1418
R. Zarnboni, A. Giacornelli,F. Malatesta, and A. lndelli
Kinetic and Salt Effects of the Ethyloxalate-Hydroxide Ion Reaction Roberto Zamboni, * Ambrogio Giacomelll, Francesco Malatesta, lstituto di Chimica Analitica ed Elettrochimica, Universitadi Pisa, Pisa, Italy
and Antonio lndelli lstituto di Tecnologie Chimiche Speciali, Universitadi Bologna, Bologna, Italy (Received December 1, 1975) Publication costs assisted by the Consiglio Nationale delle Ricerche
The rate of the reaction of the ethyloxalate ion with hydroxide ion has been measured at eight different temperatures, using a potentiometric method, in the presence of twelve different salts. Where a direct comparison is possible, an agreement with published results within 4% was obtained, despite the difference in conditions (method and concentrations). The Olson-Simonson effect was shown to be predicted by a numerical integration of the Poisson-Boltzmann equation (IPBE). The effects of alkali metal sulfates and tri- and tetrametaphosphates, however, cannot be quantitatively predicted neither by IPBE nor by the Mayer theory in its simplest form (DHLL Bz),using a single d parameter, but only by DHLL Bz using two different d , which are the same for all the salts. Different methods give values of ho, extrapolated at I = 0 for the alkali metal salts, which are in good agreement with each other, and fit very accurately an Arrhenius plot. Results for the alkaline earth metal ions require the additional assumption of ion pair formation, whereas results for the tetraalkylammonium ions cannot be predicted by any of the above theories or assumptions. It is concluded that nonelectrostatic effects are present, and this is confirmed by the values of AH* and AS*. For Mgz+ inner-sphere interactions appear to be the most likely explanation, whereas for tetraalkylammonium ions the effects appear to be related to the structure-forming properties of these electrolytes.
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Is it well known that elementary electrostatic theories, such as the Debye-Huckel theory, are inadequate to interpret the details of salt effects in reactions between ions, particularly when multivalent anions and/or cations or tetraalkylammoIt has been shown qualitatively that nium salts are the use of more sophisticated electrostatic theories can account for some of the most striking a n ~ m a l i e s . ~ We - ~considered it useful to make a careful study of a rather simple reaction in a variety of conditions in order to test the extent to which different theories can account for all the results, using as few adjustable parameters as possible. The reaction of ethyloxalate with hydroxide ion was chosen because it has already been studied with remarkable ~ a r e , 3 ,so ~ , that ~ a cpmparison with previous results is possible. Experimental Section Materials. Most of the chemicals were Erba RP. Sodium, potassium, magnesium, strontium, and barium nitrates were recrystallized three times from conductivity water and dried under vacuum. Potassium sulfate was recrystallized three times from conductivity water and then dried in a furnace a t 350 OC. Sodium tri- and tetrametaphosphates were prepared and purified as described in preceding papems Stock solutions of these salts were prepared by weighing, and the various concentrations were obtained by dilution. The tetraalkylammonium nitrates were prepared by adding to the corresponding bromides an amount of silver nitrate slightly lower than stoichiometric. The resulting solutions, after concentration under vacuum, were passed through an ionexchange column filled with Dowex 50 resin in the nitrate form. The solutions were dried and the solid salts were dissolved in hot conductivity water and crystallized a t low temperature. This was possible only for tetramethyl- and tetraethylammonium nitrates, while the n-tetrapropyl- and ntetrabutylammonium nitrates, because of their high solubilities in water, were precipitated a t low temperature from acThe Journal of Physical Chemistry, Vol. 80, No. 13, 1976
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etone and benzene, respectively. The solutions of tetraalkylammonium nitrates were standardized by passing them through an ion-exchange column filled with Dowex 50 in acid form and the effluents were titrated with 0.1 N NaOH using brom thymol blue as an indicator. Potassium ethyloxalate was prepared by Nielsen's method9 and crystallized from absolute ethanol. It was found to be 99.5% pure by complete hydrolysis and titration. Kinetics. All the kinetic measurements were carried out at 4, 10, 15, 20, 25, 30, 35, and 40 OC,in thermostats which assured a temperature constancy of better than 0.005 OC. A tenfold excess of potassium ethyloxalate was used, so that the reaction was pseudo-first order. The disappearance of OH- was followed potentiometrically using a glass electrode Metrohm EA 121, which can be used without appreciable error up to a p H of 12. Since the reaction, even in the absence of added salts, did not involve a remarkable change in ionic strength, or in general of the ionic environment, the logarithm of [OH-] is a strictly linear function of the potential, with a Nernstian slope. This has been confirmed in preliminary runs using different additions of NaOH. In all the following runs the concentration of potassium ethyloxalate was M and that of NaOH was M. The electrode potential was measured and recorded using a Metrohm potentiograph E 436. A solution of potassium ethyloxalate was put in a quartz flask and freed from COz by bubbling with pure nitrogen. After thermal equilibration the appropriate quantity of COz-free NaOH solutions was added with a microsyringe. Continuous stirring was maintained by means of a magnetic bar during all runs. The kinetic measurements were all repeated from two to four times.
Results The rate constants at 25 O C , where a direct comparison is possible, agree with the results already p ~ b l i s h e dwithin ~ , ~ 4%, despite the different experimental method, and the different
1419
Salt Effects in Ion Reactions 1.0
0.5
-1.8
e
2.0
-1.8
: '"
b
b4
A
1
-1.8
- 3.0
t
*
- 2.0
1.9
E
f+2%
-1.8
- 1.0
li.8 I I
I
0.05
0.10
Figure 1. Values of log k for the reaction in the presence of NaN03 (squares), KN03 (circles), Na3P309(triangles), and (CH&NN03 (diamonds): filled symbols, present results; _half-filledsymbols, ref 6; open symbols,ief 3. Blank values of log kareJ.716 for I = 0.001 10 (present results), 1.732 for I = 0.001 30,8and 1.717 for I = 0.001 67.3 Temperature, 25 "C.
reactant concentrations. This is shown in Figure 1,where some examples of the rate constants plotted against the square root of the ionic strength at 25 "C are reported. The agreement between our results in the presence of alkaline earth metal nitrates and the results of Hopp6 and Pruej in the presence of the corresponding chlorides is somewhat poorer. This can probably be attributed to the conditions chosen by Hop$ and Prue (high concentration of reactants, in comparison with the added salts), and to their sampling technique, which was not well suited for high reaction rates. In the presence of sodium or potassium the substitution of NO,?- for C104- or C1- does not appear to have a remarkable influence. On the contrary, the tetraethylammonium nitrate shows an effect which is strikingly different from that of the corresponding perchlorate.GA lack of additivity of the action of the cation and the anion of the added salts has been observed in the persulfate iodide reaction: KC104 is as effective as KCl, whereas NaC104 is more effective than NaCLIO In the presence of multivalent anions S042-,P30g3-, or P4012~-,the well-known "Olson-Simonson e f f e ~ t is " ~observed, that is the rate does not depend upon the ionic strength, but only upon the concentration of the cations. This is in agreement with what had been already r e p ~ r t e d . ~ , ~ Alkaline earth cations have a large accelerating effect increasing in the order Ba2+ < Sr2+ < Mg2+at all temperatures. A summary of the results is reported in Tables I, 11, and 111, which give activation parameters for the reactions in the presence of different salts. These values were obtained from rate measurements at eight temperatures from 4 to 40 "C.
Discussion One of the main features of the reaction between anions, that is, the Olson-Simonson effect, has been shown by Scatchard2 to be in agreement with the prediction of the Mayer theory.ll These results have been later confirmed for very high dilution^.^,^ On the other hand, the Mayer theory, in the simple form called by Friedman DHLL B2,l2fails to predict
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Figure 2. Natural logarithm of the ratio of the activity coefficients for the reaction A- 4- B2- + C3-, in the presence of salts of different valence types, MX, MpY, M3W, and M4Z, calculated by means of IPBE for a distance of closest approach of 3.0 A and plotted against I"* (lower curves, lower and right scales) or against [M+] ' I 2 (upper and left scale; 0 , 0, A, and * for the presence of MX, MpY, M3W, and M4Z,respectively): Dash-and-dot line, Debye-Huckel limiting slope; temperature, 25 "C.
the activity coefficients and the molar enthalpies at I > 0.01.13J4The numerical integration of the Poisson-Boltzmann equation (IPBE) gives satisfactory results for the same problems up to I = 0.1.13J4 Therefore we tried first of all to see whether IPBE is able to predict the Olson-Simonson effect. The calculations, performed with the usual method,13have shown that for reactions between two anions, independently of the charge of the activated complex, from -2 to -7, the Olson-Simonson effect is predicted in a range of ionic strength and of distance of closest approach ( d ) much wider than for the Mayer theory. Figure 2 shows one example for a reaction between a univalent and a bivalent anion, for a' = 3.0 A. In contrast, the decrease in rate when increasing the charge of the ions of the same sign as that of the reactants for small d (or for low values of the dielectric constant) is not predicted by IPBE. Such effect, which is opposite to the predictions of the Brdnsted-Debye-Scatchard equation, has been some times observed experimentally,14-lG and is predicted by the Mayer t h e ~ r y .On ~ , the ~ other hand, for very large values of d (beyond 6 8, and much more for highly charged activated complexes) IPBE, as well as DHLL B2,3,4predicts a gradual disappearance of the Olson-Simonson effect and a dependence upon the ionic strength. We conclude that the failure of the simple Debye-Huckel theory appears to be due mainly to the neglecting of the higher terms of the series development of the exponential, rather than to the well-known internal inconsistency of the Poisson-Boltzmann e q ~ a t i 0 n . l ~ Both IPBE and DHLL B2 predict the Olson-Simonson effect even in the presence of quadrivalent anions. This is not in agreement with our experimental data which show that in
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The Journal of Physical Chemistry, Vol. 80, No. 13, 1976
R. Zamboni, A. Giacomelli, F. Malatesta, and A. lndelli
1420
TABLE I: A*
(kJ mol-') and A@ (J mol-' deg-') for the Reaction in the Presence of Alkali Metal Salts"
0.105 AH*
35.2 f 0.2 -131.0 f 0.4 35.2 f 0.2 -130.2 f 0.4 35.3 f 0.2 -128.9 f 0.6 36.0 f 0.3 -125.9 f 1.0 37.3 f 0.3 -121.1 f 1.1
AS* 0.145 AH* AS 0.203 AH* AS* 0.247 AH* AS* 0.285 AH* AS*
*
a
35.7 f 0.1 -128.9 f 0.4 36.0 f 0.1 -127.3 f 0.5 36.6 f 0.2 -124.7 f 0.5 36.5 f 0.2 - 124.5 f 0.5 36.1 f 0.2 -125.2 f 0.6
35.4 f 0.4 -130.1 f 1.3 35.8 f 0.5 -128.4 f 1.5 35.5 f 0.2 -128.4 f 0.8 35.2 f 0.4 -129.4 f 1.1 35.3 f 0.3 -128.6 f 1.0
35.3 f 0.3 -130.8 f 1.0 35.9 f 0.3 -128.4 f 0.8 35.0 f 0.3 -130.8 f 0.7 34.8 f 0.3 -131.0 f 0.9 34.8 f 0.3 -130.6 f 0.9
35.2 f 0.2 -131.5 f 0.8 35.8 f 0.4 -129.2 f 1.1 34.9 f 0.3 -131.8 f 0.9 34.3 f 0.2 -133.5 f 0.7 34.2 f 0.2 -133.8 f 0.5
The values of AH* and AS* for the reaction in absence of added salt ( I l l 2 = 0.033) are 35.2 f 0.4 and 132.1 f 1.3, respectively.
TABLE 11: AI@ ( k J mol-') and A@ ( J mol-' deg-I) for the Reaction in the Presence of Tetraalkylammonium Nitrates
0.105 AH* AS* 0.145 AH* AS* 0.203 AH* AS* 0.247 AH* AS* 0.285 AH* AS
*
35.3 f 0.2 -131.0 f 0.5 36.2 f 0.1 -127.6 f 0.4 36.1 f 0.4 -127.6 f 1.1 35.2 f 0.2 -130.5 f 0.4 34.9 f 0.5 -131.1 f 1.5
36.7 f 0.4 -126.6 f 1.2 37.7 f 0.4 -123.5 f 1.2 38.8 f 0.2 -120.1 f 0.5 39.0 f 0.2 -120.0 f 0.5 39.0 f 0.3 -120.7 f 0.8
TABLE 111: A* ( k J mol-') and A@ ( J mol-' deg-I) for the Reaction in the Presence of Alkaline Earth Metal Salts 1112
0.052 AH* AS* 0.071 AH* AS* 0.084 AH* AS* 0.105 AH* AS* 0.145 AH* AS* 0.203 AH* AS* 0.247 AH* AS* 0.285 AH* AS*
Mg(N03h
Sr(N03)~
Ba(NOd2
45.7 f 0.2 -88.7 f 0.5 45.8 f 0.4 -83.5 f 1.4 45.4 f 0.3 -82.8 f 0.8 46.0 f 0.4 -78.0 f 1.3 45.1 f 0.3 -78.3 f 1.1 42.3 f 0.2 -85.1 f 0.7 39.4 f 0.4 -94.0 f 1.3 39.1 f 0.6 -94.6 f 1.8
38.7 f 0.2 -116.1 f 0.5 39.0 f 0.4 -112.2 f 1.1 39.2 f 0.2 -109.5 f 0.4 40.6 f 0.4 -102.6 f 1.3 40.5 f 0.5 -100.1 f 1.6 40.5 f 0.5 -97.3 f 1.6 41.7 f 0.5 -91.6 f 1.7 40.5 f 0.5 -94.8 f 1.5
41.3 f 0.4 -102.2 f 1.3 42.1 f 0.6 -96.8 f 1.9 42.7 f 0.8 -92.0 f 2.5 42.7 f 0.5 -90.7 f 1.7 42.6 f 0.3 -89.7 f 0.8
the presence of Na4P4O12 the rate is definitely smaller than in the presence of NaN03, at constant [Na+].A good agreement can instead be obtained by using DHLL B2 applied to a slightly different model, which involves two values of d , one for the pairs of ions containing the activated complex and the other for all the other pairs of ions. For instance, in the case of alkali metal salts, assuming d = 3.7 for the first value and d = 3.0 for the second, a good quantitative agreement can be found for the reaction rate in the presence of NaN03, KN03, K2S04, Na~P309,and Na4P4012 a t 25 "C (Figure 3).18 For other temperatures, the agreement is almost as good as at 25 OC for 15,20, and 30 "C and deteriorates progressively both at higher and lower temperatures. Probably this is due to the oversimplifying assumption of constant d's. Even at 4 and 40 O C , however, the agreement is good enough to obtain
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The Journal of Physical Chemistry, Vol. 80, No. 13, 1976
37.3 f 0.3 -124.6 f 0.8 37.5 f 0.3 -123.4 f 0.7 38.0 f 0.5 -122.1 f 1.4 38.3 f 0.4 -121.0 f 1.4 38.5 f 0.4 -120.8 f 1.4
35.2 f 0.3 -131.8 f 1.0 35.6 f 0.3 -129.9 f 0.9 35.9 f 0.3 -128.8 f 0.9 35.7 f 0.6 -129.3 f 2.1 35.4 f 0.6 -130.6 f 1.8
a reliable extrapolation at zero ionic strength, using all results obtained with the five mentioned salts. The values of log ko/T at 4,10,15,20,25,30,35, and 40 "C are respectively 3.745 f 0.010, 3.889 f 0.007, 3.000 f 0.004, 3.111 f 0.005, 3.208 f 0.003,3.313 f 0.005,3.417 f 0.006, and 3.510 f 0.008, and are a strictly linear function of 1 / T with a standard deviation of f0.003. The corresponding AH* and AS* values are 35.22 f 0.07 kJ mol-' and -132.7 f 0.3 J mol-' deg-l, respectively. Practically the same values of ko can be obtained using IPBE in the presence of NaN03, KN03, KzS04,and Na~P309, assuming d = 4.0 A. A reasonable fit, giving again the same values of ko, can also be obtained using the simple Br4nsted-Debye-Scatchard equation if individual values of d are assumed for each different salt, namely, d = 3.0 for NaN03 and KNOB,5 for K2S04,7 8, for Na3P309, and 15 for Na4P4O12 (the fit in the last case is somewhat poorer). The latter treatment cannot be considered more than an empirical curve-fitting method, in our opinion, but it is reassuring that practically the same extrapolations are obtained. In contrast, none of the three treatments succeeds in accounting for the rates obtained in the presence of tetraalkylammonium salts, particularly for the tetraethyl- and tetrapropylammonium nitrates. Since for 1-1 electrolytes the three theories give very similar iresults, the discrepancy must be attributed to the presence of nonelectrostatic effects. A similar conclusion can be reached for the effects of the alkaline earth metal ions. Both DHLL B2 and IPBE predict a large accelerating effect of such cations, but neither succeeds in giving a quantitative agreement with the experimental values, no matter what averaged is assumed for IPBE, or what combination of d's is assumed for DHLL B2. The results of the two treatments are remarkably different, though, so that, in principle, the possibility that a different type of electrostatic treatment gives a good results cannot be ruled out. However, further support for the existence of a nonelectrostatic effect will be given below.
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1421
Salt Effects in Ion Reactions
C
011
0:2
Figure 4. Values of log k/T from experimentat measurements ( * for the blank solution, 0 ,A,and in the presence of Mg(NO&, Sr(N03)2, and Ba(N03)2,respectively) and comparison with calculated values: dashed lines, DHLL B2 previsions for ‘a = 2.5 A (curve a) and 3.0 8, (curve b); continuous lines, IPBE for ‘a = 2.0 8, (curve c) and 2.6 8, (curve d); dotted lines, DHLL B2 with the assumption of different reaction paths (see text). Temperature, 25 O C .
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Figure 3. Values of log k/Tas a function of the total cation’s concentration in the blank solution (”) and in the presence of NaN03 (O),K N 0 3 (+I, K2S04 (0),Na3P309 (A), and Na4P4012 (0).As a comparison, DHLL 4-B2 calculated curves for a model which assumes 3.7 A as a distance of closest approach with the activated Eomplex and 3.0 8, in all the other cases, by starting from log ko/T = 3.208 for I = 0: continuous, dashed, dotted, and dash-anddot lines for the presence of 1-1, 1-2, 1-3, and 1-4 electrolytes, respectively. Temperature, 25 O C .
electrostatic effects is the dependence upon the concentration of the activation parameters.17 The Br4nsted-Debye-Scatchard equation predicts that an increase in the rate due to salt effects is accompanied by an increase in the activation enthalpy and an increase (eventually, a less negative value) of the activation entropy. This occurs because when increasing the temperature the product D T decreases, and in all cases this product appears in the denominator of the different A good agreement can be obtained using DHLL B2 and equations expressing the logarithms of the activity coeffithe assumption of different reaction paths, involving the ascients. Therefore all electrostatic treatments should give the sociation of the alkaline earth metal cations with the reactants same trend. Not all the apparent exceptions have a clear (see Figure 4). One path could involve the ion pair MOH+ and meaning, because of the experimental errors. In the case of the Etox-, another OH- and MEtox+, and a third MOH+ and alkali metal salts, for instance, this prediction cannot be MEtox+. The reaction rate, u , was therefore calculated from considered unfulfilled. In the case of the alkaline earth cations and particularly of u = [OH-] [ E t ~ ~ - ] f - l ~ Mg2+,at very low concentration, up to 0.02 M, there is a conX [kof-a-l + hi[M]f+z + h2[MI2f+z} (1) tinuous increase of the activation entropy, but from 0.02 onwhere [OH-], [Etox-1, and [MI are the concentrations of hyward the increase in the rate is entirely due to a decrease in droxide, ethyloxalate, and alkaline earth metal ions calculated AH*, AS* becoming progressively more negative. Of course, from the association constants, f + 2 , f- 1, and f-2 are the DHLL if the ion pair assumption is considered as the explanation of B2 activity coefficients of a +2, -1, and -2 charged ion, and the increase in the rate due to the magnesium salt, the activation parameters contain the enthalpy and entropy of fork l and kz are rate constants including the association conmation of the ion pairs. However if the ion pairs were elecstants of alkaline earth metal ion with OH- (K’) and with Etox- (K”), kl = kl’K’ k1”K” and k2 = k2’K’K”. hl and trostatic in nature, the above rule for the comprehensive activation parameters should be obeyed in any case. This can k2 were calculated by least squares with given values of & (or be taken as another indication that Mg2+ forms inner-sphere ti's), K’, and K”. A wide range of association constants and & compounds with the reactants or the activated complex. values leads to very similar fits to the experimental data, in spite the fact that very different values of hl and h2 are obNonelectrostatic effects must be present in the case of tained. No particular meaning can therefore be given to the tetraalkylammonium salts as well, and particularly for various rate and association constants. Figure 4 shows the tetraethyl- and tetrapropylammonium salts, which at high results of the calculation performed at 25 “C by assuming for concentration show consistently larger activation energies and less negative entropies than the alkali metal salts, despite the 8’s the same values as previously used for the alkali-metal ions case (6 = 3.7 and 3.0 .&) and arbitrarily setting K” equal to K’. fact that the reaction rate is smaller. On the other hand, for T o K’ were attributed the values of 380,6.6, and 4.35 M-l for the tetramethylammonium nitrate, there is apparently a slight MgZf, Sr2+,and Ba2+,respectively,lg a choice which also is decrease in 1H*and a slight decrease in AS* when increasing rather arbitrary. The respective values of hl (Mw2s-l) and kz the concentration. Finally for tetraethyl-, tetrapropyl-, and (M-3 s-l) are 3.91 X lo3 and 1.32 X lo6 for Mg2+,7.63 X 102 tetrabutylammonium nitrates the decrease in the rate at and 8.54 X lo3for Sr2+,and 5.48 X lo2and 2.77 X 10Sfor Ba2+, higher concentrations is accompanied by an increase in the whereas ko has the previously extrapolated value of 0.482 M-l activation energy. Although the experimental errors could s-l. A similar treatment was adopted by Hopp6 and P r ~ e , ~invalidate some of these trends, it appears unlikely that they using the simple Br4nsted-Debye-Scatchard equation. The are not real. That these nonelectrostatic effects are due to the fact that, even using DHLL B2, ionic association must be influence of these ions on the structure of the water is very invoked suggests that inner-sphere interactions are involved. likely. More specific effects, however, cannot be ruled out. The A critical, although difficult, test for the existence of nonfact that the highest AH* are observed for EtdNNO3 and for
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The Journal of Physical Chemistry?Vol. 80, No. 13, 1976
K. C. Chang and E. Grunwald
1422
n-PrdNNO3 is perhaps related to the similarity in size with the leaving ethyl group. References a n d Notes (1) (2) (3) (4) (5) (6)
8.Perlmutter-Hayman, frog. React. Kinet., 6, 240 (1971).
G. Scatchard, Natl. Bur. Stand. (U.S.), Circ., No. 534, 185 (1953). A. indeili, Gazz. Chim. /tal., 92, 365 (1962). A. indeili and R. De Santis, J. Chem. fhys., 55,481 1 (1971). J. I. Hoppe and J. E. Prue, J. Chem. SOC.,1775 (1957). A. Indeiii, V. Bartocci, F. Ferranti, and M. G. Lucarelli, J. Chem. fhys., 44, 2069 (1966); A. Indelli. J. fhys. Chem., 65, 972 (1961). (7) A. R. Olson and T. R. Simonson, J. Chem. fhys., 17, 1167 (1949). (8) A. Indelii. Ann. Chim. (Rome), 43, 845 (1953); 46, 717 (1956): 47, 588 (1957); T.Moeller, lnorg. Synth., 5, 98 (1957).
R. F. Nielsen, J. Am. Chem. SOC.,58, 206 (1936). A indelli and J. E. Prue, J. Chem. Soc., 107 (1959). J. E. Mayer, J. Chem. fhys., 18, 1426 (1950). J. C. Rasaiah and H. L. Friedmann, J. Chem. Phys., 48, 2742 (1968). A. indelli and F. Malatesta, Gazz. Chim. ltal., 103, 421 (1973). A. lndelli and F. Malatesta, Gazz. Chim. ltal., 103, 435 (1973). R. M. Heaiy and M. L. Kilpatrich, J. Am. Chem. SOC.,77, 5258 (1955). V. Carassiti, C. Dejak, and I Mazzei, Ann. Chim. (Rome),50, 979 (1960). A. lndelli and E. S.Amis, J. Am. Chem. SOC.,82, 332 (1960). A rationalizationof such mcdei can be done by taking into account that the activated complex is a species intrinsically different from the other ions. it is not surprising therefore that its "radius" cannot be assumed to be identical with that of the other ions, even in a gross approximation which neglects the individual difference between "normal" ions. Obviously, no treatment of this kind can be made using IPBE. (19) J. N. Butler, "Ionic Equilibrium. A Mathematical Approach", Addison-Wesiy, Reading, Mass.. 1964, p 469.
(9) (10) (11) (12) (13) (14) (15) (16) (17) (18)
Water Participation in Proton-Transfer Reactions of Glycine and Glycine Methyl Ester' K. C. Chang and Ernest Grunwald* Chemistry Department, Brandeis Unlversity, Waltham, Massachusetts 02154 (Received January 15, 1976)
NH-proton transfer reactions with water participation of glycine and methyl esters of glycine, alanine, and phenylalanine were studied by dynamic NMR methods. For glycine and its methyl ester, kinetic analysis (and comparison with total rates of NH-proton exchange due to Sheinblatt and Gutowsky) reveals that processes which are second order in substrate proceed partly with and partly without water participation. There is evidence for intramolecular proton transfer between the NH3+ and C02- groups in glycine zwitterion, and for bifunctional proton transfer between the zwitterion and the uncharged amino acid. Rate constants are reported for proton transfer between ammonia and a series of carboxylic acids. A precise and convenient pulse sequence for NMR T1 measurement is described. The kinetics of "3-proton exchange of glycine in aqueous solution has been studied comprehensively by Sheinblatt and Gutowsky (SG).2These authors examined the CHz-proton resonance and thus measured the total rate of NH-proton exchange. Because of the insight one can gain into solvation phenomena by studying proton exchange with water particip a t i ~ n we , ~ now report a complimentary study of "3-toHOH proton exchange. Total rates of NH3-proton transfer between glycine and water have also been studied by 15N NMR,4s5and by relaxation ~ p e c t r o m e t r y . ~ , ~ In the present work, exchange rates were deduced from measurements of (l/Tz - 1/T1) of the H2O NMR or, for fast exchange, of the collapsed H20-NH3 NMR. The technique and rate calculations are familiar from previous publications.8 Rate measurements were made at five glycine concentrations ranging from 0.04 to 0.20 M, and in the pH range 3.9-6.3. All in all, -60 independent solutions were measured at 25 " C and subjected to kinetic analysis. The kinetics is fully consistent with that established by SG. The rate law is shown as follows: 3/7"
= kA + kg[R*]
+ kc[OH-] + k ~ [ R - l
(1)
where 7" = mean time a proton resides on -NHs during one cycle of proton exchange between R* and water; R+ = H3N+CHzCOzH;Rf = H3NfCH&02-; Ro = HzNCHzCOzH; R- = HzNCHzC02-. The factor 3 allows for the fact that there are three NH protons per molecule of the reactant, R'. The Journal of Physical Chemistry, Vol. 80, No. 13, 1976
The new rate constants for proton exchange with water, and comparable rate constants reported by SG2 for total NHproton exchange, are listed in Table I. The ratio in each case measures the fraction fw of reaction with water participation. These rate constants will now be discussed briefly. kA. As pointed out by SG, the major contribution to kA is made by an intramolecular reaction, R* Ro. Our rate constant is in good agreement with that of SG. Unfortunately, fw is indeterminate in this case because the COzH proton in Ro is in rapid exchange with water proton^.^^^ kg. Table I shows three possible reactions which might account for the observed kinetics. A fourth possibility, 2R* s 2R0, has a pK of 10.8 and would be undetectably slow. Reactions I11 and IV, which yield R- R+, involve just one functional group in each of the reactants, while the symmetrical process V involves two functional groups. Reaction I11 involves the functional groups NH3+ and "2. Rate constants for such processes in the direction of negative AGO are rarely greater than 2 X lo9 s-l M-1.9J0 Thus a plausible upper limit to the contribution of 111to kB is estimated to be 100 s-l M-l. Reaction IV involves proton transfer between NH3+ and COz-. Rate constants for the analogue of the reverse reaction have been measured11,12and are listed in Table 11. The values are generally less than lo9 s-l M-l, and a plausible upper limit to the contribution of IV to k B is thus estimated to be 50 s-l M-1. Thus reaction V appears to make a significant contribution. Jencks and Hand13 recently con-
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