J . Phys. Chem. 1991, 95, 345-352
345
Thermodynamics of Protonatlon of Ketones and Esters and Energies of Hydration of Their Conjugate Acids Alessandro Bagno,+ Vittorio Lucchini,* and Gianfranco Scorrano*.+ Centro CNR Meccanismi di Reazioni Organiche, Dipartimento di Chimica Organica, via Marzolo 1, 351 31 Padoua, Italy, and Dipartimento di Scienze Ambientali, Vniuersita' di Venezia, Dorsoduro 21 37, 301 23 Venezia, Italy (Received: March 30, 1990; In Final Form: June 19, 1990)
The protonation equilibria of methyl ketones RCOCH, (R = Me, Et, i-Pr, c-CJHS,c-C6HII,t-Bu, Ph, pMe0c6H4),symmetrical ketones RCOR (R = Et, i-Pr, c-C3HS,Ph), benzaldehyde, and methyl esters RCOOCH, (R = Me, Et, i-Pr, c-C3H5,c-C6HII, I-Bu, Ph) have been investigated in aqueous sulfuric acid in the temperature range 25-90 OC. The protonation parameters m* and pKBH+ have been derived, along with enthalpies and entropies of ionization. Structural effects on basicity are discussed in terms of the competition between internal and external stabilization by solvation. Solution- and gas-phase data have been combined to obtain the solvent effect on free energies and enthalpies of ionization and differential free energies (-bRAC',(BH+)) and enthalpies of hydration of the protonated bases. The values of m* for compounds with a similar substitution pattern around the basic site are correlated with -6RAC',,(BH+). Enthalpies of hydration of BH+ are compared with values predicted by empirical correlations.
Introduction An accurate knowledge of the protonation behavior of weak bases is required both for structure-reactivity correlations and for the detailed kinetic analysis of acid-catalyzed reactions.' However, difficulties derived from the failure of the analytical technique or of data treatment procedures have led in the past to a good deal of confusion and a spread of values for the protonation parameters of many weak bases2 A general formulation and treatment of protonation equilibria in nonideal acidic solutions is now available as the Bunnett-Olsen, Marziano-Cimino-Passerini, or Cox-Yates The protonation equilibrium of any weak base B is described in terms of its protonation constant pKBHt(hereafter pK) and a parameter ( m * ) that depends on the sensitivity of the equilibrium to the large medium changes occurring in the wide range of acid concentration needed to observe its protonation. As m* is related to the solvation of the protonated base," the trends in m* and pK observed in a series of structurally related bases can be interpreted in terms of the competition between internal (substituents) and external (solvation) factors on the stabilization of the positive charge in the protonated base.4 Difficulties in the interpretation of experimental data may still lead to questionable pK values, however. In fact, while the protonation parameters of many functional groups have been satisfactorily defined,2s4carbonyl bases have been repeatedly investigated with very confusing results (see, e.g., acetone, with published pK values ranging from -0.24 to -7.2).6 Both techniques normally employed for the study of protonation equilibria ( U V and NMR) are affected to some extent by the medium effects that accompany acid concentration changes, and various methods have been devised to correct for these.7 In two previous papers,'~~we examined the shortcomings of UV and NMR spectroscopies as applied to the study of ketones and esters. We found the following: (a) UV spectra of aliphatic ketones in aqueous sulfuric acid suffer from the combined influence of medium effects and side reactions, and the spectral changes thus produced cannot be separated from those due to protonation in any simple way; (b) N M R chemical shifts are also affected by medium effects, which must be corrected for by an appropriate internal standard. We concluded that UV spectroscopy generally is an inadequate technique for the study of aliphatic ketones. We could also show by NMR that simple aliphatic ketones' and estersB undergo the formation of a hydrogen-bonded complex with H30+, while aromatic esters behave regularly.8 The three species (free ketone, protonated ketone, and hydrogen-bonded complex) interconvert rapidly on the NMR time scale, and the observed
chemical shifts vary with medium acidity in a complex way. We were able to devise a procedure' for separating the contributions of the two equilibria and therefore for calculating the parameters for the pure protonation process. The determination of all the thermodynamic parameters of the protonation equilibrium (AGO, AH', AS'), both in solution and in the gas phase, allows one to describe the protonation process in more detail, to study the solvent influence on it, and to evaluate the energetics of ion s o l ~ a t i o n . Thus, ~ obtaining accurate protonation equilibrium data in solution is still of great importance, especially now that several methods for calculating hydration energies for neutral and charged species have been propo~ed,~*I~ and reliable data are needed for comparison. The m* values of carbonyl compounds span an unusually large interval (from 0.35 for acetone7 to 1.1 1 for 2,4,4'-trimethoxybenzophenone4), showing that the solvation of the protonated base is very sensitive to the nature of the substituents; thus this class of compounds is well suited for studying the competition between internal and external stabilization. To this purpose, we have studied a number of ketones and esters at various temperatures.
Results We have studied the protonation equilibria of eight methyl ketones RCOCH3 (R = Me, Et, i-Pr, c-C3H5, c-C6HII,t-Bu, Ph, p-MeOC6H4),four symmetrical ketones RCOR (R = Et, i-Pr, ( I ) Stewart, R. The Proton: Applications to Organic Chemistry; Academic Press: New York, 1985. (2) Bagno, A.; Scorrano, G.J . Am. Chem. SOC.1988, 110, 4577. (3) Cox, R. A.; Yates, K. Can. J . Chem. 1983, 61, 2225. (4) Bagno, A.; Scorrano, G.; More O'Ferrall, R. A. Reu. Chem. Intermed. 1987, 7, 313. (5) Cox, R. A. Acc. Chem. Res. 1987, 20, 27. (6) (a) Arnett, E. M.; Quirk, R. P.;Burke, J. J. J . Am. Chem. SOC.1970, 92, 1260. (b) Hine, J. J. Am. Chem. Soc. 1971, 93, 3701. (c) Levi, A,; Modena, G.; Scorrano, G. J. Am. Chem. SOC.1974, 96,6585. (d) McClelland, R. A.; Reynolds, W. F. Can. J. Chem. 1976,54,718. (e) Cox, R. A,; Smith, C. R.; Yates, K. Can. J. Chem. 1979, 57, 2952. (7) Bagno, A.; Lucchini, V.; Scorrano, G. Bull. Soc. Chim. Fr. 1987,563,
and references therein. (8) Bagno, A.; Scorrano, G.; Lucchini, V. C a n . Chim. Ira/. 1987, 117,475. (9) (a) Taft, R. W.; Wolf, J. F.; Beauchamp, J. L.; Scorrano, G.; Arnett, E. M. J. Am. Chem. SOC.1978,100, 1240. (b) Taft, R. W. Prog. Phys. Org. Chem. 1983, 14,247. (10) Ford, G. P.; Scribner, J. D. J. Org. Chem. 1983,48, 2226. (11) Meot-Ner (Mautner), M. J . Phys. Chem. 1987, 91, 417. (12) (a) Kang, Y. K.; Nemethy, G.; Scheraga, H.A. J . Phys. Chem. 1987, 91, 4105. (b) Ibid. 4109. (c) Ibid. 4118. (13) (a) Alagona, G.; Ghio, C.; Kollman, P. J . Am. Chem. Soc. 1986, 108, 185. (b) Jorgensen, W. L.; Briggs, J. M.; Gao, J. J. Am. Chem. Soc. 1987, 109,6857. (c) Rashin, A. A.; Namboodiri, K. J. Phys. Chem. 1987, 91,6003.
'Centro CNR Meccanismi di Reazioni Organiche. 'Universita' di Venezia.
0022-3654/91/2095-0345%02.50/0 0 1991 American Chemical Society 8-
346 The Journal of Physical Chemistry, Vol. 95, No. I . I991
c-C3HS,Ph), benzaldehyde, and seven methyl esters RCOOCH, (R = Me, Et, i-Pr, c-C3H5,c-C6HII,t-Bu, Ph) in aqueous sulfuric acid in the temperature range 25-90 OC. Ionization ratios I = [BH+]/[B] were calculated7 from eq 1 , where 0 is the observed quantity (chemical shift or absorbance) I = (0- OB)/(oBHt - 0 )
(1)
and fitted to eq 2, where log C H t and X are functions of acid log I - log C H t = m*X + pK (2) concentration and temperature. X (excess acidity) is defined as X = log (TB'YHt/yB*Ht), where B* is a reference base and the
coefficient^,^^^ and can be calculated as described elsewhere." m* and pK are obtained from the linear regression (2). Experimental points were eventually fitted to the calculated curve expressed by eq 3, obtaining a value for the sum yi terms are molar activity
+
@"' = (0, IOBH+)/( 1
+ I)
(3)
of squared errors U = Ci(OiCalC - OjcxP)z, and plotted against the acidity function Ho to detect systematic errors.' In most cases, the experimental points showed systematic deviations from the theoretical curve (eq 3). In these cases, equilibrium 4 was assumed to take place,7 where BCH+ is the hyBCH+ z B + H+
+ H,O
(4)
drogen-bonded complex between the base B and H30+. A complexation ratio I, was defined as I, = [BCH+]/[B], and the experimental chemical shifts were dissected into the individual contributions I and I, with a previously described method.7 The parameters of equilibrium 4 were derived from eq 5, where a, log I, - log C H t - log a, = m*J pK, (5)
+
is the activity of water and the theoretical curve is now expressed by eq 6, which gives a new value for the sum of squared errors, 0 = (0,+ I,OBCHt + I ~ B H + ) /+( ~IC I ) (6)
+
U'. The ratio 6U = lOO(U - U?/U is independent of the absolute value and number of data and thus may be used to compare the extent of deviation from eq 3 among different data sets. If no appreciable decrease of U occurred after fitting to the equations for hydrogen bonding (6U I0), the data set was considered free from other effects and treated according to eqs 2 and 3 only. As previously disc~ssed,~ the hydrogen-bonding treatment requires a good estimate of the three observables appearing in eq 6, which are eventually optimized. Since most measurements were carried out by NMR spectrometry, hereafter we will refer to the observable 0 as the chemical shift, thus defining the limiting parameters of eq 1 and 6 as us, uBH+,and uBCH+. (Of course, similar empressions apply to data obtained by UV, with the absorbance A in place of 0.) Since it was found that the values of V~ and V B H t are very little dependent on substituent and temperature, the same can be reasonably expected also for vBCH+,and care was taken that the optimization process converged to values lying in a comparable range. 'H NMR spectroscopy was used for all substrates except (cC,H&CO (also UV), Ph2C0, and PhCHO (UV). N M R chemical shifts were referred to internal trimethylammonium ~ u l f a t e . The ~ monitored groups were the CH3 signal CY to the carbonyl group for methyl ketones and esters and the a-CH, signal for dialkyl ketones except (i-Pr)zCO and (c-C3H,),C0 (0-CH, and /3-CH2, respectively). Temperature Variation of Acid Molarity, X Function, and Water Activity. To correctly use eqs 2 and 5 at any temperature, values of cH+.X,and log a, as a function of acid concentration and temperature must be supplied. To this purpose, the following procedures were applied (see Experimental Section for details). Acid molarity (Co) as a function of percent by weight (p) and temperature (t) is given by eq 7, where d is the density of the Co@,t) = IOpd(t)/Ma
(7)
solution in g mL-' and M a is the molecular weight of sulfuric acid.
Bagno et al. The temperature variation of the density of the sulfuric acid-water system is about 4% going from 0 to 100 O C for the most concentrated solutions.14 The temperature change of Co (which is proportional to d) was therefore expressed as suggested by Kell:I5 co(t) o: d ( t ) = (ao + a,t + ... + ast5)/(1 + bt) (8) The term C H t was evaluated as the root of eq 9, which states the mass and charge balance in the sulfuric acid-water system,' with the appropriate values for K, and the dissociation constant of HS04- (K,). CoKa/(K, + CHt) + K,/cH+ + co - C H t = 0 (9) Because of its physical meaning (see above), the X function at any acid concentration is expected to change with temperature according to X(T) = a/T b (10)
+
where T i s the absolute temperature. It has been proposed that X be defined so that X ( T ) = 298.15X0/T, where Xo is the value of the function at 25 0C.16 This implies b = 0 in eq 10. We tested this assumption with the available data7J7of Ho+ log C H t (= -X):,'* but in no case was b found negligible with respect to a/T. Thus no constraint was set on b. The activity of water was calculated from the extensive thermodynamic data by Giauque et aI.l9 for the sulfuric acid-water system in the 0-100% acid concentration range. The temperature dependence of In ywwas evaluated from partial molar heat contents and heat capacities. Aliphatic Ketones. As we previously reported, the chemical shifts of simple methyl ketones RCOCH, (R = Me, Et, i-Pr, t-Bu) as a function of the medium acidity do not strictly follow the sigmoid curve described by eqs 2 and 3; the deviations have been attributed to the presence of the hydrogen-bonding equilibrium, (eqs 4-6) and can be visually seen to decrease on going from R = Me to t-Bu; di-tert-butyl ketone displays no d e ~ i a t i o n . ~ For most bases studied in this work, an increase in temperature leads to a decreased deviation from the ideal sigmoid curve (smaller 6U values). The effect is more marked with increased size of R; thus, 6U for acetone is 88 and remains approximately constant with temperature, while for pinacolone it drops from 50 to 29 on going from 25 to 80 OC. In no case it was found to appreciably increase. However, in some cases ((i-Pr)$O, cC3H,COMe, (c-C3H,),C0, and c-C6HIICOMe(only at 60 and 90 "C)) both treatments (2) and (5) give the same results and 6U is zero at all temperatures. Therefore, we have applied the hydrogen-bonding treatment (eqs 4-6) to all substrates for which a positive 6U was obtained, and the normal one (eqs 1-3) to the others. Dicyclopropyl ketone was studied also by UV spectroscopy for comparison. The medium effect on the spectra was corrected by means of principal component analysis (PCA).-?" Values similar to NMR data were obtained (see Table I) despite the occurrence of a noticeable medium effect. The optimized chemical shifts and absorbances and protonation parameters at all temperatures are collected in Table 1. Aromatic Ketones. Some aromatic carbonyl compounds were also studied, both to further elucidate the behavior of this class of compounds and to obtain new data about the influence of the hydrogen-bonding equilibrium. The low solubility of these compounds in dilute sulfuric acid caused some difficulty in the detection of NMR signals; acetophenone can also form emulsions in such media, which may go visually undetected. As a conse(14) Infernational Critical Tables; McGraw-Hill: New York, 1930; Vol. 3, pp 56-57. (15) Kell, G. S . J . Chem. Eng. Data 1975, 20, 97. (16) (a) Cox, R. A.; Goldman, M. F.; Yates, K. Can. J . Chem. 1979,57, 2960. (b) Cox, R. A.; Yates K. Can. J . Chem. 1984,62, 2155. (17) Johnson, C. D.; Katritzky, A. R.; Shapiro, S. A. J . Am. Chem. SOC. 1969, 91. 6654. (18) Lucchini, V.; Modena, G.;Scorrano, G.; Cox, R. A,; Yates, K.J . Am. Chem. SOC.1982, 104, 1958. (19) Giauque, W. F.; Hornung, E. W.; Kunzler, J. E.; Rubin, T. R. J . Am. Chem. SOC.1960,82, 62. (20) Edward, J. T.;Wong, S . C . J . Am. Chem. SOC.1977, 99, 4229.
Protonation of Ketones and Esters
The Journal of Physical Chemistry, Vol. 95, No. I, 1991 341
quence, if relatively high concentrations of the substrate are employed, the measured chemical shift is not constant, as expected, in the low acidity region, but may show some random or systematic deviations that disappear, decreasing the substrate concentration. The resonances of aromatic and methoxy protons vary with acidity in an unpredictable way, probably the result both of substantial changes in electron density at the neighboring carbons and of inadequacy of the trimethylammonium ion as standard for these signal^.^ Furthermore, no aryl signals could be detected in the range 28-55% H2S04,where they overlap with the solvent peak. For all these reasons, the equilibria were followed monitoring the a-methyl signals. The protonation parameters are collected in Table I. Acetophenone and p-methoxyacetophenone behave rather differently. Similarly to dialkyl ketones, the former has 6U = 71 at 25 OC, with little change with temperature, while the latter is well behaved with 6U = 0. Benzaldehyde was studied by U V spectroscopy; the medium effect on the spectra was corrected for by the Davis-Geissman using the absorbance differences between 249 and 295 nm. Complexation occurs to an extent similar to acetophenone (6U = 71). Since benzophenone has only aromatic protons and has a very small solubility in aqueous media, it could only be studied by UV analysis, correcting the medium effect by PCA as before; 6U was ==O at all temperatures. The protonation parameters of PhCOMe, PhCHO, and Ph2C0 are in good agreement with those reported by Edward and Wong20 Esters. As reported above and in a previous work? the behavior of esters is generally similar to that of ketones. The deviations are comparable (e.g., 91 for methyl acetate and 79 for methyl pivalate at 25 "C). The changes with temperature are not very regular, which may reflect the lower accuracy of these measurements (at temperatures higher than 25 OC signal detection is difficult because of fast hydrolysis). Methyl cyclohexanecarboxylate shows an intermediate behavior, with a lower 6U ( 1 0-20). Methyl cyclopropanecarboxylate and methyl benzoate are well behaved at all temperatures (for the latter, care must be again taken to avoid the formation of emulsions). All data are reported in Table I. Thermodynamic Parameters and Energies of Hydration. pK values at the various temperatures were fitted to van't Hoff plots (pK vs 103/7"), yielding the enthalpies of ionization, AH'. The pK values at 298 K were converted to free energies of ionization (AGO = 2.303RT(pK)), allowing the calculation of the entropies of ionization at 298 K as ASo= (AHo- A G o ) / T (the same values were obtained from the intercepts of the van't Hoff plots). All data are collected in Table I. Values of AGO, published gas-phase basicities (GB):2 and free energies of hydration (AGoaq(i))of neutrals23 were eventually combined by means of a thermodynamic cycle9to yield the solvent effect on the free energies of ionization (6,,AGo) and the differential free energies of hydration for the free and protonated bases, relative to NH3 and NH,+ (6RAC0,,(B) and 6RAGoa,(BH+), respectively) :9 6,,ACo = AGOg - AGO,, 6RAGoaq(B)= AGoa,(B) - AGoaq(NH3)
(11)
(12)
-bRAGoaq(BH+)= -(AGoaq(BH+) - AGoaq(NH4+))= ba,AGo - 6RAGo,,(B) (13) The analogous enthalpy terms (?iaqAHo,6,AH0,,(B),
and
Davis, C. T.;Geissman, T.A. J. Am. Chem. SOC.1954, 76, 3507. (a) Lias, S. G.:Liebman, J. F.; Levin, R. D. J. Phys. Chem. ReJ Data 3,695. (b) Personal communication from R. W. Taft, quoted in the 1987 informal update of the same compilation. (23) (a) Hine, J.; Mookerjee, P. K. J. Org. Chem. 1975.40, 292; (b) Lam, S.Y.; Benoit, R . L. Can. J . Chem. 1974, 52, 718. (c) Cabani, S.;Gianni, P.; Mollica, V.;Lepori, L. J . Solul. Chem. 1981, 10, 563. (d) Abraham, M. H.; Whiting, G. S.; Fuchs, R.; Chambers, E. J . J . Chem. SOC.,Perkin Trans. 2 1990, 291.
6RAHoa (BH+)) were calculated in the same way from AH', proton affinities (PA):2 and enthalpies of s o l ~ t i o n ?Entropy ~~~~~ terms were calculated from all the above data. These values are reported in Tables I1 and 111, which also include some literature data for comparison. Given the accuracies of the published values of GB and PA (2 kcal mol-'), AGoaq(B),and AHo,,(B) (0.2 kcal mol-') and of the parameters for aqueous ionization in Table I, we can give typical errors on -bRAGoaq(BH+)and -6,AH", (BH') as ca. 2 kcal mol-' and on 6RASoaq(BH+)as ca. 9 cal m o b K-I.
Discussion Physical Significance of m*. The m* parameter appearing as the slope of the linear regression (2) is defined by the empirical relationship (14).3.4 log
(YBYH+/YBH+)
= m* log
(YB*YH+/YB*H+)
= m*X
(14)
Activity coefficients represent free energies of transfer from the standard state to nonideal acidic solutions and are dominated by solvation effects4 It is easy to show that if YBH+/YB >> I , then 0 < m* < 1 and, conversely, m* > 1 when yBH+/yB = 1. The first case applies to when there is a large difference In solvation between B and BH+, on a scale relative to B*/B*H+ (an aniline), and the second when the solvation of B and BH+ is s i ~ n i l a r . ~ . ~ ~ Since solvation energies of charged species are higher than of neutral^,"^ m* essentially measures the hydrogen-bonding solvation of the protonated base. Accordingly, experimental values of m* increase, for different protonation sites, as ROH < R 2 0 < R 2 C 0 < R2S0 < ArNH, < R2S < Ar3C+.4 Structural and Temperature Effects on Basicity and Solvation. The reliability of the p R s of the weak bases studied in this work was affected, in the first place, by the nonlinear fitting procedure, which is approximate in itself.7 Therefore we emphasize that small differences, especially in the thermodynamic parameters, must be considered with caution. Moreover, while the m* and pK values show a good correlation with the nature of the substituents (see below), the corresponding data for the hydrogen-bonding equilibrium (m*c and pKc) are much more erratic and are largely dependent on the chemical shift vBCH+, which cannot be estimated with the desirable accuracy, as the nature of the species involved cannot be firmly e~tablished.~For these reasons, we will not attempt a discussion of these parameters but will consider only the values of [BCH+], Le., the largest calculated percentage of complex. We propose that this parameter might be taken as an indicator of the strength of hydrogen bonding between the solvent and the free base B. The data of m* and pK at 25 OC define a trend consistent with, and supporting, previous finding^.^ With the exception of cyclopropyl derivatives (see below), m* increases with alkyl substitution at the carbonyl group; aryl substituents cause a further marked increase. The same trend is found for esters, though with higher values. These data are consistent with the interpretation of m* as a solvation parameter$~7~8*'8~25 which expresses the strength of the hydrogen-bonding interaction between the solvent and BH+. The trend for pRs is similar, the compounds becoming less basic in the standard state (lower pK) with increased s u b ~ t i t u t i o n . ~ ~ ~ Thus the general effect of alkyl substitution at the basic site is 2-fold: it decreases external stabilization of BH+ through solvation (as probed by m*) in favor of internal stabilization through interaction with the substituents! If the changes are small (as within a family of bases), these two effects may largely compensate each other; the practical consequence is that ketones and esters with the same substitution pattern are ionized (in terms of log I ) to a similar extent.8 The hydrogen-bonding equilibrium, though difficult to analyze in detail, also follows a recognizable pattern. In fact, the values of [BCH+] among different compounds at the same temperature show that substituent branching near the basic site cause a marked decrease of this phenomenon, to the point that all dialkyl ketones (24) Della Gatta, G.; Stradella, L.; Venturello, P. J . Soluf. Chem. 1981, 10, 209.
(25) Perdoncin, G.;Scorrano, G . J. Am. Chem. SOC.1977, 99, 6983.
348
Bagno et ai.
The Journal of Physical Chemistry, Vol. 95, No. 1 , 1991
TABLE I: Chemical Shifts,' Protonation Data, and Thermodynamic Parametersbfor Ketones and Esters [BCH']'
bU
AGO
AHo
Aso
(0.03) (0.02) (0.02) (0.02)
43 42 39 25
88 87 88 87
-4.18 (0.04)
0.2 (0.3)
14.3 (0.8)
-2.16 -2.16 -2.31 -2.17
Ethyl Methyl Ketone 0.40 (0.01) -3.48 (0.03) 0.41 (0.01) -3.45 (0.02) 0.43 (0.01) -3.37 (0.02) 0.42 (0.01) -3.19 (0.01)
24 16 16 16
89 86 80 80
-4.75 (0.04)
-3.0 (0.2)
6.2 (0.6)
-2.31 -2.31 -2.47 -2.58
Isopropyl Methyl 0.42 (0.01) -3.63 0.42 (0.01) -3.52 0.46 (0.01) -3.57 0.47 (0.01) -3.49
21 18 14
81 73 72 64
-4.96 (0.06)
-0.9 (0.4)
13.4 (1.2)
-4.03 (0.11)
0.9 (0.6)
17.2 (1.8)
m**
~gcHi-'
m*,g
pK,r
3.53 3.62 3.35 3.39
-34. I -35.3 -35.6 -33.4
0.40 0.39 0.39 0.38
-1.65 -1.55 -1.58 -1.81
Acetone 0.35 (0.01) -3.06 0.34 (0.01) -2.91 0.36 (0.01) -2.89 0.39 (0.01) -3.03
-42.17 -42.30 -42.36 -42.84
I .66 1.64 1.51 1.80
-29.9 -26.9 -29.2 -3 1.6
0.43 0.38 0.44 0.40
25.4 39.8 60.1 79.9
-40.81 -41.10 -41.51 -41.83
2.18 2.23 I .64 I .60
-30.0 -30.1 -25.9 -25.4
0.45 0.44 0.46 0.47
25.0 40.0 60.0 90.0
-61.20 -62.31 -62.92 -63.88
-1 4.96 -16.52 -1 7.43 -1 8.29
25.0 40.0 60.0 90.0
-42.00 -42.26 -41.57 -42.55
0.64 0.2 1 I .05 I .32
-32.4 -36.3
0.35 0.42
-1.92 -1.87
24.6 39.9 60.3 80.1
-40.46 -40.68 -41.10 -41.58
3.12 2.90 2.76 2.65
-33.4 -30.5 -32.7 -28.5
0.40 0.36 0.34 0.34
-2.19 -2.21 -2.08 -2.26
tert-Butyl Methyl 0.40 (0.01) -3.48 0.42 (0.01) -3.54 0.44 (0.01) -3.46 0.46 (0.01) -3.44
24.9 40.4 60.6 80.3
-45.94 -46.47 -48.49 -50.12
64.66 63.73 62.80 61.64
-34.0 -34.9 -33.4 -36.2
0.56 0.55 0.54 0.56
-2.19 -2.16 -2.17 -2.14
Acetophenonek 0.60 (0.02) -3.87 (0.08) 0.62 (0.01) -3.83 (0.07) 0.63 (0.01) -3.76 (0.05) 0.66 (0.01) -3.70 (0.05)
25.2
-55.24
26.12
25.0 40.0 60.0 90.0
-31.97 -32.66 -33.36 -35.07
27.38 26.90 26.48 24.99
25.0 40.0 60.0 90.0
-166.81 -166.83 -166.77 -167.14
- I 34.1 7 -1 34.35 -1 34.35 -1 34.76
25.0 25.0 40.0 60.0 90.0
-164.55 0.038' -163.47 -164.13 -164.67
25.0 40.0 60.0
-0.031 -0.028 -0.032
25.0 40.0 60.0
0.44 0.40 0.40
-0.84 -0.82 -0.80
0.2 0.2 0.2
0.47 0.57 0.55
-2.38 -2.96 -2.99
Benzaldehyde' 0.54 (0.01) -4.48 (0.05) 0.56 (0.01) -4.52 (0.06) 0.60 (0.01) -4.52 (0.05)
25 18
25.0 40.0 60.0
46.35 46.28 46.10
87.34 88.71 87.63
52. I 50. I 51.9
0.57 0.49 0.50
-2.41 -2.02 -2.05
Methyl Acetate 0.46 (0.01) -3.90 (0.03) 0.39 (0.01) -3.35 (0.02) 0.47 (0.01) -3.66 (0.04)
7p
vgd
25.2 40.1 60.0 80.2
-40.64 -40.80 -41 .I4 -41.32
25.0 39.7 60.3 80.1
vgH+'
pKh
Ketone (0.04) (0.04) (0.03) (0.03)
Cyclopropyl 0.38 (0.01) 0.48 (0.01) 0.52 (0.01) 0.55 (0.01)
Methyl Ketond -2.95 (0.07) -3.23 (0.06) -3.24 (0.03) -3.15 (0.05)
Cyclohexyl 0.42 (0.01) 0.42 (0.01) 0.41 (0.01) 0.45 (0.01)
Methyl Ketone -3.55 (0.03) -3.32 (0.03) -3.25 (0.05) -3.26 (0.03) Ketone (0.04) (0.04) (0.04) (0.04)
11
20 32
71 78
-4.84 (0.04)
-1.9 (0.3)
9.4 (1.0)
18 12 13 9
50 41 32 29
-4.74 (0.05)
-0.5 (0.5)
14.3 (1.4)
45 40 34 36
71 73 71 72
-5.28 (0.11)
-1.5 (0.7)
12.6 (2.2)
-5.29 (0.04)
-2.3 (0.3)
9.9 (0.9)
-5.80 (0.06)
-1.8 (0.5)
13.5 (1.6)
-3.27 (0.05)
1 . 1 (0.4)
14.8 (1.4)
-6.43 (0.11)
-3.3 (1.2)
10.5 (3.7)
-6.11 (0.07)
0.5 (0.9)
22.3 (2.9)
11
71 30 17
46 41 38
91 86 76
-5.32 (0.04)
-4.5 (0.6)
1.8 (2.1)
p-Methoxyacetophenonek 0.60 (0.02) -3.17 (0.07) -1 8.0 -20.6 -19.7 -20.0
0.39 0.48 0.62 0.57
-1.89 -2.34 -3.16 -2.59
Diethyl Ketone' 0.44 (0.01) -3.88 (0.03) 0.45 (0.01) -3.74 (0.03) 0.46 (0.01) -3.68 (0.03) 0.50 (0.01) -3.56 (0.03)
-4.33 (0.10) 30 24 15 22
82 69 38 59
Diisopropyl Ketone' 0.45 (0.01) -4.25 (0.04) 0.47 (0.01) -4.18 (0.06) 0.49 (0.01) -4.14 (0.05) 0.52 (0.01) -4.00 (0.06)
-82.60 0.355' -82.1 5 -82.06 -82.75
Dicyclopropyl Ketone' 0.41 (0.01) -2.40 (0.04) 0.40 (0.02) -2.60 (0.08) 0.44 (0.01) -2.46 (0.04) 0.47 (0.01) -2.55 (0.06) 0.53 (0.01) -2.57 (0.05)
0.87 0.87 0.84
Benzophenone' 0.67 (0.01) -4.71 (0.08) 0.67 (0.02) -4.56 (0.09) 0.70 (0.01) -4.45 (0.05)
The Journal of Physical Chemistry, Vol. 95, No. 1, 1991 349
Protonation of Ketones and Esters
TABLE I (Continued) F
rigd
vgH+'
uBCHJ
m*cg
PKC*
[BCH+]' bU
pKh
m*h
As'
AH'
AGO
Methyl Propionate 25.0 40.0 60.0
46.79 46.39 47.15
88.79 89.19 88.66
56.1 55.7 52.3
0.54 0.49 0.60
-2.57 -2.23 -2.64
0.51 (0.01) 0.53 (0.01) 0.55 (0.01)
25.0 40.0 60.0
70. I2 70.55 70.20
132.46 132.95 133.06
80.2 76.0 79.4
0.52 0.56 0.49
-2.50 -2.66 -2.50
0.57 (0.01) 0.54 (0.01) 0.57 (0.01)
25.0 40.0 60.0
71.28 69.67 69.52
120.43 121.12 122.01
25.0 40.0 60.0
45.3 49.0 47.7
88.2 88. I 88.3
62.9 56.0 55.6
0.52 0.54 0.57
-2.69 -3.3 I -3.14
0.59 (0.01) 0.62 (0.01) 0.60 (0.01)
25.0 40.0 60.0
46.56 46.78 46.67
89.93 90.06 89.89
58.2 56.0 55.9
0.62 0.67 0.70
-3.13 -3.22 -3.16
0.77 (0.02) 0.77 (0.01) 0.78 (0.02)
25.0 40.0 60.0
6 I .60 61.35 61.53
98.22 98.34 99.40
-4.37 (0.04) -4.27 (0.03) -4.12 (0.03)
30 32 33
80 82 88
-5.96 (0.06)
-3.3 (0.6)
9.0 (2.0)
32 28 20
64 75 55
-6.70 (0.10)
-1.4 (0.8)
16.2 (2.5)
-6.50 (0.12)
-4.9 (0.8)
4.6 (2.5)
Methyl Isobutyratd -4.91 (0.07) -4.45 (0.03) -4.54 (0.03)
Methyl Cyclopropanecarboxylate' 0.58 (0.01) 0.55 (0.01) 0.56 (0.01)
-4.76 (0.09) -4.34 (0.04) -4.21 (0.02)
Methyl Cyclohexanecarboxylate -4.92 (0.07) -4.97 (0.03) -4.60 (0.06)
22 7 11
18 9 7
-6.71 (0.10)
-5.0 (1.2)
6.5 (3.7)
30 33 38
79 89 87
-8.95 (0.21)
-7.0 (2.4)
6.3 (7.7)
-9.62 (0.22)
-13.4 (3.3)
-12.6 (10.6)
Methyl Pivalate -6.56 (0.15) -6.21 (0.08) -5.99 (0.12)
Methyl Benzoate 0.82 (0.02) 0.81 (0.02) 0.75 (0.03)
-7.05 (0.16) -6.68 (0.12) -6.00 (0.20)
In hertz at 60 MHz, except where noted. Positive values denote downfield shifts from trimethylammonium ion. bThermodynamicparameters for aqueous ionization process. Energies in kcal mol-', entropies in cal mol-' K-I; errors from least-squares regression follow in brackets. cf0.4 'C. "Chemical shift of the free base. 'Chemical shift of the protonated base. fChemical shift of the hydrogen-bonded complex (optimized values). 8ComDlcxation Darameters lea 5). Protonation aarameters: errors are from least-squares regression (eq 2). Maximum calculated percent of hydrogen-bonded complex. ;9b MHz. 200 MHz: 'UV absorbances
(including di-tert-butyl ketone7) except diethyl ketone are not affected at all. The results for acetophenone (in contrast to p-methoxyacetophenone and methyl benzoate) show that the lone pairs on the carbonyl group may still be available for hydrogen bonding dcspite the conjugation with the phenyl ring. Thus the major factors influencing this equilibrium are steric hindrance and, to some extent, charge delocalization. None of the cyclopropyl derivatives is affected by the hydrogen-bonding cquilibrium. While the result for dicyclopropyl ketone is comparable with diisopropyl ketone and in line with steric considerations, methyl cyclopropyl ketone and methyl cyclopropanecarboxylate behave differently from their i-Pr analogues. This can be explained considering that these compounds prefer almost exclusively an eclipsed s-cis conformation in which the carbonyl group is above the plane of the ringz6 Because of this peculiar conformation, the effective steric hindrance around the carbonyl group is much higher than expected, and there might also be a stabilizing interaction of the ring with the oxygen lone pairs. This rigid conformation is retained in the protonated form too;26since the acidic proton is normally cis to the less hindered side of the carbonyl it is also constrained to lie in a position particularly exposed to the solvent, which accounts for the increased aqueous basicity and solvation relative to a similar substitution pattern, such as i-Pr and c-C6HII. For dicyclopropyl ketone the conformation is similar, with the carbonyl group s-cis to both rings.z8 However, despite the expected rigidity of this system, there is a 100-fold increase of aqueous basicity relative to diisopropyl ketone, which is entirely due to the enthalpy term (see Table I), the origin of which is not clear. Anyway, it is interesting to note that all data indicate that the cyclopropyl ring does not particularly stabilize the protonated base. (26) Richey, H. G., Jr. In Carbonium Ions; Olah, G . A,, Schleyer, P. v. R.. Eds.; Wiley-Interscience: New York, 1972; pp 1201-1294. (27) Olah, G. A.; White, A. M.; OBrien, D. H . Chem. Rev. 1970, 70, 561. (28) (a) Aroney. M.J.; Calderbank, K. E.; Stootman, H. J. J . Chem. Soc., Perkin Trans. 2 1973. 1365. (b) Okazawa, N.; Sorensen, T. S.Can. J. Chem. 1978, 56, 2355. (c) Nease, A. 8.; Wurrey, C. J. J . Phys. Chem. 1979,83, 2135. (c) Schrumpf, G.;Alshuth, T. Spectrochim. Acra 1987, 43A, 939. (d) Diallo, A. 0.;Waters, D. N. Spectrochim. Acta 1988, 44A, 1109.
-8
-
-10
-
i"
-12.
