Metal Ions on Oxidative Conversion of Folic Acid - American Chemical

Jan 12, 2010 - sulfonamide (bromamine-T; BAT) was studied in alkaline medium. The detailed kinetic and mechanistic investigations of all four catalyze...
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Ind. Eng. Chem. Res. 2010, 49, 1550–1560

Catalysis and Mechanistic Studies of Ru(III), Os(VIII), Pd(II), and Pt(IV) Metal Ions on Oxidative Conversion of Folic Acid C. H. Vinod Kumar,† R. V. Jagadeesh,*,‡ K. N. Shivananda,§ Y. Sree Sandhya,‡ and C. Naga Raju*,† Department of Chemistry, Sri Venkateswara UniVersity, Tirupati-517502, India, DiVision of Organic Chemistry, School of AdVanced Sciences, VIT UniVersity, Vellore-632014, India, and Schulich Faculty of Chemistry, Technion - Israel Institute of Technology, Technion City, Haifa-32000, Israel

The homogeneous catalysis of Ru(III), Os(VIII), Pd(II), and Pt(IV) in the oxidative conversion of folic acid (FA) to pterin-6-carboxylic acid, p-aminobenzoic acid, and glutamic acid by sodium N-bromo-p-toluenesulfonamide (bromamine-T; BAT) was studied in alkaline medium. The detailed kinetic and mechanistic investigations of all four catalyzed oxidation reactions were studied at 313 K. The stoichiometries and oxidation products of the four catalyzed reactions were found to be the same, but their kinetic patterns and oxidation mechanisms were different. The reactions were studied at different temperatures, and the activation parameters were evaluated for each catalyzed reaction. Under identical sets of experimental conditions, the kinetics of all four catalyzed reactions were compared with those of uncatalyzed reactions, revealing that the catalyzed reactions are 6-22 times faster. The catalytic efficiency of these catalysts follows the order Os(VIII) > Ru(III) > Pt(IV) > Pd(II). This trend can be attributed to the different d electronic configuration of the catalysts. Based on the observed experimental results, a detailed mechanistic interpretation and the related kinetic model were derived for each catalyst. 1. Introduction Oxidation reactions are important in the synthesis of organic compounds, biomolecules, and pharmaceuticals because these reactions create new functional groups or modify existing functional groups in a molecule.1,2 Several methods are available for the oxidation of organic molecules using different oxidants ranging from metal oxidants to atmospheric O2. However, there is still a need to develop environmentally friendly methodologies and introduce safe, cost-effective, and stable reagents for the oxidation of organic molecules. The development of new processes for selective oxidations with environmentally friendly oxidants has potential practical applications in organic synthesis. In this regard, we used sodium N-bromo-p-toluene sulfonamide or bromamine-T (p-CH3C6H4SO2NBrNa.3H2O; abbreviated as BAT), a member of the organic haloamine family, as an oxidant in NaOH medium for the selective oxidative conversion of folic acid into pterin-6-carboxylic acid, p-aminobenzoic acid, and glutamic acid using platinum-group metal ions as catalysts. The use of platinum-group metal ions as catalysts in redox reactions has become vital in recent years. The uses of such platinum-group metal catalysts reveal mechanistic details of redox reactions, providing great advantages in the interpretation of reaction. Hence, studies on the use of platinum-group metal ions either alone or in binary mixtures as catalysts in many redox reactions have been gaining interest.3-5 Because of the catalytic applications of metal ions and our interest in exploring the detailed mechanistic picture of metal ion-FA reactions, we performed studies on the oxidation kinetics of FA using Ru(III), Os(VIII), Pd(II), and Pt(IV) catalysts. Bromamine-T, the bromine analogue of chloramine-T (CAT), is gaining importance as a mild oxidant and has been found to * To whom correspondence should be addressed. E-mail: rvjdeesh@ yahoo.com (R.V.J.), [email protected] (C.N.R.). † Sri Venkateswara University. ‡ VIT University. § Technion - Israel Institute of Technology.

be a better oxidizing agent than the chloro compound. Although the mechanistic aspects of many of haloamine reactions have been well-documented,6-13 similar studies on bromine analogues are sparse. In view of these facts, there is a considerable scope for studies with BAT to gain better insight into the speciation of BAT reaction models and to understand BAT redox chemistry in solutions. Folic acid (FA) is widely distributed in nature and is present in many animal and plant tissues and microorganisms.14,15 Pterin-6-carboxylic acid and its derivatives have been mainly used as fluorescent agents. A derivative of pterin-6-carboxylic acid, 7,8-dihydropterin-6-carboxylic acid, is a natural product that is used as the light emitter in Luminodesmus bioluminescence.16 para-Aminobenzoic acid (PABA) is a B complex vitamin that is synthesized in the body. PABA is used in the formation of FA and the metabolism of protein.17 It is an antioxidant that helps to protect skin from sunburn and cancer. Glutamic acid is also a component of folic acid and a precursor of glutathione. Glutamic acid might play a role in the normal function of the heart and the prostate. An extensive literature survey reveals no reports on the oxidation of FA by any oxidants viewed from kinetic and mechanistic aspects. Hence, there is a need to understand the oxidation mechanism of this vitamin, so that a study could shed some light on the fate of the substrate in the biological systems. Therefore, this is an impetus for the present kinetic study, as the substrate FA is an important vitamin. The main objectives of the present investigation were to do the following: (i) develop appropriate experimental conditions for the oxidation process, (ii) study the catalytic activities of catalysts, (iii) elucidate plausible mechanisms, (iv) deduce appropriate rate laws, (v) ascertain the various reactive species, (vi) assess the relative rates of catalysts, (vii) understand the catalytic redox chemistry of FA in the presence of platinumgroup metal ions, (viii) determine the catalytic efficiencies of catalysts, and (ix) compare the reactivity with that under uncatalyzed conditions.

10.1021/ie900340k  2010 American Chemical Society Published on Web 01/12/2010

Ind. Eng. Chem. Res., Vol. 49, No. 4, 2010

2. Experimental Section 18

2.1. Materials. Bromamine-T was prepared using chloramine-T. To a solution of CAT (20 g in 400 mL of water), about 4 mL of liquid bromine was added dropwise with constant stirring at room temperature to yield dibromamine-T (DBT). The solid DBT was filtered under suction, washed thoroughly with ice-cold water until all the absorbed bromine had been removed, and then vacuum-dried for 24 h. About 20 g of DBT thus obtained was dissolved in 30 mL of 4 mol dm-3 NaOH with constant stirring at room temperature, and the resultant aqueous solution was cooled in ice. The pale yellow crystals of BAT that formed were filtered under suction, washed quickly with the minimum amount of ice-cold water, and dried over P2O5. An aqueous solution of the oxidant was standardized by iodometric procedure and preserved in brown bottles to prevent photochemical deterioration. Folic acid (Lancaster) was of acceptable purity and was used as received. An aqueous solution of the compound was prepared by dissolving FA in 5.0 × 10-4 mol dm-3 of NaOH and was employed for the kinetic study. Solutions of ruthenium trichloride (Merck), palladium chloride (Arora-Matthey), and platinum chloride (SD Fine Chemicals Ltd.) were prepared in 2 mM HCl, whereas osmium tetroxide (BDH) was prepared in 2 mM NaOH. Allowance for the amount of acid/alkali present in the catalyst solutions was made in preparing the reaction mixtures for kinetic runs. Mass spectrometry (MS) data were obtained on a 17A Shimadzu gas chromatograph with a QP-5050 Shimadzu mass spectrometer. Solvent isotope studies were performed with D2O (99.4%) supplied by Bhabha Atomic Research Center, Mumbai, India. Reagent-grade chemicals and doubly distilled water were used. 2.2. Kinetic Measurements. The reactions were carried out under pseudo-first-order conditions with a known excess in the initial concentration of FA ([FA]0) over the initial concentration of BAT ([BAT]0) at 313 K. The reactions were carried out in stoppered Pyrex boiling tubes whose outer surfaces were coated black to eliminate photochemical effects. For each run, requisite amounts of solutions of FA, NaOH, and catalyst and water (to keep the total volume constant for all runs) were introduced into the tube and thermostatted at 313 K until thermal equilibrium was attained. A measured amount of BAT solution, also thermostatted at the same temperature, was rapidly added with stirring to the mixture in the tube. The progress of the reaction was monitored by the iodometric determination of unreacted BAT in aliquots (5 mL each) of the reaction mixture withdrawn at different intervals of time. The course of the reaction was studied for at least 2.5 h. The pseudo-first-order rate constants (k′) calculated from the linear plots of log [BAT] vs time were reproducible to within (5%. 2.3. Reaction Stoichiometry. Reaction mixtures containing varying ratios of BAT to FA in the presence of 0.01 mol dm-3 NaOH and 1.0 × 10-5 mol dm-3 catalyst were equilibrated at 313 K for 24 h. Estimation of unreacted BAT in the reaction mixture showed that 1 mol of FA consumed 3 mol of BAT in all three catalyzed reactions. Accordingly, the following stoichiometric equation was formulated

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Table 1. Effect of Varying Reactant Concentrations on the Reaction Rate at 313 Ka k′ (×104 s-1)

103[BAT]0 (mol dm-3)

102[FA]0 (mol dm-3)

Ru(III)

Os(VIII)

Pd(II)

Pt(IV)

0.50 1.00 2.00 4.00 8.00 2.00 2.00 2.00 2.00 2.00

2.00 2.00 2.00 2.00 2.00 0.50 1.00 2.00 4.00 8.00

8.20 8.3 8.25 8.29 8.19 3.16 5.00 8.25 11.0 18.2

15.23 15.15 15.2 15.28 15.22 3.75 7.50 15.2 30.0 60.5

4.12 4.09 4.10 4.15 4.14 2.11 3.13 4.10 6.40 8.12

6.20 6.30 6.25 6.21 6.22 6.3 6.20 6.25 6.19 6.20

a [NaOH] ) 1.00 × 10-2 mol dm-3, [catalyst] ) 1.00 × 10-5 mol dm-3, I ) 0.30 mol dm-3 [in the case of Pt(IV) catalysis].

2.4. Oxidation Procedure and Product Analysis. Reaction mixtures containing 1 mmol of FA and 3 mmol of BAT in the presence of 1 mmol of NaOH and 1 µmol of catalyst under stirred conditions were allowed to progress for 6-7 h at 313 K in round-bottomed flasks. After completion of the reaction, the reaction products were neutralized with HCl and subjected to thin-layer chromatographic analysis, which revealed the formation of three oxidation products of FA, namely, pterin-6carboxylic acid, p-aminobenzoic acid, and glutamic acid, along with the reduction product of BAT, p-toluenesulfonamide. Pterin-6-carboxylic acid, p-aminobenzoic acid, and p-toluenesulfonamide were extracted with ether and separated by column chromotography. The two products were recrystallized by CH2Cl2 and subjected to mass spectral analysis. Separation of glutamic acid was achieved by the liophilization technique and confirmed by MS analysis. 3. Results The oxidation kinetics of FA by BAT was investigated at several initial concentrations of the reactants in NaOH medium in the presence of Ru(III), Os(VIII), Pd(II), and Pt(IV) catalysts at 313 K. All reactions were carried out under pseudo-firstorder conditions wherein [FA]0 . [BAT]0. 3.1. Effect of Varying Reactant Concentrations on the Rate. Under the experimental condition [FA]0 . [BAT]0, with other reaction conditions constant, the order with respect to the concentration of BAT was found to be unity in all cases, as indicated by the linearity of the plots of log [BAT] versus time (r > 0.9901). Further, the linearity of these plots with constant slopes obtained at various initial BAT concentrations indicate a first-order dependence of the reaction rate on the BAT concentration. The pseudo-first-order rate constants (k′) obtained are presented in Table 1. Under the same experimental conditions, the rate of the reaction increased with initial concentration of FA ([FA]0) (Table 1) for Ru(III)-, Os(VIII)-, and Pd(II)catalyzed reactions, and plots of log k′ versus log [FA] (Figure 1) were found to be linear (r > 0.9915) with slopes of 0.6, 1, and 0.5 for Ru(III), Os(VIII), and Pd(II) catalysis, respectively, indicating fractional-order dependences on the FA concentration for Ru(III) and Pd(II) catalysis and a first-order dependence on the FA concentration for Os(VIII). In contrast, the order with respect to the FA concentration was found to be zero in the case of Pt(IV) catalysis (Table 1). 3.2. Effect of Varying NaOH and Catalyst Concentrations on the Rate. The rate of the reaction increased with increasing NaOH concentration (Table 2) in all cases. Double-logarithmic plots of rate versus NaOH concentration (r > 0.9981; Figure 2) showed that the orders with respect to OH- concentration were

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Figure 4. Plot of log k′ vs 1/T for catalyzed reactions. Figure 1. Plot of log k′ vs log [FA]. Table 2. Effect of Varying NaOH and Catalyst Concentrations on the Reaction Rate at 318 Ka k′ (×104 s-1)

102[NaOH] (mol dm-3)

105[catalyst] (mol dm-3)

Ru(III)

Os(VIII)

Pd(II)

Pt(IV)

0.25 0.50 1.00 2.00 4.00 1.00 1.00 1.00 1.00 1.00

1.00 1.00 1.00 1.00 1.00 0.25 0.50 1.00 2.00 4.00

3.63 5.26 8.25 12.3 18.62 3.21 5.01 8.25 15.1 25.1

6.40 10.0 15.2 24.0 36.0 5.10 8.52 15.2 27.1 50.2

2.00 3.01 4.10 6.12 9.12 1.01 2.00 4.10 8.25 16.5

2.72 4.21 6.25 9.19 13.9 1.57 3.15 6.25 12.6 25.1

a [BAT]0 ) 2.00 × 10-3 mol dm-3, [FA]0 ) 2.00 × 10-2 mol dm-3, I ) 0.30 mol dm-3 [in the case of Pt(IV) catalysis].

Figure 2. Plot of log k′ vs log [NaOH].

Figure 3. Plot of log k′ vs log [catalyst].

less than unity (0.5-0.62), indicating fractional-order dependences on OH- concentration for all four catalyzed reactions. The rate of reaction increased with increasing concentration of catalyst (Table 2). The orders with respect to the Ru(III) and Os(VIII) concentrations were found to be fractional (0.75-0.82), whereas those with respect to the Pd(II) and Pt(IV) concentrations were found to be unity. Plots of log k′ versus log [catalyst] are presented in Figure 3.

3.3. Effect of Varying p-Toluenesulfonamide Concentration on the Rate. Addition to the reaction mixture of ptoluensulfonamide [PTS; (1.0-8.0) × 10-3 mol dm-3], a reduction product of BAT, did not affect the rate significantly. This indicates that PTS is not involved in any step prior to the rate-limiting step in the schemes proposed. 3.4. Effect of Varying Halide Ion Concentration on the Rate. The rate remained unchanged with the addition of Cl- or Br- ions in the form of NaCl or NaBr [(1.0-8.0) × 10-2 mol dm-3]. These results showed that halide ions do not play any role in the reaction. 3.5. Effect of Solvent Isotope on the Rate. Solvent isotope studies were performed using D2O for all four catalyzed reactions. Studies of the reaction rate in D2O medium for Ru(III)-, Os(VIII)-, Pd(II)-, and Pt(IV)-catalyzed reactions gave k′(H2O) values of 8.2 × 10-4, 15.2 × 10-4, 4.1 × 10-4, and 6.21 × 10-4 s-1, respectively, and k′(D2O) values of 9.87 × 10-4, 17.89 × 10-4, 5.13 × 10-4, and 7.56 × 10-4 s-1, respectively. The solvent isotope effect, k′(H2O)/k′(D2O) was found to be 0.83, 0.85, 0.8, and 0.82 for the four cases. 3.6. Effect of Varying Ionic Strength of the Medium on the Rate. An increase in ionic strength (I) of the reaction system through addition of NaClO4 showed a negligible effect on the reaction rate in the cases of Ru(III), Os(VIII), and Pd(II) catalysis. However, the rate of the reaction increased with increasing ionic strength in the case of Pt(IV) catalysis, giving rate constants at 0.1, 0.2, 0.3, 0.4, and 0.5 mol dm-3 ionic strengths of 13 × 10-4, 8.62 × 10-4, 6.21 × 10-4, 4.21 × 10-4, and 3.21 × 10-4 s-1, respectively. Plots of log k′ versus I1/2 were found to be linear (r > 0.9928) with a positive slope of 1.64 for Pt(IV) catalysis. Hence, the ionic strength of the medium was maintained at a constant concentration of 0.30 mol dm-3 of NaClO4 only in the case of Pt(IV) catalysis for kinetic runs in order to swamp the reaction. 3.7. Effect of Varying Temperature on the Rate. The reactions were studied at different temperatures (303-323 K) with other experimental conditions kept constant. From the linear Arrhenius plots of log k′ versus 1/T (r > 0.9931; Figure 4), values of activation parameters for the composite reaction were evaluated for each catalyst (Table 3). 3.8. Test for Free Radicals. Addition of reaction mixtures to the aqueous acrylamide monomer solution did not initiate any polymerization, indicating the absence of in situ formation of free-radical species in the reaction sequence. Control experiments were also performed under the same reaction conditions but without the oxidant, BAT. 4. Discussion Pryde and Soper,19 Morris et al.,20 and Bishop and Jennings21 have shown the existence of similar equilibria for N-metalloN-arylhalosulfonamides in aqueous media. Like CAT, bromamine-T acts as an oxidizing agent in both acidic and alkaline

Ind. Eng. Chem. Res., Vol. 49, No. 4, 2010 Table 3. Temperature Dependence on Rate and Values of Activation Parameters for the Oxidation of FA by BAT in the Presence and Absence of Catalysta

[Ru(H2O)6]

temperature (K)

Ru(III)

Os(VIII)

Pd(II)

Pt(IV)

uncatalyzed

303 308 313 318 323

4.20 6.41 8.25 11.3 16.0

8.59 12.6 15.2 20.0 25.0

1.58 2.61 4.10 6.41 9.00

2.81 4.11 6.25 9.13 13.0

0.20 0.31 0.70 1.26 1.85

52.3 49.7 92.0 –145

37.6 35.0 87.9 –187

69.6 66.6 98.0 –95.7

61.1 58.4 96.6 –112

-

+ OH h [Ru(H2O)5(OH)]

2+

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+ H2O

(9)

k′ (×104 s-1)

Ea (kJ mol–1) ∆Hq (kJ mol–1) ∆Gq (kJ mol–1) ∆Sq (J K–1 mol–1)

3+

94.0 91.2 104 –31.8

[BAT]0 ) 2.00 × 10-3 mol dm-3, [FA]0 ) 2.00 × 10-2 mol dm-3, [NaOH] ) 1.00 × 10-2 mol dm-3, [catalyst] ) 1.00 × 10-5 mol dm-3, I ) 0.30 mol dm-3 [in the case of Pt(IV) catalysis].

Similar equilibria have been reported in Ru(III)-catalyzed oxidations of several other substrates with various oxidants in alkaline media.4,35,36 The existence of a complex between Ru(III) and FA was evidenced by the UV-visible spectra of Ru(III) and a Ru(III)-FA mixture, in which a shift of the Ru(III) band from 352 to 340 nm was observed, indicating the formation of a complex. Complex formation between a metal ion (M) and a substrate (S) is given by the equilibrium K

M + nS y\z (MSn)

(10)

a

solutions. Depending on the pH, BAT exhibits the following equilibria in aqueous solutions19-21 (where Ts ) p-CH3C6H4SO2) TsNBrNa h TsNBr- + Na+

(2)

TsNBr- + H+ h TsNHBr

(3)

2TsNHBr h TsNH2 + TsNBr2

(4)

TsNHBr + H2O h TsNH2 + HOBr

(5)

TsNBr2 + H2O h TsNHBr + HOBr

(6)

HOBr h H+ + OBr-

(7)

HOBr + H+ h H2OBr+

(8)

The possible oxidizing species in acidified BAT solutions are TsNHBr, TsNBr2, HOBr, and possibly H2OBr+. In alkaline solutions of BAT, TsNBr2 does not exist, and the possible oxidizing species are TsNBr- and OBr- anions, which could be transformed into the more reactive species TsNHBr and HOBr during the course of the reaction in alkali retarding steps. Several workers have observed the retarding effect of OH- ions on the rates of haloamine reactions with a number of substrates22-28 and have suggested that the reactivity of weakly alkaline solutions of haloamines is due to the formation of the conjugate acid TsNHBr from TsNBr- in an OH- retarding step. In the present investigation, however, the fact that OH- ions increase the rate of reaction clearly indicates that TsNBr- is the reactive oxidizing species. In earlier works,29-31 the positive influence of OH- ions on the rate of haloamine reactions with a number of substrates was observed, and RNX- (R ) Ts or PhSO2; X ) Cl or Br) was suggested to be the reactive oxidizing species. In the present investigations, the fact that the reaction is accelerated by OH- ions clearly indicates that the anion TsNBr- is the most likely oxidizing species involved in the oxidation of FA by BAT in all four catalyzed reactions. 4.1. Mechanism and Rate Law of Ru(III) Catalysis. Ruthenium(III) chloride acts as a catalyst in many redox reactions, particularly in alkaline media.32-34 Under the experimental conditions, [OH-] . [Ru(III)], and the fact that increasing OH- concentration increases the rate suggests that ruthenium(III) is mostly present as the hydroxylated species [Ru(H2O)5OH]2+, whose formation is given by the equilibrium

where M and MSn are two metal species with different extinction coefficients. For equilibrium reaction 10, Ardon37 derived the following relation 1/∆A ) 1/[S]n(1/∆E[MTotal]K) + 1/∆E[MTotal]

(11)

where K is the formation constant of the complex, [S] is the concentration of substrate (FA in this case), ∆E is the difference in extinction coefficient between the two metal species M and MS, [M]Total is the total concentration of metal species, and ∆A is the absorbance difference between the Ru(III) + FA mixture and Ru(III) alone. Equation 11 is valid when the concentration of substrate ([S]) is much higher than the total metal concentration ([M]Total). According to eq 11, a plot of 1/∆A vs 1/[S] or 1/[S]2 (i.e., n ) 1 or 2) should be linear in the case of 1:1- or 1:2-type complex formation between M and S. The ratio of the intercept to the slope of this linear plot gives the value of K. Ruthenium(III) in NaOH medium containing FA showed an absorption peak at 340 nm (λmax for the complex). Complex formation studies were performed at a wavelength of 340 nm. In a set of experiments, solutions were prepared by taking different amounts of FA (0.005-0.08 mol dm-3) at constant amounts of RuCl3 (1.0 × 10-5 mol dm-3) and NaOH (0.01 mol dm-3) at 313 K. The absorbance of these solutions was measured at 340 nm. The absorbance of Ru(III) in alkaline medium was also measured at same wavelength (340 nm). The difference in these absorbance values (with and without FA) gave the differential absorbance, ∆A. A plot of 1/∆A vs 1/[FA] was linear (r ) 0.9921) with an intercept suggesting the formation of a 1:1 complex between FA and Ru(III) catalyst. Further, a plot of log(1/∆A) vs log(1/[FA]) was also linear (r ) 0.9801). From the slope and intercept of the plot 1/∆A vs 1/[FA], the value of Scheme 1 K1

TsNHBr + OH- y\z TsNBr- + H2O (i) fast K2

FA + Ru(III) y\z X

(ii) fast

K3

X + TsNBr- y\z XI

(iii) fast

k4

XI 98 XII

(iv) slow and rds k5

XII + 2TsNBr- 98 products

(v) fast

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the formation constant of the complex, K, was deduced and found to be 5.5 × 102. In view of the above experimental results, the mechanism presented in Scheme 1 is proposed for the Ru(III)-catalyzed oxidation of FA by BAT in alkaline medium. In Scheme 1, X, XI, and XII are the intermediate species and whose structures are shown in Scheme 2, where a detailed mechanistic interpretation of the Ru(III)-catalyzed FA-BAT reaction in alkaline medium is depicted. The total concentration of BAT is [BAT]t, so [BAT]t ) [TsNHBr] + [TsNBr-] + XI

(12)

By substituting for [TsNHBr] and [TsNBr-] from steps i and iii of Scheme 1 into eq 12 and solving for [XI], we obtain [XI] )

K1K2K3[BAT]t[FA][Ru(III)][OH-] [H2O] + K1[OH-] + K1K2K3[FA][Ru(III)]

(13) From the slow and rate-determining step (rds) (step iv) of Scheme 1, the rate can be expressed as Scheme 2. Ru(III)-Catalyzed Oxidation of FA by Bromamine-T

rate ) -d[BAT]t /dt ) k4[XI]

(14)

By substituting for [XI] from eq 12 into eq 14, the following rate law was obtained rate )

K1K2K3K4[BAT]t[FA][Ru(III)][OH-] [H2O] + K1[OH-] + K1K2K3[FA][Ru(III)]

(15) The proposed mechanism (Scheme 1) and the derived rate law (eq 15) are consistent with each other and are substantiated by the following experimental facts: For reactions involving fast pre-equilibrium H+- or OH--ion transfer, the rate increases in D2O because D3O+ and OD- are 2-3 times stronger acids and stronger bases,38,39 respectively, than H3O+ and OH- ions. In the present studies, the observed solvent isotope effect of k′(H2O)/k′(D2O) < 1 is due to the greater basicity of OD- compared to OH-. The magnitude of the rate increase in D2O is small [k′(H2O)/k′(D2O) ) 0.83] compared to the expected value of a factor of 2-3, which can be attributed to the fractional-order dependence of the rate on the concentration of OH-. The negligible influence of variations in the ionic

Ind. Eng. Chem. Res., Vol. 49, No. 4, 2010 -

-

OsO4 + OH + H2O h [OsO4(OH)(H2O)]

Table 4. Values of Catalytic Constant (KC) at Different Temperatures and Activation Parameters Calculated Using KC Valuesa temperature (K)

Ru(III)

Os(VIII)

Pd(II)

Pt(IV)

303 308 313 318 323

2.25 3.44 4.26 5.67 7.99

10.5 15.4 18.2 23.5 29.6

13.7 20.0 34.0 51.5 71.5

26.0 38.0 55.5 78.5 111

40.1 37.4 67.0 –101

67.9 65.3 75.2 –7.40

58.4 55.6 73.3 –34.1

1555

(17)

[OsO4(OH)(H2O)]- + OH- h [OsO4(OH)2]2- + H2O (18)

strength and additions of PTS and halide ions on the rate of the reaction and also on the activation parameters is in good agreement with the mechanism proposed and the rate law derived. It was felt reasonable to compare the reactivity of BAT toward FA in the absence of Ru(III) catalyst under identical experimental conditions in order to evaluate the catalytic efficiency of Ru(III). The reactions were carried out at different temperatures (303-323 K), and from plots of log k′ vs 1/T (r > 0.9915), activation parameters were evaluated for the uncatalyzed reactions (Table 3). However, the Ru(III)-catalyzed reactions were found to be about 12 times faster than the uncatalyzed reactions. The catalyst Ru(III) forms a complex (X) with FA, which increases the reducing property of the substrate compared to the case without Ru(III). Moelwyn-Hughes40 derived the following equation to relate catalyzed and uncatalyzed reactions

The complexes [OsO4(OH)(H2O)]- and [OsO4(OH)2]2- can be reduced to [OsO2(OH)4]2- with octahedral geometry. These octahedral complexes of osmium are less likely to form highercoordination species with the oxidant/substrate. It is more realistic to postulate OsO4 with a tetrahedral geometry as the catalyst species.43 The formation of a complex between Os(VIII) and BAT was evidenced by UV-vis spectra of both Os(VIII) and Os(VIII)-BAT in which a shift of the Os(VIII) peak from 319 to 300 nm was observed, indicating formation of a complex. According to eq 11, a plot of 1/∆A vs 1/[BAT] (r ) 0.9932) with a nonzero intercept suggests the formation of a 1:1 complex between Os(VIII) and BAT. Further, a plot of log(1/∆A) vs log(1/[BAT]) was also linear (r ) 0.9878). From the slope and intercept of the plot of 1/∆A vs 1/[BAT], the value of formation constant of the complex, K, was found to be 8.29 × 102. Furthermore, the first-order dependence of the rate on the concentrations of both BAT and FA and the fractional-order dependence on the concentrations of each of OH- and Os(VIII) indicate that an intermediate complex is formed from the catalyst and oxidant and then interacts with the substrate. In light of these considerations, an Os(VIII)-catalyzed FA-BAT oxidation mechanism was formulated, as shown in Scheme 3. In Scheme 3, XIII and XIV represent the intermediate species, whose structures are shown in Scheme 4, where a detailed mechanistic interpretation of the Os(VIII)-catalyzed FA-BAT reaction in alkaline medium is illustrated. The total effective concentration of BAT is

k1 ) k0 + KC[catalyst]x

[BAT]t ) [TsNHBr] + [TsNBr-] + [XIII]

Ea (kJ mol–1) ∆Hq (kJ mol–1) ∆Gq (kJ mol–1) ∆Sq (J K–1 mol–1)

48.5 45.6 71.1 –86.8

[BAT]0 ) 2.00 × 1003 mol dm-3; [FA]0 ) 2.00 × 10-2 mol dm-3; [NaOH] ) 1.00 × 10-2 mol dm-3; [catalyst] ) 1.00 × 10-5 mol dm-3; I ) 0.30 mol dm-3 (in the case of Pt(IV) catalysis). a

(16)

where k1 is the observed pseudo-first-order rate constant obtained in the presence of Ru(III) catalyst, k0 is the pseudo-first-order constant for the uncatalyzed reaction, KC is the catalytic constant, and x is the order of the reaction with respect to Ru(III). In the present investigations, the x value was found to be 0.75. Then, the value of KC was calculated using eq 16. The value of KC was obtained at different temperatures (303-323 K), and KC was found to vary with temperature. Further, a plot of log KC vs 1/ T was linear (r > 0.9932), and values for the energy of activation and other activation parameters for the Ru(III) catalyst were computed, as summarized in Table 4. The proposed mechanism is supported by the observed moderate values of energy of activation and other thermodynamic parameters. The fairly high positive values of the free energy and enthalpy of activation suggest that the transition state is highly solvated, whereas the fairly high negative entropy of activation (Table 3) indicates the formation of a rigid associated transition state. Addition of the reduction product of BAT, PTS, did not alter the rate, indicating its noninvolvement in the preequilibrium with the oxidant. Addition of halide ions had no effect on the rate, indicating that halide ions play no role in the reaction. All of these observations support the proposed mechanism. 4.2. Mechanism and Rate Law of Os(VIII) Catalysis. Osmium tetroxide is used as a homogeneous catalyst for redox reactions with different oxidants in alkaline media.41,30 Osmium is stable in its 8+ oxidation state and exists in the following equilibria4,42-45

(19)

By substituting for [TsNHBr] and [TsNBr-] from steps i and ii of Scheme 3 into eq 19 and solving for [XIII], we obtain [XIII] )

K6K7[BAT]t[Os(VIII)][OH-] [H2O] + K6[OH-] + K6K7[Os(VIII)][OH-]

(20) From the slow step of Scheme 3, the rate can be expressed as rate ) -d[BAT]t /dt ) k8[XIII][FA]

(21)

By substituting for [XIII] from eq 20 into eq 21, the following rate law was obtained Scheme 3 K6

TsNHBr + OH- y\z TsNBr- + H2O (i) fast K7

TsNBr- + Os(VIII) y\z XIII

(ii) fast

k8

XIII + FA 98 XIV

(iii) slow and rds k9

XIV + 2TsNBr- 98 products

(iv) fast

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Ind. Eng. Chem. Res., Vol. 49, No. 4, 2010

rate )

K6K7K8[BAT]t[FA][Os(VIII)][OH-] -

-

[H2O] + K6[OH ] + K6K7[Os(VIII)][OH ]

(22) Scheme 3 and the rate law in eq 22 explain the obtained experimental results and are supported by the following facts: Solvent isotope studies showed that k′(H2O)/k′(D2O) < 1. This is generally correlated with the fact that OD- ion is a stronger base than OH-, and in base-catalyzed reactions, enhancement of rate in D2O medium is expected and was observed in the present study.38,39 The negligible influence of variations in the ionic strength and addition of PTS and halide ions on the rate of the reaction and the activation parameters is in good agreement with the mechanism proposed and the rate law derived. The reactivity of BAT toward FA in the absence of Os(VIII) catalyst was compared with the Os(VIII)-catalyzed reaction under identical experimental conditions. Rate constants revealed that the Os(VIII)-catalyzed reactions are 22-fold faster than uncatalyzed reactions (Table 3). The value of KC was determined (from eq 16) at different temperatures, and from a plot of log KC vs 1/T (r ) 0.9968), values of activation parameters for Os(VIII) catalyst were computed (Table 4). Scheme 4. Os(VIII)-Catalyzed Oxidation of FA by Bromamine-T

4.3. Mechanism and Rate Law of Pd(II) Catalysis. Palladium(II) chloride catalysis has been observed during various redox reactions.4,29 Pd(II) exists as different complex species in alkaline solutions,4,29,46 and the possible Pd(II) complex species are [Pd(OH)Cl3]2-, [Pd(OH)2Cl2]2-, [Pd(OH)3Cl]2-, and [Pd(OH)4]2-. The species [Pd(OH)3Cl]2- and [Pd(OH)4]2- are not commonly found because they are insoluble. Further, the rate increases with increasing OH- concentration, and there was no effect of Cl- concentration on the rate of reaction, which clearly rules out [Pd(OH)Cl3]2- as the reactive species. Hence, [Pd(OH)2Cl2]2- complex ion was assumed to be the reactive species in the present study. UV-vis spectra of FA, BAT, and a mixture of the two showed that a complex is formed between FA and BAT. Absorption maxima in alkaline medium appear at 283 nm for FA, 222 nm for BAT, and 270 nm for their mixture. The hypsochromic shift of 13 nm from 283 to 270 nm for FA suggests that complexation occurs between FA and BAT. On the basis of these experimental findings, the probable mechanism presented in Scheme 5 is proposed for BAT oxidation of FA by Pd(II) catalysis in alkaline medium. In Scheme 5, XV and XVI represent the intermediate species whose structures are shown in Scheme 6 where a detailed

Ind. Eng. Chem. Res., Vol. 49, No. 4, 2010

rate ) -d[BAT]t /dt ) k12[X ][Pd(II)]

Scheme 5

V

K10

TsNHBr + OH- y\z TsNBr- + H2O (i) fast

1557

(25)

By substituting for [XV] from eq 24 into eq 25, the following rate law was derived

K11

TsNBr- + FA y\z XV V

k12

VI

-

X + Pd(II) 98 X X

VI

(ii) fast

rate )

(iii) slow and rds

k13

+ TsNBr 98 products

(iv) fast

Scheme 6. Pd(II)-Catalyzed Oxidation of FA by Bromamine-T

K10K11K12[BAT]t[FA][Pd(II)][OH-] [H2O] + K10[OH-] + K10K11[FA][OH-]

(26)

The rate law in eq 26 satisfies all of the experimental observations. The increase in the reaction rate in D2O medium supports the proposed mechanism because it is well-known that OD- is a stronger base than OH-38,39 ion and, hence, exerts a stronger accelerating influence on the reaction. The proposed mechanism is also supported by the activation parameters, as well as the negligible effects of PTS, halide ions, and ionic strength on the rate of reaction. Further, the energies of activation of Pd(II)catalyzed and uncatalyzed reactions were computed, and it was found that Pd(II)-catalyzed reactions were nearly 6 times faster than the uncatalyzed reactions (Table 3). The values of KC and the activation parameters for Pd(II) catalyst were determined (Table 4) by plotting log KC vs 1/T (r ) 0.9971). 4.4. Mechanism and Rate Law of Pt(IV) Catalysis. Pt(IV) catalysis is well-reported3,47,48 in a variety of redox reactions. Chloroplatinic acid [H2(PtCl6)] is the starting material in platinum(IV) catalysis. In aqueous solution, chloroplatinic acid ionizes49 as follows H2[PtCl6] h [PtCl6]2- + 2H+

(27)

In alkaline medium (pH > 8), [PtCl6]2- is in equilibrium with [PtCl5(OH)]2- 48,49 [PtCl6]2- + OH- h [PtCl5(OH)]2- + Cl-

(28)

Further ligand (Cl-) replacement from [PtCl5(OH)]2- has also been reported48,49 [PtCl5(OH)]2- + OH- h [PtCl4(OH)2]2- + Cl-

mechanistic interpretation of Pd(II)-catalyzed FA-BAT reaction in alkaline medium is presented. The total effective concentration of BAT is

(29)

However, the dihydroxy platinum(IV) species is quite unstable50 in aqueous solutions, and therefore, under the present experimental conditions, [PtCl5(OH)]2- acts as the reactive species of Pt(IV) in alkaline medium. The formation of a complex between Pt(IV) and oxidant was evidenced from UV-vis spectra of both Pt(IV) and Pt(IV)-BAT, in which a shift of the Pt(IV) signal from 360 to 345 nm was observed, indicating formation of a complex. According to eq 11, a plot of 1/∆A vs 1/[BAT] (r ) 0.9972) with a nonzero intercept suggests the formation of a 1:1 complex between Pt(IV) and BAT. Further, a plot of log 1/∆A vs log 1/[BAT] was also linear (r ) 0.9876). From the slope Scheme 7

[BAT]t ) [TsNHBr] + [TsNBr-] + [XV]

(23)

By substituting for [TsNHBr] and [TsNBr- ] from steps i and ii of Scheme 5 into eq 23 and solving for [XV ], one obtains [XV] )

K10K11[BAT]t[FA][OH-]

K14

TsNHBr + OH- y\z TsNBr- + H2O (i) fast k15

TsNBr- + Pt(IV) 98 XVII

(ii) slow and rds

k16

XVII + FA 98 XVIII

[H2O] + K10[OH-] + K10K11[FA][OH-]

(24) From the slow step of Scheme 5, the rate can be expressed as

(iii) fast

k17

XVIII + TsNBr- 98 products

(iv) fast

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and intercept of the plot of 1/∆A vs 1/[BAT], the value of the formation constant of the complex, K, was found to be 3.29 × 102. Based on the experimental results, it is likely that TsNBracts as the reactive oxidant species in the present case as well. Considering the above facts and allof the observed experimental data, the reaction mechanism presented in Scheme 7 can be suggested for the Pt(IV)-catalyzed oxidation of FA by BAT in alkaline medium. In Scheme 7, XVII and XVIII represent the intermediate species whose structures are presented in Scheme 8, where a detailed mechanistic interpretation of the Pt(IV)-catalyzed FA-BAT reaction in alkaline medium is elucidated. The total effective concentration of BAT is [BAT]t ) [TsNHBr] + [TsNBr-]

(30)

By substituting for [TsNHBr] from step i of Scheme 7 into eq 30 and solving for [TsNBr-], we obtain [TsNBr-] )

K14[BAT]t[OH-] [H2O] + K14[OH-]

(31)

Scheme 8. Pt(IV)-Catalyzed Oxidation of FA by Bromamine-T

From the slow step of Scheme 7 (step ii), the rate can be expressed as rate ) -d[BAT]t /dt ) k15[TsNBr-][Pt(IV)]

(32)

By substituting for [TsNBr-] from eq 31 into eq 32, the following rate law was obtained rate )

K14K15[BAT]t[Pt(IV)][OH-] [H2O] + K14[OH-]

(33)

The rate law in eq 33 is in agreement with the observed experimental results. The observed isotope effects are in accordance with the theory of Collins and Bowman.38 The negligible influences of halide ions in the medium and added p-toluenesulfonamide, the rate of the reaction, and also the observed activation parameters further corroborate the suggested mechanism. Catalytic constants and activation parameters with reference to Pt(IV) catalyst have been computed (Table 4). Pt(IV)catalyzed reactions were found to be 9 times faster than the uncatalyzed reactions (Table 3). The effect of ionic strength (I) on the reaction rates in this work was described according to the theory of Brønsted and

Ind. Eng. Chem. Res., Vol. 49, No. 4, 2010 51

Bjerrum, which postulates the reaction through the formation of an activated complex. According to this theory, the effect of ionic strength on the rate of reaction involving two ions is given by the relationship log k' ) log k0 + 1.02ZAZBI1/2

(34)

where ZA and ZB are the valencies of ions A and B, respectively, and k and k0 are the rate constants in the presence and absence of the added electrolyte, respectively. In the present case, a primary salt effect was observed, as the rate increased with increasing ionic strength of the medium,51 supporting the involvement of negative ions in the rate-limiting step (Scheme 8). A Debye-Hu¨ckel plot (log k′ against I1/2) gave a straight line with a slope of 1.64. In the present system, two negative ions of the catalyst species and a negative ion of the oxidant species are involved in the rate-limiting step (Scheme 8), and the expected slope of 2 was not found. This might be due to the fact that the ionic strength employed is beyond the formal Debye-Hu¨ckel limiting law. Alternatively, formation of ion pairs could occur in concentrated solutions, as suggested by Brønsted and Bjerrum.51 The kinetics of all four catalyzed reactions were compared with those of uncatalyzed reactions under similar experimental conditions, and it was found that the catalyzed reactions are 6-22 times faster. For the catalyzed reactions, it is seen from the Table 3 that the activation energy is highest for the slowest reaction and vice versa. From an inspection of the rate constants and the activation energies (Table 3), the relative reactivity of these catalysts in the oxidation of FA by BAT in alkaline medium is in the order Os(VIII) > Ru(III) > Pt(IV) > Pd(II). This finding can be attributed to the d-electronic configuration of the metal ions. Osmium, having a d0 electronic configuration, has greater catalytic efficiency to oxidize the substrate compared to the other metal ions used in the present study. Thus, the catalytic efficiency decreases as the number of electrons in the d orbital increases. Pd(II), having a d8 electronic configuration, is expected to have the lowest catalytic efficiency among the catalysts used. It is likely that, during the course of the reaction, the metal ions momentarily undergo reduction when the oxidant/ substrate/oxidant-substrate complex is attached to the metal ions and that, after this, the metal ions return to their original valence state, as shown in Schemes 2, 4, 6, and 8. Ru(III) and Pt(IV), having d5 and d6 electronic configurations, respectively, exhibit intermediate catalytic efficiencies in the present study. Hence, based on the d-electronic configuration of the metal ions, the reactivity decreases as the number of electrons in the d orbital increases, as d0 [Os(VIII)] > d5 [Ru(III)] > d6 [Pt(IV)] > d8 [Pd(II)]. Therefore, the observed reactivity in the present study during the oxidation of FA in the presence of these catalysts is based on the d-electronic configuration of the metal ions and follows the order Os(VIII) > Ru(III) > Pt(IV) > Pd(II). 5. Conclusions Ru(III), Os(VIII), Pd(II), and Pt(IV) serve as effective catalysts for the oxidative conversion of folic acid to pterin-6carboxylic acid, p-aminobenzoic acid, and glutamic acid using bromamine-T in alkaline medium. The stoichiometries and oxidation products of the catalyzed reactions were found to be the same, but their kinetic patterns and oxidation mechanisms were different. Under comparable experimental conditions, the oxidation kinetics of FA with the four catalysts obeys the following underlying rate laws

1559

- z

rateRu(III) ) k[BAT]t[FA] [Ru(III)] [OH ] x

y

rateOs(VIII) ) k[BAT]t[FA][Os(VIII)]x[OH-]y ratePd(II) ) k[BAT]t[FA]x[Pd(II)][OH-]y ratePt(IV) ) k[BAT]t[FA]0[Pt(IV)][OH-]x In all cases above, x, y, z < 1. A comparison of the kinetics of oxidation of FA catalyzed by these catalysts, under identical experimental conditions, was found give the following order of reactivity: Os(VIII) > Ru(III) > Pt(IV) > Pd(II). The trend in the catalytic activities can be explained in terms of the different d-electronic configuration of the catalysts. Further, under identical experimental conditions, the kinetics of all four catalyzed reactions were compared with those of uncatalyzed reactions, and the catalyzed reactions were found to proceed 6-22 times faster than the uncatalyzed reactions. Based on the observed experimental results, detailed mechanistic interpretations and the related kinetic modeling were elucidated for each catalyst. Literature Cited (1) Hudlicky, M. Oxidations in Organic Chemistry; ACS Monograph 186; American Chemical Society: Washington, DC, 1990; p 838. (2) Russel, G. A.; Janzen, E. G.; Bemis, A. G.; Geels, E. J.; Moye, A. J.; Mak, S.; Strom, E. T. Oxidation of Hydrocarbons in Basic Solution. In SelectiVe Oxidation Processes; Fields, E. K., Ed.; American Chemical Society: Washington, DC, 1965; Vol. 51, pp 112-172. (3) Bhat, K. R.; Jyothi, K.; Thimme Gowda, B. Mechanism of Ru(III) Catalyzed Tl(III) Oxidation of Di- and Tri-substituted Phenols in Aqueous Acetic Acid. A Kinetic Study. Oxid. Commun. 2002, 25, 117, and references therein. (4) Cotton, F. A.; Wilkinson, G.; Murillo, C. A.; Bochmann, M. AdVanced Inorganic Chemistry, 6th ed.; John Wiley and Sons Inc.: New York, 1999. (5) Griffith, W. P. The Chemistry of Rare Platinum Metals; InterScience: New York, 1967. (6) Puttaswamy; Jagadeesh, R. V.; Vaz, N. Oxidation of Some Catecholamines by Sodium N-chloro-p-toluenesulfonamide in Acid Medium: A Kinetic and Mechanistic Approach. Cent. Eur. J. Chem. 2005, 3 (2), 326. (7) Puttaswamy; Jagadeesh, R. V.; Gowda, N. M. M. Oxidation of Metronidazole with Sodium N-bromo-p-toluenesulfonamide in Acid and Alkaline Media: A Kinetic and Mechanistic Study. Int. J. Chem. Kinet. 2005, 35, 700. (8) Rao, V. R. S.; Venkappayya, D.; Aravamudan, G. Stability Characteristics of Aqueous Chloramine-T Solutions. Talanta 1970, 17, 770. (9) Campbell, M. M.; Johnson, G. Chloramine T and Related NHalogeno-N-metallo Reagents. Chem. ReV. 1978, 78, 65. (10) Bremner, D. H. Synthesis of R-Aminoaldehydes from Enamines Review. Synth. Reagents 1986, 6, 9. (11) Benerji, K. K.; Jayaram, B.; Mahadevappa, D. S. Mechanistic Aspects of Oxidation by N-Metallo-N-haloarylsulfonamides. J.Sci. Ind. Res. 1987, 46, 65. (12) Rangappa, K. S.; Raghavendra, M. P.; Mahadevappa, D. S.; Channegowda, D. Sodium N-Chlorobenzenesulfonamide as a Selective Oxidant for Hexosamines in Alkaline Medium: A Kinetic and Mechanistic Study. Ind. Eng. Chem. Res. 1998, 65, 531. (13) Puttaswamy; Anuradha, T. M.; Ramachandrappa, R.; Made Gowda, N. M. Oxidation of Isoniazide by N-Haloarenesulfonamides in Alkaline Medium: A Kinetic and Mechanistic Study. Int. J. Chem. Kinet. 2000, 32 (4), 221. (14) Agarwal, G. R.; Agarwal, K. Agarwal’s Text Book of Biochemistry, (Physiological Chemistry), 8th ed.; Goel Publishing House: Meerut, India, 1995; p 477. (15) Nelson, D. L.; Cox, M. M. Lehninger Principles of Biochemistry, 3rd ed.; MacMillan Worth Publishers: New York, 2000; p 640. (16) Masaki, K.; Akira, K.; Sathorn, S.; Kazushi, K.; Minoru, I.; Osamu, S. 7,8 Dihydropterin-6-carboxylic Acid as Light Emitter of Luminous Millipede, Luminodesmus sequoiae. Bioorg. Med. Chem. Lett. 2001, 11, 1037.

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(17) Devlin, T. M., Ed. Textbook of Biochemistry with Clinical Correlations, 2nd ed.; John Wiley and Sons Inc.: New York, 1982; pp 476477. (18) Nair, C. G. R.; Kumari, R. L.; Indrasenan, P. Bromamine-T as a New Oxidimetric Titrant. Talanta 1978, 25, 525. (19) Pryde, D. R.; Soper, F. G. The Direct Interchange of Chlorine in the Interaction of p-Toluenesulphonamide and N-Chloroacetamide. J. Chem. Soc. 1931, 1510. (20) Morris, J. C.; Salazar, J. R.; Wineman, M. A. Equilibrium Studies on Chloro Compounds: The Ionization Constant of N-Chloro-p-toluenesulfonamide. J. Am. Chem. Soc. 1948, 70, 2036. (21) Bishop, E.; Jennings, V. J. Titrimetric Analysis with ChloramineT: The Status of Chloramine-T as a Titrimetric Reagent. Talanta 1958, 1, 197. (22) Hardy, F. F.; Johnston, J. P. The Interactions of N-Bromo-NSodiobenzesulfonamide (Bromamine-B) with p-Nitrophenoxide Ion. J. Chem. Soc., Perkin Trans. 2 1973, 742. (23) Higuchi, T.; Ikeda, K.; Hussain, A. Mechanism and Thermodynamics of Chlorine Transfer Among Organochlorinating Agents. Part II. Reversible Disproportination of Chloramine-T. J. Chem. Soc. (B) 1967, 546. (24) Ruff, F.; Kucsman, A. Acta Chim. Acad. Sci. Hung. 1969, 62, 438. (25) Mushran, S. P.; Sanehi, R.; Agrawal, M. C. Z. Naturforsch. 1972, 27B, 1161. (26) Mahadevappa, D. S.; Rangappa, K. S.; Gowda, N. M. M.; Gowda, B. T. Kinetic and Mechanistic Studies of Oxidation of Arginine, Histidine, and Threonine in Alkaline Medium by N-chloro-N-sodio-p-toluenesulfonamide. Int. J. Chem. Kinet. 1982, 14, 1183. (27) Meenakshisundaram, S.; Sockalingam, R. M. Os(VIII)-Catalysed Oxidation of Sulfides by Sodium Salt of N-Chlorobenzenesulfonamide. J. Mol. Catal. A: Chem. 2000, 160, 269. (28) Puttaswamy; Jagadeesh, R. V. Mechanistic Investigations of Oxidation of Isatins by Sodium N-chlorobenzenesulfonamide in Alkaline Medium: A Kinetic Study. Cent. Eur. J. Chem. 2005, 3 (3), 482. (29) Kondarasaiah, M. H.; Ananda, S.; Puttaswamy; Gowda, N. M. M. Palladium(II)-Catalyzed Oxidation of Primary Amines by Bromamine-T in Alkaline Medium: A Kinetic and Mechanistic Study. Synth. React. Inorg. Met-Org. Chem. 2003, 33 (7), 1145. (30) Puttaswamy; Jagadeesh, R. V. Kinetics of Oxidation of Pantothenic Acid by Chloramine-T in Perchloric Acid and in Alkaline Medium Catalyzed by OsO4: A Mechanistic Approach. Int. J. Chem. Kinet 2005, 37, 201. (31) Puttaswamy; Jagadeesh, R. V. Mechanistic Studies of Oxidation of Thiols to Disulfides by Sodium N-Chloro-p-toluenesulfonamide in an Alkaline Medium: A Kinetic Approach. Ind. Eng. Chem. Res. 2006, 45, 1563. (32) Singh, H. S.; Singh, R. K.; Singh, S. M.; Sisodia, A. K. Kinetics and Mechanism of the Ruthenium(III) Chloride Catalyzed Oxidation of Butan-2-ol and Methyl-1-Propanol by the Hexacyanoferrate(III) Ion in an Aqueous Alkaline Medium. J. Phys. Chem. 1977, 81 (11), 1044. (33) Bilehal, D. C.; Kulkarni, R. M.; Nandibewoor, S. T. Kinetics and Mechanistic Study of the Ruthenium(III) Catalyzed Oxidative Deamination and Decarboxylation of L-Valine by Alkaline Permanganate. Can. J. Chem. 2001, 79, 1926.

(34) Mahesh, R. T.; Bellaki, M. B.; Nandibewoor, S. T. Spectral and Mechanistic Study of the Ruthenium(III) Catalysed Oxidation of Gabapentin (Neurontin) by Heptavalent Manganese: A Free Radical Intervention. Cat. Lett. 2004, 9, 91. (35) Balado, A. M.; Galan, B. C.; Martin, F. J. P. Anal. Quim. 1992, 88, 170. (36) Radhakrishnamurthy, P. S.; Panda, H. P. Bull. Soc. Kinet. Ind. 1980, 2 (1), 6. (37) Ardon, M. Oxidation of Ethanol by Ceric Perchlorate. J. Chem. Soc. 1957, 1811. (38) Collins, C. J.; Bowman, N. S. Isotope Effects in Chemical Reactions, Van Nostrand Reinhold: New York. 1970; p. 267. (39) Wiberg, K. B. Chem. ReV. 1955, 55, 713. Physical Organic Chemistry, Wiley: New York. 1964. (40) Moelwyn-Hughes, E. A. Kinetics of Reaction in Solutions; Oxford University Press: Oxford, U.K., 1947; pp 297-299. (41) Abbar, J. C.; Lamani, S. D.; Nandibewoor, S. T. Mechanistic Investigation of Uncatalyzed and Osmium(VIII) Catalyzed Oxidation of Guanidine by Ag(III) Periodate Complex in Aqueous Alkaline Medium: A Comparative Kinetic Approach. Ind. Eng. Chem. Res. 2009, 48, 7550. (42) Griffith, W. P. Osmium and Its Compounds. Q. ReV. Chem. Soc. 1965, 19, 254. (43) Mayell, J. S. Oxidation of Olefins by Ferricyanide Using Osmium Tetroxide Catalyst. Ind. Eng. Chem. Res. 1968, 7 (2), 129. (44) Mackay, K. M.; Mackay, R. A. Introduction to Modern Inorganic Chemistry, 4th ed.; Prentice Hall: Englewood Cliffs, NJ, 1989; p 259. (45) Rangappa, K. S.; Ramachandra, H.; Mahadevappa, D. S.; Made Gowda, N. M. Osmium(VIII) Catalyzed Kinetics and Mechanism of Indoles Oxidation with Aryl-N-haloamines in Alkaline Medium. Int. J. Chem. Kinet. 1996, 28 (4), 265. (46) Szabo, K. J. Benzoquinone-Induced Stereoselective Chloride Migration in (η3-Allyl) Palladium Complexes. A Theoretical Mechanistic Study Complemented by Experimental Verification. Organometallics 1998, 17 (9), 1677. (47) Sen Gupta, K. K.; Begum, B. A Reactivities of Osmium(VIII), Iridium(IV) and Platinum(IV) Towards Glycolaldehyde. Transition Met. Chem. 1998, 23, 295. (48) Tripathi, R.; Kambo, N.; Upadyay, S. K. Platinum(IV) Complexes of Reducing Sugars in an Alkaline Medium and Their Resistance to Reaction withN-Bromosuccinimide. A Kinetic Study. Transition Met. Chem. 2004, 29, 861. (49) Sen Guptha, K. K.; Sen, P. K. Kinetics of the Oxidation of Hydroxylamine by Platinum(IV). J. Inorg. Nucl. Chem. 1977, 39, 1651. (50) Grinberg, A. A. The Chemistry of Complex Compounds; Pergamon Press: Oxford, U.K., 1962; p 279. (51) Laidler, K. J. Chemical Kinetics, 2nd ed.; McGraw-Hill: New York, 1965; pp 219-222.

ReceiVed for reView March 1, 2009 ReVised manuscript receiVed November 28, 2009 Accepted December 19, 2009 IE900340K