T. I. CROWELL, J. E. HICKS,AND C. C. LAI
2116
The Rate of Reaction of Chloroacetate Ion with Thiocyanate in Concentrated Solutions
by Thomas I. Crowell, Jackson E. Hicks, and Ching Chih Lai Cobb Chemical Laboratory, University of Vireinia, Charlottesville, Virginia (Received August 19, 1966)
The S N reaction ~ of thiocyanate ion with chloroacetate ion was studied kinetically a t 1 :1 electrolyte molarities (c) of 0.0075-7.7 and a t several temperatures. The rate constants for the potassium salts at 25" are correlated by the equation 6 log IC ( M - 1 sec-1) = cl"/(l c"') 0.059~ 0.272. The activation energy is 18.9 kcal/mole a t both c = 2.1 and c = 5.5. Differences in the effects of the cations, Kf and Na+, or the anions, ClCH&OO-, SCN-, and C1-, are not large.
+
+
+
Introduction The high nucleophilic reactivity of thiocyanate ion in solution together with the low melting point and high solubility of KSCN led us to believe that an organic displacement reaction could be studied over a wide range of concentrations, from dilute aqueous solution to the pure fused salt. Reasoning that the organic compounds most soluble in these media would be ionic,' we selected the reaction of chloroacetate ion with thiocyanate. SCNClCHzCOO- + NCSCHzCOOC1- (1)
+
+
Although limited solubility of potassium chloroacetate a t room temperature and rapid decomposition a t high temperature discouraged measurements of the most concentrated solutions, we obtained data at ionic strengths up to 7.7 M and temperatures from 25 to 62.5".
Experimental Section Potassium chloroacetate and sodium chloroacetate were prepared by neutralizing reagent grade chloroacetic acid to the phenolphthalein end point with the corresponding hydroxide. Reaction mixtures were made up by conventional methods, titrated with iodate, and checked with silver nitrate as in previous studies with n-alkyl bromides.2 I n very dilute solutions, because of the difficulties previously encountered in studying the thiocyanate-bromoacetate reaction, reaction 1 was followed by analysis for chloride, in order The Journal o j Physical Chemistry
+
to eliminate all errors caused by interference of products with the iodate titration. A gravimetric modification of the method of Backer and Van Mels4 was used. A sample of the reaction mixture containing 0.5 mmole of SCN- was acidified with nitric acid, and silver nitrate solution was added to precipitate silver chloride and silver thiocyanate. Chloroacetic acid did not react under these conditions. After collecting and washing the precipitate in a sintered-glass crucible, 15 ml of concentrated nitric acid was added, and the crucible was placed inside a covered beaker and warmed on the hot plate until oxidation of the silver thiocyanate was complete. The silver chloride remaining was washed, dried a t 110" for 3 hr, and weighed. I n experiments run a t higher temperatures, the reactants were pipeted into the compartments of a divided erlenmeyer f l a ~ k ,warmed ~ in the thermostat, and mixed by swirling. After the desired time interval, the reaction mixture was cooled and titrated. The rate constant so obtained was multiplied by dz6/d,to correct for the decrease in molarity caused by thermal expansion of the liquid. (1) We now know that hydroxyl groups in sufficient number cause high solubility of organic compounds in fused KSCN (175O) or KSCN-NaSCN (130'); see T. I. Crowell and P. Hillery, J . Org. Chem., 30, 1339 (1965). (2) T. I. Crowell, J . Am. Chem. Soc., 75, 6046 (1953). (3) V. K. La Mer and J. Greenspan, ibid., 54, 2739 (1932). (4) H. J. Backer and W. H. Van Mels, R ~ c Trav. . Chim., 49, 363 (1930). (5) V. K. La Mer and M. E. Kamner, J. Am. Chem. SOC.,57, 2662 (1935).
REACTION OF CHLOROACETATE IONIN KSCN SOLUTIONS
In contrast to alkyl bromides, which consumed 99% of the theoretical quantity of thiocyanate ion,2 chloroacetate ion appeared to proceed only about 97% to completion in the more concentrated potassium thiocyanate solutions where it was feasible to observe the final stages of the reaction. (The chloroacetic acid used yielded 99.9 rt 0.301, of theoretical chloride ion upon alkaline hydrolysis ; hydrolysis in these kinetic runs was negligible.) It was found that the product, prepared as described below, reacted slightly with the titrating solution, causing a 1-4% error in the calculated thiocyanate concentration. This behavior of the reaction product is similar to that of the Bunte salt formed by the action of thiosulfate ion on alkyl halides, though error can be minimized in the latter case by avoiding an excess of iodine in the titratione6 The reaction product, potassium thiocyanoacetate, was isolated by evaporating reaction mixtures to dryness and extracting with boiling methanol. The crystals so obtained showed a sharp infrared peak at 4.6 p . The compound was also prepared by dropwise addition of 50 g of ethyl chloroacetate to a boiling solution of 33 g of sodium thiocyanate in 150 ml of ethanol. The precipitated sodium chloride was filtered off, water was added to the filtrate, and the organic layer was separated and dried over magnesium sulfate. The ethyl thiocyanoacetate fraction boiling at 117-119" (12 mm) showed the sharp RSCN peak at 4.6 1.1. Rearrangement to ethyl isothiocyanoacetate was effected in poor yield by heating 2.5 hr at 250" under an atmosphere of nitrogen. The product, bp 75-80" (0.9 mm) showed a broad RNCS band at 4.64.9 k . Both esters could be hydrolyzed by adding just enough potassium hydroxide solution to cause the compound to dissolve in boiling water. Evaporation of the solution obtained from the unrearranged ester produced crystals of potassium thiocyanoacetate. The infrared spectrum of this solid was identical with that of the reaction product except for a small peak a t 4.9 p .
Results The second-order rate constants, k , are plotted logarithmically in Figure 1 against the square root of the electrolyte molarity, c. This electrolyte concentration was just that of the reactants except when potassium chloride or sodium chloride was added. Six categories are shown on the graph. The solid circles represent a set of runs' in which the initial potassium chloroacetate concentration [KA] was fixed at 0.1 M and the potassium thiocyanate concentration [KSCN] was varied to give the c value shown. I n the runs8 indicated by open circles, the concentration ranges of [KAJ and [KSCN] were 0.390-1.608 and 0.508-6.88 M ,
2117
1.4
1.2
1.0
* g +o.8
0.6
0.4
0.2 0
0.4
0.8
1.2
1.6
2.0
2.4
2.8
C1/%
Figure 1. Log 2rs.cl/z at 250, Experimental points as explained in the text, curve calculated by eq 2.
respectively. The ratio [KSCN]/ [KA] varied from 1.2 to 8.2. A set of five experiments, starting with [KA] = 0.131 M, [KSCN] = 0.202 M, and [KCl] = 0.1, 1.0, 2.0, 3.0, and 4.0 M is shown by squares. The reactions followed by gravimetric chloride analysis (Table I) are plotted as shaded circles o in Figure 1. The k value reported4 for initially equal concentrations of KA and KSCN, 0.1 M , is shown by X. Two runs in this concentration range followed by iodate titration, uncorrected, are denoted by crosses The quantity c in Figure 1 is the total electrolyte molarity, which is equal to the molar ionic strength sinceonly 1:1electrolytes were employed.
+.
Table I: Kinetic Runs in Dilute Solution lo",
[KAli,
M 0.00530 0.01059 0.02118 0.0424 0.1059
M
IKSCNli, M
8ec -1
0.00215 0.00430
2.9 2.73
0.00860
2.76
0.0172 0.0430
3.07 3.58
-1
10'koorq
M
-1
8ec -1
2 . 2 5 f0 . 1 5 2.42 f 0.10 2 . 6 0 i0 . 1 0 2 . 9 7 f 0.10 3 . 5 5 0.09
*
Although the hydrolysis of chloroacetate ion is very slow at 25" and pH 6 (1% complete in 3 months), it was appreciable in the reactions of Table I. The (6) T.I. Crowell and L. P. Hammett, (1948).
J. Am. C h m . Soc., 70, 3444
(7) J. E. Hicks, M.S. Thesis, University of Virginia, 1960. (8) C. C.Lai, M.S. Thesis, University of Virginia, 1965.
Volume 70, Number 7 July I966
T. I. CROWELL, J. E. HICKS,AND C. C. LAI
2118
specific hydrolysis rate was determined to be 1.3 f 0.2 X sec-l. The k values in Table I were corrected for this hydrolysis by subtracting from the measured chloride the quantity calculated to be formed by hydrolysis in that time interval. The corrected chloride concentration could be used in the simple second-order integrated rate equation because the runs with the larger corrections went only a few per cent toward completion. For example, after 3.69 X loe sec, a 100-ml sample from the second run in Table I produced 6.1 mg of silver chloride, or 0.00043 mole/l. Hydrolysis accounted for 0.00005 mole/l., leaving 0.00038 M as the chloride concentration produced by thiocyanate displacement. The correct value of [KSCN] a t this time is then 0.00392, and to use 0.01059 - 0.00038 = 0.01021 as [KA] (rather than the correct value of 0.01059 - 0.00043 = 0.01016) causes an error of only 0.5% and eliminates the necessity of using more elaborate differential equations. Possibly the measured hydrolysis rate is that of a rate-controlling internal displacement to form the alactone. This species could react with thiocyanate, if present, adding a first-order component of 1.3 X sec-'. The correction described above would, however, still yield the desired rate of the external S N ~ attack by thiocyanate. The estimate of reliability of the runs of Table I, given in the last column, is based on the deviations from linearity in the second-order plots, weighing errors, and in the first run, the uncertainty in the rate constant for hydrolysis. The uncertainty in k above c = 0.2 is generally =k5%. The rate constants for c = 3.70 and c = 5.46 were reproduced within 1% by different operators using new solutions. Because of the possible effect of the products on the titration, however, the error in these experiments is estimated at *2%. At the highest concentration, c = 7.70, the small change in thiocyanate ion concentration due to its eightfold excess caused a greater uncertainty. The effect of temperature on reaction rate at two different ionic strengths is shown in the Arrhenius plots in Figure 2. The activation energy E, calculated from the slopes is 18.9 f 0.3 kcal/mole at both c = 2.05 and c = 5.46. (The respective rates at 25" are 8.6 X and 2.08 X 10-5 M-' sec-l and the entropies of activation, calculated using the value AH* = E, RT = 18.3 kcal/mole, are -20.3 and -18.5 cal/deg mole.)
Discussion The rate constants for reaction 1 of the potassium salts a t 25" are correlated by eq 2. If b is set equal to 0.059, the curve drawn in Figure 1 is obtained. The The Journal of Physical Chemistry
k
+ + cl/a
log - = - bc ko 1 CLI1 ordinate is chosen as ;I' in order to show both the approach to the Debye-Huckel limiting slope of 1.0 (dotted line) at low concentrations and the practically linear portion of the curve from cl/* = 1 to 2.5. The average deviations of the points from the curve is 0.017 log unit, omitting the points a t c'/' = 0.273 and 7.70. The extrapolated value of log k in infinitely dilute solution is 0.272. Figure 3 shows the correlation of log k with c"~. In the absence of satisfactory activity coefficient data for the reactants and the activated complex, eq 2 is semiempirical. (It must be realized that a t the highest concentrations studied, the average distance between centers of ions is about 5 A; the radius of 3.0 L
I
2.6
2.2
4 bo
+ 1.8 W
1.4
1.0
I
I
3.0
3.05
I
I 3.1
I
I
3.15 3.2 10t/T.
3.25
3.3
3.35
Figure 2. Arrhenius plots at c = 2.05 (lower line) and c = 5.46 (upper line).
0
0.2
0.4
0.6
0.8
1.0
1.2
1.4
1.6
1.8
2.0
2/:. Figure 3. The experimental points are the same as in Figure 1 (omitting 0, and x), and the values calculated by eq 2 (dotted line) are plotted against c":, showing the degree to which they are approximated by 6 log k = 0.610c'/' 0.225 (solid line).
+,
+
+
REACTION OF CHLOROACETATE ION IN KSCN SOLUTIONS
[HPO],M .
Figure 4. Correlrhion of rate with water concentration.
K + is 1.3 A; SCN- is about 7 A; and there are only five HzO molecules per positive ion.) However, in view of recent interest in kinetics in concentrated electrolyte solutionsg and the paucity of examples, it seems worthwhile to present the data even though complete theoretical interpretation is lacking. The substitution of chloride for chloroacetate ion increases the rate constant only 5 4 % in the higher concentration range, barely outside experimental error. This same substitution actualy partially occurs during
21 19
the reaction and no general upward or downward trend in the second-order plots was observed. Preliminary results indicate that sodium ion depresses the rate about 10% in comparison with potassium. The values of AH* and AS*show that there are no large parallel changes in these quantities in the medium concentration range. It would require more precise data to determine whether the increase in k with increasing electrolyte concentration is indeed chiefly associated with larger AS* values. This is approximately the case in dilute solutions: differentiation of the ionic atmosphere part of AF* predicts contributions to AS*/R and AH*/RT in the ratio 1.58/0.56 as the ionic strength is increased in water at 25”.1° Finally, we note that k increases linearly as the actual water concentration of the reaction mixtures decreases, as shown in Figure 4. The same relationship does not hold in the solutions containing KCl, however.
Acknowledgment. We are grateful for the support of the U. S. Army Research Office (Durham).
(9) (a) B. Perlmutter-Hayman and G. Stein, J . Chem. Phys., 40, 848 (1964); (b) B. Perlmutter-Hayman and Y. Weissmann, J . Phys. Chem., 68, 3307 (1964); (c) J. W. Gryder, J . Chem. Phys., 37, 718 (1962); C. W. Davies, P r o p . Reactbion Kinetics, 1, 163 (1961). (10) S. Glasstone, K. J. Laidler, and H . Eyring, “The Theory of Rate Processes,” McGraw-Hill Book Co., Inc., New York, N. Y., 1941, pp 436,437.
Volume 70,Number 7 Julg 1966
