THE JOURNAL OF
PHYSICAL CHEMISTRY
Registered in U.8.Patent Ofice @ Copyright, 1969, by the American Chemical Society
VOLUME 73, NUMBER 6 JUNE 1969
Oxidation Kinetics of Thioglycolic Acid by Ferricyanide Ion in Acid Medium by R. C. Kapoor, 0.P. Kachhwaha, and B. P. Sinha Department of Chemistry, University of Jodhpur, Jodhpur, Rajasthan, I n d i a
(Received February 20,1968)
The oxidation kinetics of thioglycolic acid by ferricyanide ion in acid medium has been described. The two compounds react in equimolar ratio and in excess of thioglycolic acid; the disappearanceof ferricyanide follows the second-order rate law. The order in thioglycolic acid is unity, and the rate is inversely proportional to the initial concentration of ferricyanide ion. An increase in hydrogen ion concentration also decreases the rate, though in a somewhat complex manner. The effect of dithiodiglycolic acid, ferrocyanide ion, and the salt effect has been investigated. A tentative mechanism of the oxidation scheme has also been given.
Introduction The oxidation of a thiol to its disulfide is a process of fundamental importance in biochemistry, since the sulfhydryl groups are involved in biochemical processes such as enzyme catalysis, cell division, and radiation injury. While extensive studies relating the rate of oxygen uptake by solutions of cysteine, glutathione, and thioglycolic acid have been made,l kinetic studies on the oxidation of thiols to disulfide with different oxidants are relatively few. Of late, some significant paper^^-^ have appeared on this aspect of these compounds. Mention may be made of the searching investigations by Bohning and Weiss2 on the oxidation kinetics of 3mercaptopropionic acid by ferricyanide ion and those by Kolthoff, et al.,3$4on n-octyl mercaptan and 2mercaptoethanol using the same oxidant. I n present investigations, the oxidation kinetics of thioglycolic acid by ferricyanide ion in an acid medium have been described. Ferricyanide ion was employed as an oxidant owing to a variety of advantages. Firstly, it is a mild oxidant and does not, therefore, extend the oxidation beyond the formation of corresponding disulfide. Secondly, it abstracts only one electron from the reductant and thus gives a simple step in any oxidation scheme. Moreover, on reduction it forms a stable product, vix., ferrocyanide.
Experimental Section E. Merck analytical reagent grade potassium ferricyanide was used, and the solutions were prepared by exact weighing. Thioglycolic acid (assay 99%) and dithiodiglycolic acid were provided by Evans Chemetics Inc. The solution of thioglycolic acid was standardized with a standard solution of analytical reagent grade iodine after acidifying with hydrochloric acidus Fresh solutions were prepared for every run. All other reagents such as potassium chloride, potassium ferrocyanide, potassium nitrate, hydrochloric acid, etc., were of Analar grade. All solutions were prepared in doubly distilled and boiled water. A Hilger Uvispek spectrophotometer Ytodel H 700-8 fitted with thermostatic arrangement was employed to obtain the absorbance spectra. The polarograms were taken with a Toshniwal manual polarograph. The kinetics of the reaction were followed colorimetrically using a Klett-Summerson (1) N. Kharasch, “Organic Sulfur Compounds,” Vol. 1, Pergamon Press, London, 1961, p 97. (2) J. J. Bohning and K. Weiss, J . A m e r . Chem. Soc., 82, 4724 (1960). (3) I. M. Kolthoff, P. J. Meehan, M. S. Tsao, and Q. W. Choi, J . P h y s . Chem., 66, 1233 (1962). (4) E.J. Meehan, I. M. Kolthoff, and H. Kakiuchi, ibid., 66, 1238 (1962). (5) H.Kramer, J . Ass. O$lc. Agr. Chem., 35, 285 (1952).
1627
R. C. KAPOOR, 0. P. KACHHWAHA, AND B. P. SINHA
1628 photoelectric colorimeter. A blue filter No. 42, whose approximate spectral range lies between 400 and 465 mp, was used throughout as the light filter. I n this wavelength region ferricyanide ion is found to absorb strongly, whereas the absorbance due to ferrocyanide ion is negligibleS6 Thioglycolic acid, and its oxidation product in the present case, viz., dithiodiglycolic acid, also do not show any detectable absorbance in this wavelength region, as is shown by their absorbance spectra (Figure 1). It was found that Beer’s law is obeyed throughout the concentration range of ferricyanide employed in these investigations. All the constituents of the reaction mixture except potassium ferricyanide were kept in a Pyrex flask coated black from outside to avoid any photochemical action. This was kept in a thermostat regulated at 30 f 0.1’. The flask containing ferricyanide solution was also kept immersed in the same thermostat. The contents of both flasks were flushed with nitrogen. When both solutions attained the desired temperature, the requisite quantity of ferricyanide solution was pipetted into the reaction flask. The stop watch was started when half of the ferricyanide solution had passed out of the pipet. An aliquot was rapidly transferred in a Klett’s tube kept at the same temperature. The scale reading was quickly taken, and the concentration of ferricyanide left was computed from a concentration-optical density graph.
Wave length in
m
r
Figure 1. Absorption spectra of dithiodiglycolic acid and thioglycolic acid: a, 1.00 X loF3M dithiodiglycolic acid, pH 1.65; b, 1.00 X 10-3 M thioglycolic acid, pH 1.65.
Stoichiometry The stoichiometry of the reaction was established p~larographically,~ whereby the characterization of the final product of oxidation was also achieved. It was found that the cathodic wave given by a known sample of the disulfide dimer (Figure 2) coincided with the wave given by the final product of oxidation. The stoichiometry was also confirmed from colorimetric measurements in the following manner. Reaction mixtures with excess of ferricyanide were allowed to go to completion and analyzed for ferricyanide thereafter. These results also gave a 1: 1 ratio according to
2RSH
+ 2F’e(CN)63-
---f
RSSR
+ 2 F e ( c N ) ~ ~+- 2H+
Results A number of runs with 10- to 20-fold excess of thioglycolic acid were studied. A plot of l/(a - x) against time yielded a straight line for each run (Figure 3). On the basis of these observations, it can be concluded that the order in ferricyanide is two. The values of the pseudo-second-order rate constant were obtained from the slopes. When the logarithms of the values so obtained were plotted against the logarithms of the corresponding concentrations of thioglycolic acid, a straight line having a slope of 1.02 was obtained (Figure 4). Since in all the variations of thioglycolic acid the The Journal of Physical Chemistry
-
1.0
I
Figure 2. Polarograms of (a) known thioglycolic acid (2.50 X 10-4 M), (b) thioglycolic acid left in the reaction mixture, (c) known dithiodiglycolic acid (1.25 X 10-4 M) to be compared with, and (d) dithiodiglycolic acid formed in the mixture. (6) A. W. Adamson, J. Phys. Chem., 5 6 , 858 (1962). (7) 0. P. Kachhwaha, B. P. Sinha, and R. C. Kapoor, Proceedings of the Symposium on Electrode Processes, Jodhpur, India, 1966, p 84.
OXIDATION KINETICS OF THIOGLYCOLIC ACID
1629
The rate was also found to depend on the initial concentration of ferricyanide ion taken. To determine this relationship a number of runs with different initial concentrations of ferricyanide ion, while keeping other parameters constant, were studied. The values of second-order rate constants and also the first-order rate constants (calculated from the initial rate measurements) are recorded in Table I.
Table I [HSCHZCOOH],4.00 X 10-3 M Temp, 30" [HCl], 0.02 M I, 0.054 M
Figure 3. Second-order diagram for six different experiments M; [HCl], 2.00 (temperature, 30'; [K,Fe(CN)e], 2.50 X X 10+ M; I, 0.054 M): (1) [HSCH2COOH],2.50 X M; (2) 3.00 x 10-3 M ; (3) 3.50 x 10-3 M; (4) 4.00 x 10-3 M ; (5) 4.50 X M ; (6) 5.00 X M.
[Fe(CN)sS-]o x 104, mol 1. -1
Secondorder rate constant k2, 1. mol-' min-1
Initial firstorder rate Constant kl, min-1
ka[Fe(CN)a*-l x 102
4.00 3.50 3.00 2.50 2.00
603 695 805 1025 1180
0.201 0.216 0.223 0.225 0.233
24.1 24.3 24.2 25.6 23.0
It is found that the rate is inversely proportional to the initial concentration of ferricyanide as shown by almost constant values recorded in the last column of Table I. It can also be seen that the values of the initial first-order constant are also practically constant, the reason for which is given in the Discussion. The effect of hydrogen ion concentration on the rate was studied in the presence of different concentrations of hydrochloric acid, keeping ionic strength constant with the help of potassium chloride. The values of the second-order constant at different acid concentrations are recorded in Table 11. With regard to the quantitative relationship, an inverse proportionality between the rate and square of the hydrogen ion concentration
Table I1 [K3Fe(CN)6],2.50 X Temp, 30' [HSCH2COOH],2.50 I, 0.054 M
Figure 4. Dependence of the rate on concentration of thioglycolic acid (temperature, 30"; [HCl], 2.00 X M; I, 0.054 M).
concentration of ferricyanide ion remains constant, the slope gives the order in thioglycolic acid.
M
x
10-3 M
IHtl
ka, 1. mol-' min-1
kz[H f 1
kz[Ht12 x 102
0.012 0.014 0.016 0.020 0.024 0.028 0,032 0.036
1783 1580 1420 650.9 500 446.8 408 255
21.40 22.12 22.72 13.02 12.00 12.51 13.06 9.18
25 68 30.97 36.35 25.90 28.80 35.03 41.78 33.05 I
Volume 73. Number 6 June 2969
1630
R. C. KAPOOR, 0. P. KACHHWAHA,
appears most likely (Table 11). It may be remarked that in oxidation of thiols the dependence of the rate of hydrogen ion concentration is generally observed to be complex.' This point' will be taken up again in the Discussion. Ferrocyanide ion was found to retard the reaction (Table 111), while dithiodiglycolic acid which is the final product of the oxidation of thioglycolic acid had no influence on the rate.
h,
[Fe(CNW-]o
x
104,
mol 1. - 1
k2( [Fe(CN)fia-Io f
1. mol-' min-1
[KaFe(CN)s], 2.5 X 10-4 M [HCI], 0.02 M [HSCH2COOHl,2.5 x M Temp, 30" h, I,
1. mol-'
mol
min-1
K C1
0.059 0.064 0.069 0.074 0.079
704 770 836 927 1010
KNOa
0.059 0,064 0.069 0.074 0.079
690 770 834 915 958
NaCl
0 059 0.064 0.069 0.074
760 880 1018 1184
MgSOi
0.058 0.062 0.066 0.070 0.074
742 850 970 1104 1235
Salt
2.00 2.50 4.00 5.00 8.00 10.00 12.00
370 312 210 157 141 113 110
10'
16.7 15.6 13.7 11.8 14.8 14.1 16.0
Table IV shows the values of the rate constant in the presence of various salts at different concentrations. A change in ionic strength changes the reaction rate. The salt effect is positive in all the cases studied. It can also be seen that for the same ionic strengths, sodium ions have a more pronounced effect on the rate than potassium ions.
taken as the reactive one, which for the sake of convenience will be denoted by RS2-. The various steps proposed are
+ Fe(CN)s3- Fe(Cr\T)e4- + RS.- ( 2 ) RS.- + Fe(CN)e3- "2,Fe(CN)B4- + R8.t (3) IC 1
RS2-
k-
RSk
Discussion Based on the experimental observations, the following mechanism is proposed. The first and second dissociation constants reportedB for thioglycolic acid are 2.10 X and 2.10 X lo-" at 25", respectively. The magnitude, particularly that of the second one, is fairly small. In acid medium, therefore, one would normally expect the acid to remain practically in undissociated form. However, at any instant there will be some ionized species present according to the equilibrium Kl
HSCHzCOOH
I
[Fe(CN)s'-]o)
x
HSCHzCOO-
+ H+
-SCHzCOO-
+ H+
1
+ RS2- 2RSSR2-
(4)
The above mechanism is based on an initial reversible electron transfer, similar to that proposed by Watersg and Fieser.lo It can be seen that in the above mechanism we have proposed RSe- and RS* as the reactive species. These two species are formed due to the abstraction of two electrons from a sulfur atom. The possibility of the existence of the free radical RSa- is evidenced by the ability of the system to initiate the polymerization of olefins. This is also supported by the observations of Kolthoff and Meehan.I1 The existence of an intermediate transient species of the type RSh has also been suggested by Basford and Huennekens.I2
(1)
Because the final product of the oxidation of thioglycolic acid by ferricyanide ion is its disulfide (via, dithiodiglycolic acid) it is obvious that the -SH group provides the reaction site. Ferricyanide, being a mild oxidant, does not attack the -COOB group. I n the mechanism, therefore, the species-SCH2COO- has been The Journal of Physical Chemistry
B. P. SINHA
Table IV
Table I11 [KsFe(CN)a], 2.50 x 10-4 M Temp, 30" [HSCH&OOH], 2.50 X 10-3 M [HCl], 0.02 M I , 0.054 M
AND
(8) E. Larson, 2.Anorg. Allg. Chem., 172, 376 (1928).
(9) C. G.Haynes, A. H. Turner, and W. A. Waters, J . Chem. SOC., 2823 (1956). (10) L. F. Fieser, J. Amer. Chem. Soc., 52, 5204 (1930). (11) I. M. Kolthoff and E. J. Meehan, J . Polym. Sci., 11, 71 (1953). (12) R. E.Ba:rford and F. M.Huennekens, J . Amer. Chem. Soc., 77, 3873 (1955).
1631
OXIDATION KINETICS OF THIOGLYCOLIC ACID At the stationary state, the concentration of the free radical RS. - will become practically constant; i e . d[RS*-] dt
centration of ferricyanide, Dhe above equation reduces to
=o
Using this value we have
where k” is now the experimentally determined pseudosecond-order constant. I n the initial stages the concentration of ferrocyanide will be very small compared to ferricyanide; hence, the rate eq 8 will assume the form
-d [Fe(CN)&‘-]
k’ -d [ F e ( C N ) P ] = [Fe(CN)&-][HRSH]
From steps 2, 3, and 4 we have
-
dt
dt
k-1
= kt[Fe(CN)&-](when [HRSH]
where kr is the experimentally determined first-order constant. Obviously kt should be independent of initial concentration of ferricyanide taken. This can be seen from the values of the initial first-order rate constants recorded in Table I. The values are practically constant although the concentration of ferricyanide is to 4.0 X M . The above varied from 2.0 X scheme explains all our observations. However, it is also quite likely that the hydrogen ion may combine with the ferrocyanide ion to form ferrocyanic acid according to the equilibrium
From step 1 [RS2-1 =
>> [Fe(CN)a3-])
KlKz [HRSH] [H+lz
Substituting this value of [RS2-] into eq 6
K’
H + f Fe(CN)64-
HFe(CN)63-
The equilibrium constant K’ is reportedla to be 5.40 X When we substitute the value of [Fe(CN)64-] from this equilibrium into eq 7, we have
-d[Fe(CN)8-] dt
2K’k1kzK1Kz[HRSH 1 [Fe(CN)63-]2 [H+]{~ - I [ H F ~ ( C N ) ~ ~K’kz[H+] -] [Fe(CN)63-] ]
+
(11) The existence of such a state of affairs is bound to make the relationship rather cloudy.
(9)
+
since the quantity [Fe(CN)64-] [Fe(CN)e3-] is equal to the initial concentration of ferricyanide taken, [Fe(CNla3-10. When [HRSH] >> [ F ~ ( C N ) B ~for - ] , any given con-
Acknowledgment. One of the authors (0. P. K.) wishes to thank the Council of Scientific and Industrial Research, New Delhi, for the award of a research fellowship. Thanks are also due to Evans Chemetics Inc. for the gift samples of thioglycolic acid and dithiodiglycolic acid. (13) I. M. Kolthoff, J. Phya. Chem., 39, 955 (1935).
Volume 73, Number 6 June 1969
