Kinetics and mechanism of oxidation of DL-methionine by hexacyanoferrate(III) in aqueous alkaline medium

Article abstract of DOI:10.1080/17415993.2010.481796

The title reaction was studied spectrophotometrically. The reaction was found to be first order in [HCF] and fractional order each in [dl-methionine] and [OH^-]. Initially added products did not alter the rate of reaction. The variation of ionic

Full text of DOI:10.1080/17415993.2010.481796

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Kinetics and mechanism of oxidation of
DL-methionine by hexacyanoferrate(III)
in aqueous alkaline medium
K. Sharanabasamma a & Suresh M. Tuwar a
a Department of Chemistry , Karnatak Science College , Dharwad,
580001, Karnataka, India
Published online: 26 May 2010.
To cite this article: K. Sharanabasamma & Suresh M. Tuwar (2010) Kinetics and mechanism of
oxidation of DL-methionine by hexacyanoferrate(III) in aqueous alkaline medium, Journal of Sulfur
Chemistry, 31:3, 177-187, DOI: 10.1080/17415993.2010.481796
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Journal of Sulfur Chemistry
Vol. 31, No. 3, June 2010, 177–187
Kinetics and mechanism of oxidation of DL-methionine
by hexacyanoferrate(III) in aqueous alkaline medium
K. Sharanabasamma and Suresh M. Tuwar*
Department of Chemistry, Karnatak Science College, Dharwad 580001, Karnataka, India
(Received 20 January 2010; final version received 28 March 2010)
The title reaction was studied spectrophotometrically. The reaction was found to be first order in [HCF]
and fractional order each in [dl-methionine] and [OH−]. Initially added products did not alter the rate of
reaction. The variation of ionic strength had unaltered the rate of reaction whereas dielectric constant of the
medium had an influence on kobs values.A suitable mechanism was proposed in which oxidant was involved
in the formation of adduct (K2) which is formed by reacting with dl-methionine in a prior equilibrium
step. It decomposes in a slow step (k) to form a free radical generated on dl-methionine. This free radical
leads to methionine sulfoxide in the fast step by reacting with another molecule of hexacyanoferrate(III).
The effect of temperature on rate of reaction was also studied. Activation parameters were evaluated. The
rate law was derived as:
kK1K2[dl-methionine][OH−]
kobs
=
.
1 + K1[OH−] + K2[dl-methionine] + K1K2[dl-methionine][OH−]
The reaction constants are calculated.
Keywords: kinetics; oxidation; dl-methionine; hexacyanoferrate(III); mechanism
1. Introduction
Hexacyanoferrate(III), HCF is a stable octahedral complex anion. It is a mild oxidizing agent as it
has the moderate (1) reduction potential (0.41V). It is mainly used (1) as an oxidizing agent in the
oxidation of variety of organic compounds (oxygenated, sulfated, nitrogenated, unsaturated, etc.).
Recently Leal et al. (2) have reviewed the oxidation studies of hexacyanoferrate(III). It reveals
that the order in oxidant and reductant is found to be one each in uncatalyzed reactions, whereas
in Ru(III) catalysis, the rate was found (3, 4) to be independent of [HCF]. It also emphasizes that
HCF exists as an ion pair form by pairing with Na+ or K+ present in the reaction mixture.
dl-Methionine (dl-M) is a sulfur-containing essential α-amino acid which is not synthesized
in the body and must be supplemented through the food. It contributes to supply mineral sulfur
improving tone and pliability of the skin, conditioning the hair and strengthening nails, and pro-
tecting the cells from airborne pollutants. It contributes to other compounds including S-adenosyl
methionine (SAM), which transfer the active labile methyl group and mineral sulfur to over 100
*Corresponding author. Emails: sureshtuwar@yahoo.co.in; sm.tuwar@gmail.com
ISSN 1741-5993 print/ISSN 1741-6000 online
© 2010 Taylor & Francis
DOI: 10.1080/17415993.2010.481796
http://www.informaworld.com

178 K. Sharanabasamma and S.M. Tuwar
biochemical reactions for normal brain function, and l-cystine, which is a component of glu-
tathione, an important antioxidant molecule in the body. The active transfer of the methyl group
is due to the interaction of adenosine triphosphate and a liver enzyme such as phosphatase or
dehydrogenase, with a high-energy S-methyl bond.
It is known that dl-methionine has three coordinating canters viz. O, N and S. Its oxidation path
is well established (5–9) with various oxidants. In all such cases (10–12) it is reported that, ‘S’
moiety is more susceptible for oxidation. However, the other α-amino acids such as l-histidine,
phenyl alanine, leucine, glycine and valine are oxidized (13, 14) by HCF to yield keto acids and
ammonia. In case (15) of oxidation of d-proline and l-methionine by Mn(III) (16) and HCF
catalyzed by Os(VIII) (17), the products and other kinetic results are found to be substantially
different from the present investigation.Apart from this and the fact that in the present investigation
the added ferrocyanide did not alter the rate of reaction which is uncommon feature in ferricyanide
oxidations. Hence, the title study is undertaken to establish the oxidation path of dl-methionine
by HCF in an aqueous alkaline medium.
2. Results
2.1. Stoichiometry
Ten different sets of concentrations of dl-methionine and HCF were kept at constant [OH−] and
ionic strength in an inert atmosphere at 28◦C for over 48 h in a closed vessel. The remaining
HCF was estimated spectrophotometrically by measuring the absorbance at 420 nm. The results
indicated that two moles of HCF were consumed by one mole of dl-methionine as in Equation (1).
O
S
OH
OH
S
+ 2[Fe(CN)6]3– + 2OH–
O
CH3
O
CH3
NH2
NH2
+ 2[Fe(CN)6]4– + H2O
(1)
The oxidation product of dl-methionine was found to be methionine sulfoxide, which was
characterized by the spot test (18) and also confirmed as described below. After completion of the
reaction, the sodium carbonate was added with vigorous stirring along with drop wise addition of
benzyl chloride solution to give a precipitate of N-benzoyl methionine sulfoxide, whose identity
was confirmed (19) by its melting point (183◦C). The product, methionine sulfoxide, was precipi-
tated by adding an acetone–ethanol mixture of 1:1 volume ratio to the reaction solution previously
brought to pH 4.0. It was subjected to IR analysis. It had shown a strong absorption at 1045 cm−1
,
in addition to the normal characteristic bands of the ionized carboxylic group and of an amine
salt occurring between 3130 and 2500 cm−1. The appearance of a band at 1045 cm−1 indicates
that dl-methionine is oxidized to sulfoxide without affecting any other part of the carbon chain
leading to the formation of keto acid or aldehyde.
2.2. Effect of varying oxidant concentration
The effect of [HCF] on the reaction rate was studied by varying the [HCF] from 7.0 × 10−5
to 7.0 × 10−4 mol dm−3 at constant concentrations of dl-methionine, alkali and ionic strength
and temperature (Table 1). It was observed that the slope of the plot of log (conc.) versus time

Journal of Sulfur Chemistry 179
Table 1. Effect of variation of [HCF], [dl-M] and [OH−] on the oxidation of
dl-methionine by hexacyanoferrate(III) in aqueous alkaline medium at 28 ◦C,
I = 1.0 mol−1 dm−3
.
[HCF] × 104
[dl-M] ×102
[OH−]
kobs × 104
kcal × 104a
(mol dm−3
)
(mol dm−3
)
(mol dm−3
)
(s−1
)
(s−1
)
0.70
1.0
2.0
3.0
4.0
5.0
7.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
0.50
0.80
1.0
2.0
5.0
7.0
4.0
4.0
4.0
4.0
4.0
4.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.22
1.20
1.21
1.20
1.22
1.23
1.20
0.411
0.552
0.628
0.917
1.41
1.23
1.23
1.23
1.23
1.23
1.23
1.23
0.399
0.551
0.630
0.921
1.42
1.0
1.69
1.62
0.10
0.15
0.20
0.40
0.60
0.80
0.340
0.421
0.509
0.722
0.932
1.12
0.338
0.416
0.520
0.718
0.934
1.13
Note: aThe kcal are calculated by using k, K1 and K2 as 2.44 × 10−4 s−1, 2.9 and
55.3 dm3 mol−1, respectively, in rate law (7).
remained constant for most of the different [HCF] values. It was also confirmed by the linearity of
the first-order plot after a small initial period. Hence, the order in [HCF] was considered as unity.
2.3. Effect of varying substrate concentration
The dl-methioninewasvariedintheconcentrationrangefrom5.0 × 10−3 to7.0 × 10−2 mol dm−3
while the concentrations of all other reactants were held constant. It was observed that the rate
constant of reaction was increased with increase in dl-methionine concentration (Table 1). The
order of dl-methionine concentration was calculated from the slope of plot of log kobs versus
log [dl-methionine] and found to be 0.5.
2.4. Effect of alkali concentration
The effect of OH− concentration on the reaction rate was studied in the concentration range
0.1–1.0 mol dm−3 while the concentrations of all other reactions were held constant. It was noticed
that there was an increase in ‘kobs’with increasing the concentration of OH− (Table 1). The slope
of the plot of log kobs versus log[OH−] lead the order 0.55.
2.5. Effect of ionic strength and solvent polarity
The effect of ionic strength was studied by varying [NaCl] in the reaction medium. The ionic
strength of the reaction medium was varied from 0.2 to 1.5 mol dm−3 at constant [HCF], [dl-
methionine] and [OH−]. It was found that the ionic strength had no influence on the rate of the
reaction.
The relative permittivity (D) effect was studied by varying the t-butyl alcohol–water content in
the reaction mixture with all other conditions being kept constant. The D values were computed

180 K. Sharanabasamma and S.M. Tuwar
(20) from the values of pure liquids. In the reaction, as ‘D’ decreases the kobs values increased.
The graph of log kobs versus 1/D was found to be linear with a positive slope. Earlier the reaction
between t-butyl alcohol and oxidant was studied. It was observed that there was no reaction
between solvent and oxidant.
2.6. Effect of initially added products
The added products, [Fe(CN)6]−4 and methionine sulfoxide, did not alter the rate of reaction in
the studied concentration range of 8.0 × 10−5–7.0 × 10−4 mol dm−3 at constant concentrations
of all reactants and conditions.
2.7. Intervention of free radicals
The polymerization study reveals that the oxidation of dl-methionine was followed with an
intervention of free radicals.
2.8. Effect of temperature
Kinetics were also carried out at various temperatures by varying [dl-methionine] at constant
concentrations of HCF, OH− and ionic strength. The rate of reaction increased with increasing
temperatures. The values of kobs at different temperatures were tabulated in Table 2. The energy of
activation was determined from the slope of the Arrhenius plot of log kobs versus 1/T . The other
activation parameters ꢀH =, ꢀS=, ꢀG= and log A were also calculated (Table 3). Apart from
this, the equilibrium constants for the adduct formation (discussed in the mechanism) at various
temperatures were evaluated from the plot of 1/kobs versus 1/[dl-methionine] plot (Figure 1).
From Van’t Hoff’s plot of log K versus 1/T , the thermodynamic quantities ꢀH, ꢀS and ꢀG
were also calculated for the equilibrium of the adduct formation (see discussion section). Both
activation parameters and thermodynamic parameters are tabulated in Table 3.
Table 2. Effect of temperature on rate for the oxidation of dl-methionine
by hexacyanoferrate(III) in aqueous alkali and equilibrium constant of
adduct. [HCF] = 4.0 × 10−4, [dl-M] = 4.0 × 10−2, [OH−] = 1.0 mol dm−3
,
I = 1.0 mol−1 dm−3
.
Temperature (K)
kobs × 104 (s−1
)
K2 (dm3 mol−1
)
301
306
311
316
1.21
1.64
2.40
3.41
55.3
64.9
76.9
88.0
Table 3. Thermodynamics and activation parameters for the oxidation of
dl-methionine by hexacyanoferrate(III) in aqueous alkaline medium.
Activation parameters
Values
Thermodynamic parameters
Values
Ea (kJ mol−1
)
55.7
2
2
1
2
ꢀH== (kJ mol−1
)
53.2
−42
66
ꢀH (kJ mol−1
ꢀS (J K−1 mol−1
ꢀG (kJ mol−1
)
−24
+112
−57
1
4
2
ꢀS =(J K−1 mol−1
)
)
ꢀG (kJ mol−1
log A
)
)
6.0 0.2

Journal of Sulfur Chemistry 181
30
25
20
15
10
5
301K
306K
311K
[[
316K
0
0
50
100
150
200
250
1/[DL-M] dm3 mol–1
Figure 1. Verification of rate law (7) on the oxidation of dl-methionine by hexacyanoferrate(III) in aqueous alkaline
medium for various temperatures at different [dl-methionine].
Each fractional order in [OH−] and [dl-methionine] is an implicit fact to support the pre-
equilibrium before forming the adduct; the adduct is formed by interacting with dl-methionine and
HCF in pre-equilibrium before the rate determining step, and the polymerization study strongly
supports the free radical intervention generated from the dl-methionine oxidation by alkaline
HCF. These results can be accommodated in the mechanism as shown in Scheme 1.
Scheme 1. Mechanism of oxidation of dl-methionine by alkaline hexacyanoferrate(III).
dl-Methionine has three donor atoms namely ‘N’ from amino moiety, ‘O’ from the carboxylic
group and ‘S’of the thionyl group. It is a known fact that ‘N’and ‘O’of α-amino acid are involved

182 K. Sharanabasamma and S.M. Tuwar
in coordination with metal ions by donating a pair of electrons. However, in the present study
aqueous alkali is used as reaction medium. Hence, their involvement in the adduct formation with
HCF can be ruled out. Moreover, the product of dl-methionine is found to be its sulfoxide. Hence,
it can be ascertained that sulfur atom donates a pair of electrons to the HCF in the formation of
adduct.
It is also known that HCF is an inert complex and exchange of CN− with dl-methionine is quite
unusual as CN− is a labile ligand. Nevertheless, HCF is reported (21, 22) to be involved in an outer
sphere electron exchange with thio compounds by forming intermediate complex of [Fe(CN)5–
CNSO3]5− particularly with SO32− or S2O23−. Lancaster and Murray (23) have suggested that
many oxidations by transitional metal complex of weak oxidants may proceed via intermediate
complexes as described in the case of the [Fe(CN)6]3−–SO32− system. In all such cases, the thio
compounds are bonded with one of the CN− ligands through CN–SO32−. This CN–SO32− facilitates
the transfer of electron from sulfur to metal ion through the CN group. The analogous complexes
in the present study may also be possible as sulfur is present in the α-amino acid. This is evidenced
from the plot of log (absorbance) versus time (Figure 2), which is a curve at the initial stage and
subsequently of constant slope, i.e. the reaction occurs in two consecutive stages (Figure 2). In the
first stage of the curve, the outer sphere complex may be rapidly formed between dl-methionine
and HCF followed by its oxidation in a linear portion of the second stage as shown in Figure 2.
Time (s)
0
500
1000
1500
2000
2500
3000
3500
–0.35
–0.4
–0.45
–0.5
–0.55
–0.6
Figure 2. log (Absorbance) versus time plot for the evidence of formation of adduct during the oxidation of
dl-methionine by HCF in alkali at 28 ◦C.

Journal of Sulfur Chemistry 183
A similar adduct formation between dl-methionine and cis-diaquaethylenediamineplatinum(II)
perchlorate has been reported (7). The probable structure of the adduct can be represented as
4–
CN
CN
NC
NC
NC
CH3
Fe
CNS
CH2
CH2
CH
H2N
COO
The mechanism as written in Scheme 1 suggests that the oxidation takes place when the adduct
is decomposed by giving free radicals generated from –S– as HCF is a single equivalent oxidant.
The polymerization study also envisages the formation of free radicals. In the subsequent fast step,
this free radical is further oxidized by reacting with another molecule of HCF to yield sulfoxide
in the fast step. The rate law for the above mechanism can be derived as
d[HCF]
−
= k[complex]
dt
Rate = kK1K2[Fe(CN6)]3−[dl-methionine]f [OH−].
(2)
However,
3−
[Fe(CN)6] = [Fe(CN)6]3− + [complex],
T
f
= [Fe(CN)6]3−[K2(CN)6]3−[dl-methionine]−,
f
f
f
= [Fe(CN)6]3−{1 + K2}[dl-methionine]−,
f
f
3−
[Fe(CN)6]
3−
T
[Fe(CN)6]
=
.
(3)
(4)
f
1 + K2[dl-methionine]−
f
From the equilibrium (i), the [OH−]f can be given as
[OH−]T
[OH−]f =
.
1 + K1[dl-methionine]f
In view of low concentration of dl-methionine used in the present study the term K1[dl-
methionine]f is negligibly small compared with ‘1’ in the denominator of Equation (4). Hence,
the denominator tends to unity.
Similarly,
∴ [OH−]f = [OH−]T
[dl-Methionine]T = [dl-methionine]f + [dl-methionine]− + [complex].
The [dl-methinonine]f can be calculated from the equilibria (i) and (ii):
[dl-methionine]T
[dl-Methionine]f =
.
3−
1 + K1[OH−]f + K1K2[Fe(CN)6] [OH−]f
f
In view of the low [Fe(CN)6]3− used in the experiment, the term K1K2 [Fe(CN)6]3− [OH−] in
the denominator of RHS can be neglected.

184 K. Sharanabasamma and S.M. Tuwar
Thus,
[dl-methionine]T
1 + K1[OH−]f
[dl-Methionine]f =
.
(5)
(6)
On substituting Equations (3)–(5) in Equation (2), the following equation results:
3−
kK1K2[Fe(CN)6] [dl-methionine]T [OH−]T
T
Rate =
,
1 + K1[OH−]K2[dl-methionine] + K1K2[dl-methionine][OH−]
kK1K2[dl-methionine]T [OH−]T
kobs
=
.
1 + K1[OH−]f + K2[dl-methionine]f + K1K2[dl-methionine]f [OH−]f
For the verification of rate law, the subscripts ‘T ’ and ‘f ’ are omitted and hence, Equation (6)
becomes
kK1K2[dl-methionine]T [OH−]
kobs
=
.
(7)
1 + K1[OH−] + K2[dl-methionine] + K1K2[dl-methionine][OH−]
Equation (7) is rearranged into Equation (8),
1
1
1
1
1
=
+
+
+
.
(8)
kobs
kK1K2[dl-methionine][OH−] kK2[dl-methionine] kK1[OH−]
k
The mechanism as in Scheme 1 and rate law (7) are verified in the form of Equation (8) by
plotting the graphs of 1/kobs versus 1/[OH−] and 1/[dl-M]. They should be linear and found
so in Figures 1 and 3. From the slopes and intercepts of such plots, the values of k, K1 and K2
are calculated as 2.44 × 10−4 s−1, 2.9 and 55.3 dm3 mol−1, respectively, for 28◦ C. Further, these
values are used in rate law (7) at different experimental conditions as in Table 1 to regenerate kobs
.
The regenerated values are found to be in close agreement with those of experimentally observed
values. This fortifies the mechanism and rate law (7).
The outer sphere mechanism of oxidation of dl-methionine by HCF is evidenced by the large
value of energy of activation 55.7 kJ mol−1. This outer sphere complex may be more ordered than
its reactants as it is evidenced by the negative value of ꢀS= (−42 JK−1 mol−1). However, the
independency of kobs on ionic strength variation cannot be accounted. This may be due to the
direct interaction of the CN group of HCF with ‘S’ of dl-methionine, otherwise, there would be
an ionic strength effect on rate of reaction. This direct interaction between CN and ‘S’ is also
evident from the relatively small value of log A (6.0).
An increase in rate with decrease in the dielectric constant of the medium is contradictory to
the expected direction (24) as ionic species are involved in the prior equilibrium step of the slow
step as shown in Scheme 1. This may be due to more solvation of the activated complex at low
dielectric constant of the reaction media rather than that of a higher one where the reactants are
more solvated.
The [dl-methionine] was varied at four different temperatures at constant concentrations of
oxidant, OH− and ionic strength, and from the slopes and intercepts of plot of 1/kobs versus
1/[dl-methionine], the K2 values were determined for all the temperatures varied and used to
plot Van’t Hoff’s graph (i.e. log K2 versus 1/T ). From the slopes and intercept of such plots,
the thermodynamic quantities like enthalpy change (ꢀH) of the second equilibrium step (K2) of
the reaction, entropy of the reaction (ꢀS) and free energy of the reaction (ꢀG) were calculated
(Table 3). These values are compared with the values obtained for the slow step of the reaction
and show that the adduct forms fairly fast (ꢀG = −57 kJ mol−1); and it can be concluded that the
adduct formationisthermodynamically controlled ratherthan the kinetics.The ꢀH, −24 kJ mol−1
indicates that the adduct formation is exothermic and highly reactive as evidenced by the large
positive value of ꢀS (+112 JK−1 mol−1).

Journal of Sulfur Chemistry 185
40
35
30
25
20
15
10
5
0
0
2
4
6
8
10
12
1/ [OH–]dm3 mol–1
Figure 3. Verification of rate law (7) for the oxidation of dl-methionine by hexacyanoferrate(III) in aqueous alkali at
28 ◦C.
3. Conclusion
The mechanism of oxidation of dl-methionine by alkaline HCF was proposed as an outer-sphere
election transfer through an infused ligand of CN and S of dl-methionine. The complex formation
was evidenced by the rapid fall of absorbance of HCF in the presence of dl-methionine at the
initial stage of the reaction (Figure 2).
The oxidation product of dl-methionine was identified as methionine sulfoxide which is unlike
the formation of l-methionine from d-methionine in the metabolism via the formation of keto-
acid followed by re-amination. The thermodynamic parameters of formation of adduct between
HCF and dl-methionine was found to be thermodynamically controlled as evidenced by the large
negativevalueofꢀG(−57 kJ mol−1). However, thisadductmayberelativelylessactivecompared
with the adduct formed (17) between HCF and Os(VIII) in the oxidation of dl-methionine by
Os(VIII) catalyzed HCF in alkali in which the reaction occurred 1,20,000 times faster than the
uncatalyzed one even in the lower concentration of dl-methionine.
4. Experimental
All chemicals used were of reagent grade. Millipore water was used throughout the study. An
aqueous solution of HCF was prepared by dissolving K3[Fe(CN)6] (BDH) in water and was
standardized (25) iodometrically. dl-methionine, a colorless crystalline compound (E-Merck),
was used without further purification for the preparation of aqueous stock solution. Aqueous

186 K. Sharanabasamma and S.M. Tuwar
solutions of K4[Fe(CN)6] and dl-methionine sulfoxide were used to study the effect of products
on the rate of reaction. Aqueous solutions of NaOH and NaCl were used to maintain the [OH−]
and ionic strength, respectively.
4.1. Kinetics measurements
The reaction was initiated by mixing HCF solution with dl-methionine which also contained
the required amounts of NaOH and NaCl to maintain constant concentration of alkali and ionic
strength, respectively. The progress of the reaction was monitored by measuring the absorbance
of HCF using a Hitachi U-3010 spectrophotometer (Tokyo, Japan) at its absorption maximum
λmax = 420 nm in which all other species in the reaction medium at this wave length were found to
have negligible interferences. Earlier, the obedience of HCF to Beer’s law at 420 nm was studied
in the concentration range, 5.0 × 10−5–8.0 × 10−4 mol dm−3 and the molar extinction coefficient
(ε) was found to be 1060 40 dm3 mol−1 cm−1. The pseudo-first order rate constant kobs were
calculated from the slopes of log[HCF] versus time plots, which were linear up to 80% completion
of the reaction in most of the variations of concentrations of oxidant, reductant and alkali except
at the initial stage (Figure 2). Orders with respect to each reactant are determined from the slopes
of plots of log kobs versus log(conc.) except in [HCF]. The results were reproducible within 4%.
4.2. Polymerization study
The intervention of free radicals generated during the oxidation of dl-methionine with a single
equivalent oxidant, [HCF], was expected. This possibility of intervention of free radicals was
tested by adding a free radical scavenger, acrylonitrile, while the reaction was in progress. On
diluting the reaction mixture with methyl alcohol after the reaction was complete, a copious
precipitate of the polymer resulted, indicating that the oxidation occurred via the intervention of
a free radical. Earlier, it was ascertained that there was no precipitate formed either with the HCF
in alkali, dl-methionine in alkali or with OH− alone with methyl alcohol.
References
(1) Thyagarajan, B.S. Chem. Rev. 1958, 58, 439–460.
(2) Leal, J.M.; Garcia, B.; Domigo, P.L. Coord. Chem. Rev. 1998, 173, 173–178.
(3) Singh, B.; Singh, J.P. Natl Acad. Sci. Lett. (India) 1982, 5, 323–327.
(4) Singh, B.; Singh, A.K. J. Indian Chem. Soc. 1983, 60, 704–707.
(5) Natile, G.; Bordignon, E.; Cattalini, L. Inorg. Chem. 1976, 15, 246–248.
(6) Meenakshisundaram, S.; Vinothini, R. Croat. Chem. Acta 2003, 76, 75–79.
(7) Ghosh, S.; Sengupta, P.; De, G.S. Transit. Met. Chem. 1999, 24, 59–65.
(8) Read, J.F.; Wyand, A.E.H. Transit. Met. Chem. 1998, 23, 755–759.
(9) Sharanabasamma, K.; Salunke, M.S.; Tuwar, S.M. J. Sol. Chem. 2008, 37, 1217–1225.
(10) Olatunji, M.A.; McAuley, A. Can. J. Chem. 1977, 55, 3335–3338.
(11) Olatunji, M.A.; McAuley, A. Can. J. Chem. 1977, 55, 3328–3332.
(12) McAuley, A.; McCann, J.P. J. Chem. Soc. Dalton Trans. 1975, 783–788.
(13) Upadhyay, S.K.; Agrawal, M.C. J. Indian. Chem. Soc. 1981, 58, 871–875.
(14) Laloo, D.; Mahanti, M.K. Oxidation Commun. 1987, 109, 205–216 and references therein.
(15) Upadhyay, S.K. Int. J. Chem. Kinet. 1983, 15, 669–671.
(16) Chidankumar, C.S.; Chandraju, S.; Netkal, M.; Gowda, M. Synth. React. Inorg. Me. 2009, 39, 645–649.
(17) Jose, T.P.; Nandibewoor, S.T.; Tuwar, S.M. J. Sulfur Chem. 2006, 27, 25–36.
(18) Feigl, F. Spot Tests in Organic Analysis; Elsevier: New York, 1975.
(19) Goswami, K.B.; Chandra, G.; Srivastava, S.N. J. Indian Chem. Soc. 1981, 5, 252–255.
(20) Lide, D.R. Handbook of Chemistry and Physics, 73rd ed.; CRC: London, 1992, pp 8–51.
(21) Swinehart, J.H. J. Inorg. Nucl. Chem. 1967, 29, 2313–2320.

Journal of Sulfur Chemistry 187
(22) Das Gupta, G.; Mohanti, M.K. Afinidad 1984, 11, 506–511.
(23) Lancaster, J.M.; Murray, R.S. J. Chem. Soc. A 1971, 2755–2758 and references therein.
(24) Laidler, K.J. Chemical Kinetics, 3rd ed.; Pearson Education: Singapore, 2004, p 195.
(25) Mendham, J.; Denney, R.C.; Barnes, J.D.; Thomas, M.J.K. Vogel’s Text Book of Quantitative Chemical Analysis, 6th
ed.; Pearson Education: Delhi, India, 2003, p 466.