Mechanistic Investigations of Uncatalyzed and Ruthenium(III) Catalyzed Oxidation of Vanillin by Periodate in Aqueous Alkaline Medium

Article abstract of DOI:10.1007/s10953-015-0341-1

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 mol·dm^-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 (K _C) 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.

Full text of DOI:10.1007/s10953-015-0341-1

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 molÁdm-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 cat-
alytic 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 Á Periodate Á Kinetics Á Oxidation Á Catalysis Á 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
stnandibewoor@yahoo.com
1
P. G. Department of Studies in Chemistry, Karnatak University, Dharwad 580003, India
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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 mecha-
nism 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 docu-
mented. 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 molÁdm-3) was prepared by dissolving
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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 molÁdm–3 HCl. The concentration was determined [19, 20] by EDTA titra-
tion. 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 molÁdm–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] [ [pe-
riodate] 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 mea-
suring 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.
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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 molÁdm–3) and KNO3 (0.01 molÁdm–3) in the uncatalyzed reaction,
and with a constant amount of 1.0 9 10–5 molÁdm–3 Ru(III) in the catalyzed reaction, with
constant ionic strength of 0.31 molÁdm–3, were kept for 8 h in a closed vessel under
nitrogen atmosphere at 298 K and then analyzed. Under the condition [IOÀ4 ] [ [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
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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 molÁdm-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 [IOÀ4 ] 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 molÁdm-3
[IOÀ4 ] 9 104
[Van] 9 102
[OH-] 9 10
kU 9 104 (s-1
)
(molÁdm-3
)
(molÁdm-3
)
(molÁdm-3
)
Found
Calculated
0.5
1.0
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
2.0
2.0
2.0
2.0
2.0
2.0
0.4
1.0
1.5
2.0
3.0
4.0
2.0
2.0
2.0
2.0
2.0
2.0
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
1.89
1.91
1.90
1.88
1.91
1.91
0.70
1.35
1.63
1.91
2.30
2.60
0.41
0.65
0.89
1.30
1.66
1.91
1.90
1.90
1.90
1.90
1.90
1.90
0.59
1.10
1.46
1.69
2.00
2.20
0.40
0.70
0.99
1.25
1.43
1.69
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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 molÁdm-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).
[Van]0.5(mol dm-3)
(A)
0
0.05
0.1
0.15
0.2
3.50E-03
3.00E-03
2.50E-03
2.00E-03
1.50E-03
1.00E-03
5.00E-04
0.00E+00
3.50E-03
3.00E-03
2.50E-03
2.00E-03
1.50E-03
1.00E-03
5.00E-04
0.00E+00
0.00E+00 1.00E-02 2.00E-02 3.00E-02 4.00E-02 5.00E-02
[Van](mol dm-3)
[Van]0.6(mol dm-3)
(B)
0
0.02 0.04 0.06 0.08 0.1 0.12 0.14 0.16
3.50E-03
3.00E-03
2.50E-03
2.00E-03
1.50E-03
1.00E-03
5.00E-04
0.00E+00
3.50E-03
3.00E-03
2.50E-03
2.00E-03
1.50E-03
1.00E-03
5.00E-04
0.00E+00
0.00E+00 1.00E-02 2.00E-02 3.00E-02 4.00E-02 5.00E-02
[Van](mol dm-3)
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]
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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 molÁdm-3 at constant concentrations of Van, IOÀ4 and at ionic strength
0.31 molÁdm-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 molÁdm-3
,
at constant concentrations of IOÀ4 , 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 IOÀ3 , did not have any sig-
nificant effect on the rate of reaction. Thus, from the observed experimental results, the
experimental rate law for uncatalyzed reaction is given as
Rate = kU [IO4À]1.0 [Van]0.5 [OH-]0.6
And the rate law for the Ru(III) catalyzed reaction is given as
Rate = kC [IO4À]1.0 [Van]0.6 [OH-]-0.6 [Ru(III)]1.0
.
.
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 molÁdm-3 at constant [IO4À], [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 DwVw þ DBVB
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 experi-
mental 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
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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 con-
stants 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
x
kT ¼ kU þ KC½RuðIIIފ
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 investi-
gation, the x value for standard runs was found to be unity. Then, the value of Kc was
calculated by using the equation:
kT À kU
kC
Kc ¼
¼
x
x
½Ru(III)Š
½Ru(III)Š
where kT - kU = kC.
The values of Kc evaluated at different temperatures were found to vary with temper-
atures. 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
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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 (dm3Ámol-1Ás-1) catalyzed
288
293
298
303
1.80
2.30
3.16
4.53
1.89
2.36
2.90
3.22
(B)
Parameters
Uncatalyzed values
Catalyzed values
Ea (kJÁmol-1
)
44.7 2.3
42.3 2.1
26.3 1.3
23.8 1.2
DH# (kJÁmol-1
)
DS# (JÁK-1Ámol-1
)
-170
91.1 2.7
4.34
8
-99
8
DG#(298K) (kJÁmol-1
)
53.2 2.9
8.1 2.2
log10
A
2
(C)
Temperature (K)
K1 (dm3Ámol-1) uncatalyzed
K2 (dm3Ámol-1) uncatalyzed
288
293
298
303
1.82
1.38
1.05
0.83
198
222
241
274
(D)
Temperature (K)
K3 9 102 (molÁdm-3) catalyzed
K4 (dm3Ámol-1) catalyzed
288
293
298
303
1.5
3.3
364
240
50.4
18.5
4.3
11.2
(E)
Thermodynamic
quantities
Values from K1
uncatalyzed
Values from K2
uncatalyzed
Values from
K3 catalyzed
Values from
K4 catalyzed
DH (kJÁmol-1
)
-39.5 1.5
15.2 0.5
91.2 4.2
282 16
7.8 0.5
-152
8
DS (JÁK-1Ámol-1
)
-132
9
97.0
4
-476 20
-9.7 1.2
DG298 (kJÁmol-1
)
-1.16 0.03
13.3 5.5
[Periodate] = 5.0 9 10-4 molÁdm-3
,
[Van] = 2.0 9 10-2 molÁdm-3
,
[OH-] = 0.3 molÁdm-3
,
[Ru(III)] = 1.0 9 10-5 molÁdm-3, and I = 0.31 molÁdm-3
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1215
Table 4 Values of the catalytic
Temperature (K)
KC 9 102
constant (Kc) at different tem-
peratures and activation parame-
ters calculated using Kc values:
288
293
298
303
70.2
104
171
187
[IOÀ4 ] = 5.0 9 10-4 molÁdm-3
,
,
[Van] = 2.0 9 10-2 molÁdm-3
[OH-] = 0.3 molÁdm-3
[Ru(III)] = 1.0 9 10-5
molÁdm-3, and
,
Parameters
Values
I = 0.3 molÁdm-3
Ea (kJÁmol-1
)
50
2
DH# (kJÁmol-1
)
47.5 3.5
DS# (JÁK-1Ámol-1
)
-44
60
2
3
1
DG
(kJÁmol-1
)
#‘
(298K)
log10
A
11
H4IOÀ6 because it is involved [24, 25] in the following equillibria, depending on the pH of
the solution:
H5IO6 ꢀ H4IOÀ6 þ Hþ
H4IO6À ꢀ H3IO62À þ Hþ
H3IO26À ꢀ H2IO63À þ Hþ
ð2Þ
ð3Þ
ð4Þ
The species H4IOÀ6 exists near pH 7.0. Hence, under the alkaline conditions employed
in the present system, the main species is expected to be trihydrogenparaperiodate
(H3IO26À).
The observed fractional order in alkali concentration may be understood in terms of
H2IO36À as the main species in alkaline medium with the following equilibrium, which is
also supported by earlier work [26]:
K1
H3IO62À þ OHÀ ꢀ H2IO36À þ H2O
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 [IOÀ4 ], [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
IOÀ3 (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):
Rate
k1K1K2½VanŠ ½OHÀŠ
T
T
kU
¼
¼
ð6Þ
1þK1½OHÀŠþK1K2½VanŠ½OHÀŠ
which explains all the observed kinetic orders of different species.
½H3IO26À
Š
123

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J Solution Chem (2015) 44:1205–1223
K1
3-
2-
OH-
H2IO6
H2O
H3IO6
CHO
OH
K2
3-
Complex (C1)
H2IO6
OCH3
COOH
k1
2H+
5-
H2IO6
H2O
Complex (C1)
Slow
OCH3
OH
Scheme 1 Proposed mechanism for the reaction between Van and IO4À
The rate law (Eq. 6) can be arranged in the following form, which is suitable for
verification:
1
1
1
1
¼
þ
þ
ð7Þ
kU k1K1K2½VanŠ½OHÀŠ k1K2½VanŠ k1
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 dm3Ámol-1, 241.5 dm3Ámol-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 dm3Ámol-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 JÁK-1Ámol-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
1217
(A)
4.00E+04
3.50E+04
3.00E+04
2.50E+04
2.00E+04
1.50E+04
1.00E+04
5.00E+03
0.00E+00
288
293
298
303
-5.00E+010.00E+00 5.00E+01 1.00E+02 1.50E+02 2.00E+02 2.50E+02 3.00E+02
1/[Van] (dm3mol-1)
(B)
5.00E+04
4.50E+04
4.00E+04
3.50E+04
3.00E+04
2.50E+04
2.00E+04
1.50E+04
1.00E+04
5.00E+03
0.00E+00
288
293
298
303
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/
U versus 1/[Van], and (b) 1/kU versus [OH-] at four different temperatures (conditions as given in Table 1)
k
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 kJÁmol-1) of Scheme 1
with that of DH# (42.3 kÁJÁmol-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 experi-
mental 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
H3IO62À from H2IO63À 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,
K3
2-
H2IO6
OH-
3-
H3IO6
H2O
CHO
K4
[Ru(H2O)5OH]2+
Complex (C2)
OCH3
OH
COOH
k2
[Ru(H2O)5OH]2+
4-
2H+
2-
H3IO6
H3IO6
H2O
Complex (C2)
Slow
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
½H3IO26À
k2K1K3½VanŠ ½OHÀŠ ½RuðIIIފ
T
T
T
¼ kC ¼ kT À kU
¼
ð8Þ
1 þ K3½VanŠ þ K1½OHÀŠ þ K1K3½VanŠ½OHÀŠ
Š
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:
½RuðIIIފ
½OHÀŠ
½OHÀŠ
1
1
T
T
¼
þ
þ
þ
T
ð9Þ
kC
k2K3K4½VanŠ
k2K3
k2K4½VanŠ
k2
T
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-1Ádm-3
,
50.4 dm3Ámol-1 and 2.90 9 103 dm3Ámol-1Ás-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 kJÁmol-1) of Scheme 2 with
that of DH# (23.8 kJÁmol-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 molÁdm-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 con-
clusion is supported by earlier observations [34, 35]. The moderate values of DH=
(23.8 kJÁmol-1) and DS= (-98.9 JÁK-1Ámol-1) are both favorable for an electron transfer
process, which is in agreement with earlier work [36]. The negative value of DS=
(–98.9 JÁK-1Ámol-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

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J Solution Chem (2015) 44:1205–1223
(A) 3.00E-02
2.50E-02
2.00E-02
1.50E-02
1.00E-02
5.00E-03
0.00E+00
288
293
298
303
0
50
100
150
200
250
300
1/[Van] (dm3mol-1)
(B) 1.20E-02
1.00E-02
8.00E-03
6.00E-03
4.00E-03
2.00E-03
0.00E+00
288
293
298
303
0
0.1
0.2
0.3
0.4
[OH-] (mol dm-3 )
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 peri-
odate 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. Ther-
modynamic 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:
Â
Ã
d H2IO36À
Rate ¼ À
K2 ¼
¼k1½c1Š
ð10Þ
ð11Þ
dt
½C1Š
Â
Ã
½Van] H2IO36À
½C1Š ¼ K2½Van] H2IO63À
Â
Ã
Â
Ã
H2IO36À
Â
Ã
K1 ¼
H3IO26À ½OHÀŠ
H2IO63À ¼ K1 H3IO62À ½OHÀŠ
½C1Š ¼ K1K2½VanŠ H3IO26À ½OHÀŠ
Rate ¼ k1K1K2½VanŠ H3IO26À ½OHÀŠ
Â
Ã
Â
Ã
Â
Ã
Â
Ã
ð12Þ
The total concentration of [Van]T is given by,
½VanŠ ¼ ½VanŠ þ C1
where T and f refer to total and free concentrations.
T
f
½VanŠ ¼ ½VanŠ þ K1K2½VanŠ ½H3O26ÀŠ½OHÀŠ
T
f
T
¼ ½VanŠfð1 þ K1K2½H3IO26ÀŠ½OHÀŠÞ
½VanŠ
T
½Van]f ¼
1 þ K1K2½H3IO26ÀŠ½OHÀŠ
Â
Ã
In view of the low concentration of H3IO26À used, the second term in denominator is
neglected.
½VanŠ ¼ ½VanŠ
Similarly, the concentration of OHÀ is
ð13Þ
ð14Þ
f
T
½OHÀŠ ¼ ½OHÀŠ
f
T
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J Solution Chem (2015) 44:1205–1223
Similarly,
½H3IO63ÀŠ ¼ ½H3IO62ÀŠ þ ½H3IO36ÀŠ þ C1
T
f
f
¼ ½H3IO26ÀŠ þ K1½H3IO62ÀŠ ½OHÀŠ þ K1K2½Van][H3IO26ÀŠ ½OHÀŠ
f
f
f
¼ ½H3IO26ÀŠfð1 þ K1½OHÀŠ þ K1K2½Van]½OHÀŠÞ
½H3IO26À
Š
½H3IO26ÀŠ ¼
ð15Þ
ð16Þ
T
f
ð1 þ K1½OHÀŠ þ K1K2½Van]½OHÀŠÞ
Substituting Eqs. 13, 14, and 15 in Eq. 12 we get,
k1K1K2½VanŠ ½H3IO26ÀŠ½OHÀŠ
T
T
Rate =
ð1 þ K1½OHÀŠ þ K1K2½Van]½OHÀŠÞ
k1K1K2½VanŠ ½OHÀŠ
Rate
½H3IO26À
T
T
ku¼
¼
ð1 þ K1½OHÀŠ þ K1K2½Van]½OHÀŠÞ
Š
Similarly the rate law for catalysed reaction can be derived.
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