3142
A. V. WILLI,C. M. WON,AND P. VILK
Kinetics and Mechanism of the Decarboxylation of Anthranilic Acid in Aqueous Solution’
by Alfred V. Willi, Chong Min Won, and Paul Vilk2 The College of Pharmaceutical Sciences, Columbia University, N e w York, N e w York 100$6 (Received January 17, 1968)
The kinetics of decarboxylation of anthranilic acid (H-4) has been studied at 70 and 85” in weakly acidic solutions (pH 4.7-1.4) of a constant ionic strength of f i = 0.1 N , and at 85” in 0.1-3.9 N aqueous hydrochloric acid. The results in weakly acidic solutions supply evidence for two parallel reactions of H30+ with HA and of H30+ with A-. The mechanisms of both pathways involve predominant rate-determining proton transfer from &of to the carbon atom in the position 1 of the aromatic ring. I n strongly acidic solutions, above [H30+]= 0.5 N , the pseudo-first-order rate constant, k , decreases with increasing [H30+],which indicates a change of mechanism. A quantitative treatment of the pH dependence of k has been carried out on the basis of the stationary state principle and with the aid of the Hammett acidity function. For sufficiently high [HaO+],there is a linear relationship between l / k and b with a high precision. It may be derived solely from a mathematical analysis of the experimental rate data that in weakly acidic solutions the total rate is mainly governed by the rates of the proton-transfer steps; while with increasing acidity of the solution, a gradual change takes place to complete rate control by the C-C bond-cleavage steps. Introduction The decarboxylation kinetics of aromatic acids with electron-releasing substituents has been studied in some detail by several research worker^.^-^ On the basis of the empirical rate law (eq 1) found in aqueous solutions, the reactions could be formulated either as first-order decompositions of the unionized acid HA or as secondorder reactions of the anion A- with the hydroxonium ion6t6 rate = ~ H A [ H A=] k~*[A-][Hao+]
(1)
In the examples of p-aminosalicylic and p-aminobenzoic acids, decarboxylation proceeds via two parallel pathways according to the rate equations rate = /cH~A[H~A+] -I- ~ H A [ H A=] ~ H ~ * [ H A I [ H ~ 4O +~H*[A-][HL)+I I (2) Experimental evidence for mechanism A - S E ~ was supplied with the aid of data on substituent effects, general acid catalysis, and solvent isotope effects. It was concluded that the slow step involves a proton transfer from H30+to the carbon atom in position 1 of the aromatic ring of either the acid, HA, or the anion, A-.6j8 The two-step nature of the mechanism was first established by Lynn and bourn^,^ who measured carbon-13 isotope effects in the decarboxylation of 2,4dihydroxybenzoic acid in various acetate buff ers. A new development was indicated by the findings of Dunn, Leggate, and Scheffler,’ who observed that the decarboxylation rates of p-methoxy- and p-methylanthranilic acids decrease with increasing [Ha0+] for hydroxonium ion concentrations above 0.1 N . In this pH range, however, the substrate is not involved in any further acid-base equilibrium beyond the addition of a The Journal of Physical Chemistry
proton to the amino group. Similar results have been reported recently by Los, Rekker, and Tonsbeel? for a few 2-substituted 4-aminobenzoic acids. These findings support a change of the slow step in the mechanism, from rate-determining proton transfer to rate-determining C-C bond c l e a ~ a g e . ~ , ~ Though experimental results are available for paminobenzoic acid as well as for a few substituted anthranilic no previous work has been published on the decarboxylation kinetics of unsubstituted anthranilic acid in aqueous solutions of various pH values. Since the absence of a carbon isotope effect has been established for this reaction in 1 N (=0.5 M ) , aqueous sulfuric acid,1° it was desirable to obtain information on (1) The major portion of this material has been taken from a thesis submitted by C. M. Won to the College of Pharmaceutical Sciences, Columbia University, New York, N. Y., in partial fulfillment of the requirements for the Master of Sciences degree, 1967. (2) Undergraduate research participant at the College of Pharmaceutical Sciences, Columbia University, New York, N. Y., 1967. (3) For a review of earlier work, see A. V. Willi, “SLurekatalytische Reaktionen der organischen Chemie-Kinetik und Mechanismen,” F. Vieweg und Sohn, Braunschweig, West Germany, 1965. (4) W. M. Schubert, J . Amer. Chem. Soc., 71, 2639 (1949); W.M. Schubert, J. Donahue, and J. D. Gardner, ibid., 76, 9 (1954); W. M. Schubert, R. E. Zahler, and J. Robins, ibid., 77, 2293 (1956). (6) W. M. Schubert and J. D. Gardner, J . Amer. Chem. Soc., 75, 1401 (1963). (6) A. V. Willi and J. F. Stocker, Helv. Chim. Acta, 37, 1113 (1964); A. V. Willi, ibid., 40, 1053 (1957); 43, 644 (1960); 2.Naturforsch., l3a, 997 (1958); Trans. Faraday SOC., 5 5 , 433 (1959); 2. Physik. Chem. (Frankfurt), 27, 221 (1961). (7) G. E. Dunn, P. Leggate, and I. E. Scheffler, Can. J . Chem., 43, 3080 (1966). (8) R. F. Rekker and W. Th. Nauta, J . Med. Pharm. Chem., 2, 281 (1960); J. M. Los, R. F. Rekker, and C. H. T. Tonsbeek, Rec. Trav. Chim. Pays-Bas, 86,622 (1967). (9) K. R. Lynn and A. N. Bourns, Chem. I n d . (London), 782 (1963). (10) W. H. Stevens, J. M. Pepper, and M. Lounsbury, Can. J . Chem., 30,629 (1952).
DECARBOXYLATION OF ANTHRANILIC ACIDIN AQUEOUS SOLUTION the rate equation in order to supply sufficient evidence to establish a mechanism. It was expected that a change of mechanism in strongly acidic solutions would occur also in this example and that this reaction might be well suited for a complete mathematical analysis of the kinetic data in the region of the change of mechanism. A detailed kinetic study of the decarboxylation of anthranilic acid in aqueous solution has been carried out and the results are reported in this article. Experiments have been done in weakly acidic solutions at a constant ionic strength of p = 0.1 N , as well as in 0.1-3.9 N aqueous hydrochloric acid solutions.
Experimental Section The ionization constant K1 (definition below) of anthranilic acid at 70 and 85” and at a constant ionic strength of 0.1 N was determined by potentiometric measurements in a thermostated cell without liquid junction with a glass electrode Metrohm Type H and a silver-sjlver chloride electrode. pcH values” were measured with a Metrohm Type E388 pH meter. Kinetic measurements were carried out as described previouslye by following the decrease of the uv absorption at 310 rand 315 mp. The concentration of the substrate in the kinetic solutions (in the ampoules) was 3 X M , The solutions contained 0.9% ethanol which had been used to dissolve the organic acid. Samples of the solutions were neutralized with NaOH solution, were buffered to a pH of 7.8, and were diluted fivefold before measuring the optical densities (D)in 5-cm quartz cells with a Beckman DU spectrophotometer. (Neutralization and buffering were omitted in the experiments with solutions in dilute hydrochloric acid ( 2 0 . 1 N ) , and optical densities were measured at 275 and 280 mp.) Ampoules had been treated with 2 N hydrochloric acid for 4 weeks before use. A stream of nitrogen was passed through the kinetic solutions for the “s1ow” experiments with acetate buffers, before filling them into ampoules. In all kinetic experiments, log (D - D,) decreased linearly with time, indicating that the reaction followed a first-order course at a constant pH. The reactions were followed through 2-3 half-lives. D, was obtained from the optical density after 10 half-lives; usually there was good agreement with the theoretical value (obtained from the uv spectra of aniline). For the very slow reactions, the theoretical value of D, was used. Pseudo-first-order rate constants were calculated from the experimental data with the aid of the least-squares method. Their estimated precision is within *2%. Results Ionization Constants. It is necessary for a discussion of the pH dependence of the rate to have available
3143
values of the ionization constants K1 and KO of anthranilic acid [H30 1 [HAI/ [ H A +
+
[H&+l [A-l/[HAl
1
KO Ki
Determinations of these constants are to be done at the same temperature and ionic strength as the kinetic measurements. Results for Kl from potentiometric determinations are K1
= (2.53 f 0.12) X
K~ = (2.81
f
0.10)
x
10-5
(70°,p
= 0.1
(850, cc
=
N)
0.1 N )
(The accuracy of the pK determinations at high temperatures is discussed in the Appendix.) K Ois in the order of magnitude of Under these circumstances it is not possible to obtain reliable values with the potentiometric method. Rate Constants. Experimental results of pseudofirst-order rate constants, k, for various hydroxonium ion concentrations at a constant ionic strength of p = 0.1 N are given in Table I. The data exhibit a similar pH dependence as found in the decarboxylation of p-aminobenzoic acid: k increases parallel to an increasing concentration of the cation, HzA+, in the acid-base equilibrium of the amino acid. The experimental values can be fitted into eq 3, which has been derived from rate equation 2. (The pseudo-first-order
rate constant is defined by the equation k = rate/ [Atotall,with [Atotsll= [ H A + ] [HA1 [A-I.) Empirical values for the parameters of eq 3 are given in Table 11. KO, ~ H = ~ A K o k ~ ~ and * , HA = KlkHAare obtained from the rate data by a method of successive approximations, which has been described previously. Pseudo-first-order rate constants are calculated from these parameters with the aid of eq 3 and are compared with the experimental data in Table I. The agreement is satisfactory. Approximate Arrhenius parameters for l ~ ”and ~ k~~ are reported in Table I11 and a comparison is done with the corresponding results obtained for the decarboxylation reactions of 4-aminobenzoic and 2,4-dihydroxybenzoic acids. Table IV contains results of kinetic measurements in acetate buffers. If buffer ratio and ionic strength are kept constant, the rate constant, k, does not depend on the buffer concentration. Consequently, there is no indication of general acid catalysis for this buffer system. Results for k in strongly acidic solutions are reported in Table V. The first-order rate constant decreases by
+
+
(11) R. G. Bates, “Determination of pH,” John Wiley and Sons, Ino., New York, N. Y., 1964.
Volume 78, Number 9 September 1968
A. V. WILL^, C. M. WON,AND P. VILK
3144
Table V : Decarboxylation of Anthranilic Acid in Aqueous Hydrochloric Acid Solutions a t 84.97'
Table I : Pseudo-First-Order Rate Constants for the Decarboxylation of Anthranilic Acid in Water ( p = 0.1 N KCl)
CHCl
---lOek, Obsd
sec-1Calcd
Temp, OC
[HsO+I
70.07 rt 0.04
3.89 x 2.90 X 1 . 9 1 x 10-2 9.87 x 10-3 4.92 x 10-3 2.95 x 10-3 7.53 x lo-@ 2.51 x
1.39 1.35 1.25 0.81 0.70 0.63 0.30 0.20
1.44 1.36 1.22 0.97 0.75 0.63 0.30 0.20
84.97 rt 0.04
3.98 x 2.98 X 1.98 x 9.89 x 4.93 x 2.96 x 2.05 x
7.57 6.95 5.97 4.22 3.38 2.99 0.78
7.53 6.90 6.02 4.52 3.44 2.85 0.77
10-2 10-3 10-8 10-3 10-6"
" I n acetate buffer.
Table 11: Equilibrium and Rate Constants in Water ( p = 0.1 N) 104
knHA, I O ~ ~ H A , Temp, o c
lO'Ko
70.07 84.97
1.57 2.29
kHiA,
kHA,
sec -1
sec -1
1.88 X 1.08 X
eec-1 mol-'
4.05 X lo-' 1.82 X
sec-1 mol-]
1.20 4.72
1.60 6.48
Table I11 : Arrhenius Parameters for the Decarboxylation in Water ( p = 0.1 N) -&k-
-k"A--,
Substrate
log A
E,,'kcal
log A
Ea, koa1
Anthranilic acid" p-Aminobenzoic acidb 2,4-Dihydroxybenzoic acidb
10.3 10.3
22.4 21.7
12.8 12.5 13.06
22.9 22.8 23.3
' This work.
...
...
Reference 6.
Table IV : Decarboxylation of Anthranilic Acid in Aqueous Acetate Buffers ( w = 0.1 N KCl) Temp, OC
~~C'HOA ~O ~~ , C C N ~ O A ~ 107k, , N N aec-1
70.07 f 0.04
1 2 3 3 6
1 2 3 1 2
2.10 2.08 2.04 3.01 3.05
84.97rt0.04
2 3
2 3
7.88 7.70
The Journal of Physical Chemistry
(=["I),
N
0.100 0.478 0.991 1.92 2.94 3.91
ha, N
lo%, 8ec -1
IO-EY, sec
0.105 0.60 1.59 4.68 11.0 24.6
7.28 7.54 6.58 4.00 2.16 1.11
1.13 1.28 1.50 2.49 4.62 9.01
a factor of 7 when the hydrochloric acid concentration is increased from 0.5 to 3.9 N
Discussion The apparent absence of general catalysis by acetic acid does not generally exclude the possibility of a ratedetermining proton transfer. General catalysis may remain undetectable if the BrZnsted a coefficient is close to unity. A well-known example of such a case is the acid-catalyzed hydration of alkenes.12 An QI value of 0.9 was found in the decarboxylation of paminosalicylic acid. It may be expected that QI is even higher f o r less reactive substrates (which indicates more complete proton transfer in the transition state). General catalysis by acetic acid was still detectable in the decarboxylation of 2-hydroxy-1-naphthoic acid and p-aminobenzoic acid.6 (However, the rate increases with increasing buffer concentration were not very much beyond the limits of experimental error.) It appears that in the decarboxylation of anthranilic acid, general acid catalysis has passed below the limit of detectability, possibly as a consequence of weak steric hindrance caused by the 0-NHz group. The rate maximum of the anthranilic acid decarboxylation is at a pH value between 0.2 and 0.5 (Table V), which is far away from the isoelectric point, l/z (pKo pK1) = 3.10. Consequently, the observed rate-pH profile cannot be explained solely by assuming either rate-determining formation or rate-determining decomposition of a tautomer of HA, such as X (see below), whose rate of formation or equilibrium concentration, respectively, would be at a maximum value at the isoelectric point. The same conclusion has been drawn already by Dunn, et al.,? for the decarboxylation of p-methoxyanthranilic acid. Actually, the observed validity of eq 3, with ~ H = ~ ~ A H ~> ~ k H AK , inothe pH region above 1.5, is also sufficient to exclude such a simple mechanism. Therefore, the presence of the rate maximum in the high-acidity region must indicate a change of the rate-determining step. It has been p o s t ~ l a t e dthat ~ ~ in ~ ~the ~ two reaction
+
(12) F. G. Ciapetta and M. Kilpatrick, J . Amer. Chem. Soc., 70, 639 (1948); R. W. Taft, E. L. Purlee, P. Riesz, and C. A. deFazio, ibid., 77, 1684 (1955). For further references, see ref 3, Chapter VI.
DECARBOXYLATION OF ANTHRANILIC ACIDIN AQUEOUS SOLUTION pathways corresponding to rate equation 2, the plex intermediates X and H X + are formed
u
com-
3145
the first factor of eq 4. For [H+] 2 1 N , the first factor attains a constant value because [H+] >> KO>> K1 and k1OK0 > k1-K1 fl
Since it may be expected that COO- is a much better leaving group than COOH, the intermediate H X + presumably undergoes a fast acid-base reaction to form X, which then splits off carbon dioxide rapidly.? A complete scheme of all possible reaction steps is HA
+ H30+
krQ k-iQ
1 tg, A-
+ H30+
H X + -% product
1 IKQ*
krk-I
kii-
X +product
-
+ COOH+
+
h 0 K o (if [H+] 2 1 N)
According to the experimental data (Table V), k decreases if [H+] is increased to l N and higher. Consequently, the derivative of the second factor must be negative. It follows from eq 5 that
- k11-k-IO < 0 kIIO/k-IO < k I I -/k-1-
k-I-kIIO
For the reasons discussed above, both ratios of rate constants cannot be on the same side with respect to unity, except if both are relatively close to unity. Consequently
+ C02
As has been shown p r e v i o u ~ l y , ~eq~ 4 ~ ~may ~ ~ ~be' ~ derived by applying the principle of the stationary state to the Eium of concentrations of X and H X +
K.1-
=
+ ~II"[H+I/Ko* + (k-1' + ~IIO)[H+I/KO*(4)
~ I I -
~ I I -
The first factor in this equation is essentially identical with the right side of eq 3. The second factor takes care of the decrease of the total rate due to return from the intermediates. If the two conditions ~ I I O >> and k11- >> k-I- are fulfilled simultaneously, the second factor will be equal to unity. It will be equal to a constant for low values of [H+]if k11- q k-x-. If on the other hand, the conditions ~ I I O IO and ~ I I - k - ~ - . (However, these extreme cases are not the only possibilities.) The first derivative with respect to [H+] of the second factor, f2, of eq 4 is equal to
The sign of df2/d[H+] does not depend on [H+]; it is determined by the sign of the expression in the parentheses in the numerator. For hydroxonium ion concenN , the second factor of eq 4 trations below 4 X appears t o be constant, since the pH dependence of the experimental rate constants, k , can be represented by
kIIO/k-IO kII-/k-I-
q1 51
(64 (Ob)
As ~ I I and " k-1O become more important with increasing [Hf], eq 6a and 6b indicate a partial or complete change of the rate-determining step from proton transfer to the carbon at low [H+] to cleavage of the bond between the ring and carboxylic (or carboxylate) group at high [H+]. For a quantitative treatment of the rate data in strongly acidic solutions, it is advantageous to split off ~II-/(~IILI-) from the second factor of eq 4. Furthermore, since in strongly acidic solutions K1
