2-DI1VIETEYLAMJNOETHaNETHIOL
HYDROCHLORIDE
Clustering could affect both ion distribution and ion transport. Rlodel calculations were recently carried out by Simons and Kedem* for an array of rectangular pores In an ion-exchange medium, not assuming homogeneous ion distribution. From these calculations it appears that the distribution coefficient is quite sensitive to pore size, at given average charge density, but the relation between salt rejection and salt distribution is only slightly influenced. It is thus possible
3641 that in practice the correlation given in eq 6 is valid in systems for which not all the assumptions of the TRlS model are justified.
Acknowledgments. The authors thank the Office of Saline Wat,er for a grant to support work on hyperfiltration membranes (OSW Contract ?lo. 14-01-0001961). (8) R. Simons w d 0. Kedem, t o be published.
Oxidation Kinetics of 2-Dimethylaminoethanethiol Hydrochloride by Ferricyanide Ion in Acid Medium by pt. M. Chohan, B. P. Sinha, and R. C. Kapoor” Depo,rtment of Chemistry, University of Jodhpur, Jodhpur, Rajasthan, I n d i a
(Received April 34, 1979)
‘The oxidation kinetics of 2-dimethylaminoethanethiol hydrochloride in aqueous hydrochloric acid medium have been described. The stoichiometry is found to be 1: 1 and the substrate is shown to be oxidized to the $:orresponding disulfide. The order is unity both in ferricyanide and thiol hydrochloride. The rate decreases 3n increasing the initial concentration of ferricyanide and hydrogen ion concentration. Addition of ferrocyanide, disulfide, and neutral salt produces no change in rate. A tentative reaction scheme is suggested.
Introduction I n a recent publication‘ the oxidation kinetics of 2mercaptoethylamine hydrochloride by the ferricyanide ion in aqueous hydrochloric acid medium were described. Besides, this oxidant has also been used earlier for oxidation of a number of other thiols.2-6 The present investigations stems from an extention of the work being canried out in this laboratory on ferricyanideoxidation of compounds containing a sulfhydryl group There is a marked divergence in the oxidation scheme of the cwo thiol hydrochlorides largely due to the role of ferrocyanide.
Experimental Section The sample of 2-dxmethylaminoethanethiol hydrochloride (assay 98%) was from Evans Chemetics Inc., 1J.S.A. An approximate solution of this compound was first prepared in air-free double-distilled water. This was subscquently standardized with a standard EoPution of iodine in presence of hydrochloric acids7 A standard solution of potassium ferricyanide (E. Merck reagent grade) was prepared by direct weighing. A fresh solution was prepared for each run. 2,2’-Dimethylaminodiethyl disulfide dihydrochloride (Le., the disulfidr of 1he present thiol hydrochloride) was obtained
by oxidizing the latter with dimethyl sulfoxide.8 The indicator solution of 3,3‘-dimethylnaphthidinedisulfonic acid (B.D.H., L.R.) was prepared in dilute ammonia solution. All other chemicals used were of analytical grade. The kinetics of the reaction wer2 followed by estimating the product ferrocyanide volumetrically with tjhe aid of a standard solution of zinc sulfate as done earlier.g The overall reaction is 3Zn2+
+ 2K4[Fe(CnT)s]
lizZns[Fe(CN)6]2
+ 6K+
(1) R. C. Kapoor, R. K . Chohan, and B. P. Sinha, J . Phys. Chem., 75, 2036 (1971). (2) 0. P. Kaohhwaha, B. P. Sinha, and R. C. Kapoor, I n d i a n J . Chem., 8 , 806 (1970). (3) R. C. Kapoor, 0. P. Kachhwaha, and B. P. Sinha, J . P h y s . Chem., 73, 1627 (1969). (4) I. M . Kolthoff, P. J. Meehan, M. S.Tsao, and Q. W. Choi, ibid., 66, 1233 (1962). ( 5 ) E. J. Meehan, I. M . Kolthoff, and H. Kakinchi, ihid., 66, 1238 (1962). (6) J. J. Bohning and K. Weiss, J . A m e r . Chem. Soc., 82, 4724 (1960). (7) H. Kramer, J . A s s . Ofic. Agr. Chem., 35,285 (1952).
( 8 ) W. T.V. Epstein and F. W .Sweat, Chem. Rea., 67,247 (1967). (9) A . I. Vogel, “Quantitative Inorganic Analysis,” Longmans, Green and Go., New York, N. Y., 1962, p 402.
T h e Journal of Physical Chemistry, Val. 7 6 , N o . 24, 1972
R. K. CHOHAN, B. P. SINHA,AND R,. 6.KAPOOR
3642
Results Order of Reaction. A number of runs having equimolar concentration of ferricyanide and thiol hydrochloride were made in hydrochloric acid and sodium chloride buffer. A plot of [l/(a - x ) ] us. time (where a denotes the initial concentration of the two reactants and II: is the decrease in their concentration a t time t ) gave a straight line for nearly 70y0 of the reaction. The second-order velocity coefficients in all these runs were calculated by using the modified equation
L
@'4
T o
\ zzo
240
;bo
280
sbo
WAVE LENGTHIN m p
Figure 1. Absorption spectra of (a) "L-dimethylaminoethanethiolhydrochloride (1 X N ) , (b) 2,2'-dimethylaminodiethyl disulfide dihydrochloride (1.0 X M ) , and (e) the disulfide produced in the reaction mixture.
The absorption spectra of 2-dimethylaminoethanethiol hydrochloride and its disulfide were recorded using a Hilger Uvispeck spectrophotometer Model H 700-8 fitted with ti thermostatic ariangement. Polymerization Test. Vincyl cyanide and the acidified thiol were kept in a conical flask and the vessel was flushed with nitrogen for about 15 min. Potassium ferricyanide solurion which had been previously flushed with nitrogen was then added and the nitrogen was again bubbled through the mixture. Each experiment was accompanied by a blank control. The reaction greatly catdyzcd the polymerization of vinyl cyanide and a thick polymer was obtained in each case. Stoichiometry~ Here again, both analytical and polarograp2iic methodslO were employed to determine the stoichiometry. The latter method helped not only in the estimation but also the characterization of one of the final products of oxidation, vis., the disulfide. The spectrophot>ometric studies also led to a similar conclusion, 2.6 , the reaction leads to the formation of disulfide as the sole oxidation product of thiol hydrochloride. 'This is demonstrated by Figure 1, where absorption spectra of thiol hydrochloride, the disulfide, and that of a reaction mixture, which had been allowed suEcient time, are recorded. All these methods led to one and the same finding regarding the stoichiometry being 1 : 1. The overall reaction is therefore
SRSII
+ 2~e(cn7)~3-
obtained on solving the differential equation dxjdt = k2(a - x ) between ~ limits tl, tS and xl,z2. The results of a typical run are presented in Table I. A good agreement was found between the bimolecular rate coefficients calculated from the above formula and the ones obtained graphically. It may be mentioned here that the above equation was used because of the difficulty in exact experiment,al estimation of ferrocyanide formed a t the commencement of the reaction. A standard rate constant is k,
=
6.98 X
( 2 x40-2)
where numerator in parentheses denotes the volume of reaction mixture withdrawn for estimation and denominator signifies the concentration as required by factor v/s for conversion of an arbitrary second-order rate constant (kobsd) into the standard second-order rate constant (IC,). Hence IC, = 1.74 * 0.04 A - l min-l and k , (graphically) = 1.73 i: 0.04 J 4 - l min-l. A11 rate constant recorded hereafter are standard rate constant unless otherwise mentioned. Runs mere also made a t constant ionic strength by changing the initial concentration of ferricyanide in one series of experiments and thiol hydrochloride in the other. In a,ll such variations, plots of log [ ( b - x)/(a .- z ) ] against time gave a straight line (where a is initial concentration of thiol hydrochloride and b thal of ferricyanide). For variations in the initial concentration of ferricyanide, these plots are shown in Figure 2 . The secondorder rate coefficients for all such variations as calculated from the slopes ( k , = ( 2 X 303)/(b - a) X slope X 8/s) are recorded in Table 11. The variations in the magnitude of second-order rate coefficients is due to the effect of initial concentration of ferricyanide as the other substrates concentration wields no such influence, The ionic strength was kept constant throughout with the help of sodium chloride. The order in ferricyanide ion was also confirmed by isolating it in a number of runs by keeping thc concentration of thiol hydrochloride in 10-20-fold excess over
--f
RSSR
+ 2 F e ( c N ) ~ ~+- 2H+
The Journal of Physical Chmistry, Vol. 76, No. 94, 1979
(10) 0. P. Kaohhwaha, B. P. Sinha, and R. 6. Kapoor, Proc. Sump. Electrode Processes, 1966, 84 (1969).
3643
~-DIMETHYLAMI.NQETHANETHIQL HYDROCHLORIDE
Time, mi8
Volume of 0.02 M
Volume of
Volume of 0.005 M ZnSOc required, ml
ferrocyanide produced (4,ml
ferricyanide left, ml
4 .7 7 * 95 10.45 12.4 14. i 15.4
0.783 1.32 1.74 2.06 2.35 2.56 2.76 5.92
4.21 3.67 3.25 2.93 2.65 2.43 2.23 2.07
5
10 18 20 25 30 35 40
16.0,
17.55
0.02
AV =
Ivl kobsd X
108
7.00 6.97 6.91 7.01 6.94 7.01 7.00
6.98
x
0.11
10
MI [HS(CH2)2N(CH3)z*HCl] = [K3Fe(CN)e]=-= 2.0 X MI [HCl] = 5.0 X 10-2 M, [KCl] = 11.0 X 10-2M, 2.0 X ? = 0.3 M, temperature 35". a
%lk
1.0
Q
[HCl]
=
>( 10-2
M, I
=
30
/
I
2.18 2.03 2.55 2.21 2.19 2.39 1.72 1.31 1.14 0.98
2.0 1.0 3.0 2.5 3.0 3.5 2.0 2.0 2.0 2.0 5.0
as
HNWU
Figure 2. Second-order plot of four different runs (temperature 35", [thiol hydrochloride] = 2.0 X 10-2 M, M I [K3Fe(CN)6]= (a) 2.5 X IO-* M ; [HCl] = 5.0 X (b) 3.0 X M, (c) 3.5 X MI (d) 4.0 X lo-* M, I = 0.3 M .
Table 110
2.0 1.0 3.0 2.0 2.0 2.0 2.5 3.0 3.5 4.0
ao
IS
10-3
0.3 M ,temperature 35".
-s_l-
the former. The reactions showed a pseudo-first-order kinetics in ferricyanide. These constants are recorded in Table 111. When the pseudo-first-order rate coefficients are divided by the concentration of thiol hydlrochlioride the quotients are found to be practically constant which incidentally confirms the first-order kinetics in the substrate. The total order of the reaction was further verified by an application of Vant IIoff's differential method. Initial rates were measured corresponding to the different initial concentrations C Table LII5 ENS(CHzjzN(CHaj%.HCl]X 10%M
Pseudo-first-order constant, ki X 102 min-l
10.0
1.40
15.0 20.0 25.0 30.0
2.07 2.84 3.56 4.31
[KSI:IFe(CN)8] .= 1.0 X 10-2 MI [HCl] temperature 35".
kl/C
14.0 13.8 14.2 14.1 14.3 =
2.0 MI 1
=
2.4 MI
Figure 3. Relationship between initial rate and corresponding different initial concentrations of reactants (temperature 35", [HCl] = 5.0 X MI I = 0.3 M).
of the two reactants. A plot of log (dz/dl) against log C (where C denotes the concentration of both ferricyanide and thiol hydrochloride) gave a etraight line with a slope equal to 1.8-2 (Figure 3). As remarked earlier, the rate of oxidation is also dependent on the initial concentration of ferricyanide. The increase in the initial concentration produced retardation in rate. This is further demonetrated by the results shown in Table IV, but no simple relationship could be observed between the two. On the contrary, the variation of initial concentration of thiol hydrochloride produced no effect when all other parameters were kept constant (Table I1 and also on the basis of practically constant values of pseudo-first-order rate coefficient when variation of thiol hydrochloride in isolation is followed). The rate was found to be highly dependent on hydrogen ion concentration. The effect was studied in the presence of different concentration The Journal
of
Physical Chemistry,Vol. 76,N o , 24, 1972
3644
R. K. CHOHAN, B. P. SINHA,AND R. 6. K A P ~ O R
Figure 4. Dependence of the rate on hydrogen ion concentration (temperature 35", [K8Fe(CN),] = 2.0 X 10-8 M , [thiol hydrocliloride] = 2.0 X 10-* M , I = 0.3 M ) .
of hydrochiorx acid. The ionic strength was maintained constant in the usual manner. An inverse linear relationship between the rate and [E+]was found to exist as shown PXB Figure 4.
Figure 5 . Dependence of the rate on dielectric constant (temperature 35', [K,Fe(CN),) = 2.0 X M , [thiol hydrochloride] = 2.0 X M , [HCIJ = 5.0 X lo-' M , I = 0.3 M ) .
Table Vu
[KCN] X 103 $1
0.0 4.0
-____^_--___
Table IV0
2.19 2.04 2.04
7.0 Pseudo-first-order rate constant,
[IC 103 .W
k~ X 10%min-*
26.0
2.34 2.46 2.64 2.95
15.0
10.0 8.0
[HS(CX1[,)z$\T(GHi)z.WCl]= 2.0 X 10-aM, [BCl] = 2.0 M , I = 2.4 M , temperature 35".
The rate of reaction remained uninfluenced by ionic strength, though some specific ion effect appeared likely. Thc xralues of second-order rate coefficients were found to be 2.17 f 0.07 for NaCl and 1.61 & 0.20 for KCl a t reactant concentrations 2.0 X 10-2 M , [HCI 1 = 5.10 >< 10-2 M a t 35' while the ionic strength was varied from 0,3 to 0.6 M . Likewise addition of the disulfide arid ferrocyanide produced no effect which shows that the reaction products are not involved in any reversible step of the likely reaction scheme. The latter observation is unlike that in 2-mercaptoethylamine hydrochloride where an antocatalysis due to ferrocyanide (loc. ( i t . ) ~ 7 a sobserved. Runs were also made in the presence of potassium cyanide, The addition showed no effect up to a conwritration of 10.0 >( 10-3 &I beyond which an increase in the rate was observed as shown in Table V. The effect of dielectric cclnstant (0) variation was also investigated. ?'u+x~ clistilled ethanol was used for the purpose. h plot of log icz against 1/D gave a straight line (Figure a.,, . T h e Journal. o,f P42~sicaIChemistry, Vol. 76, No. 34, 1979
Second-order rate constant, kz, M - 1 min-*
Second-order [KCNl X
lo3
rate constant. -l4 k2, M - 1 min-1
10.0 20.0 30.0
2.03 2.52 3.14
[&Fe(CN)6] = 2.0 X M , [HS(Crr,),$\T(CW,),.WCl] = 2.0 X M , [HCl] = 5.0 X X , I = 0.3 M,temperature 35".
Runs were also made at different temperatures ranging between 30 and 50" and at intervals of 5'. Obedience to the Arrhenius equation was observed both in aqueous and water-ethanol media. Various thermodynamic parameters determined with the aid of usual expressions are recorded in Table VI.
Table VI D
AH*,
73.1 62.4
kcal mol-*
23.9 f 0 . 5 21.1 & 0 . 5
AS+, eu
1 1 . 2 1+- 1 . 1 8.5 f 1.1
A@,
kea1 mol-1
20.6 1+- 0.6 20.8 rt 0.6
Discussion As remarked earlier, the reaction under consideration i s different from the oxidation of a somewhat similar thiol hydrochloride, namely, 2-mercaptoethylamine hydrochloride, by the same oxidant, as this shows no autocatalytic effect due t o ferrocyanide. It is, therefore, natural to assume that the mechanisms in the two cases are not identical. However, it can be assumed that the present thiol hydrochloride will also exist mostly in the ionized form in aqueous medium with its cationic part as €IS-(CE&X(CH,)&l+, i.e., +RSH,
2-DIiVETHYLAMl NOETHANETHIOL HYDROCHLORIDE
3645
and not -S(CEz),N(GH&H+, i.e., +RS-,due to the strongiy acidic medium employed in the present study JFhich precluded the ionization of the weak -SH group. Since disulfide it) the sole final product of oxidation, undoubtedly the sulfhydryl group (-SH) provides the site of attack. For the sake of convenience, the above specieci is referred later as +RSH (where +R represents the remaining part of the cation). A simple reaction scheme of type A may explain the simple second-order kinetics but, admittedly, would fail to account for some observations such as the effect of initial concentration of ferricyanide, cyanide ion effect, etc.
radicals and subsequent dimerization of free radicals to give rise to product generally exhibit moderately large negative values of entropy of a ~ t i v a t i 0 n . l4 l ~ A reaction involving the electron transfer for the formation of free radicals via intermediate complex formation will include terms both for electron transfer and for the lengthening of the bond between the metal and the ligand for the calculation of activation parameters.lj The system involving stronger metal-sulfur bond may give an overall positive value of AS*^ This suggests that both the reaction mechanisms A and B are equally responsible for the observatione. However entropy criterion has t o be exercised with caution and preferably in conjugation with other factors. In the reaction under consideration an inverse linear relationdiip between the rate and [€I+] has been observed, whereas by and large, this rela tionship has been found to be complexle in such oxidations. In Scheme A such a relationship can be explained if the dissociation step of the sulfhydryl group in the thiol hydrochloride molecule is also included. This, however, appears unlikely in the acidic medium under consideration. I n fact, it is difficult to state with any amount of certainty the extent of involvement of proton vis-a-vis ferricyanide and ferrocyanide ion, and a simple relationship is probably more of an accidental nature. Ferricyanic acid is a strong acid, but ferrocyanic acid is strong only for the first two protons. According to Yekrasov and Z O ~ O V ,below ~’ pH 4 only a negligible fraction of iron(I1) is present as Fe(CX)e4- and mainly HFe(CN)63-or H2Fe(CN)6*-will exist. The plot of In k against the reciprocal of the dielectric constant is a straight line with a negative slope. This is in agreement with Brclnsted-Christiansen-Scatchard equation and indicates the presence of similarly charged reactants. 18 This indicates that the rate of decomposition of the complex (eq 4)is due to the reaction between anionic complex and cyanide anion. Variation of the thermodynamic parameters, like energy of activation with dielectric constant, is in agreement with the equation relating dielectric constant and activation energy. l a The participation of free radicals in the reaction is demonstrated by the initiation of polymerization of vinyl cyanide by such a system, Kolthoff and Nee-
(.4) 1 ~ e ( ~ ; ” i ) ~+3 -+RSR R_ IIFe(CiV)63-
+ ‘RS.
-‘-RS. + .SR+ fase +RSSR+
(1)
(2)
which gives -d[Fe(CN)B3--l/dt = ~ t [ F e ( c N ) ~[RSH] ~-] (3) where k is the experimentally determined second-order rate constant. In order to explain the decrease in rate observed on increasing the initial Concentration of ferricyanide it may be assumed that this behavior may be due to formation of a relatively less reactive 1:1 complexes {Fe(CN)SSR+j3- (reduction in the effective concentration of ferricyanide due t o formation of this complex). The complex of the type shown will have transient existence. It would combine with the cyanide ion to give rise to the products. The rate of this reaction would be very slow as it is governed by two factors: first very low concentration of this complex and second slow electron transfer from sulfur to iron. It has been shown, that an electron-withdrawing group adjacent to the sulfur atom (R+ in the present case) decreases the rate of oxidation.ll
(B) {Fe(CN)b(SR+)\3-
+ CY- --+
nlow
Fe(CW)64-
+I@.
+ +RS. (4)
+ .SR+% +RSSR+
This reaction mechanism is similar to one described by I
