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The kinetics of peroxydisulphate ion oxidation of glycolic acid in aqueous medium
L/~org. sue/. Chem., 1976, Vol, 38, pp, 2249-2250. Pergamon Press. Printed in Great Britain
THE KINETICS OF PEROXYDISULPHATE ION OXIDATION OF GLYCOLIC ACID IN AQUEOUS MEDIUM S. N. SHUKLA Chemistry Department, Lucknow University, Lucknow, (U.P.), India
(Firstreceived14 April 1975; in revised/otto20 September 1975) Abstract--The peroxydisulphate ion oxidation of giycolic acid has been studied kinetically under various conditions using different concentrations of reactants and of a neutral salt and at different temperatures. The reaction has been found to be first order and the thermodynamic constants have been calculated. The product, formic acid, has been identified and a probable mechanism for the reaction is proposed. INTRODUCTION The kinetics of oxidation of glycolic acid has been studied by several workers using different oxidising agents, e.g. chromic acid[l], Ce(SO4)212,3], Ce(CIO4)2, and Ce(NO3)4 [4]. The present communication deals with the kinetics of oxidation of glycolic acid by peroxydisulphate ion.
findings have also been reported by several other workers using peroxydisulphate as the oxidant [7]. The values of the energy of activation, entropy of activation and the frequency factor have been found to be 27.08 kcal, 2.5 e.u. and 5.87 x 10-u sec -t respectively. The following mechanistic steps are suggestkl. Table 1.
EXPERIMENTAL The glycolic acid used was from Riedel and the K2S2Os was of AR grade. The other chemicals employed were also of AR grade. The rate of reaction was studied in a constant temperature bath (-+0.01°C) by containing the reactants in thoroughly cleaned bottles. The progress of the reaction was followed by estimating the resultant change in hydrogen ion concentration as a function of time, which yielded the necessary kinetic data. The rate constant was calculated using the integration method. An average of six determinations gave the mean rate constant and the error is between 2 and 3%. Other routine details are given in an earlier paper of the anthor[5].
Ka
Kb
Temperature
Ke
Kd
I. 1I+
45°c
0.62
0.999
1.08
0.~6
~O°c
0.65
0.919
1.12
t~K2SzOI = 0.075 M, glycolic acid = 0.05 M; eO)K2S20$ = 0.075 M, glycolic acid = 0.07 M and 0.06 M at 45"C. (°K2S2Os= 0.07 M and 0.06 M, giycolic acid = 0.05 M at 45"C.(e~Inpresence of MgSO,, 0.0625 M and 0.050 M with the other conditions as in (a).
The following set of reactions is suggested for the peroxydisulphate ion oxidation of the substrate.
RESULTS
Identification of products The formation of formic acid as the end product was confirmed by the chromotropic acid test[6] (giving a brilliant violet colour after reduction with Zn and H2SOD. DISCUSSION The oxidation of glycolic acid by peroxydisulphate ion follows first order kinetics. The variation of the rate constant under different conditions is recorded in Table I. The rate increases with an increase in the peroxydisulphate concentration and decreases as the glycolic acid concentration is increased. The addition of MgSO, produces only a small and irregular change in the rate constant, suggesting the absence of salt effect as is to be anticipated for a free radical chain mechanism. Similar OH
$202 - K,, 2SOd
(1)
SO4' + H20 K~, "OH + HSO,-
(2)
HCH(OH)COOH + "OH r~) HCH(OH)CO0' + H20 +P, HCH(OH)CO0" + S2Os2- r, ~ P2 + SO42- dp S047
(4) SO: + HCH(OH)COO"
:0: H
H
I I
H - - C " + CO2 H
H
H + H:O(II)
(I) :O: H~C
II
/
:O: + H'
H
Formaldehyde
II
toj ~ H - - C - - O H ['OH]
(3)
Formic acid
(v)
(Iv) 2249
+ H20
H (III)
~' , P2 + SO4 2-
(5)
S. N. SHUKLA
2250
P~ and P2 are the products derived from the oxidisable substrate. Hence the peroxydisulphate ion is decomposed by steps 1 and 4 leading to the rate expression d(S:Os 2-) dt
K~[S20 2-] + K,[HCH(OH)COO'][S:Os2-] = (Kt + K4[HCH(OH)COO'])[S2Os2-].
By the steady state hypothesis the free radical concentration is constant, thus Ko = (K~ + K4[HCH(OH)COO']). If Ko is the observed velocity constant, then the observed rate law is
d(S20s) dt = K°[S2Os2-]" This rate law conforms to first order kinetics. Acknowledgement--Author is thankful to Dr. B. P. Yadav for suggestions in carrying out the present work. REFERENCES 1. G. V. Bokare and A. A. Deshpande, Z. Phys. Chem. (Leipzig) 227(1/2), 14 (1964). 2. G. V. Bokare, R. Dayal and P. Nath, Z. Physik. Chem. (Leipzig) 227(1/2), 19 (1964). 3. K. S. Gupta and S. Aditya, J. Electrochem. Soc. Japan 32(4), 200 (1964). 4. A. Macauley, J. Chem. Soc. 4054 (1965). 5. S. N. Shuklaand B. P. Yadava, Curt. Sci. 41(22),811 (1972). 6. F. Feigel,Spot Tests, pp. 340-341.Elsevier,New York (1956). 7. D. A. House, Chem. Rer. 62, 185 (1962).
THE KINETICS OF PEROXYDISULPHATE ION OXIDATION OF GLYCOLIC ACID IN AQUEOUS MEDIUM S. N. SHUKLA Chemistry Department, Lucknow University, Lucknow, (U.P.), India
(Firstreceived14 April 1975; in revised/otto20 September 1975) Abstract--The peroxydisulphate ion oxidation of giycolic acid has been studied kinetically under various conditions using different concentrations of reactants and of a neutral salt and at different temperatures. The reaction has been found to be first order and the thermodynamic constants have been calculated. The product, formic acid, has been identified and a probable mechanism for the reaction is proposed. INTRODUCTION The kinetics of oxidation of glycolic acid has been studied by several workers using different oxidising agents, e.g. chromic acid[l], Ce(SO4)212,3], Ce(CIO4)2, and Ce(NO3)4 [4]. The present communication deals with the kinetics of oxidation of glycolic acid by peroxydisulphate ion.
findings have also been reported by several other workers using peroxydisulphate as the oxidant [7]. The values of the energy of activation, entropy of activation and the frequency factor have been found to be 27.08 kcal, 2.5 e.u. and 5.87 x 10-u sec -t respectively. The following mechanistic steps are suggestkl. Table 1.
EXPERIMENTAL The glycolic acid used was from Riedel and the K2S2Os was of AR grade. The other chemicals employed were also of AR grade. The rate of reaction was studied in a constant temperature bath (-+0.01°C) by containing the reactants in thoroughly cleaned bottles. The progress of the reaction was followed by estimating the resultant change in hydrogen ion concentration as a function of time, which yielded the necessary kinetic data. The rate constant was calculated using the integration method. An average of six determinations gave the mean rate constant and the error is between 2 and 3%. Other routine details are given in an earlier paper of the anthor[5].
Ka
Kb
Temperature
Ke
Kd
I. 1I+
45°c
0.62
0.999
1.08
0.~6
~O°c
0.65
0.919
1.12
t~K2SzOI = 0.075 M, glycolic acid = 0.05 M; eO)K2S20$ = 0.075 M, glycolic acid = 0.07 M and 0.06 M at 45"C. (°K2S2Os= 0.07 M and 0.06 M, giycolic acid = 0.05 M at 45"C.(e~Inpresence of MgSO,, 0.0625 M and 0.050 M with the other conditions as in (a).
The following set of reactions is suggested for the peroxydisulphate ion oxidation of the substrate.
RESULTS
Identification of products The formation of formic acid as the end product was confirmed by the chromotropic acid test[6] (giving a brilliant violet colour after reduction with Zn and H2SOD. DISCUSSION The oxidation of glycolic acid by peroxydisulphate ion follows first order kinetics. The variation of the rate constant under different conditions is recorded in Table I. The rate increases with an increase in the peroxydisulphate concentration and decreases as the glycolic acid concentration is increased. The addition of MgSO, produces only a small and irregular change in the rate constant, suggesting the absence of salt effect as is to be anticipated for a free radical chain mechanism. Similar OH
$202 - K,, 2SOd
(1)
SO4' + H20 K~, "OH + HSO,-
(2)
HCH(OH)COOH + "OH r~) HCH(OH)CO0' + H20 +P, HCH(OH)CO0" + S2Os2- r, ~ P2 + SO42- dp S047
(4) SO: + HCH(OH)COO"
:0: H
H
I I
H - - C " + CO2 H
H
H + H:O(II)
(I) :O: H~C
II
/
:O: + H'
H
Formaldehyde
II
toj ~ H - - C - - O H ['OH]
(3)
Formic acid
(v)
(Iv) 2249
+ H20
H (III)
~' , P2 + SO4 2-
(5)
S. N. SHUKLA
2250
P~ and P2 are the products derived from the oxidisable substrate. Hence the peroxydisulphate ion is decomposed by steps 1 and 4 leading to the rate expression d(S:Os 2-) dt
K~[S20 2-] + K,[HCH(OH)COO'][S:Os2-] = (Kt + K4[HCH(OH)COO'])[S2Os2-].
By the steady state hypothesis the free radical concentration is constant, thus Ko = (K~ + K4[HCH(OH)COO']). If Ko is the observed velocity constant, then the observed rate law is
d(S20s) dt = K°[S2Os2-]" This rate law conforms to first order kinetics. Acknowledgement--Author is thankful to Dr. B. P. Yadav for suggestions in carrying out the present work. REFERENCES 1. G. V. Bokare and A. A. Deshpande, Z. Phys. Chem. (Leipzig) 227(1/2), 14 (1964). 2. G. V. Bokare, R. Dayal and P. Nath, Z. Physik. Chem. (Leipzig) 227(1/2), 19 (1964). 3. K. S. Gupta and S. Aditya, J. Electrochem. Soc. Japan 32(4), 200 (1964). 4. A. Macauley, J. Chem. Soc. 4054 (1965). 5. S. N. Shuklaand B. P. Yadava, Curt. Sci. 41(22),811 (1972). 6. F. Feigel,Spot Tests, pp. 340-341.Elsevier,New York (1956). 7. D. A. House, Chem. Rer. 62, 185 (1962).