Kinetics andMechanism of Electron Transfer to Heptavalent Manganese by-Aspartic Acid in Alkaline Aqueous and Micellar Media

e kinetics and mechanism of the electron transfer of -Aspartic acid (Asp) by Mn (VII) in alkaline medium has been studied spectrophotometrically over the range 2.0 ≤ 10[Asp] ≤ 5.0mol dm; 0.01 ≤ [OH] ≤ 0.05mol dm; 298 ≤ TT ≤ 318K and II I 0.05mol dm (KNO3). e reaction exhibits �rst-order dependence in [MnO4 ]TT but shows fractional-order dependence in both [Asp]TT and [OH ]TT.e 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.e products were characterized by spectral analysis. A mechanism involving free radicals is proposed. Asp binds MnO4 − to form a complex that subsequently decomposes to products. Activation parameters ΔHH (kJmol) and ΔSS (JKmol) for the reaction are 5.62 ± 0.35 and −227.65 ± 1.1, respectively. e negative value of ΔSS indicates that oxidation occurs via inner sphere mechanism.


Introduction
Study of oxidation of amino acids has received considerable attention due to the importance of degradation of these compounds in biological systems.e Aspartic acid is a nonessential amino acid which is found abundantly in plant proteins especially in sprouting seeds.is 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 [1].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][3][4][5][6][7][8][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.

Experimental
All chemicals used were of reagent grade.Double distilled water was used throughout the work.Stock solution of -Aspartic acid (SRL Chemicals) was prepared by dissolving the appropriate amount of the sample in double distilled water.e stock solution of KMnO 4 (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 [10].NaOH and KNO 3 were used to maintain the required alkalinity and ionic strength, respectively.e 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  pH system-361 equipped with a combination of glass Ag/AgCl/Cl − (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 UVvisible spectrophotometer equipped with a CE-2024 thermoelectric controller.Ten-millimeter quartz Suprasil cuvettes were used.IR spectra were taken in Varian FTIR spectrophotometer (USA).

Kinetic Measurements
e oxidation of -Aspartic acid by KMnO 4 was followed under pseudo-�rst-order conditions where -Aspartic acid was taken excess over MnO 4 − at 25 ∘ C ± 0.1 ∘ C. e reaction was initiated by mixing the required quantities of previously thermostated solutions of MnO 4 − and -Aspartic acid, which also contained de�nite quantities of NaOH and KNO 3 to maintain required ionic strength.e progress of the reaction was followed by measuring the decrease in absorbance at 525 nm with time using a conventional mixing technique. ∞ was measured aer the completion of the reaction (approximately aer 24 hours of mixing) when the absorbance became almost constant.e 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.e correlation coefficients ( 2 ) of the plots used to determine  obs were found to be 0.99.e pseudo-�rst-order rate constant ( obs ) was calculated by the least squares method from the above relationship.e redox reactions were followed for about 3 half lives.e reported rate data represented as an average duplicate runs were reproducible to within ±3%.

Results
e electron transfer reaction between -Aspartic acid and alkaline Mn(VII) has been studied over the range 2.0 ≤ 10 3 [Asp] ≤ 5.0 mol dm −3 ; 0.01 ≤ [OH − ] ≤ 0.05 mol dm −3 ; 298 ≤ T ≤ 318 K and  = 00 mol dm −3 (KNO 3 ).e reaction orders were determined using the slopes of log obs versus log[MnO 4 − ] plots by varying the concentration of the reductant and OH − while keeping other factors constant.�ith �xed concentrations of -Aspartic acid 10 × 10 −2 mol dm −3 , and alkali, 0 × 10 −2 mol dm −3 at constant ionic strength, 0.05 mol dm −3 , the permanganate concentration varied from 10 × 10 −4 mol dm −3 to 30 × 10 −4 mol dm −3 .e linearity of plots of log(absorbance) versus time, for different concentrations of permanganate, indicates that the order in [Mn(VII)] is unity (Figure 1).e -Aspartic acid concentration was varied in the range of 20 × 10 −3 to 0 × 10 −3 mol dm −3 at constant alkali and permanganate concentrations and constant ionic strength of 0.05 mol dm −3 at 298 K. e  obs values increased with an increase in -Aspartic acid over the concentration range shown in Figure 3.At low concentration of -Aspartic acid, the reaction was of �rst order and at high concentration of -Aspartic acid, the reaction was independent of -Aspartic acid.

Effect of Alkali Concentration
. e effect of alkali concentration on the reaction was studied at constant ionic strength of 0.05 mol dm −3 at 25 ∘ C. e [OH − ] was varied in the range of 0.01 to 0.05 mol dm −3 .e rate constant increased with an increase in alkali concentration (Figure 4), indicating a fractional-order dependence of the rate on alkali concentration.

Effect of Ionic
Strength.e effect of ionic strength was studied by varying the potassium nitrate concentration from 0.05 to 0.5 mol dm −3 at constant concentration of permanganate, -Aspartic acid, and alkali.Increasing ionic strength had no effect on the rate constant.

4.3.
Effect of Temperature.e kinetics was also studied at �ve different temperatures with varying concentrations of -Aspartic acid and alkali, keeping other conditions constant.e rate constants were found to increase with the increase in temperature.e rate of the slow step was obtained from the slopes and intercepts of 1/ obs versus 1/[-Aspartic acid] and 1/ obs versus 1/[OH − ] plots at �ve different temperatures (298-318 K).

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 [11].

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].e 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 [14].During the study, the concentration of sodium dodecyl sulfate (SDS) varied keeping the concentrations of -Aspartic acid, MnO 4 − , 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.e equilibrium constant values for the formation of complex (C) with permanganate for -Aspartic acid and cysteine are comparable indicating equal probability of formation of the complex.

Stoichiometry and Product Analysis. e reaction between -Aspartic acid and MnO 4
− in alkaline medium has a stoichiometry of 1 : 4. e 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.e FT-IR spectra of -Aspartic acid and its product complex are similar.e C=N stretching band appears at 1619 cm −1 which is absent in -Aspartic acid.e other signi�cant bands appear at 1411 cm −1 (for symmetric carboxylate stretching) 1520 cm −1 (N-H bending).e other reaction products are identi�ed as ammonia (Nessler�s reagent test), CO 2 (lime water test), and manganate (MnO 4 2− ).

Discussion
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 [15].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 MnO 4 − 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 aer keeping for a long time the solution decolorizes and forms a brown precipitate.is 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 [MnO 4 ⋅ OH] 2− [16,17].In the second step, [MnO 4 ⋅ OH] 2− combines with -Aspartic to form an intermediate complex.e variable order with respect to -Aspartic is most probably due to the complex formation between oxidant and -Aspartic prior to the slow step.A plot of 1/ obs 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).e cleavage of this complex (C) is assigned as the slowest step, leading to the formation of an -Aspartic radical intermediate and Mn(VI).e radical intermediate reacts with another Mn(VII) species, [MnO 4 ⋅ OH] 2− , to give the �nal products: Mn(VI), 3-oxopropanic acid and NH 3 (Scheme 1).e effect of the ionic strength and dielectric constant on the rate is consistent with the involvement of a neutral molecule in the reaction.e suggested structure of complex (C) is given in Scheme 1.
From Scheme 1, the rate law can be derived as follows: e total [MnO 4 − ] can be written as [OH − ] (mol dm −3 ) 10 where "" and "" stand for "total" and "free", respectively.Similarly, total [OH − ] can be calculated as In view of the low concentration of MnO 4 − and -Aspartic acid used in the experiment, in (3), the terms  1 [MnO  [SDS] (mol dm −3 ) 10  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 veri�ed by rearranging to (9) According to (9), other conditions being constant, plots of 1/ obs versus 1/[Asp] and 1/ obs versus 1/[OH − ] should be linear (Figures 3 and 4).e slopes and intercepts of such plots lead to values of  1 ,  2 , and  (Table 2).With these values, the rate constants were calculated under different experimental conditions.e thermodynamic quantities for the �rst and second equilibrium steps of Scheme 1 can be evaluated.e [-Aspartic acid] and [OH − ] (Table 1) were varied at �ve different temperatures.van�t Hoff �s plots of log 1 versus 1/T and log  2 versus 1/T gave the values of enthalpy of reaction Δ ∘ , entropy of reaction Δ ∘ , and free energy of reaction Δ ∘ , calculated for the �rst and second equilibrium steps (Table 2).A comparison of the later values (from  2 ) 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 ratedetermining step is fairly fast and involves low activation energy [20,21].e moderate values of Δ ∘ were both favorable for the electron transfer processes.e values of Δ ∘ , 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 [22].e negative values of Δ ∘ indicate that the complex (C) is more ordered than the reactants [23,24].e 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].

Conclusion
It is noteworthy that the oxidant species MnO 4 − 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.e rate constant of the slowest step and other equilibrium constants involved in the mechanism were evaluated and activation parameters were calculated.e proposed mechanism is consistent with product, mechanistic, and kinetic studies.
2[-Aspartic acid] (mol dm −3 ) 10 3  obs (s −1 ) Activation and thermodynamic parameters for the oxidation of -Aspartic acid by alkaline KMnO 4 with respect to the slow step of the reaction (Scheme 1).