Edited by RALPHK. BIRDWHISTELL
textbook forum
University of West Florida Pensacola, FL 32504
On a Common Graphical Flaw in the Carboxylic Acid/Alcohol Relative Acidity Argument Ronald W. McClard Reed College, Portland, OR 97202
One question that is always covered in introductory organic chemistry courses, and occasionally in first-year general chemistry courses, is why are carboxylic acids so much more acidic than aliphatic alcohols. I t is generally concluded that the simple explanation rests in the ability of the carboxylate ion to participate in resonance stabilization-a process apparently unavailable to an alkoxide ion. The importance of resonance in stabilizing anions has, however, been challenged quite recently (1). Solvation effects are rarely discussed even though the reactions are generally regarded as occurring in aqueous solution. Aside from whatever is the correct explanation for the acidity of carboxylic acids relative to alcohols, a problem arises in textbooks and lectures when the "energy profiles" for the dissociations of acid and alcohol are presented for the students' consideration. Figure 1depicts a common representation of the free energy vs. the progress of reaction. Thermodynamic acidity is correctly illustrated, of course, by the difference in height between the protonated (left) and unprotonated (right) states. The lower position (lower apparent free energy) of acetate relative to the isopropyloxide ion1 results in a less for acetic acid, consistent with the facts positive AGionization thus far. Often this difference is attributed incorrectly (as in Fig. 1) to the stabilization of acetate relative to a common transition-state energy for the alcoh01.~
This implies, of course, that the AGI for the recombination of an acetate ion with a proton is much less favorble than for an alkoxide ion. Although the picture in Figure 1 is a neat explanation, it ignores a large body of long-established, primarily kinetic information regarding proton transfer reactions ( 3 ) .If one plots the log of the dissociation rate constant of proton for a variety of mostly N and 0 acids versus the
lsopropyl alcohol is the best analogy to acetic acid for the reasons discussed by Siggel and Thomas ( l). See, for example, Morrison and Boyd (2).
Figure 1. Hypothetical free energy diagram for the dissociation of acetic acid and isopropanol. The bracket indicates the supposed stabilization of the acetate ion by resonance.
416
Journal of Chemical Education
Progress
-
Progress pK a of acid Figure 2. Relationship between the rate of proton dissociation and pKa for a variety of acids. Data are taken from Bender et al. (4) and references therein. The acids (in order of increasing pKa) are H30+, H2S04, HF, CH3COOH, imidazolium ion, H2S, (CH~)~NH+, glucose, and H20.
pKa of the acid, a straight line is obtained except where the rate approaches the diffusion-controlled limit for a strong acid like H 3 0 +(Fig. 2). That indicates that recombination of H+ and anion (or even uncharged base) is essentially isoenergetic and possibly a diffusion-controlled process. Indeed many reactions are even known to involve zero or negative free energy of activation (4). The table lists a number of such reactions; acetate ion and a wide variety of bases react at virtually diffusion-controlfed rates without regard to their pKa. Thus the barrier for the recombination of a proton with a carboxylate or alkoxide ion should be depicted as only the smallest bump. In order to be in accord with the evidence outlined above, it is then more reasonable to draw the diagram like that in Figure 3 when explaining the difference in acidity between alcohols and carboxylic acidse3Mention of the stabilization of acetate (relative to the solvated protonated species) by better solvation of the ion by water is made because of the recent work on gas-phase acidities of organic and inorganic molecules (see for example, refs 5-9). Absolute acidities expressed as AGO (298 K) for acetic acid and isopropanol are 341.5 kcal/mol (10) and 367.5 kcallmol (7), respectively. Although the differences between the absolute acidities of all candidate molecules are greatly attenuated by solvent water (7), it is clear that acetic acid is stronger than would be predicted and that isopropanol is only slightly less acidic than would be predicted (7). Thus it is reasonable to conclude that the lower position of the acetate ion on the free energy profile, relative to the alkoxide ion, is at least partly due to solvation. The resonance argument, although traditional and well accepted, has been cast into doubt. In the gas phase at least, charge on the isopropyloxide ion is actually dispersed slightly better than it is for acetate (1).Similarly, the acidity of phenol is often attributed to the ability of the phenoxide ion to produce four species in resonance. Delocalization of the negative charge into the aromatic ring, which is
The majority of texts depict this properly; however, reasoning for it is almost never provided.
-
Figure 3. Revised hypothetical free energy diagram for the dissociation of acetic acid and isopropanol. The bracket on the right indicates the relative stabilization of the acetate ion by solvation (3, and the bracket on the left indicates possible destabilization of acetic acid in the un-ionized form ( 1).Note that the barrier for recombination of proton and base (anion) is depicted as near zero for the reasons discussed in this paper.
Rates of Recombination of Bases with Proton at 298 K and Zero Ionic Strength Base
a
k,.,+(M-l~-l)~
pKa (conj. acid)
Data and their sources can be found in Eigen (4.
poorly solvated by water, is the probable cause of decreased acidity of phenol in H 2 0 (7, and references therein). Resonance stabilization of the phenoxide ion has also been questioned (1). Certainly the conveyance to chemistry students of the correct explanation for the phenomenon of acidity is very important and a graphical approach is highly useful. The picture suggested by Figure 3, and the considerations presented sometime ago by Calder and Barton (11) are potentially useful components of a complete discussion of this subject. Literature Cited 1. Siggel, M. R.; Thomas, T . D. J. Am. Chem. SOC.1986,108,4360. 2. Morrison, R. T.; Boyd, R. N . Organic Chemistry, 4th ed.; Allyn and Bacon: Boston; p 793. 3. Eigen, M. Angew. Chem. Intern. Ed. Engl. 1964,3,1. 4. Bender, M. L.; Bergeron, R. J.; Komiyama, M. The Bioorganic Chemistry of Enzyme Catalysis;Wiley:New York, 1984; p 20 (Fig. 2.1). 5. McMahon, T . B.; Kebarle, P. J. Am. Chem. SOC.1977,99,2222. 6. Mackay, G. I.; Bohme, D. K. J. Am. Chem. SOC.1978,100,327. 7. Bartmess, J. E.; Scott, J. A.; McIver, R. T., Jr. J. Am. Chem. SOC.1979,101,6056. 8. Fujio, M.; McIver, R. T., Jr.; Taft,R. W . J. Am. Chem. SOC.1981,103,4017. 9. Locke, M. J.; McIver, R. T., Jr. J. Am. Chem. SOC.1983,105,4226. 10. Cumming, J. B.; Kebarle, P. Can. J. Chem. 1978,56,1. 11. Calder,G. V.; Barton, T . J. J. Chem. Educ. 1971,48,338.
Volume 64
Number 5
May 1987
417
