Research Article | Open Access
M. Sanjana, A. K. Patnaik, S. K. Badamali, P. Mohanty, "Kinetics and Mechanism of Electron Transfer to Heptavalent Manganese by DL-Aspartic Acid in Alkaline Aqueous and Micellar Media", Journal of Chemistry, vol. 2013, Article ID 724505, 7 pages, 2013. https://doi.org/10.1155/2013/724505
Kinetics and Mechanism of Electron Transfer to Heptavalent Manganese by DL-Aspartic Acid in Alkaline Aqueous and Micellar Media
The kinetics and mechanism of the electron transfer of DL-Aspartic acid (Asp) by Mn (VII) in alkaline medium has been studied spectrophotometrically over the range [Asp] mol dm−3; mol dm−3; K and mol dm−3 (KNO3). The reaction exhibits first-order dependence in but shows fractional-order dependence in both and . The reaction was studied in the presence of sodium dodecyl sulfate (SDS); an increase in the rate with the increase in the micellar concentration was observed. The products were characterized by spectral analysis. A mechanism involving free radicals is proposed. Asp binds to form a complex that subsequently decomposes to products. Activation parameters ° (kJ mol−1) and ° (JK−1 mol−1) for the reaction are and , respectively. The negative value of ° indicates that oxidation occurs via inner sphere mechanism.
Study of oxidation of amino acids has received considerable attention due to the importance of degradation of these compounds in biological systems. The Aspartic acid is a nonessential amino acid which is found abundantly in plant proteins especially in sprouting seeds. This can be manufactured in the body from oxaloacetic acid. It is of paramount importance in the metabolism during construction of amino acids and biochemicals in the citric acid cycle . Among the biochemicals that are synthesized from aspartic acid are asparagine, arginine, lysine, methionine, threonine, isoleucine, and several nucleotides.
Although the kinetics of oxidation of aspartic acid has been studied using different oxidants [2–9], the current work is an attempt to understand the redox chemistry of permanganate oxidation in alkaline media as well as in micellar media and to derive a plausible mechanism.
All chemicals used were of reagent grade. Double distilled water was used throughout the work. Stock solution of DL-Aspartic acid (SRL Chemicals) was prepared by dissolving the appropriate amount of the sample in double distilled water. The stock solution of KMnO4 (Merck) was prepared by using double distilled water and was stored in a dark place. It was standardized against oxalic acid by following the literature method . NaOH and were used to maintain the required alkalinity and ionic strength, respectively. The solution of sodium dodecyl sulfate was prepared by dissolving calculated amount of SDS (Merck) in double distilled water.
During kinetic investigation, the pH was maintained using a SYSTRONICS μpH system-361 equipped with a combination of glass Ag/AgCl/ (3 M NaCl) electrode. It was calibrated with standard buffers of pH 4.0, 7.0, and 9.0 (Merck). Absorbance was recorded with a Cecil CE-7200 UV-visible spectrophotometer equipped with a CE-2024 thermoelectric controller. Ten-millimeter quartz Suprasil cuvettes were used. IR spectra were taken in Varian FTIR spectrophotometer (USA).
3. Kinetic Measurements
The oxidation of DL-Aspartic acid by was followed under pseudo-first-order conditions where DL-Aspartic acid was taken excess over at 25°C ± 0.1°C. The reaction was initiated by mixing the required quantities of previously thermostated solutions of and DL-Aspartic acid, which also contained definite quantities of NaOH and to maintain required ionic strength. The progress of the reaction was followed by measuring the decrease in absorbance at 525 nm with time using a conventional mixing technique. was measured after the completion of the reaction (approximately after 24 hours of mixing) when the absorbance became almost constant. The plot of ln() versus was found to be linear as indicated in the following equation: where and are absorbance of the reaction mixture at time, , and at equilibrium, respectively. The correlation coefficients () of the plots used to determine were found to be 0.99. The pseudo-first-order rate constant () was calculated by the least squares method from the above relationship. The redox reactions were followed for about 3 half lives. The reported rate data represented as an average duplicate runs were reproducible to within ±3%.
The electron transfer reaction between DL-Aspartic acid and alkaline Mn(VII) has been studied over the range 2.0 ≤ [Asp] ≤ 5.0 mol ; 0.01 ≤  ≤ 0.05 mol ; 298 ≤ T ≤ 318 K and mol ().
The reaction orders were determined using the slopes of log versus log plots by varying the concentration of the reductant and while keeping other factors constant. With fixed concentrations of DL-Aspartic acid mol , and alkali, mol at constant ionic strength, 0.05 mol , the permanganate concentration varied from mol to mol . The linearity of plots of log(absorbance) versus time, for different concentrations of permanganate, indicates that the order in [Mn(VII)] is unity (Figure 1). The DL-Aspartic acid concentration was varied in the range of to mol at constant alkali and permanganate concentrations and constant ionic strength of 0.05 mol at 298 K. The values increased with an increase in DL-Aspartic acid over the concentration range shown in Figure 3. At low concentration of DL-Aspartic acid, the reaction was of first order and at high concentration of DL-Aspartic acid, the reaction was independent of DL-Aspartic acid.
4.1. Effect of Alkali Concentration
The effect of alkali concentration on the reaction was studied at constant ionic strength of 0.05 mol at 25°C. The  was varied in the range of 0.01 to 0.05 mol . The rate constant increased with an increase in alkali concentration (Figure 4), indicating a fractional-order dependence of the rate on alkali concentration.
4.2. Effect of Ionic Strength
The effect of ionic strength was studied by varying the potassium nitrate concentration from 0.05 to 0.5 mol at constant concentration of permanganate, DL-Aspartic acid, and alkali. Increasing ionic strength had no effect on the rate constant.
4.3. Effect of Temperature
The kinetics was also studied at five different temperatures with varying concentrations of DL-Aspartic acid and alkali, keeping other conditions constant. The rate constants were found to increase with the increase in temperature. The rate of the slow step was obtained from the slopes and intercepts of 1/ versus 1/[DL-Aspartic acid] and 1/ versus 1/[OH−] plots at five different temperatures (298–318 K).
4.4. Test for Free Radical
To test for the involvement of free radicals, acrylonitrile was added to the reaction mixture, which was then kept for 24 h under nitrogen atmosphere. Addition of methanol, resulted in the precipitation of a polymer, suggesting the involvement of free radicals in the reaction. However, the blank experiments with reactants in presence of acrylonitrile did not respond to positive test for free radical formation. Initially added acrylonitrile decreased the rate of reaction .
4.5. Effect of Ionic Surfactants
It is well established that most of the micellar reactions involving an ionic or neutral reactants are believed to take place either inside the stern layer or at interface between the micellar surface and bulk solvent water [12, 13]. The effect of the ionic and nonionic micelles on the reaction rates of bimolecular reactions is due to the association through electrostatic/hydrophobic and hydrogen bonding interactions between the reactants within a small volume of the self-assemblies . During the study, the concentration of sodium dodecyl sulfate (SDS) varied keeping the concentrations of DL-Aspartic acid, , temperature and ionic strength constant (Table 3). With the increase in the concentration of SDS, the rate tends to attend a limiting value at high surfactant concentration indicating a micellar binding of the substrate. The equilibrium constant values for the formation of complex (C) with permanganate for DL-Aspartic acid and cysteine are comparable indicating equal probability of formation of the complex.
4.6. Stoichiometry and Product Analysis
The reaction between DL-Aspartic acid and in alkaline medium has a stoichiometry of 1 : 4. The main reaction product is 3-oxopropanoic acid. It was treated with 2,4-DNP and kept in refrigerator for 24 hours, a yellow solid was separated and recrystallized with ethanol. It was characterized by FT-IR. The FT-IR spectra of DL-Aspartic acid and its product complex are similar. The C=N stretching band appears at 1619 cm−1 which is absent in DL-Aspartic acid. The other significant bands appear at 1411 cm−1 (for symmetric carboxylate stretching) 1520 cm−1 (N–H bending). The other reaction products are identified as ammonia (Nessler’s reagent test), CO2 (lime water test), and manganate ().
Under the experimental conditions at pH > 12, the reduction product of Mn(VII), that is, Mn(VI), is stable, and no further reduction is initially observed . During this reaction, color changes from violet Mn(VII) to dark green Mn(VI) through blue Mn(IV). It is clear from Figure 2 that the absorbance of decreases at 525 nm, while increases at 630 and 430 nm are due to Mn(VI). As the reaction proceeds, a yellow turbidity slowly develops, and after keeping for a long time the solution decolorizes and forms a brown precipitate. This suggests that the initial products might have undergone further oxidation resulting in a lower oxidation state of manganese. It appears that the alkali combines with permanganate to give [16, 17]. In the second step, combines with DL-Aspartic to form an intermediate complex. The variable order with respect to DL-Aspartic is most probably due to the complex formation between oxidant and DL-Aspartic prior to the slow step. A plot of 1/ versus 1/[Asp] (Figure 3) shows an intercept in agreement with complex formation. Further evidence for complex formation was obtained from the UV-vis spectra of reaction mixture. Two isosbestic point were observed for this reaction (Figure 2), indicating the presence of an equilibrium before the slow step of the mechanism [18, 19]. In our proposed mechanism, in the complex one electron is transferred from aspartic acid to Mn(VII). The cleavage of this complex (C) is assigned as the slowest step, leading to the formation of an DL-Aspartic radical intermediate and Mn(VI). The radical intermediate reacts with another Mn(VII) species, , to give the final products: Mn(VI), 3-oxopropanic acid and (Scheme 1). The effect of the ionic strength and dielectric constant on the rate is consistent with the involvement of a neutral molecule in the reaction. The suggested structure of complex (C) is given in Scheme 1.
From Scheme 1, the rate law can be derived as follows: The total  can be written as where “t” and “” stand for “total” and “free”, respectively. Similarly, total [OH−] can be calculated as In view of the low concentration of and DL-Aspartic acid used in the experiment, in (3), the terms  and [Asp] can be neglected.
Thus, Similarly, Substituting (3), (5), and (6) in (2), we get Equation (8) is consistent with the observed orders with respect to different species, which can be verified by rearranging to (9) According to (9), other conditions being constant, plots of 1/ versus 1/[Asp] and 1/ versus 1/[OH−] should be linear (Figures 3 and 4). The slopes and intercepts of such plots lead to values of , , and k (Table 2). With these values, the rate constants were calculated under different experimental conditions. The thermodynamic quantities for the first and second equilibrium steps of Scheme 1 can be evaluated. The [DL-Aspartic acid] and [OH−] (Table 1) were varied at five different temperatures. van’t Hoff’s plots of log versus 1/T and log versus 1/T gave the values of enthalpy of reaction , entropy of reaction , and free energy of reaction , calculated for the first and second equilibrium steps (Table 2). A comparison of the later values (from ) with those obtained for the slow step of the reaction shows that they mainly refer to the rate-limiting step, supporting the fact that the reaction before the rate-determining step is fairly fast and involves low activation energy [20, 21]. The moderate values of were both favorable for the electron transfer processes. The values of , that is within the expected range for radical reactions, have been ascribed to the nature of electron pairing and unpairing processes and the loss of degrees of freedom formerly available to the reaction upon the formation of a rigid transition state . The negative values of indicate that the complex (C) is more ordered than the reactants [23, 24]. The enthalpy of activation and a relatively low value of entropy and a higher rate constant of the slow step indicate that the oxidation most probably occurs via inner-sphere mechanism [25, 26].
It is noteworthy that the oxidant species required the pH 12, below which the system gets disturbed and the reaction proceeds further to give a more reduced state of Mn, that is, Mn(IV) which slowly develops yellow turbidity. In this reaction, the role of pH is crucial. The rate constant of the slowest step and other equilibrium constants involved in the mechanism were evaluated and activation parameters were calculated. The proposed mechanism is consistent with product, mechanistic, and kinetic studies.
- H. A. Krebs, The Nobel Lecture, 1953.
- M. Akram, M. Altaf, and Kabir-ud-Din, “Oxidation of aspartic acid by water soluble colloidal MnO2 in absence and presence of ionic and nonionic surfactants,” Indian Journal of Chemistry A, vol. 46, no. 9, pp. 1427–1431, 2007.
- A. Arrizabalaga, F. J. Andrés-Ordax, M. Y. Fernández-Aránguiz, and R. Peche, “Kinetics and mechanism of the oxidation of L-α-amino-n-butyric acid by permanganate in acid medium,” International Journal of Chemical Kinetics, vol. 28, no. 11, pp. 799–805, 1996.
- A. Arrizabalaga, F. I. Andrés-Ordax, M. Y. Fernández-Aránguiz, and R. Peche, “Kinetic studies on the permanganic oxidation of amino acids. Effect of the length of amino acid carbon chain,” International Journal of Chemical Kinetics, vol. 29, no. 3, pp. 181–185, 1997.
- G. Ionita, V. E. Sahini, G. Semenescu, and P. Ionita, “Kinetics of oxidation of amino acids by some free stable hydrazyl radicals,” Acta Chimica Slovenica, vol. 47, no. 1, pp. 111–119, 2000.
- J. M. Pullar, M. C. M. Vissers, and C. C. Winterbourn, “Glutathione oxidation by hypochlorous acid in endothelial cells produces glutathione sulfonamide as a major product but not glutathione disulfide,” The Journal of Biological Chemistry, vol. 276, no. 25, pp. 22120–22125, 2001.
- Sonbol and H. Ridha, Journal of SaudiChemical Society, vol. 7, p. 187, 2003.
- N. Nalwaya, A. Jain, and B. L. Hiran, “Oxidation of some α-amino acids by pyridinium bromochromate in an aquo-acetic acid medium—a kinetic and mechanistic study,” Kinetics and Catalysis, vol. 45, no. 3, pp. 345–350, 2004.
- R. B. Chougale, R. G. Panari, and S. T. Nandibewoor, “Kinetics and mechanism of alkaline permanganate oxidation of L(+) aspartic acid,” Oxidation Communications, vol. 22, no. 2, pp. 298–307, 1999.
- A. I. Vogel, Text Book of Quantitative Chemical Analysis, Longman, Essex, UK, 5th edition, 1989.
- S. Bhattacharya and P. Banerjee, “Kinetic studies on the electron transfer between azide and nickel(IV) oxime imine complexes in aqueous solution,” Bulletin of the Chemical Society of Japan, vol. 69, no. 12, pp. 3475–3482, 1996.
- E. H. Cordes, “Kinetics of organic reaction in micelles,” Pure and Applied Chemistry, vol. 50, no. 7, pp. 617–625, 1978.
- E. H. Cordes and R. B. Dunlap, “Kinetics of organic reactions in micellar systems,” Accounts of Chemical Research, vol. 2, no. 11, pp. 329–337, 1969.
- J. H. Fendler and E. J. Fendler, Catalysis in Micellar and Macromolecular Systems, Academic Press, New York, NY, USA, 1975.
- P. L. Timmanagoudar, G. A. Hiremath, and S. T. Nandibewoor, “Permanganate oxidation of chromium(III) in aqueous alkaline medium: a kinetic study by the stopped-flow technique,” Transition Metal Chemistry, vol. 22, no. 2, pp. 193–196, 1997.
- R. G. Panari, R. B. Chougale, and S. T. Nandibewoor, “Oxidation of mandelic acid by alkaline potassium permanganate. A kinetic study,” Journal of Physical Organic Chemistry, vol. 11, no. 7, pp. 448–454, 1998.
- K. A. Thabaj, S. D. Kulkarni, S. A. Chimatadar, and S. T. Nandibewoor, “Oxidative transformation of ciprofloxacin by alkaline permanganate—a kinetic and mechanistic study,” Polyhedron, vol. 26, no. 17, pp. 4877–4885, 2007.
- R. Chang, Physical Chemistry with Applications to Biological Systems, MacMillan, New York, NY, USA, 1981.
- D. N. Sathyanarayana, Electronic Absorption Spectroscopy and Related Techniques, Universities Press, Andhra Pradesh, India, 2001.
- K. S. Rangappa, M. P. Raghavendra, D. S. Mahadevappa, and D. Channegowda, “Sodium N-chlorobenzenesulfonamide as a selective oxidant for hexosamines in alkaline medium: a kinetic and mechanistic study,” Journal of Organic Chemistry, vol. 63, no. 3, pp. 531–536, 1998.
- D. C. Bilehal, R. M. Kulkarni, and S. T. Nandibewoor, “Kinetics and mechanistic study of the ruthenium(III) catalyzed oxidative deamination and decarboxylation of L-valine by alkaline permanganate,” Canadian Journal of Chemistry, vol. 79, no. 12, pp. 1926–1933, 2001.
- C. Walling, Free Radicals in Solution, Academic Press, New York, NY, USA, 1957.
- K. S. Rangappa, N. Anitha, and N. M. M. Gowda, “Mechanistic investigations of the oxidation of substituted phenethyl alcohols by manganese(III) sulfate catalyzed by ruthenium(III) in acid solution,” Synthesis and Reactivity in Inorganic and Metal-Organic Chemistry, vol. 31, no. 8, pp. 1499–1518, 2001.
- Z. D. Bugarić, S. T. Nandibewoor, M. S. A. Hamza, F. Heinemann, and R. van Eldik, “Kinetics and mechanism of the reactions of Pd(II) complexes with azoles and diazines. Crystal structure of [Pd(bpma)(H2O)](ClO4)2·2H2O,” Dalton Transactions, no. 24, pp. 2984–2990, 2006.
- K. W. Hicks, “Kinetics of the permanganate ion-potassium octacyanotungstate(IV) reaction,” Journal of Inorganic and Nuclear Chemistry, vol. 38, no. 7, pp. 1381–1383, 1976.
- S. A. Farokhi and S. T. Nandibewoor, “Kinetic, mechanistic and spectral studies for the oxidation of sulfanilic acid by alkaline hexacyanoferrate(III),” Tetrahedron, vol. 59, no. 38, pp. 7595–7602, 2003.
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