1928
PAUL ALLEN,JR.,AND LEOREICH
Vol. 64
TIIF, DISPROPORTIONATION OF p-TOLUENESULFINIC ACID IN AQUEOUS SOLUTION BY PAULALLEN,JR.,AND LEOREICH’ Department of Chemistry and Chemical Engineering, Stevens Institute of Technobgy, Hoboken, New Jei sev Received June 67,l g 6 0
A kinetic study has been made of the disproportionation of p-toluenesulfinic acid in water over a range of acidity and temperature; iodide ion served as a catalyst. The disproportionation has an initial unsteady state period under certain conditions and then proceeds by a reaction second order in sulfinic acid and, a t low pH, approxiaately first order in hydrogen ion. A mechanism is postulated and an equation derived relating rate constant to hydrogen ion concentration. The action of iodide ion as a catalyst is discussed.
assigned to this class of compounds. Three Introduction The thermal disproportionation of sulfinic acids formulas have been advocated 0 0 0 0 to sulfonic acids and thiolsulfonates has been known since the latter part of the nineteenth century.2 However, the kinetics and mechanism of this type of reaction have not been reported in the l i t e r a t ~ r e . ~ I I1 I11 Because of this, a study involving the kinetics and mechanism of the reaction in aqueous solution was Without going into the arguments suffice it to say undertaken. that the predominant chemical and optical evidence The disproportionation occurs according to the favors structure II.8,g equation Sulfbic Acids.-The exact structure of sulfinic acids has also been a matter of controversy. Two 3RS02H +RSOzSR + RSOaH HzO (1) forms may exist The stoichiometry of this reaction has been well 0 0 established by various ~ o r k e r s . ~ ~One ~ ~ investi5 // // gator4 found that the presence of hydroquinone or R S 4 H and RS-H catechol had no effect on the course of the reaction, indicating that free radicals are probably not inIV V volved as int,ermediates. When an attempt was made to carry out the reaction in the non-polar The predominance of structure IV has recently been medium, diethyl ether, very little, if any, product favored on grounds of infrared and ultraviolet could be isolated.6 However, Bredereck3succeeded in carrying out the disproportionation in nonSulfenic Acids.-In the thermal disproportionaaqueous media. Hilditch7 discovered that the tion of sulfinic acids, it has been assumed that reaction rate could be materially increased by hy- the reaction occurs via an unstable intermediate driodic acid in small amounts and others also used sulfenic acid,2ds6 RSOH. No free sulfenic acids this acid.’,* The thiolsulfonates formed during have been is01ated’~J~and various attempts to the disproportionation were, in general, very in- prepare them have invariably led to sulfinic acids, soluble in water or mixtures of alcohol and water, disulfides, etc. and usually separated as oils. Despite their instability, sulfenic acids have been One of the main difficulties involved in a study of assumed by numerous authors as intermediates in the thermal disproportionation of sulfinic acids is many reactions, because only in this way has it been the possible existence of structural isomers of sul- possible to predict the products actually obtained. l 4 h i c acids and thiosulfonates. Only relatively reIn the mechanism proposed, in the present work, cently has the problem of which isomers predomi- for the disproportionation of sulfinic acids, strucnate been fairly well resolved. tures I1 and IV of thiolsulfonates and sulfinic acids. Thio1sulfonates.-For many years, there has respectively, have been assumed as the predominant been discussion about the correct structure to be forms, and sulfenic acids have been considered to be involved as reactive intermediates. (1) Abstracted from a thesis submitted by Leo Reich to Stevens
+
[nstitute of Technology in partial fulfillment of the requirements for the degree of Doctor of Philosophy, June 1959. (2) (a) R. Otto and 0. Gruber, Ann., 146, 10 (1868); (b) C. Panly and R. Otto, Ber., 10,2184 (1877); (c) R. Otto, ibid., 16, 121 (1882); ( d ) E. Fromm and J. S. Palms, ibid., 39, 3308 (1906). (3) During the course of this study, H. Bredereck, et 52.. Angew. Chem., 70, 260 (1958),postulated s mechanism for such a reaction in lion-aqueous media, similar to the one independently proposed in the present study, for the reaction in polar media (4) J. von Braun and K. Weissbach. Bw.,68, 2836 (1930). (5) L. Horner and 0. H. Basedow, Ann., 618, 108 (1958). ( 6 ) C. 9. Marvel and R. S Johnson, J . Org. Chem., 18,822 (1948). (7) T.P. Hilditch, J . Chem. Sac., 1091 (1910); 1091 (1911); C. M. Here and 8. Smiles, ibzd., 2359 (1924); J. Cymerman and J. L. Lowe, zbzd., 1666 11949). (8) S. Smiles and D. T. Gibson, ibid., 176 11924).
(9) H. Gilman, L. E. Smith and H. H. Parker, J . Am. Chem. Soc., 47, 851 (1925); L. D.Small, J. H. Bailey and C. J. Cavallito, ibid., 71,3565 (1949); G.Leandri and A. Tundo, Ann. Chim. (Rome),47,575 (1957): J. Cymerman and J. B. Willis, J . C k m . Soc., 1332 (1951). (10) H. Bredereck, et al., Ber., 88, 438 (1955). (11) 5. Detoni and D. Had% J . Chem. Soc., 3163 (1955); but see E. N. Guryanova and Y. K. Syrkin, Zhur. frz. Khim., 93, 105 (1949). (12) W.Autenrieth and H. Hefner, BeT., 68,2153 (1928); T.Zincke and J. Baeumer, Ann., 416. 86 (1916); T. Zincke and I(. Eismayer, Ber., 61, 751 (1918). (13) K. Fries has reaorted (Ber., 46, 2965 (1912)) the isolation of one sulfenic aoid, 1-anthraquinonesulfenic acid, but this has been questioned (P. N. Rylsnder, J . Oro. Chem., 91, 1296 (1956)). (14) (a) 0. Hineberg, Bar., 36, 107 (1903): (b) N. Kharasch. S. J. Potempa and H. L. Wehrmeister, Chem. Reua., 89, 269 (1946); also refs. 2s. 2d. 4.
Dec., 1960
DISPROPORTIONATION OF ~TOLVENESULFINIC ACID
Experimental Sodium p-toluenesulfinate, Eastman Kodak Co., was extracted with hot benzene, dissolved in tm ethanol-water mixture and heated with decolorizing carbon, Norit. After filtration and cooling, the solution deposited white crystals of sodium sulfinate dihydrate. The free acid was not prepared, but for each experiment a fresh aqueous solution of the salt was acidified with sulfuric acid to the desired pH. Hydriodic acid was added to supply iodide ion when desired.16 About 200 ml. of solution was placed in a 3-necked flask immersed in a constant temperature bath and air was removed by passing dried nitrogen isto the flask above the surface of the solution. An hour was allowed for the flask to reach the bath temperature. From time to time, aliquots were removed with a pipet for analysis; the time at which the first sample was withdrawn was taken as zero time for the calculations. After the removal of a sample, nitrogen was again passed over the solution. The analysis of sullinic acid was essentially by the method of + l e d e ; the sample was run into ice and water, made &&ne and titrated, using excess permanganate and then arsenious acid solution, about 0.05 N. A test showed that p-toluenethiolsulfonate did not interfere with the titration. The pH remained essentially constant during a run. As disproportionation proceeded, thiolsulfonate settled out. To determine if it decomposed under the experimental conditions employed, thereby affecting the permanganate titration, samples of thiolsulfonate were heated at 851and 90' with acid and iodide ion. There was no effect on permanganate up to about 35 hours of heating, On ion er heating there was a slight effect, but small enough to%e neglected.
TABLE I DISPROPORTIONATION OF ~-TOLUEXESULFINIC ACID IN AQUEOUS SOLUTION Run no.
1 2 3 4 5
Initial concn. of sulfinic acid, moles/l.
0.070
Temp., OC.
Iodide ooncn.,
2.6 2.4
7. 2.2
3 2.0 2-1.8 x 1.6 E& 2 1.4 1.2 1.0 0.9 0 5 10 15 20 25 30 35 40 45 50 55 60 Time, hours. Fig. 1.-Reciprocal of net ml. of permanganate used in titration us. elapsed time.
2.4 2.2 I
-
2.0
I
$
1.8 1.6
$ .
1.4 I
ri
kr,
0.34 .46 .37 .065 .33 .35 .36 .090 .43 .74 .51 .80
.32
I
X
mg./100 ml. 1Jhr. mole
1.05 5.7 0.95 14.2 0.90 5.7 ,074 1.05 .079 0.0 1.05 .081 2.8 6 .037 1.05 5.7 ,044 1.05 5.7 7 ,065 5.7 8 1.55 ,071 9 0.70 5.7 5.7 .30 10 .060 11 ,060 28.5 .85 12 .20 5.7 .075 1.05 .070 5.7 13" 14 0.40 5.7 .074 0.0 .069 15 1.55 16 85.0 .073 1.05 5.7 17 ,075 0.65 5.7 0.40 .073 5.7 18 20 5.7 .078 90.0 1.05 21 0.60 .074 5.7 22 5.7 .079 0.38 24 5.7 .069 1.55 In run 13, 0.7 g. of sodium sulfate per 100 present. .071
79.9
PH
1929
0.8-
K 0
10
30 40 50 60 70 Time, hours. Fig. 2.-(Same as Fig. 1). 20
80
90
rate constant IC, was estimated from the linear portions of the plots, and obtained using the expression 1/T = 1/TO Nkrtl2V (2) where TO=net ml. of permanganate used in titration, a t zero time; T = net ml. of permanganate used in titration after the elapsed time, t ; A' = the normality of the permanganate solution; B = the volume of the aliquot used.
+
Results The experimental results have been summarized in Table I. The table shows the effect of acidity and iodide ion concentration on the second-order rate constant a t various temperatures and initial sulfinic acid concentrations. As the pH decreases, the rate constant increases (the iodide ion concentration being maintained at 1.23 5.7 mg./100 ml. of solution for most of the runs). 0.092 ml. was From runs 1 and 5, it can be seen that when the pH is kept constant (1.05), and the iodide ion concentration varied from 2.8 to 5.7 mg./100 ml., the rate The data obtained indicated that the reaction could best be described as second order in SUWC acid. Plots of 1 / T constant remains essentially constant (0.34 and 0.33). us. time, t , afforded linear relationships, or linear relation- When the iodide concentration is raised to 14.2, ships preceded by initial curvature. The second-order there is only a relatively small increase in the rate constant (runs 2 and 3), as compared with the (15) Iodide ion was added in order t o obtain reasonable reaction increase in the iodide ion concentration. A still rates. Without it, conversions obtained were very low, even after larger increase in the iodide concentration results in reaction times of about one week. a still relatively small increase in the value of the rate (16) P. Allen, Jr., J . o?o.Chem., 7, 23 (1942). .66 .00 .36 .56 .83 .44 .81
1930
PAULALLEN,JR., AND LEOREICH
3.4
of 0.32 as compared with values of 0.34 and 0.35 (runs 1 and 6) in which no sulfate was present. As the reaction temperature increases, the secondorder reaction rate constants increase. Some of the second-order plots show initial curvatures, Figs. l and 2, followed by linear relationships. However, under different experimental conditions, only linear relationships are obtained for the duration of the reaction, (Figs. 3 and 4). The diminution or disappearance of the initial curvature results from increased acidity and in some cases from an increase in the iodide ion concentration.
3.2 3 .0
2.8 2.6
7 2.4
-6
2
2.2
x' 2.0
Discussion
--. h r(
1.8
1.6 1.4
0
10
20
30
40
50
60 70 80 90
Time,hours. Fig. 3.--(Smne as Fig. 1). 7 C
3.2
3.0 2.8 2.6
- 2.4
i
2*2
t,'
2 2.0
L
Vol. 64
The disappearance of the initial curvature in the various second-order plots indicates that there is an initial unsteady state but when conditions are favorable to the rapid establishment of a high steady-state concentration of reactive intermediates, the initial unsteady state condition and also the initial curvature vanish. The duration of the unsteady state condition is found to depend on the iodide concentration. However, the rate constant IC,, obtained from the linear portion of the second-order plot, does not change to any great extent. Thus, when the concentration was 2.8 mg./100 ml. (run 5 ) , the duration of the unsteady state period was about 40 hours, but when the iodide ion concentration was raised to 5.7, the duration decreased to about 15 hours (run 1). However, in both cases, the rate constant ?remained essentially constant (0.33 and 0.34). This indicates that a t the iodide concentration employed for most of the work, 5.7 mg./100 ml., the iodide has only a small effect, if any, on any rate-determining step in the disproportionation. Thus, it is apparent that the role of the iodide ionhs as:a'pure:catalyst .I7 It is postulated that the iodide participates in an equilibrium reaction involving the isomerization of one form of sulfinic acid to another
a
X F* 1.8
:
1.6
Thus, the role of the iodide is to increase the rate of attainment of this equilibrium. A low iodide concentration would mean a long unsteady state
1.4 1.2
(17) An indication t h a t the iodide ion functions as a pure catalyst even at the higher iodide concentrations, 14.2 and 28.5, may be seen from the following treatment. Let UE aeaume t h a t the rate constant observed at 79.Q0.kr,may be expressed by the equation
0
5
15 20 25 30 35 Time, hours. Fig.4.-(Same as Fig.1). 10
40
45
constant (runs 3 and 11) : it should be noted that at these higher iodide concentrations, the pH's were lower and would thus contribute toa higher of IC,. However, when no iodide ion is Present (runs 4 and 15), the reaction rate is extremely slow. In order to determine the influence of effects was added to on the reaction the reaction mixture Of Nn l3. The. effect was slight since the resulting rate constant had a Value
where f(1) is assumed to be Uome function of the iodide ion concentration. It will later be shown that such an equation (without the f(1) term, equation 22) fib most of the data (especially at lower p H ' s ) a t the iodide conoentration of 6.7 mn./100 ml. Bv utilizing the values in Table I for runs 2, 6, 0 and 11. the foilowing {alues of %I) are obtained for each of these runs, respectively, 1.49. 1.29, 1.37 and 1.45; the average value of f(I) ie 1.40. If an experimental deviation of about 10% ia aeaumed for h,then the deviation obtained for f(1) for the runs employing relatively high and relatively low iodide concentrations. would tie well within the experimental error. (In unpubliehed work in this Laboratory. on t h e decomposition of n-dodecanesutfinio acid, a variation in iodide concentration from 0.6 to 4.8 ma./100 ml.. at constant PH, had a nedipible effect on the rate constant, 2.6 to 2.4.)
Dec., 1960
~ I S P R O P O l < T I O N R T I O NOF
p-TOLUENESULFINIC ACIU
period, a longer time to reach equilibrium. Absence of iodide ion should lead to an extremely slow attainment of the maximum equilibrium concentra0
// \\
tion of RS-H
1031
1.2
1 .o
(structure V), resulting in a very
i
0
low rate constant (runs 4, 15, 24). On the other hand, with a relatively high iodide concentration, rapid attainment of the equilibrium should result, leading to a diminution or disappearance of the unsteady state period, as observed. The postulation of an equilibrium involving structures IV and V is made in order to account for the thiolsulfonate structure I1 which was the only form that could be isolated for this type of disproportionation (see the Introduction). The other form of sulfinic acid, IV, would lead to the formation of ai1 anhydride or disulfoxide structure (I or 111). The former would not be expected to form in an acidic, polar medium, and would therefore, a t best, have only a transient existence. The same can be said for the latter structure, since it has never been isolated. The isomerization of one form of sulfinic acid to the other in the presence of iodide ion is attributed to the availability of electrons of the latter. It is found that the hydrogen ion concentration affects the rate constant considerably, (Fig. 5). With increase in pH, the rate constant decreases; a t a pH of 1.05, k, is 0.34, but a t pH 1.55 (that expected from the suifinic acid alone), k, is 0.09. It can be observed in Fig. 5 that a t relatively low pH's, as the hydrogen ion concentration increases, k, increascs in a linear manner, a t a particular temperature. However, this relationship does not hold a t the relatively high pH's (greater than 1.0). Thus, it is apparent that hydrogen ions play an important role in the formation of reactive polar intermediates in the disproportionation reaction. That no salt effect is involved on changing the hydrogen ion concentration is shown by the slight influence on k, of the addition of sodium sulfate (run 13). The Reaction Scheme.-The following reaction mechanism is postulated
g
-z ->
0.8
5 0.6 Y
a
8 al
3 0.4 0.2
0 0 0.2 0.4 0.6 Hydrogen ion concn., moles/l. Fig. 5.-Rate constant us. hydrogen ion concentration, at different temperatures and a t the iodide concentration of 5.7 mg./100 ml. 0
RS=O+
/- kl
+ RS-\\
HzO fast
+ [RSOSOzR] --+
0 0
0
// ki' R&O+ + RS-H +II+ + [RSOSOzR] O \ 0
-+ HZ0 RS-H // fast 0
P
RS--H
+ II+
K4
RS-011
+ RSOaIl
(IO)
+ 11'
(11)
0
1/ K // RS-0- H + I_
+
P
-4RS-SR r)
fast
n
\\
+ 1110
(12)
'6
It is assumed that the concentrations of both forms of the sulfenic acid increase until steady-state concentrations are reached, Le. 0
/
d(RS-I-I)/dt
=
0 and d(RS-OII)/dt
= 0 (13)
Also, the rate of disproportionation, designated simply as Rate, can be expressed by Rate = d[RSO,SR]/dt =
- 1/3 d[RSOzH ]/dt
PAULALLEN,JR.,AND LEOREICH
1932 1
Vol. 64
o
r
!
1 2
1.4-
1.2
-
In the experimental determination of unreacted sulfinic acid, the total sulfinic content [RSOZH1 was determined by permanganate titration and must therefore be accounted for. Thus
1.0
5o.8
,
9
0.6 0.4
0.2
L
_r
01
0.2 0.4 0.6 Hydrogen ion concn., molesfl. Fig. 6.--Rate constant divided by 02 us. hydrogen ion concentration.
0
(19)
and hence, the expression for rat,e becomes Rate = CdCdG
Assuming C4 has a very low value (see the Introduction), equation 14 becomes, finally
(14)
Rate = ( A
From steady-state considerations, we may also write
r
1
01
Rate = k , [RS=O+] RS -
+
Equation 15 may be coiiverted into
Letting C1 = kl,Cz = k,'/K1, a i d
r
o 7
we obtain
Let tin g
+ C2[Hfl)
+ B[H+]) ([H+l[H+l +K)
[
