J Solution Chem (2015) 44:1205–1223 DOI 10.1007/s10953-015-0341-1
Mechanistic Investigations of Uncatalyzed and Ruthenium(III) Catalyzed Oxidation of Vanillin by Periodate in Aqueous Alkaline Medium Deepa G. Patil1 • Prashant A. Magdum1 Sharanappa T. Nandibewoor1
•
Received: 28 July 2014 / Accepted: 28 December 2014 / Published online: 9 June 2015 Ó Springer Science+Business Media New York 2015
Abstract The kinetics and mechanism of uncatalyzed and ruthenium(III) catalyzed oxidation of vanillin (Van) by periodate were studied in alkaline medium at 298 K, and at constant ionic strength of 0.3 moldm-3. The reaction exhibits 1:1 stoichiometry ([Van]:[periodate]). The reaction shows first-order kinetics in [periodate] and [Ru(III)] and less than unit order with respect to [Van] and [OH-]. The ionic strength and dielectric constant of the medium did not affect the rate significantly. The main products were identified by spot tests, melting temperature and FT-IR. From the effect of temperature on the reaction rate, the Arrhenius and activation parameters have been calculated. The catalytic constant (KC) was also calculated for Ru(III) catalysis at different temperatures. Plausible mechanisms have been proposed and rate laws explaining the experimental results are derived. Kinetic studies suggest that the active species of periodate and Ru(III) were [H2IO6]3- and [Ru(H2O)5OH]2?, respectively. The reaction constants involved in the different steps of the mechanism were calculated. The activation parameters with respect to the slow step of the mechanism, along with the corresponding thermodynamic quantities, were determined and discussed. Keywords
Vanillin Periodate Kinetics Oxidation Catalysis Ruthenium(III)
1 Introduction Periodate oxidations have been reported to play an important role in biological compounds [1–3]. Although periodate is a less potent oxidant in alkaline than in acidic media, it is widely employed as a diol cleaving reagent [4]. In an alkaline medium, periodate is known to exist as several different species involving multiple equilibria [5] and this prompted us to investigate its active form in the reaction system.
& Sharanappa T. Nandibewoor
[email protected] 1
P. G. Department of Studies in Chemistry, Karnatak University, Dharwad 580003, India
123
1206
J Solution Chem (2015) 44:1205–1223
3-Methoxy-4-hydroxy benzaldehyde, commercially called P vanillin (Van) and in general vanillin, occurs in nature as a glucoside, which hydrolyzes to form vanillin and a sugar. Van is also used in the preparation of perfume and as a catalyst to polymerize methyl methacrylate [6]. Van is a naturally occurring compound found in vanilla beans and may also be released to the environment as a byproduct in the decay process of plant material. If released to air, the vapor pressure of 0.01599 Pa at 298 K indicates Van will exist solely as a vapor in the ambient atmosphere. Van shows up in all sorts of products— in foods and flavoring, obviously, but also in fragrances, cosmetics and aromatherapy. Transition metals are known to catalyze many oxidation–reduction reactions since they involve multiple oxidation states. In recent years, the use of transition metal ions such as osmium, ruthenium, palladium, chromium and iridium, either alone or as binary mixtures, as catalysts in various redox processes, has attracted considerable interest [7]. Ru(III) acts as a catalyst in the oxidation of many organic and inorganic substrates [8, 9]. The catalyzed mechanism can be quite complicated due to the formation of different intermediate complexes, and different oxidation states of Ru. Although the mechanism of catalysis depends on the nature of the substrate, oxidant and on experimental conditions, it has been shown that metal ions act as catalysts by one of these different paths such as the formation of complexes with reactants or oxidation of the substrate itself or through the formation of free radicals [10]. In earlier reports, it was observed that Ru(III) forms a complex with the substrate, which becomes oxidized by the oxidant to form a Ru(IV)-substrate complex followed by the rapid decomposition to generate Ru(III) [11]. In other reports, formation of a Ru(III)-substrate complex with further cleavage in a concerted manner giving rise to a Ru(I) species, which gets rapidly oxidized by the oxidant to generate the catalyst has been observed [12]. In some other reports, it is observed that Ru(III) forms a complex with substrate and is oxidized by the oxidant with the regeneration of the catalyst [13]. Hence, understanding the role of Ru(III) in catalyzed reactions is important. We have observed that Ru(III) in micro amounts catalyzes the oxidation of Van by periodate in alkaline medium. Such oxidation studies may throw some light on the mechanism of conversion of the compounds in biological systems. In order to understand the active species of oxidant and catalyst, to calculate the activity of the catalyst and to propose the appropriate mechanism, the title reaction was investigated in detail. An understanding of the mechanism allows the chemistry to be interpreted, understood and predicted. Some work on oxidation of Van by hexacyanoferrate(III) [14], diperiodatoargantate(III) [15], cerium(IV) [16], bismuth(V) [17] and diperiodatonickelate(IV) [18] has been documented. A literature survey reveals that there are no reports either on uncatalyzed or on Ru(III) catalyzed oxidation of Van by periodate in alkaline medium. In order to investigate the redox chemistry and in view of potential pharmaceutical importance of Van, to determine the active species of oxidant and catalyst in such media, and to propose the appropriate mechanisms of the reactions on the basis of kinetic and spectral results, the title reaction was investigated in detail.
2 Experimental Section 2.1 Chemicals Used All chemicals used were of reagent grade and double distilled water was used throughout this study. The stock solution of periodate (0.02 moldm-3) was prepared by dissolving
123
J Solution Chem (2015) 44:1205–1223
1207
2.30 g of potassium metaperiodate (S.D. Fine Chem.&90 % pure) in 500 cm3 water and the solution was used after keeping it for 24 h. The concentration of the solution was verified by titration with standard sodium thiosulfate (sodium thiosulfate, C99.99 % trace metals basis) iodometrically at neutral pH, maintained by adding 5 cm3 of 5 % potassium dihydrogen phosphate (potassium phosphate monobasic AR, 99.5–100.5 %) and 5 cm3 of 5 % dipotasssium hydrogen phosphate (sodium phosphate dibasic, BioXtra, C99.0 %) solution, using starch solution as the indicator. A solution of vanillin (Sigma-Aldrich) was prepared by dissolving an appropriate amount of recrystallized sample in double distilled water. The purity of the Van sample was checked by comparing its melting temperature (355 K) with the literature data (355–356 K). The required concentration of Van was obtained from its stock solution. A standard solution of Ru(III) was prepared by dissolving ‘‘RuCl3’’ (S.D. Fine Chem. Extra pure) in 0.20 moldm–3 HCl. The concentration was determined [19, 20] by EDTA titration. The concentrations obtained were the averages of triplicate measurements that agreed within ±2 % error. A sodium thiosulfate (Thomos Baker Chemicals Ltd.) solution was prepared in water. It was standardized [20] against potassium iodate as follows—to the potassium iodide solution containing 1 moldm–3 sulfuric acid, a known volume of standard potassium iodate solution was added. The liberated iodine was titrated against sodium thiosulfate solution using the starch indicator. Potassium hydroxide and potassium nitrate were employed to maintain the required alkalinity and ionic strength, respectively. Potassium dihydrogen phosphate (Thomas Baker Chemicals Ltd.) and potassium iodide (S.D. Fine Chem.) were used in the iodometric determination of periodate at neutral pH.
2.2 Instruments Used For product analysis a Nicolet 5700 FT-IR spectrometer (Thermo, USA) was used. For pH measurements an ELICO pH meter model LI 120 (Hyderabad, India) was used.
2.3 Kinetic Measurement The kinetics were followed under pseudo-first-order rate conditions where [Van] [ [periodate] in both uncatalyzed and catalyzed reactions at 298 K, unless specified. The reaction was initiated by mixing periodate with the Van solution that contained the required concentrations of KNO3 and KOH. The reaction in the presence of catalyst Ru(III) was initiated by mixing periodate with the Van solution which also contained the required concentrations of KNO3, KOH and Ru(III) catalyst. The reaction was followed by measuring the decrease in concentration of periodate titrimetrically, using sodium thiosulfate, at regular intervals of time, and the reaction was quenched by adding ice pieces before pipetting the reaction mixture. In view of the modest concentration of OH- used in the reaction medium, attention was also directed to the potential effect of the reaction vessel surface on the kinetics. Use of polythene/acrylic wares gave the same results, indicating that the surface did not have any significant effect on the reaction rates. The reaction was followed to more than 90 % completion. Plots of log10 (concentration) versus time lead to the first-order rate constants (kU or kC). The plots were linear up to 85 % completion of reaction, and the rate constants were reproducible within ±5 %. Regression analysis of the experimental data, to obtain the regression coefficient r and the standard deviation S of points from the regression line, was performed with the Microsoft Office Excel 2003 program.
123
1208
J Solution Chem (2015) 44:1205–1223
3 Results 3.1 Stoichiometry and Product Analysis Different sets of mixtures containing varying ratios of periodate to Van, with constant amounts of OH- (0.3 moldm–3) and KNO3 (0.01 moldm–3) in the uncatalyzed reaction, and with a constant amount of 1.0 9 10–5 moldm–3 Ru(III) in the catalyzed reaction, with constant ionic strength of 0.31 moldm–3, were kept for 8 h in a closed vessel under nitrogen atmosphere at 298 K and then analyzed. Under the condition [IO 4 ] [ [Van], periodate in the reaction was estimated iodometrically at neutral pH. The results show the same 1:1 stoichiometry for both uncatalyzed and catalyzed reactions as given in Eq. 1:
ð1Þ The main oxidation product, vanillic acid, was isolated by acidifying the reaction mixture followed by ether extraction. The product identified by spot tests [18–21]. To a small portion of product in a test tube, 25 % sodium carbonate and 4 % of formaldehyde were added and gently heated in a water bath. Formation of a white precipitate along with the liberation of CO2 indicated the presence of vanillic acid. The nature of the vanillic acid was confirmed by the IR spectrum, which showed a carbonyl ([C=O) stretch at 1744 cm-1 and an O–H stretching at 3412.31 cm-1 for a carboxylic acid (Fig. 1). In contrast, for Van the O–H stretching was observed at 3177.50 cm-1. The shift to higher frequency is probably due to the presence of the OH group of carboxylic acid and the OH group attached to the ring. A broad intense peak at 1650.80 cm–1 is retained in the product due to absorption by the –OCH3 group. It was also confirmed by its melting temperature 479 K (literature melting temperature 479–481 K) using a melting point apparatus. It was observed that vanillic acid did not undergo further oxidation under the present kinetic conditions.
Fig. 1 FTIR spectrum of vanillic acid
123
J Solution Chem (2015) 44:1205–1223
1209
3.2 Reaction Orders As the periodate oxidation of Van in alkaline medium proceeds with a measurable rate in the absence of Ru(III), the catalyzed reaction is understood to occur in parallel paths with contributions from both the catalyzed and uncatalyzed paths. Thus, the total rate constants (kT) is equal to the sum of the rate constants of the catalyzed (kC) and uncatalyzed (kU) reactions, so kC = kT - kU. Hence the reaction orders for various species were determined from the slopes of plots of log10 (kU or kC) versus the respective concentration of species, except for [periodate] for which no variation of kU or kC was observed, as expected for the reaction conditions. The reaction orders were determined from the slopes of log10 kC versus log10 (concentration) plots by varying the concentrations of Van, OH- and catalyst Ru(III), in turn while keeping the other concentrations constant.
3.3 Effect of [Periodate] In the absence and presence of Ru(III) catalyst, the periodate concentration was varied in the range of 0.5 9 10–4 to 5.0 9 10-4 moldm-3 at fixed [Van], [OH-] and ionic strength. The lack of variation in the pseudo-first order rate constants at various concentration of KIO4 indicates the order in [IO 4 ] is unity (Table 1) for kU and (Table 2) for kC. This was also confirmed from the linearity of log10 (concentration) versus time plots up to 80 % completion of the uncatalyzed and Ru(III) catalyzed reactions.
Table 1 Effect of the variation of periodate, Van and OH- concentrations on the oxidation of Van by periodate in aqueous alkaline medium at 298 K and I = 0.31 moldm-3 4 [IO 4 ] 9 10 -3 (moldm )
[Van] 9 102 (moldm-3)
0.5
2.0
1.0
2.0
2.0
[OH-] 9 10 (moldm-3)
kU 9 104 (s-1) Found
Calculated
3.0
1.89
1.90
3.0
1.91
1.90
2.0
3.0
1.90
1.90
3.0
2.0
3.0
1.88
1.90
4.0
2.0
3.0
1.91
1.90
5.0
2.0
3.0
1.91
1.90
5.0
0.4
3.0
0.70
0.59
5.0
1.0
3.0
1.35
1.10
5.0
1.5
3.0
1.63
1.46
5.0
2.0
3.0
1.91
1.69
5.0
3.0
3.0
2.30
2.00
5.0
4.0
3.0
2.60
2.20
5.0
2.0
0.3
0.41
0.40
5.0
2.0
0.6
0.65
0.70
5.0
2.0
1.0
0.89
0.99
5.0
2.0
1.5
1.30
1.25
5.0
2.0
2.0
1.66
1.43
5.0
2.0
3.0
1.91
1.69
123
[OH-] 9 10 (moldm-3)
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
0.3
0.6
1.0
1.5
2.0
3.0
3.0
3.0
3.0
3.0
3.0
4 [IO 4 ] 9 10 (moldm-3)
0.5
1.0
123
2.0
3.0
4.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
4.0
3.0
2.0
1.5
1.0
0.4
2.0
2.0
2.0
2.0
2.0
2.0
[Van] 9 102 (moldm-3)
3.0
2.0
1.0
0.5
0.3
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
[Ru(III)] 9 105 (moldm-3)
5.00
3.53
1.90
0.80
0.55
1.90
2.63
3.16
4.00
5.92
8.07
3.13
2.60
1.90
1.60
1.38
0.70
1.90
1.91
1.90
1.92
1.89
1.90
kT 9 103 (s-1)
1.91
1.91
1.91
1.91
1.91
1.91
1.66
1.30
0.89
0.65
0.41
2.60
2.30
1.91
1.53
1.15
0.76
1.91
1.88
1.90
1.90
1.91
1.89
kU 9 104 (s-1)
4.80
3.33
1.70
0.60
0.35
1.70
2.46
3.03
3.91
5.85
8.03
2.87
2.37
1.70
1.44
1.26
0.62
1.70
1.72
1.71
1.73
1.69
1.71
Found
kC 9 103 (s-1)
5.50
3.66
1.80
0.91
0.55
1.80
2.50
3.20
4.30
5.80
8.50
2.70
2.00
1.80
1.50
1.20
0.60
1.80
1.80
1.80
1.80
1.80
1.80
Calculated
Table 2 Effect of the variation of periodate, Van, OH- and Ru(III) concentrations on the Ru(III) catalyzed oxidation of Van by periodate in aqueous alkaline medium at 298 K and I = 0.31 moldm-3
1210 J Solution Chem (2015) 44:1205–1223
J Solution Chem (2015) 44:1205–1223
1211
3.4 Effect of [Van] In both cases [Van] was varied in the range 0.4 9 10-2 to 4.0 9 10-2 moldm-3 at 298 K keeping all other reactant concentrations and conditions constant. The kU and kC values increased with the increase in concentration of Van, indicating an apparent less than unit order dependence on [Van] under the experimental conditions and concentration range used (Table 1, uncatalyzed; Table 2, Ru(III) catalyzed). This was also confirmed by the plots of kU versus [Van]0.5 and kC versus [Van]0.6, which are linear unlike the direct plot of kU versus [Van] and kC versus [Van] (Fig. 2; r C 0.991, S B 0.001 for uncatalyzed; r C 0.986, S B 0.008 for Ru(III) catalyzed).
0.5
-3
[Van] (mol dm )
-1
0.05
0.1
0.15
0.2 3.50E-03
3.00E-03
3.00E-03
2.50E-03
2.50E-03
2.00E-03
2.00E-03
1.50E-03
1.50E-03
1.00E-03
1.00E-03
5.00E-04
5.00E-04
-1
ku(s )
0 3.50E-03
ku(s )
(A)
0.00E+00 0.00E+00 0.00E+00 1.00E-02 2.00E-02 3.00E-02 4.00E-02 5.00E-02 -3
[Van](mol dm ) 0.6
-3
[Van] (mol dm ) 0.02 0.04 0.06 0.08
0.1
0.12 0.14 0.16 3.50E-03
3.00E-03
3.00E-03
2.50E-03
2.50E-03
2.00E-03
2.00E-03
1.50E-03
1.50E-03
1.00E-03
1.00E-03
5.00E-04
5.00E-04
k (s )
0 3.50E-03
-1
-1
k (s )
(B)
0.00E+00 0.00E+00 0.00E+00 1.00E-02 2.00E-02 3.00E-02 4.00E-02 5.00E-02 -3
[Van](mol dm )
Fig. 2 a Plot of kU versus [Van]0.5 and kU versus [Van], b plot of kC versus [Van]0.6 and kC versus [Van]
123
1212
J Solution Chem (2015) 44:1205–1223
3.5 Effect of [OH2] The effect of OH- in the absence and presence of Ru(III) was studied in the range of 0.03 to 0.3 moldm-3 at constant concentrations of Van, IO 4 and at ionic strength 0.31 moldm-3. The rate increased with increasing concentration of alkali in the case of the uncatalyzed reaction, whereas the rate decreased with increase in concentration of alkali for the catalyzed reaction (i.e., Table 1 for kU and Table 2 for kC).
3.6 Effect of [Ru(III)] The Ruthenium(III) concentration was varied from 3.0 9 10-6 to 3.0 9 10-5 moldm-3, at constant concentrations of IO 4 , Van, OH and ionic strength. The order in [Ru(III)] was found to be unity from the linearity of the plot kC versus [Ru(III)] (r C 0.995, S B 0.0018) (Table 2).
3.7 Effect of Initially Added Products In both the cases initially added products, vanillic acid and IO 3 , did not have any significant effect on the rate of reaction. Thus, from the observed experimental results, the experimental rate law for uncatalyzed reaction is given as 1.0 [Van]0.5 [OH-]0.6. Rate = kU [IO 4] And the rate law for the Ru(III) catalyzed reaction is given as 1.0 Rate = kC [IO [Van]0.6 [OH-]-0.6 [Ru(III)]1.0. 4]
3.8 Effect of Ionic Strength (I) and Dielectric Constant of the Medium (D) The ionic strength of the uncatalyzed and Ru(III) catalyzed reaction media was varied from 0.31 to 0.5 moldm-3 at constant [IO 4 ], [Van] and [OH ]. It was found that increasing of ionic strength had no significant effect on the rate of reaction in both cases of uncatalyzed and catalyzed reactions. The relative permittivity (D) effect was studied at constant concentrations of reactants and with other conditions constant, where t-butyl alcohol was varied from 5 to 25 % (V/V) in the reaction medium. Attempts to measure the relative permitivities were not successful. However, they were computed from the values of pure liquids [22]. The approximate dielectric constants of the reaction medium at various composition of t-butyl alcohol–water (V/V) were estimated by using the following equation: D ¼ Dw Vw þ DB VB where Vw and VB are volume fractions and Dw and DB are dielectric constants of water and t-butyl alcohol. There was no reaction of the solvent with the oxidant under the experimental conditions. The rate constants, kU and kC did not change with changes in the dielectric constant of the medium.
3.9 Test for Free Radicals (Polymerization) To test for free radicals in the presence of Ru(III) catalyst, the reaction mixture containing acrylonitrile scavenger was kept for 24 h in an inert atmosphere. On diluting the reaction
123
J Solution Chem (2015) 44:1205–1223
1213
mixture with methanol, no precipitate was observed, indicating that there is no intervention of free radicals in the reaction.
3.10 Effect of Temperature (T) The influence of temperature on the rate of reaction was studied for both uncatalyzed and catalyzed reactions at different temperatures (288, 293, 298 and 303 K) under varying concentrations of Van and OH- while keeping other conditions constant. The rate constants were found to increase with increase in temperature. The rate constant (k1) of the slow step of the uncatalyzed reaction was obtained from the slopes and intercept of plots of 1/kU versus 1/[Van] and 1/kU versus 1/[OH-] at the four different temperatures, and were used to calculate the activation parameters. The energy of activation corresponding to these constants was evaluated from an Arrhenius plot of log10 k1 versus 1/T (r C 0.990, S B 0.00012) and the activation parameters obtained are tabulated in Table 3. Similarly, the rate constant (k2) of the slow step of catalyzed reaction mechanism was obtained from the intercept of the plots of [Ru(III)]/kC versus 1/[Van] and [Ru(III)]/kC versus [OH-] at different temperatures. The values are given in Table 3. The energy of activation for the rate determining step was obtained from the plot of log10 k2 versus 1/T (r C 0.981, S B 0.058) and this and other activation parameters calculated for reaction are presented in Table 3.
3.11 Catalytic Activity It has been pointed out by Moelwyn-Hughes [23] that, in the presence of a catalyst, the uncatalyzed and catalyzed reactions proceed simultaneously so that kT ¼ kU þ KC ½RuðIIIÞx Here ‘kT’ is observed pseudo-first-order rate constant, ‘kU’ is the pseudo-first-order rate constant for the uncatalyzed reaction, kC is for the catalyzed reaction, ‘Kc’ is the catalytic constant and x the order of the reaction with respect to [Ru(III)]. In the present investigation, the x value for standard runs was found to be unity. Then, the value of Kc was calculated by using the equation: Kc ¼
kT kU kC x ¼ ½Ru(III) ½Ru(III)x
where kT - kU = kC. The values of Kc evaluated at different temperatures were found to vary with temperatures. Further, the plot of log10 kC versus 1/T was linear; values of the energy of activation and other activation parameters with reference to the catalyst were computed, and are summarized in Table 4.
4 Discussion The activity of periodate as an oxidizing agent varies greatly as a function of pH and is capable of subtle control. In acidic solution it is one of the most powerful oxidizing agents known, whereas in alkaline solution it is slightly less so. However, in aqueous alkaline media and in the pH ranges employed in the present study, periodate does not exist as
123
1214
J Solution Chem (2015) 44:1205–1223
Table 3 Activation parameters and thermodynamic quantities for the oxidation of Van by periodate in aqueous alkaline medium with respect to the slow step of Scheme 1 and Scheme 2 (uncatalyzed and catalyzed): (A) effect of temperature; (B) activation parameters; (C) effect of temperature on calculated K1 and K2 values; and (D) effect of temperature on calculated K3 and K4 values; (E) thermodynamic quantities using K1, K2, K3 and K4 (A) Temperature (K)
k1 9 104 (s-1) uncatalyzed
k2 9 103 (dm3mol-1s-1) catalyzed
288
1.80
1.89
293
2.30
2.36
298
3.16
2.90
303
4.53
3.22
(B) Parameters
Uncatalyzed values
Catalyzed values
Ea (kJmol-1)
44.7 ± 2.3
26.3 ± 1.3
DH# (kJmol-1)
42.3 ± 2.1
23.8 ± 1.2
DS# (JK-1mol-1)
-170 ± 8
DG#(298K) (kJmol-1)
91.1 ± 2.7
log10 A
4.34 ± 2
-99 ± 8 53.2 ± 2.9 8.1 ± 2.2
(C) Temperature (K)
K1 (dm3mol-1) uncatalyzed
K2 (dm3mol-1) uncatalyzed
288
1.82
198
293
1.38
222
298
1.05
241
303
0.83
274
(D) K3 9 102 (moldm-3) catalyzed
Temperature (K)
K4 (dm3mol-1) catalyzed
288
1.5
364
293
3.3
240
298
4.3
50.4
303
11.2
18.5
(E) Thermodynamic quantities
Values from K1 uncatalyzed
Values from K2 uncatalyzed
Values from K3 catalyzed
Values from K4 catalyzed
DH (kJmol-1)
-39.5 ± 1.5
15.2 ± 0.5
91.2 ± 4.2
-152 ± 8
DS (JK-1mol-1)
-132 ± 9
97.0 ± 4
282 ± 16
-476 ± 20
DG298 (kJmol-1)
-1.16 ± 0.03
13.3 ± 5.5
7.8 ± 0.5
-4
-3
-2
moldm , [Van] = 2.0 9 10 [Periodate] = 5.0 9 10 [Ru(III)] = 1.0 9 10-5 moldm-3, and I = 0.31 moldm-3
123
-3
moldm ,
-9.7 ± 1.2 -
[OH ] = 0.3 moldm-3,
J Solution Chem (2015) 44:1205–1223 Table 4 Values of the catalytic constant (Kc) at different temperatures and activation parameters calculated using Kc values: -4 moldm-3, [IO 4 ] = 5.0 9 10 [Van] = 2.0 9 10-2 moldm-3, [OH-] = 0.3 moldm-3, [Ru(III)] = 1.0 9 10-5 moldm-3, and I = 0.3 moldm-3
1215
Temperature (K)
KC 9 102
288
70.2
293
104
298
171
303
187
Parameters
Values
Ea (kJmol-1)
50 ± 2
DH# (kJmol-1)
47.5 ± 3.5
DS# (JK-1mol-1)
-44 ± 2
-1 DG#‘ (298K)(kJmol )
60 ± 3
log10 A
11 ± 1
H4 IO 6 because it is involved [24, 25] in the following equillibria, depending on the pH of the solution: þ H5 IO6 H4 IO 6 þH
ð2Þ
2 þ H4 IO 6 H3 IO6 þ H
ð3Þ
3 þ H3 IO2 6 H2 IO6 þ H
ð4Þ
The species H4 IO 6 exists near pH 7.0. Hence, under the alkaline conditions employed in the present system, the main species is expected to be trihydrogenparaperiodate (H3 IO2 6 ). The observed fractional order in alkali concentration may be understood in terms of H2 IO3 6 as the main species in alkaline medium with the following equilibrium, which is also supported by earlier work [26]: K1
3 H3 IO2 6 þ OH H2 IO6 þ H2 O
ð5Þ
4.1 Mechanism for the Uncatalyzed Reaction Based on the experimental results, a mechanism is proposed for which all the observed orders in each constituent such as [IO 4 ], [Van] and [OH ] may be well accommodated (Scheme 1). The less than unit order in [Van] presumably results from formation of a complex (C1) between periodate species and Van prior to the formation of the products. The complex decomposes in a slow step to give the products, vanillic acid and IO 3 (V) species. The plot of 1/kobs versus 1/[Van] proves complex formation between the oxidant and reductant. Scheme 1 leads to the following rate equation. The Michaelis–Menten plot shows complex formation between periodate and Van, which explains the less than unit order dependence on [Van]. Scheme 1 leads to the rate law (Eq. 6): kU ¼
Rate k1 K1 K2 ½VanT ½OH T ¼ 1þK1 ½OH þK1 K2 ½Van½OH ½H3 IO2 6
ð6Þ
which explains all the observed kinetic orders of different species.
123
1216
J Solution Chem (2015) 44:1205–1223
H3IO62-
OH-
K1
H2IO63-
H2O
CHO
H2IO63-
K2
Complex (C1)
OCH3 OH
Complex (C1)
COOH
H2O
k1
H2IO65-
Slow
2H+
OCH3 OH
Scheme 1 Proposed mechanism for the reaction between Van and IO 4
The rate law (Eq. 6) can be arranged in the following form, which is suitable for verification: 1 1 1 1 ¼ þ þ kU k1 K1 K2 ½Van½OH k1 K2 ½Van k1
ð7Þ
According to Eq. 7, other conditions being constant, the plots of 1/kU versus 1/[Van] (r C 0.951, S B 0.00005) and 1/kU versus 1/[OH-] (r C 0.986, S B 0.00006) should be linear with an intercept supporting the Van–periodate complex, which is verified in Fig. 3. From the intercepts and slopes of such plots the values of K1, K2 and k1 were calculated as 1.05 dm3mol-1, 241.5 dm3mol-1 and 3.16 9 10-4 s-1, respectively. These reaction constants are in good agreement with earlier work [27]. These constants were used to calculate the rate constants which were then compared with the experimental values and found to be reasonable agreement with each other as given in Table 1, which supports Scheme 1. The dielectric constant and ionic strength did not vary significantly enough to impact the reaction rates. Variation of the ionic strength from 0.31 to 0.5 dm3mol-1 has a marginal impact on the activity coefficients and therefore on the reaction rate, which also varies due to the presence of various ions in Scheme 1. A high negative value of DS# (–170.0 JK-1mol-1) suggests that the intermediate complex (C1) is more ordered than the reactants [28]. The thermodynamic quantities for the different equilibrium steps in Scheme 1 can be evaluated as follows. The [Van] and [OH-] concentrations (as in Table 1) were varied at different temperatures. From the slopes and intercepts, the values of K1 and K2 were calculated at different temperatures and these values are given in Table 3. The van’t Hoff’s plots were made for the variation of K1 and K2 with temperature [i.e. log10 K1 versus 1/T (r C 0.999, S B 0.430)] and [log10 K2 versus 1/T (r C 0.993, S B 0.318)], and the values of the enthalpy of reaction DH, entropy of reaction DS and Gibbs energy of reaction
123
J Solution Chem (2015) 44:1205–1223
(A)
1217
4.00E+04 3.50E+04
288
1/kU (s)
3.00E+04 2.50E+04
293
2.00E+04
298
1.50E+04
303 1.00E+04 5.00E+03 0.00E+00 -5.00E+010.00E+00 5.00E+01 1.00E+02 1.50E+02 2.00E+02 2.50E+02 3.00E+02 3
-1
1/[Van] (dm mol )
(B) 5.00E+04 4.50E+04
288
1/kU (s)
4.00E+04
293
3.50E+04 3.00E+04
298
2.50E+04 2.00E+04
303
1.50E+04 1.00E+04 5.00E+03 0.00E+00
0
5
10
15
20
25
30
35
40
1/[OH-] (dm3mol-1) Fig. 3 Verification of the rate law (Eq. 7) for the uncatalyzed oxidation of Van by periodate. Plots of (a) 1/ kU versus 1/[Van], and (b) 1/kU versus [OH-] at four different temperatures (conditions as given in Table 1)
DG were calculated for the first and second equilibrium steps. These values are given in Table 3. A comparison of the DH value for the second step (15.2 kJmol-1) of Scheme 1 with that of DH# (42.3 kJmol-1) 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 is fairly fast and involves a low activation energy [29, 30].
4.2 Mechanism for Ru(III) Catalyzed Reaction Ru(III) chloride acts as an efficient catalyst in many redox reactions, particularly in an alkaline medium [31]. It is interesting to identify probable Ru(III) chloride species in alkaline media. In the present study it is quite probable that the [Ru(H2O)5OH]2? species may assume the general form [Ru(III)(OH)x]3-x. The x value should always be less than
123
1218
J Solution Chem (2015) 44:1205–1223
six because there are no definite reports of any hexahydroxy ruthenium species. The remainder of the coordination sphere would then be filled by water molecules. At higher pH, electronic spectral studies have confirmed that Ru(III) chloride exists in the hydrated form as [Ru(H2O)6]3? [32]. Metal ions of the form [Ru(H2O)6]3? are also known to exist as [Ru(H2O)5OH]2? in alkaline media and are most likely mononuclear species. Hence, under the conditions employed, e.g. [OH-] [Ru(III)], Ru(III) is mostly present as the hydroxylated species, [Ru(H2O)5OH]2?. Similar species have been reported for the Ru(III) catalyzed oxidation of several other substrates with various oxidants in alkaline media [30, 33]. In earlier reports of Ru(III) catalyzed oxidation, it was observed that there is a fractional order dependence with respect to [substrate] and [Ru(III)], and unit order with respect to [oxidant]; Ru(III) forms a complex with the substrate [11]. This gets oxidized by the oxidant to form a Ru(IV)-substrate complex followed by rapid redox decomposition to regenerate Ru(III). In another case, if the process shows a zeroth-order dependence with respect to [oxidant], first order with respect to [Ru(III)] and a fractional order with respect to [substrate], then this involves the formation of a Ru(III)-substrate complex [12]. It undergoes further cleavage in a concerted manner giving rise to a Ru(I) species, which is rapidly oxidized by the oxidant to regenerate the catalyst. In some other reports, it is observed that Ru(III) forms a complex with a substrate and is oxidized by the oxidant with the regeneration of the catalyst [13]. Hence, the study of behavior of Ru(III) in a catalyzed reaction is significant. In the Ru(III) catalyzed reaction, the rate was first-order dependent on [Ru(III)] and [periodate], an apparent order of less than unity in [Van], and a negative fractional order dependence on [alkali]. No effect of added products was observed. Based on the experimental results, a mechanism is proposed as given in Scheme 2 for which all the observed orders in each constituent, [periodate], [Van], [Ru(III)], and [OH-], may be accommodated. In the first pre-equilibrium step of Scheme 2, the decrease in rate of reaction with increase in [OH-] can be explained in terms of prevailing equilibrium for formation of 3 H3 IO2 6 from H2 IO6 hydrolysis as given in Scheme 2. In the second equilibrium step, the Ru(III) species combines with one mole of Van to give an intermediate complex. This complex reacts with one mole of periodate species in a slow step to give the products,
H2IO63-
K3
H2O
H3IO62-
OH-
CHO
[Ru(H2O)5OH]2+
K4
Complex (C2)
OCH3 OH
Complex (C2)
COOH
H3IO62-
H2O
k2
[Ru(H2O)5OH]2+
Slow
H3IO64- 2H+
OCH3 OH
Scheme 2 Detailed scheme for the Ru(III) catalyzed oxidation of Van by alkaline periodate
123
J Solution Chem (2015) 44:1205–1223
1219
vanillic acid, with regeneration of Ru(III) species and iodate species by two equivalent changes of periodate in a single step, as no intervention of free radicals has been observed. The Michaelis–Menten plot proved complex formation between catalyst and substrate, which explains the less than unit order in [Van]. Such a complex between a catalyst and substrate has also been observed in other studies [32]. The rate law (Eq. 8) for Scheme 2 can be derived as: Rate k2 K1 K3 ½VanT ½OH T ½RuðIIIÞT ¼ k k ¼ ¼ k C T U 1 þ K3 ½Van þ K1 ½OH þ K1 K3 ½Van½OH ½H3 IO2 6
ð8Þ
This explains all the observed kinetic orders of different species. The rate law Eq. 8 can be rearranged to be Eq. 9, which is suitable for verification: ½RuðIIIÞ ½OH T ½OH T 1 1 ¼ þ þ þ kC k2 K4 ½VanT k2 k2 K3 K4 ½VanT k2 K3
ð9Þ
According to Eq. 9, other conditions being constant, the plots of [Ru(III)]/kC versus [OH-] (r C 0.998, S B 0.0023) and [Ru(III)]/kC versus 1/[Van] (r C 0.989, S B 0.0008) should be linear and are found to be so in Fig. 4. From the intercepts and slopes of such plots, the reaction constants K3, K4, and k2 were calculated as 0.043 mol-1dm-3, 50.4 dm3mol-1 and 2.90 9 103 dm3mol-1s-1, respectively. These constants were used to calculate the rate constants and compared with the experimental kC values and are found to be in reasonable agreement with each other, which supports Scheme 2. The thermodynamic quantities for the different equilibrium steps in Scheme 2 were evaluated as follows. The [Van] and [OH-] concentrations (Table 2) were varied at the four different temperatures. From the slopes and intercepts, the values of K3 and K4 were calculated at these different temperatures. A van’t Hoff’s plot was made for the variation of K3 and K4 with temperature [i.e., log10 K3 versus 1/T (r C 0.958, S B 0.042); log10 K4 versus 1/T (r C 0.956, S B 0.163)] and the values of the enthalpy of reaction DH, entropy of reaction DS, and Gibbs energy of reaction DG were calculated. These values are also given in Table 3. A comparison of DH value of first step (91.2 kJmol-1) of Scheme 2 with that of DH# (23.8 kJmol-1), 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 is fairly slow and involves a high activation energy [29]. The dielectric constant and ionic strength did not vary enough to impact the reaction rates. Variation of the ionic strength from 0.31 to 0.5 moldm-3 had a marginal impact on the activity coefficients and therefore on the reaction rate. This may also be due to the presence of various ions present in Scheme 2. The observed higher rate constant for the slow step indicates that the oxidation occurs via an inner-sphere mechanism. This conclusion is supported by earlier observations [34, 35]. The moderate values of DH= (23.8 kJmol-1) and DS= (-98.9 JK-1mol-1) are both favorable for an electron transfer process, which is in agreement with earlier work [36]. The negative value of DS= (–98.9 JK-1mol-1) indicates that the complex (C2) is more ordered than the reactants. The activation parameters evaluated for the catalyzed and uncatalyzed reactions explain the catalytic effect on the reaction. The catalyst Ru(III) forms a complex (C2) with the substrate, which enhances the reducing property of the substrate more than that in the absence of catalyst. Further, the catalyst, Ru(III), modifies the reaction path by lowering the energy of activation.
123
1220
J Solution Chem (2015) 44:1205–1223
-3
[Ru(III)]/k (mol dm s)
(A) 3.00E-02 2.50E-02
288
2.00E-02
293 298
1.50E-02
303 1.00E-02
5.00E-03
0.00E+00
0
50
100
150
200 3
250
300
-1
1/[Van] (dm mol )
[Ru(III)]/k (mol dm
-3
(B)
1.20E-02 1.00E-02
288
8.00E-03
293
6.00E-03
298
4.00E-03
303
2.00E-03 0.00E+00
0
0.1
0.2 -
0.3
0.4
-3
[OH ] (mol dm ) Fig. 4 Verification of the rate law (Eq. 8) for the Ru(III) catalyzed oxidation of Van by periodate. Plots of (a) [Ru(III)]/kC versus 1/[Van], and (b) [Ru(III)/kC versus [OH-] at four different temperatures (conditions as given in Table 2)
5 Conclusion A comparative study of the uncatalyzed and Ru(III) catalyzed oxidation of Van by periodate was made. Oxidation products were identified and found to be same for both cases. Among the various species of periodate in alkaline medium, [H2IO6]3- is considered as the active species for the uncatalyzed reaction. The active form of Ru(III) was found to be [Ru(H2O)5OH]2?. The catalyzed reaction was about ten fold faster than uncatalyzed reaction. It became apparent that the role of pH in the reaction medium is crucial. Thermodynamic activation parameters of individual steps in the mechanisms were evaluated for
123
J Solution Chem (2015) 44:1205–1223
1221
the uncatalyzed and Ru(III) catalyzed reactions at different temperatures. The activation parameters with reference to the catalyst were also computed. The description of the mechanisms is consistent with all the experimental evidences including kinetic, spectral and product studies.
Appendix Derivation of rate law for the uncatalyzed reaction According to Scheme 1: d H2 IO3 6 ¼k1 ½c1 Rate ¼ dt K2 ¼
ð10Þ
½C 1 3 ½Van] H2 IO6
½C1 ¼ K2 ½Van] H2 IO3 6
ð11Þ
H2 IO3 6 K1 ¼ H3 IO2 6 ½OH ¼ K1 H3 IO2 H2 IO3 6 6 ½OH ½C1 ¼ K1 K2 ½Van H3 IO2 6 ½OH Rate ¼ k1 K1 K2 ½Van H3 IO2 6 ½OH
ð12Þ
The total concentration of [Van]T is given by, ½VanT ¼ ½Vanf þ C1 where T and f refer to total and free concentrations. ½VanT ¼ ½Vanf þ K1 K2 ½VanT ½H3 O2 6 ½OH ¼ ½Vanf ð1 þ K1 K2 ½H3 IO2 6 ½OH Þ
½Van]f ¼
½VanT 1 þ K1 K2 ½H3 IO2 6 ½OH
used, the second term in denominator is In view of the low concentration of H3 IO2 6 neglected. ½Vanf ¼ ½VanT
ð13Þ
Similarly, the concentration of OH is ½OH f ¼ ½OH T
ð14Þ
123
1222
J Solution Chem (2015) 44:1205–1223
Similarly, 2 3 ½H3 IO3 6 T ¼ ½H3 IO6 f þ ½H3 IO6 f þ C1 2 2 ¼ ½H3 IO2 6 f þ K1 ½H3 IO6 f ½OH þ K1 K2 ½Van][H3 IO6 f ½OH ¼ ½H3 IO2 6 f ð1 þ K1 ½OH þ K1 K2 ½Van]½OH Þ
½H3 IO2 6 f ¼
½H3 IO2 6 T ð1 þ K1 ½OH þ K1 K2 ½Van]½OH Þ
ð15Þ
Substituting Eqs. 13, 14, and 15 in Eq. 12 we get, Rate =
ku ¼
k1 K1 K2 ½VanT ½H3 IO2 6 ½OH T ð1 þ K1 ½OH þ K1 K2 ½Van]½OH Þ
Rate k1 K1 K2 ½VanT ½OH T ¼ 2 ½H3 IO6 ð1 þ K1 ½OH þ K1 K2 ½Van]½OH Þ
ð16Þ
Similarly the rate law for catalysed reaction can be derived.
References 1. Tiziani, S., Sussich, F., Cesaro, A.: The kinetic of periodate oxidation of carbohydrates. 2. Polymeric substrates. Carbohydr. Res. 338, 1083–1095 (2003) 2. Bassil, D., Makris, D.P., Kefalas, P.: Oxidation of caffeic acid in the presence of L-cysteine: isolation of 2-S-cysteinylcaffeic acid and evaluation of its antioxidant properties. Food Res. Int. 38, 395–402 (2005) 3. Balakrishnan, B., Lesieur, S., Labarre, D.: Jayakrishnan: periodate oxidation of sodium alginate in water and in ethanol–water mixture: A comparative study. Carbohydr. Res. 340, 1425–1429 (2005) 4. Burton, C.A.: In: Wiberg, K.B. (ed.) Oxidation in Organic Chemistry, Part A, p. 368. Academic Press, New York (1965) 5. Cotton, F.A., Wilkinson, G.: Advanced Inorganic Chemistry a Comprehensive Text, 5th edn, p. 569. Wiley, New York (1988) 6. Imoto, M., Maeda, T., Ouchi, T.: Studies on vinyl polymerization. Part 372. Polymerization of methyl methacrylate initiated with vanillin. Chem. Lett. 2, 153–156 (1978) 7. Das, A.K.: Kinetic and mechanistic aspects of metal ion catalysis in cerium(IV) oxidation. Coord. Chem. Rev. 213, 307–325 (2001) 8. Shivananda, K.N., Lakshmi, B., Jagadeesh, R.V., Puttaswamy, Mahendra, K.N.: Ru(III)-catalyzed oxidation of acetanilide, o-chloroacetanilide, p-chloroacetanilide and p-methylacetanilide by sodium periodate at pH 4. Appl. Catal: Gen. 326, 202–212 (2007) 9. Tandon, P.K., Mahrotra, A.M., Srivastava, M., Dhunsia, M., Singh, S.B.: Ruthenium(III) catalysis in the reaction of hexacynosferrate and iodide ions in perchloric acid medium. Transit. Metal Chem. 32, 991–999 (2007) 10. Singh, A.K., Srivastava, S., Srivastava, J., Srivastava, R., Singh, P.: Studies in kinetics and mechanism of oxidation of D-glucose and D-fructose by alkaline solution of potassium iodate in the presence of Ru(III) as homogenious catalyst. J. Mol. Catal. A 278, 72–81 (2007) 11. Sandu, S., Sethuram, B., Rao, T.N.: Ruthenium (III) catalysed oxidation of chalcones by Ce(IV) in H2SO4–HOAc mixtures: a kinetic study. J. Indian Chem. Soc. 60, 198–200 (1983) 12. Panda, H.P., Sahu, B.D.: Kinetics and mechanism of ruthenium(III) catalyzed oxidation of substituted cinnamic acids by cerium(IV). Indian J. Chem. 28A, 323–324 (1989) 13. Malode, S.J., Abbar, J.C., Nandibewoor, S.T.: Mechanistic aspects of uncatalysed and ruthenium(III) catalysed oxidation of dl-ornithine monohydrochloride by silver(III) periodate complex in aqueous alkaline medium. Inorg. Chim. Acta 363, 2430–2442 (2010)
123
J Solution Chem (2015) 44:1205–1223
1223
14. Timy, P.J., Nandibewoor, S.T., Tuwar, S.M.: Kinetics and mechanism of the oxidation of vanillin by hexacyanoferrate(III) in aqueous alkaline medium. J. Solution Chem. 35, 51–62 (2006) 15. Deepak, S.M., Chimatadar, S.A., Nandibewoor, S.T.: Oxidation of vanillin by a new oxidant diperiodatoargentate(III) in aqueous alkaline medium. Ind. Eng. Chem. Res. 46, 1459–1464 (2007) 16. Satapathy, P.K., Baral, D.K., Aswar, A.S., Mohanty, P.: Kinetics and mechanism of oxidation of vanillin by cerium(IV) in aqueous perchlorate medium. Indian J. Chem. Tech. 20, 271–275 (2013) 17. Mishra, P.: Kinetics and mechanisms of oxidation of 4-hydroxy-3-methoxy benzaldehyde (vanillin) by Bi(V) in aqueous alkaline medium. Int. J. Pharm. Tech. Res. 1, 1234–1240 (2009) 18. Kathari, C., Pol, P., Nandibewoor, S.T.: The kinetics and mechanism of oxidation of vanillin by diperiodatonickelate(IV) in aqueous alkaline medium. Turk. J. Chem. 26, 229–236 (2002) 19. Reddy, C.S., Vijay Kumar, T.: kinetic and mechanistic study of ruthenium(III) catalysed and uncatalysed oxidation of oxalic acid by acid bromate. Indian J. Chem. 34A, 615–620 (1995) 20. Kamble, D.L., Chougale, R.B., Nandibewoor, S.T.: Kinetic and mechanism of uncatalyzed and Ru(III) catalyzed oxidation of allyl alcohol by N-bromosuccinimide in aqueous alkaline medium. Indian J. Chem. 35A, 865–869 (1996) 21. Feigl, F.: Spot Tests in Organic Analysis, p. 332. Elsevier, New York (1975) 22. Lide, D.R. (ed.): CRC Hand Book of Chemistry and Physics, 73rd edn. CRC Press, London (1992) 23. Moelwyn-Hughes, E.A.: Kinetics of Reaction in Solutions, p. 297. Oxford University Press, London (1947) 24. Crouthamel, C.E., Hayes, A.M., Martin, D.S.: Ionisation and hydration equilibria of periodic acid. J. Am. Chem. Soc. 73, 82–87 (1951) 25. Bailar Jr., J.C., Emeleus, H.J., Nyholm, S.R., Trotman-Dikenson, A.F.: Comprehensive Inorganic Chemistry, vol. 2. Pergamon Press, Oxford (1975) 26. Tuwar, S.M., Nandibewoor, S.T., Raju, J.R.: Oxidation of allyl alcohol by diperiodatonickelate(IV) in aqueous alkaline medium. J. Indian Chem. Soc. 69, 651–653 (1992) 27. Magdum, P.A., Nandibewoor, S.T.: Mechanistic investigations of uncatalysed and ruthenium(III) catalysed oxidation of pharmaceutically important D-sorbitol by periodate in aqueous alkaline medium. World J. Pharm. Res. 3, 910–931 (2014) 28. Lewis, E.S.: Investigation of rates and mechanism of reactions in techniques of chemistry. In: Weissberger, A. (ed.) Techniques of Chemistry, vol. 6, pp. 421–427. Wiley, New York (1974) 29. Rangappa, K.S., Raghavendra, M.P., Mahadevappa, D.S., Channegouda, D.: Sodium N-chloro-ptoluenesulfonamide as a selective oxidant for hexosamine in alkaline medium: a kinetic and mechanistic study. J. Org. Chem. 63, 531–536 (1998) 30. Bilehal, D.C., Kulkarni, R.M., Nandibewoor, S.T.: Kinetics and mechanistic study of ruthenium(III) catalysed oxidative deamination and decarboxylation of L-valine by alkaline permanganate. Can. J. Chem. 79, 1926–1933 (2001) ` 31. Jagadeesh, R.V., Puttaswamy, : Ru(III), Os(VIII), Pd(II), and Pt(IV) catalysed oxidation of glycyl A glycine by sodium N-chloro-p-toluenesulfonamide: Comparative mechanistic aspects and kinetic modeling. J. Phys. Org. Chem. 21, 844–858 (2008) 32. Cotton, F.A., Wilkinson, G., Murrillo, C.A., Bochmann, M.: Advanced Inorganic Chemistry, 6th edn, p. 1012. Wiley, New York (1999) 33. Sataraddi, S.R., Nandibewoor, S.T.: Oxidation of D-glucose by Silver (III) periodate complex in the presence of Ru(III)/Os(VIII) as a homogeneous catalyst: a comparative mechanistic study (Stopped flow technique). J. Solution Chem. 42, 897–915 (2013) 34. Martinez, M., Pitarque, M., Eldik, R.V.: Outer-sphere redox reactions of [CoIII(NH3)5(HxPyOz)](m-3)— complexes. A temperature and pressure-dependence kinetic study on the influence of the phosphorous oxo anions. J. Chem. Soc. Dalton Trans. 2665–2669 (1996) 35. Farokhi, S.A., Nandibewoor, S.T.: Kinetic, mechanistic and spectral studies for the oxidation of sulfanilic acid by alkaline hexacyanoferrate(III). Tetrahedron 59, 7595–7602 (2003) 36. Gunagi, S.D., Nandibewoor, S.T., Chimatadar, S.A.: Oxidation of acyclovir by a cuprate(III) periodate complex in aqueous alkaline media: a kinetic and mechanistic approach. J. Solution Chem. 41, 777–792 (2012)
123
