Acidity of Organic Compounds and Resonance CALVIN A. VANDER WERF University of Kansas, Lawrence, Kansas 0 . NCE he has weighed all the many arguments favonng each of the various acid-base concepts and has made a timid decision in favor of one or another, the teacher who is seriously attempting a correlation and integration of organic chemistry by application of the electronic theory is left squarely on the hprns of a second dilemma. How can he best relate acidity to structure; i. e., bow can he teach his students to predict by inspection of the formula whether a given compound will manifest acid properties and, if so, to what extent and a t what point in the molecule? If the Bronsted definition of acids is the one adopted, this problem resolves itself into an analysis of those structural characteristics which confer acid properties upon a hydrogen atom in a molecule. At the present time, this question is approached by sponsors of the electronic theory from two complementary viewpoints. According to the first argument, an acid hydrogen is one which is bound to another atom or a group which has a high tendency to attract the electrons of the covalence binding the hydrogen (la, 2, 3). I n other words, an acid hydrogen is a positive or electrophilic hydrogen, a center of low electron density in which the effective nuclear charge is relatively high. In the general case of the type molecule Y-0-H, for example, unless Y is an atom or group of exceedingly low electronegativity, the hydrogen atom, because of its attachment to highly electronegative oxygen, will exhibit definite acid properties which become more marked as the electron withdrawal tendency of Y is increased. Thus, on the basis of relative electronegativities, it may be concluded not only that the hydrogen of the hydroxyl group in primary alcohols will show acid character, but also that it will be more acidic than the hydrogen of secondary or tertiary alcohols, but less acidic than the hydrogen of water. In the case of phenol, the resonance forms in which one of the unshared electron pairs of the oxygen is shifted to increase the covalence between the oxygen and the ring carbon atom (+T effect), contribute considerably to the structure of the molecule and enhance the acidity of the hydroxyl hydrogen by placing a positive charge on the oxygen atom. Introduction on the ring of an electron acceptor such as the nitro group, particularly a t the ortho or para position, serves to decrease still further the electron density about the oxygen, thereby producing an additional increase in acidity. The carboxylic acids, according to the argument, are even stronger because the electron withdrawal from the hydroxyl ,group . - due to the resonance about the doubly bbnded oxygen (-T effect) is greater. I n the acetone molecule, resonance (-T effect) 40 about the oxygen atom produces a low electron density a t the carbonyl carbon atom and a slight positive charge is transmitted to each of the methyl carbons. As a result, the hydrogen atoms of acetone are weakly acidic. When the effect is doubled, as in the case of acetylacetone, the acidity becomes much more marked. Examples of the application of this point of view to a host of additional organic compounds commonly considered as weak acids are too familiar to warrant further discussion. Careful study of the principle leads to the valid assignment of acidic properties to many organic compounds ordinarily not classified as Bronsted acids. In chloroform, to mention only one, the electron affinity induced in the hydrogen atom by the electronegative chlorine atoms should render the former somewhat acidic. Such, indeed, is probably the case, for in the addition of chloroform to acetone in the presence of solid sodium hydroxide, the most feasible mechanism would involve the removal of the proton from the chloroform molecule by the hydroxyl ion. Attack by the resulting trichloromethyl carbanion on the electrophilic carbonyl carbon followed by the transfer of the proton to the oxygen atom would com~1;ltethe reaction. Ac ,,ardingto the second viewpoint, the tendency for, a q . m l e to release a proton is heightened when thr, ior.! .suiting from the discharge of a proton is stabilized' dr~eto resonance. This is especially true if the reso-,.r-.- '-e energy of the ion is greater than that of the ionogm, due either to the fact that the ion resonates among, moie structures than does the ionogen or to the fact, that the resonance forms of the ion are more nearly, egi..yalent. .,;lauy examples of the application of this principle bC: h to inorganic and to organic compounds have been d t ~ dpreviously. As Pauling (4) has pointed out, resi'nance is complete in the carbonate, nitrate, and bot, 'F ions and somewhat inhibited in the acids. The in? sed strength of nitric acid over nitrous acid would be I.--ticted on the basis of the fact that the nitrate ion epresented by three equivalent resonance forms, the trite ion by only two. In the teaching of beginnir students, this concept may be simplified cousic .bly by the explanation that the distribution of the ,gegative charge of the ion over two or more atoms red&ces the electron density on each atom; hence, the attl,ction for the proton is diminished, an effect which is I: 1 ;eased with increasing distribution of the negative chq;e. This broad idea may be extended to correlate facts the relative acidity of many com>, 8. c,?nceming . p o ~ . ,.,:~ such : as the oxygen acids of the halogens, for exak,-2, in which the acid strength increases with a? incr&e in the number of oxygen atoms over whici 'I the negative charge on the ion may be distributed, even though trne resonance is not involved. As a matter of fact, this general concept provides a simple explanation of the well-known rule that an oxide or hydroxide of a given element in a higher valence state is more acidic than an oxide or hydroxide of that element in a lower valence state. As applied to organic compounds, the fact that the stabilization by resonance of the anion resulting from the loss of a proton contributes to the strength of the ionogen as an acid is well illustrated by the carboxylic acids. Although the acids themselves, as well as the carboxylate ions, resonate between two structures, only in the case of the ion are the two structurrs equivalent; hence, the ion is stabilized more than the ionogen, and, as a result, a carboxylic acid is more acidic than an alcohol. The phenols are likewise stronger acids than the alcohols; here, again, the explanation lies in the fact that the resonance forms of the phenolate ion are more nearly equivalent than those of the unionized molecule. The comparatively strong acid character of the phenylammonium ion (weak basic strength of aniline) may be attributed to the fact that, whereas the ion is restricted to two structures, the molecule of the conjugate hase, aniline, resonates among five structures. By similar application of the principle, acid properties of such compounds as 1,3-diketones,@-ketonicesters, amides, nitroalkanes, pyrazole and imidaz.de, the barbiturates, and many others find a logicL1 explanation. Conversely, the basic strength of guanidinc n a y \e attributed to the fact that the guanidiqiur ion esonates among three equivalent structures, wi ..as guanidine itself resonates among three structures a c k h .>I are not equivalent. Even a hasty comparison of the two points of v'-n ~egardingacidity reveals that the "ion resonan&" heory cannot be applied to molecules in which ihe electronic displacements are of the inductive type oi'v. The question of the acidic nature of a large bodq":if ,rganic compounds, therefore, is approached oi.,y From the "positive hydrogen" viewpoint. It is pr&ably for this reason, for purposes of simplicity, tht teachers of organic chemistry who are attemptin? an application of the electronic theory, appear t r be showing an increasing tendency to neglect the ion resonance" idea. In the author's opinion, thi~''~'lnd is to be deplored for two reasons, the first of w@ is the fact that many interesting cases are encoun.' 5d in which the "positive hydrogen" viewpoint sheds 'le light on the problem of the comparative acid s t d 'dh of two different compounds hut where the "ion rl.bnance" approach suggests a clear-cut decision. It wr :ld be difficult, for example, to predict the relative .:id strength of toluene and benzene reasoning only 1 ,m the "positive hydrogen" principle. The "ion resonat.seM point of view dictates that a lateral hydrogen in t +me in should be more acidic than a nuclear hydro, r, pzene; that is actually the case, as shown hy me tlaI. ?n reactions (5). Further indication of the potential acidity of the methyl hydrogen in toluene is offered by thereported (6) condensations of fi-nitrotoluene in the presence of sodium ethylate and by the well-known reactivity to aldehydic reagents of the methyl group in ortho- and para- di- and tri-nitrotoluenes. In all such reactions of the Claisen and Knoevenagel type, the first step is postdated (16) as the removal by the basic catalyst of the acid hydrogen atom to form a carbanion which in the second step adds to the electrophilic system involved in the particular synthesis. Of the many other cases where the "ion resonance" approach permits a much more clear-cut decision concerning acidic nature than the "positive hydrogen" explanation, one of the most interesting is that of cyclopentadiene and many of its derivatives. The "surprising" acidity of the simple diene, as revealed by its positive Zerewitinoff test, its reaction with metallic potassium, and its condensations with aldehydes and ketones in the presence of sodium ethoxide, may be attributed to the fact that the ion resultingfrorn the loss of a proton in cyclopentadiene resonates among five structures all equivalent; in other words, the excess electron density in the ion is distributed equally over five carbon atoms. Teaching of the. "ion resonance" principle is important, in the second place, because there are a host of organic compounds in which the correct decisions concerning relative acidities may be reached only by the combined application of both points of view. This statement may seem a t first to be mere academic hair splitting. "What difference," a skeptic may ask, "is to be found between the concept that a substance is an acid because it possesses an electrophilic hydrogen which should show a pronounced tendency to accept the unshared electron pair of a base and the notion that a substance is an acid because the ion resulting from its discharge of a proton is stable and has little tendency to retrieve the proton?" The answer lies in the fact that acid strength is determined not by the nature of the ionogen, alone, or by the nature of the ion, alone, but rather by the free energy relationship between the two (4). Many examples can be cited to show that the application of the "positive hydrogen" principle alone does not lead to consistent results in predicting relative addities-comparisons of the hydrogen halides, of the primary, secondary, and tertiary amines, and of the alcohols and mercaptans, to mention but a few. In all of these cases involving electron displacements of the inductive type only, otherfactorsfrequently undetermined or not clearly understood are involved. There are many compounds exhibiting tautomeric displacements for which the "positive hydrogen" approach alone resdts in erroneous predictions concerning acidities; in most of these cases, however, an extension of the "ion resonance" concept obviates the difficplty. Viewed in the light of the "positive hydrogen" principle alone, the problem of naming the acid hydrogen in simple aldehydes is a poser. In the aldehyde molecule, a certain degree of acidity should result est number of atoms. Application of this principle is essentially tantamount to the explanation and prediction concerning the exact course of reactions of the type mentioned on the basis of the relative stabilities of the possible intermediates postulated by the mechanism. Conversely, comparisons of d i e r e n t compounds which form identical ions upon loss of a proton lead to interesting interpretations of many reactions. In the base-catalyzed allylic shift of 6,y unsaturated acids, esters, and lactones to the corresponding a,6 isomers, for example, the &st step appears to be the removal of a proton from the a-carbon atom. The resulting ion may be represented by three dzerent resonance structures showing the unshared electron pair (negative charge) on the a-carbon, the ?-carbon, and the oxygen atoms, respectively. Exactly the same ion would result from the loss of a proton by the 7-carbon atom of the corresponding a,g unsaturated compound. That ion, however, is more highly stabilized with respect to the 6,y unsaturated compound than it is with the a,@unsaturated isomer because the latter exists in more resonance forms than the former; in other words, the 6,y compound is a stronger acid than its a,6 isomer and its rearrangement may be looked upon simply as the formation of a weaker from a stronger acid under conditions favoring protolysis. Further evidence of the more acidic nature of the @,yunsaturated compounds is not lacking; A~~Y-angelicalactone condenses readily with benzaldehyde in the presence of aniline, for example, whereas the a,6 isomer does not (8). This principle may be extended to cover a wide variety of parallel cases. Prediction of the course a reaction should take on the basis of the relative stability of intermediates postulated by the mechanism is by no means valid only in the case of compounds which may be considered weak Bronsted acids. The principle has wide application in many types of reactions, a few of which are cited as examples. It would be expected that the alkaline hydrolysis of diethyl ethoxymethylenemalonate (7) would be initiated by the attack of the hydroxyl ion a t one of the electrophilic carbons of that compound. The intermediate formed by the addition of the hydroxyl ion to the carbon bearing the ethoxyl group is an ion of symmetrical codguration with complete resonance. It is not a t all anomalous, therefore, that the ethoxyl group should suffer direct and facile replacement by the hydroxyl group. Likewise, a comparison of the various ions which would be obtained by the attachment of the carbanion of malonic ester to d i ferent carbon atoms in the 6-vinylacrylic ester molecule suggests that attachment a t the 8-carbon atom should lead to the most stable intermediate, for only in this ion would the various resonance structures all I I - El=&, the ion resulting from the loss of a proton a t show displacement of the extra electron pair toward the group A is more stable, due to more extensive res- the primary electron acceptor, the carbonyl oxygen. onance, than an ion resulting from the loss of a proton It is to be expected, then, that the condensation (7) at any other point in the molecule. In other words, the of &vinylacrylic ester with malonic ester would occur negative charge on the ion formed upon removal of a in the 1,6 manner (Michael condensation). Consideraproton from A is most widely distributedover thegreat- tion of the entire principle of vinylogy from the standfrom the primary electron displacement, the -T effect at the carbonyl oxygen, which reduces the electron density on the carbonyl carbon; the positive charge on the latter is of course transmitted inductively to the a-carbon atom and on down the chain with progressively diminishing intensity. The thinking student who reasons that the effect should not be greater than the cause must conclude that the electron density of the hydrogen atom attached to the carbonyl carbon is lower than that of a hydrogen atom attached to the a carbon. The hydrogen of the carbonyl carbon is potentially acidic, i t is true, as indicated by the basecatalyzed condensation of formaldehyde to sugars and by the benzoin condensation of benzaldehyde. For simple aldehydes possessing a hydrogen atom on the a carbon, however, the first step of the mechanism postulated to explain such reactions as the aldol condensation and halogenation in basic solution is the removal of the proton by the basic catalyst, not from the carbonyl carbon, but from the a-carbon atom. This apparent ambiguity is readily explainable by the fact that only one reasonable electronic structure may be written for the ion resulting from the loss of the carbonyl hydrogen atom, whereas the loss of an a hydrogen produces a stabilized resonance hybrid of the keto-en01 type. This principle may be extended broadly to cover the behavior of a host of organic compounds which are said to react according to the principle of vinylogy (7). I n the base-catalyzed condensation of ethyl crotonate with ethyl oxalate, for example, the reaction must be initiated by the removal of a proton from the 7-carbon atom. Invariably, the inquiring student, again troubled by an effect greater than its cause, points out that the hydrogen on the a-carbon atom should be more "positive." He is not to be satisfied by the statement that vinylogous systems behave as they do because they are vinylogous; the empirical answer that, other things being equal, the tautomeric effect will operate more readily the more extended the conjugation offers little consolation to the inquisitive mind. It may be very clearly shown, however, that the ion formed by the loss of a proton from the 7-carbon atom of ethyl crotonate is stabilized with respect to the ion which would b e formed by a loss of the proton from the a-carbon atom, due to the f a d that the former may be represented by one more resonance structure than the latter. As a matter of f a d , this may be stated as a general rule designed to explain the behavior of compounds which illustrate the principle of vinylogy in reactions such as the Claisen, Knoevenagel, and malonic ester type condensations in which a carbanion is postulated as an intermediate: in compounds of the type A-(C--C)). 8 8 point of this general concept should prove both interesting and profitable to the individual teacher. Remick (9), in his excellent critical review of the status of the electronic theory as applied to organic chemistry, has indicated the inconsistencies in the present attemptsto account theoretically for the high p/o ratio usually found when benzene derivatives containing an ortho-para directing substituent are further substituted. A clue to the correct explanation may, perhaps, be found in the fact that the mechanism of para substitution in a benzene derivative possessing a substituent which can increase its covalence with the carbon of the ring (+T effect) requires the intermediate formation of an ion or group possessing the p-benzoquinoid structure, whereas for ortho substitution the intermediate is of the less stable o-benzoquinoid type. LITERATURE CITED (1) a. WATSON, H. B., "Modern Theories of Organic Chemistry." 2nd ed., Oxford Press, London, 1941, pp. 29-31; b. pp. 152-7. (2) HI,,, G.A. AND L. KELLEY,-organic chemistry," TI,^ Blakiston Company, Philadelphia. 1943, p. 615. (3) BRANCH~G. E. K., D. L.YABROFF.AND B BETIMAN,''The dissoc~ationconstants of chlorophenyl and phenetyl boric acids." I.Am. Chem. Soc.. 56, 939 (1934). (4) PAUL~NG, L., "The Nature of the Chemical Bond," 2nd ed., Cornell University Press, Ithaca, 1940, pp. 201-14. ( 5 ) G ~ ~ H., M-organic ~ ~ , chemistry," 2nd ed., johnwiley and Sons, Inc.. New York, 1943, Vol. I , p. 534. (6) REIssERr, A.. "Einwirkung von Oxalester und Natriumithylat auf Nit:otaluole. Synthese nitrirter Phenylbrenztraubensiuren, Ber.. 30, 1030-53 (1897). (7) PUSON. 1-27 R. (1935). C., "The principle of vinylagy," Chem. Revs., 16, (8) T ~ E L E'Yleber , ungesittigte y-Lactane,". Ann., 319, 147 (1901). (9) REMICK,A. E.,"El~troni~Interpretationsof Organic Chem- istrv." Tohn WilevandSons. Inc.. New York. 1943. D. 104. (10) FIESER; i.F., ref. (5). P. 159.
