OXIDATION OF THIOUREA AND THIOACETAMIDE nism of reaction from the UFe-UF4 interaction to noninteraction. Acknowledgment. Thanks are due to A h . S. Tsujimura, A h . A. Takahashi, and Mr. G. Fujisawa for their
1497
aid in the construction of the thermobalance, and to Dr. S. Suxuki and other members of the Analytical Section, Japan Atomic Fuel Corp., for their analysis of the water content in the sample of uranium tetrafluoride.
Oxidation of Thiourea and Thioacetamide by Alkaline Hexacyanoferrate(111) by M. C. Agrawal and S. P. Mushran Chemical Laboratories, University of Allahabad, Allahabad, I n d i a
(Received Jtdy 28, 1967)
Kinetics of the oxidation of thiourea and thioacetamide with alkaline hexacyanoferrate(II1) has been studied in aqueous solution. The reaction with thiourea has been found to be first order with respect to thiourea, hexacyanoferrate(II1), and hydroxyl ion, while that with thioacetamide shows zero-order dependence on Fe(CN)6*- and first-order dependence on both thioacetamide and OH- ion. Potassium chloride concentrations showed accelerating effect on the reaction rate, while the effect of ferrocyanide ion concentration was negligible. Increase in dielectric constant (D)enhanced the reaction rate and the plot of log k against l / D showed a linear relation. A suitable mechanism for the oxidation processes has been proposed.
Introduction Hexacyanoferrate(II1) is well known for its oxidative power in alkaline media and can be described as one of the efficient electron-abstracting reagents. Though uncatalyxed oxidations by hexacyanoferrate(II1) ion are, in general, fairly fast, osmium(VII1) ion has been observed to exert a positive catalytic influence, and several inorganic and organic substances may directly be titrated against hexacyanoferrate(II1) in its presence. The kinetics of the oxidation of 3-mercaptopropionic acid has been studied by Bohning and Weiss,2 while Kolthoff and coworkers* have investigated the mechanism of oxidation of 2-mercaptoethanol by hexacyanoferrate(II1) in acid medium. In another publication, Kolthoff and coworkers4 have studied the oxidation of n-octyl mercaptan by alkaline hexacyanoferrate(II1) in acetone-water medium. Recently, Gorin and Godwin6 have suggested that the oxidation of some mercaptans is catalyzed by metal ions. I n spite of the above work, the literature on the kinetics of oxidation reactions of organic sulfur compounds by hexacyanoferrate(II1) appears to be scanty. In the present communication, the kinetics of the oxidation of thiourea and thioacetamide by alkaline hexacyanoferrate(II1) has been investigated in some detail, in order to shed some further light on oxidation mechanisms involving hexacyanoferrate(II1) ion. The
reactions have been studied in presence of sodium carbonate-bicarbonate buffer, to avoid the effects of any possible pH variations during the progress of the reactions. Influence of several variable factors such as dielectric constant and the concentrations of potassium chloride, potassium hexacyanoferrate(I1) , and hydroxyl ion has been investigated to ascertain the exact nature of the oxidation processes.
Experimental Section Materials. (i) Aqueous
potassium hexacyanoferrate(II1) was prepared from an AnalaR BDH sample and the concentration was checked by iodometryS6 (ii) Aqueous thiourea was prepared from a p r o analysi E. Merck reagent. (iii) Aqueous thioacetamide was also prepared from a pro analysi E. nferck sample. (iv) All other reagents were of analytical grade. (v) Doubly distilled water was employed for preparing solutions and for diluting where necessary, and in every case glass vessels of Jena Geratglas were (1) (2) (3) 66,
F. Solymosi, M a g y . Kem. Folyoirat, 63,294 (1957). J. J. Bohning and K. Weiss, J . Am. Chem. SOC.,82,4724 (1960). E. J. Meehan, I. M. Kolthoff, and H. Kakiuchi, J . Phys. Chem., 1238 (1962).
(4) I. M. Kolthoff, E. J. Meehan, M. S. Tsao, and Q. W. Choi, ibid., 66, 1233 (1962).
(6) 0. Gorin and W. E. Godwin, J . Catalysis, 5 , 2 7 9 (1966). (6) A. Berka, J. Vulterin, and J. Z$ka, “Newer Redox Titrants,” Pergamon Press Inc., New York, N. Y., 1965. Volume 71,Number 6 M a y 1068
1498 used. The reaction vessel was coated from outside with black Japan to exclude photochemical effects. Procedure. The reactions were carried out at constant pH 11.0 using standard carbonate-bicarbonate buffer. The kinetics of the oxidation of thiourea and thioacetamide were followed by estimating hexacyanoferrate(II.1) lyith the progress of time. Samples (10 ml) of the reaction mixture were withdrawn at definite intervals and the amount of unconsumed hexacyanoferrate(II1) was estimated colorimetrically using a Klett-Summerson photoelectric colorimeter with blue filter no. 42 (transmission 400-450 mp). It was ascertained that hexacyanoferrate(I1) did not absorb in this region and also that Beer's law was valid with aqueous hexacyanoferrate(II1). The absorption cell mas chilled before adding an aliquot portion of the reaction mixture in order to quench the react)ion at lower temperatures. I n general, the procedure gave satisfactory results but in a few cases the reaction mixture turned turbid. This probably was due to the precipitation of sulfur at higher concentrations of the reducing substrate. Hence, concentrations of thiourea and thioacetamide were chosen where the reaction mixtures did not develop any turbidity.
Results Stoichiometry. Reaction mixtures with slight excess of potassium hexacyanoferrate(II1) were maintained at 50°, pH 11.0, for 24 hr, when the reaction was complete. The hexacyanoferrate(I1) formed equivalent to thiourea and thioacetamide was determined titrimetrically using ceric sulfate. The amounts of reducing substrates consumed and of hexacyanoferrate(11) formed showed that the over-all reactions are
+ 100H- + 8 F e ( c N ) ~ += ?;HzCOM12 + Sodz-+ 8Fe(CN)64- -k 5Hz0 CH3CSNH2+ llOH- + 8 F e ( C N ) P = CH3COO- + + KH3 + 8Fe(CK)64- + 5Hz0
KHzCSXHz
Kinetics. E$ect of Hexacyanoferrate(III) Concentration. To evaluate the order of the reactions with respect to the oxidant, the reactions were carried out at different concentrations of hexacyanoferrate(II1). The kinetic analysis was mainly applied to the initial stages of the reaction, approximately up to the half-life of the reaction. For a particular concentration of the oxidant, the reaction with thiourea followed first-order disappearance to hexacyanoferrate(II1) at all concentrations of thiourea (Figure 1). The reaction with thioacetamide showed zero-order dependence to hexacyanoferrate(II1) which in later stages tended to become first order in hexacyanoferrate(II1) (Figure 2). Since in this case the concentrations of the sulfur compound are low, the apparent decrease in zero-order The Journal of Physical Chemistry
nI. C. AGRAWAL AND S.P. M U S H R A N
50
1
40
30
N 0
c
X
7 %
20
F CIJ
10
0
10
30
20
Time in minutes
40
Figure 1. Reaction of hexacyanoferrate(111), [Fe(CN)63-] = 4 X M , pH 11.0, temperature 40'. Concentration of NH2CSNHz: (1) 3.3 x 10-8 M , ( 2 ) 4.0 x 10-3 M , (3) 4.5 x 10-3 M , (4) 5.0 x M , ( 5 ) 5.*5 x M , (6) 6.0 x M.
rate constant may in part be ascribed to the change in the concentration of thioacetamide during the reaction. Further, change in concentration of hexacyanoferrate(II1) showed no effect on the first-order and zero-order rate constants with respect to the oxidant for reactions with thiourea and thioacetamide, respectively (Table I). This confirms that the oxidation of thiourea has first-order dependence to hexacyanoferrate(lI1) , while that of thioacetamide is independent of the hexacyanoferrate(II1) concentration. Table I : Hexacyanoferrate(II1) Dependence Thioacetamide Thiourea
10~lc,xptllc
104[Fe(CN)8a-l,a ivf
10akexpti,b
mole I. - 1 min-1
3.0 3.5 4.0 4.5
14 . Y 16.0 16.0 16.3
min - 1
a p H 11.0 and temperature 3R'. [Thioacetamide] = 2 X M.
3.64 3.87 3.73 3.82
* [Thiourea] = 4 X
l o w sM .
OXIDATION OF THIOUREA AND THIOACETAMIDE
1499 obtained as 4.30 1. mole-' min-' and 18.9 X min-l, respectively, at 35' and pH 11.0. Efect of Alkali Concentration. The oxidation of thiourea and thioacetamide is very susceptible to changes in the alkali concentrations. The oxidations are directly proportional to the concentration of alkali, and the average values JCexptl/ [OH-] for the oxidations of thiourea and thioacetamide were calculated as 16.9 1. mole-' min-l and 6.94 X min-', respectively (Table 111).
Table I11: Alkali Dependence bxptl/
104[NaOH],a
M
4.0 6.0 8.0 10.0 I
10
I
I
I
40
20 30 Time in minufes
E$ect of Changing the Reductant Concentration. I n order to determine the order of the reaction with respect to the reducing substrate, the reactions were studied a t different concentrations of thiourea and thioacetamide. The first-order and zero-order constants with respect to hexacyanoferrate(II1) are observed to be directly proportional l,o the concentrations of thiourea and thioacetamide, respectively. Thus, both the reactions are first order in the reducing substrate (Table 11). The over-all bimolecular and unimolecular rate constants for the oxidi+tion of thiourea and thioacetamide, respectively, show fairly concordant values at all concentrations of the reactants. Average values of kexptl/ [thiourea] and kex,tt/ [thioacetamide] have been
Table I1 : Thiourea and Thioacetamide Dependence
10"thio-
M
4.0
.
4 5
5.0 5.5 6.0 b
16.0 18.9 22.8 24.5 27.0
4.0 4.2 4.5 4.4 4.5
1.50 2.00 2.50 3.00
...
6.90 9.85 13.4 17.2
[OH-], 1. mole-' min-1
10skexptltc mole L - 1 min -1
10*kexptl/ [OH-], min-1
17.2 16.4 16.7 17.2
2.76 4.20 5.50 6.97
6.90 7.00 6.87 6.97
[Fe(CN)83-] = 4 x lO-4M, temperature 35'. b [Thiourea] 4 X 10-3 M . 0 [Thioacetamide] = 4 X M.
5
=
Figure 2. Reaction of hexacyanoferrate(III), [Fe(CN)63-] = 4 X 10-4 M, p H 11.0, temperature 40'. Concentration of CHaCSNH2: (1) 1.5 X lo-* M, (2) 2.0 x 10-4 M , (3) 2.5 x 10-4 M , (4) 3.0 x 10-4.
kexptl/ [thiourea], 104[thioaoet108kexptl,b 1. mole-1 min-1 min-1 M
108kexptilb min-1
On increasing the concentration of alkali twofold, the rate constants are almost doubled showing firstorder dependence to alkali. Injluence of Dielectric Constant of the Medium. The oxidations of thiourea and thioacetamide by alkaline hexacyanoferrate(II1) were studied in different methanol-water mixtures. The results are represented in Table IV.
Table IV : Influence of Dielectric Constant kz,C
Db
1. mole-1 min-1
10akr,d min-1
74.83 72.76 70.68 68.62 66.52 64.38
4.56 3.50 2.67 1.91 1.13 0.61
23.7 18.4 14.3 9.49 7.08
% methanola
0
5 10 15 20 25
...
[Fe(CN)ea-] = 4 X M, pH 11.0, temperature 35'. Dielectric constant. [Thiourea] = 4 X M. [Thioacetamide] = 4 X 10-4 M . a
1Oakexptl/ [thio10%expti,c acetmole 1.-' amide], min-1 min-1
2.80 3.73 4.92 5.61
18.7 18.7 19.7 18.7
...
...
a [Fe(CN)Ba-] = 4 X 10-4 M, temperature 3 5', p H 11.0. First-order rate constants. c Zero-order rate constants.
It is evident that the decrease in the dielectric constant of the medium results in a decrease in rate constants. The plots of the reciprocal of dielectric constant against log k gave straight lines having negative slopes (Figures 3 and 4). The data, however, could not be correlated with the theoretical expressions for ion-ion and ion-molecule reactions. Volume 72, Number 5
May 1968
1500
M. C. AQRAWAL
-al 15.0.
I
13.8
['/D
I
I
14.6
13.4
x
The influence of several other factors on the reaction rates has also been investigated. Addition of potassium chloride showed a marked accelerating effect on the oxidation of thiourea. Effect of temperature was also studied and the energies and the entropies of activation are 11.65 kcal and -27.28 eu (thiourea) and 16.70 kcal and -21.36 eu (thioacetamide). Platinum(IV) appears to be inactive as a catalyst but the osmium(VII1) ion has strong catalytic influence on both of these reactions.
Discussion The alkali dependence of the oxidation processes involved clearly indicates that the oxidation of thiourea and thioacetamide takes place through an intermediate enolic anion of the reducing substrate formed with hydroxyl ions. I n accordance with this view, it has been shown by Waters, et al.,I that substances like benzaldehyde, which cannot enolize or do not form an intermediate anion, are not oxidized by alkaline hexacyanoferrate(III), whereas substances like acetone and propanal, which exist in the enolic form, are easily oxidizable. Thus, taking into consideration the kinetic data obtained, the following scheme for the oxidation of thiourea and thioacetamide is proposed on the basis of the mechanism proposed by Speakman and Waters' for the oxidation of carbonyl compounds by alkaline hexacyanoferrate(II1). The Journal of Physical Chemistry
S. P. MUSHRAN
[I/. x lo3]
x io3]
Figure 3. Effect of dielectric constant, [Fe(CN)e3-] = 4 10-4 M , [NHzCSNHZ] = 4 X 10-8 M , temperature 35".
AND
Figure 4. Effect of dielectric constant, [ F e ( C N ) p ] = 4 X lo-' M , [CH&SNH2] = 4 X 10-4 M , temperature 3S0.
The first step (1) involves the formation of the enolic intermediate anion of the substrate which is essentially reversible in nature and may be fast or slow depending upon the nature of the reductant RCSNHz
+ OH-
ki
+ HzO
(1)
+ F e ( C N ) P kn_ complex (x)
(2)
RCSNHk- 1
This is followed by
RCSNHcomplex (x)
+ F e ( C N ) P -+ ka
2Fe(CN)e4-
+ products
(3)
I n step 2 the intermediate anion forms a complex with the oxidant which is subsequently followed by a fast step 3 which gives the products of the reaction. Step 2 may be fast or slow depending upon the stability of the intermediate enolic anion obtained in step 1. Step 3 is a relatively fast step and gives the products of the reaction. I n view of the stoichiometry step 3 cannot form the products of the reaction, and it seems likely that step 3 generates intermediate products which react further in a sequence of reactions which are very rapid relative to steps 1-3. Now under the steady-state conditions, the rates of formation and destruction of the intermediate anion must be equal. Hence (7) P. T. Speakman and W. A. Waters, J. Chem. SOC.,40 (1965).
OXIDATION OF THIOUREA AND THIOACETAMIDE d dt
- [RCSNH-] =
kj [RCSNHz] [OH-]
1501 d dt
-
-- [Fe(CN)8-]
k-l[RCSNH-] - kz[RCSNH-] [ F e ( c N ) ~ ~ -=l 0
or [RCSNH-] =
ki [RCSNHz] [OH-] k-I kz[Fe(CN)&-l
+
(4)
Also the loss of hexacyanoferrate(II1) may be represented as d -- [Fe(CN)B3-]= kz[RCSNH-] [Pe(CN),P] dt k3[complex (x)][Fe(CN)ea-] ( 5 )
+
or d dt
-- [Fe(CN)63-]= 21c,kZ[RCSNHJ [OH-] [Fe(CN)e3-] k-1 4- k z [ F e ( c N ) ~ ~ - ]
(6)
According to the above-derived rate law, if the ratedetermining step is ( 2 ) as we propose for the oxidation will become much larger than kz of thiourea, then and the rate law will approximate to d dt
-- [F~(C)N)B~-] = 2*[NH2CSNH2]
[OH-] [Fe(CN)a3-] (7)
ffi-1
and if the rate-determining step is (l),as we propose for the oxidation of thioacetamide, then rate law 6 will reduce to
=
2k1 [ C H G W H Z [OH] J
(8)
The derived rate laws (7) and (8) reveal that the oxidation of thiourea by alkaline hexacyanoferrate(II1) would have first-order dependence on thiourea, hexacyanoferrate(II1) , and hydroxyl ions, while the oxidation of thioacetamide would be independent of the concentration of the oxidant. The kinetic data obtained agree well with derived rate law equations and the mechanism proposed. For the oxidation of thioacetamide by hexacyanoferrate(III), however, at lower concentration of the oxidant step 2 will become slow and rate determining. The rate law equation in that case will correspond to ( 7 ) as for thiourea and the reaction will have first-order dependence to hexacyanoferrate(II1). Figure 2 also represents similar results. According to above rate laws, the rate-determining step involves interaction between two negatively charged ions or between a negatively charged ion and a polar molecule and thus should correspond to a negative entropy change. This has been found to be true from the experimental data recorded earlier. For the reactions involving the interaction between similarly charged ions or between a negatively charged ion and dipole, the plots of log k against the reciprocal of dielectric constant would be straight lines having negative slopes. In the present investigations also similar results have been obtained (Figures 3 and 4),which further show that the experimental observations agree well with the mechanism proposed for the oxidation of these sulfur compounds by hexacyanoferrate(II1) in an alkaline medium. Acknowledgment. The authors thank the Council of Scientific and Industrial Research, New Delhi, India, for a Senior Research Fellowship to M. C. A.
Volume ?'8!Number 6
M a y 1968
