The pk, values of simple aldehydes determined by kinetics of chlorination

Transcription

1 The pk, values of simple aldehydes determined by kinetics of chlorination J. PETER GUTHRIE' AND JOHN COSSAR Department of Chemistry, University of Western Ontario, London, Ont., Canada N6A 5B7 Received June J. PETER GUTHRIE and JOHN COSSAR. Can. J. Chem. 64, 2470 (1986). From the kinetics of chlorination of acetaldehyde and isobutyraldehyde we have been able to determine the pk, values of the free aldehydes (corrected for covalent hydration) as 16.9 t 0.5 and 15.7 t 0.5, respectively. These values are in reasonable accord with other work and show that our kinetic method for ketone pk,'s can be extended to aldehydes. For isobutyraldehyde the transition from rate-determining halogenation to rate-determining enolization occurs at [OCl-] values around M for hydroxide concentrations of 0.1 to 1. For acetaldehyde the transition would only occur at much higher [OCI-1. J. PETER GUTHRIE et JOHN COSSAR. Can. J. Chem. 64, 2470 (1986). En se basant sur la cinetique de la chloration de I'acCtaldChyde et de I'isobutyraldthyde, on a pu dcterminer que les valeurs des pk, des aldehydes libres (corrigees pour I'hydratation covalente) sont respectivement 16,9 t 0,5 et 15,7? 0,5. Ces valeurs sont en bon accord avec d'autres travaux et ces concordances dkmontrent que notre m6thode cinetique pour dcterminer des pk, de cctones peut &tre Ctendue aux aldehydes. Dans le cas de l'isobutyraldchyde, la transition entre une halogination qui determine la vitesse de la reaction et une Cnolisation qui dctermine la vitesse de la rcaction se produit a des valeurs de [OCl-] qui se situent autour de M pour des concentrations d'ions hydroxydes allant de 0,l a 1 M. Dans le cas de I'acCtaldehyde, la transition ne se produirait qu'8 des valeurs beaucoup plus Clevtes de [OCI-1. [Traduit par la revue] Introduction We have reported (1-3) a method for determining the pka of a simple ketone by analysis of the kinetics of alkaline 0 0- I I R1R2CH-C-R3 R1R2C=C-R3 + H+ halogenation. The kinetics are normally carried out with the halogenating agent in excess, under pseudo-first order conditions, so that u = kob, [ketone] kobs = k2 [halogen] total k2 = k20 + k2[oh-] The basis of the method is the evaluation of the term in the rate law which is first order in carbonyl compound and first order in halogenating agent, but zero order in base; this is expressed as k20. Such terms were first detected by Bartlett (4, 5). We have shown (1, 2) that these terms represent diffusion-controlled reaction of the enolate with the hypohalous acid, i.e. the conjugate base of one reagent and the conjugate acid of the other. The other term in the rate law, k2-, represents reaction of the enolate ion with the hypohalite ion. The pka of the carbonyl compound is then easily calculated from: [I] log k20 = log kdiff + p~~~~~ - pkak where kdiff is the rate constant for a diffusion-controlled reaction, p~ahox is the acid dissociation constant for the hypohalous acid, and p ~ a K is the acid dissociation constant for the keto tautomer of the carbonyl compound. We have examined the question of the best value to use for the rate constant of a diffusion-controlled process involving an unsymmetrical reagent such as a carbon-carbon double bond and proposed that the best value is log kdiff = (2). A value of this magnitude, kdiff = (2.8 * 0.2) X 1 O9 M- ' s ', has been reported recently for the reaction of bromine with acetophenone en01 (6). '~uthor to whom correspondence should be addressed I Recently Kresge has reported a pk, of i 0.04 for isobutyraldehyde (7), based on measurements of the rates of ketonization of the enolate and of base-catalyzed enolization. Kresge has also determined a pk, of i 0.06 for acetaldehyde (8, 9). In addition to our work on acetone (1-3) we have been investigating a number of other ketones in work now nearing completion, and have examined the question of whether our method can be applied to simple aldehydes. We wish now to report a determination of the pk, of acetaldehyde, and a less detailed examination of the behavior of isobutyraldehyde showing that our method is consistent with the results of Kresge and co-workers (7, 10). Results When the observed pseudo-first order rate constants for the chlorination of acetaldehyde with hypochlorite in excess are plotted vs. [OCl- 1, good linear relations are obvious; see Fig. 1. For some of the sets of data the best straight lines appeared to have non-zero intercepts. When straight lines arc fitted to the data by least squares it is found that the intercepts are small and show little sign of any simple dependence upon the hydroxide concentration. For some of the data sets it was noticed that the point at highest [OClp] falls below the line determined by the other points. This suggests that the first signs of curvature are being detected; the data do not, however, define a curve. The best data treatment appears to be the use of straight lines constrained to pass through the origin. As shown in Fig. 1, such lines give a reasonable description of the data, and correspond to the expected pattern. The slopes of these plots constitute the apparent second order rate constants for the chlorination reaction: when they are plotted against [OHp], as shown in Fig. 2, the trend is monotonic, but non-linear. It is, however, necessary to correct for both addition of water and ionization of the hydrate, as shown in Scheme 1. Good values for Kh (1 1) and Kh- ( )~ are available and permit calculation of corrected rate constants 'we thank Professor Kresge for con~municatlnp this result prior to its publication. 3 ~ h - = K,/K,; pk, = (12, 13).

2 p 1 GUTHRIE AKD COSSAR e k - p p L ----, E 4 a 1213 [[lcl-l FIG 1 Kinetics of chlorination of acetaldehyde at various hydrox- FIG. 3. Kinetics of chlorination of isobutyraldehyde at various lde concentrat~ons Successive offsets of 0 5 in lo3 k have been added hydroxide concentrations. The lines were determined by least squares to the rate constants so that the different data sets do not overlap (-) fitting: (E) points included in the least squares fitting; (0) points [OH-] = 0 026, (-.-.) [OH-] = 0 054, (-..-) [OH-] = 0 106, (--) omitted in the least squares fitting. [OH-] = 0 243, (--) [OH-] = (---) [OH-] = L p 0 a a 5! 0 i3h ' FIG 2 Pseudo-second order rate constants for the chlorination of acetaldehyde (A) kz,,,, the slopes derived from Flg 1, (0) kzlo,, the same values corrected for hydration and hydrate anion formation as described In the text corresponding to reaction of the free aldehyde. When these kc,, values are plotted against [OHp], as shown in Fig. 2, an excellent straight line results. with a slope of and an intercept of t This leads to a carbon pk, for the aldehyde of 16.8 t 0.5. One process which might lead to non-zero intercepts for plots of kobs VS. [OCl-] at fixed [OH-], corresponding to hypochlorite independent consumption of acetaldehyde, would be aldolization. However, the equilibrium constant for this process is only 400 M-' (14) and it would not proceed to a significant extent at lop4 M acetaldehyde. When the observed pseudo-first order rate constants for the chlorination of isobutyraldehyde are plotted vs. [OCl-] at fixed [OHp] a non-linear dependence is obvious, see Fig. 3, with leveling off at high [OCl-1. The data were fitted to eq. [2] by non-linear least squares. For each of two sets of data one point was clearly divergent from the line defined by the others; in the final calculations these two divergent points were omitted. The theoretical lines based on the least-squares parameters are shown in Fig. 3. In analyzing the hydroxide and hypochlorite dependence of these kinetics we make use of the mechanism shown in Scheme 1. In terms of this mechanism, the observed rate constant is given by eq. [3] [3] kobs={kl[ohpl/{l +Kh(l + Kh-[OH-l)))/ (1 + kpl/{k20~,/~~0c'[o~-] + k2)[oc1-1) Kh, the equilibrium constant for hydration of isobutyraldehyde, and Kh-, the equilibrium constant for ionization of the hydrate, are available from the work of Hine et al. (15). The quantity kl[ohp] can be extracted from the al values using these equilibrium constants. At M OHp k, [OH-] is t ; at M OH- kl[oh-] is ; at

3 CAN. J. CHEM. VOL TABLE 1. Kinetics of chlorination of aldehydesa [Aldehyde] X 10" [OH-] (a) Acetaldehyde (b) Isobutyraldehyde ' ' ' ' ~ ~ "All in aqueous solution at 25'C, ionic strength = 1.0 M (KC]). Data were fitted to y = A, + A, exp (-A,t) y = A + A, exp (-A3t) + A4t. Quantities in parentheses are estimated standard deviations of the parameters calculated by the least-squares program. The calculated standard deviations for A, and A, were C0.001 in all cases except the last entry in the table (where A, = 2.439) for which both were ca All reactions in cells with 1 cm path unless noted. b10 cm path length. '5 cm path length. d2 cm path length. or M OH- it is From these values kl = 0.145? was calculated, which is in excellent agreement with Kresge's value of (10). The apparent second order rate constant, which would be obtained directly at very low [OCl-1, can be calculated as al/a2. The values so obtained are t 0.85 ([OHp] = 0.034); 19.1 i 0.7 ([OH-] = 0.106); and 106 i 13.8 ([OHp] = 0.992); seefig. 4. Uncertainties in the apparent second order rate constants were calculated including the covariance between al and a2 (16). After correcting for hydration and hydrate anion formation we obtain 19.6 t 1.4, , and , respectively. Least-squares fitting gives k2 = (12.17 i 2.43) i 27.O) [OH-]. From the intercept we calculate (2) a pk, value of Kresge et al. reported (7). Discussion It is a striking observation that the kinetics of alkaline chlorination of acetaldehyde show almost no signs of a tendency to undergo a change in rate-determining step with increasing [OCl-] in the range covered by our experiments, while for isobutyraldehyde the change in rate-determining step is quite apparent, and indeed becomes nearly complete at low [OH-]. A major reason for the difference lies in the carbon acidities of the two carbonyl compounds, and in the fact that the rates of

4 GUTHRIE AND COSSAR base-catalyzed proton abstraction are in the reverse order compared to acidity. The pk, values are 16.8 and 15.7, and the rates of proton abstraction are 1.17 M-' s-' (9)2 and M-' s-' for acetaldehyde and isobutyraldehyde. As a consequence the rate of reprotonation of acetaldehyde enolate by water is about 100 times faster than for isobutyraldehyde enolate. Since partitioning between the two available paths determines which step is rate determining, a larger rate constant for reprotonation of the enolate will require a higher [OClp] to cause a change in rate-determining step, if the rates of halogenation of the enolates are the same. This is the case at low [OHp], where diffusion-controlled reaction with HOCl dominates the kinetics of halogenation. At high [OHp] the rate of halogenation is dominated by the contribution from reaction of OC1- with the enolate. These rate constants are not expected to be the same, and are well below the diffusion-controlled limit. Surprisingly, however, it seems unavoidable to conclude that the relative rates of the reactions of the enolates with OC1- are in the unexpected sense, with isobutyraldehyde being the more reactive. Thus the more hindered and the more stable enolate is the more reactive. The rate constants for the reactions of the enolates with OClp are 1.3 X lo3 for acetaldehyde and 9.5 x lo3 for isobutyraldehyde. A possible explanation for this unexpected order might be a transition state for the reaction of enolate with hypochlorite ion which resembles a charge transfer complex between the alkene and hypochlorite ion. The change in absorbance for each reaction provides a measure of the stoichiometry. Although there is considerable scatter in the values for both aldehydes, it seems clear that for acetaldehyde the stoichiometry is greater than 2 but less than 3. There is no clear dependence of the observed stoichiometry upon the hypochlorite concentration. For isobutyraldehyde the observed values scatter about the expected value of 1. Observation of a stoichiometry less than 3 for acetaldehyde suggests that for this compound as for acetone (3) hydrolysis of the intermediate partially chlorinated aldehydes is a major side reaction. Our method for the determination of pk, values of simple carbonyl compounds has now been shown to work for aldehydes as well as ketones. The major disadvantage of the method is that it relies upon an imprecisely known value of the rate constant for a diffusioc-controlled reaction. Although the uncertainties FIG. 4. Pseudo-second order rate constants for the chlorination of isobutyraldehyde: (A) k2app, the values derived from the data in Fig. 1; (0) kz,,,, the same values corrected for hydration and hydrate anion formation as described in the text. The line is calculated by weighted least squares. in this value will diminish as more absolute determinations become available, it will probably be found to vary for different compounds and/or reactions even though it is very insensitive to the chemical nature of the reaction occurring after diffusion. The major advantage of the method is that it uses standard, widely available, apparatus and inexpensive chemicals. Experimental Materials Isobutyraldehyde, Aldrich 98%, was distilled at atmospheric pressure and the middle fraction boiling at 62-63OC was used for kinetics. A stock solution, M, was prepared using glass-distilled water as solvent. Acetaldehyde dimethyl acetal, Eastman, was distilled at atmospheric pressure and the middle fraction boiling at 63-64'C was used for kinetics. Weighed amounts of the acetal were hydrolyzed in M HC1 (4 h at room temperature). neutralized to ph 8 (ph meter) using 1 M NaOH, and then made to a known volume with water. Two stock solutions, M and M, were used. Solutions containing sodium hydroxide and sodium hypochlorite were made up and titrated as previously described (2, 3). The ionic strength was held at 1 M using KC1. Methods Kinetics procedures were as previously described (2, 3). Reactions were initiated by injecting small volumes of an aldehyde stock solution into alkaline hypochlorite solution; the volume change was always less than 1 %. Calculations. Data were fitted to theoretical equations by non-linear least squares (16, 17), using computer programs written for the purpose. Acknowledgments We thank the Natural Sciences and Engineering Research Council of Canada, and the Academic Development Fund of the University of Western Ontario for financial support of this work. 1. J. P. GUTHRIE. J. COSSAR. and A. KLYM. J. Am. Chem. Soc. 104, 895 (1982).

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