1926
Kinetics and mechanistic study of the ruthenium(III) catalyzed oxidative deamination and decarboxylation of L-valine by alkaline permanganate Dinesh C. Bilehal, Raviraj M. Kulkarni, and Sharanappa T. Nandibewoor
Abstract: The kinetics of ruthenium(III) catalyzed oxidation of L-valine by permanganate in alkaline medium at a constant ionic strength has been studied spectrophotometrically. The reaction between permanganate and L-valine in alkaline medium exhibits 2:1 stoichiometry (KMnO4:L-valine). The reaction shows first-order dependence on the concentration of permanganate and ruthenium(III) and less than unit-order dependence on the concentrations of L-valine and alkali. The reaction rate increases both with an increase in ionic strength and a decrease in solvent polarity of the medium. Initial addition of reaction products did not significantly affect the rate. A mechanism involving the formation of a complex between catalyst and substrate has been proposed. The activation parameters were computed with respect to the slowest step of the mechanism. Key words: oxidation,
L-valine,
catalysis, ruthenium(III), kinetics.
Résumé : Faisant appel à des méthodes spectrophotométriques, on a étudié la cinétique de l’oxydation de la L-valine par le permanganate, catalysée par le ruthénium(III), en milieu alcalin et à force ionique constante. En milieu alcalin, la réaction entre le permanganate et la L-valine présente une stoechiométrie de 2:1 (KMnO4:L-valine). La réaction présente une dépendance du premier ordre avec les concentrations de permanganate et de ruthénium(III) et cette dépendance est inférieure à l’unité pour les concentrations de L-valine et de base. La vitesse de réaction augmente avec une augmentation de la force ionique et avec une diminution dans la polarité du solvant. L’addition initiale de produits réactionnels n’affecte pas la vitesse de façon significative. On propose un mécanisme impliquant la formation initiale d’un complexe entre le catalyseur et le substrat. On a calculé les paramètres d’activation par rapport à l’étape la plus lente du mécanisme. Mots clés : oxydation,
L-valine,
[Traduit par la Rédaction]
catalyse, ruthénium(III), cinétique.
Bilehal et al.
1933
Introduction Potassium permanganate is widely used as an oxidizing agent in synthetic as well as in analytical chemistry and also as a disinfectant. The reactions with permanganate are governed by pH of the medium. Among the six oxidation states of manganese from 2+ to 7+, permanganate Mn(VII) is the most potent oxidation state in both acid and alkaline medium. The manganese chemistry involved in these multistep redox reactions is an important source of information as the manganese intermediates are relatively easy to identify when they have sufficiently long life times and the oxidation states of the intermediates permit useful conclusions as to the possible reaction mechanism including the nature of intermediates.
Received February 2001. Published on the NRC Research Press Web site at http://canjchem.nrc.ca on December 14, 2001. D.C. Bilehal, R.M. Kulkarni, and S.T. Nandibewoor.1 P.G. Department of Studies in Chemistry, Karnatak University, Dharwad 580 003, India. 1
Corresponding author (fax: 0836-747-884; e-mail:
[email protected]).
Can. J. Chem. 79: 1926–1933 (2001)
The oxidation by permanganate ion finds extensive applications in organic syntheses (1–7), especially since the advent of phase transfer catalysis (3, 4, 6), which permits the use of solvents such as methylene chloride and benzene. Kinetic studies are important sources of mechanistic information on the reactions, as demonstrated by the results referring to unsaturated acids in both aqueous (1, 3, 7) and nonaqueous media (8). During the oxidation by permanganate, it is evident that permanganate is reduced to various oxidation states in acidic, alkaline, and neutral media. Furthermore, the mechanism by which the multivalent oxidant oxidizes a substrate depends not only on the substrate but also on the medium (9) used for the study. In strongly alkaline medium, the stable reduction product (10, 11) of permanganate ion is manganate ion, MnO24 - . No mechanistic information is available to distinguish between a direct one-electron reduction to Mn(VI) (Scheme 1) and a mechanism in which a hypomanganate is formed in a two-electron reduction followed by a rapid oxidation of the hypomanganate ion (12) (Scheme 2). L-valine is one of the essential basic amino acids, which is essential in the nutrition of mammals. Its role is crucial in
DOI: 10.1139/cjc-79-12-1926
© 2001 NRC Canada
Bilehal et al.
1927
Scheme 1.
k 1¢
Mn(VII) + S
® Mn(VI) + S×
Fig. 1. Spectral changes during the ruthenium (III) catalyzed oxidation of L-valine by alkaline permanganate (scanning time interval = 1 min). [MnO-4 ] = 2.0 × 10–4, [L-valine] = 2.0 × 10–3, [Ru(III)] = 7.5 × 10–8, [OH–] = 0.30, I = 0.50 mol dm–3.
k 2¢
Mn(VII) + S×® Mn(VI) + Products Where, S = substrate; k2¢ >> k1¢ Scheme 2.
k 3¢
Mn(VII) + S
® Mn(V) + Products k 4¢
Mn(VII) + Mn(V)
® 2 Mn(VI)
Where, S = substrate; k4¢ >> k3¢ the development of organs, especially in children. It also finds applications in medicine and pharmaceuticals. Amino acids have been oxidized by a variety of oxidizing agents (13). Although several types of organic (14) and inorganic (10, 11) substrates are oxidized by permanganate in aqueous alkaline medium, there are only a few reports on the oxidation of amino acids by aqueous alkaline permanganate. The kinetic investigation of the oxidation of amino acids has been carried out under different experimental conditions (15). In many cases it was reported that amino acids undergo oxidative decarboxylation. Jayaprakash Rao et al. (16) suggested that the oxidation of amino acids by two-electron oxidants such as diperiodatoargentate(III) in alkaline medium involves a two-electron transfer, which produces an imino acid intermediate in the rate determining step. This intermediate subsequently undergoes hydrolysis to yield the keto acids. One electron oxidants such as ceric sulphate (17), peroxomonosulphate (18), and hexacyanoferrate(III) (19) also oxidize amino acids to keto acids. But other studies with amino acids report the oxidation products as the corresponding aldehydes (20). Thus, the study of amino acids becomes important because of their biological significance and their selectivity towards the oxidants. Ruthenium(III) is known to be an efficient catalyst in several redox reactions particularly in alkaline medium (21). The mechanism of catalysis can be quite complicated due to the formation of different intermediate complexes, free radicals, and different oxidation states of ruthenium. The kinetics of fast reactions between ruthenate(VII), RuO-4 , and manganate(VI), i.e., MnO24 - , have been studied (22) and the reaction is presumed to proceed via an outer-sphere mechanism. The rapid exchange between MnO24 - and MnO-4 has been studied in detail by a variety of techniques (23). The uncatalyzed reaction between L-valine and permanganate in alkaline medium has been studied previously (24). A micro amount of ruthenium(III) is sufficient to catalyze the reaction in alkaline medium and a variety of mechanisms are possible. Thus, to explore the mechanism of oxidation by permanganate ion in strongly aqueous alkaline medium and to check the selectivity of L-valine towards permanganate in catalyzed system, we have selected ruthenium(III) as a cata-
lyst. The present study deals with the title reaction to investigate the redox chemistry of permanganate and L-valine in presence of a catalysis.
Experimental Materials and methods Stock solution of L-valine (Sisco Research Labs, India) was prepared by dissolving the appropriate amount of sample in doubly distilled water. The solution of potassium permanganate (BDH) was prepared and standardized against oxalic acid (25). Potassium manganate solution was prepared as described by Carrington and Symons (26).The solution was standardized by measuring the absorbance on a Hitachi 15020 spectrophotometer with a 1 cm quartz cell at 608 nm (e = 1530 ± 20 dm3 mol–1 cm–1). The ruthenium(III) solution was prepared by dissolving a known weight of RuCl3 (S.D. Fine Chemicals) in 0.20 mol dm–3 HCl. Mercury was added to the ruthenium(III) solution to reduce any Ru(IV) formed during the preparation of ruthenium(III) stock solution and kept for a © 2001 NRC Canada
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Can. J. Chem. Vol. 79, 2001 H
[1]
CH3
C H 3C
H CH-COOH + 2MnO 4– + 2OH–
Ru(III)
NH2
CH3
C
2–
CHO + 2MnO 4 +
H3C NH3 + CO2 + H2O
day. The ruthenium(III) concentration was assayed (27) by EDTA titration. All other reagents were of analytical grade and their solutions were prepared by dissolving the requisite amounts of the samples in doubly distilled water. NaOH and NaClO4 were used to provide the required alkalinity and to maintain the ionic strength, respectively. Kinetic procedure All kinetic measurements were performed under pseudofirst-order conditions with [L-valine]:[MnO-4 ] ³ 10:1 at a constant ionic strength of 0.50 mol dm–3. The reaction was initiated by mixing previously thermostated solutions of MnO-4 , and L-valine, which also contained the necessary quantities of Ru(III), NaOH, and NaClO4 to maintain the required alkalinity and ionic strength, respectively. The temperature was uniformly maintained at 26 ± 0.1°C. The course of reaction was followed by monitoring the decrease in the absorbance of MnO-4 in a 1 cm quartz cell of a Hitachi model 150-20 spectrophotometer at its absorption maximum of 526 nm as a function of time. The application of Beer’s law to permanganate at 526 nm had been verified, giving e = 2083 ± 50 dm3 mol–1 cm–1 (lit. (10) e = 2200 dm3 mol–1 cm–1). The first-order rate constants (kobs) were evaluated by plots of log [MnO-4 ] vs. time. The first-order plots in almost all cases were linear to 80% completion of the reaction and kobs were reproducible within ±5%. During the course of measurements, the solution changed from violet to blue and then to green. The spectrum of the green solution was identical to that of MnO24 - . It is probable that the blue colour originated from the violet of permanganate and the green from the manganate, excluding the accumulation of hypomanganate. It is also evident from the Fig. 1 that the absorbance of permanganate decreases at 526 nm, whereas the absorbance of manganate increases at 608 nm. The effect of dissolved oxygen on the rate of reaction was checked by preparing the reaction mixture and following the reaction in an atmosphere of nitrogen. No significant difference between the results obtained under the nitrogen and in the presence of air was observed. In view of the ubiquitous contamination of basic solutions by carbonate, the effect of carbonate on the reaction was also studied. Added carbonate had no effect on the reaction rate. However, fresh solutions were used when conducting the experiments.
Results Stoichiometry and product analysis The reaction mixtures containing an excess permanganate concentration over L-valine and constant concentration of Ru(III) (0.30 mol dm–3), NaOH, and an adjusted ionic strength of 0.50 mol dm–3 was allowed to react for 2 h at
26 ± 0.1°C. After completion of the reaction, solid KI was added followed by acidification by 10% H2SO4. The remaining MnO-4 was then titrated against standard sodium thiosulphate (28). The results indicated that two mol of MnO-4 are consumed by one mol of L-valine as given by eq. [1]. The main reaction products were identified as aldehyde by spot test (29), ammonia by Nessler’s reagent (30), and manganate by its visible spectrum. CO2 was qualitatively detected by bubbling N2 gas through the acidified reaction mixture and passing the liberated gas through a tube containing lime water (31). The quantitative estimation of aldehyde by 2,4-DNP derivative (32), yielded nearly 74%. The nature of the aldehyde, was confirmed by its IR spectrum (33), which showed a carbonyl stretching at 1729 cm–1 and a band at 2928 cm–1 due to aldehydic C—H stretching, thus confirming the presence of isobutylaldehyde. It was further observed that the aldehyde does not undergo further oxidation under the present kinetic conditions. The test for corresponding acid was negative. Reaction orders The reaction orders were determined from the slopes of log kobs vs. log (concentration) plots by varying the concentration of reductant, catalyst, and alkali in each, while keeping others constant. The oxidant, potassium permanagate, concentration was varied in the range of 5.0 × 10–5 to 5.0 × 10–4 mol dm–3, and the fairly constant kobs values indicate that the order with respect to [MnO-4 ] was one. This was also confirmed by varying the concentration of MnO-4 , which did not show any change in pseudo-first-order constants (kobs) values as shown in Table 1. The substrate L-valine was varied in the range of 5.0 × 10–4 to 5.0 × 10–3 mol dm–3 at 26°C; the order with respect to [L-valine] was found to be less than unity. The catalyst Ru(III) was varied in the range of 1.0 × 10–8 to 1.0 × 10–7 mol dm–3 at 26°C; the order with respect to [Ru(III)] was found to be unity. The effect of alkali on the reaction was studied at constant concentrations of L-valine, Ru(III), and potassium permanganate and at a constant ionic strength of 0.50 mol dm–3 at 26°C. The rate constant increased with the increase in the concentration of alkali, indicating an apparent less than unit order dependence on [alkali] as given in Table 1. Effect of initially added products The externally added products such as manganate, ammonium hydroxide, and aldehyde did not show any significant effect on the rate of the reaction. Effect of ionic strength and dielectric constant The effect of ionic strength was studied by varying the sodium perchlorate concentration from 0.3 to 2.0 mol dm–3 at constant concentrations of permanganate, L-valine, and © 2001 NRC Canada
Bilehal et al.
1929 Table 1. Effects of [L-valine], [MnO-4 ], [Ru(III)], and [OH–] on ruthenium(III) catalyzed oxidation of –3 L-valine by KMnO4 at 26°C; I = 0.50 mol dm and error ±4%. kobs × 10 [L-valine] × 103 (mol dm–3)
[OH–] (mol dm–3)
[Ru(III)] × 108 (mol dm–3)
Expt.
Calcd.
0.5 1.0 2.0 3.0 5.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0
2.0 2.0 2.0 2.0 2.0 0.5 1.0 2.0 3.0 5.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0
0.3 0.3 0.3 0.3 0.3 0.3 0.3 0.3 0.3 0.3 0.05 0.1 0.2 0.3 0.5 0.3 0.3 0.3 0.3 0.3
7.5 7.5 7.5 7.5 7.5 7.5 7.5 7.5 7.5 7.5 7.5 7.5 7.5 7.5 7.5 1.0 2.5 5.0 7.5 10.0
3.60 3.70 3.75 3.70 3.60 2.45 3.03 3.75 4.07 4.16 1.05 1.76 2.80 3.75 4.68 0.51 1.31 2.60 3.75 5.06
3.67 3.67 3.67 3.67 3.67 2.38 3.11 3.67 3.91 4.12 1.03 1.81 2.92 3.47 4.61 0.49 1.22 2.45 3.67 4.90
I
1.25
1.00
Table 2. Thermodynamic activation parameters for the ruthenium(III) catalyzed oxidation of L-valine by aqueous alkaline permanganate with respect to slow step of Scheme 3.
1/2
0.75
0.50
0.25
0 0.3
Activation parameters –1
0.5
0.91 0.7 0.89
3 +log kobs
3+ log kobs
0.93
0.9
0.87
0.85 1.30
(s–1)
[MnO-4 ] × 104 (mol dm–3)
Fig. 2. Plot of log kobs vs. I1/2 and log kobs vs. 1/D. 1.50 0.95
3
1.35
1.40
1.45
1.50
1.55
1.1 1.60
1 / D x 10-2
alkali. It was found that the rate constant increased with increase in concentration of NaClO4 and the plot of log kobs vs. I1/2 was linear (Fig. 2). The effect of dielectric constant (D) was studied by varying the tert-butanol–water content in the reaction mixture with all other conditions being constant. Attempts to measure the relative permittivities were not successful. However, they were computed from the values of pure liquids (34). The solvent did not react with the oxidant under the experimental conditions. The rate constants (kobs) increased with a
Ea (kJ mol ) DH‡ (kJ mol–1) DS‡ (J K–1 mol–1) DG‡ (kJ mol–1)
Values 72 69 75 46
(±4) (±4) (±4) (±3)
decrease in the dielectric constant of the medium. The plot of log kobs vs. 1/D was linear with positive slope as shown in Fig. 2. Test for free radicals The reaction mixture was mixed with acrylonitrile monomer and kept for 2 h in an inert atmosphere. On diluting with methanol a white precipitate was formed, indicating the intervention of free radicals in the reaction. Effect of temperature The rate of the reaction was measured at four different temperatures with varying the concentration of OH–, keeping other conditions constant. The rate was found to increase with increase in temperature. The rate constant (k) of the slow step of Scheme 3 were obtained from the intercept of the plots of [Ru(III)]/kobs vs. 1/[OH–] for different temperatures. The energy of activation corresponding to these constants were evaluated from the plot of log k vs. 1/T and other activation parameters for reaction were calculated and are given in Table 2. © 2001 NRC Canada
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Can. J. Chem. Vol. 79, 2001
Scheme 3. O
K1
MnO4– + OH –
Mn
O
O
O
R
K2
[Ru(H2O)5OH]2+
CH-COO- +
2–
OH
Complex ( C ) + H2O
NH2
O (C)+
O
2–
OH
k
Mn
R- CH. + MnO42- + [Ru(H2O)4OH]2+ + HCO3–
Slow O
O
O R- CH. +
NH2
OH
2– fast
Mn
O
O
NH2
R- CHO + MnO42- + NH3
O
H where R = C--CH 3 CH3
The probable structure of complex ( C ) is
H2O H2O H2O
OH Ru
+ OC(O)CH(NH2)R
OH2
Discussion Permanganate ion (MnO-4 ) is a powerful oxidant in an aqueous alkaline medium, and because it exhibits many oxidation states, the stoichiometric results and pH of the reaction media play an important role. Under the prevailing experimental conditions at pH > 12, the reduction product of Mn(VII) is stable and further reduction of Mn(VI) might be stopped (10, 11). Diode Array Rapid Scan Spectrophotometric (DARSS) studies have shown that at pH > 12, the product of Mn(VII) is Mn(VI) and no further reduction was observed as reported by Simandi et al. (10). However, on prolonged standing, the green Mn(VI) is reduced to Mn(IV) under our experimental conditions. The permanganate in alkaline medium exhibits various oxidation states, such as Mn(VII), Mn(V), and Mn(VI). The
colour of the solution changed from violet to blue and further to green excluding the accumulation of hypomanganate. The violet colour originates from pink of permanganate and blue from hypomanganate is observed during the course of the reaction. The colour change of KMnO4 solution from the violet Mn(VII) ion to the dark green Mn(VI) ion through the blue Mn(IV) ion has been observed. It is interesting to identify the probable species of ruthenium(III) chloride in alkaline medium. Electronic spectral studies (21) have confirmed that ruthenium(III) chloride exists in hydrated form as [Ru(H2O)6]3+. In the present study it is quite probable that the species [Ru(H2O)5OH]2+ might assume the general form [Ru(III)(OH)x]3–x. The value of x would always be less than six because there are no definite reports of any hexahydroxy species of ruthenium. The remainder of the coordination sphere will be filled by water mol© 2001 NRC Canada
Bilehal et al. Fig. 3. Plot of [Ru(III)]/ kobs vs. 1/[L-valine] and [Ru(III)]/kobs vs. 1/[OH–] (conditions as in Table 1).
ecules. Hence under the experimental conditions [OH–] >> [RuIII], RuIII is mostly present as the hydroxylated species [Ru(H2O)5OH]2+. The reaction between permanganate and L-valine in alkaline medium has a stoichiometry of 1:2 with a first-order dependence on the concentration of MnO-4 and Ru(III) and less than unit order dependence on both theconcentration of alkali and L-valine. No effect of added products such as aldehyde and ammonia was observed. It is known that L-valine exists as deprotonated form in basic medium (24). The results suggest that first alkali combines with permanganate to form an alkali–permanganate species [MnO4·OH]2– in a preequilibrium step (35). L-valine in the deprotonated form reacts with ruthenium(III) species to form a complex (C). Complex C reacts with the alkali–permanganate species in a slow step to form a free radical derived from L-valine, which further reacts with another permanganate species in a fast step to yield the products. The experimental results can be accommodated as seen in Scheme 3. Spectroscopic evidence for complex formation was not successful, since there is no change in the lmax (except some hyperchromicity is observed) for the mixture of L-valine and ruthenium(III) compared with that of L-valine or ruthenium(III) itself. This might be due to the involvement of weak interactions. However, the evidence for complex formation is obtained by kinetic studies (i.e., from the Michaelis–Menten plot). The plot of [Ru(III)]/kobs vs. 1/[L-valine] is linear with an intercept supporting the Ru(III)–L-valine complex. Such type of substrate–catalyst complex formation has been reported previously (36). The observed modest enthalpy of activation, relatively low value of the entropy of activation and higher rate constant for the slow step of the mechanism, indicates that oxidation presumably occurs by an inner-sphere mechanism. This conclusion is supported by earlier work (37). Since Scheme 3 is in accordance with generally well-accepted principle of non-complementary oxidations taking place in a sequence of one-electron steps, the reaction would involve a radical intermediate. Since per-
1931
manganate is a one-electron oxidant in alkaline medium, the reaction between substrate and oxidant would give rise to a radical intermediate. A free radical scavenging experiment revealed such a possibility. This type of radical intervention in the oxidation of amino acids has also been observed earlier (20). The thermodynamic quantities for the first equilibrium step in Scheme 3 and activation parameters for the limiting step in Scheme 3 can be evaluated as follows: the hydroxyl ion concentration as in Table 1 was varied at four different temperatures and the K1 value was determined at each temperature. The values of K1 (dm3 mol–1) were obtained as 3.2, 5.0, 6.8, and 8.1 at 26, 30, 35, and 40°C, respectively. A Vant Hoff’s plot was made for the variation of K1 with temperature (i.e., log K1 vs. 1/T) and the values of the enthalpy of the reaction (DH), entropy of the reaction (DS), and free energy of reaction (DG), were calculated as 47.9 kJ mol–1, –74.3 J K–1 mol–1, and 70.6 kJ mol–1, respectively. Similarly the L-valine concentration as in Table 1 was varied at four different temperatures and the K2 value was determined at each temperature. The values of K2 (dm3 mol–1) were obtained as 2.2 × 103, 3.5 × 103, 4.7 × 103, and 6.4 × 103 at 26, 30, 35, and 40°C, respectively. A plot was made for the variation of K2 with temperature (i.e., log K2 vs. 1/T) and the values of DH, DS, and DG, were calculated as 51.9 kJ mol–1, –6.1 J K–1 mol–1, and 53.8 kJ mol–1, respectively. A comparison of the latter values with those obtained for the slow step of the reaction shows that these values mainly refer to the rate limiting step, supporting the fact that the reaction before the rate determining step are fairly rapid and involves only little activation energy. Scheme 3 leads to the following rate law, and can be derived as follows: [2]
Rate = k obs = [MnO-4 ] k K1K2 [L -val][ Ru(III)][OH- ] 1 + K1[OH- ] + K2 [L -val] + K1K2 [OH- ][L -val]
The terms (1 + K1K2[OH –][Ru(III)][MnO-4 ]), (1 + K1[MnO-4 ]), and (1 + K2[Ru(III)]) also should be in the denominator of eq. [2], but in view of low concentration of MnO-4 and ruthenium(III) used, they approximate to unity. Thus the above eq. [2] can be rearranged to the following form, which is used for the verification of the rate law. [3]
[Ru(III)] 1 = kobs k K1K2 [L -val][OH- ] +
1 1 1 + + k K2 [L -val] k K1[OH ] k
According to eq. [3], the plots of [Ru(III)]/kobs vs. 1/[L-val] and [Ru(III)]/kobs vs. 1/[OH–] should be linear, which is verified in Fig. 3. The slopes and intercepts of such plots lead to the values of k, K1, and K2 at 26°C of 12.2 (±0.6) × 104 dm3 mol–1 s–1, 3.2 (±0.1) dm3 mol–1, and 2.2 (±0.1) × 103 dm3 mol–1, respectively. Using these values, the rate constants under different experimental conditions were calculated by eq. [2] and compared with experimental data. There is a good agreement between them, which supports the © 2001 NRC Canada
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Can. J. Chem. Vol. 79, 2001
Scheme 3. The value of K1 is in good agreement with earlier work (38). The effect of ionic strength on the rate can be understood essentially on the basis of ionic species as in Scheme 3. The effect of solvent on the reaction kinetics has been described detail in the literature (39). In the present study the rate determining step involves the reaction between two ions and so eq. [4] is applicable: [4]
ln k = ln k¥ – ZAZB e2/kBTrABD
where k¥ is the rate constant in a medium of infinite dielectric constant, rAB is the sum of the ionic radii, ZA and ZB are the charges on the two ions, and D is the dielectric constant of the medium. The observed linear plot of log kobs vs. 1/D with positive slope (Fig. 2) is in accordance with eq. [4] as ZA and ZB have opposite charges (Scheme 3). The values of DH‡ and DS‡ were both favourable for electron transfer process. The value of DS‡, within the range for radical reaction, has been ascribed to the nature of electron pairing and unpairing processes and to the loss of degree of freedom formerly available to the reactants upon the formation of a rigid transition state (40).
Conclusions It is interesting that the oxidant species [MnO-4 ] requires a pH > 12, below which the system becomes disturbed and the reaction will proceed further to give a reduced product of the oxidant as Mn(IV), which slowly develops yellow turbidity. Hence, it becomes apparent that in carrying out this reaction the role of pH in a reaction medium is crucial. It is also noteworthy that under the conditions studied the reaction occurs in two successive one-electron reductions (Scheme 3) rather than two-electron in a single step (Scheme 2). A micro amount of Ru(III) is sufficient to catalyze the title reaction. It is interesting to note that uncatalyzed reaction follows a two path mechanism, while catalyzed reaction follows a single path. The description of the mechanism is consistent with all the experimental evidence including both kinetic and product studies.
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