J. F. KINTOEAND F. J. JOHNSTON
854
Halogen Displacement Reactions of Chloro- and Bromoacetic Acids in Water and Dioxane-Water Solutions'
by J. F. Hinton and F. J. Johnston Department of Chemistry, University of Georgia, Athens, Georgia
(Received August 24, 1964)
The reactions involving halogen displacenient in chloroacetic acid by broiiiide and in bronioacetic acid by chloride have been studied in aqueous and in 50% dioxane solutions. The corresponding reactions involving chloro- and bronioacetic acids with both hydrogen chloride and hydrogen broiiiide were studied in a 95% dioxane solvent. I n each case, rates were first order with respect to the haloacetic acid and to the total halide or halogen acid concentrations. Activation enthalpies and entropies were determined. For the reactions involving different halogen atoms, equilibrium constants and reaction enthalpies and entropies were evaluated froni the kinetic parameters. I n the 95% dioxane system, equilibrium constants measured directly were in good agreement with those determined from the rate constant ratio. Activation parameters in the 95% dioxane system were markedly different froni those in aqueous solution and are best explained in ternis of a protonation equilibrium involving the haloacetic acid and the halogen acid.
Introduction Several halogen-exchange and replacement reactions involving halide ions and a-halogen acids in aqueous solutions have been previously described. Wagner2 studied the reactions between chloroacetic acid and iodide and between iodoacetic acid and chloride. O l s 0 1 1 ~and ~ ~ co-workers have measured exchange and replacement reaction rates in systems consisting of chloro- and bromosuccinic acids and the corresponding halide ions and in systems consisting of phenylchloroand bromoacetic acids and halide ions. The latter reactions involved substitution on an asyiniiietric carbon arid were followed by optical activity measurements. We have516used radioactive tracers to follow the exchange reactions between chloro-, bromo-, and iodoacetic acids and the corresponding halide ions. I n every case, these reactions were second order and evidently involved a direct nucleophilic displaceiiient process. When replacement of one halogen by another was involved, enthalpy and entropy changes evaluated froni the activation paranieters were consistent with available therniochemical data. We have studied the kinetics of the following reactions in the indicated solvent systems The Journal of Physical Chemistry
CHzCICOOH
+ Br-
water
__--
--+
50% dioxane-water
+ C1-
(Af)
+ Br-
(A,)
+ HC1
(B)
+ HBr
(C)
+ HCI
(Df)
CHzBrCOOH CHzBrCOOH
+ C1-
water -----+ 50% dioxane-water
CHzCICOOH CHzCICOOH
+
HC1 95% dioxane-water -3
CHzClCOOH
CH~B~COOH + H B 95% ~ dioxane-water
+
CHzBrCOOH CHzCICOOH + H B 95% ~
dioxane-water
--+
CHzBrCOOH
(1) (a) This paper was abstracted from the Ph.D. dissertation of James F. Hinton, University of Georgia, Aug. 1964; (b) this work has been performed under A.E.C. Contract AT(40-1)2826. (2) C . Wagner, 2. physik. Chem., 115, 121 (1925). (3) A. R. Olson and F. A. Long, J . Am. Chem. Soc., 58, 393 (1936). (4) A. R . Olson and M.J. Young, ibid.,58, 1157 (1936). (5) R . A. Kenney and F. J. Johnston, .J. P h y s . Chem., 63, 1426 (1959). (6) J. F. Hinton and F. J. Johnston, ibid., 67, 2557 (1963).
HALOGEN DISPLACEMENT REACTIONS OF CHLOROAND BROMOACETIC ACIDS
Radioactive tral-ers mere used to follow both the exchange reactions and the replacement reactions. Formation of molecular iodine under the conditions of these reactions prevented our obtaining meaningful data on the corresponding systems in which iodide or hydrogen iodide were involved. From the rate expressions arid activation parameters for the forward and reverse replacement processes, enthalpies and entropies for the over-all reaction were evaluated. This report summarizes our rate studies and presents thermochemical data for replacement reactions in water and dioxane-water solutions.
Experimental Reagents. Baker Analyzed chloroacetic acid was used after fractional recrystallization from benzene. Eastnian White Label brotnoacetic acid was used without further purification. Eastman White Label p dioxane was refluxed over potassiuni hydroxide for approximately 24 hr. and distilled. The final fraction used in the solvent mixture had a refractive index of 1.4198 a t 25". Solvent conipositions are expressed as per cent by volume. Chlorine-36 labeled hydrochloric acid as obtained from Oak Ridge was diluted with reagent hydrochloric acid. Sodium Ix-oniide or hydrobromic acid solutions were spiked with potassium bromide-82 obtained from Oak Ridge. Procedure. The experimental procedures are similar to those described previously. I n general, reactant aliquots were quenched in ice-water, and halide separation was accomplishrd through silver nitrate titration. The HC1-CH,ClCOOH systeni in 95y0 dioxane was studied at higher temperatures, and individual, sealed cells were used. I n every case, background and separationinduced reactions were negligible. Chlorine-36 activities mere measured by Geiger-Mueller counting of liquid samples, and bromine-82 samples were counted in a scintillation system. Comparisons of reacted samples containing the bromine activity were made with a standard stock sample within time intervals such that decay corrections were unnecessary. Stock samples of the unseparated reaction mixture were adjusted to the same solvent conditions as the reacted samples so that counting rates were compared under similar density conditions. St:tndard deviations for the net counting rates were always within 27". In the aqueous systenis hydrolysis of the chloro- and bromoacetic acids was evident upon prolonged exposure at the temperatures used in these experiments.
855
This prevented our studying the replacement rates under conditions close to equilibrium. In general, at the concentrations used, hydrolysis was not detectable, and the reverse reaction was negligible to a t least 10% and as high as 30% replacement. Under these conditions, second-order rate constants were obtained using the expression k
=
~
1 In b- ( a ( a - b)t u ( b -
5) 2)
For the replacement react,ions, radioadvity was used as a direct measure of concentration. Rates for the exchange react,ions which were studied in the 95% dioxane systems obeyed the familiar exchange rate law
For a process first order in each of the exchanging species, this expression becomes k
=
In (1 - F ) ( a b)t
+
(3)
In these equations a and b have their usual significance as initial reactant concentrations, and F is given by the ratio
F =
(specific activity, haloacetic acid), (specific activity, haloacetic acid), =
I n the 95% dioxane s y s t e m hydrolysis of the haloacetic acids was not detectable in any of the experiments. Equilibrium constants were evaluated from the ratios of the rate constants for the forward the reverse reactions, K = Icr/k,. Enthalpies and entropies for the reactions were obtained as AH = AHr* - AH,* and A S = AS{* - AS,*.
Results and Discussion Rates of reactions A{ through I), were first order with respect to the total concentration of each of the exchanging species under all conditions studied. Our experiments were performed over reactant concentration ranges of from 0.03 to 0.48 M in haloacetic arid and from 0.02 to 0.09 dd in the halogen acid. The widest range covered at a given temperature in the aqueous and 50% dioxane systems was from 0.0755 to 0.364 '11 for the haloacetic acid arid from 0.0200 to 0.0947 J/ for the halide. In the 95oJ, dioxane system, the widest range covered at a given temperature was from 0.091 1 to 0.2109 M for the haloacetic acid and from 0.0197 to 0.0420 M for the halogen acid. Standard deviations
J. F. HINTOX AND F. J. JOHNSTON
856
for the second-order rate constants corresponded with few exceptions to variations within 2%. The maximum deviation corresponded to a variation of slightly over 6%. The aqueous systems were 0.006 to 0.6 M in nitric acid in order to ensure the presence of the haloacetic acid in the molecular form. No effect on the rate constants resulted from variation of the nitric acid concentration over this range. I n the 50 and 95% dioxane systems the rate constants were the same in the presence or absence of nitric acid. The rate constants for the several systenis studied are summarized in Tables I through V.
Table IV : Summary of Rate Constants for the Reaction CHzClCOOH HC1* CHzCl*COOH HC1 in 957, Dioxane-Water Solutions
+
Temp., OK.
k. 1. mole-] set.-'
333.7 343.6 355.1 371.2
1 . 0 5 x 10-5 1.83 x 10-5 3.67 x 10-5 8.72 X
Table V : Summary of Rate Constants for the Reactions CHzBrCOOH HBr* CHzBr*COOH HBr in 95% Dioxane-Water Solutions
+
Table I : Summary of Rate Constants for the Reactions
+ Br- e CHIBrCOOH + C1- in k, ki
CHzCICOOH
OK.
k f ,1. mole-' 8ec. - 1
k,, 1. mole-' 8ec.-1
2.14 X 5.56 X 2.10 x 10-5 5.86 X
1.25 X 3.17 x 10-5 1 . 2 0 x 10-4 3.36 x 10-4
304.1 311.9 323,7 333.7
+
Temp., OK.
k, 1. mole-' 8ec.-t
305.9 315,O 323.6
1 . 8 5 x 10-3 4 . 2 3 x 10-3 8.37 x 10-3
Aqueous Solution Temp.,
+
these systems are characteristic of a process involving the haloacetic aci'd molecule and halide ion. Activation parameters were obtained by a leastsquares fitting of the rate constants t,o t,he expression
Table I1 : Summary of Rate Constants for the Reactions CHZCICOOH
kr
+ Br-
CHzBrCOOH kr
+ C1- in
50y0 Dioxane-Water Solutions Temp., OK.
k f ,1. mole-1 8ec.-l
k,, 1. mole-' 8ec.-l
323.6 333,7 343.4
4.39 x 10-6 1 . 1 5 x 10-4 2 . 7 5 x 10-4
3.49 x 10-4 8.90 X 2.14 x 10-3
Table 111: Summary of Rate Constants for the Reactions CHzCICOOH
kr
+ HBr
CHzBrCOOH kr
+ HC1 in
Standard deviations for the activation enthalpies were within kO.15 kcal. mole-' in the aqueous and 50% dioxane systems and within 2~0.20kcal. mole-' in the 95% dioxane systems. For the entropies of activation the corresponding deviations were within i0.4 and 2Z0.5 e.u. The activation parameters for the replacement reactions in aqueous solutions are summarized in Table VI and compared with those for the exchange reactions which we have previously reported. The results parallel very closely those obtained by Olson with the halosuccinic acids and phenylhaloacetic acid.
95% Dioxane-Water Solutions Temp., OK.
313 323 333 343
7 6 7 4
ki. 1. mole-' sec.-l
2 6 1 4
36 x 10-5 50 X 79 x 10-4 14 x 10-4
kr,1. mole-' 8ec.-I
I 3 6 1
50 x 10-4 12 x 10-4 55 X 29 x 10-4
Conductance measurements by Owen and Waters' on hydrochloric acid solutions which are &yo by weight dioxane indicate little ion association. Conductance measurements on hydrobromic acid in the 50% by volume dioxane salient show a similar result. We have assumed, therefore, that the rate Constants in The Journal o j Physical Chemietry
Table VI : Activation Parameters for the Reactions CHIXCOOH Y- +CHzYCOOH X- in Aqueous Solution
+
+
X
c1 c1 LI
Br Br See ref. 5.
Y
AH*, kcal. mole-'
c1. Br c1 Brb See ref. 6.
23.85 22.06 21.97 18.98
(7) B. Owen and B. Waters, J . Am. Chem.
AS*,e.u.
-10.5 -11.5 - 8.8
-13 4
SOC.,
60,2371 (1938).
HALOGEN DISPLACEMENT REACTIONS OF CHLOROA N D BROMOACETIC ACIDS
I n the 5oQ/, dioxane solutions the enthalpy of activation for reaction Ai is 19.70 kcal. mole-’, and the entropy of activation is -17.8 e.u. The corresponding values for reaction A, are 19.39 kcal. mole-’ and - 14.7 e.u. The exchange reactions were not studied in the 50% system. The equilibrium constant for the reaction with chloroacetic acid and bromide ions as reactants may be obtained from the ratio of forward to reverse reaction rate constants. I n water this value is 0.174, and in 5OLr, dioxane it is 0.129. These quantities were essentially constant over the temperature range studied. In water AH for the reaction is approximately zero, and in 50yo dioxane it is endothermic by 0.3 f 0.2 kcal. mole-’. The entropy changes in the two solvents are -2.7 and --3.1 e.u. Standard enthalpies of chloride and bromide ions in water are -40.0 and -28.9 kcal. mole-’, respectively.8 The corresponding entropies are 13.2 and 19.3 e.u. Neglecting concentration effects, the over-all zero enthalpy change for the reaction leads to an enthalpy difference in aqueous solution between chloroacetic and bromoacetic acids of 11.1 kcal. mole-’. From the entropy change for the reaction in wai,er, one obtains for the corresponding entropy difference between the two haloacetic acids - 3.4 e.u. Thermochemical data for the dioxane systems were not available, and similar comparisons were not made. Activation parameters for reactions B through D in the 95% dioxane systems are summarized in Table VII. The values are markedly different from those in
Table VI1 : Activation Parameters for the Reactions CHzXCOOH H Y +CHzYCOOH H X in 95y0 Dioxane Solutions
+
+
X
Y
c1 c1
c1
Br Br
Br
c1 Br
AH*,kcsl.
mole-’
13.22 20 11 14 89 16.40
AS*,e.u.
-42.0 -16 0 -28 7 -17 4
aqueous systems. The most striking feature involves the very low activation enthalpies and large negative entropy terms associated with the HCl systems. Owen and Waters’ give a value of 2 X lo-* as the dissociation constant a t 25’ for HCI in 82y0 by weight dioxane. Conductivities of HC1 and HBr solutions in g5yo dioxane were not measureable with the type of apparatus available to us. It must be assumed that in this system both halogen acids exist essentially in the molecular form. A reaction mechanism involving
857
the haloacetic acid molecule and a halide ion, the latter concentration depending upon a dissociation of the halogen acid, is not consistent with the second-order character of the rate expression. A mechanism involving a direct exchange reaction between a haloacetic acid molecule and a halogen acid molecule cannot be discounted and would result in a rate expression consistent with our data. An energetically more satisfying mechanism, however, is one which involves a protonation equilibrium between the haloacetic acid and halogen acid. Using the chlorine exchange reaction as an example ki
+ HC1 E CH&IC(OH2)+ + C1CHzClC(OH)z+ + C1- k“, CH2ClCOOH + HC1
CHzClCOOH
ki’
(E) (F)
The exchange rate is k2K [CH&lCOOH] [HCI] where K is the equilibrium constant for reaction E. According to this picture the observed second-order rate constants are given by kzK. The apparent enthalpy of activaAHE where AHz* is tion is then AH*,,bsd = AHz* the activation enthalpy for reaction F, and AHE is the enthalpy change for reaction E. The more highly exothermic the protonation or dissociation reaction, the smaller will be the measured activation enthalpy, A reliable estimate of AHE for either HCl or HBr is difficult to make. The corresponding reaction in HCl H30+(aq) C1-(aq), is estiwater, H20 mateds to be exothermic by approximately 25 kcal. mole-’. For HBr, the dissociation is estimated by the same process to be exothermic by approximately 14 kcal. mole-’. While both enthalpy changes will be numerically smaller in dioxane, it is quite reasonable that the process involving HCl will be more exothermic than that with HBr. The activation enthalpies for reactions involving HC1 will then be lowered to a greater extent than those with HBr, a result consistent with our observations. Entropies of activation for the reactions will be of the form A S * o ~ s d = ASZ* 4-ASE. In water, dissociation entropies for acids are of the order of -20 e.u. with values for larger ions being less negative. Again, the values in 95yo dioxane are not easily estimated. Our observed activation entropies are consistent, at least in direction, with these considerations. The independence of the rate constants from the concentration of added acid in aqueous and SOYo di-
+
+
-
+
~
(8) G. N. Lewis and
M . Randall, “Thermodyanics,” 2nd Ed., Mc-
Graw-Hill Book Co., Inc., New York, N . T., 1961, p. 400. (9) A. A. Frost and R. G. Pearson, “Kinetics and Mechanism,” 2nd Ed., John Wiley and Sons, Inc., New Tork, N Y.,1961, p. 132.
Volume 69,Number S
M a r c h 1966
K . C. Hou
858
oxane solutions precludes the possibility that such a protonation mechanism occurs in these solvents. The equilibrium constant for the reaction with chloroacetic acid and HBr as reactants in 95% dioxane may be obtained from the rate constant ratio. The values for the temperatures studied are listed in Table VIII.
Table VI11 : T h e Equilibrium Constant for the Reaction CH,ClCOOH HBr +CHzBrCOOH HC1 in
+
+
95Tc Dioxane Solution Temp.,
OK.
313 7 323.6 333.7 343 4
K = ki/k,
K (measd.)
0.157 0.209 0.272 0 318
0.213 0,272
...
AND
H. B. PALMER
Since no hydrolysis of the haloacetic acids was detected in this solvent system upon prolonged exposure to the temperatures a t which the replacement reactions were studied, it was possible to obtain a direct measurement of the equilibrium constant. The results of two such experiments are listed in column 3 of Table VIII. The agreement with the results obtained from the kinetic data is quite satisfactory. The enthalpy change for the reaction in 95% dioxane is 5.22 kcal. mole-', and the entropy change is 12.7 e.u. Thermochemical data are not available in this solvent to allow comparison with that obtained from kinetics.
Acknowledgment. We appreciate helpful discussions with Dr. R. C. Lamb concerning this work.
The Kinetics of Thermal Decomposition of Diacetylene in a Flow System'
by K. C. Hou and H. B. Palmer Department of Fuel Technology, College of Mineral IndustrGs, T h e Pennsylvania State University, University P a r k , Pennsylvania (Received Auguet 87, 1964)
Diacetylene decomposes in a manner very similar to acetylene. Between 973 and 1223"K., the decomposition is second order in the diacetylene concentration. The rate constant behaves with increasing temperature in a way that implies a transition from long-chain toward nonchain decomposition. The long-chain rate constant is approximately k = lo1*exp(-40 kcal./RT) cc./mole sec. It is suggested that 40 kcal. may represent the approximate energy of a low-lying triplet state of diacetylene.
Diacetylerie C4Hz,has been reported to be a significant product of the pyrolysis of acetylene in shock waves2a and in a flow systeni.2b'c I t also is found as a product of the thermal decomposition of bensene3j4; i n fact, Kinriey and Slyshdconcluded that it is a primary product (with hydrogen and acetylene) of benzene pyrolysis. Although our own more recent work4 indicates that this conclusion was not quite correct, diacetylene IS, nevertheless, a significant product and niay be an ilnportant intermediate in the processes T h e Journal of Physical Chemistry
leading to carbon formation. Therefore, a study of its thermal decomposition seems worthwhile. (1) This work has been supported by a grant from the J. M.Huber Corp. (2) (a) C. F. Aten and E. F. Greene, Combust. F l a m i , 5 , 55 (1961); (b) H. B. Palmer and F. L. Dormish, J . P h y s . Chem., 68, 1553 (1964); (c) F. L. Dormish. R I S . Thesis, Department of Fuel Technology, Pennsylvania State University, 1962. (3) C. R. Kinney and R. S. Slysh, Proc. Conf. Carbon, 4th, B u f a l o , 1967, 301 (1960). (4) K. C. HOUand H. B. Palmer, J . P h y s . Chem., 69, 863 (1965)
