37, 345-353 (1988)
Kinetics and Mechanisms of the Oxidation of Oxalic and Malonic Acids by Chromium(VI) in Acidic Sulfate Media SAMY
S. ANIS, MOHAMMED Department
A. MANSOUR, AND SHAKER LABIB STEFAN’ Ain Shams University,
Received October 5, 1987; accepted January 4, 1988
The reduction kinetics of Cr(VI) by oxalic and malonic acids in sulfuric acid solution was studied by measurement of the reduction of Cr(V1) solution. The reaction was found to be of first order in Cr(VJ) and malonic acid, and second order in oxalic acid. Activation parameters (hE* and A,S*) were evaluated and found to be 7.23 kcal mol-‘, -44.98 e.u. and 13.34 kcal mol-‘, -34.78 e.u. for oxalic and malonic acids, respectively. Mechanisms which involve complex formation between the two reactants as a rate-determining step prior to their decomposition were suggested. The oxidation of oxalic acid has been studied with different oxidants like Mn(V1) (I, 2), Tl(II1) (3), Ce(IV) (4), and V(V) (5,6) in acid medium. Likewise, the oxidation of malonic acid was also studied with different oxidants as V(V) (7) and Ce(IV) (8). In the present paper we report on the oxidation of oxalic and malonic acids by the chromium(V1) ion. The mechanism involves the formation of coordinated complexes between the reactive species of the oxidant and the reductant which then decompose slowly yielding Cr(V) and an organic radical. These oxidations exhibit acid catalysis and are of first order with respect to Cr(V1) and malonic acid and of second order in oxalic acid. EXPERIMENTAL
The chemicals used were either BDH, Analar, or E. Merck G.R. grade. Potassium chromate and oxalic acid (or malonic acid) solutions of known concentration were brought to the same temperature in a thermostat and then mixed. The course of the reaction was followed iodometrically by withdrawing aliquots of the reaction mixture at known intervals of time and quenching the reaction by pouring in an excess of KI/H2S04 mixture. The liberated iodine was estimated by titrating against standard sodium thiosulfate solution. For spectrophotometric studies, a Unicam SP 1800 spectrophotometer was used. KINETICS AND MECHANISMS A. Oxalic
The present reaction is of first order in the concentration of Cr(VI), second order in oxalic acid (H,A) concentration, and directly proportional to the first ’ To whom correspondence should be addressed. 345 0026-265X/88 $1SO Copyrisbt AlI rights
63 1% by Academic F’ress, Inc. of reproduction in any fom reserved.
power of the hydrogen ion concentration. The oxidations of oxalic acid by Cr(V1) were all carried out in solutions where oxalic acid was present in a large excess over Cr(VI), and hence pseudo-first-order rate constants were obtained from the plots of [email protected]
) versus time (Fig. 1). The reaction followed first-order kinetics with respect to Cr(VI), which is confirmed by constant values of kobsobtained upon varying Cr(V1) concentrations (Table 1). The role of the hydrogen ion could be explained by the reactive speciesCr(H,0)6+ and H,A according to the equilibrium Cr(H20)6+ * Cr(OH)‘+ + HS H2A z$ HA- + Hi When H+ is increased, the equilibrium shifts to the left causing an increase in Cr(H20)6+ and H,A concentrations and also an increase in the rate of reaction which agrees with our experimental results. Taking all these facts into consideration leads to the formation of the mechanisms Cr(H20)6+ 3 Cr(OH)‘+ + H+
Cr(H20)a+ + H2A 2 [Cr(H20)(H2A)]
D-(H2O)(H2A>I + Hd zs K4 B-U-W)W2&
Time ,min., FIG.
plots at 35°C.
TABLE 1 Observed First-Order Rate Constants at Varying Temperatures and Hydrogen Ion Concentrations
H+ 0.052 0.020
k,(min - ‘)
k,(min - ‘) 0.219 0.184
0.380 0.317 0.205 0.155 0.025
1.778 x 10-3
7.080 x lo-’ 1.580 x 1O-3
0.014 6.310 x 1O-3 2.820 x 1O-3 1.580 x lo-’ 0.011
k,(min - ‘) 0.314 0.322 0.190 0.061 0.032 0.240” 0.23b
Note. Cr(VI) = 5 X 10e3 M; oxalic acid = 0.25 M. a Cr(V1) = 2.5 X 10e3 M. b Cr(VI) = 7.5 x 1O-3 M.
[WHAVW2&1 &r(V) R’ + Cr(V1) -0(V) fast fast 2 Cr(V) + 2L -2
Cr(III) + 2 Product
Cr3+ was the product of the oxalic acid oxidation by Cr(VI), being confirmed by the spectra of the product which is analogous to Cr3+ spectra at wavelengths of 420 and 580 nm (Fig. 2). Rocek et al. (9-12) showed that in the mechanism for the oxalic acid oxidation by Cr(V1) none of the chromium passes through a Cr(IV) intermediate. He also found that Cr(V) is only reduced by oxalate. A derivation gives the rate law
- 4WW dt
_ W& -
W’liH~A12KrWI)1 F-J+1+ K1
W&4 [H +ID-Ml2 [H+] + Kt
According to Eq. (9), plots of kzd,, versus [H+l- ’ should yield a straight line relationship as shown in Fig. 3 (Table 1). Also, plots of kobs.versus [H2A12 should yield a straight line relationship and pass through the origin as shown in Fig. 4 (Table 2). Second order with respect to oxalic acid was also obtained by Jones and Waters (5) in the oxidation of oxalic acid by vanadium(V).
FIG. 2. The visible spectra of (3) reaction product of CrWI) and malonic acid, (2) reaction product of Cr(V1) and oxalic acid, and (1) Cr(II1) solution.
FIG. 3. Variation of kil with [H+]-’ acid] = 0.25 M.
at different temperatures.
[CtiVI)I = 5 X 10e3hf; [O.dk
FIG. 4. Variation of k, with [oxalic acid]’ at different [H+l. [Cr(VI)I = 5 x 10m3M; temp. = 25°C.
Reaction rates were measured over the temperature range 25-3X and the energy of activation AE* was evaluated using the Arrhenius equation and was found to be 7.23 kcal mol- ‘. The entropy of activation was found to be -44.98 e.u. at 35°C. B. Malonic Acid
The kinetics of oxidation of malonic acid by Cr(V1) was found to be of first order with respect to Cr(V1) and malonic acid. The first-order rate constant was found to vary directly with the first power of the hydrogen ion concentration. TABLE 2 Observed First-Order Rate Constants at Varying Oxalic Acid Concentrations H+ = 0.052 M Oxalic acid WI
4 (min-‘) 0.07 0.16 0.22 0.34
0.15 0.20 0.25 0.30 Note. Cr(V1) = 5
H+ = 0.020 M
Oxalic acid VW
k, (min- ‘)
0.10 0.15 0.20 0.25
0.020 0.048 0.122 0.184
low3 M; temperature = 25°C.
Malonic acid concentration was present in a large excess over Cr(V1) and firstorder plots were obtained from the plot of log&-x) versus time (Fig. 1). First-order kinetics were seen with respect to Cr(V1) by obtaining constant values of kohs.upon varying Cr(V1) concentration (Table 3). As in the oxidation of oxalic acid, the reactive specieswere Cr(H,0)6f and H,A. These reactive species were postulated by Oswal and Bakore (13) in the oxidation of malonic acid by chromic acid. The mechanism suggested involves formation of the intermediate formed between the reactive species of the two reactants Cr(H,0)6+ and H,A which then decompose with a rate-determining step, Cr(H20)6+ $! Cr(OH)S+ + H+
Cr(H20)6f + H2A 3 [Cr(H20)6+ . &Al
[Cr(H20)6+ * H2A]--r(V) fast R’ + Cr(VI) -Cr(V) fast 2Cr(V) + 2L -2
Cr(II1) + 2F’roduct
As in oxalic acid, Cr3+ is the product of the oxidation of malonic acid by Cr6’ as shown in Fig. 2. From the above mechanism, the observed first-order rate constant is
TABLE 3 Rate Constants at Varying Temperatures
k, x 10’
k, x lo3
k, x 10’
0.050 0.079 0.120 0.159 0.079
2.88 5.47 8.90 12.95 5.50”
0.040 0.089 0.126 0.209
3.91 8.92 12.60 20.00
0.045 0.063 0.132 0.251
0.065 0.092 0.018 0.046
Note. Cr(V1) = 5 X lo-’ a Cr(V1) = 1 X lo-’ M.
iU; malonic acid = 0.875 M.
k&3 [H+lFWd =
According to Eq. (16), plots of kobs. versus [H+] give a straight line relationship as shown in Fig. 5 (Table 3). Also, plots of kobs. versus [H,A] yielded a straight line relationship as shown in Fig. 6 (Table 4). Similar observations were obtained by Oswal and Bakore (13) in the oxidation of malonic acid by chromic acid. Reaction rates were measured over 35-45”C and the energy of activation was found to be 13.34 kcal mol - ‘, while the entropy of activation was found to be - 34.78 e.u. at 35°C. The values obtained for AE* are in good agreement with that obtained by Oswal and Bakore (13) in the oxidation of malonic acid by chromic acid (AE = 12.9 kcal mol-‘). DISCUSSION The suggested results involve formation of complexes between Cr(VI) and the reductant as intermediates during the reaction. This is essential for the transfer of an electron from the reductant to Cr(VI) with the simultaneous breaking of the carbon bond of the acid. Previous studies indicated that the formation of cyclic intermediates favors the movement of an electron to the positively charged metal l
i 36 34 32 30 28 26 24 22 20 18 16 14 12 10 8 6 4 2
FIG. 5. Variation = 0.875 M.
of k, with
= 5 X 10e3 IU; [malonic
c 3 9 0.0200.018
FIG. 6. Variation of k, with malonic acid concentration at 40°C. [H+] = 0.141; [Cr(VI)] = 5 x 1O-3 M.
center (5,8). This constitutes the first step in the reaction chain and is followed by a number of rapid steps to form the final products. The suggested intermediates are given by the formula
Ho Cr ‘0
Oxalic acid forms five-membered ring intermediates which react with a further molecule of oxalic acid, whereby decomposition of the complex appears to be of second order with respect to oxalic acid. Malonic acid forms six-membered ring intermediates. The higher reactivity of oxalic acid than that of malonic acid in the reduction of Cr(V1) (as shown in Table 5 from the values of the observed secondorder rate constants) was due to the second molecule of oxalic acid attached to TABLE 4 Observed First-Order Constants at Varying Malonic Acid Concentrations Malonic acid (M) 0.595 0.714 0.833 0.984
k, X [email protected]
‘) 10.36 13.24 16.12 18.42
Note. Cr(V1) = 5 x 10e3 M; temperature = 40°C; H+ = 0.141.
TABLE 5 Activation and Kinetic Parameters Reductant Oxalic acid MaIonic acid
kobs,at 35°C (M-l min-‘) 1so at H+ = 0.018 3.29 x 10m3at H+ = 0.050
Al? (kcal mol - ‘)
-44.98 - 34.78
Cr(VI) in the intermediate complex facilitating the carbon-carbon bond fission through the inductive effect of the two carboxyl groups of oxalic acid. Comparison of the values of the energy of activation for oxalic and malonic acids showed them to be in good agreement with the values of their observed second-order rate constants (Table 5). The entropy of activation also shows a high negative value which was attributed to the loss of translational and vibrational degrees of freedom during the formation of a cyclic transition state. If cyclic intermediates are formed from noncyclic reactants, a decrease in the values of AS* is expected (24). On the basis of the formation of the above intermediates, one can expect that the oxidation of succinic acid by Cr(V1) does not proceed due to the formation of a seven-memberedring as the intermediate complex because it is a very unstable ring. This conclusion was confirmed by Yadav and Bhagwat (8) who showed that succinic acid is not oxidized by ceric sulfate. REFERENCES R. M. Trans. N.Y. Acad. Sci., 1951, 13, 314-318. Ladbury, J. W.; Cullis, C. F. Chem. Rev., 1958, 58, 403-438. Srinivasan, V. S.; Venkatasubramanaian, N. Indian J. Chem., 1973, 11, 702-103. Kansal, B. D.; Nepal, Singh. J. Indian Chem. Sot. LV, 1978, 304-307, 618-620. Jones, J. R.; Waters, W. A. J. Chem. Sot., l%l, 47574761. Kurihari, H.; Nozaki, T. J. Chem. Japan, 1962, 83, 708-711. Kemp, T. J.; Waters, W. A. J. Chem. Sot., 1964, 3101-3106. Yadav, R. L.; Bhagwat, W. V. J. Indian Chem. Sot., 1964, 41(5), 389-393. Hasan, F.; Rocek, J. J. Amer. Chem. Sot., 1972, 94, 90729081. Srinivasan, V.; Rocek, J. J. Amer. Chem. Sot., 1974, %, 127-133. Hasan, F.; Rocek, J. Tetrahedron, 1974, 30, 21-24. Haran, F.; Rocek, J. J. Org. Chem., 1974, 39, 2612-2615. Oswal, S. L.; Bakore, G. V. Indian J. Chem., 1980, 19, 211-213. Gould, E. S. Mechanism and Structure in Organic Chemistry, p. 233. Holt, Rinehart & Winston, New York. 1962.
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