Indian Journal of Chemistry Vol. 48A, June 2009, pp. 797-800
Kinetics and mechanism of oxidation of formic and oxalic acids by bis(pyridine) silver permanganate Jayshree Banerjia, Laszlo Kotaib & Kalyan K Banerjic, * a
Department of Biotechnology, Meerut Institute of Engineering and Technology, Meerut 250 005, India b Chemical Research Center, Hungarian Academy of Sciences, H-1025, Pusztaszeri u. 59-67, Budapest, Hungary c Faculty of Science, National Law University, Mandore, Jodhpur 342 304, India Email:
[email protected] Received 23 January 2009; revsied and accepted 5 May 2009
Kinetics and mechanism of oxidation of formic and oxalic acids by bis(pyridine) silver permanganate in aqueous acetic acid have been studied. The main product of oxidation is carbon dioxide. The reaction is first order each with respect to the permanganate and the reductant. The reaction is catalyzed by hydrogen ions. The oxidation of α-deuterioformic acid exhibits a substantial primary kinetic isotope effect (kH/kD = 3.69 at 298 K). Suitable mechanisms have been proposed. Keywords: Kinetics, Reaction mechanisms, Oxidations IPC Code: Int. Cl.8 C07B 33/00
Selective oxidation of organic compounds is an important reaction in synthetic organic chemistry. Inorganic permanganate salts are drastic and nonselective oxidants. A number of different permanganate derivatives have been reported1 as relatively mild and selective oxidizing reagent. We have recently reported the thermally induced intra-molecular redox reaction of bis(pyridine) silver permanganate (BPSP)2. BPSP has been widely used in organic chemistry as an oxidant3. Oxidation of oxalic acid by permanganate derivatives has been a fascinating subject of mechanistic studies for more than half a century4,5. However, there seems to be no report on the kinetics and mechanism of oxidation by BPSP. We report herein, the kinetics of oxidation of oxalic and formic acids by BPSP. Experimental BPSP and α-deuterioformic acid (DCO2H or DFA) were prepared by the reported method6,7. Oxalic acid (OA) and formic acid (FA) were commercial products and were used as such. Perchloric acid has been used as a source of hydrogen ions. FA solutions were
standardized by alkalimetry. Acetic acid was refluxed with acetic anhydride and chromic oxide for 3 h and then fractionated. To determine the stoichiometry, an excess of BPSP (× 10 or greater) was reacted with the organic acid in 100 cm3 of 1:1 acetic acid-water (v/v) and the amount of residual BPSP after the completion of reaction was measured spectrophotometrically at 490 nm. In another set of reaction, a known excess of the reductant was treated with BPSP in aqueous acetic acid solution and the residual reductant was determined by oxidation titrations. The results showed that permanganate ion and the reductant react in the ratio of 1:2.5. No quantitative determination of carbon dioxide formed was carried out. The evolution of carbon dioxide was detected by lime water test. The reactions were followed under pseudo-first order conditions by keeping a large excess (×15 or greater) of the organic acid over BPSP. The temperature was kept constant to ±0.1oC. The solvent was 1:1 acetic acid-water (v/v), unless specified otherwise. The reactions were followed by monitoring the decrease in the concentration of BPSP spectrophotometrically at 490 nm for up to 80% of the reaction. No other reactant or product had any significant absorption at this wavelength. The pseudo-first-order rate constants, kobs, were evaluated from the linear plots of log [BPSP] against time (r2 = 0.995-0.999). Duplicate kinetic runs showed that the rate constants were reproducible to within ±3%. Preliminary experiments showed that the reaction is not sensitive to changes in ionic strength; therefore no attempt was made to keep ionic strength constant. Results and discussion The oxidation of organic acids leads to the formation of carbon dioxide. The stoichiometric determination indicated the following overall reactions: 5(COOH)2 + 2MnO4- + 6H+ → 10CO2 + 8H2O + 2Mn2+ … (1) 5HCOOH + 2MnO4- + 6H+ → 5CO2 + 8H2O + 2Mn2+
… (2)
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The reactions are of first order with respect to BPSP. Further, the values of kobs are independent of the initial concentration of BPSP. The order of reaction with respect to the reductant is less than one (Table 1). A plot between 1/[reductant] against 1/kobs is linear (r2 > 0.9990) with an intercept on the rate ordinate. Thus the reaction exhibits Michaelis-Menten type kinetics with respect to the organic acids. This indicates the following overall mechanism and rate law at constant hydrogen ion concentration. K1 Reductant + BPSP [Intermediate]
… (3)
k1 [Intermediate] → Products
… (4)
- d[BPSP] / dt = k1 K1 [reductant] [BPSP] / 1 + K1 [reductant] … (5) The variation in the reductant concentration was studied at different temperatures. The values of the overall formation constants and rate constants of the decompositions of the intermediate complex were evaluated from the double reciprocal plots. The thermodynamic and activation parameters were calculated from the values of the equilibrium constants Table 1Rate constants for the oxidation of oxalic and formic acids by BPSP. {[H+] = 0.20 mol dm-3; T = 298 K} 3
10 [BPSP] (mol dm-3)
4
[Organic acid] (mol dm-3)
Table 2Formation constants and thermodynamic parameters of the organic acids-BPSP complexes. {[H+] = 0.02 mol dm-3}
OA FA DFA
K1 (dm3 mol -1) at temp. (K) =
∆H ∆S ∆G (kJ mol−1) (J mol1 K−1) (kJ mol−1)
288
298
308k
318
6.83 5.18 5.23
4.00 3.82 3.85
2.32 2.55 2.62
1.34 -43.8±0.7 1.72 -30.6±1.1 1.69 -31.1±1.2
Table 3Rate constants of decomposition and activation parameters of organic acid-BPSP complexes. {[H+] = 0.02 mol dm-3} Subst. 104 k1 (s-1) at temp. (K) =
−1
10 kobs (s ) Oxalic acid Formic acid 1.00 0.02 2.00 3.73 1.00 0.04 3.67 6.97 1.00 0.08 6.42 12.2 1.00 0.12 8.57 16.6 1.00 0.16 10.4 20.0 1.00 0.20 11.8 23.1 1.00 0.25 13.2 25.5 1.00 0.30 14.5 28.0 2.00 0.20 12.1 23.0 4.00 0.20 11.8 23.3 6.00 0.20 11.6 22.7 8.00 0.20 11.9 23.0 1.00 0.20 12.0a 22.9a a Contained 0.001 mol dm-3 acrylonitrile.
Subst.
and rate constants of the decomposition at different temperatures respectively (Tables 2 and 3). The reaction is catalyzed by hydrogen ions (Table 4). A plot of rate versus hydrogen ion concentration is a curve concave to the rate axis and makes an intercept on the rate axis. This indicates that there are more than one protonated species, all species are reactive and the protonation constants are small8. The dependence on hydrogen ions has the form kobs = a + b[H+] + c [H+]2. The values of the constants are recorded in the Table 4. Similar dependence on acidity was observed in the oxidation of several organic compounds by bis (2,2'-bipyridyl) copper(II) permanganate9. It implies that two catalytic protons are added in equilibria. The oxidation of organic acids by BPSP, in an atmosphere of nitrogen, failed to induce the polymerization of acrylonitrile. Further, rate of oxidation was not affected by the addition of acrylonitrile (Table 1). To further confirm the absence of free radicals in the reaction pathway, the reaction was carried out in the presence of 0.05 mol dm-3 of 2,6-di-t-butyl-4-methylphenol (butylated hydroxyltoluene or BHT). It was observed
−128±2 −84±4 −86±4
-5.9±0.5 -5.7±0.9 -5.7±1.0
288
298
308
∆H*
∆S*
∆G*
318 (kJ mol ) (J mol K ) (kJ mol−1) −1
1
−1
OA 1.31 2.65 5.78 12.6 55.0±1.4 −110±4 87.6±1.1 FA 2.75 5.26 10.0 19.2 46.7±0.7 −132±2 86.0±0.6 DFA 0.74 1.43 2.75 5.33 47.5±0.7 −140±2 89.2±0.6 kH/kD 3.71 3.69 3.64 3.60 Table 4Effect of hydrogen ion concentration on the reaction rate. {[FA] = 0.40 mol dm-3; [OA] = 0.08 mol dm-3; T = 298 K} [H+] ( mol dm-3)
0.05 0.10 0.20 0.40 0.50 1.00 1.50 1.80 2.00 2.50 a b c r2
104 kobs (s-1) FA
OA
6.01 6.31 6.97 8.62 9.70 16.1 24.3 31.1 35.4 48.7 5.68 × 10-4 5.66 × 10-4 4.61 × 10-4 0.9991
5.73 5.73 6.42 8.24 9.33 15.5 24.2 30.5 34.2 47.6 5.06 × 10-4 6.22 × 10-4 4.29 × 10-4 0.9990
NOTES
799
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INDIAN J CHEM, SEC A, JUNE 2009
that BHT was recovered unchanged, almost quantitatively. Thus, a one-electron oxidation giving rise to free radials is unlikely. The reaction was studied in the presence of added Mn(II) sulphate. There is no effect of Mn(II). Thus intermediate valence states of manganese play no role up to the rate-determining step of the reaction. To ascertain the importance of the cleavage of the α-C-H bond in the rate-determining step, the oxidation of DFA was studied. The results recorded in Tables 2 and 3, show that the formation constants of the intermediate complexes of FA and DFA have almost identical values. Thus, a deuterium substitution does not affect the complex formation. The decomposition of the intermediate, however, exhibits the presence of a substantial primary kinetic isotope effect (kH/kD = 3.69 at 298 K). The oxidation was studied in solvents containing different amounts of acetic acid and water. The rate of oxidation increases with an increase in the amount of acetic acid in the solvent. This may be attributed to the change in the acidity of the medium with a change in the amount of acetic acid. Hammett’s acidity function, H0, for low concentration of perchloric acid in a series of acetic acid-water mixtures has been determined. It was observed that the acidity increases as the concentration of acetic acid increases. The studied reaction is acid-catalyzed and with an increase in the acidity of the solution, the rate is expected to increase. A one-electron oxidation, giving rise to free radicals, is not likely to be operative in this reaction, in view of the failure to induce polymerization of acrylonitrile and recovery of unchanged BHT. The presence of a substantial kinetic isotopic effect confirms that a α-C-H bond is cleaved in the rate-determining step in the oxidation of formic acid. The presence of Michaelis-Menten kinetics indicates the formation of an intermediate at pre-equilibrium. The intermediate may well be a mixed anhydride. Formation of similar
intermediate has been proposed earlier also in the oxidation of these acids by Cr(VI) compounds10. Mechanisms depicted in Schemes 1 and 2 account for all the experimental data. The non-linear transition state implied in the anhydride mechanism is supported by the relatively lower value of the kinetic isotope effect. Initially Mn(VII) is reduced to Mn(V), which is likely to undergo further reduction to yield finally Mn(II). Such a sequence of reactions in Mn(VII) oxidations is well known11. The mechanisms depicted in schemes lead to the following rate equation assuming that K1 >> K2 and K1 >> K3. Rate = [OA] [BPSP] (K1k1 + K2k2 [H+] + K3k3 [H+]2)/ (1 + K1[OA]) … (6) OA represents either formic acid or oxalic acid. References 1 Wentzell P D, Wang J-H, Loucks L F & Miller K M, Can J Chem, 76 (1998), 1144 and refs. cited therein. 2 Kéki S & Beck M T, React Kinet Catal Lett, 44 (1991) 75. 3 Firouzabadi H, Sardarian A R, Naderi M & Vessal B, Tetrahedron, 40 (1984) 5001. 4 Kotai L, Fodor J, Jakab E, Sajo I, Szabo P, Lonvi F, Valyon J, Gacs I, Argay G & Banerji K K, Trans Met Chem, 31 (2006) 30. 5 Chandra R, Sarkar A & Biswas N, Proc Indian Natn Sci Acad, Part A., 60 (1994) 565; Anjaneyulu A S R, Umasundari P & Sastry Ch V M, Indian J Chem, 25B (1986) 955. 6 Kotai L, Sajo, I, Fodor J, Szabo P, Jakab E, Argay G, Holly S, Gacs I & Banerji K K Trans Met Chem 30 (2005) 939. 7 Wiberg K B & Stewart R, J Am Chem Soc, 78 (1956) 1214. 8 Gupta Y K & Gupta K S, J Chem Educ, 61 (1984) 972. 9 Satsangi B K, Kothari S & Banerji K K, Proc Indian Acad Sci (Chem Sci), 108 (1996) 421 and refs. cited therein. 10 Purohit P, Kumbhani S, Shastri I, Banerji K K & Sharma P K, Indian J Chem, 47A (2008)671; Bishnoi G, Malani N, Sindal R S & Sharma P K, Oxid Commun, 30 (2007) 607. 11 Wiberg K B & Richardson W H, Oxidation in Organic Chemistry, Part A (Academic Press, New York) 1965.
