KIXETICS O F T H E REACTIOX BETWEEN POTASSIUM PERSULFATE AND SODIUhl THIOPHENOLXTE BY T. B. DOCGLAS
Apparently no work has been done on the kinetics of the oxidation of a mercaptan to the corresponding disulfide. In this paper are given the results of a study of the velocity of the reaction between potassium persulfate and sodium thiophenolate in alkaline solution. The kinetics of a similar reaction, that between a persulfate and an inorganic iodide, have been investigated.* Although the reaction is stoichiometrically trimolecular, the results indicate that the reaction whose velocity was being measured was bimolecular, and consequently that the process took place in two steps, one slow and the other relatively fast.
Experimental Procedure Baker's C. P. potassium persulfate was used. This salt was carefully analyzed by reaction with ferrous sulfate and back-titration with permanganate, the average value obtained being 97.6 percent purity, the contamination probably being KHSO, from decomposition. A11 standard solutions of the salt were prepared on the basis of this figure. The thiophenol first used was prepared by the reduction of Eastman's benzene sulfonyl chloride, and collection of the fraction boiling 166'-169~ (uncorr.). The second supply of thiophenol was prepared by the treatment of phenyl magnesium bromide with sulfur, and collection of the fraction boiling 168'-169' (uncorr.). The standard solutions of I i 2 P 2 0 9 were prepared by weighing the salt, and these solutions were used within a day after preparation to avoid the error due to s l o decomposition. ~ In preparing both reaction-solutions, only water which had previously been boiled could be used, sinc,e the sodium thiophenolate, if dissolvrd in ordinary water, was found to decompose at the rate of j percent per hour because of oxidation by dissolved oxygen. ('are was taken not to agitate the solutions and thus introduce air, in the course of their preparation. The sodium thiophenolate solutions were prepared by dissolving approximately twice the calculated amount of thiophenol in sodium hydroxide of the proper titre, removing the solution by filtration or decantation, and diluting with sodium hydroxide of the same titre just before use, in accordance with a standardization of the thiophenolate solution just performed. This was necessary because of a slow but inevitable drop of titre of the thiophenolate solution due to dissolved oxygen. A standardization of this solution im-
__
.4. van Kiss and
L. von Zombory: Rec. Trav. chim., 46,
225
(192;).
POTASSIUM PERSULFATE AND SODIUM THIOPHENOLATE
3 28I
mediately after the velocity measurement gave data from which to calculate the correction on the thiophenolate concentration in the reaction-mixture, since this concentration was slowly decreasing, during the course of the slow reaction, from that calculated from oxidation by persulfates alone. Titration of the thiophenolate solutions was performed with 0.05 N aqueous iodine after diluting the thiophenolate with water to about 0.005 S and acidifying it (no precipitation of thiophenol taking place at this dilution). An attempt to titrate the thiophenol in alkaline solution with iodine, acidify, and then back-titrate with thiosulfate showed that in alkaline solution iodine oxidizes thiophenol further than to disulfide and in no way which could be followed quantitatively. The concentration of excess NaOH in the thiophenolate solution, definite for each series of determinations, fell between 0.005 N and 0 . 2 5 N. One determination was made on each individual reaction-mixture. A quantity of the persulfate solution and a quantity of the thiophenolate solution each containing 0.001equivalent of reactant were set into a thermostat a t 2 j " ( i 0 . 1 ~ and, ) ~ after the temperature of the thermostat had been reached, the persulfate was poured into the thiophenolate, rapid mixing being effected by vigorous swirling of the flaak. The reaction was arrested by poliring the mixture into a liter of water, and, after acidification and addition of starch, was titrated with iodine. A correction of 0.40 ml of 0.05 N iodine, amounting to 2 to 7 percent of the total titration-reading, was found necessary, because of the large amount of water present.
Results The velocity constants are calculated on the basis of a bimolecular and also a trimolecular reaction. The values obtained are given in Table I. Discussion The data show clearly that the reaction whose velocity is being measured is bimolecular. Hence the total reaction is really composed of two reactions, one being much more rapid than the other. Attempting an atomistic picture, one conceives that by a slow reaction there is formed a very reactive intermediate product, which disappears approximately as fast as it is formed. Such a case might be expressed by the following equations:
+
Slow . . . . .SpOs-CsHsS- --+ S208--( z C ~ H E+ S CsHcS.SCsH5;) F a s t . . . . . . .&Os---
f C&S-
+ ?SO,--
+ CcHsS; + CsH5S.
As the concentration of OH- is increased, the velocity constant becomes larger. Hence the OH- acts catalytically. Also, the velocity constant varies with varying concentration of the reactants. The author proposes the following hypothesis as a possible explanation of the mechanism of this catalysis.
3282
T. B. DOUGLAB
TABLE I Reaction Velocities in Oxidation of Sodium Thiophenolate by Potassium Persulfate a t 25'C (The concentration of CsHSSXa indicated below is the initial one. Concentrations of the two reactants are equivalent unless otherwise stated. Calculation of the velocity constant#sgiven below is based on the corrected initial concentrations referred to above. These corrections are not noted here.) Progress of Reaction Time percent minutes
A.
20
28 49
k
bimolecular trimolecular
CBHbSNa= 0.0025 M
a. NaOH = I2
k
0.01
C.
M
3.40 3.55 3.35 I20 3.25 Average: 3 , 4 0 16 28 47
25
c. NaOH = I2
I3 I8 21
29 30 39 51 53
a. 12
26 40 69
b. I1
3.50
I450 I590 1610 1920
d. KaOH =
1750
I4 I8 31
3.50 3,70 3.60 120 3.65 Average : 3 .60
23 38
16 32
53 69
60
0.01
3.75 120 3.75 Average : 3 . 7 5 * At 25.2'.
950
4.60 4,80 4.60 4.55
i100
1330 1570 2160
M*
0.25
7 12
4.55 3.95
60
I . 50
c. C&Sr\la
1750 I 890
1910
s.
2070
25
2080
4' 56 71
2150
2290
980 890 370
NaOH =
=
0.01
0.01
M
M*
4.15 4.25 4.30 4.00
8 16 30 60
490 570 710
880
Average : 4 . 2 0 b. 760
870
950 1530
SaOH =
770 860
0.025
8 I7 45
30 48 69
M 630 800 1030
5.20
5.50 4.90 Average : 5 . 2 0
D.
M*
3,60 3.75 3.85
4.40
18.10
B. CeHaSNa = 0.005 ?vl NaOH = o 005 M 8 19 36
I20
I560
M
4.05 4.23 3.95 25 4.20 40 4.00 40 4.25 44 3.93 I20 3.55 I21 3.70 Average : 4 . 0 0
k
M
0.02j
7 16 32 60
27
43 58 73
k
bimolecular trimolecular
Average : 4 . 6 0
0.0~5
14 I4 25
NaOH = 7
KaOH
I3
b. NaOH = 0.0113 M I8
Progress of Reaction Time percent minutes
K2S208
=
0.01
M
C6HbSNa= 0.01M (approx.) NaOH = 0 . 0 9 %1 (approx.)
1000
I2
2
1220
22
4
1580
42
IO
72
40
6.75 6.70 6.05 4.05
1550 1640 1730 1840
POTASSIUM PERSULFATE A S D SODIUM THIOPHENOLATE
3283
The reaction depends ultimately on the collision of thiophenolate ions with persulfate ions. Let us suppose that temporary combination takes place, and in a reaction so mobile as to remain in equilibrium:
(a). . . . . . . . . . . . . .C6H5Sfor which:
+ SzOs--F', (C6H5S).(&08)---]
this complex being so unstable that its formation does not appreciably affect the concentrations of the original ions. Such a com[CsHsS-] nor [SzOs--], plex would very likely decompose automatically in one of two ways: ( I ) into the ions out of which it was formed, according to the reverse of reaction (a), or ( 2 ) in a manner involving exchange of electrons] such as: ( b ) .. . . . . . . .(C6H5S).(SzO8)--- + CsHsS
+ S*Og---,
following which we must suppose a secondary reaction: (c). . . . . . . . .SzOg---
f CsHsS-
---f
2SOp--
+ CcHsS.
Let us suppose either reaction (b) or reaction (c) sufficiently slow to contribute to the measured velocity. If reaction (b) were the rapid one, thiophenolate would disappear in large amount almost immediately. Since this is contrary to experimental fact, we shall consider (b) as the slow, and (c) as the comparatively fast, reactions. (To assume both reactions slow would greatly complicate the derivations.) In order to attempt to explain the catalytic action of OH-, let us suppose that a complex ion is formed as follows: (d) . . . . . . . . .CcHsS-
+ OH-
(C,HSS).(OH)--,
in mobile equilibrium] for which:
assuming this complex also to be so unstable that its formation does not appreciably affect [ C 6 H S ] ,but to be so highly reactive that approximately all direct oxidation of thiophenolate takes place through oxidation of this complex. Reaction (c) should then be corrected to read: (c')
. . . . SzO8--- + (CCH,S).(OH)-- + 2 SOa--
+ C6H5S + OH-.
There is a third possibility in the disappearance of the hypothetical ion (CeHSS).(SzOs)---: (e). . . . . . . . .(CsH5S).(SzOs)---
+ (C~HSS).(OH)-CeHsS + z S O ~ - - + OH-. ---f
2
3284
T. B. DOUGLAS
The measured reaction velocity, thus assumed to be dependent on simultaneous reactions (b) and (e), will be:
[ (CsHsS). (SzO8)---] [ (C6HbS).(OH)--].
Substituting the equivalents of [ (C&S). ( SZO~)---] and [ (CGH5S). (OH)--] from equations ( I ) and ( z ) , respectively, and factoring, we have:
- d[CsHaS-j dt
=
( K I ki
+ K1 Kz kz [ C s H S ] [OH-]) { [CsHsS-] [S,O,--]}.
Expressed more simply : - d[C6H5S-1 = { k I dt
+ kII [C6HsS-] [OH-]} { [szOs--] [CsH&]
+
1.
The expression kI kII [C6HSS-; [OH-j is the modified velocity constant, and will subsequently be represented by k. An attempt has been made to apply this modification to the data obtained. Each experimentally determined value for k is associated with the average concentration of total thiophenolate of the interval of concentration change of that particular determination, since, [OH-] remaining constant
TABLE 11 Comparison of the Calculated and Experimental Values for the Velocity Constants of the Reaction between Persulfate and Thiophenolate at 25'C (The average concentration of C6HsSxa,as described above, is indicated below .) NaOH
k
k
moles per liter
(calculated)
(found)
CsH,S?rTa 0.01
0.01 13
0.025
0.0022
P*f
3.50 3.55 3.95
3.40 3 ' 50 4.00
C6H5SNa= 0.0040 M 0.01
3 ' 50 3.75
3.60 3.75
0.02j
4.50
4.60
0.005
CGH5SNa= 0.0076 M 0.01
4.20
4.20
0.025
5.60
5.20
POTASSIUM PERSULFATE AND SODIUN THIOPHENOLATE
3285
during a single determination, k becomes a first-degree function of [CsH5S-] in accordance with the foregoing hypothesis. Calculation of the different. values for kI of the modified velocity constant gives a mean value of 3.25. A value for kI cannot easily be determined directly, because it is very difficult to dissolve thiophenol in very weakly alkaline aqueous solutions. Table II affords a comparison between the calculated and experimental values of k. The value used for kII in the calculations is the one which gives the best agreement between the calculated and experimental values for the constant. I n all these calculations the activity coefficient of each ion concerned is assumed to be unity. This is intrinsically, of course, not true. From Table I the experimental velocity constant is seen to fall off very sharply during the course of the reaction when the OH- concentration is high. S o possible explanation of this behavior is offered except to state that oxidation further than to the disulfide stage may take place, with consequent faster lowering of the persulfate concentration than anticipated.
Summary The velocity of the reaction between potassium persulfate and sodium thiophenolate in alkaline solution has been studied, and, although the reaction is stoichiometrically trimolecular, the results indicate that the measured reaction is bimolecular. The values obtained for the velocity constant of the simple velocity equation are found to vary in a more or less regular manner with varying hydroxide concentration or with varying concentration of the reactants. A hypothesis involving increased activation through complex ion formation has been proposed as a possible explanation of this anomaly. t7niuersity of North Carolina,
Chapel Hill, S . C.
