or 1-methylimidazole, bind a second such axial ligand 10 to 30 times more avidly than the first to give the thermodynamically and kinetically (d6, S = 0) stable hemochrome species, a process that is avoided by hemoglobins. That is, the equilibrium constant for the following disproportionation reaction is greater than unity,
$Fe—N + Fe—N \rightleftharpoons N—Fe—N +Fe \tag{4.33}$
except for bulky ligand N, such as 2-methylimidazole and 1,2-dimethylimidazole, for which the five-coordinate species predominates at room temperature even with a mild excess of ligand:100
$\tag{4.34}$
For iron(III)-porphyrinato complexes, strong-field ligands lead to low-spin (S = $\frac{1}{2}$) complexes. A pair of identical weak-field ligands, such as tetrahydrofuran, leads to intermediate-spin (S = $\frac{3}{2}$) species. Five-coordinate species are, with few exceptions, high-spin (S = $\frac{5}{2}$), with all five 3d electrons in separate orbitals. Spin equilibria S = $\frac{1}{2} \rightleftharpoons$ S = $\frac{5}{2}$ and S = $\frac{3}{2} \rightleftharpoons$ S = $\frac{5}{2}$ are not unusual. Specific examples of these spin systems are given in Table 4.4.65,91 Higher oxidation states are found in some other hemoproteins. Fe(V)-porphyrin systems actually occur as Fe(IV)-porphyrin cation radical species, and Fe(I)-porphyrin systems exist as Fe(II)-porphyrin anion radical species.
Substantial structural changes occur upon the addition of ligands and upon changes in spin state. In one mechanism of cooperativity these changes are the "trigger" (metrical details are deferred until the next section). Spectral changes in the UV-visible region are observed also (Figure 4.17)10 and may be monitored conveniently to evaluate the kinetic and thermodynamic parameters of ligand binding to hemoglobin.
Table 4.4 - Oxidation and spin states of iron porphyrins and their biological occurrences
a) Could be placed in FeIII column
b) Could be placed in FeIII column with spin = 0.
c) Non-linear Fe—NCS moiety. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.08%3A_General_Aspects_of_the_Chemistry_of_Iron.txt |
There are many ways in which ligand affinity may be perturbed (Figure 4.25). It is convenient to divide these into two groups, referred to as distal and proximal effects.163 Proximal effects are associated with the stereochemistry of the metalloporphyrinato moiety and the coordination of the axial base, and thus their influence on O2 and CO affinity is indirect. Distal effects pertain to noncovalent interactions of the metal-porphyrinato skeleton and the sixth ligand (O2, CO, etc.) with neighboring solvent molecules, with substituents, such as pickets or caps, on the porphyrin, and with the surrounding protein chain. The distal groups that hover over the O2-binding site engender the most important distal effects. For convenience, the effects of crystal packing and the protein matrix on porphyrin conformation will be discussed among the proximal effects, although as nonbonded interactions they properly are distal effects.
To a first approximation, the effects of substituents on the porphyrin ring, as transmitted through bonds to the metal center, do not perturb the ligand binding properties as much as do distal effects.170 Thus substituents, such as vinyl and propionic-acid groups on protoporphyrin IX and o-pivalamidophenyl pickets, are ignored; one porphyrin is much like another. At the end of this subsection the various ways ligand affinity may be modulated will be summarized in an augmented version of Figure 4.3.
Proximal Effects
Few molecules have had their conformational properties characterized as exhaustively as have metalloporphyrins.
a. Porphyrin Conformation and M-Np Separations
The cyclic aromatic 24-atom porphyrinato skeleton offers a tightly constrained metal-binding site. The conformation of least strain is planar, and the radius of the hole of the dianion is close to 2.00 Å,110 leading to metal-porphyrinato nitrogen-atom separations, M-Np, of 2.00 Å if the metal is centered in the square plane defined by the four porphyrinato nitrogen atoms. Small deviations from planarity are generally observed and attributed to crystal packing effects; large deviations may be induced by bulky substituents on the porphyrin skeleton, especially at the meso positions, by the crystal matrix,65 or by the highly anisotropic protein matrix. The 2.00 Å radius hole neatly accommodates low-spin (S = 0) and intermediatespin (S = 1) iron(II), low-spin (S = $\frac{1}{2}$) iron(III), and cobalt(II) and cobalt(III) ions.91* With few exceptions the metal is centered in or above the central hole for mononuclear porphyrin species; only rarely do M-Np bonds show a significant (though still small) scatter about their mean value.
* In order to accommodate smaller ions, such as nickel(II), the porphyrin skeleton may contract by ruffling, with little loss of aromaticity; like a pleated skirt the pyrrole rings rotate alternately clockwise and counterclockwise about their respective M-Np vectors. This distortion leaves the four porphyrinato nitrogen atoms, Np, still coplanar. Alternatively, the porphyrin skeleton may buckle to give a saddle conformation; the Np atoms may acquire a small tetrahedral distortion in this process. M-Np bonds as short as 1.92 Å have been observed. Metals with one or two electrons in their 3dx2-y2 orbital have a radius larger than 2.00 Å. In order to accommodate them in the plane of the porphyrin, the porphyrin skeleton expands. M-Np separations as long as 2.07 Å may occur with the metal still centered in the plane of the Np atoms.110
b. M • • • Porph Displacement
For five-coordinate complexes the magnitude of the displacement of the metal from the plane of the four nitrogen atoms, M • • • porph, is a consequence of the electronic configuration of ML5 complexes. Of course, the effect is augmented if the 3dx2-y2 orbital (directed along M-Np bonds, Figure 4.16) is occupied. Compare a displacement of 0.14 Å for Co(TPP)(1,2-Me2Im) (no 3dx2-y2 occupancy)111 with 0.43 Å for Fe(PF)(2-MeIm) (3dx2-y2 occupied).169 For six-coordinate complexes where the two axial ligands, L1 and L2, are different, the M • • • porph displacement usually reflects relative trans influences.
Generally, displacement of the metal from the plane of the porphyrinatonitrogen atoms is within 0.04 Å of the displacement from the 24-atom mean plane of the entire porphyrin skeleton. On occasions this second displacement may be much larger, for example in Fe(TPP)(2-MeIm), where it is 0.15 Å larger110 than it is for Fe(PF)(2-MeIm). This effect is called doming, and it is usually attributed to crystal packing forces. Interaction of the porphyrin with protein side chains leads to considerable doming or folding of the heme in vertebrate hemoglobins.
c. M-L Separations
The metal-axial ligand separations, M-L (when more than one, L1 denotes the heterocyclic axial base), are dependent on the nature of the ligand, L. When L1 and L2 are different, the M-L separations are sensitive to the relative trans influences of L1 and L2 as well as to steric factors. For example, for Fe(TPP)(1-Melm)2, the Fe—NIm bond length is 2.016(5) Å,110 whereas for Fe(TPP)(1-Melm)(NO) it is 2.180(4) Å.111For sterically active ligands, such as 2-methylimidazole compared to 1-methylimidazole (4.34), the longer Co—NIm bond occurs for the 2-Melm ligand because of steric clash between the 2-methyl group and the porphyrin.111It is possible that combinations of intrinsic bonding and steric factors may give rise to a double minimum and two accessible axial ligand conformations (Figure 4.26). This situation seems to occur in the solid state for Fe(PF)(2-Melm)(O2)•EtOH, where a short Fe—NIm and a long Fe—O bond are observed both from the structure revealed by single-crystal x-ray diffraction methods and by EXAFS data. On the other hand, for solvate-free Fe(PF)(2-MeIm)(O2) and for Fe(PF)(1,2-Me2Im)(O2), the EXAFS patterns are interpreted in terms of a short Fe—O and long Fe—Im bond.171
d. The Angle $\phi$
This parameter is the minimum angle that the plane of the axial base (e.g., pyridine, substituted imidazole, etc.) makes with a plane defined by the Np, M, and L1 atoms (Figure 4.25).65 If there are two axial ligands, e.g., 1-methylimidazole and O2, then, as before, the angle the axial base makes is denoted $\phi_{1}$ and the other angle $\phi_{2}$. For a linear CO ligand bound perpendicularly to the porphyrin plane, $\phi_{2}$ is undefined. Note that the orientation of the second ligand is influenced by distal effects.
When $\phi$ = 0, the axial base eclipses a pair of M-Np bonds; contacts with the porphyrin are maximized. When $\phi$ = 45°, contacts are minimized. Unless the axial base has a 2-substituent, however, the contacts are not excessively close for any value of $\phi$. With a 2-methyl substituent, the contacts are sufficiently severe that the M-NL1 vector is no longer perpendicular to the porphyrin plane, and the imidazole group is rotated so that the M-NL1 vector no longer approximately bisects the imidazole C—N—C bond angle, as illustrated Figure 4.25.110,172
Distal Effects
Distal effects arise from noncovalent interactions of the coordinated dioxygen, carbon monoxide, or other ligand with its surroundings. The protein matrix, the pickets, and the caps are functionally equivalent to an anisotropic solvent matrix that contains a variety of solutes. The limits of this simplification are illustrated in the following example. The electronically similar cobalt meso-, deutero-, and protoporphyrin IX complexes bind dioxygen with similar affinities under identical solvent conditions. When they are embedded in globin, larger differences in affinity and changes in cooperativity are observed.170 These effects are attributed to the slightly different nestling of the porphyrin molecules in the cleft in hemoglobin or, in the generalization introduced, to slightly different solvation effects.
Interaction of the coordinated O2 or CO molecule with solvent molecules or with the protein has a profound influence on kinetics and thermodynamics (see Figure 4.24, and Tables 4.2 and 4.5). As discussed earlier, there is accumulation of negative charge on the dioxygen ligand. The possibility then arises for stabilization of coordination through hydrogen bonding or dipolar interactions with solute molecules,175 porphyrin substituents (such as amide groups in the picket-fence porphyrins176 and some species of strapped porphyrins161), or with protein residues* (such as histidine).167,177-179
Destabilization of coordinated ligands and lowered affinity can result if the coordinated ligand is unable, through steric clash, to achieve its optimum stereochemistry or if the closest neighboring groups are electronegative, as are the ether and ester linkages on capped porphyrins.31,180 We will describe in detail in the next subsection (III.C) the fascinating variety of means by which ligand binding is modulated by distal amino-acid residues.
* For Glycera CoMbO2 no change in EPR parameters occurs on substituting D2O for H2O.168 No hydrogen bond between O2 and a distal group comparable in strength to that in whale CoMbO2 was inferred.
Approximate Contribution of Proximal and Distal Effects to Ligand Affinity
Dissimilar systems may show similar affinities for a ligand as a result of a different mix of the proximal and distal effects enumerated above. These effects are not all of equal magnitude, and an attempt is made here to show the increment in free energy that occurs if the effect is manifest in the deoxy or liganded state of Figure 4.3. Increasing the free energy of the deoxy state while holding that of the liganded state constant leads to an increase in affinity. The reference state is gaseous Fe(TPP)(1-MeIm). The magnitude and sign of these effects are shown in Figure 4.27. For the coordination of alkylisocyanide molecules to hemoglobin, the steric effects of different alkyl groups have been quantified.35 Lowered affinity occurs with increasing alkyl chain length, with the exception of methyl isocyanide. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.09%3A_General_Structural_Features_that_Modulate_Ligand_Activity.txt |
The binding of dioxygen is normally a reversible process:
$M + O_{2} \rightleftharpoons MO_{2} \tag{4.22}$
Under some circumstances, such as in the presence of added nucleophiles and protons, coordinated dioxygen is displaced as the superoxide anion radical, O2-, leaving the metal center oxidized by one electron and unreactive to dioxygen:49,50
$MO_{2} \rightleftharpoons M^{+} + O_{2}^{-} \tag{4.23}$
For hemoglobin there exists a flavoprotein reductase system, comprising a reduced pyridine nucleotide (e.g., NADH), cytochrome b5 reductase, and cytochrome b5 , that reduces the ferric iron back to the ferrous state, so that it may coordinate dioxygen again.1,51 In addition, all aerobically respiring organisms and many air-tolerant anaerobes contain a protein, superoxide dismutase, that very efficiently catalyzes the dismutation of superoxide ion to dioxygen and hydrogen peroxide:52-54
$2O_{2}^{-} + 2H^{+} \rightarrow O_{2} + H_{2}O_{2} \tag{4.24}$
However, the physiological effects of the superoxide moiety remain controversial.53,54 Finally, there is a third enzyme, the hemoprotein catalase, that converts the toxic hydrogen peroxide into water and dioxygen:1
$2H_{2}O_{2} \rightarrow O_{2} + 2H_{2}O \tag{4.25}$
This topic is discussed further in Chapter 5.
4.11: Nature of the Metal-Dioxygen Linkage in Biological Systems
Many techniques have been used to probe the geometry and electronic structure of the metal-dioxygen moiety in biological systems and in synthetic models. The results form the basis of any understanding of the factors that determine and modulate oxygen affinity.
Oxyhemocyanin and Oxyhemerythrin
By resonance Raman techniques, the O—O stretch is observed at 744 cm-1 in oxyhemocycanin and at 844 cm-1 in oxyhemerythrin.140-142 Dioxygen is therefore coordinated as peroxo species. By use of unsymmetrically labeled dioxygen, 18O—16O, it was established that dioxygen coordinates symmetrically:141
$\tag{4.43}$
Carbon monoxide binds to only one of the CuI centers in deoxyhemocyanin, through the C atom,* apparently blocking the second CuI site.15 Similar behavior is also seen for the nitrosyl adduct.
* A report143 that CO binds to the copper center through the O atom, an unprecedented mode, has been challenged.144
On the other hand, in an experiment parallel to that just described for hemocyanin, dioxygen was found to bind asymmetrically in oxyhemerythrin:142
$\tag{4.45}$
The existence and location of the proton, which cannot be proven in the crystal structure of deoxyhemerythrin or of oxyhemerythrin, are inferred from a model system for the former that contains a hydroxo group bridging two high-spin, weakly antiferromagnetically coupled Fell centers (J = -10 cm-1).88 For oxyhemerythrin (and for one conformation of methydroxohemerythrin), a small change in the position of the symmetric Fe—O—Fe mode is observed when H2O is replaced by D2O.146 The strong antiferromagnetic coupling observed for methemerythrin and oxyhemerythrin (-J ~ 100 cm-1)147 is uniquely consistent with a bridging oxo moiety between a pair of FeIII centers.80 Finally, a Bohr effect (release or uptake of protons) is absent in oxygen binding to hemerythrin.16,18 These observations are consistent with a $\mu$-oxo group slightly perturbed by hydrogen bonding to a coordinated hydroperoxo species. An important role for the protein in hemerythrin is to assemble an asymmetric diiron(II) species. Only a few of the myriad of known $\mu$-oxodiiron(III) complexes are asymmetric,81 and the synthesis of realistic asymmetric models remains a challenge.
Deoxyhemocyanin (CuI d10) and deoxyhemerythrin (Fell d6) are colorless. In the oxygenated derivatives there is considerable charge transfer between the coordinated peroxo groups and the metal centers. This phenomenon makes the essentially d-d metal transitions more intense than those for the simple aquated Fe3+ or Cu2+ ions, and permits facile measurement of oxygen-binding curves. The spectral changes accompanying oxygenation are shown in Figures 4.20 and 4.21.5
Nitric oxide binds to deoxyhemocyanin, to deoxyhemerythrin, and to the mixed valence FeIII • • • Fell semimethemerythrin.148 Carbon monoxide binds to neither form of hemerythrin: apparently the other ligands have insufficiently strong fields to stabilize the low-spin state for which electron density would be available for back donation into the CO $\pi$* orbitals.
Oxyhemoglobin
The O—O stretch that is observed by difference infrared techniques at around 1105 cm-1 for oxyhemoglobin and oxymyoglobin149 clearly categorizes the dioxygen moiety as a superoxo species; that is, the order of the O—O bond is about 1.5. Considerable ink has been spilled about the nature of the Fe—O2 fragment since Pauling's original suggestion150 in 1948 that dioxygen binds to iron in an end-on bent fashion:
$\tag{4.46}$
He subsequently reaffirmed this geometry, and proposed that hydrogen bonding between the coordinated dioxygen and the distal imidazole H—N group was important in stabilizing the Fe—O2 species.137 In an alternative model Weiss proposed that a low-spin FeIII center (S = $\frac{1}{2}$) was very strongly antiferromagnetically coupled to a superoxide anion radical (S = $\frac{1}{2}$).151 A triangular peroxo mode has also been advanced.152,153 The problem has been how to resolve the observed diamagnetism of oxyhemoglobin92,98 with UV-visible, x-ray absorption, and resonance Raman spectroscopic characteristics154 that are distinctly different from those of Fell systems (such as carbonmonoxyhemoglobin and low-spin six-coordinated hemochromes, such as Fe(Porph)(Py)2) and from unambiguously FeIII systems (such as chloromethemoglobin or cyanomethemoglobin).
Any adequate theoretical treatment must also explain how iron-porphyrin systems can bind not only O2, but also CO, NO, alkyl isocyanides, and alkyl-nitroso moieties. A simple qualitative model presented by Wayland and coworkers129,155 conveniently summarizes ligand-binding geometries of cobalt and iron porphyrins. Although a reasonable quantitative theoretical consensus exists for 1:1 cobalt-dioxygen species, the same cannot be said yet for irondioxygen systems.
A Simple Model for the Electronic Structure of Liganded Hemoglobins
Why does dioxygen bind to iron and cobalt porphyrins in an end-on bentbond fashion as in (4.37) and (4.46)? Why does carbon monoxide bind in a linear manner (Equation 4.40)? Why are six-coordinate dioxygen and carbonmonoxide adducts more stable than five-coordinate ones? A unified picture of ligand binding that addresses these questions is important in understanding properly the specific case of dioxygen binding to hemoglobin and related systems.
The splitting of the metal d orbitals for a four-coordinate metalloporphyrin is shown in the center of Figure 4.22. These orbitals contain some porphyrin character and are antibonding with respect to metal-porphyrin bonds. As shown in Figure 4.16, the primary effect of a single $\sigma$-donor axial ligand, such as pyridine or 1-methylimidazole, is to elevate the energy of the antibonding dz2 and lower the energy of the dx2-y2 orbital and hence lead to a high-spin species in place of the intermediate-spin four-coordinate one. Thus, for simplicity in highlighting interaction of the metal center with the diatomic $\sigma$-donor: $\pi$-acid ligands CO, NO, and O2, the perturbations wrought by primarily $\sigma$-donor ligands, such as 1-methylimidazole, are omitted. For the corresponding cobalt(II) compound, there is an additional electron. The diatomic ligands of interest share a qualitatively similar molecular orbital scheme. The filling of orbitals for CO is shown on the left-hand side. Dioxygen, which is shown on the right-hand side, has two more electrons than CO; these occupy the doubly degenerate $\pi$* orbitals. Quantitative calculations show that the energy of the $\pi$* orbitals decreases monotonically from CO to NO to O2, indicating increasing ease of reduction of the coordinated molecule, a feature that has not been included in the diagram. Only those interactions of molecular orbitals that have appropriate symmetry and energy to interact significantly with the metal d orbitals are shown.
Two extremes are shown in Figure 4.22 for the interaction of a diatomic molecule A-B with the metal center: a linear geometry on the left and a bent geometry on the right. A side-on geometry is omitted for the binding of O2 to a CoIl or Fell porphyrin, since this would lead to either an MIll side-on superoxo or an MIV peroxo species; both these modes of coordination to these metals are currently without precedent.
Linear diatomic metal bonding maximizes the metal-d$\pi$ to ligand-p$\pi$* bonding. When a ligand coordinates in a bent manner, axial symmetry is destroyed, and the degeneracy of the ligand p$\pi$* orbitals is lifted. One p$\pi$* orbital is now oriented to combine with the metal dz2 orbital to form a $\sigma$ bond, and the other is oriented to combine with dxz and dyz orbitals to form a $\pi$ bond. A bent geometry for the diatomic molecule will result when either or both of the metal dz2 or the ligand p$\pi$* orbitals are occupied, since this geometry stabilizes the occupied dz2 orbital in the five-coordinate complex. Thus O2 binds in a strongly bent manner to CoIl and Fell porphyrins; NO binds in a strongly bent manner to CoIl porphyrins; CO binds in a linear fashion to Fell porphyrins.
The interaction of NO with Fell porphyrins and CO with CoIl porphyrins—the resultant species are formally isoelectronic—is more complicated. The degree of bending seen in Fell(TPP)(NO) is midway between the two extremes.111For CO the higher-energy p$\pi$* orbitals lead to a greater mismatch in energy between the dz2 and p$\pi$* orbitals, and less effective $\sigma$ bonding. In EPR experiments the odd electron is found to be localized in a molecular orbital with about 0.87 metal dz2 character for the five-coordinate Co—CO adduct, as expected for a nearly linear geometry.129 On the other hand, for the Fe—NO adduct the metal dz2 character of the odd electron is about 0.4 to 0.5;155 a somewhat bent geometry (140°) is observed in the crystal structure of Fe(TPP)(NO). Because the CO ligand is a very weak $\sigma$ donor, the Co—CO species exists only at low temperatures.
Only qualitative deductions can be made from this model about the extent of electron transfer, if any, from the metal onto the diatomic ligand, especially for dioxygen. The higher in energy the metal dz2 orbital is with respect to the dioxygen p$\pi$* orbitals, the closer the superoxo ligand comes to being effectively a coordinated superoxide anion. With an additional electron, the dioxygen ligand in Co—O2 complexes can acquire greater electron density than it can in Fe—O2 complexes.
From the diagram it may be inferred that a ligand with very strong $\pi$-acceptor properties will lower the energy of the dxz and dyz orbitals through strong (dxz, dyz)-$\pi$* interaction. The resultant energy gap between these two orbitals and the other three metal d orbitals may be sufficient to overcome the energy involved in spin-pairing, and hence lead to five-coordinate low-spin species, as happens for complexes containing phosphines and carbon monosulfide.99,132 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.10%3A_Hazards_of_Life_with_Dioxygen.txt |
As noted above, a variety of other $\alpha$-donor or $\pi$-acceptor ligands will bind to the active sites of biological oxygen carriers.
Carbon Monoxide
As documented in Table 4.2, carbon monoxide (CO) generally binds more strongly to hemoglobin than does dioxygen, hence causing carbon-monoxide poisoning. In addition to being readily available from car exhausts and tobacco smoke to convert oxyhemoglobin to carbonmonoxyhemoglobin, CO is produced in the catabolism of heme molecules.117 Thus under even the most favorable of conditions, about 3 percent of human hemoglobin is in the carbonmonoxy form. When CO binds to a single metal atom in nonbiological systems, without exception it does so through the carbon atom and in a linear manner:56
$Fe—C=O \leftarrow Fe^{II} +\;^{-}C=O^{+} \tag{4.40}$
Model systems for carbonmonoxy (also called carbonyl) hemoglobin show a geometry similar to that of the Fe—C$\equiv$O group, linear or nearly so and essentially perpendicular to the porphyrin plane.110,118-121 The biochemical literature is littered with reports that this is not the geometry adopted by CO in binding to hemoglobins.122-128 We will return to this topic later in this chapter, since the physiological consequences are potentially important.
Carbon monoxide binds weakly as a $\sigma$-donor ligand to four-coordinate cobalt(II) systems.129 Despite a bout of artifactual excitement,130 CO has never been observed to bind significantly to five-coordinate CoIl systems with a nitrogenous axial base to yield octahedral six-coordinate species.131 The sulfur analogue thiocarbonyl (CS), although not stable as a free entity, binds very strongly to iron-porphyrin species in a linear manner.132
Nitric Oxide
Nitric oxide (NO) binds to hemes even more strongly than CO (and hence O2),10 so strongly, in fact, that the Fe—N1m bond is very weak and easily ruptured.11,111,133 Attachment to the metal is via the nitrogen atom; however, the geometry of attachment is sensitive to the $\pi$ basicity of the metalloporphyrin, and ranges from linear to strongly bent. In binding to ColI the NO ligand is effectively reduced to NO-, with concomitant oxidation of CoIl to CoIII:111
$\tag{4.42}$
In much the same way that cobalt-dioxygen systems are paramagnetic (S = $\frac{1}{2}$) and amenable to EPR studies, iron-nitric oxide (also called iron nitrosyl) species are also paramagnetic and isoelectronic with cobalt-dioxygen species. The unpaired spin is localized mostly on the NO group.
Isocyanide and Nitroso Species
In contrast to the dioxygen, carbon-monoxide, and nitric-oxide ligands, the isocyanide and nitroso functions bear an organic tail. Moreover, nitroso ligands are isoelectronic with dioxygen.
$\tag{4.42}$
Thus, in principle, not only may the steric bulk of the ligand be varied, in order to probe the dimensions35 of the dioxygen-binding pocket,* but also the $\sigma$-donor/$\pi$-acceptor properties of the ligands may be varied by appropriate substituents on the aryl ring.
Isocyanide groups may bind to metals in a variety of ways. For 1:1 adducts (Figure 4.19), the isocyanide group is approximately linear, although some flexibility seems to exist in a bis(t-butylisocyanide)iron(Il)tetraphenylporphyrinato complex.135 For zerovalent metals with much electron density available for donation into ligand $\pi$* orbitals, the isocyanide ligand has been observed to bend at the N atom.136 One prediction exists that an isocyanide ligand binds in this manner to hemoglobin.137
For 1:1 adducts of nitroso ligands, side-on,138 O-, and N-ligated modes are possible (Figure 4.19). No O-nitroso complexes have been definitively characterized by diffraction methods. For hemoglobin the N-nitroso mode is likely, since this is the mode found for the nitrosoalkane in Fe(TPP)(amine)(RNO).139
To date isocyanide ligands have not achieved their potential as probes of the geometry of the ligand-binding pocket in hemoglobin, partly because we lack structural data on the preferred geometry of attachment of these ligands in a sterically uncongested environment.
* In fact, this classic experiment of St. George and Pauling established, for the first time and before any crystallographic data were available, that the heme group and the ligand-binding site in hemoglobin reside at least partway inside the protein, rather than on the surface.134 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.12%3A_Other_Ligands_for_Biological_Oxygen_Carriers.txt |
For hemoglobin, the majority of ligands around the iron center are provided by a fairly rigid macrocycle, the protoporphyrin IX dianion (Figure 4.1), that by itself enforces a square-planar stereochemistry. Thus, the task of assembling synthetic analogues of the active site in hemoglobin is simplified. Essentially, any square-planar, tetradentate ligand containing at least a couple of nitrogen atoms will suffice: to this end a variety of other porphyrins have been used, as well as Schiff-base and non-porphyrinato nitrogen-containing macrocycles that serve to delineate the role of the porphyrin in dioxygen binding. Tetraphenylporphyrin, in place of the naturally occurring porphyrins, has served as the basis of numerous model systems. It is easily synthesized and derivatized (see below). Its fourfold symmetry precludes formation of chemical isomers that may arise if substitution on the asymmetric, naturally occurring porphyrins is attempted. Moreover, its derivatives can be crystallized. Finally, with the porphyrin meso positions occupied by phenyl groups, the molecule is less susceptible to photoinduced oxidation.
In order to obtain five-coordinate species, putative models for a deoxy hemoglobin, access to one side of the porphyrin must be blocked to the coordination of a second axial base (Reaction 4.33), but still must be accessible for the binding of small molecules (O2, NO, CO, etc). Or second, an axial base may be attached covalently to the porphyrin, to give so-called "tail-under" or "chelated" porphyrins. Here the chelate effect ensures an effectively 100 percent five-coordinate complex with only a 1:1 stoichiometric ratio of axial base to porphyrin. These two approaches are illustrated in Figure 4.14. A third means is to incorporate a sterically bulky substituent in the 2-position of the base, such as a methyl group, to give 2-methylimidazole (4.34). The formation of the hemochrome Fe(porph)(2-MeIm)2, where the iron atom is in the center of the porphyrin ring, is strongly disfavored relative to the 1-methylimidazole analogue, because the 2-methyl substituents clash with the porphyrin ring. The axial base needs to be a strong $\sigma$ donor, such as imidazole or pyridine, in order to increase affinity at the iron (or cobalt) center for dioxygen (Figure 4.18).
Steric hindrance on one side, or on both, provides a pocket for small molecules to bind and, for O2, prevents the bimolecular contact of two iron(II)-porphyrinato species that would lead to irreversible oxidation (Reaction 4.29). A picturesque collection of substituted porphyrins has been synthesized. Some of these are illustrated in Figure 4.23.31,72,160-164 The only system that has led to crystalline dioxygen complexes stable at room temperature is the "picket-fence" porphyrin.72* A derivative of this, the "pocket" porphyrin,165 and various "capped" porphyrins,166 provide binding sites with steric hindrance even to small diatomic molecules.
In the next section the structures of various derivatives of hemoglobin and its models are presented, and the relationship of structure to ligand-binding properties is examined. Although there is now a wealth of thermodynamic data available from model systems, attention is focused primarily on those for which structural data are also available.
* The pivalamido pickets (—NH—CO—C(CH3)3) of the picket-fence porphyrin are sufficiently bulky that their free rotation is sterically hindered. Thus the various atropisomers-$\alpha$$\alpha$$\alpha$$\alpha$, $\alpha$$\alpha$$\alpha$$\beta$, $\alpha$$\beta$$\alpha$$\alpha$, $\alpha$$\alpha$$\beta$$\beta$, where $\alpha$ denotes picket "up" and $\beta$ denotes picket "down"—can be separated chromatographically on the basis of their different polarities. The tail-under picket-fence porphyrin is derived from the $\alpha$$\alpha$$\alpha$$\beta$ atropisomer.72 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.13%3A_Requirements_for_a_Model_System_for_Hemoglobin.txt |
In order for dioxygen transport to be more efficient than simple diffusion through cell membranes and fluids, it is not sufficient that a metalloprotein merely binds dioxygen. Not only is there an optimal affinity of the carrier for dioxygen, but also, and more importantly, the carrier must bind and release dioxygen at a rapid rate. These thermodynamic and kinetic aspects are illustrated in Figure 4.3, a general diagram of energy vs. reaction coordinate for the process
$M + O_{2} \xrightleftharpoons[k_{-1}]{k_{1}} MO_{2}\tag{4.2}$
where M is an oxygen carrier, for example hemocyanin or a simple nonbiological metal complex. Thermodynamic or equilibrium aspects are summarized by $\Delta$G in Figure 4.3. As illustrated there, $\Delta$G is negative, and thus the forward reaction, dioxygen binding, is spontaneous. The equilibrium constant (K) is given by
$K = \frac{a(MO_{2})}{a(M)a(O_{2})} , \tag{4.3}$
where $a$ is the activity (i.e., effective concentration) of the component. The equilibrium constant is related to the change in free energy by
$\Delta G^{o} = -RT\ln K \ldotp \tag{4.4}$
The rate of the forward reaction ($k_1$) is related to $\Delta G_1^*$; the rate of the reverse reaction ($k_{-1}$) is related to $\Delta G^*_{-1}$. Provided that oxygen binding is effectively a single-step process, then
$K = \frac{k_{1}}{k_{-1}} \ldotp \tag{4.5}$
Usually the rates of the forward and reverse reactions are related by the empirical Arrhenius expression to quantities termed the activation energies (E1 and E-1) of the reactions, where
$k_{1} = A_{1}^{\frac{-E_{1}}{RT}}\; and\; k_{-1} = A_{-1}^{\frac{-E_{-1}}{RT}} \ldotp \tag{4.6}$
These quantities are experimentally accessible through the change in rate as a function of temperature.
Kinetic Factors
It is of little benefit to the organism if its dioxygen carrier, such as hemoglobin, binds and releases O2 at such slow rates that O2 is not delivered faster than it would be by simple diffusive processes. Thus, a binding rate within a couple of orders of magnitude of the rate of diffusion, together with the high carrying capacity of O2 that high concentrations of oxygen carrier enable (noted earlier), and a pumping system ensure adequate O2 supplies under all but the most physiologically stressful conditions.
Whereas measurements of equilibrium give little or no molecular information, rather more molecular information may be inferred from kinetic data. The processes of binding and release can be examined by a variety of techniques, with timescales down to the picosecond range. The temperature behavior of the rates gives information on the heights of energy barriers that are encountered as dioxygen molecules arrive at or depart from the binding site. The quantitative interpretation of kinetic data generally requires a molecular model of some sort. It is because of this multibarrier pathway that the equilibrium constant measured as k1/k-1 (Equation 4.5) may differ substantially from the thermodynamically measured value (Equation 4.3).
The simple Adair scheme outlined above is readily adapted to cater to kinetic data.
Dioxygen Reactions
Most biological conversions involving dioxygen require enzymatic catalysis. It is reasonable then that metals found in the proteins involved in the transport and storage of O2 also frequently appear, with minor modification of ligands, in enzymes that incorporate oxygen from dioxygen into some substrate. Dioxygen, in this case, is not only coordinated, but also activated and made available to the substrate. In the family of proteins with heme groups, hemoglobin is a dioxygen carrier and cytochrome P-450 is an oxygenase. A similar differentiation in function is also found for hemocyanin and tyrosinase from the family of proteins with a dinuclear copper complex at the active site. Note that not all enzymes that mediate the incorporation of oxygen from O2 into some substrate coordinate and activate dioxygen. For example, lipoxygenase probably catalyzes the conversion of a 1,4-diene to a 1,3-diene-4-hydroxyperoxy species by activation of the organic substrate. The active site does not resemble that of any known oxygen-carrier protein. This topic is discussed more fully in Chapter 5. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.14%3A_Requirements_for_Effective_Oxygen_Carriers.txt |
In our survey of the dioxygen chemistry of iron and copper species in earlier subsections, three general functions for the protein matrix became apparent: provision of ligand(s) in an appropriate stereochemistry; protection of the metal-dioxygen moiety from oxidation and competitive ligands; and modulation of dioxygen affinity through nonbonded interactions with distal groups.
Provision of Ligands to the Metal
In the hemoglobin family the heme group is anchored in a cleft in the globin chain by an imidazole ligand from a histidine residue (the proximal histidine). The other (distal) side of the heme plane is more or less open to accommodate a small sixth ligand (see Figure 4.2).
For hemerythrin and hemocyanin the requirements of the protein chain are more severe. In contrast to the hemoglobin family, all but two of the ligands are provided by the protein chain, and in addition the metal ions are encapsulated as a pair. The exogenous ligands for hemerythrin are a $\mu$-(hydr)oxo moiety and dioxygen or anions (depending on oxidation state); for hemocyanin, the identity of the second exogenous ligand, if there is one at all, is still unclear. Although hemerythrin has a distinctive cofacial bioctahedral structure (Figure 4.10)46-48 that would appear to be very difficult to assemble in the absence of the protein, it turns out that with a variety of tridentate ligands the ($\mu$-oxo)bis($\mu$-carboxylato)diiron(III) core may be assembled rather easily.82,156,157 Thus, this core appears to be a thermodynamically very stable structural motif. Such a synthesis has been termed "self-assembly" and appears to be a common phenomenon in biological systems.158 The low-temperature assembly of bis-copper(II)-$\mu$-peroxo complexes (models for oxyhemocyanin) from mononuclear copper(l) compounds provides other examples of this phenomenon.103f,g
Protection of the Metal-Dioxygen Moiety
The immobilization of the heme group inside the protein prevents (i) the bimolecular contact of an FeO2 species with an Fell species (Reaction 4.29b), the key step in the irreversible oxidation of FeII porphyrins; (ii) the facile access of nucleophiles that would cause autoxidation (Reactions 4.30 and 4.31); (iii) the oxygenase activity (Reaction 4.32) that is the normal function of other hemoproteins, such as cytochrome P-450, horseradish peroxidase, catalase, etc.; and (iv) the self-oxygenase activity that has been observed in some iron(II) systems that bind dioxygen, activating it for destruction of the ligand itself. Avoiding these last two fates also appears to be very important in the active site of hemocyanin. Finally (v), the globin chain serves to restrain the binding of the distal histidine to give a six-coordinate hemochrome (Reaction 4.33), at least at room temperature.159 Thus, unoxygenated hemoglobin is held in a five-coordinate state, allowing a rapid rate of oxygen binding and greater oxygen affinity—hemochromes such as Fe(TPP)(Py)2 are impervious to oxygenation and subsequent oxidation.
Modulation of Ligand-binding Properties
The protein chain in hemoglobin may place restraints on the iron-to-proximal histidine bond. On the other side of the heme, the distal histidine and occluded water molecules may hydrogen-bond to the coordinated dioxygen and force ligands to adopt geometries that are different from those observed in the absence of steric hindrances. The conformation of the porphyrin skeleton may also be perturbed by the protein chain. Clearly, it is the protein chain that bestows the property of cooperativity on oligomeric oxygen carriers. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.15%3A_Role_of_the_Protein_in_Effecting_Biological_Oxygen_Transport.txt |
Dioxygen is a powerful oxidant, capable of oxidizing all but the noble metals and of converting many low-valent metal complexes to higher-valent states. As will be detailed in this section, the binding of dioxygen to metals is most usefully considered as an oxidative addition process. The nature of the interaction is determined by the metal, its oxidation state, and its ligands that modulate the redox properties of the metal center. In biological and nonbiological oxygen carriers, several factors allow reversible binding of O2 to occur, even though this process is metastable with respect to (irreversible) oxidation of the metal, or its ligands, or other species that may be present. Later in this section the bioinorganic chemistry of iron, copper, and cobalt is described. For a wider perspective on the coordination chemistry of these metals, see comprehensive texts on inorganic chemistry.56-58
Many techniques have been used to probe the metal-dioxygen moiety. A summary of these techniques, key concepts, and results is presented in Table 4.3.59-61 UV-visible spectroscopy usually characterizes the oxidation state of the metal and in favorable cases the number, geometry, and ligand field strength of ligands. The O—O and M—O stretching modes may be investigated with infrared spectroscopy, provided that the complex is not a centrosymmetric dimer, for then the O—O stretch for the $\mu$-dioxygen species is infrared-inactive. Resonance Raman techniques complement infrared spectroscopy. Not only are the selection rules different in Raman spectroscopy, but a suitable choice of the irradiating wavelength (to coincide approximately with an M-L electronic transition) can amplify those vibrational modes that are coupled, or in resonance, with the electronic transition. This technique is particularly suited as a probe of the metal-ligand environment of metalloproteins, since the many solely protein vibrational modes disappear into background noise. Geometric information on the orientation of the CO moiety with respect to the heme normal has been obtained by examining polarization behavior of infrared bands following photolysis of the Fe—CO bond by linearly polarized light.
Spin and oxidation states of mononuclear iron-porphyrin systems may be assigned directly from magnetic susceptibility measurements and indirectly from Mössbauer spectroscopy. Variable temperature susceptibility measurements are particularly useful for detecting dinuclear systems that share at least one ligand in common if there is antiferromagnetic (or ferromagnetic) coupling of the electron spin of one metal center with that of a second.
Definitive characterization of the stereochemistry is usually provided by x-ray diffraction data when single crystals are available. In general, the level of resolution and precision available from protein crystal structures leads to tantalizing uncertainties over the geometry of the M—O2 species and of the structural changes occurring on oxygenation that are the origin of cooperativity. Precise structural data are more readily obtained from small-molecule model systems. The relevance of these to biological systems is established through congruence of spectroscopic and functional properties. X-ray diffraction techniques also provide important information on the environment beyond the immediate surroundings of the metal center: this information is usually unobtainable from other techniques, although recent developments in two-dimensional NMR spectroscopy can provide this information for diamagnetic systems. Limited information may be obtained with the use of spin labels or, if the metal center is paramagnetic, with EPR techniques.
Two other techniques that selectively probe the immediate environment of the metal center are EXAFS (Extended X-ray Absorption Fine Structure)60 and XANES (X-ray Absorption Near-Edge Structure).61 The former may yield information on the number and type of bonded atoms and their radial separation from the metal center. The latter technique may reveal the oxidation state and, in principle, may yield geometric information, although in its present state of development some interpretations are contentious. Both techniques have the advantage of not requiring crystalline material. The structural information is more reliable if definitive model systems are available for comparison.
X-ray (and, less frequently, neutron) diffraction techniques on single crystals give absolute structural information* and thus provide the basis for interpretation of data obtained from these other techniques that yield relative structural information.
* In favorable situations, sophisticated NMR techniques have been applied successfully to detennine the polypeptide folding (e.g., in metallothionein).55
Table 4.3: Techniques used to probe the active sites of oxygen carriers
Technique Abbrev. Description of Technique Description of Results
Nuclear Magnetic Resonance NMR Quantized orientation of nuclear spin in a magnetic field. Energy separations sampled with radio-frequency radiation. Identification of histidine by deuterium exchange (N—H vs. N—D) at or near metal, especially if paramagnetic.
Electron Paramagnetic Resonance EPR Quantized orientation of electron spin in a magnetic field. Energy separations samples with X- or Q-band microwave radiation. Location of unpaired electron density from hyperfine splitting by metals or atoms with nuclear spin.
Magnetic Susceptibility Strength of interaction of sample with magnetic field. Solid state or solution state by Evans' NMR method. Identification of spin state, spin-equilibria, and spin coupling (ferro- or antiferromagnetic); identificaiton of FeIII—O—FeIII moiety.
Infrared Spectroscopy IR Vibrational modes involving change in dipole moment. Classificaiton of O—O moiety (superoxo vs. peroxo). Identification of v(M—O) and v(M—O—M) modes, etc.
Raman and Resonance Raman R, RR Vibrational modes involving a change in polarizability. For RR enchancement of modes coupled with electronic transition excited by laser light source. Complementary to v(O—O) and v(M—O) especially in metalloproteins. In porphyrins, oxidation and spin state.
UV-visible spectroscopy UV-Vis Valence electron transition Electronic state of metal from d-d transitions. Identification of unusual ligands, e.g., CU(II)—SR, FEIII—OPh, FeIII—O—FeIII. Single crystals and polarized light give geometrical information.
X-ray Photoelectron Spectroscopy XPS (ESCA) Inner-shell electron transitions. Oxidation state of metal.
Mössbauer Spectroscopy Excitation of nuclear spin by $\gamma$ rays. Oxidation and spin state. Antiferromagnetic coupling (Fe only).
X-ray Single-crystal Diffraction Fourier transform of diffraction data reveals location of electron density. Precise three-dimensional structure, bond distances and angles for small molecules. Lower resolution and precision for proteins.
Extended X-ray Absorption Fine Structure EXAFS Backscattering of x-rays produces interference fringes on absorption curve at energies just greater than metal absorption edge (K$\beta$ transition) Number, type, and radial distance of ligand donor atoms bonded to the metal.
X-ray Absorption Near-edge Structure XANES Similar to EXAFS except that absorption is monitored at energies near and below the absorption edge. As for EXAFS. May give geometric information | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.16%3A_Selected_Chemistry_of_Dioxygen_Iron_Copper_and_Cobalt.txt |
Upon binding a second axial ligand, the iron center together with the axial base move toward the plane of the porphyrin, initiating a change in spin state from high-spin to low-spin when the sixth ligand is O2 or CO or any other strong ligand with an even number of valence electrons. Given these general features, what are the structural differences between systems that bind O2 with high affinity and those that bind O2 with low affinity? The answers to this question are relevant to understanding at the molecular level the mechanism of cooperativity, where a low-affinity conformation, the T state, and a high-affinity conformation, the R state, are in dynamic equilibrium in one tetrameric molecule. In looking at crystallographic data one sees a particular conformation frozen in the crystal, usually the one of lowest free energy among many in equilibrium in the solution state. The R $\rightleftharpoons$ T equilibrium for hemoglobin is moderately rapid, at 4 x 103 s-1; hemocyanin also switches quaternary conformations with a similar rate constant.4
Human hemoglobins are a heterogeneous group. Many mutants are known, and several have been structurally characterized. A structural alteration that affects the equilibrium between R and T states has a marked effect on ligand affinity and cooperativity in hemoglobin. If a specific amino-acid substitution destabilizes the T state, then the transition to the R state will occur earlier in the ligation process, and the hemoglobin will have an increased oxygen affinity. Hemoglobin Kempsey is an example. In this mutant an aspartic acid on the $\beta$ chain is replaced by asparagine. Conversely, if the R state is destabilized, then the hemoglobin will have a lowered oxygen affinity. Hemoglobin Kansas is an example. Here an asparagine on the $\beta$ chain has been replaced by threonine.9
Structural Changes in Normal-affinity Systems
It was proposed earlier that the molecule Fe(PF)(1-MeIm)(O2) in the solid or solution state was a fair approximation to the reference gas-phase molecule. The axial base, although not oriented for minimization of contacts with the porphyrin (i.e., $\phi$ = 45°), is well-removed from an eclipsing orientation where $\phi$ = 0 ± 10°; the Fe atom is centered in the plane of a highly planar porphyrin; the O2 ligand is oriented for minimization of contacts with the porphyrin, and its geometry is largely unconstrained by distal groups (the pickets); no groups are hydrogen-bonded to the axial base. The major difference from the reference state is that there is a significant attractive interaction between the electronegative dioxygen moiety and the amide groups on the pickets, and a smaller repulsive interaction with the picket t-butyl groups.176
For the CO adduct, contacts with the pickets are all at ideal van der Waals' separations and the Fe—CO moiety is free to assume its normal linear geometry. For CO binding the reference molecule is again the carbonyl adduct of the iron picket-fence porphyrinato molecule. In contrast to O2 binding, there are no specific distal effects, such as hydrogen bonding, by which CO affinity may be increased; there remain many ways, as with O2 binding, by which CO affinities may be reduced. Thus, the CO binder with highest affinity is the iron picket-fence porphyrin.
The O2 affinities of myoglobin, R-state hemoglobin, and the Fe(PF)(1-MeIm) system are similar. However, the means by which this is achieved are different, and this difference is reflected most clearly in the kinetics of binding and release of O2, which for Mb are much slower. The similarities and differences are summarized in Table 4.9, which is culled from Tables 4.2, 4.4, and 4.5.
Table 4.9 - Comparison of the picket-fence porphyrin system with Mb.
a) Solution state studies on Fe(PF-Im). Structural details on Fe(PF)(1-Melm)(O2) and Fe(PF)(2-Melm).
Characteristic Mb
Fe(PF-Im)a
Fe(PF)(1-MeIm)
Fe(Poc-PF)
(1-MeIm)
P1/2(O2), Torr 0.7 0.58 0.36
P1/2(CO), Torr 0.018 0.000022 0.0015
kon(O2), $\mu$M-1s-1 15 430 2.2
koff(O2), s-1 10 2,900 9
kon(CO), $\mu$M-1s-1 0.50 36 0.58
koff(CO), s-1 0.015 0.0078 0.0086
Solvent H2O/PO43- toulene toulene
Local Environment
His (H2O)
polar, protic
H-N(amide)
polar, aprotic
H-N (amide)
phenyl
O2 • • • distal, Å 2.97(18)
NH 4.06(5)
CH3 2.67(6)
CO • • • distal, Å (calc.) 2.7
NH 5.0
CH3 3.3
Fe-NIm(deoxy), Å
2.22 2.095(6)
Fe-NIm(oxy), Å 2.07(6) 2.07(2)
Fe-O, Å 1.83(6) 1.75(2)
$\phi_{Im}$ (deoxy), deg 19° 22.8°
$\phi_{Im}$ (oxy), deg 20°
$\phi_{O2}$, deg ~0° 45°
Fe • • • porph(deoxy), Å 0.42 0.43
Fe • • • porph(oxy), Å 0.18 -0.03
For the cobalt-dioxygen derivative, the putative hydrogen bonding between the dioxygen and the amide groups of the pickets assumes greater importance because the coordinated dioxygen is substantially more negative. Again the picket-fence porphyrin, being structurally characterized, is the reference system. Although no Co picket-fence porphyrin structures have been determined, the structures may be predicted with confidence from the iron analogues together with related structures of ColI and CollI tetraphenylporphyrinato systems.*
* For CoHbO2, single-crystal EPR spectra have been interpreted in terms of a nearly triangularly coordinated O2,197 although a crystal structure of CoMbO2 shows a bent CoOO group.198 There is no precedent for this triangular arrangement in any Colll-superoxo (O21-) system, whereas there are many for angularly coordinated O2 in electronically not dissimilar square-planar Schiff-base systems.66 Regardless of geometry, the picket amide • • • O2 contacts do not change substantially.
Structural Changes in Low-affinity Systems
The 2-methyl substituent on 2-methylimidazole is not sterically active in the five-coordinate structures Fe(PF)(2-MeIm) and Fe(TPP)(2-MeIm), since the iron atom is displaced from the plane of the porphyrin by the expected amount and the Fe—NIm bond is unstretched and similar to that in deoxyhemoglobin (low O2 affinity) and deoxyMb (higher O2 affinity). Moreover, resonance Raman measurements also indicate little strain$\dagger$ in this bond.200 In other words, there is no "tension at the heme," a key concept in early discussions of cooperativity before structures on model systems and high-resolution, refined protein structures became available.11a On moving into the plane of the porphyrin upon oxygenation, the 2-methyl substituent prevents the Fe-imidazole group from achieving its optimum geometry with the iron at the center of the porphyrin hole, as seen in the structure of Fe(PF)(1-MeIm)(O2). Thus, the sterically active 2-methyl substituent leads to lowered O2 (and CO) affinity relative to the 1-methyl analogue. In metrical terms the lowered affinity is reflected in an increase in the sum of the axial bond lengths from 1.75 + 2.07 = 3.82 Å to 1.90 + 2.11 = 4.01 Å.
In the crystal structure of Fe(C2Cap)(1-MeIm)(CO) the cap is about 5.6 Å from the porphryin plane.121d Hence, in the crystal structures of the free base H2(C2Cap)201a and FeCI(C2Cap)201b species, in which the cap is screwed down to approximately 4.0 Å from the porphyrin plane, considerable conformational rearrangement of the cap and the four chains attaching it to the porphyrin is needed to provide room for a small ligand such as CO. This is even more pronounced in a Co(C3Cap) complex where the cap is only 3.49 Å from the mean porphyrin plane.202 Thus not only is affinity for CO lowered, but some additional discrimination against it is induced, since a linear, perpendicular coordination creates considerable strain energy elsewhere in the molecule.
For the pocket porphyrin (Figure 4.23), structural data are available on the carbonyl adduct.121bThe CO ligand is unable to achieve the linear perpendicular geometry seen in the high-affinity picket-fence porphyrin derivative, Fe(PF)(1-MeIm)(CO),110and distortion of the porphyrin core is greater. In the pocket-porphyrin system, O2 affinity is unaffected, but CO affinity is lowered.
The crystal structure of partially oxygenated hemoglobin, [$\alpha$-FeO2]2[$\beta$-Fe]2,191a reveals that the quaternary structure, except in the immediate vicinity of the $\alpha$ hemes, which have O2 coordinated, resembles that of T-state deoxyhemoglobin rather than R-state liganded hemoglobin. In accord with the low affinity of T-state hemoglobin, the Fe—NIm bonds for the six-coordinate $\alpha$-hemes at 2.37 Å are significantly longer than those in fully oxygenated R-state oxyhemoglobin, [$\alpha$-FeO2]2[$\beta$-FeO2]2 in the notation above, (1.94 ($\alpha$-hemes) and 2.06 Å ($\beta$ hemes)) and that found in oxymyoglobin (2.07 Å). In contrast to the R-state structure and oxyMb, the a-hemes are folded as seen in the deoxy parent, leaving the Fe still substantially displaced (0.2 Å) from the plane of the four pyrrole nitrogen atoms. The deoxyhemoglobin T-state quaternary structure also has been observed in two other partially liganded hybrid hemoglobins, [$\alpha$-FeCO]2[$\beta$-Mn(II)]2203 and [$\alpha$-Ni]2[$\beta$-FeCO]2.191d Again, structural changes upon coordination do not propagate beyond the immediate vicinity of the liganded heme to the critical $\alpha_{1} \beta_{2}$ interfaces.
Note that although the crystal structure of hemoglobin A reveals that access to the binding site for the $\beta$ chains is blocked by groups at the entrance to the cavity above the iron center, this does not prevent facile access to the binding site; the rate of O2 binding is slowed by a factor of only five. A similar situation occurs also for vertebrate myoglobins.
The large structural differences that exist between deoxy (T) and oxy (R) hemoglobin and the much smaller differences between deoxy (T) and partially liganded (T) hybrid hemoglobin are shown in Figure 4.32.203
Because of the steric hindrance afforded by the distal histidine, all biological systems have low affinity for CO relative to the picket-fence porphyrins, with the exception of mutants where the distal histidine has been replaced by glycine. Thus low affinity to CO is associated primarily with the inability of the Fe-CO group to achieve its preferred linear geometry perpendicular to the porphyrin.
Low-affinity O2 binding in the hemoglobins appears to be associated with the inability of the Fe-proximal histidine unit to move into the plane of the porphyrin and less so to distal effects, such as a cavity too small to accommodate the coordinated ligand. The blocked access to the site affects the kinetics but not necessarily the thermodynamics of ligand binding, as evidenced by the structure of T-state [$\alpha$-FeNi]2[$\beta$-FeCO]2191d Some similarities between the structures and properties of partially oxygenated (T-state) [$\alpha$-FeO2]2[$\beta$-Fe]2 hemoglobin and Fe(PF)(2-MeIm)(O2) are provided in Table 4.10. In the synthetic systems low O2 affinity can be induced by 2-methyl substituents—a restraint on the movement of the Fe-imidazole moiety analogous to that provided by the protein chain. A second means is by distal effects, such as caps and straps.
Table 4.10 - Comparison of the low-affinity picket-fence porphyrin system with low-affinity (T-state) partially liganded hemoglobin.
a) Ligand binding to Fe(PF)(1,2-Me2Im)
Characteristic
HbAT
[$\alpha$-FeO2]2[$\beta$-Fe]2
Fe(PF)(2-MeIm)
Fe(PF)(1,22-Me2Im)a
P1/2(O2), Torr 46, first O2 38
P1/2(CO), Torr ~0.7, 1st CO 0.0089
kon(O2), $\mu$M-1s-1 2.9($\alpha$) 106
koff(O2), s-1 183($\beta$) 46,000
kon(CO), $\mu$M-1s-1 0.099 1.4
koff(CO), s-1 0.09 0.14
Solvent Tris buffer, pH 7, no 2,3-DPG toluene
Local Environment
histidine
polar
H-N(amide)
polar, aprotic
O2 • • • distal, Å ?
NH 3.88
CH3 2.77(3)
CO • • • distal, Å (calc.)
NH 4.9
CH3 3.5
Fe-NIm(deoxy), Å
2.13(6) (average of $\alpha$ and $\beta$ 2.095(6)
Fe-NIm(oxy), Å 2.24(10) 2.107(4)
Fe-O, Å 1.82(4) 1.898(7)
$\phi_{Im}$ (deoxy), deg 21(3) (average of $\alpha$ and $\beta$ 22.8
$\phi_{Im}$ (oxy), deg ~6 22.2
$\phi_{O2}$, deg ? 45
Fe • • • porph(deoxy), Å 0.38(5) (average of $\alpha$ and $\beta$ 0.43
Fe • • • porph(oxy), Å 0.19(5) 0.09
$\dagger$ Shortly (10- 9-10-12 s) after a ligand dissociates, a large difference in v(Fe—NIm) between R and T structures is observed, prior to relaxation to the equilibrium R and T conformations.199
Structural Changes in High-affinity Systems
Few structural data are available for high-affinity oxygen carriers. The crystal structures of two leghemoglobin derivatives, a monomeric myoglobin-like oxygen carrier found in the nitrogen-fixing nodules of legumes, are known at 2.0 and 3.3 Å.204,205 The binding pocket appears more open, perhaps allowing H2O to enter and partake in stronger hydrogen bonding than that offered by the distal imidazole. Consistent with this notion is the more rapid rate of autoxidation observed for oxyleghemoglobin. Aplysia oxymyoglobin, which lacks a distal histidine, also autoxidizes rapidly,204 although a distal arginine further along the helix E, E10Arg, fulfills the role of the distal histidine by hydrogen bonding to the sixth ligand, at least in the fluoride derivative, met-MbF.
Although no structural data are available, a tenfold increase in O2 affinity was observed between an ester-strapped porphyrin, offering no hydrogen-bonding possibilities, and its conformationally very similar amide analogue. O2 • • • amide hydrogen bonding was demonstrated by means of NMR shift data (Zn and Fe—CO complexes vs. the Fe—O2 complex) and from infrared spectroscopy, which showed shifted amide N—H absorptions.166
The specific structural features that lead to the extraordinarily high affinity for O2 and low affinity for CO in hemoglobin Ascaris remain unidentified. This high affinity is due to an extremely slow dissociation rate of O2 of only 0.1s-1; in most hemoglobins the rate is about 10 to 2,500 s-1 (Table 4.2). Dioxygen binding is thus close to irreversible. Figure 4.27 shows that hydrogen bonding to the coordinated dioxygen ligand, unrestrained motion of the Fe-proximal histidine group into the plane of the porphyrin, hydrogen bonding to the proximal histidine, and, in the deoxy form, compression of the Fe—NIm bond and decrease in the out-of-plane displacement of the Fe atom will all increase O2 affinity over that of a system where these effects are absent.
When hydrogen bonding is impossible, as in various synthetic systems (Table 4.5) as well as hemoglobin Glycera and Mb(E7His → Gly), O2 affinity is much lower than when hydrogen bonding can occur (see Table 4.6), especially for the cobalt analogues. But caution is needed in the absence of complete structural information: the lowered affinity of Aplysia hemoglobin had been attributed to the lack of a distal histidine and its attendant hydrogen-bonding capabilities. However, the crystal structure reveals that an arginine residue, normally directed out into the solution, is capable of folding back into the ligand-binding pocket and of hydrogen bonding to ligands at the sixth site. In oxyhemerythrin the hydrogen bonding of the coordinated hydroperoxy group to the oxo bridge linking the two iron atoms (Figure 4.10B), described in Section II.F.1, may not only increase the stability of oxyhemerythrin,146 but also facilitate electron transfer that occurs in dioxygen binding.205 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.17%3A_Stereochemical_Changes_Upon_Ligation.txt |
The interaction of ligands, such as dioxygen, with metal complexes, such as iron-porphyrinato systems, and the means by which this interaction is characterized, have been covered in broad outline in the previous sections. As noted earlier, the affinities of hemoglobins for carbon monoxide and dioxygen span a wide range (see Table 4.2 and Figure 4.24). In this section the active site is examined in much finer detail than before in order to develop relationships between perturbations in structure and affinity (and hence function)—so called structure-function relationships. The reference point is the somewhat hypothetical situation where the dioxygen binder is in the gas phase and independent of interactions with solvent molecules, solute molecules, and itself, and where dioxygen, carbon monoxide, and other small molecules may bind without steric constraints—in other words, a state where intrinsic affinity is measured. In this section attention is focused exclusively on the hemoglobin family and on iron- and cobalt-porphyrinato systems. In recent years structural data on hemoglobin, myoglobin, and their derivatives have become available with a precision that permits meaningful comparison with the more precisely determined model or synthetic systems. In addition, the various hemoglobins and myoglobins, and especially the naturally occurring mutants of hemoglobin A (human Hb), have provided a sort of poor man's site-directed mutagenesis. Now the techniques of molecular biology permit the site of mutation to be selected, the altered gene to be inserted into E. coli, and the mutant protein to be expressed in large (mg) quantities. With the conditions for crystallization of hemoglobins now well-established, we can discover quite rapidly what structural perturbations are caused by the substitution of one amino acid for another, and can relate these to the perturbations in properties, such as cooperativity, dioxygen affinity, and kinetics of ligand binding.
The principles enunciated here are applicable generally to hemerythrin and hemocyanin; however, we currently lack the thermodynamic and especially structural data we would like to have for these systems.
Ligand Affinities in Hemoglobins and Their Models
The O2 affinities in biological carriers span five orders of magnitude, which at room temperature corresponds to a difference in the free energy of oxygen binding
$\delta \Delta G = -RTln(K_{max}/K_{min}) = -RTln(P_{1/2min}/P_{1/2max}) \tag{4.47}$
of about 6.0 kcal/mol. This wide range of O2 and CO affinities has not yet been paralleled in synthetic systems; the values for O2 affinity do not exceed those for R-state human hemoglobin. A selection of values from model systems is given in Table 4.5.23,31,160-165 For the flat-open porphyrin system (Figure 4.23) the dioxygen ostensibly binds in an unconstrained manner, but is actually subject to solvent influences. In order to obtain thermodynamic constants on these "unhindered" systems, one must gather data at several low temperatures and then extrapolate to room temperature, or obtain them from kinetic measurements, K = kon/koff, at room temperature.
For the picket-fence porphyrins, dioxygen binds in a protected pocket that is deep enough to accommodate it and to prevent the dimerization that leads to irreversible oxidation, provided that there is a slight excess of base to ensure full saturation of the coordination sites on the unprotected face of the porphyrin.72 Thus the picket-fence, the capped, and the bis-pocket porphyrins reversibly bind dioxygen at room temperature with little oxidation over many cycles. This stability facilitated isolation of crystals of a synthetic iron-dioxygen species of the picket-fence porphyrin. The capped porphyrin offers a more highly protected site. The low affinity these latter systems have for dioxygen indicates that the binding cavity is so small that repulsive steric interactions between coordinated dioxygen and the cap are unavoidable. The left-hand side of Figure 4.24 depicts on a logarithmic scale the range of O2 affinities. Each power of 10 corresponds to around 1.2 kcal/mol at 25 °C.
The right-hand side of Figure 4.24 illustrates the range of affinities for CO binding. For many synthetic systems the CO affinities are orders of magnitude greater than in the biological systems that have an O2 affinity similar to the synthetic; for example, see the entries for the picket-fence porphyrin. Comparison of the left- and right-hand sides of Figure 4.24 reveals that the strongest O2 binder, hemoglobin Ascaris, is one of the weakest CO binders. The O2 affinity of the picket-fence porphyrins is very similar to that of myoglobin, but, as will be detailed shortly, one cannot infer from this that the binding sites are strictly comparable. Indeed, similar affinities have been observed with a non-porphyrin iron complex.121,162 Moreover, if the CO affinity of myoglobin paralleled that of the picket-fence porphyrins, some 20 percent of myoglobin (and hemoglobin) would be in the carbonmonoxy form (in contrast to the approximately 3 percent that occurs naturally), a level that could render reading this section while chewing gum physically taxing.117
Table 4.5 - Thermodynamics and kinetics of ligand binding to synthetic oxygen carriers at 20-25 ºC
a) When available P1/2 are from thermodynamic measurements, otherwise from kon/koff, where solubility of O2 in toluene is 1.02 x 105 M/Torr and of CO in toluene is 1.05 x 10-2; solubilities in benzene are very similar.
b) Some koff are calculated from K(O2), kon(CO), and M.
Carrier
P1/2(O2)
Torr
$\Delta$H
kcal/mol
Dioxygen
$\Delta$S
eu
Binding
kon
$\mu$M-1s-1
koff
s-1
P1/2(CO)
Torr
Carbon
$\Delta$H
kcal/mol
Monoxide
$\Delta$S
eu
Binding
kon
$\mu$M-1s-1
koff
s-1
Toluene/benzene solvent
Picket fence, pocket
Fe(PF-Im) 0.58 -16.3 -40 430 2,900 0.000022 36 0.0078
Fe(PF)(1,2-Me2Im) 38 -14.3 -42 106 46,000 0.0089 1.4 0.14
Fe(Poc-PF)(1-MeIm) 0.36 2.2 9 0.0015 0.58 0.0086
Fe(Poc-PF) (1, 2-Me2lm) 12.6 -13.9 -28 1.9 280 0.067 0.098 0.055
Fe(Bis-Poc) (1, 2-Me2Im) 508 -14.4 -47 0.0091
Cap
Fe(C2Cap) (1-MeIm) 23 -10.5 -28 0.0054 0.95 0.05
Fe(C2Cap) (1, 2-Me2Im) 4,000 -9.7 -36 0.20
Strapped
Fe(7, 7-CP) (1, 5-Cy2Im) 1.4 65 1,000 0.00091 6 0.05
Fe(6, 6-CP) (1, 5-Cy2Im) 700 0.1 800 0.17 0.03 0.05
Flat open
Fe(PPIX-Im) 5.6 62 4,200 0.00025 11 0.025
Bis-strapped
Fe(Amide-Im) 0.29 310 620 0.000017 40 0.067
Fe(Amide-Py) 2.0 360 5,000 0.00009 35 0.03
Fe(Ether-Py) 18 300 40,000 0.0001 68 0.069
H2O, alkylammonium micelles, pH 7.3
Fe(PPIX-Im) 1.0 -14.0 -3.5 26 4.7 0.002 -17.5 -34 3.6 0.009
Fe(MPIX-Im) 0.57 22 23 0.0013 11 0.019
Fe(MPIX-Py) 12.2 1 380 0.0021 12 0.035
There is a convenient index to summarize the extent to which CO (or O2) binding is discriminated against for a given iron-porphyrin system. M is defined as the ratio of O2 affinity (as P1/2) to CO affinity for a particular system and experimental conditions:
$M = \frac{P_{1/2}(O_{2})}{P_{1/2}(CO)} \tag{4.48}$
From Figure 4.24 and from Tables 4.2 and 4.5 the M values calculated may be somewhat arbitrarily divided into three classes: those where M > 2 x 104 (good CO binder); those where 2 x 102 < M < 2 x 104; and those where M < 2 x 102 (good O2 binder). An analogous parameter, N, may be defined to summarize the differences in the O2 affinity between an iron-porphyrin system and its cobalt analogue:
$M = \frac{P_{1/2}(O_{2}—Co)}{P_{1/2}(O_{2}—Fe)} \ldotp \tag{4.49}$
For the picket-fence porphyrins and for vertebrate hemoglobins N is in the range 10 to 250, whereas for the flat-open porphyrins and for some hemoglobins that lack a distal histidine (e.g., hemoglobin Glycera and hemoglobin Aplysia), N is at least an order of magnitude larger, indicating for these latter species that the cobalt analogue binds O2 relatively poorly167,168 (see Table 4.6).
Note that whereas the O2 binding of the picket-fence porphyrins is similar to that for myoglobin, the kinetics of the process are very different; the synthetic system is more than an order of magnitude faster in k1 and k-1 (often also referred to as kon and koff). On the other hand, O2 binding to the pocket porphyrin is similar to that for the biological system. The factors by which ligand affinities are modulated, generally to the benefit of the organism, are subtle and varied, and their elucidation requires the precise structural information that is currently available only from x-ray diffraction experiments. Figure 4.25 shows the structural features of interest that will be elaborated upon in the next subsections.110,169
Table 4.6 - Relative affinities (M) of iron-porphyrinato systems for O2 and CO, and relative affinities (N) for O2 of iron and cobalt-porphyrinato systems.
Compound
P1/2(Fe—CO)
Torr
P1/2(Fe—O2)
Torr
M
P1/2(Fe—O2)/P1/2(Fe—CO)
P1/2(Co—O2)
Torr
N
P1/2(Co—O2)/P1/2(Fe—O2)
H2O, pH 7
Whale Mb 0.018 0.51 28 57 110
Whale Mb
(E7His→Gly)
0.0049 6.2 1,300
Aplysia Mb 0.013 2.7 200 50 x CoMb > 1,000
Glycera Mb 0.00089 5.2 5,800 50 x CoMb > 1,000
Fe(PPIX-Im) 0.002 1.0 500
Toluene/Benzene
Fe(PF-Im)/ Co(PF) (1-MeIm) 0.000022 0.58 27,000 140 240
M(PF)(1, 2-Me2Im) 0.0089 38 4,300 900 24
M(Bis-Poc)- (1, 2-Me2Im) 0.0091 508 55,800
Fe(PPIX-Im)/ Co(PPIX) (1-Melm) 0.00025 5.6 22,000 18,000 3,200
M(C2-Cap)(1-Melm) 0.0054 23 4,200 140,000 6,100 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.18%3A_Structural_Basis_of_Ligand_Affinities_of_Oxygen_Carriers.txt |
Stereochemistry of the Active Site
Before the advent of techniques that enabled the preparation and stabilization of oxyhemoglobin crystals, key information on the probable structure of oxyhemoglobin and thence on the mechanism of cooperativity was extrapolated from structures of methemoglobin derivatives11 and from various five- and six-coordinate cobalt-porphyrinato complexes.110,112,113 The structures of these met derivatives have proved to be similar to that of oxyhemoglobin, at least in the stereochemistry of the metalloporphyrinato species and for the protein tertiary and quaternary structure as well.
Two synthetic iron-dioxygen adducts built from the picket-fence porphyrins have been structurally characterized.172,187 The high, effectively fourfold symmetry of the binding pocket in these systems results in fourfold disorder of the angularly coordinated dioxygen molecule, and precludes the precise and accurate measurements of the Fe—O—O angle and O—O separation that are grist to the theoretical mills.* Figure 4.28B illustrates the stereochemistry for one conformer. Subsequently, the structures of several dioxygen adducts of biological oxygen carriers have been determined.183,188-191 Although the dioxygen moiety is usually ordered, the precision is tantalizingly just less than that needed to decide whether the apparently more-linear geometry seen for oxyerythrocruorin183 and oxyhemoglobin190 is significantly different from that for oxymyoglobin188 and therefore attributable to the water molecule or imidazole that is hydrogenbonded to the coordinated dioxygen ligand. Nonetheless, several interesting differences emerge.
The axial base in oxymyoglobin and oxyhemoglobin is almost eclipsed; that is, $\phi_{1}$ ≈ 0°. The axial base has moved from a tilted position in deoxyhemoglobin to a symmetric one in oxyhemoglobin. In the absence of steric constraints, the iron atom is essentially in the center of the porphyrin plane for Fe(PF)(1-MeIm)(O2), oxymyoglobin, and oxyhemoglobin. For the 2-methyl analogue, Fe(PF)(2-MeIm)(O2), the iron remains significantly out of the plane, as also appears to occur for oxyerythrocruorin.
In the structure of Fe(TPP)(Py)(CO), Figure 4.30, a model for carbonyl hemoglobins, the iron atom is in the plane and the Fe—C$\equiv$O bond is linear and perpendicular, as expected.118 Not so for carbonyl hemoglobins, where the blob of electron density that is identified with the coordinated carbon monoxide lies substantially off the normal to the porphyrin. We return to this point shortly. In general, with the exception of the coordinated ligand, the structures of sixcoordinate low-spin hemoglobins, whether FeII or FeIII, are similar. Indeed, the refined structures of oxy- and carbonmonoxyhemoglobin are superimposable within experimental uncertainties, except in the immediate vicinity of the diatomic ligand. Some metrical details are given in Table 4.8.11c ,118,121,172,183,187,188,190-192
Table 4.8 - Metrical details of selected liganded hemoglobins and their modelsa a) See Figure 4.25 for definition of symbols. b) Alternative interpretation of EXAFS data.
Compound
Resol.
Å
Fe-Np
Å
Fe-porph
Å
Doming
Å
Fe-L1
Å
Fe-L2
Å
Fe-XY
deg.
$\varnothing_{1}$
deg.
$\varnothing_{2}$
deg.
Tilt
deg.
Dioxygen Adducts
Fe(PF)(1-MeIm)(O2) 1.98(1) -0.03 0.02 2.07(2) 1.75(2) 131(2) 20 45 0
Fe(PF)(2-MeIm)(O2) 1.996(4) 0.09 0.02 2.107(4) 1.898(7) 129(1) 22 45 7
MbO2 1.6 1.95(6) 0.18(3) 0.01 2.07(6) 1.83(6) 115(4) 1 ~0 4
EXAFS 2.02(2) 2.06(2) 1.80(2)
123(4)
148(8)b
ErO2 1.4 2.04 0.38 -0.08 2.1 1.8 150 7 3
HbO2 $\alpha$ 2.1 1.99(5) 0.12(8) 0.04 1.94(9) 1.66(8) 153(7) 11 0 3
HbO2 $\beta$ 1.96(6) -0.11(8) 0.11 2.07(9) 1.87(3) 159(12) 27 45 5
EXAFS 1.99(2) 2.05(2) 1.82(2)
122(4)
143(8)b
[$\alpha$-FeO2]2[$\beta$-Fe]2 $\alpha$ 2.1 2.04(4) 0.19(5) 0.17(5) 2.24(10) 1.82(4) 153(4) 6 11
[$\alpha$-FeO2]2[$\beta$-Fe]2 $\beta$ 0.3 2.2
Carbonmonoxy Adducts
MbCO 1.5 1.97(3) 0.00 0.03 2.2 1.9 140 30
EXAFS 2.01(2) 2.20(2) 1.93(2)
127(4)
145(8)b
ErCO 1.4 2.01 -0.11 -0.10 2.1 2.2 161(9) 7 1
HbCO $\alpha$ 2.2 2.02 -0.10 1.95 1.83 175(15)
HbCO $\beta$ 2.03 -0.10 2.20 1.70 171(15)
[$\alpha$-Ni][$\beta$-FeCo] $\alpha$ 2.6
[$\alpha$-Ni][$\beta$-FeCo] $\beta$ 0.15 0.12 2.23(5)
Fe(TPP)(Py)(CO) 2.02(3) -0.02 -0.02 2.10(2) 1.77(2) 179(2) 45 ~0
EXAFS 2.02(2) 2.09(2) 1.81(2)
138(6)
180(11)b
Fe(poc)(1,2-Melm)(CO) 1.973(8) 0.001 2.079(5) 1.768(7) 172.5(6)
Fe(C2Cap)(1-Melm)(CO)
1.990(7)
1.988(13)
0.01
0.02
2.043(6)
2.041(5)
1.742(7)
1.748(7)
172.9(6)
175.9(6)
Interactions of Coordinated Ligands with Distal Groups
Without exception to date (but see footnote 1 in Reference 168), in structurally characterized oxyhemoglobins, the coordinated dioxygen ligand is hydrogen-bonded to the distal histidine or to a water molecule—even though theoretical calculations show that hydrogen bonding would destabilize M—O2 moieties.192 This universal observation of hydrogen bonding in these biological systems is consistent with notions that electron density accumulates on the dioxygen molecule upon coordination. Given the errors associated with atomic positions (at best, ±0.20 Å) the x-ray crystallographic evidence could be equivocal, since hydrogen atoms on the distal imidazole are not observed. There are at least three lines of evidence that support the existence of a specific O2 • • • HN interaction. First, the EPR spectrum of cobalt oxyhemoglobin indicates that the coordinated dioxygen is hydrogen-bonded to something.177-179 Second, and more directly, in the neutron-diffraction structure of oxymyoglobin,189 where hydrogen and especially deuterium nuclides scatter strongly, the imino hydrogen or deuteron was located on the nitrogen atom closest to the coordinated dioxygen, as illustrated in Figure 4.31A. In contrast, in the neutron-diffraction structure of carbonmonoxymyoglobin, the alternative imidazole tautomer was observed (Figure 4.3IB).125,189 The absence of hydrogen bonding of the distal imidazole residue with the coordinated CO molecule is consistent with other lines of evidence that there is little accumulation of electron density on the carbonyl ligand. Third, but less directly, genetically engineered mutants have been produced in which the distal histidine has been replaced by glycine—sperm whale Mb E7His → Gly, and Hb$\alpha$E7His → Gly and HbA $\beta$E7His → Gly.35b,192 For the myoglobin mutant, the O2 binding rate constant at room temperature increases by an order of magnitude, but the dissociation rate constant increases by two orders of magnitude, leading to a decrease in affinity of more than an order of magnitude, as derived from kon/koff. This leads to an estimate of the free energy associated with hydrogen bonding of
$\Delta G = -RTlog \bigg[\frac{P_{1/2}(O_{2})\cdotp Native}{P_{1/2}(O_{2})\cdotp Mutant} \bigg] = 1.5 kcal/mol \ldotp$
In addition, this mutant myoglobin autooxidizes rapidly compared to the native one. On the other hand, the affinity for CO is greatly increased, leading to a value of M for the mutant of 1300, compared to 16 for the native. Thus the distal histidine stabilizes a coordinated O2 ligand by hydrogen bonding and destabilizes a coordinated CO ligand by steric clash.
A similar discrimination is seen for the $\alpha$ chains of the hemoglobin mutant in the binding of the fourth O2 or CO molecule. For the $\beta$ chains little difference is seen relative to the native protein: hydrogen bonding between the distal histidine and the coordinated dioxygen ligand appears to be much weaker in $\beta$ chains, as evidenced by longer N(H) • • • O separations than those seen in the $\alpha$ chains. Comparison of the crystal structures of the native and mutant $\alpha$2($\beta$E7His → Gly)2 structures reveals negligible changes in the distal environment, except for that occasioned by the replacement of —CH2—C3N2H3 (histidine side chain) by —H (glycine side chain).
Studies of hemoglobin mutants where the nonpolar distal residue $\beta$ValE11 (—CH(CH3)2) is replaced by alanine (—CH3), isoleucine (—CH(CH3)CH2CH3), and leucine (—CH2CH(CH3)2) reveal that this valine offers steric hindrance to oxygen binding in the T state.
Whereas the angularly coordinated O2 ligand fits comfortably around the distal histidine, a perpendicular and linear CO moiety cannot. Either the distal histidine rotates out of the way, or the CO tilts off axis, or the Fe—C$\equiv$O group bends, or some combination of these occurs. Notwithstanding the absence of bent M—CO moieties in the inorganic literature, reports of strongly bent M—CO groups appear in the biochemical literature.122-127 The controversy is illustrative of the synergistic interplay of data from models and proteins, and the importance of examining a problem with a miscellany of techniques. The molecular orbital model of ligand-metal interactions presented in Figure 4.22 does not preclude a bent M—C$\equiv$O moiety on symmetry grounds. Groups related to CO can bend; the normally linearly coordinated SCN- moiety has been observed194 to become strongly bent under severe steric stress, with an Fe—N—CS of 140°. Unfortunately, the resolution in protein crystal structures is not sufficient to distinguish unequivocally a linear tilted stereochemistry from a bent one or from a combination of tilt and bend. Studies by the XANES technique have been interpreted in terms of a bent Fe—C$\equiv$O moiety (150°)
$\tag{4.50}$
both in MbCO127 and in the CO adduct of a chelated heme in micelles, the latter being an especially surprising result. From EXAFS data on a number of carbonyl adducts, two interpretations were offered: linear or moderately bent (150°) FeCO moieties for unhindered model systems, and moderately bent or strongly bent (130°) FeCO moieties for hindered synthetic and biological systems.195 In the crystal structure of MbCO at 1.5 Å resolution,122 the CO group is disordered, and Fe—C$\equiv$O angles of 120° and 140° were proposed, although the alternative model of tilted, nearly linear Fe—C$\equiv$O stereochemistry could not be eliminated, and is indeed far more likely to account for the off-axis nature of the oxygen position. Vibrational spectroscopy confirms the existence of two major configurations, and indicates a third minor configuration of the Fe—C$\equiv$O moiety in MbCO.196 An elegant infrared study of the polarization of reattached carbon-monoxide molecules following photolysis of MbCO by linearly polarized light at 10 K gave tilt angles of the CO vector with respect to the heme normal of 15(3)°, 28(2)°, and 33(4)° for the three conformational substates;196b the former two values were confirmed in a similar study at room temperature.196c Note that these studies do not yield the tilt of the Fe—C bond to the heme normal.
In three synthetic compounds with severe steric hindrance, the extent of bending and tilting of the Fe—CO moiety is small. In one nonporphyrinic system the Fe—C$\equiv$O group is bent by 9.4(5)° and tilted by 4.2°.121a In Fe(PocPF)(1,2-Me2Im)(CO) the Fe—C$\equiv$O angle is 172.5(6)° and modest tilting of the Fe—CO group and substantial buckling of the porphyrin ring are apparent.121b In Fe(C2Cap)(1-MeIm)(CO) the two independent Fe—C$\equiv$O angles are 172.9(6)° and 175.9(6)° and modest tilting of the Fe—CO group is again apparent.121d
From a detailed analysis of the force constants describing the vibrational spectroscopy for the Fe—CO moiety, values of 171° for the Fe—CO angle, 9.5° for the tilt, and 11° for porphyrin buckling were calculated for MbCO.121c These results are particularly important, for in a model complex very closely related to Fe(Poc-PF)(1,2-Me2Im)(CO), just mentioned, an EXAFS study195 suggested an Fe—C$\equiv$O bond angle of 127(4)°; that same study ascribed an Fe—C$\equiv$O bond angle of around 130° to MbCO. The structure of carbonmonoxyhemoglobin, Hb(CO)4, now is interpreted in terms of a nearly linear tilted geometry.192 Clearly the geometry of attachment of CO to hemoglobins is perturbed by the surroundings of the ligand-binding site and hence the affinity of hemoglobins for CO is also perturbed. Unfortunately, a clear resolution of the geometry of the Fe—CO moiety in MbCO does not exist yet. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.19%3A_Structures_Relevant_to_Liganded_Hemoglobins.txt |
The equilibrium constant K was defined in Equation (4.3) in terms of the activity ai of component i. The ai may be expressed as a function of concentration as
$a_{i} = \gamma_{i}[i], \tag{4.7}$
where for species i, $\gamma$i is its activity coefficient and [i] is its concentration (strictly molality, but usually as molarity in mol L-1). At infinite dilution $\gamma$i = 1. Provided that the charge and size of species M and MO2 are similar and that O2 forms an ideal solution, then the activities of Equation (4.3) may be approximated by concentrations to give the expression
$K_{c} = \frac{[MO_{2}]}{[M][O_{2}]} \ldotp \tag{4.8}$
However, Equation (4.8) does not permit a direct comparison of the oxygen-binding behavior of one species in some solvent with that of a second in some other solvent. First, for a given partial pressure of dioxygen, the concentration of O2 in the solution varies considerably with temperature and from one solvent to another. Second, reliable measurements of oxygen solubilities are not always available, and it is only relatively recently that oxygen electrodes have been developed to measure directly oxygen concentrations (strictly, activities). However, oxygen-binding measurements are normally made with a solution of M in equilibrium with gaseous dioxygen. At equilibrium the molar Gibbs' free energies (chemical potentials) of the dissolved and gaseous dioxygen are identical—if they are not, gaseous O2 would dissolve, or dissolved O2 would be released. Thus the solvent-dependent quantity [O2] in Equation (4.8) may be replaced by the solvent-independent quantity P(O2), the partial pressure of dioxygen. Under almost all experimental conditions the quantity P(O2) is a very good approximation to the gas-phase activity (fugacity) of dioxygen; hence we obtain for the equilibrium constant*
$K_{p} = \frac{[MO_{2}]}{[M]P(O_{2})} \ldotp \tag{4.9}$
* There has been considerable discussion as to whether Kc (4.8) or Kp (4.9) should be used to compare dioxygen binding under different solvent conditions21-23 We believe that the latter is more appropriate, since for a system al equilibrium, the chemical potential of gaseous O2 must be identical with that of dissolved O219 On the other hand, the concentration of O2 varies from one solvent to another.
It is very convenient to express the affinity as the partial pressure of dioxygen required for half-saturation of the species M, P1/2(O2). Under such conditions, [M] = [MO2], one obtains
$P_{1/2}(O_{2}) = 1/K_{p}, \tag{4.10}$
where P1/2(O2) is usually given in Torr or mm Hg.* As will be detailed shortly, values for P1/2(O2) are typically in the range 0.5 to 40 Torr.
The dioxygen affinity is composed of enthalpic ($\Delta$H) and entropic ($\Delta$S) components, with
$\Delta G^{o} = -RTlnK = \Delta H^{o} - T \Delta S^{o} \ldotp \tag{4.11}$
Within a family of oxygen carriers the values of $\Delta$Sº and $\Delta$H° are usually similar. Large deviations (such as a change of sign) are therefore indicative of a change in the nature of the oxygen-binding process.
* Many authors use the symbol P50 (corresponding to 50% saturation) for P1/2.
a. Non-cooperative Dioxygen Binding
If the oxygen-binding sites M are mutually independent and noninteracting, as in moderately dilute solutions of monomeric molecules, then the concentration of species MO2 as a function of the partial pressure of O2 is generally well fit by a Langmuir isotherm.20 Here a plot of the fractional saturation of dioxygen binding sites, $\theta$, where
$\theta = \frac{[MO_{2}]}{[M] + [MO_{2}]} = \frac{K_{p}P(O_{2})}{1+K_{p}P(O_{2})} \tag{4.12}$
versus P(O2) gives the hyberbolic curve labeled "non-cooperative" in Figure 4.4A.9 Alternatively,24 a plot of log ($\theta$/(1 - $\theta$)) versus log (P(O2)), the so-called "Hill plot," gives a straight line with a slope of unity and an intercept of -log P1/2(O2) (Figure 4.4B). A differential form is shown as the dotted line in Figure 4.4C. Such binding, where the dioxygen sites are independent of each other, is termed non-cooperative.
b. Cooperative Dioxygen Binding
Many dioxygen-binding proteins are not independent monomers, with only one dioxygen-binding site, but oligomeric species with the protein comprising two or more similar subunits. The subunits may be held together by van der Waals' forces or by stronger interactions, such as hydrogen bonds or salt bridges, or even by covalent bonds. For example, most mammalian hemoglobins are tetramers, consisting of two pairs [$\alpha \beta$]2 of myoglobin-like subunits denoted as $\alpha$ and $\beta$. Either none, one, two, three, or all four sites may be occupied by dioxygen. This situation is illustrated schematically in Figure 4.5, which also shows the statistical weighting of each level of saturation, treating the $\alpha$ and $\beta$ subunits as identical. Thus the binding or release of dioxygen at one site may affect the affinity and kinetics of ligand binding and release at a neighboring site. As a result, the saturation curve becomes sigmoidal in shape, as illustrated in Figure 4.4A. The dioxygen binding is cooperative. When cooperativity is positive, the affinity of a vacant site is increased by occupancy of an adjacent one.
This behavior, where the binding of one molecule influences the binding of successive molecules of the same kind, is referred to as a homotropic allosteric interaction. A heterotropic allosteric interaction occurs when the interaction with the protein of a second unlike molecule, for instance, an organic polyphosphate for human hemoglobins, influences the binding of the first molecule (e.g., dioxygen). Such molecules are often termed allosteric effectors. A commonly observed heterotropic allosteric interaction is the Bohr effect, named after the biologist Christian Bohr, father of physicist Niels Bohr. This effect, which relates the change in partial pressure of O2 to a change in pH at constant saturation of binding sites ($\theta$), is related thermodynamically to the Haldane effect, which relates the number of protons released (#H+) with a change in $\theta$ at constant pH (Equation 4.13). A very large Bohr effect, where O2 affinity decreases sharply with pH, is often called the Root effect.25a It is physiologically important for fish such as trout, probably in maintaining buoyancy, but its molecular basis in trout hemoglobin IV remains to be discovered.25b
$\bigg[\frac{\partial (\#H^{+})}{\partial \theta}\bigg]_{pH} = \bigg[\frac{\partial (log P(O_{2}))}{\partial pH}\bigg]_{\theta} \tag{4.13}$
The degree of cooperativity can be characterized in a number of ways. By means of a Hill plot of log ($\theta$/(1 - $\theta$)) versus log (P(O2)), the limiting slopes (which should be unity) at high O2 pressure and low O2 pressure may be extrapolated as shown in Figure 4.4B to log ($\theta$/(1 - $\theta$)) = 0, where $\theta$ = 0.5. Two limiting values for P1/2(O2) are obtained, one characterizing the regime of high partial pressure of dioxygen, where the O2 affinity is high (for the case illustrated of positive cooperativity). The other P1/2(O2) value characterizes the regime of low partial pressure of dioxygen, where affinity is relatively low. This difference in affinities can be converted into a difference between the free-energy change upon O2 binding in the low-affinity state (KpT) and the high-affinity state (KpR) [the designations T and R will be described in subsection d]:
$\delta \Delta G^{o} = -RTln \left(\dfrac{K_{p}^{T}}{K_{p}^{R}}\right) \ldotp \tag{4.14}$
A second way to characterize cooperativity involves fitting the oxygen-binding data at intermediate saturation (0.2 < $\theta$ < 0.8)—that is, about the inflection point in a Hill plot—to the Hill equation
$\frac{\theta}{1-\theta} = K_{p} P^{n}(O_{2})$
or
$log \left(\dfrac{\theta}{1-\theta}\right) = -log(P_{1/2}(O_{2})) + nlog(P(O_{2})) \ldotp \tag{4.15}$
The Hill coefficient (n) is an empirical coefficient that has a value of unity for non-cooperative binding, where Equation (4.15) reduces to the Langmuir isotherm, Equation (4.12). Any number greater than unity indicates positive cooperativity. If O2 binding is an all-or-nothing affair, where dioxygen binding sites are either all occupied or all vacant, n equals the number of subunits in the molecule. The fit is only approximate, since the Hill plot is only approximately linear about the inflection point, as may be seen in Figure 4.4B. A more precise value of n may be obtained by plotting the slope in the Hill plot (n') as a function
$n' = \frac{d\bigg[log\left(\dfrac{\theta}{1-\theta}\right)\bigg]}{d[log(P(O_{2}))]} \tag{4.16}$
of log (P(O2)) (Figure 4.4C). The maximum value of n' is taken as the Hill coefficient n.9 Note that the maximum in this first-derivative plot of the binding curve will occur at P1/2(O2) only if the Hill plot is symmetric about its inflection point. For tetrameric hemoglobins, a maximum Hill coefficient of around 3.0 is seen, and for hemocyanins n may be as high as 9. These values, like P1/2(O2) values, are sensitive to the nature and concentrations of allosteric effectors.
c. Benefits of Cooperative Ligand Binding
In general, oxygen-carrier proteins, being oligomeric, coordinate dioxygen cooperatively, whereas oxygen-storage proteins, being monomeric, do not. Oligomerization and cooperative binding confer enormous physiological benefits to an organism. The first benefit derives directly from oligomerization. Oxygen carriers either form small oligomers that are encapsulated into cells or erythrocytes (such hemoglobins are referred to as intracellular hemoglobins) or associate into large oligomers of 100 or more subunits. Such encapsulation and association reduce by orders of magnitude the number of independent particles in the blood, with consequent reductions in the osmotic pressure of the solution and in strain on vascular membranes.
The second benefit derives from cooperative binding of ligands and the abilities of heterotropic allosteric effectors to optimize exquisitely the oxygen-binding behavior in response to the external and internal environment. The situation is illustrated in general terms in Figure 4.6.9 Most organisms that require O2 live in an environment where the activity of O2 corresponds to about 21 percent of an atmosphere, that is, to about 160 Torr, although usually the effective availability, because of incomplete exchange of gases in the lungs, for example, is around 100 Torr. The concentration of O2 in vertebrate tissues at rest is equivalent to a partial pressure of about 35-40 Torr dioxygen; lower values obtain at times of exertion. Now consider a noncooperative oxygen binder with an affinity expressed as P1/2(O2) of 60 Torr (Figure 4.6, curve a). Then, at 100 Torr the fractional saturation $\theta$ is 0.625. In other words, in a realm of high O2 availability, only 62.5 percent of the oxygen-binding capacity is used, which is not particularly efficient if the organism wished to climb Mt. Everest, where the partial pressure of O2 is less than half that at sea level. In the tissues, where P(O2) = 40 Torr, the fractional saturation is about 40 percent. Thus, only about one third of the coordinated dioxygen is released to the tissues, and total effciency is only 22.5 percent. Consider now a noncooperative oxygen carrier with a much higher affinity, P1/2(O2) = 1.0 Torr (Figure 4.6, curve b). If we assume the same ambient pressure of O2 in the tissues, the fractional saturation is 97.6 percent. Note that at 100 Torr of O2 the carrier is 99.0 percent saturated. In other words, only about 1.4 percent of the available oxygen is delivered.
With an oligomeric protein that binds dioxygen cooperatively, the problem of inefficient and inflexible oxygen delivery disappears. For example, the tetrameric protein hemoglobin has a mean affinity for O2 of P1/2(O2) ≈ 26 Torr at 37 °C and pH 7.4. If hemoglobin bound O2 noncooperatively, then the hyberbolic binding curve (c) in Figure 4.6 would represent the O2 binding. Instead, the observed binding follows curve (d). Since the partial pressure of dioxygen in the lungs and arterial blood of vertebrates is around 100 Torr, but in the tissues and venous blood it is around 40 Torr, then at these pressures a typical myoglobin (P1/2(O2) = 1 Torr) remains effectively saturated. On the other hand, about 25 percent of the available dioxygen can be delivered, even in the absence of myoglobin. With venous blood remaining 75 percent oxygenated, hemoglobin has substantial capacity to deliver more O2 at times of exertion or stress when P(O2) in the tissues falls below 40 Torr.
The net result is that whole blood, which contains about 15 g of hemoglobin per 100 mL, can carry the equivalent of 20 mL of O2 (at 760 Torr) per 100 mL, whereas blood plasma (no hemoglobin) has a carrying capacity of only 0.3 mL of O2 per 100 mL.9
Oxygen binding in vivo is modulated by allosteric effectors that through interaction with the protein change the affinity and degree of cooperativity. For hemoglobin A (adult human hemoglobin), naturally occurring allosteric effectors include the proton, carbon dioxide, and 2,3-diphosphoglycerate (2,3-DPG). Increasing concentrations of these species progressively lower the affinity of free hemoglobin A, thereby enhancing the release of coordinated O2 (Figure 4.6, curve e). For example, 2,3-DPG is part of a subtle mechanism by which dioxygen is transferred from mother to fetus across the placenta. The subunits comprising fetal hemoglobin and adult hemoglobin are slightly different. In the absence of allosteric effectors (referred to as stripped hemoglobin), the oxygen-binding curves are identical. However, 2,3-DPG binds less strongly to fetal hemoglobin than to adult hemoglobin. Thus fetal hemoglobin has a slightly higher affinity for dioxygen, thereby enabling dioxygen to be transferred. The proton and carbon dioxide are part of a short-term feedback mechanism. When O2 consumption outpaces O2 delivery, glucose is incompletely oxidized to lactic acid (instead of CO2). The lactic acid produced lowers the pH, and O2 release from oxyhemoglobin is stimulated (Figure 4.6, curve e). The CO2 produced in respiration forms carbamates with the amino terminals, preferentially of deoxy hemoglobin.
$R-NH_{2} + CO_{2} \rightleftharpoons R-NH-COO^{-} + H^{+}$
Thus hemoglobin not only delivers O2 but also facilitates removal of CO2 to the lungs or gills, where CO2 is exhaled.
d. Models for Cooperativity
The binding of Oz to hemoglobin can be described as four successive equilibria:
$Hb + O_{2} \xrightleftharpoons {K^{(1)}} Hb(O_{2}) \qquad P_{1/2}^{(1)}(O_{2}) = 123[46] Torr$
$Hb(O_{2})_{1} + O_{2} \xrightleftharpoons {K^{(2)}} Hb(O_{2}) _{2} \qquad P_{1/2}^{(2)}(O_{2}) = 30[16] Torr \tag{4.17}$
$Hb(O_{2})_{2} + O_{2} \xrightleftharpoons {K^{(3)}} Hb(O_{2}) _{3} \qquad P_{1/2}^{(3)}(O_{2}) = 33[3.3] Torr$
$Hb(O_{2})_{3} + O_{2} \xrightleftharpoons {K^{(4)}} Hb(O_{2}) _{4} \qquad P_{1/2}^{(4)}(O_{2}) = 0.26[0.29] Torr$
(0.6 mM hemoglobin A, bis(Tris) buffer, pH 7.4, 0.1 M CI-, 2 mM 2,3-DPG, 25 °C. The values in square brackets are affinities in Torr measured in the absence of 2,3- DPG.)
This simple scheme proposed by Adair26 assumes that each of the four binding sites is identical. The P1/2(O2) values given come from fitting the binding curve to this scheme.27 When 2,3-DPG is removed, the affinity of hemoglobin for the first three molecules of Oz is substantially increased, and the degree of cooperativity is lowered (values in square parentheses). For progressively stronger binding, the following inequalities, reflecting the proper statistical weighting illustrated in Figure 4.5, should hold:
$\frac{1}{4}K^{(1)} > \frac{4}{6}K^{(2)} > \frac{6}{4}K^{(3)} > \frac{4}{1}K^{(4)} \tag{4.18}$
The $\frac{6}{4}$ ratio, for example, reflects the six equivalent forms of the doubly and the four equivalent forms of the triply ligated species. In other words, relative to a noncooperative system, at low O2 availability dioxygen release is facilitated; at high O2 availability dioxygen binding is facilitated. The scheme is readily extended to higher orders of oligomerization.
A simple model for analyzing cooperative ligand binding was proposed by Monod, Wyman, and Changeux in 1965, and is usually referred to as the MWC two-state concerted model.28 Molecules are assumed to be in equilibrium between two conformations or quaternary structures, one that has a low ligand affinity and a second that has a high ligand affinity. The low-affinity conformation is often designated the T or tense state, and the high-affinity conformation the R or relaxed state. The equilibrium between the two conformations is characterized by the allosteric constant
$L_{0} = \frac{[R_{0}]}{[T_{0}]} \tag{4.19}$
where the subscript denotes the unliganded Rand T states. The free-energy change upon binding a ligand to the R state, irrespective of saturation, is assumed to be a constant, and the associated equilibrium constant is designated KR; a third constant, KT, characterizes binding to the T state. Figure 4.7 illustrates this model, and introduces the terminology conventionally used.
To a reasonable approximation, the cooperative binding of dioxygen can be summarized by these three parameters, L0, KR, and KT. The Adair constants may be expressed in terms of these parameters:
$K^{(1)} = \frac{(1+L_{0}C)K_{T}}{1+L_{0}}, \qquad K^{(2)} = \frac{(1+L_{0}C^{2})K_{T}}{1+L_{0}}, \tag{4.20}$
$K^{(3)} = \frac{(1+L_{0}C^{3})K_{T}}{1+L_{0}C^{2}}, \qquad K^{(4)} = \frac{(1+L_{0}C^{4})K_{T}}{1+L_{0}C^{3}},$
where C = KR/KT. The fractional saturation is given as
$\theta = \frac{\alpha(1+\alpha)^{3} + L_{0}\alpha C(1+\alpha C)^{3}}{(1+\alpha)^{4} + L_{0}\alpha C(1+\alpha C)^{4}}, \tag{4.21}$
where a = KT[X], and [X] is the concentration of the free ligand (e.g., O2) in the same units (M or Torr) in which KT is expressed. Figure 4.7B illustrates how the allosteric parameters, C = KR/KT and L0 = [R0]/[T0], are extracted from a plot of saturation (as log [$\theta$/(1 - $\theta$)]) versus partial pressure of dioxygen (as log [P(O2)]). Notice how the two-state model (Figure 4.7B) matches very closely the form of the binding curve for hemoglobin (Figure 4.4B). Equations (4.20) and (4.21) may be generalized to an oligomer with n subunits. In the case of hemoglobin, Perutz and coworkers,11 through the determination of the crystal structures of a variety of hemoglobin derivatives, have given subsequently a sound structural basis to the MWC model of two basic quaternary states (see below).
A more exact treatment of ligand-binding data would allow for different affinities for different binding sites (called subunit heterogeneity) and different intrinsic affinities for ligand binding to the R-state conformation compared with the T-state conformation, for each level of ligand saturation—that is, for tertiary structure change within subunits upon ligation. This more exact treatment requires 25 separate equilibrium constants. Statistical thermodynamical approaches exist.29 These explicitly incorporate the different types of subunit interactions that structural studies have revealed, and give improved fits to oxygen-binding data and to the Bohr effect. The key element of two basic quaternary states is preserved, at least for dioxygen binding.29b
For some modified hemoglobins, for example [$\alpha$-Fe(II)2[$\beta$-Mn(III)]2, where in the $\beta$ subunits the heme iron is replaced by Mn(III), there is now strong evidence for three quaternary states,29c with the singly and several of the doubly ligated species having an energy state intermediate between the T (unliganded) and R (fully, triply, and the other doubly liganded) states.
4.21: Ch. 4 References and Abbreviations
Abbreviations
1-MeIm 1-methylimidazole
1,2-Me2Im 1, 2-dimethylimidazole
2-MeIm 2-methylimidazole
2,3-DPG 2,3-diphosphoglycerate
3,4-Me2-Py 3,4-dimethylpyridine
4-t-Bu-Py 4-t-butylpyridine
4-CN-Py 4-cyanopyridine
4-NH2-Py 4-aminopyridine
ai activity of component i
Arg arginine
B general axial base ligand that binds to a metalloporphyrin
Ch chlorocruorin
E1, E-1 activation energy
EDTA ethylenediaminetetraacetic acid
electrochemical potential at unit activity and fugacity
E°' electrochemical potential at physiological pH (7.4) and unit pressure (1 atm = 760 Torr)
Er ertythrocruorin
EtOH ethanol
EXAFS extended x-ray absorption fine structure
G Gibbs free energy
H enthalpy
H2(6,6-CP) cyclophane-strapped porphyrin (see Figure 4.23)
H2(Poc-PF) pocket picket-fence porphyrin
H2Amide-Im basket-handle porphyrin: imidazole base covalently attached to porphyrin with amide straps; amide straps on other side of porphyrin (see Figure 4.23)
H2Bis-Poc meso-tetrakis (2,4,6-triphenylphenyl) porphyrin
H2bzacen N,N-ethylenebis (benzoylacetoninime)
H2C2Cap capped porphyrin; see Figure 4.23
H2Ether-Py basket-handle porphyrin: as for H2Amide-Im except ether straps and pyridine base (see Figure 4.23)
H2MPIX-Im mesoporphyrin IX dimethylester with imidazole covalently attached to porphyrin (see Figure 4.23)
H2PF picket fence porphyrin, meso-tetrakis ($\alpha$,$\alpha$,$\alpha$,$\alpha$-o-pivalamidephenyl) porphyrin (see Figure 4.23)
H2PPIX protoporphyrin IX dimethylester
H2PPIX-Im protoporphyrin IX dimethylester with imidazole covalently attached to porphyrin (see Figure 4.23)
Hb hemoglobin
HbA human adult hemoglobin
HbF human fetal hemoglobin
Hc hemocyanin
His histidine
His203 position on polypeptide chain (203) of a histidine residue
Hr hemerythrin
Im imidazole
k rate constant
Kp, Kc equilibrium constant: concentration of gas expressed in terms of pressure (P) and molarity (M), respectively
Li allosteric constant: equilibrium constant for conformational change of protein with i ligands bound
M general metalloprophyrinato species
M molarity, moles/L
M general metal complex
Mb myoglobin
Me methyl group
met oxidized (e.g. met Hb)
OEP 2,3,7,8,12,13,17,18-octaethylporphyrinato
P pressure, usually in Torr or mm Hg
THF tetrahydrofuran
TPP 5,10,15,20-tetraphenylporphyrinato | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/04%3A_Biological_and_Synthetic_Dioxygen_Carriers/4.20%3A_Thermodynamic_Factors.txt |
I. Introduction1
The major pathway of dioxygen use in aerobic organisms is four-electron reduction to give two molecules of water per dioxygen molecule:2
$O_{2} + 4H^{+} + 4e^{-} \rightarrow 2 H_{2}O \qquad E^{o} = +0.815 V \tag{5.1}$
This reaction represents the major source of energy in aerobic organisms when coupled with the oxidation of electron-rich organic foodstuffs, such as glucose. Biological oxidation of this type is called respiration, and has been estimated to account for 90 percent or more of the dioxygen consumed in the biosphere. It is carried out by means of a series of enzyme-catalyzed reactions that are coupled to ATP synthesis, and the ATP produced is the major source of energy for the organism. The actual site of the reduction of dioxygen in many organisms is the enzyme cytochrome c oxidase.2
Another use of dioxygen in aerobic organisms is to function as a source of oxygen atoms in the biosynthesis of various molecules in metabolic pathways, or in conversions of lipid-soluble molecules to water-soluble forms for purposes of excretion. These reactions are also enzyme-catalyzed, and the enzymes involved are either monooxygenase or dioxygenase enzymes, depending on whether one or both of the oxygen atoms from dioxygen are incorporated in the final organic product. Many of these enzymes are metalloenzymes.2-4
The advantages of life in air are considerable for an aerobic organism as compared to an anaerobic organism, mainly because the powerful oxidizing power of dioxygen can be controlled and efficiently converted to a form that can be stored and subsequently used.5 But aerobic metabolism has its disadvantages as well. The interior of a living cell is a reducing environment, and many of the components of the cell are fully capable thermodynamically of reacting directly with dioxygen, thus bypassing the enzymes that control and direct the beneficial reactions of dioxygen.6 Luckily, for reasons that are discussed below, these reactions generally are slow, and therefore represent minor pathways of biological dioxygen consumption. Otherwise, the cell would just burn up, and aerobic life as we know it would be impossible. Nevertheless, there are small but significant amounts of products formed from nonenzymatic and enzymatic reactions of dioxygen that produce partially reduced forms of dioxygen, i.e., superoxide, O2-, and hydrogen peroxide, H2O2 , in aerobic cells. These forms of reduced dioxygen or species derived from them could carry out deleterious reactions, and enzymes have been identified that appear to protect against such hazards. These enzymes are, for superoxide, the superoxide dismutase enzymes, and, for peroxide, catalase and the peroxidase enzymes. All of these enzymes are metalloenzymes.2-4
Much of the fascination of the subject of biological reactions of dioxygen stems from the fact that the mechanisms of the biological, enzyme-catalyzed reactions are clearly quite different from those of the uncatalyzed reactions of dioxygen or even those of dioxygen reactions catalyzed by a wide variety of nonbiological metal-containing catalysts.7 Investigators believe, optimistically, that once they truly understand the biological reactions, they will be able to design synthetic catalysts that mimic the biological catalysts, at least in reproducing the reaction types, even if these new catalysts do not match the enzymes in rate and specificity. To introduce this topic, therefore, we first consider the factors that determine the characteristics of nonbiological reactions of dioxygen.
Contributors and Attributions
• Joan Selverstone Valentine (University of California, Los Angeles, Department of Chemistry and Biochemistry)
05: Dioxygen Reactions
Description of the Enzymes
Catalase and peroxidase are heme enzymes that catalyze reactions of hydrogen peroxide.94,95 In catalase, the enzymatic reaction is the disproportionation of hydrogen peroxide (Reaction 5.82) and the function of the enzyme appears to be prevention of any buildup of that potentially dangerous oxidant (see the discussion of dioxygen toxicity in Section III).
$2H_{2}O_{2} \xrightarrow{catalase} 2H_{2}O + O_{2} \tag{5.82}$
Peroxidase reacts by mechanisms similar to catalase, but the reaction catalyzed is the oxidation of a wide variety of organic and inorganic substrates by hydrogen peroxide (Reaction 5.83).
$H_{2}O_{2} + AH_{2} \xrightarrow{peroxidase} 2H_{2}O + A \tag{5.83}$
(The catalase reaction can be seen to be a special case of Reaction 5.83 in which the substrate, AH2, is hydrogen peroxide.) Some examples of peroxidases that have been characterized are horseradish peroxidase, cytochrome c peroxidase, glutathione peroxidase, and myeloperoxidase.94,95
X-ray crystal structures have been determined for beef-liver catalase80 and for horseradish peroxidase96 in the resting, high-spin ferric state. In both, there is a single heme b group at the active site. In catalase, the axial ligands are a phenolate from a tyrosyl residue, bound to the heme on the side away from the active-site cavity, and water, bound to heme within the cavity and presumably replaced by hydrogen peroxide in the catalytic reaction. In horseradish peroxidase, the axial ligand is an imidazole from a histidyl residue. Also within the active-site cavity are histidine and aspartate or asparagine side chains that appear to be ideally situated to interact with hydrogen peroxide when it is bound to the iron. These residues are believed to play an important part in the mechanism by facilitating O—O bond cleavage (see Section VI.B below).
Three other forms of catalase and peroxidase can be generated, which are referred to as compounds I, II, and III. Compound I is generated by reaction of the ferric state of the enzymes with hydrogen peroxide. Compound I is green and has spectral characteristics very similar to the FeIV(P•-)(O)+ complex prepared at low temperatures by reaction of ferric porphyrins with single-oxygenatom donors (see Section V.C.1.a). Titrations with reducing agents indicate that it is oxidized by two equivalents above the ferric form. It has been proposed (see 5.84) that the anionic nature of the tyrosinate axial ligand in catalase may serve to stabilize the highly oxidized iron center in compound I of that enzyme,80 and furthermore that the histidyl imidazole ligand in peroxidase may deprotonate, forming imidazolate,52,97 or may be strongly hydrogen bonded,98 thus serving a similar stabilizing function for compound I in that enzyme.
$\tag{5.84}$
Reduction of compound I by one electron produces compound II, which has the characteristics of a normal ferryl-porphyrin complex, analogous to 2, i.e., (L)FeIV(P)(O). Reaction of compound II with hydrogen peroxide produces compound III, which can also be prepared by reaction of the ferrous enzyme with dioxygen. It is an oxy form, analogous to oxymyoglobin, and does not appear to have a physiological function. The reactions producing these three forms and their proposed formulations are summarized in Reactions (5.85) to (5.88).
$Fe^{III}(P)^{+} + H_{2}O_{2} \rightarrow Fe^{IV}(P^{-})(O)^{+} + H_{2}O \tag{5.85}$
$ferric\; form \quad \qquad \qquad Compound\; I \qquad \qquad$
$Fe^{IV}(P^{\cdotp -})(O)^{+} + e^{-} \rightarrow Fe^{IV}(P)(O) \tag{5.86}$
$Compound\; I \qquad \qquad Compound\; II \quad$
$Fe^{IV}(P)(O) + H_{2}O_{2} \rightarrow Fe(P)O_{2} + H_{2}O \tag{5.87}$
$Compound\; II \qquad \qquad Compound\; III \qquad \qquad$
$Fe{II}(P) + O_{2} \rightarrow Fe(P)O_{2} \tag{5.88}$
$ferrous\; form \qquad Compound\; III$
Mechanism
The accepted mechanisms for catalase and peroxidase are described in Reactions (5.89) to (5.94).
$Fe^{III}(P)^{+} + H_{2}O_{2} \rightarrow Fe^{III}(P)(H_{2}O_{2})^{+} \rightarrow Fe^{IV}(P^{\cdotp -})(O)^{+} + H_{2}O \tag{5.89}$
$\qquad \qquad \qquad \qquad \qquad \qquad \qquad \qquad \qquad \qquad \qquad \qquad \qquad Compound\; I$
catalase: $Fe^{IV}(P^{\cdotp -})(O)^{+} + H_{2}O_{2} \rightarrow Fe^{III}(P)^{+} + H_{2}O + O_{2} \tag{5.90}$
$Compound\; I \qquad \qquad \qquad \qquad \qquad \qquad \qquad \qquad$
peroxidase: $Fe^{IV}(P^{\cdotp -})(O)^{+} + AH_{2} \rightarrow Fe^{IV}(P)(O) +HA^{\cdotp} + H^{+} \tag{5.91}$
$Compound\; I \qquad \qquad \qquad \qquad Compound\; II$
$Fe^{IV}(P)(O) + AH_{2} \rightarrow Fe^{III}(P)^{+} +HA^{\cdotp} + OH^{-} \tag{5.92}$
$Compound\; II \qquad \qquad \qquad \qquad \qquad \qquad \qquad \qquad$
$2HA^{\cdotp} \rightarrow A + AH_{2} \tag{5.93}$
or
$2HA^{\cdotp} \rightarrow HA - AH \tag{5.94}$
In the catalase reaction, it has been established by use of H218O2 that the dioxygen formed is derived from hydrogen peroxide, i.e., that O—O bond cleavage does not occur in Reaction (5.90), which is therefore a two-electron reduction of compound I by hydrogen peroxide, with the oxo ligand of the former being released as water. For the peroxidase reaction under physiological conditions, it is believed that the oxidation proceeds in one-electron steps (Reactions 5.91 and 5.92), with the final formation of product occurring by disproportionation (Reaction 5.93) or coupling (Reaction 5.94) of the one-electron oxidized intermediate.94,95
Comparisons of Catalase, Peroxidase, and Cytochrome P-450
The proposal that these three enzymes all go through a similar high-valent oxo intermediate, i.e., 3 or compound I, raises two interesting questions. The first of these is why the same high-valent metal-oxo intermediate gives two very different types of reactions, i.e., oxygen-atom transfer with cytochrome P-450 and electron transfer with catalase and peroxidase. The answer is that, although the high-valent metal-oxo heme cores of these intermediates are in fact very similar, the substrate-binding cavities seem to differ substantially in how much access the substrate has to the iron center. With cytochrome P-450, the substrate is jammed right up against the location where the oxo ligand must reside in the high-valent oxo intermediate. But the same location in the peroxidase enzymes is blocked by the protein structure so that substrates can interact only with the heme edge. Thus oxidation of the substrate by electron transfer is possible for catalase and peroxidase, but the substrate is too far away from the oxo ligand for oxygen-atom transfer.99,124
The second question is about how the the high-valent oxo intermediate forms in both enzymes. For catalase and peroxidase, the evidence indicates that hydrogen peroxide binds to the ferric center and then undergoes heterolysis at the O—O bond. Heterolytic cleavage requires a significant separation of positive and negative charge in the transition state. In catalase and peroxidase, analysis of the crystal structure indicates strongly that amino-acid side chains are situated to aid in the cleavage by stabilizing a charge-separated transition state (Figure 5.14).
In cytochrome P-450, as mentioned in Section V.C.1, no such groups are found in the hydrophobic substrate-binding cavity. It is possible that the cysteinyl axial ligand in cytochrome P-450 plays an important role in O—O bond cleavage, and that the interactions found in catalase and peroxidase that appear to facilitate such cleavage are therefore not necessary. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/05%3A_Dioxygen_Reactions/5.01%3A_Catalase_and_Peroxidase.txt |
Thermodynamics
The reduction potential for the four-electron reduction of dioxygen (Reaction 5.1) is a measure of the great oxidizing power of the dioxygen molecule.8 However, the reaction involves the transfer of four electrons, a process that rarely, if ever, occurs in one concerted step, as shown in Reaction (5.2).
$O_{2} \xrightarrow{e^{-}} O_{2}^{-} \xrightarrow{e^{-}, 2H^{+}} H_{2}O_{2} \xrightarrow{e^{-}, H^{+}} H_{2}O + OH \xrightarrow{e^{-}, H^{+}} 2 H_{2}O \tag{5.2}$
$dioxygen \xrightarrow{e^{-}} superoxide \xrightarrow{e^{-}, 2H^{+}} hydrogen\; peroxide \xrightarrow{e^{-}, H^{+}} water + hyrdoxyl\; radical \xrightarrow{e^{-}, H^{+}} water$
Since most reducing agents can transfer at most one or two electrons at a time to an oxidizing agent, the thermodynamics of the one- and two-electron reductions of dioxygen must be considered in order to understand the overall mechanism.
In aqueous solution, the most common pathway for dioxygen reduction in the absence of any catalyst is one-electron reduction to give superoxide. But this is the least favorable of the reaction steps that make up the full four-electron reduction (see Table 5.1) and requires a moderately strong reducing agent. Thus if only one-electron pathways are available for dioxygen reduction, the low reduction potential for one-electron reduction of O2 to O2- presents a barrier that protects vulnerable species from the full oxidizing power of dioxygen that comes from the subsequent steps. If superoxide is formed (Reaction 5.3), however, it disproportionates quite rapidly in aqueous solution (except at very high pH) to give hydrogen peroxide and dioxygen (Reaction 5.4). The stoichiometry of the overall reaction is therefore that of a net two-electron reduction (Reaction 5.5). It is thus impossible under normal conditions to distinguish one-electron and two-electron reaction pathways for the reduction of dioxygen in aqueous solution on the basis of stoichiometry alone.
$2O_{2} + 2e^{-} \rightarrow 2 O_{2}^{-} \tag{5.3}$
$2O_{2}^{-} + 2 H^{+} \rightarrow H_{2}O_{2} + O_{2} \tag{5.4}$
$O_{2} + 2 e^{-} + 2 H^{+} \rightarrow H_{2}O_{2} \tag{5.5}$
Table 5.1: Standard reduction potentials for dioxygen species in water. a) The standard state used here is unit pressure. If unit activity is used for the standard state of O2, the redox potentials for reactions of that species must be adjusted by +0.17 V.8,9
Reaction E°, V vs. NHE, pH 7, 25 °C
$O_{2} + e^{-} \rightarrow O_{2}^{-}$ -0.33a
$O_{2}^{-} + e^{-} + 2 H^{+} \rightarrow H_{2}O_{2}$ +0.89
$H_{2}O_{2} + e^{-} + H^{+} \rightarrow H_{2}O + OH$ +0.38
$OH + e^{-} + H^{+} \rightarrow H_{2}O$ +2.31
$O_{2} + 2 e^{-} + 2 H^{+} \rightarrow H_{2}O_{2}$ +0.281a
$H_{2}O_{2} + 2 e^{-} + 2 H^{+} \rightarrow 2 H_{2}O$ +1.349
$O_{2} + 4 H^{+} + 4 e^{-} \rightarrow 2 H_{2}O$ +0.815a
The thermodynamics of dioxygen reactions with organic substrates is also of importance in understanding dioxygen reactivity. The types of reactions that are of particular interest to us here are hydroxylation of aliphatic and aromatic C—H bonds and epoxidation of olefins, since these typical reactions of oxygenase enzymes are ones that investigators are trying to mimic using synthetic reagents. Some of the simpler examples of such reactions (plus the reaction of H2 for comparison) are given in the reactions in Table 5.2. It is apparent that all these reactions of dioxygen with various organic substrates in Table 5.2 are thermodynamically favorable. However, direct reactions of dioxygen with organic substrates in the absence of a catalyst are generally very slow, unless the substrate is a particularly good reducing agent. To understand the sluggishness of dioxygen reactions with organic substrates, we must consider the kinetic barriers to these reactions.
Table 5.2: Examples of hydroxylation and epoxidation reactions.
Reaction $\Delta$H in kcal/mol Reference
$CH_{4(g)} + \frac{1}{2} O_{2(g)} \rightarrow CH_{3}OH_{(g)}$ -30 10
$C_{6}H_{6(g)} + \frac{1}{2} O_{2(g)} \rightarrow C_{6}H_{5}OH_{(g)}$ -43 11,12
$C_{6}H_{5}OH_{(g)} + \frac{1}{2} O_{2(g)} \rightarrow C_{6}H_{5}(OH)_{2(g)}$ -42 12,13
$C_{2}H_{4(g)} + \frac{1}{2} O_{2(g)} \rightarrow C_{2}H_{4}O_{(g)}$ -25 10
$C_{5}H_{5}N_{(g)} + \frac{1}{2} O_{2(g)} \rightarrow C_{5}H_{5}NO_{(g)}$ -13 14
$H_{2(g)} + \frac{1}{2} O_{2(g)} \rightarrow H_{2}O_{(g)}$ -58 10
Kinetics
The principal kinetic barrier to direct reaction of dioxygen with an organic substrate arises from the fact that the ground state of the dioxygen molecule is triplet, i.e., contains two unpaired electrons.15,16 Typical organic molecules that are representative of biological substrates have singlet ground states, i.e., contain no unpaired electrons, and the products resulting from their oxygenation also have singlet ground states. Reactions between molecules occur in shorter times than the time required for conversions from triplet to singlet spin. Therefore the number of unpaired electrons must remain the same before and after each elementary step of a chemical reaction. For these reasons, we know that it is impossible for Reaction (5.6) to go in one fast, concerted step.
$\frac{1}{2} \;^{3}O_{2} + \;^{1}X \rightarrow \;^{1}XO \tag{5.6}$
$\qquad \uparrow \uparrow \qquad \downarrow \uparrow \quad \qquad \downarrow \uparrow$
The arrows represent electron spins: $\downarrow \uparrow$ represents a singlet molecule with all electron spins paired; $\uparrow \uparrow$ represents a triplet molecule with two unpaired electrons; and $\uparrow$ (which we will see in Reaction 5.13) represents a doublet molecule, also referred to as a free radical, with one unpaired electron. The pathways that do not violate the spin restriction are all costly in energy, resulting in high activation barriers. For example, the reaction of ground-state triplet dioxygen, i.e.,3O2, with a singlet substrate to give the excited triplet state of the oxygenated product (Reaction 5.7) is spin-allowed, and one could imagine a mechanism in which this process is followed by a slow spin conversion to a singlet product (Reaction 5.8).
$\frac{1}{2} \;^{3}O_{2} + \;^{1}X \rightarrow \;^{3}XO \tag{5.7}$
$\qquad \uparrow \uparrow \qquad \downarrow \uparrow \qquad \quad \uparrow \uparrow$
$\;^{3}XO \xrightarrow{slow} \;^{1}XO \tag{5.38}$
$\; \uparrow \uparrow \qquad \qquad \downarrow \uparrow$
But such a reaction pathway would give a high activation barrier, because the excited triplet states of even unsaturated molecules are typically 40-70 kcal/mol less stable than the ground state, and those of saturated hydrocarbons are much higher.17 Likewise, a pathway in which O2 is excited to a singlet state that then reacts with the substrate would be spin-allowed (Reactions 5.9 and 5.10). The high reactivity of singlet dioxygen, generated by photochemical or chemical means, is well-documented.18,19 However, such a pathway for a reaction of dioxygen, which is initially in its ground triplet state, would also require a high activation energy, since the lowest-energy singlet excited state of dioxygen is 22.5 kcal/mol higher in energy than ground-state triplet dioxygen.15,16
$\;^{3}O_{2} + 22.5\; kcal/mol \rightarrow \;^{1}O_{2} \tag{5.9}$
$\uparrow \uparrow \qquad \qquad \qquad \qquad \qquad \downarrow \uparrow$
$\frac{1}{2} \;^{1}O_{2} + \;^{1}X \rightarrow \;^{1}XO \tag{5.10}$
$\quad \downarrow \uparrow \qquad \downarrow \uparrow \quad \qquad \downarrow \uparrow$
Moreover, the products of typical reactions of singlet-state dioxygen with organic substrates (Reactions 5.11 and 5.12, for example) are quite different in character from the reactions of dioxygen with organic substrates catalyzed by oxygenase enzymes (see Section V):
$\tag{5.11}$
$\tag{5.12}$
One pathway for a direct reaction of triplet ground-state dioxygen with a singlet ground-state organic substrate that can occur readily without a catalyst begins with the one-electron oxidation of the substrate by dioxygen. The products of such a reaction would be two doublets, i.e., superoxide and the oneelectron oxidized substrate, each having one unpaired electron (Reaction 5.13). These free radicals can diffuse apart and then recombine with their spins paired (Reaction 5. 14).
$\;^{3}O_{2} + \;^{1}X \rightarrow \;^{2}O_{2}^{-} + \;^{2}X^{+} \tag{5.13}$
$\uparrow \uparrow \qquad \downarrow \uparrow \quad \quad \uparrow \qquad \uparrow$
$\;^{2}O_{2}^{-} + \;^{2}X^{+} \rightarrow \;^{2}O_{2}^{-} + 2X^{+} \rightarrow \;^{1}XO_{2} \tag{5.14}$
$\uparrow \qquad \uparrow \qquad \qquad \uparrow \qquad \downarrow \qquad \qquad \downarrow \uparrow$
Such a mechanism has been shown to occur for the reaction of dioxygen with reduced flavins shown in Reaction (5.15).20
$\tag{5.15}$
However, this pathway requires that the substrate be able to reduce dioxygen to superoxide, a reaction that requires an unusually strong reducing agent (such as a reduced flavin), since dioxygen is not a particularly strong one-electron oxidizing agent (see Table 5.1 and discussion above). Typical organic substrates in enzymatic and nonenzymatic oxygenation reactions usually are not sufficiently strong reducing agents to reduce dioxygen to superoxide; so this pathway is not commonly observed.
The result of these kinetic barriers to dioxygen reactions with most organic molecules is that uncatalyzed reactions of this type are usually quite slow. An exception to this rule is an oxidation pathway known as free-radical autoxidation.
Free-Radical Autoxidation
The term free-radical autoxidation describes a reaction pathway in which dioxygen reacts with an organic substrate to give an oxygenated product in a free-radical chain process that requires an initiator in order to get the chain reaction started.21 (A free-radical initiator is a compound that yields free radicals readily upon thermal or photochemical decomposition.) The mechanism of free radical autoxidation is as shown in Reactions (5.16) to (5.21).
Initiation: $X_{2} \rightarrow 2X \cdotp \tag{5.16}\] $X \cdotp + RH \rightarrow XH + R \cdotp \tag{5.17}$ Propagation:$R \cdotp + O_{2} \rightarrow ROO \cdotp \tag{5.18}\]
$\downarrow \qquad \uparrow \uparrow \qquad \quad \uparrow$
$ROO \cdotp + RH \rightarrow ROOH + R \cdotp \tag{5.19}$
Termination: R \cdotp + ROO \cdotp \rightarrow ROOR \tag{5.20}\]
$2 ROO \cdotp \rightarrow ROOOOR \rightarrow O_{2} + ROOR \tag{5.21}$
(plus other oxidized products, such as ROOH, ROH, RC(O)R, RC(O)H). This reaction pathway results in oxygenation of a variety of organic substrates, and is not impeded by the spin restriction, because triplet ground-state dioxygen can react with the free radical R• to give a free-radical product ROO•, in a spin-allowed process (Reaction 5.18). It is a chain reaction, since R• is regenerated in Reaction (5.19), and it frequently occurs with long chain lengths prior to the termination steps, resulting in a very efficient pathway for oxygenation of some organic substrates, such as, for example, the oxidation of cumene to give phenol and acetone (Reaction 5.22).22
$\tag{5.22}$
When free-radical autoxidation is used for synthetic purposes, initiators are intentionally added. Common initiators are peroxides and other compounds capable of fragmenting readily into free radicals. Free-radical autoxidation reactions are also frequently observed when no initiator has been intentionally added, because organic substrates frequently contain peroxidic impurities that may act as initiators. Investigators have sometimes been deceived into assuming that a metal-complex catalyzed reaction of dioxygen with an organic substrate occurred by a nonradical mechanism. In such instances, the reactions later proved, upon further study, to be free-radical autoxidations, the role of the metal complex having been to generate the initiating free radicals.
Although often useful for synthesis of oxygenated derivatives of relatively simple hydrocarbons, free-radical autoxidation lacks selectivity and therefore, with more complex substrates, tends to give multiple products. In considering possible mechanisms for biological oxidation reactions used in vivo for biosynthesis or energy production, free-radical autoxidation is not an attractive possibility, because such a mechanism requires diffusion of highly reactive free radicals. Such radicals, produced in the cell, will react indiscriminately with vulnerable sites on enzymes, substrates, and other cell components, causing serious damage.6 In fact, free-radical autoxidation is believed to cause certain deleterious reactions of dioxygen in biological systems, for example the oxidation of lipids in membranes. It is also the process that causes fats and oils to become rancid (Reaction 5.23).23,24
$\tag{5.23}$
How do Enzymes Overcome These Kinetic Barriers?
We see then the reasons that uncatalyzed reactions of dioxygen are usually either slow or unselective. The functions of the metalloenzymes for which dioxygen is a substrate are, therefore, to overcome the kinetic barriers imposed by spin restrictions or unfavorable one-electron reduction pathways, and, for the oxygenase enzymes, to direct the reactions and make them highly specific. It is instructive to consider (1) how these metalloenzymes function to lower the kinetic barriers to dioxygen reactivity, and (2) how the oxygenase enzymes redirect the reactions along different pathways so that very different products are obtained. The first example given below is cytochrome c oxidase. This enzyme catalyzes the four-electron reduction of dioxygen. It overcomes the kinetic barriers to dioxygen reduction by binding dioxygen to two paramagnetic metal ions at the dioxygen binding site, thus overcoming the spin restriction, and by reducing dioxygen in a two-electron step to peroxide, thus bypassing the unfavorable one-electron reduction to form free superoxide. The reaction occurs in a very controlled fashion, so that the energy released by dioxygen reduction can be used to produce ATP. A second example is provided by the catechol dioxygenases, which appear to represent substrate rather than dioxygen activation, and in which dioxygen seems to react with the substrate while it is complexed to the paramagnetic iron center. Another example given below is the monooxygenase enzyme cytochrome PASO, which catalyzes the reaction of dioxygen with organic substrates. It binds dioxygen at the paramagnetic metal ion at its active site, thus overcoming the spin restriction, and then carries out what can be formally described as a multielectron reduction of dioxygen to give a highly reactive high-valent metal-oxo species that has reactivity like that of the hydroxyl radical. Unlike a free hydroxyl radical, however, which would be highly reactive but nonselective, the reaction that occurs at the active site of cytochrome P-450 can be highly selective and stereospecific, because the highly reactive metal-oxo moiety is generated close to a substrate that is bound to the enzyme in such a way that it directs the reactive oxygen atom to the correct position. Thus, metalloenzymes have evolved to bind dioxygen and to increase while controlling its reactivity. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/05%3A_Dioxygen_Reactions/5.02%3A_Chemistry_of_Dioxygen.txt |
Background
Two families of metalloproteins are excellent catalysts for the disproportionation of superoxide (Reaction 5.95).
$2O_{2}^{-} + 2 H^{+} \xrightarrow{SOD} O_{2} + H_{2}O_{2} \tag{5.95}$
These are (1) the copper-zinc superoxide dismutases, CuZnSOD,100-102 found in almost all eukaryotic cells and a very few prokaryotes, and (2) the manganese and iron superoxide dismutases, MnSOD and FeSOD, the former found in the mitochondria of eukaryotic cells, and both found in many prokaryotes.103 Recent studies of bacterial104 and yeast105 mutants that were engineered to contain no superoxide dismutases demonstrated that the cells were unusually sensitive to dioxygen and that the sensitivity to dioxygen was relieved when an SOD gene was reintroduced into the cells. These results indicate that the superoxide dismutase enzymes playa critical role in dioxygen metabolism, but they do not define the chemical agent responsible for dioxygen toxicity (see Section III).
Enzymatic Activity
Several transition-metal complexes have been observed to catalyze superoxide disproportionation; in fact, aqueous copper ion, Cu2+, is an excellent SOD catalyst, comparable in activity to CuZnSOD itself!37 Free aqueous Cu2+ would not itself be suitable for use as an SOD in vivo, however, because it is too toxic (see Section III) and because it binds too strongly to a large variety of cellular components and thus would not be present as the free ion. (Most forms of complexed cupric ion show much less superoxide dismutase activity than the free ion.) Aside from aqueous copper ion, few other complexes are as effective as the SOD enzymes.
Two mechanisms (Reactions 5.96 to 5.99) have been proposed for catalysis of superoxide disproportionation by metal complexes and metalloenzymes.37
Mechanism I: $M^{n+} + O_{2}^{-} \rightarrow M^{(n-1)+} + O_{2} \tag{5.96}\] $M^{(n-1)+} + O_{2}^{-} \rightarrow M^{n+}(O_{2}^{2-}) \xrightarrow{2H^{+}} M^{n+} + H_{2}O_{2} \tag{5.97}$ Mechanism II:$M^{n+} + O_{2}^{-} \rightarrow M^{n+} (O_{2}^{-}) \tag{5.98}\]
$M^{n+} (O_{2}^{-}) + O_{2}^{-} \rightarrow M^{n+}(O_{2}^{2-}) \xrightarrow{2 H^{+}} M^{n} + H_{2}O_{2} \tag{5.99}$
$+O_{2}$
In Mechanism I, which is favored for the SOD enzymes and most redox-active metal complexes with SOD activity, superoxide reduces the metal ion in the first step, and then the reduced metal ion is reoxidized by another superoxide, presumably via a metal-peroxo complex intermediate. In Mechanism II, which is proposed for nonredox metal complexes but may be operating in other situations as well, the metal ion is never reduced, but instead forms a superoxo complex, which is reduced to a peroxo complex by a second superoxide ion. In both mechanisms, the peroxo ligands are protonated and dissociate to give hydrogen peroxide.
Analogues for each of the separate steps of Reactions (5.96) to (5.99) have been observed in reactions of superoxide with transition-metal complexes, thereby establishing the feasibility of both mechanisms. For example, superoxide was shown to reduce CuII(phen)22+ to give Cul(phen)2+ (phen = 1,10-phenanthroline),106 a reaction analogous to Reaction (5.96). On the other hand, superoxide reacts with CuII(tet b)2+ to form a superoxo complex107 (a reaction analogous to Reaction 5.98), presumably because CuII(tet b)2+ is not easily reduced to the cuprous state, because the ligand cannot adjust to the tetrahedral geometry that CuI prefers.53
$\tag{5.100}$
Reaction of superoxide with a reduced metal-ion complex to give oxidation of the complex and release of hydrogen peroxide (analogous to Reaction 5.97) has been observed in the reaction of FeIIEDTA with superoxide.108 Reduction of a CoIII superoxo complex by free superoxide to give a peroxo complex (analogous to Reaction 5.99) has also been observed.109
If a metal complex can be reduced by superoxide and if its reduced form can be oxidized by superoxide, both at rates competitive with superoxide disproportionation, the complex can probably act as an SOD by Mechanism I. Mechanism II has been proposed to account for the apparent catalysis of superoxide disproportionation by Lewis acidic nonredox-active metal ions under certain conditions.37 However, this mechanism should probably be considered possible for redox metal ions and the SOD enzymes as well. It is difficult to distinguish the two mechanisms for redox-active metal ions and the SOD enzymes unless the reduced form of the catalyst is observed directly as an intermediate in the reaction. So far it has not been possible to observe this intermediate in the SOD enzymes or the metal complexes.
Structure
The x-ray crystal structure of the oxidized form of CuZnSOD from bovine erythrocytes shows a protein consisting of two identical subunits held together almost entirely by hydrophobic interactions.100-102 Each subunit consists of a flattened cylindrical barrel of $\beta$-pleated sheet from which three external loops of irregular structure extend (Figure 5.15). The metal-binding region of the protein binds CuII and ZnII in close proximity to each other, bridged by the imidazolate ring of a histidyI side chain. Figure 5.16 represents the metal-binding region. The CuII ion is coordinated to four histidyl imidazoles and a water in a highly distorted square-pyramidal geometry with water at the apical position. The ZnII ion is coordinated to three histidyl imidazoles (including the one shared with copper) and an aspartyl carboxylate group, forming a distorted tetrahedral geometry around the metal ion.
One of the most unusual aspects of the structure of this enzyme is the occurrence of the bridging imidazolate ligand, which holds the copper and zinc ions 6 Å apart. Such a configuration is not unusual for imidazole complexes of metal ions, which sometimes form long polymeric imidazolate-bridged structures.
$\tag{5.101}$
However, no other imidazolate-bridged bi- or polymetallic metalloprotein has yet been identified.
The role of the zinc ion in CuZnSOD appears to be primarily structural. There is no evidence that water, anions, or other potential ligands can bind to the zinc, so it is highly unlikely that superoxide could interact with that site. Moreover, removal of zinc under conditions where the copper ion remains bound to the copper site does not significantly diminish the SOD activity of the enzyme.110 However, such removal does result in a diminished thermal stability, i.e., the zinc-depleted protein denatures at a lower temperature than the native protein, supporting the hypothesis that the role of the zinc is primarily structural in nature.111
The copper site is clearly the site of primary interaction of superoxide with the protein. The x-ray structure shows that the copper ion lies at the bottom of a narrow channel that is large enough to admit only water, small anions, and similarly small ligands (Figure 5.17). In the lining of the channel is the positively charged side chain of an arginine residue, 5 Å away from the copper ion and situated in such a position that it could interact with superoxide and other anions when they bind to copper. Near the mouth of the channel, at the surface of the protein, are two positively charged lysine residues, which are believed to play a role in attracting anions and guiding them into the channel.112 Chemical modification of these lysine or arginine residues substantially diminishes the SOD activity, supporting their role in the mechanism of reaction with superoxide.100-102
The x-ray structural results described above apply only to the oxidized form of the protein, i.e., the form containing Cull. The reduced form of the enzyme containing CuI is also stable and fully active as an SOD. If, as is likely, the mechanism of CuZnSOD-catalyzed superoxide disproportionation is Mechanism I (Reactions 5.96-5.97), the structure of the reduced form is of critical importance in understanding the enzymatic mechanism. Unfortunately, that structure is not yet available.
Enzymatic Activity and Mechanism
The mechanism of superoxide disproportionation catalyzed by CuZnSOD is generally believed to go by Mechanism I (Reactions 5.96-5.97), i.e., reduction of CuII to CuI by superoxide with the release of dioxygen, followed by reoxidation of CuI to Cull by a second superoxide with the release of HO2- or H2O2. The protonation of peroxide dianion, O22-, prior to its release from the enzyme is required, because peroxide dianion is highly basic and thus too unstable to be released in its unprotonated form. The source of the proton that protonates peroxide in the enzymatic mechanism is the subject of some interest.
Reduction of the oxidized protein has been shown to be accompanied by the uptake of one proton per subunit. That proton is believed to protonate the bridging imidazolate in association with the breaking of the bridge upon reduction of the copper. Derivatives with CoII substituted for ZnII at the native zinc site have been used to follow the process of reduction of the oxidized CuII form to the reduced CuI form. The CoIl in the zinc site does not change oxidation state, but acts instead as a spectroscopic probe of changes occurring at the native zinc-binding site. Upon reduction (Reaction 5.102), the visible absorption band due to CoII shifts in a manner consistent with a change occurring in the ligand environment of CoII. The resulting spectrum of the derivative containing CuI in the copper site and CoIl in the zinc site is very similar to the spectrum of the derivative in which the copper site is empty and the zinc site contains ColI. This result suggests strongly that the imidazolate bridge is cleaved and protonated and that the resulting imidazole ligand is retained in the coordination sphere of CoIl (Reaction 5.102).101
$\tag{5.102}$
The same proton is thus an attractive possibility for protonation of peroxide as it is formed in the enzymatic mechanism (Reactions 5.103 and 5.104).
$\tag{5.103}$
$\tag{5.104}$
Attractive as this picture appears, there are several uncertainties about it. For example, the turnover of the enzyme may be too fast for protonation and deprotonation of the bridging histidine to occur.113 Moreover, the mechanism proposed would require the presence of a metal ion at the zinc site to hold the imidazole in place and to regulate the pKa of the proton being transferred. The observation that removal of zinc gives a derivative with almost full SOD activity is thus surprising and may also cast some doubt on this mechanism. Other criticisms of this mechanism have been recently summarized.102
Studies of CuZnSOD derivatives prepared by site-directed mutagenesis are also providing interesting results concerning the SOD mechanism. For example, it has been shown that mutagenized derivatives of human CuZnSOD with major differences in copper-site geometry relative to the wild-type enzyme may nonetheless remain fully active.114 Studies of these and similar derivatives should provide considerable insight into the mechanism of reaction of CuZnSOD with superoxide.
Anions as Inhibitors
Studies of the interaction of CuZnSOD and its metal-substituted derivatives with anions have been useful in predicting the behavior of the protein in its reactions with its substrate, the superoxide anion, O2-.101,102 Cyanide, azide, cyanate, and thiocyanate bind to the copper ion, causing dissociation of a histidyl ligand and the water ligand from the copper.115 Phosphate also binds to the enzyme at a position close to the CuII center, but it apparently does not bind directly to it as a ligand. Chemical modification of Arg-141 with phenylglyoxal blocks the interaction of phosphate with the enzyme, suggesting that this positively charged residue is the site of interaction with phosphate.116
Electrostatic calculations of the charges on the CuZnSOD protein suggest that superoxide and other anions entering into the vicinity of the protein will be drawn toward and into the channel leading down to the copper site by the distribution of positive charges on the surface of the protein, the positively charged lysines at the mouth of the active-site cavity, and the positively charged arginine and copper ion within the active-site region.112 Some of the anions studied, e.g., CN-, F-, N3-,and phosphate, have been shown to inhibit the SOD activity of the enzyme. The source of the inhibition is generally assumed to be competition with superoxide for binding to the copper, but it may sometimes result from a shift in the redox potential of copper, which is known to occur sometimes when an anion binds to copper.100,101
Metal-ion Substitutions
1. SOD Activity
In the example described above, studies of a metal-substituted derivative helped in the evaluation of mechanistic possibilities for the enzymatic reaction. In addition, studies of such derivatives have provided useful information about the environment of the metal-ion binding sites. For example, metal-ion-substituted derivatives of CuZnSOD have been prepared with CuII, CuI, ZnII, AgI, Nill, or CoII bound to the native copper site, and with Znll, Cull, CuI, CoII, HgII, Cdll, Nill, or AgI bound to the native zinc site.100,101,117 The SOD activities of these derivatives are interesting; only those derivatives with copper in the copper site have a high degree of SOD activity, whereas the nature of the metal ion in the zinc site or even its absence has little or no effect.100,101
2. Spectroscopy
Derivatives of CuZnSOD are known with CuII ion bound either to the native copper site or to the native zinc site. The electronic absorption spectra of these derivatives indicate that the ligand environments of the two sites are very different. Copper(II) is a d9 transition-metal ion, and its d-d transitions are usually found in the visible and near-IR regions of the spectrum.53 Copper(II) complexes with coordinated nitrogen ligands are generally found to have an absorption band between 500 and 700 nm, with an extinction coefficient below 100 M-1cm-1. Bands in the absorption spectra of complexes with geometries that are distorted away from square planar tend to be red-shifted because of a smaller d-d splitting, and to have higher extinction coefficients because of the loss of centrosymmetry. Thus the optical spectrum of CuZnSOD with an absorption band with a maximum at 680 nm (14,700 cm-1; see Figure 5.18A) and an extinction coefficient of 155 M-1cm-1 per Cu is consistent with the crystal structural results that indicate that copper(II) is bound to four imidazole nitrogens and a water molecule in a distorted square-pyramidal geometry. Metal-substituted derivatives with CuII at the native copper site but with CoII, CdII, HgII, or NiII substituted for ZnII at the native zinc site all have a band at 680 nm, suggesting that the substitution of another metal ion for zinc perturbs the copper site very little, despite the proximity of the two metal sites. The absorption spectra of native CuZnSOD and these CuMSOD derivatives also have a shoulder at 417 nm (24,000 cm-1; see Figure 5.18A), which is at lower energy than normal imidazole-to-CuII charge-transfer transitions, and has been assigned to an imidazolate-to-CuII charge transfer, indicating that the imidazolate bridge between CuII and the metal ion in the native zinc site is present, as observed in the crystal structure of CuZnSOD. Derivatives with the zinc site empty, which therefore cannot have an imidazolate bridge, are lacking this 417 nm shoulder.
Small but significant changes in the absorption spectrum are seen when the metal ion is removed from the zinc site, e.g., in copper-only SOD (Figure 5.18B). The visible absorption band shifts to 700 nm (14,300 cm-1), presumably due to a change in ligand field strength upon protonation of the bridging imidazolate. In addition, the shoulder at 417 nm has disappeared, again due to the absence of the imidazolate ligand.
The spectroscopic properties due to copper in the native zinc site are best observed in the derivative AgICuSOD, which has AgI in the copper site and CuII in the zinc site (see Figure 5.18C), since the d10 AgI ion is spectroscopically silent. In this derivative, the d-d transition is markedly red-shifted from the visible region of the spectrum into the near-IR, indicating that the ligand environment of CuII in that site is either tetrahedral or five coordinate. The EPR properties of CuII in this derivative are particularly interesting (as discussed below).
The derivative with Cull bound at both sites, CuCuSOD, has a visible-near IR spectrum that is nearly a superposition of the spectra of CuZnSOD and AgICuSOD (see Figure 5.19), indicating that the geometry of Cull in each of these sites is little affected by the nature of the metal ion in the other site.
EPR spectroscopy has also proven to be particularly valuable in characterizing the metal environments in CuZnSOD and derivatives. The EPR spectrum of native CuZnSOD is shown in Figure 5.20A. The g$\parallel$ resonance is split by the hyperfine coupling between the unpaired electron on Cull and the I = $\frac{3}{2}$ nuclear spin of copper. The A$\parallel$ value, 130 G, is intermediate between the larger A$\parallel$ typical of square-planar CuII complexes with four nitrogen donor ligands and the lower A$\parallel$ observed in blue copper proteins (see Chapter 6). The large linewidth seen in the g$\perp$ region indicates that the copper ion is in a rhombic (i.e., distorted) environment. Thus, the EPR spectrum is entirely consistent with the distorted square-pyramidal geometry observed in the x-ray structure.
Removal of zinc from the native protein to give copper-only SOD results in a perturbed EPR spectrum, with a narrower g$\perp$ resonance and a larger A$\parallel$ value (142 G) more nearly typical of Cull in an axial N4 environment (Figure 5.20B). Apparently the removal of zinc relaxes some constraints imposed on the geometry of the active-site ligands, allowing the copper to adopt to a geometry closer to its preferred tetragonal arrangement.
The EPR spectrum due to CuII in the native ZnII site in the AgICuSOD derivative indicates that CuII is in a very different environment than when it is in the native copper site (Figure 5.20C). The spectrum is strongly rhombic, with a low value of A$\parallel$ (97 G), supporting the conclusion based on the visible spectrum that copper is bound in a tetrahedral or five-coordinate environment. This type of site is unusual either for copper coordination complexes or for copper proteins in general, but does resemble the Cull EPR signal seen when either laccase or cytochrome c oxidase is partially reduced (see Figure 5.21).
Partial reduction disrupts the magnetic coupling between these CuII centers that makes them EPR-silent in the fully oxidized protein.
The EPR spectrum of CuCuSOD is very different from that of any of the other copper-containing derivatives (Figure 5.22) because the unpaired spins on the two copper centers interact and magnetically couple across the imidazolate bridge, resulting in a triplet EPR spectrum. This spectrum is virtually identical with that of model imidazolate-bridged binuclear copper complexes.101
Electronic absorption and EPR studies of derivatives of CuZnSOD containing CuII have provided useful information concerning the nature of the metal binding sites of those derivatives. 1H NMR spectra of those derivatives are generally not useful, however, because the relatively slowly relaxing paramagnetic CuII center causes the nearby proton resonances to be extremely broad. This difficulty has been overcome in two derivatives, CuCoSOD and CuNiSOD, in which the fast-relaxing paramagnetic CoIl and NiII centers at the zinc site interact across the imidazolate bridge and increase the relaxation rate of the CuII center, such that well-resolved paramagnetically shifted 1H NMR spectra of the region of the proteins near the two paramagnetic metal centers in the protein can be obtained and the resonances assigned.118,119
The use of 1H NMR to study CuCoSOD derivatives of CuZnSOD in combination with electronic absorption and EPR spectroscopies has enabled investigators to compare active-site structures of a variety of wild-type and mutant CuZnSOD proteins in order to find out if large changes in active-site structure have resulted from replacement of nearby amino-acid residues.120 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/05%3A_Dioxygen_Reactions/5.03%3A_Copper-zinc_Superoxide_Dismutase.txt |
Most of the O2 consumed by aerobic organisms is used to produce energy in a process referred to as "oxidative phosphorylation," a series of reactions in which electron transport is coupled to the synthesis of ATP and in which the driving force for the reaction is provided by the four-electron oxidizing power of O2 (Reaction 5.1). (This subject is described in any standard text on biochemistry and will not be discussed in detail here.) The next to the last step in the electrontransport chain produces reduced cytochrome c, a water-soluble electron-transfer protein. Cytochrome c then transfers electrons to cytochrome c oxidase, where they are ultimately transferred to O2. (Electron-transfer reactions are discussed in Chapter 6.)
Cytochrome c oxidase is the terminal member of the respiratory chain in all animals and plants, aerobic yeasts, and some bacteria.44-46 This enzyme is always found associated with a membrane: the inner mitochondrial membrane in higher organisms or the cell membrane in bacteria. It is a large, complex, multisubunit enzyme whose characterization has been complicated by its size, by the fact that it is membrane-bound, and by the diversity of the four redox metal sites, i.e., two copper ions and two heme iron units, each of which is found in a different type of environment within the protein. Because of the complexity of this system and the absence of detailed structural information, spectroscopic studies of this enzyme and comparisons of spectral properties with O2-binding proteins and with model iron-porphyrin and copper complexes have been invaluable in its characterization.
Spectroscopic Characterization
1. Models
Iron-porphyrin complexes of imidazole are a logical starting point in the search for appropriate spectroscopic models for heme centers in metalloproteins, since the histidyl imidazole side chain is the most common axial ligand bound to iron in such enzymes. Iron-porphyrin complexes with two axial imidazole ligands are known for both the ferrous and ferric oxidation states.47
$\tag{5.32}$
Monoimidazole complexes of iron porphyrins are also known for both the ferrous and the ferric oxidation states. The design of these model complexes has been more challenging than for six-coordinate complexes because of the high affinity of the five-coordinate complexes for a sixth ligand. In the ferrous complex, five coordination has been achieved by use of 2-methylimidazole ligands, as described in Chapter 4. The ferrous porphyrin binds a single 2-methylimidazole ligand, and, because the Fell center is raised out of the plane of the porphyrin ring, the 2-methyl substituent suffers minimal steric interactions with the porphyrin. However, the affinity of the five-coordinate complex for another 2-methylimidazole ligand is substantially lower, because the FeII must drop down into the plane of the porphyrin to form the six-coordinate complexes, in which case the 2-methyl substitutents on both axial ligands suffer severe steric interactions with the porphyrin.51 Using this approach, five-coordinate monoimidazole complexes can be prepared. They are coordinatively unsaturated, and will bind a second axial ligand, such as O2 and CO. They have been extensively studied as models for O2-binding heme proteins such as hemoglobin and myoglobin. Monoimidazole ferrous porphyrins thus designed are high-spin d6 with four unpaired electrons. They are even-spin systems and EPR spectra have not been observed.
Five-coordinate monoimidazole ferric-porphyrin complexes have also been prepared in solution52 by starting with a ferric porphyrin complex with a very poorly coordinating anion, e.g., FeIIIP(SbF6). Addition of one equivalent of imidazole results in formation of the five-coordinate monoimidazole complex (Reaction 5.33).
$Fe^{III}(TPP)(SbF_{6}) + ImH \rightarrow [Fe^{III}(TPP)(ImH)]^{+} + SbF_{6}^{-} \tag{5.33}$
When imidazole is added to ferric-porphyrin complexes of other anionic ligands, e.g., CI-, several equivalents of imidazole are required to displace the more strongly bound anionic ligand; consequently, only six-coordinate complexes are observed (Reaction 5.34).
$FE^{III}(TPP)Cl + 2 ImH \rightarrow [Fe^{III}(TPP)(ImH)_{2}]^{+} + Cl^{-} \tag{5.34}$
Monoimidazole ferric porphyrins are coordinatively unsaturated, readily bind a second axial ligand, and thus are appropriate models for methemoglobin or metmyoglobin. The five-coordinate complexes are high-spin d5, but usually become low-spin upon binding another axial ligand to become six-coordinate.
2. Spectroscopy of the Enzyme
The oxidized form of cytochrome c oxidase contains two CuII and two FeIII heme centers. It can be fully reduced to give a form of the enzyme containing two CuI and two FeII heme centers.44-46 The heme found in cytochrome c oxidase is different from that found in other heme proteins. It is heme a, closely related to heme b, which is found in hemoglobin, myoglobin, and cytochrome P-450, but has one of the vinyl groups replaced by a farnesyl substituent and one of the methyl groups replaced by a formyl substituent (see 5.35).
Each of the four metal centers has a different coordination environment appropriate to its function. Cytochrome a and CuA appear solely to carry out an electron-transfer function without interacting directly with dioxygen. Cytochrome a3 and CuB appear to be part of a binuclear center that acts as the site for dioxygen binding and reduction. A schematic describing the probable nature of these four metal sites within cytochrome oxidase is given in Figure 5.3 and a description of the evidence supporting the formulation of each center then follows.
Cytochrome a in both oxidation states has spectral characteristics that are entirely consistent with a low-spin ferric heme center with two axial imidazole ligands. In its oxidized form, it gives an EPR spectrum with g values44-46 similar to those obtained with model ferric-porphyrin complexes with two axial imidazole ligands 49,50 (see above, Section IV.B.1). Moreover, addition of cyanide anion to the oxidized enzyme or CO to the reduced enzyme does not perturb this center, indicating that cyanide does not bind to the heme, again consistent with a six-coordinate heme. The absence of ligand binding is characteristic of six-coordinate heme sites found in electron-transfer proteins and suggests strongly that cytochrome a functions as an electron-transfer center within cytochrome c oxidase.
CuA is also believed to act as an electron-transfer site. It has quite remarkable EPR spectroscopic characteristics, with g values at g = 2.18, 2.03, and 1.99, and no hyperfine splitting,44-46 resembling more an organic free radical than a typical Cull center (see Figures 5.2 and 5.4A). ENDOR studies of yeast cytochrome c oxidase containing 2H-cysteine or 15N-histidine (from yeast grown with the isotopically substituted amino acids) showed shifts relative to the unsubstituted enzyme, indicating that both of these ligands are bound to CuA.54 But the linear electric-field effect of CuA did not give the patterns characteristic of CuII-histidine complexes, indicating that the unpaired electron is not on the copper ion.55 The current hypothesis about this center is that copper is bonded in a highly covalent fashion to one, or more likely two, sulfur ligands, and that the unpaired electron density is principally on sulfur, i.e., [Cull--SR$\leftrightarrow$CuI- • SR]. Copper-thiolate model complexes with spectroscopic properties similar to CuA have never been synthesized, presumably because such complexes are unstable with respect to disulfide bond formation, i.e, 2 RS• $\rightarrow$ RS-SR. In the enzyme, RS• radicals are presumably constrained in such a way that they cannot couple to form disulfide bonds.
The other heme center, cytochrome a3, does bind ligands such as cyanide to the FellI form and carbon monoxide to the Fell form, indicating that it is either five-coordinate or that it has a readily displaceable ligand. Reaction with CO, for example, produces spectral changes characteristic of a five-coordinate ferrous heme binding CO to give the six-coordinate carbonmonoxy product analogous to MbCO. The cytochrome a3 site is therefore an excellent candidate for O2 binding within cytochrome oxidase.
The EPR spectrum of fully oxidized cytochrome c oxidase might be expected to give signals corresponding to two CuII centers and two ferric heme centers. In fact, all that is observed in the EPR spectrum of the oxidized enzyme is the typical low-spin six-coordinate ferric heme spectrum due to cytochrome G and the EPR signal attributed to CuA (see Figure 5.4A). The fact that signals attributable to cytochrome a3 and CuB are not observed in the EPR spectrum led to the suggestion that these two metal centers are antiferromagnetically coupled.44-46 The measured magnetic susceptibility for the isolated enzyme was found to be consistent with this hypothesis, suggesting that these two metal centers consist of an S = $\frac{1}{2}$ CuII antiferromagnetically coupled through a bridging ligand to a high-spin S = $\frac{5}{2}$ FeIII to give an S = 2 binuclear unit.45 EXAFS measurements indicating a copper-iron separation of 3-4 Å as well as the strength of the magnetic coupling suggest that the metal ions are linked by a single-atom ligand bridge, but there is no general agreement as to the identity of this bridge.45,46
The cytochrome a3-CuB coupling can be disrupted by reduction of the individual metal centers. In this fashion, a g = 6 ESR signal can be seen for cytochrome a3 or g = 2.053, 2.109, and 2.278 signals for CuB. Nitric-oxide binding to CuB also decouples the metals, allowing the g = 6 signal to be seen56 (see Figure 5.4B). Mössbauer spectroscopy also indicates that cytochrome a3 is high-spin in the oxidized as well as the reduced state.57 ENDOR studies suggest that CuB has three nitrogens from imidazoles bound to it with water or hydroxide as a fourth ligand.58 Studies using 15N-labeled histidine in yeast have demonstrated that histidine is a ligand to cytochrome a3.59 All of these features have been incorporated into Figure 5.3.
Mechanism of Dioxygen Reduction
1. Models
Before we consider the reactions of cytochrome c oxidase with dioxygen, it is instructive to review the reactions of dioxygen with iron porphyrins and copper complexes. Dioxygen reacts with ferrous-porphyrin complexes to make mononuclear dioxygen complexes (Reaction 5.36; see preceding chapter for discussion of this important reaction). Such dioxygen complexes react rapidly with another ferrous porphyrin, unless sterically prevented from doing so, to form binuclear peroxo-bridged complexes60,61 (Reaction 5.37). These peroxo complexes are stable at low temperature, but, when the temperature is raised, the O—O bond cleaves and two equivalents of an iron(IV) oxo complex are formed (Reaction 5.38). Subsequent reactions between the peroxo-bridged complex and the FeIV oxo complex produce the$\mu$-oxo dimer (see Reactions 5.39-5.40).
$3Fe^{II}(P) + 3O_{2} \rightarrow 3Fe(P)(O_{2}) \tag{5.36}$
$3Fe(P)(O_{2}) + 3Fe^{II}(P) \rightarrow 3(P)Fe^{III}—O—O—Fe^{III}(P) \tag{5.37}$
$(P)Fe^{III}—O—O—Fe^{III}(P) \rightarrow 2Fe^{IV}(P)(O) \tag{5.38}$
$2Fe^{IV}(P)(O) + 2(P)Fe^{III}—O—O—Fe^{III}(P) \rightarrow 2(P)Fe^{III}—O—Fe^{III}(P) + 2Fe(P)(O_{2}) \tag{5.39}$
$2Fe(P)(O_{2}) \rightarrow 2 Fe^{II}(P) + 2O_{2} \tag{5.40}$
$4Fe^{II}(P) + O_{2} \rightarrow 2(P)Fe^{III}—O—Fe^{III}(P) \tag{5.41}$
The reaction sequence (5.36) to (5.40) thus describes a four-electron reduction of O2 in which the final products, two oxide, O2-, ligands act as bridging ligands in binuclear ferric-porphyrin complexes (Reaction 5.41).
Copper(l) complexes similarly react with dioxygen to form peroxo-bridged binuclear complexes.62 Such complexes do not readily undergo O—O bond cleavage, apparently because the copper(III) oxidation state is not as readily attainable as the Fe(IV) oxidation state in an iron-porphyrin complex. Nevertheless, stable peroxo complexes of copper(II) have been difficult to obtain, because, as soon as it is formed, the peroxo complex either is protonated to give free hydrogen peroxide or is itself reduced by more copper(l) (Reactions 5.42 to 5.46).
$2Cu^{I} + O_{2} \rightarrow Cu^{II}—O—O—Cu^{II} \tag{5.42}$
$Cu^{II}—O—O—Cu^{II} + 2H^{+} \rightarrow 2 Cu^{II} + H_{2}O_{2} \tag{5.43}$
$2Cu^{I} + H_{2}O_{2} + 2H^{+} \rightarrow 2 Cu^{II} + 2 H_{2}O \tag{5.44}$
or Cu^{II}—O—O—Cu^{II} + 2 Cu^{I} + 4 H^{+} \rightarrow 4 Cu^{II} + 2 H_{2}O \tag{5.45}\]
$4 Cu^{I} + O_{2} + 4 H^{+} \rightarrow 4 Cu^{II} + 2 H_{2}O \tag{5.46}$
Recently, however, examples of the long-sought stable binuclear copper(II) peroxo complex have been successfully synthesized and characterized, and interestingly enough, two entirely different structural types have been identified, i.e., $\mu$-1,2 and $\mu$-$\eta^{2}$:$\eta^{2}$ dioxygen complexes63,64 (see 5.47)
$\tag{5.47}$
2. Mechanistic Studies of the Enzyme
A single turnover in the reaction of cytochrome c oxidase involves (1) reduction of the four metal centers by four equivalents of reduced cytochrome c, (2) binding of dioxygen to the partially or fully reduced enzyme, (3) transfer of four electrons to dioxygen, coupled with (4) protonation by four equivalents of protons to produce two equivalents of water, all without the leakage of any substantial amount of potentially harmful partially reduced dioxygen byproducts such as superoxide or hydrogen peroxide.44-46 At low temperatures, the reaction can be slowed down, so that the individual steps in the dioxygen reduction can be observed. Such experiments are carried out using the fully reduced enzyme to which CO has been bound. Binding of CO to the Fell heme center in reduced cytochrome c oxidase inhibits the enzyme and makes it unreactive to dioxygen. The CO-inhibited derivative can then be mixed with dioxygen and the mixture cooled. Photolysis of metal-CO complexes almost always leads to dissociation of CO, and CO-inhibited cytochrome c oxidase is no exception. Photolytic dissociation of CO frees the Fell heme, thereby initiating the reaction with dioxygen, which can then be followed spectroscopically.44-46 Dioxygen reacts very rapidly with the fully reduced enzyme to produce a species that appears to be the dioxygen adduct of cytochrome a3 (Reaction 5.48). Such a species is presumed to be similar to other mononuclear oxyheme derivatives. The dioxygen ligand in this species is then rapidly reduced to peroxide by the nearby CuB, forming what is believed to be a binuclear $\mu$-peroxo species (Reaction 5.49). These steps represent a two-electron reduction of dioxygen to the peroxide level, and are entirely analogous to the model reactions discussed above (Reactions 5.36 to 5.46), except that the binuclear intermediates contain one copper and one heme iron. The $\mu$-peroxo FellI - (O22-) - Cull species is then reduced by a third electron, resulting in cleavage of the O—O bond (Reaction 5.50). One of the oxygen atoms remains with iron in the form of a ferryl complex, i. e., an FeIV oxo, and the other is protonated and bound to copper in the form of a CuII aquo complex.65 Reduction by another electron leads to hydroxo complexes of both the FeIII heme and the CuII centers (Reaction 5.51).65 Protonation then causes dissociation of two water molecules from the oxidized cytochrome a3-CuB center (Reaction 5.52).
$(cyt\; a_{3})\overbrace{Fe^{II} \quad Cu_{B}^{I}} + O_{2} \rightarrow (cyt\; a_{3})\overbrace{Fe^{III}(O_{2}^{-}) \quad Cu_{B}^{I}} \tag{5.48}$
$(cyt\; a_{3})\overbrace{Fe^{III}(O_{2}^{-}) \quad Cu_{B}^{I}} \rightarrow (cyt\; a_{3})\overbrace{Fe^{III}-(O_{2}^{2-}) \quad Cu_{B}^{II}} \tag{5.49}$
$(cyt\; a_{3})\overbrace{Fe^{III}-(O_{2}^{2-})-Cu_{B}^{II}} + e^{-} + 2H^{+} \rightarrow (cyt\; a_{3})\overbrace{Fe^{IV}=O \quad H_{2}O-Cu_{B}^{II}} \tag{5.50}$
$(cyt\; a_{3})\overbrace{Fe^{IV}=O \quad H_{2}O-Cu_{B}^{II}} +e^{-} \rightarrow (cyt\; a_{3})\overbrace{Fe^{III}-(OH^{-}) \quad (HO^{-})-Cu_{B}^{II}} \tag{5.51}$
$(cyt\; a_{3})\overbrace{Fe^{III}-(OH^{-}) \quad (HO^{-})-Cu_{B}^{II}} + 2H^{+} \rightarrow (cyt\; a_{3})\overbrace{Fe^{III} \quad Cu_{B}^{II}} + 2 H_{2}O \tag{5.52}$
Several important questions remain to be resolved in cytochrome c oxidase research. One is the nature of the ligand bridge that links cytochrome a3 and CuB in the oxidized enzyme. Several hypotheses have been advanced (imidazolate, thiolate sulfur, and various oxygen ligands), but then discarded or disputed, and there is consequently no general agreement concerning its identity. However, EXAFS measurements of metal-metal separation and the strength of the magnetic coupling between the two metal centers provide evidence that a single atom bridges the two metals.45,46 Another issue, which is of great importance, is to find out how the energy released in the reduction of dioxygen is coupled to the synthesis of ATP. It is known that this occurs by coupling the electron-transfer steps to a proton-pumping process, but the molecular mechanism is unknown.46 Future research should provide some interesting insights into the mechanism of this still mysterious process. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/05%3A_Dioxygen_Reactions/5.04%3A_Cytochrome_c_Oxidase.txt |
Background
Before we consider the enzymatically controlled reactions of dioxygen in living systems, it is instructive to consider the uncontrolled and deleterious reactions that must also occur in aerobic organisms. Life originally appeared on Earth at a time when the atmosphere contained only low concentrations of dioxygen, and was reducing rather than oxidizing, as it is today. With the appearance of photosynthetic organisms approximately 2.5 billion years ago, however, the conversion to an aerobic, oxidizing atmosphere exposed the existing anaerobic organisms to a gradually increasing level of oxidative stress.25,26 Modern-day anaerobic bacteria, the descendants of the original primitive anaerobic organisms, evolved in ways that enabled them to avoid contact with normal atmospheric concentrations of dioxygen. Modern-day aerobic organisms, by contrast, evolved by developing aerobic metabolism to harness the oxidizing power of dioxygen and thus to obtain usable metabolic energy. This remarkably successful adaptation enabled life to survive and flourish as the atmosphere became aerobic, and also allowed larger, multicellular organisms to evolve. An important aspect of dioxygen chemistry that enabled the development of aerobic metabolism is the relatively slow rate of dioxygen reactions in the absence of catalysts. Thus, enzymes could be used to direct and control the oxidation of substrates either for energy generation or for biosynthesis. Nevertheless, the balance achieved between constructive and destructive oxidation is a delicate one, maintained in aerobic organisms by several means, e.g.: compartmentalization of oxidative reactions in mitochondria, peroxisomes, and chloroplasts; scavenging or detoxification of toxic byproducts of dioxygen reactions; repair of some types of oxidatively damaged species; and degradation and replacement of other species.6
The classification "anaerobic" actually includes organisms with varying degrees of tolerance for dioxygen: strict anaerobes, for which even small concentrations of O2 are toxic; moderate anaerobes, which can tolerate low levels of dioxygen; and microaerophiles, which require low concentrations of O2 for growth, but cannot tolerate normal atmospheric concentrations, i.e., 21 percent O2, 1 atm pressure. Anaerobic organisms thrive in places protected from the atmosphere, for example, in rotting organic material, decaying teeth, the colon, and gangrenous wounds. Dioxygen appears to be toxic to anaerobic organisms largely because it depletes the reducing equivalents in the cell that are needed for normal biosynthetic reactions.6
Aerobic organisms can, of course, live in environments in which they are exposed to normal atmospheric concentrations of O2. Nevertheless, there is much evidence that O2 is toxic to these organisms as well. For example, plants grown in varying concentrations of O2 have been observed to grow faster in lower than normal concentrations of O2.27 E. coli grown under 5 atm of O2 ceased to grow unless the growth medium was supplemented with branched-chain amino acids or precursors. High concentrations of O2 damaged the enzyme dihydroxy acid dehydratase, an important component in the biosynthetic pathway for those amino acids.28 In mammals, elevated levels of O2 are clearly toxic, leading first to coughing and soreness of the throat, and then to convulsions when the level of 5 atm of 100 percent O2 is reached. Eventually, elevated concentrations of O2 lead to pulmonary edema and irreversible lung damage, with obvious damage to other tissues as well.6 The effects of high concentrations of O2 on humans is of some medical interest, since dioxygen is used therapeutically for patients experiencing difficulty breathing, or for those suffering from infection by anaerobic organisms.6
Biological Targets
The major biochemical targets of O2 toxicity appear to be lipids, DNA, and proteins. The chemical reactions accounting for the damage to each type of target are probably different, not only because of the different reactivities of these three classes of molecules, but also because of the different environment for each one inside the cell. Lipids, for example, are essential components of membranes and are extremely hydrophobic. The oxidative damage that is observed is due to free-radical autoxidation (see Reactions 5.16 to 5.21), and the products observed are lipid hydroperoxides (see Reaction 5.23). The introduction of the hydroperoxide group into the interior of the lipid bilayer apparently causes that structure to be disrupted, as the configuration of the lipid rearranges in order to bring that polar group out of the hydrophobic membrane interior and up to the membrane-water interface.6 DNA, by contrast, is in the interior of the cell, and its exposed portions are surrounded by an aqueous medium. It is particularly vulnerable to oxidative attack at the base or at the sugar, and multiple products are formed when samples are exposed to oxidants in vitro.6 Since oxidation of DNA in vivo may lead to mutations, this type of damage is potentially very serious. Proteins also suffer oxidative damage, with amino-acid side chains, particularly the sulfur-containing residues cysteine and methionine, appearing to be the most vulnerable sites.6
Defense and Repair Systems
The biological defense systems protecting against oxidative damage and its consequences are summarized below.
1. Nonenzymatic Oxidant Scavengers
Some examples of small-molecule antioxidants are $\alpha$-tocopherol (vitamin E; 5.24), which is found dissolved in cell membranes and protects them against lipid peroxidation, and ascorbate (vitamin C; 5.25) and glutathione (5.26), which are found in the cytosol of many cells. Several others are known as well.6,29
$\tag{5.24}$
$\tag{5.25}$
$\tag{5.26}$
2. Detoxification Enzymes
The enzymatic antioxidants are (a) catalase and the various peroxidases, whose presence lowers the concentration of hydrogen peroxide, thereby preventing it from entering into potentially damaging reactions with various cell components (see Section VI and Reactions 5.82 and 5.83), and (b) the superoxide dismutases, whose presence provides protection against dioxygen toxicity that is believed to be mediated by the superoxide anion, O2- (see Section VII and Reaction 5.95).
Some of the enzymatic and nonenzymatic antioxidants in the cell are illustrated in Figure 5.1.
3. Systems for Sequestration of Redox-active Metal Ions
Redox-active metal ions are present in the cell in their free, uncomplexed state only in extremely low concentrations. They are instead sequestered by metal-ion storage and transport proteins, such as ferritin and transferrin for iron (see Chapter 1) and ceruloplasmin for copper. This arrangement prevents such metal ions from catalyzing deleterious oxidative reactions, but makes them available for incorporation into metalloenzymes as they are needed.
In vitro experiments have shown quite clearly that redox-active metal ions such as Fe2+/3+ or Cu+/2+ are extremely good catalysts for oxidation of sulfhydryl groups by O2 (Reaction 5.27).30
$4RSH + O_{2} \xrightarrow{M^{n+}} 2RSSR + 2H_{2}O \tag{5.27}$
In addition, in the reducing environment of the cell, redox-active metal ions catalyze a very efficient one-electron reduction of hydrogen peroxide to produce hydroxyl radical, one of the most potent and reactive oxidants known (Reactions 5.28 to 5.30).31
$M^{n+} + Red^{-} \rightarrow M^{(n-1)+} + Red \tag{5.28}$
$M^{(n-1)+} + H_{2}O_{2} \rightarrow M^{n+} + OH^{-} + HO \cdotp \tag{5.29}$
$Red^{-} + H_{2}O_{2} \rightarrow Red + OH^{-} + HO \cdotp \tag{5.30}$
$(Red^{-} = reducing\; agent)$
Binding those metal ions in a metalloprotein usually prevents them from entering into these types of reactions. For example, transferrin, the iron-transport enzyme in serum, is normally only 30 percent saturated with iron. Under conditions of increasing iron overload, the empty iron-binding sites on transferrin are observed to fill, and symptoms of iron poisoning are not observed in vivo until after transferrin has been totally saturated with iron.32 Ceruloplasmin and metallothionein may playa similar role in preventing copper toxicity.6 It is very likely that both iron and copper toxicity are largely due to catalysis of oxidation reactions by those metal ions.
4. Systems for the Repair or Replacement of Damaged Materials
Repair of oxidative damage must go on constantly, even under normal conditions of aerobic metabolism. For lipids, repair of peroxidized fatty-acid chains is catalyzed by phospholipase A2, which recognizes the structural changes at the lipid-water interface caused by the fatty-acid hydroperoxide, and catalyzes removal of the fatty acid at that site. The repair is then completed by enzymatic reacylation.6 Although some oxidatively damaged proteins are repaired, more commonly such proteins are recognized, degraded at accelerated rates, and then replaced.6 For DNA, several multi-enzyme systems exist whose function is to repair oxidatively damaged DNA.6 For example, one such system catalyzes recognition and removal of damaged bases, removal of the damaged part of the strand, synthesis of new DNA to fill in the gaps, and religation to restore the DNA to its original, undamaged state. Mutant organisms that lack these repair enzymes are found to be hypersensitive to O2, H2O2, or other oxidants.6
One particularly interesting aspect of oxidant stress is that most aerobic organisms can survive in the presence of normally lethal levels of oxidants if they have first been exposed to lower, nontoxic levels of oxidants. This phenomenon has been observed in animals, plants, yeast, and bacteria, and suggests that low levels of oxidants cause antioxidant systems to be induced in vivo. In certain bacteria, the mechanism of this induction is at least partially understood. A DNA-binding regulatory protein named OxyR that exists in two redox states has been identified in these systems.33 Increased oxidant stress presumably increases concentration of the oxidized form, which then acts to turn on the transcription of the genes for some of the antioxidant enzymes. A related phenomenon may occur when bacteria and yeast switch from anaerobic to aerobic metabolism. When dioxygen is absent, these microorganisms live by fermentation, and do not waste energy by synthesizing the enzymes and other proteins needed for aerobic metabolism. However, when they are exposed to dioxygen, the synthesis of the respiratory apparatus is turned on. The details of this induction are not known completely, but some steps at least depend on the presence of heme, the prosthetic group of hemoglobin and other heme proteins, whose synthesis requires the presence of dioxygen.34
Molecular Mechanisms of Dioxygen Toxicity
What has been left out of the preceding discussion is the identification of the species responsible for oxidative damage, i.e., the agents that directly attack the various vulnerable targets in the cell. They were left out because the details of the chemistry responsible for dioxygen toxicity are largely unknown. In 1954, Rebeca Gerschman formulated the "free-radical theory of oxygen toxicity" after noting that tissues subjected to ionizing radiation resemble those exposed to elevated levels of dioxygen.35 Fourteen years later, Irwin Fridovich proposed that the free radical responsible for dioxygen toxicity was superoxide, O2-, based on his identification of the first of the superoxide dismutase enzymes.36 Today it is still not known if superoxide is the principal agent of dioxygen toxicity, and, if so, what the chemistry responsible for that toxicity is.6
There is no question that superoxide is formed during the normal course of aerobic metabolism,121 although it is difficult to obtain estimates of the amount under varying conditions, because, even in the absence of a catalyst, superoxide disproportionates quite rapidly to dioxygen and hydrogen peroxide (Reaction 5.4) and therefore never accumulates to any great extent in the cell under normal conditions of pH.37
One major problem in this area is that a satisfactory chemical explanation for the purported toxicity of superoxide has never been found, despite much indirect evidence from in vitro experiments that the presence of superoxide can lead to undesirable oxidation of various cell components and that such oxidation can be inhibited by superoxide dismutase.38 The mechanism most commonly proposed is production of hydroxyl radicals via Reactions (5.28) to (5.30) with Red- = O2-, which is referred to as the "Metal-Catalyzed Haber-Weiss Reaction". The role of superoxide in this mechanism is to reduce oxidized metal ions, such as Cu2+ or Fe3+, present in the cell in trace amounts, to a lower oxidation state.37 Hydroxyl radical is an extremely powerful and indiscriminate oxidant. It can abstract hydrogen atoms from organic substrates, and oxidize most reducing agents very rapidly. It is also a very effective initiator of free-radical autoxidation reactions (see Section II.C above). Therefore, reactions that produce hydroxyl radical in a living cell will probably be very deleterious.6
The problem with this explanation for superoxide toxicity is that the only role played by superoxide here is that of a reducing agent of trace metal ions. The interior of a cell is a highly reducing environment, however, and other reducing agents naturally present in the cell such as, for example, ascorbate anion can also act as Red- in Reaction (5.28), and the resulting oxidation reactions due to hydroxyl radical are therefore no longer inhibitable by SOD.39
Other possible explanations for superoxide toxicity exist, of course, but none has ever been demonstrated experimentally. Superoxide might bind to a specific enzyme and inhibit it, much as cytochrome oxidase is inhibited by cyanide or hemoglobin by carbon monoxide. Certain enzymes may be extraordinarily sensitive to direct oxidation by superoxide, as has been suggested for the enzyme aconitase, an iron-sulfur enzyme that contains an exposed iron atom.122 Another possibility is that the protonated and therefore neutral form of superoxide, HO2, dissolves in membranes and acts as an initiator of lipid peroxidation. It has also been suggested that superoxide may react with nitric oxide, NO, in the cell producing peroxynitrite, a very potent oxidant.123 One particularly appealing mechanism for superoxide toxicity that has gained favor in recent years is the "Site-Specific Haber-Weiss Mechanism."40,41 The idea here is that traces of redox-active metal ions such as copper and iron are bound to macromolecules under normal conditions in the cell. Most reducing agents in the cell are too bulky to come into close proximity to these sequestered metal ions. Superoxide, however, in addition to being an excellent reducing agent, is very small, and could penetrate to these metal ions and reduce them. The reduced metal ions could then react with hydrogen peroxide, generating hydroxyl radical, which would immediately attack at a site near the location of the bound metal ion. This mechanism is very similar to that of the metal complexes that cause DNA cleavage; by reacting with hydrogen peroxide while bound to DNA, they generate powerful oxidants that react with DNA with high efficiency because of their proximity to it (see Chapter 8).
Although we are unsure what specific chemical reactions superoxide might undergo inside of the cell, there nevertheless does exist strong evidence that the superoxide dismutases play an important role in protection against dioxygen-induced damage. Mutant strains of bacteria and yeast that lack superoxide dismutases are killed by elevated concentrations of dioxygen that have no effect on the wild-type cells. This extreme sensitivity to dioxygen is alleviated when the gene coding for a superoxide dismutase is reinserted into the cell, even if the new SOD is of another type and from a different organism.42,43
Summary of Dioxygen Toxicity
In summary, we know a great deal about the sites that are vulnerable to oxidative damage in biological systems, about the agents that protect against such damage, and about the mechanisms that repair such damage. Metal ions are involved in all this chemistry, both as catalysts of deleterious oxidative reactions and as cofactors in the enzymes that protect against and repair such damage. What we still do not know at this time, however, is how dioxygen initiates the sequence of chemical reactions that produce the agents that attack the vulnerable biological targets in vivo. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/05%3A_Dioxygen_Reactions/5.05%3A_Dioxygen_Toxicity.txt |
Metal-containing monooxygenase enzymes are known that contain heme iron, nonheme iron, or copper at their active sites.2 For most of these enzymes, there is only limited information about the nature of the active site and the mode of interaction with dioxygen or substrates. But there are three monooxygenase enzymes that strongly resemble well-characterized reversible dioxygen-carrying proteins (see preceding chapter), suggesting that dioxygen binding to the metalloenzyme in its reduced state is an essential first step in the enzymatic mechanisms, presumably followed by other steps that result in oxygenation of substrates. The enzymes are:
1. cytochrome P-450,73 a heme-containing protein whose active site resembles the dioxygen-binding sites of myoglobin or hemoglobin in many respects, except that the axial ligand to iron is a thiolate side chain from cysteine rather than an imidazole side chain from histidine;
2. tyrosinase,74 which contains two copper ions in close proximity in its active site and which has deoxy, oxy, and met states that closely resemble comparable states of hemocyanin in their spectroscopic properties; and
3. methane monooxygenase,75,76 which contains two nonheme iron ions in close proximity and which resembles hemerythrin in many of its spectroscopic properties.
In addition to these three, there are also monooxygenase enzymes containing single nonheme iron77 or copper ions,78 or nonheme iron plus an organic cofactor such as a reduced pterin at their active sites.79 Just as with the dioxygenase enzymes, we do not know how similar the mechanisms of the different metal-containing monooxygenase enzymes are to one another. The enzyme for which we have the most information is cytochrome P-450, and we will therefore focus our discussion on that system. Speculations about the mechanisms for the other systems are discussed at the end of this section.
1. Cytochrome P-450
Cytochrome P-450 enzymes are a group of monooxygenase enzymes that oxygenate a wide variety of substrates.73 Examples of such reactions are:
1. hydroxylation of aliphatic compounds (Reaction 5.59);
2. hydroxylation of aromatic rings (Reaction 5.60);
3. epoxidation of olefins (Reaction 5.61);
4. amine oxidation to amine oxides (Reaction 5.62);
5. sulfide oxidation to sulfoxides (Reaction 5.63); and
6. oxidative dealkylation of heteroatoms (for example, Reaction 5.64).
$\tag{5.59}$
$\tag{5.60}$
$\tag{5.61}$
$\tag{5.62}$
$\tag{5.63}$
$Ph—O-CH_{3} \longrightarrow Ph—OH + HCHO \tag{5.64}$
Some of these reactions have great physiological significance, because they represent key transformations in metabolism, as in lipid metabolism and biosynthesis of corticosteroids, for example.73 Cytochrome P-450 is also known to catalyze the transformation of certain precarcinogens such as benzpyrene into their carcinogenic forms.73
Many of the P-450 enzymes have been difficult to characterize, because they are membrane-bound and consequently relatively insoluble in aqueous solution. However, cytochrome P-450cam, which is a component of the camphor 5-monooxygenase system isolated from the bacterium Pseudomonas putida, is soluble and has been particularly useful as the subject of numerous spectroscopic and mechanistic studies, as well as several x-ray crystallographic structure determinations.80 This enzyme consists of a single polypeptide chain, mainly $\alpha$-helical, with a heme b group (Fe-protoporphyrin IX) sandwiched in between two helices, with no covalent attachments between the porphyrin ring and the protein. One axial ligand complexed to iron is a cysteinyl thiolate. In the resting state, the iron is predominantly low-spin FellI, probably with a water as the other axial ligand. When substrate binds to the resting enzyme, the spin state changes to high-spin, and the non-cysteine axial ligand is displaced. The enzyme can be reduced to an FeII state, which is high-spin, and resembles deoxyhemoglobin or myoglobin in many of its spectroscopic properties. This ferrous form binds dioxygen to make an oxy form or carbon monoxide to make a carbonyl form. The CO derivative has a Soret band (high-energy $\pi$-$\pi$* transition of the porphyrin ring) at 450 nm, unusually low energy for a carbonyl derivative of a heme protein because of the presence of the axial thiolate ligand. This spectroscopic feature aids in the isolation of the enzyme and is responsible for its name.
a. "Active Oxygen"
Camphor 5-monooxygenase is a three-component system, consisting of cytochrome P-450cam and two electron-transfer proteins, a flavoprotein, and an iron-sulfur protein (see Chapters 6 and 7). The role of the electron-transfer proteins is to deliver electrons to the P-450 enzyme, but these may be replaced in vitro by other reducing agents. The reaction sequence is in Figure 5.10.
For cytochrome P-450, the question that is possibly of greatest current interest to the bioinorganic chemist is just what mechanism enables activation of dioxygen and its reaction with substrate. It seems clear that dioxygen binds to the ferrous state of the enzyme-substrate complex, and that the resulting oxy ligand, which presumably is similar to the oxy ligand in oxyhemoglobin and oxymyoglobin, is not sufficiently reactive to attack the bound substrate. The oxy form is then reduced and the active oxidant is generated, but the nature of the active oxidant has not been deduced from studies of the enzyme itself, nor has it been possible to observe and characterize intermediates that occur between the time of the reduction and the release of product. Three species are potential candidates for "active oxygen," the oxygen-containing species that attacks the substrate, in cytochrome P-450. They are:
1. a ferric peroxo, 1a, or hydroperoxo complex, 1b, formed from one-electron reduction of the oxy complex (Reaction 5.65);
2. an iron(IV) oxo complex, 2, formed by homolytic O—O bond cleavage of a ferric hydroperoxo complex (Reaction 5.66); and
3. a complex at the oxidation level of an iron(V) oxo complex, 3, formed by heterolytic O—O bond cleavage of a ferric hydroperoxo complex (Reaction 5.67).
The hydroxyl radical, HO•, although highly reactive and capable of attacking P-450 substrates, is considered to be an unlikely candidate for "active oxygen" because of the indiscriminate character of its reactivity.
$Fe^{II}P + O_{2} \rightarrow FePO_{2} \xrightarrow{e^{-}} [Fe^{III}P(O_{2}^{2-})]^{-} \xrightarrow{H^{+}} Fe^{III}P(O_{2}H^{-}) \tag{5.65}$
$\qquad \qquad \qquad \qquad \qquad \bf{1a} \qquad \qquad \qquad \qquad \bf{1b}$
$Fe^{III}P(O_{2}H^{-}) \rightarrow Fe^{IV}P(O) + HO \cdotp \tag{5.66}$
$\bf{1b} \qquad \qquad \qquad \bf{2}$
$Fe^{III}P(O_{2}H^{-}) \rightarrow [Fe^{V}(P^{2-})(O)^{+} \leftrightarrow Fe^{IV}(P^{-})(O)^{+}] +HO^{-} \tag{5.67}$
$\bf{1b} \qquad \qquad \qquad\qquad \qquad \qquad \bf{3} \qquad \qquad \qquad$
(P2- = porphyrin ligand; P- = one-electron oxidized porphyrin ligand)
An iron(V) oxo complex (or a related species at the same oxidation level), 3, formed via Reaction (5.67), is the favored candidate for "active oxygen" in cytochrome P-450.81 This conclusion was initially drawn from studies of reactions of the enzyme with alkylhydroperoxides and single-oxygen-atom donors. Single-oxygen-atom donors are reagents such as iodosylbenzene, OIPh, and periodate, IO4-,capable of donating a neutral oxygen atom to an acceptor, forming a stable product in the process (here, iodobenzene, IPh, and iodate, IO3-). It was discovered that ferric cytochrome P-450 could catalyze oxygenation reactions using organic peroxides or single-oxygen-atom donors in place of dioxygen and reducing agents. Usually the same substrates would give the identical oxygenated product. This reaction pathway was referred to as the "peroxide shunt" (see Figure 5.10). The implication of this discovery was that the same form of "active oxygen" was generated in each reaction, and the fact that single-oxygen-atom donors could drive this reaction implied that this species contained only one oxygen atom, i.e., was generated subsequent to O—O bond cleavage. The mechanism suggested for this reaction was Reactions (5.68) and (5.69).
$Fe(III)P^{+} + OX \rightarrow \textbf{3} + X \tag{5.68}$
$\textbf{3} + substrate \rightarrow Fe(III)P^{+} + substrate(O) \tag{5.69}$
b. Metalloporphyrin Model Systems
Studies of the reactivities of synthetic metalloporphyrin complexes in oxygen-transfer reactions and characterization of intermediate species observed during the course of such reactions have been invaluable in evaluating potential intermediates and reaction pathways for cytochrome P-450. Logically, it would be most desirable if one could mimic the enzymatic oxygenation reactions of substrates using iron porphyrins, dioxygen, and reducing agents. However, studies of such iron-porphyrin-catalyzed reactions have failed to produce meaningful results that could be related back to the P-450 mechanism. This is perhaps not surprising, since the enzyme system is designed to funnel electrons into the iron-dioxygen-substrate complex, and thus to generate the active oxidant within the confines of the enzyme active site in the immediate proximity of the bound substrate. Without the constraints imposed by the enzyme, however, iron porphyrins generally will either (1) catalyze the oxidation of the reducing agent by dioxygen, leaving the substrate untouched, or (2) initiate free-radical autoxidation reactions (see Section II.C). A different approach was suggested by the observation of the peroxide shunt reaction (Reactions 5.68 and 5.69) using organic peroxides or single-oxygen-atom donors, and the earliest successful studies demonstrated that Fe(TPP)Cl (TPP = tetraphenylporphyrin) would catalyze the epoxidation of olefins and the hydroxylation of aliphatic hydrocarbons by iodosylbenzene81(Reactions 5.70 and 5.71).
$\tag{5.70}$
$\tag{5.71}$
Reactions (5.70) and (5.71) were postulated to occur via an iron-bound oxidant such as 3 in Reaction (5.67). This hypothesis was tested by studying the reaction of dioctyl Fe(PPIX)CI with iodosylbenzene, which resulted in 60 percent hydroxylation at positions 4 and 5 on the hydrocarbon tail (see 5.72), positions for which there is no reason to expect increased reactivity except for the fact that those particular locations are predicted from molecular models to come closest to the iron center when the tail wraps around the porphyrin molecule.82
Visible absorption spectra of porphyrin complexes are due largely to $\pi$-$\pi$* transitions of the porphyrin ligand. The bright green color is unusual for iron-porphyrin complexes, which are usually red or purple. (However, this green color has been seen for compound I of catalase and peroxidases; see Section VI below.) The unusually long-wavelength visible absorption bands that account for the green color result from the fact that the porphyrin ring has been oxidized by one electron. Similar visible absorption bands can be seen, for example, in other oxidized porphyrin complexes, such as CoIlI(P•-)+, formed by two-electron oxidation of ColI(P2-)(see 5.73).84
$\tag{5.73}$
Oxidized porphyrin ligands also give characteristic proton NMR spectra, which are seen for the green porphyrin complex as well.81,83
Magnetic measurements indicate that the green porphyrin complex contains three unpaired electrons. Detailed analysis of the Mössbauer spectra has indicated that the two unpaired electrons on the FeIV ion are strongly ferromagnetically coupled to the unpaired electron on the porphyrin, accounting for the resulting S = $\frac{5}{3}$ state.81,83
$\tag{5.74}$
Studies of the reactions of this species with P-450-type substrates demonstrate that this species is reactive enough to make it an attractive candidate for" active oxygen" in the enzymatic mechanism.81,83
Synthetic analogues for two of the other candidates for "active oxygen" have also been synthesized and their reactivities assessed. For example, FellI and MnIII-porphyrin peroxo complexes analogous to 1a in Reaction (5.65) have been synthesized. The x-ray crystal structure of the Mn complex shows that the peroxo ligand is bound to the metal in a triangular, side-on fashion (see 5.75). The Fe complex is believed to have a similar structure.85,86
$\tag{5.75}$
Studies of this species indicate that 1a in Reaction (5.65) would not have the requisite reactivity to be a candidate for "active oxygen" in the cytochrome P-450 mechanism, since it will not even oxidize triphenylphosphine, PPh3, to triphenylphosphine oxide, OPPh3 , one of the more facile oxygenation reactions known.87 Attempts to examine the protonated form, 1b in Reaction (5.65), however, indicate that it is highly unstable, and its reactivity has not yet been thoroughly examined.87 FeIV-oxo-porphyrin complexes analogous to 2 in Reaction (5.66) have also been prepared in solution and characterized by NMR.60,61 Such complexes will react with PPh3 to give OPPh3, but are relatively unreactive with olefins and totally unreactive with saturated hydrocarbons. Thus 2 is also ruled out as a candidate for "active oxygen" in P-450 mechanisms.
These reactivity studies, and the observation of the peroxide shunt described above, indicate that FeV(P2-)(O)+ or FelV(P-)(O)+ is the most likely candidate for "active oxygen." These two formulations are, of course, isoelectronic, and it is tempting to conclude that the latter is the more likely formulation of the enzymatic intermediate. However, it is important to remember that the model systems lack the axial cysteinylligand present in cytochrome P-450. The effect of the relatively easily oxidized sulfur ligand on the electron distribution within that intermediate is not known, since model systems for high-valent iron-oxo complexes containing axial thiolate ligands have not been synthesized.
The mechanism of reactions of the high-valent oxo complex 3 in Reaction (5.67) with a variety of substrates is an area of active interest.81,88 Such studies are generally carried out by generation of the species in situ from the reaction of a ferric porphyrin with a single-oxygen-atom donor, such as a peracid or iodosylbenzene.89 In hydroxylation reactions of aliphatic hydrocarbons, the initial step appears to be abstraction of a hydrogen atom from the substrate to form a substrate radical and an FeIV hydroxide complex held together in a cage created by the enzyme active site so that they cannot diffuse away from each other (Reaction 5.76). This step is then followed by recombination of the OH fragment with the substrate radical to make the hydroxylated product (Reaction 5.77). This mechanism is referred to as the "oxygen rebound mechanism."83
$\tag{5.76}$
$\tag{5.77}$
The radical character of the intermediates formed in this reaction is supported by the observation that such reactions carried out using synthetic porphyrins and single-oxygen-atom donors in the presence of BrCCl3 give substantial amounts of alkyl bromides as products, a result that is consistent with radical intermediates and inconsistent with either carbanion or carbonium-ion intermediates.83
In the enzymatic reactions themselves, there is also strong evidence to support a stepwise mechanism involving free-radical intermediates. For example, cytochrome P-450cam gives hydroxylation of d-camphor only in the 5-exo position, but deuterium-labeling studies show that either the 5-exo or the 5-endo hydrogen is lost (Reaction 5.78).88
$\tag{5.78}$
Such results are obviously inconsistent with a concerted mechanism in which the oxygen atom would be inserted into the 5-exo C—H bond in one step; so there would be no chance for the hydrogens in the two positions to exchange. (Remember that alcohol protons exchange rapidly with water and therefore are not expected to remain deuterated when the reaction is carried out in H2O.)
The crystal structure of reduced cytochrome P-450cam with CO bound to the iron and the substrate camphor bound90 adjacent to it has been examined and compared with the crystal structure of the oxidized enzyme with camphor bound. The former is expected to be similar in structure to the less-stable oxy complex. The comparison shows that the substrate camphor is closer to the iron center in the oxidized enzyme. It is therefore possible that a similar movement of the substrate occurs during the catalytic reaction after either a 5-exo or a 5-endo hydrogen is abstracted, and that the new position of the camphor molecule then restricts the hydroxylation step to the 5-exo position. It is interesting to note that the 5-exo position on the camphor that is hydroxylated is held in very close proximity to the FeIII center, and therefore to the presumed location of the oxo ligand in the high-valent oxo intermediate in the structure of the ferric enzyme plus camphor derivative (Figure 5.12). Crystal structures of the ferric form of cytochrome P-450cam with norcamphor and adamantanone bound in place of camphor have also been determined.90 These alternative substrates are smaller than camphor, and appear to fit more loosely than camphor. It is therefore reasonable to assume that they "rattle around" to a certain extent in the substrate binding site, which probably accounts for the less-specific pattern of hydroxylation observed for these alternative substrates.
Mechanisms for olefin epoxidations catalyzed either by the enzyme or by model porphyrin complexes are not as well understood as those for hydroxylation of aliphatic hydrocarbons. Some of the possibilities that have been proposed88,91 are represented schematically in Figure 5.13.
c. O—O Bond Cleavage
The evidence is persuasive that the "active oxygen" species that attacks substrate in cytochrome P-450 is a high-valent iron-oxo complex. However, the mechanism of formation of that species in the catalytic reaction with dioxygen is less well-understood. Heterolytic O—O bond cleavage of a ferric porphyrin hydroperoxide complex, 1b (Reaction 5.67), is the logical and anticipated route, but it has not yet been unequivocally demonstrated in a model complex.92,93 The catalase and peroxidase enzymes catalyze heterolytic O—O bond cleavage in reactions of hydrogen peroxide, but in them the active sites contain amino-acid side chains situated to facilitate the developing charge separation that occurs in heterolytic cleavage (see Section VI). The crystal structure of cytochrome P-450cam shows no such groups in the activesite cavity, nor does it give any clue to the source of a proton to protonate the peroxide ligand when it is produced.80 Also, we have little experimental evidence concerning possible roles that the cysteinyl sulfur axial ligand might play in facilitating O—O bond cleavage. These issues remain areas of active interest for researchers interested in cytochrome P-450 mechanisms.
Other Metal-containing Monooxygenase Enzymes
As mentioned above, much less is known about the structural characteristics and mechanisms of the nonheme metal-containing monooxygenase enzymes. From the similarities of the overall stoichiometries of the reactions and the resemblance of some of the enzymes to dioxygen-binding proteins, it is likely that the initial steps are the same as those for cytochrome P-450, i.e., dioxygen binding followed by reduction to form metal-peroxide or hydroperoxide complexes. It is not obvious that the next step is the same, however (i.e., O—O bond cleavage to form a high-valent metal-oxo complex prior to attack on substrate). The problem is that such a mechanism would generate metal-oxo complexes that appear to contain metal ions in chemically unreasonable high-oxidation states, e.g., FeV, CuIII , or CuIV (Reactions 5.79-5.81).
$(Fe^{III} - OOH)^{2+} \rightarrow (Fe^{V}O)^{3+} + OH^{-} \tag{5.79}$
$(Cu^{II} - OOH)^{+} \rightarrow (Cu^{IV}O)^{2+} +OH^{-} \tag{5.80}$
$(Cu^{II} - OO - Cu^{II})^{2+} \rightarrow 2 (Cu^{III}O)^{+} \tag{5.81}$
An alternative mechanism is for the peroxide or hydroperoxide ligand to attack the substrate directly; i.e., O—O bond cleavage could be concerted with attack on substrate. Another possibility is that the oxygen atom is inserted in a metalligand bond prior to transfer to the substrate. Neither of these alternative mechanisms has been demonstrated experimentally. These various possibilities remain to be considered as more information about the monooxygenase enzymes becomes available. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/05%3A_Dioxygen_Reactions/5.06%3A_Monooxygenases.txt |
Background
The oxygenase enzymes catalyze reactions of dioxygen with organic substrates in which oxygen atoms from dioxygen are incorporated into the final oxidized product.2-4 These enzymes can be divided into dioxygenases, which direct both atoms of oxygen into the product (Reaction 5.53), and monooxygenases, where one atom of oxygen from dioxygen is found in the product and the other has been reduced to water (Reaction 5.54):
$Dioxygenase:\; substrate + \;^{*}O_{2} \rightarrow substrate(^{*}O)_{2} \tag{5.53}$
$Monoxygenase:\; substrate + \;^{*}O_{2} + 2 H^{+} + 2e^{-} \rightarrow substrate(^{*}O) + H_{2}\;^{*}O \tag{5.54}$
Dioxygenases
Dioxygenase enzymes are known that contain heme iron, nonheme iron, copper, or manganese.66,67 The substrates whose oxygenations are catalyzed by these enzymes are very diverse, as are the metal-binding sites; so probably several, possibly unrelated, mechanisms operate in these different systems, For many of these enzymes, there is not yet much detailed mechanistic information. However, some of the intradiol catechol dioxygenases isolated from bacterial sources have been studied in great detail, and both structural and mechanistic information is available.66,67 These are the systems that will be described here.
1. Intradiol Catechol Dioxygenases
The role of these nonheme iron-containing enzymes is to catalyze the degradation of catechol derivatives to give muconic acids (Reaction 5.55, for example). The enzymes are induced when the only carbon sources available to the bacteria are aromatic molecules. The two best-characterized members of this class are catechol 1,2-dioxygenase (CTD) and protocatechuate 3,4-dioxygenase (PCD).
$\tag{5.55}$
a. Characterization of the Active Sites
Even before the x-ray crystal structure of PCD was obtained, a picture of the active site had been constructed by detailed spectroscopic work using a variety of methods. The success of the spectroscopic analyses of these enzymes is a particularly good example of the importance and usefulness of such methods in the characterization of metalloproteins. The two enzymes referred to in Reaction (5.55) have different molecular weights and subunit compositions,66 but apparently contain very similar active-site structures and function by very similar mechanisms. In both, the resting state of the enzyme contains one FeIII ion bound at the active site. EPR spectra show a resonance at g = 4.3, characteristic of high-spin FeIII in a so-called rhombic (low symmetry) environment,66 and the Mössbauer parameters are also characteristic of high-spin ferric.66-68 Reactions with substrate analogues (see below) cause spectral shifts of the iron chromophore, suggesting strongly that the substrate binds directly to the iron center in the course of the enzymatic reaction.
It is straightforward to rule out the presence of heme in these enzymes, because the heme chromophore has characteristic electronic-absorption bands in the visible and ultraviolet regions with high extinction coefficients, which are not observed for these proteins. Likewise, the spectral features characteristic of other known cofactors or iron-sulfur centers are not observed. Instead, the dominant feature in the visible absorption spectrum is a band with a maximum near 460 nm and a molar extinction coefficient of 3000 to 4000 M-1cm-1 per iron (see Figure 5.5).
This type of electronic absorption spectrum is characteristic of a class of proteins, sometimes referred to as iron-tyrosinate proteins, that contain tyrosine ligands bound to iron(III) in their active sites, and which consequently show the characteristic visible absorption spectrum due to phenolate-to-iron(III) charge-transfer transitions. This assignment can be definitively proven by examination of the resonance Raman spectrum, which shows enhancement of the characteristic tyrosine vibrational modes (typically ~1170, 1270, 1500, and 1600 cm-1) when the sample is irradiated in the charge-transfer band described above. Ferric complexes of phenolate ligands may be seen to give almost identical resonance Raman spectra (see Figure 5.6).
These bands have been assigned as a C—H bending vibration and a C—O and two C—C stretching vibrations of the phenolate ligand.69 In addition, NMR studies of the relaxation rates of the proton spins of water indicate that water interacts with the paramagnetic FeIII center in the enzyme. This conclusion is supported by the broadening of the FeIII EPR signal in the presence of H217O, due to interaction with the I = $\frac{5}{2}$ nuclear spin of 17O. Thus numerous spectroscopic studies of the catechol dioxygenases led to the prediction that the high-spin ferric ion was bound to tyrosine ligands and water. In addition, EXAFS data, as well as the resemblance of the spectral properties to another, better characterized iron-tyrosinate protein, i.e., transferrin (see Chapter 1), suggested that histidines would also be found as ligands to iron in these proteins.66,67
Preliminary x-ray crystallographic results on protocatechuate 3,4-dioxygenase completely support the earlier predictions based on spectroscopic studies.70 The FeIlI center is bound to two histidine and two tyrosine ligands and a water, the five ligands being arranged in a trigonal bipyramidal arrangement, with a tyrosine and a histidine located in axial positions, and with the equatorial water or hydroxide ligand facing toward a cavity assumed to be the substrate-binding cavity. The cavity also contains the positively charged guanidinium group of an arginine side chain, in the correct position to interact with the negatively charged carboxylate group on the protocatechuate substrate (see Figure 5.7).
b. Mechanistic Studies
As mentioned above, substrates and inhibitors that are substrate analogues bind to these enzymes and cause distinct changes in the spectral properties, suggesting strongly that they interact directly with the FeIII center. Nevertheless, the spectra remain characteristic of the FeIIl oxidation state, indicating that the ferric center has not been reduced. Catecholates are excellent ligands for FeIII (see, for example, the catecholate siderophores, Chapter 1) and it might therefore be assumed that the catechol substrate would bind to iron using both oxygen atoms (see 5.56).
$\tag{5.56}$
However, the observation that phenolic inhibitors p-X-C6H4-OH bind strongly to the enzymes suggested the possibility that the substrate binds to the iron center through only one oxygen atom (see 5.57).
These results contradict an early hypothesis that the mode of substrate binding, i.e., monodentate versus bidentate, might be a crucial factor in activating the substrate for reaction with dioxygen.67
Spectroscopic observations of the enzymes during reactions with substrates and substrate analogues have enabled investigators to observe several intermediates along the catalytic pathway. Such studies have led to the conclusion that the iron center remains high-spin FeIII throughout the entire course of the reaction. This conclusion immediately presents a problem in understanding the nature of the interaction of dioxygen with the enzyme, since dioxygen does not in general interact with highly oxidized metal ions such as FeIII. The solution seems to be that this reaction represents an example of substrate rather than dioxygen activation.
Studies of the oxidation of ferric catecholate coordination complexes have been useful in exploring mechanistic possibilities for these enzymes.71 A series of ferric complexes of 3,5-di-t-butyl-catechol with different ligands L have been found to react with O2 to give oxidation of the catechol ligand (Reaction 5.58)
All these studies of the enzymes and their model complexes have led to the mechanism summarized in Figure 5.9.66 In this proposed mechanism, the catechol substrate coordinates to the ferric center in either a monodentate or a bidentate fashion, presumably displacing the water or hydroxide ligand. The resulting catechol complex then reacts with dioxygen to give a peroxy derivative of the substrate, which remains coordinated to FeIll. The subsequent rearrangement of this peroxy species to give an anhydride intermediate is analogous to well-characterized reactions that occur when catechols are reacted with alkaline hydrogen peroxide.72 The observation that both atoms of oxygen derived from O2 are incorporated into the product requires that the ferric oxide or hydroxide complex formed in the step that produces the anhydride does not exchange with external water prior to reacting with the anhydride to open it up to the product diacid.
It is interesting to consider how the intradiol dioxygenase enzymes overcome the kinetic barriers to oxidations by dioxygen, and why this particular mechanism is unlikely to be applicable to the monooxygenase enzymes. The first point is that the ferric catechol intermediate is paramagnetic, with resonance forms that put unpaired electron density onto the carbon that reacts with dioxygen. The spin restriction is therefore not a problem. In addition, the catechol ligand is a very good reducing agent, much more so than the typical substrates of the monooxygenase enzymes (see next section). It is possible, therefore, that the reaction of dioxygen with the ferric catechol complex results in a concerted two-electron transfer to give a peroxy intermediate, thus bypassing the relatively unfavorable one-electron reduction of O2. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/05%3A_Dioxygen_Reactions/5.07%3A_Oxygenases.txt |
Contributors and Attributions
• Harry B. Gray (California Institute of Technology, Beckman Institute)
• Walther R. Ellis, Jr. (University of Utah, Department of Chemistry)
06: Electron Transfer
Three types of oxidation-reduction (redox) centers are found in biology: protein side chains, small molecules, and redox cofactors. The first class is frequently overlooked by mechanistic enzymologists. The sulfhydryl group of cysteine is easily oxidized to produce a dimer, known as cystine:
$2R-SH \xrightarrow[-2H^{+}]{-2e^{-}} R-S-S-R \tag{6.1}$
This type of interconversion is known to occur in several redox proteins, including xanthine oxidase, mercuric ion reductase, and thioredoxin. Other enzyme systems display spectral evidence pointing to the presence of a protein-based radical in at least one intermediate. EPR spectroscopy provides a powerful tool in studying such systems; the observation of a g = 2.0 signal that cannot be attributed to impurities or an organic redox cofactor is generally taken to be evidence for a protein-based radical. Radicals localized on tyrosine (e.g., in photosystem II and the B2 subunit of ribonucleotide reductase1) and tryptophan (e.g., in yeast cytochrome c peroxidase2) have been unambiguously identified using EPR techniques together with protein samples containing isotopically labeled amino acids (e.g., perdeuterated Tyr) or single amino-acid mutations (e.g., Trp → Phe).
A variety of small molecules, both organic and inorganic, can function as redox reagents in biological systems. Of these, only the nicotinamide and quinone coenzymes are found throughout the biosphere. Nicotinamide adenine dinucleotide (NAD) and nicotinamide adenine dinucleotide phosphate (NADP) participate in a wide variety of biological redox reactions. The 4-position of the pyridine ring is the reactive portion of both molecules (Figure 6.1). Both typically function as 2-electron redox reagents.
In contrast, quinones may function as either 1- or 2-electron carriers:
$Q \xrightleftharpoons{e^{-},H^{+}} QH \cdotp \xrightleftharpoons{e^{-},H^{+}} QH_{2} \tag{6.2}$
Free-radical semiquinone (QH•) intermediates have been detected by EPR spectroscopy in some electron transfers. Coenzyme Q, also called ubiquinone because it occurs in virtually all cells, contains a long isoprenoid tail that enables it to diffuse through membranes rapidly. This quinone derivative, which occurs in both free and protein-bound forms, is called ubiquinol when reduced (Figure 6.2). Other types of quinones are less frequently found in cells.
Metalloproteins containing a single type of redox cofactor can be divided into two general classes: electron carriers and proteins involved in the transport or activation of small molecules. Adman3 has identified some of the factors that seem to be characteristic of electron-transfer proteins (these proteins are sometimes called "electron transferases"): (a) possession of a suitable cofactor to act as an electron sink; (b) placement of the cofactor close enough to the protein surface to allow electrons to move in and out; (c) existence of a hydrophobic shell adjacent to, but not always entirely surrounding, the cofactor; (d) small structural changes accompanying electron transfer; and (e) an architecture that permits slight expansion or contraction in preferred directions upon electron transfer.
Proteins that function as electron transferases typically place their prosthetic groups in a hydrophobic environment and may provide hydrogen bonds (in addition to ligands) to assist in stabilizing both the oxidized and the reduced forms of the cofactor. Metal-ligand bonds remain intact upon electron transfer to minimize inner-sphere reorganization4 (discussed in Section III). Many of the complex multisite metalloenzymes (e.g., cytochrome c oxidase, xanthine oxidase, the nitrogenase FeMo protein) contain redox centers that function as intramolecular electron transferases, shuttling electrons to/from other metal centers that bind exogenous ligands during enzymatic turnover.
There are four classes3,5 of electron transferases, each of which contains many members that exhibit important structural differences: flavodoxins, blue copper proteins, iron-sulfur proteins, and cytochromes.
The flavodoxins6 are atypical in that they contain an organic redox cofactor, flavin mononucleotide (FMN; see Figure 6.3). These proteins have molecular weights in the 8-13 kDa range, and are found in many species of bacteria and algae. The FMN cofactor is found at one end of the protein, near the molecular surface, but only the dimethylbenzene portion of FMN is significantly exposed to the solvent (Figure 6.4). FMN can act as either a 1- or a 2-electron redox center. In solution, the semiquinone form of free FMN is unstable, and disproportionates to the quinone (oxidized) and hydroquinone (reduced) forms. Hence, free FMN functions in effect as a 2-electron reagent. FMN in flavodoxins, on the other hand, can function as a single-electron carrier. This is easily discerned by comparing reduction potentials for free and protein-bound FMN (Table 6.1). Clearly, the protein medium is responsible for this drastic alteration in oxidation-state stability. From an NMR study7 of the M. elsdenii flavodoxin quinone/semiquinone and semiquinone/hydroquinone electron self-exchange rates, it was concluded that the latter is approximately 300 times faster than the former, in keeping with the view that the physiologically relevant redox couple is semiquinone/hydroquinone.
Table 6.1 - Reduction potentials of FMN couples.
Abbreviations: Q, quinone; SQ, semiquinone; HQ, hydroquinone.
E°Q/SQ E°SQ/HQ
Free FMN -238 mV -172 mV
C.M.P. flavodoxin -92 mV -399 mV
The blue copper proteins are characterized by intense S(Cys) → Cu chargetransfer absorption near 600 nm, an axial EPR spectrum displaying an unusually small hyperfine coupling constant, and a relatively high reduction potential.4,8-10 With few exceptions (e.g., photosynthetic organisms), their precise roles in bacterial and plant physiology remain obscure. X-ray structures of several blue copper proteins indicate that the geometry of the copper site is approximately trigonal planar, as illustrated by the Alcaligenes denitrificans azurin structure (Figure 6.5).11,12 In all these proteins, three ligands (one Cys, two His) bind tightly to the copper in a trigonal arrangement. Differences in interactions between the copper center and the axially disposed ligands may significantly contribute to variations in reduction potential that are observed12 for the blue copper electron transferases. For example, E°' = 276 mV for A. denitrificans azurin, whereas that of P. vulgaris plastocyanin is 360 mV. In A. denitrificans azurin, the Cu-S(Met) bond is 0.2 Å longer than in poplar plastocyanin, and there is a carbonyl oxygen 3.1 Å from the copper center, compared with 3.8 Å in plastocyanin. These differences in bond lengths are expected to stabilize Cull in azurin to a greater extent than in plastocyanin, and result in a lower E°' value for azurin.
The iron-sulfur proteins play important roles13,14 as electron carriers in virtually all living organisms, and participate in plant photosynthesis, nitrogen fixation, steroid metabolism, and oxidative phosphorylation, as well as many other processes (Chapter 7). The optical spectra of all iron-sulfur proteins are very broad and almost featureless, due to numerous overlapping charge-transfer transitions that impart red-brown-black colors to these proteins. On the other hand, the EPR spectra of iron-sulfur clusters are quite distinctive, and they are of great value in the study of the redox chemistry of these proteins.
The simplest iron-sulfur proteins, known as rubredoxins, are primarily found in anaerobic bacteria, where their function is unknown. Rubredoxins are small proteins (6 kDa) and contain iron ligated to four Cys sulfurs in a distorted tetrahedral arrangement. The E°' value for the FeIII/II couple in water is 770 mV; that of C. pasteurianum rubredoxin is -57 mV. The reduction potentials of iron-sulfur proteins are typically quite negative, indicating a stabilization of the oxidized form of the redox couple as a result of negatively charged sulfur ligands.
The [2Fe-2S] ferredoxins (10-20 kDa) are found in plant chloroplasts and mammalian tissue. The structure of Spirulina platensis ferredoxin15 confirmed earlier suggestions, based on EPR and Mössbauer studies, that the iron atoms are present in a spin-coupled [2Fe-2S] cluster structure. One-electron reduction (E°' ~ -420 mV) of the protein results in a mixed-valence dimer (Equation 6.3):
$[Fe_{2}S_{2}(SR)_{4}]^{2-} \xrightleftharpoons[-e^{-}]{+e^{-}} [Fe_{2}S_{2}(SR)_{4}]^{3-} \tag{6.3}$
$Fd_{ox} \qquad \qquad \qquad Fd_{red}$
$2Fe(III) \qquad \qquad Fe(II) + Fe(III)$
The additional electron in Fdred is associated with only one of the iron sites, resulting in a so-called trapped-valence structure.16 The [Fe2S2(SR)4]4- cluster oxidation state, containing two ferrous ions, can be produced in vitro when strong reductants are used.
Four-iron clusters [4Fe-4S] are found in many strains of bacteria. In most of these bacterial iron-sulfur proteins, also termed ferredoxins, two such clusters are present in the protein. These proteins have reduction potentials in the -400 mV range and are rather small (6-10 kDa). Each of the clusters contains four iron centers and four sulfides at alternate comers of a distorted cube. Each iron is coordinated to three sulfides and one cysteine thiolate. The irons are strongly exchange-coupled, and the [4Fe-4S] cluster in bacterial ferredoxins is paramagnetic when reduced by one electron. The so-called "high-potential ironsulfur proteins" (HiPIPs) are found in photosynthetic bacteria, and exhibit anomalously high (~350 mV) reduction potentials. The C. vinosum HiPIP (10 kDa) structure demonstrates that HiPIPs are distinct from the [4Fe-4S] ferredoxins, and that the reduced HiPIP cluster structure is significantly distorted, as is also observed for the structure of the oxidized P. aerogenes ferredoxin. In addition, oxidized HiPIP is paramagnetic, whereas the reduced protein is EPR-silent.
This bewildering set of experimental observations can be rationalized in terms of a "three-state" hypothesis (i.e., [4Fe-4S(SR)4]n- clusters exist in three physiological oxidation states).17 This hypothesis nicely explains the differences in magnetic behavior and redox properties observed for these iron-sulfur proteins (Equation 6.4):
$[4Fe-4S(SR)_{4}]^{-} \xrightleftharpoons[-e^{-}]{+e^{-}} [4Fe-4S(SR)_{4}]^{2-} \xrightleftharpoons[-e^{-}]{+e^{-}} [4Fe-4S(SR)_{4}]^{3-} \tag{6.4}$
$HiPIP_{ox} \qquad \qquad \qquad HiPIP_{red} \qquad \qquad \qquad Ferredoxin_{red}$
$Ferredoxing_{ox}$
The bacterial ferredoxins and HiPIPs all possess tetracubane clusters containing thiolate ligands, yet the former utilize the -2/-3 cluster redox couple, whereas the latter utilize the -1/-2 cluster redox couple.
The protein environment thus exerts a powerful influence over the cluster reduction potentials. This observation applies to all classes of electron transferases—the factors that are critical determinants of cofactor reduction potentials are poorly understood at present but are thought18 to include the low dielectric constants of protein interiors (~4 for proteins vs. ~78 for H2O), electrostatic effects due to nearby charged amino-acid residues, hydrogen bonding, and geometric constraints imposed by the protein.
As a class, the cytochromes19-22 are the most thoroughly characterized of the electron transferases. By definition, a cytochrome contains one or more heme cofactors. These proteins were among the first to be identified in cellular extracts because of their distinctive optical properties, particularly an intense absorption in the 410-430 nm region (called the Soret band). Cytochromes are typically classified on the basis of heme type. Figure 6.6 displays the three most commonly encountered types of heme: heme a possesses a long phytyl "tail" and is found in cytochrome c oxidase; heme b is found in b-type cytochromes and globins; heme c is covalently bound to c-type cytochromes via two thioether linkages. Cytochrome nomenclature presents a real challenge! Some cytochromes are designated according to the historical order of discovery, e.g., cytochrome c2 in bacterial photosynthesis. Others are designated according to the $\lambda_{max}$of the $\alpha$ band in the absorption spectrum of the reduced protein (e.g., cytochrome c551).
Cytochromes c are widespread in nature. Ambler23 divided these electron carriers into three classes on structural grounds. The Class I cytochromes c contain axial His and Met ligands, with the heme located near the N-terrninus of the protein. These proteins are globular, as indicated by the ribbon drawing of tuna cytochrome c (Figure 6.7). X-ray structures of Class I cytochromes c from a variety of eukaryotes and prokaryotes clearly show an evolutionarily conserved "cytochrome fold," with the edge of the heme solvent-exposed. The reduction potentials of these cytochromes are quite positive (200 to 320 mV). Mammalian cytochrome c, because of its distinctive role in the mitochondrial electron-transfer chain, will be discussed later.
Class II cytochromes c (E°' ~ -100 mV) are found in photosynthetic bacteria, where they serve an unknown function. Unlike their Class I cousins, these c-type cytochromes are high-spin: the iron is five-coordinate, with an axial His ligand. These proteins, generally referred to as cytochromes c' , are four-$\alpha$-helix bundles (Figure 6.8). The vacant axial coordination site is buried in the protein interior.
Finally, Class III cytochromes c, also called cytochromes c3, contain four hemes, each ligated by two axial histidines. These proteins are found in a restricted class of sulfate-reducing bacteria and may be associated with the cytoplasmic membrane. The low molecular weights of cytochromes c3 (~14. 7 kDa) require that the four hemes be much more exposed to the solvent than the hemes of other cytochromes (see Figure 6.9), which may be in part responsible for their unusually negative (-200 to -350 mV) reduction potentials. These proteins possess many aromatic residues and short heme-heme distances, two properties that could be responsible for their anomalously large solid-state electrical conductivity.24 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/06%3A_Electron_Transfer/6.01%3A_Biological_Redox_Components.txt |
Electron transfers are key steps in many enzymatic reactions involving the oxidation or reduction of a bound substrate. Relevant examples include cytochrome c oxidase (O2 → 2H2O) and nitrogenase (N2 → 2NH3). To reinforce the claim that electron-transfer steps are of widespread importance, several other redox systems, representative of diverse metabolic processes, will be mentioned here.
Xanthine oxidase (275 kDa; $\alpha_{2}$ dimer) catalyzes the two-electron oxidation37-39 of xanthine to uric acid (Equation 6.7).
The first step in the biosynthesis of DNA involves the reduction of ribonucleotides (Equation 6.8) catalyzed by ribonucleotide reductase.40 The E. coli enzyme is an $\alpha_{2} \beta_{2}$ tetramer composed of a B1 protein (160 kDa) and a B2 protein (78 kDa). The B1 protein (a dimer) contains redox-active dithiol groups, binding sites for ribonucleotide substrates, and regulatory binding sites for nucleotide diphosphates. Protein B2, also a dimer, possesses a phenolate radical (Tyr-122) that is stabilized by an antiferromagnetically coupled binuclear iron center (Figure 6.18). This radical is essential for enzyme activity, and is ~10 Å from the protein-B1/protein-B2 interface. Hence it cannot directly participate in an H-atom abstraction from the substrate (bound to protein B1). Instead, the x-ray structure of the B2 protein41 suggests that a long-range electron transfer from the Tyr radical to a residue (perhaps Trp-48) on the B1 protein is operative during enzyme turnover.
$\tag{6.8}$
Most of the presently known metal-containing mono- and dioxygenases are multicomponent, requiring the involvement of additional proteins (electron transferases) to shuttle electrons from a common biological reductant (usually NADH or NADPH) to the metallooxygenase. Cytochrome P-450, whose substrate oxidation chemistry was discussed in detail in Chapter 5, serves as an excellent example. Figure 5.10 presented a catalytic cycle for cytochrome P-450-dependent hydroxylations42 that begins with substrate (RH) binding to the ferric enzyme (RH is camphor for Pseudomonas putida cytochrome P-450). To hydroxylate the camphor substrate, the monooxygenase must be reduced via the electron-transport chain in Equation (6.9).
$\tag{6.9}$
The ferredoxin reductase receives two electrons from NADH and passes them on, one at at time, to putidaredoxin, a [2Fe-2S] iron-sulfur protein. Thus, two single-electron-transfer steps from reduced putidaredoxin to cytochrome P-450 are required to complete one enzyme turnover.
The activity of the enzyme appears to be regulated at the first reduction step.43 In a 1:1 putidaredoxin-cytochrome P-450 complex, the reduction potential of putidaredoxin is -196 mV, but that of cytochrome P-450 is -340 mV in the absence of camphor; reduction of the cytochrome P-450 is thus thermodynamically unfavorable (k ~ 0.22 s-1). Upon binding camphor, the reduction potential of cytochrome P-450 shifts to -173 mV, and the electron-transfer rate in the protein complex accordingly increases to 41 s-1. "Costly" reducing equivalents are not wasted, and there are no appreciable amounts of noxious oxygen-reduction products when substrate is not present.
In the third step, molecular oxygen binds to the camphor adduct of ferrous cytochrome P-450. This species, in the presence of reduced putidaredoxin, accepts a second electron, and catalyzes the hydroxylation of the bound camphor substrate. The turnover rate for the entire catalytic cycle is 10-20 s-1, and the second electron-transfer step appears to be rate-determining.44
The bulk of the interest in electron-transfer reactions of redox proteins has been directed toward questions dealing with long-range electron transfer and the nature of protein-protein complexes whose structures are optimized for rapid intramolecular electron transfer. Before we undertake a discussion of these issues, it is worth noting that studies of the reactions of redox proteins at electrodes are attracting increasing attention.45-47 Direct electron transfer between a variety of redox proteins and electrode surfaces has been achieved. Potential applications include the design of substrate-specific biosensors, the development of biofuel cells, and electrochemical syntheses. An interesting application of bioelectrochemical technology is the oxidation of p-cresol to p-hydroxybenzaldehyde (Figure 6.19).48 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/06%3A_Electron_Transfer/6.02%3A_Coupling_Electron_Transfers_and_Substrate_Activation.txt |
Overview
Measurements of the rates of oxidation-reduction reactions began in the late 1940s. A great deal of the early experimental work was carried out by inorganic chemists, and by the 1970s the reactivity patterns of many complexes had been uncovered.49-51 Chemists studying the mechanisms of metalloprotein electrontransfer reactions frequently seek parallels with the redox behavior of less-complicated inorganic complexes.
In examining biological electron transfers, it is important to remember that metalloproteins are more than just metal ions in disguise. Virtually every property of a protein (excluding its amino-acid sequence) depends on the solution pH. Redox proteins are very large polyelectrolytes whose redox prosthetic groups are typically buried in the protein interior. One important distinction between redox reactions of proteins and redox reactions of small transition-metal complexes is the magnitude of the electron donor-to-acceptor distance. The relevant distance for small molecules, unlike redox proteins, is generally taken to be van der Waals contact. Within the last ten years, it has been convincingly demonstrated that electrons can "tunnel" at significant rates across distances of 15 Å or more in protein interiors.52-58
Experimental investigation of the factors that control the rates of biological redox reactions has not come as far as the study of the electron transfers of metal complexes, because many more variables must be dealt with (e.g., asymmetric surface charge, nonspherical shape, uncertain details of structures of proteins complexed with small molecules or other proteins). Many experimental approaches have been pursued, including the covalent attachment of redox reagents to the surfaces of metalloproteins.
Self-exchange and Cross Reactions
The simplest reactions in solution chemistry are electron self-exchange reactions (Equation 6.10), in which the reactants and products are the same (the asterisk is used to identify a specific isotope).
$\;^{\ast}A_{ox} + A_{red} \rightarrow \;^{\ast}A_{red} + A_{ox} \tag{6.10}$
The only way to establish chemically that a reaction has taken place is to introduce an isotopic label. There is no change in the free energy ($\Delta$G° = 0) for this type of reaction. As will become evident later on, the reason why these types of reactions are studied is because self-exchange rates and activation parameters are needed to interpret redox reactions in which a net chemical change occurs. The experimental measurement58 of self-exchange rates is tedious and usually only results in an order-of-magnitude estimate of the rate constant (as inferred from the experimental timescale; see Table 6.2). Most of the protein self-exchange rates reported to date have been measured by NMR line-broadening studies. Other potentially useful methods, such as Mössbauer spectroscopy and EPR, have not been widely used.
Table 6.2 - Experimental timescales in seconds
Laser Flash Photolysis $\geq 10^{-14}$
Pulse Radiolysis $\sim 10^{-9}$
Mössbauer Spectroscopy (57Fe) $10^{-9} - 10^{-6}$
EPR (transition metals) $10^{-9} - 10^{-8}$
Temperature-jump Spectrometry $\geq 10^{-8}$
NMR (1H) $\sim 10^{-5}$
Chemical Mixing $\geq 10^{-3}$
An elegant example of the measurement of an electron self-exchange rate of a redox protein was reported by Dahlin et al.59 The copper ion of stellacyanin was removed and then replaced with either 63Cu or 65Cu. Oxidized [63Cu] stellacyanin was allowed to react with reduced [65Cu] stellacyanin for various times (10 ms to 7 min) at 20 °C, after which the reaction was quenched by lowering the solution temperature to -120°C using a rapid-freeze apparatus:
$\;^{63}Cu^{2+} + \;^{65}Cu^{+} \rightarrow \;^{63}Cu^{+} + \;^{65}Cu^{2+} \tag{6.11}$
Subtle differences in the EPR spectra (Figure 6.20) of the two isotopic forms of stellacyanin (due to a small difference in the nuclear magnetic moments of the two isotopes) were used to monitor the progress of the reaction, yielding a rate constant of 1.2 x 105 M-1s-1.
Much more common are cross reactions (Equation 6.12), where Aox is the oxidized reactant, Bred is the reduced reactant, Ared is the reduced product, and Box is the oxidized product.
$A_{ox} + B_{red} \rightarrow A_{red} + B_{ox} \tag{6.12}$
For these reactions, $\Delta$G° $\neq$ 0. The experimental measurement of cross-reaction rates is generally more straightforward than the measurement of self-exchange rates. Either the reactants are simply mixed together, or a thermodynamically unstable system is generated rapidly (via pulse radiolysis, flash photolysis, or temperature-jump relaxation) to initiate the redox reaction. Absorption spectroscopy has almost always been used to monitor the progress of protein cross reactions. The primary goal of theory, as will become evident, is to provide a relationship between $\Delta$G° and $\Delta$G$\ddagger$ for cross reactions.
Both self-exchange and cross reactions can be broadly classified as inner-sphere or outer-sphere reactions. In an inner-sphere reaction, a ligand is shared between the oxidant and reductant in the transition state. An outer-sphere reaction, on the other hand, is one in which the inner coordination shells of both the oxidant and the reductant remain intact in the transition state. There is no bond breaking or bond making, and no shared ligands between redox centers. Long-range electron transfers in biology are all of the outer-sphere type. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/06%3A_Electron_Transfer/6.03%3A_Electron-transfer_Rates.txt |
The simplest electron transfer occurs in an outer-sphere reaction. The changes in oxidation states of the donor and acceptor centers result in a change in their equilibrium nuclear configurations. This process involves geometric changes, the magnitudes of which vary from system to system. In addition, changes in the interactions of the donor and acceptor with the surrounding solvent molecules will occur. The Franck-Condon principle governs the coupling of the electron transfer to these changes in nuclear geometry: during an electronic transition, the electronic motion is so rapid that the nuclei (including metal ligands and solvent molecules) do not have time to move. Hence, electron transfer occurs at a fixed nuclear configuration. In a self-exchange reaction, the energies of the donor and acceptor orbitals (hence, the bond lengths and bond angles of the donor and acceptor) must be the same before efficient electron transfer can take place.
The incorporation of the Franck-Condon restriction leads to the partitioning60-65 of an electron-transfer reaction into reactant (precursor complex) and product (successor complex) configurations. The steps in Equations \ref{6.13} to \ref{6.15} go from reactants to products: $K$ is the equilibrium constant for the formation of the precursor complex [Aox, Bred], and $k_{et}$ is the forward electron transfer rate to produce the successor complex [Ared, Box].
$A_{ox} + B_{red} \xrightleftharpoons{K} [A_{ox},\; B_{red}] \label{6.13}$
$[A_{ox},\; B_{red}] \xrightarrow{k_{et}} [A_{red},\; B_{ox}] \label{6.14}$
$[A_{red},\; B_{ox}] \xrightarrow{fast} A_{red} + B_{ox} \label{6.15}$
Marcus pioneered the use of potential energy diagrams as an aid in describing electron-transfer processes.60 For the sake of simplicity, the donor and acceptor are assumed to behave like collections of harmonic oscillators. Instead of two separate potential energy surfaces being used for the reactants, they are combined into a single surface that describes the potential energy of the precursor complex as a function of its nuclear configuration (i.e., the sum of the translational, rotational, and vibrational degrees of freedom of the reactant molecules and the molecules in the surrounding solvent-3N coordinates, where N is the number of nuclei present). Similarly, a single potential energy (3N-dimensional) surface is used to describe the potential energy of the successor complex as a function of its nuclear configuration. It has become conventional to simplify such potential energy diagrams by using one-dimensional slices through the reactant and product surfaces in order to visualize the progress of a reaction, as illustrated in Figure 6.21.
The intersection of the reactant and product surfaces (point S) represents the transition state (or "activated complex"), and is characterized by a loss of one degree of freedom relative to the reactants or products. The actual electron-transfer event occurs when the reactants reach the transition-state geometry. For bimolecular reactions, the reactants must diffuse through the solvent, collide, and form a precursor complex prior to electron transfer. Hence, disentangling the effects of precursor complex formation from the observed reaction rate can pose a serious challenge to the experimentalist; unless this is gone, the factors that determine the kinetic activation barrier for the electron-transfer step cannot be identified with certainty.
The surfaces depicted in Figure 6.21 presume that the electrons remain localized on the donor and acceptor; as long as this situation prevails, no electron transfer is possible. Thus some degree of electronic interaction, or coupling, is required if the redox system is to pass from the precursor to the successor complex. This coupling removes the degeneracy of the reactant and product states at the intersection of their respective zero-order surfaces (points S in Figure 6.21) and leads to a splitting in the region of the intersection of the reactant and product surfaces (Figure 6.22). If the degree of electronic interaction is sufficiently small, first-order perturbation theory can be used to obtain the energies of the new first-order surfaces, which do not cross. The splitting at the intersection is equal to 2HAB, where HAB is the electronic-coupling matrix element.
The magnitude of $H_{AB}$ determines the behavior of the reactants once the intersection region is reached. Two cases can be distinguished. First, $H_{AB}$ is very small; for these so-called "nonadiabatic" reactions, there is a high probability that the reactants will "jump" to the upper first-order potential energy surface, leading to very little product formation. If the electronic interaction is sufficiently large, as it is for "adiabatic" reactions, the reactants will remain on the lower first-order potential energy surface upon passage through the transition-state region.
The term adiabatic (Greek: a-dia-bainein, not able to go through) is used in both thermodynamics and quantum mechanics, and the uses are analogous. In the former, it indicates that there is no heat flow in or out of the system. In the latter, it indicates that a change occurs such that the system makes no transition to other states. Hence, for an adiabatic reaction, the system remains on the same (i.e., lower) first-order electronic surface for the entire reaction. The probability of electron transfer occurring when the reactants reach the transition state is unity. The degree of adiabaticity of the reaction is given by a transmission coefficient, $\kappa$, whose value ranges from zero to one. For systems whose HAB is sufficiently large (>kBT, where kB is the Boltzmann constant), $\kappa$ = 1. This situation occurs when the reacting centers are close together, the orbital symmetries are favorable, and no substantial changes in geometry are involved. The transmission coefficient is generally very small ($\kappa$ < 1) for electron-transfer reactions of metalloproteins, owing to the long distances involved. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/06%3A_Electron_Transfer/6.04%3A_Electron-Transfer_Theory.txt |
Electron-transfer reactions play key roles in a great many biological processes, including collagen synthesis, steroid metabolism, the immune response, drug activation, neurotransmitter metabolism, nitrogen fixation, respiration, and photosynthesis. The latter two processes are of fundamental significance-they provide most of the energy that is required for the maintenance of life. From the point of view of global bioenergetics, aerobic respiration and photosynthesis are complementary processes (Figure 6.10). The oxygen that is evolved by photosynthetic organisms is consumed by aerobic microbes and animals. Similarly, the end products of aerobic respiratory metabolism (CO2 and H2O) are the major nutritional requirements of photosynthetic organisms. The global C, H, and O cycles are thus largely due to aerobic respiration and photosynthesis.
The extraction of energy from organic compounds, carried out by several catabolic pathways (e.g., the citric-acid cycle), involves the oxidation of these compounds to CO2 and H2O with the concomitant production of water-soluble reductants (NADH and succinate). These reductants donate electrons to components of the mitochondrial electron-transfer chain, resulting in the reduction of oxygen to water:
$\frac{1}{2} O_{2} + NADH + H^{+} \rightarrow H_{2}O + NAD^{+} \tag{6.5}$
In aerobic organisms, the terminal oxidant is, of course, oxygen. However, some species of bacteria respire anaerobically and are able to use inorganic oxyanions (nitrate or sulfate) as terminal oxidants. The translocation of protons across the inner mitochondrial membrane accompanies the electron transfers that ultimately lead to the reduction of O2; these protons, in turn, activate ATP synthase, which catalyzes the phosphorylation of ADP to ATP (a process known as oxidative phosphorylation). Because the hydrolysis of ATP is very exoergonic (i.e., $\Delta$G < 0), the newly synthesized ATP is used as a molecular energy source to drive thermodynamically unfavorable reactions to completion.
The rediscovery of cytochromes by Keilin25 in 1925 led him to propose that the reduction of O2 is linked to the oxidation of reduced substrates by a series of redox reactions, carried out by cellular components collectively referred to as the respiratory electron-transport chain. Progress toward a molecular understanding of these redox reactions has been painfully slow. Most of the components are multisubunit proteins that reside in the inner mitochondrial membrane (Figure 6.11). These proteins (Complexes I-IV) are quite difficult to purify with retention of in vivo properties, and they do not crystallize well.
The components26-28 of the respiratory chain contain a variety of redox cofactors. Complex I (NADH-Q reductase; > 600 kDa) contains five iron-sulfur clusters and FMN. Complex II (succinate-Q reductase; 150 kDa) contains several iron-sulfur clusters, FAD (flavin adenine dinucleotide), and cytochrome b568. Complex III (ubiquinol-cytochrome C reductase; 250 kDa) contains a [2Fe-2S] iron-sulfur center and cytochromes b562, b566, and c1. Complex IV (cytochrome C oxidase; 200 kDa) contains at least two copper ions and cytochromes a and a3; Q denotes coenzyme Q, which may be bound to hydrophobic subunits of Complexes I, II, and/or III in vivo. Cytochrome c (cyt c in Figure 6.11) is a water-soluble protein (12.4 kDa) that is only peripherally associated with the inner mitochondrial membrane; it has been so thoroughly studied that it is generally regarded as the prime example of an electron transferase.
More than 20 redox centers are involved in the electron-transport chain. Figure 6.12 depicts a simplified view of the flow of electrons from NADH to O2 via this series of electron carriers. Electron flow through Complexes I, III, and IV is associated with the release of relatively large amounts of energy, which is coupled to proton translocation by these complexes (and therefore ATP production). The redox potentials of the electron carriers thus appear to playa role in determining the pathway of electron flow through the electron-transport chain.
Approximately 50 percent of the surface area of the inner mitochondrial membrane is lipid bilayer that is unoccupied by membrane proteins and through which these proteins, in principle, are free to diffuse laterally. Kinetic (laser photobleaching and fluorescence recovery) and ultrastructural (freeze-fracture electron microscopy) studies29,30 indicate that Complexes I-IV diffuse independently and laterally over the inner membrane, whereas cytochrome c diffuses in three dimensions (i.e., through the intramembrane space). Respiratory electron transport has been shown to be a diffusion-coupled kinetic process.29,30 The term "electron-transport chain" is thus somewhat misleading, because it implies a degree of structural order that does not exist beyond the level of a given protein complex.
In view of these observations, why are all of the electron transfers associated with mitochondrial respiration required? For example, why is cytochrome c needed to shuttle electrons in Figures 6.11 and 6.12 when the cofactor reduction potentials of Complex III are more negative than those of Complex IV? Evidently, factors other than $\Delta$G° are of importance—these will be discussed in Sections III and IV.
Photosynthesis could be viewed as the most fundamental bioenergetic process. Biological reactions are driven by an energy flux, with sunlight serving as the energy source. Photosynthesis31-36 is the process by which radiant solar energy is converted into chemical energy in the form of ATP and NADPH, which are then used in a series of enzymatic reactions to convert CO2 into organic compounds. The photosynthetic algae that appeared on Earth two million years ago released oxygen into the atmosphere and changed the environment from a reducing to an oxidizing one, setting the stage for the appearance of aerobically respiring organisms.
Photosynthesis is initiated by the capture of solar energy, usually referred to as "light harvesting." A large number of organic pigments, including chlorophylls, carotenoids, phycoerythrin, and phycocyanin (in green plants and algae) are clustered together in pigment-protein complexes called photosystems. These pigments collectively absorb most of the sunlight reaching the Earth—their absorption spectra are displayed in Figure 6.13.
Light is transformed into chemical energy in pigment-protein complexes called reaction centers. The concentration of reaction centers within a photosynthetic cell is too small to offer a suitable absorption cross section for sunlight. Hence, hundreds of these lightharvesting pigments function as molecular antennas; an x-ray structure35 of one subunit of a bacteriochlorophyll-protein complex is displayed in Figure 6.14.
Absorption of a photon by an antenna pigment promotes the pigment into an electronically excited state, which can return to the ground state by a variety of relaxation processes, including fluorescence or resonance transfer of excitation energy to a nearby pigment at picosecond rates. As much as 100 ps may elapse between the photon absorption and the arrival of the light energy at a reaction center. During this time, the energy may "migrate" in a random-walk fashion among hundreds of pigments.
The energy of the excited state is converted into electrochemical potential energy at the reaction center, which contains a primary electron donor P that transfers an electron to a nearby acceptor Al within the same protein (and P becomes oxidized to P+):
$P A_{1}A_{2}A_{3} \cdotp \cdotp \cdotp \xrightarrow{h_{\nu}} P^{\ast}A_{1}A_{2}A_{3} \cdotp \cdotp \cdotp \rightarrow P^{+} A_{1}^{-}A_{2}A_{3} \cdotp \cdotp \cdotp \tag{6.5}$
This charge separation is of paramount importance. The key problem is maintaining the charge separation, which involves minimization of the energy-wasting back reaction. Reaction centers contain an ordered array of secondary electron acceptors (A1, A2, A3•••) that optimize the $\Delta$G° that occurs at each step:
$P^{+} A_{1}^{-}A_{2}A_{3} \cdotp \cdotp \cdotp \rightarrow P^{+} A_{1}A_{2}^{-}A_{3} \cdotp \cdotp \cdotp \rightarrow P^{+} A_{1}A_{2}A_{3}^{-} \cdotp \cdotp \cdotp \tag{6.6}$
Thus, the back reaction is circumvented by optimizing forward electron transfers that rapidly remove electrons from A1-. As the acceptors are separated by greater and greater distances from P+, the probability of the back electron transfer to P+ decreases. Put another way, the overlap of P+ and each acceptor orbital decreases in the order P+/A1- > P+/A2- > P+/A3-.
Photosynthetic bacteria contain only one type of reaction center (l00 kDa). The solution of the x-ray structure (at 2.9 Å resolution) of the Rps. viridis reaction center was reported36 in 1984, providing conclusive proof that electrons can "tunnel" over 10-20 Å distances through protein interiors. The reaction-center protein contains many cofactors (Figure 6.15): two bacteriochlorophylls (BChl) in close proximity (the so-called "special pair"), two further bacteriochlorophylls that are spectroscopically identical, two bacteriopheophytins (BPh), two quinones (QA and QB), and one iron center. (QB was lost during isolation of the Rps. viridis reaction center and thus does not appear in Figure 6.15.) The reaction center contains an approximate two-fold rotation axis. Despite this strikingly high symmetry in the reaction center, one pathway of electron flow predominates, as the cartoon in Figure 6.16 indicates. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/06%3A_Electron_Transfer/6.05%3A_Energy_Storage_and_Release.txt |
Electronic Coupling
The electron-transfer reactions that occur within and between proteins typically involve prosthetic groups separated by distances that are often greater than 10 Å. When we consider these distant electron transfers, an explicit expression for the electronic factor is required. In the nonadiabatic limit, the rate constant for reaction between a donor and acceptor held at fixed distance and orientation is:70-73
$k_{et} = \bigg[ \frac{H_{AB} \;^{2}}{\hbar}\left(\dfrac{\pi}{\lambda RT}\right)^{1/2} \bigg]^{\frac{-(\lambda + \Delta G^{o})^{2}}{4 \lambda RT}}\ldotp \tag{6.27}$
The electronic (or tunneling) matrix element HAB is a measure of the electronic coupling between the reactants and the products at the transition state. The magnitude of HAB depends upon donor-acceptor separation, orientation, and the nature of the intervening medium. Various approaches have been used to test the validity of Equation (6.27) and to extract the parameters HAB and $\lambda$. Driving-force studies have proven to be a reliable approach, and such studies have been emphasized by many workers.73,74
In the nonadiabatic limit, the probability is quite low that reactants will cross over to products at the transition-state configuration.72 This probability depends upon the electronic hopping frequency (determined by HAB) and upon the frequency of motion along the reaction coordinate.75 In simple models, the electronic-coupling strength is predicted to decay exponentially with increasing donor-acceptor separation (Equation 6.28):72,76
$H_{AB} = (H_{AB}^{o})^{\frac{- \beta}{2} (\textbf{d} - \textbf{d}^{o})} \tag{6.28}$
In Equation (6.28), HAB° is the electronic coupling at close contact (d°), and $\beta$ is the rate of decay of coupling with distance (d). Studies of the distance dependence of electron-transfer rates in donor-acceptor complexes, and of randomly oriented donors and acceptors in rigid matrices, have suggested 0.8 $\leq \beta \leq$ 1.2 Å-1.73,74,77,78
Analysis of a large number of intramolecular electron-transfer rates has suggested a $\beta$ value of 1.4 Å-1 for protein reactions (Figure 6.24).79,80 Assigning a single protein $\beta$ implies that the intervening medium is homogenous. At best this is a rough approximation, because the medium separating two redox sites in a protein is a heterogenous array of bonded and nonbonded interactions.81-86 Beratan and Onuchic have developed a formalism that describes the medium in terms of "unit blocks" connected together to form a tunneling pathway.84-86 A unit block may be a covalent bond, a hydrogen bond, or a through-space jump, each with a corresponding decay factor. Dominant tunneling pathways in proteins are largely composed of bonded groups (e.g., peptide bonds), with less favorable through-space interactions becoming important when a through-bond pathway is prohibitively long (Figure 6.25).84 The tunneling pathway model has been used successfully in an analysis of the electron-transfer rates in modified cytochromes c (Section IV.D.1).
1. Binding Sites on the Plastocyanin Molecular Surface
Plastocyanin cycles between the CuII and CuI oxidation states, and transfers electrons from cytochrome f to the P700 component of photosystem I in the chloroplasts of higher plants and algae.87-89 The low molecular weight (10.5 kDa) and availability of detailed structural information90 have made this protein an attractive candidate for mechanistic studies, which, when taken together,87,91-94 point to two distinct surface binding sites (i.e., regions on the plastocyanin molecular surface at which electron transfer with a redox partner occurs). The first of these, the solvent-exposed edge of the Cu ligand His-87 (the adjacent site A in Figure 6.26), is ~6 Å from the copper atom and rather nonpolar. The second site (the remote site R in Figure 6.26) surrounds Tyr-83, and is much farther (~15 Å) from the copper center. Negatively charged carboxylates at positions 42-45 and 59-61 make this latter site an attractive one for positively charged redox reagents.
Bimolecular electron-transfer reactions are typically run under pseudo-first-order conditions (e.g., with an inorganic redox reagent present in ~15-fold excess):
$Rate = k[plastocyanin][complex] = k_{obs}[plastocyanin] \ldotp \tag{6.29}$
For some reactions [e.g., Co(phen)33+ oxidation of plastocyanin (CuI)] the expected linear plot of kobs vs. [complex] is not observed. Instead, the rate is observed to saturate (Figure 6.27).95 A "minimal" model used to explain this behavior involves the two pathways for electron transfer shown in Equation (6.30).
$\tag{6.30}$
Surprisingly, the rate ratio kremote/kadjacent is 7.
Calculations81 indicate that, despite the significant differences in distances, HAB for the remote site is ~15 percent of HAB for the adjacent site. This figure is much higher than would be expected from distance alone, suggesting that the value of the decay parameter $\beta$ in Equation (6.28) depends strongly on the structure of the intervening medium.
Modified Metalloproteins
Chemical modification of structurally characterized metalloproteins by transition-metal redox reagents has been employed52,53,96-98 to investigate the factors that control long-range electron-transfer reactions. In these semisynthetic multisite redox systems, the distance is fixed, and tunneling pathways between the donor and acceptor sites can be examined.
1. Ruthenium-modified Myoglobin
Sperm-whale myoglobin can be reacted with (NH3)5Ru(OH2)2+ and then oxidized to produce a variety of ruthenated products,52,99-101 including a His-48 derivative whose Ru $\leftrightarrow$ Fe tunneling pathway is depicted in Figure 6.28.
Electrochemical data (Table 6.5) indicate that the (NH3)5Ru3+ group does not significantly perturb the heme center, and that equilibrium (i.e., kobs = k1 + k-1) should be approached when a mixed-valent intermediate is produced by flash-photolysis techniques:
$(NH_{3})_{5}Ru^{3+}-Mb(Fe^{3+}) \xrightarrow[e^{-}]{fast} (NH_{3})_{5}Ru^{2+}-Mb(Fe^{3+}) \xrightleftharpoons[k_{-1}]{k_{1}} (NH_{3})_{5}Ru^{3+}-Mb(Fe^{2+}) \tag{6.31}$
This kinetic behavior was observed,52 and both the forward (k1) and reverse (k-1) reactions were found to be markedly temperature-dependent: k1 = 0.019 s-1 (25 °C), $\Delta$H1$\ddagger$ = 7.4 kcal/mol, k-1 = 0.041 s-1) (25°C), $\Delta$H1$\ddagger$ = 19.5 kcal/mol. X-ray crystallographic studies102 indicate that the axial water ligand dissociates upon reduction of the protein. This conformational change does not control the rates, since identical results were obtained when a second flash-photolysis technique99 was used to generate (NH3)5Ru3+-Mb(Fe2+) in order to approach the equilibrium from the other direction.
Table 6.5 - Thermodynamic parameters for the reduction of (NH3)5Ru3+ and the heme site in native and modified myoglobin (Mb).a
a) pH 7.0 $\mu$ = 0.1 M phosphate buffer.
Thermodynamic Parameter Native Mb Fe3+/2+ Modified
Fe3+/2+
Mb
(NH3)5Ru3+/2+
E°, mV vs. NHE (25 °C) 58.8 ± 2 65.4 ± 2 85.8 ± 2
$\Delta$ G°, kcal mol-1 (25 °C) -1.26 ± 0.05 -1.51 ± 0.05 -1.98 ± 0.05
$\Delta$S°, e.u. -39.2 ± 1.2 -37.6 ± 1.2 4.2 ± 1.2
$\Delta$H°, kcal mol-1 (25 °C) -13.0 ± 0.4 -12.7 ± 0.4 -0.7 ± 0.4
Cyanogen bromide has been used103 to modify the six-coordinate metmyoglobin heme site, causing the coordinated water ligand to dissociate. The CNBr-modified myoglobin heme site is thus five-coordinate in both oxidation states. As expected, the self-exchange rate increased from ~1 M-1s-1 to ~104 M-1s-1.
Recent efforts in modeling biological electron transfers using chemically modified redox proteins104-106 point the way toward the design of semisynthetic redox enzymes for catalytic applications. An intriguing example, termed flavohemoglobin, was produced by reaction of hemoglobin with a flavin reagent designed to react with Cys-93 of the $\beta$-chain (i.e., the hemoglobin molecule was modified by two flavin moieties).107 The resulting derivative, unlike native hemoglobin, accepts electrons directly from NADPH and catalyzes the para-hydroxylation of aniline in the presence of O2 and NADPH.
Protein-protein Complexes
In physiologically relevant precursor complexes, both redox centers are frequently buried in protein matrices. Characterization of such protein-protein complexes is clearly important, and several issues figure prominently:
1. What are the "rules" that govern complex formation? How important are protein-dipole/protein-dipole interactions, intermolecular hydrogen bonding, and hydrophobic interactions?
2. Are the water (and small solute) molecules associated with protein surfaces "squeezed" out of the interfacial region upon complex formation?
3. Within a given complex, is there a high degree of structural order, or do the proteins retain some independent mobility?
Most of our knowledge about the structures of protein-protein complexes comes from crystallographic studies108-110 of antigen-antibody complexes and multisubunit proteins; such systems generally exhibit a high degree of thermodynamic stability. On the other hand, complexes formed as a result of bimolecular collisions generally are much less stable, and tend to resist attempts to grow x-ray-quality crystals; the high salt conditions typically used in protein crystallizations often lead to dissociation of such complexes.
1. Cytochrome b5-cytochrome c
One of the most widely studied protein-protein complexes is that formed between mammalian cytochrome b5 and cytochrome c. Using the known x-ray structures of both proteins, Salemme111generated a static computer graphics model of this electron-transfer complex by docking the x-ray structures of the individual proteins. Two features of this model and its revision112 by molecular dynamics simulations (Figure 6.29 See color plate section, page C-12.) are noteworthy: (1) several Lys residues on cytochrome c and carboxylate-containing groups on cytochrome b5 form "salt bridges" (i.e., intermolecular hydrogen bonds); and (2) the hemes are nearly coplanar and are ~17 Å (Fe-Fe) apart. This distance was confirmed by an energy-transfer experiment113 in which the fluorescence of Zn-substituted cytochrome c was quenched by cytochrome b5. Spectroscopic studies114,115 have verified the suggestion that these proteins form a 1:1 complex at low ionic strength (Figure 6.30). In addition, chemical modification116 and spectroscopic analyses117-119 are all in agreement with the suggestion111,112 that the complex is primarily stabilized by electrostatic interactions of the (-NH3+ ••• -O2C—) type. The effect of ionic strength on the reduction of cytochrome c by cytochrome b5 is also in accord with this picture:120 lowering the ionic strength increases the reaction rate, as expected for oppositely charged molecules.
2. Hybrid Hemoglobins
A common52,55,57,121,122 experimental strategy for studying electron transfers between proteins uses a metal-substituted heme protein as one of the reactants. In particular, the substitution of zinc for iron in one of the porphyrin redox centers allows facile initiation of electron transfer through photoexcitation of the zinc porphyrin (ZnP). The excited zinc porphyrin, 3ZnP* in Equation (6.32), may decay back (kd ~ 102 s-1) to the ground state or transfer an electron to an acceptor.
$\tag{6.32}$
The ZnP+ cation radical produced in the kf step is a powerful oxidant; back electron transfer (kb) will thus occur and regenerate the starting material.
The reactions shown in Equation (6.32) have been investigated in mixedmetal [Zn, Fe] hemoglobins.123-125 A hemoglobin molecule can be viewed as two independent electron-transfer complexes, each consisting of an $\alpha_{1}$-$\beta_{2}$ subunit pair (Figure 6.31), since the $\alpha_{1}$-$\alpha_{2}$, (\beta_{1}\)-(\beta_{2}\), and $\alpha_{1}$-(\beta_{1}\) distances are prohibitively long (> 30 Å).
Both [$\alpha$(Zn), $\beta$(Fe)] and [$\alpha$(Fe), $\beta$(Zn)] hybrids have been studied. The ZnP and FeP are nearly parallel, as in the cytochrome b5-cytochrome c model complex. Long-range electron transfer (3ZnP* → Fe3+) between the $\alpha_{1}$ and $\beta_{2}$ subunits has been observed (the heme-edge/heme-edge distance is ~20 Å). The driving force for the forward electron-transfer step is ~0.8 eV, and kf (see Equation 6.32) is ~100 s-1 at room temperature, but decreases to ~9 s-1 in the low-temperature region (Figure 6.32). Below 140-160 K the vibrations that induce electron transfer "freeze out"; nuclear tunneling is usually associated with such slow, temperature-independent rates. A complete analysis of the full temperature dependence of the rate requires a quantum-mechanical treatment126,127 of $\lambda_{i}$ rather than that employed in the Marcus theory. It is interesting to note that the heme b vinyl groups (see Figure 6.6) for a given [$\alpha_{1}$(Fe), $\beta_{2}$Zn)] hybrid point toward each other and appear125 to facilitate electron transfer. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/06%3A_Electron_Transfer/6.06%3A_Long-range_Electron_Transfer_in_Proteins_%28Part_1%29.txt |
Cytochrome c
Cytochrome c occupies a prominent place in the mitochondrial electron-transport chain. Its water solubility, low molecular weight (12.4 kDa), stability, and ease of purification have allowed many experiments, which, when taken together, present a detailed picture of the structure and biological function of this electron carrier.128-133
X-ray structures134 of oxidized and reduced tuna cytochrome c are very similar; most of the differences are confined to changes in the orientations of the side chains of some surface-exposed amino acids and sub-Ångström adjustments of some groups in the protein interior. Upon reduction, the heme active site becomes slightly more ordered (Figure 6.33). Two-dimensional NMR studies135-137 confirm this interpretation of the x-ray data, and further establish that the crystal and solution structures of cytochrome c differ in only minor respects.
Cytochrome c exhibits several pH-dependent conformational states. In particular, an alkaline transition with a pKa ~ 9.1 has been observed for ferricytochrome c. This transition is believed to be associated with the dissociation of Met-80; the reduction potential decreases dramatically,138 and the 695-nm absorption band, associated with a sulfur → iron charge-transfer transition, disappears. The 2H NMR resonance due to (2H3C-) Met-80 in deuterium-enriched ferricytochrome c disappears from its hyperfine-shifted upfield position without line broadening, and reappears coincident with the (2H3C-)Met-65 resonance.139 In contrast, ferrocytochrome c maintains an ordered structure over the pH range 4 to 11.140 The heme iron in ferricytochrome c remains low-spin throughout this transition, and a new strong-field ligand must therefore replace Met-80. It has been suggested that an e-amino nitrogen of a nearby Lys provides the new donor atom, but this has not been confirmed. However, it is clear that reduction of ferricytochrome c at alkaline pH values below 11 causes a drastic conformational change at the heme site. The unknown sixth ligand must be displaced by Met-80 in order for the reduced protein to assume a structure similar to the one at neutral pH. This structural change is accompanied by a decrease in the rate of reduction of ferricytochrome c by hydrated electrons,141 as expected.
How does the protein control the reduction potential of the iron center in cytochrome c? Factors that appear to play a role include the nature of the axial ligands, the stability and solvent accessibility of the heme crevice, and the hydrophobicities of the amino acids that line the heme crevice. These issues have been addressed theoretically142,143 and experimentally144-149 using cytochrome c variants engineered by protein semisynthesis or site-directed mutagenesis. Results for horse heart cytochrome c are set out in Table 6.6. Point mutations at either of positions 78 or 83 do not significantly alter E°'; however, the double mutant (Thr-78 → Asn-78; Tyr-83 → Phe-83) exhibits a substantially lower redox potential. Evidently, the results of such changes are not necessarily additive; great care must be taken in drawing conclusions about structure-function relations in engineered proteins. Finally, the ~310 mV difference between the values for the heme octapeptide and the native protein (the axial ligands are the same in both) provides a dramatic illustration of protein environmental effects on the redox potential: shielding the heme from the solvent is expected to stabilize Fell and therefore result in an increase in E°'.
Table 6.6: Reduction potentialsa of horse heart cytochrome c. a) pH 7.0, 25 °C. b) In 2M N-acetyl-DL-methionine.
Cytochrome E°' (mV vs. NHE) Reference
Native 262 138
Met-80 → His-80 41 144
Tyr-67 → Phe-67 225 145
Thr-78 → Asn-78 264 145
Tyr-83 → Pro-83 266 145
Thr-78, Tyr-83 → Asn-78, Pro-83 235 145
Heme octapeptideb -50 146
During the last fifteen years, much has been learned about the interaction of cytochrome c with its redox partners.128-133 Cytochrome c is a highly basic protein (pI = 10.05); lysine residues constitute most of the cationic amino acids. Despite the indication from the x-ray structures that only ~1 percent of the heme surface is solvent-exposed, the asymmetric distribution of surface charges, particularly a highly conserved ring of Lys residues surrounding the exposed edge of the heme crevice, led to the suggestion that electron-transfer reactions of cytochrome c (and other Class I cytochromes as well) occur via the exposed heme edge.
Chemical modification of the surface Lys residues of cytochrome c has afforded opportunities to alter the properties of the surface $\varepsilon$-amino groups suspected to be involved in precursor complex formation. Margoliash and coworkers133,150,151 used a 4-carboxy-2,6-dinitrophenol (CDNP) modification of the Lys residues to map out the cytochrome c interaction domains with various transition-metal redox reagents and proteins. These experiments have shown that cytochrome c interacts with inorganic redox partners near the exposed heme edge.
Numerous studies129,152,153 of cytochrome c with physiological reaction partners are in accord with electrostatic interactions featured in the model cytochrome c/cytochrome b5 complex discussed earlier. Similar types of interactions have been proposed for cytochrome c/flavodoxin154 and cytochrome c/cytochrome c peroxidase complexes.155 (Recent x-ray crystal structure work155b has shed new light on this problem.) Theoretical work156 additionally suggests that electrostatic forces exert torques on diffusing protein reactants that "steer" the proteins into a favorable docking geometry. However, the domains on cytochrome c for interaction with physiological redox partners are not identical, as Figure 6.34 illustrates.
Reactions between cytochrome c and its physiological redox partners at low ionic strength generally are very fast, ~108 M-1s-1, even though the thermodynamic driving force may be as low as 20 mV, as it is for the reduction of cytochrome a in cytochrome c oxidase. Such rates are probably at the diffusioncontrolled limit for such protein-protein reactions.157,158 A more detailed understanding of these reactions will require studies that focus on the dynamical (rather than static) features of complexes of cytochrome c with other proteins. For example, there is evidence159 that a cytochrome c conformational change in the vicinity of the heme edge accompanies the formation of the complex with cytochrome c oxidase. Studies of the influence of geometry changes on activation energies52,60,160 are of particular importance in elucidating the mechanisms of protein-protein reactions.
1. Ruthenium-modified Cytochrome c
Intramolecular electron transfer in cytochrome c has been investigated by attaching photoactive Ru complexes to the protein surface.98,161 Ru(bpy)2(CO3) (bpy = 2,2'-bipyridine) has been shown to react with surface His residues to yield, after addition of excess imidazole (im), Ru(bpy)2(im)(His)2+. The protein-bound Ru complexes are luminescent, but the excited states (*Ru2+) are rather short lived ($\tau \leq$ 100 ns). When direct electron transfer from *Ru2+ to the heme cannot compete with excited-state decay, electron-transfer quenchers (e.g., Ru(NH3)63+) are added to the solution to intercept a small fraction (1-10%) of the excited molecules, yielding (with oxidative quenchers) Ru3+. If, before laser excitation of the Ru site, the heme is reduced, then the Fe2+ to Ru3+ reaction (ket) can be monitored by transient absorption spectroscopy. The ket values for five different modified cytochromes have been reported: (Ru(His- 33), 2.6(3) x 106; Ru(His-39), 3.2(4) x 106; Ru(His-62), 1.0(2) x 104; Ru(His- 72), 9.0(3) x 105; and Ru(His-79), > 108 s-1).162,163
According to Equation (6.27), rates become activationless when the reaction driving force (- $\Delta$G°) equals the reorganization energy (A), The driving force (0.74 eV) is approximately equal to the reorganization energy (0.8 eV) estimated for the Ru(bpy)2(im)(His)-cyt c reactions.161 The activationless (maximum) rates (kmax) are limited by HAB2, where HAB is the electronic matrix element that couples the reactants and products at the transition state. Values of kmax and HAB for the Fe2+ to Ru3+ reactions are given in Table 6.7.
Table 6.7: Electron-transfer parameters163 for Ru(bpy)2(im)(His-X)-cytochromes c.a) C = covalent bond, H = hydrogen bond, S = space jump.
X kmax (s-1) HAB (cm-1) [Fe2+—Ru3+] d (Å) neffa $\sigma$ 1(Å)
79 > 1.0 x 108 > 0.6 4.5 8 (8C) 11.2
39 3.3 x 106 0.11 12.3
14.0 (11C)
(1H)
19.6
33 2.7 x 106 0.097 11.1
13.9 (11C)
(1H)
19.5
72 9.4 x 105 0.057 8.4
17.6 (7C)
(1S)
24.6
62 1.0 x 104 0.006 14.8
20.6 (16C)
(2H)
28.8
Calculations that explicitly include the structure of the intervening medium81- 86,164-169 have been particularly helpful in developing an understanding of distant electronic couplings. As discussed in Section IV.A, the couplings in proteins can be interpreted in terms of pathways comprised of covalent, H-bond, and through-space contacts. An algorithm has been developed85,170 that searches a protein structure for the best pathways coupling two redox sites (the pathways between the histidines (33, 39, 62, 72, 79) and the heme are shown in Figure 6.35). A given coupling pathway consisting of covalent bonds, H-bonds, and through-space jumps can be described in terms of an equivalent covalent pathway with an effective number of covalent bonds (neff). Multiplying the effective number of bonds by 1.4 Å/bond gives a-tunneling lengths ($\sigma$1) for the five pathways (Table 6.7) that correlate well with the maximum rates (one-bond limit set at 3 x 1012 s-1; slope of 0.71 Å-1) (Figure 6.35). The 0.71 Å-1 decay accords closely with related distance dependences for covalently coupled donor-acceptor molecules.73,77
Bacterial Photosynthetic Reaction Centers
Photosynthetic bacteria produce only one type of reaction center, unlike green plants (which produce two different kinds linked together in series), and are therefore the organisms of choice in photosynthetic electon-transfer research.171- 176 As indicated in Section I.B, the original reaction center structure (Figure 6.15) lacked a quinone (QB). Subsequent structures for reaction centers from other photosynthetic bacteria177,178 contain this quinone (Figure 6.36 See color plate section, page C-13.). The Rps. sphaeroides reaction center contains ten cofactors and three protein subunits. (Note that the Rps. viridis structure contains a cytochrome subunit as well.) The cofactors are arrayed so that they nearly span the 40-Å-thick membrane (Figure 6.37 See color plate section, page C-13.). The iron atom is indicated by the red dot near the cytoplasmic side of the membrane (bottom). In spite of the near two-fold axis of symmetry, electron transfer proceeds along a pathway that is determined by the A branch. In particular, BChlB and BPheB do not appear to play an important role in the electron transfers.
It was demonstrated long ago that (BChl)2 is the primary electron donor and that ubiquinone (or metaquinone) is the ultimate electron acceptor. Transient flash photolysis experiments indicate that several electron-transfer steps occur in order to translocate the charge across the membrane (Figure 6.38). Curiously, the high-spin ferrous iron appears to play no functional role in the QA to QB electron transfer.179 In addition, the part played by BChlA is not understood—it may act to promote reduction of BPheA via a superexchange mechanism.180,181 Cytochromes supply the reducing equivalents to reduce the special pair (BChl)2+.
Estimated rate constants for the various electron-transfer steps, together with approximate reduction potentials, are displayed in Figure 6.39. For each step, the forward rate is orders of magnitude faster than the reverse reaction. The rapid rates suggest that attempts to obtain x-ray structures of intermediates (especially the early ones!) will not be successful. However, molecular dynamics methods are being explored in computer simulations of the structures of various intermediates.182,183 Within a few years we may begin to understand why the initial steps are so fast. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/06%3A_Electron_Transfer/6.07%3A_Long-range_Electron_Transfer_in_Proteins_%28Part_2%29.txt |
In classical transition-state theory, the expression for the rate constant of a bimolecular reaction in solution is
$k = (\kappa \nu_{n})^{\frac{-\Delta G^{\ast}}{RT}}, \tag{6.16}$
where $\nu_{n}$, the nuclear frequency factor, is approximately 1011 M-1s-1 for small molecules, and $\Delta$G* is the Gibbs-free-energy difference between the activated complex and the precursor complex. This theoretical framework provides the starting point for classical electron-transfer theory. Usually the transmission coefficient $\kappa$ is initially assumed to be unity. Thus, the problem of calculating the rate constant involves the calculation of $\Delta$G*, which Marcus partitioned into several parameters:
$\Delta G^{\ast} = w^{r} + \left(\dfrac{\lambda}{4}\right) \left(1 + \dfrac{\Delta G^{o\; \prime}}{\lambda}\right)^{2}, \tag{6.17}$
$\Delta G^{o\; \prime} = \Delta G^{o} + w^{p} - w^{r} \ldotp \tag{6.18}$
Here wr is the electrostatic work involved in bringing the reactants to the mean reactant separation distance in the activated complex, and wp is the analogous work term for dissociation of the products. These terms vanish in situations where one of the reactants (or products) is uncharged. $\Delta$G° is the Gibbs-free-energy change when the two reactants and products are an infinite distance apart, and $\Delta$G°' is the free energy of the reaction when the reactants are a distance r apart in the medium; $\Delta$G° is the standard free energy of the reaction, obtainable from electrochemical measurements (the quantity - $\Delta$G° is called the driving force of the reaction).
The reorganization energy $\lambda$ is a parameter that contains both inner-sphere ($\lambda_{i}$) and outer-sphere ($\lambda_{o}$) components; $\lambda = \lambda_{i} + \lambda_{o}$. The inner-sphere reorganization energy is the free-energy change associated with changes in the bond lengths and angles of the reactants. The $\lambda_{i}$ term can be evaluated within the simple harmonic-oscillator approximation:
$\lambda_{i} = \left(\dfrac{1}{2}\right) \sum_{j} k_{j} (\Delta x_{j})^{2}, \tag{6.19}$
where kj values are normal-mode force constants, and the $\Delta$xj values are differences in equilibrium bond lengths between the reduced and oxidized forms of a redox center.
The outer-sphere reorganization energy reflects changes in the polarization of solvent molecules during electron transfer:
$\lambda_{o} = e^{2} \bigg[\left(\dfrac{1}{2r_{A}}\right) + \left(\dfrac{1}{2r_{B}}\right) - \left(\dfrac{1}{d}\right) \bigg] \bigg[\left(\dfrac{1}{D_{op}}\right) - \left(\dfrac{1}{D_{s}}\right) \bigg] ; \tag{6.20}$
d is the distance between centers in the activated complex, generally taken to be the sum of the reactant radii rA and rB; Dop is the optical dielectric constant of the medium (or, equivalently, the square of the refractive index); and Ds is the static dielectric constant. This simple model for the effect of solvent reorganization assumes that the reactants are spherical, and that the solvent behaves as a dielectric continuum. (Sometimes the latter approximation is so rough that there is no correspondence between theory and experiment.)
Variations in $\lambda$ can have enormous effects on electron-transfer rates, Some of the possible variations are apparent from inspection of Equation (6.20). First, $\lambda_{o}$ decreases with increasing reactant size. Second, the dependence of the reaction rate on separation distance attributable to $\lambda_{o}$occurs via the $\frac{1}{d}$ term. Third, $\lambda\_{o}$ decreases markedly as the solvent polarity decreases. For nonpolar solvents, Ds $\simeq$ Dop $\simeq$ 1.5 to 4.0. It is significant to note that protein interiors are estimated to have Ds $\simeq$ 4, whereas, Ds $\simeq$ 78 for water. An important conclusion is that metalloproteins that contain buried redox cofactors need not experience large outer-sphere reorganization energies.
The key result of Marcus theory is that the free energy of activation displays a quadratic dependence on $\Delta$G° and $\lambda$ (ignoring work terms). Hence, the reaction rate may be written as
$k_{et} = (\nu_{n} \kappa)^{\frac{-(\lambda + \Delta G^{o})^{2}}{4 \lambda RT}}\ldotp \tag{6.21}$
For intramolecular reactions, the nuclear frequency factor ($\nu_{n}$) is ~1013 s-1. One of the most striking predictions of Marcus theory follows from this equation: as the driving force of the reaction increases, the reaction rate increases, reaching a maximum at - $\Delta$G° = $\lambda$; when - $\Delta$G° is greater than $\lambda$, the rate decreases as the driving force increases (Figure 6.23).
Two free-energy regions, depending on the relative magnitudes of - $\Delta$G° and $\lambda$, are thus distinguished. The normal free-energy region is defined by - $\Delta$G° < $\lambda$A. In this region, $\Delta$G* decreases if - $\Delta$G° increases or if $\lambda$ decreases. If - $\Delta$G° = $\lambda$, there is no free-energy barrier to the reaction. In the inverted region, defined by - $\Delta$G° > $\lambda$, $\Delta$G* increases if $\lambda$ decreases or if - $\Delta$G° increases. Another widely used result of Marcus theory deals with the extraction of useful kinetic relationships for cross reactions from parameters for self-exchange reactions. Consider the cross reaction, Equation (6.22), for which the rate
$A_{1}(ox) + A_{2}(red) \rightarrow A_{1}(red) + A_{2}(ox) \tag{6.22}$
and equilibrium constants are k12 and K12, respectively. Two self-exchange reactions are pertinent here:
$A_{1}(ox) + A_{1}(red) \rightarrow A_{1}(red) + A_{1}(ox) \tag{6.23a}$
$A_{2}(ox) + A_{2}(red) \rightarrow A_{2}(red) + A_{2}(ox) \tag{6.23b}$
These reactions are characterized by rate constants k11 and k22, respectively. The reorganization energy ($\lambda_{12}$for the cross reaction can be approximated as the mean of the reorganization energies for the relevant self-exchange reactions:
$\lambda_{12} = \frac{1}{2} (\lambda_{11} + \lambda_{22}) \tag{6.24}$
Substitution of Equation (6.24) into Equation (6.17) leads to the relation
$\Delta G_{12}^{\ast} = \frac{1}{2}(\Delta G_{11}^{\ast} + \Delta G_{22}^{\ast}) + \frac{1}{2}\Delta G_{12}^{\ast}(1 + \alpha), \tag{6.25a}$
where
$\alpha = \frac{\Delta G_{12}^{\ast}}{4(\Delta G_{11}^{\ast} + \Delta G_{22}^{\ast})}\ldotp \tag{6.25b}$
When the self-exchange rates k11 are corrected for work terms or when the latter nearly cancel, the cross-reaction rate k12 is given by the Marcus cross relation,
$k_{12} = (k_{11}k_{22}K_{12}f_{12})^{\frac{1}{2}}, \tag{6.26a}$
where
$ln f_{12} = \frac{(ln K_{12})^{2}}{4\; ln \left(\dfrac{k_{11}k_{22}}{\nu_{n}^{2}}\right)}\ldotp \tag{6.26b}$
This relation has been used to predict and interpret both self-exchange and cross-reaction rates (or even K12, depending on which of the quantities have been measured experimentally. Alternatively, one could study a series of closely related electron-transfer reactions (to maintain a nearly constant $\lambda$12) as a function of $\Delta$G12; a plot of In k12 vs. In K12 is predicted to be linear, with slope 0.5 and intercept 0.5 In (k11k22). The Marcus prediction (for the normal free-energy region) amounts to a linear free-energy relation (LFER) for outer-sphere electron transfer.
Cross Reactions of Blue Copper Proteins
Given the measured self-exchange rate constant for stellacyanin (k11 1.2 x 105 M 1s-1), the Marcus cross relation (Equation 6.26a) can be used to calculate the reaction rates for the reduction of CuII-stellacyanin by Fe(EDTA)2- and the oxidation of Cul-stellacyanin by Co(phen)33+. E°(Cu2+/+) for stellacyanin is 0.18 V vs. NHE, and the reduction potentials and self-exchange rate constants for the inorganic reagents are given in Table 6.3.66,67 For relatively small $\Delta$E° values, f12 is ~1; here a convenient form of the Marcus cross relation is log k12 = 0.5[log k11 + log k22 + 16.9$\Delta$E12°]' Calculations with k11, k22, and $\Delta$E12° from experiments give k12 values that accord quite closely with the measured rate constants.
$Cu^{II}St + Fe(EDTA)^{2-} \rightarrow Cu^{I}St + Fe(EDTA)^{-}$
$k_{12}(calc.) = 2.9 \times 10^{5} M^{-1}s^{-1} \qquad (\Delta E_{12}^{o} = 0.06 V)$
$k_{12}(obs.) = 4.3 \times 10^{5} M^{-1}s^{-1} \qquad \qquad \qquad \qquad \qquad$
$Cu^{I}St + Co(phen)_{3}^{3+} \rightarrow Cu^{II}St + Co(phen)_{3}^{2+}$
$k_{12}(calc.) = 1.4 \times 10^{5} M^{-1}s^{-1} \qquad (\Delta E_{12}^{o} = 0.19 V)$
$k_{12}(obs.) = 1.8 \times 10^{5} M^{-1}s^{-1} \qquad \qquad \qquad \qquad \qquad$
Table 6.3 - Reduction potentials and self-exchange rate constants for inorganic reagents.
Reagent E°(V vs. NHE) k22 (M-1s-1)
Fe(EDTA)-/2- 0.12 6.9 x 104
Co(phen)33+/2+ 0.37 9.8 x 101
The success of the Marcus cross relation with stellacyanin indicates that the copper site in the protein is accessible to inorganic reagents. The rate constants for the reactions of other blue copper proteins with inorganic redox agents show deviations from cross-relation predictions (Table 6.4).68 These deviations suggest the following order of surface accessibilities of blue copper sites: stellacyanin > plastocyanin > azurin. Rate constants for protein-protein electron transfers also have been subjected to cross-relation analysis.69
Table 6.4 - Reactions of blue copper proteins with inorganic reagents.
a) M-1s-1
Protein Reagent k12 (obs.)a $\Delta$E12° k11 (obs.)a k11 (calc.)a
Stellacyanin Fe(EDTA)2- 4.3 x 105 0.064 1.2 x 105 2.3 x 105
Co(phen)33+ 1.8 x 105 0.186 1.2 x 105 1.6 x 105
Ru(NH3)5py3+ 1.94 x 105 0.069 1.2 x 105 3.3 x 105
Plastocyanin Fe(EDTA)2- 1.72 x 105 0.235 ~103 - 104 7.3 x 101
Co(phen)33+ 1.2 x 103 0.009 ~103 - 104 1.1 x 104
Ru(NH3)5py3+ 3.88 x 103 -0.100 ~103 - 104 4.9 x 104
Azurin Fe(EDTA)2- 1.39 x 103 0.184 2.4 x 106 2.8 x 10-2
Co(phen)33+ 2.82 x 103 0.064 2.4 x 106 7.0 x 103
Ru(NH3)5py3+ 1.36 x 103 -0.058 2.4 x 106 1.1 x 103 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/06%3A_Electron_Transfer/6.08%3A_Marcus_Theory.txt |
Transition-metal/sulfide sites, especially those containing iron, are present in all forms of life and are found at the active centers of a wide variety of redox and catalytic proteins. These proteins include simple soluble electron-transfer agents (the ferredoxins), membrane-bound components of electron-transfer chains, and some of the most complex metalloenzymes, such as nitrogenase, hydrogenase, and xanthine oxidase. In this chapter we first review the chemistry of the Fe-S sites that occur in relatively simple rubredoxins and ferredoxins, and make note of the ubiquity of these sites in other metalloenzymes. We use these relatively simple systems to show the usefulness of spectroscopy and model-system studies for deducing bioinorganic structure and reactivity. We then direct our attention to the hydrogenase and nitrogenase enzyme systems, both of which use transition-metalsulfur clusters to activate and evolve molecular hydrogen.
Contributors and Attributions
• Edward I. Stiefel (Exxon Research and Engineering Company)
• Graham N. George (Exxon Research and Engineering Company)
07: Ferrodoxins Hydrogenases and Nitrogenases - Metal-Sulfide Proteins
Iron sulfide proteins involved in electron transfer are called ferredoxins and rubredoxins.* The ferredoxins were discovered first, and were originally classified as bacterial (containing Fe4S4 clusters) and plant (containing Fe2S2 clusters) ferredoxins. This classification is now recognized as being not generally useful, since both Fe2S2 and Fe4S4 ferredoxins are found in plants,14,15 animals,2,6,16 and bacteria.4 Ferredoxins are distinguished from rubredoxins by their possession of acid-labile sulfide; i.e., an inorganic S2- ion that forms H2S gas upon denaturation at low pH. Rubredoxins have no acid-labile sulfide, and generally have a single iron in a more or less isolated site. Despite their lack of acid-labile sulfide, rubredoxins are included in this chapter because they have sequences much like those of the ferredoxins, and because their simple mononuclear Fe2+ and Fe3+ sites provide convenient illustrations of key structural and spectroscopic features.
In most ferredoxins, and in all rubredoxins, the protein ligands are cysteines, which provide four thiolate donors to the 1Fe, 2Fe, or 4Fe center. Additionally, the existence of 3Fe centers and of Fe-S sites that contain a second metal (i.e., heteronuclear clusters) make the Fe-S class a broad and multifunctional one.
Simple cytochromes and simple iron-sulfide proteins are similar, in that both can undergo one-electron transfer processes that are generally uncoupled from proton-, atom-, or group-transfer processes. Some of these proteins, such as cytochrome c3 from Desulfovibrio with four hemes17 or ferredoxin from Clostridium pasteurianum with two Fe4S4 centers,6 can transfer more than one electron, because they have multiple copies of a one-electron transfer group. The cytochromes were discovered in 1886 by McMunn,18 and their role in metabolism was discovered in the 1920s by Keilin (Chapter 6). The intense optical absorbance of these heme-containing proteins contributed singularly to their discovery and biochemical characterization. In contrast, the iron-sulfur proteins, although red to red-brown, absorb far more weakly in the visible region than do the cytochromes. Their presence is sometimes obscured by the cytochromes, and their frequent air instability made their initial recognition and isolation more difficult. It was not until the early 1960s that discoveries by several research groups19 led to the isolation, recognition, and characterization of the ferredoxins. The use of EPR spectroscopy and its application to biological systems had a profoundly stimulating effect on the field (see below).
Although cytochromes were discovered first, the ferredoxins are likely to be the older proteins from an evolutionary perspective.20 Ferredoxins have relatively low-molecular-weight polypeptide chains, require no organic prosthetic group, and often lack the more complex amino acids. In fact, the amino-acid composition in clostridial ferredoxin is close to that found in certain meteorites.21
The various Fe-S sites found in electron-transfer proteins (ferredoxins) are also found in many enzymes,6,11,22.23 where these centers are involved in intraor interprotein electron transfer. For example, sulfite reductase contains a siroheme and an Fe4S4 center,24 which are strongly coupled and involved in the sixelectron reduction of SO32- to H2S. Xanthine oxidase (see Figure 7.1) has two identical subunits, each containing two different Fe2S2 sites plus molybdenum and FAD sites. In xanthine oxidase, the Mo(VI) site carries out the two-electron oxidation of xanthine to uric acid, being reduced to Mo(IV) in the process.25 The Mo(VI) site is regenerated by transferring electrons, one at a time, to the Fe2S2 and flavin sites, thereby readying the Mo site for the next equivalent of xanthine. Although the Fe2S2 sites do not directly participate in substrate reactions, they are essential to the overall functioning of the enzyme system. The Fe2S2 centers in xanthine oxidase play the same simple electron-transfer role as the Fe2S2 ferredoxins play in photosynthesis.
Structurally, all the iron-sulfur sites characterized to date are built up of (approximately) tetrahedral iron units (see Figure 7.2). In rubredoxins the single iron atom is bound in tetrahedral coordination by four thiolate ligands provided by cysteine side chains. In two-iron ferredoxins the Fe2S2 site consists of two tetrahedra doubly bridged through a pair of sulfide ions, i.e., Fe2($\mu_{2}$-S)2, with the tetrahedral coordination of each Fe completed by two cysteine thiolates. In four-iron or eight-iron ferredoxins, the 'thiocubane' Fe4S4 cluster consists of four tetrahedra sharing edges with triply bridging S2- ions, i.e., Fe4($\mu_{3}$-S)4, with each Fe completing its tetrahedron by binding to a single cysteine thiolate. Finally, for Fe3S4 clusters, which are now being found in more and more proteins, the well-established structure has one triply bridging and three doubly bridging sulfide ions, Fe3($\mu_{3}$-S)($\mu_{2}$S)3. The Fe3S4 unit can be thought of as derived from the 'thiocubane' Fe4S4 unit by the removal of a single iron atom.
In what follows we will introduce these structures in the order 1Fe, 2Fe, 4Fe, and 3Fe. For each, we will first discuss the physiological role(s) of the particular proteins, then the structural features, followed by the spectroscopic properties and model systems.
* For review articles, see References 1-11. For a discussion of nomenclature, see References 12 and 13. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/07%3A_Ferrodoxins_Hydrogenases_and_Nitrogenases_-_Metal-Sulfide_Proteins/7.01%3A_Iron-sulfur_Proteins_and_Models.txt |
Rubredoxin Model Systems
The simple mononuclear tetrahedral site of Rd has been chemically modeled in both its reduced and its oxidized forms. The bidentate o-xylyl-$\alpha$,$\alpha$'-dithiolate ligand forms bis complexes of Fe(II) and Fe(III) that have spectroscopic features quite similar to those of the protein.62,63 The preparative procedure is relatively straightforward (Equation 7.1).
$\tag{7.1}$
The UV-visible-NIR spectra, Mössbauer spectra, and magnetic susceptibility differ only slightly from those of oxidized and reduced rubredoxins.
The monodentate benzenethiolate (thiophenolate) ligand, C6H5S-, similarly forms the ferrous Fe(SC6H5)42- complex.64,65 Although for some time it was felt that the oxidized form, Fe(SC6H5)4-, was inherently unstable, the sterically hindered monothiolate ligand 2,3,5,6-tetramethylbenzenethiolate was found to form66-68 a stable, quite symmetric Fe(III) tetrathiolate anion. Armed with this information, the preparation of the tetrakis(benzenethiolate) Fe(III) complex was reinvestigated, and the complex successfully synthesized67 (Equation 7.2).
$\tag{7.2}$
The Fe(III) and Fe(II) tetrathiolate species now serve as excellent structural models for the Fe sites of both oxidized and reduced Rd.69
The structural parameters for the oxidized rubredoxin analogues are very similar to those of the oxidized Rd iron site. The reduced complexes reveal a lengthening of the average Fe-S bond from 2.27 to 2.36 Å, consistent with the change in oxidation state from ferric to ferrous. The addition of an electron has a more profound structural effect in this single-iron center than in some of the multiiron clusters, where electrons are more delocalized.
Clearly, for the single-Fe sites, the dominant structural feature is their near-tetrahedral tetrathiolate coordination. The dominant electronic structural feature is the presence of high-spin Fe3+ and Fe2+ sites. The important mode of chemical reactivity is a simple one-electron transfer. Each of these features carries over to the 2Fe, 4Fe, and 3Fe sites discussed below.
Fe2S2 Ferredoxins
The simple 2Fe-2S proteins are sometimes referred to as "plant" or "plant-type" ferredoxins. The protein from spinach, which serves as an electron acceptor in the photosynthetic apparatus,14,15,50,70 was among the first to be wellcharacterized and widely studied, and could be considered the prototypical 2Fe-2S ferredoxin, However, 2Fe-2S proteins are also well-known in bacteria.4 The protein from the cyanobacterium (blue-green alga) Spirulina platensis has been structurally elucidated by x-ray crystallography.47 Putidaredoxin, from Pseudomonas putida, which serves as a donor to the P-450 camphor monooxygenase system, has been extensively studied.28 Fe2S2 centers are also well-established in mammalian proteins. Adrenodoxin29 serves as the electron donor to the P-450 monooxygenase system that carries out the 11-$\beta$-hydroxylation of steroids, The so-called "Rieske proteins" are found in the bc1 complex of mitochondria47 as well as in the bd complex of the photosynthetic apparatus of plants.71 In addition, Fe2S2 centers are well-known constituents of such redox proteins as xanthine oxidase,25,72 CO oxidase,25 succinate dehydrogenase,73-75 and putidamonooxin.76 Table 7.1 lists some of the Fe2S2 proteins and their properties.
The x-ray crystal structure of only the single 2Fe-2S protein mentioned above has been determined;70a the 2Fe-2S ferredoxin from the blue-green alga Spirulina platensis6,22,47,77,78 shows significant sequence identity with chloroplast ferredoxins typical of higher plants.79,80 As Figure 7.8 shows, the Fe2S2 unit in this 11-kDa protein is bound by Cys-41, Cys-46, Cys-49, and Cys-79, The binuclear iron cluster is found in a largely hydrophobic region of the protein, but is within 5 Å of the protein surface.6 The sulfur atoms of the cluster, both inorganic and cysteinyl, are hydrogen-bonded to six peptide NH groups and one serine OH group, which presumably stabilize the cluster/protein complex, The serine involved in the H-bonding, Ser-40, is conserved in all plant and algal 2Fe-2S ferredoxins sequenced, which implies that it plays a crucial structural or functional role.
The structure of the 2Fe-2S core in Figure 7.2 reveals a tetrahedron of S ligands surrounding each Fe atom. The two tetrahedra share an edge defined by the two bridging sulfide ions, and the core structure is designated Fe2($\mu_{2}$-S)2. Fe-S distances and angles cannot be measured accurately in the structure at the present 2.5-Å resolution;70a so we will later discuss these details in terms of model compounds.
The Fe2S2 center shows nicely how spectroscopy can be used to deduce the structure of an active site. Indeed, in this case the now well-established active-site structure was deduced by a combination of chemical, spectroscopic, and magnetic methods, and the site was successfully modeled long before the first protein crystallographic study was reported. The presence of acid-labile, inorganic sulfide is a key feature of both the Fe2S2 and the Fe4S4 centers. The 1:1 stoichiometry between iron and acid-labile sulfide was eventually established analytically for Fe2S2 centers.9-11 Care must be taken to ensure that both the protein and its active-site complement are homogeneous. Although protein homogeneity is usually established by electrophoretic methods, these methods may not distinguish between pure proteins and those with absent or incomplete active centers. Fortunately, absorption at 420 nm is due solely to the Fe2S2 cluster, whereas the 275-nm absorption is dominated by the protein. Therefore a good criterion for active-site saturation and homogeneity is the ratio of the absorbances at 420 and 275 nm, A420 nm/A275 nm, which is ~0.48 for pure spinach ferredoxin.81 Once homogeneous protein is obtained, the Fe2S2 composition of the "plant" ferredoxins can be correctly deduced analytically.
The Fe2S2 center displays two redox states that differ by a single electron. The potential range for the couple is -250 to -420 mV, revealing the highly reducing nature of the ferredoxin. The correct structure of the Fe2S2 center was first proposed in 1966 based on EPR studies.82 The reduced state of the cluster shows a rhombic EPR signal with g values of 1.88, 1.94, and 2.04 (Figure 7.6B) characteristic of an S = $\frac{1}{2}$ center. The oxidized state is EPR-silent. The weakness of the sulfur ligand field causes the iron atoms to be high-spin. But how can two sulfur-ligated iron atoms, each with a tendency to be high-spin, produce a state with a single unpaired electron?
The individual Fe atoms in the Fe2S2 cluster resemble those in rubredoxin quite closely. The two redox states of the Fe2S2 protein correspond to an Fe3+-Fe3+ and an Fe3+-Fe2+ pair, respectively, as shown in Figure 7.9. In the all-ferric oxidized state, the two Fe3+ sites are antiferromagnetically coupled; i.e., the spins of the five d electrons on the two iron atoms are oppositely aligned, such that their pairing produces an effective S = 0, diamagnetic ground state. In the reduced form, a single unpaired electron is present, because the S = $\frac{5}{2}$ Fe3+ and S = 2 Fe2+ sites are antiferromagnetically coupled, leaving one net unpaired spin and an S = $\frac{1}{2}$ ground state. The profound difference between the electronic properties of rubredoxin and Fe2S2 ferredoxin arises because the latter has two Fe atoms in close proximity, which allows for their magnetic coupling.
Strong support for the spin-coupling model in Fe2S2 ferredoxins comes from a detailed analysis of their absorption and circular dichroism spectra.83 As with rubredoxin (see Figure 7.5), we expect no low-energy spin-allowed d-d bands for the ferric site in either the oxidized or the reduced state. Indeed, the oxidized state containing all Fe3+ shows no low-energy bands; the reduced state containing a single Fe2+ displays low-energy, low-intensity bands in the region 4,000-9,000 cm-1, in close analogy to the situation in reduced rubredoxin. The combined EPR and optical spectra leave little doubt about the structural assignment: two coupled high-spin ferric ions in the oxidized state, and coupled high-spin ferric and ferrous ions in the reduced state. Moreover, the spectra are consistent only with a localized model, i.e., one in which the Fe(II) site is associated with a single iron.83,83a The Fe2S2 site is inherently asymmetric, and inequivalence of the Fe(llI) sites is spectroscopically detectable in the all-ferric oxidized form.84 In fact, the localized valence trapping is present in reduced model compounds that contain no ligand asymmetry.
Mössbauer spectra provide additional and striking confirmation of the structural assignment. The spectrum of the oxidized ferredoxin (Figure 7.7) resembles strongly that of oxidized rubredoxin, indicating the presence of high-spin Fe3+, even though the net spin is zero. In the reduced form, the Mössbauer spectrum involves the superposition of signals from a high-spin Fe2+ and a high-spin Fe3+, i.e., a reduced and an oxidized rubredoxin, respectively. Clearly, the simplest interpretation of this result consistent with the S = $\frac{1}{2}$ spin state required by the EPR is the localized Fe2+-Fe3+ antiferromagnetic coupling model discussed above.
NMR studies of oxidized Fe2S2 proteins reveal broad isotropically shifted resonances for the CH2 protons of the cysteine ligands.85 Despite the coupling of the irons, the net magnetism at room temperature is sufficient to lead to large contact shifts (-30 to -40 ppm downfield from TMS). The assignment of the resonance was confirmed with the synthesis and spectroscopic analysis of model compounds.86 Extensive NMR studies of the Fe2S2 proteins have been reported.87,87a
Resonance Raman spectra of Fe2S2 sites88,89,90 reveal many bands attributable to Fe-S stretching. Detailed assignments have been presented for the four bridging and four terminal Fe-S modes. A strong band at 390 cm-1, which shifts on 34S sulfide labeling, is assigned to the A1g "breathing" mode; another band at 275 cm-1 is assigned to B3u symmetry in point group D2h.57,88 Spectroscopic differences in the terminal, Fe-S(Cys) stretches between plant ferredoxins and adrenodoxin (which also differ somewhat in redox potential) seem to reflect different conformations of the cysteine ligands in the two classes. Evidence for asymmetry of the iron atoms is found in the intensity of the resonance enhancement of certain modes.
Rieske Centers
Within the class of Fe2S2 ferredoxins there is a subclass called the Rieske proteins, or the Rieske centers.47,91,92 The Rieske iron-sulfur centers are found in proteins isolated from mitochondria and related redox chains.47,92 In addition, the phthalate dioxygenase system from Pseudomonas cepacia93,94 contains one Fe2S2 Rieske center as well as one additional nonheme Fe atom. Although the Rieske centers appear to contain an Fe2S2 core, there is extensive evidence for nonsulfur ligands coordinated to at least one of the Fe atoms. The proposed model in Scheme (7.3) has two imidazole ligands bound to one Fe atom. The nitrogen atoms are seen in ENDOR (Electron Nuclear Double Resonance) experiments,93 and are manifest in EXAFS spectra, which are consistent with the presence of a low-Z (atomic number) ligand bound to iron.94 The potentials for the Rieske proteins range from +350 to -150 mV,47 in contrast to the plant-type Fe2S2 centers, which range from -250 to -450 mV. The strong dependence of redox potential on pH95 suggests a possible role in coupling protonand electron-transfer processes.
$\tag{7.3}$
Fe2S2 Models
Although spectroscopic studies led to the correct deduction of the structure of the Fe2S2 core, the synthesis of model compounds containing this core provided unequivocal confirmation. The model compounds allowed detailed structural analysis unavailable for the proteins. Moreover, by using a uniform set of peripheral ligands, properties inherent to the Fe2S2 core could be discerned.
The Fe2S2 core has been synthesized by several routes86,96,96a,b,c,d (see Figure 7.10). For example, the reaction of Fe(SR)42-, the ferrous rubredoxin model, with elemental sulfur produces the complex Fe2S2(SR)42-. In this reaction the sulfur presumably oxidizes the Fe2+ to Fe3+, being reduced to sulfide in the process. The Fe2S2 core has been prepared with a variety of peripheral S-donor ligands. Metrical details for Fe2S2(SC6H4-p-CH3)42- are given in Table 7.3. Notable distances are the Fe-S (bridging) distance of 2.20 Å, the Fe-S (terminal) distance of 2.31 Å, and the Fe-Fe distance of 2.69 Å.
Table 7.3: Structural parameters for $Fe_2S_2(SC_6H_4-p-CH_3)_4^{2-}$. a) Data from Reference 211.
Atomsa Distance Å Atomsa Anglea
Fe-Fe 2.691 (1) Fe-S-Fe 75.3
Fe-S1 (bridge) 2.200 (1) S-Fe-S 104.6
Fe-S2 (bridge) 2.202 (1) S-Fe-S 115.1
Fe-S3 2.312 (1) S-Fe-S 105.4
To date, all analogue systems structurally characterized contain the Fe3+-Fe3+ fully oxidized form. Attempts to isolate the Fe3+-Fe2+ form have so far failed. However, the mixed-valence Fe2S2 form can be generated and trapped by freezing for spectroscopic examination.97,98 Mössbauer spectroscopy reveals the presence of distinct Fe2+ and Fe3+ ions, as found in the proteins, clearly showing that "trapped" valence states are an inherent characteristic of the Fe2S22+ core and are not enforced by the protein.97,98
The existence of noncysteine-bound Fe2S2 cores in Rieske-type proteins has led to attempts to synthesize complexes with oxygen and nitrogen ligands.99-101 Characterized species include Fe2S2(OC6H5)42-, Fe2S2(OC6H4-p-CH3)42-, Fe2S2(C4H4N)42-, and Fe2S2(L)22-, where L is a bidentate ligand.
$\tag{7.4}$
The potentially tridentate ligand
$\tag{7.5}$
acts in a bidentate fashion, binding through S and O but not N.
No Fe2S2 complexes containing mixed S,N terminal ligands, such as those suggested for the Rieske site, have been prepared. The Se2- bridged analogue has been prepared for some of the complexes.102,103 | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/07%3A_Ferrodoxins_Hydrogenases_and_Nitrogenases_-_Metal-Sulfide_Proteins/7.02%3A_Iron-sulfur_Proteins_and_Models_%28Part_2%29.txt |
Fe4S4 Ferredoxins (including HiPIPs)
We now turn our attention to proteins containing the Fe4S4 center. Historically, within this class a strong distinction was made between the "ferredoxins," which are low-potential (as low as -600 mV in chloroplasts) iron-sulfur proteins, and the "HiPIPs" = High Potential Iron Proteins, which have positive redox potentials (as high as +350 mV in photosynthetic bacteria). Although the HiPIP designation is still useful, proteins of both high and low potential are considered ferredoxins, whose key defining feature is the presence of iron and acid-labile sulfide.13
The Fe4S4 proteins participate in numerous electron-transfer functions in bacteria, and in some organisms (such as Clostridium) are the immediate electron donors for the nitrogenase and/or hydrogenase enzymes. The function of the HiPIPs seems obscure at present. In addition, Fe4S4 centers have been shown or postulated to occur in numerous microbial, plant, and mammalian redox enzymes, including nitrate reductase,104 sulfite reductase,24 trimethylamine dehydrogenase,105 succinate dehydrogenase,73,106 hydrogenase, and, possibly, in altered forms, nitrogenase. Table 7.1 lists some of the Fe4S4 ferredoxins and their properties.
In the Fe2S2 ferredoxins, combined spectroscopic, analytical, and model-system work led to an unequivocal assignment of the structural nature of the active site long before the crystallography was done. In contrast, for Fe4S4 systems and in particular the 8Fe-8S = 2Fe4S4 systems from bacteria, the initial chemical suggestions were fallacious, and even the number and stoichiometry of the clusters were in doubt. In these cases, crystallography provided the definitive structural information.
The first indication of the presence of the "thiocubane" structure came in 1968, when a 4-Å resolution study107 indicated a compact cluster of potentially tetrahedral Fe4 shape in the HiPIP from Chromatium vinosum. This finding did not lead to much excitement, since it was not yet appreciated that HiPIPs and ferredoxins were structurally similar. In 1972, the high-resolution structure solution of both Chromatium HiPIP108 and the 8Fe ferredoxin from Peptococcus aerogenes (formerly Microbacter aerogenes)101 confirmed the presence of virtually identical thiocubane clusters in the two proteins.108 Moreover, the structures for both oxidized and reduced HiPIP were deduced, and these revealed that the Fe4S4 cluster remained intact during the redox interconversion.109 Subsequently, four-iron clusters have been crystallographically confirmed in an Fe4S4 ferredoxin from Bacillus thermoproteolyticus,110,110a in Azotobacter vinelandii ferredoxin I (also previously called Shethna Fe-S protein II), which also contains a 3Fe-4S cluster,111,112 in the active form of aconitase,113 and in sulfite reductase, where the cluster is probably bridged by cysteine sulfur to a siroheme.
In all the proteins characterized to date, the Fe4S4 clusters adopt the thiocubane structure,108 which is discussed at greater length in the section on models. The clusters are usually bound to their proteins by four cysteine residues. As shown in Figure 7.11, in the P. aerogenes protein the two Fe4S4 clusters are bound by cysteines numbered 8, 11, 14, 18, 35, 38, 41, and 45.101,114 The presence of the Cys-x-x-Cys unit is again apparent. However, this sequence seems prominent in all Fe-S proteins, and so is not specific for a particular Fe-S site. At first glance one might expect one cluster to be bound by cysteines 8,11,14, and 18, the other by cysteines 35, 38, 41, and 45. Actually, one cluster is bound by cysteines 8, 11, 14, and 45, the other by cysteines 35, 38, 41, and 18. The binding of a given cluster by cysteine residues from different portions of the polypeptide chain apparently helps stabilize the tertiary structure of the protein and brings the two clusters into relatively close proximity, the center-center distance being 12 Å.114
The C. pasteurianum protein displays weak magnetic coupling, which leads to an unusual EPR spectrum115 consistent with the 12-Å cluster-cluster separation. However, the redox potentials for the two sites seem virtually identical at -412 mV, thus allowing the 8Fe ferredoxin to deliver two electrons at this low redox potential.115 Significant sequence identity indicates the likelihood that other 8Fe ferredoxins, such as the well-studied one from C. pasteurianum,116-118 have quite similar structures.
The thiocubane unit of Fe4S4 proteins can exist in proteins in at least three stable oxidation states. This so-called three-state modeI74,109,119,120 contrasts dramatically with the situation for Rd(1Fe), Fe2S2, and Fe3S4 systems, in which only two oxidation states are accessible through simple electron transfer for each center. For the thiocubane structure, the three accessible states can be designated Fe4S43+, Fe4S42+, and Fe4S4+, corresponding to [Fe(III)2Fe(II)], [Fe(III)2Fe(II)2], and [Fe(III)Fe(II)3] valence-state combinations, respectively. It is crucial to note that, in sharp contrast to the Fe2S2 and Fe3S4 sites, the oxidation states are not localized in the Fe4S4 clusters. In most cases, each Fe atom behaves as if it had the same average oxidation level as the other Fe atoms in the cluster. The redox interconversion of the Fe4S4 sites is shown in Figure 7.12. The Fe4S43+ $\rightleftharpoons$ Fe4S42+ couple is the high-potential redox couple characteristic of HiPIPs; the Fe4S42+ $\rightleftharpoons$ Fe4S4+ couple is responsible for the low-potential process characteristic of the classical ferredoxins. In any given protein under physiological conditions, only one of the two redox couples appears to be accessible and functional.
Both the Fe4S4+ and the Fe4S43+ states of the thiocubane cluster are paramagnetic and display characteristic EPR spectra (Figure 7.6C,D). The Fe4S43+ site in reduced ferredoxins46,48,49,119 displays a rhombic EPR signal (Figure 7.6C) with g = 1.88, 1.92, and 2.06. The oxidized form of low-potential ferredoxins is EPR-silent, and attempts to "superoxidize" it to achieve the Fe4S43+ state invariably lead to irreversible cluster decomposition, probably through a 3Fe-4S structure. The Fe4S43+ signal is usually referred to as the HiPIP signal (Figure 7.6D) and shows distinct g values at 2.04(g$\perp$) and 2.10(g$\parallel$); it is present in oxidized HiPIP but absent in reduced HiPIP.46,119 Reduction of HiPIP to a "super-reduced" state apparently occurs under partially denaturing conditions in aqueous DMSO.108 The observed axial EPR signal with g = 1.94 and 2.05 is assigned to the Fe4S4+ state characteristic of reduced ferredoxins. This result108 is consistent with structural and spectroscopic identity of the HiPIP and Fd sites, as required by the three-state model of the Fe4S4 proteins (Figure 7.12).
In Fe4S4 centers at each level of oxidation, electronic transitions give rise to characteristic visible and UV spectra, although the delocalized nature of the electronic states makes detailed assignment difficult. MCD spectra of clusters in the three states of oxidation are clearly distinguishable from each other and from MCD of Fe2S2 clusters.43,119 MCD, magnetic susceptibility, and Mössbauer spectra provide evidence that the S = $\frac{1}{2}$ state, whose EPR signal is so distinct in reduced ferredoxins, may coexist at higher T with S = $\frac{3}{2}$ and perhaps even higher spin states. Indeed, recent studies with model systems121,122 and theoretical treatments123,124 clearly support the ability of the Fe4S4 cluster to display a number of spin states that are in labile equilibria, which are influenced, perhaps quite subtly, by local structural conditions. The iron protein of nitrogenase also displays this behavior.
The Mössbauer spectra of Fe4S4 centers of ferredoxins reveal the equivalence of the Fe sites, and quadrupole splittings and isomer shifts at averaged values for the particular combination of oxidation states present.3,51,52 Representative spectra are shown in Figure 7.7. Magnetic coupling is seen for the paramagnetic states.
Resonance Raman spectra (and IR spectra) have been extensively investigated in C. pasteurianum ferredoxin and in model compounds.57,125 Selective labeling of either thiolate sulfur or sulfide sulfur with 34S allows modes associated with the Fe4S4 core to be distinguished from modes associated with the FeSR ligands. The band at 351 cm-1 is assigned to Fe-SR stretching, and Fe4S4 modes occur at 248 and 334 cm-1 in reduced ferredoxin from C. pasteurianum. There is little difference between the oxidized and reduced spectra, although an extra band at 277 cm-1 seems present in the oxidized protein. The Fe4Se4 substituted protein has also been studied.125
As in the 1Fe and 2Fe proteins, 1H NMR spectra reveal resonances from contact-shifted-CH2-groups of cysteinyl residues.125a However, unlike the other proteins, where all states are at least weakly magnetic, only the reduced ferredoxin and the oxidized HiPIP states show contact shifts.87a,125a,b,c
EXAFS studies on proteins and on model compounds clearly identify the Fe-S distance of ~2.35 Å and an Fe-Fe distance of 2.7 Å. These distances, as expected, vary only slightly with state of oxidation.125d
Fe4S4 Models
Judging from the ease with which models of Fe4S4 are prepared under a variety of conditions and their relative stability, the Fe4S42+ core structure seems to be a relatively stable entity, a local thermodynamic minimum in the multitude of possible iron-sulfide-thiolate complexes. The initial preparation and structural characterization126,127 of the models showed that synthetic chemistry can duplicate the biological centers in far-simpler chemical systems, which can be more easily studied in great detail.
The general synthetic scheme for Fe-S clusters is shown in Figure 7.13. Many different synthetic procedures can be used to obtain complexes with the Fe4S4 core.126-138,138a,b The multitude of preparative procedures is consistent with the notion that the Fe4S42+ core is the most stable entity present and "spontaneously self-assembles" when not limited by stoichiometric constraints.
The thiocubane structure can be viewed as two interpenetrating tetrahedra of 4Fe and 4S atoms. The 4S tetrahedra are the larger, since the S-S distance is ~3.5 Å, compared with the Fe-Fe distance of ~2.7 Å. The S4 tetrahedron encloses ~2.3 times as much volume as does the Fe4 tetrahedron.128 Key distances and angles for Fe4S4(SCH2C6H5)42- given in Table 7.4 are extremely similar to those found in oxidized ferredoxin and reduced HiPIP centers in proteins.127
Table 7.4 - Structural parameters for Fe4S4(SCH2C6H5)42-.
a) Data from References 126 and 127.
Atomsa Average Distances Number of Bonds Type
Fe(1)-S(3) 2.310 (3) 8 Sulfide
Fe(1)-S(2) 2.239 (4) 4 Sulfide
Fe(1)-S(5) 2.251 ( 3) 4 Thiolate
Fe(1)-Fe(2) 2.776 (10) 2
Fe-Fe(other) 2.732 (5) 4
Atomsa Average Distances Number of Bonds Type
Fe-S-Fe 73.8 (3) 12
S-Fe-S 104.1 (2) 12 Sulfide-Fe-Sulfide
S-Fe-S 111.7-117.3 12 Sulfide-Fe-Thiolate
The idealized symmetry of Fe4S42+ model systems is that of a regular tetrahedron, i.e., Td. Though the distortion of the cube is quite pronounced, all known examples of the Fe4S2+ core show distortion, which lowers the symmetry at least to D2d. In most Fe4S2+ core structures, this distortion involves a tetragonal compression, which leaves four short and eight long Fe-S bonds.
Complexes with non-S-donor peripheral ligands have been prepared and studied. The halide complexes Fe4S4X42- (X = CI-, Br-, I-) have been prepared, and serve as useful starting points for further syntheses.129-133 The complex Fe4S4(OC6H5)42- can be prepared134 from the tetrachloride (or tetrathiolate) thiocubane by reaction with NaOC6H5 (or HOC6H5). There are a few examples of synthetic Fe4S42+ cores in which the peripheral ligands are not identical. For example, Fe4S4Cl2(OC6H5)22- and Fe4S4Cl2(SC6H5)22- have structures characterized by D2d symmetry.135 The complexes Fe4S4(SC6H5)2[S2CN(C2H5)2]2- and Fe4S4(SC6H4OH)42- are similarly asymmetric, containing both four- and five-coordinate iron.136-138 The presence of five-coordinate iron in the Fe4S4 cluster is notable, since it offers a possible mode of reactivity for the cluster wherever it plays a catalytic role (such as in aconitase). Complexes with Fe4Se2+ and Fe4Te42+ cores have also been prepared.138c,d
One structural analysis of Fe4S4(SC6H5)43-, which contains the reduced Fe4S4+ core, revealed a tetragonal elongation139 in the solid state. In contrast, analysis of Fe4S4(SCH2C6H5)43- revealed a distorted structure possessing C2v symmetry.102 It would appear that the Fe4S4+ clusters maintain the thiocubane structure, but are nevertheless highly deformable. Interestingly, when the solidstate C2v structure, Fe4S4(SCH2C6H5)43-, is investigated in solution, its spectroscopic and magnetic behavior change to resemble closely those of the Fe4S4(SC6H5)43- cluster,140 which does not change on dissolution. The simplest interpretation assigns the elongated tetragonal structure as the preferred form for Fe4S4+ cores with deformation of sufficiently low energy that crystal packing (or, by inference, protein binding forces) could control the nature of the distortions in specific compounds.128 The elongated tetragonal structure has four long and eight short bonds in the core structure. The terminal (thiolate) ligands are 0.03-0.05 Å longer in the reduced structure, consistent with the presence of 3Fe(ll) and 1Fe(III) in the reduced form, compared to 2Fe(II) and 2Fe(III) in the oxidized form. There is no evidence for any valence localization.128
The oxidized Fe4S43+ core defied isolation and crystallization in a molecular complex prior to the use of sterically hindered thiolate ligands. With 2,4,6 tris(isopropyl)phenylthiolate, the Fe4S4L4- complex could be isolated and characterized.141 The structure is a tetragonally compressed thiocubane with average Fe-S and Fe-SR distances 0.02 and 0.04 Å shorter than the corresponding distances in the Fe4S4L42- complex. Again, there is no evidence for Fe inequivalence or more profound structural distortion in this 3Fe(III)-1Fe(II) cluster. Clearly, the Fe4S4 clusters have highly delocalized bonding.
Evidence from model systems using sterically hindered thiolate ligands indicates the existence of an Fe4S44+, i.e., all-ferric fully oxidized cube.142 The existence of the complete series Fe4S4[(Cy)3C6H2S]4n (Cy =cyclohexyl; n = 0, -1, -2, -3) is implied by reversible electrochemical measurements. Clearly, five different states of the Fe4S4 core—including the (at least) transient fully oxidized state and the all-ferrous fully reduced state—may have stable existence. Although only the central three states have been shown to exist in biological contexts, one must not rule out the possible existence of the others under certain circumstances.
Recently, specifically designed tridentate ligands have been synthesized that bind tightly to three of the four Fe atoms in the thiocubane structure.143,143a,b The remaining Fe atom can then be treated with a range of reagents to produce a series of subsite-differentiated derivatives and variously bridged double-cubane units. These derivatives illustrate the potential to synthesize complexes that mimic the more unusual features of Fe4S4 centers that are bound specifically and asymmetrically to protein sites. The recently synthesized complex ion [(Cl3Fe4S4)2S]4-, containing two Fe4S4 units bridged by a single S2- ligand, illustrates the potential coupling of known clusters into larger aggregates.143c
The model-system work has made an important contribution to our understanding of the Fe4S4 centers. The existence of three states, the exchange of ligands, the redox properties, the metrical details of the basic Fe4S4 unit, and the subtleties of structural distortion can each be addressed through the study of models in comparison with the native proteins. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/07%3A_Ferrodoxins_Hydrogenases_and_Nitrogenases_-_Metal-Sulfide_Proteins/7.03%3A_Iron-sulfur_Proteins_and_Models_%28Part_3%29.txt |
Core Extrusion/Cluster Displacement Reactions
Synthetic model-system work led to the realization that the cluster cores can exist outside the protein and undergo relatively facile ligand-exchange reactions.128 This behavior of the purely inorganic complexes allowed core extrusion reactions128,144 to be developed. The basic assumption behind these reactions is that the cluster core retains its integrity when it is substituted by low-MW thiolates, especially aryl thiolates, which replace the cysteinyl ligands that bind it to the protein. In order to free the cluster from the protein, one must at least partially denature the protein, usually by using ~80 percent aqueous solution of a polar aprotic solvent, such as DMSO or HMPA. The resulting inorganic clusters can be identified and quantified by measurement of their characteristic electronic absorption or NMR spectra. An alternative approach involves transferal of the unknown cluster in question to an apoprotein that binds a cluster of known type.145
Since the 1Fe, 2Fe, and 4Fe sites are each usually bound to the protein by four cysteine residues, it is perhaps not surprising that there have been reports58 of interconversion of cluster types bound to a given protein. Specifically, in 90 percent aqueous DMSO, the single Fe site in rubredoxin from C. pasteurianum is converted to an Fe4S4 cluster by the addition of sodium sulfide, ferrous chloride, and ferric chloride in ratio 4:2:1. Presumably the spacing and geometric disposition of the cysteines are suitable to bind a single Fe or the Fe4S4 cluster, which is readily formed under the reaction conditions. Another example of cluster rearrangement involves the three-iron center discussed below that does not extrude as an Fe3S4 center. Rather, at least under certain conditions, the Fe3Sx center rearranges to form Fe2S2 centers.146 The facile interconversion of the Fe clusters demonstrated the lability of Fe-S systems, and indicates that caution must be exercised in interpreting the results of cluster-displacement reactions.
Fe3S4 Centers
Three-iron centers are a comparatively recent finding,119,146 and the full scope of their distribution is not yet known. Although they have now been confirmed in dozens of proteins, it often remains uncertain what physiological role these centers play. Indeed, since Fe3S4 centers can be produced as an artifact upon oxidation of Fe4S4 centers, it has been suggested that 3Fe centers may not be truly physiological, and could be side products of aerobic protein isolation. This caveat notwithstanding, the 3Fe sites are being found in more and more proteins and enzymes. Their physiological raison d' être may be more subtle than that of their 1, 2, and 4Fe cousins; we should certainly try to find out more about them. Some proteins containing Fe3S4 centers are listed in Table 7.1.
The 3Fe center was first recognized147 in the protein ferredoxin I from the anaerobic nitrogen-fixing bacterium Azotobacter vinelandii. The protein is called Av FdI for short. It is instructive to sketch historically the evolution of our understanding of this protein. Av ferredoxin I was reported to have 6 to 8 Fe atoms and was first thought to resemble the clostridial 8Fe ferredoxins. However, unlike the clostridial protein, the Av FdI clusters appeared to have two quite different redox couples at +320 and -420 mV. Although it might have been thought that this protein contained one HiPIP-type and one Fd- or "ferredoxin"-type Fe4S4 cluster, the protein as isolated had an EPR signal with g = 2.01, which differed significantly (Figure 7.6) from that of an oxidized HiPIP or a reduced Fd.148 Cluster extrusion reactions also seemed to indicate the presence of an unusual cluster type.149
Fortunately, the protein was crystallized, and could be studied by x-ray diffraction. Unfortunately, the initial conclusions150 and subsequent revisions151,152 of the crystal-structure analysis have proven to be wrong, teaching us in the process that protein x-ray crystallography, taken alone, does not always provide definitive results. Specifically, the first crystallographic report suggested the presence of a conventional Fe4S4 cluster and a smaller packet of electron density that was assigned as a 2Fe-2S center.150 However, upon further refinement, and following the formulation of the 3Fe center by Mössbauer spectroscopy,147 a 3Fe-3S center was identified and refined.151,152 The "refined" Fe3S3 center was a six-membered alternating iron-sulfide ring with an open, almost flat, twistedboat conformation (Figure 7. 14). The Fe-Fe separation of 4.1 Å and the structural type was unprecedented, and did not agree with the results of resonance Raman spectroscopy,153 with x-ray absorption spectroscopy on the native protein or on samples from which the Fe4S4 center was removed,103 or even with stoichiometry, which eventually led to the reformulation154 of the cluster as Fe3S4. The x-ray absorption studies (EXAFS) clearly led to the assignment of a 2.7-Å Fe-Fe distance for the 3Fe cluster.103
In parallel with the studies on Av FdI, two additional proteins played key roles in the resolution of the nature of the Fe3 cluster. These are FdII from Desulfovibrio gigas and aconitase from beef heart. Each contains (under certain conditions) only Fe3S4 sites, thus enabling more definitive structural, stoichiometric, and spectral information to be acquired. Studies on these proteins using EXAFS,155 Mössbauer,52,156,157 EPR,146 and resonance Raman (to which we will return briefly) clearly favor the closed structure shown in Figure 7.14C. Indeed, x-ray crystallography on aconitase by the same group that did the initial x-ray work on Av FdII revealed the compact structure in agreement with the spectroscopy.20
In 1988 the structural error in the crystallography of Av FdI was found by two groups, and a new refinement in a corrected space group led to a structure in agreement with the spectroscopy.111,112 The Fe3S4 cluster has the apoFe thiocubane structure, with each iron atom bound to the protein by a single cysteinyl thiolate. Clearly, even x-ray crystallography is potentially fallible, and its findings must be critically integrated with the data from other techniques in arriving at full structural definition of metalloenzyme sites.
The studies on the ferredoxins from D. gigas present an interesting lesson on the lability of the Fe-S cluster systems. Two distinct proteins from D. gigas, FdI and FdII, contain the same polypeptide chain (6 kDa) in different states of aggregation,158,159 Whereas FdI is a trimer containing three Fe4S4 clusters, FdII is a tetramer that contains four Fe3S4 clusters. The ferredoxins differ in their redox potentials and appear to have different metabolic functions in D. gigas. The oxidation of D. gigas FdI with Fe(CN)63- leads to FdII, and treatment of FdII with iron salts leads to FdI. The D. gigas system reveals the lability and interconvertibility of Fe-S clusters. The recently reported160 crystal structure of D. gigas FdII shown in Figure 7.15 confirms the partial (apoFe) thiocubane Fe3S4 center. The iron atoms are ligated by three cysteinyl residues from protein side chains. The cube missing an iron is now firmly established as a viable structural type.
The aconitase system presents yet another fascinating story.159a Aconitase is a key enzyme in the Krebs cycle, catalyzing the conversion of citrate and isocitrate through the intermediacy of cis-aconitate, as shown in Equation (7.6).
$\tag{7.6}$
This is a hydrolytic nonredox process, and for some time it was thought that aconitase was a simple Fe2+ protein wherein the ferrous iron was involved in the Lewis-acid function of facilitating the hydrolytic reaction. Indeed, aconitase is inactive when isolated from mitochondria, and requires the addition of Fe2+ to achieve activity.
Surprisingly, the isolated aconitase was found by analysis and Mossbauer spectroscopy to possess an Fe3S4 site in its inactive form.161 Low-resolution crystallographic study supports the presence of an apo-Fe thiocubane, Fe3S4 structure in aconitase.162 Resonance Raman163 and EXAFS103,155 studies clearly fingerprint the Fe3S4 cluster. The current hypothesis for aconitase activation involves the Fe3S4 thiocubane fragment reacting with Fe2+ to complete the cube, which is the active form of the enzyme.164 Recent crystallographic studies113 confirm the presence of a complete cube in the activated aconitase. The dimensions and positioning of the Fe3S4 and Fe4S4 centers in the cube are virtually identical. The added Fe2+ iron atom is ligated by a water (or hydroxide) ligand, consistent with the absence of any cysteine residues near the exchangeable iron. Since this water (or a hydroxide) is also present in the Fe3S4 system, one wonders whether a small ion such as Na+ might be present in the Fe3S4 aconitase system.
Questions of detailed mechanism for aconitase remain open. ENDOR spectroscopy165 shows that both substrate and water (or OH-) can bind at the cluster. Does one of its Fe-atom vertices play the Lewis-acid role necessary for aconitase activity? Is the Fe3S4 $\rightleftharpoons$ Fe4S4 conversion a redox- or iron-activated switch, which works as a control system for the activity of aconitase? These and other questions will continue to be asked. If aconitase is indeed an Fe-S enzyme with an iron-triggered control mechanism, it may be representative of a large class of Fe3/Fe4 proteins. Other hydrolytic enzymes containing similar Fe-S centers have recently been reported.166,166a
Spectroscopically, the Fe3S4 center is distinct and clearly distinguishable from 1Fe, 2Fe, and 4Fe centers. The center is EPR-active in its oxidized form, displaying a signal (Figure 7.6) with g = 1.97, 2.00, and 2.06 (D. gigas FdII).158 The Mössbauer spectrum (Figure 7.7) shows a single quadrupole doublet with $\Delta$Q = 0.53 nm/s and isomer shift of 0.27 nm, suggesting the now familiar highspin iron electronic structure.158,159 In its reduced form, the center becomes EPR-silent, but the Mössbauer spectrum now reveals two quadrupole doublets of intensity ratio 2:1. The suggestion of the presence of a 3Fe center was first made based on this observation.147 The picture of the reduced Fe3S4 state that has emerged involves a coupled, delocalized Fe2+/Fe3+ unit responsible for the outer doublet, with a single Fe3+ unit responsible for the inner doublet of half the intensity. The oxidized state contains all Fe3+ ions, which are coupled in the trinuclear center.
EXAFS studies were consistent and unequivocal in finding an Fe-Fe distance of ~2.7 Å in all putative Fe3S4 proteins.103,155,167 Resonance Raman spectra compared with those of other proteins and of model compounds with known structures168,169 for other metals also favored the structure of Figure 7.14B. Clearly, what has been termed the spectroscopic imperative170 has been crucial in the successful elucidation of the 3Fe structure. An interesting excursion has led to isolation of what are presumed to be ZnFe3S4, CoFe3S4, and NiFe3S4 thiocubane structures by adding Zn2+, Co2+, or Ni2+, respectively, to proteins containing reduced Fe3S4 cores.170a,b,c These modified proteins provide interesting electronic structural insights and, potentially, new catalytic capabilities.
Fe3 Model Systems
To date, the Fe3S4 center is the only structurally characterized biological iron-sulfide center that does not have an analogue in synthetic Fe chemistry. In fact, the closest structural analogue21,170d is found in the Mo3S44+ or V3S43+ core in clusters such as Mo3S4(SCH2CH2S)32-, whose structure is shown in Figure 7.16.
The resonance Raman spectrum169 of this complex bears a close resemblance to that of the D. gigas ferredoxin II. Since the vibrational bands responsible for the resonance Raman spectrum are not strongly dependent on the electronic properties, it is not surprising that an analogue with a different metal can be identified using this technique.
In synthetic Fe chemistry, although there is no precise structural analogue, it is instructive to consider three types of trinuclear and one hexanuclear center in relation to the three-iron biocenter.
The trinuclear cluster Fe3S[(SCH2)C6H4]32- is prepared171,172 by the reaction of FeCI3, C6H4(CH2SH)2, Na+OCH3-, and p-CH3-C6H4-SH in CH3OH. As shown in Figure 7.17A, this cluster has, like the biocluster, a single triply bridging sulfide ion but, unlike the biocluster, it uses the sulfur atoms of the ethane 1,2-dithiolate as doubly bridging, as well as terminal, groups. The inorganic ring Fe3(SR)3S63- (X = CI, Br) has a planar Fe3(SR)3 core, which resembles the now-discredited structure for Av FdI173 (Figure 7. 14A).
The complex Fe3S4(SR)43- is prepared174,175 by reaction of Fe(SR)42- with sulfur. The x-ray-determined structure reveals two tetrahedra sharing a vertex with the linear Fe-Fe-Fe array shown in Figure 7.17B. This complex has distinctive EPR, Mössbauer, and NMR spectra that allow it to be readily identified174,175,175a Interestingly, after the complex was reported, a study of (denatured) aconitase at high pH (>9.5) revealed that the thiocubane fragment Fe3S4 site in that enzyme rearranged to adopt a structure that was spectroscopically almost identical with that of the linear complex.176 Although the state may not have any physiological significance, it does show that Fe-S clusters different from the common (conventional) ones already discussed could be important in proteins under certain physiological conditions or in certain organisms; i.e., iron centers first identified synthetically may yet prove to be present in biological systems.
A synthetic cluster that displays features related to the biological Fe3S4 cluster is the hexanuclear cluster Fe6S9(SR)24- shown in Figure 7.18. This cluster contains two Fe3S4 units bridged through their diiron edges by a unique quadruply bridging S2- ion ($\mu_{4}$-S2-) and by two additional $\mu_{2}$S2- bridges. The inability of synthetic chemists to isolate an Fe3S4 analogue may indicate that in proteins this unit requires strong binding. Significant sequestration by the protein may be needed to stabilize the Fe3S4 unit against oligomerization through sulfide bridges or, alternatively, rearrangement to the stable Fe4S4 center.
Fe-S Chemistry: Comments and New Structures
The first successful model system for an iron-sulfur protein was an analogue of the Fe4S4 system, i.e., the system with the largest presently established biological Fe cluster. The reactions used to synthesize the cluster shown in Figure 7.13 are said to involve self-assembly, meaning that starting materials are simply mixed together, and thermodynamic control causes the cluster to assemble in its stable form. Interestingly, the Fe4S4-containing proteins, such as those of C. pasteurianum, are considered to be among the most ancient of proteins. Perhaps on the anaerobic primordial Earth, Fe-S clusters self-assembled in the presence of protein ligands to form the progenitors of the modem ferredoxins.
Much progress has been made in synthetic chemistry, and it is clear that both understanding and control of Fe-S chemistry are continuing to grow. New preparations for known clusters continue to be found, and new clusters continue to be synthesized, Although many of the new clusters appear to be abiological, we should not ignore them or their potential. They add to our understanding of Fe-S chemistry in general, and serve as starting points in the study of heteronuclear clusters, There is also the distinct possibility that one or more of these synthetic clusters represent an existing biological site that has not yet been identified in an isolated system.
Among the "nonbiological" structures that have been synthesized are complexes with Fe6S63+/2+ cores177,178 including the thioprismanes,108,109,177-179 the octahedron/cube Fe6S83+ cores,180,181 the Fe6S92- cores176a,b discussed above, the adamantane-like Fe6(SR)104- complexes related to Zn and Cu structures in metallothioneins,182 basket Fe6S62+/Fe6S6+ cores,182a,b,c,d monocapped prismatic Fe7S63+ cores,183 the cube/octahedron Fe8S65+ cores,179 and the circular Na+-binding Fe18S3010- unit.183a,b Representative ions are shown in Figure 7.19. Some of these cores are stabilized by distinctly nonbiological phosphine ligands. Nevertheless, one should not a priori eliminate any of these structures from a possible biological presence. Indeed, recently a novel, apparently six-iron protein from Desulfovibio gigas has been suggested183c to have the thioprismane core structure first found in model compounds.
Detection of Fe-S Sites
Several recent reviews have concentrated on the ways in which the various Fe-S centers can be identified in newly isolated proteins.5,6,184 It is instructive to summarize the central techniques used in the identification of active sites. Optical spectra are usually quite distinctive, but they are broad and of relatively low intensity, and can be obscured or uninterpretable in complex systems. MCD spectra can give useful electronic information, especially when the temperature dependence is measured. EPR spectra, when they are observed, are distinctive, and are usually sufficiently sharp to be useful even in complex systems. Mössbauer and resonance Raman spectroscopies have each been applied with good effect when they can be deconvoluted, and NMR and magnetic susceptibility have given important information in some simple, lower-MW protein systems. X-ray absorption spectra, especially EXAFS, give accurate Fe-S and Fe-Fe distances when a single type of Fe atom is present. Analytical and extrusion data complement the spectroscopic and magnetic information. Extrusion data must be viewed with considerable caution, because of possible cluster-rearrangement reactions. Even x-ray crystallography has led to incorrect or poorly refined structures. In general, no one technique can unequivocally identify a site except in the very simplest systems, and there is continued need for synergistic and collaborative application of complementary techniques to a given system.
Redox Behavior
Figure 7.20 shows the ranges of redox behavior known for Fe-S centers. Clearly, the Fe-S systems can carry out low-potential processes. The rubredoxins cover the mid-potential range, and the HiPIPs are active in the high-potential region. The lack of extensive Fe-S proteins in the positive potential region may reflect their instability under oxidizing conditions and their preemption by Mn, Cu, or heme-iron sites (such as in cytochrome c), which function in this region. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/07%3A_Ferrodoxins_Hydrogenases_and_Nitrogenases_-_Metal-Sulfide_Proteins/7.04%3A_Iron-sulfur_Proteins_and_Models_%28Part_4%29.txt |
In making the transition from the relatively simple electron-transfer proteins to the far more complex Fe-S-containing enzymes, we must recognize that the difference in our degree of understanding is enormous. Not only are the catalytic proteins ten to twenty times larger than the redox proteins, but they often also have several subunits and multiple copies of prosthetic groups. Moreover, very few crystal structures are known for the redox enzymes, and none is known at high resolution. In the absence of three-dimensional protein structural information, we do not know the arrangement, relative separation, or orientation of the prosthetic groups. Finally, studies on model systems have not yet approached the sophisticated state that they have for the structurally known centers.
In general, multicomponent redox enzyme systems appear to be organized in two distinct ways to effect their substrate reactions. In the first mode, the enzyme is designed to bring the oxidant and reductant together so that they may directly interact. For example, oxygenases bring O2 and an organic molecule together, and activate one or both of these reactants to cause them to react directly with each other. This mode can be called proximation, as the reactants are brought near each other by the enzyme catalyst.
In contrast to proximation, many redox enzymes keep the oxidant and reductant well-separated, and use rapid (usually long-distance) internal electron transfer to bring electrons from the reductant to the oxidant. We can term their mode of action electrochemical. The oxidant and reductant are separated spatially. The enzyme provides the "anode" site to interact with the reductant, the "cathode" site to interact with the oxidant, and the wire to allow electronic flow between the "anode" and "cathode" sites. Hydrogenases and nitrogenases adopt the electrochemical mode of redox activation. In hydrogenases, the electron acceptor, even if it must formally take up hydrogen (e.g., NAD+→ NADH), does not interact at the same site as the H2. There is no direct transfer of H - from H2 to NAD+. Rather, H2 reduces the enzyme at one site, and NAD+ or other acceptors, such as methylene blue, retrieve the electrons at other sites following internal electron transfer. For nitrogenase, the redox partners are even more removed, as a separate protein, the Fe protein, delivers electrons to the FeMo protein, that eventually end up at the FeMoco site ready to reduce N2 to NH3. These enzymes work much like electrochemical cells.
Hydrogenase and Nitrogenase
Hydrogenase is the enzyme responsible for the uptake or evolution of H2. Nitrogenase is the enzyme that catalyzes the ATP-dependent reduction of N2 to NH3, with concomitant evolution of H2.
The relationship between H2 and N2 in biology is intricate.184 Metabolically, H2 use and N2 use are tightly coupled in many nitrogen-fixing organisms, with H2 serving indirectly as the reductant for N2. Moreover, H2 and N2 react in related ways with various transition-metal complexes, which are at present the closest (albeit quite imperfect) models of the enzyme active sites. The biological fixation of molecular nitrogen is dependent on iron-sulfide proteins that also contain molybdenum or vanadium. The biological production or uptake of H2 depends on the presence of iron-sulfide proteins, which often also contain nickel and sometimes selenium. Spectroscopic and model-system studies, which have played such a key role in advancing the understanding of simple Fe-S sites, are now helping to foster an understanding of these more complex enzyme sites, although we have much yet to learn about structure and mechanism in these enzymes. The remainder of this chapter seeks to convey the state of our rapidly evolving knowledge.
Hydrogenases
1. Physiological Significance
Molecular hydrogen, H2, is evolved by certain organisms and taken up by others. For either process, the enzyme responsible is called hydrogenase. The raison d' être for hydrogenases in particular organisms depends on the metabolic needs of the organism. Properties of some representative hydrogenases are given in Table 7.5.184,184a-g
Hydrogenases are found in a wide variety of anaerobic bacteria, such as the eubacterial C. pasteurianum and Acetobacterium woodii and the archaebacterial Methanosarcina barkerii. Interestingly, C. pasteurianum sometimes evolves H2 during its growth on sugars. This H2 evolution is required for continued metabolism, since it allows the organism to recycle (reoxidize) cofactors that are reduced in the oxidation of sugars (or their metabolic descendants, lactate or ethanol). In effect, H+ is acting as the terminal oxidant in clostridial metabolism, and H2 is the product of its reduction. In contrast, methanogens such as methanosarcina take up H2 and in effect use it to reduce CO2 to CH4 and other carbon products. Clearly, either H2 uptake or H2 evolution may be important in particular anaerobic metabolic contexts. The hydrogenases of the anaerobic sulfate-reducing bacteria of the genus Desulfovibrio have been particularly well-studied (see Table 7.5).
Table 7.5 - Properties of some representative hydrogenases.
Organism (designation) MW (subunits) Approximate Composition Reference
Clostridium pasteurianum
(Hydrogenase I)
60,000 (1)
12Fe
22Fe
191
194
Clostridium pasteurianum
(Hydrogenase II)
53,000 (1)
8Fe
17Fe
191
194
Acetobacterium woodii 15,000 Fe 192
Megasphaera elsdenii 50,000 (1) 12Fe 188
Desulfovibrio vulgaris
(periplasmic)
56,000 (2) 12Fe 171
Desulfovibrio gigas 89,000 (2) 11Fe, 1Ni 363
Desulfovibrio africanus 92,000 11Fe, 1Ni 364
Methanobacterium thermoautotrophicum 200,000 Fe, 1Ni 145
Methanosarcina barkeri 60,000 8-10Fe, 1Ni 77
Methanococcus vanielli 3400,000 Fe, Ni, 1Se, FAD 365
Desulfovibrio baculatus 85,000 (2) 12Fe, 1Ni, 1Se 365
In nitrogen-fixing organisms, H2 is evolved during the nitrogen-fixation process, and hydrogenase is present to recapture the reducing equivalents, which can then be recycled to fix more nitrogen. In N2-fixing organisms, such "uptake" hydrogenases can make an important contribution to the overall efficiency of the nitrogen-fixation process. In fact, certain species of rhizobia lacking the hydrogen-uptake system (hup- strains) can be made more efficient by genetically engineering the hup activity into them.
Aerobic bacteria such as Azotobacter vinelandii, Alcaligenes eutrophus, and Nocardia opaca, and facultative anaerobes, such as Escherichia coli and various species of Rhizobium and Bradyrhizobium (the symbionts of leguminous plants), also contain hydrogenase, as do photosynthetic bacteria such as Chromatium vinosum, Rhodobacter capsulatus (formerly Rhodopseudomonas capsulata), and Anabaena variabilis (a filamentous cyanobacterium). The thermophilic hydrogen oxidizer Hydrogenobacter thermophilus, which grows in alkaline hot springs above 70 °C, obviously has a critical requirement for hydrogenase.
In certain aerobic organisms, such as hydrogenomonas, H2 and O2 are caused to react (but not directly) according to the Knallgas reaction185
$2 H_{2} + O_{2} \rightarrow 2 H_{2}O \qquad \Delta G^{o} = -54.6 kcal/mol \tag{7.7}$
These organisms break up this thermodynamically highly favorable redox process by using intermediate carriers, thereby allowing the large negative free-energy change to be captured in biosynthetic capacity.
Hydrogenases seem to be especially prevalent in anaerobic, nitrogen-fixing, and photosynthetic organisms. However, although hydrogenases are obviously found widely among prokaryotes, unlike nitrogenase, their domain is not restricted to prokaryotes. Eukaryotic green algae such as Chlorella fusca and Chlamydomonas reinhardtii possess hydrogenase. The anaerobic protozoan Trichomonas vaginalis, which lacks typical aerobic organelles, such as mitochondria and peroxisomes, has an organelle called a hydrogenosome, whose function is to oxidize pyruvate to acetate, producing H2, via hydrogenase, in the process.
The various hydrogenase enzymes are all transition-metal sulfide proteins. However, before we discuss these enzymes, we turn briefly to the dihydrogen molecule and its physical and chemical properties.
2. Dihydrogen: The Molecule
Diatomic H2 has a single H—H bond formed by overlap of the two 1s orbitals of the two hydrogen atoms. In molecular orbital terms, this overlap forms bonding $\sigma$ and antibonding $\sigma$* orbitals, shown in the energy-level diagram in Figure 7.21A and displayed spatially in Figure 7.21B. The H-H distance is 0.74 Å, and the bond dissociation energy is 103.7 kcal/mole. The isotopes of hydrogen 1H1, 2H1 = 2D1, 3H1 = 3T1 are called protium (a designation seldom used), deuterium, and tritium, respectively. Deuterium, at natural abundance of 0.015 percent, is a stable isotope with nuclear spin I = 1, whereas both 1H and 3T have nuclear spin I = $\frac{1}{2}$. NMR has been fruitfully applied to all the hydrogen isotopes, including tritium. Tritium is radioactive, decaying to 3He2 by $\beta^{-}$ emission with a half-life of 12 years. The nuclear properties of deuterium and tritium make them useful as labels to probe structure and mechanism in hydrogen-containing compounds. "Exchange" reactions involving the formation of HD or HT have played a significant role in mechanistic studies of both hydrogenase and nitrogenase.
In molecular hydrogen, the existence of nuclear-spin energy levels is responsible for the distinction between ortho and para hydrogen, which correspond to the triplet and singlet (i.e., parallel and antiparallel) orientations, respectively, of the two nuclei in H2. Because of the coupling of the rotational and spin levels, ortho and para hydrogen differ in specific heat and certain other properties. The correlated orientation of the nuclear spins in para H2 has recently been shown to constitute a powerful mechanistic probe, wherein NMR may be used to trace the relative fate of the two H nuclei in the original molecule.186,187 Although this technique has not yet been applied to any enzyme systems, hydrogenase is known to catalyze the interconversion of the ortho and para forms of H2 (as does the hydrogenase analogue Pd).
Dihydrogen is a reducing agent. The H2/H+ couple at [H+] = 1 M defines the zero of the potential scale. At pH = 7 the hydrogen half-reaction
$H_{2} \rightarrow 2 H^{+} + 2 e^{-} \tag{7.8}$
has Eo' = -420 mV. Dihydrogen is therefore one of the strongest biological reductants.
Although many hydrogenases are reversible, some "specialize" in the uptake of H2. One hydrogenase has been reported188 to specialize in the evolution of H2. This "specialization" seems curious, since it appears to contradict the notion of microscopic reversibility, and seems to violate the rule that catalysts increase the speed of both forward and backward reactions without changing the course (direction) of a reaction. In fact, there is no contradiction or violation, since the overall reactions catalyzed by the various types of hydrogenases are fundamentally different. The electron acceptor in uptake hydrogenases differs from the electron donor/acceptor in the reversible hydrogenases. The difference involves structure and, more importantly, redox potential. The reaction catalyzed by the uptake hydrogenase involves an acceptor of such high positive redox potential that its reaction with H2 is essentially irreversible. The enzyme appears to be designed so that it can transfer electrons only to the high potential acceptor.
A selection of hydrogenases from various organisms is given in Table 7.5. All hydrogenases contain Fe-S centers. The hydrogenases from more than 20 organisms189 have been found to contain Ni by analysis and/or spectroscopy. Many more Ni hydrogenases are likely to be found, given the nutritional requirements189 for hydrogenase synthesis or growth on H2. Hydrogenases may be cytoplasmic (as in C. pasteurianum), membrane-bound (as in E. coli), or located in the periplasmic space (as in Desulfovibrio vulgaris). The isolation of hydrogenases is sometimes complicated by their air sensitivity or membranebound nature. Many hydrogenases have now been isolated and studied in detail; they can be divided into two categories, the iron hydrogenases and the nickeliron hydrogenases. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/07%3A_Ferrodoxins_Hydrogenases_and_Nitrogenases_-_Metal-Sulfide_Proteins/7.05%3A_Multisite_Redox_Enzymes.txt |
3. Iron Hydrogenases
The iron hydrogenases189a generally have higher activities than the NiFe enzymes, with turnover numbers approaching 106 min-1. Iron hydrogenases from four genera of anaerobic bacteria have been isolated: Desulfovibrio,190 Megasphaera,188 Clostridium,191 and Acetobacterium.192 Of these, the enzymes from Desulfovibrio vulgaris, Megasphaera elsdenii, Clostridium pasteurianum (which contains two different hydrogenases), and Acetobacterium woodii192 have been well-characterized (especially the D. vulgaris and C. pasteurianum enzymes). Although Acetobacterium and Clostridium are closely related, the other organisms are only distant cousins.193 Nevertheless, their hydrogenases display significant similarities; all contain two different types of iron-sulfur cluster, called F and H clusters,194 and carbon monoxide is a potent inhibitor (although this has not been reported for the M. elsdenii enzyme). The F clusters are thought to be of the Fe4S4+/2+ thiocubane type, and give S = $\frac{1}{2}$ EPR signals when the enzyme is in the reduced form. On the other hand, the H cluster, which is thought to be the hydrogen-activating site, gives an EPR signal only when the enzyme is in the oxidized form. The H-cluster EPR signals of all the enzymes are quite similar (g = 2.09, 2.04, 2.00), and are quite unlike the signals from other oxidized iron-sulfide clusters (such as Fe3S4 clusters and HiPIPs), in that they are observable at relatively high temperatures (> 100 K). Inhibition of the D. vulgaris and both C. pasteurianum enzymes by carbon-monoxide yields a photosensitive species that has a modified H-cluster EPR signal.195,196
The two different hydrogenases of C. pasteurianum, called hydrogenase I and II, have both been quite extensively studied, and can be regarded as prototypical iron-only hydrogenases. Hydrogenase I is active in catalyzing both H2 oxidation and H2 evolution, whereas hydrogenase II preferentially catalyzes H2 oxidation.188 The two enzymes differ in their iron contents: hydrogenase I contains about 20 iron atoms, 16 of which are thought to be involved in four F clusters,194 while the remainder presumably constitute the H cluster, which may contain six Fe atoms.194 Hydrogenase II contains about 14 iron atoms as two F clusters and one H cluster.194 These estimates of iron content result from a recent reappraisal of the metal contents (based on amino-acid analysis) that indicated a rather higher Fe content than previously realized.188,194 It is important to note that much of the spectroscopic work, which will be discussed below, was initially interpreted on the basis of the earlier, erroneous, iron analysis. Of particular interest is the possibility (first suggested197 for the D. vulgaris enzyme) that the H cluster contains six iron atoms.
Carbon-monoxide treatment of the D. vulgaris and both C. pasteurianum enzymes yields a photosensitive species that has a modified H cluster EPR signal.195,196 Interestingly, the C. pasteurianum enzymes also form complexes with O2, in a process that is distinguishable from the deactivation of hydrogenase by O2, which results from a much more prolonged exposure to O2 than that required to form the O2 complex. The O2 complexes have (photosensitive) EPR signals much like those of the CO complex.196 It is important to note that although CO, when in excess, is a potent inhibitor of the enzymes, the hydrogenase I-CO complex is actually quite active.196 With hydrogenase II, the CO complex is dissociated on exposure to H2, restoring the "active" enzyme.196 The EPR spectrum of reduced hydrogenase I is typical of (interacting) Fe4S4+ clusters, and integrates to 3 or 4 spins/protein.194,198 Electrochemical studies199 show that these clusters possess indistinguishable reduction potentials. Recently, MCD and EPR spectroscopies have been used to demonstrate the presence of significant quantities of an S = $\frac{3}{2}$ species in reduced hydrogenase I. This signal apparently integrates to about one spin per molecule, and probably originates from an S = $\frac{3}{2}$ state of an Fe4S4 cluster.198 No information is yet available on the reduction potential of the S = $\frac{3}{2}$ species. However, based on analogy with the nitrogenase iron protein,200 we might expect the S = $\frac{3}{2}$ form to have electrochemistry indistinguishable from the S = $\frac{1}{2}$ form. EPR signals with high g values (g = 6.1 and 5.0) have also been observed in C. pasteurianum hydrogenase I, and in the D. vulgaris enzyme.198,201 Since there is some uncertainty about the nature and origin of these signals,198 we will not discuss them further. The F clusters of hydrogenase II, on the other hand, give two different EPR signals that integrate to one spin each per protein molecule, and that correspond to sites with different redox potentials.194,198,199 This suggests that hydrogenase II contains two different F clusters, called F and F' (note that the presence of F' was in fact first suggested by Mössbauer spectroscopy202). The EPR spectrum from the F' cluster is unusually broad. The H-cluster EPR signals of active hydrogenase I and II are quite similar and have essentially identical redox potentials.199
The redox behavior of the F and H centers in hydrogenases I and II is nicely consistent with their respective modes of function. As shown in Figure 7.22, the F clusters are presumed to transfer electrons intermolecularly with the external electron carrier and intramolecularly with the H center. In the reversible hydrogenase I, the F and H centers have the same redox potentials (about -400 mV at pH 8), similar to that of the hydrogen electrode (-480 mV at pH 8). Thus, electrons may flow in either direction when a mediator such as methyl viologen (E°' = -440 mV at pH 8) is used as the external electron acceptor (methyl viologen is the 4,4'-bipyridinium ion). On the other hand, for hydrogenase II, the F clusters have E°' (pH 8) = -180 mV and E°' > -300 mV for F and F', respectively. In hydrogenase II, therefore, electrons can only move favorably from H2 to the H cluster, through F' and F, and then to a lower-potential acceptor, such as methylene blue [for which E°' (pH 8) = 11 mV].
Mössbauer spectroscopic studies of both hydrogenase I and II have been reported.202,203 Our discussion focuses primarily on the H cluster. The results are similar for the two enzymes, but better defined for hydrogenase II because of the smaller number of clusters. The H cluster apparently contains only two types of iron, in the ratio of 2:1, with quadrupole splittings reminiscent of Fe3S4 clusters. The oxidized cluster is confirmed to be an S = $\frac{1}{2}$ system, also reminiscent of Fe3S4 clusters, and the reduced H cluster is an S = 0 system; this contrasts with reduced Fe3S4 clusters, which have S = 2. In agreement with the Mössbauer studies, ENDOR spectroscopy of 57Fe-enriched protein indicates at least two different types of iron in the H cluster, with metal hyperfine couplings of about 18 and 7 MHz in hydrogenase II. Rather-less-intense ENDOR features were also observed at frequencies corresponding to couplings of about 11 and 15 MHz.204 The H-cluster EPR signals of hydrogenase II and hydrogenase I change on binding carbon monoxide. Although the signals of the uncomplexed enzymes are quite similar, the signals of the CO-bound species are very different (note, however, that C. pasteurianum CO-bound hydrogenase I has EPR similar to that of the CO-bound D. vulgaris enzyme), When produced with 13C-enriched CO, the EPR signal of hydrogenase II shows resolved 13C hyperfine coupling (Aav = 33 MHz) to a single 13C nucleus, indicating that only a single CO is bound, presumably as a metal carbonyl. A slightly smaller coupling of 20 MHz was obtained using ENDOR spectroscopy for the corresponding species of hydrogenase I.205
Recent ESEEM spectroscopy of hydrogenase I indicates the presence of a nearby nitrogen, which may be a nitrogen ligand to the H-cluster. This nitrogen possesses an unusually large nuclear electric quadrupole coupling and a rather novel structure, involving an amide amino-acid side chain connected to an H-cluster sulfide via a bridging proton ligand, has been suggested for it.206 Although the nitrogen in question must come from a chemically novel species, the proposed proton bridge might be expected to be exchangeable with solvent water. The ENDOR-derived result205 that there are no strongly coupled exchangeable protons in the oxidized H cluster may argue against such a structure.
Rather weak MCD198,207 and resonance Raman spectra208 have also been reported for iron hydrogenases. The lack of an intense MCD spectrum198,207 contrasts markedly with results for other biological FeS clusters. The resonance Raman spectra of hydrogenase I resemble, in some respects, spectra from Fe2S2 sites.208 These results further emphasize that the hydrogenase H clusters are a unique class of iron-sulfur clusters. Perhaps most tantalizing of all are the recent EXAFS results on hydrogenase II.209 It is important to remember that for enzymes with multiple sites, the EXAFS represents the sum of all sites present (i.e., the iron of the F, F', and H clusters). Despite this complication, useful information is often forthcoming from these experiments. The EXAFS of oxidized hydrogenase I showed both iron-sulfur and iron-iron interactions; the latter, at about 2.7 Å, is at a distance typical of Fe2S2, Fe3S4, and Fe4S4 clusters, and thus is not unexpected. The reduced enzyme, however, gave an additional, long Fe-Fe interaction at 3.3 Å. This Fe-Fe separation is not found in any of the FeS model compounds reported to date.209 The appearance of the 3.3-Å interaction indicates a change in structure on reduction of the H cluster, again revealing a cluster of unique structure and reactivity. The large structural change of this H cluster on H2 reduction is likely to have significant mechanistic implications.
4. Nickel-iron Hydrogenases
The presence of nickel in hydrogenases has only been recognized relatively recently. Purified preparations of the active enzymes were the subject of quite intensive studies for years before the Ni content was discovered by nutritional studies (see Reference 189 for a history). Some workers even tried (in vain) to purify out "impurity" EPR signals that were later found to be from the Ni. In contrast to the Fe hydrogenases discussed in the previous section, the Ni enzymes possess a variety of compositions, molecular weights, activation behavior, and redox potentials.189,210,211 As Table 7.5 shows, some of the Ni hydrogenases contain selenium, likely in the form of selenocysteine, some contain flavin (FMN or FAD), and all contain iron-sulfur centers, but in amounts ranging from 4 to 14 iron atoms per Ni atom.
Among the different Ni hydrogenases there is a common pattern of protein composition, to which many, but not all, seem to conform (especially those enzymes originating from purple eubacteria). There are two protein subunits, of approximate molecular masses 30 and 60 kDa, with the nickel probably residing in the latter subunit. The hydrogenase of the sulfate-reducing bacterium Desulfovibrio gigas is among the best investigated, and we will concentrate primarily on this enzyme. D. gigas hydrogenase contains a single Ni, two Fe4S4 clusters, and one Fe3S4 cluster. Of primary interest is the Ni site, which is thought to be the site of H2 activation.189,210,211
EPR signals attributable to mononuclear Ni [as shown by enrichment with 61Ni (I = $\frac{3}{2}$)] have been used in numerous investigations of the role of Ni in hydrogenases.189,210,211 Three major Ni EPR signals are known, which are called Ni-A, Ni-B, and Ni-C. The principal g values of these signals are: 2.32, 2.24, and 2.01 for Ni-A; 2.35, 2.16, and 2.01 for Ni-B; and 2.19, 2.15, and 2.01 for Ni-C. Of these, Ni-C is thought to be associated with the most active form of the enzyme (called active); the other two are thought to originate from less-active enzyme forms.189,210
In the enzyme as prepared (aerobically) the Ni-A EPR is characteristically observed. On hydrogen reduction the Ni-A EPR signal disappears, and the enzyme is converted into a higher-activity form (Ni-B arises from reoxidation of this form). Further progressive reduction of the enzyme gives rise to the Ni-C EPR signal, which also finally disappears. These redox properties show that Ni-C arises from an intermediate enzyme oxidation state. Although the Ni-A and Ni-B EPR signals189,210-212 almost certainly originate from low-spin Ni(III), the formal oxidation state of Ni-C is rather less certain. Both an Ni(I) site189,210 and an Ni(III) hydride213 have been suggested, with the former alternative currently favored because of the apparent absence of the strong proton hyperfine coupling expected for the latter. In the fully reduced enzyme, Ni-C is converted to an EPR-silent species. This has variously been suggested to be Ni(0), Ni(II), or an Ni(II) hydride.214 One possible reaction cycle189,210 is shown in Figure 7.23.
Information on the coordination environment of the nickel has been obtained from both x-ray absorption spectroscopy and EPR spectroscopy. The Ni K-edge EXAFS of several different hydrogenases,215-218 and EPR spectroscopy of 33S enriched Wolinella succinogenes hydrogenase,219 clearly indicate the presence of sulfur coordination to nickel. A recent x-ray absorption spectroscopic investigation of the selenium-containing D. baculatus hydrogenase, using both Ni and Se EXAFS, suggests seienocysteine coordination to Ni.216
ESEEM spectroscopy of the Ni-A and Ni-C EPR signals220,221 indicate the presence of 14N coupling, which probably arises from a histidine ligand to Ni. Interestingly, Ni-C, but not Ni-A, shows coupling to a proton that is exchangeable with solvent water. Although this coupling is too small to suggest a nickel-hydride (consistent with conclusions drawn from EPR), the proton involved could be close enough to the Ni to playa mechanistic role.
Despite the extensive studies reported to date, there are still many unanswered questions about the mechanism of the NiFe hydrogenases, which remain as exciting topics for future research. Despite our lack of detailed knowledge of enzyme mechanism, it is nevertheless not premature to seek guidance from inorganic chemistry.
5. Insights from Inorganic Chemistry
Recent years have brought insights into the way dihydrogen can be bound at a transition-metal site. Unexpectedly, it has been shown that molecular H2 forms simple complexes with many kinds of transition-metal sites.222-224 This finding contrasts with the classical situation, in which H2 interacts with a transition-metal site by oxidative addition to form a dihydride complex.23The H—H bond is largely maintained in the new/nonclassical structures. The dihydrogen and dihydride complexes can exist in simple equilibrium,224 as in Equation (7.9).
$\tag{7.9}$
The bonding of dihydrogen to a metal occurs via the $\sigma^{b}$ orbital of the H—H bond acting as a donor, with the $\sigma$* level of H2 acting as a weak acceptor. If the back donation is too strong, sufficient electron density will build up in the $sigma$* level to cause cleavage of the H—H bond, leading to the formation of a dihydride. Dihydrogen complexes therefore require a delicate balance, in which the metal coordination sphere facilitates some back-bonding, but not too much.
The proclivity of a metal center to form H2 complexes can be judged by the stretching frequency of the corresponding N2 complexes: N2 can usually displace H2 from the H2 complex to form an N2 complex without changing the remainder of the coordination sphere. If v(N$\equiv$N) is between 2060 and 2160 cm-1, H2 complexes form upon replacement of N2. If v(N$\equiv$N) is less than 2060 cm-1, indicative of electron back-donation from the metal center, a dihydride complex should form. For example, v(N$\equiv$N) = 1950 cm-1 in MoN2(PCy3)5 and MoH2(PCy3)5 is a dihydride complex, but v(N$\equiv$N) = 2090 cm-1 in Mo(Ph2PCH2CH2PPh2)CO(N2) and Mo(Ph2PCH2CH2PPh2)H2(CO)2 is a dihydrogen complex. By comparison, Mo(Et2PCH2CH2PEt2)2(CO)(N2) has v(N$\equiv$N) = 2050 cm-1 and forms a dihydride complex, Mo(Et2PCH2CH2-PEt2)2H2(CO). The correlation between v(N$\equiv$N) and the type of hydrogen complex formed seems quite useful.
Since Fe-S and Ni-S sites are implied for hydrogenase, the reactivity of transition-metal/sulfide systems with H2 may also be relevant. Interestingly, H2 can react with metal-sulfide systems at S instead of at the metal site. For example,225 (Cp')2Mo2S4 reacts with H2 to form (Cp')2Mo2(SH)4. Here the dihydrogen is cleaved without any evidence for direct interaction with the metal center, and the resulting complex contains bridging SH groups and no direct metal-H bonding.225 In recent work,226 the binuclear rhodium-sulfur complex {RhS[P(C6H5)2CH2CH2]3CH}2 was reported to react with two equivalents of dihydrogen to yield the complex {Rh(H)(SH)[P(C6H5)2CH2CH2]3CH}2, in which two SH groups bridge the two Rh centers, each of which contains a single hydrido ligand. Figure 7.24 illustrates the possibilities for hydrogen activation. Each of these types of reactivity must be considered as possibilities for the hydrogen activation process of hydrogenase.
Recently, a great deal of attention has been given to the chemistry of nickel-sulfur systems, inspired in part by the results showing that many hydrogenases are nickel-sulfur proteins.227-230 A particularly interesting finding is that Ni thiolates can react with O2 to produce sulfinate complexes.231,232 The oxygenated thiolate can be regenerated, thus providing a potential model for the O2 inactivation of Ni hydrogenases. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/07%3A_Ferrodoxins_Hydrogenases_and_Nitrogenases_-_Metal-Sulfide_Proteins/7.06%3A_Multisite_Redox_Enzymes_%28Part_2%29.txt |
Nitrogenases
Nitrogen fixation is a key reaction of the biological nitrogen cycle.233 Fixed nitrogen, in which N is in molecules other than N2, is frequently the limiting factor in plant growth.234 Since natural systems often cannot provide enough fixed nitrogen for agriculture or animal husbandry, industrial processes have been developed to "fix nitrogen" chemically. The major process in use, often referred to as ammonia synthesis, is the Haber-Bosch process, in which N2 and H2 are reacted at temperatures between 300-500°C and pressures of more than 300 atm, using catalysts (usually) based on metallic iron. 235 Hundreds of massive chemical plants are located throughout the world, some producing more than 1,000 tons of NH3/day. In contrast, in the biological process, N2 is reduced locally as needed at room temperature and ~0.8 atm by the enzyme system called nitrogenase (variously pronounced with the accent on its first or second syllable).
1. The Scope of Biological Nitrogen Fixation
Biological nitrogen fixation occurs naturally only in certain prokaryotic organisms (sometimes called diazotrophs). Although the majority of bacterial species are not nitrogen fixers, the process of nitrogen fixation has been confirmed in at least some members of many important phylogenetic groups. Nitrogen fixation occurs in strict anaerobes such as Clostridium pasteurianum, in strict aerobes such as Azotobacter vinelandii, and in facultative aerobes such as Klebsiella pneumoniae. Much of the established biochemistry of N2 reduction has been gleaned from studies of these three species. However, nitrogen fixation has a far broader range, occurring in archaebacterial methanogens such as Methanobacillus omelianskii, which produce methane, and eubacterial methanotrophs such as Methylococcus capsulatus, which oxidize methane. Photosynthetic organisms ranging from the purple bacterium Rhodobacter capsulatus (formerly Rhodopseudomonas capsulata) to the cyanobacterium (blue-green alga) Anabaena cylindrica fix nitrogen. Nitrogen fixation occurs mostly in mesophilic bacteria (existing between 15 and 40 °C), but has been found in the thermophilic archaebacterial methanogen Methanococcus thermolithotrophicus at 64 °C.
Many organisms fix N2 in nature only symbiotically. Here the most studied systems are species of Rhizobium and Bradyrhizobium, which fix nitrogen in the red root nodules of leguminous plants such as soybeans, peas, alfalfa, and peanuts. The red color inside the nodules is due to leghemoglobin, a plant O2-binding protein analogous to animal myoglobins and hemoglobins (Chapter 4). Other symbioses include that of blue-green algae such as Anabaena azollae with Azola (a water fern); actinomycetes such as Frankia with trees such as alder; and Citrobacter freundii, living in the anaerobic hind gut of termites. The distribution of nitrogenase clearly points to its adaptability as a metabolic option for species occupying widespread ecological niches.
The absence of nitrogen fixation in eukaryotes therefore seems somewhat puzzling. There would appear to be no fundamental limitation to the existence of nitrogen fixation in higher organisms. Indeed, nif genes have been transferred to yeast, where they work effectively under anaerobic conditions. Furthermore, the problem of the simultaneous presence of nitrogen fixation and aerobiosis has been solved effectively by aerobic bacteria such as Azotobacter, Gleocapsa, and Anabaena. Indeed, the lack of a fundamental limitation has encouraged researchers to propose the construction of nonsymbiotic nitrogen-fixing plants (whose niche to date is limited to the grant proposal).
Due to the mild conditions under which it occurs, the biological nitrogen-fixation process may seem inherently simpler than the industrial one. However, it is not; the biological process displays a complexity235,235a that belies the simplicity of the chemical conversion of N2 → 2NH3. Genetic analysis reveals that at least twenty genes are required for nitrogen fixation in the bacterium Klebsiella pneumoniae.236.237 These nif genes (illustrated in Figure 7.25) specify proteins that are involved in regulation (nif A and L), pyruvate oxidation/flavin reduction (nif J), electron transfer (nif F for flavodoxin), the subunits of the structural proteins of the nitrogenase (nif H, D, K), Fe-S cluster assembly (nif M) and biosynthesis of the iron-molybdenum cofactor, FeMoco (nif N, B, E, Q, V, H).238 The last two functions specify proteins that are responsible for the incorporation of unusual transition-metal sulfide clusters into the nitrogenase proteins. These clusters have allowed nitrogenase to be studied by biophysical and bioinorganic chemists to establish aspects of its structure and mechanism of action.
We will first discuss the N2 molecule and focus on its reduction products, which are the presumed intermediates or final product of nitrogen fixation. We then present what has been called239,240 the "Dominant Hypothesis" for the composition, organization, and function of molybdenum-based nitrogenases. Until 1980, it was thought that molybdenum was essential for nitrogen fixation. However, work starting in 1980 led finally in 1986 to the confirmation of vanadiumbased nitrogen fixation. The newly discovered vanadium-based nitrogenases differ in reactivity from the Mo-based enzyme in having "alternative" substrate specificity. The distinct reaction properties of the different nitrogenases point to the importance of the study of alternative substrate reactions in probing the mechanism of nitrogen fixation.
2. Dinitrogen: The Molecule and Its Reduced Intermediates
The N2 molecule has a triple bond with energy 225 kcal/mole, a v(N$\equiv$N) stretch of 2331 cm-1, and an N$\equiv$N distance of 1.098 Å. The stable isotopes of nitrogen are 14N(I = 1) with natural abundance of 99.64 percent and 15N(I = $\frac{1}{3}$) with an abundance of 0.36 percent.
The challenge to which nitrogenase rises is to break and reduce at a reasonable rate the extremely strong N$\equiv$N triple bond. The kinetic inertness of N2 is highlighted by the fact that carrying out reactions "under nitrogen" is considered equivalent to doing the chemistry in an inert atmosphere. Despite this kinetic inertness, thermodynamically the reduction of N2 by H2 is a favorable process,
$N_{2} + 3 H_{2} \rightarrow 2 NH_{3} \qquad \Delta G^{o} = -3.97 kcal/mole \tag{7.10}$
and at pH = 7 the reaction
$N_{2} + 8 H^{+} + 6 e^{-} \rightarrow 2 NH_{4}^{+} \qquad E^{o} \; ' = -280 mV \tag{7.11}$
has an E°' value that makes it easily accessible to biological reductants such as the low-potential ferredoxins discussed earlier in this chapter.
What, then, is the cause of the kinetic inertness of the N2 molecule? The thermodynamically favorable reduction of N2 to 2NH3 is a six-electron process. Unless a concerted 6e-, 6H+ process can be effected, intermediates between N2 and NH3 must be formed. However, all the intermediates on the pathway between N2 and NH3 are higher in energy than either the reactants or the products. The EO' values for the formation of N2H2 (= diimine, diazene, diamide) or N2H4 (hydrazine) are estimated241 as
$N_{2} + 2 H^{+} + 2 e^{-} \rightarrow N_{2}H_{2} \qquad E^{o} \; ' \sim -1000 to -1500 mV \tag{7.12}$
$N_{2} + 5 H^{+} + 4 e^{-} \rightarrow N_{2}H_{5}^{+} \qquad E^{o} \; ' = -695 mV \tag{7.13}$
Clearly, these potentials are sufficiently negative that the normal biological reductants cannot effect the reaction. The difficulty of reaching these intermediates is indicated in Figure 7.26.
Several factors may allow this barrier to be overcome. First, the six-electron reduction might be carried out in a concerted or near-concerted manner to avoid the intermediates completely. Alternatively, the intermediates could be compIexed at metal centers to stabilize them to a greater extent than either the reactants or products. Finally, the formation reaction for the unfavorable intermediate could be coupled with ATP hydrolysis or with the evolution of dihydrogen, each a favorable process, so that the overall process is favorable. Which of the above strategies is used by nitrogenase is unknown, but it seems likely that some combination of the last two of these is used to effect the difficult reduction of N2 to NH3. To probe the possibilities, a variety of complexes of N2, diazenes, and hydrazines has been prepared and chemically characterized, and these are discussed toward the end of this section.
3. The Dominant Hypothesis for Molybdenum Nitrogenase239,240,242,243
The action of the Mo-nitrogenase enzyme involves the functioning of two separately isolatable component proteins, as sketched in Figure 7.27A. The larger of the two proteins, sometimes incorrectly244 designated245 dinitrogenase has, in the past, been called molybdoferredoxin, azofermo, or component I. More often this protein is called the MoFe or FeMo protein ([MoFe] or [FeMo]). The smaller protein, formerly called azoferredoxin or component II, is sometimes incorrectly244 referred to245 as dinitrogenase reductase.* This protein is properly called the Fe protein or [Fe]. A useful nomenclature for discussions of kinetics and comparative biochemistry designates the FeMo protein as Xyl, where X and y are the first letters of the first and second name of the bacterial source, respectively. For example, Cpl is the FeMo protein from Clostridium pasteurianum. Similarly, for the Fe protein the designation Xy2 is given; for example, the Fe protein of Azotobacter vinelandii is called Av2. This system will be used where appropriate to distinguish the protein source. Properties of representative Mo nitrogenases are given in Table 7.6.
Table 7.6: Properties of some representative nitrogenases.
Organism Component MW Metal Content Reference
Azotobacter vinelandii
[MoFe]
[Fe]
234,000
64,000
2Mo, 34-38Fe, 26-28S
3.4Fe, 2.8S
366, 367
Azotobacter chrococcum
[MoFe]
[Fe]
227,000
65,400
2Mo, 22Fe, 20S
4Fe, 3.9S
366, 368
Clostridium pasteurianum
[MoFe]
[Fe]
221,800
55,000
2Mo, 24Fe, 24S
4Fe, 4S
366, 369
Klebsiella pneumoniae
[MoFe]
[Fe]
229,000
66,800
2Mo, 20Fe, 20S 366
Anabaena cylindrica
[MoFe]
[Fe]
223,000
60,000
2Mo, 20Fe, 20S 370
Rhodospirillum rubrum
[MoFe]
[Fe]
215,000
60,000
2Mo, 25-30Fe, 19-22S 371
The schematic diagram in Figure 7.27 shows some of the compositional and functional relationships of the nitrogenase proteins. The iron protein contains two identical subunits of MW ~ 30 kDa. The subunits are products of the nif H gene.246 A single Fe4S4 center is present in the protein and appears to be bound between the subunits.246a A recent x-ray structure246b of the iron protein confirms this picture. As shown in Figure 7.27A, the single Fe4S4 center is located at one end of the molecule, in the only region of significant contact between the two subunits. In vivo, the Fe protein is reducible by ftavodoxin or ferredoxin. In vitro, artificial reductants such as dithionite or viologens are generally used. The single Fe4S4 center undergoes a single one-electron redox process, wherein the reduced form is EPR-active and the oxidized form is diamagnetic. As such, this center resembles four-iron-cluster-containing ferredoxins. Its redox potential is dependent on the ATP or ADP level in the solution. For example, Cp2 (the Fe protein from Clostridium pasteurianum) shows E°' = -294 mV in the absence and -400 mV in the presence of MgATP.247 Two equivalents of MgATP and MgADP each bind to [Fe].
Until recently there was a major mystery over the number of Fe4S4 centers in [Fe] as deduced by EPR quantitation of the Fe4S4 centers compared to the number derived analytically or by extrusion experiments.239,248 However, it has now9,200,249,250 been clearly established that the single Fe4S4 center can exist in this protein in two spin states, S = $\frac{1}{2}$ and S = $\frac{3}{2}$ Only that part of the EPR signal corresponding to the S = $\frac{1}{2}$ form, with its g values near 2, was considered in earlier spin quantitations. When the S = $\frac{3}{2}$ center, with g values between 4 and 6, is included, the EPR spin integration shows one paramagnetic site per Fe4S4 unit. Model systems121,122 and theoretical studies123,124,251 strongly support the ability of Fe4S4 to exist in various spin states. During enzyme turnover, the single Fe4S4 of the Fe-protein center transfers electrons to the FeMo protein in one-electron steps. There is no evidence for any difference in redox behavior between the S = $\frac{1}{2}$ and S = $\frac{3}{2}$ states of the protein.200
The Fe protein binds two molecules of MgATP.252 The recent structure246b,378 suggests that a cleft between the two subunits may serve as the ATP binding site. As the enzyme system turns over, a minimum of two molecules of MgATP are hydrolyzed to MgADP and phosphate in conjunction with the transfer of each electron to the FeMo protein.253 The ATP/2e- ratio is generally accepted to have a minimum value of 4. Higher numbers represent decreased efficiency, often attributed to "futile cycling," where back electron transfer from [FeMo] to [Fe] raises the effective ratio.248,253 Except for an as-yet-unconfirmed report of reduction by thermalized electrons produced by pulse radiolysis,254 there is no evidence that the FeMo protein can be reduced to a catalytically active form without the Fe protein present.
Even though [Fe] must be present for catalysis to take place, the Dominant Hypothesis239 designates [FeMo] as the protein immediately responsible for substrate reduction, and genetic/biochemical evidence supports this view. The FeMo protein contains an ($\alpha_{2} \beta_{2}$ subunit structure, where $\alpha$ and $\beta$ are coded by the nif D and nif K genes,15,229 respectively. The overall molecular weight of about 230 kDa reflects the 50- to 60-kDa MW of each of its four subunits. In addition to protein, a total of 30 Fe, 2 Mo, and 30 S2-, all presumed to be in the form of transition-metal sulfide clusters, add relatively little to the molecular weight, but are presumed to be major parts of the active centers of the protein. Figure 7.27, which is highly schematic, displays the cluster types in accord with the Dominant Hypothesis.
* The nomenclature proposal245 that [FeMo] be designated as dinitrogenase and [Fe] as dinitrogenase reductase, although sometimes used in the literature, is incorrect or, at best, premature.244 The suggested nomenclature implies that both [FeMo] and [Fe] are enzymes. However, neither protein can function catalytically in the absence of the other. [FeMo] will not reduce N2 or C2H2 or evolve H2 in the absence of [Fe]. The iron protein will not hydrolyze MgATP in the absence of [FeMo]. Nitrogen fixation requires the simultaneous presence of both proteins. Although mechanistic considerations255 point to [FeMo] as the substrate binding and reducing protein, and [Fe] as the ATP binding locus, catalytic reactions characteristic of this enzyme system have never been consumated by one protein in the absence of the other (but see later for the uptake of H2). In this chapter, we use the [FeMo] and [Fe] designations in accord with most workers in the field.
4. Protein Purity and Active Sites
It has been almost 20 years since the first relatively pure preparations of nitrogenase became available. Indeed, homogeneous preparations are a sine qua non for progress in our understanding of the chemical nature and reactivity of any active site. In metalloproteins, there are two levels of homogeneity. The first involves purity with respect to the protein/subunit composition. This type of purity is achieved by conventional protein-purification techniques, and can be monitored by gel electrophoresis under native and denaturing conditions. In the language of polymer science, the macromolecular portion of the protein can be said to be monodisperse, corresponding to a single molecular weight for the polypeptide chain(s). However, even if the protein is homogeneous by this criterion, it may be inactive or only partially active because it does not have a full complement of active metal sites. The metal sites may be empty, filled with the wrong metals, or otherwise imperfect. Often the apo or inactive enzyme has chromatographic, electrophoretic, and centrifugal behavior very much like that of the holo protein, and therefore copurifies with it. Therefore, purification to electrophoretic homogeneity is only the first step. It is then necessary to ensure the chemical homogeneity of the active site. Very often activity is the major criterion for the approach toward such purity; i.e., the most homogeneous preparations are usually those in which the activity is highest. Several studies done on preparations that lacked active-site homogeneity were, as a result, not meaningful.
The two types of centers present in the nitrogenase FeMo protein are designated P clusters and FeMoco (or M) centers. Both types of centers display unique spectroscopic properties, but only FeMoco continues to display most of those properties when it is extracted from the protein.
5. FeMoco
The presence of the FeMo cofactor within the FeMo protein of nitrogenase, i.e., the M center, is revealed through spectroscopic and redox studies.239 In the resting state of [FeMo], as isolated in the presence of dithionite, the M center has a distinct S = $\frac{3}{2}$ EPR signal, which is discussed below (see Figure 7.28).
When the enzyme is turning over the EPR signal essentially disappears, leaving an EPR-silent state in which the FeMoco site is super-reduced to what is presumed to be its catalytically active form. In addition, a third state in which the S = $\frac{3}{2}$ EPR signal disappears is produced upon oxidation under non-turnover conditions. Thus the M center within the protein shows three states of oxidation, and these appear to have been reproduced in the FeMoco extracted from the protein:255a
$FeMoco\; (oxidized) \xrightarrow{e^{-}} FeMoco\; (reduced) \xrightarrow{e^{-}} FeMoco\; (super-reduced) \tag{7.14}$
The detailed characterization of the FeMoco site has involved parallel studies of the site within the protein and in its extracted form. The authentication of the extracted FeMoco involves the production and use of mutant organisms that make an inactive FeMo protein that contains all subunits and P clusters, but lacks the FeMoco sites.172,256 A mutant of Azotobacter vinelandii called UW-45 (UW = University of Wisconsin) was first used to assay for isolated FeMoco.257 Since several genes are involved in specifying FeMoco biosynthesis, mutants lacking these genes produce FeMo protein either lacking FeMoco or having a defective version of FeMoco. Mutants such as Nif B- of Klebsiella pneumoniae172 lack cofactor, and an inactive "apo" protein can be isolated from them.
The breakthrough in this field257 came in 1976, when FeMoco was extracted from [FeMo] into N-methylformamide258 after the protein was acidified and then neutralized, The acidification removes most of the acid-labile P clusters, and partially denatures the protein. Reneutralization precipitates the protein (near its isoelectric point) and the precipitated denatured protein can then be extracted.
It has been shown that FeMoco can be extracted into many organic solvents,10,257,259-259b provided proper combinations of cations and anions259a are present in the solvent. The role of the cation is to balance the charge of the negatively charged cofactor. The role of the anion is to displace the cofactor from anion-exchange columns, such as DEAE cellulose or TEAE cellulose, to which the cofactor and/or its protein source had been adsorbed. The ability to dissolve cofactor in such solvents as CH3CN, acetone, THF, and even benzene should facilitate attempts at further characterization and crystallization.259,259a
The biochemical authenticity of FeMoco has been assayed by its ability to activate the FeMo protein from the cofactor-less mutant organism.258 The stoichiometry of the cofactor is MoFe6-8S7-10, with the variability likely due to sample inhomogeneity. The extracted cofactor resembles the M-center unit spectroscopically and structurally as shown in Table 7.7, The differences are presumed to result from differences in the peripheral ligands of the metal-sulfide center between the protein and the organic solvent.260
Strong evidence to support FeMoco as the site of substrate binding and reduction comes from the study of nif V mutants.261-263 (The V designation is somewhat unfortunate, as nif V has nothing to do with vanadium.) The Nif V mutants do not fix nitrogen in vivo, and have altered substrate specificity in vitro. Dihydrogen evolution by isolated nif V nitrogenase is inhibited by CO, in contrast to the wild type, where H2 evolution is insensitive to CO. FeMoco can be extracted from the nif V protein and used to reactivate the FeMoco-deficient mutants, such as nif B or UW-45. Remarkably, the reconstituted FeMo protein has CO-sensitive H2 evolution, which is characteristic of nif V; i.e., the nif V phenotype is a property of FeMoco and not of the protein.263 This result clearly implicates the FeMoco site as an important part of the substrate reactions of the nitrogenase enzyme complex.
Recently, a heat-stable factor called the V-factor has been discovered that restores the wild-type phenotype when added to nif V mutants during in vitro FeMoco assembly reactions.264 The V-factor has been shown to be homocitrate (see Scheme 7.15) and 14C labeling strongly suggests that homocitrate (or a part of it) is a component of the cofactor center. Interestingly, the far more metabolically common citrate appears to be present in the nif V mutant.265a Replacement of homocitrate by analogues that differ in structure or stereochemistry yields modified FeMoco sites that have altered substrate specificities.265b Thus, as is true for many cofactors (e.g., heme = porphyrin + iron; B12 = corrin + cobalt; F430 = corphin + nickel; Moco, the molybdenum cofactor = Mo +molybdopterin), both inorganic and organic components are present in FeMoco.
$\tag{7.15}$
The biosynthesis of the cofactor and its insertion into [FeMo] apparently requires the presence of [Fe] and ATP in A. vinelandii.266,266a Whether this involves redox or conformational change in [FeMo] induced by [Fe] is unknown, but the fact that inactive versions of [Fe] are effective would seem to favor the nonredox mechanism. An attractive idea266 is that [Fe]•MgATP binds to [FeMo], producing a state that is conformationally accessible for cofactor insertion.
Recently, site-directed mutagenesis studies266b,c have shown that cysteine residues are involved in binding FeMoco to the subunits of [FeMo]. Moreover, these studies again implicate FeMoco in the substrate-reducing site.
6. The P-clusters
Evidence has been presented229 for the presence of four Fe4S4-like clusters (designated as P-clusters) in [FeMo]. The P-clusters are, however, by no means ordinary Fe4S4 clusters, and may not be Fe4S4 clusters at all. P-clusters are manifest239,248 in electronic absorption and, especially, MCD and Mössbauer spectra of [FeMo]. These spectra are clearly not conventional; i.e., they are not like those found in ferredoxins and have not yet been seen in model compounds. In their oxidized forms, the P-clusters are high-spin, probably S = t$\frac{7}{2}$ according to EPR studies.267 Mössbauer spectra reveal decidedly inequivalent Fe populations,268,269 indicating that the putative Fe4S4 clusters are highly distorted or asymmetric. The four P-clusters do not appear to behave identically under many circumstances, and it is clear that they form at least two subsets. There is open disagreement over the redox behavior of these sets.239,270,271 Furthermore, an additional Mössbauer signal sometimes designated as S may also be part of the P-cluster signal.268
Although spectroscopic studies of the P-clusters do not unequivocally reveal their structural nature, extrusion of these clusters from the protein leads to the clear identification of three or four Fe4S4 clusters.248,272 As discussed previously, the extrusion technique has inherent uncertainties, because it may be accompanied by cluster rearrangement. Nevertheless, the experimental result does support the Dominant Hypothesis, which designates the P centers as highly unusual Fe4S4 clusters.* The P-clusters are thought to be involved in electron storage and transfer, and presumably provide a reservoir of low-potential electrons to be used by the M center (FeMoco) in substrate reduction. Attractive as it may seem, there is no direct evidence to support this notion.
Table 7.7 - Comparison of the FeMo protein and isolated FeMoco.a
a) Distance in Å with number of atoms in parentheses.
b) From 287; earlier study reported in 286.
c) Data from 373; earlier study reported in 372.
d) Data from 290.
e) Data from 291.
f) Data from 288, 289.
FeMo protein (M center)
FeMoco (in NMF)
EPR
g' values 4.27 4.8
3.79 3.3
2.01 2.0
EXAFSa
Mo-S 2.36 (4)b 2.37 (3.1)c
Mo-Fe 2.69 (3)b 2.70 (2.6)c
Mo-O or N 2.18 (1)b 2.10 (3.1)c
Fe-S 2.25 (3.4)d 2.20 (3.0)e
Fe-Fe
2.66 (2.3)d 2.64 (2.2)e
2.66 (2.3)d 3.68 (0.8)e
Fe-Mo 2.76 (0.4)d 2.70 (0.8)e
Fe-O or N 1.81 (1.2)d
XANES
MoO3S3 fits bestf
* Recent x-ray crystallographic results show that, if Fe4S4 clusters are present, they are very close together in two pairs,379,380 which may account for their unusual properties. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/07%3A_Ferrodoxins_Hydrogenases_and_Nitrogenases_-_Metal-Sulfide_Proteins/7.07%3A_Multisite_Redox_Enzymes_%28Part_3%29.txt |
7. EPR, ENDOR, and ESEEM Studies
The FeMoco or M center has been identified spectroscopically within the FeMo protein;239,248,273 it has a distinctive EPR signal with effective g values of 4.3, 3.7, and 2.01, and originates from an S = $\frac{3}{2}$ state of the M center. The signal arises from transitions within the ±$\frac{1}{2}$ ground-state Kramers doublet of the S = $\frac{3}{2}$ system (D = +5.1 cm-1, E/D = 0.04). The isolated cofactor (FeMoco) gives a similar EPR signal, but with a rather larger rhombicity (E/D = 0.12). Spectra from the C. pasteurianum nitrogenase and cofactor are shown in Figure 7.28, and comparative data are given in Table 7.7. The M-center EPR signal has proved useful in characterizing the nature of the site, especially when more sophisticated magnetic resonance techniques, such as ENDOR or ESEEM, are used.
Extensive ENDOR investigations274,275,275a have been reported using protein samples enriched with the stable magnetic isotopes 2H, 33S, 57Fe, 95Mo, and 97Mo. The 57Fe couplings have been investigated in the most detail. Individual hyperfine tensors of five coupled 57Fe nuclei are discernible, and were evaluated by simulation of the polycrystalline ENDOR spectrum.275 The data from 33S and 95Mo were analyzed in less detail; 33S gave a complex ENDOR spectrum, evidently with quite large hyperfine couplings, although no quantification was attempted because of the complexity of the spectrum.274 On the other hand, 95Mo was shown to possess a small hyperfine coupling, indicating that the molybdenum possesses very little spin density (although the quantitative aspects of the conclusions of the 95Mo ENDOR study have recently been shown to be in error276).
Although no nitrogen splittings were reported in any of the ENDOR studies, evidence for involvement of nitrogen as a cluster component has been forthcoming from ESEEM spectroscopy.277-279 14N modulations are observed in the ESEEM of the M center. The observed 14N is not from the substrate (N2), or from an intermediate or product of nitrogen fixation, because enzyme turnover using 15N as a substrate does not change the ESEEM spectrum. The isolated cofactor (FeMoco) does not show the modulation frequencies observed for the M center in the protein. These experiments suggest that the M-center 14N ESEEM arises from a nitrogen atom that is associated with the M center, and probably from an amino-acid side chain (most likely a histidine) ligated to the cluster.279 Recent evidence from site-directed mutagenesis of the Azotobacter vinelandii protein280 provides strong support for the presence of histidine ligation, and points specifically to His-195 of the $\alpha$ subunit as the N ligand.
8. Mössbauer Studies
Extensive Mossbauer investigations of nitrogenase271,281-283 and FeMoco283a have been reported. Unlike EPR and EPR-based spectroscopies, which can be used to investigate only the EPR-active S = $\frac{3}{2}$ oxidation state, all three available M-center oxidation states are accessible to Mössbauer spectroscopy. The fully reduced site was found to be diamagnetic with S = 0 (but see Reference 284), whereas the oxidized site was found to have S $\geq$ 1. The zero-field spectrum of reduced C. pasteurianum nitrogenase is shown in Figure 7.29; the spectrum is comprised of four quadrupole doublets, one of which was concluded to originate from the M site.282 Mössbauer spectra taken in the presence of applied magnetic fields were used to deduce the presence of four types of 57Fe hyperfine coupling; these were called sites A1, A2, and A3, which have negative hyperfine couplings, and B sites, which have positive hyperfine couplings. The A sites were quantitated as a single Fe each; the B sites were estimated to contain three irons. These conclusions were largely confirmed and extended by later ENDOR investigations,274 although the B sites were resolved as two inequivalent, rather than three equivalent, sites. ENDOR is rather more sensitive to the nature of the hyperfine couplings than Mössbauer, although it cannot usually be used to count numbers of exactly equivalent sites. Thus the number of iron atoms in the M center is minimally five, although larger numbers cannot be excluded. Note also that some of the quantitative aspects of the earlier Mossbauer investigations have been criticized.285
9. X-ray Absorption Studies
One of the early triumphs of biological x-ray absorption spectroscopy was the deduction that the nitrogenase M center is an Mo-Fe-S cluster.286 (It is also worth noting that nitrogenase was the first enzyme to be studied by x-ray absorption spectroscopy.) Early work on lyophilized protein samples indicated the presence of two major contributions to the Mo K-edge EXAFS, which were attributed to Mo-S ligands, plus a more distant Mo-Fe contribution.286 Subsequently, these conclusions have been confirmed and extended, using samples in solution and with much more sensitive detection systems.
Most EXAFS studies to date have been on the molybdenum K-edge of the protein or of FeMoco, and indicate a very similar Mo environment in both (Table 7.7, Figure 7.30). A consensus of the best available analyses287 indicates that Mo is coordinated by three or four sulfur atoms at 2.4 Å, one to three oxygens or nitrogens at 2.2 Å, with approximately three nearby iron atoms at 2.7 Å. Of these, the EXAFS evidence for the oxygen/nitrogen contribution is weakest. However, comparison of Mo K-edge288 and Mo L-edge XANES289 spectra with model compounds indicates strong similarities with MoFe3S4 thiocubane model compounds possessing MoS3O3 coordination, and provides some support for the presence of O/N ligands.
The iron EXAFS of FeMoco has been independently examined by two groups.290,291 Both groups agree that the iron is coordinated largely to sulfur at about 2.2 Å, with more distant Fe-Fe interactions at about 2.6 Å. They differ, however, concerning the presence of short (1.8 Å) Fe-O interactions. Such interactions were apparently observed in the earlier study,290 but not in the later study.291 One possible explanation for this discrepancy is that the short Fe-O interactions of the earlier study were due to extraneous iron coordinated to solvent, contaminating the FeMoco preparation.291 A final resolution of this discord must, however, await the results of further experiments. Interestingly, a long Fe-Fe interaction at 3.7 Å was also observed in the later study.291
Largely on the basis of the Mo K-edge EXAFS results and model studies discussed below, several proposals for the structure of the M center have been put forward. These are illustrated in Figure 7.31.
The MoFe proteins from Clostridium pasteurianum292 and from Azotobacter vinelandii293 have been crystallized. For the former protein, crystals of space group P21 are obtained, with two molecules per unit cell of dimensions 70 x 151 x 122 Å. There is good evidence for a molecular two-fold axis, which presumably relates equivalent sites in the two $\alpha \beta$ dimers that make up the protein molecule.294 Preliminary refinement reveals that the two FeMoco units per protein are about 70 Å apart and the four P clusters are grouped in two pairs.
Single crystal EXAFS studies295 have provided important structural information on the molybdenum site. For different crystal orientations (relative to the polarized x-ray beam), the amplitude of the Mo-Fe EXAFS changes by a factor of 2.5, but the Mo-S EXAFS changes only slightly. Analysis of the anisotropy of the Mo-Fe EXAFS using the available crystallographic information294 is consistent with either a tetrahedral MoFe3 geometry such as that found in thiocubanes (Figure 7.32) or a square-based pyramidal MoFe4 arrangement of metals. This interpretation tends to rule out some of the structural proposals shown in Figure 7.33. The observed orientation-dependence of the iron amplitudes is too small for clusters containing a linear or planar arrangement of iron and molybdenum (e.g., Figure 7.33B,C), and too large for arrangements that involve regular disposition of iron about molybdenum. Moreover, the lack of anisotropy of the sulfur EXAFS (which was apparently not considered in the original interpretation295) argues against an MoS3 (O/N)3 model that has molybdenum coordinated by sulfur atoms that bridge only to Fe atoms disposed to one side of the molybdenum. Significant anisotropy for the Mo-S EXAFS (of opposite polarization, and smaller than that for Mo-Fe) would be expected for such an arrangement of sulfur atoms. However, the cubane model of Figure 7.33, which provides the best model of both geometric and electronic structure, remains viable if one of the nonbridging ligands to molybdenum is a sulfur atom (rather than oxygen or nitrogen) with a bond length similar to that of the bridging sulfides.
10. Substrate Reactions
The two-component Mo-nitrogenase enzyme catalyzes the reduction of N2 to 2NH4+ as its physiological reaction. Concomitant with the reduction of N2, H2 evolution occurs, with electrons supplied by the same reductants that reduce N2. The limiting stoichiometry appears to be
$N_{2} + 10 H^{+} + 8 E^{-} \rightarrow 2 NH_{4}^{+} + H_{2} \tag{7.16}$
If N2 is omitted from the assay, all the electrons go to H2 evolution. Indeed, to a first approximation the rate of electron flow through nitrogenase is independent of whether the enzyme is producing only H2, producing both NH4+ and H2, or reducing most of the alternative substrates. As displayed in Table 7.8, many alternative substrates are known for this enzyme.240,243,296 The most important of these from a practical perspective is acetylene, C2H2, which is reduced by the Mo nitrogenase exclusively to ethylene, C2H4. Acetylene can completely eliminate H2 evolution by nitrogenase. Many of the substrates in Table 7.8 have a triple bond. Indeed, the only triple-bonded molecule not reduced by nitrogenase is CO, which nevertheless inhibits all substrate reactions, but not H2 evolution (in the wild type). Triple-bonded molecules such as acetylene (H—C$\equiv$C—H) are useful probe molecules for related reactivity as discussed below for simple inorganic systems. All substrate reductions involve the transfer of two electrons or multiples thereof (i.e., 4,6,8 . . .). Multielectron substrate reductions may involve the stepwise execution by the enzyme of two-electron processes. Further, about as many protons as electrons are usually transferred to the substrate. One way of viewing the nitrogenase active site is that it can add the elementary particles (H+ and e-) of H2 to the substrate. This may have mechanistic implications.297
It is potentially fruitful to pursue the intimate connection between H2 and the N2 binding site in nitrogenase. It has been shown unequivocally298,299 that one H2 is evolved for each N2 "fixed" even at 50 atm of N2, a pressure of N2 well above full saturation. Moreover, H2 is a potent inhibitor of N2 fixation, and under D2, HD is formed, but only in the presence of N2. These complex relationships between N2 and H2(D2) have elicited a variety of interpretations.255,300-302
Recently, it has been demonstrated that the FeMo protein alone acts as an uptake hydrogenase.303 Dihydrogen in the presence of [FeMo] causes the reduction of oxidizing dyes such as methylene blue or dichlorophenolindophenol in the absence of Fe protein. This is the only known catalytic reaction displayed by the FeMo protein alone. The hydrogen evolution and uptake by [FeMo] suggest that understanding hydrogen interaction with transition-metal/sulfur centers may be crucial to understanding the mechanism of nitrogenase action.
Table 7.8 - Table 7.8 Nitrogenase substrate reactions.296,374-376
Two-electron Reductions
$2e^{-} + 2H^{+} \rightarrow H_{2}$C_{2}H_{2} + 2e^{-} + 2 H^{+} \rightarrow C_{2}H_{4}$$N_{3}^{-} + 2e^{-} + 3H^{+} \rightarrow NH_{3} + N_{2}$$N_{2}O + 2e^{-} + 2H^{+} \rightarrow H_{2}O + N_{2}$Four-electron Reductions$HCN + 4e^{-} + 4H^{+} \rightarrow CH_{3}NH_{2}RNC + 4e^{-} + 4H^{+} \rightarrow RNHCH_{3}$Six-electron Reductions$N_{2} + 6e^{-} + 6H^{+} \rightarrow 2NH_{3}HCN + 6e^{-} + 6H^{+} \rightarrow CH_{4} + NH_{3}HN_{3} + 6e^{-} + 6H^{+} \rightarrow NH_{3}+ N_{2}H_{4}RNC + 6e^{-} + 6H^{+} \rightarrow RNH_{2} + CH_{4}RCN + 6e^{-} + 6H^{+} \rightarrow RCH_{3} + NH_{3}NCNH_{2} + 6e^{-} + 6H^{+} \rightarrow CH_{3}NH_{2} + 2NH_{3}NO_{2}^{-} + 6e^{-} + 6H^{+} \rightarrow NH_{3}$Multielectron Reductions$RNC \rightarrow (C_{2}H_{6} , C_{3}H_{6} , C_{3}H_{8}) + RNH_{2}NCNH_{2} + 8e^{-} + 8H^{+} \rightarrow CH_{4} + 2NH_{3}
11. The Role of ATP
ATP hydrolysis appears to be mandatory, and occurs during electron transfer from [Fe] to [FeMo]. Dissociation of [Fe] and [FeMo] following electron transfer is probably the rate-limiting step in the overall turnover of the enzyme.255 The fact that reductant and substrate levels do not affect turnover rates is consistent with this finding.
The role of ATP on a molecular level remains one of the great mysteries of the mechanism of nitrogen fixation. As discussed above, the overall thermodynamics of N2 reduction to NH3 by H2 or by its redox surrogate flavodoxin or ferredoxin is favorable. The requirement for ATP hydrolysis must therefore arise from a kinetic necessity. This requirement is fundamentally different from the need for ATP in other biosynthetic or active transport processes, wherein the free energy of hydrolysis of ATP is needed to overcome a thermodynamic limitation.
What is the basis for the kinetic requirement of ATP hydrolysis in nitrogen fixation? To answer this question, we again look at the potential reduction products of the N2 molecule. Of these, only N2H2 (diimide, three potential isomers), N2H4 (hydrazine and its mono and dications), and NH3 (and its protonated form, NH4+) are isolable products. (In the gas phase, other species such as N2H, N2H3, or NH2 also have a "stable" existence.) In the presence of H2, only the formation of ammonia is thermodynamically favored (Figure 7.26). Clearly, the formation of the intermediate species in the free state cannot occur to any reasonable extent. However, this does not mean that nitrogenase must form NH3 directly without the formation of intermediates. It is possible for these reactive intermediates to be significantly stabilized by binding to a metal-sulfur center or centers.
Detailed kinetic studies255,304 have suggested a scheme in which intermediates with bound and probably reduced nitrogen are likely to be present. Rapid quenching experiments in acid solution lead to the detection of hydrazine during nitrogenase turnover.305 Likewise, studies of inhibition of N2 fixation by H2 and the formation of HD under D2 have been interpreted in terms of a bound diimide intermediate.306,307 Although a bound "dinitrogen hydride" is likely to be present, its detailed structure remains unknown. | textbooks/chem/Inorganic_Chemistry/Book3A_Bioinorganic_Chemistry_(Bertini_et_al.)/07%3A_Ferrodoxins_Hydrogenases_and_Nitrogenases_-_Metal-Sulfide_Proteins/7.08%3A_Multisite_Redox_Enzymes_%28Part_4%29.txt |
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