Kinetics of oxidation of pyruvate ion by hexachloroiridate(IV)

.L inorg, nacL Chem. VoL 42, pp. 1177-1179 Pergamon Press Ltd.. I~0. Printed in Great Britain

KINETICS OF OXIDATION OF PYRUVATE ION BY HEXACHLOROIRIDATE(IV) KALYAN K. SEN GUPTA,* UMA CHATTERJEE, SURAJIT MAITI, PRATIK K. SEN and B. BHATTACHARYA Department of Chemistry, Jadavpur University, Calcutta-700032, India

(Received 1 August 1979; receivedfor publication 22 October 1979) Abstract--The kinetics of oxidation of pyruvate ion by lr(IV) have been studied spectrophotometrically in sodium acetate-acetic acid buffer. The reaction is first order with respect to (substrate) and (Ir(IV)). Hydrogen ion inhibits the rate of reaction. The addition of different salts accelerates the rate of reaction. The rate also increases with the increase in dielectric constant. The different activation parameters have been evaluated. The existence of a free radical intermediate has been demonstrated.

INTRODUCTION Pyruvic acid plays an important role in the degradation of sugars[la], in Kreb's cycle and in cellular respiration [1 b]. The kinetics and mechanism of the oxidation of pyruvic acid by transition metal ions [2-5] have been studied. These reactions exhibit a remarkable variety of mechanistic behaviour. There have been very few reports on the kinetics of oxidations by Ir(IV) and majority of them are concerned with the reactions of inorganic complex compounds[6a]. Very few studies on the kinetics of the oxidation of organic compounds by hexachloroiridate(IV) have been reported[7-10] and no attempt seems to have been made to study the reaction between pyruvate ion and hexachloroiridate(IV).

Stoichiometry and product. Identification of acetic acid as the reaction product[5] when hexachloroiridate(IV) is present in excess over pyruvate ion (in HCIO4) requires the stoichiometry. CH3COCOOH+ 2 Ir(IV) + H20"-) CHsCOOH+21r(III)+CO2 +2H +. (1)

EXPERIMENTAL

Reagents. Sodium pyruvate (Biochemical grade) was used without further purification. Sodium chloroiridate (JohnsonMathey) was used. Standardisation of Ir(IV) was effected spectrophotometrically [I 1]. Sodium perchlorate was prepared by neutralizing perchloric acid with sodium hydroxide. Sodium chloride and potassium chloride were of analytical reagent grades. All other reagents were of highest purity available. Dioxane was purified by distilling twice over metallic sodium. All solutions were prepared in doubly distilled water. Absorption spectra. The absorption spectra of Ir(IV) solutions were recorded in the visible region. An absorption maximum at 488 nm and a minimum at 460 nm were in good agreement with the published value[12]. Solutions in dilute acid were stable to aquation over a day. Absorption contributions by Ir(IIl) in the UV-visible region are relatively small and could be neglected. This is supported by the work of Paulson and Garner[13] who showed that IrCl63- or its aquo product Ir(OH2)CI52- is transparent in the region of maximum absorption of IrCl62-. Kinetic measurement. A large excess of pyruvate ion was used for all kinetic measurements. The reactions were followed spectrophotometrically as has been mentioned in an earlier communication[6b]. The rate of decrease of Ir(IV) was followed to at least 50% reduction of the Ir(IV). The monitoring wave length was 488 nm. The reactions were studied at 30°C unless otherwise mentioned. The pseudo first order rate constant was calculated from the slope of the logarithm of the absorbance vs time plot. Generally more than 6 experimental points were taken in each ran, The rate constants were reproducible to within _ 3%. *Author for correspondence.

Potentiometric titration of the reaction mixture (in HCIO4) indicates that one mole of substrate is oxidized to one mote of acetic acid. The degree of hydrolysis of iridium products[10] was determined as follows. The reduced iridium was reoxidized by passing chlorine gas (dry and HCI free) for a long time. The optical densities of the solution were measured at 460 and 487 nm. The extinction coefficients of IrCl62- are 2080 and 4050 and those of IrCI5OH2- are 3100 and 2345 mol-~ cm-l respectively. It has been observed that more than 90%IRC163- aquated to IrCIs(OH2)2after 5 hr although the aquation of IrCl63- is insignificant during the kinetic experiments (< 5%). Polymerization test. A polymerization test on sodium pyruvate and hexachloroiridate(IV) was made in 40% (w/v) of acrylamide. I ml of the substrate (2.0 × l0 ~M) was added to a mixture of 2ml of Ir(IV) (l.19xl0-3M) and 1.25ml of acrylamide. The acidity and total volume was adjusted to 1.0 x 10-4 M and 10 ml respectively. Addition of dioxane to the mixture gave a white suspension. Control experiments in which either substrate or Ir(IV) was excluded gave no detectable polymer in dioxane.

RESULTS

Effect of reactant concentration The reactions were studied at various concentrations of Ir(IV) but at constant substrate concentration, pH and temperature, viz. 2.0x 10-2M, 3.42 and 30°C respectively. The values of kobsO are (3.2-+ 0.1)× 10-4(sec -t) in the concentration range of (1.0-4.0)× 10-4M of Ir(IV). The observed pseudo first order rate constants are found to be independent of the initial Ir(IV) concentration indicating that the reaction is first order with respect to Ir(IV). The reaction was also studied overwide range of substrate concentrations when the initial It(IV) and pH were held constant. The quotient, kobsd/(substrate) at different concentrations are recorded in Table 1. From inspection of the data it is clear that the kinetics of oxidation is much simpler at lower substrate concentrations. The plot of kob~d against (substrate) is linear indicating further that the order with respect to substrate (<2.0 x 10-2 M) is one. The reaction seems to be cam-

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1178

KALYAN K. SEN GUPTA et al. Table I. Effect of substrate concentration at different temperatures (Ir(IV))= I. 19x 10-4 M, pH = 3.42 (Substrate) × 102M kobsdX liP (sec-1) kobsd/(Substrate)x 102 (lit tool-i sec-i) [ (Substrate) x 102M (b) ~ kobsd × liP(sec -~) [ kobJ(Substrate) × 102 (lit tool-j sec-I) f (Substrate) × 102M (c) kobsd× liP (seC-~) kobsd/(Substrate)× 102 (lit mol-t sec-I) (Substrate) × 102M (d) kobsa× liP (sec-~) kobsd/(Substrate)× 102 (lit tool-I sec-~) (a)

0.25 0.41 1.64

0.5 0.815 1.63

1.0 1.60 1.60

1.5 2.60 1.73

2.0 3.20 1.60

3.0 6.39 2.13

5.0 14.54 2.91

0.25 0.51 2.04

0.5 1.0 2.0

1.0 2.0 2.0

1.5 3.30 2.2

2.0 4.20 2.1

3.0 9.08 2.69

5.0 16.2 3.25

0.25 0.66 2.64

0.5 1.20 2.40

1.0 2.60 2.60

1.5 4.0 2.66

2.0 5.8 2.90

3.0 10.1 3.37

5.0 20.7 4.14

0.25 0.96 3.84

0.5 1.85 3.90

1.0 3.70 3.70

1.5 5.80 3.86

2.0 8.0 4.0

3.0 14.54 4.84

5.0 20.06 5.01

(a), (b), (c) and (d) denote experiments at temperature 30, 34.4, 39 and 43.5°C respectively. plicated at higher substrate concentrations (>2.0x 10-2M) when the order with respect to substrate becomes (1.4_+0.1).

Effect o[ acidity The effect of changing acidity on the pseudo first order rate constants has been studied by the addition of sodium acetate-acetic acid buffer. The ionic strength was held constant (I = 0.20M) by the addition of sodium perchlorate. Hydrogen ion inhibits the rate of reaction (Table 2) and the plot of log kobsd against log(H +) is linear over the entire acid range studied. The slope of the plot is - 0.4.

Table 4. Effect of solvent on the pseudo first order rate constant (Substrate) = 2.0 x 10-2 M; 0r(IV)) = 1.19× 10-4 M; pH = 3.94 % Dioxane (v/v)

0

10

20

25

~14

76.7 7.62

67.7 5.02

58.65 3.79

54.1 3.03

kob~l X liP (sec-I)

pseudo first order rate constant decreases with the decrease in the dielectric constant (~) of the medium. The decrease in rate was about 60% in 25% dioxane (v/v), so that the solvent effect also indicates that the reaction takes place between ions of the same charge.

Effect o[ salts The effects of the addition of salts such as sodium perchlorate, sodium chloride, and potassium chloride were investigated. The salts produced a marked increase in the pseudo first order rate constants. The results are shown in Table 3. The pseudo first order rate constant increases linearly with the increase in salt concentrations. The positive salt effect indicates a mechanism in which the rate determining step is presumed to involve two anions as reactants.

Effect of solvent The solvent effect was studied by the addition of dioxane. The results presented in Table 4 shows that the Table 2. Effect of pH on the pseudo first order rate constant (Substrate) = 2 x 10-2 M; (!r(IV))= 1.19x 10-4 M; I = 0.20M pH 3.42 3.72 4.05 4.27 4.45 4.63 kob~xliP(sec-I) 4.93 5.76 9.03 10.03 13.09 15.27

Table 3. Effect of salt concentrations on the pseudo first order rate constant 0r(IV)) = 1.19x 10-4 M, (Substrate) = 2.0 x 10-2 M, pH= 3.42; ¢ kobsdX liP (sec-1) (Salt) × I0 M

NaCIO4

NaCl

KCI

0 0.5 1.0 1.5 2.0 2.5 3.0

3.20 4.26 5.91 6.82 8.44 9.98 --

3.20 4.57 5.99 7.38 8.61 10.41 11.58

3.20 4.93 7.04 8.67 10.01 14.18

Activation parameters The oxidation of pyruvate ion by hexachloroiridate(IV) was studied at four different temperatures. The second order rate constants (average of 4 determinations) are (1.64-+0.06)×10 -2, (2.07_0.08)×10 -2, (2.64_+ 0.1) × 10-2 and (3.86_+0.1) × 10-2 dm 3 mol-J sec -I at 30, 34.4, 39.0 and 43.5°C respectively. The energy of activation has been calculated from a log k2 (k2 = second order rate constant) vs lIT plot. The values of Ea and AS" are 50---4kJmol -~ and - 132-+ 14JK -~ mol -~ respectively. The negative entropy of activation is to be expected if reactants are ions of the same charge. DISCUSSION The reaction is too slow to be studied at higher acidities (> 10-1M) but it takes place at a measurable rate when the acidity is < 10-3 M. Hence, the reaction was studied in the pH range 3.42-4.63. The dissociation constant of pyruvic acid is 3.2 x 10-3 at 25°C. Moreover, pyruvic acid is partially hydrated [2] in aqueous solution and hexachloroiridate(IV) is known to be stable towards substitution or hydrolysis over a wide range of acidities [10]. The reaction is not retarded by the hexachloroiridate(III) produced. This indicates that the rate determining step is not preceded by an equilibrium involving Ir(III). The formation of acetate ion and finally acetic acid from pyruvate ion is expected when hexachloroiridate(IV) oxidizes hydrated form of pyruvate ion. It is, therefore, suggested that hexachloroiridate(IV) ion attacks the hydrated form of pyruvate ion (PA-) to give a radical ion (PA=) and Ir(III) in the slow step. Although Ir(II) is known[17], it is very unlikely to be involved, so that it is rather difficult to suppose that Ir(IV) acts as a

Kinetics of oxidation of pyruvate ion by hexachloroiridate(IV) two equivalent oxidant especially when the tests with acrylamide have proved the existence of an organic radical during the reaction. We suppose, therefore, that the radical ion picks up another IrC162 and by a fast step gives the products of oxidation (step 3). The oxidation of pyruvate ion at low substrate concentrations (~<2.0× 10 2 M) may, therefore, take place according to the steps (2) and (3). Since most of the experiments were carried out at pH 3.42 and the dissociation constant of acetic acid is 1.8 × 10 ~ at 25°C, acetate ion will be protonated to give acetic acid in solution. The steps of the reaction may be, CH3C'(OH).,COO- + IrC162 (PA-)

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1 : 1 complex reacts further with another substrate to give 1 : 2 complex which then decomposes or the 1 : 1 complex

oxidizes another substrate molecule to give products. It has earlier been shown[7] that oxidation of glyoxylate ion by Ir(IV) proceeds through the formation of 1:1 intermediate complex by a fast step followed by its decomposition in the rate determining step. Thus the present reaction seems to be kinetically different from the oxidation of glyoxylate ion. This may be due to the fact that unlike glyoxylate ion oxidation, the oxidation of pyruvate ion takes place by the rupture of a C-C bond. Acknowledgements--Thanks are due to U.G.C. for financial assistances to U.C. and S.M.

.~,ow

CH~C(OH)O'COO +IrCk 3 - + H ÷

(2)

(PA:) REFERENCES

CH3C(OH)O'COO- + IrCk: PA:

fast

CH3COOH+ IrC163 + CO:.

(3)

The electron transfer from hydrated pyruvate ion to the oxidant and subsequent reduction of hexachloroiridate(IV) to hexachloroiridate(III) may take place in the following manner. CH3C(OH)2COO- + lrC162 . , CH3C(OH)O'COO- + HCI.IrC152-. HCI.IrCI~2

) H + + IrC163 .

(4) (5)

The oxidation of pyruvate ion at low substrate concentrations may also take place through the formation of an intermediate complex as the rate determining step (6). In the intermediate complex Ir(IV) exhibits a co-ordination number of 7. Such complexes are well known[15, 16]. However, the complex decomposes to give the products of the reaction according to the step (7). The steps of the reaction may be P A - + I r C k 2-

(PA IrCl6)3-

slow •(PAirCl6) 3-

(6)

f~st ~ Products.

(7)

The complex order with respect to substrate concentrations at > 2.0 x 10-2 M may be due to the fact that the

1. (a) E. S. G. Barron, Trends in Physiologyand Biochemistry, p. 471. Academic Press, New York (1952); (b) A. N. Nesmeyanov and N. A. Nesmeyanov, Fundamentals of Organic Chemistry, VoL 2, p. 135. Mir, Moscow (1977). 2. K. K. Sen Gupta and T. Sarkar, Tetrahedron 31,123 (1975). 3. K. K. Sen Gupta and S. Aditya, Anal. Chem. Acta 29, 483 (1963). 4. K. K. Sen Gupta, T. Sarkar, S. Sen Gupta and H. R. Chatterjee, Ind. J. Chem. 14A, 583 (1976). 5. K. K. Sen Gupta and H. R. Chatterjee, Inorg. Chem. 17, 2429 0978). 6. (a) W. P. Griffith, The Chemistry of the Rarer Platinum Metals, p. 285. Interscience, London (1967); (b) K. K. Sen Gupta and P. K. Sen, Inorg. Chem. 18, 979 (1979). 7. K. K. Sen Gupta, U. Chatterjee and P. K. Sen, Indian J. Chem. 16A, 767 (1978). 8. S. P. Mushran, M. C. Agarwal and K. C. Gupta, Indian J. Chem. 10, 642 (1972). 9. E. Pelizzetti, E. Mentasti and C. Baiocchi, J. Phys. Chem. 80, 2979 (1976). 10. (a) R. Cecil and J. S. Littler, J. Chem. Soc. (B), 1420 (1968); (b) R. Cecil, J. S. Littler and G. Easton, J. Chem. Soc. (B), 626 (1970). 11. S. Payne, Analyst ~ , 698 (1970). 12. A. G. Sykes and P. W. F. Throneley, J. Chem. Soc. (A), 2, 232 (1970). 13. I. A. Poulson and C. S. Garner, J. Am. Chem. Soc. $4, 822 (1962). 14. G. Akertof and O. A. Short, J. Am. Chem. Soc. 58, 1241 (1936). 15. F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, p. 524. Interscience, New York (1962). 16. A. E. MarteU and M. Calvin, Chemistry o[ the Metal Chelates, p. 212. Prentice Hall, Englewood Cliffs, New Jersey (1962). 17. Ref. [6], p. 244.