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Chemistry is unique among the physical and life sciences in one very important respect. It can be manipulated extensively to man’s design. That is, molecular structures can be designed and then constructed by choosing appropriate chemical reactions. This is chemical synthesis, which has been developed to such a degree that the economies and indeed the living standards of the industrialized nations have come to depend on it. Not everyone agrees that the present state of civilization in the industrialized nations is a way station to the millenium. But whether one agrees or not, there is no question that chemical synthesis has played an enormous role in making possible the accessories of modem life.
Chemical synthesis is not a science that can be taught or learned by any well-defined set of rules. Some classify synthesis as more art than science because, as with all really creative endeavors, to be very successful requires great imagination conditioned by a wealth of background knowledge and experience. The problems of synthesis basically are problems in design and planning. Given the objective of synthesizing a specific organic compound, there always is a variety of ways that the objective can be achieved, either from the same or from different starting materials. What we hope to do here is to show how one can go about developing efficient syntheses from available starting materials. However, practice in planning syntheses is imperative to obtaining a good grasp of the principles and problems involved. This will be up to you; no one else can do it for you. Practice also will help greatly to convert short-term memories of organic reactions to longer-term memories through repeated review and demonstrated relevance.
Methodology
In almost all syntheses the target compound is defined precisely, both as to structure and stereochemistry. Regardless of whether the synthesis is destined to be carried out on an industrial scale or on a laboratory scale, careful planning is required. The usual methodology for the planning stage involves two, not wholly independent, steps. First, one considers the various possible ways the desired carbon skeleton can be constructed, either from smaller molecules or by changes in some existing skeleton. Second, means are considered for generation of desired functional groups on the desired carbon skeleton. In many cases, the desired functional groups can be generated as a consequence of the reactions whereby the desired skeleton itself is generated. Alternative syntheses almost always are possible and one should proceed on the notion that the first sequence one thinks of is unlikely to be the best.
The choice of the best route usually is made by considering:
1. The availability of the starting materials
2. The cost of the starting materials and the equipment needed
3. The simplicity of the various steps and the scale of synthesis
4. The number of separate steps involved
5. The yield in each step
6. The ease of separation and purification of the desired product from by-products and stereoisomers
These considerations are dealt with in the following sections and in subsequent chapters.
Starting Materials
Availability of the starting materials obviously is a limiting factor in any synthetic operation. As far as laboratory-type synthesis is concerned, “availability” means that the starting materials either may be bought “off the shelf” or may be prepared easily by standard methods from other inexpensive and available compounds. For large-scale industrial syntheses, the limiting factor usually is the cost of the starting materials, including the energy required.
But in some cases the limiting factor may be problems in disposal of the byproducts. Costs will vary according to geographical location and will fluctuate widely, as with crude oil costs, so as to cause obsolescence and constant change in the chemical industry. However, it is worth remembering that the cheapest organic starting materials available are methane, ethene, ethyne, propene, butenes, benzene, and methylbenzene (toluene). Any chemical that can be prepared easily in high yield from one of these hydrocarbons is likely to be relatively inexpensive, readily available, and useful as a starting material in more involved syntheses.
The Yield Problem
Among the factors considered in choosing among several possible synthetic routes is: Which gives the best yield? The definition of yield and its distinction from another useful term, conversion, should be clearly understood. To help you understand, consider a specific example, the bromination of 2-methyl-propane to give tert-butyl bromide as the desired product. This type of reaction is carried on best with an excess of hydrocarbon to avoid polysubstitution (Section 4-5), and if we use such an excess of hydrocarbon, bromine will be the limiting reagent. This means simply that the amount of the desired product that could be formed is determined, or limited, by the amount of bromine used:
Suppose we start with one mole of hydrocarbon and 0.2 mole of bromine and, after a specified reaction time, 0.1 mole of bromine has reacted. If only the desired product were formed, and there were no other losses of hydrocarbon or bromine,
$\% \: \text{conversion} = \frac{\text{moles of limiting reagent reacted}}{\text{moles of limiting reagent initially present}} = \frac{0.1}{0.2} \times 100 = 50\%$
If there are no losses in isolating the product or in recovering unused starting material, then
$\% \: \text{yield} = \frac{\text{moles of product}}{\text{moles of limiting reagent initially present}} = \frac{0.1}{0.1} \times 100 = 100\%$
Now suppose all of the 0.2 mole of bromine reacts, 0.08 mole of the desired product can be isolated, and 0.7 mole of hydrocarbon is recovered. Under these circumstances, the percent conversion is $100\%$, because all of the bromine has reacted. The yield can be figured in different ways depending on which starting material one wishes to base the yield. Based on bromine (which would be logical because bromine is the more expensive reagent) the yield of tert-butyl bromide is $\left( 0.08/0.2 \right) \times 100 = 40\%$. However, one also could base the yield of tert-butyl bromide on the unrecovered hydrocarbon, and this would be $\left[ 0.08/\left( 1.0-0.7 \right) \right] \times 100 = 27\%$.
In a multistep synthesis, the overall percent yield is the product of the fractional yields in each step times 100 and decreases rapidly with the number of steps. For this reason, a low-yield step along the way can mean practical failure for the overall sequence. Usually, the best sequence will be the one with the fewest steps. Exceptions arise when the desired product is obtained as a component of a mixture that is difficult to separate. For example, one could prepare 2-chloro-2-methylbutane in one step by direct chlorination of 2-methyl-butane (Section 4-5A). But because the desired product is very difficult to separate from the other, isomeric monochlorinated products, it is desirable to use a longer sequence that may give a lower yield but avoids the separation problem. Similar separation problems would be encountered in a synthesis that gives a mixture of stereoisomers when only one isomer is desired. Again, the optimal synthesis may involve a longer sequence that would be stereospecific for the desired isomer.
One way of maximizing the yield is to minimize the number of sequential steps and, whenever possible, to use parallel rather than sequential reactions. For example, suppose that we wish to synthesize a compound $\ce{ABCDEF}$ by linking together $\ce{A}$, $\ce{B}$, $\ce{C}$, $\ce{D}$, $\ce{E}$, and $\ce{F}$. The sequential approach would involve at least five steps as follows:
$\ce{A} \overset{\ce{B}}{\rightarrow} \ce{AB} \overset{\ce{C}}{\rightarrow} \ce{ABC} \overset{\ce{D}}{\rightarrow} \ce{ABCD} \overset{\ce{E}}{\rightarrow} \ce{ABCDE} \overset{\ce{F}}{\rightarrow} \ce{ABCDEF}$
If each of these steps proceeds in $90\%$ yield, the overall yield would be $\left( 0.90 \right)^5 \times 100 = 59\%$.
One possible parallel approach would involve synthesis of the fragments $\ce{ABC}$ and $\ce{DEF}$ followed by the combination of these to $\ce{ABCDEF}$:
There are still at least five reaction steps, but only three sequential steps; and if each of these proceeds in $90\%$ yield, the overall yield would be $\left( 0.90 \right)^3 \times 100 = 73\%$. The parallel approach is especially important in the synthesis of polymeric substances such as peptides, proteins, and nucleic acids in which many subunits have to be linked.
Finally, product yields are very dependent on manipulative losses incurred in each step by isolating and purifying the synthetic intermediates. The need to minimize losses of this kind is critically important in very lengthy syntheses.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/13%3A_Polyfunctional_Compounds_Alkadienes_and_Approaches_to_Organic_Synthesis/13.06%3A_Approaches_to_Planning_Practical_Organic_Syntheses.txt |
According to the suggested approach to planning a synthesis, the primary consideration is how to construct the target carbon skeleton starting with smaller molecules (or, alternatively, to reconstruct an existing skeleton). Construction of a skeleton from smaller molecules almost always will involve formation of carbon-carbon bonds. Up to this point we have discussed only a few reactions in which carbon-carbon bonds are formed and these are summarized in Table 13-4. Other important reactions that can be used to enlarge a carbon framework will be discussed in later chapters.
Table 13-4: Some Carbon-Carbon Bond-Forming Reactions with Section Reference
The most logical approach to planning the synthesis of a particular carbon framework requires that one work backward by mentally fragmenting the molecule into smaller pieces that can be “rejoined” by known $\ce{C-C}$ bond-forming reactions. The first set of pieces in turn is broken into smaller pieces, and the mental fragmentation procedure is repeated until the pieces correspond to the carbon skeletons of readily available compounds. There almost always will be several different backward routes, and each is examined for its potential to put the desired functional groups at their proper locations. In almost all cases it is important to use reactions that will lead to pure compounds without having to separate substances with similar physical properties.
Example
A typical synthesis problem would be to devise a preparation of cis-2-octene, given the restrictions that the starting materials have fewer than eight carbons, and that we use the $\ce{C-C}$ bond-forming reactions we have discussed up to now. The reasoning involved in devising an appropriate synthesis with the given restrictions will be outlined for this example in detail.
First, we can see that the carbon skeleton of the desired product can be divided to give the following combinations of fragments:
Next, we have to decide what reaction or reactions would be useful to put these fragments together to reform the $\ce{C_8}$ chain. If we look at the list of available reactions in Table 13-4$^4$ for $\ce{C-C}$ bond formation, we can rule out 1, 2, 4, 8, and 9: 1 and 4 because the reactions are unsuitable for making unbranched chains; 8 and 9 because they make rings, not chains; and 2 because it does not work well in the absence of activating groups. Reaction 3 could be used to combine $\ce{C_1}$ and $\ce{C_7}$ units to give $\ce{C_8}$, as by the radical addition of $\ce{CBrCl_3}$ to 1-heptene:
Reactions 6 and 7 could also be used to make $\ce{C_8}$, 6 by linking $\ce{C_7}$ to $\ce{C_1}$ and 7 by putting together two $\ce{C_4}$ units.$^5$
Reaction 5 could be useful for all of the possible ways of dividing $\ce{C_8}$. Some of the possible combinations are:
This does not exhaust the possibilities because, as 5b and 5e show, Reaction 5 can be used to make the same $\ce{C_8}$ compound from different sets of starting materials.
We now have to consider how to convert the $\ce{C_8}$ materials that we might make into cis-2-octene. The possibilities are:
Of these, d is the obvious choice for first consideration because it has its functionality, a single triple bond, between the same two carbons we wish to have joined by a cis double bond in the product. Now, we have to ask if there are reactions that will convert $\ce{-C \equiv C}-$ to cis-. Two possibilities were mentioned previously - hydrogenation of a triple bond with the Lindlar catalyst (Section 11-2B) and hydroboration followed by treatment with propanoic acid (Section 11-6D):
Either of these two reactions provides a simple and straightforward way of converting 2-octyne to cis-2-octene, so a satisfactory answer to the original problem is
You can see that even with having available only seven $\ce{C-C}$ bond-forming reactions and two ways of converting $\ce{-C \equiv C}- \rightarrow$ , considerable amount of logical screening is required to eliminate unsuitable possibilities. The skilled practitioner makes this kind of diagnosis quickly in his head; at the outset you will find it useful to write the steps in your screening in the same way as we have done for this example.
$^4$You may wish to review the sections cited for each reaction to be sure you understand the judgments we make here as to the suitability of particular reactions for the purpose at hand.
$^5$Reaction 7 could also be used to make $\ce{C_8}$ by other combinations, such as of $\ce{C_5}$ and $\ce{C_3}$, but these would give undesirable mixtures of products. Thus, $\ce{C-C-C-C \equiv CH} + \ce{HC \equiv C-C} \rightarrow \ce{C-C-C-C \equiv C-C \equiv C-C} + \ce{C-C \equiv.C-C \equiv C-C} + \ce{C-C-C-C \equiv C-C \equiv C-C-C-C}$
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/13%3A_Polyfunctional_Compounds_Alkadienes_and_Approaches_to_Organic_Synthesis/13.07%3A_Building_the_Carbon_Skeleton.txt |
As we have seen in the previous section, it may be easy to construct the carbon skeleton of the target compound of a synthesis, but with a reactive functional group at the wrong carbon. Therefore it is important also to have practice at shifting reactive entry points around to achieve the final desired product. We shall illustrate this form of molecular chess with reactions from previous post. A typical problem may be to devise syntheses for achieving the following conversions:
If you proceed as in the previous section, you will see that the starting material and products have the same number of carbons and the same general bonding arrangement of those carbons. Getting the functionality to the right carbons is now the problem. If you review the reactions discussed up to now, you will find that the only good way of getting a reactive group at the end of a chain starting with a reactive group in the middle of the chain is borane isomerization (Section 11-6C). The borane, \(6\), can be obtained from the starting material, \(5\), by hydroboration (Section 11-6) and, on heating, \(6\) will be converted to \(7\):
Production of 4-methyl-1-pentanol from the substituted alkylborane, \(7\), can be achieved by oxidation (Section 11-6D):
The three steps - hydroboration, isomerization, and oxidation - thus constitute a reasonable synthesis of the first desired compound.
The second desired product is a little more tricky because the isomerization of \(6\) to \(7\) cannot be stopped at the alkylborane, \(8\):
The best procedure to get the desired product is to generate the 1-alkene from the borane with 1-decene (Section 11-6C) and then add hydrogen bromide by a polar mechanism (Section 10-4). Incursion of radical-chain addition must be avoided because it would give 1-bromo-4-methylpentane (Section 10-7):
A very brief summary of the transformations that we have studied so far, which do not change the carbon skeleton, is given in Table 13-5 along with appropriate section references. In using this table, it is necessary to check the specific sections to be sure the reaction is applicable to the conversion that you wish to achieve and to determine the proper conditions for the reaction.
Table 13-5: Summary of Useful Synthetic Transformations
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
13.09: Construction of Ring Systems by Cycloaddition
Another example of a synthesis problem makes use of the cycloaddition reactions discussed here. Consider the synthesis of bicyclo[2.2.1]heptane, $9$, from compounds with fewer carbons.
Whenever a ring has to be constructed, you should consider the possibility of cycloaddition reactions, especially [4 + 2] cycloaddition by the Diels-Alder reaction. A first glance at $9$, written in the usual sawhorse-perspective formula, might lead to overlooking the possibility of constructing the skeleton by [4 + 2] addition, because the compound seems only to be made up of five-membered rings. If the structure is rewritten as $10$, the six-membered ring stands out much more clearly:
If we now try to divide the six-membered ring into [2] and [4] fragments, we find that there are only two different ways this can be done:
The left division corresponds to a simple [4 + 2] cycloaddition, whereas the right division corresponds to a complex reaction involving formation of three ring bonds at once. Actual Diels-Alder reactions require diene and dienophile starting materials, and two possibilities, using 1,3-cyclopentadiene as the diene and ethene or ethyne as dienophile, follow:
Either of the products can be converted to bicyclo[2.2.1]heptane by hydrogenation (Table 13-5):
Neither ethene nor ethyne is a very good dienophile but [4 + 2] cycloadditions of either with 1,3-cyclopentadiene go well at temperatures of $160$-$180^\text{o}$ because 1,3-cyclopentadiene is a very reactive diene. Achieving the overall result of addition of ethene or ethyne to a less reactive diene could necessitate a synthetic sequence wherein one of the reactive dienophiles listed in Table 13-1 is used to introduce the desired two carbons, and the activating groups are subsequently removed. An example follows:
Reactions that can be used to remove a $\ce{-CO_2H}$ group will be discussed in Chapter 18.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/13%3A_Polyfunctional_Compounds_Alkadienes_and_Approaches_to_Organic_Synthesis/13.08%3A_Introducing_Functionality.txt |
One of the major problems in organic synthesis is the suppression of unwanted side reactions. Frequently the desired reaction is accompanied by reaction at other parts of the molecule, especially when more than one functional group is present. Functional groups usually are the most reactive sites in the molecule, and it may be difficult or even impossible to insulate one functional group from a reaction occurring at another. Therefore any proposed synthesis must be evaluated at each step for possible side reactions that may degrade or otherwise modify the structure in an undesired way. To do this will require an understanding of how variations in structure affect chemical reactivity. Such understanding is acquired through experience and knowledge of reaction mechanism and reaction stereochemistry.
To illustrate the purpose and practice of protecting groups in organic synthesis, let us suppose that the synthesis of cis-2-octene, which we outlined in Section 13-7, has to be adapted for the synthesis of 5-octyn-1-ol. We could write the following:
However, the synthesis as written would fail because the alkyne is a weaker acid than the alcohol (Section 11-8), and the alkynide anion would react much more rapidly with the acidic proton of the alcohol than it would displace bromide ion from carbon:
The hydroxyl group of 4-bromo-1-butanol therefore must be protected before it is allowed to react with the alkynide salt. There are a number of ways to protect hydroxyl groups, but one method, which is simple and effective, relies on the fact that unsaturated ethers of the type are very reactive in electrophilic addition reactions (Section 10-4). An alcohol readily adds to the double bond of such an ether in the presence of an acid catalyst:
The protected compound is a much weaker acid than the alkyne, and the displacement reaction can be carried out with the alkynide salt without difficulty. To obtain the final product, the protecting group must be removed, and this can be done in dilute aqueous acid solution by an $S_\text{N}1$ type of substitution (Sections 8-7D and 8-7E):
Some Common Protecting Groups in Organic Synthesis
Hydroxyl $\left( \ce{OH} \right)$ protecting groups in Organic Synthesis
Protection of alcohols:
Acetyl $\left( \ce{Ac} \right)$ – Removed by acid or base.
Benzoyl $\left( \ce{Bz} \right)$ – Removed by acid or base, more stable than $\ce{Ac}$ group.
Benzyl ($\ce{Bn}$, $\ce{Bnl}$) – Removed by hydrogenolysis. $\ce{Bn}$ group is widely used in sugar and nucleoside chemistry.
$\beta$-Methoxyethoxymethyl ether (MEM) – Removed by acid.
Dimethoxytrityl, [bis-(4-methoxyphenyl)phenylmethyl] (DMT) – Removed by weak acid. DMT group is widely used for protection of 5′-hydroxy group in nucleosides, particularly in oligonucleotide synthesis.
Methoxymethyl ether (MOM) – Removed by acid.
Methoxytrityl [(4-methoxyphenyl)diphenylmethyl, MMT) – Removed by acid and hydrogenolysis.
p-Methoxybenzyl ether (PMB) – Removed by acid, hydrogenolysis, or oxidation.
Methylthiomethyl ether – Removed by acid.
Pivaloyl $\left( \ce{Piv} \right)$ – Removed by acid, base or reductant agents. It is substantially more stable than other acyl protecting groups.
Tetrahydropyranyl (THP) – Removed by acid.
Trityl (triphenylmethyl, $\ce{Tr}$) – Removed by acid and hydrogenolysis.
Silyl ether (most popular ones include trimethylsilyl (TMS), tert-butyldimethylsilyl (TBDMS), tri-iso-propylsilyloxymethyl (TOM), and triisopropylsilyl (TIPS) ethers) – Removed by acid or fluoride ion. (such as $\ce{NaF}$, TBAF (Tetra-n-butylammonium fluoride, HF-Py, or HF-NEt3)). TBDMS and TOM groups are used for protection of 2′-hydroxy function in nucleosides, particularly in oligonucleotide synthesis.
Methyl Ethers – Cleavage is by TMSI in DCM or MeCN or Chloroform. An alternative method to cleave methyl ethers is BBr3 in DCM
Ethoxyethyl ethers (EE) – Cleavage more trivial than simple ethers e.g. 1N Hydrochloric acid
Amine protecting groups in Organic Synthesis
Protection of amines:
Carbobenzyloxy (Cbz) group – Removed by hydrogenolysis
p-Methoxybenzyl carbonyl (Moz or MeOZ) group – Removed by hydrogenolysis, more labile than Cbz
tert-Butyloxycarbonyl (BOC) group (Common in solid phase peptide synthesis) – Removed by concentrated, strong acid. (such as HCl or CF3COOH)
9-Fluorenylmethyloxycarbonyl (FMOC) group (Common in solid phase peptide synthesis) – Removed by base, such as piperidine
Acetyl (Ac) group is common in oligonucleotide synthesis for protection of N4 in cytosine and N6 in adenine nucleic bases and is removed by treatment with a base, most often, with aqueous or gaseous ammonia or methylamine. Ac is too stable to be readily removed from aliphatic amides.
Benzoyl (Bz) group is common in oligonucleotide synthesis for protection of N4 in cytosine and N6 in adenine nucleic bases and is removed by treatment with a base, most often with aqueous or gaseous ammonia or methylamine. Bz is too stable to be readily removed from aliphatic amides.
Benzyl (Bn) group – Removed by hydrogenolysis
Carbamate group – Removed by acid and mild heating.
p-Methoxybenzyl (PMB) – Removed by hydrogenolysis, more labile than Benzyl
3,4-Dimethoxybenzyl (DMPM) – Removed by hydrogenolysis, more labile than p-methoxybenzyl
p-methoxyphenyl (PMP) group – Removed by Ammonium cerium(IV) nitrate (CAN)
Tosyl (Ts) group – Removed by concentrated acid (HBr, H2SO4) & strong reducing agents (sodium in liquid ammonia or sodium naphthalenide)
Other Sulfonamides (Nosyl & Nps) groups – Removed by samarium iodide, tributyltin hydride
Carbonyl protecting groups in Organic Synthesis
Protection of carbonyl groups:
Acetals and Ketals – Removed by acid. Normally, the cleavage of acyclic acetals is easier than of cyclic acetals.
Acylals – Removed by Lewis acids.
Dithianes – Removed by metal salts or oxidizing agents.
Carboxylic acid protecting groups in Organic Synthesis
Protection of carboxylic acids:
Methyl esters – Removed by acid or base.
Benzyl esters – Removed by hydrogenolysis.
tert-Butyl esters – Removed by acid, base and some reductants.
Silyl esters – Removed by acid, base and organometallic reagents.
Orthoesters – Removed by mild aqueous acid to form ester, which is removed according to ester properties.
Oxazoline – Removed by strong hot acid (pH < 1, T > 100 °C) or alkali (pH > 12, T > 100 °C), but not e.g. LiAlH4, organolithium reagents or Grignard (organomagnesium) reagents
Phosphate protecting groups in Organic Synthesis
2-cyanoethyl – removed by mild base. The group is widely used in oligonucleotide synthesis.
Methyl (Me) – removed by strong nucleophiles e.c. thiophenole/TEA.
Terminal alkyne protecting groups in Organic Synthesis
propargyl alcohols in the Favorskii reaction,
silyl groups, especially in protection of the acetylene itself
Orthogonal protection in Organic Synthesis
Orthogonal protection is a strategy allowing the deprotection of multiple protective groups one at a time each with a dedicated set of reaction conditions without affecting the other. It was introduced in the field of peptide synthesis by Robert Bruce Merrifield in 1977. As a proof of concept orthogonal deprotection is demonstrated in a photochemical transesterification by trimethylsilyldiazomethane utilizing the kinetic isotope effect:
Due to this effect the quantum yield for deprotection of the right-side ester group is reduced and it stays intact. Significantly by placing the deuterium atoms next to the left-side ester group or by changing the wavelength to 254 nm the other monoarene is obtained.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/13%3A_Polyfunctional_Compounds_Alkadienes_and_Approaches_to_Organic_Synthesis/13.10%3A_Protecting_Groups_in_Organic_Synthesis.txt |
According to the suggested approach to planning a synthesis, the primary consideration is how to construct the target carbon skeleton starting with smaller molecules (or, alternatively, to reconstruct an existing skeleton). Construction of a skeleton from smaller molecules almost always will involve formation of carbon-carbon bonds. Up to this point we have discussed only a few reactions in which carbon-carbon bonds are formed and these are summarized in Table 13-4. Other important reactions that can be used to enlarge a carbon framework will be discussed in later chapters.
Table 13-4: Some Carbon-Carbon Bond-Forming Reactions with Section Reference
The most logical approach to planning the synthesis of a particular carbon framework requires that one work backward by mentally fragmenting the molecule into smaller pieces that can be “rejoined” by known $\ce{C-C}$ bond-forming reactions. The first set of pieces in turn is broken into smaller pieces, and the mental fragmentation procedure is repeated until the pieces correspond to the carbon skeletons of readily available compounds. There almost always will be several different backward routes, and each is examined for its potential to put the desired functional groups at their proper locations. In almost all cases it is important to use reactions that will lead to pure compounds without having to separate substances with similar physical properties.
Example
A typical synthesis problem would be to devise a preparation of cis-2-octene, given the restrictions that the starting materials have fewer than eight carbons, and that we use the $\ce{C-C}$ bond-forming reactions we have discussed up to now. The reasoning involved in devising an appropriate synthesis with the given restrictions will be outlined for this example in detail.
First, we can see that the carbon skeleton of the desired product can be divided to give the following combinations of fragments:
Next, we have to decide what reaction or reactions would be useful to put these fragments together to reform the $\ce{C_8}$ chain. If we look at the list of available reactions in Table 13-4$^4$ for $\ce{C-C}$ bond formation, we can rule out 1, 2, 4, 8, and 9: 1 and 4 because the reactions are unsuitable for making unbranched chains; 8 and 9 because they make rings, not chains; and 2 because it does not work well in the absence of activating groups. Reaction 3 could be used to combine $\ce{C_1}$ and $\ce{C_7}$ units to give $\ce{C_8}$, as by the radical addition of $\ce{CBrCl_3}$ to 1-heptene:
Reactions 6 and 7 could also be used to make $\ce{C_8}$, 6 by linking $\ce{C_7}$ to $\ce{C_1}$ and 7 by putting together two $\ce{C_4}$ units.$^5$
Reaction 5 could be useful for all of the possible ways of dividing $\ce{C_8}$. Some of the possible combinations are:
This does not exhaust the possibilities because, as 5b and 5e show, Reaction 5 can be used to make the same $\ce{C_8}$ compound from different sets of starting materials.
We now have to consider how to convert the $\ce{C_8}$ materials that we might make into cis-2-octene. The possibilities are:
Of these, d is the obvious choice for first consideration because it has its functionality, a single triple bond, between the same two carbons we wish to have joined by a cis double bond in the product. Now, we have to ask if there are reactions that will convert $\ce{-C \equiv C}-$ to cis-. Two possibilities were mentioned previously - hydrogenation of a triple bond with the Lindlar catalyst (Section 11-2B) and hydroboration followed by treatment with propanoic acid (Section 11-6D):
Either of these two reactions provides a simple and straightforward way of converting 2-octyne to cis-2-octene, so a satisfactory answer to the original problem is
You can see that even with having available only seven $\ce{C-C}$ bond-forming reactions and two ways of converting $\ce{-C \equiv C}- \rightarrow$ , considerable amount of logical screening is required to eliminate unsuitable possibilities. The skilled practitioner makes this kind of diagnosis quickly in his head; at the outset you will find it useful to write the steps in your screening in the same way as we have done for this example.
$^4$You may wish to review the sections cited for each reaction to be sure you understand the judgments we make here as to the suitability of particular reactions for the purpose at hand.
$^5$Reaction 7 could also be used to make $\ce{C_8}$ by other combinations, such as of $\ce{C_5}$ and $\ce{C_3}$, but these would give undesirable mixtures of products. Thus, $\ce{C-C-C-C \equiv CH} + \ce{HC \equiv C-C} \rightarrow \ce{C-C-C-C \equiv C-C \equiv C-C} + \ce{C-C \equiv.C-C \equiv C-C} + \ce{C-C-C-C \equiv C-C \equiv C-C-C-C}$
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/13%3A_Polyfunctional_Compounds_Alkadienes_and_Approaches_to_Organic_Synthesis/13.11%3A_Building_the_Carbon_Skeleton.txt |
The general term of "organohalogen" refers to compounds with covalent carbon-halogen bonds. Substances such a bromomethane (\(\ce{CH_3Br}\)) and chloroethene (\(\ce{CH_2=CHCl}\)), are examples of organohalogen compounds, whereas others such as the methylammonium chloride salt, which have no carbon-halogen bonds, are not. This chapter is only concerned with compounds that have covalent carbon-halogen bonds.
• 14.1: Prelude to Organohalogen and Organometallic Compounds
There is wide diversity in the nature of organohalogen compounds but we have restricted this chapter to alkyl, cycloalkyl, alkenyl, alkynyl, and aryl halides. Some of the chemistry of the carbon-halogen bonds already will be familiar to you because it involves the addition, substitution, and elimination reactions discussed in previous chapters. We will amplify these reactions and consider nucleophilic substitution by what are called the addition-elimination and elimination-addition mechanisms.
• 14.2: Physical Properties of Organohalogen and Organometallic Compounds
The physical properties of haloalkanes are much as one might expect. Volatility decreases: (a) with increasing molecular weight along a homologous series, (b) with increasing atomic number of the halogen, and (c) with the structure of the alkyl group in the order such that tertiary << secondary << primary for isomeric halides.
• 14.3: Spectroscopic Properties
Organohalogen compounds give rise to strong absorptions in the infrared arising from stretching vibrations of the carbon-halogen bond. The frequency of absorption decreases as the mass of the halogen increases.
• 14.4: Alkyl Halides
The important chemistry of alkyl halides includes the nucleophilic (SN) displacement and elimination (E) reactions. Recall that tertiary alkyl halides normally are reactive in ionization SN1 reactions, whereas primary halides, and to a lesser extent secondary halides, are reactive in SN2 reactions, which occur by a concerted mechanism with inversion of configuration.
• 14.5: Alkenyl and Alkynyl Halides
The most readily available alkenyl halide is chloroethene (vinyl chloride), which can be prepared by a number of routes.
• 14.6: Cycloalkyl Halides
The cycloalkyl halides, except for cyclopropyl halides, have physical and chemical properties that are similar to those of the open-chain secondary halides and can be prepared by the same types of reactions. All the cycloalkyl halides undergo SN2 reactions rather slowly and, with nucleophiles that are reasonably basic, E2 reactions can be expected to predominate.
• 14.7: Aryl Halides
Aryl halides have a halogen directly bonded to a carbon of an aromatic ring. The simple aryl halides generally are resistant to attack by nucleophiles. However, considerable activation is produced by strongly electron-attracting substituents provided these are located in either the ortho or para positions, or both.
• 14.8: Polyhalogenated Alkanes and Alkenes
Polychlorination of methane yields the di-, tri-, and tetrachloromethanes cheaply and efficiently. These substances have excellent solvent properties for nonpolar and slightly polar substances. Chloroform once was used widely as an inhalation anesthetic. However, it has a deleterious effect on the heart and is oxidized slowly by atmospheric oxygen to highly toxic carbonyl dichloride.
• 14.9: Organometallic Compounds from Organohalogen Compounds
One of the more important reactions of organohalogen compounds is the formation of organometallic compounds by replacement of the halogen by a metal atom. Carbon is positive in carbon-halogen bonds and becomes negative in carbon-metal bonds, and therefore carbon is considered to be reduced in formation of an organometallic compound.
• 14.10: Properties of Organometallic Compounds
How carbon-metal bonds are formed depends on the metal that is used. Conditions that are suitable for one metal may be wholly unsuited for another. Some organometallic compounds react very sluggishly even toward acids, whereas others react avidly with water, oxygen, carbon dioxide, and almost all solvents but the alkanes themselves. Reactivity increases with increasing polarity of the carbon-metal bond, which is determined by the electropositivity of the metal.
• 14.11: Preparation of Organometallic Compounds
The reaction of a metal with an organic halide is a convenient method for preparation of organometallic compounds of reasonably active metals such as lithium, magnesium, and zinc. Ethers, particularly diethyl ether and oxacyclopentane (tetrahydrofuran), provide inert, slightly polar media in which organometallic compounds usually are soluble. Care is necessary to exclude moisture, oxygen, and carbon dioxide, which would react with the organometallic compound.
• 14.12: Organomagnesium Compounds
For many years the most important organometallic compounds for synthetic purposes have been the organomagnesium halides, or Grignard reagents. They are named after Victor Grignard, who discovered them and developed their use as synthetic reagents, for which he received a Nobel Prize in 1912. As already mentioned, these substances customarily are prepared in dry ether solution from magnesium turnings and an organic halide.
• 14.13: Organomagnesium and Organolithium Compounds in Synthesis
The most important synthetic use of Grignard reagents and organolithium reagents is to form new carbon-carbon bonds by addition to polar multiple bonds, particularly carbonyl bonds. An example is the addition of methyl-magnesium iodide to methanal. With suitable variations of the carbonyl compound, a wide range of compounds can be built up from substances containing fewer carbon atoms per molecule.
• 14.E: Organohalogen and Organometallic Compounds (Exercises)
These are the homework exercises to accompany Chapter 14 of the Textmap for Basic Principles of Organic Chemistry (Roberts and Caserio).
Thumbnail: Structure of 1-Chlorobenzene.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
14: Organohalogen and Organometallic Compounds
There is wide diversity in the nature of organohalogen compounds but, of necessity, we have restricted this chapter to alkyl, cycloalkyl, alkenyl, alkynyl, and aryl halides. Some of the chemistry of the carbon-halogen bonds already will be familiar to you because it involves the addition, substitution, and elimination reactions discussed in previous chapters. To some extent, we will amplify these reactions and consider nucleophilic substitution by what are called the addition-elimination and elimination-addition mechanisms. Subsequently, we will discuss the formation of carbon-metal bonds from carbon-halogen bonds. The latter type of reaction is of special value because compounds that have carbon-metal bonds are potent reagents for the formation of carbon-carbon bonds, as we will show later in this chapter.
Although large numbers of organohalogens are known, very few of them occur naturally. Thyroid hormones (e.g., thyroxine) that contain iodine are exceptions; other organohalogens are found as mold metabolites (such as griseofulvin) and in marine organisms:
Almost all of the organohalogen compounds in use today are synthetic in origin. You may wonder why, if nature doesn’t choose to make them, man elects to do so. The main interest to us here is that they are very useful intermediates for the synthesis of a wide range of other compounds. However, vast quantities of synthetic halogen compounds, particularly polyhalogen compounds, are used as pesticides, cleaning solvents, anaesthetics, aerosol propellants, refrigerants, polymers, and so on. The wisdom of this massive use of materials that are foreign to our natural environment gradually is being reevaluated as the long-term detrimental effects of many of these chemicals become known. For example, many of the chlorinated hydrocarbons such as DDT, Chlordane, and Lindane, which have been used very widely as insecticides, now are at least partially banned because of concern for their long-term effects on nontarget species, including man.
Sometimes the long-term effects are quite unexpected and difficult to predict. For example, millions of kilograms of \(\ce{CF_2Cl_2}\), which is used as a propellant, have been released into the atmosphere from aerosol cans. This compound appears to be wholly free of direct adverse physiological effects. However, as the substance diffuses into the upper atmosphere, it is slowly decomposed by sunlight to produce chlorine atoms. Serious danger then is possible because chlorine atoms are known to catalyze the decomposition of ozone, and it is the ozone layer in the upper atmosphere that absorbs most of the sun’s ultraviolet radiation that is strongly harmful to life.
Contributors
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/14%3A_Organohalogen_and_Organometallic_Compounds/14.01%3A_Prelude_to_Organohalogen_and_Organometallic_Compounds.txt |
The physical properties of haloalkanes are much as one might expect. Volatility decreases: (a) with increasing molecular weight along a homologous series, (b) with increasing atomic number of the halogen, and (c) with the structure of the alkyl group in the order such that tertiary $<$ secondary $<$ primary for isomeric halides. These trends are apparent from the physical properties listed in Table 14-1, which includes data for simple halogen derivatives of alkanes, alkenes, alkynes, and arenes.
Table 14-1: Physical Properties of Organic Halides
The boiling points of many halogen compounds are similar to hydrocarbons of the same molecular weight, but there are some conspicuous exceptions. lodomethane, for example, has about the same molecular weight as decane (MW 142), but the boiling point of iodomethane is $132^\text{o}$ lower than that of decane. Likewise, fluorocarbons (e.g., tetrafluoromethane, $\ce{CF_4}$, MW 88, bp $-129^\text{o}$) are far more volatile than hydrocarbons of similar weights (e.g., hexane, $\ce{C_6H_{14}}$, MW 86, bp $69^\text{o}$).
In general, halogen compounds are insoluble in water but are readily soluble in organic solvents and, with the exception of some fluoro and monochloro compounds, they are more dense than water. Aryl halides are fairly pleasant smelling liquids, but arylmethyl (benzylic) halides of structure $\ce{ARCH_2X}$ are irritating to the eyes, skin, and nasal passages. Toxicity varies, but the chlorinated hydrocarbons such as $\ce{CCl_4}$ (“carbon tet”) and $\ce{CHCl_2-CHCl_2}$ are quite toxic and should be used with care.
Contributors
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
14.03: Spectroscopic Properties
Organohalogen compounds give rise to strong absorptions in the infrared arising from stretching vibrations of the carbon-halogen bond. The frequency of absorption decreases as the mass of the halogen increases. For monohaloalkanes the absorptions useful for identification are those of $\ce{C-F}$ at $1100$-$1000 \: \text{cm}^{-1}$ and $\ce{C-Cl}$ at $850$-$550 \: \text{cm}^{-1}$. The $\ce{C-Br}$ and $\ce{C-I}$ absorptions are below $690 \: \text{cm}^{-1}$ and therefore are out of range of most commercial spectrophotometers. Because these bands are in the fingerprint region or far infrared, it is difficult to infer the presence of halogen in a molecule solely from its infrared spectrum.
Apart from fluorine, the magnetic properties of halogen nuclei do not complicate proton or $\ce{^{13}C}$ nuclear magnetic resonance spectra of organohalogen compounds. But fluorine $\left( \ce{^{19}F} \right)$ has a spin of 1/2 and causes spin-spin splitting of the resonances of neighboring magnetic nuclei ($\ce{^{13}C}$, $\ce{^1H}$, and other $\ce{^{19}F}$ nuclei). Proton chemical shifts are influenced strongly by the presence of halogen, which serves to deshield neighboring protons by electronegativity effects (see Section 9-10E).
The mass spectra of chlorine- and bromine-containing compounds clearly show the abundance ratios of the stable isotopes $\ce{^{35}Cl}$:$\ce{^{37}Cl} =$ 3:1 and $\ce{^{79}Br}$:$\ce{^{81}Br} =$ 1:1 in the molecular ions and those ionic fragments which contain halogens (Section 9-11).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/14%3A_Organohalogen_and_Organometallic_Compounds/14.02%3A_Physical_Properties_of_Organohalogen_and_Organometallic_Compounds.txt |
The important chemistry of alkyl halides, $\ce{RX}$, includes the nucleophilic $\left( S_\text{N} \right)$ displacement and elimination $\left( E \right)$ reactions discussed in Chapter 8. Recall that tertiary alkyl halides normally are reactive in ionization $\left( S_\text{N}1 \right)$ reactions, whereas primary halides, and to a lesser extent secondary halides, are reactive in $S_\text{N}2$ reactions, which occur by a concerted mechanism with inversion of configuration (Sections 8-4 to 8-7).
Elimination competes with substitution in many $S_\text{N}$ reactions and can become the major pathway at high temperatures or in the presence of strong base. Elimination $\left( E_2 \right)$, unlike displacement $\left( S_\text{N}2 \right)$, is insensitive to steric hindrance in the alkyl halide. In fact, the $E_2$ reactivity of alkyl halides is tert $\ce{RX} >$ sec $\ce{RX} >$ prim $\ce{RX}$, which is opposite to their $S_\text{N}2$ reactivity.
Several useful reactions for the synthesis of alkyl halides that we already have encountered are summarized below with references to the sections that supply more detail:
A summary of these and some other reactions for the synthesis of alkyl halides or organohalogen compounds is given in Table 14-5.
Allylic (2-Propenyl) Halides
Halogen compounds in which the carbon-halogen bond is adjacent to a double bon, as in are known as allylic halides. The simplest example is 3-chloropropene, $\ce{CH_2=CHCH_2Cl}$, which is made on a large scale by the radical chlorination of propene at $400^\text{o}$:
Most of the 3-chloropropene prepared in this manner is converted to other important compounds. For example, addition of hypochlorous acid gives a mixture of dichloropropanols, which on treatment with base gives a substance known commercially as "epichlorohydrin":
The ring closure reaction with $\ce{Ca(OH)_2}$ is an internal $S_\text{N}2$ reaction. Hydroxide ion converts the alcohol to an alkoxide ion that acts as a nucleophile in displacing the neighboring chlorine:
The epichlorohydrice so produced is used primarily to make epoxy resins (see Section 29-5E), although some of it is hydrolyzed to glycerol:
A general method for preparing allylic halides is by addition of hydrogen halides to conjugated dienes. This reaction usually produces a mixture of 1,2- and 1,4-addition products (see Section 13-2):
A second general method involves the bromination of alkene with N-bromosuccinimide (the Wohl-Ziegler reaction). A radical-chain reaction takes place between N-bromosuccinimide (NBS) and alkenes, which commonly is initiated by light, peroxides, or other catalysts, and yields allylic bromides:
This reaction, like the chlorination of propene, is highly selective in that the so-called allylic $\ce{C-H}$ is attacked preferentially.
From bond energies (Table 4-6) we know that the weakest $\ce{C-H}$ bonds of propene are to the allylic hydrogens, $\ce{H_2C=CHCH_2-H}$. Therefore, in the first step of radical-chain chlorination of propene, an allylic hydrogen is removed by a chlorine atom (Equation 14-1). The allylic $\ce{C-H}$ bonds are weaker than the alkenic $\ce{C-H}$ bonds because of the extra stabilization of the radical obtained on hydrogen abstraction (Equation 14-1). Two equivalent valence-bond structures ($1a$ and $1b$) can be written for the 2-propenyl radical; the electron delocalization enhances the stability of the radical (see Section 6-5C):
In the second step of the chain reaction (Equation 14-2) the propenyl radical can form a carbon-halogen bond at either end by abstracting a halogen atom from the halogenating agent:
The $\ce{Cl} \cdot$ atom produced now can participate in Reaction 14-1, thereby continuing the chain. With propene the intermediate radical gives the same product, 2-propenyl chloride, irrespective of whether a chlorine atom is transferred to the 1- or 3-carbon. However, the radical formed by removal of an allylic hydrogen from 2-butene gives a mixture of products:
$S_\text{N}$ Reactions of Allylic Halides
The carbon-halogen bonds of allylic halides are especially reactive in both $S_\text{N}1$ and $S_\text{N}2$ reactions (Table 14-6). The reasons for the enhanced $S_\text{N}1$ reactivity have been discussed previously (Section 8-7B). For example, the ease with which 1-chloro-2-butene ionizes compared to 1-chlorobutane is attributed to the stability of the 2-butenyl cation, which is distributed between $\ce{C_1}$ and $\ce{C_3}$, and the nucleophile (water) attacks at both positions to give mixtures of products. The same results are obtained if one starts with 3-chloro-1-butene because the same cation is formed:
Benzylic (Phenylmethyl) Halides
Reactivities comparable to allylic halides are found in the nucleophile displacement reactions of benzylic halides by $S_\text{N}1$ and $S_\text{N}2$ mechanisms (Table 14-6). The ability of the benzylic halides to undergo $S_\text{N}1$ reactions clearly is related to the stability of the resulting benzylic cations, the electrons of which are extensively delocalized. Thus, for phenylmethyl chloride,
When the halogen substituent is located two or more carbons from the aryl group as in 2-phenylethyl bromide, $\ce{C_6H_5CH_2CH_2Br}$, the pronounced activating effect evident in benzylic halides disappears, and the reactivity of the halides is essentially that of a primary alkyl halide (e.g., $\ce{CH_3CH_2CH_2Br}$).
Benzylic halides can be prepared by the same radical-halogenating agents that give allylic halides from alkenes. These include $\ce{Cl_2}$, $\ce{Br_2}$, N-bromosuccinimide (Section 14-3A), $\ce{SO_2Cl_2}$, and tert-butyl hypochlorite:
The benzylic $\ce{C-H}$ bond is weaker and more restrictive then primary alkane $\ce{C-H}$ bonds because of the stabilization of benzylic radicals (see Table 4-6).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/14%3A_Organohalogen_and_Organometallic_Compounds/14.04%3A_Alkyl_Halides.txt |
The most readily available alkenyl halide is chloroethene (vinyl chloride), which can be prepared by a number of routes:
The most economical commercial preparation is high-temperature chlorination of ethene. A useful modification of this process uses hydrogen chloride in place of chlorine. An oxidizing agent is required to raise the oxidation state of chlorine in $\ce{HCl}$ to that of $\ce{Cl_2}$; molecular oxygen is used for this purpose along with cupric salts as catalysts.
General methods of preparation for alkenyl and alkynyl halides are listed in Table 14-5. By the alkynyl halides we mean 1-halo-alkynes. One interesting method by which they may be prepared employs 1-alkynes with hypohalites:
This kind of reaction does not proceed with either alkanes or alkenes.
Uses of Alkenyl Halide
Chloroethene is produced in vast quantities for the production of polymers (polyvinyl chloride) and copolymers:
These polymers commonly are described as PVC plastics or less specifically as “vinyl.” They are materials that may be either flexible or rigid according to what they are mixed with, and they are used in the manufacture of many familiar articles such as plastic curtains, rainwear, floor tile, synthetic leather goods, upholstery, table mats, phonograph records, insulation, plastic pipes, tubing, and packaging materials.
Recently, it has been found that persons working in plants that manufacture and use chloroethene have an unusually high incidence of an unusual type of liver cancer. As a result, strict safety regulations and pollution standards have been set for plants where chloroethene is made or used. The once widespread use of chloroethene as a propellant for aerosol cans has been curtailed. Polyvinyl chloride itself seems to be quite safe, but there are possible problems with its incorporation into interior building materials, clothing, and upholstery because heat, such as fire, causes polyvinyl chloride to decompose, thereby producing hydrogen chloride as one decomposition product. In closed areas the toxicity of hydrogen chloride gas may be as serious a hazard as the fire itself. Other polymers may give off similarly toxic products on strong heating.
Chemical Properties
The outstanding chemical characteristic of alkenyl halides is their general inertness in $S_\text{N}1$ and $S_\text{N}2$ reactions. Thus chloroethene fails to react with silver nitrate in ethanol (i.e., low $S_\text{N}1$ reactivity), fails to react with potassium iodide in acetone (i.e., low $S_\text{N}2$ reactivity), and only reacts slowly with sodium hydroxide to give ethyne (low $E2$ reactivity). The haloalkynes, such as $\ce{RC \equiv C-Cl}$, are similarly unreactive.
It is not surprising that $\ce{=C-X}$ and $\ce{\equiv C-X}$ bonds are hard to break heterolytically. In general, $\ce{C-X}$ bonds are strong in alkenyl halides (cf. Table 4-6) and this property tends to make them less reactive than alkyl halides. Furthermore, double- and triple-bonded carbons are more strongly electron-attracting than saturated $sp^3$ carbons, which is the reason why 1-alkynes and alkenes are stronger acids (Section 11-8) than alkanes. Consequently it is easier to break a $\ce{\equiv C-H}$ bond in the sense $\ce{C}^\ominus \ce{H}^\oplus$ than as $\ce{\equiv C}^\oplus \ce{H}^\ominus$. It also will be more difficult to ionize a carbon-halogen bond to $\ce{C}^\oplus \ce{X}^\ominus$ if the carbon is unsaturated. Therefore ethenyl and ethynyl cations, such as $\ce{CH_2=CH}^\oplus$ and $\ce{HC \equiv C}^\oplus$, are difficult to generate from the corresponding halides. Superior leaving groups are required, such as trifluoromethanesulfonate, $\ce{-OSO_3CF_3}$ (Section 8-7C):
The reason for the lack of $S_\text{N}2$ reactivity in ethenyl or ethynyl halides may be that the attacking nucleophile is unable to react by the concerted inversion mechanism that invariably is observed with alkyl halides:
Nevertheless, substitution of the halogen does occur under some circumstances. In such cases, the nucleophile first adds to the multiple bond, and in a subsequent step the halogen leaves as halide ion. This is an “addition-elimination” mechanism, of which we will have more examples later:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
14.06: Cycloalkyl Halides
The cycloalkyl halides, except for cyclopropyl halides, have physical and chemical properties that are similar to those of the open-chain secondary halides and can be prepared by the same types of reactions (Table 14-5). All the cycloalkyl halides undergo $S_\text{N}2$ reactions rather slowly and, with nucleophiles that are reasonably basic ($^\ominus \ce{OH}$, $^\ominus \ce{OC_2H_5}$, $^\ominus \ce{C \equiv N}$, etc.), $E2$ reactions can be expected to predominate (Table 14-6). The rate of carbocation formation leading to $S_\text{N}1$ and $E1$ reactions is sensitive to ring size but, except for the small-ring halides, the carbocation reactions are normal in most other respects.
The cyclopropyl halides are exceptional in that their behavior is much more like alkenyl halides than like secondary alkyl halides. Thus cyclopropyl chloride undergoes $S_\text{N}1$ and $S_\text{N}2$ reactions much less rapidly than isopropyl or cyclohexyl chlorides. A relationship between the reactivity of cyclopropyl chloride and chloroethene is not surprising in view of the general similarity between cyclopropane rings and double bonds (Section 12-5). This similarity extends to cyclopropylmethyl derivatives as well. Cyclopropylmethyl chloride is reactive in both $S_\text{N}1$ and $S_\text{N}2$ reactions in much the same way as 3-chloropropene:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/14%3A_Organohalogen_and_Organometallic_Compounds/14.05%3A_Alkenyl_and_Alkynyl_Halides.txt |
Aryl halides have a halogen directly bonded to a carbon of an aromatic ring. Examples are bromobenzene, fluorobenzene, and 2,4-dichloromethylbenzene:
Some of the methods by which alkyl halides are prepared do not work for aryl halides because it is difficult to form $\ce{C}$-halogen bonds at aromatic ring carbons by nucleophilic displacement reactions. The most common ways of forming $\ce{C}_\text{aryl}$-halogen bonds are by substitution of $\ce{C}_\text{aryl} \ce{-H}$ by electrophilic halogenating agents (e.g., $\ce{Br_2}$ or $\ce{Cl_2}$),
and by replacement of $\ce{C-NH_2}$ by $\ce{C}-$halogen. These reactions are listed in Table 14-5 and will be discussed in more detail in Chapters 22 and 23.
Nucleophilic Aromatic Displacement Reactions
The carbon-halogen bonds of aryl halides are like those of alkenyl halides in being much stronger than those of alkyl halides (see Table 4-6). The simple aryl halides generally are resistant to attack by nucleophiles in either $S_\text{N}1$ or $S_\text{N}2$ reactions (Table 14-6). However, this low reactivity can be changed dramatically by changes in the reaction conditions and the structure of the aryl halide. In fact, nucleophilic displacement becomes quite rapid (a) when the aryl halide is activated by substitution with strongly electron-attracting groups such as $\ce{NO_2}$, and (b) when very strongly basic nucleophilic reagents are used.
Addition-Elimination Mechanism of Nucleophilic Substitution
Although the simple aryl halides are inert to the usual nucleophilic reagents, considerable activation is produced by strongly electron-attracting substituents provided these are located in either the ortho or para positions, or both. For example, the displacement of chloride ion from 1-chloro-2,4-dinitrobenzene by dimethylamine occurs readily in ethanol solution at room temperature. Under the same conditions chlorobenzene completely fails to react; thus the activating influence of the two nitro groups amounts to a factor of at least $10^8$:
A related reaction is that of 2,4-dinitrofluorobenzene with the amino groups of peptides and proteins, and this reaction provides a means for analysis of the N-terminal amino acids in polypeptide chains. (See Section 25-7B.)
In general, the reactions of activated aryl halides closely resemble the $S_\text{N}2$-displacement reactions of aliphatic halides. The same nucleophilic reagents are effective (e.g., $\ce{CH_3O}^\ominus$, $\ce{HO}^\ominus$, and $\ce{RNH_2}$); the reactions are second order overall (first order in halide and first order in nucleophile); and for a given halide the more nucleophilic the attacking reagent, the faster the reaction. However, there must be more than a subtle difference in mechanism because an aryl halide is unable to pass through the same type of transition state as an alkyl halide in $S_\text{N}2$ displacements.
The generally accepted mechanism of nucleophilic aromatic substitution of aryl halides carrying activating groups involves two steps that are closely analogous to those briefly described in Section 14-4 for alkenyl and alkynyl halides. The first step involves attack of the nucleophile $\ce{Y}^\ominus$ at the carbon bearing the halogen substituent to form an intermediate carbanion $4$ (Equation 14-3). The aromatic system is destroyed on forming the anion, and the carbon at the reaction site changes from planar ($sp^2$ bonds) to tetrahedral ($sp^3$ bonds).
In the second step, loss of an anion, $\ce{X}^ominus$ or $\ce{Y}^\ominus$, regenerates an aromatic system, and, if $\ce{X}^\ominus$ is lost, the overall reaction is nucleophilic displacement of $\ce{X}$ by $\ce{Y}$ (Equation 14-4).
In the case of a neutral nucleophilic reagent, $\ce{Y}$ or $\ce{HY}$, the reaction sequence would be the same except for the necessary adjustments in the charge of the intermediate:
Why is this reaction pathway generally unfavorable for the simple aryl halides? The answer is that the intermediate $4$, which we can express as a hybrid of the valence-bond structures $4a$-$4c$, is too high in energy to be formed at any practical rate. Not only has $4$ lost the aromatic stabilization of the benzene ring, but its formation results in transfer of negative charge to the ring carbons, which themselves are not very electronegative:
However, when strongly electron-attracting groups are located on the ring at the ortho-para positions, the intermediate anion is stabilized by delocalization of electrons from the ring carbons to more favorable locations on the substituent groups. As an example, consider the displacement of bromine by $\ce{OCH_3}$ in the reaction of 4-bromonitrobenzene and methoxide ion:
The anionic intermediate formed by addition of methoxide ion to the aryl halide can be described by the valence-bond structures $5a$-$5d$. Of these structures $5d$ is especially important because in it the charge is transferred from the ring carbons to the oxygen of the nitro substituent:
Substituents in the meta positions have much less effect on the reactivity of an aryl halide because delocalization of electrons to the substituent is not possible. No formulas can be written analogous to $5c$ and $5d$ in which the negative charges are both on atoms next to positive nitrogen, $\overset{\ominus}{\ce{C}} \overset{\oplus}{\ce{-N}-} \overset{\ominus}{\ce{O}}$ and $\overset{\ominus}{\ce{O}} \overset{\oplus}{\ce{-N}-} \overset{\ominus}{\ce{O}}$,
In a few instances, stable compounds resembling the postulated reaction intermediate have been isolated. One classic example is the complex $7$ (isolated by J. Meisenheimer), which is the product of the reaction of either the methyl aryl ether 6$6$ with potassium ethoxide, or the ethyl aryl ether $8$ and potassium methoxide:
Elimination-Addition Mechanism of Nucleophilic Aromatic Substitution. Arynes
The reactivities of aryl halides, such as the halobenzenes, are exceedingly low toward nucleophilic reagents that normally effect displacements with alkyl halides and activated aryl halides. Substitutions do occur under forcing conditions of either high temperatures or very strong bases. For example, chlorobenzene reacts with sodium hydroxide solution at temperatures around $340^\text{o}$ and this reaction was once an important commercial process for the production of benzenol (phenol):
In addition, aryl chlorides, bromides, and iodides can be converted to areneamines $\ce{ArNH_2}$ by the conjugate bases of amines. In fact, the reaction of potassium amide with bromobenzene is extremely rapid, even at temperatures as low as $-33^\text{o}$ with liquid ammonia as solvent:
However, displacement reactions of this type differ from the previously discussed displacements of activated aryl halides in that rearrangement often occurs. That is, the entering group does not always occupy the same position on the ring as that vacated by the halogen substituent. For example, the hydrolysis of 4-chloromethylbenzene at $340^\text{o}$ gives an equimolar mixture of 3- and 4-methylbenzenols:
Even more striking is the exclusive formation of 3-methoxybenzenamine in the amination of 2-chloromethoxybenzene. Notice that this result is a violation of the principle of least structural change (Section 1-1H):
The mechanism of this type of reaction has been studied extensively, and much evidence has accumulated in support of a stepwise process, which proceeds first by base-catalyzed elimination of hydrogen halide $\left( \ce{HX} \right)$ from the aryl halide - as illustrated below for the amination of bromobenzene:
Elimination
The product of the elimination reaction is a highly reactive intermediate $9$ called benzyne, or dehydrobenzene, which differs from benzene in having two less hydrogen and an extra bond between two ortho carbons. Benzyne reacts rapidly with any available nucleophile, in this case the solvent, ammonia, to give an addition product:
Addition
The rearrangements in these reactions result from the attack of the nucleophile at one or the other of the carbons of the extra bond in the intermediate. With benzyne the symmetry is such that no rearrangement would be detected. With substituted benzynes isomeric products may result. Thus 4-methylbenzyne, $10$, from the reaction of hydroxide ion with 4-chloro-1-methylbenzene gives both 3- and 4-methylbenzenols:
In the foregoing benzyne reactions the base that produces the benzyne in the elimination step is derived from the nucleophile that adds in the addition step. This need not always be so, depending on the reaction conditions. In fact, the synthetic utility of aryne reactions depends in large part of the success with which the aryne can be generated by one reagent but captured by another. One such method will be discussed in Section 14-10C and involves organometallic compounds derived from aryl halides. Another method is to generate the aryne by thermal decomposition of a 1,2-disubstituted arene compound such as $11$, in which both substituents are leaving groups - one leaving with an electron pair, the other leaving without:
When $11$ decomposes in the presence of an added nucleophile, the benzyne intermediate is trapped by the nucleophile as it is formed. Or, if a conjugated diene is present, benzyne will react with it by a [4 + 2] cycloaddition. In the absence of other compounds with which it can react, benzyne will undergo [2 + 2] cycloaddition to itself:
Uses for Aryl Halogen Compounds
As with most organic halides, aryl halides most often are synthetic intermediates for the production of other useful substances. For example, chlorobenzene is the starting aryl halide for the synthesis of DDT; it also is a source of benzenol (phenol, Section 14-6C) which, in turn, has many uses (Section 26-1).
Several aromatic chloro compounds are used extensively as insecticides, herbicides, fungicides, and bactericides. They also have acquired much notoriety because in some instances their indiscriminate usage has led to serious problems. For example, hexachlorophene is an external bactericide that until recently was used in cosmetic preparations such as soaps, deodorants, and so on. Its use has been discontinued because of compelling evidence that it can be absorbed through the skin in amounts that are dangerous, if not lethal, for infants and small children.
Other pesticides, notably DDT and the herbicides 2,4-D and 2,4,5-T have been partially banned for different reasons.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/14%3A_Organohalogen_and_Organometallic_Compounds/14.07%3A_Aryl_Halides.txt |
Useful Compounds
Polychlorination of methane yields the di-, tri-, and tetrachloromethanes cheaply and efficiently:
These substances have excellent solvent properties for nonpolar and slightly polar substances. Chloroform once was used widely as an inhalation anesthetic. However, it has a deleterious effect on the heart and is oxidized slowly by atmospheric oxygen to highly toxic carbonyl dichloride (phosgene, $\ce{COCl_2}$). Commercial chloroform contains about $1\%$ ethanol, which destroys any $\ce{COCl_2}$ formed by oxidation.
Carbon tetrachloride commonly was employed as a cleaning solvent, although its considerable toxicity entails considerable hazard when used indiscriminately. It has been used as a fire-extinguishing fluid for petroleum fires, but its toxicity and tendency to form still more toxic carbonyl dichloride makes it undesirable for confined areas. The common laboratory practice of removing traces of water from solvents with metallic sodium should not be applied to halogenated compounds; carbon tetrachloride-sodium mixtures are shock sensitive and can detonate.
Trichloroethene ("Tri-Clene", bp $87^\text{o}$) is a widely used dry-cleaning solvent. It can be prepared from either ethene or ethyne:
Compared with monohaloalkanes, polyhalogen compounds have quite different reactivities and behavior toward nucleophiles and bases. Thus dichloromethane reacts with hydroxide ion by an $S_\text{N}2$ mechanism much less readily than methyl chloride. The chloromethanol formed then undergoes a rapid $E2$ elimination to give methanal (formaldehyde), a substance that exists in water largely as dihydroxymethane:
Trichloromethane (chloroform) reacts quite differently with base than does chloromethane or dichloromethane - as will be described in the following section.
$\alpha$ Elimination. Carbenes
Trihalomethanes, such as trichloromethane (chloroform), are quite reactive toward strong base. The base, such as hydroxide, removes the hydrogen of $\ce{HCCl_3}$ as a proton much more rapidly than it attacks the carbon in the $S_\text{N}2$ manner. The carbanion so formed, $\ce{Cl_3C}^\ominus$, is unstable and loses chloride ion to form a highly reactive neutral intermediate, $:\ce{CCl_2}$, called dichlorocarbene:
This intermediate has only six valence electrons around carbon and therefore is strongly electrophilic. In aqueous solution it reacts rapidly to form carbon monoxide and methanoate (formate) ion:
The formation of $:\ce{CCl_2}$ from $\ce{HCCl_3}$ by the reactions of Equation 14-6 results in the elimination of $\ce{HCl}$ - the leaving groups, $\ce{H}$ and $\ce{Cl}$, both originating from the same carbon atom. Such reactions are not uncommon and are called $\alpha$ eliminations or 1,1 eliminations to distinguish them from $E1$ and $E2$ reactions, which are $\beta$ eliminations or 1,2 eliminations. Still other possibilities are reactions such as $\gamma$ or 1,3 eliminations, but these take on the character of internal $S_\text{N}2$ reactions and will not be considered in detail here.
The product of $\alpha$ elimination is a neutral species that resembles a carbocation in having only six carbon valence electrons. The simplest carbene is $:\ce{CH_2}$, methylene. Carbenes are highly reactive, so much so that they cannot be isolated. Their involvement in reactions usually has to be inferred from the nature of the products or the reaction kinetics. The characteristic carbene reactions involve forming an electron-pair bond to the carbene carbon by reacting with $\sigma$ bonds, $\pi$ bonds, or unshared pairs $\left( n \right)$. Some of these reactions are illustrated here for methylene $:\ce{CH_2}$.$^1$
with $\sigma$ bonds (insertion):
with $\pi$ bonds ([2 + 1] cycloaddition):
with unshared pairs (dimerization, addition):
Carbenes are much more reactive toward carbon-carbon double bonds than toward single bonds. Without doubt the most useful feature of $\alpha$ elimination is that it provides a practical route to cyclopropanes and cyclopropenes by [2 + 1] cycloaddition of carbenes to double or triple bonds. These additions are stereospecific suprafacial additions if they involve singlet carbenes, but can give mixtures with triplet carbenes:
Carbene precursors are compounds that have or acquire good leaving groups (e.g., halide ions). Thus, halogen compounds frequently are carbene sources. Trihalomethanes are the oldest known sources of dihalocarbenes; but there are other methods for generating carbenes, and some of these are listed for reference in Table 14-2 (see also Section 14-10C). There is a question as to whether a “free” carbene actually is formed in some of these reactions, particularly those involving metals, but for our purposes we will classify them as routes to carbenes or carbenelike species.
Table 14-2: $\alpha$-Elimination Reactions Producing Carbene Intermediates$^a$
Many carbenes, like carbocations, rearrange to more stable structures by the migration of a neighboring group to the electron-deficient carbon. Thus phenylmethylcarbene rearranges to ethenylbenzene (styrene):
Fluorochloromethanes
Replacement of either one or two of the chlorines of carbon tetrachloride by fluorine can be achieved readily with antimony trifluoride containing some antimony pentachloride. The reaction stops after two chlorines have been replaced. The antimony trifluoride can be regenerated continuously from the antimony chloride by addition of anhydrous hydrogen fluoride:
Both products are useful as refrigerants, particularly for household refrigerators and air-conditioning units, under the trade name Freon. Difluorodichloromethane (Freon 12) also is employed as a propellant in aerosol bombs, shaving-cream dispensers, and other such containers. It is nontoxic, odorless, nonflammable, and will not react with hot concentrated mineral acids or metallic sodium. This lack of reactivity is generally characteristic of the difluoromethylene group, provided the fluorines are not located on an unsaturated carbon. Attachment of a fluorine atom to a carbon atom bonded to one or more chlorine atoms tends greatly to reduce the reactivity of the chlorines toward almost all types of reagents. Possible environmental problems associated with these substances were discussed in the introduction to this chapter.
Fluorocarbons
During World War II, plastics and lubricating compounds of unusual chemical and thermal stability were required for many applications, in particular for pumping apparatus used to separate $\ce{^{235}U}$ from $\ce{^{238}U}$ by diffusion of corrosive uranium hexafluoride through porous barriers. It was natural to consider the use of substances made only of carbon and fluorine (fluorocarbons) for such purposes, and considerable effort was spent on methods of preparing compounds such as $\ce{-(CF_2)}-_n$. Today, many such substances are in common use. These often are called "perfluoro-" compounds, which indicates that all available hydrogens of the parent compound are replaced by fluorine. Thus perfluorocyclohexane is $\ce{(CF_2)_6}$. A widely used perfluorocarbon is the plastic material $\ce{-(CF_2)}-_n$, which is produced in quantity by radical polymerization of tetrafluoroethene:
The product ("Teflon") is a solid, chemically inert substance that is stable to around $300^\text{o}$. It makes excellent electrical insulation and gasket materials. It also has self-lubricating properties, which are exploited in the preparation of low-adhesion surfaces (such as "nonstick" fry pans) and light-duty bearings.
Tetrafluoroethene can be made on a commercial scale by the following method:
The latter reaction involves difluorocarbene $\left( :\ce{CF_2} \right)$:
In the presence of peroxides, tetrafluoroethene polymerizes to the long-chain polymer. If peroxides are excluded, [2 + 2] cycloaddition occurs in high yield to give octafluorocyclobutane (see Section 13-3D):
Similar cycloaddition reactions occur with chlorotrifluoroethene and 1,1-dichloro-2,2-difluoroethene.
Radical polymerization of chlorotrifluoroethene gives a useful polymer (Kel-F) that is similar to Teflon.
An excellent elastomer of high chemical resistance (Viton) can be made by copolymerizing hexafluoropropene with 1,1-difluoroethene. The product is stable to $300^\text{o}$ and is not attacked by hot concentrated nitric acid. Although expensive, it is unrivaled among elastomers for chemical durability under extreme conditions.
Properties of Fluorocarbons
The fluorocarbons have extraordinarily low boiling points relative to the hydrocarbons of comparable molecular weight. As seen in Figure 14-3, their boiling points are nearly the same or even lower than those of the alkanes or cycloalkanes with the same number of carbons. Thus octafluorocyclobutane or boils $17^\text{o}$ lower than cyclobutane, despite an almost fourfold greater molecular weight!
Fluorocarbons are very insoluble in most polar solvents and are only slightly soluble in alkanes in the kerosene range. The higher-molecular-weight fluorocarbons are not even miscible in all proportions with their lower-molecular-weight homologs.
The physiological properties of organofluorine compounds vary widely. Dichlorodifluoromethane and the saturated fluorocarbons appear to be completely nontoxic. In contrast, perfluoro-2-methylpropene is exceedingly toxic, more so than the war gas, carbonyl dichloride $\left( \ce{COCl_2} \right)$. Sodium fluoroethanoate $\left( \ce{CH_2FCO_2Na} \right)$ and 2-fluoroethanol are toxic fluorine derivatives of oxygen-containing organic substances. The fluoroethanoate salt is sold commercially as a rodenticide. Interestingly, sodium trifluoroethanoate is nontoxic.
Fluorocarbon derivatives have another interesting and potentially useful property. They dissolve large quantities of oxygen. This fact, combined with their nontoxicity, has led to their use as blood replacements in heart surgery on experimental animals. Mice can live totally immersed in oxygen-saturated liquid fluorocarbons.
$^1$Life with carbenes is substantially complicated by the fact that there are two different forms (singlet and triplet) of $:\ce{CH_2}$ and presumably of all other carbenes. The two forms of $:\ce{CH_2}$ differ considerably in their reactivity. One is the singlet, which has its unshared electrons paired, while the other is the triplet with the same electrons unpaired. For $:\ce{CH_2}$, the singlet form is the less stable and more reactive, whereas with $:\ce{CCl_2}$, the triplet is the less stable and more reactive.
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/14%3A_Organohalogen_and_Organometallic_Compounds/14.08%3A_Polyhalogenated_Alkanes_and_Alkenes.txt |
One of the more important reactions of organohalogen compounds is the formation of organometallic compounds by replacement of the halogen by a metal atom. Carbon is positive in carbon-halogen bonds and becomes negative in carbon-metal bonds, and therefore carbon is considered to be reduced in formation of an organometallic compound (Section 11-1):
This transformation is of value because it makes an electrophilic carbon into a nucleophilic carbon. Organometallic compounds are a convenient source of nucleophilic carbon. A typical example of their utility is the way the achieve addition of nucleophilic carbon to carbonyl groups with formation of carbon-carbon bonds:
In this chapter we will restrict our discussion of organometallic compounds to the alkyl and aryl compounds of magnesium and lithium, and the sodium and potassium salts of 1-alkynes. These substances normally are derived directly or indirectly from organohalogen compounds and are used very widely in organic synthesis. Organometallic compounds of transition metals and of boron are discussed in Chapters 11 and 31.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
14.10: Properties of Organometallic Compounds
How carbon-metal bonds are formed depends on the metal that is used. Conditions that are suitable for one metal may be wholly unsuited for another. Some organometallic compounds react very sluggishly even toward acids, whereas others react avidly with water, oxygen, carbon dioxide, and almost all solvents but the alkanes themselves. Reactivity increases with increasing polarity of the carbon-metal bond, which is determined by the electropositivity of the metal. Strongly electropositive metals, such as sodium and potassium, form largely ionic bonds to carbon, as we have mentioned in the case of alkynide salts, $\ce{RC \equiv C}^\ominus \ce{Na}^\oplus$ (Section 11-8). Estimates of the ionic character of various carbon-metal bonds are given in Table 14-3, and it will be seen that organosodium and organopotassium compounds have the most ionic bonds and they are, in fact, among the most reactive organometallic compounds known. Many organosodium and organopotassium compounds burn spontaneously when exposed to air and react violently with water and carbon dioxide. As might be expected from their saltlike character, they are nonvolatile and do not dissolve readily in nonpolar solvents. In contrast, the more covalent, less ionic, organometallic compounds, such as $\ce{(CH_3)_2Hg}$, are far less reactive; they are stable in air, quite volatile, and dissolve in nonpolar solvents.
Table 14-3: Percent Ionic Character of Carbon-Metal Bonds$^a$
Bond Percent Ionic Chateracter Bond Percent Ionic Chateracter Bond Percent Ionic Chateracter
C-K 51% C-Mg 35% C-Sn 12%
C-Na 47% C-Al 22% C-Pb 12%
C-Li 43% C-Zn 18% C-Hg 9%
C-Ca 43% C-Cd 15%
Source: L. Pauling, The Nature of the Chemical Bond, Cornell University Press, Ithaca, N.Y. 3rd, 1960, Chap. 3
All of these compounds must be handled with great care because some are dangerously reactive and others are very toxic. They seldom are isolated from the solutions in which they are prepared, but are used immediately in other reactions.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/14%3A_Organohalogen_and_Organometallic_Compounds/14.09%3A_Organometallic_Compounds_from_Organohalogen_Compounds.txt |
Metals with Organic Halides
The reaction of a metal with an organic halide is a convenient method for preparation of organometallic compounds of reasonably active metals such as lithium, magnesium, and zinc. Ethers, particularly diethyl ether and oxacyclopentane (tetrahydrofuran), provide inert, slightly polar media in which organometallic compounds usually are soluble. Care is necessary to exclude moisture, oxygen, and carbon dioxide, which would react with the organometallic compound. This can be accomplished by using an inert atmosphere of nitrogen or helium.
The reactivity order of the halides is $\ce{I} \: > \: \ce{Br} \: > \: \ce{Cl} \: \gg \ce{F}$. Whereas magnesium and lithium react well with chlorides, bromides, and iodides, zinc is satisfactory only with bromides and iodides. Mercury only reacts when amalgamated with sodium. Sodium and potassium present special problems because of the high reactivity of alkylsodium and alkylpotassium compounds toward ether and organic halides. Alkane solvents usually are necessary.
Alkenyl, alkynyl, and aryl halides, like alkyl halides, can be converted to the corresponding magnesium and lithium compounds. However, the reaction conditions, such as choice of solvent, can be critical. Bromoethene, for instance, can be converted to ethenylmagnesium bromide in good yield if the solvent is oxacyclopentane [tetrahydrofuran, $\ce{(CH_2)_4O}$]:
The more reactive allylic and benzylic halides present a problem - not so much in forming the organometallic derivative as in keeping it from reacting further with the starting halide. An often unwanted side reaction in the preparation of organometallic compounds is a displacement reaction, probably of the $S_\text{N}2$ type:
This problem can be lessened greatly by using a large excess of magnesium and dilute solutions of the allylic halide to minimize the coupling reaction.
The same difficulty also occurs in the preparation of alkylsodium compounds. The starting halide $\ce{RX}$ couples with $\ce{RNa}$ (to give $\ce{R-R}$ and $\ce{NaX}$) or is converted to an alkene. These reactions appear to involve radical intermediates undergoing combination and disproportionation (Section 10-8C):
In the absence of metallic sodium, ethylsodium probably still reacts with ethyl bromide by a radical reaction rather than $S_\text{N}2$ or $E2$. This happens because $\ce{CH_3CH_2^-}$ tends to lose an electron easily and can act like metallic sodium to donate an electron to $\ce{CH_3CH_2Br}$ to form an ethyl radical and itself become an ethyl radical:
Reactions between the resulting radicals then produce butane, ethane, and ethene.
The point at which one can expect $S_\text{N}2$ and $E2$ reactions to go faster than radical formation as the structures of the halides and the nature of the metal are changed is not yet clearly defined. However, it is becoming increasingly evident that there are substitution reactions of "unactivated" aryl halides that proceed without rearrangement by way of radical intermediates. The key step in these reactions is donation of an electron to one of the unfilled $\pi$ orbitals of the ring and subsequent ejection of a halide ion:
Such a mechanism probably is involved in the formation of organometallic compounds from aryl halides and metals.
Some Other Preparations of Organometallic Compounds
Brief descriptions follow of less general but very useful methods of forming organometallic compounds (also see Table 14-7). In each of these preparations the solvent must be inert to all of the organometallic compounds involved.
Halogen-Metal Exchange
$\ce{RBr} + \ce{R'Li} \rightleftharpoons \ce{RLi} + \ce{R'Br}$
The equilibrium in these reactions favors formation of the organometallic compound with the metal attached to the more electronegative $\ce{R}$ group. The method is mainly used in the preparation of organolithium compounds derived from unreactive halides such as aryl, ethenyl, or ethynyl halides. These halides do not always react readily with lithium metal, but may react well with butyllithium:
Displacement of One Metal by Another
$\ce{R_2Hg} + 2 \ce{Na} \rightleftharpoons 2 \ce{RNa} + \ce{Hg}$
Here the equilibrium is such that the $\ce{R}$ group favors attachment to the more electropositive metal.
Organometallic Compounds with Metal Halides
$\ce{RMgCl} + \ce{HgCl_2} \rightleftharpoons \ce{RHgCl} + \ce{MgCl_2}$
$\ce{RLi} + \ce{CuI} \rightleftharpoons \ce{RCu} + \ce{LiI}$
The equilibrium favors the products with $\ce{R}$ connected to the less electropositive metal so the reaction tends to form a less reactive organometallic compound from a more reactive one.
Organometallic Compounds from Acidic Hydrocarbons
Some organometallic compounds are prepared best by the reaction of a strong base or an alkyl metal derivative with an acidic hydrocarbon, such as an alkyne:
An especially important example is that of 1,3-cyclopentadiene, which is acidic because its conjugate base (cyclopentadienide anion) is greatly stabilized by electron delocalization. The anion is formed easily from the hydrocarbon and methyllithium:
Organometallic Compounds from Polyhalogen Compounds
Diorganometallic compounds cannot be prepared from dihalides if the halogens are separated by three $\ce{C-C}$ bonds or less because elimination or other reactions usually predominate. With active metals and 1,1-, 1,2-, or 1,3-dihalides, the following reactions normally occur:
When the halogens are at least four carbons apart a diorganometallic compound can be formed:
Carbenes, $\ce{R_2C} :$ (Section 14-7B) are produced by $\alpha$ eliminations from polyhalogen compounds with organometallic reagents. The first step is halogen-metal exchange and this is followed by elimination of metal halide:
Elimination reactions of this type can be useful in synthesis for the formation of carbon-carbon bonds. For example, if dibromocarbene is generated in the presence of an alkene, it will react by cycloaddition to give a cyclopropane derivative:
A related example is the generation of benzyne from 1-bromo-2-fluorobenzene with magnesium in oxacyclopentane (tetrahydrofuran). If the temperature is kept around $0^\text{o}$, 2-fluorophenylmagnesium bromide is formed. At higher temperatures, magnesium halide is eliminated and benzyne results:
If a diene is present, the benzyne will react with it by a [4 + 2] cycloaddition as in the following example:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
14.12: Organomagnesium Compounds
For many years the most important organometallic compounds for synthetic purposes have been the organomagnesium halides, or Grignard reagents. They are named after Victor Grignard, who discovered them and developed their use as synthetic reagents, for which he received a Nobel Prize in 1912. As already mentioned, these substances customarily are prepared in dry ether solution from magnesium turnings and an organic halide:
Chlorides often react sluggishly and, in addition, may give an unwelcome precipitate of magnesium chloride, which, unlike magnesium bromide and iodide, is only very slightly soluble in ether. Organomagnesium fluorides eluded preparation until quite recently.
Although we usually write the structure of a Grignard reagent as $\ce{RMgX}$, in which $\ce{X}$ is a halogen, the structure of the reagent in ether solution is more complex. There is a rapidly established equilibrium between the organomagnesium halide $\left( \ce{RMgX} \right)$ and the corresponding dialkylmagnesium $\left( \ce{RMgR} \right)$:
$2 \ce{RMgX} \rightleftharpoons \ce{R_2Mg} + \ce{MgX_2}$
Both of these species, $\ce{RMgX}$ and $\ce{R_2Mg}$, are reactive, and in ether solvents are solvated by coordination of the ether oxygen to magnesium. They further associate as dimers or higher polymers in solution. Although it is an oversimplification to regard a Grignard reagent as $\ce{RMgX}$, most of the reactions can be rationalized easily by this simple structure.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/14%3A_Organohalogen_and_Organometallic_Compounds/14.11%3A_Preparation_of_Organometallic_Compounds.txt |
Additions to Carbonyl Groups. Synthesis of Alcohols
The most important synthetic use of Grignard reagents and organolithium reagents is to form new carbon-carbon bonds by addition to polar multiple bonds, particularly carbonyl bonds. An example is the addition of methyl-magnesium iodide to methanal:
The yields of addition products in reactions of this kind are generally high. The adducts have metal-oxygen bonds that can be broken readily by acid hydrolysis to give the organic product. Grignard reagents seldom add to carbon-carbon multiple bonds (however, see Section 14-12D).
With suitable variations of the carbonyl compound, a wide range of compounds can be built up from substances containing fewer carbon atoms per molecule. The products formed when several types of carbonyl compounds react with Grignard reagents are listed in Table 14-4. The sequence of reactions starting with an organic halide, $\ce{RX}$, amounts to the addition of $\ce{R-H}$ across a carbonyl bond.$^2$
Primary alcohols can be prepared by the addition of $\ce{RMgX}$ or $\ce{RLi}$ to methanal, $\ce{CH_2=O}$,
Alcohols of formula $\ce{RCH_2CH_2OH}$ can be prepared by addition of $\ce{RMgX}$ to oxacyclopropane (oxirane):
Table 14-4: Products from the Reaction of Grignard Reagents $\left( \ce{RMgX} \right)$ with Carbonyl Compounds
Secondary alcohols are obtained from aldehydes, whereas ketones give tertiary alcohols:
Hydrolysis of the intermediate $\ce{R-OMgX}$ compound is achieved best with aqueous ammonium chloride solution. Addition of water gives an unpleasant mess of $\ce{Mg(OH)_2}$ whereas addition of strong acids such as $\ce{HCl}$ or $\ce{H_2SO_4}$ can lead to side reactions of dehydration and so on, especially with tertiary alcohols (Section 8-9C):
Tertiary alcohols also are obtained from both acyl halides, $\ce{RCOCl}$, and esters, $\ce{RCO_2R}$, by the addition of two moles of Grignard reagent. The first mole of $\ce{RMgX}$ adds to the carbonyl bond to give the adducts $13$ or $14$:
However, these first-formed adducts are unstable and decompose to a ketone, $\ce{CH_3COR}$, and magnesium salts, $\ce{MgXCl}$ or $\ce{MgXOC_2H_5}$. The ketone usually cannot be isolated, but reacts rapidly with more $\ce{RMgX}$ ultimately to give a tertiary alcohol:
Organolithium compounds behave very much like Grignard reagents, but with increased reactivity. They offer advantages over the magnesium compounds when the $\ce{R}$ group or the carbonyl compound is highly branched. For instance, isopropyllithium adds in good yield to 2,4-dimethyl-3-pentanone, whereas isopropylmagnesium bromide fails completely to give the normal addition product:
Failure of Grignard reagents to add in the normal way generally is because reactions by alternative paths occur more rapidly. If the Grignard reagent has a hydrogen on the carbon adjacent to the point of attachment of $\ce{-MgX}$ (i.e., a $\beta$ hydrogen), then reduction can occur, with the effect of adding $\ce{H_2}$ to the carbonyl group.
Side reactions - reduction
Furthermore, if the carbonyl compound has a hydrogen located on the carbon next to the carbonyl group, the Grignard reagent can behave as a base and remove this hydrogen as a proton. The result is that the compound becomes an enolate salt and $\ce{RMgX}$ becomes $\ce{RH}$.
Side reactions - enolization
Apparently, the complicating side reactions observed with $\ce{RMgX}$ are not nearly as important with $\ce{RLi}$. The reasons for this difference are not well understood.
Synthesis of Carboxylic Acids
The reaction of carbon dioxide with Grignard reagents initially gives a magnesium salt of a carboxylic acid, $\ce{RCO_2MgX}$:
This salt, which has a carbonyl group, in principle could add a second $\ce{RMgX}$. However, further addition is usually slow, and for most practical purposes the reaction stops at this stage. If the reaction can go further, the worst way to run it is by bubbling $\ce{CO_2}$ into the Grignard solution. This exposes the first-formed $\ce{RCO_2MgX}$ to excess $\ce{RMgX}$ and may lead to further addition reactions. The easy way to avoid this problem is to pour the $\ce{RMgX}$ solution onto powdered dry ice (solid $\ce{CO_2}$). Hydrolysis of the product (here a stronger acid than $\ce{NH_4Cl}$ is required) generates the carboxylic acid, $\ce{RCO_2H}$:
Synthesis of Ketones
Although organomagnesium compounds are not sufficiently reactive to add to carboxylate anions, alkyllithium compounds add quite well. A useful synthesis of methyl ketones involves the addition of methyllithium to the lithium salt of a carboxylic acid:
Other methods begin with acid chlorides or esters and attempt to add only one mole of $\ce{RMgX}$:
The disadvantage of using Grignard reagents for this purpose is that they add very rapidly to the ketone as it is formed. There are two ways in which this disadvantage can be minimized. First, one can add the Grignard solution to an excess of acid chloride solution (the so-called “inverse addition” procedure) to keep the concentration of $\ce{RMgX}$ in the reaction mixture low, and hope that the reaction will stop at the ketone stage. However, this device seldom works very well with acid chlorides. Better results can be obtained with $\ce{RMgX}$ and $\ce{R'CON(CH_3)_2}$. The second method is to use a less reactive organometallic compound - one that will react with $\ce{RCOCl}$ but not with $\ce{R_2C=O}$. One easy way to do this is to add cadmium chloride to the Grignard solution, whereby an organocadmium compound, $\ce{R_2Cd}$, is formed (cf. Section 14-10B, Method 3). In the presence of magnesium halides, $\ce{R_2Cd}$ reacts moderately rapidly with acid chlorides, but only slowly with ketones. The addition therefore can be arrested at the ketone stage:
$2 \ce{RMgCl} + \ce{CdCl_2} \rightarrow \ce{R_2Cd} + 2 \ce{MgCl_2}$
Alkylcopper compounds, $\ce{R-Cu}$, also are selective reagents that react with acid chlorides to give ketones, but do not add to esters, acids, aldehydes, or ketones. The $\ce{R-Cu}$ compounds can be prepared from$\ce{CuI}$ and the alkyllithium. With an excess of the alkyllithium, the alkylcopper is converted to $\ce{R_2CuLi}$:
1,4 Additions to Unsaturated Carbonyl Compounds
A conjugated alkenone, , can react with an organometallic reagent by a normal 1,2 addition across the carbonyl group, or by 1,4 addition to the conjugated system.
1,4 addition
On hydrolysis, the 1,4 adduct first yields the corresponding enol, but this is normally unstable and rearranges rapidly to the ketone. The final product therefore corresponds to addition of $\ce{R-H}$ across the carbon-carbon double bond:
Organomagnesium and organolithium compounds can add both 1,2 and 1,4 to alkenones, and the relative importance of each mode of addition depends on the structure of the reactants. This sort of dual behavior can be a nuisance in synthetic work because it leads to separation problems and low yields. Organocopper compounds are a great help in this situation because they show a very high selectivity for 1,4 addition and add to unsaturated ketones in excellent yield:
Oxygen, Sulfur, and Halogens
Grignard reagents react with oxygen, sulfur, and halogens to form substances containing $\ce{C-O}$, $\ce{C-S}$, and $\ce{C-X}$ bonds, respectively:
These reactions are not often important for synthesis because the products, $\ce{ROH}$, $\ce{RSH}$, and $\ce{RX}$, can be obtained more conveniently and directly from alkyl halides by $S_\text{N}1$ and $S_\text{N}2$ displacement reactions, as described in Chapter 8. However, when both $S_\text{N}1$ and $S_\text{N}2$ reactions are slow or otherwise impractical, as for neopentyl derivatives, the Grignard reactions can be very useful:
Also, oxygenation of a Grignard reagent at low temperatures provides an excellent method for the synthesis of hydroperoxides:
To prevent formation of excessive amounts of the alcohol, inverse addition is desirable (i.e., a solution of Grignard reagent is added to ether through which oxygen is bubbled rather than bubbling oxygen through a solution of the Grignard reagent).
Table 14-5: Methods of Preparation of Organic Halides
Table 14-6: Reactivities of Organohalogen Compounds in Displacement and Elimination Reactions
Table 14-7: Methods of Preparation of Organometallic Compounds
$^2$It is not possible to add $\ce{RH}$ to directly because $\Delta G^0$ generally is somewhat unfavorable [$+5 \: \text{kcal}$ for $\ce{CH_4} + \ce{(CH_3)_2C=O} \rightarrow \ce{(CH_3)_3COH}$]. How we get around this unfavorable equilibrium in practice provides an interesting example of how energy can be (and is) squandered to achieve some particular desired result; for example, the reaction $\ce{CH_3CH_3} + \ce{CH_3CHO} \rightarrow \ce{CH_3CH_3CH(CH_3)OH}$ has $\Delta H^0 = -12 \: \text{kcal}$ but $\Delta G^0 = + 0.5 \: \text{kcal}$. A possible sequence is
The overall result is the expenditure of $10 + 76 + 71 = 157 \: \text{kcal}$ to achieve a reaction that itself has $\Delta H^0 = -12 \: \text{kcal}$, but an unfavorable $\Delta G^0$. ($\ce{Li}$ is used in this example rather than $\ce{Mg}$ because the heat of formation of $\ce{C_2H_5MgBr}$ is not available.)
Contributors
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/14%3A_Organohalogen_and_Organometallic_Compounds/14.13%3A_Organomagnesium_and_Organolithium_Compounds_in_Synthesis.txt |
The physical, chemical and spectroscopic properties of alcohols are relative to it’s chemical structures. Alcohols are compounds of the general formula ROH, where R is any alkyl or substituted alkyl group. The simple ethers, ROR, do not have O-H bonds, and most of their reactions are limited to the substituent groups. Before turning to the specific chemistry of alcohols and ethers, we remind you that the naming alcohols and ethers is summarized in naming alcohols, phenols and Naming Ethers.
• 15.1: Prelude to Alcohols and Ethers
All carbohydrates and their derivatives, including nucleic acids, have hydroxyl groups. Some amino acids, most steroids, many terpenes, and plant pigments have hydroxyl groups. These substances serve many diverse purposes for the support and maintenance of life. The reactions involving the hydrogens of alcoholic OH groups are expected to be similar to those of water, HOH, the simplest hydroxylic compound. Alcohols, ROH, can be regarded in this respect as substitution products of water.
• 15.2: Physical Properties of Alcohols; Hydrogen Bonding
Comparison of the physical properties of alcohols with those of hydrocarbons of comparable molecular weight shows several striking differences, especially for those with just a few carbons. Alcohols are substantially less volatile, have higher melting points, and greater water solubility than the corresponding hydrocarbons, although the differences become progressively smaller as molecular weight increases.
• 15.3: Spectroscopic Properties of Alcohols
The hydrogen-oxygen bond of a hydroxyl group gives a characteristic absorption band in the infrared that is considerably influenced by hydrogen bonding. For example, in the vapor state (in which there is essentially no hydrogen bonding), ethanol gives an infrared spectrum with a fairly sharp absorption band at 3700 cm−1, owing to a free or unassociated hydroxyl group.
• 15.4: Preparation of Alcohols
Many of the common laboratory methods for the preparation of alcohols have been discussed in previous post or will be considered later; thus to avoid undue repetition we shall not consider them in detail at this time. Included among these methods are hydration and hydroboration, addition of hypohalous acids to alkenes , hydrolysis of alkyl halides and of allylic and benzylic halides, addition of Grignard reagents to carbonyl compounds, and the reduction of carbonyl compounds.
• 15.5: Chemical Reactions of Alcohols. Reactions Involving the O-H Bond
Several important chemical reactions of alcohols involve only the oxygen-hydrogen bond and leave the carbon-oxygen bond intact. An important example is salt formation with acids and bases. Alcohols, like water, are both weak bases and weak acids. The acid ionization constant of ethanol is slightly less than that of water.
• 15.6: Reactions Involving the C-O Bond of Alcohols
Alkyl halide formation from an alcohol and a hydrogen halide affords an important example of a reaction wherein the C−O bond of the alcohol is broken.
• 15.7: Oxidation of Alcohols
According to the scale of oxidation levels established for carbon, primary alcohols are at a lower oxidation level than either aldehydes or carboxylic acids. With suitable oxidizing agents, primary alcohols in fact can be oxidized first to aldehydes and then to carboxylic acids.
• 15.8: Polyhydric Alcohols
Polyhydric alcohols in which the hydroxyl groups are situated on different carbons are relatively stable, and, as we might expect for substances with multiple polar groups, they have high boiling points and considerable water solubility, but low solubility in nonpolar solvents.
• 15.9: Unsaturated Alcohols - Alkenols
The simplest unsaturated alcohols, ethenol (vinyl alcohol), is unstable with respect to ethanal and has never been isolated. Other simple unsaturated alkenols (enols) also rearrange to carbonyl compounds. However, ether and ester derivatives of enols are known and can be prepared by the addition of alcohols and carboxylic acids to alkynes. The esters are used to make many commercially important polymers.
• 15.10: Protection of Hydroxyl Groups
By now it should be apparent that hydroxyl groups are very reactive to many reagents. This is both an advantage and a disadvantage in synthesis. To avoid interference by hydroxyl groups, it often is necessary to protect (or mask) them by conversion to less reactive functions. In the case of alcohols the hydroxyl group may be protected by formation of an ether, an ester, or an acetal.
• 15.11: Types and Reactions of Simple Ethers
Substitution of the hydroxyl hydrogens of alcohols by hydrocarbon groups gives compounds known as ethers. These compounds may be classified further as open-chain, cyclic, saturated, unsaturated, aromatic, and so on.
• 15.12: Cyclic Ethers
Ring compounds containing nitrogen, oxygen, sulfur, or other elements as ring atoms generally are known as heterocyclic compounds, and the ring atoms other than carbon are the hetero atoms.
• 15.E: Alcohols and Ethers (Exercises)
These are the homework exercises to accompany Chapter 15 of the Textmap for Basic Principles of Organic Chemistry (Roberts and Caserio).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
15: Alcohols and Ethers
The physical, chemical and spectroscopic properties of alcohols are relative to it’s chemical structures. Alcohols are compounds of the general formula ROH, where R is any alkyl or substituted alkyl group. The hydroxyl group (OH groups) is the characteristic functional group of alcohols and is one of the most important functional groups of naturally occurring organic molecules. All carbohydrates and their derivatives, including nucleic acids, have hydroxyl groups. Some amino acids, most steroids, many terpenes, and plant pigments have hydroxyl groups. These substances serve many diverse purposes for the support and maintenance of life. One extreme example is the potent toxin tetrodotoxin, which is isolated from puffer fish and has obvious use for defense against predators. This compound has special biochemical interest, having six different hydroxylic functions arranged on a cagelike structure:
Stick figure of the tetrodotoxin molecule, \(C_{11}H_{17}N_3O_8\). (Public Domain; Ayacop).
On the more practical side, vast quantities of simple alcohols —methanol, ethanol, 2-propanol, 1-butanol —and many ethers are made from petroleum-derived hydrocarbons. These alcohols are widely used as solvents and as intermediates for the synthesis of more complex substances.
The reactions involving the hydrogens of alcoholic OH groups are expected to be similar to those of water, HOH, the simplest hydroxylic compound. Alcohols, ROH, can be regarded in this respect as substitution products of water. However, with alcohols we shall be interested not only in reactions that proceed at the O-H bond but also with processes that result in cleavage of the C-O bond, or changes in the organic group R.
The simple ethers, ROR, do not have O-H bonds, and most of their reactions are limited to the substituent groups. The chemistry of ethers, therefore, is less varied than that of alcohols. This fact is turned to advantage in the widespread use of ethers as solvents for a variety of organic reactions, as we already have seen for Grignard reagents. Nonetheless, cyclic ethers with small rings show enhanced reactivity because of ring strain and, for this reason, are valuable intermediates in organic synthesis.
Before turning to the specific chemistry of alcohols and ethers, we remind you that the naming alcohols and ethers is summarized in naming alcohols, phenols and Naming Ethers.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/15%3A_Alcohols_and_Ethers/15.01%3A_Prelude_to_Alcohols_and_Ethers.txt |
Comparison of the physical properties of alcohols with those of hydrocarbons of comparable molecular weight shows several striking differences, especially for those with just a few carbons. Alcohols are substantially less volatile, have higher melting points, and greater water solubility than the corresponding hydrocarbons (see Table 15-1), although the differences become progressively smaller as molecular weight increases.
Table 15-1: Comparison of Physical Properties of Alcohols and Hydrocarbons
The reason for these differences in physical properties is related to the high polarity of the hydroxyl group which, when substituted on a hydrocarbon chain, confers a measure of polar character to the molecule. As a result, there is a significant attraction of one molecule for another that is particularly pronounced in the solid and liquid states. This polar character leads to association of alcohol molecules through the rather positive hydrogen of one hydroxyl group with a correspondingly negative oxygen of another hydroxyl group:
This type of association is called “hydrogen bonding,” and, although the strengths of such bonds are much less than those of most conventional chemical bonds, they are still significant (about $5$ to $10 \: \text{kcal}$ per bond). Clearly then, the reason alcohols have higher boiling points than corresponding alkyl halides, ethers, or hydrocarbons is because, for the molecules to vaporize, additional energy is required to break the hydrogen bonds. Alternatively, association through hydrogen bonds may be regarded as effectively raising the molecular weight, thereby reducing volatility (also see Section 1-3).
The water solubility of the lower-molecular-weight alcohols is pronounced and is understood readily as the result of hydrogen bonding with water molecules:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
15.03: Spectroscopic Properties of Alcohols
The hydrogen-oxygen bond of a hydroxyl group gives a characteristic absorption band in the infrared but, as we may expect, this absorption is considerably influenced by hydrogen bonding. For example, in the vapor state (in which there is essentially no hydrogen bonding), ethanol gives an infrared spectrum with a fairly sharp absorption band at $3700 \: \text{cm}^{-1}$, owing to a free or unassociated hydroxyl group (Figure 15-2a). In contrast, this band is barely visible at $3640 \: \text{cm}^{-1}$ in the spectrum of a $10\%$ solution of ethanol in carbon tetrachloride (Figure 15-2b). However, there is a relatively broad band around $3350 \: \text{cm}^{-1}$, which is characteristic of hydrogen-bonded hydroxyl groups. The shift in frequency of about $300 \: \text{cm}^{-1}$ arises because hydrogen bonding weakens the $\ce{O-H}$ bond; its absorption frequency then will be lower. The association band is broad because the hydroxyl groups are associated in aggregates of various sizes and shapes. This produces a variety of different kinds of hydrogen bonds and therefore a spectrum of closely spaced $\ce{O-H}$ absorption frequencies.
In very dilute solutions of alcohols in nonpolar solvents, hydrogen bonding is minimized. However, as the concentration is increased, more and more of the molecules become associated and the intensity of the infrared absorption band due to associated hydroxyl groups increases at the expense of the free-hydroxyl band. Furthermore, the frequency of the association band is a measure of the strength of the hydrogen bond. The lower the frequency relative to the position of the free hydroxyl group, the stronger is the hydrogen bond. As we shall see in Chapter 18 the hydroxyl group in carboxylic acids $\left( \ce{RCO_2H} \right)$ forms stronger hydrogen bonds than alcohols and accordingly absorbs at lower frequencies (lower by about $400 \: \text{cm}^{-1}$, see Table 9-2).
The infrared spectra of certain 1,2-diols (glycols) are interesting in that they show absorption due to intramolecular hydrogen bonding. These usually are fairly sharp bands in the region $3450 \: \text{cm}^{-1}$ to $3570 \: \text{cm}^{-1}$, and, in contrast to bands due to intermolecular hydrogen bonding, they do not change in intensity with concentration. A typical example is afforded by cis-1,2-cyclopentanediol:
Besides the $\ce{O-H}$ stretching vibrations of alcohols, there is a bending $\ce{O-H}$ vibration normally observed in the region $1410$-$1260 \: \text{cm}^{-1}$. There also is a strong $\ce{C-O}$ stretching vibration between $1210 \: \text{cm}^{-1}$ and $1050 \: \text{cm}^{-1}$. Both these bands are sensitive to structure as indicated below:
The influence of hydrogen bonding on the proton nmr spectra of alcohols has been discussed previously (Section 9-10E). You may recall that the chemical shift of the $\ce{OH}$ proton is variable and depends on the extent of association through hydrogen bonding; generally, the stronger the association, the lower the field strength required to induce resonance. Alcohols also undergo intermolecular $\ce{OH}$ proton exchange, and the rate of this exchange can influence the line-shape of the $\ce{OH}$ resonance, the chemical shift, and the incidence of spin-spin splitting, as discussed in more detail in Sections 9-10E and 9-10I. Concerning the protons on carbon bearing the hydroxyl group, that is, , they are deshielded by the electron-attracting oxygen atom and accordingly have chemical shifts some $2.5$-$3.0 \: \text{ppm}$ to lower fields than alkyl protons.
Perhaps you are curious as to why absorptions are observed in the infrared spectrum of alcohols that correspond both to free and hydrogen-bonded hydroxyl groups, whereas only one $\ce{OH}$ resonance is observed in their proton nmr spectra. The explanation is that the lifetime of any molecule in either the free or the associated state is long enough to be detected by infrared absorption but much too short to be detected by nmr. Consequently, in the nmr one sees only the average $\ce{OH}$ resonance of the nonhydrogen-bonded and hydrogen-bonded species present. The situation here is very much like that observed for conformational equilibration (Section 9-10C).
The longest-wavelength ultraviolet absorption maxima of methanol and methoxymethane (dimethyl ether) are noted in Table 9-3. In each case the absorption maximum, which probably involves an $n \rightarrow \sigma^*$ transition, occurs about $184 \: \text{nm}$, well below the cut-off of the commonly available spectrometers.
The mass spectra of alcohols may not always show strong molecular ions. The reason is that the $\ce{M^+}$ ions readily fragment by $\alpha$ cleavage. The fragment ions are relatively stable and are the gaseous counterparts of protonated aldehydes and ketones:
Ethers also fragment by $\alpha$ cleavage:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
15.04: Preparation of Alcohols
Many of the common laboratory methods for the preparation of alcohols have been discussed in previous post or will be considered later; thus to avoid undue repetition we shall not consider them in detail at this time. Included among these methods are hydration (Section 10-3E) and hydroboration (Section 11-6D), addition of hypohalous acids to alkenes (Section 10-4B), $S_\text{N}1$ and $S_\text{N}2$ hydrolysis of alkyl halides (Sections 8-4 to 8-7) and of allylic and benzylic halides (Sections 14-3B and 14-3C), addition of Grignard reagents to carbonyl compounds (Section 14-12), and the reduction of carbonyl compounds (Sections 16-4E and 16-5). These methods are summarized in Table 15-2.
Some of the reactions we have mentioned are used for large-scale industrial production. For example, ethanol is made in quantity by the hydration of ethene, using an excess of steam under pressure at temperatures around $300^\text{o}$ in the presence of phosphoric acid:
A dilute solution of ethanol is obtained, which can be concentrated by distillation to a constant-boiling point mixture that contains $95.6\%$ ethanol by weight. Dehydration of the remaining few percent of water to give “absolute alcohol” is achieved either by chemical means or by distillation with benzene, which results in preferential separation of the water. Ethanol also is made in large quantities by fermentation, but this route is not competitive for industrial uses with the hydration of ethene. Isopropyl alcohol and tert-butyl alcohol also are manufactured by hydration of the corresponding alkenes.
Table 15-2: General Methods of Preparation of Alcohols
The industrial synthesis of methyl alcohol involves hydrogenation of carbon monoxide. Although this reaction has the favorable $\Delta H^0$ value of $-28.4 \: \text{kcal mol}^{-1}$, it requires high pressures and high temperatures and a suitable catalyst; excellent conversions are achieved using zinc oxide-chromic oxide as a catalyst:
Various methods of synthesis of other alcohols by reduction of carbonyl compounds will be discussed in Section 16-4E.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/15%3A_Alcohols_and_Ethers/15.02%3A_Physical_Properties_of_Alcohols_Hydrogen_Bonding.txt |
Acidic Properties
Several important chemical reactions of alcohols involve only the oxygen-hydrogen bond and leave the carbon-oxygen bond intact. An important example is salt formation with acids and bases. Alcohols, like water, are both weak bases and weak acids. The acid ionization constant $\left( K_\text{a} \right)$ of ethanol is about $10^{-18}$, slightly less than that of water. Ethanol can be converted to its conjugate base by the conjugate base of a weaker acid such as ammonia $\left( K_\text{a} \sim 10^{-35} \right)$, or hydrogen $\left( K_\text{a} \sim 10^{-38} \right)$. It is convenient to employ sodium metal or sodium hydride, which react vigorously but controllably with alcohols:
The order of acidity of various liquid alcohols generally is water $>$ primary $>$ secondary |(>\) tertiary $\ce{ROH}$. By this we mean that the equilibrium position for the proton-transfer reaction (Equation 15-1) lies more on the side of $\ce{ROH}$ and $\ce{OH}^\ominus$ as $\ce{R}$ is changed from primary to secondary to tertiary; therefore, tert-butyl alcohol is considered less acidic than ethanol:
$\ce{ROH} + \ce{OH}^\ominus \rightleftharpoons \ce{RO}^\ominus + \ce{HOH} \tag{15-1}$
However, in the gas phase the order of acidity is reversed, and the equilibrium position for Equation 15-1 lies increasingly on the side of $\ce{RO}^\ominus$ as $\ce{R}$ is changed from primary to secondary to tertiary. tert-Butyl alcohol is therefore more acidic than ethanol in the gas phase. This seeming contradiction appears more reasonable when one considers what effect solvation (or the lack of it) has on equilibria expressed by Equation 15-1. In solution, the larger anions of alcohols, known as alkoxide ions, probably are less well solvated than the smaller ions, because fewer solvent molecules can be accommodated around the negatively charged oxygen in the larger ions:
Acidity of alcohols therefore decreases as the size of the conjugate base increases. However, “naked” gaseous ions are more stable the larger the associated $\ce{R}$ groups, probably because the larger $\ce{R}$ groups can stabilize the charge on the oxygen atom better than the smaller $\ce{R}$ groups. They do this by polarization of their bonding electrons, and the bigger the group, the more polarizable it is. (Also see Section 11-8A, which deals with the somewhat similar situation encountered with respect to the relative acidities of ethyne and water.)
Basic Properties
Alcohols are bases similar in strength to water and accept protons from strong acids. An example is the reaction of methanol with hydrogen bromide to give methyloxonium bromide, which is analogous to the formation of hydroxonium bromide with hydrogen bromide and water:
Nucleophilic Properties - Ether Formation
Alkoxide ion formation is important as a means of generating a strong nucleophile that will readily form $\ce{C-O}$ bonds in $S_\text{N}2$ reactions. Thus ethanol reacts very slowly with methyl iodide to give methyl ethyl ether, but sodium ethoxide in ethanol solution reacts quite rapidly:
In fact, the reaction of alkoxides with alkyl halides or alkyl sulfates is an important general method for the preparation of ethers, and is known as the Williamson synthesis. Complications can occur because the increase of nucleophilicity associated with the conversion of an alcohol to an alkoxide ion always is accompanied by an even greater increase in eliminating power by the $E2$ mechanism. The reaction of an alkyl halide with alkoxide then may be one of elimination rather than substitution, depending on the temperature, the structure of the halide, and the alkoxide (Section 8-8). For example, if we wish to prepare isopropyl methyl ether, better yields would be obtained if we were to use methyl iodide and isopropoxide ion rather than isopropyl iodide and methoxide ion because of the prevalence of $E2$ elimination with the latter combination:
Potassium tert-butoxide is an excellent reagent to achieve $E2$ elimination because it is strongly basic and so bulky as to not undergo $S_\text{N}2$ reactions readily.
Nucleophilic Properties. Ester Formation
An ester may be thought of as a carboxylic acid in which the acidic proton has been replaced by some organic group, $\ce{R}$,
Esters can be prepared from carboxylic acids and alcohols provided an acidic catalyst is present,
or they can be prepared from acyl halides and alcohols or carboxylic anhydrides and alcohols:
These reactions generally can be expressed by the equation $+ \ce{ROH} \rightarrow$ $+ \ce{HX}$ which overall is a nucleophilic displacement of the $\ce{X}$ group by the nucleophile $\ce{ROH}$. However, the mechanism of displacement is quite different from the $S_\text{N}2$ displacements of alkyl derivatives, $\ce{R'X} + \ce{ROH} \rightarrow \ce{R'OR} + \ce{HX}$, and closely resembles the nucleophilic displacements of activated aryl halides (Section 14-6B) in being an addition-elimination process.
Acyl halides have a rather positive carbonyl carbon because of the polarization of the carbon-oxygen and carbon-halogen bonds. Addition of a nucleophilic group such as the oxygen of an alcohol occurs rather easily.
addition
The complex $1$ contains both an acidic group and a basic group , so that a proton shifts from one oxygen to the other to give $2$, which then rapidly loses hydrogen chloride by either an $E1$- or $E2$-type elimination to form the ester.
elimination
A similar but easily reversible reaction occurs between alcohols and carboxylic acids, which is slow in either direction in the absence of a strong mineral acid. The catalytic effect of acids, such as $\ce{H_2SO_4}$, $\ce{HCl}$, and $\ce{H_3PO_4}$ is produced by protonation of the carbonyl oxygen of the carboxylic acid, thereby giving $3$. This protonation greatly enhances the affinity of the carbonyl carbon for an electron pair on the oxygen of the alcohol (i.e., $3 \rightarrow 4$).
protonation step
addition step
Subsequently, a proton is transferred from the $\ce{OCH_3}$ to an $\ce{OH}$ group of $4$ to give $5$. This process converts the $\ce{OH}$ into a good leaving group $\left( \ce{H_2O} \right)$. When $\ce{H_2O}$ leaves, the product, $6$, is the conjugate acid of the ester. Transfer of a proton from $6$ to a base such as $\ce{H_2O}$ or $\ce{HSO_4^-}$ completes the reaction, giving the neutral ester and regenerating the acid catalyst.
proton transfer and elimination
Although a small amount of strong acid catalyst is essential in the preparation of esters from acids and alcohols, the amount of acid catalyst added must not be too large. The reason for the “too much of a good thing” behavior of the catalyst can be understood from the basic properties of alcohols (Section 15-4B). If too much acid is present, then too much of the alcohol is converted to the oxonium salt:
Clearly, formation of the methyloxonium ion can operate only to reduce the nucleophilic reactivity of methanol toward the carbonyl carbon of the carboxylic acid.
Another practical limitation of esterification reactions is steric hindrance. If either the acid or the alcohol participants possesses highly branched groups, the positions of equilibrium are less favorable and the rates of esterification are slow. In general, the ease of esterification for alcohols, $\ce{ROH}$, by the mechanism described is primary $\ce{R}$ $>$ secondary $\ce{R}$ $>$ tertiary $\ce{R}$ with a given carboxylic acid.
As mentioned, esterification is reversible, and with ethanol and ethanoic acid the equilibrium constant for the liquid phase is about 4 $\left( \Delta G^0 = -0.8 \: \text{kcal} \right)$ at room temperature, which corresponds to $66\%$ conversion to ester:
The reaction may be driven to completion by removing the ester or water or both as they are formed.
Nucleophilic Properties - Hemiacetal, Hemiketal, and Acetal Formation
The structural unit,
possesses both an alkoxyl $\left( \ce{OR} \right)$ and a hydroxyl $\left( \ce{OH} \right)$ group on the same carbon. This arrangement, although often unstable, is an important feature of carbohydrates such as glucose, fructose, and ribose. When the grouping is of the type
,
it is called a hemiacetal, and if it is
,
with no hydrogen attached to the carbon, it is called a hemiketal:
Each of these compounds has several other hydroxyl groups, but only one of them is a hemiacetal or hemiketal hydroxyl. Be sure you can identify which one.
The acetal function has two alkoxy $\left( \ce{OR} \right)$ groups and a hydrogen on the same carbon, , whereas the ketal function has the same structure but with no hydrogen on the carbon. These groupings also are found in carbohydrates and in carbohydrate derivatives, and are called glycosido functions (see Chapter 20).
For our present purposes, we are interested in the ways in which hemiacetals, acetals, hemiketals, and ketals are formed. Hemiacetals and hemiketals can be regarded as products of the addition of alcohols to the carbonyl groups of aldehydes and ketones. Thus methanol adds to ethanal to give a hemiacetal, 1 -methoxyethanol:
Acetals and ketals result from substitution of an alkoxy group for the $\ce{OH}$ group of a hemiacetal or hemiketal. Thus methanol can react with 1-methoxyethanol to form the acetal, 1,1-dimethoxyethane, and water:
The reactions of alcohols with aldehydes and ketones are related to the reactions of alcohols with acids (esterification) discussed in the preceding section. Both types involve addition of alcohols to carbonyl groups, and both are acid-catalyzed.
Acid catalysis of formation, like ester formation, depends on formation of the conjugate acid of the carbonyl compound. This is expected to enhance the positive (electrophilic) character of the carbonyl carbon so that the nucleophilic alcohol can add readily to it:
The hemiacetal can react further, also with the aid of an acidic catalyst. Addition of a proton can occur in two ways, to give $7$ or $8$:
The first of these, $7$, has $\ce{CH_3OH}$ as a leaving group and reverts back to the conjugate acid of ethanal. This is the reverse of acid-catalyzed hemiacetal formation:
The second of these,$8$, has $\ce{H_2O}$ as a leaving group and can form a new entity, the methoxyethyl cation, $9$:
The ion $9$ resembles and can be expected to behave similarly by adding a second molecule of alcohol to the electrophilic carbon. The product, $10$, is then the conjugate acid of the acetal and loses a proton to give the acetal:
Formation of hemiacetals and acetals, as well as of hemiketals and ketals, is reversible under acidic conditions, as we already have noted for acid-catalyzed esterification. The reverse reaction is hydrolysis and the equilibrium for this reaction can be made favorable by having an excess of water present:
The position of equilibrium in acetal and hemiacetal formation is rather sensitive to steric hindrance. Large groups in either the aldehyde or the alcohol tend to make the reaction less favorable. Table 15-3 shows some typical conversions in acetal formation when 1 mole of aldehyde is allowed to come to equilibrium with 5 moles of alcohol. For ketones, the equilibria are still less favorable than for aldehydes, and to obtain reasonable conversion the water must be removed as it is formed.
Table 15-3: Conversion of Aldehydes to Acetals with Various Alcohols (1 Mole of Aldehyde to 5 Moles of Alcohol)
Percent Conversion to Acetal
Aldehyde Ethanol Cycloehexane 2-Propoanol tert-Butyl Alcohol
$CH_3CHO$ 78% 56% 43% 23%
$(CH_3)_2CHCHO$ 71% - 23% -
$(CH_3)_3CCHO$ 56% 16% 11% -
$C_6H_5CHO$ 39% 235 13% -
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/15%3A_Alcohols_and_Ethers/15.05%3A_Chemical_Reactions_of_Alcohols._Reactions_Involving_the_O-H_Bond.txt |
Electrophilic Properties of Alcohols
Alkyl halide formation from an alcohol and a hydrogen halide affords an important example of a reaction wherein the $\ce{C-O}$ bond of the alcohol is broken:
The reaction is reversible and the favored direction depends on the water concentration. Primary bromides often are prepared best by passing dry hydrogen bromide into the alcohol heated to just slightly below its boiling point.
Halide formation proceeds at a useful rate only in the presence of strong acid, which can be furnished by excess hydrogen bromide or, usually and more economically, by sulfuric acid. The alcohol accepts a proton from the acid to give an alkyloxonium ion, which is more reactive in subsequent displacement with bromide ion than the alcohol (by either $S_\text{N}2$ or $S_\text{N}1$ mechanisms) because $\ce{H_2O}$ is a better leaving group than $\ce{OH}^\ominus$:
or
Whether the displacement reaction is an $S_\text{N}1$ or $S_\text{N}2$ process depends on the structure of the alcohol. In general, the primary alcohols are considered to react by $S_\text{N}2$ and the secondary and tertiary alcohols by $S_\text{N}1$ mechanisms.
Hydrogen chloride is less reactive than hydrogen bromide toward primary alcohols, and a catalyst (zinc chloride) may be required. A solution of zinc chloride in concentrated hydrochloric acid (Lucas reagent) is a convenient reagent to differentiate between primary, secondary, and tertiary alcohols with less than eight or so carbons. Tertiary alcohols react very rapidly to give an insoluble layer of alkyl chloride at room temperature. Secondary alcohols react in several minutes, whereas primary alcohols form chlorides only on heating. The order of reactivity is typical of $S_\text{N}1$ reactions. Zinc chloride probably assists in the breaking of the $\ce{C-O}$ bond of the alcohol much as silver ion aids ionization of $\ce{RCl}$ (Section 8-7D):
Thionyl chloride, $\ce{O=SCl_2}$, is useful for the preparation of alkyl chlorides, especially when the use of strongly acidic reagents, such as zinc chloride and hydrochloric acid, is undesirable. Thionyl chloride can be regarded as the acid chloride of sulfurous acid, $\ce{O=S(OH)_2}$, and like most acid chlorides the halogen is displaced readily by alcohols. Addition of 1 mole of an alcohol to 1 mole of thionyl chloride gives an unstable alkyl chlorosulfite, which generally decomposes on mild heating to yield the alkyl chloride and sulfur dioxide:
Chlorides can be prepared in this way from primary and secondary, but not tertiary, alcohols. In practice, an equivalent of a weak base, such as pyridine (azabenzene), is added to neutralize the hydrogen chloride that is formed. If the acid is not removed, undesirable degradation, elimination, and rearrangement reactions may occur.
The thionyl chloride reaction apparently can proceed from the alkyl chlorosulfite stage by more than one mechanism: an ionic $S_\text{N}2$ chain reaction with chloride ion,
or an $S_\text{N}1$-like ionization and collapse of the resulting $\ce{R}^\oplus \ce{Cl}^\ominus$ ion pair to give $\ce{RCl}$:
or
Obviously, the greater the $S_\text{N}2$ reactivity associated with the better the $S_\text{N}2$ reaction will go and, conversely, if $\ce{R}^\oplus$ is formed easily from the $S_\text{N}1$ reaction is likely to be favored.
Other halides that are useful in converting alcohols to alkyl halides are $\ce{PCl_5}$, $\ce{PCl_3}$, $\ce{PBr_3}$, and $\ce{PI_3}$, which are acid halides of phosphorus oxyacids. As with thionyl chloride, a weak base often is used to facilitate the reaction. The base acts to neutralize the acid formed, and also to generate bromide ion for $S_\text{N}$ reactions:
Sulfate and Sulfonate Esters
It is possible to prefer esters of sulfuric acid by the reaction of an alcohol with the acid:
The reaction is closely related to alkyl halide formation under strongly acidic conditions, whereby conversion of the alcohol to an oxonium salt is a first step:
Conversion of the oxonium hydrogen sulfate to the ester probably proceeds by an $S_\text{N}2$ mechanism with primary alcohols and an $S_\text{N}1$ mechanism with tertiary alcohols:
An alternative mechanism, which operates either in $100\%$, or in fuming sulfuric acid (which contains dissolved $\ce{SO_3}$), is addition of sulfur trioxide to the $\ce{OH}$ group:
The sodium salts of alkyl hydrogen sulfate esters have useful properties as detergents if the alkyl group is large, $\ce{C_{12}}$ or so:
The mechanism of detergent action will be considered in more detail in Chapter 18.
In principle, dialkyl sulfates could be formed by an $S_\text{N}2$ reaction between an alkyloxonium salt and an alkyl sulfate ion:
Indeed, if methanol is heated with fuming sulfuric acid, dimethyl sulfate, $\ce{CH_3O(SO_2)OCH_3}$, is obtained; but other alcohols are better converted to dialkyl sulfates by oxidation of the corresponding dialkyl sulfites formed by the reaction of 1 mole of thionyl chloride $\left( \ce{SOCl_2} \right)$ with 2 moles of the alcohol:
The reason that dialkyl sulfates seldom are prepared by direct reaction of the alcohol with $\ce{H_2SO_4}$ is that the mono esters react rapidly on heating to eliminate sulfuric acid and form alkenes, as explained in Section 15-5C.
Sulfonic acids, $\ce{R-SO_2-OH}$ or $\ce{Ar-SO_2-OH}$, are oxyacids of sulfur that resemble sulfuric acid, $\ce{HO-SO_2-OH}$, but in which sulfur is in a lower oxidation state.
Sulfonate esters are useful intermediates in displacement reactions (Section 8-7C) and provide a route for the conversion of an alcohol, $\ce{ROH}$, to $\ce{RX}$ by the sequence:
Sulfonate esters usually are prepared through treatment of the alcohol with the acid chloride (sulfonyl chloride) in the presence of pyridine (azabenzene):
Dehydration of Alcohols with Strong Acids
In the reaction of an alcohol with hot concentrated sulfuric acid, the alcohol is dehydrated to an alkene:
This is the reverse of acid-catalyzed hydration of alkenes discussed previously (Section 10-3E) and goes to completion if the alkene is allowed to distill out of the reaction mixture as it is formed. One mechanism of dehydration involves proton transfer from sulfuric acid to the alcohol, followed by an $E2$ reaction of hydrogen sulfate ion or water with the oxonium salt of the alcohol:
Alternatively, the alkyl hydrogen sulfate could be formed and eliminate sulfuric acid by an $E2$ reaction:
At lower temperatures the oxonium salt or the alkyl hydrogen sulfate may react by an $S_\text{N}$ displacement mechanism with excess alcohol in the reaction mixture, thereby forming a dialkyl ether. Although each step in the reaction is reversible, ether formation can be enhanced by distilling away the ether as fast as it forms. Diethyl ether is made commercially by this process:
Most alcohols also will dehydrate at fairly high temperatures in the presence of solid catalysts such as silica gel or aluminum oxide to give alkenes or ethers. The behavior of ethanol is reasonably typical of primary alcohols and is summarized in the following equations:
$\ce{C-O}$ Bond Cleavage of Tertiary Alcohols
Tertiary alcohols react with sulfuric acid at much lower temperatures than do most primary or secondary alcohols. The reactions typically are $S_\text{N}1$ and $E1$ by way of a tertiary carbocation, as shown here for tert-butyl alcohol and sulfuric acid:
2-Methylpropene can be removed from the reaction mixture by distillation and easily is made the principal product by appropriate adjustment of the reaction conditions. If the 2-methylpropene is not removed as it is formed, polymer and oxidation products become important. Sulfuric acid often is an unduly strenuous reagent for dehydration of tertiary alcohols. Potassium hydrogen sulfate, copper sulfate, iodine, phosphoric acid, or phosphorus pentoxide may give better results by causing less polymerization and less oxidative degradation which, with sulfuric acid, results in the formation of sulfur dioxide.
The $S_\text{N}1$-$E1$ behavior of tertiary alcohols in strong acids can be used to advantage in the preparation of tert-butyl ethers. If, for example, a mixture of tert-butyl alcohol and methanol is heated in the presence of sulfuric acid, the tertiary alcohol reacts rapidly but reversibly to produce 2-methylpropene by way of the tert-butyl cation. This cation can be trapped by the methanol to form tert-butyl methyl ether. High yields of ethers can be obtained in this way:
Carbocation Rearrangements
Rearrangement of the alkyl groups of alcohols is very common in dehydration, particularly in the presence of strong acids, which are conducive to carbocation formation. Typical examples showing both methyl and hydrogen migration follow:
The key step in each such rearrangement is isomerization of a carbocation, as discussed in Section 8-9B. Under kinetic control, the final products always correspond to rearrangement of a less stable carbocation to a more stable carbocation. (Thermodynamic control may lead to quite different results, Section 10-4A.)
In the dehydration of 3,3-dimethyl-2-butanol, a secondary carbocation is formed initially, which rearranges to a tertiary carbocation when a neighboring methyl group with its bonding electron pair migrates to the positive carbon. The charge is thereby transferred to the tertiary carbon:
Phosphate Esters
Phosphoric acid $\left( \ce{H_3PO_4} \right)$ often is used in place of sulfuric acid to dehydrate alcohols. This is because phosphoric acid is less destructive; it is both a weaker acid and a less powerful oxidizing agent than sulfuric acid. Dehydration probably proceeds by mechanisms similar to those described for sulfuric acid (Section 15-5C) and very likely through intermediate formation of a phosphate ester:
The ester can eliminate $\ce{H_3PO_4}$, as sulfate esters eliminate $\ce{H_2SO_4}$, to give alkenes:
The chemistry of phosphate esters is more complicated than that of sulfate esters because it is possible to have one, two, or three alkyl groups substituted for the acidic hydrogens of phosphoric acid:
Also, phosphoric acid forms an extensive series of anhydrides (with $\ce{P-O-P}$ bonds), which further diversify the number and kind of phosphate esters. The most important phosphate esters are derivatives of mono-, di-, and triphosphoric acid (sometimes classified as ortho-, pyro-, and meta-phosphoric acids, respectively):
The equilibrium between the esters of any of these phosphoric acids and water favors hydrolysis:
However, phosphate esters are slow to hydrolyze in water (unless a catalyst is present). The difference in kinetic and thermodynamic stability of phosphate esters toward hydrolysis is used to great effect in biological systems.
Of particular importance is the conversion of much of the energy that results from photosynthesis, or from the oxidation of fats, carbohydrates, and proteins in cells into formation of phosphate ester bonds $\left( \ce{C-O-P} \right)$ or phosphate anhydride bonds $\left( \ce{P-O-P} \right)$. The energy so stored is used in other reactions, the net result of which is hydrolysis:
The substance that is the immediate source of energy for many biological reactions is adenosine triphosphate (ATP). Although this is a rather large and complex molecule, the business end for the purpose of this discussion is the triphosphate group. Hydrolysis of this group can occur to give adenosine diphosphate (ADP), adenosine monophosphate (AMP), or adenosine itself:
(The phosphate groups are repesented here as the major ionized form present at pH $\cong 7$ in solutions of ATP.)
All of these hydrolysis reactions are energetically favorable $\left( \Delta G^0 < 0 \right)$, but they do not occur directly because ATP reacts slowly with water. However, hydrolysis of ATP is the indirect result of other reactions in which it participates. For example, as we showed in Section 15-4D, equilibrium for the direct formation of an ester from a carboxylic acid and an alcohol in the liquid phase is not very favorable (Equation 15-2). However, if esterification can be coupled with ATP hydrolysis (Equation 15-3), the overall reaction (Equation 15-4) becomes much more favorable thermodynamically than is direct esterification.
The ATP hydrolysis could be coupled to esterification (or other reactions) in a number of ways. The simplest would be to have the ATP convert one of the participants to a more reactive intermediate. For esterification, the reactive intermediate is an acyl derivative of AMP formed by the displacement of diphosphate from ATP:
The acyl AMP is like an acyl chloride, $\ce{RCOCl}$, in having a leaving group (AMP) that can be displaced with an alcohol:
The net result of the sequence in Equations 15-5 and 15-6 is esterification in accord with Equation 15-4. It is not a catalyzed esterification because in the process one molecule of ATP is converted to AMP and diphosphate for each molecule of ester formed. The AMP has to be reconverted to ATP to participate again. These reactions are carried on by cells under the catalytic influence of enzymes. The adenosine part of the molecule is critical for the specificity of action by these enzymes. Just how these enzymes function obviously is of great interest and importance.
If the role of phosphate esters, such as ATP, in carrying out reactions such as esterification in aqueous media under the influence of enzymes in cells is not clear to you, think about how you would try to carry out an esterification of ethanol in dilute water solution. Remember that, with water in great excess, the equilibrium will be quite unfavorable for the esterification reaction of Equation 15-2. You might consider adding $\ce{CH_3COCl}$, for which the equilibrium for ester formation is much more favorable (Section 15-4D). However, $\ce{CH_3COCl}$ reacts violently with water to form $\ce{CH_3CO_2H}$, and this reaction destroys the $\ce{CH_3COCl}$ before it has much chance to react with ethanol to give the ester. Clearly, what you would need is a reagent that will convert $\ce{CH_3CO_2H}$ into something that will react with ethanol in water to give the ester with a favorable equilibrium constant and yet not react very fast with water. The phosphate esters provide this function in biochemical systems by being quite unreactive to water but able to react with carboxylic acids under the influence of enzymes to give acyl phosphates. These acyl phosphates then can react with alcohols under the influence of other enzymes to form esters in the presence of water.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/15%3A_Alcohols_and_Ethers/15.06%3A_Reactions_Involving_the_C-O_Bond_of_Alcohols.txt |
According to the scale of oxidation levels established for carbon (see Table 11-1), primary alcohols $\left( \ce{RCH_2OH} \right)$ are at a lower oxidation level than either aldehydes $\left( \ce{RCHO} \right)$ or carboxylic acids $\left( \ce{RCO_2H} \right)$. With suitable oxidizing agents, primary alcohols in fact can be oxidized first to aldehydes and then to carboxylic acids.
primary alcohols
Unlike the reactions discussed previously in this chapter, oxidation of alcohols involves the alkyl portion of the molecule, or more specifically, the $\ce{C-H}$ bonds of the hydroxyl-bearing carbon (the $\alpha$ carbon). Secondary alcohols, which have only one such $\ce{C-H}$ bond, are oxidized to ketones, whereas tertiary alcohols, which have no $\ce{C-H}$ bonds to the hydroxylic carbon, are oxidized only with accompanying degradation into smaller fragments by cleavage of carbon-carbon bonds.
secondary alcohols
tertiary alcohols
Industrial Oxidation of Alcohols
Conversion of ethanol to ethanal is carried out on a commercial scale by passing gaseous ethanol over a copper catalyst at $300^\text{o}$:
At room temperature this reaction is endothermic with an equilibrium constant of about $10^{22}$. At $300^\text{o}$ conversions of $20\%$-$50\%$ per pass can be realized and, by recycling the unreacted alcohol, the yield can be greater than $90\%$.
Another commercial process uses a silver catalyst and oxygen to combine with the hydrogen, which makes the net reaction substantially exothermic:
In effect, this is a partial combustion reaction and requires very careful control to prevent overoxidation. In fact, by modifying the reaction conditions (alcohol-to-oxygen ratio, temperature, pressure, and reaction time), the oxidation proceeds smoothly to ethanoic acid:
Reactions of this type are particularly suitable as industrial processes because they generally can be run in continuous-flow reactors, and can utilize a cheap oxidizing agent, usually supplied directly as air.
Oxidation of Alcohols
Chromic Acid Oxidation
Laboratory oxidation of alcohols most often is carried out with chromic acid $\left( \ce{H_2CrO_4} \right)$, which usually is prepared as required from chromic oxide $\left( \ce{CrO_3} \right)$ or from sodium dichromate $\left( \ce{Na_2Cr_2O_7} \right)$ in combination with sulfuric acid. Ethanoic (acetic) acid is a useful solvent for such reactions:
The mechanism of the chromic acid oxidation of 2-propanol to 2-propanone (acetone) has been investigated very thoroughly. It is a highly interesting reaction in that it reveals how changes of oxidation state can occur in a reaction involving a typical inorganic and a typical organic compound. The initial step is reversible formation of an isopropyl ester of chromic acid. This ester is unstable and is not isolated:
The subsequent step is the slowest in the sequence and appears to involve attack of a base (water) at the alpha hydrogen of the chromate ester concurrent with elimination of the $\ce{HCrO_3^-}$ group. There is an obvious analogy between this step and an $E2$ reaction (Section 8-8A):
The transformation of chromic acid $\left( \ce{H_2CrO_4} \right)$ to $\ce{H_2CrO_3}$ amounts to the reduction of chromium from an oxidation state of $+6$ to $+4$. Disproportionation of Cr(IV) occurs rapidly to give compounds of Cr(III) and Cr(VI):
or
The $E2$ character of the ketone-forming step has been demonstrated in two ways. First, the rate of decomposition of isopropyl hydrogen chromate to 2-propanone and $\ce{H_2CrO_3}$ is strongly accelerated by efficient proton-removing substances. Second, the hydrogen on the $\alpha$ carbon clearly is removed in a slow reaction because the overall oxidation rate is diminished sevenfold by having a deuterium in place of the $\alpha$ hydrogen. No significant slowing of oxidation is noted for 2-propanol having deuterium in the methyl groups:
Carbon-deuterium bonds normally are broken more slowly than carbon-hydrogen bonds. This so-called kinetic isotope effect provides a general method for determining whether particular carbon-hydrogen bonds are broken in slow reaction steps.
Primary alcohols are oxidized by chromic acid in sulfuric acid solution to aldehydes, but to stop the reaction at the aldehyde stage, it usually is necessary to remove the aldehyde from the reaction mixture as it forms. This can be done by distillation if the aldehyde is reasonably volatile:
Unsaturated alcohols can be oxidized to unsaturated ketones by chromic acid, because chromic acid usually attacks double bonds relatively slowly:
However, complications are to be expected when the double bond of an unsaturated alcohol is particularly reactive or when the alcohol rearranges readily under strongly acidic conditions. It is possible to avoid the use of strong acid through the combination of chromic oxide with the weak base azabenzene (pyridine). A crystalline solid of composition $\ce{(C_5H_5N)_2} \cdot \ce{CrO_3}$ is formed when $\ce{CrO_3}$ is added to excess pyridine at low temperatures. (Addition of pyridine to $\ce{CrO_3}$ is likely to give an uncontrollable reaction resulting in a fire.)
The pyridine-$\ce{CrO_3}$ reagent is soluble in chlorinated solvents such as dichloromethane, and the resulting solutions rapidly oxidize to at ordinary temperatures:
The yields usually are good, partly because the absence of strong acid minimizes degradation and rearrangement, and partly because the product can be isolated easily. The inorganic products are insoluble and can be separated by filtration, thereby leaving the oxidized product in dichloromethane from which it can be easily recovered.
Permanganate Oxidation
Permanganate ion, $\ce{MnO_4^-}$, oxidizes both primary and secondary alcohols in either basic or acidic solution. With primary alcohols the product normally is the carboxylic acid because the intermediate aldehyde is oxidized rapidly by permanganate:
Oxidation under basic conditions evidently involves the alkoxide ion rather than the neutral alcohol. The oxidizing agent, $\ce{MnO_4^-}$, abstracts the alpha hydrogen from the alkoxide ion either as an atom (one-electron transfer) or as hydride, $\ce{H}^\ominus$ (two-electron transfer). The steps for the two-electron sequence are:
In the second step, permanganate ion is reduced from Mn(VII) to Mn(V). However, the stable oxidation states of manganese are $+2$, $+4$, and $+7$; thus the Mn(V) ion formed disproportionates to Mn(VII) and Mn(IV). The normal manganese end product from oxidations in basic solution is manganese dioxide, $\ce{MnO_2}$, in which $\ce{Mn}$ has an oxidation state of $+4$ corresponding to Mn(IV).
In Section 11-7C we described the use of permanganate for the oxidation of alkenes to 1,2-diols. How is it possible to control this reaction so that it will stop at the diol stage when permanganate also can oxidize to ? Overoxidation with permanganate is always a problem, but the relative reaction rates are very much a function of the pH of the reaction mixture and, in basic solution, potassium permanganate oxidizes unsaturated groups more rapidly than it oxidizes alcohols:
Biological Oxidations
There are many biological oxidations that convert a primary or secondary alcohol to a carbonyl compound. These reactions cannot possibly involve the extreme pH conditions and vigorous inorganic oxidants used in typical laboratory oxidations. Rather, they occur at nearly neutral pH values and they all require enzymes as catalysts, which for these reactions usually are called dehydrogenases.
An important group of biological oxidizing agents includes the pyridine nucleotides, of which nicotinamide adenine dinucleotide ($\ce{NAD}^\oplus$, $13$) is an example:
This very complex molecule functions to accept hydride $\left( \ce{H}^\ominus \right)$ or the equivalent $\left( \ce{H}^\oplus + 2 \ce{e^-} \right)$ from the $\alpha$ carbon of an alcohol:
The reduced form of $\ce{NAD}^\oplus$ is abbreviated as $\ce{NADH}$ and the $\ce{H}^\ominus$ is added at the 4-position of the pyridine ring:
Some examples follow that illustrate the remarkable specificity of this kind of redox system. One of the last steps in the metabolic breakdown of glucose (glycolysis; Section 20-10A) is the reduction of 2-oxopropanoic (pyruvic) acid to $L$-2-hydroxypropanoic (lactic) acid. The reverse process is oxidation of $L$-lactic acid. The enzyme lactic acid dehydrogenase catalyses this reaction, and it functions only with the $L$-enantiomer of lactic acid:
A second example, provided by one of the steps in metabolism by way of the Krebs citric acid cycle (see Section 20-10B), is the oxidation of $L$-2-hydroxy-butanedioic ($L$-malic) acid to 2-oxobutanedioic (oxaloacetic) acid. This enzyme functions only with $L$-malic acid:
All of these reactions release energy. In biological oxidations much of the energy is utilized to form ATP from ADP and inorganic phosphate (Section 15-5F). That is to say, electron-transfer reactions are coupled with ATP formation. The overall process is called oxidative phosphorylation.
Another important oxidizing agent in biological systems is flavin adenine dinucleotide, $\ce{FAD}$. Like $\ce{NAD}^\oplus$, it is a two-electron acceptor, but unlike $\ce{NAD}^\oplus$, it accepts two electrons as $2 \ce{H} \cdot$ rather than as $\ce{H}^\ominus$. The reduced form, $\ce{FADH_2}$, has the hydrogens at ring nitrogens:
Contributors
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/15%3A_Alcohols_and_Ethers/15.07%3A_Oxidation_of_Alcohols.txt |
The simplest example of an alcohol with more than one hydroxyl group is methanediol or methylene glycol, $\ce{HOCH_2OH}$. The term “glycol” indicates a diol, which is a substance with two alcoholic hydroxyl groups. Methylene glycol is reasonably stable in water solution, but attempts to isolate it lead only to its dehydration product, methanal (formaldehyde):
This behavior is rather typical of gem-diols (gem $=$ geminal, that is, with both hydroxyl groups on the same carbon atom). The few gem-diols of this kind that can be isolated are those that carry strongly electron-attracting substituents such as the following:
Polyhydric alcohols in which the hydroxyl groups are situated on different carbons are relatively stable, and, as we might expect for substances with multiple polar groups, they have high boiling points and considerable water solubility, but low solubility in nonpolar solvents:
1,2-Diols are prepared from alkenes by oxidation with reagents such as osmium tetroxide, potassium permanganate, or hydrogen peroxide (Section 11-7C). However, ethylene glycol is made on a commercial scale from oxacyclopropane, which in turn is made by air oxidation of ethene at high temperatures over a silver oxide catalyst (Section 11-7D).
Ethylene glycol has important commercial uses. It is an excellent permanent antifreeze for automotive cooling systems because it is miscible with water in all proportions and a $50\%$ solution freezes at $-34^\text{o}$ $\left( -29^\text{o} \text{F} \right)$. It also is used as a solvent and as an intermediate in the production of polymers (polyesters) and other products (Chapter 29).
The trihydric alcohol, 1,2,3-propanetriol (glycerol), is a nontoxic, water-soluble, viscous, hygroscopic liquid that is used widely as a humectant (moistening agent). It is an important component of many food, cosmetic, and pharmaceutical preparations. At one time, glycerol was obtained on a commercial scale only as a by-product of soap manufacture through hydrolysis of fats, which are glyceryl esters of long-chain alkanoic acids. The major present source is by synthesis from propene (Section 14-3A). The trinitrate ester of glycerol (nitroglycerin) is an important but shock-sensitive explosive:
Dynamite is a much safer and more controllable explosive, and is made by absorbing nitroglycerin in porous material such as sawdust or diatomaceous earth. Dynamite has largely been replaced by cheaper explosives containing ammonium nitrate as the principal ingredient.
Glycerol, as a constituent of fats and lipids, plays an important role in animal metabolism.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
15.09: Unsaturated Alcohols - Alkenols
The simplest unsaturated alcohols, ethenol (vinyl alcohol), is unstable with respect to ethanal and has never been isolated (see Sections 10-5A and 13-5B):
Other simple unsaturated alkenols (enols) also rearrange to carbonyl compounds. However, ether and ester derivatives of enols are known and can be prepared by the addition of alcohols and carboxylic acids to alkynes. The esters are used to make many commercially important polymers (Chapter 29):
The enol of 2-oxopropanoic acid (pyruvic acid) is of special biological interest because the phosphate ester of this compound is, like ATP (Section 15-5F), a reservoir of chemical energy that can be utilized by coupling its hydrolysis $\left( \Delta G^0 = -13 \: \text{kcal} \right)$ to thermodynamically less favorable reactions:
In fact, the ester can be utilized to synthesize ATP from ADP; that is, it is a phosphorylating agent, and a more powerful one than ATP:
Acidity of Enols
Enols usually are unstable and are considerably more acidic than saturated alcohols. This means that the conjugate bases of the enols (the enolate anions) are more stable relative to the enols themselves than are alkoxide ions relative to alcohols. (Enolate anions are important reagents in the chemistry of carbonyl compounds and will be discussed in detail in Chapter 17.)
The important factor here is delocalization of the negative charge on oxygen of enolate anions, as represented by the valence-bond structures $14a$ and $14b$:
Because acidity depends on the difference in energy of the acid and its conjugate base, we must be sure that the stabilization of the enolate anion by electron delocalization represented by $14a$ and $14b$ is greater than the analogous stabilization of the neutral enol represented by $15a$ and $15b$:
The rules for evaluating valence-bond structures (Section 6-5B) tell us that the stabilization will be greatest when there are two or more nearly equivalent low-energy electron-pairing schemes. Inspection of $14a$ and $14b$ suggests that they will be more nearly equivalent than $15a$ and $15b$ because, although $14b$ and $15b$ have a negative charge on the carbon, in $15b$ the oxygen has a positive charge. Another way of putting it is that $15b$ represents an electron-pairing scheme with a charge separation, which intuitively is of higher energy than $15a$ with no charge separation. Structures corresponding to $14b$ and $15b$ are not possible for saturated alkanols or their anions, hence we can see that enols should be more acidic than alcohols.
Ascorbic acid (Vitamin C) is an example of a stable and quite acidic enol, or rather an enediol. It is a di-acid with p$K_\text{a}$ values of 4.17 and 11.57:
Other important examples of stable enol-type compounds are the aromatic alcohols, or phenols. The $K_\text{a}$’s of these compounds are about $10^{-10}$, some $10^8$ times larger than the $K_\text{a}$’s for alcohols.
The chemistry of these compounds, including their stability as enols, is discussed in Chapter 26.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/15%3A_Alcohols_and_Ethers/15.08%3A_Polyhydric_Alcohols.txt |
By now it should be apparent that hydroxyl groups are very reactive to many reagents. This is both an advantage and a disadvantage in synthesis. To avoid interference by hydroxyl groups, it often is necessary to protect (or mask) them by conversion to less reactive functions. The general principles of how functional groups are protected were outlined and illustrated in Section 13-9. In the case of alcohols the hydroxyl group may be protected by formation of an ether, an ester, or an acetal.
Ether Formation
A good protecting group is one that does everything you want it to do when you want it to. It must be easily put into place, stable to the reagents from which protection is required, and easily removed when desired. For this reason simple ethers such as methyl or ethyl ethers usually are not suitable protecting groups because they cannot be removed except under rather drastic conditions (Section 15-10).
More suitable ethers are phenylmethyl and trimethylsilyl ethers:
Both of these ethers are prepared easily by nucleophilic displacements (Equations 15-7 and 15-8) and can be converted back to the parent alcohol under mild conditions, by catalytic hydrogenation for phenylmethyl ethers (Equation 15-9), or by mild acid hydrolysis for trimethylsilyl ethers (Equation 15-10):
Ester Formation
Esters are formed from the alcohol and acyl halide, anhydride, or acid (Section 15-4D). The alcohol can be regenerated easily by either acid or base hydrolysis of the ester:
Acetal Formation
We have seen that alcohols can be converted reversibly to acetals under acidic conditions (Section 15-4E). The acetal function is a very suitable protecting group for alcohols under basic conditions, but is not useful under acidic conditions because acetals are not stable to acids:
An excellent reagent to form acetals is the unsaturated cyclic ether, \(16\). This ether adds alcohols in the presence of an acid catalyst to give the acetal \(17\):
The 3-oxacyclohexene (dihydropyran) protecting group can be removed readily by treating the acetal, \(17\), with aqueous acid:
An example of the use of \(16\) to protect an \(\ce{OH}\) function is given in Section 13-10.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
15.11: Types and Reactions of Simple Ethers
Substitution of the hydroxyl hydrogens of alcohols by hydrocarbon groups gives compounds known as ethers. These compounds may be classified further as open-chain, cyclic, saturated, unsaturated, aromatic, and so on. For the naming of ethers, see Sections 7-3 and 15-11A.
The most generally useful methods of preparing ethers already have been discussed (Sections 8-7C, 8-7E, 15-4C, and 15-5C). These and some additional special procedures are summarized in Table 15-4.
Table 15-4: General Methods of Preparation of Ethers
In general, ethers are low on the scale of chemical reactivity because the carbon-oxygen bond is not cleaved readily. For this reason ethers frequently are employed as inert solvents in organic synthesis. Particularly important in this connection are diethyl ether, diisopropyl ether, tetrahydrofuran, and 1,4-dioxane. The mono- and dialkyl ethers of 1,2-ethanediol, 3-oxa-1,5-pentanediol, and related substances are useful high-boiling solvents. Unfortunately, their trade names are not very rational. Abbreviated names are in common use, such as “polyglymes,” “Cellosolves,” and “Carbitols.” For reference, Cellosolves are monoalkyl ethers of 1,2-ethanediol; Carbitols are monoalkyl ethers of 3-oxa-1,5-pentanediol; polyglymes are dimethyl ethers of 3-oxa-1,5-pentanediol or 3,6-dioxa-1,8-octanediol and are called diglyme and triglyme, respectively.
The spectroscopic properties of ethers are unexceptional. Like alcohols, they have no electronic absorption beyond $185 \: \text{nm}$; the important infrared bands are the $\ce{C-O}$ stretching vibrations in the region $1000$-$1230 \: \text{cm}^{-1}$; their proton nmr spectra show deshielding of the alpha hydrogens by the ether oxygen $\left( \delta_{HC_{\alpha}OC} \sim 3.4 \: \text{ppm} \right)$. The mass spectra of ethers and alcohols are very similar and give abundant ions of the type ($\ce{R} = \ce{H}$ or alkyl) by $\alpha$-cleavage (see Section 15-2).
Unlike alcohols, ethers are not acidic and usually do not react with bases. However, exceptionally strong basic reagents, particularly certain alkali-metal alkyls, will react destructively with many ethers:
Ethers, like alcohols, are weakly basic and are converted to highly reactive salts by strong acids (e.g., $\ce{H_2SO_4}$, $\ce{HClO_4}$, and $\ce{HBr}$) and to relatively stable coordination complexes with Lewis acids (e.g., $\ce{BF_3}$ and $\ce{RMgX}$):
Dimethyl ether is converted to trimethyloxonium fluoroborate by the combination of boron trifluoride and methyl fluoride:
Both trimethyl- and triethyloxonium salts are fairly stable and can be isolated as crystalline solids. They are prepared more conveniently from the appropriate boron trifluoride etherate and chloromethyloxacyclopropane (epichlorohydrin).
Trialkyloxonium ions are much more susceptible to nucleophilic displacement reactions than are neutral ether molecules. The reason is that $\ce{ROR}$ is a better leaving group than $\ce{RO}^\ominus$. In fact, trimethyloxonium salts are among the most effective methylating reagents known:
Ethers can be cleaved under strongly acidic conditions by intermediate formation of dialkyloxonium salts. Hydrobromic and hydroiodic acids are especially useful for ether cleavage because both are strong acids and their anions are good nucleophiles. Tertiary alkyl ethers are very easily cleaved by acid reagents:
Ethers are susceptible to attack by halogen atoms and radicals, and for this reason they are not good solvents for radical reactions. In fact, ethers are potentially hazardous chemicals, because in the presence of atmospheric oxygen radical-chain formation of peroxides occurs, and peroxides are unstable, explosion-prone compounds. This process is called autoxidation and occurs not only with ethers but with many aldehydes and hydrocarbons. The reaction may be generalized in terms of the following steps involving initiation (1), propagation (2 and 3), and termination (4).
The initiation and termination steps can occur in a variety of ways but it is the chain-carrying steps, 2 and 3, that effect the overall destruction of the compound. Commonly used ethers such as diethyl ether, diisopropyl ether, tetrahydrofuran, and 1,4-dioxane often become seriously contaminated with peroxides formed by autoxidation on prolonged storage and exposure to air and light. Purification of ethers frequently is necessary before use, and caution always should be exercised in the last stages of distilling them, because the distillation residues may contain dangerously high concentrations of explosive peroxides.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/15%3A_Alcohols_and_Ethers/15.10%3A_Protection_of_Hydroxyl_Groups.txt |
Nomenclature of Cyclic Ethers
Ring compounds containing nitrogen, oxygen, sulfur, or other elements as ring atoms generally are known as heterocyclic compounds, and the ring atoms other than carbon are the hetero atoms. Over the years, the more common heterocyclic compounds have acquired a hodge-podge of trivial names, such as ethylene oxide, tetrahydrofuran, and dioxane. Systematic naming of ring compounds is necessary but, unfortunately, several competing systems have been developed. We prefer the simplest procedure, which is to name the simple heterocycles as oxa, aza, and thia-derivatives of cycloalkanes. However, this procedure has not been accepted (or adopted) universally and we are obliged to deal with the usages in existing literature. Having lived through at least two cycles of drastic changes in these matters, the authors hope that the simple procedure will prevail in the long run, but the long run is still ahead.
We summarize here the rules of the so-called Hantzsch-Widman nomenclature system for heterocycles, which currently is the fashionable procedure, although relegated to second-class status by a recent, very practical approach to organic nomenclature.$^1$
1. Ring size is denoted by the stem, ir, et, ol, in, ep, oc, on, or ec for 3-, 4-, 5-, 6-, 7-, 8-, 9-, or 10-membered rings, respectively.
2. The kind of hetero atom present is indicated by the prefix, oxa, thia, or aza for oxygen, sulfur, or nitrogen, respectively; the prefixes dioxa, dithia, or diaza denote two oxygen, sulfur, or nitrogen atoms. When two or more different hetero atoms are present, they are cited in order of preference: oxygen before sulfur before nitrogen, as in the prefixes oxaza for one oxygen and one nitrogen, and thiaza for one sulfur and one nitrogen.
3. The degree of unsaturation is specified in the suffix. A list of appropriate suffixes and their stems according to ring sizes is given in Table 15-5. Notice that the suffix changes slightly according to whether the ring contains nitrogen.
Table 15-5: Stems, Suffix, and Ring Size of Heterocyclic Compounds
4. Numbering of the ring starts with the hetero atom and proceeds around the ring so as to give substituents (or other hetero atoms) the lowest numbered positions. When two or more different hetero atoms are present, oxygen takes precedence over sulfur and sulfur over nitrogen for the number one position. Examples follow to illustrate both the heterocycloalkane and the Hantzsch-Widman systems. Trivial names also are included.
Although Hantzsch-Widman system works satisfactorily (if you can remember the rules) for monocyclic compounds, it is cumbersome for polycyclic compounds. In the case of oxiranes it is simplest for conversational purposes to name them as oxides of the cycloalkenes or epoxy derivatives of the corresponding cycloalkanes. The oxabicycloalkane names seem preferable for indexing purposes, particularly because the word “oxide” is used in many other connections.
Reactivity of Cyclic Ethers - Oxacyclopropanes (Oxiranes)
Oxacyclopropane (oxirane), the simplest cyclic ether, is an outstanding exception to the generalization that most ethers are resistant to cleavage. Like cyclopropane, the three-membered ring is highly strained and readily opens under mild conditions. Indeed, the importance of oxacyclopropane as an industrial chemical lies in its readiness to form other important compounds. The major products derived from it are shown in Figure 15-5.
The lesser known four-membered cyclic ether, oxacyclobutane (oxetane), $\ce{(CH_2)_3O}$, also is cleaved readily, but less so than oxacyclopropane. Oxacyclopentane (oxolane, tetrahydrofuran) is a relatively unreactive water-miscible compound with desirable properties as an organic solvent. It often is used in place of diethyl ether in Grignard reactions and reductions with lithium aluminum hydride.
Preparation of Oxacyclopropane
Three-membered cyclic ethers are important as reactive intermediates in organic synthesis. Like the cyclopropanes, the vicinal$^2$ disubstituted compounds have cis and trans isomers:
The most important method of preparation involves oxidation, or "epoxidation", of an alkene with a peroxycarboxylic acid, $\ce{RCO_3H}$. This reaction achieves suprafacial addition of oxygen across the double bond, and is a type of electrophilic addition to alkenes:
Oxacyclopropanes also can be prepared from vicinal chloro- or bromoalcohols and a base. This is an internal $S_\text{N}2$ reaction and, if the stereochemistry is correct, proceeds quite rapidly, even if a strained ring is formed:
Ring-Opening Reactions of Oxacyclopropanes
Unlike most ethers, oxacyclopropanes react readily with nucleophilic reagents. These reactions are no different from the nucleophilic displacements previously encountered in Chapter 8, except that the leaving group, which is the oxygen of the oxide ring, remains a part of the original molecule. The stereochemistry is consistent with an $S_\text{N}2$ mechanism because inversion of configuration at the site of attack occurs. Thus cyclopentene oxide yields products with the trans configuration:
Acidic conditions also can be used for the cleavage of oxacyclopropane rings. An oxonium ion is formed first, which subsequently is attacked by the nucleophile in an $S_\text{N}2$ displacement or forms a carbocation in an $S_\text{N}1$ reaction. Evidence for the $S_\text{N}2$ mechanism, which produces inversion, comes not only from the stereochemistry but also from the fact that the rate is dependent on the concentration of the nucleophile. An example is ring opening with hydrogen bromide:
The same kind of mechanism can operate in the formation of 1,2-diols by acid-catalyzed ring-opening with water as the nucleophile:
Some acid-catalyzed solvolysis reactions of oxacyclopropanes appear to proceed by $S_\text{N}1$ mechanisms involving carbocation intermediates. Evidence for the $S_\text{N}1$ mechanism is available from the reactions of unsymmetrically substituted oxacyclopropanes. For example, we would expect the conjugate acid of 2,2-dimethyloxacyclopropane to be attacked by methanol at the primary carbon by an $S_\text{N}2$ reaction and at the tertiary carbon by an $S_\text{N}1$ reaction:
Because both products actually are obtained, we can conclude that both the $S_\text{N}1$ and $S_\text{N}2$ mechanisms occur. The $S_\text{N}1$ product, the tertiary ether, is the major product.
Metal Complexes of Cyclic Polyethers
We have emphasized the contrasts in properties between the ionic compounds such as sodium chloride, and the nonpolar organic compounds, such as the alkanes and arenes. There are great difficulties in dissolving the extreme types of these substances in a mutually compatible medium for carrying on chemical reactions, as, for example, in $S_\text{N}$ reactions of organic halides with alkali-metal salts (Sections 8-3 and 8-7F). The essence of the problem is that electrostatic forces in ionic crystals of inorganic salts are strong, and nonpolar solvents simply do not have the solvating power for ions to make dissolution of the crystals a favorable process. However, it has long been known that polyethers, such as the "glymes" (Section 15-10), are able to assist in the dissolution of ionic compounds through their ability to solvate metal cations by providing multiple complexing sites:
In 1967, C. J. Pedersen reported the synthesis of a series of cyclic polyethers, which he called "crown ethers", that have shown great promise for brining together what traditionally have been regarded as wholly incompatible substances - even achieving measurable solubilities of salts such as $\ce{NaCl}$, $\ce{KOH}$, and $\ce{KMnO_4}$ in benzene. The crown ethers can be regarded as cyclic "glymes" and are available by $S_\text{N}2$-type cyclization reactions:
The crown ethers and many modifications of them (especially with nitrogen replacing one or more of the oxygens) function by coordinating with metal cations and converting them into less polar entities that are more stable in solution, even in a nonpolar solvent, than they are in the crystal.
Many of the crown ethers have considerable specificity with regard to the metal with which they complex. Ring size as well as the number and kind of hetero atoms are very important in this connection. 18-Crown-6 is especially effective for potassium:
An important application for the crown ethers in synthetic work is for solubilization of salts such as $\ce{KCN}$ in nonpolar solvents for use in $S_\text{N}2$ displacements. If the solvent has a low anion-solvating capability, then the reactivity of the anion is enhanced greatly. Consequently many displacement reactions that proceed slowly at elevated temperatures will proceed at useful rates at room temperatures, because the energy of "desolvating" the anion before it undergoes $S_\text{N}2$ displacement is low (Section 8-7F). For example, potassium fluoride becomes a potent nucleophilic reagent in nonpolar solvents when complexed with 18-crown-6:
Acetals and Ketals as Ethers
The grouping $\ce{C-O-C-O-C}$ is characteristic of an acetal or a ketal (see Section 15-4E), but it also can be regarded as an ether with two ether links to one carbon. Compared to other ethers (except for the oxacyclopropanes), substances with the $\ce{C-O-C-O-C}$ group are very active toward acidic reagents, as pointed out in connection with their formation from alcohols (Section 15-4E) and their sue as protecting groups for the $\ce{OH}$ function (Section 15-9C).
$^1$J. H. Fletcher, O. C. Dermer, and R. B. Fox (Editors), "Nomenclature of Organic Compounds, Principles and Practice." Advances in Chemistry Series, No. 126, American Chemical Society, Washington, D.C., 1974.
$^2$Vicinal means substituted on adjacent carbons.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/15%3A_Alcohols_and_Ethers/15.12%3A_Cyclic_Ethers.txt |
The carbonyl group is a structural feature of many different types of compounds. It is present in carbon dioxide and in methanal, which represent respectively the high and low extremes in the level of oxidation of a carbonyl carbon. In between, there are carbonyl compounds ranging from aldehydes and ketones to carboxylic acids and their derivatives (esters, amides, anhydrides, and acyl halides). In this and succeeding chapters we describe the chemistry of these compounds with the intent of emphasizing the similarities that exist between them. The differences turn out to be more in degree than in kind. Even so, it is convenient to discuss aldehydes and ketones separately from carboxylic acids and, following some general observations about the carbonyl group, this chapter mainly is concerned with aldehydes and ketones.
• 16.1: Prelude to Aldehydes and Ketones
The carbonyl group, −C=O, is a structural feature of many different types of compounds. It is present in carbon dioxide and in methanal, which represent respectively the high and low extremes in the level of oxidation of a carbonyl carbon In between, there are carbonyl compounds ranging from aldehydes and ketones to carboxylic acids and their derivatives (esters, amides, anhydrides, and acyl halides).
• 16.2: The Carbonyl Bond
The carbonyl bond is both a strong bond and a reactive bond. The bond energy varies widely with structure. Methanal has the weakest bond (166 kcal) and carbon monoxide the strongest (237.3kcal). Irrespective of these variations, the carbonyl bond not only is significantly stronger but also is more reactive than a carbon-carbon double bond.
• 16.3: Physical Properties
The polarity of the carbonyl group is manifest in the physical properties of carbonyl compounds. Boiling points for the lower members of a series of aldehydes and ketones are 50-80o higher than for hydrocarbons of the same molecular weight. The water solubility of the lower-molecular-weight aldehydes and ketones is pronounced, which is the consequence of hydrogen-bonding between the water and the electronegative oxygen of the carbonyl group.
• 16.4: Spectroscopic Properties
A carbonyl group in a compound can be positively identified by the strong infrared absorption band in the region 1650-1850cm−1, which corresponds to the stretching vibration of the carbon-oxygen double bond. The position of the band within this frequency range depends on the molecular environment of the carbonyl group. As a result, we frequently can tell from the band position whether the structure is an aldehyde, ketone, carboxylic acid, ester, amide, or anhydride.
• 16.5: Typical Carbonyl-Addition Reactions
We turn now to discuss a few specific addition reactions of the carbonyl groups of aldehydes and ketones. We shall not attempt to provide an extensive catalog of reactions, but will try to emphasize the principles involved with especially important reactions that are useful in synthesis.
• 16.6: Catalytic Hydrogenation
The simplest large-scale procedure for reduction of aldehydes and ketones to alcohols is by catalytic hydrogenation since the product can be obtained simply by filtration from the catalyst and then distillation. The common catalysts are nickel, palladium, copper chromite, or platinum activated with ferrous ion. Hydrogenation of aldehyde and ketone carbonyl groups is much slower than of carbon-carbon double bonds so more strenuous conditions are needed.
• 16.7: Reduction of Carbonyl Compounds to Hydrocarbons
There are several methods of reducing carbonyl groups to hydrocarbons. In some cases, a three-step sequence of conventional reactions may be useful with alcohol and alkene intermediates. Alternative approach include the Clemmensen and Wolff-Kishner reductions.
• 16.8: Oxidation of Carbonyl Compounds
Aldehdyes are oxidized easily by moist silver oxide or by potassium permanganate solution to the corresponding acids. The mechanism of the permanganate oxidation has some resemblance to the chromic acid oxidation of alcohols. The oxidation of benzenecarbaldehyde with peroxybenzenecarboxylic acid is an example of a reaction of wide applicability in which aldehydes are oxidized to carboxylic acids, and ketones are oxidized to esters. The reaction is known as a Baeyer-Villiger oxidation.
• 16.9: Protection of Carbonyl Groups
There are few reactions of aldehydes and ketones that do not in some way affect the carbonyl function. For this reason, it may be necessary to protect the carbonyl function when it is desirable to avoid reaction at this function.
• 16.10: Preparative Methods for Aldehydes and Ketones
A number of useful reactions for the preparation of aldehydes and ketones, such as ozonization of alkenes and hydration of alkynes, have been considered in previous chapters. Only a few rather general methods that we have not discussed will be taken up here.
• 16.11: General Methods for the Preparation of Aldehydes and Ketones
• 16.E: Carbonyl Compounds I (Exercises)
These are the homework exercises to accompany Chapter 16 of the Textmap for Basic Principles of Organic Chemistry (Roberts and Caserio).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
16: Carbonyl Compounds I- Aldehydes and Ketones. Addition Reactions of the Carbonyl Group
The carbonyl group, \(\ce{-C=O}\), is a structural feature of many different types of compounds. It is present in carbon dioxide and in methanal, which represent respectively the high and low extremes in the level of oxidation of a carbonyl carbon:
In between, there are carbonyl compounds ranging from aldehydes and ketones to carboxylic acids and their derivatives (esters, amides, anhydrides, and acyl halides). The naming of these compounds is described in Sections 7-4 to 7-7.
At the upper end of the oxidation scale, along with \(\ce{CO_2}\), are the carbonic acid derivatives such as carbonic esters, amides, halides, and carbonate salts, and isocyanates:
In this and succeeding chapters we describe the chemistry of these compounds with the intent of emphasizing the similarities that exist between them. The differences turn out to be more in degree than in kind. Even so, it is convenient to discuss aldehydes and ketones separately from carboxylic acids and, following some general observations about the carbonyl group, this chapter mainly is concerned with aldehydes and ketones.
Apart from \(\ce{CO_2}\) and metal carbonates, the most abundant carbonyl compounds of natural origin are carboxylic esters and amides. These occur as fats and lipids, which are esters of long-chain alkanoic acids, and as proteins, which are polyamides of natural amino acids. The same structural features are found in certain synthetic polymers, in particular the polyesters (e.g., Dacron) and the polyamides (e.g., nylon 6):
Compared to carboxylic and carbonic acid derivatives, the less highly oxidized carbonyl compounds such as aldehydes and ketones are not so widely-spread in nature. That is not to say that they are unimportant. To the contrary, aldehydes and ketones are of great importance both in biological chemistry and in synthetic organic chemistry. However, the high reactivity of the carbonyl group in these compounds enables them to function more as intermediates in metabolism or in synthesis than as end products. This fact will become evident as we discuss the chemistry of aldehydes and ketones. Especially important are the addition reactions of carbonyl groups, and this chapter is mostly concerned with this kind of reaction of aldehydes and ketones.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/16%3A_Carbonyl_Compounds_I-_Aldehydes_and_Ketones._Addition_Reactions_of_the_Carbonyl_Group/16.01%3A_Prelude_to_Aldehydes_and_Ketones.txt |
Comparison with Carbon-Carbon Double Bonds
The carbonyl bond is both a strong bond and a reactive bond. The bond energy varies widely with structure, as we can see from the carbonyl bond energies in Table 16-1. Methanal has the weakest bond $\left( 166 \: \text{kcal} \right)$ and carbon monoxide the strongest $\left( 237.3 \: \text{kcal} \right)$. Irrespective of these variations, the carbonyl bond not only is significantly stronger but also is more reactive than a carbon-carbon double bond. A typical difference in stability and reactivity is seen in hydration:
The equilibrium constant for ethene hydration is considerably greater than for methanal hydration, largely because the carbon-carbon double bond is weaker. Even so, methanal adds water rapidly and reversibly at room temperature without need for a catalyst. The corresponding addition of water to ethene occurs only in the presence of strongly acidic catalysis (Section 10-3E, Table 15-2).
Table 16-1: Carbonyl Bond Energies
Structure and Reactivity
The reactivity of the carbonyl bond is primarily due to the difference in electronegativity between carbon and oxygen, which leads to a considerable contribution of the dipolar resonance form with oxygen negative and carbon positive:
The polarity of the carbonyl bond facilitates addition of water and other polar reagents relative to addition of the same reagent to alkene double bonds. This we have seen previously in the addition of organometallic compounds $\overset{\delta \ominus}{\ce{R}} \overset{\delta \oplus}{\ce{-MgX}}$ and $\overset{\delta \ominus}{\ce{R}} \overset{\delta \oplus}{\ce{-Li}}$ to carbonyl compounds (Section 14-12A). Alkene double bonds are normally untouched by these reagents:
Likewise, alcohols add readily to carbonyl compounds, as described in Section 15-4E. However, we must keep in mind the possibility that, whereas additions to carbonyl groups may be rapid, the equilibrium constants may be small because of the strength of the carbonyl bond.
Further Considerations of Reactivity
The important reactions of carbonyl groups characteristically involve addition at one step or another. For the reactions of organometallic reagents and alcohols with carbonyl compounds (Chapters 14 and 15), you may recall that steric hindrance plays an important role in determining the ratio between addition and other, competing reactions. Similar effects are observed in a wide variety of other reactions. We expect the reactivity of carbonyl groups in addition processes to be influenced by the size of the substituents thereon, because when addition occurs the substituent groups are pushed back closer to one another. In fact, reactivity and equilibrium constant decrease with increasing bulkiness of substituents, as in the following series (also see Table 15-3):
Strain effects also contribute to reactivity of cyclic carbonyl compounds. The normal bond angles around a carbonyl group are about $120^\text{o}$:
Consequently if the carbonyl group is on a small carbocyclic ring, there will be substantial angle strain and this will amount to about $120^\text{o} - 60^\text{o} = 60^\text{o}$ of strain for cyclopropanone,
and $120^\text{o} - 90^\text{o} = 30^\text{o}$ of strain for cyclobutanone (both values being for the $\angle \ce{C-C-C}$ at the carbonyl group). Addition of a nucleophile such as $\ce{CH_3OH}$ (cf. Section 15-4E) to these carbonyl bonds creates a tetrahedral center with less strain in the ring bonds to $\ce{C_1}$:
Thus the hemiketal from cyclopropanone will have $109.5^\text{o} = 60^\text{o} = 49.5^\text{o}$, and that from cyclobutanone $109.5^\text{o} - 90^\text{o} = 19.5^\text{o}$ of strain at $\ce{C_1}$. This change in the angle strain means that a sizable enhancement of both the reactivity and equilibrium constant for addition is expected. In practice, the strain effect is so large that cyclopropanone reacts rapidly with methanol to give a stable hemiketal from which the ketone cannot be recovered. Cyclobutanone is less reactive than cyclopropanone, but more reactive than cyclohexanone or cyclopentanone.
Electrical effects also are important in influencing the ease of addition to carbonyl groups. Electron-attracting groups facilitate the addition of nucleophilic reagents to carbon by increasing its positive character:
Thus compounds such as the following add nucleophilic reagents readily:
$^1$An electrical dipole results when unlike charges are separated. The magnitude of the dipole, its dipole moment, is given by $e \times r$, where $e$ is the magnitude of the charges and $r$ is the distance the charges are separated. Molecular dipole moments are measured in debye units $\left( \text{D} \right)$. A pair of ions, $\overset{\oplus}{\ce{C}}$ and $\overset{\ominus}{\ce{O}}$, as point charges at the $\ce{C=O}$ distance of $1.22 \: Å$, would have a dipole moment of $5.9 \: \text{D}$. Thus, if the dipole moment of a carbonyl compound is $2.7 \: \text{D}$, we can estimate the "$\%$ ionic character" of the bond to be $\left( 2.7/5.9 \right) \times 100 = 46\%$. The analysis is oversimplified in that the charges on the atom are not point charges and we have assumed that all of the ionic character of the molecule is associated with the $\ce{C=O}$ bond. One should be cautious in interpreting dipole moments in terms of the ionic character of bonds. Carbon dioxide has no dipole moment, but certainly has polar $\ce{C=O}$ bonds. The problem is that the dipoles associated with the $\ce{C=O}$ bonds of $\ce{CO_2}$ are equal and opposite in direction to each other and, as a result, cancel. Thus, $\overset{\delta \ominus}{\ce{O}} \overset{\delta \oplus}{\ce{-C}} \overset{\delta \ominus}{\ce{-O}}$ has no net dipole moment, even though it has highly polar bonds.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/16%3A_Carbonyl_Compounds_I-_Aldehydes_and_Ketones._Addition_Reactions_of_the_Carbonyl_Group/16.02%3A_The_Carbonyl_Bond.txt |
The polarity of the carbonyl group is manifest in the physical properties of carbonyl compounds. Boiling points for the lower members of a series of aldehydes and ketones are $50$-$80^\text{o}$ higher than for hydrocarbons of the same molecular weight; this may be seen by comparing the data of Table 16-2 (physical properties of aldehydes and ketones) with those in Table 4-1 (physical properties of alkanes).
Table 16-2: Physical Properties of Aldehydes and Ketones
The water solubility of the lower-molecular-weight aldehydes and ketones is pronounced (see Table 16-2). This is to be expected for most carbonyl compounds of low molecular weight and is the consequence of hydrogen-bonding between the water and the electronegative oxygen of the carbonyl group:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
16.04: Spectroscopic Properties
Infrared Spectra
A carbonyl group in a compound can be positively identified by the strong infrared absorption band in the region $1650$-$1850 \: \text{cm}^{-1}$, which corresponds to the stretching vibration of the carbon-oxygen double bond. The position of the band within this frequency range depends on the molecular environment of the carbonyl group. As a result, we frequently can tell from the band position whether the structure is an aldehyde, ketone, carboxylic acid, ester, amide, or anhydride. The date of Table 16-3 show typical infrared absorption frequencies for specific types of carbonyl compounds. Thus aldehydes and ketones absorb at slightly lower frequencies (longer wavelengths) than carboxylic esters and anhydrides. We usually find that absorption shifts to lower frequencies $\left( \sim 20 \: \text{cm}^{-1} \right)$ when the carbonyl group is conjugated with other multiple bonds, as in aromatic ketones, $\ce{C_6H_5COCH_3}$.
Table 16-3: Characteristic Infrared Absorption Frequencies of Carbonyl Compounds$^a$
Aldehydes can be distinguished from ketones by a band at $2720 \: \text{cm}^{-1}$ which is characteristic of the $\ce{C-H}$ stretching vibration of an aldehyde function:
This band is unusually low in frequency for a $\ce{C-H}$ stretching vibration; although the band is rather weak, it occurs in a region of the spectrum where other absorptions generally are absent so it can be identified with no special difficulty.
Electronic Absorption Spectra
Aldehydes and ketones absorb ultraviolet light in the region $275$-$295 \: \text{nm}$, and the result is excitation of an unshared electron on oxygen to a higher energy level. This is the $n \rightarrow \pi^*$ transition discussed in Section 9-9. A more intense $\pi \rightarrow \pi^*$ transition occurs about \180\)-$190 \: \text{nm}$, which corresponds to excitation of an electron from a $\pi$-bonding orbital to a $\pi$-antibonding orbital. Neither of these absorptions is especially useful for specific identification unless the carbonyl group is conjugated, in which case the $n \rightarrow \pi^*$ and $\pi \rightarrow \pi%+^*$ bands occur at longer wavelengths (by $30$-$40 \: \text{nm}$). For example, if you suspect that a compound is an alkenone from its infrared spectrum, you easily could tell from the $\lambda_\text{max}$ of the $n \rightarrow \pi^*$ and $\pi \rightarrow \pi^*$ absorptions of the compound whether it is a conjugated alkenone. The absorption frequency would be expected around $320 \: \text{nm}$ and $220 \: \text{nm}$ (see Figure 9-20).
Mass Spectra
Aldehydes and ketones generally give moderately intense signals due to their molecular ions, $\ce{M^+}$. Thus the determination of the molecular weight of a ketone by mass spectroscopy usually is not difficult. Furthermore, there are some characteristic fragmentation patterns that aid in structural identification. These are:
• $\alpha$ cleavage
• transfer of $\gamma$ hydrogen with $\beta$ cleavage (McLafferty rearrangement)
NMR Spectra
The character of the carbonyl bond gives rise to very low-field nmr absorptions for the proton of an aldehyde group $\left( \ce{-CH=O} \right)$. As Table 9-4 shows, these absorptions are some $4 \: \text{ppm}$ to lower fields than alkenyl hydrogens (). Some of this difference in shift can be ascribed to the polarity of the carbonyl group $\overset{\delta \oplus}{\ce{C}} \ce{=} \overset{\delta \ominus}{\ce{O}}$, which reduces electron density around the aldehyde hydrogen (see Section 9-10E). The effect appears to carry over in much smaller degree in the $\alpha$ positions, and protons of the type are about $0.3 \: \text{ppm}$ to lower fields than those of .
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/16%3A_Carbonyl_Compounds_I-_Aldehydes_and_Ketones._Addition_Reactions_of_the_Carbonyl_Group/16.03%3A_Physical_Properties.txt |
We turn now to discuss a few specific addition reactions of the carbonyl groups of aldehydes and ketones. We shall not attempt to provide an extensive catalog of reactions, but will try to emphasize the principles involved with especially important reactions that are useful in synthesis.
Grignard reagents, organolithium compounds, and the like generally add to aldehydes and ketones rapidly and irreversibly, but the same is not true of many other reagents; their addition reactions may require acidic or basic catalysts; the adducts may be formed reversibly and with relatively unfavorable equilibrium constants. Also, the initial adducts may be unstable and react further by elimination. (We recommend that you review Section 15-4E to see examples of these points.) To organize this very large number of addition reactions, we have arranged the reactions according to the nucleophile that adds to the carbonyl carbon. The types of nucleophiles considered here form $\ce{C-C}$, $\ce{C-O}$, $\ce{C-N}$, $\ce{C}$-halogen, $\ce{C-S}$, and $\ce{C-H}$ bonds. A summary is given in Table 16-4.
Table 16-4: Addition Reactions of Aldehydes and Ketones
Addition of Carbon Nucleophiles
Cyanohydrin Formation
Hydrogen cyanide adds to many aldehydes and ketones to give hydroxylnitriles, usually called "cyanohydrins":
The products are useful in synthesis - for example, in the preparation of cyanoalkenes and hydroxy acids:
An important feature of cyanohydrin formation is that it requires a basic catalyst. In the absence of base, the reaction does not proceed, or is at best very slow. In principle, the basic catalyst may activate either the carbonyl group or hydrogen cyanide. With hydroxide ion as the base, one reaction to be expected is a reversible addition of hydroxide to the carbonyl group:
However, such addition is not likely to facilitate formation of cyanohydrin because it represents a competitive saturation of the carbonyl double bond. Indeed, if the equilibrium constant for this addition were large, an excess of hydroxide ion could inhibit cyanohydrin formation by tying up the ketone as the adduct $1$.
Hydrogen cyanide itself has no unshared electron pair on carbon and does not form a carbon-carbon bond to a carbonyl carbon. However, a small amount of a strong base can activate hydrogen cyanide by converting it to cyanide ion, which can function as a carbon nucleophile. A complete sequence for cyanohydrin formation follows:
The second step regenerates the cyanide ion. Each step of the reaction is reversible but, with aldehydes and most nonhindered ketones, formation of the cyanohydrin is reasonably favorable. In practical syntheses of cyanohydrins, it is convenient to add a strong acid to a mixture of sodium cyanide and the carbonyl compound, so that hydrogen cyanide is generated in situ. The amount of acid added should be insufficient to consume all the cyanide ion, therefore sufficiently alkaline conditions are maintained for rapid addition.
Addition of Ylide Reagents
There are a number of rather interesting substances for which we can write important dipolar valence-bond structures of the type . The important factor with these structures is that the negative end of the dipole is carbon with an unshared electron pair. The positive end of the dipole can be several kinds of atoms or groups, the most usual being sulfur, phosphorus, or nitrogen. Some examples (each written here as a single dipolar valence-bond structure) are:
The systematic naming of these substances is cumbersome, but they have come to be known as ylides. The genesis of this name may seem obscure, but it is an attempt to reconcile the presence of a $\ce{C-X}$ $\sigma$ bond, which is covalent and nonpolar as in alkyl derivatives, as well as an ionic bond as in metal halides. Hence, the combination yl-lide.$^2$
As we might expect from the dipolar structure, ylides can behave as carbon nucleophiles to form carbon-carbon bonds by addition to the carbonyl groups of aldehydes and ketones:
However, the further course of reaction depends on the type of ylide used. In the case of phosphorus ylides, the overall reaction amounts to a very useful synthesis of alkenes by the transfer of oxygen to phosphorus and carbon to carbon, as summarized in Equation 16-1. This is called the Witting reaction:
Reactions with sulfur ylides proceed differently. The prodcuts are oxacylcopropanes (oxiranes) - not alkenes. The addition step proceeds as with the phosphorus ylides, but the negatively charged oxygen of the dipolar adduct then displaces the sulfonium group as a neutral sulfide. This is an intramolecular $S_\text{N}2$ reaction similar to the formation of oxacyclopropanes from vicinal chloroalcohols (Section 15-11C):
As for the nitrogen ylides, a useful reagent of this type is diazomethane, $\ce{CH_2N_2}$. Diaxomethane can react with carbonyl compounds in different ways, depending on what happens to the initial adduct $2$. Oxacyclopropanes are formed if the nitrogen is simply displaced (as $\ce{N_2}$) by oxygen (Path $a$, Equation 16-2). Ketones of rearranged carbon framework result if nitrogen is displaced $as \(\ce{N_2}$) by $\ce{R}^\ominus$ which moves over to the $\ce{CH_2}$ group (Path $b$, Equation 16-2):
Diazoketones, $\ce{RCOCHN_2}$, are formed if there is a good leaving group, such as halogen, on the carbonyl (Equation 16-3). Under these circumstances the reactant is an acid halide, not an aldehyde or ketone:
Addition of Oxygen and Sulfur Nucleophiles
Alcohols, Thiols, Water
We already have discussed additions of alcohols and, by analogy, thiols $\left( \ce{RSH} \right)$ to carbonyl compounds (see Section 15-4E). We will not repeat this discussion here except to point out that addition of water to the carbonyl group of an aldehyde is analogous to hemiacetal formation (Section 15-4E) and is catalyzed both by acids and bases:
The equilibrium for hydrate formation depends both on steric and electrical factors. Methanal is $99.99\%$ hydrated in aqueous solution, ethanal is $58\%$ hydrated, and 2-propanone is not hydrated significantly. The hydrates seldom can be isolated because they readily revert to the parent aldehyde. The only stable crystalline hydrates known are those having strongly electronegative groups associated with the carbonyl (see Section 15-7).
Hydrogen Sulfite (Bisulfite) Addition to Carbonyl Compound
Several carbonyl additions have characteristics similar to those of cyanohydrin formation. A typical example is the addition of sodium hydrogen sulfite, which proceeds readily with good conversion in aqueous solution with most aldehydes, methyl ketones, and unhindered cyclic ketones to form a carbon-sulfur bond. No catalyst is required because sulfite is an efficient nucleophilic agent. The addition step evidently involves the sulfite ion - not hydrogen sulfite ion:
The addition products often are nicely crystalline solids that are insoluble in excess concentrated sodium hydrogen sulfite solution. Whether soluble or insoluble, the addition products are useful for separating carbonyl compounds from substances that do not react with sodium hydrogen sulfite.
Polymerization of Aldehydes
A reaction closely related to acetal formation is the polymerization of aldehydes. Both linear and cyclic polymers are obtained. For example, methanal in water solution polymerizes to a solid ling-chain polymer called paraformaldehyde or "polyoxymethylene":
This material, when strongly heated, reverts to methanal; it therefore is a convenient source of gaseous methanal. When heated with dilute acid, paraformaldehyde yields the solid trimer, 1,3,5-trioxycyclohexane (mp $61^\text{o}$). The cyclic tetramer is also known.
Long-chain methanal polymers have become very important as plastics in recent years. The low cost of paraformaldehyde is highly favorable in this connection, but the instability of the material to elevated temperatures and dilute acids precludes its use in plastics. However, the "end-capping" of polyoxymethylene chains through formation of esters or acetals produces a remarkable increase in stability, and such modified polymers have excellent properties as plastics. Delrin (DuPont) is a stabilized methanal polymer with exceptional strengths and ease of molding.
Ethanal (acetaldehyde) polymerizes under the influence of acids to the cyclic trimer, "paraldehyde", and a cyclic tetramer, "metaldehyde". Paraldehyde has been used as a relatively nontoxic sleep-producing drug (hypnotic). Metaldehyde is used as a poison for snails and slugs, "Snarol". Ketones do not appear to form stable polymers like those of aldehydes.
Nitrogen Nucleophiles
Reactions of $\ce{RNH_2}$ Derivatives with Carbonyl Compounds
A wide variety of substances with $\ce{-NH_2}$ groups react with aldehydes and ketones by an addition-elimination sequence to give compounds and water. These reactions usually require acid catalysts:
Table 16-5 summarizes several important reactions of this type and the nomenclature of the reactants and products.
Table 16-5: Products from Reactions of Carbonyl Compounds with $\ce{RNH_2}$ Derivatives
Clearly, if the unshared electron pair on the nitrogen of $\ce{RNH_2}$ is combined with a proton, Equation 16-4, it cannot attack the carbonyl carbon to give the aminoalkanol as in Equation 16-5. So at high acid concentration (low pH) we expect the rate and the equilibrium for the overall reaction to be unfavorable. In more dilute acid, the rate picks up because there is more free $\ce{RNH_2}$ in solution. Dehydration of the aminoalkanol (Equation 16-6) is acid catalyzed; this reaction normally is fast at pH values smaller than 3-4. Therefore, the slow step at pH $<$ 4 is addition of $\ce{RNH_2}$ to the carbonyl group as per Equation 16-5. As the pH is increased above 4, the addition becomes progressively faster because less $\ce{RNH_2}$ is tied up as $\ce{RNH_3^+}$. However, then the dehydration step, Equation 16-6, decreases in rate because it requires an acid catalyst. At pH 6 (remember that going from pH 4 to pH 6 is a 100-fold decrease in $\ce{H}^\oplus$ concentration), dehydration is the slow step, and at higher pH values it finally becomes too slow to give a useful overall rate of reaction. This sequence of changes in rate and equilibria has been shown to account precisely for rate vs. pH curves such as in Figure 16-4.
Dehydration of $\ce{(CH_3)_2CHNR(OH)}$ to $\ce{(CH_3)_2C=NR}$ involves acid catalysis in very much the same way as in acetal formation (Section 15-4E):
Addition of Ammonia to Aldehydes
Ammonia adds readily to many aldehydes. For example,
The aldehyde-ammonia adducts usually are not very stable. They readily undergo dehydration and polymerization. 1-Aminoethanol, for example, gives a cyclic trimer composition $\ce{C_6H_{15}N_3} \cdot 3 \ce{H_2O}$, mp $97^\text{o}$, with structure $4$:
Methanal and ammonia react by a different course with the formation of a substance known as "hexamethylenetetramine":
The product is a high-melting solid (mp $> \: 230^\text{o}$ d.) and its structure has been established by x-ray diffraction (Section 9-3). In fact, it was the first organic substance whose structure was determined in this way. The high melting point is clearly associated with the considerable symmetry and rigidity of the cage structure:
The corresponding all-carbon compound, adamantane (Section 12-8), also has a high melting point $\left( 268^\text{o} \right)$:
Treatment of hexamethylenetetramine with nitric acid gives the high explosive "cyclonite", which often is designated as RDX and was widely used in World War II:
Then methanal and ammonia that split off the cage structure during the reaction with nitric acid need not be wasted. In the large-scale manufacture of cyclonite, a combination of nitric acid, ammonium nitrate, and ethanoic anhydride is used, which results in full utilization of the methanal and ammonia:
$\ce{C_6H_{12}N_4} + 4 \ce{HNO_3} + 2 \ce{NH_4NO_3} + 6 \ce{(CH_3CO)_2O} \rightarrow 2 \ce{C_3H_6O_6N_6} + 12 \ce{CH_3CO_2H}$
Enamines
Secondary amino compounds of the type $\ce{R_2N-H}$ add to aldehyde and ketone carbonyl groups in an acid-catalyzed reaction in much the same way as do $\ce{RNH_2}$ compounds - with one important difference. The product contains the structural unit $\ce{C=C-N}$ rather than $\ce{C-C=N}$; and because there is a carbon-carbon double bond, such a substance is called an enamine (alkene $+$ amine). An example is:
The course of this reaction can be understood if we notice that loss of $\ce{OH}$ from the initial product leads to an immonium ion, $5$, that cannot lose a proton and form a $\ce{C=N}$ bond:
However, if there is a hydrogen on a carbon attached to the immonium carbon, it is possible for such a hydrogen to be lost as a proton with concurrent formation of the neutral enamine:
Enamine formation, like many other carbonyl addition reactions, is readily reversible, and the carbonyl compound can be recovered by hydrolysis with aqueous acids. For this reason, to obtain a good conversion of carbonyl compound to enamine, it usually is necessary to remove the water that is formed by distilling it away from the reaction mixture.
Enamines are useful synthetic intermediates for the formation of carbon-carbon bonds, as we will discuss in greater detail in Section 17-4B.
Enamines generally are unstable if there is a hydrogen on nitrogen. They rearrange to the corresponding imine. This behavior is analogous to the rearrangement of alkenols to carbonyl compounds (Section 10-5A):
Hydrogen-Halide Addition to Carbonyl Groups and Replacement of Carbonyl by Halogen
Addition of hydrogen halides to carbonyl groups usually is so easily reversible as to preclude isolation of the addition products:
However, many aldehydes react with alcohols in the presence of an excess of hydrogen chloride to give $\alpha$-chloro ethers:
In carrying out laboratory syntheses of $\alpha$-chloro ethers, gaseous hydrogen chloride is passed into a mixture of the alcohol and aldehyde. Aqueous $\ce{HCl}$ is not useful because the excess water gives an unfavorable equilibrium. $\alpha$-Chloro ethers are highly reactive compounds that very readily undergo $S_\text{N}2$ as well as $S_\text{N}1$ and $E1$ reactions. Two simple examples, methoxychloromethane (chloromethyl methyl ether) and chloromethoxychloromethane (bis-chloromethyl ether), have been put under severe restrictions as the result of tests that show they have strong chemical carcinogenic properties.
Replacement of the carbonyl function by two chlorines occurs with phosphorus pentachloride in ether:
This reaction is useful in conjunction with $E2$ elimination to prepare alkenyl halides, allenes, and alkynes. Cycloalkenyl halides are easily prepared, but because of angle strain the cycloalkynes and cycloallenes with fewer than eight atoms in the ring cannot be isolated (see Section 12-7):
Replacement of a carbonyl group by gem-fluorines$^3$ can be achieved with molybdenum hexafluoride or sulfur tetrafluoride. Sulfur tetrafluoride converts carboxyl functions to trifluoromethyl groups:
Hydride as a Nucleophile. Reduction of Carbonyl Compounds
Metal and Boron Hydrides
In recent years, inorganic hydrides such as lithium aluminum hydride, $\ce{LiAlH_4}$, and sodium borohydride, $\ce{NaBH_4}$, have become extremely important as reducing agents of carbonyl compounds. These reagents have considerable utility, especially with sensitive and expensive carbonyl compounds. The reduction of cyclobutanone to cyclobutanol is a good example, and you will notice that the net reaction is the addition of hydrogen across the carbonyl double bond, $\overset{ 2 \left[ \ce{H} \right]}{\longrightarrow}$ ,
With the metal hydrides, the key step is transfer of a hydride ion to the carbonyl carbon of the substance being reduced.
The hydride transfer is analogous to the transfer of $\ce{R}^\ominus$ from organometallic compounds to carbonyl groups (Section 14-12A).
Lithium aluminum hydride is best handled like a Grignard reagent, because it is soluble in ether and is sensitive to both oxygen and moisture. (Lithium hydride is insoluble in organic solvents and is not an effective reducing agent for organic compounds.) All four hydrogens on aluminum can be utilized:
The reaction products must be decomposed with water and acid as with the Grignard complexes. Any excess lithium aluminum hydride is decomposed by water and an acid with evolution of hydrogen:
$\ce{LiAlH_4} + 2 \ce{H_2SO_4} \rightarrow \frac{1}{2} \ce{Li_2SO_4} + \frac{1}{2} \ce{Al_2(SO_4)_3} + 4 \ce{H_2}$
Lithium aluminum hydride usually reduces carbonyl groups without affecting carbon-carbon double bonds. It is, in addition, a good reducing agent for carbonyl groups of carboxylic acids, esters, and other acid derivatives, as will be described in Chapter 18.
Sodium borohydride is a milder reducing agent than lithium aluminum hydride and will reduce aldehydes and ketones, but not acids or esters. It reacts sufficiently slowly with water in neutral or alkaline solution that reductions which are reasonably rapid can be carried out in water solution without appreciable hydrolysis of the reagent:
$\ce{NaBH_4} + 4 \ce{CH_2=O} + 3 \ce{H_2O} \rightarrow 4 \ce{CH_3OH} + \ce{NaOB(OH)_2}$
Borane (as $\ce{BH_3}$ in tetrahydrofuran or dimethyl sulfide) is an even milder reducing agent than $\ce{BH_4^+}$ for the carbonyl group of aldehydes and ketones. This difference in reactivity can be used to advantage when selective reduction is necessary. For example, borohydride reduces a ketone carbonyl more rapidly than a carbon-carbon double bond, whereas borane reduces the carbon-carbon double bond more rapidly than carbonyl:
A useful comparison of the reactivities of boranes and metal hydrides toward various types of multiple bonds is given in Table 16-6.
Table 16-6: Comparison of Products and Reactivities of Functional Groups for Reduction with Borane and Metal Hydrides$^{a,b}$
The Cannizzaro Reaction
A characteristic reaction of aldehydes without $\alpha$ hydrogens is the self oxidation-reduction they undergo in the presence of strong base. With methanal as an example,
If the aldehyde has $\alpha$ hydrogens, other reactions usually occur more rapidly.
The mechanism of this reaction, usually called the Cannizzaro reaction,$^4$ combines many features of other processes studied in this chapter. The first step is reversible addition of hydroxide ion to the carbonyl group:
A hydrogen can be transferred as hydride ion to methanal from the hydroxyalkoxide ion, thereby reducing the methanal to methanol:
Reduction with Aluminum Alkoxides
Hydride transfer similar to that of the Cannizzaro reaction also may be achieved from a $\ce{C-H}$ grouping in an alkoxide ion corresponding to a primary or secondary, but not a tertiary, alcohol. This is expected to be a reversible reaction, because the products are another alkoxide and another carbonyl compound:
To utilize this equilibrium process as a practical reduction method requires rather special conditions. It is preferable to use an aluminum alkoxide, $\ce{Al(OR)_3}$, rather than a sodium alkoxide, $\overset{\oplus}{\ce{Na}} \overset{\ominus}{\ce{O}} \ce{R}$, to ensure that the reaction mixture is not too strongly basic. (Carbonyl compounds, particularly aldehydes, are sensitive to strong bases.) The overall reaction may be written
for which the alkoxide is derived from 2-propanol. The advantage of this method is that the reaction can be driven essentially to completion by distilling out the 2-propanone as it is formed. The reduction product subsequently can be obtained by acid hydrolysis of the aluminum alkoxide:
Biological Reactions
These have been discussed already in the context of the reverse reactions - oxidation of alcohols (Section 15-6C).
$^2$Pronounced variously as ill/id, yill/id, ill/ide, yill/ide. The dipolar structures usually written for ylides are an oversimplified representation of the bonding in these substances.
$^3$Gem is an abbreviation for geminal (twinned) and is a common conversational designation for arrangements having two identical substituents on one carbon.
$^4$Named after its discoverer, the same Cannizzaro who, in 1860, made an enormous contribution to the problem of obtaining self-consistent atomic weights (Section 1-1).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/16%3A_Carbonyl_Compounds_I-_Aldehydes_and_Ketones._Addition_Reactions_of_the_Carbonyl_Group/16.05%3A_Typical_Carbonyl-Addition_Reactions.txt |
The simplest large-scale procedure for reduction of aldehydes and ketones to alcohols is by catalytic hydrogenation:
The advantage over most other kinds of reduction is that usually the product can be obtained simply by filtration from the catalyst, then distillation. The common catalysts are nickel, palladium, copper chromite, or platinum activated with ferrous ion. Hydrogenation of aldehyde and ketone carbonyl groups is much slower than of carbon-carbon double bonds so more strenuous conditions are required. This is not surprising, because hydrogenation of carbonyl groups is calculated to be less exothermic than that of carbon-carbon double bonds:
It follows that it is generally difficult to reduce a carbonyl group in the presence of a carbon-carbon double bond by hydrogenation without also saturating the double bond. Other reducing agents are more selective (Section 16-4E).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
16.07: Reduction of Carbonyl Compounds to Hydrocar
There are several methods of transforming to . In some cases, the following three-step sequence of conventional reactions may be useful:
This route requires a hydrogen $\alpha$ to the carbonyl function and may give rearrangement in the dehydration step (Sections 8-9B and 15-5E). Alternatively, the hydroxyl can be converted to a better leaving group (halogen or sulfonate ester), which then may be displaced by $\ce{H}^\ominus$ (as $\ce{LiAlH_4}$; see Table 16-6):
More direct methods may be used, depending on the character of the $\ce{R}$ groups of the carbonyl compound. If the $\ce{R}$ groups are stable to a variety of reagents there is no problem, but with sensitive $\ce{R}$ groups not all methods are equally applicable. When the $\ce{R}$ groups are stable to acid but unstable to base, the Clemmensen reduction with amalgamated zinc and hydrochloric acid is often very useful.
The mechanism of the Clemmensen reduction is not well understood. It is clear that in most cases the alcohol is not an intermediate, because the Clemmensen conditions do not suffice to reduce most alcohols to hydrocarbons.
When the $\ce{R}$ groups of the carbonyl compound are stable to base but not to acid, the Huang-Minlon modification of the Wolff-Kishner reduction usually gives good results. This procedure involves heating the carbonyl compound in a high-boiling polar solvent, such as 1,2-ethanediol, with hydrazine and potassium hydroxide and driving the reaction to completion by distilling out the water formed:
When the carbonyl compound is sensitive to both acids and bases, or for other reasons gives poor yields in both the Clemmensen and Wolff-Kishner reductions, a recourse may be reduction of the corresponding thioacetal or thioketal with hydrogen-saturated Raney nickel (Section 11-2B):
Thioketals, unlike ordinary ketals, are formed readily from ketones and thiols $\left( \ce{RSH} \right)$ in the presence of acid catalysts. The desulfurization procedure usually goes well, but the product is rather difficult to separate by extraction from the large excess of Raney nickel required for optimum yields.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/16%3A_Carbonyl_Compounds_I-_Aldehydes_and_Ketones._Addition_Reactions_of_the_Carbonyl_Group/16.06%3A_Catalytic_Hydrogenation.txt |
Aldehdyes are oxidized easily by moist silver oxide or by potassium permanganate solution to the corresponding acids. The mechanism of the permanganate oxidation has some resemblance to the chromic acid oxidation of alcohols (Section 15-6B):
Many aldehydes are oxidized easily by atmospheric oxygen in a radical-chain mechanism. Oxidation of benzenecarbaldehyde to benzenecarboxylic acid has been studied particularly well and involves formation of a peroxy acid as an intermediate. Reaction is initiated by a radical $\ce{R} \cdot$ which breaks the relatively weak aldehyde $\ce{C-H}$ bond $\left( 86 \: \text{kcal} \right)$.
• initiation
The benzenecarbonyl radical, $\ce{C_6H_5} \overset{\cdot}{\ce{C}} \ce{O}$, then propagates a chain reaction.
• propagation
The peroxy acid formed then reacts with benzenecarbaldehyde to give two molecules of carboxylic acid:
The oxidation of benzenecarbaldehyde with peroxybenzenecarboxylic acid (Equation 16-8) is an example of a reaction of wide applicability in which aldehydes are oxidized to carboxylic acids, and ketones are oxidized to esters.
The reaction, which is known as the Baeyer-Villiger oxidation, has synthetic utility, particularly for the oxidation of ketones to esters because ketones normally are difficult to oxidize without degrading the structure to smaller fragments. Two examples of the Baeyer-Villiger reaction follow:
The mechanism of the Baeyer-Villiger oxidation has been studied extensively and is of interest because it involves a rearrangement step in which a substituent group $\left( \ce{R} \right)$ moves from carbon to oxygen. The reaction sequence is shown in Equations 16-9 through 16-11:
In the first step, Equation 16-9, the peroxy acid adds to the carbonyl group. The adduct has several oxygen atoms on which protons can reside, and there will be rapid shifts of protons between these oxygens. However, at some stage the structure will be appropriate to allow elimination of a molecule of carboxylic acid, $\ce{R'CO_2H}$, Equation 16-10. The resulting intermediate has an electron-deficient oxygen atom with only six valence electrons. As with carbocations and borane complexes (Sections 8-9B, 15-5E, 11-6E, and 16-9D,G), a neighboring $\ce{R}$ group can move over with its bonding electron-pair to the electron-deficient (oxygen) atom, Equation 16-11. You will notice that for aldehydes, the aldehyde hydrogen migrates in preference to the alkyl or aryl group. In the other examples given, a cycloalkyl migrates in preference to a methyl group, and aryl in preference to methyl.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
16.09: Protection of Carbonyl Groups
There are few reactions of aldehydes and ketones that do not in some way affect the carbonyl function. For this reason, it may be necessary to protect the carbonyl function when it is desirable to avoid reaction at this function. For example, you may plan to synthesize 4-cyclohexylidene-2-butanone by way of a Wittig reaction (Section 16-4A), which would involve the following sequence:
This synthesis would fail in the second step because the phenyllithium would add irreversibly to the carbonyl group. To avoid this, the carbonyl group would have to be protected or blocked, and the most generally useful method of blocking is to convert the carbonyl group to a ketal, usually a cyclic ketal:
With the carbonyl group suitably protected, the proposed synthesis would have a much better chance of success:
Notice that the carbonyl group is regenerated by acid hydrolysis in the last step.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/16%3A_Carbonyl_Compounds_I-_Aldehydes_and_Ketones._Addition_Reactions_of_the_Carbonyl_Group/16.08%3A_Oxidation_of_Carbonyl_Compounds.txt |
A number of useful reactions for the preparation of aldehydes and ketones, such as ozonization of alkenes and hydration of alkynes, have been considered in previous chapters. These and other methods of preparation are summarized in Tables 16-7 and 16-8 at the end of the chapter. Only a few rather general methods that we have not discussed will be taken up here.
Oxidation of 1,2-Diols and Alkenes
Aldehydes and ketones often can be prepared by oxidation of alkenes to 1,2-diols (Sections 11-7C and 11-7D), followed by oxidative cleavage of the 1,2-diols with lead tetraethanoate or sodium periodate. For example,
Cleavage of glycols with these reagents proceeds according to the following stoichiometry:
Oxidation of Primary Alcohols and Related Compounds
In Chapter 15 primary alcohols, $\ce{RCH_2OH}$, were shown to be readily oxidized to aldehydes, $\ce{RCHO}$, and secondary alcohols, $\ce{R_2CHOH}$, to ketones, $\ce{R_2CO}$, by inorganic reagents such as $\ce{CrO_3}$ and $\ce{KMnO_4}$. However, it is a problem to avoid overoxidation with primary alcohols because of the ease with which aldehydes are oxidized to acids, $\ce{RCHO} \rightarrow \ce{RCO_2H}$. A milder oxidant is methylsulfinylmethane [dimethyl sulfoxide, $\ce{(CH_3)_2S=O}$], and this reagent can be used to prepare aldehydes from alcohols by way of an intermediate such as the ester or halide in which the $\ce{OH}$ group is converted to a better leaving group:
Whichever method is employed, the key step is the formation of an alkoxysulfonium salt, $7$, by a displacement reaction involving dimethyl sulfoxide as an oxygen nucleophile. (Notice that the $\ce{S=O}$ bond, like the $\ce{C=O}$ bonds, is strongly polarized as $\overset{\oplus}{\ce{S}} \ce{-} \overset{\ominus}{\ce{O}}$.)
In the examples listed in Equations 16-12 through 16-15, the $\ce{X}$ group is $\ce{Br}$, $\ce{-OSO_2R'}$, $\ce{-O_2CCF_3}$, and , respectively. In the nest step a sulfur ylide, $8$, is formed from the reaction of a base with $7$, but the ylide evidently is unstable and fragments by an internal $E2$ reaction to form an aldehyde:
Reduction of Carboxylic Acids to Aldehydes
Conversion of a carboxylic acid to an aldehyde by direct reduction is not easy to achieve, because acids generally are difficult to reduce, whereas aldehydes are easily reduced. Thus the problem is to keep the reaction from going too far.
The most useful procedures involve conversion of the acid to a derivative that either is more easily reduced than an aldehyde, or else is reduced to a substance from which the aldehyde can be generated. The so-called Rosenmund reduction involves the first of these schemes; in this procedure, the acid is converted to an acyl chloride, which is reduced with hydrogen over a palladium catalyst to the aldehyde in yields up to $90\%$. The rate of reduction of the aldehyde to the corresponding alcohol is kept at a low level by poisoning the catalyst with suflur:
Metal hydrides, such as lithium aluminum hydride, also can be used to reduce derivatives of carboxylic acids (such as amids and nitriles see Table 16-6) to aldehydes. An example follows:
Rearrangements of 1,2-Diols
Many carbonyl compounds can be synthesized by acid-catalyzed rearrangements of 1,2-diols (a type of reaction often called the "pinacol-pinacolone" rearrangement).
The general characteristics of the reaction are similar to those of carbocation rearrangements (Section 8-9B). The acid assists the reaction by protonating one of the $\ce{-OH}$ groups to make it a better leaving group. The carbocation that results then can undergo rearrangement by shift of the neighboring $\ce{R}$ group with its pair of bonding electrions to give a new, thermodynamically more stable species with a carbon-oxygen double bond (see Section 16-7).
The prototype of this rearrangement is the conversion of pinacol to pinacolone as follows:
Rearrangements of Hydroperoxides
An important method of preparing carbonyl (and hydroxy) compounds, especially on an industrial scale, is through rearrangements of alkyl hydroperoxides:
The peroxides can be made in some cases by direct air oxidation of hydrocarbons,
and in others by sulfuric acid-induced addition of hydrogen peroxide (as $\ce{H-O_2H}$) to double bonds:
(Notice that hydrogen peroxide in methanoic acid behaves differently toward alkenes in producing addition of $\ce{HO-OH}$, Section 11-7D.) The direct air oxidation of hydrocarbons is mechanistically similar to that of benzenecarbaldehyde (Section 16-7).
The rearrangements of hydroperoxides are acid-catalyzed and are analogous to carbocation rearrangements except that positive oxygen (with only six valence electrons) instead of positive carbon is involved in the intermediate stage:
In principle, either phenyl or methyl could migrate to the positive oxygen, but only phenyl migration occurs in this case. The rearrangement reaction is closely related to the Baeyer-Villiger reaction (Section 16-7).
Aldehydes by Hydroformylation of Alkenes
This reaction is important for a number of reasons. It is an industrial synthesis of aldehydes from alkenes by the addition of carbon monoxide and hydrogen in the presence of a cobalt catalyst. A prime example is the synthesis of butanal from propene, in which 2-methylpropanal also is formed:
As you can see, the reaction formally amounts to the addition of methanal as $\ce{H-CHO}$ to the alkene double bond. Because one additional carbon atom is introduced as a "formyl" $\ce{CHO}$ group, the reaction often is called hydroformylation, although the older name, oxo reaction, is widely used.
Hydroformylation to produce aldehydes is the first step in an important industrial route to alcohols. The intermediate aldehydes are reduced to alcohols by catalytic hydrogenation. Large quantities of $\ce{C_4}$-$\ce{C_8}$ alcohols are prepared by this sequence:
The history of the oxo reaction is also noteworthy. It was developed originally in Germany in the years following World War I. At that time, the German chemical industry was faced with inadequate supplies of petroleum. Many German chemists therefore turned to research on ways by which hydrocarbons could be synthesized from smaller building blocks, particularly carbon monoxide and hydrogen derived from coal. The success achieved was remarkable and led to alkane and alkene syntheses known as the Fischer-Tropsch process:
This reaction in turn led to the discovery that aldehydes were formed by the further addition of carbon monoxide and hydrogen to alkenes, and was further developed as the oxo process for production of alcohols. The combination $\ce{CO} + \ce{H_2}$ often is called "synthetic gas". It is prepared by the reduction of water under pressure and at elevated temperatures by carbon (usually coke), methane, or higher-molecular-weight hydrocarbons:
Carbonylation of Alkylboranes
The aldehyde synthesis by hydroformylation of alkenes described in the preceding section can be achieved indirectly using boron hydrides. An oversimplified expression of this reaction is
The overall reaction is quite complex but involves a rearrangement similar to that described for the hydroboration-oxidation of alkenes (Section 11-6E). The first step is hydroboration of the alkene to a trialkylborane. When the trialkylborane is exposed to carbon monoxide, it reacts (carbonylates) to form a tetracovalent boron, $9$:
The complex $9$ is unstable and rearranges by transfer of an alkyl group from boron to the electron-deficient carbonyl carbon to give $10$:
Now, if a metal-hydride reducing agent, such as $\ce{LiAlH_4}$, is present, the carbonyl group of $10$ is reduced and $11$ is formed:
The reduction product, $11$, can be converted to an aldehyde by oxidation with aqueous hydrogen peroxide, provided the pH is carefully controlled. (Remember, aldehydes are unstable in strong base.)
You may have noticed that only one of the three alkyl groups of a trialkylborane is converted to an aldehyde by the carbonylation-reduction-oxidation sequence. To ensure that carbonylation takes the desired course without wasting the starting alkene, hydroboration is achieved conveniently with a hindered borane, such as "9-BBN", $12$. With $12$, only the least-hindered alkyl group rearranges in the carbonylation step:
Carbonylation of alkylboranes also can produce ketones. The conditions are similar to those in the aldehyde synthesis except that the hydride reducing agent is omitted. By omitting the reducing agent, a second boron-to-carbon rearrangement can occur. Oxidation then produces a ketone:
Rearrangement will continue a third time (ultimately to produce a tertiary alcohol) unless movement of the alkyl group remaining on boron in $13$ is prevented by steric hindrance.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
16.11: General Methods for the Preparation of Alde
Table 16-7: General Methods for the Preparation of Aldehydes and Ketones\(^a\)
Table 16-8: General Methods for the Preparation of Ketones\(^a\)
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/16%3A_Carbonyl_Compounds_I-_Aldehydes_and_Ketones._Addition_Reactions_of_the_Carbonyl_Group/16.10%3A_Preparative_Methods_for_Aldehydes_and_Keton.txt |
Some of the most useful reactions of carbonyl compounds involve carbon-hydrogen bonds adjacent to the carbonyl group. Such reactions, which can be regarded as the backbone of much synthetic organic chemistry, usually result in the replacement of the hydrogen by some other atom or group. The important examples we will consider in this chapter are halogenation, alkylation, and aldol reactions of aldehydes and ketones.
• 17.1: Prelude to Enols and Enolate Anions, Unsaturated, and Polycarbonyl Compounds
Some of the most useful reactions of carbonyl compounds involve carbon-hydrogen bonds adjacent to the carbonyl group. Such reactions, which can be regarded as the backbone of much synthetic organic chemistry, usually result in the replacement of the hydrogen by some other atom or group. The important examples we will consider in this chapter are halogenation, alkylation, and aldol reactions of aldehydes and ketones.
• 17.2: Enolization of Aldehydes and Ketones
Transformation of a carbonyl compound to an enol at a useful rate normally requires either a basic catalyst or an acidic catalyst and, of course, at least one hydrogen on the αα carbon. The features of each type of catalysis follow.
• 17.3: Halogenation of Aldehydes and Ketones
Halogenation of saturated aldehydes and ketones usually occurs exclusively by replacement of hydrogens alpha to the carbonyl group. The reagents that commonly are used to halogenate carbonyl compounds are those that are used to halogenate alkanes. However, the mechanism of the two types of halogenation normally are very different. When one attempts E2 reactions with α-halo ketones using strong bases such as alkoxides, an interesting rearrangement may occur called the Favorskii rearrangement
• 17.4: Nucleophilic Addition Reactions of Enolate Anions
A most important property of enolate anions, at least as far as synthesis is concerned, is their excellent nucleophilicity, which enables them to add to double bonds and to participate in nucleophilic substitution. When the addition is to a carbonyl double bond, it is called an aldol addition. Additions of enolate anions to carbon-carbon double bonds usually are classified as Michael additions.
• 17.5: Nucleophilic Substitution with Enolate Anions
The synthetic chemistry of enolate anions is centered on their nucleophilic and basic properties. Accordingly these ions participate in SN2 reactions with suitable alkyl compounds. However, there are a number of complicating factors to consider. First, the basic conditions needed to form the enolate ions often lead to side reactions such as aldol addition and E2 elimination of RX compounds.
• 17.6: α,β-Unsaturated Aldehydes and Ketones
The most generally useful preparation of α,β-unsaturated carbonyl compounds is by dehydration of aldol addition products. There are many addition reactions of α, β -unsaturated aldehydes, ketones, and related compounds that are the same as the carbonyl addition reactions described previously. Others are quite different and result in addition to the alkene double bond.
• 17.7: Ketenes
Substances with cumulated carbonyl and carbon-carbon double bonds are called ketenes and, as may be expected, have interesting and unusual properties. Ketene has a boiling point of −56° and normally would be stored under pressure in steel cylinders. Ketenes in general are useful reagents for acylating alcohols, and amines, because the reactions involve additions; there are no by-products to be separated:
• 17.8: 1,2-Dicarbonyl Compounds
Most of the 1,2-dicarbonyl compounds are yellow. Ethanedial is unusual in being yellow in the liquid state, but green in the vapor state. It has very reactive aldehyde groups and is employed in the manufacture of plastics and as a component of embalming fluids to harden proteins by linking together their amino groups through imine formation.
• 17.9: 1,3-Dicarbonyl Compounds
Much of the chemistry of 1,3-dialdehydes, aldehyde ketones, and diketones already has been mentioned in this chapter. The reactions discussed in this chapter that depend on the formation of enolate anions (i.e., halogenation, aldol addition, and alkylation) often proceed smoothly and under milder conditions with 1,3-diketones than with monoketones. This is because the 1,3-diketones are stronger acids and therefore can form the enolate anion with weaker bases.
• 17.10: 1,4-Dicarbonyl Compounds
Most of the reactions of the 1,4-dicarbonyl compounds are the conventional reactions expected for isolated carbonyl groups. These reactions are reasonably general and also can be used to prepare compounds with oxygen and sulfur in five-membered rings.
• 17.11: Tricarbonyl Compounds
The properties of tricarbonyl compounds are for the most part as expected, except when the three groups are contiguous to one another, as in diphenylpropanetrione. With such compounds, the central carbonyl group is highly reactive; it is lost, as carbon monoxide, in the presence of acidic catalysts such as aluminum chloride, and adds water readily to give a monohydrate.
• 17.12: Cyclopropanones and Cyclopropenones
ones deserve special comment, not because of their practical importance, but because of their novel behavior and reactivity. No unambiguous synthesis of cyclopropanones was known prior to 1965, and the older textbooks usually contained statements such as "cyclopropanones apparently cannot exist". However, they had been postulated as intermediates in various reactions, but until recently had defied isolation and identification.
• 17.E: Carbonyl Compounds II (Exercises)
These are the homework exercises to accompany Chapter 17 of the Textmap for Basic Principles of Organic Chemistry (Roberts and Caserio).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
17: Carbonyl Compounds II- Enols and Enolate Anions. Unsaturated and Polycarbonyl Compounds
Some of the most useful reactions of carbonyl compounds involve carbon-hydrogen bonds adjacent to the carbonyl group. Such reactions, which can be regarded as the backbone of much synthetic organic chemistry, usually result in the replacement of the hydrogen by some other atom or group, as in the sequence $\ce{H-C-C=O} \rightarrow \ce{X-C-C=O}$. The important examples we will consider in this chapter are halogenation, alkylation, and aldol reactions of aldehydes and ketones, illustrated here for 2-propanone:
Although these reactions lead to many diverse products depending on the reagents and conditions, they have one feature in common - they proceed by way of the enol or the enolate anion of the parent carbonyl compound:
Therefore, to understand the nature of these reactions we first must understand the conditions that convert aldehydes and ketones to their enol forms or the anions of those enol forms.
Contributors
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/17%3A_Carbonyl_Compounds_II-_Enols_and_Enolate_Anions._Unsaturated_and_Polycarbonyl_Compounds/17.01%3A_Prelude_to_Enols_and_Enolate_Anions_Unsat.txt |
Transformation of a carbonyl compound to an enol at a useful rate normally requires either a basic catalyst or an acidic catalyst and, of course, at least one hydrogen on the $\alpha$ carbon. The features of each type of catalysis follow.
Enolization in Basic Solution. $\ce{C-H}$ Acidity of Carbonyl Compounds
With a basic catalyst such as hydroxide ion, the first step in enolization is removal of a proton from the $\alpha$ position to give the enolate anion $1$:
Normally, $\ce{C-H}$ bonds are highly resistant to attack by basic reagents, but removal of a proton alpha to a carbonyl group results in the formation of a considerably stabilized anion with a substantial proportion of the negative charge on oxygen, as represented by the valence-bond structure $1a$. Carbonyl compounds such as 2-propanone therefore are weak acids, only slightly weaker than alcohols (compare the p$K_\text{a}$ values for some representative compounds in Table 17-1).$^1$
Table 17-1: $\ce{C-H}$ and $\ce{O-H}$ Acidities of Some Representative Compounds$^a$
Two carbonyl groups greatly increase the acidity. For example, 2,4-pentanedione (acetylacetone, $2$) has a p$K_\text{a} \cong$ 9, which is comparable to the $\ce{O-H}$ acidity of phenols (see Table 17-1). The reason is that the enolate anion $3$ has the charge largely shared by the two oxygen atoms (cf. $3b$ and $3c$). As a result, the enolate anion $3$ is stabilized more with respect to the ketone than the enolate anion from 2-propanone is stabilized relative to 2-propanone:
Enol Formation from Enolate Anions
You will notice from Structures $1a$ and $1b$ that because the negative charge of the enolate anion is distributed on both oxygen and carbon, the ion can, in principle, combine with a proton at either site. If the enolate ion adds a proton to oxygen, the enol is formed; if it adds a proton to carbon, the ketone is formed:
Ions of this type, which can react at either of two different sites, often are called ambident ions.
In fact, enolate anions add a proton at oxygen at least $10^{10}$ times faster than at carbon; the proton also is removed from oxygen much faster than from carbon. Thus the enolate anion of 2-propanone is in rapid equilibrium with the enol, but is converted back and forth to the ketone only slowly (Equation 17-1).
Another important point is that, although enolization by way of enolate anions requires a basic catalyst, both an acid and a base are necessary: a base to form the enolate anion; an acid to donate a proton to the anion to form the enol. If there is no acid available that is strong enough to donate a proton to the anion, then only the enolate anion is formed:
Enolization in Acid Solution
Catalysis of the enolization of 2-propanone by acids involves first, oxonium-salt formation and second, removal of an $\alpha$ proton with water or other proton acceptor (base):
This sequence differs from enolization induced by basic catalysis (as discussed in Section 17-1B) in that the enol is formed directly and not subsequent to the formation of the enolate anion. The proton addition to the carbonyl oxygen greatly facilitates proton removal from the $\alpha$ carbon because of the electron-attracting power of the positively charged oxygen. Nevertheless, this last step is the rate-determining for enolization in acid solution.
Stabilities of Enols
The equilibrium position between a simple ketone and its enol usually lies far on the side of the ketone (see Table 17-2). However, there are some interesting and important exceptions to this generalization. For instance, the influence of two carbonyl groups on the enol content is very striking, as we can see from the fact that $85\%$ of 2,4-pentanedione is the enol form at equilibrium:
Table 17-2: The Enol Content of Some Carbonyl Compounds
The enol form of 2,4-pentanedione (and of related dicarbonyl compounds of the type ) not only is stabilized by electron-delocalization, as shown in Structures $4a$ and $4b$, but by hydrogen-bonding of the acidic hydrogen between the two oxygens:
Of course, such stabilization is not possible for the keto form.
An extreme example of the stabilization of an enol by electron delocalization is benzenol (phenol), which exists $100\%$ in the enol form. In this case the extra stability of the benzene ring is the important factor:
In the succeeding sections of this chapter we will discuss several important reactions that take place by way of enols or enolate anions.
$^1$The important difference between 2-propanone and ethanol as acids is that the rate of establishment of equilibrium with 2-propanone or similar compounds where ionization involves breaking a $\ce{C-H}$ bond is very much slower than the corresponding reaction with $\ce{O-H}$ bonds.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/17%3A_Carbonyl_Compounds_II-_Enols_and_Enolate_Anions._Unsaturated_and_Polycarbonyl_Compounds/17.02%3A_Enolization_of_Aldehydes_and_Ketones.txt |
Synthesis of $\alpha$-Halo Ketones
Halogenation of saturated aldehydes and ketones usually occurs exclusively by replacement of hydrogens alpha to the carbonyl group:
The reagents that commonly are used to halogenate carbonyl compounds are those that are used to halogenate alkanes (e.g. $\ce{Cl_2}$, $\ce{Br_2}$, $\ce{SO_2Cl_2}$, and N-bromoamides; see Sections 4-4 and 14-3). However, the characteristics of the two types of halogenation normally are very different. 2-Propanone has been particularly well studied, and the important features of the halogenation of this compound are summarized as follows:
1. 2-Propanone reacts easily with chlorine, bromine, and iodine.
2. 2-Propanone reacts at the same rate with each halogen. Indeed, the rate of formation of the 1-halo-2-propanone is independent of the concentration of the halogen, even at very low halogen concentrations.
3. The halogenation of 2-propanone is catalyzed by both acids and bases. The rate expressions for formation of 1-halo-2-propanone in water solution are:
• at moderate $\ce{OH}^\ominus$ concentrations
$v = k \left[ \ce{CH_3COCH_3} \right] \left[ \ce{OH}^\ominus \right]$
• at moderate $\ce{H}^\oplus$ concentrations
$v = k' \left[ \ce{CH_3COCH_3} \right] \left[ \ce{H}^\oplus \right]$
The ratio of $k$ to $k'$ is 12,000, which means that hydroxide ion is a much more effective catalyst than is hydrogen ion.
The hydroxide ion is a much more effective catalyst than is hydrogen ion
To account for the role of the catalysts and the independence of the rate from the halogen concentration, the ketone necessarily must be slowly converted by the catalysts to something that can react rapidly with halogen to give the products. This something is either the enol or the enolate anion of 2-propanone:
As long as the first step is slow compared with the steps of Equations 17-2 and 17-3, the overall rate of reaction will be independent of both the concentration of halogen and whether it is chlorine, bromine, or iodine (cf. Section 4-4C).
The reaction of either the enol or the enolate anion (Equations 17-2 or 17-3) with $\ce{Br_2}$ resembles the first step in the electrophilic addition of halogens to carbon-carbon multiple bonds (Section 10-3A). However, the second step, addition of the nucleophilic halide, if it occurs at all, does not produce any stable product:
Unsymmetrical ketones, such as 2-butanone, can form two different enols that will react with halogens to give isomeric halo ketones:
The composition of the product mixture will depend on the relative rates of formation of the isomeric enols, provided that the halogenation step is not a reversible reaction. Barring any serious steric effects that influence the rate of reaction, the more rapidly formed enol generally is the more thermodynamically stable enol.
The Haloform Reaction
The previous discussion of the halogenation of ketones is incomplete in one important respect concerning base-induced halogenation. That is, once an $\alpha$-halo ketone is formed, the other hydrogens on the same carbon are rendered more acidic by the electron-attracting effect of the halogen and are replaced much more rapidly than the first hydrogen:
The result is that, if the monobromoketone is desired, the reaction is carried out best with an acidic catalyst rather than a basic catalyst. A further complication in the base-catalyzed halogenation of a methyl ketone is that the trihaloketone formed is attacked readily by base, thereby resulting in cleavage of a carbon-carbon bond:
This sequence is called the haloform reaction because it results in the production of chloroform, bromoform, or iodoform, depending upon the halogen used. The haloform reaction is a useful method for identification of methyl ketones, particularly when iodine is used, because iodoform is a highly insoluble, bright-yellow solid. The reaction also is very effective for the synthesis of carboxylic acids when the methyl ketone is more available than the corresponding acid:
Because the haloform reaction is fast, in some cases it can be used to prepare unsaturated acids from unsaturated ketones without serious complications caused by addition of halogen to the double bond:
A reaction somewhat similar to the cleavage of haloforms with hydroxide occurs with ketones that do not have $\alpha$-hydrogens through the action of sodium amide:
This reaction, called the Haller-Bauer reaction, has utility for the preparation of amides of the types $\ce{ArCONH_2}$ and tert-$\ce{RCONH_2}$, and, through hydrolysis, the corresponding carboxylic acids.
Reactions of $\alpha$-Halo Ketones
The halogen of an $\alpha$-halo aldehyde or an $\alpha$-halo ketone is exceptionally unreactive in $S_\text{N}1$-displacement reactions, but is exceptionally reactive in $S_\text{N}2$ displacements, compared with the halogen of alkyl halides having comparable potential steric effects. Similar behavior is observed with $\alpha$-halo carboxylic acids and is discussed in Chapter 18.
In some circumstances, the production of a 2-halo alcohol by reduction of the carbonyl group of an $\alpha$-halo ketone with metal hydrides is a useful synthetic reaction:
When one attempts $E2$ reactions with $\alpha$-halo ketones using strong bases such as alkoxides, an interesting rearrangement pathway may occur called the Favorskii rearrangement. In this reaction, the $\alpha$-halo ketone is converted to an ester. For example, 2-chlorocyclohexanone is converted to the methyl ester of cyclopentanecarboxylic acid by treatment with sodium methoxide in ether:
The mechanism of this reaction has been the subject of many investigations.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/17%3A_Carbonyl_Compounds_II-_Enols_and_Enolate_Anions._Unsaturated_and_Polycarbonyl_Compounds/17.03%3A_Halogenation_of_Aldehydes_and_Ketones.txt |
The Aldol Addition
A most important property of enolate anions, at least as far as synthesis is concerned, is their excellent nucleophilicity, which enables them to add to double bonds and to participate in nucleophilic substitution. When the addition is to a carbonyl double bond, it is called an aldol addition (Equation 17-4). Additions of enolate anions to carbon-carbon double bonds usually are classified as Michael additions (Equation 17-5), and these are discussed in Sections 17-5B and 18-9D. The principles of $S_\text{N}$ nucleophilic reactions of enolate anions (Equation 17-6) will be considered in Section 17-4, and their synthetic applications in detail in Chapter 18.
The products of aldol addition are $\beta$-hydroxy aldehydes (ald-ols) or $\beta$-hydroxyl ketones (ket-ols). A typical example is the reaction of ethanal with base and, if the conditions are reasonably mild, the product is 3-hydroxybutanal:
The overall reaction corresponds to a dimerization of ethanal, that is, an addition of one ethanal molecule to another with formation of a new carbon-carbon bond. The synthetic value of the reaction lies in the fact that it can be used to build large molecules from smaller molecules (see Section 13-7).
Formation of the enolate anion, $7$, by removal of an $\alpha$ hydrogen by base is the first step in the aldol addition:
The anion then adds to the carbonyl group of a second molecule of ethanal in a manner analogous to the addition of other nucleophiles to carbonyl groups (e.g., cyanide ion, Section 16-4A). The adduct so formed, $8$, rapidly adds a proton to the alkoxide oxygen to form the aldol, 3-hydroxybutanal. This last step regenerates the basic catalyst, $\ce{OH}^\ominus$:
Ambident Nature of Enolate Ions in Aldol Addition
The two possible valence-bond structures of the enolate anion, $7a$ and $7b$, show that the anion should act as an ambident nucleophile - a nucleophile with nucleophilic properties associated with both carbon and oxygen. The addition step in the aldol reaction therefore may be expected to take place in either of two ways: The anion could attack as a carbon nucleophile to form a carbon-carbon bond, $8$, leading ultimately to the aldol, $9$, or it might attack as an oxygen nucleophile to form a carbon-oxygen bond, thereby leading to the hemiacetal, $10$. By this reasoning, we should obtain a mixture of products $9$ and $10$. However, the aldol $9$ is the only one of these two possible products that can be isolated:
Why is only one of these products formed? To understand this, you must recognize that aldol reactions are reversible and therefore are subject to equilibrium rather than kinetic control (Section 10-4A). Although the formation of $10$ is mechanistically reasonable, it is not reasonable on thermodynamic grounds. Indeed, while the overall $\Delta H^0$ (for the vapor) calculated from bond energies is $-4 \: \text{kcal mol}^{-1}$ for the formation of the aldol, it is $+20.4 \: \text{kcal mol}^{-1}$ for the formation of $10$.$^2$ Therefore, the reaction is overwhelmingly in favor of the aldol as the more stable of the two possible products.
Position of the Equilibrium in Aldol Additions
The equilibrium constant is favorable for the aldol addition of ethanal, as in fact it is for most aldehydes. For ketones, however, the reaction is much less favorable. With 2-propanone (acetone) only a few percent of the addition product "diacetone alcohol", $11$ is present at equilibrium:
The 2-propanone is boiled and the hot condensate from the reflux condenser flows back over solid barium hydroxide contained in the porous thimble and comes to equilibrium with the addition product $11$. The barium hydroxide is retained by the porous thimble and the liquid phase returns to the boiler where the 2-propanone, which boils $110^\text{o}$ below the temperature at which $11$ boils, is selectively vaporized and returns to the reaction zone to furnish more adduct.
The key step in aldol addition requires an electron-pair donor (nucleophile) and an electron-pair acceptor (electrophile). In the formation of 3-hydroxybutanal or $11$, both roles are played by one kind of molecule, but there is no reason why this should be a necessary condition for reaction. Many kinds of mixed aldol additions are possible.
Consider the combination of methanal and 2-propanone. Methanal cannot form an enolate anion because it has no $\alpha$ hydrogens. However, it is expected to be a particularly good electron-pair acceptor because of freedom from steric hindrance and the fact that it has an unusually weak carbonyl bond ($166 \: \text{kcal}$ compared to $179 \: \text{kcal}$ for 2-propanone). In contrast, 2-propanone forms an enolate anion easily but is relatively poor as the electrophile. Consequently the addition of 2-propanone to methanal should and does occur readily:
The problem is not to get addition, but rather to keep it from going too far. Indeed, all six $\alpha$ hydrogens of 2-propanone can be replaced easily by $\ce{-CH_2OH}$ groups:
A commercially important mixed addition involves ethanal and an excess of methanal in the presence of calcium hydroxide. Addition occurs three times and the resulting trihydroxymethylethanal (which has no $\alpha$ hydrogens) undergoes a "crossed Cannizzaro" reaction with more methanal to give a tetrahydroxy alcohol known as "pentaerythritol::
Pentaerythritol is used widely in the preparation of surface coatings and in the formation of its tetranitrate ester, pentaerythrityl tetranitrate [PETN, $\ce{C(CH_2ONO_2)_4}$], which is an important high explosive.
Dehydration of Aldol Addition Products
An important property of aldol addition products is the ease with which they eliminate water in the presence of either acids or bases. For example, when 3-hydroxybutanal is heated in the basic solution in which it is formed (by aldol addition of ethanal), 2-butenol results:
The ease of dehydration compared with simple alcohols is related to the fact that the product is a conjugated alkenone. The stabilization energy of the conjugated system makes the equilibrium constant for dehydration especially favorable. In many cases the aldol adduct is only an intermediate in aldol reactions because it dehydrates more rapidly than it can be isolated. Such is most often the case when the dehydration product is a polyunsaturated conjugated aldehyde or ketone. 2-Propanone and bezenecarbaldehyde (benzaldehyde), for instance, give the unsaturated ketone $12$ in cold aqueous sodium hydroxide solution:
Although the equilibrium for aldol addition may be unfavorable, when dehydration of the aldol product is rapid, $\ce{C-C}$ bond formation may be pushed to completion by conversion of the aldol to the $\alpha$,$\beta$-unsaturated ketone.
The mechanism of base-catalyzed dehydration of aldols involves formation of an enolate anion by removal of a proton from the $\ce{C2}$ or alpha carbon and subsequent elimination of the hydroxyl group as hydroxide ion:
This last step is one of the rare examples in which the leaving group is $\ce{OH}^\ominus$. Generally, hydroxide is a poor leaving group in substitution ($S_\text{N}1$ or $S_\text{N}2$) or elimination ($E1$ or $E2$) reactions (see Section 8-7C).
Dehydration of aldols to $\alpha$,$\beta$-unsaturated carbonyl compounds usually is achieved best with acidic catalysts. An example is the dehydration of the aldol from 2-propanone to give 4-methyl-3-penten-2-one:
If this reaction were attempted under basic conditions, extensive reversion of the aldol to 2-propanone would occur (see Section 17-3C). Under acidic conditions, however, the process is a straightforward proton transfer to oxygen followed by elimination of water and proton transfer from carbon:
The Use of Aldol Addition Reactions in Synthesis
Aldol reactions provide a valuable synthetic method for forming carbon-carbon bonds. They can be adapted to extend the length of a carbon chain, to form cyclic compounds, and to provide intermediates that can be transformed into more useful materials. An important feature of these intermediates is that functional groups useful for later reactions are located close to or on the carbons of the newly formed $\ce{C-C}$ bond. There is an almost bewildering number of variations on the aldol reaction and we shall not mention all of them. The main thing to recognize in all of these reactions is that the acceptor molecule always is a carbonyl compound, best an aldehyde, sometimes a ketone, even an ester (see Section 18-8E). The donor molecule is some type of carbanion; usually, but not always, an enolate anion. However, any substance that has a $\ce{C-H}$ acidity in the p$K_\text{a}$ range of 25 or less can be converted easily to a carbanion, which in principle may serve as the donor in aldol additions. Examples are listed in Table 17-1 and include not only aldehydes and ketones but esters, nitriles, and nitro compounds. The use of a nitroalkane in aldol addition is shown in the following sequence. The use of esters as the donor is discussed further in Section 18-8E.
Cyclic products can be formed by aldol additions provided the donor carbanion and acceptor carbonyl are part of the same molecule. For example, consider how the synthesis of 3-methyl-2-cyclohexanone could be achieved from acyclic substances. The carbon-carbon bond formed in this process of aldol addition closes the ring and ultimately becomes the double bond in the conjugated system when the aldol product undergoes dehydration. Working backwards, we have the sequence
and the starting material for the synthesis therefore is 2,6-heptanedione. Because $\Delta G^0$ for the formation of aldol products is not very favorable, cyclizations involving aldol reactions usually will not proceed to give strained carbocyclic rings.
The industrial importance of aldol reactions is in the synthesis of alcohols, especially 1-butanol and 2-ethyl-1-hexanol:
Notice that the combination of hydroformylation (Section 16-9F), aldol addition, dehydration, and hydrogenation takes a simple alkene (propene) to an alcohol with more than twice as many carbons.
A Biological Aldol Addition
One of the reactions in the metabolism of carbohydrates by the glycolic pathway is a type of aldol addition. In this reaction $D$-fructose (as the 1,6-diphosphate ester) is formed from $D$-glyceraldehyde and 1,3-dihydroxypropanone (both as monophosphate esters). The process is readily reversible and is catalyzed by an enzyme known as aldolase:
It seems likely that this reaction could occur in quite the same way as in the laboratory aldol reactions discussed so far, because the enolate anion of the donor molecule (dihydroxypropanone) is not expected to be formed in significant amount at the pH of living cells. In fact, there is strong evidence that the enzyme behaves as an amino $\left( \ce{ENH_2} \right)$ compound and reacts with the carbonyl group of dihydroxypropanone to form an imine, analogous to the reactions described in Section 16-4C:
This implies that the imine form of dihydroxypropanone is a key intermediate in the overall aldol-type addition.
How can the imine behave as the carbon donor in addition to the aldehyde carbonyl of glyceraldehyde 3-phosphate? It is unlikely to do so directly, but it can rearrange to an enamine which, as we will explain in Section 17-4B, can act as a carbon nucleophile:
Attack of the nucleophilic carbon of the enamine at the aldehyde carbonyl of glyceraldehyde 3-phosphate forms the aldol of the imine which, on hydrolysis, gives the aldol and regenerates the enzyme:
By using the neutral enamine as the carbon nucleophile rather than an enolate anion, the biological system avoids the need for strongly basic reaction conditions in aldol addition.
$^2$This value probably is too large by $3$ to $4 \: \text{kcal}$, because resonance stabilization of alkoxyalkanes has been ignored in this calculation.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/17%3A_Carbonyl_Compounds_II-_Enols_and_Enolate_Anions._Unsaturated_and_Polycarbonyl_Compounds/17.04%3A_Nucleophilic_Addition_Reactions_of_Enolat.txt |
Alkylation of Ketones
The synthetic chemistry of enolate anions is centered on their nucleophilic and basic properties. Accordingly these ions participate in $S_\text{N}2$ reactions with suitable alkyl compounds:
However, there are a number of complicating factors to consider. First, the basic conditions needed to form the enolate ions often lead to side reactions such as aldol addition and $E2$ elimination of $\ce{RX}$ compounds. Aldol addition is minimized if the carbonyl compound is a ketone with a structure unfavorable for aldol addition or if all of the carbonyl compound is converted to its enolate. To convert all of a simple carbonyl to its enolate usually requires a very strong base, such as $\ce{NH_2^+}$ in an aprotic solvent or liquid ammonia. Because the enolate anion itself is a strong base, best results are obtained when the halide, $\ce{RX}$, does not undergo $E2$ reactions readily.
The second complication arises if the alkyl compound reacts with both carbon and oxygen of the nucleophilic enolate anion. The carbon product is the result of "$\ce{C}$-alkylation", whereas the oxygen product is the result of "$\ce{O}$-alkylation":
The possibility of the enolate anion acting as if its charge were effectively concentrated on carbon or on oxygen was discussed previously in connection with aldol addition (Section 17-3B). However, the situation there was quite different from the one here, because aldol addition is easily reversible, whereas alkylation is not. Furthermore, while the aldol reaction involving $\ce{C-O}$ bond formation is unfavorable $\left( \Delta H^0 = + 20 \: \text{kcal mol}^{-1} \right)$ compared to $\ce{C-C}$ bond formation $\left( \Delta H^0 = - 4 \: \text{kcal mol}^{-1} \right)$, both $\ce{O}$- and $\ce{C}$-alkylation of the anion have $\Delta H^0 < 0$.
Whether $\ce{C}$- or $\ce{O}$-alkylation predominates depends on kinetic control (Section 10-4A). It is not a simple matter to predict which of the two positions of the enolate will be more nucleophilic, and in fact, mixtures of products often are obtained in distributions that depend on the solvent used, the temperature, the nature of $\ce{X}$, and the nature of the base employed to form the anion. $\ce{O}$-Alkylation tends to occur with ketones of high enol content (which usually means that the enolate anion will have especially high charge density on oxygen) and with alkylating agents possessing a high degree of $S_\text{N}2$ reactivity.
There is another correlation that seems to have validity in many situations, at least where kinetic control is dominant; namely, the freer (less associated) the ambident anions is from its cation, the more likely is the electrophile to attack the atom of the anion with the highest negative charge. Thus $\ce{O}$-alkylation of the sodium enolate of 2-propanone is favored in aprotic solvents that are good at solvating cations [such as $\ce{(CH_3)_2SO}$, Section 8-7F].
In the alkylation of unsymmetrical ketones, formation of more than one enolate anion is possible, and when this occurs, mixtures of products are obtained. Thus,
However, when one of the possible enolate anions is especially stabilized, either by conjugation or by strong electron-withdrawing groups, that enolate usually is the dominant form and only one product is formed. Thus 2,4-pentanedione is methylated at $\ce{C3}$, not at $\ce{C1}$:
Alkylation of Enamines
Enamines (Section 16-4C), like enolate anions, have two reactive positions and, in principle, can give either $\ce{N}$- or $\ce{C}$-alkylation.
Both products may be formed, but they can be separated readily because, on treatment with dilute acid, only the $\ce{C}$-alkylation product hydrolyzes to a ketone. Generally, the alkylated ketone is the desired product:
Alkylation of enamines therefore is a feasible, and sometimes much more useful, alternative to the direct alkylation of ketones because it proceeds under less strongly basic conditions. The sequence starts with conversion of a ketone to an enamine, $\overset{\ce{RNH_2}}{\longrightarrow}$ $+ \ce{H_2O}$, followed by $\ce{C}$-alkylation of the enamine, $\overset{\ce{RX}}{\longrightarrow}$ $+ \ce{X}^\ominus$, and ends with hydrolysis to the alkylated ketone, $+ \ce{H_2O} \rightarrow$ $+ \ce{R_2NH_2^-}$. A typical example of the use of enamines for alkylation of a ketone follows:
Several important biological reactions utilize enamine intermediates as carbon nucleophiles in $\ce{C-C}$ bond-forming reactions. One example is discussed in Section 17-3F.
Alkylation of Sulfur-Stabilized Carbanions
The chemistry of carbanions stabilized by groups other than carbonyl functions is closely analogous to the chemistry of enolate anions. We have seen that $\ce{C-H}$ acidity of compounds with the structural feature can be significant (p$K_\text{a}$ of 25 or less) when $\ce{X}$ is an atom or group that can effectively delocalize the negative charge on carbon in . Typical $\ce{X}$ groups are $\ce{C=O}$, $\ce{C \equiv N}$, $\ce{PR_3^+}$, $\ce{SR_2^+}$, $\ce{SO_2R}$, and $\ce{SR}$. Consequently, we can expect that carbanions of the type , when formed, will resemble enolate anions and will undergo addition reactions to $\ce{C=O}$ and $\ce{C=C}$, and will be alkylated with halides of good $S_\text{N}2$ reactivity. In fact, the reactions of ylides discussed in Section 16-4A are examples of the addition of phosphorus-, sulfur-, and nitrogen-stabilized carbanions to carbonyl groups.
Sulfur in its higher oxidation states (e.g., sulfone, $\ce{-SO_2}-$) is especially effective in stabilizing adjacent carbanion centers. However, from a synthetic standpoint there are disadvantages to the sulfone grouping in that the better stabilized carbanions also are the least reactive, and subsequent removal of the sulfone grouping can be difficult. A good balance between carbanion stability, carbanion reactivity, and ease of $\ce{C-S}$ bond cleavage is present in the structures $\ce{RS-CH_2-SR}$ and . This is illustrated below for a strikingly simple concept for preparing cyclobutanone, in which the ring carbons are derived from methanal and 1,3-dibromopropane:
To achieve this synthesis, the methanal first is converted to a thioketal, which then is partially oxidized to give $13$. Treatment of $13$ with a strong base converts it to the carbanion, which can be readily alkylated. By using 1,3-dibromopropane and two equivalents of base, a double displacement forms the cyclic product, $14$. The sulfur groups of $14$ can be removed easily by acid hydrolysis to give cyclobutanone:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/17%3A_Carbonyl_Compounds_II-_Enols_and_Enolate_Anions._Unsaturated_and_Polycarbonyl_Compounds/17.05%3A_Nucleophilic_Substitution_with_Enolate_An.txt |
Structure and Spectral Properties
The most generally useful preparation of $\alpha$,$\beta$-unsaturated carbonyl compounds is by dehydration of aldol addition products, as described in Section 17-3D. Conjugation of the carbonyl group and double bond has a marked influence on spectroscopic properties, particularly on ultraviolet spectra, as the result of stabilization of the excited electron states, which for $\pi \rightarrow \pi^*$ transitions can be described in terms of important contributions of polar resonance structures (see Sections 9-9B and 16-3B):
Such resonance is much less important in the ground state but is still sufficiently important to account for the moderate differences in dipole moments between saturated and $\alpha$,$\beta$-unsaturated aldehydes and ketones; for example,
The effect of conjugation also is reflected in infrared carbonyl frequencies (Section 16-3) and NMR spectra. With respect to the latter, it is found that the protons on the $\beta$ carbon of $\alpha$,$\beta$-unsaturated carbonyl compounds usually come at $0.7$ to $1.7 \: \text{ppm}$ lower fields than ordinary alkenic protons. The effect is smaller for the $\alpha$ protons.
Addition Reactions
There are many addition reactions of $\alpha$,$\beta$-unsaturated aldehydes, ketones, and related compounds that are the same as the carbonyl addition reactions described previously. Others are quite different and result in addition to the alkene double bond. Organometallic compounds are examples of nucleophilic reagents that can add to either the alkene or the carbonyl bonds of conjugated ketones (Section 14-12D). Hydrogen cyanide behaves likewise and adds to the carbon-carbon double bond of 3-butene-2-one, but to the carbonyl group of 2-butenal:
All of these reactions may be classified as nucleophilic additions, but when addition occurs at the alkene bond, the orientation always is such that the nucleophile adds at the $\beta$ carbon. An example is the addition of methanol catalyzed by sodium methoxide:
Nucleophilic reagents normally do not attack carbon-carbon double bonds (Section 10-6). The adjacent carbonyl function therefore must greatly enhance the reactivity of the double bond toward such reagents. This enhancement is not surprising when it is realized that the attack of a nucleophile produces a stabilized enolate anion:
The products are formed from the enolate intermediate by proton transfer to either carbon or oxygen. If the proton adds to oxygen the enol is formed, which is unstable with respect to the ketone and ultimately will rearrange:
Reactions of this type are referred to in a variety of terms, many of which are rather confusing and nondescriptive. They sometimes are classified as 1,4-additions, implying that addition occurs across the terminal protons of the conjugated system. A synonymous term is conjugate addition. When the nucleophile is a carbanion, the reaction is called a Michael addition. Thus, by this definition, Equation 17-7 represents a Michael addition. Another, perhaps more typical, example is the addition of an enolate to a conjugated ketone:
Michael-type additions, like aldol additions, are useful for the formation of carbon-carbon bonds.
Electrophilic addition of hydrogen halides to $\alpha$,$\beta$-unsaturated aldehydes and ketones places the halogen on the $\beta$ carbon. This orientation is opposite to that observed for related additions to conjugated dienes:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/17%3A_Carbonyl_Compounds_II-_Enols_and_Enolate_Anions._Unsaturated_and_Polycarbonyl_Compounds/17.06%3A_-Unsaturated_Aldehydes_and_Ketones.txt |
Preparation of Ketenes
Substances with cumulated carbonyl and carbon-carbon double bonds, , are called ketenes and, as may be expected, have interesting and unusual properties. Ketene itself, $\ce{CH_2=C=O}$, and its monosubstitution products, $\ce{RCH=C=O}$ ($\ce{R} =$ alkyl or aryl), are called aldoketenes, whereas disubstituted ketenes, $\ce{R_2C=C=O}$, are called ketoketenes.
There are relatively few general methods for preparing ketenes. The simplest procedure is to treat an $\alpha$-bromoacyl bromide with zinc, but the yields usually are not very good:
Several special methods are available for the preparation of ketene itself. The most convenient laboratory preparation is to pass 2-propanone vapor over a coil of resistance wire heated electrically to a dull red heat; air is excluded to avoid simple combustion:
The weakest bonds are the $\ce{C-C}$ bonds and, at $750^\text{o}$, fragmentation yields a methyl radical and an ethanoyl radical:
Transfer of a hydrogen atom (i.e., disproportionation) gives methane and ketene. Industrially, ketene is best prepared by dehydration of ethanoic acid:
Reactions of Ketenes
Ketene has a boiling point of $-56^\text{o}$ and normally would be stored under pressure in steel cylinders. However, this is not possible because ketene is unstable with respect to formation of a dimer known as "diketene":
The dimer also is a highly reactive substance with such unusual characteristics that its structure was not firmly established until 1956, almost 48 years after it first was prepared.
Ketenes in general are useful reagents for acylating alcohols, $\ce{ROH}$, and amines, $\ce{RNH_2}$, because the reactions involve additions; there are no by-products to be separated:
Ketenes also can be used for the synthesis of cyclobutane derivatives through [2 + 2] cycloadditions with suitably active alkenes (Section 13-3D):
Diketene is very useful in synthesis, particularly through its reactions with alcohols and amines to give derivatives of 3-oxobutanoic acid:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
17.08: 12-Dicarbonyl Compounds
Some typical and important members of this class have structures as follows:
Most of the 1,2-dicarbonyl compounds are yellow. Ethanedial is unusual in being yellow in the liquid state, but green in the vapor state. It has very reactive aldehyde groups and is employed in the manufacture of plastics and as a component of embalming fluids to harden proteins by linking together their amino groups through imine formation:
Ethanedial undergoes an internal Cannizzaro reaction with alkali to give hydroxyethanoic (glycolic) acid:
An analogous reaction occurs with diphenylethanedione, which results in carbon-skeleton rearrangement. This is one of the few carbon-skeleton rearrangements brought about by basic reagents, and is known as the "benzilic acid rearrangement".
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/17%3A_Carbonyl_Compounds_II-_Enols_and_Enolate_Anions._Unsaturated_and_Polycarbonyl_Compounds/17.07%3A_Ketenes.txt |
Much of the chemistry of 1,3-dialdehydes, aldehyde ketones, and diketones already has been mentioned in this chapter and is well illustrated in the properties of 2,4-pentanedione,
The liquid ketone exists $85\%$ in the enol form and is moderately acidic. The $K_\text{a}$ in water is $\cong 10^{-9}$. The enol form is stabilized significantly by both electron delocalization and hydrogen bonding. The amount of enol present at equilibrium depends on the solvent, and is smallest in hydrogen-bonding solvents and largest in nonpolar solvents such as carbon tetrachloride.
The reactions discussed in this chapter that depend on the formation of enolate anions (i.e., halogenation, aldol addition, and alkylation) often proceed smoothly and under milder conditions with 1,3-diketones than with monoketones. This is because the 1,3-diketones are stronger acids and therefore can form the enolate anion with weaker bases. The principal synthetic methods for preparing 1,3-dicarbonyl compounds will be discussed in Chapter 18.
With 2,4-pentanedione, polyvalent metal cations often form very stable and only slightly polar enolate salts, better known as metal chelates. Cupric ion forms a particularly stable dark-blue chelate:
The beryllium chelate of 2,4-pentanedione is another example of a stable chelate; it melts at $108^\text{o}$, boils at $270^\text{o}$, and is soluble in many organic solvents. By replacing the methyl groups of 2,4-pentanedione with tert-butyl groups, a diketone is obtained which, with many metals including transition and rare-earth metals, forms complexes that often are highly soluble in nonpolar organic solvents. The interior of these chelates is saltlike but the exterior is hydrocarbonlike and nonpolar, which accounts for the substantial solubility in nonpolar solvents.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
17.10: 14-Dicarbonyl Compounds
Most of the reactions of the 1,4-dicarbonyl compounds are the conventional reactions expected for isolated carbonyl groups. An important exception is formation of azacylclopentadiene (pyrrole) derivatives from 1,4-dicarbonyl compounds and ammonia or primary amines:
These reactions are reasonably general and also can be used to prepare compounds with oxygen and sulfur in five-membered rings.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
17.11: Tricarbonyl Compounds
The properties of tricarbonyl compounds are for the most part as expected, except when the three groups are contiguous to one another, as in diphenylpropanetrione. With such compounds, the central carbonyl group is highly reactive; it is lost, as carbon monoxide, in the presence of acidic catalysts such as aluminum chloride, and adds water readily to give a monohydrate:
We shall consider the hydrate of the cyclic triketone, \(18\), known as "ninhydrin", later in connection with amino acids:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
17.12: Cyclopropanones and Cyclopropenones
Cyclopropanones deserve special comment, not because of their practical importance (they have no commercial value at this time), but because of their novel behavior and reactivity. No unambiguous synthesis of cyclopropanones was known prior to 1965, and the older textbooks usually contained statements such as "cyclopropanones apparently cannot exist". However, they had been postulated as intermediates in various reactions (see, for example, the Favorskii rearrangement, Section 17-2C), but until recently had defied isolation and identification. The problem is that the three-ring ketone is remarkably reactive, especially towards nucleophiles. Because of the associated relief of angle strain, nucleophiles readily add to the carbonyl group without the aid of a catalyst and give good yields of adducts from which the cyclopropanone is not easily recovered:
To avoid destructive side reactions, cyclopropanones have to be prepared at low temperatures in the absence of nucleophiles. A good example is the synthesis of cyclopropanone itself from ketene and diazomethane (see Section 16-4A):
When seemingly simple organic structures defy isolation, this usually stimulates many theoretical and experimental studies in an effort to rationalize anomalous behavior. In the case of cyclopropanone, the possibility was considered that the molecule might preferably exist as an open-chain dipolar structure rather than as the cyclic ketone:
Although the spectral properties of cyclopropanones and the easy formation of hydrates and hemiketals are inconsistent with the dipolar form, some reactions of cyclopropanones do indicate that the ring carbons are much more electrophilic than in other cyclic or acyclic ketones. For example, nucleophilic ring opening often occurs easily:
Also, both [3 + 4] cycloadditions of cyclopropanone to dienes and [3 + 2] additions to carbonyl groups have been observed. These reactions seem easiest to understand if cyclopropanone can behave as if it had, or could be converted to, a dipolar open-chain structure:
Cyclopropenone has been prepared by a route that illustrates the value of the acetal grouping in protecting ketone groups (Section 16-8):
Cyclopropenone undergoes many interesting reactions - one example is Diels-Alder addition, the product of which in methanol solution is a hemiketal. That the hemiketal is favored for the adduct, but not for cyclopropenone, indicates that the double bond of cyclopropenone has a considerable effect on the reactivity of the carbonyl group.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/17%3A_Carbonyl_Compounds_II-_Enols_and_Enolate_Anions._Unsaturated_and_Polycarbonyl_Compounds/17.09%3A_13-Dicarbonyl_Compounds.txt |
Almost all of the basic types of reactions now have been covered: addition, elimination, substitution, and rearrangement by polar, radical, and concerted mechanisms. Indeed, if you have been looking for similarities, you will have seen that most of the reactions discussed in the preceding three chapters are variations on basic types we have discussed earlier. Furthermore, most of the basic structural effects that determine chemical reactivity also have been covered in previous chapters: bond energies, steric hindrance, electronegativity, electron delocalization, hydrogen bonding, solvation, and conformational influences.
You might well ask what is left. The answer is, a great deal - but now we will be concerned mostly with putting concepts together, moving from the simple to the complex. For example, in this chapter we will be trying to understand the ways that carboxylic acids, which possess the $\ce{-COOH}$ functional group, are similar to and different from alcohols, which have the $\ce{-OH}$ group, and aldehydes and ketones, which have $\ce{C=O}$ bonds.
• 18.1: Prelude to Carboxylic Acids and Their Derivatives
Almost all of the basic types of reactions now have been covered: addition, elimination, substitution, and rearrangement by polar, radical, and concerted mechanisms. Indeed, if you have been looking for similarities, you will have seen that most of the reactions discussed in the preceding three chapters are variations on basic types we have discussed earlier.
• 18.2: Physical Properties of Carboxylic Acids
Carboxylic acids show a high degree of association through hydrogen bonding. Carboxylic acids have substantially higher melting points and boiling points of acids relative to alcohols, aldehydes, ketones, and chlorides can be attributed to the strength and degree of hydrogen bonding. Hydrogen bonding also is responsible for the high water solubility of the simple carboxylic acids with less than five carbons; water molecules can solvate the carbonyl group through hydrogen bonds.
• 18.3: Some Chemical Properties of Carboxylic Acids
Most of the reactions of carboxylic acids belong to one of four principal classes, depending on the point in the molecule where the reaction occurs. (1) Reactions involving the -O-H bond, (2) Reactions at the carbonyl bond, (3) Decarboxylation, and (4) Substitution on the R group.
• 18.4: Reactions at the Carbonyl Carbon of Carboxylic Acids
Many important reactions of carboxylic acids involve attack on the carbon of the carbonyl group by nucleophilic species. These reactions frequently are catalyzed by acids, because addition of a proton or formation of a hydrogen bond to the carbonyl oxygen makes the carbonyl carbon more vulnerable to nucleophilic attack.
• 18.5: Decarboxylation of Carboxylic Acids
The decarboxylation of RCO2H to give RH and CO2 can be calculated from bond energies and the stabilization energy of the carboxyl group to have ΔH°=−7 kcal/mol. This does not mean that the reaction goes easily. Special structural features are required. The simple aliphatic carboxylic acids do not lose carbon dioxide on heating, but when there are strongly electron-attracting groups attached to the αα carbon, decarboxylation often proceeds readily at 100 - 150°.
• 18.6: Reactions at the $\alpha$ Carbons of Carboxylic Acids
The halogen of an α-haloalkanoic acid is replaced readily by nucleophilic reagents. Thus a variety of αα -substituted carboxylic aids may be prepared by reactions that are analogous to SN2 substitution of alkyl halides. However, The SN1 reactivity of α-haloalkanoic acids is particularly low.
• 18.7: Functional Derivatives of Carboxylic Acids
A functional derivative of a carboxylic acid is a substance formed by replacement of the hydroxyl group of the acid by some other group, X , such that it can be hydrolyzed back to the acid.
• 18.8: Reactions at the Carbonyl Carbon of Acid Derivatives
Hydrolysis of most acid derivatives to the parent acids is acid- or base-catalyzed. Carboxylic acid derivatives also react with organomagnesium and organolithium compound
• 18.9: Reactions at the $\alpha$ Carbons of Carboxylic Acid Derivatives
Many important synthetic reactions in which C−CC−C bonds are formed involve esters and are brought about by basic reagents. This is possible because the αα hydrogens of an ester are weakly acidic, and a strong base, such as sodium ethoxide, can produce a significant concentration of the ester anion at equilibrium. The acidity of αα hydrogens is attributed partly to the electron-attracting inductive effects of the ester oxygens, and partly to resonance stabilization of the resulting anion.
• 18.10: Reactions of Unsaturated Carboxylic Acids and Their Derivatives
Unsaturated carboxylic acids of the type RCH=CH(CH2)nCOOH usually exhibit the properties characteristic of isolated double bonds and isolated carboxyl groups when n is large and the functional groups are far apart. As expected, exceptional behavior is found most commonly when the groups are sufficiently close together to interact strongly. We shall emphasize those properties that are exceptional.
• 18.11: Dicarboxylic Acids
Acids in which there are two carboxyl groups separated by a chain of more than five carbon atoms (n>5) for the most part have unexceptional properties, and the carboxyl groups behave more or less independently of one another. However, when the carboxyl groups are closer together the possibilities for interaction increase; we shall be interested primarily in such acids.
• 18.12: Methods of Preparation of Carboxylic Acids and Their Derivatives
• 18.E: Carboxylic Acids and Their Derivatives (Exercises)
These are the homework exercises to accompany Chapter 18 of the Textmap for Basic Principles of Organic Chemistry (Roberts and Caserio).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
18: Carboxylic Acids and Their Derivatives
Almost all of the basic types of reactions now have been covered: addition, elimination, substitution, and rearrangement by polar, radical, and concerted mechanisms. Indeed, if you have been looking for similarities, you will have seen that most of the reactions discussed in the preceding three chapters are variations on basic types we have discussed earlier. Furthermore, most of the basic structural effects that determine chemical reactivity also have been covered in previous chapters: bond energies, steric hindrance, electronegativity, electron delocalization, hydrogen bonding, solvation, and conformational influences.
You might well ask what is left. The answer is, a great deal - but now we will be concerned mostly with putting concepts together, moving from the simple to the complex. For example, in this chapter we will be trying to understand the ways that carboxylic acids, which possess the $\ce{-COOH}$ functional group, are similar to and different from alcohols, which have the $\ce{-OH}$ group, and aldehydes and ketones, which have $\ce{C=O}$ bonds.
Subsequently we will look at acids that also possess $\ce{OH}$ or $\ce{NH_2}$ substituent groups (or both) and develop a rationale for the behavior of these combinations in terms of effects we already have discussed. Insofar as possible, you should try to do this yourself whenever you encounter a substance with a new set of combinations or functional groups on its molecules. You often will be in error (as many experts will be), because even in you take account of all of the structural effects, as well as the possible reactions or interactions, the overall result of these frequently is very difficult to judge in advance. In one case, steric hindrance may dominate, in another, electron delocalization, and so on. Still, trying to assess the effects and possible reactions leads to understanding and recognition of what the alternatives are, even if the resultant of them is difficult to assess.$^1$ Continuing study can be expected to develop an instinct for what is "good" chemistry and what is not.
We have described previously the acidic properties of several types of compounds: alkynes, alkenes, and alkanes (Sections 11-8 and 13-5B); halides (Section 14-7B); alcohols (Section 15-4A); and carbonyl compounds (Section 17-1A). Now we come to compounds that we actually call acids - the carboxylic acids, $\ce{RCO_2H}$. Are these acids different in kind, or only in degree, from other acidic compounds discussed before? This is not a simple question and deserves some thought. In the most widely used sense, acids are proton donors but, as we have seen, their abilities to donate a proton to water vary over an enormous range: $\ce{CH_4}$ has a $K_a$ of $< 10^{-40}$, whereas $\ce{HI}$ has a $K_a$ of $\sim 10^9$. This represents a difference in ionization energies of more than $70 \: \text{kcal mol}^{-1}$. The differences in $K_a$ are only differences in degree, because examples are available of acids with $K_a$ values in all parts of the range of $K_a$ values. An important difference in kind was mentioned in Section 17-1B, namely, that acids with the same $K_a$ values can differ greatly in the rates at which they give up a proton to a given base, such as water. Carbon acids, in which the proton comes from a $\ce{C-H}$ bond, may react more than $10^{10}$ times slower than an oxygen acid with the same $K_a$ in which the proton is give up from an $\ce{O-H}$ bond.
Tradition reserves the use of the name "acid" for substances that transfer protons measurably to water. Thus the carboxylic acids stand out from alkynes, halides, alcohols, and simple aldehydes and ketones in giving water solutions that are "acidic" to indicator papers or pH meters as the result of proton transfers from the carboxylic groups:
$\ce{RCO_2H} + \ce{H_2O} \rightleftharpoons \ce{RCO_2^-} + \ce{H_3O^+}$
Even so, carboxylic acids are not very strong acids and, in a $1 \: \text{M}$ water solution, a typical carboxylic acid is converted to ions to the extent of only about $0.5\%$.
The nomenclature of carboxylic acids and their derivatives was discussed in Section 7-6. Many carboxylic acids have trivial names and often are referred to as "fatty acids". This term applies best to the naturally occurring straight-chain saturated and unsaturated aliphatic acids, which, as esters, are constituents of the fats, waxes, and oils of plants and animals. The most abundant of these fatty acids are palmitic, stearic, oleic, and linoleic acids. They occur as glycerides, which are esters of the trihydric alcohol, 1,2,3-propanetriol (glycerol):
Fats or glycerides belong to a class of biomolecules known as lipids. The traditional definition of a lipid is a water-insoluble organic compound that can be extracted from cells and tissues by nonpolar solvents (chloroform, ether, benzene). Compounds that meet this definition are substantially hydrocarbonlike, although they may differ widely in structure. They include not only esters of long-chain fatty acids but steroids (Section 30-4), terpenes (Section 30-3), and prostaglandins (Section 30-7). These nonpolar substances serve a variety of different biological functions, but lipids derived from fatty acids are most important for energy storage (Sections 18-8F and 20-10) and as components of membranes. Phosphoglycerides represent an important class of membrane lipids derived from glycerol. The hydroxyl groups of glycerol are esterified with two fatty-acid chains and one phosphate-ester residue. One of the phosphate-ester groups carries a highly polar $\ce{-N(CH_3)_2^+}$ group, the importance of which is indicated in Section 18-2F:
Hydrolysis of fats with alkali (e.g., sodium hydroxide) yields salts of the fatty acids, and those of the alkali metals, such as sodium or potassium, are useful as soaps:
The properties of salts of long-chain carboxylic acids that make them useful as soaps will be discussed in Section 18-2F.
General methods for the preparation of carboxylic acids are summarized in Table 18-5, at the end of the chapter.
$^1$The major problem with assessing the resultant to be expected from opposing factors in chemical reactions is that relatively small energy differences can cause great differences in which product is favored. For an equilibrium such as $\ce{A} \rightleftharpoons \ce{B}$ at $25^\text{o} \text{C}$, a $5.5 \: \text{kcal mol}^{-1}$ change in $\Delta G^0$ (Section 4-4A) can cause the equilibrium to shift from $99\%$ in favor of $\ce{A}$ to $99\%$ in favor of $\ce{B}$.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/18%3A_Carboxylic_Acids_and_Their_Derivatives/18.01%3A_Prelude_to_Carboxylic_Acids_and_Their_Derivatives.txt |
Hydrogen Bonding
Carboxylic acids show a high degree of association through hydrogen bonding. We have encountered such bonding previously with alcohols; however, acids form stronger hydrogen bonds than alcohols because their $\ce{O-H}$ bonds are more strongly polarized as $\ce{-} \overset{\delta \ominus}{\ce{O}} \ce{-} \overset{\delta \oplus}{\ce{H}}$. Furthermore, carboxylic acids are able to form hydrogen bonds to the negative oxygen of the carbonyl dipole rather than just to the oxygen of another hydroxyl group. Carboxylic acids in the solid and liquid states mostly exist as cyclic dimers, and these dimeric structures persist to some extent even in the vapor state and in dilute solution in hydrocarbon solvents:
Table 18-1: Physical Properties of Representative Carboxylic Acids
Hydrogen bonding also is responsible for the high water solubility of the simple carboxylic acids with less than five carbons; water molecules can solvate the carbonyl group through hydrogen bonds. Nonetheless, as the chain length of the hydrocarbon residue $\ce{R}$ increases, the solubility decreases markedly, because the proportion of polar to nonpolar groups becomes smaller.
Spectra of Carboxylic Acids
The infrared spectra of carboxylic acids provide clear evidence for the hydrogen bonding discussed in the preceding section. This is illustrated in Figure 18-2, which shows the spectrum of ethanoic acid in carbon tetrachloride solution, together with those of ethanol and ethanal for comparison.
The spectrum of ethanol has two absorption bands that are characteristic of the $\ce{OH}$ bond; one is a sharp band at $3640 \: \text{cm}^{-1}$, which corresponds to free or unassociated hydroxyl groups, and the other is a broad band centered on $3350 \: \text{cm}^{-1}$ due to hydrogen-bonded groups. The spectrum of ethanoic acid shows no absorption from free hydroxyl groups but, like that of ethanol, has a very broad intense absorption ascribed to associated $\ce{OH}$ groups. However, the frequency of absorption, $3000 \: \text{cm}^{-1}$, is shifted appreciably from that of ethanol and reflects stronger hydrogen bonding than in ethanol. The absorption due to the carbonyl group of ethanoic acid $\left( 1740 \: \text{cm}^{-1} \right)$ is broad, but is at about the same position as the carbonyl absorption in ethanal.
The carboxyl function does absorb ultraviolet radiation, but the wavelengths at which this occurs are appreciably shorter than for carbonyl compounds such as aldehydes and ketones, and, in fact, are out of the range of most commercial ultraviolet spectrometers. Some idea of how the hydroxyl substituent modifies the absorption properties of the carbonyl group in carboxylic acids can be seen from Table 18-2, in which are listed the wavelengths of maximum light absorption $\left( \lambda_\text{max} \right)$ and the extinction coefficients at maximum absorption $\left( \epsilon_\text{max} \right)$ of several carboxylic acids, aldehydes, and ketones.
Table 18-2: Wavelengths for Maximum Ultraviolet Absorption of Some Carboxylic Acids, Aldehydes, and Ketones $\left( n \rightarrow \pi^* \right)$
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/18%3A_Carboxylic_Acids_and_Their_Derivatives/18.02%3A_Physical_Properties_of_Carboxylic_Acids.txt |
Most of the reactions of carboxylic acids belong to one of four principal classes, depending on the point in the molecule where the reaction occurs.
a. Reactions involving the $\ce{O-H}$ bond - these include acid dissociation and solvolytic reactions.
b. Reactions at the carbonyl bond - most of which involve attack by a nucleophile $: \ce{Nu}$ on the carbonyl carbon with subsequent cleavage of a $\ce{C-O}$ bond. Examples are esterification, acyl chloride formation, and reduction with hydrides.
c. Decarboxylation - these are reactions in which the $\ce{R-C}$ bond is broken in such a way that $\ce{CO_2}$ is lost and $\ce{R-H}$ is formed.
d. Substitution on the $\ce{R}$ group - substitutions for hydrogen or halogen at the 2-carbon are especially important.
We will emphasize the way in which the chemistry of carboxylic acids in each of these categories can be correlated with the principles outlined in previous chapters.
Dissociation of Carboxylic Acids. The Resonance Effect
Compared with mineral acids such as hydrochloric, perchloric, nitric, and sulfuric acids, the carboxylic acids, $\ce{CH_3(CH_2)}_n \ce{CO_2H}$, are weak. The extent of dissociation in aqueous solution is relatively small, the acidity constants, $K_\text{a}$, being approximately $10^{-5}$ (see Table 18-1).
Even though the carboxylic acids are weak acids, they are many orders of magnitude stronger than the corresponding alcohols, $\ce{CH_3(CH_2)}_n \ce{CH_2OH}$. Thus the $K_\text{a}$ of ethanoic acid, $\ce{CH_2CO_2H}$, is $10^{11}$ times larger than that of ethanol, $\ce{CH_3CH_2OH}$.
The acidity of the carboxyl group arises, at least in part, from the polar nature of the carbonyl group, the polarity of which can be ascribed to contributions of the structure
For a carboxyl group, these structures and an additional possibility are shown by $1a$, $1b$, and $1c$:
Although the uncharged structure, $1a$, is of major importance, structures $1b$ and $1c$ make significant contributions. The stabilization is substantial and carboxylic acids are more stable than would be expected, from summing up their bond energies, by fully $18 \: \text{kcal mol}^{-1}$.
The stabilization energy of the carboxylate anion is substantially greater than that of the acid, because the anion is a resonance hybrid of two energetically equivalent structures, $2a$ and $2b$, whereas the acid is represented by a hybrid of nonequivalent structures, $1a$ through $1c$:
The rules for resonance stress that the greatest stabilization is expected when the contributing structures are equivalent (Section 6-5B). Therefore we can conclude that the resonance energy of a carboxylate anion should be greater than that of the corresponding acid. Consequently we can say that there is a "driving force" (a gain in stability) that promotes the dissociation of carboxylic acids. The fact that alcohols are far weaker acids than carboxylic acids may be attributed to the lack of stabilization of alkoxide ions compared to that of carboxylate anions. The difference in energy corresponding to the dissociation of a carboxylic acid (Equation 18-1) relative to that of an alcohol (Equation 18-2) actually amounts to about $15 \: \text{kcal mol}^{-1}$:
The Inductive Effect and Acid Strengths
You may have noticed that there are considerable differences between the strengths of some of the acids listed in Table 18-1. Methanoic acid and almost all the substituted ethanoic acids are stronger than ethanoic acid. In fact, trifluoroethanoic acid is similar in strength to hydrochloric acid. The substituent groups clearly can have a profound effect on acid strength by what commonly is called the inductive effect, an effect related to the electronegativity of the substituent. The inductive effect is different from resonance effects discussed in Section 18-2A in that it is associated with substitution on the saturated carbon atoms of the chain. The inductive effect of the substituent makes the acid stronger or weaker (relative to the unsubstituted acid), depending on whether the substituent is electron-attracting or electron-donating relative to hydrogen.
The electronegativity scale (Section 10-4B) shows chlorine to be more electron-attracting than hydrogen, and chloroethanoic acid is an 80-times stronger acid than ethanoic acid itself. Substitution by more chlorines enhances the acidity. Dichloroethanoic acid is 3000 times and trichloroethanoic acid is 5000 times more acidic than ethanoic acid. Moving the position of substitution along the chain away from the carboxyl group makes the effect smaller, and 4-chlorobutanoic acid is only a two-times stronger acid than butanoic acid (Table 18-1).
The inductive effect of the substituent can be considered to be transmitted to the carboxyl group in two rather different ways. Most frequently, the substituent is regarded as causing shifts in the average distributions of the bonding electrons along the chain of atoms between it and the carboxyl proton. This produces a succession of electron shifts along the chain, which, for an electron-attracting substituent, increases the acid strength by making it more energetically feasible for the $\ce{-OH}$ hydrogen of the carboxyl group to leave as a proton:
Many other groups besides halogens exhibit electron-withdrawing acid-enhancing inductive effects. Among these are nitro $\left( \ce{-NO_2} \right)$, methoxy $\left( \ce{CH_3O} \right)$, carbonyl ($\ce{RCOR'}$, as in aldehydes, ketones, acids, esters, and amides), cyano or nitrile $\left( \ce{-C \equiv N} \right)$, and trialkylammonium $\left( \ce{R_3} \overset{\oplus}{\ce{N}} \ce{-} \right)$. Alkyl groups - methyl, ethyl, isopropyl, and so on - are the only substitutents listed in Table 18-1 that are acid-weakening relative to hydrogen (as can be seen by comparing the $K_\text{a}$ values of the longer-chain acids with those of methanoic and ethanoic acids). We may take this to mean that alkyl groups release electrons to the carboxyl group.
The Electrostatic Interpretation of Acid Strengths
The other possible mode of transmission of the polar effect of a substituent group is a purely electrostatic one, sometimes called the "field effect", in which the dipole of the substituent produces an electrostatic field at the carboxyl proton, which helps or hinders ionization depending on the way in which the dipole is oriented with respect to the carboxyl group. It is easiest to visualize how the electrostatic theory operates by considering a proton midway between two well-separated carboxylate anions and deciding with which one the proton can combine more favorably. The more favorable one will correspond to the more basic carboxylate anion and the weaker carboxylic acid. With $\ce{CH_3CO_2^-}$ and $\ce{ClCH_2CO_2^-}$ as examples, and remembering that the $\ce{Cl-C}$ bond is polarized as $\overset{\delta \ominus}{\ce{Cl}} \ce{-} \overset{\delta \oplus}{\ce{C}}$, we can write:
The proton will be attracted by both $\ce{-CO_2^-}$ groups and the differences in electrostatic energy of its combination with one or the other of the carboxylate groups will depend on the influence of the $\overset{\delta \ominus}{\ce{Cl}} \ce{-} \overset{\delta \oplus}{\ce{C}} \ce{H_2}$ dipole. The $\overset{\delta \ominus}{\ce{Cl}}$ end of the dipole will attract the proton, but the $\overset{\delta \oplus}{\ce{C}} \ce{H_2}$ end will repel it. The repulsion effect will be more important because the $\overset{\delta \oplus}{\ce{C}} \ce{H_2}$ is closer than $\overset{\delta \ominus}{\ce{Cl}}$ to the final point of attachment of the proton to the carboxylate oxygen.$^2$ Thus, the proton will go more favorably to $\ce{CH_3CO_2^-}$ than to $\ce{ClCH_2CO_2^-}$, which means that $\ce{ClCH_2CO_2H}$ is a stronger acid than $\ce{CH_3CO_2H}$.
Quantitative application of electrostatic theory to the effect of substituents on the ionization of carboxylic acids, such as chloroethanoic acid, is rendered difficult by the fact that the polarized $\overset{\delta \ominus}{\ce{Cl}} \ce{-} \overset{\delta \oplus}{\ce{C}} \ce{H_2}$ bond and the proton cannot be treated as if they were point charges in a vacuum. The intervening and surrounding matter must be taken into account. This is especially true for water solutions because a molecule of an organic acid in water is a cavity of low dielectric constant immersed in a medium of high dielectric constant. Therefore, when ionization occurs the proton goes from inside the cavity through the boundary into the water. At the same time, the nature of the cavity must change because it then contains an anion, not a neutral molecule.
What Part Does Entropy Play in the Dissociation of Carboxylic Acids?
We have discussed the influence of substituents on acid strengths of simple carboxylic acids as though the full electrostatic effect of the substituent were exerted solely on the $\Delta H$ of ionization. However, careful thermodynamic analysis of acidities in aqueous solution show that entropy effects (Section 4-4B) are very important. This may seem surprising because entropy effects ought to be small for relative acid strengths, which can be assessed by the constants for simple equilibria such as Equation 18-3, in which (1) there are the same number of molecules on each side of the equation, and (2) the constraints on the species involved hardly seem different from one side of the equation to the other:
The entropy effects associated with these equilibria have to do with the "invisible" participant, water, which is involved in an intimate way, although by convention we omit it from equations such as 18-3. Solvation of ions puts constraints on water molecules, and the same electrostatic effects that change the ease of removing the proton act to change the degree and nature of solvation, thereby requiring consideration of entropy effects on the equilibria.
If solvation entropy effects are important, how can we justify using simple electrostatic theory to account for changes in acid strengths produced by substituents? The answer lies in $\Delta G$; whatever the electrostatic effects are doing to balance between $\Delta H$ and $\Delta S$, it is $\Delta G$ that determines the equilibrium constant and $\Delta G$ quite consistently follows the predictions of simple electrostatic considerations. Furthermore, the relative acid strengths of a number of substituted ethanoic acids have been determined in the gas phase by ion-cyclotron resonance (Section 27-8), under conditions where association and solvent effects are absent (Section 11-8A). In the gas phase, entropy effects are small and the relative acidities are in the order expected from the electronegativity scale, provided on corrects for the ion-size effect that we encountered previously with respect to the gas-phase acidities of alkynes and alcohols (Section 11-8B and 15-4A). Thus fluoroethanoic acid is weaker than chloroethanoic acid in the gas phase, whereas the reverse is true in water solution. The difference may be due simply to the fact that larger ions are in general more stable than smaller ions in the gas phase.
Carboxylic Acids as Bases
In addition to their acidic properties, carboxylic acids also can act as weak bases when the carbonyl oxygen accepts a proton from a strong acid, such as $\ce{H_2SO_4}$, $\ce{HClO_4}$, or $\ce{HSbF_6}$ in $\ce{SO_2}$ (Equation 18-4). Such protonation is an important step in acid-catalyzed esterification, as discussed in Section 15-4D:
A proton also can add to the hydroxyl oxygen (Equation 18-5). The resulting conjugate acid normally is less favorable than its isomer with the proton on the carbonyl group. Nonetheless, this conjugate acid plays a role in esterification when the $\ce{R}$ group is particularly bulky and, in addition, has electron-donating properties, thereby favoring ionization to an acyl carbocation (as in Equation 18-6; see also Section 18-3A):
Salts of Carboxylic Acids as Soaps. Micelle Formation
Carboxylic acids have an important practical use in the form of their metal salts as soaps. We have mentioned how fats, which are 1,2,3-propanetriol (glyceryl) esters of long-chain acids, can be hydrolyzed with alkali to give the corresponding carboxylate salts. It has been known as far back as Roman times (Pliny) that such substances have value for cleaning purposes.$^3$ These salts have a complicated interaction with water because they are very polar at the salt end of the molecule and very nonpolar at the long-chain hydrocarbon end of the molecule. These hydrocarbon ends are not compatible with a polar solvent such as water.$^4$
When minute amounts of soaps are put into water, instead of forming simple solutions, the molecules become concentrated at the surface of the water, with the saltlike ends sticking down into the water and the hydrocarbon chains forming a layer on the surface. This arrangement greatly reduces the surface tension of the water and contributes to the startling properties of soap films and bubbles. At higher concentrations, the solutions become turbid as the result of micelle formation. Micelles are sizable aggregates of soap molecules, wherein the hydrocarbon chains form a region of low polarity that is stabilized by having the polar salt ends of the molecules in contact with the water (Figure 18-4).
The cleansing action of soap is partly due to the way soap lowers the surface tension of the water thereby helping it to penetrate into fabrics, and also to the ability of the micelles to solubilize oils and greases by taking them into their hydrocarbon regions.
A major disadvantage of the simple carboxylate soaps is that they combine with the calcium and magnesium ions normally present in most tap water to form insoluble scums, which interfere with the cleansing process. Many so-called detergents have been developed that do not have this disadvantage - an example is sodium 4-dodecanylbenzenesulfonate, whose calcium and magnesium salts are water soluble.
When carboxylate salts are put into nonpolar solvents, reversed micelles often are formed, where the polar parts of the molecules are on the inside and the nonpolar parts are on the outside.
Pronounced differences have been observed for the rates of chemical reactions in micelles as compared to pure water. For example, the solvolysis of the 1-methylheptyl sulfonate, $5$, in dilute water solution proceeds 70 times slower when sufficient sodium dodecanyl sulfate $\left( \overset{\oplus}{\ce{Na}} \overset{\ominus}{\ce{O}} \ce{SO_3C_{12}H_{25}} \right)$ is added to provide about twice as many dodecanyl sulfate ions in the micelle state as there are molecules of $5$ present:
This slowing of the solvolysis reaction by the alkyl sulfate requires that $5$ be almost completely imprisoned by the micelles, because that part of $5$ free in water would hydrolyze rapidly. An important result is in the stereochemistry of the reaction, which changes from $100\%$ inversion with optically active $5$ in pure water to only $56\%$ inversion in the micelles. Micelles of the opposite polarity, made from hexadecyltrimethylammonium bromide, $\ce{C_{16}H_{33}} \overset{\oplus}{\ce{N}} \ce{(CH_3)_3} \overset{\ominus}{\ce{Br}}$, have no effect on the rate of solvolysis of $5$.
Studies of this type have been made on a number of systems and are of great interest because of the light they may shed on the structure and function of biological membranes.
There is a close resemblance between fatty-acid salts and phospholipids in that both possess long hydrocarbon tails and a polar head. Phospholipids also aggregate in a polar medium to form micelles and continuous bilayer structures such as shown in Figure 18-5. The bilayer lipid structure is very important to the self-sealing function of membranes and their impermeability to very polar molecules.
$^2$The electrostatic energy involved in brining two charges from a distance $r_1$ to a distance $r_2$ apart is given by $\left( e_1 e_2/D \right) \left( 1/r_1 - 1/r_2 \right)$, where $e_1$ and $e_2$ are the magnitude of the charges and $D$ is the dielectric constant of the medium (Section 8-7F). If $e_1$ and $e_2$ have the same sign, the energy is positive, and with opposite signs the energy is negative. For calibration, the electrostatic energy resulting from bringing a positive charge from a large distance $\left( 1/r_1 \sim 0 \right)$, up to a negative charge at a distance of $1$ Å in a vacuum $\left( D = 1 \right)$ is $-335 \: \text{kcal mol}^{-1}$.
$^3$Until the 19th century soaps were made by boiling animal or vegetable fats with wood ashes, which contain, besides silica, considerable amounts of potassium carbonated. The resulting mixture of potassium carboxylate salts gives a "soft" soap, and this can be converted to a "hard" soap by treatment with excess $\ce{NaCl}$, which forms the less soluble sodium carboxylate salts. The $\ce{KCl}$ formed goes into the aqueous phase.
$^4$One might well wonder why soap molecules do not simply crystallize out of water solution if the hydrocarbon chains are incompatible with water. However, the crystal packing of the polar salt parts of the molecule is not likely to be very compatible with the hydrocarbon parts and, furthermore, most soaps are salts of mixtures of aliphatic acids and this hardly helps crystallization to occur.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/18%3A_Carboxylic_Acids_and_Their_Derivatives/18.03%3A_Some_Chemical_Properties_of_Carboxylic_Acids.txt |
Many important reactions of carboxylic acids involve attack on the carbon of the carbonyl group by nucleophilic species. These reactions frequently are catalyzed by acids, because addition of a proton or formation of a hydrogen bond to the carbonyl oxygen makes the carbonyl carbon more vulnerable to nucleophilic attack. The following equations illustrate how an acid-catalyzed reaction operates with a neutral nucleophile $\left( \ce{H-Nu} \right)$:
Subsequent cleavage of a $\ce{C-O}$ bond and loss of a proton yields a displacement product:
An important example of this type of reaction is the formation of esters, which was discussed previously in connection with the reactions of alcohols in Section 15-4D. Similar addition-elimination mechanisms occur in many reactions at the carbonyl groups of acid derivatives. A less obvious example of addition to carboxyl groups involves hydride ion $\left( \ce{H}^\ominus \right)$ and takes place in lithium aluminum hydride reduction of carboxylic acids (Sections 16-4E and 18-3C).
Esterification
Esters, $\ce{RCO_2R'}$, are formed from carboxylic acids and alcohols in the presence of acid catalysts. The key step in esterification is the nucleophilic attack of a neutral alcohol molecule, $\ce{R'OH}$, at the carbonyl carbon of the conjugate acid of the carboxylic acid, $\ce{RC(OH)_2^+}$, $6$:
The intermediate, $7$, either can revert to the starting materials or form a second intermediate, $8$, by proton transfer. Loss of water from $8$ leads to the conjugate acid of the ester, $9$:
The final step in formation of the ester is proton transfer from $9$ to the solvent:
All the steps in ester formation are reversible, but the equilibrium in the $\ce{C-O}$ bond-making and -breaking processes are not very favorable, and an excess of one reactant (usually the alcohol) or removal of one product (most often water) is required to give a good yield of ester.
The usefulness of direct ester formation from alcohols and acids is limited to those alcohols or acids that do not undergo extensive side reactions in the presence of strong acids. Furthermore, if the alcohol is particularly bulky the reaction usually will not proceed satisfactorily because the intermediates $7$ and $8$ (as well as the product) are rendered unstable by crowding of the substituent groups.
Bulky groups in the esterifying acid also hinder the reaction. A classic example is 2,4,6-trimethylbenzoic (mesitoic) acid, which cannot be esterified readily under normal conditions because the methyl groups ortho to the carboxyl group make the transition state for formation of the intermediate $10$ less favorable relative to the starting acid than would be the case for less hindered acids, such as ethanoic acid:
The important point is the difference in steric hindrance between the acid and the intermediate. If you make a scale model you will see that in the acid, the carboxyl group, being planar, can have reduced hindrance by turning about its bond to the ring so as to be between the methyl groups. However, no such relief is possible with $10$, in which the $\ce{-C(OH)_2OR}$ carbon is tetrahedral.
Esterification of acids with bulky substituents, such as 2,4,6-trimethylbenzoic acid, can be achieved through formation of acyl cations. This is done by simply dissolving the carboxylic acid in strong sulfuric acid, whereby the acyl cation $11$ is formed, and then pouring the solution into an excess of cold alcohol (see also Equations 18-5 and 18-6). This procedure works because it avoids the formation of a hindered tetrahedral intermediate similar to $10$ and instead forms the conjugate acid directly:
Esterification of carboxylic acids with bulky alcohols is unsatisfactory. However, tertiary alkyl esters often can be prepared by addition of the acid to the appropriate alkene using an acid catalyst:
The success of such addition reactions depends on formation of a stable carbocation from the alkene under conditions where the most reactive nucleophile present is the carboxylic acid.
Acyl Chloride Formation
Carboxylic acids react with phosphorus trichloride, phosphorus pentachloride, or thionyl chloride with replacement of $\ce{OH}$ by $\ce{Cl}$ to form acyl chlorides, $\ce{RCOCl}$:
Although detailed mechanisms have not been established, the first step is thought to be formation of an unstable mixed anhydride, which then extrudes $\ce{SO_2}$ and "collapses" with attach of chloride at the carbonyl carbon. A similar mechanism occurs in the formation of alkyl chlorides from alcohols and thionyl chloride (Section 15-5A):
Most acyl halides are stable, distillable liquids. However, methanoyl chloride, $\ce{HCOCl}$, decomposes to carbon monoxide and hydrogen chloride at room temperature.
Reduction of Carboxylic Acids
Generally, carboxylic acids are difficult to reduce either by catalytic hydrogenation or by sodium and alcohol. Nonetheless, reduction to primary alcohols proceeds smoothly with lithium aluminum hydride, $\ce{LiAlH_4}$:
The first step in lithium aluminum hydride reduction of carboxylic acids is formation of a complex aluminum salt of the acid and hydrogen:
Reduction then proceeds by successive transfers of hydride ion, (\ce{H}^\ominus\), from aluminum to carbon. The first such transfer reduces the acid salt to the oxidation level of the aldehyde; reduction does not stop at this point, however, but continues rapidly to the alcohol. Insufficient information is available to permit very specific structures to be written for the intermediates in the lithium aluminum hydride reduction of carboxylic acids. However, the product is a complex aluminum alkoxide, from which the alcohol is freed by hydrolysis:
Sodium borohydride, $\ce{NaBH_4}$, is too mild a reducing agent to transfer hydride to carboxylic acids, and one may suspect that borane, $\ce{BH_3}$, also would be ineffective. However, this is not the case and borane in oxacyclopentane (tetrahydrofuran) reduces carboxylic acids more rapidly than it adds to alkene double bonds (see Tale 16-5):
The reason for the high reactivity lies in the fact that the acid first converts the borane to an acyloxyborane, which then undergoes an intramolecular rearrangement in which the carbonyl group is reduced. Hydrolysis gives the alcohol:
Special methods are required for the direct reduction of $\ce{RCO_2H}$ to $\ce{RCHO}$. Aldehydes can be obtained directly by the slow reduction of carboxylic acids with 2,3-dimethyl-2-butylborane in oxacyclopentane solution. One hydrogen of the borane is wasted through reaction with the acidic hydrogen of the carboxyl group to give hydrogen. An example is
The borane is prepared through the addition of $\ce{B_2H_6}$ to 2,3-dimethyl-2-butene and, because of steric hindrance, only the monoalkylborane is formed (Section 11-6A):
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/18%3A_Carboxylic_Acids_and_Their_Derivatives/18.04%3A_Reactions_at_the_Carbonyl_Carbon_of_Carboxylic_Acids.txt |
The decarboxylation of $\ce{RCO_2H}$ to give $\ce{RH}$ and $\ce{CO_2}$ can be calculated from bond energies and the stabilization energy of the carboxyl group to have $\Delta H^0 = -7 \: \text{kcal mol}^{-1}$. This does not mean that the reaction goes easily. Special structural features are required. The simple aliphatic carboxylic acids do not lose carbon dioxide on heating, but when there are strongly electron-attracting groups attached to the $\alpha$ carbon, decarboxylation often proceeds readily at $100$-$150^\text{o}$. Examples include
3-Butenoic acid also undergoes decarboxylation but has to be heated to above $200^\text{o}$:
The mechanisms of thermal decarboxylation probably are not the same for all cases, but when the acid has a double-bonded function such as $\ce{O=C}$, $\ce{N=C}$, $\ce{O=N}$, or $\ce{C=C}$ attached to the $\alpha$ carbon then a cyclic elimination process appears to occur. For propanedioic acid the process is
Carboxylate radicals can be generated in several ways. One is the thermal decomposition of diacyl peroxides, which are compounds with rather weak $\ce{O-O}$ bonds:
Another method involves electrolysis of sodium or potassium carboxylate solutions, known as Kolbe electrolysis, in which carboxylate radicals are formed by transfer of an electron from the carboxylate ion to the anode. Decarboxylation may occur simultaneously with, or subsequent to, the formation of carboxylate radicals, leading to hydrocarbon radicals, which subsequently dimerize:
An example is
Decarboxylation of the silver salts of carboxylic acids in the presence of bromine or chlorine, the Hunsdiecker reaction, often is useful for the synthesis of alkyl halides:
The mechanism of this reaction seems to involve formation of carboxylate radicals through decomposition of an acyl hypobromic intermediate, $12$:
The Hunsdiecker reaction has certain disadvantages, mainly because it requires use of the pure dry silver salt, which is often difficult to prepare. With some acids, however, excellent results can be obtained using the acid itself and an excess of red mercuric oxide in place of the silver salt.
or by heating the acid with lead tetraethanoate, $\ce{Pb(O_2CCH_3)_4}$, and iodine,
A somewhat similar decarboxylation reaction with formation of an alkene can be achieved by heating a carboxylic acid with lead tetraethanoate, $\ce{Pb(O_2CCH_3)_4}$, in the presence of a catalytic amount of $\ce{Cu(OCH_3)_2}$. A useful example is
There is some competing decarboxylation of the ethanoic acid, but the conversions in this kind of reaction are usually good. The key steps in the reaction probably are exchange of carboxylic acid groups on tetravalent lead, cleavage of the $\ce{Pb-O}$ bond to give the carboxylate radical, decarboxylation, oxidation of the alkyl radical by $\ce{Cu}$ (II) to give the cation $\left[ \ce{R} \cdot + \ce{Cu} (II) \rightarrow \ce{R}^\oplus + \ce{Cu} (I) \right]$, and finally loss of a proton to form the alkene.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
18.06: Reactions at the (alpha) Carbons of Carboxylic Acids
Halogenation
Bromine reacts smoothly with carboxylic acids in the presence of small quantities of phosphorus to form alpha-bromocarboxylic acids (Hell-Volhard-Zelinsky reaction):
The reaction is slow in the absence of phosphorus, whose function appears to be to form phosphorus tribromide, which then reacts with the acid to give the acyl bromide:
Formation of the acyl bromide speeds up the reaction because acid-catalyzed enolization of the acyl bromide occurs much more readily than enolization of the parent acid. Bromine probably reacts with the enol of the acyl bromide in the same way as it reacts with the enols of ketones (Section 17-2A).
The final step is the formation of the $\alpha$-bromo acid by bromine exchange between the $\alpha$-bromoacyl bromide and the parent acid; the acyl bromide, which is necessary for continued reaction, is thus regenerated:
This bromination reaction results exclusively in alpha substitution and therefore is limited to carboxylic acids with $\alpha$ hydrogens. Chlorine with a trace of phosphorus reacts similarly but with less overall specificity, because concurrent free-radical chlorination can occur at all positions along the chain (as in hydrocarbon halogenation; see Section 4-6A).
Substitution Reactions of $\alpha$-Haloalkanoic Acids
The halogen of an $\alpha$-haloalkanoic acid is replaced readily by nucleophilic reagents such as $\ce{CN}^\ominus$, $\ce{OH}^\ominus$, $\ce{I}^\ominus$, and $\ce{NH_3}$. Thus a variety of $\alpha$-substituted carboxylic aids may be prepared by reactions that are analogous to $S_\text{N}2$ substitution of alkyl halides:
Facile $S_\text{N}2$ substitution reactions of halogens are expected from the electron-attracting characteristics of the neighboring carbonyl function, which should make the transition state for attack by a nucleophilic reagent more favorable:
Perhaps it may seem surprising that the carboxyl carbon is not attacked by the nucleophilic agents, because we have stressed earlier the susceptibility of carbonyl groups to nucleophilic reagents. No stable product results, however, from addition to the carbonyl group by the type of reagents considered here. Thus with cyanide ion the equilibrium constant for addition is unfavorable because of the associated loss of stabilization energy of the carboxyl group.
The $S_\text{N}1$ reactivity of $\alpha$-haloalkanoic acids is particularly low. This is reasonable because formation of a cationic center at the $\alpha$ carbon should be difficult, because of the positive character of the carbonyl carbon. Furthermore, little, if any, help could be expected through electron delocalization because the corresponding valence-bond structure has a positive, single-bonded oxygen:
Similar considerations apply to the $S_\text{N}1$ and $S_\text{N}2$ reactions of $\alpha$-halo aldehydes and $\alpha$-halo ketones (Section 17-2C).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
18.07: Functional Derivatives of Carboxylic Acids
A functional derivative of a carboxylic acid is a substance formed by replacement of the hydroxyl group of the acid by some other group, $\ce{X}$, such that it can be hydrolyzed back to the acid in accord with Equation 18-7:
By this definition, an amide, $\ce{RCONH_2}$, but not a ketone, $\ce{RCOCH_3}$, is a functional derivative of a carboxylic acid. Several derivatives of carboxylic acids are give in Table 18-3, and methods for preparation of these derivatives are summarized in Tables 18-6 and 18-7 at the end of the chapter.
Table 18-3: Functional Derivatives of Carboxylic Acids
The common structural feature of the compounds listed in Table 18-3 is the acyl group $\ce{RCO}-$. However, nitriles, $\ce{RC \equiv N}$, often are considered to be acid derivatives, even though the acyl group is not present as such, because hydrolysis of nitriles leads to carboxylic acids:
The chemistry of nitriles will be discussed in Section 24-5.
The two main types of reactions of carboxylic acid derivatives with which we now shall be concerned are the replacement of $\ce{X}$ by attack of a nucleophile $\ce{Nu}^\ominus$ at the carbonyl carbon with subsequent cleavage of the $\ce{C-X}$ bond (Equation 18-8), and substitution at the $\alpha$ carbon facilitated by the carbonyl group (Equation 18-9):
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/18%3A_Carboxylic_Acids_and_Their_Derivatives/18.05%3A_Decarboxylation_of_Carboxylic_Acids.txt |
Displacement Reactions
Hydrolysis of most acid derivatives to the parent acids is acid- or base-catalyzed:
However, acyl halides and anhydrides usually hydrolyze rapidly without the aid of an acidic or basic catalyst, when in solution. It is important to recognize that an insoluble acyl halide or anhydride often reacts slowly with water. Esters and amides hydrolyze much more slowly, and for useful rates require a catalyst. The hydrolysis of amides is of exceptional importance in biochemistry and will be discussed in more detail in Chapters 24 and 25.
Acid-catalyzed hydrolysis of esters is the reverse of acid-catalyzed ester formation discussed previously. Base-induced ester hydrolysis (saponification) is an irreversible reaction. The initial step is the attack of hydroxide ion at the carbonyl carbon:
The intermediate anion, $13$, so formed then either loses $\ce{OH}^\ominus$ and reverts to the original ester, or it loses $\ce{CH_3O}^\ominus$ to form the acid. The overall reaction is irreversible because once the acid is formed, it immediately is converted to the carboxylate anion, which is stabilized to such a degree that it is not attacked by the alcohol and will not reform the starting ester. Consequently, the reaction goes to completion in the direction of hydrolysis:
Ester interchange is closely related to ester hydrolysis. This is a base-catalyzed reaction that is useful to replace the alcohol group of an ester with a different alcohol group. The catalyst is alkoxide ion and the equilibrium constant is close to unity, unless the alcohols differ greatly in size. An example is
in which $\ce{RO}^\ominus$ is either $\ce{CH_3O}^\ominus$ or $\ce{CH_3CH_2O}^\ominus$.
The formation of esters from acid chlorides and anhydrides according to the following equation has been discussed:
Amides can be obtained from acyl halides, carboxylic anhydrides, or esters with amines or ammonia. The mechanisms of these reactions are very similar to the corresponding reactions of alcohols:
We will discuss this kind of reaction further in Chapters 24 and 25.
Reactions with Organometallic Compounds
The reactions of several carboxylic acid derivatives with organomagnesium and organolithium compounds were described in Section 14-12. The key step in these reactions is addition of the organometallic compound, as $\ce{R}^{\delta \ominus} \ce{M}^{\delta \oplus}$, to the carbonyl group. For a Grignard reagent,
The reaction normally does not stop at this stage; $\ce{MgXZ}$ is eliminated and the resulting ketone rapidly reacts with another molecule of organometallic compound. On hydrolysis, a tertiary alcohol is formed with at least two identical alkyl groups on the tertiary carbon:
Reduction of Acid Derivatives
Esters, chlorides, and anhydrides are reduced by lithium aluminum hydride in the same general way as the parent acids (Section 18-3C), the difference being that no hydrogen is evolved. The products after hydrolysis are primary alcohols:
Nitriles can be reduced to amines by lithium aluminum hydride. An imine salt is an intermediate product; if the reaction is carried out under the proper conditions, this salt is the major product and provides an aldehyde on hydrolysis (see Section 16-4C):
Amides are reduced to primary amines, and $\ce{N}$-substituted amides to secondary and tertiary amines:
Borane also will reduce esters, amides, and nitriles to the same products as does $\ce{LiAlH_4}$, but with reduced reactivity (Table 16-6).
Although lithium aluminum hydride and boranes are very useful reagents, they are expensive and impractical to employ on a large scale. Other methods of reduction then may be necessary. Of these, the most important are reduction of esters with sodium and ethanol (acids do not react readily),
and high-pressure hydrogenation over a copper chromite catalyst,
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/18%3A_Carboxylic_Acids_and_Their_Derivatives/18.08%3A_Reactions_at_the_Carbonyl_Carbon_of_Acid_Derivatives.txt |
The Acidic Properties of Esters with $\alpha$ Hydrogens
Many important synthetic reactions in which $\ce{C-C}$ bonds are formed involve esters and are brought about by basic reagents. This is possible because the $\alpha$ hydrogens of an ester, such as $\ce{RCH_2CO_2C_2H_5}$, are weakly acidic, and a strong base, such as sodium ethoxide, can produce a significant concentration of the ester anion at equilibrium:
The acidity of $\alpha$ hydrogens is attributed partly to the electron-attracting inductive effects of the ester oxygens, and partly to resonance stabilization of the resulting anion (Section 17-1A):
When the $\alpha$ carbon of the ester carries a second strongly electron-attracting group, the acidity of $\alpha$ hydrogen is greatly enhanced. Examples of such compounds follow:
The stabilization of the anions of these specially activated esters is greater than for simple esters because of the electron-withdrawing inductive effects of the substituents but more importantly because the negative charge can be distributed over more than two centers. Thus for the anion of ethyl 3-oxobutanoate we can regard all three of the valence-bond structures, $14a$ through $14c$, as important in contributing to the hybrid, $14$:
The anion, $14$, is sufficiently stable relative to the ester that the $K_\text{a}$ is about $10^{-11}$ in water solution (Table 17-1).
Ethyl 3-oxobutanoate exists at room temperature as an equilibrium mixture of keto and enol tautomers in the ratio of 92.5 to 7.5. The presence of enol can be shown by rapid titration with bromine, but is more evident from the proton nmr spectrum (Figure 18-6), which shows absorption of the hydroxyl, alkenyl, and methyl protons of the enol form, in addition to absorptions expected for the keto form:
Interconversion of the enol and keto forms of ethyl 3-oxobutanoate is powerfully catalyzed by bases through the anion, $15$, and less so by acids through the conjugate acid of the keto form:
Nonetheless, if contact with acidic and basic substances is rigidly excluded to the extent of using quartz equipment in place of glass (glass normally has a slightly alkaline surface), then interconversion is slow enough that it is possible to separate the lower-boiling enol from the keto form by fractional distillation under reduced pressure. The separated isomers are indefinitely stable when stored at $-80^\text{o}$ in quartz vessels.
The Claisen Condensation
One of the most useful of the base-induced reactions of esters is illustrated by the self-condensation of ethyl ethanoate under the influence of sodium ethoxide to give ethyl 3-oxobutanoate:
$\ce{CH_3CO_2C_2H_5} + \ce{H-CH_2CO_2C_2H_5} \overset{^\ominus \ce{OC_2H_5}}{\longrightarrow} \ce{CH_3COCH_2CO_2C_2H_5} + \ce{C_2H_5OH}$
This reaction, called the Claisen condensation, is interesting because, from consideration of bond and stabilization energies, it is expected to be unfavorable thermodynamically with $\Delta H^0$ (vapor) equal to $6 \: \text{kcal mol}^{-1}$. This expectation is realized in practice, and much effort has been expended to determine conditions by which practical yields of the condensation product can be obtained.
The Claisen condensation resembles both the aldol addition (Section 17-3) and carbonyl additions of acid derivatives discussed previously (Sections 16-4 and 18-7). The first step, as shown in Equation 18-10, is the formation of the anion of ethyl ethanoate which, being a powerful nucleophile, attacks the carbonyl carbon of a second ester molecule (Equation 18-11). Elimination of ethoxide ion then leads to the $\beta$-keto ester, ethyl 3-oxobutanoate (Equation 18-12):
The sum of these steps represents an unfavorable equilibrium, and satisfactory yields of the $\beta$-keto ester are obtained only if the equilibrium can be shifted by removal of one of the products.
One simple way of doing this is to remove the ethanol by distillation as it is formed; however, this may be difficult to carry to completion and, in any case, is self-defeating if the starting ester is low-boiling. Alternatively, one can use a large excess of sodium ethoxide. This is helpful because ethanol is a weaker acid than the ester enol, and excess ethoxide shifts the equilibrium to the right through conversion of the $\beta$-keto ester to the enolate salt (Equation 18-13).
Obviously, the condensation product must be recovered from the enol salt and isolated under conditions that avoid reversion to starting materials. The best procedure is to quench the reaction mixture by pouring it into an excess of cold dilute acid.
A limitation on the Claisen condensation is that although the starting ester need have only one $\alpha$ hydrogen for Reactions 18-10 through 18-12 to occur, two $\alpha$ hydrogens are necessary for a favorable equilibrium utilizing the ionization reaction of Equation 18-13. As a result, it is not surprising to find that ethyl 2-methylpropanoate fails to condense with itself in the presence of sodium ethoxide, because the condensation product has no $\alpha$ hydrogen next to the ester group:
However, it an excess of a very much stronger base than sodium ethoxide is used [such as triphenylmethylsodium, $\ce{(C_6H_5)_3C}^\ominus \ce{Na}^\oplus$], this same condensation does take place in reasonable yields. The reason is that the base is now strong enough to convert the alcohol formed in the reaction to sodium ethoxide, thus shifting the equilibrium to the right:
The overall reaction then is
Claisen condensations can be carried out between two different esters but, because there are four possible products, mixtures often result. Less difficulty is encountered if one of the esters has no $\alpha$ hydrogen and reacts readily with a carbanion according to Equations 18-11 and 18-12. The reaction then has considerable resemblance to the mixed aldol additions discussed in Section 17-3C. Among the useful esters without $\alpha$ hydrogens, and with the requisite electrophile reactivity, are those of benzenecarboxylic, methanoic, ethanedioic, and carbonic acids. Several practical examples of mixed Claisen condensations are shown in Equations 18-14 through 18-16 (all of the products exist to the extent of $10\%$ or so as the enol forms):
An important variation on the Claisen condensation is to use a ketone as the anionic reagent. This often works well because ketones usually are more acidic than simple esters and the base-induced self-condensation of ketones (aldol addition) is thermodynamically unfavorable (Section 17-3C). A typical example is the condensation of cyclohexanone with diethyl ethanedioate (diethyl oxalate):
$\alpha$-Keto esters of the type formed according to Equations 18-16 and 18-17 have synthetic utility in that they lose carbon monoxide when strongly heated:
A somewhat similar decarbonylation reaction was mentioned previously for diphenylpropanetrione (Section 17-10).
Alkylation of Ester Anions
The anions of esters such as ethyl 3-oxobutanoate and diethyl propanedioate can be alkylated with alkyl halides. These reactions are important for the synthesis of carboxylic acids and ketones and are similar in character to the alkylation of ketones discussed previously (Section 17-4A). The ester is converted by a strong base to the enolate anion, Equation 18-18, which then is alkylated in an $S_\text{N}2$ reaction with the alkyl halide, Equation 18-19. Usually, $\ce{C}$-alkylation predominates:
Esters of propanedioic (malonic) acid can be alkylated in a similar fashion (Equation 18-20):
Unfortunately, monoalkylation seldom occurs cleanly by the above sequence whenever the monoalkylation product has an $\alpha$ hydrogen located so as to permit dialkylation to occur. In practice, alkylation reactions, using one mole of ester, one mole of sodium ethoxide, and one mole of an alkyl halide (e.g., $\ce{CH_3I}$), give a mixture of the starting ester, its mono- and dialkylation products. The situation is more favorable when large alkyl groups are introduced, because then the physical properties and reactivities of the starting materials and of mono- and dialkylation products differ considerably. Usually dialkylation is inhibited by having a bulky alkyl group in the monoalkylation product.
Alkyl-substituted 3-oxobutanoic and propanedioic esters can be hydrolyzed under acidic conditions to the corresponding acids, and when these are heated they readily decarboxylate (Section 18-4). Alkyl 3-oxobutanoic esters thus yield methyl alkyl ketones, whereas alkylpropanedioic esters produce carboxylic acids:
These reactions commonly are known as the acetoacetic-ester ketone and the malonic-ester acid syntheses, respectively.
Alkyl 3-oxobutanoic esters react with concentrated alkali by a different path to reverse the Claisen condensation:
Acylation of Ester Anions
Enolate anions of esters, such as ethyl 3-oxobutanoate or diethyl propanedioate, react with acyl halides or anhydrides to give acylation products. These reactions are carried out best using sodium hydride instead of sodium ethoxide for production of the enol salt, because then no alcohol is liberated to react with the acyl halide or anhydride:
Aldol-Type Additions of Ester Anions and the Reformatsky Reaction
Addition of an ester anion to the carbonyl group of an aldehyde or ketone is a type of aldol addition (Section 17-3):
There are certain difficulties in achieving this type of aldol reaction. First, alkali-induced ester hydrolysis would compete with addition. Second, a Claisen condensation of the ester might intervene, and third, the ester anion is a stronger base than the enolate anions of either aldehydes or ketones, which means reaction could be defeated by proton transfer of the type
However, a useful synthetic reaction can be achieved in the following way. First, the ester anion is formed in the absence of water without causing a Claisen condensation or other carbonyl addition. This can be done with ethyl ethanoate by treating it with lithium bis(trimethylsilyl)amide in oxacyclopentane solution at $-80^\text{o}$:
The advantage of $\ce{LiN[Si(CH_3)_3]_2}$ as the base in this reaction is that $\overset{\ominus}{\ce{N}} \ce{[Si(CH_3)_3]_2}$ is a reasonably strong base; it is bulky, which inhibits addition to the carbonyl; and it also forms a weakly basic amine, $\ce{HN[Si(CH_3)_3]_2}$, which does not interfere in the subsequent reactions.
The solution of ethyl lithioethanoate must be kept cold and treated promptly with an aldehyde or ketone. Thus, with 2-propanone,
For the reaction to be successful, the carbonyl addition has to be faster than the proton transfer reaction, $\ce{LiCH_2CO_2C_2H_5} + \ce{CH_3COCH_3} \rightarrow \ce{CH_3CO_2C_2H_5} + \ce{LiCH_2COCH_3}$ and, at $-80^\text{o}$, this is the case. This synthesis of $\beta$-hydroxy esters is a beautiful example of how rates of competing reactions can be manipulated to obtain a high yield of a desired addition product that may not be the most thermodynamically favorable one.
A closely related synthesis of $\beta$-hydroxy esters is provided by the Reformatsky reaction. This synthesis starts with an aldehyde or ketone, $\ce{RCOR'}$, and an $\alpha$-bromo ester, such as ethyl bromoethanoate. Zinc in a nonhydroxylic solvent (usually benzene) transforms the bromo ester into an organozinc compound, which then adds to the aldehyde or ketone carbonyl. Hydrolysis produces the $\beta$-hydroxy ester:
As do aldols, $\beta$-hydroxy esters dehydrate (usually readily) to $\alpha$,$\beta$-unsaturated carbonyl compounds.
Biological Claisen Condensations and Aldol Additions. Fatty Acid Metabolism
The overall results of a Claisen condensation is the transfer of an acyl group $\left( \ce{RCO}- \right)$ from one ester molecule to another:
The reverse of the above reaction is a key step in the oxidative degradation of fatty acids. This reverse Claisen condensation (catalyzed by thiolase) involves the cleavage of a carbon-carbon bond of a $\beta$-keto ester of coenzyme A by another molecule of coenzyme A to give a new acyl derivative ($\ce{RCO-S}$CoA) and ethanoyl (acetyl) derivative ($\ce{CH_3O-S}$CoA):
After formation of the $\beta$-keto thioester, it is cleaved by CoA$\ce{SH}$, and the resulting thioester goes back into the sequence two carbons shorter than before. In this way, a fatty acid is degraded from the carboxyl end, two carbons at a time.
There are two principle pathways for utilization of the ethanoyl coenzyme A ($\ce{CH_3CO-S}$CoA) formed in each turn of the oxidation cycle of Figure 18-8. Either it is used to synthesize larger molecules such as fatty acids, steroids, and so on, as will be described in Section 30-5A, or the acyl group is oxidized to $\ce{CO_2}$ and $\ce{H_2O}$:
$\ce{CH_3CO-S}$CoA $\overset{\left[ \ce{O} \right]}{\rightarrow} 2 \ce{CO_2} + \ce{H_2O} + \ce{HS}$CoA
The oxidation of the acyl group of coenzyme A is the net outcome of the citric acid or Krebs cycle (Section 20-10B). We will be interested here in the entry point of the cycle whereby ethanoyl coenzyme A is employed in a reaction that builds the $\ce{C_6}$ chain of citric acid (3-carboxy-3-hydroxypentanedioic acid) from $\ce{C_2}$ and $\ce{C_4}$ pieces:
This reaction is quite special in that it is an aldol-type addition in which a thioester is the donor (nucleophile) and a keto acid is the acceptor (electrophile). From the discussion in Section 18-8E, you will see that reactions of this kind involving an ester as the donor and an aldehyde or ketone as the acceptor can be achieved in the laboratory only under rather special conditions. For the thioester to function as a nucleophile at the $\alpha$ carbon under the restraints imposed by having the reaction occur at the physiological pH, the catalyzing enzyme almost certainly must promote formation of the enol form of the thioester. The enol then could add to the ketone carbonyl with the assistance of a basic group on the enzyme. This kind of catalysis by enzymes is discussed in Section 25-9C.
$^5$Considerable confusion is possible because of the way in which biochemists use abbreviated names and formulas for the acyl derivatives of coenzyme A. To emphasize the vital $\ce{-SH}$ group, coenzyme A is usually written as CoA$\ce{SH}$. However, the acyl derivatives most often are called acetyl CoA and the like, not acetyl $\ce{S}$CoA, and you could well get the erroneous impression that the sulfur has somehow disappeared in forming the acyl derivative. We will include the sulfur in formulas such as $\ce{CH_3COS}$CoA, but use the customary names such as acetyl CoA without including the sulfur. To make clear that CoA does not contain cobalt, CoA is printed in this text in boldface type.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/18%3A_Carboxylic_Acids_and_Their_Derivatives/18.09%3A_Reactions_at_the_%28alpha%29_Carbons_of_Carboxylic_Acid_Derivatives.txt |
Unsaturated carboxylic acids of the type $\ce{RCH=CH(CH_2)}_n \ce{COOH}$ usually exhibit the properties characteristic of isolated double bonds and isolated carboxyl groups when $n$ is large and the functional groups are far apart. As expected, exceptional behavior is found most commonly when the groups are sufficiently close together to interact strongly, as in $\alpha$,$\beta$-unsaturated acids, $\ce{R} \overset{\beta}{\ce{C}} \ce{H=} \overset{\alpha}{\ce{C}} \ce{CO_2H}$. We shall emphasize those properties that are exceptional in the following discussion.
Migration of the Double Bond
In the presence of strong base, $\alpha$,$\beta$- and $\beta$,$\gamma$-unsaturated carboxylic acids tend to interconvert by migration of the double bond:
Ester derivatives, $\ce{RCH=CH-CH_2COOR'}$, and the corresponding unsaturated aldehydes and ketones, $\ce{RCH=CH-CH_2COR'}$, are much more prone to this type of rearrangement than are the acids.
Hydration and Hydrogen Bromide Addition
Like alkenes, the double bonds of $\alpha$,$\beta$-unsaturated acids can be brominated, hydroxylated, hydrated, and hydrobrominated, although the reactions often are relatively slow. In the addition of unsymmetrical reagents the direction of addition is opposite to that observed for alkenes (anti-Markovnikov). Thus propenoic (acrylic) acid adds hydrogen bromide and water to form 3-bromo- and 3-hydroxypropanoic acids:
These additions are analogous to the addition of halogens and halogen acids to 1,3-butadiene (Section 13-2). In the first step, a proton is transferred to the carbonyl oxygen. The resulting conjugate acid can be regarded as a resonance hybrid of structures $16a$-$16d$:
In the second step, a nucleophile (such as $\ce{Br}^\ominus$ or a water molecule) attacks an electron-deficient carbon of the hybrid $16$. Attack at the carboxyl carbon may occur but does not lead to a stable product. Attack of the nucleophile at the $\beta$ carbon, however, produces the enol form of the $\beta$-substituted acid, which then is converted rapidly to the normal carboxylic acid:
Lactone Formation
When the double bond of an unsaturated acid is farther down the carbon chain than between the alpha and beta positions, the so-called "conjugate addition" is not possible. Nonetheless, the double bond and carboxyl group frequently interact in the presence of acidic catalysts because the carbocation that results from addition of a proton to the double bond has a built-in nucleophile (the carboxyl group), which may attack the cationic center to form a cyclic ester called a lactone.
Lactone formation only occurs readily by this mechanism when a five- or six-membered ring can be formed:
Five- or six-membered lactones also are formed by internal esterification when either $\gamma$- or $\delta$-hydroxy acids are heated. Under similar conditions, $\beta$-hydroxy acids are dehydrated to $\alpha$,$\beta$-unsaturated acids, whereas $\alpha$-hydroxy acids undergo bimolecular esterification to substances with six-membered dilactone rings called lactides:
More on the Michael Addition
The foregoing examples of addition to the double bonds of unsaturated carboxylic acids all involve activation by an electrophilic species such as $\ce{H}^\oplus$. Conjugate addition also may occur by nucleophilic attack on acid derivatives, the most important being the base-catalyzed Michael addition (Section 17-5B) and 1,4-addition of organometallic compounds (Section 14-12D). In all of these reactions a nucleophilic agent, usually a carbanion, attacks the double bond of an $\alpha$,$\beta$-unsaturated compound in which the double bond is conjugated with, and activated by, a strongly electronegative unsaturated group (such as $\ce{-CN}$, $\ce{-NO_2}$, etc.). In the Michael addition, the carbanion usually is an enolate salt.
The overall reaction is illustrated here by the specific example of the addition of diethyl propanedioate (diethyl malonate) to ethyl 3-phenylpropenoate (ethyl cinnamate):
The mechanism of this kind of transformation, with diethyl propanedioate as the addend, is outlined in Equations 18-25 and 18-26. The basic catalyst required for the Michael addition (here symbolized as $\ce{B:}$) serves by forming the corresponding anion:
A variety of nucleophilic agents can be used; propanedinitrile, 3-oxobutanoate esters, and cyanoethanoate esters all form relatively stable carbanions and function well in Michael addition reactions. Obviously, if the carbanion is too stable, it will have little or no tendency to attack the double bond of the $\alpha$,$\beta$-unsaturated acid derivative.
Enamines (Sections 16-4C and 17-4B) are excellent addends in many Michael-type reactions. An example is provided by the addition of $\ce{N}$-(1-cyclohexenyl)-azacyclopentane to methyl 2-methylpropanoate:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/18%3A_Carboxylic_Acids_and_Their_Derivatives/18.10%3A_Reactions_of_Unsaturated_Carboxylic_Acids_and_Their_Derivatives.txt |
Acids in which there are two carboxyl groups separated by a chain of more than five carbon atoms $\left( n > 5 \right)$ for the most part have unexceptional properties, and the carboxyl groups behave more or less independently of one another.
However, when the carboxyl groups are closer together the possibilities for interaction increase; we shall be interested primarily in such acids. A number of important dicarboxylic acids are listed in Table 18-4 together with their physical properties, methods of manufacture, and commercial uses.
Table 18-4: Dicarboxylic Acids
Acidic Properties of Dicarboxylic Acids
The inductive effect of one carboxyl group is expected to enhance the acidity of the other. In Table 18-4 we see that the acid strength of the dicarboxylic acids, as measured by the first acid-dissociation constant, $K_1$, is higher than that of ethanoic acid $\left( K_\text{a} = 1.5 \times 10^{-5} \right)$ and decreases with increasing number of bonds between the two carboxyl groups. The second acid-dissociation constant, $K_2$, is smaller than $K_\text{a}$ for ethanoic acid (with the exception of oxalic acid) because it is more difficult to remove a proton under the electrostatic attraction of the nearby carboxylate anion (see Section 18-2C).
Thermal Behavior of Dicarboxylic Acids
The reactions that occur when dicarboxylic acids are heated depend critically upon the chain length separating the carboxyl groups. Cyclization usually is favored if a strainless five- or six-membered ring can be formed. Thus hexanedioic and heptanedioic acids decarboxylate and cyclize to cyclopentanone and cyclohexanone, respectively:
Butanedioic and pentanedioic acids take a different course. Rather than form the strained cyclic ketones, cyclopropanone and cyclobutanone, both acids form cyclic anhydrides that have five- and six-membered rings, respectively. 1,2-Benzenedicarboxylic (phthalic) and cis-1,4-butenedicarboxylic (maleic) acids behave similarly:
Because of their short chains, propanedioic and ethanedioic acids simply decarboxylate when heated (Section 18-4):
Imides from Dicarboxylic Acids
The cyclic anhydride of butanedioic acid reacts with ammonia, as may be expected for a typical anhydride; but the product, when strongly heated, forms a cyclic imide (butanimide):
1,2-Benzenedicarboxylic (phthalic) anhydride behaves similarly, giving 1,2-benzenedicarboximide (phthalimide):
Unlike amines, imides do not have basic properties in water solution; the electron pair of nitrogen is partly delocalized over the carbonyl groups, as indicated by $17a$ to $17c$. This stabilization is lost if a proton is added to nitrogen to give the conjugate acid, $18$:
Imides are, in fact, quite acidic and readily dissolve in alkali-metal hydroxide solutions to give salts. Like carboxylic acids and 1,3-dicarbonyl compounds, imides are acidic primarily because the stabilization of the anion is greater than that of the acid. This can be seen by comparison of the resonance structures that may be written for the imide, $17$, with those of the anion, $18$. Separation of positive and negative charge, as in Structures $17b$ and $17c$, increases the energy of such structures. There is no charge separation in the anion; thus $19b$ and $19c$ are more important with respect to their hybrid than are $17b$ and $17c$ to their hybrid. (You may wish to review the corresponding argument for the acidity of carboxylic acids, Section 18-2A.)
The salts of imides are useful in synthesis, as is described in Section 23-9D.
The Dieckmann Condensation
Esters of most dicarboxylic acids, except propanedioic esters, undergo Claisen condensation in much the same way as do esters of monocarboxylic acids (see Section 18-8B). However, when a strainless five- or six-membered ring can be formed, an intramolecular Claisen condensation, called the Dieckmann condensation, may take place which would result in the formation of a cyclic $\beta$-keto ester:
The Acyloin Reaction
A useful method of forming carbon-carbon bonds involves reduction of esters with sodium metal in aprotic solvents such as ether or benzene and is called the acyloin reaction:
This interesting reaction is especially useful for the synthesis of medium- and large-ring compounds from dicarboxylic esters, and is effective for ring sizes that cannot be made by the Dieckmann condensation or decarboxylation (Section 18-10B). Radical anions formed by addition of sodium to the ester groups appear to be the key intermediates for carbon-carbon bond formation. Thus, for dimethyl decanedioate:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
18.12: Methods of Preparation of Carboxylic Acids and Their Derivatives
Table 18-5: Methods of Preparation of Carboxylic Acids\(^a\)
Table 18-6: Methods of Preparation of Carboxylic Esters
Table 18-7: Methods of Preparationof Acyl Halides, Anhydrides, Amides, and Related Compounds | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/18%3A_Carboxylic_Acids_and_Their_Derivatives/18.11%3A_Dicarboxylic_Acids.txt |
The fundamentals of structure and stereochemistry have been considered in previous chapters in some detail. There are, however, practical aspects of stereochemistry that have not yet been mentioned, particularly with regard to chiral compounds. How, for instance, can a racemic mixture be separated into its component enantiomers (resolution); what methods can be used to establish the configuration of enantiomers; how can we tell if they are pure; and how do we synthesize one of a pair of enantiomers preferentially (asymmetric synthesis)? In this chapter, some answers to these questions will be described briefly.
Optical activity is an associated phenomenon of chirality and has long been used to monitor the behavior of chiral compounds. Brief mention of this was made earlier (Section 5-1C), but now the origin and measurement of optical rotation will be examined in more detail.
• 19.1: Plane-Polarized Light and the Origin of Optical Rotation
Electromagnetic radiation involves the propagation of both electric and magnetic forces. At each point in an ordinary light beam, there is a component electric field and a component magnetic field, which are perpendicular to each other and oscillate in all directions perpendicular to the direction in which the beam propagates. In plane-polarized light the component electric field oscillates as in ordinary light, except that the direction of oscillation is within a single plane.
• 19.2: Specific Rotation
Optical rotation is the usual and most useful means of monitoring enantiomeric purity of chiral molecules. Therefore we need to know what variables influence the magnitude of optical rotation.
• 19.3: Separation or Resolution of Enantiomers
Because the physical properties of enantiomers are identical, they seldom can be separated by simple physical methods, such as fractional crystallization or distillation. It is only under the influence of another chiral substance that enantiomers behave differently, and almost all methods of resolution of enantiomers are based upon this fact. We include here a discussion of the primary methods of resolution.
• 19.4: Enantiomeric Purity
Enantiomeric purity (or optical purity) is defined as the fractional excess of one enantiomer over the other. This is expressed in terms of the moles (or weights) of the two enantiomers and is equal to the ratio of the observed optical rotation and to the optical rotation of either pure enantiomer. Thus a racemic mixture has an enantiomeric purity of zero. Any other enantiomeric composition in principle can be determined provided the mixture has a measurable rotation and the rotation of the pure
• 19.5: Absolute And Relative Configuration
The sign of rotation of plane-polarized light by an enantiomer is not easily related to its configuration. This is true even for substances with very similar structures. Thus, given lactic acid with a specific rotation +3.82°, and methyl lactate with a specific rotation −8.25°, we cannot tell from the rotation alone whether the acid and ester have the same or a different arrangement of groups about the chiral center. Their relative configurations have to be obtained by other means.
• 19.6: The R,S Convention for Designating Stereochemical Configurations
The R,S or Cahn-Ingold-Prelog convention is a systematic way of denoting configuration that may eventually replace the D,L system, at least for simple compounds. To denote the configuration of a chiral center by the R ,S convention, the groups at the center are assigned an order of precedence according to a specific set of rules based on atomic numbers.
• 19.7: E,Z Notation
Many compounds cannot be described adequately by the cis-trans system. A system that is easy to use and which is based on the sequence rules already described for the R , S system is call the E,Z notation.
• 19.8: Prochirality
There is a special term for molecules that are achiral but which can be converted to molecules with chiral centers by a single chemical substitution or addition reaction. They are said to be prochiral.
• 19.9: Optical Rotatory Dispersion and Circular Dichroism
Nevertheless, much information has been obtained about structure, conformation, and configuration of organic compounds from measurements of optical rotation as a function of wavelength (i.e., optical rotatory dispersion). Like other phenomena involving interactions between electromagnetic radiation and organic molecules, as in infrared, ultraviolet, and NMR spectroscopy, optical rotatory dispersion curves often are quite sensitive to small changes in structure.
• 19.10: Asymmetric Synthesis
If one could prepare 2-hydroxypropanenitrile from ethanal and hydrogen cyanide in the absence of any chiral reagent and produce an excess of one enantiomer over the other, this would constitute an absolute asymmetric synthesis - that is, creation of preferential chirality (optical activity) in a symmetrical environment from symmetrical reagents.
• 19.11: Racemization
Optically active biphenyl derivatives are racemized if the two aromatic rings at any time pass through a coplanar configuration by rotation about the central bond. This can be brought about more or less easily by heat, unless the 2,2'-ortho substituents are very large. The way in that compounds with asymmetric carbon atoms are racemized is more complicated. One possibility would be for a tetrahedral chiral carbon attached to four groups to become planar and achiral without breaking any bonds.
• 19.E: More on Stereochemistry (Exercises)
These are the homework exercises to accompany Chapter 19 of the Textmap for Basic Principles of Organic Chemistry (Roberts and Caserio).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
19: More on Stereochemistry
Electromagnetic radiation, as the name implies, involves the propagation of both electric and magnetic forces. At each point in an ordinary light beam, there is a component electric field and a component magnetic field, which are perpendicular to each other and oscillate in all directions perpendicular to the direction in which the beam propagates. In plane-polarized light the component electric field oscillates as in ordinary light, except that the direction of oscillation is contained within a single plane. Likewise, the component magnetic field oscillates within a plane, the planes in question being perpendicular to each other. A schematic representation of the electric part of plane-polarized light and its interaction with an optical isomer is shown in Figure 19-1. The beam of polarized light, $XY$, has a component electric field that oscillates in the plane $AOD$. At the point $O$ the direction of oscillation is along $OE$. If now at $O$ the beam encounters a substance which has the power to cause the direction of oscillation of the electrical field to rotate through an angle $\alpha$ to the new direction $OE'$ in the plane $COB$, the substance is said to be optically active.
A clockwise rotation, as the observer looks towards the beam, defines the substance as dextrorotatory (i.e., rotates to the right) and the angle $\alpha$ is taken as a positive $\left( + \right)$ rotation. If the rotation is counterclockwise the substance is described as levorotatory (i.e., rotates to the left) and the angle $\alpha$ is taken as a negative $\left( - \right)$ rotation.
The question naturally arises as to why some substances interact with polarized light in this manner whereas others do not. We shall oversimplify the explanation because a rigorous treatment involves rather complex mathematics. However, it is not difficult to understand that the electric forces in a light beam impinging on a molecule will interact to some extent with the electrons within the molecule. Although radiant energy actually may not be absorbed by the molecule to promote it to higher, excited electronic-energy states (see Section 9-9A), a perturbation of the electronic configuration of the molecule can occur. One can visualize this process as a polarization of the electrons brought about by the oscillating electric field associated with the radiation.
This interaction is important to us here because it causes the electric field of the radiation to change its direction of oscillation. The effect produced by any one molecule is extremely small, but in the aggregate may be measurable as a net rotation of the plane-polarized light. Molecules such as methane, ethene and 2-propanone, which have enough symmetry so that each is identical with its reflection, do not rotate plane-polarized light. This is because the symmetry of each is such that every optical rotation in one direction is canceled by an equal rotation in the opposite direction. However, a molecule with its atoms so disposed in space that it is not symmetrical to the degree of being superimposable on its mirror image will have a net effect on the incident polarized light, because then the electromagnetic interactions do not average to zero. We characterize such substances as having chiral configurations and as being optically active.
A useful model for explanation of optical rotation considers that a beam of plane-polarized light is the vector resultant of two oppositely rotating beams of circularly polarized light. This will be clearer if we understand that circularly polarized light has a component electric field that varies in direction but not in magnitude so that the field traverses a helical path in either a clockwise or counterclockwise direction, as shown in Figure 19-2.
The resultant of the two oppositely rotating electric vectors lies in a plane, and the magnitude of the resultant varies as a sine wave, shown in Figure 19-3. This amounts to plane-polarized light.
When circularly polarized light travels through an assemblage of one kind of chiral molecules, the velocity of light observed for one direction of circular polarization is different from that for the other direction of polarization. This is eminently reasonable because, no matter how a chiral molecule is oriented, the molecule presents a different aspect to circularly polarized light rotating in one direction than to that rotating in the other direction. Consequently if the electric vectors of two circularly polarized light beams initially produce a resultant that lies in a plane, and the beams then encounter a medium in which they have different velocities, one beam will move steadily ahead of the other. This will cause a continual rotation of the plane of their resultant until they again reach a medium in which they have equal velocities.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/19%3A_More_on_Stereochemistry/19.01%3A_Plane-Polarized_Light_and_the_Origin_of_Optical_Rotation.txt |
Optical rotation is the usual and most useful means of monitoring enantiomeric purity of chiral molecules. Therefore we need to know what variables influence the magnitude of optical rotation.
The measured rotation, $\alpha$, of a chiral substance varies with the concentration of the solution (or the density of a pure liquid) and on the distance through which the light travels. This is to be expected because the magnitude of $\alpha$ will depend on the number as well as the kind of molecules the light encounters. Another important variable is the wavelength of the incident light, which always must be specified even though the sodium D line $\left( 589.3 \: \text{nm} \right)$ commonly is used. To a lesser extent, $\alpha$ varies with the temperature and with the solvent (if used), which also should be specified. The optical rotation of a chiral substance usually is reported as a specific rotation $\left[ \alpha \right]$, which is expressed by the Equations 19-1 or 19-2
For solutions:
$\left[ \alpha \right]^t_{\lambda} =\dfrac{100 \alpha}{l \times c} \tag{19-1}$
For neat liquids:
$\left[ \alpha \right]^t_{\lambda} =\dfrac{\alpha}{l \times d} \tag{19-2}$
with
• $\alpha$ is the measured rotation in degrees
• $t$ is the temperature
• $\lambda$ is the wavelength of light
• $l$ is the length in decimeters of the light path through the solution
• $c$ is the concentration in grams of sample per 100 ml of solution
• $d$ is the density of liquid in grams ml-I
For example, quinine (Section 19-3A) is reported as having $\left[ \alpha \right]_{D} = -117^\text{o} \left( c = 1.5, \: \ce{CHCl_3} \right) \left( t = 17^\text{o} \right)$, which means that it has a levorotation of 117 degrees for sodium D light $\left( 589.3 \: \text{nm} \right)$ at a concentration of $1.5 \: \text{g}$ per $100 \: \text{mL}$ of chloroform solution at $17^\text{o} \text{C}$ when contained in a tube 1 decimeter long.
Frequently, molecular rotation, $\left[ M \right]$, is used in preference to specific rotation and is related to specific rotation by Equation 19-3:
$\left[ M \right]^t_{\lambda} =\dfrac{\left[ \alpha \right]^t_{\lambda} \times M}{100} \tag{19-3}$
in which $M$ is the molecular weight of the compound. Expressed in this form, optical rotations of different compounds are directly comparable on a molecular rather than a weight basis. The effects of wavelength of the light in the polarized beam on the magnitude and sign of the observed optical rotation are considered in Section 19-9.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
19.03: Separation or Resolution of Enantiomers
Because the physical properties of enantiomers are identical, they seldom can be separated by simple physical methods, such as fractional crystallization or distillation. It is only under the influence of another chiral substance that enantiomers behave differently, and almost all methods of resolution of enantiomers are based upon this fact. We include here a discussion of the primary methods of resolution.
Chiral Amines as Resolving Agents. Resolution of Racemic Acids
The most commonly used procedure for separating enantiomers is to convert them to a mixture of diastereomers that will have different physical properties: melting point, boiling point, solubility, and so on (Section 5-5). For example, if you have a racemic or $D$,$L$ mixture of enantiomers of an acid and convert this to a salt with a chiral base having the $D$ configuration, the salt will be a mixture of two diastereomers, ($D$ acid $\cdot$ $D$ base) and ($L$ acid $\cdot$ $D$ base). These diastereomeric salts are not identical and they are not mirror images. Therefore they will differ to some degree in their physical properties, and a separation by physical methods, such as crystallization, may be possible. If the diastereomeric salts can be completely separated, the acid regenerated from each salt will be either exclusively the $D$ or the $L$ enantiomer:
Resolution of chiral acids through the formation of diastereomeric salts requires adequate supplies of suitable chiral bases. Brucine, strychnine, and quinine frequently are used for this purpose because they are readily available, naturally occurring chiral bases. Simpler amines of synthetic origin, such as 2-amino-1-butanol, amphetamine, and 1-phenylethanamine, also can be used, but first they must be resolved themselves.
Resolution of Racemic Bases
Chiral acids, such as $\left( + \right)$-tartaric acid, $\left( - \right)$-malic acid, $\left( - \right)$-mandelic acid, and $\left( + \right)$-camphor-10-sulfonic acid, are used for the resolution of a racemic base.
The principle is the same as for the resolution of a racemic acid with a chiral base, and the choice of acid will depend both on the ease of separation of the diastereomeric salts and, of course, on the availability of the acid for the scale of the resolution involved. Resolution methods of this kind can be tedious, because numerous recrystallizations in different solvents may be necessary to progressively enrich the crystals in the less-soluble diastereomer. To determine when the resolution is complete, the mixture of diastereomers is recrystallized until there is no further change in the measured optical rotation of the crystals. At this stage it is hoped that the crystalline salt is a pure diastereomer from which one pure enantiomer can be recovered. The optical rotation of this enantiomer will be a maximum value if it is "optically" pure because any amount of the other enantiomer could only reduce the magnitude of the measured rotation $\alpha$.
Resolution of Racemic Alcohols
To resolve a racemic alcohol, a chiral acid can be used to convert the alcohol to a mixture of diastereomeric esters. This is not as generally useful as might be thought because esters tend to be liquids unless they are very high-molecular-weight compounds. If the diastereomeric esters are not crystalline, they must be separated by some other method than fractional crystallization (for instance, by chromatography methods, Section 9-2). Two chiral acids that are useful resolving agents for alcohols are:
The most common method of resolving an alcohol is to convert it to a half-ester of a dicarboxylic acid, such as butanedioic (succinic) or 1,2-benzenedicarboxylic (phthalic) acid, with the corresponding anhydride. The resulting half-ester has a free carboxyl function and may then be resolvable with a chiral base, usually brucine:
Other Methods of Resolution
One of the major goals in the field of organic chemistry is the development of reagents with the property of "chiral recognition" such that they can effect a clean separation of enantiomers in one operation without destroying either of the enantiomers. We have not achieved that ideal yet, but it may not be far in the future. Chromatographic methods (Section 9-2), whereby the stationary phase is a chiral reagent that adsorbs one enantiomer more strongly than the other, have been used to resolve racemic compounds, but such resolutions seldom have led to both pure enantiomers on a preparative scale.
Other methods, called kinetic resolutions, are excellent when applicable. The procedure takes advantage of differences in reaction rates of enantiomers with chiral reagents. One enantiomer may react more rapidly, thereby leaving an excess of the other enantiomer behind. For example, racemic tartaric acid can be resolved with the aid of certain penicillin molds that consume the dextrorotatory enantiomer faster than the levorotatory enantiomer. As a result, almost pure $\left( - \right)$-tartaric acid can be recovered from the mixture:
A disadvantage of resolutions of this type is that the more reactive enantiomer usually is not recoverable from the reaction mixture.
The crystallization procedure employed by Pasteur for his classical resolution of $\left( \pm \right)$-tartaric acid (Section 5-1C) has been successful only in a very few cases. This procedure depends on the formation of individual crystals of each enantiomer. Thus if the crystallization of sodium ammonium tartrate is carried out below $27^\text{o}$, the usual racemate salt does not form; a mixture of crystals of the $\left( + \right)$ and $\left( - \right)$ salts forms instead. The two different kinds of crystals, which are related as an object to its mirror image, can be separated manually with the aid of a microscope and subsequently may be converted to the tartaric acid enantiomers by strong acid. A variation on this method of resolution is the seeding of a saturated solution of a racemic mixture with crystals of one pure enantiomer in the hope of causing crystallization of just that one enantiomer, thereby leaving the other in solution. Unfortunately, very few practical resolutions have been achieved in this way.
Even when a successful resolution is achieved, some significant problems remain. For instance, the resolution itself does not provide information on the actual configuration of the $\left( + \right)$ or $\left( - \right)$ enantiomer. This must be determined by other means (see Section 19-5). Also, it is not possible to tell the enantiomeric purity (optical purity) of the resolved enantiomers without additional information. This point is discussed further in the next section.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/19%3A_More_on_Stereochemistry/19.02%3A_Specific_Rotation.txt |
The term enantiomeric purity (or optical purity) is defined as the fractional excess of one enantiomer over the other. This is expressed in Equation 19-4 in terms of the moles (or weights) of the two enantiomers, $n_1$, and $n_2$, and is equal to the ratio of the observed optical rotation, $\alpha_\text{obs}$, and to the optical rotation of either pure enantiomer, $\alpha_0$:
$\text{enantiomeric purity of } n_1 = \dfrac{n_1-n_2}{n_1+n_2} = \dfrac{\alpha_\text{obs}}{\alpha_0} \tag{19-4}$
Thus a racemic mixture ($n_1 = n_2$) has an enantiomeric purity of zero. Any other enantiomeric composition in principle can be determined provided the mixture has a measurable rotation and the rotation of the pure enantiomer, $\alpha_0$, is known. Unfortunately, there is no simple method of calculating $\alpha_0$, in advance. In fact, specific rotations of optically pure compounds are determined most reliably from Equation 19-4 after measurement of enantiomeric purity by independent methods.
Virtually all of the methods for determining enantiomeric purity rely on the differences in chemical, physical, or spectroscopic properties of diastereomers derived from enantiomeric mixtures. We will mention here two of the most straightforward methods, based on gas-liquid chromatography and nuclear magnetic resonance.
Determination of Enantiomeric Purity by Gas Chromatography
This method amounts to a complete resolution of the type described in Section 19-3D, but on an analytical scale. For example, assume that you have a partially resolved compound, $A$, consisting of unequal amounts of the enantiomers $A_+$ and $A_-$. By reaction with a second chiral enantiomerically pure substance, $B_+$, A is converted to a mixture of diastereomers $A_+B_+$ and $A_-B_+$. Because these diastereomers are chemically and physically different, the mixture usually can be analyzed by gas-liquid chromatography (Section 9-2A). If the reaction of $B_+$ with $A_+$ and $A_-$ was quantitative, the relative areas of the two peaks eluting from the column correspond to the ratio of the diastereomers $A_+B_+/A_-B_+$, and thus to the ratio of enantiomers $A_+/A_-$, from which the enantiomeric purity of the partially resolved mixture can be calculated.
An alternative and very direct approach is to separate the enantiomers on a column in which the stationary liquid phase is a chiral compound. The diastereomeric interaction is between $A_+$ or $A_-$ and the chiral liquid phase, and may be sufficiently different to permit separation of $A_+$ from $A_-$. The ratio of the amounts of $A_+$ and $A_-$ corresponds to the enantiomeric purity.
Determination of Enantiomeric Purity by NMR Spectroscopy
The nmr chemical shifts of nuclei of enantiomeric compounds $A_+$ and $A_-$ are identical in achiral solvents. However, in a chiral solvent (enantiomerically pure) $A_+$ and $A_-$ will be effectively converted to diastereomers as the result of chiral solvation and, accordingly, their nuclei will have nonidentical chemical shifts. Provided that the shift differences are large enough to permit the resonances of one chirally solvated enantiomer to be resolved from those of the other, the ratio of enantiomers $A_+/A_-$ can be determined from the ratio of their corresponding nmr signal intensities.
Alternatively, with a pure chiral reagent the enantiomeric mixture may be converted (quantitatively) to diastereomers. The nuclei of the diastereomeric compounds are expected to have small differences in chemical shifts, even in achiral solvents, and integration of their respective signal intensities should correspond to the ratio of diastereomers, and hence to the ratio of enantiomers in the original mixture.
Application of the above-described nmr methods for the determination of enantiomeric composition is difficult, if not impossible, if the chemical-shift differences are too small $\left( 0.02 \: \text{ppm} \right)$ or if the resonances overlap extensively. This problem often can be solved by utilizing the ability of certain chelates of rare earth metals (the lanthanide metals) to complex with organic compounds, particularly with alcohols, ketones, amines, and other Lewis bases. Chemical shifts in the presence of even small amounts of lanthanide chelates usually are spread over a much wider range of field strengths than for the pure compounds. As discussed previously in Sections 9-10D and 9-10K, increasing chemical shifts can greatly simplify otherwise complex nmr spectra. The shifts are produced by these lanthanide compounds because the electrons on the metal atoms are not all paired, so that the metal atoms are paramagnetic (possess a net electron spin). In an applied magnetic field the unpaired electrons circulate around the metal atoms and produce an induced field (Figure 9-26) which, depending on the nature of the metal, can act either to increase or reduce the applied magnetic field, $H_0$.
In the complex formed between the lanthanide reagent and the organic substrate, the chemical shifts most strongly affected are those of nuclei close to the paramagnetic metal atom. The resonances of these nuclei also are broadened by the paramagnetic metal (Section 27-1), and this is undesirable. The lanthanide complexes that produce these large shift changes are called shift reagents. Most of them are chelate salts of substituted 2,4-pentanediones, especially of 2,2,6,6-tetramethyl-3,5-heptanedione (Section 17-8).
There are several useful shift reagents, usually of europium, in which the organic ligands are chiral. An example is $1$: | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/19%3A_More_on_Stereochemistry/19.04%3A_Enantiomeric_Purity.txt |
The sign of rotation of plane-polarized light by an enantiomer is not easily related to its configuration. This is true even for substances with very similar structures. Thus, given lactic acid, $\ce{CH_3CHOHCO_2H}$, with a specific rotation $+3.82^\text{o}$, and methyl lactate, $\ce{CH_3CHOHCO_2CH_3}$, with a specific rotation $-8.25^\text{o}$, we cannot tell from the rotation alone whether the acid and ester have the same or a different arrangement of groups about the chiral center. Their relative configurations have to be obtained by other means.
If we convert $\left( + \right)$-lactic acid into its methyl ester, we can be reasonably certain that the ester will be related in configuration to the acid, because esterification should not affect the configuration about the chiral carbon atom. It happens that the methyl ester so obtained is levorotatory, so we know that $\left( + \right)$-lactic acid and $\left( - \right)$-methyl lactate have the same relative configuration at the asymmetric carbon, even if they possess opposite signs of optical rotation. However, we still do not know the absolute configuration; that is, we are unable to tell which of the two possible configurations of lactic acid, $2a$ or $2b$, corresponds to the dextro or $\left( + \right)$-acid and which to the levo or $\left( - \right)$-acid:
Until 1956, the absolute configuration of no optically active compound was known. Instead, configurations were assigned relative to a standard, glyceraldehyde, which originally was chosen by E. Fischer (around 1885) for the purpose of correlating the configuration of carbohydrates. Fischer arbitrarily assigned the configuration $3a$ to dextrorotatory glyceraldehyde, which was known as $D$-$\left( + \right)$-glyceraldehyde. The levorotatory enantiomer, $3b$, is designated as $L$-$\left( - \right)$-glyceraldehyde. (If you are unsure of the terminology $D$ and $L$, or of the rules for writing Fischer projection formulas, review Sections 5-3C and 5-4.)
The configurations of many compounds besides sugars now have been related to glyceraldehyde, including $\alpha$-amino acids, terpenes, steroids, and other biochemically important substances. Compounds whose configurations are related to $D$-$\left( + \right)$-glyceraldehyde are said to belong to the $D$ series, and those related to $L$-$\left( - \right)$-glyceraldehyde belong to the $L$ series.
At the time the choice of absolute configuration for glyceraldehyde was made, there was no way of knowing whether the configuration of $\left( + \right)$-glyceraldehyde was in reality $3a$ or $3b$. However, the choice had a $50\%$ chance of being correct, and we now know that $3a$, the $D$ configuration, is in fact the correct configuration of $\left( + \right)$-glyceraldehyde. This was established through use of a special x-ray crystallographic technique, which permitted determination of the absolute disposition of the atoms in space of sodium rubidium $\left( + \right)$-tartrate. The configuration of $\left( + \right)$-tartaric acid (Section 5-5) previously had been shown by chemical means to be opposite to that of $\left( + \right)$-glyceraldehyde. Consequently the absolute configuration of any compound now is known once it has been correlated directly or indirectly with glyceraldehyde. For example,
When there are several chiral carbons in a molecule, the configuration at one center usually is related directly or indirectly to glyceraldehyde, and the configurations at the other centers are determined relative to the first. Thus in the aldehyde form of the important sugar, $\left( + \right)$-glucose, there are four chiral centers, and so there are $2^4 = 16$ possible stereoisorners. The projection formula of the isomer that corresponds to the aldehyde form of natural glucose is $4$. By convention for sugars, the configuration of the highest-numbered chiral carbon is referred to glyceraldehyde to determine the overall configuration of the molecule. For glucose, this atom is $\ce{C5}$, next to the $\ce{CH_2OH}$ group, and has the hydroxyl group on the right. Therefore, naturally occurring glucose, which has a $\left( + \right)$ rotation, belongs to the $D$ series and is properly called $D$-$\left( + \right)$-glucose:
However, the configurations of $\alpha$-amino acids possessing more than one chiral carbon are determined by the lowest-numbered chiral carbon, which is the carbon alpha to the carboxyl group. Thus, even though the natural $\alpha$-amino acid, threonine, has exactly the same kind of arrangement of substituents as the natural sugar, threose, threonine by the amino-acid convention belongs to the $L$-series, whereas threose by the sugar convention belongs to the $D$-series:
A serious ambiguity arises for compounds such as the active tartaric acids. If the amino-acid convention is used, $\left( + \right)$-tartaric acid falls in the $D$ series; by the sugar convention, it has the $L$ configuration. One way out of this dilemma is to use the subscripts $s$ and $g$ to denote the amino-acid or carbohydrate conventions, respectively. Then the absolute configuration of $\left( + \right)$-tartaric acid can be designated as either $D_s$-$\left( + \right)$-tartaric acid of $L_g$-$\left( + \right)$-tartaric acid.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/19%3A_More_on_Stereochemistry/19.05%3A_Absolute_And_Relative_Configuration.txt |
There are certain disadvantages to the $D$,$L$ system using Fischer projection formulas to denote configuration about a chiral center, and we already have seen how ambiguity arises in the case of the tartaric acids (Section 19-5). A more systematic way of denoting configuration that may eventually replace the $D$,$L$ system, at least for simple compounds, is known as the $R$,$S$ or Cahn-Ingold-Prelog convention, after its originators.
To denote the configuration of a chiral center by the $R$,$S$ convention, the groups at the center are assigned an order of precedence according to a specific set of rules based on atomic numbers. Suppose a carbon atom is bonded to four different substituents, which we will designate $A$,$B$, $C$, and $D$ and to which we assign the following priority sequences: $A$ before $B$ before $C$ before $D$. If we now view the arrangement of $A$, $B$, and $C$ from the site remote from the substituent of lowest priority, $D$, as shown in Figure 19-6, and the sequence turns out to be $A \rightarrow B \rightarrow C$ in the clockwise direction, then the configuration is said to be $R$. If the sequence $A \rightarrow B \rightarrow C$ occurs in the counterclockwise direction, the configuration is $S$. The symbols $R$ and $S$ are taken from the Latin words rectus and sinister, meaning right and left, respectively.
The understanding of $R$ and $S$ is simple; the problems are in assigning the priority sequences for actual substituents. The rules follow:
1. Priority is given to the substituent atoms that have the highest atomic number. This means that four different atoms arranged tetrahedrally about the chiral center have a priority sequence that decreases with decreasing atomic number. For example, the sequence among the halogens is $\ce{I} > \ce{Br} > \ce{Cl} > \ce{F}$, and Structure $5$ (shown here in perspective and in projection) therefore has the $R$ configuration:
For more complex substituents, priority is determined by the atomic number of the first bonded atom. The sequence $\ce{CH_3S} > \ce{CH_3O} > \ce{NH_2} > \ce{CH_3} > \ce{H}$ thus reflects the fact that atomic number decreases in the order $\ce{S} > \ce{O} > \ce{N} > \ce{C} > \ce{H}$. Structure $6$ accordingly has the $S$ configuration:
2. The first atoms in two or more substituents often are identical, in which case it is necessary to explore further and compare the atomic numbers of the second attached atoms. Precedence is given to the substituent with a second atom of higher atomic number. For example, in 2-butanol, $\ce{CH_3CH(OH)CH_2CH_3}$, two of the groups at the chiral atom have carbon as the first atom. We therefore must compare the other atoms bonded to these two carbons. It is convenient to represent the arrangement at the chiral atom as shown in $7$, where the first atoms are shown attached to the chiral center and the second atoms are listed in their priority order; thus, $\left( \ce{C}, \: \ce{H}, \: \ce{H} \right)$ for ethyl and $\left( \ce{H}, \: \ce{H}, \: \ce{H} \right)$ for methyl:
When we compare $\left( \ce{H}, \: \ce{H}, \: \ce{H} \right)$ with $\left( \ce{C}, \: \ce{H}, \: \ce{H} \right)$ in $7$, we give ethyl precedence over methyl because carbon has a higher atomic number than hydrogen. The configuration $7$ therefore must be $S$.
3. Double and triple bonds are treated as if they had duplicate or triplicate single bonds. Thus a carbonyl group, , is treated as if it were where the symbols in parentheses represent the duplicate atoms.
Let us see how this works for 3-chloro-1-pentyne, $8$:
The first-atom priority sequence is $\ce{Cl} > \ce{C}$ and $\ce{C} > \ce{H}$. We now need to order $\ce{-CH_2CH_3}$ and $\ce{-C \equiv CH}$ and, in doing this, we compare the three atoms attached to the first carbon of the ethyl group $\left( \ce{C}, \: \ce{H}, \: \ce{H} \right)$ with the three attached to the first carbon of the ethynyl group $\left[ \ce{C}, \: \left( \ce{C} \right), \: \left( \ce{C} \right) \right]$. On this basis, ethynyl comes ahead of ethyl, and the overall sequence is $\ce{Cl} > \ce{C \equiv CH} > \ce{CH_2CH_3} > \ce{H}$, so $8$ will have the $S$ configuration.
The sequence rules described thus far can be used without ambiguity in most of the examples we are likely to meet. The important thing to remember is to look at the kind of atoms attached as far out as necessary. Suppose we have to compare the aldehyde group, $\ce{-CH=O}$, with the dimethoxymethyl group, $\ce{CH(OCH_3)_2}$. The first atoms are the same $\left( \ce{C} \right)$, the second atoms are the same $\left[ \ce{O}, \: \left( \ce{O} \right), \: \ce{H} \right]$, and the difference arrives at the third-atom level where we are comparing lone pairs (priority zero) with carbons. Thus $\ce{-CH(OCH_3)_2}$, outranks $\ce{-CH=O}$.
Comparison of groups such as isopropyl and ethenyl is more difficult and requires knowing what the convention is when we have to go to the far end of a double bond. A useful way of writing these groups is as follows:
We put ethenyl ahead of isopropyl because $\left[ \left( \ce{C} \right), \: \ce{H}, \: \ce{H} \right]$ takes priority over $\left( \ce{H}, \: \ce{H}, \: \ce{H} \right)$. It is important to understand that the nonduplicated carbon is considered to be connected to the duplicated carbon as well as the two hydrogens in arriving at the connection pattern $\left( \left( \ce{C} \right), \: \ce{H}, \: \ce{H} \right)$.
The same kind of logic leads to the following sequence:
A more comprehensive list of priorities among groups is given in Table 19-1. It will be a good exercise to go through this list and work out how the priorities are established.
Table 19-1: Some Common Groups in Order of Increasing Sequence-Rule Preference
If more than one chiral center is present, the configuration at each is specified by the symbol $R$ or $S$ together with the number of the chiral atom. Thus the configuration of $\left( + \right)$-tartaric acid is known to be that designated in the name $\left( 2R,3R \right)$-$\left( + \right)$-tartaric acid:
The $R$,$S$ system is quite general and has many advantages (and a few disadvantages) compared with the $D$,$L$ notation for simple molecules. For diastereomers, it provides much clearer notations than meso, erythro,$^1$ and threo$^1$ that have been used for many years to designate the configurations of achiral and chiral diastereomers having two chiral carbon atoms:
The $R$,$S$ system can be used to designate the configuration of a molecule with no chiral carbons but with a chiral center as, for example, a chiral 1,2-diene (Section 13-5A). To do this for a 1,2-diene, the molecule is best drawn in projection, looking along the $\ce{C=C=C}$ bond with the highest-ranking group in front. The bonds in the rear will then project at $90^\text{o}$ to the bonds of the groups in front. For a 1,2-diene, ab$\ce{C=C=C}$xy, where is the highest ranking group, the possible enantomeric projections are:
We now determine the priority of the groups and then assign the configurations $R$ and $S$ as shown, provided that the highest-ranking group is in front and a $>$ b and x $>$ y. In proceeding this way, it is important to recognize that no matter what the priority is of the group b based on atomic number, b always outranks a rear group so that the priority sequence is a $\rightarrow$ b $\rightarrow$ x with $R$ clockwise and $S$ counterclockwise.
$^1$The prefixes erythro and threo are used for configurations of compounds with two differently substituted chiral carbons having similar groups on each carbon. If in the Fischer projection formula the similar groups are on the same side, the configuration is erythro. If the similar groups are on opposite sides, the configuration is threo.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/19%3A_More_on_Stereochemistry/19.06%3A_The_RS_Convention_for_Designating_Stereochemical_Configurations.txt |
The configuration about double bonds is undoubtedly best specified by the cis-trans notation when there is no ambiguity involved. Unfortunately, many compounds cannot be described adequately by the cis-trans system. Consider, for example, configurational isomers of 1-fluoro-1-chloro-2-bromo-2-iodo-ethene, \(9\) and \(10\). There is no obvious way in which the cis-trans system can be used:
A system that is easy to use and which is based on the sequence rules already described for the \(R\),\(S\) system works as follows:
1. An order of precedence is established for the two atoms or groups attached to each end of the double bond according to the sequence rules of Section 19-6. When these rules are applied to 1-fluoro-1-chloro-2-bromo-2-iodoethene, the priority sequence is:
• at carbon atom 1, \(\ce{Cl} > \ce{F}\)
• at carbon atom 2, \(\ce{I} > \ce{Br}\)
1. Examination of the two configurations shows that the two priority groups - one on each end - are either on the same side of the double bond or on opposite sides:
The \(Z\) isomer is designated as the isomer in which the top priority groups are on the same side (\(Z\) is taken from the German word zusammen - together). The \(E\) isomer has these groups on opposite sides (\(E\), German for entgegen - across).\(^2\) Two further examples show how the nomenclature is used:
This system is especially useful for oximes, which have the structural feature \(\ce{-C=N-OH}\). The two possible configurations at the double bond in the oxime of ethanal are \(11\) and \(12\):
The cis-trans notation does not work well here, and structure \(11\) has the \(Z\) configuration and \(12\) the \(E\) configuration. In the older chemical literature, these stereoisomers were designated as syn and anti forms, but these names are really no better than cis and trans.
\(^2\)It would have been simpler to remember if \(E\) stood for same side and \(Z\) for opposite side, but it is too late now.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
19.08: Prochirality
In the transformations shown in Equations 19-5 and 19-6, the organic reactants are symmetrical molecules (with no chiral centers), but the products are asymmetric molecules (each has a chiral carbon):
\(\tag{19-5}\)
\(\tag{19-6}\)
There is a special term for molecules that are achiral but which can be converted to molecules with chiral centers by a single chemical substitution or addition reaction. They are said to be prochiral.
By this definition, ethanol is a prochiral molecule. The two methylene hydrogens are enantiotopic (Section 9-10C) and substituting each separately (with, say, one deuterium) leads to a pair of enantiomers:
Prochiral molecules can be distinguished readily from more symmetrical molecules because they lack a two-fold symmetry axis passing through the prochiral center, as the following rotations show:
Organic chemists have not had much use for prochirality, but it is an important concept for biochemists following the stereochemistry of bio-organic reactions. Almost all biochemical reactions are under the control of enzymes, which function asymmetrically even on symmetrical (but prochiral) molecules. Thus it has been found that only one of the two methylene groups of citric acid, \(13\), is converted by enzymes (from rat liver) to the carbonyl of 2-oxobutanedioic acid:
The notation prochiral center is useful in molecules that already have one or more chiral centers. Development of chirality from prochirality in such cases would lead to diastereomers, as shown in the conversions of \(14\) and \(15\):
Contributors
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/19%3A_More_on_Stereochemistry/19.07%3A_EZ_Notation.txt |
Optical rotations usually are measured at just one wavelength, namely $589.3 \: \text{nm}$, simply because sodium-vapor lamps provide an especially convenient source of monochromatic light. Measurements at other wavelengths are less easily made without specialized instruments, with which relatively few laboratories are currently equipped. Nevertheless, much information has been obtained about structure, conformation, and configuration of organic compounds from measurements of optical rotation as a function of wavelength (i.e., optical rotatory dispersion).
Like other phenomena involving interactions between electromagnetic radiation and organic molecules, as in infrared, ultraviolet, and nmr spectroscopy, optical rotatory dispersion curves often are quite sensitive to small changes in structure. As an example, the rotatory dispersion curves for enantiomers of cis- and trans-10-methyl-2-decalones, $16$ and $17$, are reproduced in Figure 19-7:
Only a small positive rotation is observed for the particular enantiomers at the wavelength of the sodium line $\left( 589.3 \: \text{nm} \right)$ compared to the large, both positive and negative, rotations found at wavelengths between $270 \: \text{nm}$ and $400 \: \text{nm}$. If we measure the rotations as a function of wavelength and if, as we approach shorter wavelengths, the rotation rises to a maximum before changing sign, as it does with the trans isomer, $17$, then the compound is said to exhibit a positive Cotton effect. The opposite behavior, as with the cis isomer, $16$, is called a negative Cotton effect. The wavelength at the center point for the very rapid change in rotation for $17$ is $300 \: \text{nm}$ and corresponds to the $n \rightarrow \pi^*$ absorption maximum of the carbonyl group in the ultraviolet absorption curve of the same compound. Thus excitation of the carbonyl group by absorption of ultraviolet light and strong rotatory dispersion of polarized light are associated phenomena. In fact, when a substance exhibits a Cotton effect, not only does it transmit clockwise and counterclockwise circularly polarized light with unequal velocities (Section 19-1), it also absorbs the two forms of light unequally.
This means that the molar extinction coefficients of the two enantiomers ($\epsilon_l$ and $\epsilon_r$) are unequal in circularly polarized light. These differences in absorption ($\epsilon_l$ and $\epsilon_r$) can be measured as a function of wavelength, and the curves obtained are called circular dichroism curves. They have positive or negative signs (Cotton effect) just as for optical rotatory dispersion curves.
Most of the research on optical rotatory dispersion to date has been with optically active ketones because the carbonyl chromophore conveniently has a weak absorption band in the $300 \: \text{nm}$ region. Compounds with chromophores that absorb light strongly in the ultraviolet usually are unsatisfactory for rotatory dispersion measurements because insufficient incident light is transmitted to permit measurement of optical rotation. Weak absorption bands below about $210 \: \text{nm}$ have not been exploited because of experimental difficulties in making the necessary measurements.
Many rotatory dispersion curves have been obtained for optically active ketones derived from steroids and triterpenes, which are monocyclic, bicyclic, and open-chain compounds. Enough data have been accumulated so that the various shapes and magnitudes of the curves are recognized as characteristic of particular structural features. A good illustration is provided by the rotatory dispersion curves for the cis- and trans-8-methylhydrindan-5-ones, $18$ and $19$, which are shown in Figure 19-8:
The remarkable differences in these curves are due to changes in the environment of the carbonyl groups arising from the different configurations of the hydrogens at the ring junctions. Because the rotatory dispersion curve of the closely related structure $20a$ is very similar to that of the cis-hydrindanone, $18$, the rings labeled $A$ and $B$ in $20a$ can be inferred also to be cis oriented (see Figure 19-8):
Rotatory dispersion curves often are helpful in establishing configurations; thus the relative configurations of compounds $18$ and $20a$ must be the same because, if they were not, the two curves would resemble mirror images of one another. Therefore, if the absolute configuration of $18$ corresponds to the formula shown, then compound $20a$ has the configuration shown and not $20b$.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/19%3A_More_on_Stereochemistry/19.09%3A_Optical_Rotatory_Dispersion_and_Circular_Dichroism.txt |
If one could prepare 2-hydroxypropanenitrile from ethanal and hydrogen cyanide in the absence of any chiral reagent and produce an excess of one enantiomer over the other, this would constitute an absolute asymmetric synthesis - that is, creation of preferential chirality (optical activity) in a symmetrical environment from symmetrical reagents:
This obviously is unlikely for the given example because there is no reason for cyanide ion to have anything other than an exactly equal chance of attacking above or below the plane of the ethanal molecule, producing equal numbers of molecules of the enantiomers, $21$ and $22$. However, when a chiral center is created through reaction with a dissymmetric (chiral) reagent, we should not expect an exactly 1:1 mixture of the two possible isomers. For example, in an aldol-type addition (Section 18-8E) of a chiral ester to a prochiral ketone the two configurations at the new chiral center in the products $23$ and $24$ are not equally favored. That is to say, asymmetric synthesis is achieved by the influence of one chiral center $\left( \ce{R}^* \right)$ on the development of the second:
You will notice that the reaction products $23$ and $24$ are diastereomers, not enantiomers. Asymmetric synthesis can be achieved only when the possible transition states for reaction are diastereomeric because they then will have different energies and will lead to products at different rates. The larger the energy difference between diastereomeric transition states, the more stereochemical preference there will be for one chirality over the other.
The degree of stereochemical control displayed by the first chiral center usually depends on how close it is to the second - the more widely separated they are, the less steric control there is. Another factor is the degree of electronic control. If all the groups are very much the same electrically and sterically, not much stereochemical control is to be expected. Even when the chiral centers are close neighbors, asymmetric induction is seldom $100\%$ efficient in simple molecules. In biochemical systems, however, asymmetric synthesis is highly efficient.
The stereospecificity of living organisms is imperative to their efficiency. The reason is that it is just not possible for an organism to be so constructed as to be able to deal with all of the theoretically possible isomers of molecules with many asymmetric centers. Thus a protein molecule not uncommonly has 100 or more different asymmetric centers; such a molecule would have $2^{100}$ or $10^{30}$ possible optical isomers. A vessel with a capacity on the order of $10^7$ liters would be required to hold all of the possible stereoisomeric molecules of this structure if no two were identical. An organism so constituted as to be able to deal specifically with each one of these isomers would be very large indeed.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/19%3A_More_on_Stereochemistry/19.10%3A_Asymmetric_Synthesis.txt |
Optically active biphenyl derivatives (Section 13-5A) are racemized if the two aromatic rings at any time pass through a coplanar configuration by rotation about the central bond. This can be brought about more or less easily by heat, unless the 2,2'-ortho substituents are very large.
The way in which compounds with asymmetric carbon atoms are racemized is more complicated. One possibility would be for a tetrahedral chiral carbon attached to four groups to become planar and achiral without breaking any bonds. Theoretical calculations indicate that this is not a likely process for chiral tetravalent carbon but, as we will see, it does occur with chiral carbon and other chiral atoms that are attached to three groups:
Optically active carbonyl compounds of the type $\ce{-CHC=O}$, in which the alpha carbon is asymmetric, are racemized by both acids and bases, and from Section 17-1 we can be sure that this is related to enolization. Formation of either the enol or the enolate anion will destroy the asymmetry of the $\alpha$ carbon so that, even if only trace amounts of enol are present at any given time, eventually all of the compound will be racemized. However, the mechanism requires both that there be an $\alpha$ hydrogen and that the center of symmetry be located at this $\alpha$ carbon. Otherwise, acids and bases are ineffective in catalyzing racemization.
The racemization of an optically active secondary halide with the chiral carbon carrying the halogen (e.g., 2-chlorobutane) may occur ih solution and, usually, the more polar and better ionizing the solvent is, the more readily the substance is racemized. Ionization of the halide by an $S_\text{N}1$ process probably is responsible, and this certainly would be promoted by polar solvents (Section 8-6). All indications are that an alkyl carbocation once dissociated from its accompanying anion is planar; and, when such an ion recombines with the anion, it has equal probability of forming the $D$ and $L$ enantiomers:
Optically active halides also can be racemized by an $S_\text{N}2$ mechanism. A solution of active 2-chlorobutane in 2-propanone containing dissolved lithium chloride becomes racemic. Displacement of the chloride of the halide by chloride ion inverts configuration at the atom undergoing substitution (see Section 8-5). A second substitution regenerates the original enantiomer. Eventually, this back-and-forth process produces equal numbers of the $D$ and $L$ forms; the substance then is racemic:
Asymmetric alcohols often are racemized by strong acids. Undoubtedly, ionization takes place, and recombination of the carbocation with water leads to either enantiomer:
In contrast to halides, alcohols, and carbonyl compounds, hydrocarbons may be extremely difficult to racemize. This is particularly true for a compound with a quaternary asymmetric center, such as methylethylpropylbutylmethane, $28$, which has no "handle" to allow one to convert the asymmetric carbon to a symmetric condition by simple chemical means:
However, hydrocarbons that have a hydrogen atom at the asymmetric carbon may be racemized if they can be converted either to carbocations or to carbanions. The ease of carbanion-type racemization will depend on the acidity of the attached hydrogen and on the stereochemical stability of the intermediate carbanion that is formed. If the configuration of the carbanion intermediate inverts, racemization will result (also see Section 6-4E):
The carbocation type of racemization of an optically active hydrocarbon can occur by the exchange reaction described in Section 10-9.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/19%3A_More_on_Stereochemistry/19.11%3A_Racemization.txt |
Carbohydrates are a major class of naturally occurring organic compounds, which come by their name because they usually have, or approximate, the general formula \(\ce{C}_n \ce{(H_2O)}_m\), with \(n\) equal to or greater than three. Among the well-known carbohydrates are various sugars, starches, and cellulose, all of which are important for the maintenance of life in both plants and animals. Although the structures of many carbohydrates appear to be quite complex, the chemistry of these substances usually involves only two functional groups - ketone or aldehyde carbonyls and alcohol hydroxyl groups. The carbonyl groups normally do not occur as such, but are combined with hydroxyl groups to form hemiacetal or acetal linkages of the kind discussed in Section 15-4E. An understanding of stereochemistry is particularly important to understanding the properties of carbohydrates. Configurational and conformational isomerism play an important role. For this reason, you may wish to review Chapter 5 and Sections 12-3 and 19-5.
• 20.1: Prelude to Carbohydrates
Although the structures of many carbohydrates appear to be quite complex, the chemistry of these substances usually involves only two functional groups - ketone or aldehyde carbonyls and alcohol hydroxyl groups. The carbonyl groups normally do not occur as such, but are combined with hydroxyl groups to form hemiacetal or acetal linkages.
• 20.2: Classification and Occurrence of Carbohydrates
The simple sugars, or monosaccharides, are the building blocks of carbohydrate chemistry. They are polyhydroxy aldehydes or ketones with five, six, seven, or eight carbon atoms that are classified appropriately as pentoses, hexoses, heptoses, or octoses, respectively. They can be designated by more specific names, such as aldohexose or ketohexose, to denote the kind of carbonyl compound they represent.
• 20.3: The Structure and Properties of D-Glucose
Glucose is by far the most abundant monosaccharide; it occurs free in fruits, plants, honey, in the blood of animals, and combined in many glycosides, disaccharides, and polysaccharides. The structure and properties of glucose will be considered in greater detail than those of the other monosaccharides, not only because of its importance, but because much of what can be said about glucose also can be said about the other monosaccharides.
• 20.4: Conventions for Indicating Ring Size and Anomer Configurations of Monosaccharides
The oxide ring is six-membered in some sugars and five-membered in others, and it is helpful to use names that indicate the ring size. The five- and six-membered oxide rings bear a formal relationship to oxa-2,5-cyclohexadiene and oxa-2,4-cyclopentadiene that commonly are known as pyran and furan, respectively. For this reason, the names furanose and pyranose have been coined to denote five- and six-membered rings in cyclic sugars.
• 20.5: Derivatives of Glucose
Although we now have powerful spectroscopic methods available to determine the sizes of the oxide rings formed by the simple monosaccharides, the way in which this was done chemically for glucose highlights the difference in reactivity between ether and alcohol functions. The acid-catalyzed methylation of glucose with methanol to give two distinct glucosides corresponds to displacement of the hemiacetal hydroxyl by methoxyl to form an acetal.
• 20.6: Glycosides
Although abundant quantities of glucose and fructose are found in the free state, they and less common sugars occur widely in plants and animals combined with various hydroxy compounds. The bonding is through oxygen to the carbonyl carbon, as in the αα - and ββ -methylglucosides discussed in Section 20-4A, to give acetal or ketal structures. These substances are sometimes simply called glycosides.
• 20.7: Disaccharides
Combinations of two or more of the simple sugars through glycoside linkages give substances known as polysaccharides. They also are called oligosaccharides if made from two to ten sugar units. The simplest oligosaccharides are disaccharides made of two molecules of simple sugars.
• 20.8: Polysaccharides
The fibrous tissue in the cell walls of plants contains the polysaccharide cellulose, which consists of long chains of glucose units. A second, very widely distributed polysaccharide is starch, which is stored in the seeds, roots, and fibers of plants as a food reserve - a potential source of glucose. The chemical composition of starch varies, but there are two structurally different polysaccharides. One is a linear structure (amylose) and the other is a branched structure (amylopectin).
• 20.9: Vitamin C
The "antiscorbutic" factor of fresh fruits, which prevents the development of the typical symptoms of scurvy in humans, is a carbohydrate derivative known as vitamin C or ascorbic acid. This substance is not a carboxylic acid, but a lactone, and owes its acidic properties (and ease of oxidation) to the presence of an enediol grouping. It belongs to the L series by the glyceraldehyde convention.
• 20.10: Formation of Carbohydrates by Photosynthesis
Carbohydrates are formed in green plants by photosynthesis, which is the chemical combination, or "fixation", of carbon dioxide and water by utilization of energy from the absorption of visible light. The overall result is the reduction of carbon dioxide to carbohydrate and the formation of oxygen. If the carbohydrate formed is cellulose, then the reaction in effect is the reverse of the burning of wood, and obviously requires considerable energy input.
• 20.11: The Generation of Energy from Carbohydrate Metabolism
This section is concerned mainly with the pathway by which glucose is metabolized by the process known as glycolysis. Initially, the storage fuels or foodstuffs (fats, carbohydrates, and proteins) are hydrolyzed into smaller components (fatty acids and glycerol, glucose and other simple sugars, and amino acids). In the next stage, these simple fuels are degraded further to two-carbon fragments.
• 20.E: Carbohydrates (Exercises)
These are the homework exercises to accompany Chapter 20 of the Textmap for Basic Principles of Organic Chemistry (Roberts and Caserio).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
20: Carbohydrates
Carbohydrates are a major class of naturally occurring organic compounds, which come by their name because they usually have, or approximate, the general formula \(\ce{C}_n \ce{(H_2O)}_m\), with \(n\) equal to or greater than three. Among the well-known carbohydrates are various sugars, starches, and cellulose, all of which are important for the maintenance of life in both plants and animals.
Although the structures of many carbohydrates appear to be quite complex, the chemistry of these substances usually involves only two functional groups - ketone or aldehyde carbonyls and alcohol hydroxyl groups. The carbonyl groups normally do not occur as such, but are combined with hydroxyl groups to form hemiacetal or acetal linkages of the kind discussed in Section 15-4E.
An understanding of stereochemistry is particularly important to understanding the properties of carbohydrates. Configurational and conformational isomerism play an important role. For this reason, you may wish to review Chapter 5 and Sections 12-3 and 19-5.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
20.02: Classification and Occurrence of Carbohydrates
The simple sugars, or monosaccharides, are the building blocks of carbohydrate chemistry. They are polyhydroxy aldehydes or ketones with five, six, seven, or eight carbon atoms that are classified appropriately as pentoses, hexoses, heptoses, or octoses, respectively. They can be designated by more specific names, such as aldohexose or ketohexose, to denote the kind of carbonyl compound they represent. For example, and aldopentose is a five-carbon sugar with an aldehyde carbonyl; a ketohexose is a six-carbon sugar with a ketone carbonyl:
However, it is important to keep in mind that the carbonyl groups of sugars usually are combined with one of the hydroxyl groups in the same molecule to form a cyclic hemiacetal or hemiketal. These structures once were written as follows, and considerable stretch of the imagination is needed to recognize that they represent oxacycloalkane ring systems:
The saccharides have long and awkward names by the IUPAC system, consequently a highly specialized nomenclature system has been developed for carbohydrates. Because this system (and those like it for other natural products) is unlikely to be replaced by more systematic names, you will find it necessary to memorize some names and structures. It will help you to remember the meaning of names such as aldopentose and ketohexose, and to learn the names and details of the structures of glucose, fructose, and ribose. For the rest of the carbohydrates, the nonspecialist needs only to remember the kind of compounds that they are.
The most abundant five-carbon sugars are $L$-arabinose, $D$-ribose, 2-deoxy-$D$-ribose,$^1$ and $D$-xylose, which all are aldopentoses. Both the open-chain and cyclic structures of the $D$-aldoses up to $\ce{C_6}$ are shown in Figure 20-1.
The common six-carbon sugars (hexoses) are $D$-glucose, $D$-fructose, $D$-galactose, and $D$-mannose. They all are aldohexoses, except $D$-fructose, which is a ketohexose. The structures of the ketoses up to $\ce{C_6}$ are shown for reference in Figure 20-2. The occurrence and uses of the more important ketoses and aldoses are summarized in Table 20-1.
Table 20-1: Occurrence, Physical Properties, and Uses of Some Natural Sugars
$^1$Deoxy means lacking a hydroxyl group, and 2-deoxyribose is simply ribose without an $\ce{OH}$ group at the 2-carbon.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/20%3A_Carbohydrates/20.01%3A_Prelude_to_Carbohydrates.txt |
Configuration
Glucose is by far the most abundant monosaccharide; it occurs free in fruits, plants, honey, in the blood of animals, and combined in many glycosides, disaccharides, and polysaccharides. The structure and properties of glucose will be considered in greater detail than those of the other monosaccharides, not only because of its importance, but because much of what can be said about glucose also can be said about the other monosaccharides.
Glucose is an aldohexose, which means that it is a six-carbon sugar with a terminal aldehyde group, shown by $1$:
The carbons labeled with an asterisk in $1$ are chiral; thus there are $2^4$, or sixteen, possible configurational isomers. All are known - some occur naturally and the others have been synthesized (see Table 20-1). The problem of identifying glucose as a particular one of the sixteen possibilities was solved by Emil Fischer during the latter part of the nineteenth century, for which he was awarded the Nobel Prize in chemistry in 1902. The configurations he deduced for each of the chiral carbons, $\ce{C_2}$-$\ce{C_5}$, are shown in the projection formula $2$.$^2$ Although Fischer was aware that natural glucose could be the enantiomer of Structure $2$, his original guess at the absolute configuration proved to be correct and the configuration at $\ce{C_5}$ is the same as the configuration of the simplest "sugar" $D$-$\left( + \right)$-glyceraldehyde, $3$, (see Section 19-5). Therefore natural glucose is specifically $D$-glucose:
The complete logic of Fischer's procedures for determination of the configuration of glucose is too involved to be explained here in detail. What you will be unable to fully appreciate is the great difficulties of working with carbohydrates - that is, their considerable solubility in water, instability to strong oxidizing agents and acidic or basic reagents, reluctance to crystallize, and their tendency to decompose rather than give sharp melting points. Fortunately for Fischer, many different pentoses and hexoses already were available from the efforts of earlier investigators, and the principles of optical isomerism were well understood as the result of the work of van't Hoff.
Two of the key ideas used by Fischer can be illustrated best with aldotetroses because they have only two chiral carbons and far fewer possible structures to consider. Writing the four possibilities as the aldehyde rather than hemiacetal structures, we have $4$-$7$. Of these, $4$ and $5$ constitute a pair of enantiomers, as do $6$ and $7$. These pairs can be identified by careful oxidation of the terminal groups to give the corresponding tartaric (2,3-dihydroxybutanedioic) acids. Oxidation of both $4$ and $5$ gives meso-tartaric acid, whereas oxidation of $6$ and $7$ gives, respectively, $\left( + \right)$ and $\left( - \right)$ tartaric acids:
It should be clear from this that the configurations of $6$ and $7$ are established by being related to the respective chiral tartaric acids. However, we have no way of telling which of the tetroses represented by $4$ and $5$ is $D$ and which is $L$, because, on oxidation, they give the same achiral tartaric acid. What we need to do is relate one or the other of the chiral carbons of these tetroses to the corresponding carbon of either $6$ or $7$. One way that this can be done is by the Wohl degradation, whereby the chain length is reduced by one carbon by removing the aldehyde carbon:
As applied to $4$, $5$, $6$, and $7$, the Wohl degradation forms enantiomers of glyceraldehyde:
Here we see that $4$ and $6$ give the same enantiomer, $D$-glyceraldehyde, and therefore have the same configuration of their highest-numbered asymmetric carbon. In contrast, $5$ and $7$ give $L$-glyceraldehyde and thus must be similarly related. By this kind of procedure, the configurations of $4$ to $7$ can be established unequivocally, although, as you might imagine, it is far easier to do on paper than in the laboratory.
Knowing the configurations of the tetroses aids in establishing the configurations of the pentoses. Thus, $4$, by Kiliani-Fischer cyanohydrin synthesis,$^3$ can be converted to a mixture of two aldopentoses, $8$ and $9$, by adding a carbon at the aldehyde end of the molecule. The configurations of the two carbons at the lower end of the starting materials remain unchanged, but two diastereomeric aldopentoses are formed in the syntheses because a new chiral center is created:
Products $8$ and $9$ present a new configurational problem, but a less difficult one than before, because the configurations of two of the three chiral centers already are known. Controlled oxidation of $8$ and $9$ will give different diastereomeric 2,3,4-trihydroxypentanedioic acids, $10$ and $11$, respectively:
Of these, $11$ is achiral (meso), whereas $10$ is chiral. Therefore, by simply determining which oxidation product is optically active, and hence chiral, we can assign the configurations of $8$ and $9$. Direct comparison of these synthetic aldopentoses with the naturally occurring compounds then could be used as proof of the structure of natural aldopentoses. By this reasoning $8$ turns out to be $D$-arabinose and $9$ is $D$-ribose.
Some of the key reactions in carbohydrate chemistry involve oxidation of aldoses to carboxylic acids. There is a simple nomenclature system for these acids. In abbreviated notation, the products of oxidation at $\ce{C_1}$, $\ce{C_6}$, or both are called:
The carboxylic acids derived from glucose are therefore gluconic acid, glucuronic acid, and glucaric acid.
Hemiacetal Formation. Anomers of Glucose
Although glucose has some of the properties expected of an aldehyde, it lacks others. For example, it forms certain carbonyl derivatives (e.g., oxime and cyanohydrin), and can be reduced to the hexahydroxyhexane (sorbitol), and oxidized with bromine to gluconic acid (a monocarboxylic acid). (With nitric acid, oxidation proceeds further to give the dicarboxylic acid, $D$-glucaric acid.)
Glucose also will reduce Fehling's solution $\left[ \ce{Cu} (II) \rightarrow \ce{Cu} (I) \right]$ and Tollen's reagent $\left[ \ce{Ag} (I) \rightarrow \ce{Ag} (0) \right]$ and, for this reason, is classified as a reducing sugar.$^4$ However, it fails to give a hydrogen sulfite addition compound and, although it will react with amines $\left( \ce{RNH_2} \right)$, the products are not the expected Schiff's bases of the type $\ce{-C=NR}$. Furthermore, glucose forms two different monomethyl derivatives (called methyl $\alpha$-$D$-glucoside and methyl $\beta$-$D$-glucoside) under conditions that normally convert an aldehyde to a dimethyl acetal:
All of these reactions can be explained on the bases that the carbonyl group is not free in glucose but is combined with one of the hydroxyl groups, which turns out to be the one at $\ce{C_5}$, to form a hemiacetal, $12$ or $13$. Why are two hemiacetals possible? Because a new asymmetric center is created at $\ce{C_1}$ by hemiacetal formation, and this leads to diastereomeric forms of $D$-glucose called $\alpha$-$D$-glucose and $\beta$-$D$-glucose. In general, carbohydrate stereoisomers that differ only in configuration at the hemiacetal carbon are called anomers:
Although formulas $12$ and $13$ show the configurations at each of the chiral centers, they do not provide the crucial information for understanding the properties of glucose with respect to the arrangement of the atoms in space. Conversion of a projection formula such as $12$ or $13$ to a three-dimensional representation is not at all a trivial task. We have indicated the procedure for doing in before (Sections 5-3C and 5-5). The result of these procedures applied to $12$ and $13$ are the so-called Haworth projection formulas, $14$ and $15$, and the sawhorse conformations, $16$ and $17$:
You should be able to satisfy yourself that the configuration at $\ce{C_5}$ is the same in both Fischer and Haworth representations. This amounts to asking if $18$ and $19$ represent the same configurations:
They do, but if you have trouble visualizing this, it will be very helpful to use a ball-and-stick model to see that $18$ and $19$ are different representations of the same configuration. If you do not have models, remember that if transposition of any three groups converts one projection into the other, the formulas are identical. Thus $18$ and $19$ have the same configuration because $18$ becomes $19$ by transposition of $\ce{C_4}$ with $\ce{CH_2OH}$, then $\ce{C_4}$ with $\ce{H}$.
X-ray studies of crystalline $\alpha$- and $\beta$-$D$-glucose show that these molecules have their atoms arranged in space as correspond to $16$ and $17$. This is what we would expect from our studies of cyclohexane conformations (Sections 12-3A to 12-3D), because for the $\beta$ form, all of the substituents on the oxacyclohexane ring are in equatorial positions, and for the $\alpha$ form, all except the hydroxyl at the anomeric carbon $\left( \ce{C_1} \right)$ are equatorial.
Mutarotation of the Anomeric Forms of Glucose
Although the crystalline forms of $\alpha$- and $\beta$-$D$-glucose are quite stable, in solution each form slowly changes into an equilibrium mixture of both. The process can be observed as a decrease in the optical rotation of the $\alpha$ anomer $\left( +112^\text{o} \right)$ or an increase for the $\beta$ anomer $\left( +18.7^\text{o} \right)$ to the equilibrium value of $52.5^\text{o}$. The phenomenon is known as mutarotation and commonly is observed for reducing sugars. Both acids and bases catalyze mutarotation; the mechanism, Equation 20-1, is virtually the same as described for acid- and base-catalyzed hemiacetal and hemiketal equilibria of aldehydes and ketones (see Section 15-4E):
At equilibrium, about $64\%$ of the $\beta$ anomer and $36\%$ of the $\alpha$ anomer are present. The amount of the free aldehyde form at equilibrium is very small (about 0.024 mole percent in neutral solution). Preponderance of the $\beta$ anomer is attributed to the fact that glucose exists in solution in the chair conformation with the large $\ce{-CH_2OH}$ group equatorial. In this conformation, the hydroxyl substituent at $\ce{C_1}$ is equatorial in the $\beta$ anomer and axial in the $\alpha$ anomer; hence the $\beta$ anomer is slightly more stable. When glucose is in the aldehyde form, the hydroxyl at $\ce{C_4}$ also could add to the aldehyde carbonyl to produce a hemiacetal with a five-membered ring. This does not occur to a significant degree with glucose because the hemiacetal with the six-membered ring and many equatorial groups is more stable. With other sugars, mixtures of five- and six-membered hemiacetal or ketal rings and their respective anomers are produced in water solution.
Carbon-13 nmr spectra (Section 9-10L) provide an especially powerful tool for studying the anomeric forms of sugars. For example, with glucose the resonances of $\ce{C_1}$, $\ce{C_3}$, and $\ce{C_5}$ of the $\alpha$ anomer are seen in Figure 20-3 to be shifted substantially upfield relative to those of the $\beta$ anomer because of the axial substituent effect (Section 12-3D).
Aldose $\leftrightharpoons$ Ketose Rearrangements
In the presence of dilute base, $D$-glucose rearranges to a mixture containing the anomers of $D$-glucose $\left( \sim 64\% \right)$, $D$-fructose $\left( \sim 31\% \right)$, and $D$-mannose $\left( 3\% \right)$. This interconversion undoubtedly involves enolization of the hexoses by way of a common enediol intermediate according to Equation 20-2:
The rearrangement is of interest because the corresponding enzymatic interconversion of aldoses and ketoses is an important part of the biosynthetic, photosynthetic, and metabolic pathways, as we shall see in Section 20-9. Although the biochemical rearrangement also may proceed by way of enediol intermediates, it is highly stereospecific and yields only one of two possible stereoisomeric aldoses. For example, glucose, but not mannose, can be enzymatically interconverted with fructose as the 6-phosphate ester derivative:
$^3$The steps of the Kiliani-Fischer synthesis are:
$^4$In general, reducing sugars are hemiacetals or hemiketals and the nonreducing sugars are acetals or ketals. The difference is between the presence of the structural elements $\ce{C-O-C-O-H}$ and $\ce{C-O-C-O-C}$.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/20%3A_Carbohydrates/20.03%3A_The_Structure_and_Properties_of_D-Glucose.txt |
The oxide ring is six-membered in some sugars and five-membered in others, and it is helpful to use names that indicate the ring size. The five- and six-membered oxide rings bear a formal relationship to oxa-2,5-cyclohexadiene and oxa-2,4-cyclopentadiene that commonly are known as pyran and furan, respectively:
There is an important question as to which one of the two anomeric forms of a sugar should be designated as $\alpha$ and which one as $\beta$. The convention is simple; the $\alpha$ anomer is the one that has the same configuration of the $\ce{OH}$ at the anomeric carbon as the carbon which determines the configuration of the sugar itself:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
20.05: Derivatives of Glucose
Determination of the Oxide Ring Size
Although we now have powerful spectroscopic methods available to determine the sizes of the oxide rings formed by the simple monosaccharides, the way in which this was done chemically for glucose highlights the difference in reactivity between ether and alcohol functions. The acid-catalyzed methylation of glucose with methanol to give two distinct glucosides, methyl $\alpha$-$D$-glucoside and methyl $\beta$-$D$-glucoside, corresponds to displacement of the hemiacetal hydroxyl by methoxyl to form an acetal (see left side of Figure 20-4).
The remaining four hydroxyl groups can be methylated in basic solution by dimethyl sulfate or by methyl iodide and silver oxide in $\ce{N}$,$\ce{N}$-dimethylmethanamide, $\ce{HCON(CH_3)_2}$, solution. Hydrolysis of either of these pentamethyl glucose derivatives with aqueous acid affects only the acetal linkage and leads to a tetramethylated glucose, $20$, as shown in Figure 20-4.
The pyranose ring structure of $D$-glucose originally was established by Hirst, in 1926, by converting $D$-glucose to a tetra-$\ce{O}$-methyl-$D$-glucose and showing that this substance actually was 2,3,4,6-tetra-$\ce{O}$-methyl-$D$-glucose, $20$. The key feature of $20$ is the fact that all but the two carbons involved hemiacetal formation are protected from oxidation by being substituted with $\ce{O}$-methyl groups in place of hydroxy groups. The largest fragment isolated from oxidation of Hirst's tetra-$\ce{O}$-methyl-$D$-glucose was a trimethoxypentanedioic acid, $21$, and because the two carboxyl carbons must have been the ones originally involved in ring formation, the oxide ring must be between $\ce{C_1}$ and $\ce{C_5}$:
Reagents that specifically oxidize vicinal glycols [e.g., $\ce{NaIO_4}$, $\ce{Pb(O_2CCH_3)_4}$, and $\ce{NaBiO_3}$; Section 16-9A] are quite helpful in determining the cyclic structures of sugars. With periodate, the numbers of moles of oxidant consumed and the moles of methanoic acid and methanal produced are different for each type of ring structure. The cleavage reactions that normally are observed follow:
Reactions with Amines and Hydrazines; Osazone Formation
As we stated previously, glucose forms some, but not all, of the common carbonyl derivatives. The amount of free aldehyde present in solution is so small that it is not surprising that no hydrogen sulfite derivative forms. With amines, the product is not a Schiff's base but a glucosylamine of cyclic structure analogous to the hemiacetal structure of glucose, Equation 20-3. The Schiff's base is likely to be an intermediate that rapidly cyclizes to the glucosylamine:
The reaction of glucose with an excess of phenylhydrazine (phenyldiazane) is particularly noteworthy because two phenylhydrazine molecules are incorporated into one of glucose. Subsequent to the expected phenylhydrazone formation, and in a manner that is not entirely clear, the $\ce{-CHOH}-$ group adjacent to the original aldehyde function is oxidized to a carbonyl group, which then consumes more phenylhydrazine to form a crystalline derivative called an osazone, or specifically glucose phenylosazone:
The sugar osazones usually are crystalline and are useful for characterization and identification of sugars. Fischer employed them in his work that established the configuration of the sugars. The kind of information that can be obtained is illustrated by the following example:
Because the same phenylosazone arises from glucose, mannose, and fructose, the configurations of $\ce{C_3}$, $\ce{C_4}$, and $\ce{C_5}$ must be the same for all three sugars.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/20%3A_Carbohydrates/20.04%3A_Conventions_for_Indicating_Ring_Size_and_Anomer_Configurations_of_Monosaccharides.txt |
Although abundant quantities of glucose and fructose are found in the free state, they and less common sugars occur widely in plants and animals combined with various hydroxy compounds. The bonding is through oxygen to the carbonyl carbon, as in the $\alpha$- and $\beta$-methylglucosides discussed in Section 20-4A, to give acetal or ketal structures. These substances are sometimes simply called glycosides, but it is desirable to specify that the bonding is through oxygen by using the name $\ce{O}$-glycoside. Hydrolysis of an $\ce{O}$-glycoside gives the sugar and the hydroxy compound, called the aglycone component of the glycoside.
A specific example is glucovanillin, which can be isolated from the green fruit pods of vanilla, a climbing orchid cultivated in several tropical countries. Hydrolysis gives glucose and the aglycone, vanillin, which is the principal ingredient of vanilla flavoring. As the vanilla pods mature, a natural hydrolysis reaction proceeds to the extent that the pods may be covered with small crystals of vanillin.
The configurations of glycosides are designated by the same convention used for the sugar anomers. Thus if a glycoside of a $D$ sugar has the $D$ configuration at the anomeric carbon, it is designated as the $\alpha$-$D$-glycoside, and if it has the $L$ configuration it is called the $\beta$-$D$-glycoside (see Section 20-3). If the sugar involved in glycoside formation is glucose, the derivative is a glucoside; if fructose, a fructoside; if galactose, a galactoside; and so on. When the hydroxy compound, or aglycone, is another sugar, then the glycoside is a disaccharide, and if the sugar is already a dissacharide, the glycoside is a trisaccharide, and so on.
Among the natural products that occur as glycosides (most commonly as $\beta$-$D$-glucosides) are many plant pigments (the anthocyanins), the flavorings vanillin and amygdalin, and many steroids (such as the cardiac glycosides and saponins). The structures of some of these substances will be discussed in later chapters.
Not all glycosides are $\ce{O}$-glycosides. A group of $\ce{N}$-glycosides of special biological importance are derived from heterocyclic nitrogen bases and $D$-ribose and 2-deoxy-$D$-ribose. They commonly are known as nucleosides, or more specifically, as ribonucleosides and deoxyribonucleosides; the $\ce{N}$-glycoside linkage is always $\beta$:
The $\ce{N}$-glycoside of $D$-ribose and the nitrogen heterocycle, adenine, is adenosine:
A nucleotide is a phosphate ester of a nucleoside. The hydroxyl group at the $\ce{C_5}$ position of the pentose is the most common site of esterification. The nucleotides of adenosine are ATP, ADP, and AMP (Section 15-5F).
A dinucleotide is a combination of two nucleosides through a common phosphate ester link. Familiar examples are $\ce{NAD}^\oplus$, $\ce{NADH}$, $\ce{FAD}$, and $\ce{FADH_2}$ (Section 15-6C). Polynucleotides are polymers of nucleosides linked through phosphate ester bonds. Polynucleotides also are called nucleic acids (RNA and DNA) and are the genetic material of cells, as will be discussed in Chapter 25.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/20%3A_Carbohydrates/20.06%3A_Glycosides.txt |
General Types and Properties
Combinations of two or more of the simple sugars through glycoside linkages give substances known as polysaccharides. They also are called oligosaccharides if made from two to ten sugar units. The simplest oligosaccharides are disaccharides made of two molecules of simple sugars that can be joined with $\ce{O}$-glycoside links, and it probably is easiest to visualize these as shown in the "stripped-down" formulas, $22$ and $23$:
You should look at $22$ and $23$ carefully to be sure that you recognize the difference between them.$^5$ In $22$, sugar A is acting as a simple hydroxy compound, the aglycone of the sugar G to which it is linked by an $\ce{O}$-glycoside linkage.$^6$ Hydrolysis of $22$ at the glycoside link then will proceed as follows:
Disaccharides such as $22$ are like glucose in being reducing sugars (Section 20-2B), because Component A has the hemiacetal grouping that is opened easily to the aldehyde form in the mildly alkaline conditions used for the Tollen's and Fehling's solution oxidations. Because there is a free hemiacetal group, reducing sugars also form osazones and they mutarotate (Sections 20-4B and 20-2C).
Disaccharides of type $23$ are different in that each sugar, G and G', is acting as both a glycoside sugar and as an aglycone. The linkage between them is that of a double-barreled acetal, $\ce{-O-C-O-C-O}-$, and there is no hemiacetal grouping in the molecule. Therefore these are nonreducing sugars as far as the standard tests go. However, hydrolysis of the $\ce{O}$-glycoside linkages of $23$ does generate reducing sugars with hemiacetal carbons:
In general, we find that the nonreducing disaccharides give none of the carbonyl reactions observed for glucose, such as mutarotation and osazone formation, except when the conditions are sufficiently acidic to hydrolyze the acetal linkage.
Among the more important disaccharides are sucrose, $24$, maltose, $25$, cellobiose, $26$, and lactose, $27$:
Sucrose and lactose occur widely as the free sugars, lactose in the milk of mammals, and sucrose in fruit and plants (especially in sugar cane and sugar beet). Maltose is the product of enzymatic hydrolysis of starch, and cellobiose is a product of hydrolysis of cellulose.
To fully establish the structure of a disaccharide, we must determine (1) the identity of the component monosaccharides; (2) the type of ring junction, furanose or pyranose, in each monosaccharide, as it exists in the disaccharide; (3) the positions that link one monosaccharide with the other; and (4) the anomeric configuration ($\alpha$ or $\beta$) of this linkage.
Hydrolysis of disaccharides with enzymes is very helpful in establishing anomeric configurations, because enzymes are highly specific catalysts for hydrolysis of the different types of glycosidic linkages. For instance, $\alpha$-$D$-glucosidase (maltase) catalyzes hydrolysis of $\alpha$-$D$-glycosides more rapidly than of $\beta$-$D$-glycosides. The enzyme emulsin (found in bitter almonds) in contrast shows a strong preference for $\beta$-$D$-glycosides over $\alpha$-$D$-glycosides. Yeast invertase catalyzes hydrolysis of $\beta$-$D$-fructosides.
Structure of Sucrose
We know that sucrose consists of the two monosaccharides glucose and fructose because hydrolysis with acids or enzymes gives equal amounts of each hexose. Further, sucrose is not a reducing sugar, it forms no phenylosazone derivative, and it does not mutarotate. Therefore the anomeric carbons of both glucose and fructose must be linked through an oxygen bridge in sucrose. Thus sucrose is a glycosyl fructoside or, equally, a fructosyl glucoside.
Because sucrose is hydrolyzed by enzymes that specifically assist hydrolysis of both $\alpha$ glycosides (such as yeast $\alpha$-glucosidase) and $\beta$-fructosides (such as invertase), it is inferred that the glucose residue is present as an $\alpha$ glucoside and the fructose residue as a $\beta$ fructoside. If so, the remaining uncertainty in the structure of sucrose is the size of the rings in the glucose and fructose residues.
The size of the sugar rings in sucrose has been determined by the reactions shown in Figure 20-5. Methylation of sucrose with dimethyl sulfate in basic solution followed by hydrolysis of the octamethyl derivative gives 2,3,4,6-tetra-$\ce{O}$-methyl-$D$-glucopyranose (Section 20-4) and a tetra-$\ce{O}$-methyl-$D$-fructose. This establishes the glucose reside in sucrose as a glucopyranose. The fructose residue must be a fructofuranose because periodate oxidation of sucrose consumes three moles of periodate, whereby one mole of methanoic acid and one mole of a tetraaldehyde are formed. On bromine oxidation followed by acid hydrolysis, the tetraaldehyde gives 3-hydroxy-2-oxopropanoic acid (hydroxypyruvic acid, $\ce{HOCH_2COCO_2H}$), oxoethanoic acid (glyoxylic acid, $\ce{OCHCO_2H}$), and $D$-glyceric acid $\left( \ce{HOCH_2CHOHCO_2H} \right)$. Sucrose therefore has structure $24$, and this structure was confirmed by synthesis (R. Lemieux in 1953).
$^5$For now, we will ignore the possibility of different anomers of the disaccharide or their component sugars.
$^6$The manner in which sugars are linked together to form oligosaccharides was elucidated by W. N. Haworth, who received the Nobel Prize in chemistry in 1937 for this and other contributions to research on the structures and reactions of carbohydrates.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/20%3A_Carbohydrates/20.07%3A_Disaccharides.txt |
Cellulose
The fibrous tissue in the cell walls of plants contains the polysaccharide cellulose, which consists of long chains of glucose units, each of which is connected by a $\beta$-glucoside link to the $\ce{C_4}$ hydroxyl of another glucose as in the disaccharide cellobiose (i.e., $\beta$-1,4):
Indeed, enzymatic hydrolysis of cellulose leads to cellobiose. The molecular weight of cellulose varies with the source but is usually high. Cotton cellulose appears to have about 3000 glucose units per molecule.
The natural fibers obtained from cotton, wood, flax, hemp, and jute all are cellulose fibers and serve as raw materials for the textile and paper industries. In addition to its use as a natural fiber and in those industries that depend on wood as a construction material, cellulose is used to make cellulose acetate (for making rayon acetate yarn, photographic film, and cellulose acetate butyrate plastics), nitric acid esters (gun cotton and celluloid$^7$), and cellulose xanthate (for making viscose rayon fibers). The process by which viscose rayon is manufactured involves converting wood pulp or cotton linters into cellulose xanthate by reaction with carbon disulfide and sodium hydroxide:
The length of the chains of the cellulose decreases about 300 monomer units in the process. At this point, the cellulose is regenerated in the form of fine filaments by forcing the xanthate solution through a spinneret into an acid bath:
A few animals (especially ruminants and termites) are able to metabolize cellulose, but even these animals depend on appropriate microorganisms in their intestinal tracts to hydrolyze the $\beta$-1,4 links; other animals, including man, cannot utilize cellulose as food because they lack the necessary hydrolytic enzymes. However, such enzymes are distributed widely in nature. In fact, deterioration of cellulose materials - textiles, paper, and wood - by enzymatic degradation (such as by dry rot) is an economic problem that is not yet adequately solved. Efforts to turn this to advantage through enzymatic hydrolysis of cellulose to glucose for practical food production have not been very successful (Section 25-12).
Starch and Related Compounds
A second, very widely distributed polysaccharide is starch, which is stored in the seeds, roots, and fibers of plants as a food reserve - a potential source of glucose. The chemical composition of starch varies with the source, but in any one starch there are two structurally different polysaccharides. Both consist entirely of glucose units, but one is a linear structure (amylose) and the other is a branched structure (amylopectin).
The amylose form of starch consists of repeating 1,4-glucopyranose links as in cellulose, but unlike cellulose the linkage is $\alpha$ rather than $\beta$ (i.e., $\alpha$-1,4):
Hydrolysis by the enzyme diastase leads to maltose.
In amylopectin, amylose chains are joined by $\alpha$-1,6 linkages:
Animals also store glucose in the form of starchlike substances called glycogens. These substances resemble amylopectin more than amylose in that they are branched chains of glucose units with $\alpha$-1,4- and $\alpha$-1,6-glucoside links.
Starch is used in paper manufacture and in the textile and food industries. Fermentation of grain starches is an important source of ethanol. Hydrolysis of starch catalyzed by hydrochloric acid results in a syrupy mixture of glucose, maltose, and higher-molecular-weight saccharides. This mixture is called dextrin and is marketed as corn syrup. The hydrolysis does not proceed all the way to glucose because the $\alpha$-1,6 glucosidic link at the branch point is not easily hydrolyzed. Enzymes also catalyze hydrolysis of starch, but the enzyme $\alpha$ amylase is specific for $\alpha$-1,4 links and, like acid-catalyzed hydrolysis, gives a mixture of glucose, maltose, and polysaccharides (dextrin). The enzyme $\alpha$-1,6-glucosidase can hydrolyze the $\alpha$-1,6 links at the branch points and, when used in conjunction with $\alpha$ amylase, completes the hydrolysis of starch to glucose.
A very interesting group of polysaccharides isolated from cornstarch hydrolysates are known as cyclodextrins. One of these compounds, cyclohexaamylose, is a large doughnut-shaped molecule with a central cavity that literally can engulf a small, relatively nonpolar organic molecule and hold it in water solution, similar to a micelle (Section 18-2F). As with micelles, unusual reactivity is exhibited by the bound molecules. An example is the change in the ortho-para ratio in electrophilic substitution of methoxybenzene by hypochlorous acid, $\ce{HOCl}$, in the presence and absence of cyclohexaamylose:
Apparently the cyclohexaamylose wraps around the methoxybenzene in such a way as to protect the ortho carbons from attack by $\ce{HOCl}$ but to leave the para carbon exposed. It is this kind of specificity that we need to generate in reactions before we can claim to have synthetic reactions under control.
Other Important Polysaccharides
Many polysaccharides besides starch and cellulose are important components of animal tissues, or play a vital role in biochemical processes. One example is chitin, a celluloselike material that is the structural component of the hard shells of insects and crustaceans. The difference between chitin and cellulose is that instead of being a polymer of glucose, chitin is a polymer of 2-deoxy-2-$\ce{N}$-ethanamidoglucose ($\ce{N}$-acetyl-$\beta$-$D$-glucosamine):
Heparin is a very important and complex polysaccharide derivative that occurs in intestinal walls and has a major use as a blood anticoagulant, especially in connection with artificial kidney therapy. Heparin also has shown great promise in the treatment of patients with extensive burns, by promoting blood circulation to burn-damaged tissue. The structure of heparin can not be defined precisely because its composition depends on the source of supply. The major components of the polysaccharide chain are $D$-glucuronic acid, $L$-iduronic acid, and the same 2-deoxy-2-aminoglucose ($D$-glucosamine) that is a constituent of chitin (although in heparin it occurs as the $\alpha$ anomer):
The general construction of heparin involves the linkage of the anomeric carbons of one of the components with the 4-hydroxyl of another. A key feature of the heparin structure is the presence of sulfate groups that occur as hydrogen sulfate esters (Section 15-5B) and as sulfamido groups, $\ce{-NHSO_3H}$, on the 2-deoxy-2-amino-$D$-glucose units in the chain. Hydrogen sulfate groups also are located on the 2-hydroxyls of the $L$-iduronic acid units of the chain. In addition there are $\ce{N}$-ethanoyl groups attached to some of the 2-deoxy-2-amino-$D$-glucose nitrogens that are not connected to $\ce{-SO_3H}$.
Heparin is clearly an extraordinarily complex substance with many highly polar groups, and its mode of action as an anticoagulant is not clear. At present, because of increases in the use of artificial kidney machines, heparin is in rather short supply.
Among the plant polysaccharides are the pectins, which are used as jelling agents in the making of preserves and jellies from fruit. Also important are the alginates from seaweeds and gums from trees, which are used as stabilizers and emulsifiers in the food, pharmaceutical, cosmetic, and textile industries. The pectins principally are polysaccharides of the methyl ester of $D$-galacturonic acid, whereas the alginates are polysaccharides made up of varying proportions of $D$-mannuronic acid and $L$-guluronic acid. The plant gums are similar materials.
There are other polysaccharides besides cellulose in the cell walls of plants. These are called hemicelluloses, but the name is misleading because they are unrelated to cellulose. Those that are made of pentose sugars (mainly xylose) are most abundant. They accumulate as wastes in the processing of agricultural products, and on treatment with acids that yield a compound of considerable commercial importance, oxa-2,4-cyclopentadiene-2-carbaldehyde (furfural):
$^7$Celluloid, one of the first plastics, is partially nitrated cellulose (known as pyroxylin) plasticized with camphor.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/20%3A_Carbohydrates/20.08%3A_Polysaccharides.txt |
The "antiscorbutic" factor of fresh fruits, which prevents the development of the typical symptoms of scurvy in humans, is a carbohydrate derivative known as vitamin C or ascorbic acid. This substance is not a carboxylic acid, but a lactone, and owes its acidic properties (and ease of oxidation) to the presence of an enediol grouping. It belongs to the \(L\) series by the glyceraldehyde convention:
Most animals are able to synthesize vitamin C in their livers but, in the course of evolution, man has lost this capacity.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
20.10: Formation of Carbohydrates by Photosynthesis
Carbohydrates are formed in green plants by photosynthesis, which is the chemical combination, or "fixation", of carbon dioxide and water by utilization of energy from the absorption of visible light. The overall result is the reduction of carbon dioxide to carbohydrate and the formation of oxygen:
If the carbohydrate formed is cellulose, then the reaction in effect is the reverse of the burning of wood, and obviously requires considerable energy input.
Because of its vital character to life as we know it, photosynthesis has been investigated intensively and the general features of the process are now rather well understood. The principal deficiencies in our knowledge include just how the light absorbed by the plants is converted to chemical energy and the details of how the many complex enzyme-induced reactions involved take place.
The ingredients in green plants that carry on the work of photosynthesis are contained in highly organized, membrane-covered units called chloroplasts. The specific substances that absorb the light are the plant pigments, chlorophyll a and chlorophyll b, whose structures are shown in Figure 20-6. These highly conjugated substances are very efficient light absorbers, and the energy so gained is used in two separate processes, which are represented diagrammatically in Figure 20-7.
One photoprocess reduces nicotinamide adenine dinucleotide phosphate $\left( \ce{NADP}^\oplus \right)$ to $\ce{NADPH}$. These dinucleotides, shown below, differ from $\ce{NAD}^\oplus$ and $\ce{NADH}$ (Section 15-6C) in having a phosphate group at $\ce{C_2}$ of one of the ribose units. The oxidized form, $\ce{NADP}^\oplus$, behaves like $\ce{NAD}^\oplus$ and receives the equivalent of $\ce{H}^\ominus$ at $\ce{C_4}$ of the nicotinamide ring to form $\ce{NADPH}$:
The other important photoreaction is oxidation of water to oxygen by the reaction:
$\ce{H_2O} \rightarrow 2 \ce{H}^\oplus + \frac{1}{2} \ce{O_2} + 2 \ce{e}^\ominus$
The oxygen formed clearly comes from $\ce{H_2O}$ and not from $\ce{CO_2}$, because photosynthesis in the presence of water labeled with $\ce{^{18}O}$ produces oxygen labeled with $\ce{^{18}O}$, whereas carbon dioxide labeled with $\ce{^{18}O}$ does not give oxygen labeled with $\ce{^{18}O}$. Notice that the oxidation of the water produces two electrons, and that the formation of $\ce{NADPH}$ from $\ce{NADP}^\oplus$ requires two electrons. These reactions occur at different locations within the chloroplasts and in the process of transferring electrons from the water oxidation site to the $\ce{NADP}^\oplus$ reduction site, adenosine diphosphate (ADP) is converted to adenosine triphosphate (ATP; see Section 15-5F for discussion between the importance of such phosphorylations). Thus electron transport between the two photoprocesses is coupled to phosphorylation. This process is called photophosphorylation (Figure 20-7).
The end result of the photochemical part of photosynthesis is the formation of $\ce{O_2}$, $\ce{NADPH}$, and ATP. Much of the oxygen is released to the atmosphere, but the $\ce{NADPH}$ and ATP are utilized in a series of dark reactions that achieve the reduction of carbon dioxide to the level of a a carbohydrate (fructose). A balanced equation is
$6 \ce{CO_2} + 12 \ce{NADPH} + 12 \ce{H}^\oplus \rightarrow \ce{C_6H_{12}O_6} + 12 \ce{NADP}^\oplus + 6 \ce{H_2O}$
The cycle of reactions that converts carbon dioxide to carbohydrates is called the Calvin cycle, after M. Calvin, who received the Nobel Prize in chemistry in 1961 for his work on determining the path of carbon in photosynthesis.
Carbon enters the cycle as carbon dioxide. The key reaction by which the $\ce{CO_2}$ is "fixed" involves enzymatic carboxylation of a pentose, $D$-ribulose 1,5-phosphate.$^8$
A subsequent hydrolytic cleavage of the $\ce{C_2}$-$\ce{C_3}$ bond of the carboxylation product (this amounts to a reverse Claisen condensation; Section 18-8B) yields two molecules of $D$-3-phosphoglycerate.$^9$
In subsequent steps, ATP is utilized to phosphorylate the carboxyl group of 3-phosphoglycerate to create 1,3-diphosphoglycerate (a mixed anhydride of glyceric and phosphoric acids). This substance then is reduced by $\ce{NADPH}$ to glyceraldehyde 3-phosphate:
Two glyceraldehyde 3-phosphates are utilized to build the six-carbon chain of fructose by an aldol condensation $\left( \ce{C_3} + \ce{C_3} \rightarrow \ce{C_6} \right)$, but the donor nucleophile in this reaction is the phosphate ester of dihydroxypropanone, which is an isomer of glyceraldehyde 3-phosphate. Rearrangement of the $\ce{C_3}$ aldose to the $\ce{C_3}$ ketose (of the type described in Section 20-2D) therefore precedes the aldol addition. (For a discussion of the mechanism of the enzymatic aldol reaction, see Section 17-3F.) The fructose 1,6-diphosphate formed is then hydrolyzed to fructose 6-phosphate:
From what we have described thus far, only one atom of carbon has been added from the atmosphere, and although we have reached fructose, five previously reduced carbons were consumed in the process. Thus the plant has to get back a five-carbon sugar from a six-carbon sugar to perpetuate the cycle. Rather than split off one carbon and use that as a building block to construct other sugars, an amazing series of transformations is carried on that can be summarized by the following equations:
These reactions have several features in common. They all involve phosphate esters of aldoses or ketoses, and they resemble aldol or reverse-aldol condensations. Their mechanisms will no be considered here, but are discussed more explicitly in Sections 20-10A, 20-10B, and 25-10. Their summation is $\ce{C_6} + 3 \ce{C_3} \rightarrow 3 \ce{C_5}$, which means that fructose 6-phosphate as the $\ce{C_6}$ component reacts with a total of three $\ce{C_3}$ units (two glyceraldehyde 3-phosphates and one dihydroxypropanone phosphate) to give, ultimately, three ribulose 5-phosphates. Although the sequence may seem complex, it avoids building up pentose or hexose chains one carbon at a time from one-carbon intermediates.
The Calvin cycle is completed by the phosphorylation of $D$-ribulose 5-phosphate with ATP. The resulting $D$-ribulose 1,5-diphosphate then is used to start the cycle again by combining with carbon dioxide. There is one sixth more fructose is used to build other carbohydrates, notably glucose, starch, and cellulose.
$^8$All of the reactions we will be discussing are mediated by enzymes, and we will omit henceforth explicit mention of this fact. But it should not be forgotten that these are all enzyme-induced processes, for which we have few, if any, laboratory reagents to duplicate on the particular compounds involved.
$^9$We will henceforth, in equations, designate the various acids we encounter as the phosphate and the carboxylate anions, although this is hardly reasonable at the pH values normal in living cells. Glyceric and phosphoric acids are only partially ionized at pH 7-8. However, it would be equally unrealistic to represent the acids as being wholly undissociated.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/20%3A_Carbohydrates/20.09%3A_Vitamin_C.txt |
We will consider here the reverse process of photosynthesis, namely how carbohydrates, especially glucose, are converted to energy by being broken down into carbon dioxide and water.
A general summary of the several stages involved is shown in Figure 20-8. Initially, the storage fuels or foodstuffs (fats, carbohydrates, and proteins) are hydrolyzed into smaller components (fatty acids and glycerol, glucose and other simple sugars, and amino acids). In the next stage, these simple fuels are degraded further to two-carbon fragments that are delivered as the $\ce{CH_3C=O}$ group (ethanoyl, or acetyl) in the form of the thioester of coenzyme A, $\ce{CH_3COS}$CoA. The structure of this compound and the manner in which fatty acids are degraded has been considered in Section 18-8F, and amino acid metabolism is discussed briefly in Section 25-5C. This section is concerned mainly with the pathway by which glucose is metabolized by the process known as glycolysis.
In the conversion of glucose to $\ce{CH_3COS}$CoA, two carbons are oxidized to carbon dioxide with consumption of two oxygen molecules:
$\ce{C_6H_{12}O_6} + 2 \textbf{CoA} \ce{SH} + \left[ 2 \ce{O_2} \right] \rightarrow 2 \ce{CH_3COS} \textbf{CoA} + 4 \ce{H_2O} + 2 \ce{CO_2} \tag{20-5}$
For further oxidation to occur, the $\ce{CH_3COS}$CoA must enter the next stage of metabolism, whereby the $\ce{CH_3C=O}$ group is converted to $\ce{CO_2}$ and $\ce{H_2O}$. This stage is known variously as the citric acid cycle, the tricarboxylic acid cycle, or the Krebs cycle, in honor of H. A. Krebs (Nobel Prize, 1953), who first recognized its cyclic nature in 1937. We can write an equation for the process as if it involved oxygen:
$2 \ce{CH_3COS} \textbf{CoA} + \left[ 4 \ce{O_2} \right] \rightarrow 4 \ce{CO_2} + 2 \ce{H_2O} + 2 \textbf{CoA} \ce{SH} \tag{20-6}$
Notice that combination of the reactions of Equations 20-5 and 20-6, glycolysis plus the citric acid cycle, oxidizes glucose completely to $\ce{CO_2}$ and $\ce{H_2O}$:
$\ce{C_6H_{12}O_6} + \left[ 6 \ce{O_2} \right] \rightarrow 6 \ce{CO_2} + 6 \ce{H_2O} \tag{20-7}$
But, as you will see, none of the steps uses molecular oxygen directly. Hence there must be a stage in metabolism whereby molecular oxygen is linked to production of oxidizing agents that are consumed in glycolysis and in the citric acid cycle.
The coupling of oxygen into the metabolism of carbohydrates is an extremely complex process involving transport of the oxygen to the cells by an oxygen carrier such as hemoglobin, myoglobin, or hemocyanin. This is followed by a series of reactions, among which $\ce{NADH}$ is converted to $\ce{NAD}^\oplus$ with associated formation of three moles of ATP from three moles of ADP and inorganic phosphate. Another electron-carrier is flavin adenine dinucleotide ($\ce{FAD}$; Section 15-6C), which is reduced to $\ce{FADH_2}$ with an associated production of two moles of ATP from two moles of ADP. These processes are known as oxidative phosphorylation and can be expressed by the equations:
Oxidative phosphorylation resembles photophosphorylation, discussed in Section 20-9, in that electron transport in photosynthesis also is coupled with ATP formation.
By suitably juggling Equations 20-7 through 20-9, we find that the metabolic oxidation of one mole of glucose is achieved by ten moles of $\ce{NAD}^\oplus$ and two moles of $\ce{FAD}$:
The overall result is production of 36 moles of ATP from ADP and phosphate per mole of glucose oxidized to $\ce{CO_2}$ and $\ce{H_2O}$. Of these, 34 ATPs are produced according to Equation 20-10 and, as we shall see, two more come from glycolysis.
Glycolysis
Glycolysis is the sequence of steps that converts glucose into two $\ce{C_3}$ fragments with the production of ATP. The $\ce{C_3}$ product of interest here is 2-oxopropanoate (pyruvate):
There are features in this conversion that closely resemble the dark reactions of photosynthesis, which build a $\ce{C_6}$ chain (fructose) from $\ce{C_3}$ chains (Section 20-9). For example, the reactants are either phosphate esters or mixed anhydrides, and the phosphorylating agent is ATP:
$\ce{ROH} + \text{ATP} \rightarrow \ce{RO-PO_3^2-} + \text{ADP} + \ce{H}^\oplus$
Furthermore, rearrangements occur that interconvert an aldose and ketose,
and the cleavage of a $\ce{C_6}$ chain into two $\ce{C_3}$ chains is achieved by a reverse aldol condensation:
Also, oxidation of an aldehyde to an acid is accomplished with $\ce{NAD}^\oplus$. There is a related reaction in photosynthesis (Section 20-9) that accomplishes the reduction of an acid to an aldehyde and is specific for $\ce{NADPH}$, not $\ce{NADH}$:
First, glucose is phosphorylated to glucose 6-phosphate with ATP. Then an aldose $\rightleftharpoons$ ketose rearrangement converts glucose 6-phosphate into fructose 6-phosphate. A second phosphorylation with ATP gives fructose 1,6-diphosphate:
At this stage the enzyme aldolase catalyzes the aldol cleavage of fructose 1,6-diphosphate. One product is glyceraldehyde 3-phosphate and the other is 1,3-dihydroxypropanone phosphate. Another ketose $\rightleftharpoons$ aldose equilibrium converts the propanone into the glyceraldehyde derivative:
The next step oxidizes glyceraldehyde 3-phosphate with $\ce{NAD}^\oplus$ in the presence of phosphate with the formation of 1,3-diphosphoglycerate:
The mixed anhydride of phosphoric acid and glyceric acid then is used to convert ADP to ATP and form 3-phosphoglycerate. Thereafter the sequence differs from that in photosynthesis. The next few steps accomplish the formation of pyruvate by transfer of the phosphoryl group from $\ce{C_3}$ to $\ce{C_2}$ followed by dehydration to phosphoenolpyruvate. Phosphoenolpyruvate is an effective phosphorylating agent that converts ADP to ATP and forms pyruvate:
The net reaction at this point produces more ATP than is consumed in the phosphorylation of glucose and fructose.
What happens thereafter depends on the organism. With yeast and certain other microorganisms, pyruvate is decarboxylated and reduced to ethanol. The end result of glycolysis in this instance is fermentation. In higher organisms, pyruvate can be stored temporarily as a reduction product (lactate) or it can be oxidized further to give $\ce{CH_3COS}$CoA and $\ce{CO_2}$. The $\ce{CH_3COS}$CoA then enters the citric acid cycle to be oxidized to $\ce{CO_2}$ and $\ce{H_2O}$, as discussed in the next section:
The Citric Acid (Krebs) Cycle
Glycolysis to the pyruvate or lactate stage liberates heat, which can help keep the organism warm and produce ATP from ADP for future conversion into energy. However, glycolysis does not directly involve oxygen and does not liberate $\ce{CO_2}$, as we might expect from the overall process of the metabolic conversion of glucose to carbon dioxide and water (Equation 20-10). The liberation of $\ce{CO_2}$ occurs subsequent to pyruvate formation in a process called variously, the citric acid cycle, the Krebs cycle, or the tricarboxylic acid (TCA) cycle.
The initial step, which is not really part of the cycle, is conversion of pyruvate to ethanoyl CoA (acetyl CoA):
To achieve the oxidation of acetyl CoA on a continuing basis, intermediates consumed in certain steps must be regenerated in others. Thus we have a situation similar to that in the Calvin cycle (Section 20-9), whereby the first stage of the cycle achieved the desired reaction ($\ce{CO_2}$ formation) and the second stage is designed to regenerate intermediates necessary to perpetuate the cycle.
The entry point is the reaction between acetyl CoA and a four-carbon unit, 2-oxobutanedioic acid. An aldol-type addition of the $\ce{CH_3CO}$ group to this $\ce{C_4}$ keto acid extends the chain to a branched $\ce{C_6}$ acid (as citric acid):
Dehydration-rehydration of citrate converts it to isocitrate:
From here, oxidation of the hydroxyl function with $\ce{NAD}^\oplus$ gives a keto acid, which loses $\ce{CO_2}$ readily (Section 18-4) and affords 2-oxopentanedioate:
We now have a $\ce{C_5}$ keto acid that can be oxidized in the same way as the $\ce{C_3}$ keto acid, pyruvic acid, to give a butanedioyl CoA:
Two molecules of $\ce{CO_2}$ now have been produced and the remaining part of the citric acid cycle is concerned with regeneration of the CoA for forming acetyl CoA from 2-oxopropanoate, and also with regenerating the 2-oxobutanedioate, which is the precursor of citrate. The steps involved are
The hydrolysis of the acyl CoA in the first step is used for energy storage by conversion of guanosine diphosphate (GDP) to guanosine triphosphate (GTP):
The hydration of the trans-butenedioate (Section 10-3G) and the final oxidation reaction (Section 15-6C) have been discussed previously.
Alternative Routes in Carbohydrate Metabolism
There is an alternative route, called the pentose phosphate pathway, by which glucose enters the glycolytic sequence to pyruvate. This route achieves the oxidative decarboxylation of glucose to give ribose, as the 5-phosphate ester. The essential steps are
The net result is that three pentoses are converted into two molecules of fructose and one of glyceraldehyde $\left( 3 \ce{C_5} \rightarrow 2 \ce{C_6} + \ce{C_3} \right)$.
The relationship of the pentose-phosphate pathway to glycolysis is shown in Figure 20-11. The steps involved in the pentose shunt are readily reversible, but there are several steps in glycolysis that are not. These are the phosphorylation steps (Figure 20-9). Yet, there has to be a return route from pyruvate to glucose. This route is called gluconeogenesis and, in animals, takes place in the liver. We shall not discuss the steps in gluconeogenesis except to indicate again that they are not all the reverse of glycolysis. For comparison, the steps that differ are indicated in Figure 20-9 by dashed lines.
Why is lactate formed from pyruvate in the metabolism of glucose? Pyruvate $+ \: \ce{NADH} + \ce{H}^\oplus \rightarrow$ lactate $+ \: \ce{NAD}^\oplus$ is a dead-end path, but it does furnish the $\ce{NAD}^\oplus$ needed for glycolysis in active muscle. This route for forming $\ce{NAD}^\oplus$ is important, because in circumstances of physical exertion, the rate of production of $\ce{NAD}^\oplus$ from oxidative phosphorylation may be slower than the demand for $\ce{NAD}^\oplus$, in which case a temporary supply is available from the pyruvate $\rightarrow$ lactate reduction. The lactate so formed builds up in muscle tissue under conditions of physical exertion and is apt to cause muscles to "cramp". The excess lactate so formed ultimately is removed by being converted back to pyruvate by oxidation with $\ce{NAD}^\oplus$.
The beauty of the metabolic cycle through pyruvate, shown in summary in Figure 20-11, is the way it can be tapped at various points according to whether the organism requires ATP (from glycolysis), $\ce{NADH}$ (from pentose shunt), or $\ce{NAD}^\oplus$ (from the lactate siding).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/20%3A_Carbohydrates/20.11%3A_The_Generation_of_Energy_from_Carbohydrate_Metabolism.txt |
Molecular orbital theory is a method for determining molecular structure in which electrons are not assigned to individual bonds between atoms, but are treated as moving under the influence of the nuclei in the whole molecule. The spatial and energetic properties of electrons within atoms are fixed by quantum mechanics to form orbitals that contain these electrons. While atomic orbitals contain electrons ascribed to a single atom, molecular orbitals, which surround a number of atoms in a molecule, contain valence electrons between atoms. There are two popular approaches to the formulation of the structures and properties of organic compounds based on quantum mechanics - resonance and molecular-orbital methods. In the past, there has been great controversy as to which of these methods actually is more useful for qualitative purposes and, indeed, the adherents to one or the other could hardly even countenance suggestions of imperfections in their choice. Actually, neither is unequivocally better and one should know and use both - they are in fact more complementary than competitive.
• 21.1: Prelude to Resonance and Molecular Orbital Methods
The structural theory of organic chemistry originated and developed from the concepts of valence and the tetrahedral carbon atom. It received powerful impetus from the electronic theory of bonding, as described in Chapter 6. We now express the structures of many organic compounds by simple bond diagrams which, when translated into three-dimensional models, are compatible with most observed molecular properties.
• 21.2: Characteristics of Simple Covalent Bonds
The simplest kind of bond is that between univalent atoms in diatomic molecules, such as H2 and F2, and so on. In the gas phase the molecules are in rapid motion, colliding with one another and the walls of the container. The atoms vibrate with respect to one another, and the molecules have rotational energy as well. Despite this activity, we can assign an average equilibrium bond distance and an average bond energy for normal, unexcited molecules.
• 21.3: Comparison of the Resonance and Molecular-Orbital Methods
In this section, we will sketch the similarities and differences in the resonance (or valence-bond, VB) and molecular-orbital (MO) approaches for electron-pair bonds. Both methods normally start with atomic orbitals, but where the methods differ is in how these orbitals are used.
• 21.4: The Benzene Problem
Our task here is to see what new insight the VB and MO treatments can give us about benzene, but first we will indicate those properties of benzene that are difficult to explain on the basis of simple structure theory.
• 21.5: Application of the MO Method to 1,3-Butadiene
To treat the ππ -electron system of 1,3-butadiene by simple MO theory, we combine the four pp carbon orbitals of an atomic-orbital model. We can estimate a stabilization energy for butadiene from heats of hydrogenation, and it is useful to compare the values obtained with the calculated delocalization energy.
• 21.6: Application of MO Theory to Other Systems
Many important molecules have alternating single and double bonds (are conjugated), but have atoms that are more (or less) electron-attracting than carbon. Analysis of the electronic configuration resulting from the MO calculations accords generally with the VB hybrids. We also consider how the MO approach can be used to understand these differences in excitation energy.
• 21.7: Which Is Better- MO or VB?
The calculated energy of the electron-pair bond of the hydrogen molecule as a function of H−H intermolecular distance r by the ab initio (exact), MO, and VB procedures show that neither the MO nor the VB calculations come close to the ab initio calculation in reproducing the experimental dissociation energy or the variation of the energy with the intermolecular distance. The VB method gives a little better energy value at the minimum and the MO method gives poor results at larger values of r .
• 21.8: More on Stabilization Energies
Benzene is 36 - 38 kcal more stable than the hypothetical molecule 1,3,5-cyclohexatriene on the basis of the differences between experimental heats of combustion, or hydrogenation, and heats calculated from bond energies. We call this energy difference the stabilization energy (SE) of benzene. We have associated most of this energy difference with ππ -electron delocalization, which is the delocalization energy (DE).
• 21.9: Bond Lengths and Double-Bond Character
Bond lengths frequently are cited as evidence for, or against, electron delocalization, although some caution should be exercised in this respect. For instance, if the hybrid structure of benzene is considered to be represented by the two possible Kekule structures, then each carbon-carbon bond should be halfway between a single bond and a double bond.
• 21.10: Hückel's 4n + 2 Rule
Because the bonding molecular orbitals for $\pi$ systems will be just filled with 2, 6, or 10 electrons to give singlet states, a $\left( 4n + 2 \right)$ $\pi$-electron rule was formulated for stable configurations and a $4n$ $\pi$-electron rule for unstable configurations, where $n$ is an integer. Thus 2, 6, 10, 14, ... $\pi$ electrons will be favorable and 4, 8, 12 ... $\pi$ electrons will be unfavorable. The rule is Huckel's $4n + 2$ rule.
• 21.11: Pericyclic Reactions
There are numerous reactions in organic chemistry that proceed through cyclic transition states. They may be classified generally as pericyclic reactions. An important and familiar example is the Diels-Alder reaction, in which a conjugated diene cycloadds to an alkene or alkyne.
• 21.12: Evidence Bearing on the Mechanism of [2 + 2] Cycloadditions
We have not given you much evidence to decide why it is that some thermal [2 + 2] cycloadditions occur but not others. What is special about fluoroalkenes, allenes, and ketenes in these reactions? One possibility is that Mobius rather than the Huckel transition states are involved, but the Mobius transition states are expected to suffer from steric hindrance. It is also possible that [2 + 2] cycloadditions, unlike the Diels-Alder additions, proceed by stepwise mechanisms.
• 21.E: Resonance and Molecular Orbital Methods (Exercises)
These are the homework exercises to accompany Chapter 21 of the Textmap for Basic Principles of Organic Chemistry (Roberts and Caserio).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
21: Resonance and Molecular Orbital Methods
The structural theory of organic chemistry originated and developed from the concepts of valence and the tetrahedral carbon atom. It received powerful impetus from the electronic theory of bonding, as described in Chapter 6. We now express the structures of many organic compounds by simple bond diagrams which, when translated into three-dimensional models, are compatible with most observed molecular properties. Nonetheless, there are many situations for which ordinary structure theory is inadequate. An example is benzene (Section 1-1G), which does not behave as would be expected if it were a cyclic polyene related to alkatrienes, such as \(\ce{CH_2=CH-CH=CH-CH=CH_2}\).
There are many other substances that do not behave as predicted from the properties of simpler compounds. Some substances are more stable, some more reactive, some more acidic, some more basic, and so on, than we would have anticipated. In this chapter we shall look at the theories that explain some of these apparent anomalies. These theories will be based on quantum-mechanical arguments (Section 1-5).
There are two popular approaches to the formulation of the structures and properties of organic compounds based on quantum mechanics - resonance and molecular-orbital methods. In the past, there has been great controversy as to which of these methods actually is more useful for qualitative purposes and, indeed, the adherents to one or the other could hardly even countenance suggestions of imperfections in their choice. Actually, neither is unequivocally better and one should know and use both - they are in fact more complementary than competitive.
We have used the concepts of the resonance methods many times in previous chapters to explain the chemical behavior of compounds and to describe the structures of compounds that cannot be represented satisfactorily by a single valence-bond structure (e.g., benzene, Section 6-5). We shall assume, therefore, that you are familiar with the qualitative ideas of resonance theory, and that you are aware that the so-called resonance and valence-bond methods are in fact synonymous. The further treatment given here emphasizes more directly the quantum-mechanical nature of valence-bond theory. The basis of molecular-orbital theory also is described and compared with valence-bond theory. First, however, we shall discuss general characteristics of simple covalent bonds that we would expect either theory to explain.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/21%3A_Resonance_and_Molecular_Orbital_Methods/21.01%3A_Prelude_to_Resonance_and_Molecular_Orbital_Methods.txt |
The simplest kind of bond is that between univalent atoms in diatomic molecules, such as $\ce{H_2}$, $\ce{F_2}$, and so on. In the gas phase the molecules are in rapid motion, colliding with one another and the walls of the container. The atoms vibrate with respect to one another, and the molecules have rotational energy as well. Despite this activity, we can assign an average equilibrium bond distance $\left( r_e \right)$ and an average bond energy $\left( D_e \right)$ for normal, unexcited molecules. From ab initio calculations (Section 6-6), we learn that the energy of an $\ce{H_2}$ molecule is a function of $r$, the distance between the hydrogens, as shown in Figure 21-1.
When the distance is reduced from $r_e$, the energy increases very rapidly because of internuclear repulsion. As the separation between the atoms increases, the energy of the system increases more slowly and finally approaches that of the entirely free atoms.
The distance $r_e$, which corresponds to the bond length at minimum energy, increases with atomic number downward in a column of the periodic table as the atoms get larger. It decreases across a horizontal row of the periodic table as the electronegativity of the atoms increases and the atomic radius becomes smaller. Other things being equal, the stronger the bond is, the shorter $r_e$ will be, because a strong bond overcomes the repulsive forces between the nuclei and thus permits them to move closer together. For bonds between two carbon atoms, $r_e$ usually ranges between about $1.20$ Å and $1.55$ Å and, if Figure 21-1 (or anything similar) applies, we should not expect significant $\ce{C-C}$ bonding at internuclear distances greater than $2$ Å.
It is important to recognize that bonding occurs only if the electrons are paired (i.e., have opposite spins). The upper dashed curve of Figure 21-1 shows how the energy changes as two hydrogen atoms will parallel spins approach one another. That there is no net bonding can be understood by the Pauli principle (Section 6-1), which tells us that two electrons cannot be in the same orbital if they are unpaired.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
21.03: Comparison of the Resonance and Molecular-Orbital Methods
The Electron-Pair Bond
In this section, we will sketch the similarities and differences in the resonance (or valence-bond, VB) and molecular-orbital (MO) approaches for electron-pair bonds. Both methods normally start with atomic orbitals $1s$, $2s$, $2p$, and so on, of the types discussed in Section 6-1. Where the methods differ is in how these orbitals are used. For a bond between two atoms, the MO procedure combines (or mixes) two atomic orbitals, one from each atom, with proper account of orbital phase (Section 6-2) to obtain two molecular orbitals, one of low energy and one of higher energy. The atomic orbitals can be pure or hybrid orbitals (Sections 6-1 and 6-4). In Figure 21-2, we show the results of combining the $1s$ orbitals of hydrogen. The calculation for the most stable state proceeds by determining the energy of the system when two paired electrons are in the low-energy molecular orbital. The binding energy is the difference between the energy so calculated and the energies of the separated atoms. Because the molecular orbital extends over both atoms, the bonding electrons must be associated with both atoms.
node. (Do not confuse $+$ and $-$ amplitudes with $+$ and $-$ electronic charges.)
Remember, the MO method first combines the atomic orbitals to give molecular orbitals, then populates the molecular orbitals with electrons (no more than two paired electrons per orbital). This part of the procedure is similar to the way electrons are allocated to atomic orbitals (Section 6-1).
The VB treatment starts with the same atomic orbitals but assigns one electron to each orbital. For an electron-pair bond between two hydrogen atoms, the VB treatment in its simplest form considers two electronic configurations. One of these has electron 1 in the orbital of hydrogen 1 and electron 2 in the orbital of hydrogen 2, $\left( 1 \right)$. The other configuration, $2$, has electron 2 in the orbital of hydrogen 1 and electron 1 in the orbital of hydrogen 2:
The calculation then proceeds to predict a low-energy state and a high-energy state. These states can be regarded as hybrids of $1$ and $2$. The low-energy state, which is the one of more interest to us, usually is called a resonance hybrid.
In the VB method, each of the electrons becomes associated with both atoms through the mixing of the two configurations. A very important point here is that the calculation that mixes $1$ and $2$ leads to a six times greater binding energy than calculated for $1$ and $2$ alone. Thus in the VB treatment we combine electronic configurations (here $1$ and $2$, $\leftrightarrow$ symbolizing mixing), whereas in the MO treatment we combine atomic orbitals to get low- and high-energy molecular orbitals.
What is the Glue in These Bonds?
The forces that hold atoms together through chemical bonds are electrostatic, that is, the attraction of positively charged nuclei for negatively charged electrons. But the energy calculated for a single configuration, such as $1$, only accounts for about one sixth of the total binding. In either the VB or the MO method the electrons in an electron pair between two nuclei brought to within bonding distances are equivalent and indistinguishable. That is, we are unable to identify one electron any more than the other with a given atom. The significance of the pairing of the electrons is that is permits each electron to have maximum possible freedom to move through the orbitals of the two-atom system rather than being "localized" on particular atoms. Quantum-mechanical calculations tell us that freedom of motion of the electrons is very important. Thus, using the VB method, we calculate that fully five sixths of the binding of the hydrogen molecule is associated with the "delocalization" of the electrons between the two nuclei.
There are many compounds with structures in which electrons are delocalized over more than two atoms. Such molecules should be more stable than would be expected for molecules with the same geometry but with electron pairs constrained to be associated with just one or two atoms. We will shortly discuss some specific examples, but because most of these examples involve the delocalization of $\pi$ electrons, it is expedient to first discuss ethene as a prototype, using both the MO and VB methods.
The $\pi$ Bond in Ethene
The atomic-orbital picture of ethene (Figure 6-14) formulates the $\pi$ bond as resulting from overlap of two adjacent $p$ atomic orbitals, one from each of two $sp^2$ hybridized carbons. The $p$ orbitals are directed perpendicularly to the plane defined by the hybrid orbitals of the $\sigma$ bonds, and to a first approximation, we assume that exchange of the $\pi$ and $\sigma$ electrons between their respective orbitals does not affect the energy of the molecule. If this assumption is valid, $\pi$ bonding can be treated independently of $\sigma$ bonding. Although undoubtedly oversimplified, the VB and MO methods have been remarkably successful using this assumption. In our subsequent discussions, we shall treat the $\pi$ electrons separately from localized $\sigma$ electrons.
The $\pi$ bond of the ethene molecule can be formulated very much like the bond in the hydrogen molecule (Section 21-2A), with the difference that the bonding is achieved by the overlap of two $2p$ atomic orbitals of carbon rather than two $1s$ atomic orbitals of hydrogen.
In the MO method the mixing of the two $2p$ atomic orbitals gives two molecular orbitals. The details of the mathematics of the mixing process to give an optimum set of molecular orbitals are well beyond the scope of this book,$^1$ but the results are shown in Figure 21-3. The two $\pi$ electrons of ethene are taken as occupying the low-energy bonding orbital, while the high-energy antibonding orbital normally is empty.
How much more stable is the bonding molecular orbital relative to a pair of noninteracting $p$ atomic orbitals? It is difficult to provide a numerical answer in $\text{kcal mol}^{-1}$ that is meaningful, but we can describe the energy in symbolic terms. First, the energy of one electron in the $p$ atomic orbital of an $sp^2$-hybridized carbon, as in $3$, is taken as a standard quantity, $\alpha$, often called the Coulomb energy:
Thus, if there were no $\pi$ bonding in ethene and no repulsion between the electrons, the energy of the two electrons (one in each of the two adjacent $p$ orbitals of the carbons) would be twice the Coulomb energy, or $2 \alpha$. This would be the situation for two carbons such as $3$ that are widely separated.
The MO calculation shows that the bonding molecular orbital of ethene is more stable (of lower energy) than the nonbonding level, $\alpha$, by a quantity, $\beta$, where $\beta$ is a negative energy term (Figure 21-3). Likewise, the antibonding level is destabilized by an amount $-\beta$. For two paired electrons in the bonding molecular orbital, the $\pi$-electron energy of ethene is calculated to be $2 \left( \alpha + \beta \right) = 2 \alpha + 2 \beta$.
In the valence-bond approach, the $\pi$ bond of ethene is considered to be a hybrid of all reasonable electronic configurations of two indistinguishable paired electrons distributed between two $p$ orbitals. Each of the configurations that can be written, $4a$, $4b$, $4c$, and $4d$, have identical locations of the atomic nuclei in space:
The four valence-bond structures or configurations, $4a$-$d$, are combined mathematically to give four hybrid states, and of these, the lowest-energy one corresponds approximately to the normal state of the molecule. The calculation shows that the structures $4a$ and $4b$, which have one electron in each $p$ orbital, are the major contributors to the "hybrid" of ethene. The valence-bond structures, $4c$ and $4d$, are ionic structures, which correspond to the conventional formulas, $4e$ and $4f$:
These valence-bond structures are not important to the $\pi$ bond of the ground state of ethene, although they are important for carbonyl bonds (Section 16-1B).
$^1$There are many excellent books that cover this subject in great detail; however, the simplest introductory work is J. D. Roberts; Molecular Orbital Calculations, W. A. Benjamin, Inc., Menlo Park, Calif., 1961.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/21%3A_Resonance_and_Molecular_Orbital_Methods/21.02%3A_Characteristics_of_Simple_Covalent_Bonds.txt |
We already have alluded to the difficulties encountered in the interpretation of the structure of benzene in Sections 1-1G and 6-5. Our task here is to see what new insight the VB and MO treatments can give us about benzene, but first we will indicate those properties of benzene that are difficult to explain on the basis of simple structure theory.
Some of the Unusual Properties of Benzene
From x-ray diffraction and spectroscopic measurements, benzene is known to be a planar molecule with six carbons $1.390$ Å apart in a hexagonal ring, $5$. Six hydrogen atoms, one associated with each carbon, are located $1.09$ Å from those carbons. All $\ce{H-C-C}$ and $\ce{C-C-C}$ bond angles are $120^\text{o}$:
The 1,3,5-cyclohexatriene structure, $6$, proposed for benzene in 1866 by Kekule, has alternating single and double bonds around the ring, which would be predicted to have bond lengths of $1.48$ Å and $1.34$ Å, respectively (see Table 2-1):
The knowledge that the bond lengths are equal in the ring in benzene is a point against the Kekule formulation, but a more convincing argument is available from a comparison of the chemistry of benzene with that of 1,3,5-hexatriene, $7$:
Benzene also is more stable by about $36$-$38 \: \text{kcal mol}^{-1}$ than anticipated for the 1,3,5-cyclohexatriene structure. You will recall from earlier discussions that the heat of combustion of one mole of benzene is $38 \: \text{kcal}$ less than calculated for cyclohexatriene (see Section 6-5A). Also, the heat of hydrogenation of benzene is only $49.8 \: \text{kcal mol}^{1}$, which is $36 \: \text{kcal}$ less than expected for 1,3,5-cyclohexatriene; this estimate is based on the assumption that the heat of hydrogenation of 1,3,5-cyclohexatriene (with three double bonds) would be three times that of cyclohexane ($28.5 \: \text{kcal mol}^{-1}$, for one double bond), or $3 \times 28.2 = 85.5 \: \text{kcal mol}^{-1}$.
The extra stability of benzene relative to the hypothetical 1,3,5-cyclohexatriene can be called its stabilization energy. Most (but not all) of this stabilization may be ascribed to resonance or electron delocalization.
The Atomic-Orbital Model of Benzene
In Section 6-5 an atomic model of benzene was discussed in some detail. Each carbon in the ring was considered to form three coplanar $sp^2$-hybrid $\sigma$ bonds at $120^\text{o}$ angles. These carbon-carbon and carbon-hydrogen $\sigma$ bonds use three of the four valence electrons of each carbon. The remaining six carbon electrons are in parallel $p$ orbitals, one on each of the six carbons. Each of the $\pi$ electrons can be regarded as being paired with its immediate neighbors all around the ring, as shown by $8$:
As mentioned in Section 21-2B, delocalization of the electrons over all six centers in benzene should give a more stable electron distribution than any structure in which the electrons are localized in pairs between adjacent carbons (as in the classical 1,3,5-cyclohexatriene structure).
The simple MO and VB treatments of benzene begin with the same atomic-orbital model and each treats benzene as a six-electron $\pi$-bonding problem. The assumption is that the $\sigma$ bonds of benzene should not be very much different from those of ethene and may be regarded as independent of the $\pi$ system.
The Molecular Orbital Method for Benzene
Extension of the ideas of Section 21-2 for the MO treatment of an electron-pair bond between two nuclei to the $\pi$ bonding in benzene is fairly straightforward. What is very important to understand is that there must be more than one molecular orbital for the $\pi$ electrons because there are six $\pi$ electrons, and the Pauli principle does not allow more than two paired electrons to occupy a given orbital. In fact, combination (or mixing) os the six $2p$ orbitals of benzene, shown in $8$, gives six $\pi$ molecular orbitals. Without exception, the number of molecular orbitals obtained by mixing is always the same as the number of atomic orbitals mixed. The details of the mathematics of the mixing process to give an optimum set of molecular orbitals will not be described here,$^1$ but the results are shown in Figure 21-5. Of the six predicted molecular orbitals, three are bonding and three are antibonding. The six $\pi$ electrons are assigned to the three bonding orbitals in pairs and are calculated to have a total $\pi$-electron energy of $6 \alpha + 8 \beta$.
The calculation that leads to the results shown in Figure 21-5 is not very sophisticated. It is based on the assumption that the $\pi$ bonding between each carbon and its immediate neighbors is equal all around the ring and that bonding involving carbons more than $2$ Å apart is unimportant.
What happens if we use the MO method to calculate the $\pi$-electron energy of classical 1,3,5-cyclohexatriene? The procedure is exactly as for benzene, except that we decree that each carbon $p$ orbital forms a $\pi$ bond with only one of its neighboring $p$ orbitals. The results are shown in Figure 21-6. The $\pi$-electron energy turns out to be three times that of ethene, or $6 \alpha + 6 \beta$ compared to $6 \alpha + 8 \beta$ for benzene.
The calculated delocalization energy for benzene is the difference between these quantities, or $\left( 6 \alpha + 8 \beta \right) - \left( 6 \alpha + 6 \beta \right) = 2 \beta$. That is to say, the calculated delocalization energy is the difference between the energy of benzene with full $\pi$ bonding and the energy of 1,3,5-cyclohexatriene with alternating single and double bonds. If the electron delocalization energy $\left( 2 \beta \right)$ is equal to the stabilization energy $\left( 38 \: \text{kcal mol}^{-1} \right)$, then $\beta = 19 \: \text{kcal mol}^{-1}$. Whether this is a valid method for determining $\beta$ has been a matter of dispute for many years. Irrespective of this, the results of the calculations do account for the fact that benzene is more stable than would be expected for 1,3,5-cyclohexatriene.
However, do the results also account for the low reactivity toward the various reagents in Figure 214, such as those that donate $\ce{Br}^\oplus$ to double bonds (see Section 10-3A)? To settle this question, we have calculated the changes in $\pi$-electron energy that occur in each of the following reactions:
This means calculating the $\pi$-electron energies of all four entities and assuming that the differences in the $\sigma$-bond energies cancel between the two reactions. The result of this rather simple calculation is that attack of $\ce{Br}^\oplus$ on benzene is thermodynamically less favorable than on 1,3,5-cyclohexatriene by about $\beta$. If $\beta$ is $19 \: \text{kcal mol}^{-1}$, this is clearly a sizable energy difference, and we can conclude that the simple MO method does indeed account for the fact that benzene is attacked by $\ce{Br}^\oplus$ far less readily than is 1,3,5-cyclohexatriene.
The Valence Bond Method for Benzene
Extension of the basic ideas of the VB treatment described in Section 21-2 to the atomic-orbital model of benzene is straightforward. We can write VB structures that represent pairing schemes of electrons in the atomic orbitals as shown in $9$ through $13$:
Pairing schemes $9$ and $10$ correspond to Kekule's structures, whereas $11$, $12$, and $13$ are called "Dewar structures" because J. Dewar suggested, in 1869, that benzene might have a structure such as $14$:
The electrons are paired in the configurations represented by $11$, $12$, and $13$, but these pairing schemes are not as energetically favorable as $9$ and $10$. The reason is that the two electrons paired according to the dashed lines in $11$, $12$, and $13$ are on nuclei separated by $2.8$ Å, which is too far apart for effective bonding. The dashed lines between the distant carbons in $11$, $12$, and $13$ are significant only in that they define a pairing scheme. Such lines sometimes are said to represent "formal bonds".
We hope that it is clear from what we have said here and previously that the electron-pairing schemes $9$ through $13$ do not separately have physical reality or independent existence; indeed, the energy of the actual molecule is less than any one of the contributing structures. The double-headed arrow between the structures is used to indicate that they represent different electron-pairing schemes for a molecule and not different forms of the molecule in equilibrium with one another. When we use the resonance method in a qualitative way, we consider that the contribution of each of the several structures is to be weighted in some way that accords with the degree of bonding each would have, if it were to represent an actual molecule with the specified geometry. Thus the Kekule-type electron-pairing schemes, $9$ and $10$, are to be taken as contributing equally and predominantly to the hybrid structure of benzene - equally because they are energetically equivalent, and predominantly because they can contribute much more to the overall bonding than $11$, $12$, and $13$.
In using the resonance method, we assume that all the resonance structures contributing to a given resonance hybrid have exactly the same spatial arrangements of the nuclei but different pairing schemes for the electrons. Therefore $11$, $12$, and $13$ are not to be confused with bicyclo[2.2.0]-2,5-hexadiene, $15$, because $15$ is a known (albeit not very stable) molecule with different atom positions and therefore vastly different bond angles and bond lengths from benzene:
The electron-pairing schemes $9$ and $10$ represent the electron pairing that $15$ would have if it were grossly distorted, with each carbon at the corner of a regular hexagon and a formal bond in place of a carbon-carbon single bond. Thus $9$ and $10$ would not contribute in a significant way to the resonance hybrid of $15$.
Clearly, it is inconvenient and tedious to write the structures of the contributing forms to show the structure of a resonance hybrid. A shorthand notation is therefore desirable. Frequently, dashed rather than full lines are used where the bonding electrons are expected to be delocalized over several atoms. For benzene, $16a$ or $16b$ is quite appropriate:
Unfortunately, although these are clear and explicit renderings, they are tedious to draw. As a result, many authors use (as we will most often) a single Kekule structure to represent benzene, with the understanding that all the $\ce{C-C}$ bonds are equivalent. Other authors choose to represent benzene as a hexagon with an inscribed circle:
This is a simple notation for benzene, but is quite uninformative and even can be actively misleading with some aromatic ring systems, and thus should be used with this limitation in mind.
In calculations of the resonance energy of benzene, the five electronic configurations (valence-bond structures $9$ through $13$) are combined mathematically to give five hybrid states, and of these the lowest-energy state is assumed to correspond to the normal state of the molecule. Thus benzene is considered by this approach to be a resonance hybrid of the valence-bond structures $9$ through $13$. In this simple treatment, $9$ and $10$ are calculated to contribute about $80\%$ and $11$, $12$, and $13$ about $20\%$ to the hybrid. The actual numerical VB calculations, which are much more difficult to carry through than the corresponding MO calculations, give an energy of $Q + 2.61 J$ for benzene and $Q + 1.50 J$ for classical 1,3,5-cyclohexatriene.$^3$ The resonance or delocalization energy then is $\left( Q + 2.61 J \right) - \left( Q + 1.50 J \right) = 1.11 J$, which makes $J \sim 35 \: \text{kcal mol}^{-1}$ if the resonance energy is taken to be equal to the $38 \: \text{kcal}$ value obtained for the stabilization energy. If one carries through a simple VB calculation of the relative energy change associated with attack of $\ce{Br}^\oplus$ on benzene as compared to 1,3,5-cyclohexatriene, the value obtained is $0.63 J$, which corresponds to $22 \: \text{kcal}$. This is in excellent agreement with the $19 \: \text{kcal}$ value obtained by the MO method (Section 21-3C).
$^1$There are many excellent books that cover this subject in great detail; however, the simplest introductory work is J. D. Roberts, Molecular Orbital Calculations, W. A. Benjamin, Inc., Menlo Park, Calif., 1961.
$^2$ Note that in $9$ and also in $10$, we show a particular way of pairing the electrons. However, just as $1 \leftrightarrow 2$, and $4a \leftrightarrow 4b$, we also must consider other sets that represent exchanges of electrons across the dashed lines of $9$ and also of $10$.
$^3$$Q$ and $J$ are negative VB energy parameters that correspond roughly to the MO parameters $\alpha$ and $\beta$.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/21%3A_Resonance_and_Molecular_Orbital_Methods/21.04%3A_The_Benzene_Problem.txt |
To treat the $\pi$-electron system of 1,3-butadiene by simple MO theory, we combine the four $p$ carbon orbitals of an atomic-orbital model, such as $17$, to obtain four molecular orbitals:
We can estimate a stabilization energy for butadiene from heats of hydrogenation, and it is useful to compare the values obtained with the calculated delocalization energy. Thus the heat of hydrogenation of 1,3-butadiene is $57.1 \: \text{kcal}$, whereas that of ethene is $32.8 \: \text{kcal}$ and of propene $30.1 \: \text{kcal}$. If ethene is used as the model alkene, the stabilization energy of 1,3-butadiene is $\left( 2 \times 32.8 - 57.1 \right) = 8.5 \: \text{kcal}$, whereas with propene as the model, it would be $\left( 2 \times 30.1 - 57.1 \right) = 3.1 \: \text{kcal}$. The bond energies (Table 4-3) in combination with the heat of formation at $25^\text{o}$ $\left( 26.33 \: \text{kcal} \right)$ give a stabilization energy of $5.0 \: \text{kcal}$.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
21.06: Application of MO Theory to Other Systems
Polar Molecules
Many important molecules have alternating single and double bonds (are conjugated), but have atoms that are more (or less) electron-attracting than carbon. An example is propenal (acrolein), $18$:
With such molecules we need to take into account the fact that the $\pi$ electrons will be attracted to oxygen from carbon, because oxygen is more electronegative than carbon. With the VB method we can do this by considering ionic electron-pairing schemes, $18c$ and $18d$, along with the dienelike structures, $18a$ and $18b$. The hybrid, $18e$, is drawn to reflect the expected relative contributions of the various forms, with $18a$ being most important.
Ionic structures such as $19a$ and $19b$ need not be considered for propenal because carbon is much less electron-attracting than oxygen:
Analysis of the electronic configuration resulting from the MO calculations accords generally with the VB hybrid $18e$.
The 2-Propenyl (Allyl) Cation
An especially important type of carbocation is represented by the 2-propenyl electron-pairing schemes, $21a$ and $21b$, which correspond to the hybrid $21c$.
Because $21a$ and $21b$ are equivalent and no other single low-energy structure is possible, a sizable delocalization energy is expected. Evidence for this delocalization energy of $21c$ is available from the comparative ease of reactions involving formation of carbocation intermediates. An example is in $S_\text{N}1$ ionizations of alkenyl and alkyl halides. The ionization $\ce{CH_2=CHCH_2Br} \rightarrow \ce{CH_2=CHCH_2^+} + \ce{Br^+}$ proceeds more readily than $\ce{CH_3CH_2CH_2Br} \rightarrow \ce{CH_3CH_2CH_2^+} + \ce{Br^+}$ (for which no $\pi$-electron delocalization is possible).
MO treatment of the 2-propenyl cation begins with the atomic-orbital model $22$:
Any $\pi$ electrons will be delocalized through the orbitals of $22$, but it is not so easy to be confident that when two electrons are placed into the lowest molecular orbital the resulting electron distribution will be the same as $21c$ with half of the positive charge on $\ce{C_1}$ and half on $\ce{C_3}$.
The complete calculation gives the result shown in Figure 21-9. Here the lowest-energy molecular orbital has a higher proportion of the $p$ orbital of $\ce{C_2}$ mixed in than the $p$ orbitals of $\ce{C_1}$ and $\ce{C_3}$ - in fact, just the right amount to have $\ce{C_2}$ neutral and $\ce{C_1}$ and $\ce{C_3}$ each with $\frac{1}{2}^\oplus$ when this MO is filled with two paired electrons. The delocalization energy calculated for the cation is $\left( 2 \alpha + 2.82 \beta \right) - \left( 2 \alpha + 2 \beta \right) = 0.82 \beta$ or about $16 \: \text{kcal}$ if $\beta$ is taken to be $19 \: \text{kcal}$. Thus in every respect the simple VB and MO methods give the same representation of the 2-propenyl carbocation.
You will notice that the 2-propenyl radical and the 2-propenyl carbanion can be formulated by the same set of $\pi$ molecular orbitals (Figure 21-9) used for the carbocation by putting one or two electrons into the nonbonding MO. The delocalization energies calculated for the radical and anion are the same as for the cation. Thus $\left( 3 \alpha + 2.82 \beta \right) - \left( 3 \alpha + 2 \beta \right) = 0.82 \beta$ for the radical and $\left( 4 \alpha + 2.82 \beta \right) - \left( 4 \alpha + 2 \beta \right) = 0.82 \beta$ for the anion.
Electronic Spectra by the MO Method
Section 9-9B covers qualitative explanations of how the VB method is used to account for the lower-energy (longer-wavelength) radiation required for electron excitation of conjugated polyenes compared to nonconjugated polyenes. Thus 1,3-butadiene has a $\lambda_\text{max}$ for ultraviolet light at $217 \: \text{nm}$, whereas 1,5-hexadiene has a corresponding $\lambda_\text{max}$ at $185 \: \text{nm}$.
We will now consider how the MO approach can be used to understand these differences in excitation energy. The $\pi$-energy levels and electronic configurations for delocalized and localized 1,3-butadiene are shown in Figure 21-10 (also see Section 21-4). Because the double bonds are so far apart, the $\pi$-electron system of 1,5-hexadiene by the simple MO approach is identical with that of localized 1,3-butadiene. The calculated energy change for the lowest-energy $\pi \rightarrow \pi^*$ transition is $\left( \alpha - 0.62 \beta \right) - \left( \alpha + 0.6 \beta \right) = -1.24 \beta$ for 1,3-butadiene and $\left( \alpha - \beta \right) - \left( \alpha + \beta \right) = -2 \beta$ for 1,5-hexadiene. In each case the energy of the electron in the highest occupied $\pi$ orbital (the HOMO orbital) is subtracted from the energy that an electron would have in the lowest unoccupied $\pi^*$ orbital (the LUMO orbital). Other transitions are possible, as of an electron from the lowest occupied orbital of energy $\alpha + 1.62 \beta$ to the highest unoccupied orbital of energy $\alpha - 1.62 \beta$, but these would have far greater energies.
Qualitatively, the $\pi \rightarrow \pi^*$ transition energy is predicted to be substantially less for 1,3-butadiene than for 1,5-hexadiene. However, any attempt at a quantitative correlation is suspect, because the lowest energy $\pi \rightarrow \pi^*$ transition calculated for 1,3-butadiene is $-1.24 \beta$ and, if $\beta$ is $19 \: \text{kcal}$ (see Section 21-3C), $\lambda_\text{max}$ from Equation 9-2 should be $1214 \: \text{nm}$ instead of the observed $217 \: \text{nm}$.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/21%3A_Resonance_and_Molecular_Orbital_Methods/21.05%3A_Application_of_the_MO_Method_to_13-Butadiene.txt |
The calculated energy of the electron-pair bond of the hydrogen molecule as a function of $\ce{H-H}$ intermolecular distance $r$ by the ab initio (exact), MO, and VB procedures is shown in Figure 21-11. The results show that neither the MO nor the VB calculations come close to the ab initio calculation in reproducing the experimental dissociation energy, $D_e$, or the variation of the energy with the intermolecular distance. The VB method gives a little better energy value at the minimum and the MO method gives poor results at larger values of $r$. We can say that, as calculated by the MO method, the molecule does not "dissociate properly".
Within the calculations of the MO method, the molecule does not "dissociate properly".
Why do these calculations yield results so far from the ab initio curve? There are two reasons. First, atomic orbitals are used that are appropriate for isolated atoms, but are hardly expected to be the best orbitals for the electrons when two or more atoms are in close proximity. It is convenient to use atomic orbitals in simple calculations because they are mathematically simple, but more complicated orbitals are known to give better results. Second, neither treatment properly takes into account electron-electron repulsions. For two electrons, a term of the form $\frac{e^2}{r^2_{12}}$ (in which $e$ is the electronic charge and $r_{12}$ is the distance between the electrons) is required to describe the repulsion between electrons. The exact calculations avoid both difficulties but are so complex mathematically as to be devoid of any capability for providing qualitative understanding.
The VB method gives a slightly lower energy than the MO method at the minimum, because in the simple MO method, when we calculate the energy resulting from two electrons going into the lowest molecular orbital, we put no restraints on their being close together. As a result there is a $50\%$ probability for both electrons being simultaneously in either half of the molecular orbital. In contrast, the simple VB method combines configurations $1$ and $2$, each having just one electron per atomic orbital, and no account is taken of the possibility of either atomic orbital containing more than one electron. This is equivalent to neglecting the pairing schemes $\ce{H}^\ominus \ce{H}^\oplus \leftrightarrow \ce{H}^\oplus \ce{H}^\ominus$. Neither the VB nor the MO approximation is the best possible; the simple MO method tends to take too little account of interelectronic repulsion, whereas the VB method tends to take too much account of it. However, as can be seen in Figure 21-11, taking too much account of electron repulsion is the better approximation.
Why does an electron-pair bond calculated by the MO method not dissociate properly? We have seen that half of the time both electrons in the low-energy molecular orbital are in the vicinity of just one of the nuclei. But as the nuclei move far apart, this corresponds to a far greater energy than having only one electron in the vicinity of each nucleus, as the VB method suggests.
There is no unequivocal answer to the question as to which is the better method. Calculations by the VB method are likely to be more reliable than those by the MO method, but in practice are much more difficult to carry out. For many-electron molecules the MO procedure is simpler to visualize because we combine atomic orbitals into molecular orbitals and then populate the lower-energy orbitals with electrons. In the VB method, atomic orbitals are occupied, but the electrons of different atoms are paired to form bonds, a process that requires explicit consideration of many-electron wave functions. To put it another way, it is easier to visualize a system of molecular orbitals containing $N$ electrons than it is to visualize a hybrid wave function of $N$ electrons.
How can the MO and VB methods be improved? The answer depends on what one wants - more accurate calculations or better qualitative understanding. To improve VB calculations we need orbitals that allow the electrons to spread out over more than one atom. The GVB orbitals discussed in Section 6-6 suit this purpose and give an energy curve only slightly above the exact curve of Figure 21-11. In the GVB treatment the orbitals delocalize less as $r$ increases.
When atomic orbitals are derived for each carbon of the $\pi$-electron system of benzene by the GVB method, they are somewhat more spread out than simple carbon $p$ orbitals (Figure 21-12). Use of these orbitals in VB calculations gives excellent results with just the two pairing schemes of benzene, $9$ and $10$.
Improvement of the MO method involves better orbitals, better account of interelectronic repulsion, and introduction of mixing of different electron configurations in the molecular orbitals ("configuration interaction"). Improved MO calculations give much more accurate energies at the minimum of a plot such as Figure 21-11, but the bonds still do not dissociate properly, for the same reason as with the simple MO method.
We cannot say that either the VB or the MO method is more correct; only that one approximation may be more useful than the other in attempting to solve a particular problem. The fact is, the more each is refined, the more they appear to merge into a common procedure; but, unfortunately, in the refinement process the mathematics become so complex that qualitative understanding of what is being done tends to disappear altogether.
We cannot say that either the VB or the MO method is more correct; only that one approximation may be more useful than the other in attempting to solve a particular problem.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/21%3A_Resonance_and_Molecular_Orbital_Methods/21.07%3A_Which_Is_Better-_MO_or_VB.txt |
It was shown in Section 21-3 that benzene is $36$-$38 \: \text{kcal}$ more stable than the hypothetical molecule 1,3,5-cyclohexatriene on the basis of the differences between experimental heats of combustion, or hydrogenation, and heats calculated from bond energies. We call this energy difference the stabilization energy (SE) of benzene. We have associated most of this energy difference with $\pi$-electron delocalization, which is the delocalization energy (DE). The difference between SE and DE will be small only if our bond-energy tables are reliable and steric and strain effects are small.
The problem with bond energies is that we use bond energies that neglect changes in bond strength caused by environment. Primary, secondary, tertiary, alkenic, and alkynic $\ce{C_H}$ bonds are assumed to have equal energies; $\ce{C-C}$ single bonds are assumed to be equal, regardless of whether the other bonds to the carbon atoms in question are single or multiple; and differences in energy between double bonds that are mono-, di-, tri-, or tetra-substituted are neglected, as are changes in bond energies associated with steric strain. Bond energies are strictly applicable to molecules in which the bonds are of the normal lengths. In the case of benzene, which has $\ce{C-C}$ bonds with lengths intermediate between normal single and double bonds, there seems to be no clear agreement as to how to take the bond distances into account in computing the delocalization energy. In spite of these uncertainties the stabilization energies seem to give a good qualitative idea of the importance of electron delocalization in organic molecules.
Tables 21-1 and 21-2 give stabilization energies for several substances that are best represented as hybrid structures.
Table 21-1: Stabilization Energies (or Approximate Delocalization Energies) from Heats of Formation of Some Aromatic Compounds
Table 21-2: Stabilization Energies (SE) from Heats of Formation $\left( \Delta H^0 \right)$ of Some Conjugated Polyenes
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
21.09: Bond Lengths and Double-Bond Character
Bond lengths frequently are cited as evidence for, or against, electron delocalization, although some caution should be exercised in this respect. For instance, if the hybrid structure of benzene is considered to be represented by the two possible Kekule structures, then each carbon-carbon bond should be halfway between a single bond and a double bond. In other words, each should possess $50\%$ double-bond character. We then may expect the carbon-carbon bond lengths for benzene to be the average of single- and double-bond lengths. However, the average of the $\ce{C-C}$ bond in ethane $\left( 1.534 \: \text{Å} \right)$ and in ethene $\left( 1.337 \: \text{Å} \right)$ is $1.436 \: \text{Å}$, which does not agree well with the measured $\ce{C-C}$ bond distance for benzene of $1.397 \: \text{Å}$. The discrepancy lies largely in the assumption inherent in this crude calculation that, in the absence of resonance, all $\ce{C-C}$ single bonds are equal to $1.534 \: \text{Å}$. Clearly, this is not a valid assumption because, as we have seen, bond energies depend upon environment, and because the energy of a bond depends upon its length (see Figure 21-1), bond lengths also must vary with environment. This can be seen from the data in Table 21-3, which gives the carbon-carbon single bond lengths for several compounds. The single bonds shorten as the other bonds to carbon become progressively unsaturated, that is, as the hybridization of carbon changes from $sp^3$ to $sp$. Admittedly, some of this shortening may be ascribed to resonance, but not all.
Table 21-3: Carbon-Carbon Single-Bond Distances (Å)
If we take $1.48 \: \text{Å}$ as a reasonable $\ce{C-C}$ bond distance between two $sp^2$-hybridized carbons and $1.34 \: \text{Å}$ for $\ce{C=C}$ bonds (see Table 2-1), the average is $1.41 \: \text{Å}$, which is not much different from the $1.40 \: \text{Å}$ for the carbon-carbon bonds in benzene.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/21%3A_Resonance_and_Molecular_Orbital_Methods/21.08%3A_More_on_Stabilization_Energies.txt |
Cyclobutadiene and Cyclooctatetraene
More than 100 years ago, Kekule recognized the possible existence of other conjugated cyclic polyalkenes, which at least superficially would be expected to have properties like benzene. The most interesting of these are cyclobutadiene, $23$, and cyclooctatetraene, $24$:
For each we can write two equivalent planar VB structures, and the qualitative VB method would suggest that both compounds, like benzene, have substantial electron-delocalization energies. However, the planar structures would have abnormal $\ce{C-C=C}$ angles, and consequently at least some degree of destabilization associated with these bond angles (Section 12-7). Nonetheless, estimation of the strain energies show that while they are substantial, they are not prohibitive. Should then these molecules be stabilized by resonance in the same sense as benzene is postulated to be?
In 1911 a German chemist, R. Willstatter (Nobel Prize 1915), reported an extraordinary thirteen-step synthesis of cyclooctatetraene from a rare alkaloid called pseudopelletierine isolated from the bark of pomegranate trees. The product was reported to be a light-yellow, highly unsaturated compound that absorbed four moles of hydrogen to form cyclooctane. Numerous tries to repeat the Willstatter synthesis were unsuccessful, and in the 1930s the prevailing opinion was that the product had been misidentified. However, during the Second World War, the German chemist W. Reppe found that cyclooctatetraene can be made in reasonable yields by the tetramerization of ethyne under the influence of a nickel cyanide catalyst:
The properties of the product substantiated Willstatter's reports and it became clear that cyclooctatetraene is not like benzene.
Subsequent studies of the geometry of the molecule revealed further that it is nonplanar, with alternating single and double bonds, $25a$:
This geometry precludes the possibility of two equivalent VB structures, as for benzene, because, as you will see if you try to make a ball-and-stick model, $25b$ is highly strained and not energetically equivalent to $25a$ at all. Thus we can conclude that the delocalization energy of cyclooctatetraene is not large enough to overcome the angle strain that would develop if the molecule were to become planar and allow the $\pi$ electrons to form equivalent $\pi$ bonds between all of the pairs of adjacent carbons.
Cyclobutadiene, $23$, eluded Kekule, Willstatter, and a host of other investigators for almost 100 years. As more work was done, it became increasingly clear that the molecule, when formed in reactions, was immediately converted to something else. Finally, the will-o'-the-wisp was captured in an essentially rigid matrix of argon at $8 \: \text{K}$. It was characterized by its spectral properties (not by combustion analysis). On warming to even $35 \: \text{K}$, it dimerizes to yield $26$:
One possibility for the lack of stability$^4$ of cyclobutadiene is that the angle strain associated with having four $sp^2$ carbons in a four-membered ring is much greater than estimated. However, the stable existence of many compounds with four such $sp^2$ carbons, for example $27$ and $28$, make this argument weak, if not invalid:
Why, then, is cyclobutadiene so unstable and reactive? On this point, and also with respect to the nonaromatic character of cyclooctatetraene, the simple qualitative VB method that we have outlined is no help whatsoever. There is no way simply to look at the electron-pairing schemes $23$ and $24$ and see any difference between them and the corresponding schemes for benzene.$^5$
It is in this area that qualitative MO procedures have great success because there are general characteristics of the $\pi$ molecular orbitals of monocyclic, conjugated polyene systems that predict differences in the properties of cyclobutadiene, benzene, cyclooctatetraene, and other similar compounds that are not obvious from the simple VB method.
As a rule, for $N$ parallel atomic $p$ orbitals overlapping in the $\pi$ manner in a monocyclic array, there will be just one lowest molecular orbital, with all the atomic orbitals having the same phase. This will be seen for benzene in Figure 21-5. What is harder to understand without going through the calculations is that the higher-energy molecular orbitals for cyclic conjugated polyenes are predicted to come in successive degenerate$^6$ pairs, as shown in Figure 21-13 for $N = 3$ to $9$.
The qualitative ordering and, indeed, the numerical values of the energies of the $\pi$ molecular orbitals for a cyclic system of $N$ $p$ orbitals can be derived in a very simple way. It is necessary only to inscribe a regular polygon with $N$ sides inside a circle of radius $2 \beta$ with a corner down. For example, for $N = 5$ we get the following:
The molecular orbital energies are in units of $\beta$ at the corners of the polygon. The nonbonding level corresponds to the horizontal dashed line drawn through the center of the circle.
The data of Figure 21-13 provide a rationale for the instability of cyclobutadiene and cyclooctatetraene. For cyclobutadiene, we can calculate that four $\pi$ electrons in the lowest orbitals will lead to a predicted $\pi$-electron energy of $2 \left( \alpha + 2 \beta \right) + 2 \left( \alpha \right) = 4 \alpha + 4 \beta$, which is just the $\pi$-electron energy calculated for two ethene bonds (see Figure 21-3). The delocalization energy of the $\pi$ electrons of cyclobutadiene therefore is predicted to be zero!
Another feature of the $\pi$ system of cyclobutadiene is that the four $\pi$ electrons do not suffice to fill the three lowest orbitals and, if we apply Hund's rule (Section 6-1), the best way to arrange the electrons is as in $29$, with two unpaired electrons, which is known as a triplet state:$^7$
With the MO predictions of zero delocalization energy and an electronic configuration with unpaired electrons, we should not be surprised that cyclobutadiene readily dimerizes to give $26$ even at very low temperatures.
The energies of the molecular orbitals calculated for planar cyclooctatetraene (Figure 21-13) lead to a predicted delocalization energy of $\left( 8 \alpha + 9.64 \beta \right) - \left( 8 \alpha + 8 \beta \right) = 1.64 \beta$ $\left( \sim 31 \: \text{kcal} \right)$, which is smaller than that of benzene, even though there are eight atomic orbitals instead of six through which the electrons are delocalized. Furthermore, the lowest electronic configuration for the planar molecule is, like cyclobutadiene, predicted to be a triplet. Experimental evidence indicates that the positions of the double bonds of cyclooctatetraene shift slowly as the result of formation of the molecule in the unstable planar state. The energy input required to flatten the molecule is about $15 \: \text{kcal mol}^{-1}$:
As Huckel formulated, the $4n + 2$ rule applies only to monocyclic systems. However, as a practical matter it can be used to predict the properties of polycyclic conjugated polyenes, provided the important VB structures involve only the perimeter double bonds, as in the following examples:
Application of the $4n + 2$ rule to other $\pi$ systems, such as $30$ and $31$, is not valid because good VB structures cannot be written that involve changes in the pairing schemes of the perimeter electrons all at once.
Application of Resonance and of the $4n + 2$ Rule to Cyclic Ions
The hydrogens of the $\ce{-CH_2}-$ group of 1,3-cyclopentadiene are acidic. In fact, they are considerably more acidic than the ethyne hydrogens of the 1-alkynes (Section 11-8). This means that 1,3-cyclopentadiene is at least $10^{30}$ times more acidic than the ordinary alkanes. The reason is that loss of one of the $\ce{CH_2}$ protons of cyclopentadiene results in formation of an especially stabilized anion:
The structure of the anion may be described as a hybrid of five energetically equivalent structures, $34a$ through (34e). The unshared electron pair therefore is delocalized over five carbon atoms, and the resulting delocalized anion is much more stable than expected for any one of the equivalent localized structures:
This looks very reasonable, although the simple beauty is seemingly destroyed by the fact that the cyclopentadienyl cation is not very stable, despite the five structures, $35a$ through $35e$, that may be written for it:
Extension of these ideas to the other ring sizes of Figure 21-13 suggests that all of the following ions, which have $\left( 4n + 2 \right)$ $\pi$ electrons, should be unusually stable:
In contrast, the following should be unstable with $4n$ $\pi$ electrons and triplet electronic configurations:
These predictions indeed are borne out by many experiments, some of which we will discuss later.
$^4$It should be recognized that the term "stability" is subject to many interpretations. One criterion of stability would be whether an isolated molecule would fragment spontaneously in interstellar space, such as one would expect for a "molecule" consisting of two neon atoms $1.5 \: \text{Å}$ apart (see Figure 4-6). A different criterion would be whether a molecule could be preserved in the presence of the same or other kinds of molecules at some specified temperature. The classical criterion would be whether the substance could be isolated, put into a bottle and preserved for at least a short time. All of the existing evidence indicates that cyclobutadiene molecules would not spontaneously decompose in interstellar space, but they do react with each other extremely readily, even at low temperatures, and can be preserved only by being held isolated from one another in a rigid matrix of a very inert material, such as solid argon. "Stability" in the sense of "lack of reactivity" has to be carefully defined in terms of experimental conditions. For example, is very unstable in the presence of nucleophiles such as water or methanol, whereas it is quite stable in "super-acid solutions" where no good nucleophiles are present (Section 10-3B).
$^5$A rather simple extension of the VB method by what is called the "orbital-phase continuity principle" does permit the qualitative judgment that cyclobutadiene should be less stable than benzene [see W. A. Goddard III, J. Amer. Chem. Soc. 94, 743 (1972), for applications to many processes for which VB theory generally has been regarded as incapable of giving much insight].
$^6$Degenerate orbitals have the same energy; see Section 6-1.
$^7$The name "triplet state" is used because a system with two unpaired electrons has three different energy states in a magnetic field.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/21%3A_Resonance_and_Molecular_Orbital_Methods/21.10%3A_Huckel%27s_4n__2_Rule.txt |
Why are [4 + 2] and [2 + 2] Cycloadditions Different?
There are numerous reactions in organic chemistry that proceed through cyclic transition states. They may be classified generally as pericyclic reactions. An important and familiar example is the Diels-Alder reaction, in which a conjugated diene cycloadds to an alkene or alkyne:
This reaction has been described previously (Section 13-3A) and is an example of a [4 + 2] cycloaddition. Such reactions occur thermally (by simply heating the reactants) and appear to be entirely concerted. By this we mean that the reactants are converted to products in one step, without involving the formation of reaction intermediates. The principal evidence for the concertedness of [4 + 2] cycloadditions is the fact that they are highly stereospecific and involve suprafacial addition of both components. The configuration of substituents in the diene and the dienophile is retained in the adduct:
In contrast to the [4 + 2] cycloaddition, thermal [2 + 2] cycloadditions seldom are observed, and when they are observed, they are not stereospecific and evidently are stepwise reactions (see Section 21-11):
Why are the [4 + 2] and [2 + 2] cycloadditions different? Simple molecular orbital theory provides an elegant explanation of this difference based on the $4n + 2$ rule described in Section 21-9. To understand this, we need to look in more detail at how the $p$ orbitals of the double bonds interact in concerted addition mechanisms by suprafacial overlap, as in $36$ and $37$:
How Mr. Möbius Beat the $4n + 2$ Rule
There is a way around the $4n + 2$ rule that is not very important for substances analogous to benzene, but is quite important for cycloaddition reactions. Let us see how this works for a cyclic conjugated polyene.
From the molecular-orbital diagrams of Figures 21-5, 21-7, 21-9, and 21-14, you will see that the lowest-energy $\pi$ molecular orbital has no nodes (changes of phase). A model of such an orbital, which usually is called a Hückel orbital, can be constructed by joining the ends of a ribbon or strip of parallel $p$ orbitals, as represented on the left side of Figure 21-15. However, one could join the orbitals by making one twist in the strip, which then would give a lowest-energy orbital with one node, as on the right side of Figure 21-15. A strip with one such twist is called a Möbius strip$^8$ and has the topological property of having only one side.
If we now calculate the orbital energies for the Möbius orbitals, as was down for the normal Hückel $\pi$ orbitals in Figure 21-13, we get the results shown in Figure 21-16. From this, we see that the $4n$ situation now is favored and $4n + 2$ is unfavorable. Whereas the energies of the $\pi$ molecular orbitals in the Hückel arrangement can be obtained by inscribing a polygon in a circle with a corner down (Section 21-9A), in the Möbius arrangement the orbital energies are obtained from the polygon inscribed with a side down.
If you compare the orbital energies of the Hückel and Möbius cyclic $\pi$ systems (Figures 21-13 and 21-16), you will see that the Hückel systems have only one lowest-energy MO, whereas the Möbius systems have two. Hückel systems have an odd number of bonding orbitals (which, when full, accommodate 2, 6, 10, 14, or $4n + 2$ electrons) and the Möbius systems have an even number of bonding orbitals (which, when full, accommodate 4, 8, 12, or $4n$ electrons). The Hückel molecular orbitals have zero or an even number of nodes (see, for example, the benzene MOs, Figure 21-5); the Möbius molecular orbitals are not shown, but they have one or an odd number of nodes.
The relevance of all this may seem tenuous, especially because no example of a simple cyclic polyene with a Möbius $\pi$ system is known. However, the Möbius arrangement is relevant to cycloaddition because we can conceive of alkenes, alkadienes, and so on approaching each other to produce Möbius transition states when $4n$ electrons are involved.
For example, consider two molecules of ethene, which we showed previously would violate the $4n + 2$ rule by undergoing cycloaddition through a transition state represented by $37$. There is an alternative transition state, $38$, in which the four $p$ orbitals come together in the Möbius arrangement (with one node for minimum energy).
To achieve this arrangement the ethene molecules approach each other in roughly perpendicular planes so that the $p$ orbitals overlap suprafacially in one ethene and antarafacially in the other, as shown in $38$:
This pathway is electronically favorable, but the steric interference between the groups attached to the double bond is likely to be severe. Such repulsions can be relieved if there are no groups sticking out sidewise at one end of the double bond, as with the central carbon of 1,2-propadiene, $\ce{CH_2=C=CH_2}$, and ketene, $\ce{CH_2=C=O}$. These substances often undergo [2 + 2] cycloadditions rather readily (Section 13-3D), and it is likely that these are concerted additions occurring by the Möbius route.
A much less strained Möbius [4 + 4] transition state can be formed from two s-cis molecules of 1,3-butadiene. When 1,3-butadiene is heated by itself, a few percent of 1,5-cyclooctadiene is formed, but it is not known for sure whether the mechanism is that shown:
The principal reaction is a Diels-Alder [4 + 2] cycloaddition, with butadiene acting both as a diene and as a dienophile:
Orbital Symmetry. The Woodward-Hoffmann Rules
Much of what we have said about the electronic factors controlling whether a cycloaddition reaction can be concerted or not originally was formulated by the American chemists R. B. Woodward and R. Hoffmann several years ago, in terms of what came to be called the orbital symmetry principles, or the Woodward-Hoffmann rules. Orbital symmetry arguments are too complicated for this book, and we shall, instead, use the $4n + 2$ electron rule for normal Hückel arrangements of $\pi$ systems and the $4n$ electron rule for Möbius arrangements. This is a particularly simple approach among several available to account for the phenomena to which Woodward and Hoffmann drew special attention and explained by what they call "conservation of orbital symmetry".
Electrocyclic and Sigmatropic Rearrangements
The cycloaddition reactions that we have discussed so far in this chapter ([2 + 2], [4 + 2], etc.) have involved ring formation by bringing two unsaturated molecules together. Thus [4 + 2] addition is represented by the Diels-Alder reaction of ethene and 1,3-butadiene:
We can conceive of similar cyclizations involving only single molecules, that is, intramolecular cyclization. Such reactions are called electrocyclic rearrangements. Two examples follow to show cyclization of a diene and a triene:
Cyclization of 1,3,5-hexatriene occurs only when the central double bond has the cis configuration. The reaction is reversible at elevated temperatures because of the gain in entropy on ring opening (see Section 4-4B). The cyclobutene-1,3-butadiene interconversion proceeds much less readily, even in the thermodynamically favorable direction of ring opening. However, substituted dienes and cyclobutenes often react more rapidly.
A related group of reactions involves shifts of substituent groups from one atom to another; for example, with $\ce{H}$, alkyl, or aryl groups as $\ce{R}$:
These reactions are called sigmatropic rearrangements and, in general, they are subject to the $4n + 2$ rule and the Möbius orbital modification of it. Potential sigmatropic rearrangements can be recognized by the fact that the single bond to the migrating group $\left( \ce{R} \right)$ is "conjugated" with the $\pi$ bonds, and the group moves from a saturated $sp^3$ to an $sp^2$ carbon at a different part of the $\pi$ system.
The Stereochemistry of Electrocyclic Rearrangements
A striking feature of thermal electrocyclic reactions that proceed by concerted mechanisms is their high degree of stereospecificity. Thus when cis-3,4-dimethylcyclobutene is heated, it affords only one of the three possible cis-trans isomers of 2,4-hexadiene, namely, cis,trans-2,4-hexadiene:
We can see how this can occur if, as the ring opens, the ends of the diene twist in the same direction ($\curvearrowright \curvearrowright$ or $\curvearrowleft \curvearrowleft$, conrotatory) as indicated in the equation. You will notice that with this particular case, if conrotation occurs to the left, rather than the right, the same final product results:
The conrotatory movement of groups is typical of thermal ring openings of cyclobutenes and other rings involving $4n$ electrons.
When a cyclobutene is so constituted that conrotation cannot occur for steric reasons, then the concerted reaction cannot occur easily. Substances that otherwise might be predicted to be highly unstable often turn out to be relatively stable. An example is bicyclo[2.1.0]-2-pentene, which at first sight might seem incapable of isolation because of the possibility of immediate arrangement to 1,3-cyclopentadiene. This rearrangement does occur, but not so fast as to preclude isolation of the substance:
How can we explain the fact that this substance can be isolated? The explanation is that, if the reaction has to be conrotatory, then the product will not be ordinary 1,3-cyclopentadiene, but cis,trans-1,3-cyclopentadiene - surely a very highly strained substance. (Try to make a ball-and-stick model of it!) This means that the concerted mechanism is not favorable:
It is of great interest and importance that, with systems of $4n + 2$ electrons, the groups move in opposite directions ($\curvearrowleft \curvearrowright$ or $\curvearrowright \curvearrowleft$, disrotatory). For example,
In this case, the disrotation of the groups toward one another would lead to the cis,cis,cis product. Because this product is not formed, it seems likely that rotation of the methyl groups toward each other must be sterically unfavorable:
How can we account for the stereoselectivity of thermal electrocyclic reactions? Our problem is to understand why it is that concerted $4n$ electrocyclic rearrangements are conrotatory, whereas the corresponding $4n + 2$ processes are disrotatory. From what has been said previously, we can expect that the conrotatory processes are related to the Möbius molecular orbitals and the disrotatory processes are related to Hückel molecular orbitals. Let us see why this is so. Consider the electrocyclic interconversion of a 1,3-diene and a cyclobutene. In this case, the Hückel transition state (one having an even number of nodes) is formed by disrotation, but is unfavorable with four (that is, $4n$) electrons:
In contrast, the Möbius transition state (one having an odd number of nodes) is formed by conrotation and is favorable with four $\left( 4n \right)$ electrons:
You will notice that the ring closure of a 1,3-diene through the favorable Möbius transition state may appear to be able to form only an antibonding arrangement of the overlapping $\sigma$ orbitals, which would correspond to a high-energy cyclobutene. In fact, the normal cyclobutene would be formed, because on the way down from the transition state, the phases of the orbitals that will become the $\sigma$ bond change to give the bonding arrangement of the $\sigma$ orbitals expected for the ground state. The reverse occurs in ring opening so that this reaction also can go through the favorable Möbius transition state.
The same reasoning can be extended to electrocyclic reactions of 1,3,5-trienes and 1,3-cyclohexadienes, which involve $4n + 2$ electrons and consequently favor Hückel transition states attained by disrotation.
Summary of Rules for Predicting Thermally Favorable Pericyclic Reactions
The three principal types of pericyclic reactions are cycloaddition, electrocyclic rearrangement, and sigmatropic rearrangement:
The factors that control if and how these cyclization and rearrangement reactions occur in a concerted manner can be understood from the aromaticity or lack of aromaticity achieved in their cyclic transition states. For a concerted pericyclic reaction to be thermally favorable, the transition state must involve $4n + 2$ participating electrons if it is a Hückel orbital system, or $4n$ electrons if it is a Möbius orbital system. A Hückel transition state is one in which the cyclic array of participating orbitals has no nodes (or an even number) and a Möbius transition state has an odd number of nodes.
We summarize here a procedure to predict the feasibility and the stereochemistry of thermally concerted reactions involving cyclic transition states. The 1,2 rearrangement of carbocations will be used to illustrate the approach. This is a very important reaction of carbocations which we have discussed in other chapters. We use it here as an example to illustrate how qualitative MO theory can give insight into how and why reactions occur:
The first step of the procedure is to draw the orbitals as they are expected to be involved in the transition state. There may be several possible arrangements. There are two such arrangements, $41$ and $42$, for the rearrangement of carbocations; the dotted lines show the regions of bond-making and bond-breaking (i.e., orbital overlap):
The second step is to determine whether the transition states are Hückel or Möbius from the number of nodes. This is readily done by assigning signs to the lobes of the orbitals corresponding to their phases and counting the number of nodes that develop in the circle of overlapping orbitals. An odd number denotes a Möbius transition state, whereas an even number, including zero, denotes a Hückel transition state.
There are alternative ways of node-counting for transition states $41$ and $42$. Diagrams $43abc$ and $44abc$ represent molecular orbitals of different energies - those with more nodes having the higher energies (cf. Section 21-3C).$^9$ We show these diagrams with more than one node for the sake of completeness. It is not necessary to draw more than one such diagram to determine whether the transition state is Möbius or Hückel.
Finally, we evaluate the transition states according to the $4n$ or $4n + 2$ rule. In the example here, because only two electrons occupy the molecular orbitals, the Hückel transition state ($43a$) is the favorable one.
A bonus coming from these formulations is that the stereochemistry of the reaction can be predicted when we have predicted which transition state is the favored one. Thus the migrating group in 1,2-carbocation rearrangements should move with retention of configuration by a Hückel transition state - and this has been verified experimentally. The alternative Möbius transition state predicts inversion of the configuration of the migrating group:
You can use the procedures just outlined to determine whether any thermal reaction with a cyclic transition state is likely to be favorable. A good place to start is the Diels-Alder [4 + 2] cycloaddition, which proceeds thermally by a suprafacial (Hückel) transition state. We suggest that you apply the procedure to the Diels-Alder reaction of 1,3-butadiene and ethene, and following that, show the electrocyclic ring opening of a cyclobutene ring to be thermally favorable only by a conrotatory opening of the $\ce{C-C}$ bond.
Photochemical Pericyclic Reactions
Many pericyclic reactions take place photochemically, that is, by irradiation with ultraviolet light. One example is the conversion of norbornadiene to quadricyclene, described in Section 13-3D. This reaction would have an unfavorable [2 + 2] mechanism if it were attempted by simple heating. Furthermore, the thermodynamics favor ring opening rather than ring closure. However, quadricyclene can be isolated, even if it is highly strained, because to reopen the ring thermally involves the reverse of some unfavorable [2 + 2] cycloaddition mechanism.
Photochemical activation can be used to achieve forward or reverse cycloadditions and electrocyclic reactions that are thermodynamically unfavorable or have unfavorable concerted thermal mechanisms. Thus the thermodynamically unstable disrotatory [2 + 2] product can be obtained from 1,3-cyclopentadiene by irradiation with ultraviolet light:
The stereochemical results of electrocyclic and cycloaddition reactions carried out photochemically often are opposite to what is observed for corresponding thermal reactions. However, exceptions are known and the degree of stereospecificity is not always as high as in the thermal reactions. Further examples of photochemical pericyclic reactions are given in Section 28-2D.
$^8$Named after the mathematician A. F. Möbius.
$^9$The assignment of orbital phases must take appropriate account of molecular symmetry, and although this is easy for open-chain systems, it is much less straightforward for cyclic ones. You usually will be able to avoid this problem by always trying to set up the orbitals so that the transition state will have no nodes, or just one node at a point where a bond is being made or broken.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/21%3A_Resonance_and_Molecular_Orbital_Methods/21.11%3A_Pericyclic_Reactions.txt |
We have not given you much evidence to decide why it is that some thermal [2 + 2] cycloadditions occur but not others. What is special about fluoroalkenes, allenes, and ketenes in these reactions? One possibility is that Mobius rather than the Huckel transition states are involved, but the Mobius transition states are expected to suffer from steric hindrance (Section 21-10B). It is also possible that [2 + 2] cycloadditions, unlike the Diels-Alder additions, proceed by stepwise mechanisms. This possibility is strongly supported by the fact that these reactions generally are not stereospecific. Thus with tetrafluoroethene and trans,trans-2,5-hexadiene two products are formed, which differ in that the 1-propenyl group is trans to the methyl group in one adduct, $45$, and cis in the other, $46$:
A stepwise reaction involving a biradical intermediate accounts for the formation of both $45$ and$46$. In the biradical mechanism the first step is formation of just one $\ce{C-C}$ bond between the reactants, and this could occur in two different ways to give $47$ or $48ab$:
Of these, $48ab$ is predicted to have substantial electron delocalization because of the nearly equivalent VB structures $48a$ and $48b$. By the simple MO theory $48ab$ should have a delocalization energy of $16 \: \text{kcal mol}^{-1}$ (Section 21-5B). The biradical $47$ has no comparable electron delocalization and would be expected to be formed much less readily.
Collapse of $48$ through formation of the second $\ce{C-C}$ bond would give $45$ and an overall stereospecific addition. However, rotation around the $\ce{C-C}$ single bond of $48$ forms a different radical conformation, $49$, which would collapse to the other stereoisomer, $46$:
If the reaction is stepwise, why is it stepwise? In the first place, as we have seen (Section 21-10A), there are theoretical reasons why [2 + 2] cycloadditions may not occur in a concerted manner. Second, there are thermodynamic reasons why some alkenes undergo stepwise [2 + 2] additions and others do not. Regarding the second point, we can estimate that $2 \ce{CH_2=CH_2} \rightarrow \cdot \ce{CH_2-CH_2-CH_2-CH_2} \cdot$ has $\Delta H^0 \sim 37 \: \text{kcal}$, which is too high to achieve at a useful rate at those temperatures where the equilibrium constant is favorable for cyclobutane formation. In other words, when $K_\text{eq}$ is favorable, the rate is too slow, and when the rate is fast enough, $K_\text{eq}$ is unfavorable. In contrast, $2 \ce{CF_2=CF_2} \rightarrow \cdot \ce{CF_2-CF_2-CF_2-CF_2} \cdot$ is estimated to have $\Delta H^0 = -7 \: \text{kcal}$! This tells us that $\ce{CF_2=CF_2}$ has an abnormally low $\ce{C=C}$ $\pi$-bond energy and, in fact, $\Delta H^0$ for addition of hydrogen to one mole of tetrafluoroethene $\left( -55 \: \text{kcal} \right)$ is $22 \: \text{kcal}$ more negative than $\Delta H^0$ for ethene $\left( -33 \: \text{kcal} \right)$. If formation of $\cdot \ce{CF_2-CF_2-CF_2-CF_2} \cdot$ from $2 \ce{CF_2=CF_2}$ actually is exothermic, then it may seem surprising that $\ce{CF_2=CF_2}$ can be kept in a container without immediately reacting with itself. That it can is because fairly high-energy collisions are required to overcome the nonbonded repulsions that resist bringing the carbons close enough together to permit the formation of the biradical. Nonetheless, $\ce{CF_2=CF_2}$ generally is regarded as a hazardous and unpredictable chemical by virtue of its unusually low $\ce{C=C}$ $\pi$-bond strength.
1,2-Propadiene also appears to have the potential for much easier formation of a biradical than does ethene. Not all [2 + 2] cycloadditions proceed by biradical mechanisms, some clearly occur by stepwise reactions involving ionic intermediates.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/21%3A_Resonance_and_Molecular_Orbital_Methods/21.12%3A_Evidence_Bearing_on_the_Mechanism_of_2__2_Cycloadditions.txt |
Benzene and other aromatic hydrocarbons usually have such strikingly different properties from typical open-chain conjugated polyenes, such as 1,3,5-hexatriene, that it is convenient to consider them as a separate class of compounds called arenes. In this chapter we shall outline the essential features of the chemistry of arenes, particularly their reactions with electrophilic reagents which result in the substitution of a ring hydrogen with other functional groups. Some of the important properties of benzene were discussed in Chapter 21 in connection with the valence-bond and molecular-orbital theories, which rationalize the bonding in benzene and account for the remarkable stability and low reactivity of benzene (Section 21-3A). This chapter is especially concerned with chemical properties of benzene and its derivatives as well as related ring systems.
• 22.1: Nomenclature of Arenes
The naming of benzene derivatives is relatively straightforward. However, many benzene derivatives have acquired trivial names, and we draw your attention to a few of these below. The accepted name for the -C6H5 group as a substituent is phenyl
• 22.2: Physical Properties of Arenes
The pleasant odors of the derivatives of many arenes is the origin of the name aromatic hydrocarbons. The arenes themselves generally are quite toxic; some are carcinogenic and inhalation of their vapors should be avoided. The volatile arenes are highly flammable and burn with a luminous sooty flame, in contrast to alkanes and alkenes, which usually burn with a bluish flame leaving little carbon residue.
• 22.3: Spectral Properties of Arenes
The presence of a phenyl group in a compound can be ascertained with a fair degree of certainty from its infrared spectrum.
• 22.4: Electrophilic Aromatic Substitution
In this section we shall be mainly interested in the reactions of arenes that involve attack on the carbon atoms of the aromatic ring. We shall not elaborate now on the reactions of substituent groups around the ring.
• 22.5: Effect of Substituents on Reactivity and Orientation in Electrophilic Aromatic Substitution
In planning syntheses based on substitution reactions of substituted benzenes, it is imperative to be able to predict in advance which of the available positions of the ring are likely to be most reactive. This is now possible with a rather high degree of certainty, thanks to the work of many chemists during the past 100 years. Few, if any, other problems in organic chemistry have received so much attention over so many years.
• 22.6: Orientation in Disubstituted Benzenes
The orientation and reactivity effects of substituents discussed for the substitution of monosubstituted benzenes also hold for disubstituted benzenes, except that the directing influences now come from two groups. Qualitatively, the effects of the two substituents are additive on the reactivity. We therefore would expect 4-nitromethylbenzene to be less reactive than methylbenzene because of the deactivating effect of a nitro group.
• 22.7: IPSO Substitution
For all practical purposes, electrophilic aromatic substitution is confined to the substitution of a ring hydrogen. Attack at the substituted carbon evidently does occur, but it does not lead directly to substitution products because demethylation, unlike deprotonation, does not occur. Instead, the nitro group changes positions to the neighboring ring carbon, which then can eliminate a proton to form a substitution product.
• 22.8: Substitution Reactions of Polynuclear Aromatic Hydrocarbons
Although naphthalene, phenanthrene, and anthracene resemble benzene in many respects, they are more reactive than benzene in both substitution and addition reactions. This increased reactivity is expected on theoretical grounds because quantum-mechanical calculations show that the net loss in stabilization energy for the first step in electrophilic substitution or addition decreases progressively from benzene to anthracene.
• 22.9: Addition Reactions of Arenes
Benzenoid compounds are not readily converted to cyclohexane derivatives. Nevertheless, several addition reactions are carried out on an industrial scale. Mention was made previously of the hydrogenation of benzene to cyclohexane in the presence of a nickel catalyst.
• 22.10: Oxidation Reactions
The reagents usually employed for the oxidation of alkenes normally do not attack benzene. Although at high temperatures, benzene can be oxidized to cis-butenedioic (maleic) anhydride by air with a vanadium pentoxide catalyst. Ozonization of aromatic hydrocarbons is possible .e.g, ozonation of benzene gives ethanedial (glyoxal).
• 22.11: Sources and Uses of Aromatic Hydrocarbons
Benzene and many of its derivatives are manufactured on a large scale for use in high-octane gasolines and in the production of polymers, insecticides, detergents, dyes, and many miscellaneous chemicals. Most of the benzene and almost all of the methylbenzene and the dimethylbenzenes produced in the United States are derived from petroleum.
• 22.12: Some Conjugated Cyclic Polyenes
There are several compounds that possess some measure of aromatic character typical of benzene, but do not possess a benzenoid ring. Appropriately, they have (4n+2)(4n+2) ππ electrons and are classified as nonbenzenoid aromatic compounds
• 22.13: Fluxional Compounds
A number of compounds are known to rearrange from one structure to an entirely equivalent structure, sometimes with extraordinary facility. Such compounds are said to be fluxional to distinguish from tautomers (which usually involve rearrangements between nonequivalent structures).
• 22.E: Arenes, Electrophilic Aromatic Substitution (Exercises)
These are the homework exercises to accompany Chapter 22 of the Textmap for Basic Principles of Organic Chemistry (Roberts and Caserio).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
22: Arenes Electrophilic Aromatic Substitution
The naming of benzene derivatives was considered in Section 3-5 and is relatively straightforward. However, many benzene derivatives have acquired trivial names, and we draw your attention to a few of these below. The accepted name for the $\ce{C_6H_5}-$ group as a substituent is phenyl.
The more complex ring systems having two or more fused benzene rings have nonsystematic names and illogical numbering systems. They are described as polynuclear aromatic hydrocarbons, the three most important examples being naphthalene, anthracene, and phenanthrene. In anthracene the rings are connected in a linear manner, whereas in phenanthrene they are connected angularly:
The accepted numbering system for these hydrocarbons is as shown in the structures. The 1- and 2-positions of the naphthalene ring sometimes are designated as $\alpha$ and $\beta$, but we prefer not to use these designations. Some illustrative substitution products are:
The names that have been give to these and other more elaborate types of polynuclear aromatic hydrocarbons are for the most part distressingly uninformative with respect to their structures.$^1$
$^1$A thorough summary of names and numbering systems has been published by A. M. Patterson, L. T. Capell, and D. F. Walker, Ring Index, 2nd ed., American Chemical Society, 1960. Less complete but useful summaries are given in various handbooks of chemistry.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
22.02: Physical Properties of Arenes
The pleasant odors of the derivatives of many arenes is the origin of the name aromatic hydrocarbons. The arenes themselves generally are quite toxic; some are carcinogenic and inhalation of their vapors should be avoided. The volatile arenes are highly flammable and burn with a luminous sooty flame, in contrast to alkanes and alkenes, which usually burn with a bluish flame leaving little carbon residue.
The more common arenes and their physical properties are given in Table 22-1. They are less dense than water and are highly insoluble. Boiling points increase regularly with increasing molecular weight, but there is little correlation between melting point and molecular weight. The melting point is highly dependent on the symmetry of the compound; thus benzene melts $100^\text{o}$ higher than methylbenzene, and the more symmetrical 1,4-dimethylbenzene (para-xylene) has a higher boiling point than either the 1,2- or the 1,3-isomer. This latter fact is utilized in the separation by fractional crystallization of 1,4-dimethylbenzene from mixtures of isomers produced from petroleum.
Table 22-1: Physical Properties of Arenes
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/22%3A_Arenes_Electrophilic_Aromatic_Substitution/22.01%3A_Nomenclature_of_Arenes.txt |
Infrared Spectra
The presence of a phenyl group in a compound can be ascertained with a fair degree of certainty from its infrared spectrum. For example, in Figure 22-1 we see the infrared spectra of methylbenzene, and of 1,2-, 1,3-, and 1,4-dimethylbenzene. That each spectrum is of a benzene derivative is apparent from certain common features. The two bands near $1600 \: \text{cm}^{-1}$ and $1500 \: \text{cm}^{-1}$, although of variable intensity, have been correlated with the stretching vibrations of the carbon-carbon bonds of the aromatic ring; also, the sharp bands near $3030 \: \text{cm}^{-1}$ are characteristic of aromatic $\ce{C-H}$ bonds. Other bands in the spectra, especially those between $1650 \: \text{cm}^{-1}$ and $2000 \: \text{cm}^{-1}$, between $1225 \: \text{cm}^{-1}$ and $950 \: \text{cm}^{-1}$, and below $900 \: \text{cm}^{-1}$, have been correlated with the number and position of ring substituents. Although we shall not document all these various bands in detail, each of the spectra in Figure 22-1 is marked to show some of the correlations that have been made.
Electronic Absorption Spectra
Compared to straight-chain conjugated polyenes, aromatic compounds have relatively complex absorption spectra with several bands in the ultraviolet region. Benzene and the alkylbenzenes show two bands in which we shall be primarily interested, one near $200 \: \text{nm}$ and the other near $260 \: \text{nm}$. The $200$-$\text{nm}$ band is of fairly high intensity and corresponds to excitation of a $\pi$ electron of the conjugated system to a $\pi^*$ orbital (i.e., a $\pi \rightarrow \pi^8*$ transition). The excited state has significant contributions from dipolar structures such as $1$:
This is analogous to the absorption bands of conjugated dienes (Section 9-9B) except that the wavelength of absorption of benzenes is shorter. In fact, the $200$-$\text{nm}$ absorptions of benzene and the alkylbenzenes are just beyond the range of most commercial quartz spectrometers. However, these absorptions (which we say arise from the benzene chromophore$^2$) are intensified and shifted to longer wavelengths when the conjugated system is extended by replacement of the ring hydrogens by unsaturated groups (e.g., $\ce{-CH=CH_2}$, $\ce{-C \equiv CH}$, $\ce{-CH=O}$, and $\ce{-C \equiv N}$; see Table 22-2). The delocalized $\pi$-electron system of the absorbing chromophore now includes the electrons of the unsaturated substituent as well as those of the ring. In the specific case of ethenylbenzene the excited state is a hybrid structure composite of $2a$ and $2b$ and other related dipolar structures:
Table 22-2: Effect of Conjugation on Electronic Absorption by the Benzene Chromophore
Similar effects are observed for benzene derivatives in which the substituent has unshared electron pairs that can be conjugated with the benzene ring (e.g., $\ce{-NH_2}$, $\ce{-OH}$, $\ce{-Cl}$). An unshared electron pair is to some extent delocalized to become a part of the aromatic $\pi$-electron system in both the ground and excited states, but more importantly in the excited state. This is illustrated for benzenamine (aniline) by the following structures, which contribute to the hybrid structure:
The data of Table 22-3 show the effect on the benzene chromophore of this type of substituent - the substituent often being called an auxochrome.$^2$ This term means that, although the substituent itself is not responsible for the absorption band, it shifts the absorption of the chromophoric group, in this case the benzene ring, toward longer wavelengths. The auxochromic groups usually increase the intensity of the absorption also.
Table 22-3: Effect of Auxochromic Substituents on Electronic Absorption by the Benzene Chromophore
Table 22-4: Effects of Structure on Electronic Absorption Corresponding to the Benzenoid Band
The benzenoid band corresponds to a low-energy $\pi \rightarrow \pi^*$ transition of the benzene molecules. The absorption intensity is weak because the $\pi^*$ state involved has the same electronic symmetry as the ground state of benzene, and transitions between symmetrical states usually are forbidden. The transitions are observed in this case only because the vibrations of the ring cause it to be slightly distorted at given instants. In the valence-bond treatment this excited state of benzene is an antibonding state of the $\pi$ electrons.
The electronic spectra of polynuclear aromatic hydrocarbons such as naphthalene and anthracene, in which aromatic rings are fused together in a linear manner, resemble the spectrum of benzene except that the bands are shifted to longer wavelengths. In fact, with the four linearly connected rings of naphthacene, the benzenoid band is shifted far enough to longer wavelengths to be in the visible region of the spectrum (see Table 22-4). Accordingly, naphthacene is yellow. The next higher member, pentacene, is blue.
Compounds such as phenanthrene, chrysene, and pyrene, in which the aromatic rings are fused in an angular manner, have complex electronic spectra with considerable fine structure. The $\lambda_\text{max}$ values normally are at shorter wavelengths than those of their linear isomers.
Nuclear Magnetic Resonance Spectra
The chemical shifts of arene protons ($6.5 \: \text{ppm}$ to $8.0 \: \text{ppm}$) characteristically are toward lower magnetic fields than those of protons attached to ordinary double bonds ($4.6 \: \text{ppm}$ to $6.9 \: \text{ppm}$). The difference of about $2 \: \text{ppm}$ cannot be easily explained because the hydrogens in both types of systems are bonded to carbon through $sp^2$-$\sigma$ bonds (Sections 6-4C and 6-5A).
At least part of the chemical-shift difference between arene protons and alkene protons is the result of the special property of $\pi$ electrons in aromatic systems of circulating freely above and below the plane of the carbon nuclei, as shown in Figure 22-4. When a molecule such as benzene is subjected to a magnetic field that has a component perpendicular to the plane of the ring, the electrons circulate around the ring in such a way as to produce a local magnetic dipole in the direction opposite to the applied field. This diamagnetic shielding effect acts to reduce the applied field in the center of the ring. Therefore, if a proton could be located in the center of the ring, the applied field would have to be higher than normal to counteract the local diamagnetic field and bring the proton into resonance. A proton outside the ring is affected in the opposite way (paramagnetic deshielding effect) because, as can be seen from the diagram, such protons are located in the return path of the lines of force associated with the local field and thus are in a field greater than that arising from the external magnet alone. When the plane of the molecule is oriented parallel to the field, the diamagnetic circulation is cut off. As a result, as the molecules tumble over and over in the liquid the component of magnetization perpendicular to the plane of the ring varies rapidly. Nonetheless, a substantial net paramagnetic effect is experienced by the ring hydrogens. The resonance line positions therefore are shifted to lower magnetic fields.
Strong evidence in confirmation of the above explanation of the chemical shifts of aromatic hydrogens is provided by a study of the cyclic conjugated polyene [18]annulene, which has hydrogens both "inside" and "outside" the ring:
The inside hydrogens are strongly deshielded, coming at $1.9 \: \text{ppm}$ upfield from tetramethylsilane, while the outside hydrogens are deshielded and come at $8.8 \: \text{ppm}$ downfield from TMS. As we shall see, the ring current effect is quite general and constitutes a widely used test for aromatic character in conjugated polyene ring systems.
In general, the spin-spin splittings observed between the five protons of a phenyl group can be extremely complex. An example is afforded by nitrobenzene (Figure 22-5), which has different chemical shifts for its ortho, meta, and para hydrogens and six different spin-spin interaction constants: $J_{23}$, $J_{24}$, $J_{25}$, $J_{26}$, $J_{34}$, $J_{35}$, (the subscripts correspond to position numbers of the protons):
Such a spectrum is much too complex to be analyzed by any simple procedure. Nonetheless, nuclear magnetic resonance can be useful in assigning structures to aromatic derivatives, particularly in conjunction with integrated line intensities and approximate values of the coupling constants between the ring hydrogens, as shown below:
$^2$A chromophore is a grouping of atoms in an organic molecule that gives rise to color, or has the potential of doing so when other groups called auxochromes are present (also see Section 28-4).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/22%3A_Arenes_Electrophilic_Aromatic_Substitution/22.03%3A_Spectral_Properties_of_Arenes.txt |
Scope and Mechanism
In this section we shall be mainly interested in the reactions of arenes that involve attack on the carbon atoms of the aromatic ring. We shall not elaborate now on the reactions of substituent groups around the ring.
The principal types of reactions involving aromatic rings are substitution, addition, and oxidation. Of these, the most common type is electrophilic substitution. A summary of the more important substitution reactions of benzene is given in Figure 22-7. Many of the reagents used to achieve these substitutions will be familiar to you in connection with electrophilic addition reactions to alkenes (e.g., $\ce{Cl_2}$, $\ce{Br_2}$, $\ce{H_2SO_4}$, and $\ce{HOCl}$; Section 10-3). Electrophilic addition to alkenes and electrophilic aromatic substitution are both polar, stepwise processes, and the key step for each is attack of an electrophile at carbon to form a cationic intermediate. We may represent this type of reaction by the following general equations, in which the attacking reagent is represented either formally as a cation, $\ce{X}^\ominus$, or as a neutral but polarized molecule, $\overset{\delta \oplus}{\ce{X}}$---$\overset{\delta \ominus}{\ce{Y}}$:
electrophilic aromatic substitution (first step)
electrophilic addition to alkenes (first step)
The intermediate shown for aromatic substitution no longer has an aromatic structure; rather, it is a cation with four $\pi$ electrons delocalized over five carbon nuclei, the sixth carbon being saturated with $sp^3$-hybrid bonds. It may be formulated in terms of the following contributing structures, which are assumed to contribute essentially equally:
The importance of writing the hybrid structure with the partial charges at these three positions will become evident later. This kind of ion is referred to as a $\sigma$ complex or a benzenium ion.
The aromatic ring is regenerated from this cationic intermediate by loss of a proton from the $sp^3$-hybridized carbon. The electron pair of this $\ce{C-H}$ bond then becomes part of the aromatic $\pi$-electron system and a substitution product of benzene, $\ce{C_6H_5X}$, is formed.
electrophilic aromatic substitution (second step)
The gain in stabilization attendant on regeneration of the aromatic ring is sufficiently advantageous that this, rather than combination of the cation with $\ce{Y}^\ominus$, normally is the favored course of reaction. Herein lies the difference between aromatic substitution and alkene addition. In the case of alkenes there usually is no substantial resonance energy to be gained by loss of a proton from the intermediate, which tends therefore to react by combination with a nucleophilic reagent.
electrophilic addition to alkenes (second step)
$\overset{\oplus}{\ce{C}} \ce{H_2-CH_2X} + \ce{Y}^\ominus \rightarrow \ce{YCH_2-CH_2X}$
Nature of the Substituting Agent
It is important to realize that in aromatic substitution the actual electrophilic substituting agent, $\overset{\oplus}{\ce{X}}$ or $\overset{\delta \oplus}{\ce{X}}-\overset{\delta \ominus}{\ce{Y}}$, is not necessarily the reagent that is added to the reaction mixture. For example, nitration in mixtures of nitric and sulfuric acids is not brought about by attack of the nitric acid molecule on the aromatic compound, but by attack of a more electrophilic species, the nitronium ion, $\ce{NO_2^+}$. This ion is formed from nitric acid and sulfuric acid according to the following equation:
$\ce{HNO_3} + 2 \ce{H_2SO_4} \rightleftharpoons \ce{NO_2^+} + \ce{H_3O^+} + 2 \ce{HSO_4^-}$
The nitronium ion attacks the aromatic ring to give first a nitrobenzenium ion and then an aromatic nitro compound:
In general, the function of a catalyst (which is so often necessary to promote aromatic substitution) is to generate an electrophilic substituting agent from the given reagents. Thus it is necessary to consider carefully for each substitution reaction what the actual substituting agent may be. This problem does not arise to the same degree in electrophilic additions to alkenes, because alkenes are so much more reactive than arenes that the reagents employed (e.g., $\ce{Br_2}$, $\ce{Cl_2}$, $\ce{HCl}$, $\ce{HOCl}$, $\ce{HOBr}$, $\ce{H_3O}^\oplus$) themselves are sufficiently electrophilic to react with alkenes without the aid of a catalyst. In fact, conditions that lead to substitution of arenes, such as nitration in mixtures of nitric and sulfuric acid, often will degrade the carbon skeleton of alkenes.
Now we shall consider the individual substitution reactions listed in Figure 22-1 with regard to the nature of the substituting agent and the utility for synthesis of various classes of aromatic compounds.
Nitration
The nitronium ion, $\ce{NO_2^+}$, is the active nitrating agent in nitric acid-sulfuric acid mixtures. The nitration of methylbenzene (toluene) is a typical example of a nitration that proceeds well using nitric acid in a 1:2 mixture with sulfuric acid. The nitration product is a mixture of 2-, 3-, and 4-nitromethylbenzenes:
The presence of appreciable amounts of water in the reaction mixture is deleterious because water tends to reverse the reaction by which nitronium ion is formed:
$\ce{NO_2^+} + \ce{H_2O} \overset{\ce{HSO_4^-}}{\rightleftharpoons} \ce{HNO_3} + \ce{H_2SO_4}$
For this reason the potency of a nitric-sulfuric acid mixture can be considerably increased by using fuming nitric and fuming sulfuric acids. With such mixtures nitration of relatively unreactive compounds can be achieved. For example, 4-nitromethylbenzene is far less reactive than methylbenzene, but when heated with an excess of nitric acid in fuming sulfuric acid, it can be converted successively to 2,4-dinitromethylbenzene and to 2,4,6-trinitromethylbenzene (TNT):
There are several interesting features about the nitration reactions thus far discussed. For instance, the conditions required for nitration of 4-nitromethylbenzene would rapidly oxidize an alkene by cleavage of the double bond:
Also the mononitration of methylbenzene does not lead to equal amounts of the three possible products. The methyl substituent apparently orients the entering substituent preferentially to the 2 and 4 positions. This aspect of aromatic substitution will be discussed in Section 22-5 in conjunction with the effect of substituents on the reactivity of aromatic compounds.
Some compounds are sufficiently reactive that they can be nitrated with nitric acid in ethanoic acid. Pertinent examples are 1,3,5-trimethylbenzene and naphthalene:
Other convenient nitrating reagents are benzoyl nitrate, $\ce{C_6H_5COONO_2}$, and ethanoyl nitrate, $\ce{CH_3COONO_2}$. These reagents provide a source of $\ce{NO_2^+}$ and have some advantage over $\ce{HNO_3} \cdot \ce{H_2SO_4}$ mixtures in that they are soluble in organic solvents such as ethanenitrile or nitromethane. Having homogeneous solutions is especially important for kinetic studies of nitration. The reagents usually are prepared in solution as required from the corresponding acyl chloride and silver nitrate or from the acid anhydride and nitric acid. Such reagents are hazardous materials and must be handles with care.
Nitronium salts of the type $\ce{NO_2^+} \ce{X^-}$ are very powerful nitrating agents. The counterion, $\ce{X^-}$, must be non-nucleophilic and usually is fluoroborate, $\ce{BF_4^-}$ or $\ce{SbF_4^-}$:
Halogenation
To some degree we have oversimplified electrophilic substitution by neglecting the possible role of the 1:1 charge-transfer complexes that most electrophiles form with arenes (see Section 10-3C for discussion of analogous complexes of alkenes):
With halogens, especially iodine, complex formation is visually evident from the color of solutions of the halogen in arenes. Although complex formation may assist substitution by bringing the halogen and arene in close proximity, substitution does not necessarily occur. A catalyst usually is required, and the catalysts most frequently used are metal halides that are capable of accepting electrons (i.e., Lewis acids such as $\ce{FeBr_3}$, $\ce{AlCl_3}$, and $\ce{ZnCl_2}$). Their catalytic activity may be attributed to their ability to polarize the halogen-halogen bond in the following way:
$\overset{\delta \oplus}{\ce{Br}} \cdots \overset{\delta \ominus}{\ce{Br}} \cdots \ce{FeBr_3}$
The positive end of the dipole attacks the aromatic compound while the negative end is complexed with the catalyst. We can represent the reaction sequence as follows, with the slow step being formation of a $\sigma$ bond between $\ce{Br}^\oplus$ and the aromatic ring:
The order of reactivity of the halogens is $\ce{F_2} > \ce{Cl_2} > \ce{Br_2} > \ce{I_2}$. Fluorine is too reactive to be of practical use for the preparation of aromatic fluorine compounds and indirect methods are necessary (see Section 23-10B). Iodine usually is unreactive. It has been stated that iodination fails because the reaction is reversed as the result of the reducing properties of the hydrogen iodide that is formed:
$\ce{C_6H_6} + \ce{I_2} \overset{\rightarrow}{\longleftarrow} \ce{C_6H_5I} + \ce{HI}$
This view is not correct because, as Kekule himself showed, iodobenzene is not reduced by hydroiodic acid except at rather high temperatures.
The way to achieve direct iodination in the absence of powerful activating substituent groups is to convert molecular iodine to some more active species (perhaps $\ce{H_2OI}^\oplus$ or $\ce{I}^\oplus$) with an oxidizing agent such as nitric acid or hydrogen peroxide:
\begin{align} \ce{I_2} + 4 \ce{HNO_3} &\rightarrow 2 \ce{H_2O-I^+} + 2 \ce{NO_2} + 2 \ce{NO_3^-} \ \ce{I_2} + \ce{H_2O_2} + 2 \ce{H^+} &\rightarrow 2 \ce{H_2OI^+} \end{align}
With combinations of this kind good yields of iodination products are obtained:
Halogen substitution reactions with chlorine or bromine must be carried out with adequate protection from strong light. If such precautions are not taken, an alkylbenzene will react rapidly with halogen by a photochemical process to substitute a hydrogen of the alkyl group rather than of the aromatic ring. The reaction has a light-induced, radical-chain mechanism of the kind discussed for the chlorination of propene (Section 14-3A). Thus methylbenzene reacts with bromine when illuminated to give phenylmethyl bromide; but when light is excluded and a Lewis acid catalyst is present, substitution occurs to give principally the 2- and 4-bromomethylbenzenes. Much less of the 3-bromomethylbenzene is formed:
Benzene itself can be induced to add halogens on strong irradiation to give polyhalocyclohexanes (see Sections 21-3A and 22-9C):
Alkylation
An important method of synthesis of alkylbenzenes utilizes an alkyl halide as the alkylating agent and a metal halide, usually aluminum chloride, as catalyst:
This class of reaction is called Friedel-Crafts alkylation in honor of its discoverers, C. Friedel (a French chemist) and J. M. Crafts (an American chemist). The metal-halide catalyst functions much as it does in halogenation reactions to provide a source of a positive substituting agent, which in this case is a carbocation:
Alkylation is not restricted to alkyl halides; alcohols and alkenes may be used to advantage in the presence of acidic catalysts such as $\ce{H_3PO_4}$, $\ce{H_2SO_4}$, $\ce{HF}$, $\ce{BF_3}$, or $\ce{HF-BF_3}$. Ethylbenzene is made commercially from benzene and ethene using phosphoric acid as the catalyst. Isopropylbenzene is made similarly from benzene and propene:
Under these conditions the carbocation, which is the active substituting agent, is generated by protonation of the alkene:
\begin{align} \ce{CH_2=CH_2} + \ce{H^+} &\rightleftharpoons \ce{CH_3CH_2^+} \ \ce{CH_3CH=CH_2} + \ce{H^+} &\rightleftharpoons \ce{CH_3} \overset{+}{\ce{C}} \ce{HCH_3} \end{align}
With alcohols the electrophile can be formed by initial protonation by the acid catalyst and subsequent cleavage to a carbocation:
Limitations of Alkylation Reactions
Polysubstitution
There are several factors that limit the usefulness of alkylation reactions. First, it may be difficult to limit reaction to monosubstitution because introduction of one alkyl substituent tends to activate the ring towards a second substitution (see Section 22-5). Therefore, to obtain reasonable yields of a monoalkylbenzene, it usually is necessary to use a large excess relative to the alkylating agent:
Rearrangement of the alkylating agent
A second limitation is the penchant for the alkylating reagent to give rearrangement products. As an example, the alkylation of benzene with 1-chloropropane leads to a mixture of propylbenzene and isopropylbenzene. We may write the reaction as first involving formation of a propyl cation, which is a primary carbocation:
$\ce{CH_3CH_2CH_2Cl} + \ce{AlCl_3} \rightarrow \ce{CH_3CH_2CH_2^+} + \overset{-}{\ce{Al}} \ce{Cl_4}$
This ion either can alkylate benzene to give propylbenzene,
$\ce{C_6H_6} + \ce{CH_3CH_2CH_2^+} \rightarrow \ce{C_6H_5CH_2CH_2CH_3} + \ce{H^+}$
or it can rearrange to a more stable secondary ion by the transfer of a hydrogen from a neighboring carbon together with its bonding electron pair (i.e., 1,2-hydride shift). The positive charge is thereby transferred from $\ce{C_1}$ to $\ce{C_2}$:
Alkylation of benzene with the isopropyl cation then produces isopropylbenzene:
$\ce{C_6H_6} + \ce{CH_3} \overset{\oplus}{\ce{C}} \ce{HCH_3} \rightarrow \ce{C_6H_5CH(CH_3)_2} + \ce{H}^\oplus$
Rearrangements of this type involving carbocation intermediates often occur in Friedel-Crafts alkylations with primary and secondary alkyl groups larger than $\ce{C_2}$ and $\ce{C_3}$. Related carbocation rearrangements are discussed in Sections 8-9B and 15-5E.
Rearrangement of products
Further complications arise from the fact that the alkylation reactions sometimes are under equilibrium control rather than kinetic control. Products often isomerize and disproportionate, particularly in the presence of large amounts of catalyst. Thus 1,2- and 1,4-dimethylbenzenes (ortho- and para-xylenes) are converted by large amounts of Friedel-Crafts catalysts into 1,3-dimethylbenzene (meta-xylene):
Ethylbenzene disproportionates under the influence of excess $\ce{HF-BF_3}$ to benzene and 1,3-diethylbenzene:
Acylation
Acylation and alkylation of arenes are closely related. Both reactions were developed as the result of the collaboration between Friedel and Crafts, in 1877. The acylation reaction introduces an acyl group, $\ce{RCO}$, into an aromatic ring and the product is an aryl ketone:
The acylating reagents commonly used are carboxylic acid halides, $\ce{RCOCl}$, anhydrides, $\ce{(RCO)_2O}$, or the acid itself, $\ce{RCO_2H}$. A strong proton or other Lewis-acid catalyst is essential. The catalyst functions to generate the acyl cation:
The catalyst most commonly used with acyl halides and anhydrides is aluminum chloride:
Acylation differs from alkylation in that the reaction usually is carried out in a solvent, commonly carbon disulfide, $\ce{CS_2}$, or nitrobenzene. Furthermore, acylation requires more catalyst than alkylation, because much of the catalyst is tied up and inactivated by complex formation with the product ketone:
Unlike alkylation, acylation is controlled easily to give monosubstitution, because once an acyl group is attached to a benzene ring, it is not possible to introduce a second acyl group into the same ring. Because of this, a convenient synthesis of alkylbenzenes starts with acylation, followed by reduction of the carbonyl group with zinc and hydrochloric acid (Section 16-6). For example, propylbenzene is prepared best by this two-step route because, as we have noted, the direct alkylation of benzene with propyl chloride produces considerable amounts of isopropylbenzene and polysubstitution products:
In the acylation of alkylbenzene the product almost always is the para isomer. The synthesis of (4-tert-butylphenyl)ethanone illustrates this as well as the sequential use of alkylation and acylation reactions:
Chemists are inclined to give names to reactions that associate them either with their discoverers or with the products they give. This practice can be confusing because many named reactions (or "name reactions") which once were thought to be quite unrelated, have turned out to have very similar mechanisms. Thus we have two closely related acylation reactions: one is the Friedel-Crafts ketone synthesis, in which the electrophile is $\ce{R} \ce{-} \overset{\oplus}{\ce{C}} \ce{=O}$; and the other is the Gattermann-Koch aldehyde synthesis, in which the electrophile is $\ce{H} \ce{-} \overset{\oplus}{\ce{C}} \ce{=O}$:
The latter reaction utilizes carbon monoxide and $\ce{HCl}$ under pressure in the presence of aluminum chloride. The electrophile may be considered to be formed as follows:
$\ce{C=O} + \ce{HCl} + \ce{AlCl_3} \rightleftharpoons \ce{H} \ce{-} \overset{\oplus}{\ce{C}} \ce{=O} + \overset{\ominus}{\ce{Al}} \ce{Cl_4}$
Sulfonation
Substitution of the sulfonic acid $\left( \ce{-SO_3H} \right)$ group for a hydrogen of an aromatic hydrocarbon can be carried out by heating the hydrocarbon with a slight excess of concentrated or fuming sulfuric acid:
The actual sulfonating agent normally is the $\ce{SO_3}$ molecule, which, although neutral, has a powerfully electrophilic sulfur atom:
Sulfonation is reversible and the $\ce{-SO_3H}$ group can be removed by hydrolysis at $180^\text{o}$:
A useful alternative preparation of sulfonyl derivatives is possible with chlorosulfonic acid:
This procedure has an advantage over direct sulfonation in that sulfonyl chlorides usually are soluble in organic solvents and may be easily separated from the reaction mixture. Also, the sulfonyl chloride is a more useful intermediate than the sulfonic acid, but can be converted to the acid by hydrolysis if desired:
Sulfonation is important in the commercial production of an important class of detergents - the sodium alkylbenzenesulfonates:
The synthesis illustrates several important types of reactions that we have discussed in this and previous chapters. First, the alkyl group $\ce{R}$ usually is a $\ce{C_{12}}$ group derived from the straight-chain hydrocarbon, dodecane, which on photochlorination gives a mixture of chlorododecanes:
This mixture of chlorododecanes is used to alkylate benzene, thereby giving a mixture of isomeric dodecylbenzenes, called detergent alkylate:
Sulfonation of the detergent alkylate gives exclusively the 4-dodecylbenzenesulfonic acids, which with sodium hydroxide form water-soluble dodecylbenzenesulfonates:
In many countries it is prohibited by law to market detergents of this type, which have highly branched alkyl groups. The reason is that quaternary carbons and, to a lesser extent, tertiary carbons are not degraded readily by bacteria in sewage treatment plants:
Hydrogen Exchange
It is possible to replace the ring hydrogens of many aromatic compounds by exchange with strong acids. When an isotopically labeled acid such as $\ce{D_2SO_4}$ is used, this reaction is an easy way to introduce deuterium. The mechanism is analogous to other electrophilic substitutions:
Perdeuteriobenzene$^3$ can be made from benzene in good yield if a sufficiently large excess of deuteriosulfuric acid is used. Deuteration might appear to be competitive with sulfonation, but deuteration actually occurs under much milder conditions.
Aromatic Substitution by Electrophilic Metalation
Because metals are electropositive elements they can be considered potential electrophiles. Their reactions with arenes have been investigated most thoroughly for mercury. Benzene can be substituted with $\ce{HgX}^\oplus$ derived from a mercuric salt, $\ce{HgX_2}$, in the presence of an acid catalyst. The salt most commonly used is mercuric ethanoate, $\ce{Hg(OOCCH_3)_2}$. The catalyst is considered to function by assisting the generation of the active electrophile, $\ce{HgX}^\oplus$. Other metals that may be introduced directly into an aromatic ring in this manner include thallium and lead.
$^3$The prefix per, as in perdeuterio- or perfluoro-, means that all the hydrogens have been replaced with the named substituent, $\ce{D}$ or $\ce{F}$. Perhydro means saturated or fully hydrogenated.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/22%3A_Arenes_Electrophilic_Aromatic_Substitution/22.04%3A_Electrophilic_Aromatic_Substitution.txt |
In planning syntheses based on substitution reactions of substituted benzenes, it is imperative to be able to predict in advance which of the available positions of the ring are likely to be most reactive. This is now possible with a rather high degree of certainty, thanks to the work of many chemists during the past 100 years. Few, if any, other problems in organic chemistry have received so much attention over so many years, and there are now sufficient data on the orientation and reactivity effects of ring substituents in electrophilic substitution to permit the formation of some very valuable generalizations.
Basically, three experimental problems are involved in the substitution reactions of aromatic compounds: (1) proof of structure of the isomers that are formed; (2) determination of the percentage of each isomer formed, if the product is a mixture; and (3) measurement of the reactivity of the compound being substituted relative to some standard substance, usually benzene.
For benzenoid compounds, structures can be established by the historically important substitution method (Section 1-1F) or with the aid of correlations between spectroscopic properties and positions of substitution, as we indicated in Section 22-3. Also, it is often possible to identify the isomers by converting them to compounds of known structure. For example, trifluoromethylbenzene on nitration gives only one product, which has been shown to be the 3-nitro derivative by conversion to the known 3-nitrobenzoic acid by concentrated sulfuric acid:
The ratios of isomers formed in substitution reactions can be determined by spectroscopic means or by the analytical separation methods discussed in Section 9-2. We mainly are concerned in this chapter with the reactivity and orientation observed in aromatic substitution.
The Pattern of Orientation in Aromatic Substitution
The reaction most studied in connection with the orientation problem is nitration, but the principles established also apply for the most part to the related reactions of halogenation, sulfonation, alkylation, and acylation. Some illustrative data for the nitration of a number of mono-substituted benzene derivatives are given in Table 22-5. The table includes the percentage of ortho, meta, and para isomers formed, along with their reactivities relative to benzene. We see that there is a wide range of reactivity according to the nature of the substituent, and that the ortho, meta, and para positions are not equally reactive. Although these substituent effects may appear complex, they are related closely to substituted alkenes (Section 10-4), as will be explained in the following section.
Table 22-5: Orientation and Rate Data for Nitration of Some Monosubstituted Benzene Derivatives$^a$
Electronic Effects
It is helpful to construct an energy diagram for substitution by an electrophilic $\ce{X}^\oplus$ of a benzene derivative, $\ce{C_6H_5Y}$, in which $\ce{Y}$ is a substituent group (Figure 22-8). The rate of substitution at any one position (we have arbitrarily chosen in Figure 22-8 to compare the 3 and 4 positions) will depend on the height of the energy barrier between the reactants and the transition state. Effects that act to lower the heights of the barriers increase the rates of substitution. Because the transition state and the positively charged intermediate for aromatic substitution have much the same energy, any effect that stabilizes this intermediate is likely to lower the energy of the transition state and increase the rate of substitution. Thus under conditions of kinetic control the preferred arene substitution product, as in alkene addition, will be that derived from the most stable of the possible intermediates. Therefore the problem of predicting relative rates and orientation in aromatic substitution becomes one of deciding what factors are likely to stabilize or destabilize the various possible intermediates relative to one another and to the ground state.
We now can examine the structures of the three substitution intermediates with a view to deciding how the substituent might affect their stability. According to the valence-bond method, the positive charge in the ring is dispersed mainly on alternate carbons, as shown below.
ortho substitution
para substitution
meta substitution
The substituent $\ce{Y}$ should (and does) exert its electronic influence more strongly from the ortho and para positions than from the meta position because $\ce{Y}$ in the ortho and the para positions is close to a positively charged ring carbon. This electronic influence will be stabilizing if $\ce{Y}$ has a net electron-donating effect, and destabilizing if $\ce{Y}$ is electron withdrawing. A group can withdraw electrons relative to hydrogen if it is more electronegative than hydrogen and this is called the electron-withdrawing inductive effect (also see Section 18-2B). A group also can withdraw electrons by the resonance effect:
Accordingly, substituents fall into one of the following categories.
Meta-directing substituents
A ring substituent $\ce{Y}$ that is electron withdrawing relative to hydrogen and has no capacity to donate electrons by a resonance effect will decrease the reactivity of $\ce{C_6H_5Y}$, especially at the ortho and para positions. The result is a sluggish reaction (deactivation) with substitution occurring preferentially at the meta position. Substituents in this category are $\ce{-NO_2}$, $\ce{-CF_3}$, $\ce{-CO_2R}$, $\ce{-} \overset{\oplus}{\ce{N}} \ce{R_3}$, and so on (also see Tables 22-5 and 22-6). No groups are known that direct the electrophile to the meta position and, at the same time, make the phenyl derivative more reactive relative to benzene.
Table 22-6: Orientation and Reactivity Effects of Ring Substituents
Ortho-para directing substituents
1. A ring substituent, $\ce{-Y}$, that has an electron pair on the atom adjacent to the ring gives ortho-para substitution in preference to meta substitution. The reason is that the intermediate can be stabilized by an electron-donating resonance effect from $\ce{Y}$ that is effective from the ortho and para positions only:
This effect is made clear in the valence-bond structures for the ortho-para substitution intermediates from benzenol (phenol):
Substituents of the type $\ce{-Y}$ include $\ce{-OH}$, $\ce{-OR}$, $\ce{-SR}$, $\ce{-NH_2}$, and halogens. Most of these groups also are electron withdrawing by an inductive effect that opposes their resonance effect. However, as we saw in the case of alkene additions (Section 10-4C), even when $\ce{-Y}$ is an electronegative group, stabilization of the intermediate cation by donation of unshared electrions of $\ce{Y}$ to the adjacent positive carbon more than compensates for the polar electron-withdrawing properties of $\ce{Y}$. Electron donation thus controls the orientation. If, however, the group is strongly electron withdrawing (e.g., $\ce{-Y} = \ce{-F}$, $\ce{-Cl}$, $\ce{-Br}$, $\ce{-I}$), the reactivity of the compound $\ce{C_6H_5Y}$ may be reduced. Groups of this kind are ortho-para directing with deactivation.
But if the polar effect is not pronounced, then substitution can be powerfully assisted by the substituent. This is ortho-para direction with activation and is provided by groups such as $\ce{-OH}$, $\ce{-OR}$, $\ce{-SR}$, and $\ce{-NH_2}$. A more comprehensive list of substituents and their orientation effects is provided in Table 22-6.
2. When no important $\pi$-electron effect is possible, as with alkyl groups, the orientation effect of a substituent is controlled by its polar effect and the degree to which it polarizes the bonding electrons of the ring. Alkyl groups actually are electron donating and therefore are ortho-para directing with activation.
Steric Effects
Thus far we have made no distinction between the reactivities of the ortho and the para positions, yet they clearly are not equal. If they were equal, the ortho:para ratio would be 2:1, thereby reflecting the fact that there are two ortho positions but only one para position in monosubstituted benzenes. Most substitution reactions favor the para product, sometimes by a considerable amount (see Table 22-5). A reasonable explanation is that ortho substitution is subject to steric hindrance between the substituent and the entering group. tert-Butylbenzene, for example, gives much less ortho nitration than methylbenzene (Table 22-5), thereby suggesting that the size of the substituent is important. Also, tert-butylbenzene gives no ortho alkylation with tert-butyl chloride, suggesting that the size of the entering group is also important:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/22%3A_Arenes_Electrophilic_Aromatic_Substitution/22.05%3A_Effect_of_Substituents_on_Reactivity_and_Orientation_in_Electrophilic_Aromatic_Substit.txt |
The orientation and reactivity effects of substituents discussed for the substitution of monosubstituted benzenes also hold for disubstituted benzenes, except that the directing influences now come from two groups. Qualitatively, the effects of the two substituents are additive on the reactivity. We therefore would expect 4-nitromethylbenzene to be less reactive than methylbenzene because of the deactivating effect of a nitro group. Also, the most likely position of substiuttion should be, and is, ortho to the methyl group and meta to the nitro group:
When the two substituents have opposed orientation effects, it is not always easy to predict what products will be obtained. For example, \(\ce{N}\)-(2-methoxyphenyl)ethanamide has two powerful o,p-directing substituents, \(\ce{-OCH_3}\) and \(\ce{-NHCOCH_3}\). Nitration of this compound gives mainly the 4-nitro derivative, which indicates that the \(\ce{-NHCOCH_3}\) exerts a stronger influence than \(\ce{-OCH_3}\):
Seemingly anomalous effects of substituents are known, but such effects may be due to equilibrium control. One example is the aluminum chloride-catalyzed alkylation of benzene, which leads to the formation of a 1,3,5-trialkylbenzene in preference to the expected 1,2,4-isomer (see Section 22-4E). The preferred reaction occurs particularly readily because alkylation is reversible and because alkylation is one of the least selective of the electrophilic aromatic substitutions (considerable meta isomer is formed even under conditions where kinetic control is dominant). Equilibrium control, which favors the 1,3,5-product rather than the less stable 1,2,4-product, becomes most evident when the reaction time, the reaction temperature, and aluminum chloride concentration are increased. Another source of anomalous substituent effects is discussed in the next section.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
22.07: IPSO Substitution
For all practical purposes, electrophilic aromatic substitution is confined to the substitution of a ring hydrogen. Does this mean that an electrophile such as $\ce{NO_2^+}$ only attacks hydrogen-bearing carbons? What about substituted ring carbons?
Electrophilic attack at methyl-bearing carbons, particularly in ortho- and para-dimethylbenzenes, would appear quite reasonable because the electron-donating character of the other methyl group should activate the ring by stabilizing the intermediate ion:
Attack at the substituted (ipso) carbon evidently does occur, but it does not lead directly to substitution products because demethylation, unlike deprotonation, does not occur:
Instead, the nitro group changes positions to the neighboring ring carbon, which then can eliminate a proton to form a substitution product:
Because the product obtained indirectly (by ipso substitution) is indistinguishable from that expected by direct electrophilic attack at $\ce{C_2}$, it is not possible to say how much, if any, product is formed by the ipso route in this reaction.
In general, orientation effects in the substitution of alkylbenzenes are complicated by ipso attack. For example, in the nitration of 4-methylisopropylbenzene (para-cymene) about $10\%$ of the nitration product is 4-nitromethylbenzene:
The 4-nitromethylbenzene arises from ipso attack of $\ce{NO_2^+}$ at the isopropyl-substituted ring carbon. Unlike methyl, the isopropyl group is eliminated rapidly as propene. Can we say that the other products, $3$ and $4$, arise by direct substitution? Evidently not, because nitration at $0^\text{o}$ gives two other products, $5$ and $6$, which must be formed by ipso attack at the methyl-bearing carbon. At low temperatures, intermediate ion $7$ is attacked by the weakly nucleophilic ethanoate ion to give $5$ and $6$. Both of these adducts solvolyze rapidly in $78\%$ sulfuric acid to give $3$ only:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
22.08: Substitution Reactions of Polynuclear Aromatic Hydrocarbons
Although naphthalene, phenanthrene, and anthracene resemble benzene in many respects, they are more reactive than benzene in both substitution and addition reactions. This increased reactivity is expected on theoretical grounds because quantum-mechanical calculations show that the net loss in stabilization energy for the first step in electrophilic substitution or addition decreases progressively from benzene to anthracene; therefore the reactivity in substitution and addition reactions should increase from benzene to anthracene.
In considering the properties of the polynuclear hydrocarbons relative to benzene, it is important to recognize that we neither expect nor find that all the carbon-carbon bonds in polynuclear hydrocarbons are alike or correspond to benzene bonds in being halfway between single and double bonds.
The 1,2 bonds in both naphthalene and antracene are in fact shorter than the other ring bonds, whereas the 9,10 bond in phenanthrene closely resembles an alkene double bond in both its length and chemical reactivity.
Naphthalene
Orientation in the substitution of naphthalene can be complex, although the 1 position is the most reactive. Some examples follow.
Sometimes, small changes in the reagents and conditions change the pattern of orientation. One example is sulfonation, in which the orientation changes with reaction temperature. Another example is Friedel-Crafts acylation; in carbon disulfide the major product is the 1-isomer, whereas in nitrobenzene the major product is the 2-isomer.
Substitution usually occurs more readily at the 1 position than at the 2 position because the intermediate for 1-substitution is more stable than that for 2-substitution. The reason is that the most favorable resonance structures for either intermediate are those that have one fully aromatic ring. We can see that 1-substitution is more favorable because the positive charge can be distributed over two positions, leaving one aromatic ring unchanged. Only one resonance structure is possible for the 2-substitution intermediate that retains a benzenoid-bond arrangement for one of the rings.
1-substitution
2-substitution
Phenanthrene
The reactions of the higher hydrocarbons with electrophilic reagents are more complex than of naphthalene. For example, phenanthrene can be nitrated and sulfonated, and the products are mixtures of 1-, 2-, 3-, 4-, and 9-substituted phenanthrenes:
However, the 9,10 bond in phenanthrene is quite reactive; in fact is is almost as reactive as an alkene double bond. Addition therefore occurs fairly readily; halogenation can give both 9,10-addition and 9-substitution products by the following scheme:
Anthracene
Anthracene is even more reactive than phenanthrene and has a greater tendency to add at the 9,10 positions than to substituted. However, the addition products of nitration and halogenation readily undergo elimination to form the 9-substitution products:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/22%3A_Arenes_Electrophilic_Aromatic_Substitution/22.06%3A_Orientation_in_Disubstituted_Benzenes.txt |
Catalytic Hydrogenation
Benzenoid compounds are not readily converted to cyclohexane derivatives. Nevertheless, several addition reactions are carried out on an industrial scale. Mention was made previously of the hydrogenation of benzene to cyclohexane in the presence of a nickel catalyst:
The reaction is very important because cyclohexane is used widely as a solvent and also is oxidized to cyclohexanone, and important intermediate in the synthesis of hexanedioic (adipic) acid and azacycloheptan-2-one (caprolactam), which are used in the preparation of nylon (Section 24-3C).
Other cyclohexyl compounds are obtained by catalytic hydrogenation of the corresponding benzene derivatives. Thus cyclohexanol is obtained from benzenol, and cyclohexanamine is obtained from benzenamine (aniline):
Naphthalene can be reduced more easily than benzene. With sodium in alcohol, 1,4-dihydronaphthalene is formed. Catalytic hydrogenation gives tetralin (1,2,3,4-tetrahydronaphthalene). Further reduction to give perhydronaphthalene (decalin) can be achieved on prolonged catalytic hydrogenation at relatively high temperatures and pressures:
Phenanthrene and anthracene are reduced readily to the dihydro level by addition to the 9,10 positions. Further reduction of the terminal benzene rings is relatively difficult:
Reduction of Arenes with Metals
Catalytic hydrogenation of benzene cannot be stopped at cyclohexane or cyclohexadiene; it proceeds to cyclohexane. This is because the rate of the first addition step is much slower than of the subsequent steps:
However, benzene and its derivatives can be reduced to cyclohexadienes by solutions of metals such as $\ce{Li}$, $\ce{Na}$, $\ce{K}$, $\ce{Zn}$, and $\ce{Hg}$ in a weakly acidic solvent, such as liquid ammonia, amines, or ether-alcohol mixtures. This general type of reaction is known as the Birch reduction after the Australian chemist, A. J. Birch. With benzene, reduction with metals leads to 1,4-cyclohexadiene:
Subsequent steps include a sequence of proton- and electron-transfer steps as follows:
Substituent effects observed for this reaction are entirely consisent with those described for electrophilic substitution and addition - only reversed. That is, the reactivity of an arene in metal reductions is increased by electron-withdrawing groups and decreased by electron-donating groups. Substituents that can stabilize the anion-radical intermediate facilitate the reduction.
Reduction with metals in weakly acidic solvents is not restricted to arenes. A useful related reaction reduces alkynes to trans-alkenes, and provides a useful alternative to catalytic hydrogenation, which favors formation of cis-alkenes (Section 11-2A):
Halogen Addition
Benzene will add chlorine on irradiation with light to give the fully saturated hexachlorocyclohexane as a mixture of stereoisomers:
The reaction is commercially important because one of the isomers is a potent insecticide. The product is marketed as a mixture of isomers in which the active isomer $\left( \gamma \right)$ is optimally about $40\%$ by weight. It has a variety of trade names: Fortified, BHC, Lindane, Gammexane, Hexachlor.
Cycloaddition
In Chapter 13 we encountered the Diels-Aler reaction, which involves addition of a reactive alkene (dienophile) to the 1,4 positions of a conjugated diene. Neither benzene nor naphthalene reacts significantly with dienophiles on simple heating, but anthracene does react. Cycloaddition occurs between the 9,10 positions:
Reactions in which the transition state has a smaller volume than the reactants are speeded up by an increase in pressure. This is the case with naphthalene and cis-butenedioic anhydride. An $80\%$ yield of adduct is obtained at $100^\text{o}$ at 15,000 atmospheres pressure, whereas at one atmosphere and $100^\text{o}$, the yield is only $10\%$.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
22.10: Oxidation Reactions
The reagents usually employed for the oxidation of alkenes (e.g., \(\ce{CrO_3}\), \(\ce{KMnO_4}\), \(\ce{H_2O_2}\), \(\ce{OsO_4}\)) normally do not attack benzene. At high temperatures, benzene can be oxidized to cis-butenedioic (maleic) anhydride by air with a vanadium pentoxide catalyst. Naphthalene can be similarly oxidized to 1,2-benzenedioic (phthalic) anhydride:
Both anhydrides are prepared in this manner on a large scale for use in the production of ester polymers (Section 29-5A). Phthalic anhydride also is prepared by the oxidation of 1,2-dimethylbenzene:
Phthalic anhydride is used to make anthraquinone and to make esters of phthalic acid, which are used widely to plasticize polymers.
Ozonization of aromatic hydrocarbons is possible. Benzene itself gives ethanedial (glyoxal):
The double-bond character of the 9,10 bond in phenanthrene is particularly evident in ozonization. This bond is attacked preferentially, which leads to the formation of a dialdehyde when the ozonide is reduced with iodide ion:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
22.11: Sources and Uses of Aromatic Hydrocarbons
Benzene and many of its derivatives are manufactured on a large scale for use in high-octane gasolines and in the production of polymers, insecticides, detergents, dyes, and many miscellaneous chemicals. Prior to World War II, coal was the only important source of aromatic hydrocarbons, but during the war and thereafter, the demand for benzene, methylbenzene, and the dimethylbenzenes rose so sharply that other sources had to be found. Today, most of the benzene and almost all of the methylbenzene and the dimethylbenzenes produced in the United States are derived from petroleum.
Coal tar, which is a distillate obtained in the coking of coal (Section 4-2), is a source of an amazing number of aromatic compounds. Some of these are listed in Table 22-7, which includes nitrogen, oxygen, and sulfur compounds, as well as hydrocarbons. Although petroleum from some locations contains fairly substantial amounts of aromatic hydrocarbons, it is not a principal source for such compounds. Rather, aromatic compounds are synthesized from the \(\ce{C_6}\)-\(\ce{C_{10}}\) gasoline fraction from petroleum refining by a process referred to in the petroleum industry as catalytic re-forming or hydroforming. This involves heating a \(\ce{C_6}\)-\(\ce{C_{10}}\) fraction with hydrogen in the presence of a catalyst to modify the molecular structure of its components. Some amazing transformations take place, and the \(\ce{C_6}\)-\(\ce{C_7}\) alkanes can be converted to cycloalkanes, which, in turn, are converted to arenes. Benzene, methylbenzene (toluene), and the dimethylbenzenes (xylenes) are produced primarily in this way:
Table 22-7: Principal Compounds Obtained from Coal Tar\(^{a,b}\)
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/22%3A_Arenes_Electrophilic_Aromatic_Substitution/22.09%3A_Addition_Reactions_of_Arenes.txt |
Azulene
There are several compounds that possess some measure of aromatic character typical of benzene, but do not possess a benzenoid ring. Appropriately, they have $\left( 4n + 2 \right)$ $\pi$ electrons and are classified as nonbenzenoid aromatic compounds (Section 21-9). An example is azulene, which is isomeric with naphthalene and has a five- and a seven-membered ring fused through adjacent carbons:
As the name implies, it is deep blue. It is less stable than naphthalene, to which it isomerizes quantitatively on heating above $350^\text{o}$ in the absence of air:
Azulene has a significant polarity, with the five-membered ring negative and the seven-membered ring positive. The structure can be represented as a hybrid of neutral and ionic structures:
The polarization that has the five-membered ring negative and the seven-membered ring positive corresponds to ionic structures that have six (i.e., $4n + 2$) electrons in both the five- and seven-membered rings (Section 21-9B).
In keeping with its aromatic character and unsymmetrical charge distribution, azulene undergoes certain typical electrophilic substitution reactions at the 1 and 3 positions. Thus Friedel-Crafts acylation leads to a mixture of 1-ethanoylazulene and 1,3-diethanoylazulene:
Furthermore, in the presence of strong acids the 1 position is protonated to give a derivative of the relatively stable cycloheptatrienyl (tropylium) ion (Section 21-9B):
Cyclooctatetraene
1,3,5,7-Cyclooctatetraene (or simply cyclooctatetraene) is a bright-yellow, nonplanar, nonaromatic compound (Section 21-9A). Apparently the resonance energy of a planar structure is insufficient to overcome the unfavorable angle strain of a planar structure would have with its $\ce{C-C-C}$ bond angles of $135^\text{o}$. Cyclooctatetraene normally assumes a "tub" structure with alternating single and double bonds:
There is, however, nmr evidence that indicates that the tub form is in rapid equilibrium with a very small amount of the planar form at room temperature. There is about a $15$-$\text{kcal mol}^{-1}$ energy difference between the two forms. The dication, $\ce{C_8H_8^{2+}}$, and the dianion of cyclooctatetraene, $\ce{C_8H_8^{2-}}$, which have $\left( 4n + 2 \right)$ $\pi$ electrons, appear to exist in planar conformation.
Cyclooctatetraene can be prepared readily by polymerization of ethyne in the presence of nickel cyanide:
It could be manufactured on a large scale, but no large-scale commercial uses of the substance have yet been developed.
The chemistry of cyclooctatetraene is interesting and unusual. Particularly noteworthy is the way in which it undergoes addition reactions to form products that appear to be derived from the bicyclic isomer, bicyclo[4.2.0]2,4,7-octatriene, $8$. In fact, there is an electrocyclic equilibrium between cyclooctatetraene and $8$ (Section 21-10D) and, although the position of equilibrium lies far on the side of cyclooctatetraene, $8$ is more reactive and leads to the observed addition products:
Treatment of the bridged dichloride with strong bases causes elimination of hydrogen chloride and formation of chlorocyclooctatetraene:
The diverse ways in which cyclooctatetraene can react with a give reagent under different conditions is well illustrated by the variety of products obtained with mercuric ethanoate in ethanoic acid, methanol, and water:
Efforts to prepare "pentalene", a bridged analog of cyclooctatetraene, have not been very successful so far. A substance that appears to be a methylpentalene has been characterized at $-180^\text{o}$ by its spectral properties. On warming to $-105^\text{o}$ it forms a dimer.
Annulenes
Cyclooctatetraene is nonplanar. One reason is that the angle strain is severe in the planar form. Is it possible that larger-ring conjugated polyalkenes may have strainless planar structures? Models show that a strainless structure can be achieved with two or more of the double bonds only in trans configurations, and then only with a large enough ring that the "inside" hydrogens do not interfere with one another.
In discussing compounds of this type, it will be convenient to use the name [n]annulene to designate the simple conjugated cyclic polyalkenes, with n referring to the number of carbons in the ring - benzene being [6]annulene. The simplest conjugated cyclic polyolefin that could have a strainless planar ring containing trans double bonds, except for interferences between the inside hydrogens, is [10]annulene:
Inside-hydrogen interferences are likely to be of some importancee in all annulenes up to [30]annulene. Many annulenes have been synthesized by F. Sondheimer.
We have mentioned already (Section 22-3C) the large differences in nmr chemical shifts between the inside and outside hydrogens of [18]annulene - a substance which with 18 $\pi$ electrions should be aromatic by the $4n + 2$ rule. These differences are observed only at low temperatures. The proton nmr spectrum of [18]annulene at room temperature is a single resonance, which indicates that the inside $\left( \ce{H}_a \right)$ and outside $\left( \ce{H}_b \right)$ hydrogens are equilibrating rapidly. This can take place only if cis-trans interconversion occurs about the double bonds (marked $c$ and $t$):
At low temperatures, this equilibrium is slow enough that separate groups of resonances for the inside and outside hydrogens can be discerned in an nmr experiment (see Sections 9-10C and 27-2).
A theoretical prediction that has been borne out by experiment is that an annulene with $4n$ $\pi$ electrons should have a paramagnetic circulation of electrons - that is, opposite in direction to that shown in Figure 22-4 for benzene. For example, [16]annulene, which has $4n$ electrons, is not very stable and exists as a very rapidly interconverting mixture of two configurational isomers:
At very low temperatures $\left( -155^\text{o} \right)$, the proton nmr spectrum shows the inner hydrogens at $\delta 12.9$-$10.5$ and the outer hydrogens at $\delta 5.7$-$6.4$, which is in exactly the opposite order to the shifts with [18]annulene and the other known $\left[ 4n + 2 \right]$ $\pi$-electron annulenes.
The annulenes generally are not stable compounds, but the $\left[ 4n + 2 \right]$annulenes clearly show typical aromatic reactions. For instance [18]annulene has been converted to the nitro, ethanoyl, bromo, and carbaldehyde derivatives by electrophilic substitution reactions.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
22.13: Fluxional Compounds
A number of compounds are known to rearrange from one structure to an entirely equivalent structure, sometimes with extraordinary facility. Such compounds are said to be fluxional to distinguish from tautomers (which usually involve rearrangements between nonequivalent structures). Simple examples are the Cope rearrangement of 1,5-hexadiene,
and the electrocyclic rearrangement of bicyclo[5.1.0]-2,5-octadiene,
If the two methylenes of $9$ are bridged with two double-bonded carbons, we get the remarkable structure $10$, called "bullvalene", which rapidly interconverts amongst equivalent structures:
Equilibration of fluxional molecules must not be confused with resonance. In each electrocyclic reaction, the nuclei alter their positions as bond lengths and angles change. Interconversion of fluxional molecules also must not be confused with conformational changes (as in the interconversion of two equivalent chair forms of cyclohexane). In the interconversion of fluxional molecules, chemical bonds are broken and made; in conformational changes, no bonds are broken and no bonds are made.
Equilibration of fluxional molecules must not be confused with resonance.
A theoretical prediction is that the very-large-ring annulenes, even those with $\left[ 4n + 2 \right]$ $\pi$ electrons, may not have sufficient resonance energy to maintain equal bond lengths between the carbons and hence would be most stable with alternating double and single bonds. These substances then would be as Kekule thought benzene to be - a fluxional, equilibrating mixture of cyclohexatrienes! This does not mean that there would be no $\pi$-electron delocalization in fluxional cyclic polyenes; it means only that the VB structures would not be exactly equivalent and the MO model would have a $\sigma$-bond framework with alternating short and long $\ce{C-C}$ bonds. The theory suggests that if alternation in bond lengths occurs, then neither diamagnetic nor paramagnetic circulation of the $\pi$ electrons should be important. The synthesis and study of [26]- and [30]annulenes seems to bear out this prediction, in that with these substances there appears to be no ring-current effect on the proton chemical shifts.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/22%3A_Arenes_Electrophilic_Aromatic_Substitution/22.12%3A_Some_Conjugated_Cyclic_Polyenes.txt |
A wide variety of organic compounds contain nitrogen. In fact, the types of nitrogen compounds are so numerous and diverse that we shall be unable to consider them all. We shall give most attention to the chemistry of amines and amides in this and the following chapter, because these represent the two largest classes of nitrogen compounds.
• 23.1: Amines Compared with Alcohols
As you read the chapter you will realize a similarity between the chemistry of amines and the chemistry of alcohols. Primary amines (RNH2) and secondary amines (R2NH) are much weaker acids than alcohols (ROH) and form strongly basic anions.
• 23.2: Naturally Occurring Amines
A large and widespread class of naturally occurring amines is known as alkaloids. These are basic organic nitrogen compounds, mostly of plant origin. The structures of the plant alkaloids are extraordinarily complex, yet they are related to the simple amines in being weak nitrogen bases. The first investigator to isolate an alkaloid in pure form was F. W. A. Sertürner who described morphine as basic, salt-forming, and ammonia-like. The term "organic alkali" is derived the name alkaloid.
• 23.3: Types and Nomenclature of Amines
Amine bases are classified according to the number of alkyl or aryl groups attached to nitrogen. A further classification exists if the nitrogen is multiply bonded to carbon, as in imines and aromatic nitrogen compounds.
• 23.4: Physical Properties of Amines
The physical properties of amines depend in an important way on the extent of substitution at nitrogen. Thus primary amines, \(\ce{RNH_2}\), and secondary amines, \(\ce{R_2NH}\), are less volatile than hydrocarbons of similar size, weight, and shape, as the following examples show.
• 23.5: Spectroscopic Properties of Amines
A characteristic feature of the infrared spectra of primary and secondary amines is the moderately weak absorption at 3500cm−1 to 3300cm−1 , which corresponds to N−H stretching vibrations. Primary amines have two such bands in this region, whereas secondary amines generally show only one band. Absorption is shifted to lower frequencies by hydrogen bonding.
• 23.6: Stereochemistry of Amines
In ammonia and amines, the bonds to nitrogen are pyramidal with bond angles closer to the tetrahedral value of 109.5° than to the 90° value expected for the use of pure p orbitals of nitrogen in bond formation. We consider that the nitrogen in amines is formulated best with hybrid sp3-type orbitals; three of these orbitals are used in σ -bond formation while the fourth contains the nonbonding electron pair.
• 23.7: Amines as Bases
Perhaps the most characteristic property of amines is their ability to act as bases by accepting protons from a variety of acids.
• 23.8: Amines as Acids
Primary and secondary amines are very weak acids. The lithium salts of such amines can be prepared in ether solution by treatment of the amine with phenyllithium.
• 23.9: Amines as Nucleophiles
The unshared electrons on nitrogen play a key role in the reactions of amines. In fact, almost all reactions of amines at the nitrogen atom have, as a first step, the formation of a bond involving the unshared electron pair on nitrogen. A typical example is acylation, which is amide formation through the reaction of an acyl chloride, an anhydride, or an ester with an amine.
• 23.10: Amines with Nitrous Acid
Some of the most important reactions of amines are brought about by nitrous acid (HONO). The character of the products depends very much on whether the amine is primary, secondary, or tertiary. In fact, nitrous acid is a useful reagent to determine whether a particular amine is primary, secondary, or tertiary. With primary amines nitrous acid results in evolution of nitrogen gas; with secondary amines insoluble yellow liquids or solid N-nitroso compounds, R2N−N=O, separate; tertiary alkanamines
• 23.11: Oxidation of Amines
Nitrogen has a wide range of oxidation states in organic compounds.
• 23.12: Synthesis of Amines
There three main groups of synthesis routes for amines. The first group starts with a simple amine, or with ammonia, and builds up the carbon framework by alkylation or arylation reactions on nitrogen. The second group starts with compounds of the same carbon-nitrogen framework as in the desired amine, but with nitrogen in a higher oxidation state. The third group of reactions relies on the fact that amides usually can be converted to amines, either by reduction, hydrolysis, or rearrangement.
• 23.13: Protection of Amino Groups in Synthesis
We have mentioned previously that it may be difficult to ensure selective chemical reaction at one functional group when other functional groups are present in the same molecule. Amino groups are particularly susceptible to reactions with a wide variety of reagents, especially oxidizing reagents, alkylating reagents, and many carbonyl compounds. Therefore, if we wish to prevent the amino group from undergoing undesired reaction while chemical change occurs elsewhere in the molecule.
• 23.14: Carcinogenic Nitrogen Compounds
A number of arenamines are carcinogens. and the most dangerous examples are known to induce human bladder cancer. These chemicals were used widely in the chemical industry (mostly in azo dye manufacture) long before they were recognized as hazardous carcinogens.
• 23.E: Organonitrogen Compounds I- Amines (Exercises)
These are the homework exercises to accompany Chapter 23 of the Textmap for Basic Principles of Organic Chemistry (Roberts and Caserio).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
23: Organonitrogen Compounds I - Amines
As you read the chapter you will realize a similarity between the chemistry of amines and the chemistry of alcohols, which we discussed in Chapter 15. Primary amines $\left( \ce{RNH_2} \right)$ and secondary amines $\left( \ce{R_2NH} \right)$ are much weaker acids than alcohols $\left( \ce{ROH} \right)$ and form strongly basic anions:
Acids
Amines, like alcohols, have nonbonding electrons that impart basic and nucleophilic properties:
Bases
Nucleophiles
Also, amines and alcohols both can behave as carbon electrophiles under appropriate reaction conditions such that cleavage of $\ce{C-N}$ and $\ce{C-O}$ bonds occurs in the sense $\overset{\delta \oplus}{\ce{C}} \vdots \overset{\delta \ominus}{\ce{N}}$ and $\overset{\delta \oplus}{\ce{C}} \vdots \overset{\delta \ominus}{\ce{O}}$. However, because $\ce{-NH_2}$ and $\ce{-OH}$ both are poor leaving groups, each must be suitably activated to make this kind of reaction possible (see Section 8-7C). The $\ce{OH}$ group can be activated by addition of a proton or conversion to a sulfonate ester, $\ce{RO_3SR'}$, but these processes generally are ineffective for $\ce{RNH_2}$. The most effective activation for $\ce{RNH_2}$ is through conversion with nitrous acid, $\ce{HONO}$, to $\ce{R}- \overset{\oplus}{\ce{N}} \equiv \ce{N}$; then $\ce{N_2}$ is the leaving group (this reaction is described in more detail in Section 23-10A):
$\ce{R-OH} \overset{\ce{HBr}}{\longrightarrow} \ce{R-Br} + \ce{H_2O}$
There is, though a major difference in the way that amines and alcohols behave toward oxidizing agents. Amines genearlly show more complex behavior on oxidation because, as we shall see, nitrogen has a larger number of stable oxidation states than oxygen.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/23%3A_Organonitrogen_Compounds_I_-_Amines/23.01%3A_Amines_Compared_with_Alcohols.txt |
A large and widespread class of naturally occurring amines is known as alkaloids. These are basic organic nitrogen compounds, mostly of plant origin. The structures of the plant alkaloids are extraordinarily complex, yet they are related to the simple amines in being weak nitrogen bases. In fact, the first investigator to isolate an alkaloid in pure form was F. W. A. Sertürner who, in 1816, described morphine (Figure 23-1) as basic, salt-forming, and ammonia-like. He used the term "organic alkali" from which is derived the name alkaloid.
The structures of some of the better known plant alkaloids are shown in Figure 23-1. You will recognize some of them by name even if you have never seen their structures before. Many of the alkaloids are polycyclic structures and have other functional groups in addition to basic nitrogen. You will see that the nitrogens of alkaloids frequently are tertiary amine functions.
All of the alkaloids shown in Figure 23-1 are substances with very pronounced physiological action. Indeed, alkaloids in general have been used and abused for centuries as medicinals, drugs, and poisons. However, only in this century have their structures become known, and we are still a long way from understanding the chemistry that leads to their pronounced physiological effects. It is not even understood what function, if any, these compounds have in the host plant.
As you can see from Figure 23-1, alkaloids include compounds that may be classified as antimicrobial (quinine), as analgesics (morphine, codeine), as hallucinogens (mescaline, LSD), as stimulants (cocaine, atropine, caffeine), as topical anaesthetics (cocaine). With the possible exception of caffeine, all may be described as potentially poisonous enough to warrant great care in their use. Although some of these compounds are used as natural medicinals, an entire industry has developed in an effort to produce synthetic analogs with similar, but safer, medicinal properties. Some of the better known of these synthetic drugs are shown in Figure 23-2. They include a group of narcotic substances known as barbiturates, which are used widely as sedatives, anticonvulsants, and sleep-inducing drugs. Several representative nitrogen-containing tranquilizing drugs, synthetic stimulants, and antibiotics also are shown.
Basic nitrogen compounds similar to the plant alkaloids also occur in animals, although the description animal alkaloid seldom is used. Certain amines and ammonium compounds play key roles in the function of the central nervous system (Figure 23-3) and the balance of amines in the brain is critical for normal brain functioning. Also, many essential vitamins and hormones are basic nitrogen compounds. Nitrogen bases also are vital constituents of nucleic acid polymers (DNA and RNA) and of proteins (Chapter 25).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
23.03: Types and Nomenclature of Amines
Amine bases are classified according to the number of alkyl or aryl groups attached to nitrogen. This number is important in determining the chemical reactions that are possible at the nitrogen atom:
A further classification exists if the nitrogen is multiply bonded to carbon, as in imines and aromatic nitrogen compounds:
The nomenclature of amines was considered briefly in Section 7-8. We shall give only a short review here to focus on the main points. Amino compounds can be named either as derivatives of ammonia or as amino-substituted compounds:
To be consistent and logical in naming amines as substituted ammonias, they strictly should be called alkanamines and arenamines, according to the nature of the hydrocarbon grouping. Unfortunately, the term alkylamine is used very commonly in place of alkanamine, while a host of trivial names are used for arenamines. We shall try to indicate both the trivial and the systematic names where possible. Some typical amines, their names, and their physical properties are listed in Table 23-1. The completely systematic names give in Table 23-1 illustrate in a poignant way the difficulty one gets into by using completely systematic names, and why simpler but less systematic names continue to be used for common compounds. A good example is \(\ce{N}\),\(\ce{N}\)-dibutylbutamine versus tributylamine. The special ways of naming heterocyclic amines are mentioned previously (Section 15-11A).
Table 23-1: Typical Amines and Their Properties
Salts of amines with inorganic or organic acids are named as substituted ammonium salts, except when the nitrogen is part of a ring system. Examples are
\(^1\)Note the use of azonia to denote the cationic nitrogen in the ring, whereas aza is used for neutral nitrogen (see Section 15-11A).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/23%3A_Organonitrogen_Compounds_I_-_Amines/23.02%3A_Naturally_Occurring_Amines.txt |
The physical properties of amines depend in an important way on the extent of substitution at nitrogen. Thus primary amines, $\ce{RNH_2}$, and secondary amines, $\ce{R_2NH}$, are less volatile than hydrocarbons of similar size, weight, and shape, as the following examples show:
This is because the amines are associated through hydrogen bonding of the type $\ce{N-H} \cdots \colon \ce{N}$. Generally, $\ce{N-H} \cdots \colon \ce{N}$ bonds are somewhat weaker than those of the corresponding types, $\ce{O-H} \cdots \colon \ce{O}$ and $\ce{F-H} \cdots \colon \ce{F}$, because the electronegativity of nitrogen is less than that of oxygen or fluorine thereby making nitrogen a poorer hydrogen donor. Even so, association through hydrogen bonding is significant in amines of the type $\ce{RNH_2}$ or $\ce{R_2NH}$ as the boiling-point comparison shows. With tertiary amines, where $\ce{N-H} \cdots \colon \ce{N}$ bonding is not possible, the boiling points are much lower and are similar to those of hydrocarbons of similar branching and molecular weights:
The water solubilities of the lower-molecular-weight amines are appreciable, as can be seen from the solubility data in Table 23-1. In fact, amines are more water-soluble than alcohols of similar molecular weights. This is the result of hydrogen bonding, with amine molecules as the hydrogen acceptors and water molecules as the hydrogen donors:
Hydrogen bonds of this type are stronger than $\ce{O} \colon \cdots \ce{H-O-H}$ bonds.
Amines, especially those with significant volatility, have unpleasant odors. Some of them smell like ammonia, others smell fishy, while others are indescribably revolting. The alkanediamines of structure $\ce{H_2N(CH_3)}_n \ce{NH_2}$ are notably wretched and two are aptly called putrescine $\left( n = 4 \right)$ and cadaverine $\left( n = 5 \right)$. As you may guess from the names, these compounds are among the amines produced by bacterial decay of organic animal matter (putrefaction of protein) and are poisonous components (ptomaines) thereof.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
23.05: Spectroscopic Properties of Amines
Infrared and Ultraviolet Spectra
A characteristic feature of the infrared spectra of primary and secondary amines is the moderately weak absorption at $3500 \: \text{cm}^{-1}$ to $3300 \: \text{cm}^{-1}$, which corresponds to $\ce{N-H}$ stretching vibrations. Primary amines have two such bands in this region, whereas secondary amines generally show only one band. Absorption is shifted to lower frequencies by hydrogen bonding, but because $\ce{NH} \cdots \colon \ce{N}$ bonding is weaker than $\ce{OH} \cdots \colon \ce{O}$ bonding, the shift is not as great and the bands are not as intense as are the absorption bands of hydrogen-bonded $\ce{O-H}$ groups (see Table 9-2). Bands corresponding to $\ce{N-H}$ bending vibrations are observed around $1600 \: \text{cm}^{-1}$. Absorptions corresponding to $\ce{C-N}$ vibrations are less easily identifiable, except in the case of arenamines, which absorb fairly strongly near $1300 \: \text{cm}^{-1}$. Spectra that illustrate these effects are shown in Figure 23-4.
The ultraviolet absorptions of simple saturated amines occur at rather short wavelengths $\left( \sim 220 \: \text{nm} \right)$ and are not particularly useful for identification. These are $n \rightarrow \sigma^*$ transitions that correspond to the antibonding $\sigma$ orbital of a $\ce{C-N}$ bond.
NMR Spectra
The proton nmr spectra of amines show characteristic absorptions for $\ce{H-C-N}$ protons around $2.7 \: \text{ppm}$. The positions of the resonances of $\ce{N-H}$ protons show considerable variability as the result of differences in degree of hydrogen bonding (Section 9-10E). Sometimes the $\ce{N-H}$ resonance has nearly the same chemical shift as the resonances of $\ce{CH_3-C}$ protons (as with $\ce{N}$-ethylethanamine, Figure 23-5).
A further complication associated with $\ce{N-H}$ and $\ce{H-C-N}$ resonances is their variable chemical shift and line width in the presence of acidic substances because of a chemical exchange process of the type illustrated in Equation 23-1:
Depending on the rate at which the proton transfers of Equation 23-1 occur and the concentrations of the reactants, the chemical shift of the $\ce{N-H}$ proton will come between that of pure $\ce{(CH_3)_2NH}$ and pure $\ce{HA}$. Except at high acid concentrations, this exchange eliminates any observable coupling between the $\ce{N-H}$ proton and the $\ce{N}$-methyl protons $\left( \ce{H-C-N-H} \right)$; see Section 9-10I.
In Section 9-10L, we discussed $\ce{^{13}C}$ nmr and its many applications to structural problems. The nmr of $\ce{^{15}N}$ nuclei has similar possibilities but, because $\ce{^{15}N}$ is only $0.37\%$ of natural nitrogen and has an even smaller nuclear magnetic moment than $\ce{^{13}C}$, it is very difficult to detect $\ce{^{15}N}$ resonances at the natural abundance level.$^2$ Indeed, natural $\ce{^{15}N}$ has to be observed for about a $6 \times 10^{10}$ longer time than protons to achieve the same signal-to-noise ratio! Despite this difficulty, natural-abundance $\ce{^{15}N}$ spectra can be obtained for many compounds (even enzymes) and, in some cases, provide very useful chemical information (see Figure 24-4).
Mass Spectra of Amines
The most prominent cleavage of the parent molecular ion $\ce{M^+}$ derived from amines occurs at the $\ce{C}_\beta$-$\ce{C}_\alpha$ bond to give an imminium ion which, for ethanamine, has $m/e = 30$:
It is helpful in identifying the molecular ion of an organonitrogen compound to remember that the $m/e$ value of $\ce{M^+}$ will be an uneven number if the ion contains one or another odd number of nitrogen atoms. Thus ethanamine, $\ce{C_2H_7N}$, gives an $\ce{M^+}$ of $m/e = 45$. For all other elemental compositions of $\ce{C}$, $\ce{H}$, $\ce{O}$, or with an even number of nitrogens, the molecular ion will have an even $m/e$ value.
The cleavage reaction of Equation 23-2 reveals other useful generalizations. Whatever its source, a parent molecular ion, $\ce{M^+}$, has one unpaired electron and is properly described as an odd-electron ion (a radical cation). When a parent molecular ion fragments, it does so homolytically, as shown in Equation 23-2, and produces a radical and an ion in which the electrons are paired - an even-electron ion. The $m/e$ value of an even-electron ion is an even number for any elemental composition of $\ce{C}$, $\ce{H}$, $\ce{O}$ in combination with an odd number of nitrogens. These generalizations are summarized in Table 23-2 and can be useful in the interpretation of mass spectra.
$^2$The abundant nitrogen nucleus, $\ce{^{14}N}$, has a magnetic moment but generally gives very poor nmr spectra with very broad lines. The reason is the $\ce{^{14}N}$ usually "relaxes" rapidly, which means that its nuclear magnetic states have short lifetimes (see Section 27-1).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
23.06: Stereochemistry of Amines
In ammonia and amines, the bonds to nitrogen are pyramidal with bond angles closer to the tetrahedral value of $109.5^\text{o}$ than to the $90^\text{o}$ value expected for the use of pure $p$ orbitals of nitrogen in bond formation. We consider that the nitrogen in amines is formulated best with hybrid $sp^3$-type orbitals; three of these orbitals are used in $\sigma$-bond formation while the fourth contains the nonbonding electron pair:
A consequence of the pyramidal configuration at nitrogen is that, when the attached groups $\ce{R}_1$, $\ce{R}_2$, and $\ce{R}_3$ are nonidentical, the nitrogen becomes a chiral atom. Under these circumstances, we would expect two enantiomeric configurations:
The resolution of an acyclic chiral amine into its separate enantiomers has not been achieved yet, and it appears that the enantiomers are very rapidly interconverted by an inversion process involving a planar transition state:
With ammonia, inversion of this type occurs about $4 \times 10^{10}$ times per second at room temperature, which corresponds to the planar state being less stable than the pyramidal state by about $6 \: \text{kcal mol}^{-1}$. With aliphatic tertiary amines, the inversion rate is more on the order of $10^3$ to $10^5$ times per second. Such rates of inversion are much to great to permit resolution of an amine into its enantiomers by presently available techniques.
When the amine nitrogen is incorporated in a small ring, as in azacyclopropanes, $1$, the rate of inversion at nitrogen is markedly slower than in open-chain amines. In fact, with some oxazacyclopropanes, such as $2$, inversion does not occur rapidly at ordinary temperatures, which means that the configuration at the nitrogen persists long enough for resolution into enantiomers to be possible:
The stereochemistry of azacyclohexanes is complicated by the fact that there is a conformational change in the ring as well as inversion at the pyramidal nitrogen. Therefore it is difficult to say whether the axial-equatorial equilibrium of, for example, 1-methylazacyclohexane is achieved by ring inversion, or by nitrogen inversion, or both:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/23%3A_Organonitrogen_Compounds_I_-_Amines/23.04%3A_Physical_Properties_of_Amines.txt |
Standard Expressions of Base Strength
Perhaps the most characteristic property of amines is their ability to act as bases by accepting protons from a variety of acids:
$\ce{RNH_2} + \ce{HA} \rightleftharpoons \ce{R} \overset{\oplus}{\ce{N}} \ce{H_3} + \overset{\ominus}{\ce{A}}$
When the reference acid, $\ce{HA}$ is water, we can set up a scale of base strengths from the equilibrium constant, $K_b$, measured for the proton-transfer reaction shown in Equation 23-3:
$\ce{RNH_2} + \ce{H_2O} \overset{K_b}{\rightleftharpoons} \ce{R} \overset{\oplus}{\ce{N}} \ce{H_3} + \overset{\ominus}{\ce{O}} \ce{H} \tag{23-3}$
In many reference works, it is customary to express the strengths of organic bases not as $K_b$ but as the acid-dissociation constants, $K_a$ (or p$K_a$'s) for the corresponding conjugate acids. These $K_a$ values are then the acid constants of the corresponding ammonium ions in aqueous solution (Equation 23-4):
$\ce{R} \overset{\oplus}{\ce{N}} \ce{H_3} + \ce{H_2O} \overset{K_a}{\rightleftharpoons} \ce{RNH_2} + \overset{\oplus}{\ce{H_2O}} \tag{23-4}$
With this convention, the stronger the base, $\ce{RNH_2}$, the more the equilibrium in Equation 23-4 will lie to the left, and the smaller will be $K_a$. The relationship between $K_a$ and $K_b$ in water solution is
$K_a \times K_b = 10^{-14}$
and in terms of p$K$ values, because by definition p$K = -\text{log} K$,
p$K_a$ $+$ p$K_b$ $= 14$
Base Strengths of Alkanamines and Cycloalkanamines
The base strengths of simple alkanamines usually are around $K_b = 10^{-4}$ $\left( K_a = 10^{-10} \right)$ in water solution, and vary within perhaps a factor of 10 from ammonia to primary, secondary, and tertiary amines, as can be seen from the data in Table 23-1. Cyclohexanamine has about the same base strength as methanamine, whereas the effect on the basic nitrogen of being in a saturated ring, as in azacyclohexane, increases the base strength somewhat.
The trends that are evident, especially from basicities of amines measured in the gas phase, point to increasing basicity with the number and size of alkyl groups on the nitrogen atom.
Order of basicity (gas phase): $\ce{(CH_3)_3N} > \ce{(CH_3)_2NH} > \ce{CH_3NH_2} > \ce{NH_3}$
This is reasonable because the conjugate acids, $\ce{R_3} \overset{\oplus}{\ce{N}} \ce{H}$, are likely to be stabilized by electron-donating and polarizable alkyl groups, thereby making $\ce{R_3N}$ a stronger base. That the same trend is not evident in aqueous solution again shows the influence of the solvent on thermochemical properties (see Section 11-8A).
Generally, substituents located on saturated groups attached to nitrogen influence base strengths through their inductive effects in the same way that these substituents influence the strengths of carboxylic acids (see Section 18-2).
Base Strengths of Arenamines
Alkenamines, or enamines, $\ce{R-CH=CHNH_2}$, usually are not stable and rearrange readily to imines (Section 16-4C). An important exception is benzenamine (aniline), $\ce{C_6H_5NH_2}$, which has an amino group attached to a benzene ring. The imine structure is less favorable by virtue of the considerable stabilization energy of the aromatic ring:
From the heat of combustion of bezenamine we know that it has a $3 \: \text{kcal mol}^{-1}$ larger stabilization energy than benzene (Table 21-1). This difference in stabilization energies can be ascribed in either valence-bond or molecular-orbital theory to delocalization of the unshared pair of electrons on nitrogen over the benzene ring. The valence-bond structures are
The extra $3$-$\text{kcal mol}^{-1}$ stabilization energy of benzenamine can be accounted for in terms of the structures $4a$ to $4c$.
Benzenamine is only 1/1,000,000 as strong a base as cyclohexanamine. Most, if not all, of the difference can be accounted for by the decrease in stabilization when the unshared electron pair of nitrogen is localized in forming an $\ce{N-H}$ bond. Hence, benzenamine is stabilized more in the un-ionized state by electron delocalization, relative to cyclohexanamine, than in the ionized state, as expressed by the following equilibrium which lies far to the right:
According to the valence-bond structures, $4a$, $4b$, and $4c$, benzenamine has some degree of double-bond character between the nitrogen and the ring, and some degree of negative charge at the ortho and para positions. Accordingly, the ability of the amine nitrogen to add a proton should be particularly sensitive to the electrical effects produced by the presence of substituent groups on the aromatic ring. For example, carbonyl, nitro, cyano, and ethoxycarbonyl substituents, which can delocalize an electron pair on an adjacent carbon (see Sections 17-1A, 17-3E, and 18-8B), are expected to reduce the base strength of the amine nitrogen when substituted in the ortho or para positions. The reason is that stabilization by the substituent, as shown by structure $5$ for 4-nitrobenzenamine, is important for the free base and not for the conjugate acid, $6$:
It is simpler and common practice to discuss substituent effects on base strength in terms of the dissociation equilibria of the conjugate acids, $\ce{ArNH_3^+} + \ce{H_2O} \rightleftharpoons \ce{ArNH_2} + \ce{H_3O^+}$. Substituents that can stabilize the free base by electron delocalization or induction, as in $5$, will tend to increase the acid dissociation of $\ce{ArNH_3^+}$ (decrease base strength of $\ce{ArNH_2}$). We see this in the data of Table 23-3 for electron-withdrawing groups ($\ce{NO_2}$, $\ce{CN}$, $\ce{CF_3}$, $\ce{CH_3CO}-$), which increase acid strengths, and for electron-donating groups ($\ce{CH_3}$, $\ce{NH_2}$), which decrease acid strengths. The effect is most pronounced when the groups are at the ortho or para (2 or 4) positions.
Table 23-3: Strengths of Conjugate Acids of Monosubstituted Benzenamines in Aqueous Solution at $25^\text{o}$
Unsaturated Amines. Azarenes
Substantial differences in base strength are found between alkanamines and unsaturated amines that have the group $\ce{-C=N}-$. An example is azabenzene (pyridine, $\ce{C_5H_5N}$), which is a nitrogen analog of benzene:
Azabenzene is quite a weak base - in fact, it is 1/100,00 as strong a base as typical alkanamines. This low basicity can be ascribed to the hybridization of the nitrogen orbitals $\left( sp^2 \right)$ in azabenzene. As we indicated in Section 11-8B in connection with $\ce{C-H}$ acidity, the more $s$ character in the $\ce{C-H}$ bonding orbital, the higher the acidity. The same arguments hold for $\ce{N-H}$ bonds in the conjugate acids, $\ce{-C=} \overset{\oplus}{\ce{N}} \ce{H}-$, as the following data show:
Other examples include:
It is incorrect to assume that the basicity of unsaturated nitrogen in a $\ce{C=N}-$ group is always low. Consider, for example, the base strength of 2,2-diaminoazaethene (guanidine):
This substance is the strongest electrically neutral organonitrogen base known. The basic nitrogen is the imino $\left( sp^2 \right)$ nitrogen, which on protonation forms a particularly stable conjugate acid in which the three $\ce{NH_2}$ groups become identical because of electron delocalization:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
23.08: Amines as Acids
Primary and secondary amines are very weak acids. The lithium salts of such amines can be prepared in ether solution by treatment of the amine with phenyllithium:
The lithium salt of $\ce{N}$-ethylethanamine (diethylamine) is called lithium diethylamide,$^4$ but this nomenclature can lead to confusion with compounds of the type $\ce{RCO_2NH_2}$, which are derived from carboxylic acids and also are called amides. We choose to avoid using the name "alkali amide" for $\ce{RN} \overset{\ominus}{\ce{H}} \overset{\oplus}{\ce{Li}}$ and accordingly will refer to them as metal salts of the parent amine.
Alkanamines have acid strengths corresponding to $K_a$ values of about $10^{-33}$, which means that their conjugate bases are powerfully basic reagents. Therefore they are very effective in causing elimination reactions by the $E2$ mechanism (Section 8-8) and aromatic substitution by the aryne mechanism (Section 14-6C). The following example illustrates this property in a useful synthesis of a benzenamine from bromobenzene:
Salts of alkanamines also are useful for generating enolate salts of carbonyl compounds (Sections 17-4A and 18-8C).
$^4$The system used here names these salts as substitution products of $\ce{NH_2^-}$. Clearly, to give $\ce{LiN(C_2H_5)_2}$ the name "lithium $\ce{N}$-ethylethanamide" would be totally incorrect because $\ce{N}$-ethylethanamide is $\ce{CH_3CONHC_2H_5}$. Perhaps a better name would be lithium diethylazanide or $\ce{N}$,$\ce{N}$-diethylaminolithium.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/23%3A_Organonitrogen_Compounds_I_-_Amines/23.07%3A_Amines_as_Bases.txt |
Acylation of Amines. Synthesis of Amides
The unshared electrons on nitrogen play a key role in the reactions of amines. In fact, almost all reactions of amines at the nitrogen atom have, as a first step, the formation of a bond involving the unshared electron pair on nitrogen. A typical example is acylation, which is amide formation through the reaction of an acyl chloride, an anhydride, or an ester with an amine. The initial step in these reactions with benzenecarbonyl derivatives and methanamine as illustrative reactants is as follows:
The reaction is completed by loss of a proton and elimination of $\ce{X}^\ominus$:
The reaction is called acylation because an acyl group, $\ce{RCO}-$, is transferred to the amine nitrogen. It will be seen that these reactions are very similar to the formation of esters by acylating agents, whereby the acyl group is transferred to the oxygen of an alcohol (Section 15-4D):
A serious disadvantage to the preparation of amides through the reaction of an amine with an acyl chloride (or anhydride) is the formation of one mole of amine salt for each mole of amide:
This is especially serious if the amine is the expensive ingredient in the reaction. In such circumstances, the reaction usually is carried on in a two-phase system wit the acyl chloride and amine in the nonaqueous phase and sodium hydroxide in the aqueous phase. As the amine salt is formed and dissolves in the water, it is converted back to amine by the sodium hydroxide and extracted back into the nonaqueous phase:
$\ce{R} \overset{\oplus}{\ce{N}} \ce{H_3} \overset{\ominus}{\ce{Cl}} + \overset{\oplus}{\ce{Na}} \overset{\ominus}{\ce{OH}} \rightarrow \ce{RNH_2} + \overset{\oplus}{\ce{Na}} \overset{\ominus}{\ce{Cl}} + \ce{H_2O}$
This procedure requires an excess of acid chloride because some of it is wasted by hydrolysis.
Imine and Enamine Formation
Amines also add to the carbonyl carbon of aldehydes and ketones, but the reactions take a different course from acylation and, with ammonia or a primary amine, yield imines, $\ce{-C=N-R}$, as previously discussed in Section 16-4C.
Imines formed from ammonia and aldehydes $\left( \ce{RCH=NH} \right)$ are very unstable and readily polymerize (Section 16-4C). However, substitution of an alkyl or aryl group on the nitrogen increases the stability, and $\ce{N}$-substituted imines, $\ce{-C=N-R}$, are familiarly known as Schiff bases. They are key intermediates in a number of synthetic and biological reactions (see, for example, 17-3F) and are capable of rearrangement by reversible proton transfer that, in some respects, resembles the rearrangement of ketones to enols:
Secondary amines cannot form imines with aldehydes and ketones but may react instead to form enamines, $\ce{-C=C-NR_2}$. The formation and synthetic uses of these compounds were discussed previously (Sections 16-4C, 17-4B, and 18-9D).
Sulfonamide Formation from Amines
We have seen that amines react with acyl chlorides to give amides. A very similar reaction occurs with sulfonyl chlorides to give sulfonamides. An example is benzenesulfonyl chloride reacting with methanamine to give $\ce{N}$-methylbenzenesulfonamide:
Sulfonylation of amines can be a useful way of differentiating (chemically) between primary, secondary, and tertiary amines by what is known as the Hinsberg test. Primary and secondary amines both react with a sulfonyl chloride, but only the sulfonamide from the primary amines has an $\ce{N-H}$ hydrogen. The sulfonyl group makes this hydrogen relatively acidic and the sulfonamide therefore dissolves readily in sodium hydroxide solutions. The secondary amine does not give a base-soluble amide, whereas the tertiary amine gives no sulfonamide:
Sulfonamides have medicinal value as antibacterial agents. In fact, 4-aminobenzenesulfonamide was the first synthetic antibacterial drug in clinical use, and is effective against a large number of bacterial infections:
This substance inhibits the growth of bacteria by interfering with the synthesis of folic acid, $7$, which is an essential substance for bacteria and animals alike. However, animals acquire folic acid from a normal diet, whereas bacteria have to synthesize it. Biosynthesis of folic acid is blocked by 4-aminobenzenesulfonamide, probably because of the structural similarity of the sulfonamide to 4-aminobenzoic acid, which is a normal ingredient in the biosynthesis of folic acid. The enzyme system involved apparently substitutes the sulfonamide for the aminobenzoic acid and creates a sulfonamide-type folic acid instead of the carboxamide derivative (compare structures $7$ and $8$):
Some 10,000 structurally different sulfonamides have been synthesized as a result of the discovery of the antibacterial properties of sulfanilamide. The practice of synthesizing numerous structurally related compounds in an effort to find some that are more efficient or have fewer side effects than those already available is very important to the pharmaceutical industry. However, as is usually the case, of the many known sulfonamides only about thirty have the proper balance of qualities to be clinically useful.
Alkylation. Synthesis of Alkanamides
Ammonia and amines can function as nucleophiles in $S_\text{N}2$ displacement reactions of alkyl halides (Section 8-7E). Such processes provide syntheses of alkanamines only with those halides that are reactive in $S_\text{N}2$ but not $E2$ reactions. For example,
$\ce{NH_3} + \ce{CH_3-I} \overset{S_\text{N}2}{\longrightarrow} \ce{CH_3-NH_3^+} \ce{I^-} \tag{23-5}$
The product formed according to Equation 23-5 is an ammonium salt from which the parent amine can be recovered by neutralization with a strong base, such as sodium hydroxide:
$\ce{CH_3NH_3^+} \ce{I^-} + \ce{Na^+} \ce{^-OH} \rightleftharpoons \ce{CH_3NH_2} + \ce{Na^+} \ce{I^-} + \ce{H_2O} \tag{23-6}$
Acid-base equilibria similar to Equation 23-6 also occur between an ammonium salt and a neutral amine (Equation 23-7). This can have serious consequences in amine alkylations because it can lead to mixtures of products, whereby more than one alkyl group is bonded to nitrogen:
\begin{align} \ce{CH_3NH_3^+} \ce{I^-} + \ce{NH_3} &\rightleftharpoons \ce{CH_3NH_2} + \ce{NH_4^+} \ce{I^-} \ \ce{CH_3NH_2} + \ce{CH_3-I} &\rightarrow \ce{(CH_3)_2CH_2^+} \ce{I^-} \end{align} \tag{23-7}
Therefore we may expect the reaction of ammonia with methyl iodide to give four possible alkylation products, mono-, di-, and trimethylamines, as well as tetramethylammonium iodide:
$\ce{NH_3} \underset{-\ce{HI}}{\overset{\ce{CH_3I}}{\longrightarrow}} \ce{CH_3NH_2} \underset{-\ce{HI}}{\overset{\ce{CH_3I}}{\longrightarrow}} \ce{(CH_3)_2NH} \underset{-\ce{HI}}{\overset{\ce{CH_3I}}{\longrightarrow}} \ce{(CH_3)_3N} \overset{\ce{CH_3I}}{\longrightarrow} \ce{(CH_3)_4N}^\oplus \ce{I}^\ominus$
Despite the fact that alkylation reactions of amines generally give mixtures of products, they are of practical value on an industrial scale. The commercial synthesis of methanamines uses methanol as the methylating agent and aluminum oxide as an acidic catalyst; all three amines are formed, and are separated by distillation and extraction. The function of the catalyst is to make $\ce{OH}$ a better leaving group (Section 8-7D):
$\ce{CH_3OH} + \ce{NH_3} \underset{450^\text{o}, \: 200 \: \text{psi}}{\overset{\ce{Al_2O_3}}{\longrightarrow}} \ce{CH_3NH_2} + \ce{(CH_3)_2NH} + \ce{(CH_3)_3N} + \ce{H_2O}$
The tetraalkylammonium halides formed by complete alkylation of amines are ionic compounds that resemble alkali-metal salts. When silver oxide is used to precipitate the halide ion, tetraalkylammonium halides are converted to tetraalkylammonium hydroxides, which are strongly basic substances similar to sodium or potassium hydroxide:
$\ce{(CH_3)_4} \overset{\oplus}{\ce{N}} \: \overset{\ominus}{\ce{I}} \underset{-\ce{AgI}}{\overset{\ce{Ag_2O}, \: \ce{H_2O}}{\longrightarrow}} \ce{(CH_3)_4} \overset{\oplus}{\ce{N}} \: \overset{\ominus}{\ce{OH}}$
Higher-molecular-weight alkylammonium hydroxides decompose on heating to give alkenes. The reaction is a standard method for the preparation of alkenes and is known as the Hofmann elimination (see Section 8-8B):
In principle, the same problems of polyalkylation and $E2$ elimination exist with the amine anion as with the neutral amine - and as far as $E2$ goes, much more so.
There are nitrogen anions that are useful in alkylation reactions, but they are derived from carboxamides and sulfonamides rather than amines. Two examples are given here to illustrate the synthesis of a primary and a secondary amine (also see Section 18-10C):
Gabriel synthesis of primary amines
The success of the Gabriel synthesis depends on $\ce{N}$-alkylation being favored over $\ce{O}$-alkylation and $S_\text{N}2$ being favored over $E2$. Polar, aprotic solvents such as methylsufinylmethane, $\ce{(CH_3)_2SO}$, are useful for the Gabriel synthesis. Hydrolysis of the alkylation product often is difficult and "amide interchange" (analogous to ester interchange, Section 18-7A) with hydrazine can be an effective way to free the amine from the imide.
Sulfonamide synthesis of secondary amines
In this synthesis, the acidic properties of sulfonamides of the type $\ce{C_6H_5SO_2NHR}$ are utilized to form anions capable of alkylation by the $S_\text{N}2$ mechanism.
Arylation. Synthesis of Arenamines
In previous discussions (Section 14-6A) we stated that it is not possible to displace halogen from simple aryl halides such as bromobenzene by simple $S_\text{N}2$ reactions using amines or other weakly basic nucleophiles at ordinary temperatures:
However, arylation with such systems will occur with strong bases by the benzyne mechanism (Sections 14-6C and 23-8).
Arylation of amines by the direct displacement of aryl halides is possible when the halogen is activated by strong electron-withdrawing groups in the ortho and para positions. For examples, 2,4-dinitrobenzenamine can be prepared by heating 2,4-dinitrochlorobenzene with ammonia:
The reasons why this reaction proceeds are discussed in detail in Section 14-6B.
Arenamines as Nucleophiles. Electrophilic Aromatic Substitution
The nitrogen of arenamines is less basic and less nucleophilic than the nitrogen of alkanamines because of electron delocalization of the nitrogen lone pair, as shown for benzenamine in Section 23-7C. The polar valence-bond structures emphasize that the ring atoms, particularly the ortho and para positions, should be more nucleophilic than in benzene. Accordingly, the amino group strongly activates the ring toward attack by electrophiles. In fact, bromine reacts rapidly with benzenamine in aqueous solution to introduce three bromine substituents and form 2,4,6-tribromobenzenamine; no catalyst is required:
Weakly electrophilic reagents that do not normally attack benzene will attack the ring carbons of arenamines. Some of those reactions are described later in the chapter (Sections 23-10C and 23-10D).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/23%3A_Organonitrogen_Compounds_I_-_Amines/23.09%3A_Amines_as_Nucleophiles.txt |
Alkanamines with Nitrous Acid
Some of the most important reactions of amines are brought about by nitrous acid $\left( \ce{HONO} \right)$. The character of the products depends very much on whether the amine is primary, secondary, or tertiary. In fact, nitrous acid is a useful reagent to determine whether a particular amine is primary, secondary, or tertiary. With primary amines nitrous acid results in evolution of nitrogen gas; with secondary amines insoluble yellow liquids or solid $\ce{N}$-nitroso compounds, $\ce{R_2N-N=O}$, separate; tertiary alkanamines dissolve in and react with nitrous acid solutions without evolution of nitrogen, usually to give complex products:
Nitrous acid is unstable and always is prepared as needed, usually by mixing a solution of sodium nitrite, $\ce{NaNO_2}$, with a strong acid at $0^\text{o}$. These conditions provide a source of $^\oplus \ce{NO}$, which is transferred readily to the nucleophilic nitrogen of the amine:
With this common key step, why do amines react differently with nitrous acid depending on their degree of substitution? The answer can bee seen from the reactions that are most easily possible for the $\ce{-} \overset{\oplus}{\ce{N}} \ce{-NO}$ intermediate. Clearly, if there is a hydrogen on the positive nitrogen, it can be lost as a proton and a $\ce{N}$-nitrosamine formed:
With a secondary amine, the reaction stops here, with formation of $\ce{R_2N-NO}$, and because these substances are very weak bases, they are insoluble in dilute aqueous acids. They are characteristically yellow or orange-yellow solids or oils.
A tertiary amine$\cdot \overset{\oplus}{\ce{N}} \ce{O}$ complex, $\ce{R_3} \overset{\oplus}{\ce{N}} \ce{-NO}$, cannot lose a proton from nitrogen, but instead may lose a proton from carbon and go on to form complex products.
With a primary amine, the initially formed $\ce{N}$-nitrosamine can undergo a proton shift by a sequence analogous to interconversion of a ketone to an enol. The product is called a diazoic acid:
Some diazoic acids form salts that are quite stable, but the acids themselves usually decompose rapidly to diazonium ions:
Diazonium salts can be regarded as combinations of carbocations $\ce{R}^\oplus$ with $\ce{N_2}$ and, because of the considerable stability of nitrogen in the form of $\ce{N_2}$, we would expect diazonium salts to decompose readily with evolution of nitrogen and formation of carbocations. This expectation is realized, and diazonium salts normally decompose in this manner in water solution. The aliphatic diazonium ions decompose so rapidly that their presence can only be inferred from the fact that the products are typically those of reactions of carbocations:
With propanamine, loss of nitrogen from the diazonium ion gives the very poorly stabilized propyl cation, which then undergoes a variety of reactions that are consistent with the carbocation reactions discussed previously (see Sections 8-9B and 15-5E):
The isopropyl cation formed by rearrangement undergoes substitution and elimination like the propyl cation. About half of the products arise from isopropyl cations.
There is one exceptional reaction of the propyl cation that involves 1,3-elimination and formation of about $10\%$ of cyclopropane:
Clearly, the plethora of products to be expected, particularly those resulting from rearrangement, prevents the reaction of the simple primary amines with nitrous acid from having any substantial synthetic utility.
Arenamines with Nitrous Acid. Arenediazonium Salts
Unlike primary alkylamines, primary arenamines react with nitrous acid at $0^\text{o}$ to give diazonium salts that, in most cases, are stable enough to be isolated as crystalline $\ce{BF_4^-}$ salts. Other salts can be isolated, but some of these, such as benzenediazonium chloride, in the solid state may decompose with explosive violence.
The reason for the greater stability of arenediazonium salts compared with alkanediazonium salts appears to be related to the difficulty of forming aryl carbocations (Section 14-6A). Even the gain in energy associated with having nitrogen as the leaving group is not sufficient to make aryl cations form readily, although solvolysis of arenediazonium ions in water does proceed by an $S_\text{N}1$ mechanism:
This reaction has general utility for replacement of aromatic amino groups by hydroxyl groups. In contrast to the behavior of alkylamines, no rearrangements occur.
Generally, diazonium salts from arenamines are much more useful intermediates than diazonium salts from alkanamines. In fact, arenediazonium salts provide the only substances that undergo nucleophilic substitution reactions on the aromatic ring under mild conditions, without the necessity of having activating groups, such as nitro or cyano, in the ortho or para position. The most important reactions of this type include the replacement of the diazonium group by nucleophiles such as $\ce{Cl}^\ominus$, $\ce{Br}^\ominus$, $\ce{I}^\ominus$, $\ce{CN}^\ominus$, $\ce{NO_2^-}$, and these reactions lead to the formation of aryl halogen, cyano, and nitro compounds. Most of these reactions require cuprous ions, $\ce{Cu}$(I), as catalysts. The method is known as the Sandmeyer reaction. The following examples illustrate how a primary arenamine can be converted to a variety of different groups by way of its diazonium salt:
Aryl fluorides also may be prepared from arenamines by way of diazonium salts if the procedure is slightly modified. The amine is diazotized with nitrous acid in the usual way; then fluoroboric acid or a fluoroborate salt is added, which usually causes precipitation of a sparingly soluble diazonium fluoroborate. The salt is collected and thoroughly dried, then carefully heated to the decomposition point - the products being an aryl fluoride, nitrogen, and boron trifluoride:
$\ce{C_6H_5} \overset{\oplus}{\ce{N}} \: \overset{\ominus}{\ce{BF_4}} \overset{\text{heat}}{\longrightarrow} \ce{C_6H_5F} + \ce{N_2} + \ce{BF_3}$
This reaction is known as the Schiemann reaction. An example (which gives a better than usual yield) follows:
Later in the chapter we shall see that amines can be prepared by the reduction of nitro compounds, which permits the following sequence of reactions:
$\ce{ArH} \overset{\ce{HNO_3}}{\longrightarrow} \ce{ArNO_2} \overset{\left[ \ce{H} \right]}{\longrightarrow} \ce{ArNH_2} \overset{\ce{HONO}}{\longrightarrow} \ce{ArN_2^+} \underset{-\ce{N_2}}{\overset{\ce{CuX}}{\longrightarrow}} \ce{ArX}$
This sequence is especially useful to introduce groups or produce orientations of substituents that may not be possible by direct substitution.
The Sandmeyer group of reactions is an example of the production of nucleophilic substitution by way of radical intermediates (see Section 14-10A):
This mechanism is supported by the fact that $\ce{Cu}$(II) is important in the formation of $\ce{C_6H_5X}$. If the concentration of $\ce{Cu}$(II) is kept very low so as to slow down conversion of $\ce{C_6H_5} \cdot$ to $\ce{C_6H_5^+}$, and a compound with a reactive double bond is present, then products are formed by attack of $\ce{C_6H_5} \cdot$ on the double bond. This is called the Meerwein reaction:
$\ce{C_6H_5} \overset{\oplus}{\ce{N_2}} \: \overset{\ominus}{\ce{X}} + \ce{CH_2=CHCN} \overset{\ce{Cu} \left( \text{I} \right)}{\longrightarrow} \ce{C_6H_5CH_2-CHXCN} + \ce{N_2}$
Iodide ion appears to be a good enough reducing agent to form $\ce{C_6H_5} \cdot$ without the intervention of $\ce{Cu}$(I); considerable $\ce{I_2}$ usually is formed in the reaction:
\begin{align} \ce{C_6H_5} \overset{\oplus}{\ce{N_2}} + \overset{\ominus}{\ce{I}} &\rightarrow \ce{C_6H_5} \cdot + \ce{N_2} + \ce{I} \cdot \ \ce{C_6H_5} \cdot + \ce{I} \cdot &\rightarrow \ce{C_6H_5I} \ 2 \ce{I} \cdot &\rightarrow \ce{I_2} \end{align}
Secondary arenamines react with nitrous acid to form $\ce{N}$-nitroso compounds while tertiary arenamines undergo electrophilic substitution with $\ce{NO}^\oplus$ if they have an unsubstituted para position:
Diazo Coupling Reactions
Not all reactions of diazonium ions involve cleavage of the $\ce{C-N}$ bond. An important group of reactions of arenediazonium ions involves aromatic substitution by the diazonium ion acting as an electrophilic agent to yield azo compounds, $\ce{Ar-N=N-Ar}$:
This reaction is highly sensitive to the nature of the substituent $\ce{X}$, and coupling to benzene derivatives normally occurs only when $\ce{X}$ is a strongly electron-donating group such as $\ce{-O}^\ominus$, $\ce{-N(CH_3)_2}$, and $\ce{-OH}$. However, coupling with $\ce{X} = \ce{-OCH_3}$ may take place with particularly active diazonium ions.
Diazo coupling has considerable technical value, because the azo compounds that are produced are highly colored. Many are used as fabric dyes and for other coloring purposes. A typical example of diazo coupling is formation of 4-$\ce{N}$,$\ce{N}$-dimethylaminoazobenzene from benzenediazonium chloride and $\ce{N}$,$\ce{N}$-dimethylbenzenamine:
The product once was used to color edible fats and therefore was known as "Butter Yellow", but its use to color food is prohibited because it is reported to be a potent liver carcinogen for rats.
The pH used for diazo coupling of amines is very important in determining the nature of the products. Under near-neutral conditions the diazonium ion may attack the nitrogen of the arenamine rather than a ring carbon. In this event a diazoamino compound, a triazene, $\ce{-N=N-N}-$, is formed:
The reaction is readily reversed if the pH is lowered sufficiently.
As you see from this brief discussion of arenediazonium salts, their chemistry is complex. It is inappropriate to discuss all of their many reactions here, but a summary of the most important types of reactions is given in Table 23-4.
Table 23-4: Summary of Reactions of Arenediazonium Salts
Rearrangements of $\ce{N}$-Substituted Arenamines
A secondary arenamine behaves like a secondary alkanamine in reacting with nitrous acid to give an $\ce{N}$-nitrosamine. However, when treated with an acid the $\ce{N}$-nitrosamine rearranges:
This is one example of a group of formally related rearrangements in which a substituent, $\ce{Y}$, attached to the nitrogen of a benzenamine derivative migrates to the ortho or para positions of the aromatic ring under the influence of acid:
Rearrangement occurs most readily when $\ce{Y}$ is a strongly electron-attracting group and the $\ce{N-Y}$ bond that is broken is not as strong as the $\ce{C-Y}$ bond that is formed. A few of the many examples of this type of reaction follow:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/23%3A_Organonitrogen_Compounds_I_-_Amines/23.10%3A_Amines_with_Nitrous_Acid.txt |
Oxidation States of Nitrogen in Organic Compounds
Nitrogen has a wide range of oxidation states in organic compounds. We can arrive at an arbitrary scale for the oxidation of nitrogen in much the same way as we did for carbon (Section 11-1). We simply define elementary nitrogen as the zero oxidation state, and every atom bonded to nitrogen contributes -1 to the oxidation state if it is more electropositive than nitrogen (e.g., $\ce{H}$, $\ce{C}$, $\ce{Li}$, $\ce{B}$, $\ce{Mg}$) and +1 if it is more electronegative (e.g., $\ce{O}$, $\ce{F}$, $\ce{Cl}$). Doubly bonded atoms are counted twice, and a formal positive charge associated with nitrogen counts as +1.
To illustrate, the oxidation states of several representative compounds are as follows:
Several types of nitrogen compounds are listed in Table 23-5 to illustrate the range of oxidation states that are possible.
Table 23-5: Oxidation States of Nitrogen
Oxidation of Tertiary Amines. Amine Oxides
For the oxidation of a tertiary amine by reagents such as hydrogen peroxide, $\ce{H_2O_2}$, or peroxycarboxylic acids, $\ce{RCOOOH}$, which can supply an oxygen atom with six electrons, the expected product is an azane oxide (amine oxide). Thus $\ce{N}$,$\ce{N}$-diethylethanamine (triethylamine) can be oxidized to triethylazane oxide (triethylamine oxide):
Amine oxides are interesting for two reasons. First, amine oxides decompose when strongly heated, and this reaction provides a useful preparation of alkenes. With triethylazane oxide (triethylamine oxide), ethene is formed:
The second interesting point about amine oxides is that, unlike amines, they do not undergo rapid inversion at the nitrogen atom, and the oxides from amines with three different $\ce{R}$ groups are resolvable into optically active forms. This has been achieved for several amine oxides, including the one from $\ce{N}$-ethyl-$\ce{N}$-methyl-2-propanamine.
Oxidation of Primary and Secondary Alkanamines
Addition of an oxygen atom from hydrogen peroxide or a peroxyacid to a primary or secondary amine might be expected to yield an amine oxide-type intermediate, which then could rearrange to an azanol (hydroxylamine):
However, these oxidations usually take a more complicated course, because the azanols themselves are oxidized easily, and in the case of primary amines, oxidation occurs all the way to nitro compounds, in fair-to-good yields:
Oxidation of Aromatic Amines
We shall use benzenamine to illustrate some typical oxidation reactions of arenamines. The course of oxidation depends on the nature of the oxidizing agent and on the arenamine. With hydrogen peroxide or peroxycarboxylic acids, each of which functions to donate oxygen to nitrogen, oxidation to the azanol, the nitroso, or the nitro compound may occur, depending on the temperature, the pH, and the amount of oxidizing agent:
Oxidizing agents that abstract a hydrogen atom or hydride ion lead to more complex reactions, which often result in highly colored products. One of the best black dyes for fabric (Aniline Black) is produced by impregnating cloth with phenylammonium chloride solution and then oxidizing, first with sodium chlorate $\left( \ce{NaClO_3} \right)$ and finally with sodium dichromate $\left( \ce{Na_2Cr_2O_7} \right)$. Aniline Black probably is not a single substance, and its exact structure(s) is not known; but its formation certainly involves addition reactions in which carbon-nitrogen bonds are made. A possible structure is shown in which there are seven aniline units:
Oxidation of benzenamine with sodium dichromate in aqueous sulfuric acid solution produces 1,4-cyclohexadienedione (para-benzoquinone), which is the simplest member of an interesting class of conjugated cyclic diketones that will be discussed in more detail in Chapter 26:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/23%3A_Organonitrogen_Compounds_I_-_Amines/23.11%3A_Oxidation_of_Amines.txt |
Main Types of Synthesis
There are seemingly many different ways in which amines can be prepared. However, a careful look at these methods reveals that they fall into three main groups of reactions. The first group starts with a simple amine, or with ammonia, and builds up the carbon framework by alkylation or arylation reactions on nitrogen, as discussed in Section 23-9D:
$\ce{RX} + \ce{NH_3} \rightarrow \ce{RNH_2} + \ce{HX}$
The second group starts with compounds of the same carbon-nitrogen framework as in the desired amine but with nitrogen in a higher oxidation state. The amine then is obtained from these compounds by catalytic hydrogenation or metal-hydride reduction, as will be described in the next section:
$\ce{RNO_2} + 3 \ce{H_2} \overset{\ce{Pt}}{\longrightarrow} \ce{RNH_2} + 2 \ce{H_2O}$
The third group of reactions relies on the fact that amides usually can be converted to amines, either by reduction, hydrolysis, or rearrangement, so that any viable synthesis of amides usually is also a synthesis of amines:
These and related reactions are discussed in further detail in the following sections. For your convenience, a tabular summary of methods for the synthesis of amines appears in Tables 23-6 and 23-7.
Table 23-6: Practical Examples of the Synthesis of Amines
Table 23-7: Practical Examples of the Synthesis of Aromatic Amines
Formation of Amines by Reduction
Excellent procedures are available for the preparation of primary, secondary, and tertiary amines by the reduction of a variety of nitrogen compounds. Primary amines can be obtained by hydrogenation or by lithium aluminum hydride reduction of nitro compounds, azides, imines, nitriles, or unsubstituted amides [all possible with $\ce{H_2}$ over a metal catalyst ($\ce{Pt}$ or $\ce{Ni}$) or with $\ce{LiAlH_4}$]:
Some care must be exercised in the reduction of nitro compounds because such reductions can be highly exothermic. For example, the reaction of $1 \: \text{mol}$ $\left( 61 \: \text{g} \right)$ of nitromethane with hydrogen to give methanamine liberates sufficient heat to increase the temperature of a $25$-$\text{lb}$ iron bomb $100^\text{o}$:
Secondary and tertiary amines, particularly those with different $\ce{R}$ groups, are prepared easily by lithium aluminum hydride reduction of substituted amides (Section 18-7C).
Amines by Reductive Alkylation of Aldehydes and Ketones
A useful synthesis of primary and secondary amines that is related to the reductions just described utilizes the reaction of an aldehyde or a ketone with ammonia or a primary amine in the presence of hydrogen and a metal catalyst:
\begin{align} \ce{RCHO} + \ce{NH_3} + \ce{H_2} &\overset{\ce{Pt}}{\longrightarrow} \ce{RCH_2NH_2} + \ce{H_2O} \ \ce{R_2C=O} + \ce{CH_3NH_2} + \ce{H_2} &\overset{\ce{Pd}}{\longrightarrow} \ce{R_2CHNHCH_3} + \ce{H_2O} \end{align}
It is reasonable to suppose that the carbonyl compound first forms the imine derivative by way of the aminoalcohol (see Section 16-4C), and this derivative is hydrogenated under the reaction conditions:
Other reducing agents may be used, and the borohydride salt $\ce{Na}^\oplus \: ^\ominus \ce{BH_3(CN)}$ is convenient to use in place of $\ce{H_2}$ and a metal catalyst.
In a formal sense, the carbonyl compound is reduced in this reaction while the amine is alkylated, hence the term reductive alkylation or reductive amination.
Amines from Amides by Hydrolysis or Reduction
There are a number of ways in which an amide can be transformed into an amine. Two of these ways have been mentioned already and involve hydrolysis or reduction:
As a means of amine synthesis, both methods depend on the availability or the ease of synthesis of the corresponding amide.
Amines from Amides by the Hofmann Degradation
An interesting and general reaction for the preparation of primary amines is the Hofmann degradation, in which an unsubstituted amide is converted to an amine by bromine (or chlorine) in sodium hydroxide solution:
The mechanism of this unusual reaction first involves base-catalyzed bromination of the amide on nitrogen to give an $\ce{N}$-bromoamide intermediate:
There follows a base-induced elimination of $\ce{HBr}$ from nitrogen to form a "nitrene" intermediate, which is analogous to the formation of a carbene (Section 14-7B):
As you might expect from the structure of an acyl nitrene (only six electrons in the valence shell of nitrogen), it is highly unstable but can become stabilized by having the substituent group move as $\ce{R}^\ominus$ from carbon to nitrogen:$^5$
The rearrangement is stereospecific and the configuration at the migrating carbon is retained (see Section 21-10F). The rearrangement product is called an isocyanate and is a nitrogen analog of a ketene $\left( \ce{R_2C=C=O} \right)$; like ketenes, isocyanates readily add water. The products are carbamic acids, which are not very stable, especially in basic solution, and readily lose carbon dioxide to give the amine:
A practical example of this reaction is given in Table 23-6 together with examples of related reactions known as the Curtius and Schmidt rearrangements. The latter two probably also involve rearrangement of an acyl nitrene, this time formed by decomposition of an acyl azide:
$^5$There are several analogies for this kind of rearrangement that involve electron-deficient carbon (Sections 8-9B and 15-5E) and oxygen (Section 16-9E).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/23%3A_Organonitrogen_Compounds_I_-_Amines/23.12%3A_Synthesis_of_Amines.txt |
We have mentioned previously that it may be difficult to ensure selective chemical reaction at one functional group when other functional groups are present in the same molecule. Amino groups are particularly susceptible to reactions with a wide variety of reagents, especially oxidizing reagents, alkylating reagents, and many carbonyl compounds. Therefore, if we wish to prevent the amino group from undergoing undesired reaction while chemical change occurs elsewhere in the molecule, it must be suitably protected. There is more documented chemistry on methods of protecting amino groups than of any other functional group. This is because peptide synthesis has become very important and, as we shall see in Chapter 25, it is not possible to build a peptide of specific structure from its component amino acids unless the amino groups can be suitably protected. Therefore we now will consider the more useful protecting groups that are available - how they are introduced and how they are removed.
Protonation
It should be clear that the reactivity of amines normally involves some process in which a bond is made to the unshared electron pair on nitrogen. Therefore any reaction of an amine that reduces the reactivity of this electron pair should reduce the reactivity of the nitrogen atom. The simplest way to do this would be to convert the amine to an ammonium salt with an acid. Protonation amounts to protection of the amine function:
$\ce{RNH_2} + \ce{HX} \rightleftharpoons \ce{R} \overset{\oplus}{\ce{N}} \ce{H_3} \: \overset{\ominus}{\ce{X}}$
Examples are known in which amines indeed can be protected in this manner, but unless the acid concentration is very high, there will be a significant proportion of unprotected free base present. Also, many desirable reactions are not feasible in acid solution.
Alkylation
A related protection procedure is alkylation (Equations 23-8 and 23-9), which is suitable for primary and secondary amines:
$\ce{R'NH_2} + \ce{RX} \rightarrow \ce{R'NHR} + \ce{HX} \tag{23-8}$
$\ce{R'_2NH} + \ce{RX} \rightarrow \ce{R'_2NR} + \ce{HX} \tag{23-9}$
At first glance, you may not consider that such reactions achieve protection because there is an electron pair on nitrogen in the products. However, it a suitably bulky alkylating agent, $\ce{RX}$, is used the reactivity of the resulting alkylated amine can be reduced considerably by a steric effect. The most useful group of this type is the triphenylmethyl group $\ce{(C_6H_5)_3C}-$, which can be introduced on the amine nitrogen by the reaction of triphenylmethyl chloride ("trityl" chloride) with the amine in the presence of a suitable base to remove the $\ce{HCl}$ that is formed:
The triphenylmethyl group can be removed from the amine nitrogen under very mild conditions, either by catalytic hydrogenation or by hydrolysis in the presence of a weak acid:
Acylation
One useful way of reducing the basicity and nucleophilicity of an amine nitrogen is to convert it to an amide by treatment with an acid chloride or acid anhydride (Section 18-7):
\begin{align} \ce{RNH_2} + \ce{CH_3COCl} &\rightarrow \ce{RNHCOCH_3} + \ce{HCl} \ \ce{RNH_2} + \ce{(CH_3CO)_2O} &\rightarrow \ce{RNHCOCH_3} + \ce{CH_3CO_2H} \end{align}
The reduced reactivity is associated with the stabilization produced by the attached carbonyl group because of its ability to accept electrons from the nitrogen atom. This can be seen clearly in valence-bond structures $9a$ and $9b$, which show electron delocalization of the unshared pair of the amide function:
The stabilization energy (SE) of a simple amide grouping is about $18 \: \text{kcal mol}^{-1}$, and if a reaction occurs in which the amide nitrogen acts as an electron-pair donor, almost all of the electron delocalization of the amide group is lost in the transition state:
This loss in stabilization energy at the transition state makes an amide far less nucleophilic than an amine.
The most common acylating agents are the acyl chlorides and acid anhydrides of ethanoic acid and benzoic acid. The amine can be recovered from the amide by acid- or base-catalyzed hydrolysis:
Another useful protecting group for amines has the structure $\ce{R-O-CO}-$. It differs from the common acyl groups of the type $\ce{R-CO}-$ in that it has the alkoxycarbonyl structure rather than an alkylcarbonyl structure. The most used examples are:
The phenylmethoxycarbonyl (benzyloxycarbonyl) group can be introduced by way of the corresponding acyl chloride, which is prepared from phenylmethanol (benzyl alcohol) and carbonyl dichloride:
The tert-butoxycarbonyl group cannot be introduced by way of the corresponding acyl chloride because $\ce{(CH_3)_3COCOCl}$ is unstable. One of several alternative derivatives is the azide, $\ce{ROCON_3}$:
Although these protecting groups may seem bizarre, their value lies in the fact that they can be removed easily by acid-catalyzed hydrolysis under very mild conditions. The sequence of steps is shown in Equation 23-10 and involves proton transfer to the carbonyl oxygen and cleavage of the carbon-oxygen bond by an $S_\text{N}1$ process ($\ce{R} =$ tert-butyl) or $S_\text{N}2$ process ($\ce{R} =$ phenylmethyl). The product of this step is a carbamic acid. Acids of this type are unstable and readily eliminate carbon dioxide, leaving only the free amine (also see Section 23-12E):
The benzyloxycarbonyl group, but not the tert-butoxycarbonyl group, may be removed by catalytic hydrogenation. Again a carbamic acid is formed, which readily loses $\ce{CO_2}$:
Sulfonylation
A sulfonyl group, $\ce{RSO_2}-$, like an acyl group, $\ce{R-CO}-$ or $\ce{RO-CO}-$, will deactivate an attached nitrogen. Therefore amines can be protected by transformation to sulfonamides with sulfonyl chlorides (Section 23-9C):
$\ce{C_6H_5SO_2Cl} + \ce{RNH_2} \rightarrow \ce{C_6H_5SO_2NHR} + \ce{HCl}$
However, sulfonamides are much more difficult to hydrolyze back to the amine than are carboxamides. In peptide synthesis (Section 25-7C) the commonly used sulfonyl protecting groups are 4-methylbenzenesulfonyl or 4-bromobenzenesulfonyl groups. These groups can be removed as necessary from the sulfonamide by reduction with sodium metal in liquid ammonia:
$^6$This abbreviation is approved by the IUPAC-UIB Commission on Biochemical Nomenclature and is typical of the kind of "alphabet soup" that is making biochemistry almost completely unintelligible without a glossary of approved (and unapproved) abbreviations at hand at all times. We shall make minimum use of such designations. You will remember we already use $\ce{Z}$ for something else (Section 19-7).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/23%3A_Organonitrogen_Compounds_I_-_Amines/23.13%3A_Protection_of_Amino_Groups_in_Synthesis.txt |
Amines
Everyone who works with organic chemicals should be aware that a number of arenamines are carcinogens. The most dangerous examples (see Figure 23-8) are known to induce human bladder cancer. These chemicals were used widely in the chemical industry (mostly in azo dye manufacture) long before they were recognized as hazardous carcinogens. Voluntary action and appropriate legislation now controls the industrial uses of these substances, and there also are some controls for uses in research and teaching. It is important to be aware of the potential hazards of known carcinogens and to recognize that all chemicals, both organic and inorganic, should be treated with great respect if their thermodynamic and physiological properties are not known. Carcinogenic character is just one of many possible hazards.
Azo Compounds
Other nitrogen compounds besides amines are known to be carcinogenic. For example, certain azo dyes (see Figure 23-8) have been found to produce tumors in animals. This fact has caused concern for human health because, as we indicated in Section 23-10C, azo dyes are coloring agents that are used in many products. They certainly are not all carcinogenic, but the structural requirements for a compound to show this property may change completely the toxic properties of a chemical. For example, the \(\ce{N}\),\(\ce{N}\)-diethylamino analog of "butter yellow" (Figure 23-8) is apparently hazardous.
\(\ce{N}\)-Nitroso Compounds
We have seen that \(\ce{N}\)-nitroso compounds are formed from secondary amines and nitrous acid:
\[\ce{R_2NH} + \ce{HONO} \overset{\ce{H}^\oplus}{\longrightarrow} \ce{R_2N-NO} + \ce{H_2O}\]
\(\ce{N}\)-nitroso compounds also can be formed from carboxamides and nitrous acid:
Some of these nitrosoamines and nitrosoamides are known to be potent carcinogens for some animals, which is reason to suspect they also may be carcinogenic for humans. However, it is clear that there may be very marked differences in carcinogenic properties of a given compound for different animal species.
Why some of these substances have carcinogenic activity is a matter of chemical interest. Recall from Section 23-10 that nitrosation of amines usually leads to cleavage of a \(\ce{C-N}\) bond in the sense \(\ce{C} \vdots \colon \ce{N}\). The carbon fragment ultimately is transferred to some nucleophilic atom. In effect, this means that nitrosamines can function as alkylating agents and, in a biological system, the functions that probably would be alkylated are the nucleophilic sites along the polymeric protein or nucleic acid chains. It is not difficult to appreciate that alkylation of these substances may well disrupt the pattern of normal cell growth.
There is an unresolved problem related to the carcinogenic properties of nitroso compounds. You probably are aware (if you read the labels on food packages) that sodium nitrite is added to many packaged meat products. Sodium nitrite prevents the growth of harmful bacteria, thereby retarding spoilage, and it also enhances the appearance by maintaining the red look of fresh meat. There is a possibility that nitrite may have adverse effects on human health by nitrosating the amino and amide functions of proteins in the presence of acids. This possibility has to be balanced against the alternative threat to human health if the use of nitrite were discontinued, that of increased food spoilage. In any case, it seems clear that the amount of sodium nitrite actually used in most processing is in excess of that needed to retard bacterial decay.
There are many other chemicals that are active alkylating agents besides nitrosamines, and some are unquestionably carcinogenic (see Figure 23-9), whereas others apparently are not. In fact, it is a paradox that some of the most useful synthetic drugs in treating certain forms of cancer are alkylating agents. Several of these are shown in Figure 23-10. They all have two or more active centers in the molecule that enable them to form cross-links between protein or nucleic acid molecules.
It should be recognized that not all of the carcinogenic substances loosed on mankind are the result of modern technology. The most potent carcinogens known, which are lethal in test animals at levels of a few parts per billion, are mold metabolites called aflatoxins. These substances are complex non-nitrogenous, heterocyclic oxygen compounds, which often are formed by molds growing on cereal grains, peanuts, and so on. (See Figure 23-9)
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/23%3A_Organonitrogen_Compounds_I_-_Amines/23.14%3A_Carcinogenic_Nitrogen_Compounds.txt |
The properties of the simple amides are relevant to the chemistry of peptides and proteins, substances that are fundamental to all life as we know it. Indeed, the characteristics of peptides and proteins are primarily due to their polyamide structures. For this reason, it is important to know and understand the chemistry of simple amides.
• 24.1: Structural, Physical, and Spectral Characteristics of Amides
An important feature of the amide group is that it is planar - the carbon, oxygen, nitrogen, and the first atom of each of the R groups on carbon and nitrogen lie in the same plane. This coplanarity is induces a large dipoles with simple amides have dipole moments in the range 3.7-3.8 debye. As a consequence of the polarity of the amide group, the lower-molecular-weight amides are relatively high-melting and water-soluble, as compared to esters, amines, alcohols, and the like.
• 24.2: Amides as Acids and Bases
Amides with N−H bonds are weakly acidic. Nonetheless, amides clearly are far more acidic than ammonia, and this difference reflects a substantial degree of stabilization of the amide anion. However, amides still are very weak acids (about as weak as water) and, for practical purposes, are regarded as neutral compounds.
• 24.3: Synthesis of Amides
Amides generally are formed from acid chlorides, acid azides, acid anhydrides, and esters. It is not practical to prepare them directly from an amine and a carboxylic acid without strong heating or unless the reaction is coupled to a second reaction that "activates" the acid. Notice that esters of phenols are more reactive toward amines than esters of alcohols because phenols are stronger acids than alcohols.
• 24.4: Hydrolysis of Amides
Generally, amides can be hydrolyzed in either acidic or basic solution. The mechanisms are much like those of ester hydrolysis, but the reactions are very much slower, a property of great biological importance. Amide hydrolysis can be an important route to amines. Biological amide hydrolysis, as in the hydrolysis of peptides and proteins, is catalyzed by the proteolytic enzymes.
• 24.5: Nitriles
The carbon-nitrogen triple bond differs considerably from the carbon-carbon triple bond by being stronger and much more polar. The degree of polarity of the carbon-nitrogen triple bond is indicate by the high dipole moment of the simple nitriles, which corresponds to about 70% of the dipole moment expected if one of the bonds of the triple bond were fully ionic. It is not surprising that liquid nitriles have high dielectric constants compared to most organic liquids and are soluble in water.
• 24.6: Nitro Compounds
Nitro compounds are a very important class of nitrogen derivatives. The nitro group, −NO2, like the carboxylate anion, is a hybrid of two equivalent resonance structures with a hybrid structure that has a full positive charge on nitrogen and a half-negative charge on each oxygen. The polar character of the nitro group results in lower volatility of nitro compounds than ketones of about the same molecular weight.
• 24.7: Some Compounds with N-N Bonds
Among the organic nitrogen compounds having nitrogen above the oxidation level of ammonia are a wide variety of substances with N−N bonds. We shall mention only a very few of the more important of these substances: hydrazines, azo and diazo compounds, and azides.
• 24.E: Organonitrogen Compounds II- Amides, Nitriles, and Nitro Compounds (Exercises)
These are the homework exercises to accompany Chapter 24 of the Textmap for Basic Principles of Organic Chemistry (Roberts and Caserio).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
24: Organonitrogen Compounds II - Amides Nitriles and Nitro Compounds
Molecular Structure
The structural parameters of the amide group have been determined carefully and the following diagram gives a reasonable idea of the molecular dimensions:
An important feature of the group is that it is planar - the carbon, oxygen, nitrogen, and the first atom of each of the $\ce{R}$ groups on carbon and nitrogen lie in the same plane. The $\ce{C-N}$ bond distance of $1.34 \: \text{Å}$ is intermediate between the typical single bond $\ce{C-N}$ distance of $1.47 \: \text{Å}$ and the double bond $\ce{C=N}$ distance of $1.24 \: \text{Å}$. This and other evidence indicates that the amide group is a hybrid structure of the valence-bond forms $1a$ and $1b$, with stabilization energy of about $18 \: \text{kcal mol}^{-1}$:
Coplanarity is required if the dipolar structure $1b$ is to be significant. An appreciable dipole moment may be expected of amides and, in fact, simple amides have dipole moments in the range 3.7-3.8 debye. (For reference, the carbonyl group has a moment of about 2.7 debye, Section 16-1B.)
As a consequence of the polarity of the amide group, the lower-molecular-weight amides are relatively high-melting and water-soluble, as compared to esters, amines, alcohols, and the like. The few that are liquids, such as $\ce{N}$,$\ce{N}$-dimethylmethanamide and 1-methyl-1-aza-2-cyclopentanone, have excellent solvent properties for both polar and nonpolar substances. Therefore they are good solvents for displacement reactions of the $S_\text{N}$ type (Table 8-5).
Another very important consequence of amide structure is the extensive molecular association of amides through hydrogen bonding. The relatively negative oxygens act as the hydrogen acceptors while $\ce{N-H}$ hydrogens serve as the hydrogen donors:
Nomenclature
The naming of amides is summarized in Section 7-7D. The points to remember are that they generally are named either as (i) alkanamides, in which the prefix alkan(e) is determined by the longest carbon chain that includes the carbonyl group ($\ce{HCONH_2}$ is methanamide), or as (ii) substituted carboxamides, $\ce{RCONH_2}$, in which the name is completed by identifying the $\ce{R}$ substituent:
The degree of substitution on the amide nitrogen determines whether the amide is primary, $\ce{RCONH_2}$, secondary, $\ce{RCONHR}$, or tertiary, $\ce{RCONR_2}$. When the amide is secondary or tertiary, the symbol $\ce{N}$ (for nitrogen) must precede the name of each different group attached to nitrogen:
Infrared Spectra
Considerable information is available on the infrared spectra of amides. By way of example, the spectra of three typical amides with different degrees of substitution on nitrogen are shown in Figure 24-1.
A strong carbonyl absorption is evident in the spectra of all amides, although the frequency of absorption varies somewhat with the structure of the amide. Thus primary amides generally absorb near $1680 \: \text{cm}^{-1}$, whereas secondary and tertiary amides absorb at slightly lower frequencies. The $\ce{N-H}$ stretching frequencies of amides are closely similar to those of amines and show shifts of $100 \: \text{cm}^{-1}$ to $200 \: \text{cm}^{-1}$ to lower frequencies as the result of hydrogen bonding. Primary amides have two $\ce{N-H}$ bands of medium intensity near $3500 \: \text{cm}^{-1}$ and $3400 \: \text{cm}^{-1}$, whereas secondary amides, to a first approximation, have only one $\ce{N-H}$ band near $3440 \: \text{cm}^{-1}$. However, a closer look reveals that the number, position, and intensity of the $\ce{N-H}$ bands of monosubstituted amides depend on the conformation of the amide, which can be either cis or trans:
Normally, the trans configuration is more stable than the cis conformation for primary amides. However, for cyclic amides (lactams), in which the ring size is small, the configuration is exclusively cis:
NMR Spectra
The proton NMR resonances of the $\ce{N-H}$ protons of amides are different from any we have discussed so far. Generally, these will appear at room temperature as a broad singlet absorption, which may turn into a broad triplet at higher temperatures. A typical example is propanamide (Figure 24-2).
The broad $\ce{N-H}$ proton resonance is due to the special nuclear properties of $\ce{^{14}N}$, the predominant natural isotope of nitrogen. This is established beyond question by observation of the proton spectrum of an amide in which the $\ce{^{14}N}$ is replaced by the $\ce{^{15}N}$ isotope to give $\ce{RCO} \ce{^{15}N} \ce{H_2}$. In this case the proton lines are sharp. The details of the phenomena that lead to the broad resonances of the $\ce{N-H}$ protons in amides are discussed elsewhere;$^1$ for our purposes it should suffice to note that the $\ce{^{14}N}$ nucleus has much shorter lifetimes for its magnetic states than do protons, and the broad lines result from uncertainties in the lifetimes of the states associated with $\ce{^{14}N} \ce{-H}$ spin-spin coupling (Section 27-1). One should be prepared for absorptions of this character in amides and some other substances with $\ce{N-H}$ bonds that are not involved in rapid intermolecular proton exchanges, which, when sufficiently rapid, have the effect of averaging the magnetic effects of the $\ce{^{14}N}$ atoms to zero:
$\ce{R'NH_2} + \ce{R} \overset{\oplus}{\ce{N}} \ce{H_3} \rightleftharpoons \ce{R'} \overset{\oplus}{\ce{N}} \ce{H_3} + \ce{RNH_2}$
or
$\ce{R'NH_2} + \ce{R} \overset{\ominus}{\ce{N}} \ce{H} \rightleftharpoons \ce{R'} \overset{\ominus}{\ce{N}} \ce{H} + \ce{RNH_2}$
The situation here is analogous to that discussed for the splitting of the resonances of $\ce{O-H}$ protons of alcohols by protons on the $\alpha$ carbons (see Section 9-10I).
The NMR spectra of amides are revealing as to the structure of the amide group. For example, the spectrum of $\ce{N}$,$\ce{N}$-dimethylmethanamide shows two three-proton single resonances at $2.78 \: \text{ppm}$ and $2.95 \: \text{ppm}$, which means that at ordinary temperatures the two methyl groups on nitrogen are not in the same molecular environment:
This is a consequence of the double-bond character of the $\ce{C-N}$ bond expected from valence-bond structures $1a$ and $1b$, which leads to restricted rotation about this linkage. One of the methyl groups (A) has a different stereochemical relationship to the carbonyl group than the other methyl group (B). Groups A and B therefore will have different chemical shifts, provided that rotation about the $\ce{C-N}$ bond is slow. However, at $150^\text{o}$ the two three-proton lines are found to coalesce to a single six-proton line, which means that at this temperature bond rotation is rapid enough to make the methyl groups essentially indistinguishable (see Section 9-10C):
Most amides do not rotate freely about the $\ce{C-N}$ bond. The barrier to this kind of rotation is about $19 \: \text{kcal mol}^{-1}$, which is high enough for the nonequivalence of groups on nitrogen to be observable by spectral techniques, but not quite high enough to allow for actual physical separation of stable $E,Z$ configurational isomers.
$^1$J. D. Roberts, Nuclear Magnetic Resonance, Applications to Organic Chemistry, McGraw-Hill Book Co., New York, 1959, Chapter 5. Also see the NMR references at the end of Chapter 9 in this book.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/24%3A_Organonitrogen_Compounds_II_-_Amides_Nitriles_and_Nitro_Compounds/24.01%3A_Structural_Physical_and_Spectral_Characteristics_of_Amides.txt |
Acidity
Amides with $\ce{N-H}$ bonds are weakly acidic, the usual $K_a$ being about $10^{-16}$:
Nonetheless, amides clearly are far more acidic than ammonia $\left( K_a \sim 10^{-33} \right)$, and this difference reflects a substantial degree of stabilization of the amide anion. However, amides still are very weak acids (about as weak as water) and, for practical purposes, are regarded as neutral compounds.
Where there are two carbonyl groups to stabilize the amide anion, as in the 1,2-benzenedicarboximide (phthalimide) anion (Section 18-10C), the acidity increases markedly and imides can be converted to their conjugate bases with concentrated aqueous hydroxide ion. We have seen how imide salts can be used for the synthesis of primary amines. (Gabriel synthesis, Section 23-9D and Table 23-6).
Basicity
The degree of basicity of amides is very much less than that of aliphatic amines. For ethanamide, $K_b$ is about $10^{-15}$ ($K_a$ of the conjugate acid is $\sim 10$):
The proton can become attached either to nitrogen or to oxygen, and the choice between the assignments is not an easy one. Of course, nitrogen is intrinsically more basic than oxygen; but formation of the $\ce{N}$-conjugate acid would cause loss of all the amide stabilization energy. Addition to oxygen actually is favored, but amides are too weakly basic for protonation to occur to any extent in water solution.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
24.03: Synthesis of Amides
From Carboxylic Acids
Formation of amides from carboxylic acid derivatives already has been discussed in some detail (Section 23-9A):
$\tag{24-1}$
The ease of formation of amides by the reaction of Equation 24-1 depends a lot on the nature of the leaving group $\ce{X}$. The characteristics of a good leaving group were discussed in Sections 8-7C and 8-7D in connection with $S_\text{N}$ reactions, and similar considerations apply here. Some idea of the range of acid derivatives used in amide synthesis can be obtained from Table 24-1, which lists various $\ce{RCOX}$ compounds and the p$K_a$ values of $\ce{HX}$. As a reasonable rule of thumb, the stronger $\ce{HX}$ is as an acid, the better $\ce{X}$ is as a leaving group.
Table 24-1: Derivatives and Reactivity of Carboxylic Acids Commonly Used in Amide Formation
$\ce{RCOX} + \ce{H_2NR'} \rightarrow \ce{RCONHR'} + \ce{HX}$
Amides generally are formed from acid chlorides, acid azides, acid anhydrides, and esters. It is not practical to prepare them directly from an amine and a carboxylic acid without strong heating or unless the reaction is coupled to a second reaction that "activates" the acid. Notice that esters of phenols are more reactive toward amines than esters of alcohols because phenols are stronger acids than alcohols.
From Nitriles
The hydrolysis of nitriles is a satisfactory method for preparation of unsubstituted amides and is particularly convenient when hydrolysis is induced under mildly basic conditions by hydrogen peroxide:
For the preparation of amides of the type $\ce{R_3CNHCOR}$, which have a tertiary alkyl group bonded to nitrogen, the Ritter reaction of an alcohol or alkene with a nitrile or hydrogen cyanide is highly advantageous. This reaction involves formation of a carbocation by action of strong sulfuric acid on an alkene or an alcohol (Equation 24-2), combination of the carbocation with the unshared electrons on nitrogen of $\ce{RCN}$ (Equation 24-3), and then addition of water (Equation 24-4). We use here the preparation of an $\ce{N}$-tert-butylalkanamide as an example; $\ce{RC \equiv N}$ can be an alkyl cyanide such as ethanenitrile or hydrogen cyanide itself:
$\tag{24-2}$
$\tag{24-3}$
$\tag{24-4}$
This reaction also is useful for the preparation of primary amines by hydrolysis of the amide. It is one of the relatively few practical methods for synthesizing amines with a tertiary alkyl group on the nitrogen:
The Beckmann Rearrangement of Oximes
You may recall that ketones react with $\ce{RNH_2}$ compounds to give products with a double bond to nitrogen, $\ce{-C=NR}$ (Section 16-4C). When the $\ce{RNH_2}$ compound is azanol (hydroxylamine), $\ce{HO-NH_2}$, the product is called a ketoxime, or oxime:
Oximes rearrange when heated with a strong acid, and this reaction provides a useful synthesis of amides:
This intriguing reaction is known as the Beckmann rearrangement. It has been the subject of a number of mechanistic studies that have shown the acid or acid halide ($\ce{PCl_3}$, $\ce{C_6H_5SO_2Cl}$) makes the hydroxyl group on nitrogen into a better leaving group by forming $\ce{-OH_2^+}$ or ester intermediates:
Thereafter, a rearrangement occurs resembling the reactions of carbocations (Sections 8-9B and 15-5E). When the cleavage of the $\ce{N-O}$ bond occurs, the nitrogen atom would be left with only six valence electrons. However, as the bond breaks, a substituent $\ce{R}$ on the neighboring carbon moves with its bonding electron pair to the developing positive nitrogen (Equation 24-5):
$\tag{24-5}$
Oximes with $\ce{R}$ and $\ce{R'}$ as different groups exist as $E$ and $Z$ isomers (Section 19-7) and you will notice in Equation 24-5 that the group that migrates is the one that is trans to the leaving group. To some extent the Beckmann rearrangement is an internal $S_\text{N}2$ reaction with inversion at the nitrogen. Section 21-10F gives a theoretical treatment of this kind of reaction. The rearrangement product is a nitrilium ion, as in the Ritter reaction (Section 24-3B), which adds water to form the amide.
The synthesis of aza-2-cycloheptanone ($\varepsilon$-caprolactam) by the Beckmann rearrangement of the oxime of cyclohexanone is of commercial importance because the lactam is an intermediate in the synthesis of a type of nylon (a polyamide called "nylon-6"$^2$):
$^2$The number 6 specifies the number of carbons in each monomer unit comprising the polyamide structure. By this code, nylon-6,6 is $\ce{(-NH(CH_2)_6NHCO(CH_2)_4CO-)}_n$.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/24%3A_Organonitrogen_Compounds_II_-_Amides_Nitriles_and_Nitro_Compounds/24.02%3A_Amides_as_Acids_and_Bases.txt |
Generally, amides can be hydrolyzed in either acidic or basic solution. The mechanisms are much like those of ester hydrolysis (Section 18-7A), but the reactions are very much slower, a property of great biological importance (which we will discuss later):
As we have indicated in Section 23-12, amide hydrolysis can be an important route to amines. Hydrolysis under acidic conditions requires strong acids such as sulfuric or hydrochloric, and temperatures of about $100^\text{o}$ for several hours. The mechanism involves protonation of the amide on oxygen followed by attack of water on the carbonyl carbon. The tetrahedral intermediate formed dissociates ultimately to the carboxylic acid and the ammonium salt:
In alkaline hydrolysis the amide is heated with boiling aqueous sodium or potassium hydroxide. The nucleophilic hydroxide ion adds to the carbonyl carbon to form a tetrahedral intermediate, which, with the help of the aqueous solvent, expels the nitrogen as the free amine:
Biological amide hydrolysis, as in the hydrolysis of peptides and proteins, is catalyzed by the proteolytic enzymes. These reactions will be discussed in Chapter 25.
An indirect method of hydrolyzing some amides utilizes nitrous acid. Primary amides are converted easily to carboxylic acids by treatment with nitrous acid. These reactions are very similar to that which occurs between a primary amine and nitrous acid (Section 23-10):
Secondary amides give $\ce{N}$-nitroso compounds with nitrous acid, whereas tertiary amides do not react:
A brief summary of important amide reactions follows:
Of the many other types of organonitrogen compounds known, the more important include
Although it is impractical to discuss all of these compounds in detail, we now will discuss briefly several that have not been given much attention heretofore.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
24.05: Nitriles
The carbon-nitrogen triple bond differs considerably from the carbon-carbon triple bond by being stronger ($212 \: \text{kcal mol}^{-1}$ vs. $200 \: \text{kcal mol}^{-1}$) and much more polar. The degree of polarity of the carbon-nitrogen triple bond is indicate by the high dipole moment $\left( 4.0 \: \text{D} \right)$ of the simple nitriles $\left( \ce{RCN} \right)$, which corresponds to about $70\%$ of the dipole moment expected if one of the bonds of the triple bond were fully ionic. With this knowledge it is not surprising that liquid nitriles have rather high dielectric constants compared to most organic liquids and are reasonably soluble in water. Ethanenitrile, $\ce{CH_3CN}$, is in fact a good solvent for both polar and nonpolar solutes (Table 8-5).
Nitriles absorb with variable strength in the infrared in the region $2000 \: \text{cm}^{-1}$ to $2300 \: \text{cm}^{-1}$, due to stretching vibrations of the carbon-nitrogen triple bond.
The preparation of nitriles by $S_\text{N}2$ reactions between alkyl halides and cyanide ion has been mentioned previously (Section 8-7F) and this is the method of choice when the halide is available and reacts satisfactorily. Activated aryl or azaaryl halides similarly give nitriles with cyanide ion:
Another practical route to arenecarbonitriles involves the replacement of the diazonium group, $\ce{-} \overset{\oplus}{\ce{N}} \ce{\equiv N}$, in arenediazonium ions with cuprous cyanide (Section 23-10B). Other useful syntheses involve cyanohydrin formation (Section 16-4A) and Michael addition to conjugated alkenones (Section 17-5B).
Nitriles also can be obtained by the dehydration of the corresponding amide or aldoxime. This is a widely used synthetic method and numerous dehydrating agents have been found to be effective:
The reactions of nitriles include reduction to amines and hydrolysis to acids. Both reactions have been discussed previously (Sections 18-7C and 18-7A).
Hydrogens on the alpha carbons of nitriles are about as acidic as the hydrogens alpha to carbonyl groups; accordingly, it is possible to alkylate the $\alpha$ positions of nitriles through successive treatments with a strong base and with an alkyl halide as in the following example:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/24%3A_Organonitrogen_Compounds_II_-_Amides_Nitriles_and_Nitro_Compounds/24.04%3A_Hydrolysis_of_Amides.txt |
Physical and Spectroscopic Properties
Nitro compounds are a very important class of nitrogen derivatives. The nitro group, $\ce{-NO_2}$, like the carboxylate anion, is a hybrid of two equivalent resonance structures:
The hybrid structure has a full positive charge on nitrogen and a half-negative charge on each oxygen. This is in accord with the high dipole moments of nitro compounds, which fall between $3.5 \: \text{D}$ and $4.0 \: \text{D}$, depending upon the nature of $\ce{R}$. The polar character of the nitro group results in lower volatility of nitro compounds than ketones of about the same molecular weight; thus the boiling point of nitromethane (MW 61) is $101^\text{o}$, whereas 2-propanone (MW 58) has a boiling point of $56^\text{o}$. Surprisingly, the water solubility is low; a saturated solution of nitromethane in water is less than $10\%$ by weight, whereas 2-propanone is completely miscible with water.
Nitro groups of nitroalkanes can be identified by strong infrared bands at about $1550 \: \text{cm}^{-1}$ and $1375 \: \text{cm}^{-1}$, whereas the corresponding bands in the spectra of aromatic nitro compounds occur at slightly lower frequencies. A weak $n \rightarrow \pi^*$ transition occurs in the electronic spectra of nitroalkanes at around $270 \: \text{nm}$; aromatic nitro compounds, such as nitrobenzene, have extended conjugation and absorb at longer wavelengths $\left( \sim 330 \: \text{nm} \right)$.
Preparation of Nitro Compounds
Nitro compounds can be prepared in a number of ways, including the direct substitution of hydrocarbons with nitric acid,
$\ce{RH} + \ce{HONO_2} \rightarrow \ce{RNO_2} + \ce{H_2O}$
by displacement reactions with nitrite ions,
$\ce{RX} + \ce{NO_2^-} \rightarrow \ce{RNO_2} + \ce{X^-}$
and by oxidation of primary amines,
$\ce{RNH_2} \overset{\left[ \ce{O} \right]}{\longrightarrow} \ce{RNO_2}$
Nitration of alkanes is successful only when conducted at high temperatures in the vapor phase. Mixtures of products are invariably obtained (Section 4-6):
In contrast, direct nitration of aromatic compounds such as benzene takes place readily in the liquid phase, as discussed in Section 22-4C.
Like other electrophilic substitutions, nitration of a substituted benzene, where the substituent is electron withdrawing ($\ce{NO_2}$, $\ce{CO_2H}$, $\ce{CN}$, and so on; Table 22-6), generally produces the 1,3-isomer. To prepare the 1,4-isomer, less direct routes are necessary - the usual stratagem being to use benzene derivatives with substituent groups that produce the desired orientation on nitration and then to make the necessary modifications in these groups to produce the final product. Thus 1,4-dinitrobenzene cannot be obtained by nitration of nitrobenzene but can be prepared from benzenamine by the sequence shown in Figure 24-5. Benzenamine is converted to $\ce{N}$-phenylethanamide (acetanilide) which on nitration yields the 1,4-isomer. Hydrolysis of the amide to 4-nitrobenzenamine and replacement of amino by nitro, using nitrite ion in the presence of cuprous salts, gives 1,4-dinitrobenzene (Section 23-10B). Alternatively, the amino group of 4-nitrobenzenamine can be oxidized to a nitro group by trifluoroperoxyacetic acid. In these syntheses, $\ce{N}$-phenylethanamide is nitrated in preference to benzenamine itself because, not only is benzenamine easily oxidized by nitric acid, but the nitration reaction leads to extensive 3-substitution as the result of formation of phenylammonium ion. Another route to 4-nitrobenzenamine is to nitrate chlorobenzene and subsequently replace the chlorine by reaction with ammonia. The nitrations mentioned give mixtures of 2- and 4-isomers, but these usually are easy to separate by distillation or crystallization. The same approach can be used to synthesize 4-nitrobenzoic acid. The methyl group of methylbenzene directs nitration preferentially to the 4 position, and subsequent oxidation with chromic acid yields 4-nitrobenzoic acid:
In some cases it may be necessary to have an activating group to facilitate substitution, which otherwise would be very difficult. The preparation of 1,3,5-trinitrobenzene provides a good example; direct substitution of 1,3-dinitrobenzene requires long heating with nitric acid in fuming sulfuric acid. However, methylbenzene is converted more readily to the trinitro derivative and this substance, on oxidation and decarboxylation (Section 18-4), yields 1,3,5-trinitrobenzene:
Acylamino groups also are useful activating groups and have the advantage that the amino groups obtained after hydrolysis of the acyl function can be removed from an aromatic ring by reduction of the corresponding diazonium salt with hypophosphorous acid, preferably in the presence of copper (I) ions. An example is the preparation of 1-methyl-3-nitrobenzene from $\ce{N}$-(4-methyl-phenyl)ethanamide (aceto-para-toluidide):
Routes to aliphatic nitro compounds include the reaction of an alkyl halide (of good $S_\text{N}2$ reactivity) with nitrite ion. Suitable solvents are methylsulfinylmethane [dimethyl sulfoxide, $\ce{(CH_3)_2SO}$] and dimethyl methanamide (dimethylformamide). As will be seen from Equation 24-6, formation of the nitrite ester by $\ce{O}$- instead of $\ce{N}$-alkylation is a competing reaction:
Silver nitrite sometimes is used in preference to sodium nitrite, usually in diethyl ether as solvent:
Nitromethane can be prepared conveniently by the reaction
$\ce{ClCH_2CO_2H} + \ce{NaNO_2} \underset{-\ce{NaCl}}{\overset{\text{heat}}{\longrightarrow}} \left[ \ce{O_2NCH_2CO_2H} \right] \underset{\text{heat}}{\overset{-\ce{CO_2}}{\longrightarrow}} \ce{O_2NCH_3}$
Displacement reactions with nitrite ion do not work well with aryl halides. However, displacement of the diazonium group is a practical route to nitroarenes (the Sandmeyer reaction), as described in Section 23-10B:
$\ce{ArNH_2} \overset{\ce{HONO}}{\longrightarrow} \ce{Ar} \overset{\oplus}{\ce{N_2}} \overset{\ce{CuNO_2}}{\longrightarrow} \ce{ArNO_2}$
Reactions of Nitro Compounds
Nitro compounds are quite unstable in the thermodynamic sense; for example, the heat of decomposition of nitromethane, according to the following stoichiometry, is $67.4 \: \text{kcal mol}^{-1}$.
$\ce{CH_3NO_2} \rightarrow \frac{1}{2} \ce{N_2} + \ce{CO_2} + \frac{3}{2} \ce{H_2} \: \: \: \: \: \Delta H^0 = -67.4 \: \text{kcal mol}^{-1}$
Advantage is taken of the considerable energies and rapid rates of reactions such as this in the commercial use of nitro compounds as explosives. With some nitro compounds, such as TNT, there is a further advantage of low shock sensitivity.
TNT is not detonated easily by simple impact and even burns without exploding. However, once detonation starts, decomposition is propagated rapidly. The characteristics of reasonable handling stability and high thermodynamic potential make nitro compounds particularly useful. Other polynitro compounds that are useful as explosives include PETN (Section 17-3C), cyclonite (Section 16-4C), picric acid, and tetryl:
An important characteristic of aromatic polynitro compounds is their ability to form "charge-transfer" complexes with aromatic hydrocarbons, especially those that are substituted with alkyl groups. Complexes of 2,4,6-trinitrobenzenol (picric acid) and aromatic hydrocarbons often are crystalline solids, which are useful for the separation, purification, and identification of aromatic hydrocarbons. These substances are called "hydrocarbon picrates", but the name is misleading because they are not actually salts. Furthermore, similar complexes are formed between aromatic hydrocarbons and trinitrobenzene, which demonstrates that the nitro groups rather than the hydroxyl group are essential to complex formation. The binding in these complexes resembles that in the $\pi$ complexes of halogens with alkenes and benzene (Sections 22-4D and 10-3C) and results from attractive forces between electron-rich and electron-poor substances. The descriptive name - charge-transfer complex - suggests that the complex has VB structures involving transfer of an electron from the donor (electron-rich) molecule to the acceptor (electron-poor) molecule. The name $\pi$ complex also is used because, usually at least, one component of the complex has a $\pi$-electron system. Charge-transfer or $\pi$ complexes between polynitro compounds and aromatic hydrocarbons appear to give sandwich-type structures with the aromatic rings in parallel planes, although not necessarily centered exactly over one another:
Charge-transfer complexes are almost always more highly colored than their individual components. A spectacular example is benzene and tetracyanoethene, each of which separately is colorless, but which give a bright-orange complex when mixed. A shift toward longer wavelengths of absorption, relative to their components, is to be expected for charge-transfer complexes because of the enhanced possibility for stabilization of the excited state through electron delocalization involving both components.
Reduction of nitro compounds occurs readily with a variety of reducing agents and such reductions afford a particularly useful synthesis of aromatic amines (Section 23-12B):
The reduction of a nitro compound to an amine requires six equivalents of reducing agent:
$\ce{R-NO_2} + 6 \ce{H}^\oplus + 6 \ce{e}^\ominus \rightarrow \ce{RNH_2} + 2 \ce{H_2O}$
One would not expect such a reduction to occur in a single step. Indeed, reduction is stepwise and proceeds through a string of intermediates, which, with strong reducing agents in acid solution, have at most a transient existence. The intermediates formed successively from $\ce{RNO_2}$ by increments of two equivalents of reducing agent are nitroso compounds, $\ce{R-N=O}$, and $\ce{N}$-substituted azanols (hydroxylamines), $\ce{RNHOH}$:
$\ce{RNO_2} \underset{-\ce{H_2O}}{\overset{2 \left[ \ce{H} \right]}{\longrightarrow}} \ce{RN=O} \overset{2 \left[ \ce{H} \right]}{\longrightarrow} \ce{RNHOH} \underset{-\ce{H_2O}}{\overset{2 \left[ \ce{H} \right]}{\longrightarrow}} \ce{RNH_2}$
Thus $\ce{N}$-aryl-substituted azanols can be obtained directly from the corresponding nitro compounds with zinc and ammonium chloride solution. However, zinc and hydrochloric acid gives the amine:
The difference between these reactions is in the reduction rates associated with the acidity of the solution. Ammonium chloride is a much weaker acid than $\ce{HCl}$; the pH of ammonium chloride solutions is around 6.
Oxidation of the $\ce{N}$-arylazanols under controlled conditions yields nitroso compounds. This reaction is not unlike the oxidation of alcohols to ketones (Section 15-6B):
Reduction of aryl nitro compounds with less-powerful reducing agents, especially in alkaline media, gives what may appear to be a mysterious conglomerate of bimolecular reduction products. For example, with nitrobenzene,
All of these substances can be reduced to benzenamine with tin and hydrochloric acid. As a result, each could be, but not necessarily is, an intermediate in the reduction of nitro compounds to amines. Formation of the bimolecular reduction products is the result of base-induced reactions between nitroso compounds and azanols or amines and possibly further reduction of the initially produced substances.
Several polynitrobenzene derivatives have important herbicidal uses. Examples are $\ce{N}^3$, $\ce{N}^3$-diethyl-6-trifluoromethyl-2,4-dinitro-1,3-benzenediamine and $\ce{N}$,$\ce{N}$-dipropyl-4-trifluoromethyl-2,6-dinitrobenzenamine:
These substances when mixed with soil kill weed seedlings but not crop plants such as cotton, soybeans, and peanuts. The activity is high; normally only about $0.08 \: \text{g m}^{-2}$ is required for good weed control.
The most important reactions of nitroalkanes are those involving the $\alpha$ hydrogens of the primary and secondary compounds. For example, nitromethane is sufficiently acidic to dissolve in aqueous hydroxide solutions. The anion so produced has an electronic structure analogous to the nitrate anion:
An interesting property of this ion is that when solutions of it are acidified, an unstable, rather strongly acidic isomer of nitromethane (called the aci form) is produced, which slowly reverts to the more stable nitro form:
Similar changes take place in the acidification of the enol salt of a carbonyl compound, the principal difference being the much longer life of the aci-nitro compound compared to that of an enol of a simple ketone (see Section 17-1B).
Primary and secondary nitro compounds undergo aldol additions and Michael additions with suitable carbonyl compounds and basic catalysts:
Unfortunately, alkylation reactions analogous to the base-catalyzed alkylation of carbonyl compounds generally are not useful for the synthesis of higher nitro compounds, because $\ce{C}$-alkylation of the conjugate bases of primary nitro compounds is slower than $\ce{O}$-alkylation.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/24%3A_Organonitrogen_Compounds_II_-_Amides_Nitriles_and_Nitro_Compounds/24.06%3A_Nitro_Compounds.txt |
Among the organic nitrogen compounds having nitrogen above the oxidation level of ammonia are a wide variety of substances with $\ce{N-N}$ bonds. We shall mention only a very few of the more important of these substances: hydrazines, azo and diazo compounds, and azides.
Hydrazines
Organic hydrazines or diazanes are substitution products of $\ce{NH_2-NH_2}$ and have many properties similar to those of amines in being basic and forming acyl derivatives as well as undergoing alkylation and condensations with carbonyl compounds (Section 16-4C). Unsymmetrical hydrazines can be prepared by careful reduction of $\ce{N}$-nitrosamines. 1,1-Dimethyldiazane is prepared in this way for use as a rocket fuel:
Aromatic hydrazines are best prepared by reduction of aromatic diazonium salts (Table 23-4).
Hydrazines of the type $\ce{R-NH-NH-R}$ are easily oxidized to the corresponding azo compounds, $\ce{R-N=N-R}$. With nitrous acid, monosubstituted hydrazines are converted to azides:
$\ce{R-NH-NH_2} + \ce{HONO} \rightarrow \ce{R-N=} \overset{\oplus}{\ce{N}} \ce{=} \overset{\ominus}{\ce{N}} + 2 \ce{H_2O}$
Azo Compounds
Azo or diazo compounds possess the $\ce{-N=N}-$ grouping. Aliphatic azo compounds of the type $\ce{R-N=N-H}$ appear to be highly unstable and decompose to $\ce{R-H}$ and nitrogen. Derivatives of the type $\ce{R-N=N-R}$ are much more stable and can be prepared as mentioned above by oxidation of the corresponding hydrazines. Aromatic azo compounds are available in considerable profusion from diazo coupling reactions (Section 23-10C) and are of commercial importance as dyes and coloring materials.
A prime characteristic of azo compounds is their tendency to decompose into organic free radicals and liberate nitrogen:
$\ce{R-N=N-R} \rightarrow 2 \ce{R} \cdot + \ce{N_2}$
The ease of these reactions is usually a fairly reliable guide to the stabilities of the free radicals that result. For instance, it is found that dimethyldiazene (azomethane, $\ce{CH_3N=NCH_3}$) is stable to about $400^\text{o}$, and diphenyldiazene (azobenzene, $\ce{C_6H_5N=NC_6H_5}$) also is resistant to thermal decomposition; but, when the azo compound decomposes to radicals that have extra stability because of delocalization of the odd electron, the decomposition temperature is greatly reduced. Thus the azo compound, $2$, decomposes to radicals at moderate temperatures ($60^\text{o}$ to $100^\text{o}$), and for this reason is a very useful agent for generating radicals, such as those required for the initiation of polymerization of ethenyl compounds:
Diazo Compounds
The parent of the diazo compounds, diazomethane, $\ce{CH_2=} \overset{\oplus}{\ce{N}} \ce{=} \overset{\ominus}{\ce{N}}$, has been mentioned before in connection with ylide reactions for ring enlargement (Section 16-4A) and the preparation of methyl esters from acids (Table 18-7). It is one of the most versatile and useful reagents in organic chemistry, despite the fact that it is highly toxic, dangerously explosive, and cannot be stored without decomposition.
Diazomethane is an intensely yellow gas, bp $-23^\text{o}$, which customarily is prepared and used in diethyl ether or dichloromethane solution. It can be synthesized in a number of ways, the most useful of which employs the action of base on an $\ce{N}$-nitroso-$\ce{N}$-methylamide:
As a methylating agent of reasonably acidic substances, diazomethane has nearly ideal properties. It can be used in organic solvents; reacts very rapidly without need for a catalyst (except with alcohols, which do require an acid catalyst); the coproduct is nitrogen which offers no separation problem; it gives essentially quantitative yields; and it act as as its own indicator to show when reaction is complete. With enols, it gives $\ce{O}$-alkylation:
Besides being a methylating agent, diazomethane also is a source of $\colon \ce{CH_3}$ when irradiated with light. The carbene formed in this way is highly reactive and even will react with the electrons of a carbon-hydrogen bond to "insert" the carbon of the carbene between carbon and hydrogen. This transforms $\ce{-C-H}$ to $\ce{-C-CH_3}$:
This $\colon \ce{CH_3}$ species is one of the most reactive reagents known in organic chemistry. Diazomethane undergoes a wealth of other unusual reactions. Besides those already mentioned are the following two examples:
Arndt-Eistert synthesis ($\ce{COCl} \rightarrow \ce{-CH_2CO_2H}$, Section 16-4A)
Pyrazoline formation ([2 + 3] cycloaddition)
The Arndt-Eistert synthesis is useful for converting an acid to the next higher member of the series. Pyrazolines are important intermediates for the preparation of cyclopropanes:
Diazomethane originally was believed to possess the three-membered 1,2-diazacyclopropene ring structure, but this concept was disproved by electron-diffraction studies, which showed the linear structure to be correct:
Recently, a variety of authentic 1,2-diazacyclopropenes (sometimes called diazirines) have been prepared, and these have been found to have very different properties from the diazoalkanes. The simple 1,2-diazacyclopropenes are colorless and do not react with dilute acids, bases, or even bromine. The syntheses of these substances are relatively simple. One of several possible routes follows:
Azides
Organic azides can be prepared from hydrazines and nitrous acid (Section 24-7A) and by the reaction of sodium azide with acyl halides or with alkyl halides having good $S_\text{N}2$ reactivity:
$\ce{RBr} + \ce{N_3^-} \underset{\ce{CH_3OH}}{\overset{S_\text{N}2}{\longrightarrow}} \ce{R-N=} \overset{+}{\ce{N}} \ce{=} \overset{-}{\ce{N}} + \ce{Br^-}$
The lower-molecular-weight organic azides often are unpredictably explosive and are best handled in solution.
The use of acyl azides in the preparation of amines by the Curtius rearrangement has been discussed previously (Section 23-12E). Alkyl azides can be reduced readily by lithium aluminum hydride to amines and, if a pure primary amine is desired, the sequence halide $\rightarrow$ azide $\rightarrow$ amine may give as good or better results than does the Gabriel synthesis (Section 23-9D).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/24%3A_Organonitrogen_Compounds_II_-_Amides_Nitriles_and_Nitro_Compounds/24.07%3A_Some_Compounds_with_N-N_Bonds.txt |
The chemistry of life is largely the chemistry of polyfunctional organic compounds. The functional groups usually are of types that interact rather strongly as, for example, the hydroxyl and carbonyl functions of carbohydrates (Chapter 20). The interaction between amino and carboxyl functions of amino acids figures greatly in the present chapter. We will approach the very important chemistry of amino acids and their derivatives in three stages. First, simple a-amino acids will be considered with emphasis on how the properties of amine functions and of acid functions are modified in molecules that possess both groups. Then we shall discuss some important properties of peptides and proteins, which are substances made up of amino acids linked together by amide bonds. Attention also will be given to the chemical problems presented by enzymes, which are protein molecules able to act as efficient catalysts for specific chemical reactions, and to the role of nucleic acids in protein synthesis.
• 25.1: Types of Biologically Important Amino Acids
The amino acids that occur naturally as constituents of proteins have an amino group (NH2)(NH2) and a carboxylic acid group (CO2H)(CO2H) attached to the same carbon. They are called α -amino acids and differ only in the nature of the R group on the α carbon and, with few exceptions, they are chiral molecules with the L configuration at the chiral α carbon.
• 25.2: Acid-Base Properties of $\alpha$-Amino Acids
The behavior of glycine is reasonably typical of that of the simplest amino acids. Because glycine is neither a strong acid nor a strong base, we shall expect a solution of glycine in water to contain four species in rapid equilibrium. The proportions of these species are expected to change with pH, the cationic conjugate acid being the predominant form at low pH and the anionic conjugate base being favored at high pH
• 25.3: Physical and Spectroscopic Properties
The α-amino acids crystallize as the dipolar forms and the strong intermolecular electrical forces in the crystals lead to higher melting points than those of simple amines or monocarboxylic acids. The melting points are so high that decomposition often occurs on melting. The solubility characteristics of amino acids in water are complex because of the acid-dissociation equilibria involved, but they are least soluble at their isoelectric points.
• 25.4: Analysis of Amino Acids
In many kinds of research it is important to have simple and sensitive means for analysis of amino acids, particularly in small quantities.
• 25.5: Reactions of Amino Acids
To some degree the reactions of amino acids are typical of isolated carboxylic acid and amine functions. Thus the carboxyl function can be esterified with an excess of an alcohol under acidic conditions, and the amine function can be acylated with acid chlorides or anhydrides under basic conditions. The amine function of α-amino acids and esters reacts with nitrous acid in a manner similar to that described for primary amines.
• 25.6: Synthesis of α-Amino Acids
Many of the types of reactions that are useful for the preparation of amino acids have been discussed previously in connection with separate syntheses of carboxylic acids and amino compounds. Examples include the SN2 displacement of halogen from α-halo acids by ammonia, and the Strecker synthesis, which, in its first step, bears a close relationship to cyanohydrin formation
• 25.7: Peptides and Proteins
Amino acids are the building blocks of the polyamide structures of peptides and proteins. Each amino acid is linked to another by an amide (or peptide) bond formed between the amine group of one and the acid group of the other. In this manner a polymeric structure of repeating amide links is built into a chain or ring. The amide groups are planar and configuration about the C−NC−N bond is usually, but not always, trans.
• 25.8: Structure and Function of Proteins
The biological functions of proteins are extremely diverse. The primary structure and presence or absence of special functional groups, metals, and so on, are of paramount importance is making proteins, made from a common basis set of amino acids, so remarkably heterogeneous and exhibit such varied yet specific functions.
• 25.9: Enzymes
Virtually all biochemical reactions are catalyzed by proteins called enzymes. The catalytic power and specificity of enzymes is extraordinarily high. The reactions that they catalyze are generally enhanced in rate many orders of magnitude, often as much as 10 million, over the nonenzymatic process. Consequently enzymatic reactions may occur under much milder conditions than comparable laboratory reactions.
• 25.10: Coenzymes
Many enzymes only operate in combination with organic molecules that are actually reagents for the reaction. These substances are called coenzymes or cofactors. Some coenzymes function with more than one enzyme and are involved in reactions with a number of different substrates.
• 25.11: Enzyme Regulation
You may have wondered how the proteolytic enzymes such as trypsin, pepsin, chymotrypsin, carboxypeptidase, and others keep from self-destructing by catalyzing their own hydrolysis or by hydrolyzing each other. An interesting feature of the digestive enzymes is that they are produced in an inactive form in the stomach or the pancreas - presumably to protect the different kinds of proteolytic enzymes from attacking each other or other proteins.
• 25.12: Enzyme Technology
Because enzymes function nearly to perfection in living systems, there is great interest in how they might be harnessed to carry on desired reactions of practical value outside of living systems. The potential value in the use of enzymes (separate from the organisms that synthesize them) is undeniable, but how to realize this potential is another matter.
• 25.13: Biosynthesis of Proteins
One of the most interesting and basic problems connected with the synthesis of proteins in living cells is how the component amino acids are induced to link together in the sequences that are specific for each type of protein. There also is the related problem of how the information as to the amino-acid sequences is perpetuated in each new generation of cells. We now know that the substances responsible for genetic control in plants and animals are present in and originate from the chromosomes.
• 25.14: Chemical Evolution
A problem of great interest to those curious about the evolution of life concerns the origins of biological molecules. When and how were the molecules of life, such as proteins, nucleic acids, and polysaccharides, first synthesized? In the course of geological history, there must have been a prebiotic period when organic compounds were formed and converted to complex molecules similar to those we encounter in living systems.
• 25.E: Amino Acids, Peptides, and Proteins (Exercises)
These are the homework exercises to accompany Chapter 25 of the Textmap for Basic Principles of Organic Chemistry (Roberts and Caserio).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
25: Amino Acids Peptides and Proteins
Protein Amino Acids
The amino acids that occur naturally as constituents of proteins have an amino group $\left( \ce{NH_2} \right)$ and a carboxylic acid group $\left( \ce{CO_2H} \right)$ attached to the same carbon. They are called $\alpha$-amino acids and have the general formula
They differ only in the nature of the $\ce{R}$ group on the $\alpha$ carbon and, with few exceptions, they are chiral molecules with the $L$ configuration at the chiral $\alpha$ carbon:$^1$
The structures and names of some particularly important $\alpha$-amino acids are shown in Table 25-1. You will notice that the names in common use for amino acids are not descriptive of their structural formulas; but at least they have the advantage of being shorter than the systematic names. The abbreviations Gly, Glu, and so on, that are listed in Table 25-1 are particularly useful in designating the sequences of amino acids in proteins and peptides, as will become evident later in the chapter.
Table 25-1: Amino Acids Important as Constituents of Proteins
The nature of the substituent $\ce{R}$ varies considerably. In some amino acids, $\ce{R}$ is a hydrocarbon group, whereas in others it possesses functional groups such as $\ce{OH}$, $\ce{SH}$, $\ce{SCH_3}$, $\ce{CO_2H}$, or $\ce{NH_2}$. Amino acids that have amine or other basic functions in the $\ce{R}$ group are called basic amino acids (lysine and arginine), whereas those with acidic groups are called acidic amino acids (aspartic and glutamic acids). Three of the amino acids listed in Table 25-1 (cysteine, cystine, and methionine) contain sulfur in $\ce{-SH}$, $\ce{-S-S}-$, and $\ce{-SCH_3}$ groups. Cysteine and cystine can be interconverted readily with a wide variety of oxidizing and reducing agents according to the general reaction $2 \ce{RSH} \underset{\left[ \ce{H} \right]}{\overset{\left[ \ce{O} \right]}{\rightleftharpoons}} \ce{RSSR}$. This is an important process in the biochemistry of sulfur-containing peptides and proteins (Section 25-8A).
The $\alpha$-amino function of the common amino acids is primary $\ce{-NH_2}$ in all except proline and hydroxyproline. Several of the amino acids have aromatic $\ce{R}$ groups (phenylalanine, tyrosine, tryptophan), while histidine and tryptophan have azarene $\ce{R}$ groups.
Nonprotein Amino Acids
The most abundant amino acids are those that are protein constituents and these are always $\alpha$-amino acids. However, there are many other amino acids that occur naturally in living systems that are not constituents of proteins, and are not $\alpha$-amino acids. Many of these are rare, but others are common and play important roles in cellular metabolism. For example, 3-aminopropanoic acid is a precursor in the biosynthesis of the vitamin, pantothenic acid,$^2$
and 4-aminobutanoic acid is involved in the transmission of nerve impulses.
Homocysteine$^3$ and homoserine are among the important $\alpha$-amino acids that are not constituents of proteins. These substances are precursors in the biosynthesis of methionine.
$^1$A number of $D$-amino acids have been found to be constituents of peptides in the cell walls of bacteria.
$^2$Pantothenic acid is in turn a precursor for the synthesis of coenzyme A, which is essential for the biosynthesis of fats and lipids (Sections 18-8F and 30-5A).
$^3$The prefix homo implies an additional carbon in the longest chain.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/25%3A_Amino_Acids_Peptides_and_Proteins/25.01%3A_Types_of_Biologically_Important_Amino_Acids.txt |
The behavior of glycine is reasonably typical of that of the simplest amino acids. Because glycine is neither a strong acid nor a strong base, we shall expect a solution of glycine in water to contain four species in rapid equilibrium. The proportions of these species are expected to change with pH, the cationic conjugate acid being the predominant form at low pH and the anionic conjugate base being favored at high pH:
Spectroscopic measurements show that the equilibrium between neutral glycine and the dipolar ion favors the dipolar ion by at least 100 to 1. This is to be expected because the $\ce{H_3} \overset{\oplus}{\ce{N}}-$ group of the dipolar ion will stabilize the $\ce{-CO_2^-}$ end while the $\ce{-CO_2^-}$ group will stabilize the $\ce{H_3} \overset{\oplus}{\ce{N}}-$ end.
The acid-ionization constant of $\ce{H_3} \overset{\oplus}{\ce{N}} \ce{CH_2CO_2H}$ is $4.5 \times 10^{-3}$ (p$K_a$ $= 2.34$, Equation 25-1), which is about 25 times greater than $K_a$ for ethanoic acid (Section 18-2). This is expected because of the electron-attracting electrostatic effect of the $\ce{H_3} \overset{\oplus}{\ce{N}}-$ group. Ionization of the $\ce{H_3} \overset{\oplus}{\ce{N}}-$ group of the dipolar ion ($K_a = 2.0 \times 10^{-10}$; p$K_a = 9.60$; Equation 25-2) is oppositely affected by the electrostatic effect of the $\ce{-CO_2^-}$ group and is 10 times less than of ethanammonium ion (Section 23-7B). The manner in which the concentrations of the charged glycine species change with pH is shown in Figure 25-1. Notice that, between pH 3 and pH 8, almost all of the glycine is in the form of the dipolar ion. The pH at the center of this range, where the concentration of $\ce{H_3} \overset{\oplus}{\ce{N}} \ce{CH_2CO_2H}$ is equal to the concentration of $\ce{H_2NCH_2CO_2^-}$, is called the isoelectric point, pI, and usually corresponds to the pH at which the amino acid has minimum water solubility. Isoelectric points for the amino acids are shown in Table 25-1. The isoelectric points are the average of the p$K_a$ values for dissociation of the monocation and the dipolar ion forms of the amino acid. For glycine, pI $= \left( 2.34 + 9.60 \right) /2$.
$\text{p} K_a = \text{pH} + \text{log}_{10} \frac{\left[ \ce{H_3} \overset{\oplus}{\ce{N}} \ce{CH_2CO_2H} \right]}{\left[ \ce{H_3} \overset{\oplus}{\ce{N}} \ce{CH_2CO_2^-} \right]} = 2.34 \tag{25-1}$
$\text{p} K'_a = \text{pH} + \text{log}_{10} \frac{\left[ \ce{H_3} \overset{\oplus}{\ce{N}} \ce{CH_2CO_2^-} \right]}{\left[ \ce{H_3NCH_2CO_2^-} \right]} = 9.60 \tag{25-2}$
The pH behavior of amino acids with either acidic or basic functional groups attached to the side chains is more complicated than of simple amino acids. For example, there are three acid dissociations starting with the diconjugate acid of lysine:
The p$K_a$ values for the side-chain functions of acidic and basic amino acids are given in Table 25-1.
We already mentioned how the $\ce{H_3} \overset{\oplus}{\ce{N}}-$ group of the conjugate acid of glycine enhances the acid strength of the carboxyl group compared to ethanoic acid and how the $\ce{-CO_2^-}$ group reduces the acidity of the $\ce{H_3} \overset{\oplus}{\ce{N}}-$ group of the dipolar ion relative to ethanammonium ion. These effects will be smaller the farther away the charged group is from the ionizable group. As a result, one would predict that the carboxyl groups of aspartic acid would have different p$K_a$ values, and indeed this is so:
Similarly, the side-chain ammonium group of lysine is less acidic than that of the ammonium group close to the carboxyl group:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
25.03: Physical and Spectroscopic Properties
The $\alpha$-amino acids crystallize as the dipolar forms, $\ce{H_3} \overset{\oplus}{\ce{N}} \ce{-CHR-C} \overset{\ominus}{\ce{O_2}}$, and the strong intermolecular electrical forces in the crystals lead to higher melting points than those of simple amines or monocarboxylic acids (see Table 25-1). The melting points are so high that decomposition often occurs on melting. The solubility characteristics of amino acids in water are complex because of the acid-dissociation equilibria involved, but they are least soluble at their isoelectric points. The dipolar structures of amino acids greatly reduce their solubility in nonpolar organic solvents compared to simple amines and carboxylic acids.
The infrared spectra of $\alpha$-amino acids in the solid state or in solution do not show a carbonyl absorption band at $1720 \: \text{cm}^{-1}$ characteristic of a carboxyl group. Rather, they show a strong absorption near $1600 \: \text{cm}^{-1}$ typical of the carboxylate anion. the $\ce{N-H}$ stretch appears as a strong, broad band between $3100$-$2600 \: \text{cm}^{-1}$:
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/25%3A_Amino_Acids_Peptides_and_Proteins/25.02%3A_Acid-Base_Properties_of_%28alpha%29-Amino_Acids.txt |
The Ninhydrin and Related Tests
In many kinds of research it is important to have simple and sensitive means for analysis of amino acids, particularly in small quantities. Detection of amino acids can be achieved readily by the "ninhydrin color test", whereby an alcoholic solution of the triketone, "ninhydrin", is heated with an amino acid and produces an intense blue-violet color. The sensitivity and reliability of this test is such that 0.1 micromole of amino acid gives a color intensity reproducible to a few per cent, provided that a reducing agent such as stannous chloride is present to prevent oxidation of the colored salt by dissolved oxygen.
The color-forming reaction is interesting because most $\alpha$-amino acids give the same color irrespective of their structure.$^4$ The sequence of steps that leads to the color is as follows:
A new, very sensitive method of detection and analysis of amino acids, which is useful down to the $10^{-12}$ mole (picomole) level, depends on the formation from $\ce{RNH_2}$ and "fluorescamine", $1$, of substances that are intensely fluorescent in ultraviolet light:
Paper Chromatography
Ninhydrin (or fluorescamine) is very useful in chromatographic methods for the analysis of amino acids. One of these is paper chromatography, wherein amino acids are separated as the consequence of differences in their partition coefficients between water and an organic solvent. The aqueous phase is held stationary in the pores of the paper because of strong interaction of the water with the hydroxyl functions of the cellulose. The differences in partition coefficients show up as differences in rates of migration on the surface of moist (but not wet) paper over which there is a slow flow of a water-saturated organic solvent. We shall discuss one of several useful modes of operation. In this example, a drop of the solution to be analyzed is placed on the corner of a sheet of moist paper (often filter paper), which is then placed in an apparatus like that of Figure 25-2, arranged so that the organic solvent can migrate upward by capillarity across the paper, carrying the amino acids with it along one edge. The acids that have the greatest solubility in the organic solvent move most rapidly and when the solvent reaches the top of the paper, the paper is removed, dried, and then turned sidewise, and a different solvent allowed to migrate upward. This double migration process gives a better separation of the amino acids than a single migration and results in concentration of the different amino acids in rather well-defined spots. These spots can be made visible by first drying and then spraying the paper with ninhydrin solution. The final result is as shown in Figure 25-3 and usually is quite reproducible under a given set of conditions. The identities of the amino acids that produce the various spots are established by comparison with the behavior of known mixtures.
Analysis by thin-layer chromatography (see Section 9-2B) can be carried out in the same way as paper chromatography. The partitioning is now between a solid stationary phase (the coating on the plate) and the moving solvent front.
Ion-Exchange Chromatography
The advent of ion-exchange chromatography has revolutionized the separation and analysis of amino acids as well as that of many inorganic substances. As the name implies, it involves the exchange of ions between a stationary and a moving phase. The stationary phase is an insoluble polymer (or resin) having chains on which are located ionic functions such as sulfonate groups, $\ce{-SO_3^-}$ or quaternary ammonium groups, $\ce{-} \overset{\oplus}{\ce{N}} \ce{R_3}$. The counterions to these groups, such as $\ce{Na}^\oplus$ or $\ce{Cl}^\ominus$, are not bound to the resin and can be exchanged for other ions in the mobile phase as the mobile phase travels through the resin. A common application of this principle is in household water softeners, in which the calcium and magnesium ions in ordinary "hard" water are replaced by sodium ions from the resin (Equation 25-3). The resulting "soft" water can be freed of metal ions, if desired, by exchanging the $\ce{Na}^\oplus$ ions for protons (Equation 25-4):
$2 \left( \text{resin} \ce{-SO_3^-} \ce{Na^+} \right)+ \ce{Ca^{2+}} \rightleftharpoons 2 \left( \text{resin} \ce{-SO_3^-} \right) \ce{Ca^{2+}} + 2 \ce{Na^+} \tag{25-3}$
$\text{resin} \ce{-SO_3^-} \ce{H^+} \rightleftharpoons \text{resin} \ce{-SO_3^-} \ce{Na^+} + \ce{H^+} \tag{25-4}$
In strongly acidic solutions (pH $\sim 0$), the amine and carboxyl groups of an amino acid are completely protonated. This cationic form of the amino acid can be exchanged with the cations associated with the sulfonate groups of the resin:
$\text{resin} \ce{-SO_3^-} \ce{Na^+} + \ce{H_3} \overset{+}{\ce{N}} \ce{CRHCO_2H} \rightleftharpoons \text{resin} \ce{-SO_3^-} \ce{H_3} \overset{+}{\text{N}} \ce{CRHCO_2H} + \ce{Na^+} \tag{25-5}$
The process is reversible, and the amino acid cations can in turn be exchanged off the columns. However, different amino acids have different affinities for the resin, and these are considerably influenced by the pH of the moving phase (eluent). The basic amino acids (arginine, lysine), which form cations most readily, are more strongly held by cation-exchange resins than are acidic amino acids (aspartic and glutamic acids). There is a spectrum of affinities of the other amino acid cations for the resin between these extremes. Thus a mixture of amino acids can be separated by ion-exchange chromatography by elution with buffered aqueous solutions. The effluent from the column is mixed with ninhydrin solution and the intensity of the blue color is measured and plotted as a function of time at constant flow rates (Figure 25-4). The identity of an amino acid is determined by the volume of solvent required to elute the amino acid from the column, and the concentration is determined from the intensity of the color developed.
$^4$Proline and hydroxyproline are exceptions because neither has the necessary primary $\ce{NH_2}$ group needed for the reaction. However, these compounds do react with ninhydrin to give yellow compounds, and these colors can be used to identify them satisfactorily.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/25%3A_Amino_Acids_Peptides_and_Proteins/25.04%3A_Analysis_of_Amino_Acids.txt |
Ester and Amide Formation
To some degree the reactions of amino acids are typical of isolated carboxylic acid and amine functions. Thus the carboxyl function can be esterified with an excess of an alcohol under acidic conditions, and the amine function can be acylated with acid chlorides or anhydrides under basic conditions:
The products, however, are not indefinitely stable because the functional groups can, and eventually will, react with each other. For example, in the acylation of glycine with ethanoic anhydride, the first-formed product may cyclize to the "azlactone" if the reaction is prolonged or excess anhydride is used:
Esters of amino acids also cyclize, but they do so intermolecularly to give "diketopiperazines". These compounds are cyclic amides:
Nitrous Acid Reaction
The amine function of $\alpha$-amino acids and esters reacts with nitrous acid in a manner similar to that described for primary amines (Section 23-10A). The diazonium ion intermediate loses molecular nitrogen in the case of the acid, but the diazonium ester loses a proton and forms a relatively stable diazo compound known as ethyl diazoethanoate:
This diazo ester is formed because loss of $\ce{N_2}$ from the diazonium ion results in formation of a quite unfavorable carbocation.
Amino Acids with Aldehydes
$\alpha$-Amino acids react with aldehydes to form decarboxylation and/or deamination products. The reaction sequence is shown in Figure 25-5 and closely resembles the ninhydrin reaction (Section 25-4A). In the first step the amine condenses with the aldehyde to give an imine or Schiff base, $2$. What happens next depends on the relative rates of proton shift and decarboxylation of $2$. Proton shift produces a rearranged imine, $3$, which can hydrolyze to the keto acid $4$. The keto acid is a deamination product. Alternatively, decarboxylation can occur (see Section 18-4) and the resulting imine, $5$, can either hydrolyze or rearrange by a proton shift to a new imine, $6$. Hydrolysis of $5$ or $6$ gives an aldehyde and an amine.
There is an important biochemical counterpart of the deamination reaction that utilizes pyridoxal phosphate, $7$, as the aldehyde. Each step in the sequence is catalyzed by a specific enzyme. The $\alpha$-amino group of the amino acid combines with $7$ and is converted to a keto acid. The resulting pyridoxamine then reacts to form an imine with a different $\alpha$-keto acid, resulting in formation of a new $\alpha$-amino acid and regenerating $7$. The overall process is shown in Equation 25-6 and is called transamination. It is a key part of the process whereby amino acids are metabolized.
$\tag{25-6}$
The biochemical process occurs with complete preservation of the $L$ configuration at the $\alpha$ carbon. The same reactions can be carried out nonenzymatically using pyridoxal phosphate, but they are nonstereospecific, require metal ions as a catalyst, and give mixtures of products.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format."
25.06: Synthesis of -Amino Acids
Many of the types of reactions that are useful for the preparation of amino acids have been discussed previously in connection with separate syntheses of carboxylic acids (Chapter 18) and amino compounds (Chapter 23). Examples include the $S_\text{N}2$ displacement of halogen from $\alpha$-halo acids by ammonia,
and the Strecker synthesis, which, in its first step, bears a close relationship to cyanohydrin formation (Section 16-4A):
Other general synthetic methods introduce the $\alpha$-amino acid grouping, $\ce{H_2N-CH-CO_2H}$, by way of enolate anions. Two selected examples follow. Notice that in each a carbanion is generated and alkylated. Also the $\ce{H_2N}-$ group is introduced as a protected amide or imide group.
1. phthalimidomalonic ester synthesis
2. $\ce{N}$-formylaminomalonic ester synthesis
With those amino acids that are very soluble in water, it usually is necessary to isolate the product either by evaporation of an aqueous solution or by precipitation induced by addition of an organic solvent like alcohol. Difficulty may be encountered in obtaining a pure product when inorganic salts are coproducts of the synthesis. The best general method for removal of inorganic salts involves passage of the solutions through columns of suitable ion-exchange resins (Section 25-4C).
The products of laboratory syntheses, starting with achiral reagents, are of course racemic $\alpha$-amino acids. To obtain the natural amino acids, the $D$,$L$ mixtures must be resolved (Section 19-3).
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/25%3A_Amino_Acids_Peptides_and_Proteins/25.05%3A_Reactions_of_Amino_Acids.txt |
Classification
Amino acids are the building blocks of the polyamide structures of peptides and proteins. Each amino acid is linked to another by an amide (or peptide) bond formed between the $\ce{NH_2}$ group of one and the $\ce{CO_2H}$ group of the other:
In this manner a polymeric structure of repeating amide links is built into a chain or ring. The amide groups are planar and configuration about the $\ce{C-N}$ bond is usually, but not always, trans (Section 24-1). The pattern of covalent bonds in a peptide or protein is called its primary structure:
The distinction between a protein and a peptide is not completely clear. One arbitrary choice is to call proteins only those substances with molecular weights greater than 5000. The distinction might also be made in terms of differences in physical properties, particularly hydration and conformation. Thus proteins, in contrast to peptides, have very long chains that are coiled and folded in particular ways, with water molecules filling the voids in the coils and folds. Hydrogen bonding between the amide groups plays a decisive role in holding the chains in juxtaposition to one another, in what is sometimes called the secondary and tertiary structure.$^5$ Under the influence of heat, organic solvents, salts, and so on, protein molecules undergo changes, often irreversibly, called denaturation. The conformations of the chains and the degree of hydration are thereby altered, with the result that solubility and ability to crystallize decreases. Most importantly, the physiological properties of the protein usually are destroyed permanently on denaturation. Therefore, if a synthesis of a protein is planned, it would be necessary to duplicate not only the amino-acid sequences but also the exact conformations of the chains and the manner of hydration characteristic of the native protein. With peptides, the chemical and physiological properties of natural and synthetic materials usually are identical, provided the synthesis duplicates all of the structural and configurational elements. What this means is that a peptide automatically assumes the secondary and tertiary structure characteristic of the native peptide on crystallization or dissolution in solvents.
Representation of peptide structures of any length with conventional structural formulas is cumbersome. As a result, abbreviations are universally used that employ three-letter symbols for the component amino acids. It is important that you know the conventions for these abbreviations. The two possible dipeptides made up of one glycien and one alanine are
Notice that in the conventions used for names and abbreviated formulas the amino acid with the free amino group (the $\ce{N}$-terminal amino acid) always is written on the left. The amino acid with the free carboxyl group (the $\ce{C}$-terminal amino acid) always is written on the right. The dash between the three-letter abbreviations for the acids designates that they are linked together by an amide bond.
Determination of Amino-Acid Sequences
The general procedure for determining the primary structure of a peptide or protein consists of three main steps. First, the number and kind of amino-acid units in the primary structure must be determined. Second, the amino acids at the ends of the chains are identified, and third, the sequence of the component amino acids in the chains is determined.
The amino-acid composition usually is obtained by complete acid hydrolysis of the peptide into its component amino acids and analysis of the mixture by ion-exchange chromatography (Section 25-4C). This procedure is complicated by the fact that tryptophan is destroyed under acidic conditions. Also, asparagine and glutamine are converted to aspartic and glutamic acids, respectively.
Determination of the $\ce{N}$-terminal acid in the peptide can be made by treatment of the peptide with 2,4-dinitrofluorobenzene, a substance very reactive in nucleophilic displacements with amines but not amides (see Section 14-6B). The product is an $\ce{N}$-2,4-dinitrophenyl derivative of the peptide which, after hydrolysis of the amide linkages, produces an $\ce{N}$-2,4-dinitrophenylamino acid:
These amino-acid derivatives can be separated from the ordinary amino acids resulting from hydrolysis of the peptide because the low basicity of the 2,4-dinitrophenyl-substituted nitrogen (Section 23-7C) greatly reduces the solubility of the compound in acid solution and alters its chromatographic behavior. The main disadvantage to the method is that the entire peptide must be destroyed in order to identify the one $\ce{N}$-terminal acid.
A related and more sensitive method makes a sulfonamide of the terminal $\ce{NH_2}$ group with a reagent called "dansyl chloride". As with 2,4-dinitrofluorobenzene, the peptide must be destroyed by hydrolysis to release the $\ce{N}$-sulfonated amino acid, which can be identified spectroscopically in microgram amounts:
A powerful method of sequencing a peptide from the $\ce{N}$-terminal end is the Edman degradation in which phenyl isothiocyanate, $\ce{C_6H_5N=C=S}$, reacts selectively with the terminal amino acid under mildly basic conditions. If the reaction mixture is then acidified, the terminal amino acid is cleaved from the peptide as a cyclic thiohydantoin, $8$:
There are simple reagents that react selectively with the carboxyl terminus of a peptide, but they have not proved as generally useful for analysis of the $\ce{C}$-terminal amino acids as has the enzyme carboxypeptidase A. This enzyme catalyzes the hydrolysis of the peptide bond connecting the amino acid with the terminal carboxyl groups to the rest of the peptide. Thus the amino acids at the carboxyl end will be removed one by one through the action of the enzyme. Provided that appropriate corrections are made for different rates of hydrolysis of peptide bonds for different amino acids at the carboxyl end of the peptide, the sequence of up to five or six amino acids in the peptide can be deduced from the order of their release by carboxypeptidase. Thus a sequence such as peptide-Ser-Leu-Tyr could be established by observing that carboxypeptidase releases amino acids from the peptide in the order Tyr, Leu, Ser:
Determining the amino-acid sequences of large peptides and proteins is very difficult. Although the Edman degradation and even carboxypeptidase can be used to completely sequence small peptides, they cannot be applied successfully to peptide chains with several hundred amino acid units. Success has been obtained with long peptide chains by employing reagents, often enzymes, to selectively cleave certain peptide bonds. In this way the chain can be broken down into several smaller peptides that can be separated and sequenced. The problem then is to determine the sequence of these small peptides in the original structure. To do this, alternative procedures for selective cleavages are carried out that produce different sets of smaller peptides. It is not usually necessary to sequence completely all of the peptide sets. The overall amino-acid composition and the respective end groups of each peptide may suffice to show overlapping sequences from which the complete amino-acid sequence logically can be deduced.
The best way to show you how the overlap method of peptide sequencing works is by a specific example. In this example, we will illustrate the use of the two most commonly used enzymes for selective peptide cleavage. One is trypsin, a proteolytic enzyme of the pancreas (MW 24,000) that selectively catalyzes the hydrolysis of the peptide bonds of basic amino acids, lysine and arginine. Cleavage occurs on the carboxyl side of lysine or arginine:
Chymotrypsin is a proteolytic enzyme of the pancreas (MW 24,500) that catalyzes the hydrolysis of peptide bonds to the aromatic amino acids, tyrosine, tryptophan, and phenylalanine, more rapidly than to other amino acids. Cleavage occurs on the carboxyl side of the aromatic amino acid:
Our example is the sequencing of a peptide (P) derived from partial hydrolysis of a protein which, on complete acid hydrolysis, gave Ala, 3 Gly, Glu, His, 3 Lys, Phe, Tyr, 2 Val, and one molar equivalent of ammonia.
1. Treatment of the peptide (P) with carboxypeptidase released alanine, and with 2,4-dinitrofluorobenzene followed by hydrolysis gave the 2,4-dinitrophenyl derivative of valine. These results establish the $\ce{N}$-terminus as valine and the $\ce{C}$-terminus as alanine. The known structural elements now are
2. Partial hydrolysis of the peptide (P) with trypsin gave a hexapeptide, a tetrapeptide, a dipeptide, and one molar equivalent of lysine. The peptides, which we will designated respectively as M, N, and O, were sequenced by Edman degradation and found to have structures:
$\begin{array}{ll} \text{Gly}-\text{Ala} & \text{O} \ \text{Val}-\text{Tyr}-\text{Glu}-\text{Lys} & \text{N} \ \text{Val}-\text{Gly}-\text{Phe}-\text{Gly}-\text{His}-\text{Lys} & \text{M} \end{array}$
With this information, four possible structures can be written for the original peptide P that are consistent with the known end groups and the fact that trypsin cleaves the peptide P on the carboxyl side of the lysine unit. Thus
$\begin{array}{cc} \text{N}-\text{M}-\text{Lys}-\text{O} & \text{M}-\text{N}-\text{Lys}-\text{O} \ \text{N}-\text{Lys}-\text{M}-\text{O} & \text{M}-\text{Lys}-\text{N}-\text{O} \end{array}$
3. Partial hydrolysis of the peptide P using chymotrypsin as catalyst gave three peptides, X, Y, and Z. These were not sequenced, but their amino-acid composition was determined:
$\begin{array}{ll} \text{Gly, Phe, Val} & \text{X} \ \text{Gly, His, Lys, Tyr, Val} & \text{Y} \ \text{Ala, Glu, Gly, 2 Lys} & \text{Z} \end{array}$
This information can be used to decide which of the alternative structures deduced above is correct. Chymotrypsin cleaves the peptide on the carboxyl side of the phenylalanine and tyrosine units. Only peptide M contains Phe, and if we compare M with the compositions of X, Y, and Z, we see that only X and Y overlap with M. Peptide Z contains the only Ala unit and must be the $\ce{C}$-terminus. If we put together these pieces to get a peptide, P' (which differs from P by not having the nitrogen corresponding to the ammonia formed on complete hydrolysis) then P' must have the structure X-Y-Z:
This may not be completely clear, and it will be well to consider the logic in some detail. Peptides M and N both have $\ce{N}$-terminal valines, and one of them must be the $\ce{N}$-terminal unit. Peptide M overlaps with X and Y, and because X and Y are produced by a cleavage on the carboxyl side of Phe, the X and Y units have to be connected in the order X-Y. Because the other Val is in Y, the $\ce{N}$-terminus must be M. This narrows the possibilities to
$\begin{array}{c} \text{M}-\text{N}-\text{Lys}-\text{O} \ \text{M}-\text{Lys}-\text{N}-\text{O} \end{array}$
There are two Lys units in Z, and this means that only the sequence M-N-Lys-O is consistent with the sequence X-Y-Z, as shown:
The final piece of the puzzle is the placement of the mole of ammonia released from the original peptide on acid hydrolysis. The ammonia comes from a primary amide function:
$\ce{R-CONH_2} \overset{\ce{H_3O}^\oplus}{\longrightarrow} \ce{RCO_2H} + \ce{NH_4^-}$
The amide group cannot be at the $\ce{C}$-terminus because the peptide would then be inert to carboxypeptidase. The only other possible place is on the side-chain carboxyl of glutamic acid. The complete structure may be written as
Using procedures such as those outlined in this section more than 100 proteins have been sequenced. This is an impressive accomplishment considering the complexity and size of many of these molecules (see, for example, Table 25-3). It has been little more than two decades since the first amino acid sequence of a protein was reported by F. Sanger, who determined the primary structure of insulin (1953). This work remains a landmark in the history of chemistry because it established for the first time that proteins have definite primary structures in the same way that other organic molecules do. Up until that time, the concept of definite primary structures for proteins was openly questioned. Sanger developed the method of analysis for $\ce{N}$-terminal amino acids using 2,4-dinitrofluorobenzene and received a Nobel Prize in 1958 for his success in determining the amino-acid sequence of insulin.
Methods for Forming Peptide Bonds
The problems involved in peptide syntheses are of much practical importance and have received considerable attention. The major difficulty in putting together a chain of say 100 amino acids in a particular order is one of overall yield. At least 100 separate synthetic steps would be required, and, if the yield in each step were equal to $n \times 100\%$, the overall yield would be $\left( n^{100} \times 100\% \right)$. If the yield in each step were $90\%$, the overall yield would be only $0.003\%$. Obviously, a practical laboratory synthesis of a peptide chain must be a highly efficient process. The extraordinary ability of living cells to achieve syntheses of this nature, not of just one but of a wide variety of such substances, is truly impressive.
Several methods for the formation of amide bonds have been discussed in Sections 18-7A and 24-3A. The most general reaction is shown below, in which X is some reactive leaving group (see Table 24-1):
When applied to coupling two different amino acids, difficulty is to be expected because these same reactions can link two amino acids in a total of four different ways. Thus if we started with a mixture of glycine and alanine, we could generate for dipeptides, Gly-Ala, Ala-Gly, Gly-Gly, and Ala-Ala.
To avoid unwanted coupling reactions a protecting group is substituted on the amino function of the acid that is to act as the acylating agent. Furthermore, all of the amino, hydroxyl, and thiol functions that may be acylated to give undesired products usually must be protected. For instance, to synthesize Gly-Ala free of other possible dipeptides, we would have to protect the amino group of glycine and the carboxyl group of alanine:
Some methods of protecting amine and hydroxyl functions were discussed previously in Sections 23-13 and 15-9, respectively. A summary of some commonly used protecting groups for $\ce{NH_2}$, $\ce{OH}$, $\ce{SH}$, and $\ce{CO_2H}$ functions is in Table 25-2, together with the conditions by which the protecting groups may be removed. The best protecting groups for $\ce{NH_2}$ functions are phenylmethoxycarbonyl (benzyloxycarbonyl) and tert-butoxycarbonyl. Both groups can be removed by treatment with acid, although the tert-butoxycarbonyl group is more reactive. The phenylmethoxycarbonyl group can be removed by reduction with either hydrogen over a metal catalyst or with sodium in liquid ammonia. This method is most useful when, in the removal step, it is necessary to avoid treatment with acid:
Table 25-2: Some Amine and Carboxyl Protecting Groups Used in Peptide Syntheses
In most cases, formation of the ethyl ester provides a satisfactory protecting group for the carboxyl function.
Conversion of the carboxyl group to a more reactive group and coupling are key steps in peptide synthesis. The coupling reaction must occur readily and quantitatively, and with a minimum of racemization of the chiral centers in the molecule. This last criterion is the Achilles' heel of many possible coupling sequences. The importance of nonracemization can best be appreciated by an example. Consider synthesis of a tripeptide from three protected $L$-amino acids, A, B, and C, in two sequential coupling steps, $\text{C} \overset{\text{B}}{\rightarrow} \text{B}-\text{C} \overset{\text{A}}{\rightarrow} \text{A}-\text{B}-\text{C}$. Suppose that the coupling yield is quantitative, but there is $20\%$ formation of the $D$ isomer in the acylating component in each coupling step. Then the tripeptide will consist of a mixture of four diastereomers, only $64\%$ of which will be the desired $L$,$L$,$L$ diastereomer (Equation 25-7):
$\tag{25-7}$
This is clearly unacceptable, especially for longer-chain peptides. Nine coupling steps with $20\%$ of the wrong isomer formed in each would give only $13\%$ of the decapeptide with the correct stereochemistry.
The most frequently used carboxyl derivatives in amide coupling are azides, $\ce{RCO-N_3}$, mixed anhydrides, $\ce{RCO-O-COR'}$, and esters of moderately acidic phenols, $\ce{RCO-OAr}$ (see Table 24-1). It also is possible to couple free acid with an amine group using a diimide, $\ce{R-N=C=N-R}$, most frequently $\ce{N}$,$\ce{N'}$-dicyclohexylcarbodiimide.
The diimide reagent may be thought of as a dehydrating agent. The "elements of water" eliminated in the coupling are consumed by the diimide to form a substituted urea. The overall reaction is
This reaction takes place because diimides, $\ce{-N=C=N}-$, have reactive cumulated double-bond systems like those of ketenes, $\ce{-C=C=O}$; isocyanates, $\ce{-N=C=O}$; and isothiocyanates, $\ce{-N=C=S}$; and are susceptible to nucleophilic attack at the central carbon. In the first step of the diimide-coupling reaction, the carboxyl function adds to the imide to give an acyl intermediate, $9$. This intermediate is an activated carboxyl derivative $\ce{RCO-X}$ and is much more reactive toward an amino function than is the parent acid. The second step therefore is the aminolysis of $9$ to give the coupled product and $\ce{N}$,$\ce{N'}$-dicyclohexylurea:
After completion of a coupling reaction, and before another amino acid can be added to the $\ce{N}$-terminus, it is necessary to remove the protecting group. This must be done by selective reactions that do not destroy the peptide bonds or side-chain protecting groups. This part of peptide synthesis is discussed in Section 23-13, and some reactions useful for removal of the $\ce{N}$-terminal protecting groups are summarized in Table 25-2.
In spite of the large number of independent steps involved in the synthesis of even small peptides, each with its attendant problems of yield, racemization, and selectivity, remarkable success has been achieved in the synthesis of large peptides and certain of the smaller peptides. The synthesis of insulin (Figure 25-8) with its 51 amino acid units and 3-disulfide bridges has been achieved by several investigators. Several important hormonal peptides, namely glutathione, oxytocin, vasopressin, and thyrotropic hormone (see Figure 25-9) have been synthesized. A major accomplishment has been the synthesis of an enzyme with ribonuclease activity reported independently by two groups of investigators, led by R. Hirschman (Merck) and R. B. Merrifield (Rockefeller University). This enzyme is one of the simpler proteins, having a linear stricture of 124 amino-acid residues. It is like a peptide, not a protein, in that it assumes the appropriate secondary and tertiary structure without biochemical intervention (Section 25-7A). As a specific example of the strategy involved in peptide synthesis, the stepwise synthesis of oxytocin is outlined in Figure 25-10, using the abbreviated notation in common usage.
not to $\ce{C_1}$. Vasopressin (middle left) and oxytocin (middle right) are peptide hormones from the posterior lobe of the pituitary gland. They function primarily to raise blood pressure (vasopressin), as antidiuretic (vasopressin), and to promote contraction of uterus and lactation muscles (oxytocin). The isolation, identification, and synthesis of these hormones was accomplished by Vincent du Vigneaud, for which he was awarded the Nobel Prize in chemistry in 1965. Thyrotropin-releasing hormone (bottom) is one of several small peptide hormones secreted by the anterior lobe of the pituitary gland. These are the "master" hormones that function to stimulate hormone secretion from other endocrine glands. Thyrotropin stimulates the functioning of the thyroid gland.
Figure 25-10: Stepwise synthesis of oxytocin by the reactive ester method. In the abbreviations used here $-\text{Gly}-\ce{NH_2} = \ce{-NHCH_2CONH_2}$ and
Solid-Phase Peptide Synthesis
The overall yield in a multistep synthesis of a peptide of even modest size is very poor unless each step can be carried out very efficiently. An elegant modification of classical peptide synthesis has been developed by R. B. Merrifield, which offers improved yields by minimizing manipulative losses that normally attend each step of a multistage synthesis. The key innovation is to anchor the $\ce{C}$-terminal amino acid to an insoluble support, and then add amino-acid units by the methods used for solution syntheses. After the desired sequence of amino acids has been achieved, the peptide can be cleaved from the support and recovered from solution. All the reactions involved in the synthesis must, of course, be brought to essentially $100\%$ completion so that a homogeneous product can be obtained. The advantage of having the peptide anchored to a solid support is that laborious purification steps are virtually eliminated; solid material is purified simply by washing and filtering without transferring the material from one container to another. The method has become known as solid-phase peptide synthesis. More of the details of the solid-phase synthesis follow.
The nature of the polymer support is of great importance for a successful peptide synthesis. One that is widely used is a cross-linked polystyrene resin of the type employed in ion-exchange chromatography (Section 25-4C). It is necessary that the resin be insoluble but have a loose enough structure to absorb organic solvents. Otherwise, the reagents will not be able to penetrate into the spaces between the chains. This is undesirable because the reactions occur on the surface of the resin particles and poor penetration greatly reduces the number of equivalents of reactive sites that can be obtained per gram of resin. Finally, to anchor a peptide chain to the resin, a reactive functional group (usually a chloromethyl group) must be introduced into the resin. This can be done by a Friedel-Crafts chloromethylation reaction, which substitutes the $\ce{ClCH_2}-$ group in the 4-position of the phenyl groups in the resin:
At the start of the peptide synthesis, the $\ce{C}$-terminal amino acid is bonded through its carboxyl group to the resin by a nucleophilic attack of the carboxylate ion on the chloromethyl groups. The $\alpha$-amino group must be suitably protected, as with tert-butoxycarbonyl, before carrying out this step:
Next, the amine protecting group must be removed without cleaving the ester bond to the resin. The coupling step to a second $\ce{N}$-protected amino acid follows, with $\ce{N}$,$\ce{N'}$-dicyclohexylcarbodiimide as the coupling reagent of choice:
The peptide-bond-forming steps are repeated as many times as needed to build up the desired sequence. Ultimately, the peptide chain is removed from the resin, usually with $\ce{HBr}$ in anhydrous trifluoroethanoic acid, $\ce{CF_3CO_2H}$, or with anhydrous $\ce{HF}$. This treatment also removes the other acid-sensitive protecting groups.
The method lends itself beautifully to automatic control, and machines suitably programmed to add reagents and wash the product at appropriate times have been developed. At present, the chain can be extended by six or so amino acid units a day. It is necessary to check the homogeneity of the growing peptide chain at intervals because if any step does not proceed properly, the final product can be seriously contaminated with peptides with the wrong sequence.
In the synthesis of the enzyme ribonuclease by the Merrifield method, the 124 amino acids were arranged in the ribonuclease sequence through 369 reactions and some 12,000 individual operations of the automated peptide-synthesis machine without isolation of any intermediates.
Separation of Peptides and Proteins
In many problems of peptide sequencing and peptide synthesis it is necessary to be able to separate mixtures of peptides and proteins. The principal methods used for this purpose depend on acid-base properties or on molecular sizes and shapes.
Ultracentrifugation is widely used for the purification, separation, and molecular-weight determination of proteins. A centrifugal field, up to 500,000 times that of gravity, is applied to the solution, and molecules move downward in the field according to their mass and size.
Large molecules also can be separated by gel filtration (or gel chromatography), wherein small molecules are separated from large ones by passing a solution over a gel that has pores of a size that the small molecules can penetrate into and be trapped. Molecules larger than the pore size are carried on with the solvent. This form of chromatographic separation is based on "sieving" rather than on chemical affinity. A wide range of gels with different pore sizes is available, and it is possible to fractionate molecules with molecular weights ranging from 700 to 200,000. The molecular weight of a protein can be estimated by the sizes of the pores that it will, or will not, penetrate.
The acid-base properties, and hence ionic character, of peptides and proteins also can be used to achieve separations. Ion-exchange chromatography, similar to that described for amino acids (Section 25-4C), is an important separation method. Another method based on acid-base character and molecular size depends on differential rates of migration of the ionized forms of a protein in an electric field (electrophoresis). Proteins, like amino acids, have isoelectric points, which are the pH values at which the molecules have no net charge. At all other pH values there will be some degree of net ionic charge. Because different proteins have different ionic properties, they frequently can be separated by electrophoresis in buffered solutions. Another method, which is used for the separation and purification of enzymes, is affinity chromatography, which was described briefly in Section 9-2B.
$^5$ The distinction between secondary and tertiary structure is not sharp. Secondary structure involves consideration of the interactions and spatial relationships of the amino acids in the peptide chains that are close together in the primary structure, whereas tertiary structure is concerned with those that are far apart in the primary structure.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/25%3A_Amino_Acids_Peptides_and_Proteins/25.07%3A_Peptides_and_Proteins.txt |
The biological functions of proteins are extremely diverse. Some act as hormones that regulate various metabolic processes. An example is insulin, which regulates blood-sugar levels. Enzymes act as catalysts for biological reactions, and other proteins serve as biological structural materials - for example, collagen and elastin in connective tissue and keratin in hair. Iron-containing proteins (hemoglobin and myoglobin in mammals) and copper-containing proteins (hemocyanins in shellfish) transport molecular oxygen. Some blood proteins form antibodies, which provide resistance to disease, while the so-called nucleoproteins are important constituents of the genes that supply and transmit genetic information in cell division. Motion by means of muscle contraction and the generation and transmission of nerve impulses also involve proteins.
How can a group of compounds, made from a common basis set of amino acids, be so remarkably heterogeneous and exhibit such varied yet specific functions? Clearly, the primary structure and presence or absence of special functional groups, metals, and so on, are of paramount importance. Of complementary importance are the three-dimensional structures of proteins, which are dictated not just by the primary structure but by the way the primary structure is put together biochemically. The polypeptide chains are seldom, if every fully extended, but are coiled and folded into more or less stable conformations. As a result, amino-acid side chains in distant positions in the linear sequence are brought into close proximity, and this juxtaposition often is crucial for the protein to fulfill its specific biological function.
Three-Dimensional Structure of Proteins
The elucidation of the detailed shape of protein molecules - in fact, the spatial locations of the individual atoms in a protein - is accomplished primarily by x-ray crystallography. The three-dimensional structures of more than twenty proteins have now been established by this technique. The importance of x-ray crystallography to structural and biological chemistry has been recognized in the award of six Nobel Prizes in this area.$^6$ A number of important proteins and their properties are listed in Table 25-3.
Table 25-3: A Few Important Proteins of Known Structure
There are several other points to notice about the $\alpha$ helix shown in Figure 25-11. The amide groups are planar and normally retain the stable trans configuration in the helical structure; bond lengths and bond angles are normal, and the $\ce{-NH} \cdots \ce{O=C}-$ hydrogen bonds are nearly linear. However, the hydrogen bonds are not quite parallel to the long axis of the coil, so there are 3.6 rather than 4 amino-acid units per helical turn, and the spacing between turns is about $5.4 \: \text{Å}$. The $\alpha$ helix in proteins has right-handed turns like a right-hand screw thread.
The amino acids of the side chains lie outside the coil of the $\alpha$ helix and are in close proximity to the side chains three and four amino-acid units apart. Because of this proximity, steric hindrance between large side chains can be sufficient to reduce the stability of the normal $\alpha$ helix. When such hindrance occurs, there is a discontinuity in the helical structure, and the peptide chain may assume more random arrangements about the $\ce{C-C}_\alpha$ and $\ce{N-C}_\alpha$ bonds (see Figure 25-12), thereby allowing the molecule to fold back on itself and form new hydrogen bonds. The helical structure apparently is always interrupted at proline or hydroxyproline residues because the $\ce{N-C}_\alpha$ bonds of these amino acids are not free to rotate (they are incorporated in five-membered rings) and also because the proline and hydroxyproline amide nitrogens have no hydrogens to participate in hydrogen bonding to carbonyl groups.
Pauling and Corey recognized a second stable conformation of polypeptide chains - the extended chain or $\beta$-pleated sheet (Figure 25-13). In this conformation the chains are fully extended with trans amide configurations. In this arrangement the distance is maximized between adjacent amino-acid side chains. Hydrogen bonding of the type $\ce{-N-H} \cdots \ce{O=C}-$ is now between chains rather than between amino acids in a single chain (as in the $\alpha$ helix). This type of structure is not as common as the $\alpha$ helix and is found extensively only in silk fibroin. However, a number of proteins with a single polypeptide chain can form short sections of "antiparallel" $\beta$-pleated sheets by folding back on themselves, as illustrated in Figure 25-14.
Another very important factor in protein architecture is the disulfide $\ce{-S-S}-$ link. Remote parts of the polypeptide chain can be held close together through the oxidative coupling of two cysteine thiol groups to form a disulfide bridge:
The disulfide bridges in some proteins are between different peptide chains. Insulin, for instance, has two interchain as well as one intrachain $\ce{S-S}$ bridges (Figure 25-8).
Myoglobin and Hemoglobin
Some idea of the complexity of protein conformations can be gained from the structure of myoglobin. This protein is responsible for the storage and transport of molecular oxygen in the muscle tissue of mammals. It is a compact molecule of 153 amino-acid units in a chain that is extensively coiled as an $\alpha$ helix. There are eight regions of discontinuity in the helical structure, and in these regions the chain folds on itself as shown in Figure 15-16. Four of the eight nonhelical regions occur at proline residues; the reason for the discontinuity at the other regions is not entirely clear. With the exception of two histidine units, the interior regions of myoglobin accommodate only the nonpolar side chains; the interior, therefore, is mostly hydrocarbonlike and repellant to water and other polar molecules. In contrast, the polar side chains are on the exterior of the protein.
A number of proteins, including myoglobin, possess one or more nonpeptide components associated with specific sites on the polypeptide chain. These components are called prosthetic groups and are essential to the biological activity. When the prosthetic group is removed, the residual protein is referred to as an apoprotein.
In myoglobin the prosthetic group is a molecule of heme. The heme group belongs to a class of interesting compounds called metalloporphyrins, which are metal complexes of a highly conjugated ring system composed of four azacyclopentadiene (pyrrole) rings linked by $\ce{-CH=}$ bridges between the 2 and 5 positions. The parent compound is known as porphin. Porphyrins have highly stabilized electronic excited states and absorb visible light. As a result they usually are brightly colored compounds (e.g., chlorophyll, Figure 20-6).
The porphyrin of heme is known as protoporphyrin IX, and the associated metal is iron [as $\ce{Fe}$(II) or $\ce{Fe}$(III)]. You will notice that the porphyrin ring carries methyl, ethenyl, and propanoic acid side chains:
A major effort on the part of several eminent chemists in the early part of the century led to the elucidation of the structure of heme. The German chemist Hans Fischer successfully synthesized heme in 1929, a feat for which, in 1930, he received the Nobel Prize in chemistry. [Some years earlier (1915), Richard Willstatter received a Nobel Prize for structural studies of chlorophyll and plant pigments.]
A very important question is, how does the particular combination of protein and iron-porphyrin allow myoglobin to reversibly bind molecular oxygen? The answer to this question is not known in all its details, but it is well established that $\ce{Fe}$(II)-porphyrins will complex readily and reversibly with oxygen. There are two additional coordination sites around the iron in heme besides the four ring nitrogens. These are indicated below as the general ligands $\ce{L}$:
The disclike heme molecule fits into a cleft in the protein structure and is bound to it through one of the $\ce{L}$ coordination sites to a histidine nitrogen. The remaining coordination site on the other side of the ring is occupied by molecular oxygen. In the absence of the coordination by histidine, the porphyrin iron would be oxidized rapidly to the ferric state, which does not bind oxygen.
A number of model compounds have been synthesized which have $\ce{Fe}$(II)-porphyrin rings carrying a side chain with histidine arranged to be able to coordinate with the metal on one side. Several of these substances show promise as oxygen carriers with properties similar to myoglobin.
Hemoglobin is related to myoglobin in both its structure and function. It reversibly binds molecular oxygen which it transports in the red corpuscles of blood rather than in muscle tissue. However, hemoglobin is made up of four polypeptide chains, in contrast to myoglobin which has only one chain. Two of the hemoglobin chains are of one kind with 141 amino acid residues, called the $\alpha$ chains, and two are of another kind with 146 amino acids, called the $\beta$ chains. Each chain, or subunit, contains one heme group identical with the heme in myoglobin. The subunits are held in the hemoglobin by noncovalent interactions and provide four hemes and hence four binding sites for molecular oxygen. The $\alpha$ and $\beta$ hemes have different affinities for oxygen but function in a cooperative way to increase oxygen availability to the cells.
In spite of the fact that the $\alpha$ and $\beta$ chains of hemoglobin are nonidentical with the myoglobin chain, the three-dimensional structures of all three chains are strikingly similar; myoglobins and hemoglobins differ slightly in amino acid composition, depending on the species but the protein shape remains essentially the same.
Quaternary Structures of Proteins
Many factors contribute to the three-dimensional structures of proteins. We already have mentioned hydrogen bonding between amide groups, location and character of prosthetic groups, and disulfide bonds. Other important influences include electrostatic interactions between ionic groups ($\ce{-NH_3^+}$, $\ce{-CO_2^-}$), hydrogen-bonding involving side-chain substituents $\left( \ce{-CH_2OH} \right)$, and nonbonded interactions. Except for the disulfide linkages, most of these interactions are weak compared to covalent bond strengths, and the conformations of many proteins can be altered rather easily. In fact, several have conformations that clearly are in dynamic equilibrium under physiological conditions. Such structural flexibility may be necessary for the protein to be functional, but if the conformation is altered irreversibly - that is, if it is denatured - its biological activity usually is destroyed.
In many cases there are important interactions between protein molecules that may lead to highly organized structures such as the pleated sheet of silk fibroin (Figure 25-13) or the coiling of $\alpha$ helices, as found in $\alpha$-keratins, the fibrous proteins of hair, horn, and muscles (Figure 25-17). This sort of organization of protein molecules is called quaternary structure and is an important feature of many proteins that associate into dimers, tetramers, and so on. The tetrameric structure of hemoglobin is an important example.
$^6$The following Nobel laureates received their awards for contributions to the use of x-ray crystallography for structure determination: 1914, Max von Laue (physics), diffraction of x-rays in crystals; 1915, William Bragg and Lawrence Bragg (physics), study of crystal structure by means of x-rays; 1954, Linus Pauling (chemistry), study of structure of proteins; 1962, Max Perutz and John Kendrew (chemistry), structures of myoglobin and hemoglobin; 1962, Francis Crick, James Watson, and Maurice Wilkins (physiology and medicine), double helix of DNA; 1964, Dorothy Hodgkin (chemistry), determination of structure of vitamin B$_{12}$ and penicillin by x-ray methods. She later determined the three-dimensional structure of insulin.
Contributors and Attributions
John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, "You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format." | textbooks/chem/Organic_Chemistry/Basic_Principles_of_Organic_Chemistry_(Roberts_and_Caserio)/25%3A_Amino_Acids_Peptides_and_Proteins/25.08%3A_Structure_and_Function_of_Proteins.txt |
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