1426
R. A. KENNEYA N D E". J. JOHNSON
Vol. 63
KINETICS OF CHLORINE EXCHANGE BETWEEN HCl AND CHLOROACETIC ACID I N AQUEOUS SOLUTION1 BYR. A. KENNEYAND F. J. JOHNSTON Contribution from the Department of Chemistry, University of Louisville, Louisville, Kentucky Received February 9,18.58
I n aqueous solutions containing HC1 and CHzCICOOH, exchange of chlorine between the two species has been found to occur simultaneously with the hydrolysis of the latter. The rate of this exchange is first order with respect to each of the HC1 and chloroacetic acid concentrations and may be adequately described by the expression rate = 0.95 X 10l1 exp -24,500 f 5OO)JRTl(HC1)(CHaCICOOH)moles 1.-1 sec.-l. The corresponding entropy of activation evaluated at 80" 18 - 10.6 cal. mole-' deg.-l The simultaneous hydro1 sis reaction followed a pseudo first-order type behavior with the rate = 2.11 X l o g exp [( -24,800 f 800)/RT](CH2C1C80H) moles 1.-1 8ec.-1.
I(
Introduction I n aqueous CHzCICOOH solutions a t elevated temperatures hydrolysis occurs with the production of C1- and CH20HCOOH. This reaction has been studied over a wide range of conditions by a number of investigator^.^-' We have found, using Claa,that in the presence of HC1 exchange of the chlorine in the chloroacetic acid with the chloride occurs simultaneously with the hydrolysis. This article reports the results of a kinetic study of the simultaneous exchange and hydrolysis reactions. The mathematical treatment of isotopic exchange in such non-equilibrium systems has been clearly formulated by Luehr, Challenger and Masters.* A slightly modified version of their results as applied to the CHzCICOOH-HC1 system is given below. For the hydrolysis reaction CHzClCOOH d( C1-) dt
+ HOH +CHzOHCOOH + HCI -d( CHzClCOOH) = dt
-=E
40
(1)
+
p(t)
(2)
- p(l)
(3)
Then (Cl-)t = (C1-)O
+
or(l)dt = a
mid
(CHzC1COOH)t = (CHzC1COOH)o
-
a(t)dt = b
For the exchange reaction CHaClCOOH
+ C1-*
CHzCl*COOH
+ C1-
the exchange rate R x ( t ) = kx(C1-)m(CHzCICOOH)n = kx[a
+ P(l)I"lb
- P(0l"
(4)
Then
(1) This work was supported by a Frederick Gardner Cottrell Grant from the Research Corporation. This paper was abstracted from the Ph.D. thesis of Richard A. Kenney, Universit.y of Louisville, 1958.
(2) W. A. Drushel and G. S. Simpson, J . A m . Chem. Soc., 89, 2455 (1917). (3) E. (4) H. (5) 0. (6) L. (1946).
A. Moelwyn-Hughes, J . Chem. Soc., 101 (1932). M. Dawson and E. R. Pycock, ibid., 153 (1986). Reitz, 2. physik. Chem., 8171, 85 (1938). F. Berhenke and E. C. Britton, Ind. En#. Chem., 8 8 , 544
(7) F. Kunze. Monaf8h, 79, 254 (1948). (8) C. P. Liiehr, G. E. Challenger and B. J. Masters, J . A m . Chem. Soc., 78, 1314 (1956).
-
u
+ p ( t ) l m - ' [ b - p(L)]"-'
dt (6)
In this expression Ft is the ratio of specific activity of the chloroacetic acid at time t t o that of total chlorine in the system at exchange equilibrium. In the special case in which the exchange rate is first order with respect to each of the exchanging species, equation 6 reduces t o In (1
- F t ) = -k,(a + b)t
(7)
It is interesting to note that this expression is identical t o that for a reaction of similar order in an equilibrium system.
Experimental The chloroacetic acid used waa reagent grade material fractionally crystallized twice from benzene. Prepared solutions were stored in the dark at 0'. HCl solutions were "spiked" with H - C P obtained from the Radioactive Isotope Sales Department at Oak Ridge National Laboratory. Groups of reaction cells were prepared by first mixing the chloroacetic acid and HC1 solutions and diluting to a predetermined volume, usually 200 cc. Ten-cc. aliquots were introduced into cleaned Pyrex reaction cells and thoroughly degassed by repeated thawing and freezing on a vacuum line. The cells were then sealed off and if not reacted immediately, stored in the dark at 0" until used. No hydrolysis or exchange was detectable under these conditions of storage for a%least two weeks. Reactions were carried out by immersing the cells in an oil therrgostat controlled at 80 to rt 0.05" and at 110. to f 0.10 All cells were wrapped in aluminum foil during exposure to eliminate light catalyzed exchange or hydrolysis. Following reaction, the cells were rapidly cooled, opened, neutralized by titration with NaOH and titrated potent,iometrically with AgNOa. The solution was quickly filtered through sintered glass, the supernatant chloroacetic acid solution diluted and an aliquot taken for counting. Blank experiments showed that no exchange or hydrolysis was induced by this procedure. Chloroacetic acid fractions were counted in RCL solut'ion type Geiger tubes with an annular volume of 33 cc. necause the Cla6 p is not highly energetic (0.714 Mev.) and the wall thickness of such a tube is roughly 30 mg./ cm.*, counting efficiencies were quite low. We nevertheless felt the advantage of reproducibility made this method preferable to the several solid phase counting techniques for this radiation. Observed counting rates for the chloroacetic acid fractions of from 25 to 600 c./m. above background were obtained. These low counting rates required that extremely long counts had to be taken, in some cases as long as three hours. In all cases a sufficient number of counts was accumulated to achieve an expected standard deviation for the net counting rate of less than 2%. The total counting rate for a cell in a given set was ohtained directly by counting an aliquot of the original reactant mixture. The fraction of equilibrium exchange at time t , F t , was obtained from the expression (Specific activity of CH2C1COOH)t Ft = (Specific activity of total chlorine)k.,,
.
Sept., 1959
1427
CHLORINE EXCHANGE BETWEEN HYDROGEN CHLORIDE A N D CHLOROACETIC ACID TABLE I DATASUMMARY FOR EXPERIMENTB AT 90.0" T = 363.2"
t
(hr.)
(CHZC~COOH),~ mole/l.
(HCD, mole/l.
Countdmin. :n CHGlCOOH
F
k h (hr.-l)
kx (1. mole-' hr. -1)
Set I, 876 c./m.
0.00 1.00 2.00 4.00 8.00 21.00 37.00
0.0405 .0403 .0399 .0392 .0377 .0335 .0294
0.0186 ,0187 .0191 .0198 .0214 .0257 .0298
0.00 1.00 2.00 4.00 9.00 17.50 32.00
0.1214 .1212 .1201 .1175 .1126 ,1048 .0935
0.0186 .0188 ,0198 .0225 ,0274 ,0352 .0465
0.00 1.oo 4.00 17.00 23.00 32.00
0.0809 .0807 .0777 .0697 .0659 .0617
0.0186 .0188 .0219 .0298 .0336 ,0378
0.00 1.00 2.00 4.00 8.00 21.00 37.00
0.0405 .0404 .0400 .0391 .0374 .0336 .0292
0.0093 .0094 .0097 .0107 .0124 .0161 .0205
0 0.038 .074 .138 .271 .546 .743
0
23 44 80 152 271 324
Set 11, 896 c./m. 0 0 68 0.088 129 .168 242 ,322 407 .564 538 ,802 529 .884
...
... ... ...
... 0.829 .912 .871
x
...
... ...
0.854 X lo-'.836 .841 .815
Set 111, 873 c./m. 0
0
47 161 405 445 47 1
0.067 .236 .664 .772 ,872
... e . .
... 0.877 X lo-* .891 .848
Set V, 438 c./m. 0
12 23 44 77 149 179
0 0.034 .065 .128 .231 .503 .60G
...
*.. ... 0.861 X .993b .880 .880
0.655 .653 ,625 .671 .637 . .625
... ,658 .656 .694 .660 661 .480b
.
... 0.697 ,676 .646 .646 .646
... 0.693 .675 .689 ,661 .671 .647
Set VI, 1775 c./m. 0.0362 0.00 0.0405 0 0 ... 0.697 .0404 1.00 48 .052 ... ,0363 .688 .0401 93 .095 .0366 2.00 ... .718 .0393 ... 178 .198 .0374 4.00 ,658 ,0374 312 .365 .0393 0.882 x 9.00 .041B .655 .0350 46 1 17.00 .575 .848 550 .725 .703 .0332 .0435 24.00 .829 553 .775 .0312 .607 ,0455 32.00 ,815 Conccntrations arc quoted for 363.2"IZ. These valucs were not included in the calculation of the average.
...
a
Results and Discussion Reaction rate constants for exchange, k,, mere calculated using expression 7, L e .
- - 2.303 log ( I - Ft) s -
(a
+ b)t
lteaction rate constants for hydrolysis were calculated assuming a pseudo first-order process, Le.
Total chloroacetic acid concentrations were used in both expressions. The half-time for the exchange was much shorter than for the hydrolysis. Significant hydrolysis rate data were therefore usually obtained only after 20-30% exchange. In Table I are summarized in detail our results a t 90.0". In Table I1 are listed average rate con-
stants for exchange and hydrolysis for the sevcral temperatures. Under these experimental conditions the hydrolysis is adequately described by the pseudo first-order relationship. The consistency of the reaction rate constants for exchange over quite wide ranges of reactant concentrations indicates that this is a process first order with respect t o the chloride concentration and to the total chloroacetic acid concentration. Our results suggest either that a negligible fraction of the chloroacetic acid is present in the dissociated form or that ionic and molecular forms exchange a t essentially the same rate. Wrightg has obtained the following relationship for the temperature dependence of the ionization constant for chloroacetic acid up to 40". (9) D. D. Wright, J. Am. Chem. Soc., 66, 314 (1934).
R. A. KENNEYAND F. J. JOHNSON
1428
Vol. 63
TABLE I1 nevertheless, used this expression to estimate ionization constants at the temperatures of our SUMMARY OF REACTION RATECONSTANTS FOR EXCHANGE experiments. It was estimated that the initial AND HYDROLYSIS Temp., OK.
Set
(CHsC1COOH)o,O mole/l.
(HCl)o,'" mole/l.
k x (1.
k h (hr.-l)
mole-' hr.-l)
353.2
I1 111 VI
0.1222 ,0815 .0408
0.0187 ,0187 ,0374 Av.
3.41 X 0.232 3.16 .235 2.95 .239 3.17 X 10-8 0.235
363.2
I I1 111 V VI
0.0405 .1214 ,0809 .0405 ,0405
0.0186 ,0186 .OH6 ,0093 ,0362 Av
0.871 x .837 ,872 ,874 ,843 0.859 x
0.644 .666 ,662 .673 .675 0.664
I I1 111 V VI
0.0402 .1205 ,0803 .0402 ,0402
0.0184 .0184 .0184 .0092 .0359 Av.
2.20 x 2.17 2.14 2.34 2.10 2.19 x
1.63 1.58 1.61 1.60 1.59 1.60
373.2
.
0.0204 5.11 x 3.67 0.0190 5.35 3.66 Av. 5.25 X 3.67 5 Concentrations are quoted for the temperatures of the reaction. 383.2 VI11
XI
0.1038 0.0245
0.60
0.40
0.20
0.00 i
-F
f=fl
-0.20
degree of dissociation of chloroacetic acid in our experiments varied from 3.1% (set I1 at 353.2"K.) to 1.7% (set VI a t 373.2"K.). This, of course, decreases during a reaction because of the increasing ratio of HC1 to chloroacetic acid. In order to evaluate the relative importance of the ionic species as a participant in the exchange, a limited number of experiments were performed a t 80" in which the initial mixture was titrated to pH 7.5 with NaOH before reaction.1° The results of these experiments are summarized in Table 111. TABLE I11 353.2'K., 2020 c./m. Exchange of C1- with CHzClC00Initial pH 7.5) co.unt,J (CHSClmin. in COO-), CHzC1Time, (CI-) hr. mole/i. mole/l. COOF kx'
0 5.5 12.5
0.00899 .0455 .OF313
0.1233 .0868 ,0319
0 47.1 110.4
0 0.054 0.139
Apparent reaction rate constants, kx', calculated according to equation 7 indicate a much slower exchange rate for the ionic species than for the undissociated acid. The suggested increase of this apparent constant with increasing acidity is also consistent with this conclusion. No further experiments were performed to test the role of the ionic species in the hydrolysis reaction. The consistency of the reaction rate constant for hydrolysis suggests only that it is not markedly more important than that of the undissociated molecule. We have therefore assumed the rate constants listed in Table I1 to be characteristic of reaction by the undissociated molecule. Any error introduced into the rate constants by the use of the total chloroacetic acid concentrations in their calculation would then seem to be within the observed experimental variation. The activation energy and frequency factor for the exchange reaction were evaluated in the usual way from a plot of log k, vs. 1/T (Fig. I). It was found that k, = 0.95 X 10" exp [( -24,500 f 500)/RT] 1. mole-' sec.-l
-0.40
... 0.051 0.091
(9)
Figure 2 shows the corresponding plot for the hydrolysis rate constants from which kh = 2.11
x
loQexp [( -24,800 f 800)/RT] SeC.-'
-0.60
(10)
This rate constant, of course, includes the water concentration to an unknown exponent. Despite the slightly better fit of the hydrolysis constants I I I I I I I I I t,o the logarithmic plot, a greater uncertainty has -0.80 been ascribed to the Arrhenius activation energy 2.60 2.66 2.72 2.78 2.84 for this reaction. This is a consequence of the 1/T X 108. Fig. 1.-Plot of k. os. 1/T for the exchange reaction E, = somewhat larger expected deviation for the rate 24,500 f 500 cal./mole. constant ratios. Using the relationship" k , = (kT/h) exp [AS*/Rl 6669 53 log K = - -- 89,40136 log T exp [-AH*/RT] with AH* = E, RT, AS* for T 0.05233102' + 226.1390 the exchange reaction was calculated to be -10.6 (10) These results were obtained by Mr. Kineiro Aizawa in this Constants calculated using this equation were Laboratory and are used with his kind permission. somewhat greater than those observed experi(11) A. A. Frost and R . G. Pearson, "Kinetics and Mechanism,' inentally a t the higher temperatures. We have, John Wiley and Sons, Inc., New Tork, N. Y., 1953,p. 96.
+
-
Sept., 1959
TEMPERArURE-INTERFACIAL
TENSION O F ALKYLBENZENES AGAINST
cal. mole-' deg.-l. This evaluation was made for 80" and the standard state is one mole per liter. A calculation of the frequency factor for the exchange was made using a simple gas phase collision rate expression, i.e.
collision dimeter, was estimated to be 4.5 This estimate was based upon an ionic radius for chloride of 1.81 and upon a molecular radius of 2.7 A. for chloroacetic acid. The latter value was obtained from the expression13 r = 0.66 X 10-8V,'/a in which V , is the molar volume of chloroacetic acid. When expressed in comparable units, this calculated frequency factor is, a t 353.2"K., 2.06 X 10" 1. mole-' sec.-l, a value in reasonable agreement with the experimental. Our results indicate that the exchange takes place through a simple bimolecular displacement type reaction
1429
WATER
r
-1.20
I\
-1.40
u12, the
A.
a.12
C1-*
-2.20
+ CHZClCOOH = CHzCl*COOH + C1-
Dawson arid Pycock4 measured pseudo first-order constants for the hydrolysis a t 25 and 45'. The rate constant corresponding to k h , when expressed in the time unit we have used, was found by them to be 2.25 X st 45' Equation 10 predicts that a t this temperature, k h = 1.92 X set.-'. Our activation energy for hydrolysis, however, is considerably smaller than the 28,000 cal./mole calculated from their results a t the two temperatures. The similarity in our activation energies for exchange and hydrolysis may be fortuitous or it may suggest the interesting possibility of an acti(12) L. Pauling, "The Nature of the Chemical Bond," Cornell University Press, Ithaoa, N. Y . , 1948, p. 346. (13) E. A, Moelwyn-Hughes, "The hinetlcs of RPactioiis iii Solutions," Oxford Univeisity Press, London, 1047, I). 7.
- 2.40
L tt 1
-2.60
2.60 Fig. 2.-Plot
\
\t
I
1
2.66
I
I
2.72
I
1
2.78
I
1
1
2.84
l / T x 108. of log kh vs. 1/T for the hydrolysis reaction
E, = 24,800 f 800 cal./mol.
vated complex involving an acid molecule, a chloride ion and one or more molecules of water. The complex may then decompose by two paths, replacement of chloride being the more probable. We wish to express our appreciation to the Research Corporation for a grant in support of this WOTlC.
THE TEMPERATURE-INTERFACIAL TENSION STUDIES OF SjOilfE ALKYLBENZENES AGAINST WATER BY JOSEPH J. JASPER AND HELEN R. SEITZ Department of Chemistry, Wayne State Universzty, Detrozt, Mzchiyan Recezved February 10,1959
This is the third of a series of studies involviug the thermodynamic properties of the interfacial region betweee 8. number of homologous series and water. The present paper included the four lower alkylbepzmes over the temperature range 19.45 to 79.30'. The least squares were applied to formulate empirical e uations which express the interfacial tensions as functions of the temperature. These equations were employed t o calcaate the entropy, enthalpy and latent heat of formation per cm.2 of interface. The calculated results, together with the densities of the mutually saturated liquids and their interfacial tension valiies for four temperatures, are presented in tables.
The iiiterfacial regioii between Iwo relatively irnniiscible liquids consists of a transitional concentration gradient, the depth of which is determined by the magnitude of the intermolecular interaction between the unlike molecular species. Since this effect increases with increasing temperature, the mutual solubility of t8he two liquidh increaseb, aiid it is obvious, therefore, tliat there must be an expansion of the tmnsitionitl phwe as the wit icnl solution teniperature is approached. As a
roiibequeiice of t,his inciwdiig holubility an$ ,colirornitant increasing thickness of the transitioiial phase, it might be expected that the magnitude of the interfacial free surface energy would vary invwsely with the temperature. The linear surface tension -temperature relations of liquid-gas surfaceh :we iiot realized, however, for liquid-liquid interf:tces. Kxperiinental evidence proves thRt Ihc iu terfncid teiision dwayh decreiises more rapidly than the teniper:iture increases, giving curveb
