Oxidative Cleavage of S–S Bond During the Reduction of Tris(pyridine-2-carboxylato)manganese(III) by Dithionite in Sodium Picolinate–Picolinic Acid Buffer Medium

Article abstract of DOI:10.1002/kin.21022

The reduction of tris(pyridine-2-carboxylato)manganese(III) by dithionite has been investigated within the temperature window 288–303 K and at pH range 5.22–6.10 in sodium picolinate–picolinic acid buffer medium. The reaction obeys the following stoichiometry: S₂O^2- ₄ + 2Mn^III + 2H₂O → 2HS^O- ₃ + 2Mn^II + 2H⁺ The reaction is described in terms of a mechanism that involves an initial complex formation between S₂O₄ ^2? and [Mn^III(C₅H₄NCO₂)₃] followed by S–S bond cleavage to give 2HSO₃ ^? and [Mn^II(C₅H₄NCO₂)₂(H₂O)₂] as the products via the formation of SO₂ ^●? radical anion. Kinetics and spectrophotometric evidences are cited in favor of the suggested mechanism. Thermodynamic parameters associated with the equilibrium step and the activation parameters with the rate-determining step have been computed.

Full text of DOI:10.1002/kin.21022

Oxidative Cleavage of S–S
Bond During the Reduction
of Tris(pyridine-2-
carboxylato)manganese(III)
by Dithionite in Sodium
Picolinate–Picolinic Acid
Buffer Medium
KALYAN K. SEN GUPTA,1 NANDINI BHATTACHARJEE,1 BISWAJIT PAL2
1Department of Chemistry, Jadavpur University, Kolkata 700 032, India
2Department of Chemistry, St. Paul’s C. M. College, Kolkata 700 009, India
Received 29 September 2015; revised 7 June 2016; accepted 16 June 2016
DOI 10.1002/kin.21022
Published online in Wiley Online Library (wileyonlinelibrary.com).
ABSTRACT: The reduction of tris(pyridine-2-carboxylato)manganese(III) by dithionite has been
investigated within the temperature window 288–303 K and at pH range 5.22–6.10 in sodium
picolinate–picolinic acid buffer medium. The reaction obeys the following stoichiometry:
S2O24− + 2MnIII + 2H2O → 2HSO3− + 2MnII + 2H+
The reaction is described in terms of a mechanism that involves an initial complex formation
between S2O42− and [MnIII(C5H4NCO2)3] followed by S–S bond cleavage to give 2HSO3− and
[MnII(C5H4NCO2)2(H2O)2] as the products via the formation of SO2•− radical anion. Kinetics
and spectrophotometric evidences are cited in favor of the suggested mechanism. Thermody-
namic parameters associated with the equilibrꢀiCum step and the activation parameters with the
rate-determining step have been computed.
1–9, 2016
2016 Wiley Periodicals, Inc. Int J Chem Kinet
INTRODUCTION
Dithionite is a strong reducing agent [1–3] which con-
tains a S–S linkage. The oxidation of thiosulfate that
Correspondence to: B. Pal; e-mail: palbiswajit@yahoo.com.
2016 Wiley Periodicals, Inc.
ꢀC

2
GUPTA, BHATTACHARJEE, AND PAL
also contains S–S bond has been studied earlier and
shown to give tetrathionate by different oxidants such
as Ag(III) [4,5], Ir(IV) [6], and Mn(III) [7], and there
is no evidence that the S–S bond is cleaved. The thio-
sulfate ion [8] has the structure of S–SO3 with a S–S
6.25. The buffer Na(pic)-picH plays a dual role here.
One is to maintain the desired pH range, and the ma-
jor one is to restrict the possibility of ligand substi-
tution reaction that might have occurred between the
complex and a buffer other than this one. The ligand
substitution reaction is hindered due to the presence
of common picolinate ion both in the complex and in
the bulk of the reaction medium. Buffer solutions were
prepared by mixing calculated amounts of pyridine-2-
carboxylic acid (pKa = 5.52) with its sodium salt. The
pH of the solution was checked against standard buffer
solution with a pH meter (Elico LI 120, Hyderabad,
A. P., India). The pH of each buffer solution remained
unchanged for a substantial period of time.
Sodium dithionite was purchased from Loba
Chemie (Mumbai, India). Sodium dithionite was stan-
dardized by the reduction of Fe(III) to Fe(II) using a
mixture of thiocyanate and ferrocyanide as an indica-
tor [20]. A deep red coloration of ferric thiocyanate is
observed near the end-point of the titration. At the final
end-point, the solution rapidly turns blue. Freshly pre-
pared dithionite solutions were always used for the
experiments. Na2S2O4 solutions are quite stable at
pH >5 and in the concentration range (0.2–1.2) ×
10−3 mol dm−3 [21] at which the reactions were
studied.
The Mn(III) complex was prepared following the
procedures mentioned in the literature [22]. To an aque-
ous solution of MnCl2·6H2O (2.5 g in 25 cm3), H2O2
was added. A precipitate of MnO2 that was formed
upon addition of NaOH (1 mg in 25 cm3) was fil-
tered and washed with water. Excess of pyridine-2-
carboxylic acid was added, warmed, and stirred vig-
orously. The filtrate so formed deposited red crystals
of tris(pyridine-2-carboxylato)manganese(III) on cool-
ing. Thermogravimetric analysis indicated a loss of wa-
ter of crystallization in one step. The compound was
then analyzed for C, H, and N. Found C 49.95, H 3.3,
N 9.4% calcd for [MnIII(C5H4NCO2)3] H2O C 49.92,
H 3.2, and N 9.6%. The magnetic moment of the solid
compound has been found to be 4.92 BM.
˚
distance of 2.013 0.003 A. On the other hand, the
˚
S–S distance of dithionite is much larger (2.39 A) than
the disulfides and polysulfides, which lies in the region
˚
of 2.0–2.15 A. The long bond is believed to be due
to weakening by repulsion of the lone pairs on sulfur
atoms that account for its strong reducing property [9].
Considerable amount of work has been done on the
oxidation of dithionite by different oxidants. Prelim-
inary research concerning the voltammetric behavior
of sodium dithionite indicates that dithionite is oxi-
dized in two steps [10,11]. The kinetics of the oxida-
tion of dithionite by Fe(CN)63− and PtCl62− have been
reported [12,13]. The reduction of manganese(III)-
hematoporphyrin [14], toluidine blue [15] in aqueous
acidic medium, cytochrome P450 3A4 [16], and neu-
trophil cytochrome b-558 [17] by dithionite has been
studied. Salnikov et al. studied [18] spectroscopically
the reduction of cobalamine by dithionite and evidence
for the formation of a six-coordinated cobalamine(II)
complex has been presented. kobs increases with de-
creasing pH in the region 13.0–9.0. They have stud-
ied the reaction at 298 K. In another publication [19],
kinetics and mechanism of cyanocobalamine reduc-
tion by sodium hydroxomethane sulfinate (HMS) has
been reported. With increasing pH, the reaction rate
increases reaching a plateau in the pH range >12.0.
The character of the rate-determining step depends on
the concentration of the reducing agent when it is in
excess. The plot of 1/kobs against 1/[HMS] is linear at
pH 11.4 and temperature 298 K.
The present report deals with the ox-
idation
of
dithionite
by
tris(pyridine-2-
carboxylato)manganese(III) in picolinate–picolinic
acid buffer medium with a view to compare the results
obtained earlier [7] with sodium thiosulfate and other
inorganic reductants by this oxidant.
Absorbances were measured on a Systronics
(Ahmedabad, Gujrat, India) spectrophotometer. C, H,
and N analysis were performed by the Microanalytical
Laboratory using a Perkin-Elmer 240C,H,N analyzer
(Waltham, Massachusetts, U.S.A.). Thermogravimet-
ric studies were performed on a Shimadzu Corpora-
tion’s (Tokyo, Japan) TG 50 analyzer in normal atmo-
spheric environment. EPR spectra were recorded with
a Varian EPR spectrometer.
The cyclic voltammetric experiment was performed
under dry nitrogen using a Bio-Analytical-System
(BAS) CV 27 electrochemical analyzer and a BAS
model X-Y recorder at 298 K. A standard three
EXPERIMENTAL
Materials and Methods
The reagents employed were of the highest pu-
rity available, and all solutions were made in dou-
bly distilled water. The pH of the Mn(III) solu-
tion was recorded to be 5.71. The tris(pyridine-2-
carboxylato)manganese(III) complex is unstable either
in a high acidic or an alkaline medium. But the com-
plex is reasonably stable within the pH range of 5.14–
International Journal of Chemical Kinetics DOI 10.1002/kin.21022

REDUCTION OF TRIS(PYRIDINE-2-CARBOXYLATO)MANGANESE(III) BY DITHIONITE
3
electrode electrochemical cell consisting of a platinum
disk working electrode (BAS), a platinum wire auxil-
iary electrode, and a Ag/AgCl reference electrode was
used. The cyclic voltammogram of the manganese(III)
complex exhibited an irreversible response at 0.8 V, and
the corresponding oxidative response was observed at
−0.47 V (ꢀEp = 330 mV, E1/2 = –0.63 V).
After the kinetic experiment was over, the reac-
tion product was divided into three parts. Concen-
trated solution of Sr(NO3)2 when added to one part of
the reaction mixture at pH ꢀ 8 gave white precipitate
of SrSO3. The solubilization of this white precipitate
upon addition of dilute HCl ruled out the formation
of sulfate. To the second part of the reaction, the mix-
ture was added zinc acetate, potassium ferrocyanide,
and sodium nitroprusside solution. The white precipi-
tate {K2Zn3[Fe(CN)6]2} so formed turned red. When
sodium nitroprusside was added to the last part of the
reaction mixture in alkali medium, a pink color devel-
oped. All these confirmed that bisulfite not the sulfate
ion was formed as a reaction product. On the basis of
the stoichiometry and product formed, the reactions
may be represented stoichiometrically as
The electronic absorption spectrum of the man-
ganese(III) complex exhibits two weak shoulders at
approximately 457 and 403 nm (ε = 257 and 407 cm−1
mol−1 dm3). Since the extinction coefficients are low
in this region, the progress of the reaction was fol-
lowed spectrophotometrically at 350 nm (ε = 2710
cm−1 mol−1 dm3), where the complex absorbs appre-
ciably at the low concentration range of (1.0–6.0) ×
10−4 mol dm−3. Moreover, Beer’s law is obeyed in this
concentration range.
ꢀ
ꢁ
The kinetic studies were carried out under pseudo–
first-order conditions, keeping the concentration of
dithionite in large excess over that of the man-
ganese(III). Freshly prepared dithionite solutions were
always used to ensure reproducibility of the data. Req-
uisite volumes of the reactants at the appropriate tem-
perature were mixed, and the mixture was immediately
transferred to a cell of 1 cm path length. The cell com-
partment of the spectrophotometer was kept at constant
temperature. The rate of decrease of manganese(III)
was followed for at least two half-lives. Pseudo–first-
rate constants (kobs) were calculated from log A (where
A is absorbance) versus time plots. The maximum error
in the measurement of rate constant was within 5% de-
pending upon the experimental conditions. However,
for most experiments duplicate measurements were re-
producible to 3%.
S2O24− + 2 MnIII (C5H4NCO2)3 + 6H2O
ꢀ
ꢁ
→ 2HSO−3 + 2 MnII (C5H4NCO2)2 (H2O)2
+ 2C5H4NCO−2 + 2H+
(1)
The liberation of the picolinate ion during the re-
action was confirmed [12,23] by the yellow col-
oration formed on addition of a reagent containing
FeSO4·(NH4)2SO4·6H2O (1 g) and KF (0.5 g) in
100 cm3 of 0.1 N acetic acid to the reaction mixture.
The pH of the reaction mixture was checked before and
after the reaction, and it was found that the presence
of a slight excess of ligand does not alter the pH of
the solution. The formation of Mn(II) in the reaction
mixture was confirmed by the appearance of typical
six lines in the EPR spectrum.
Stoichiometry and Product Analysis
Test for Free Radicals
The stoichiometries of the reactions were studied keep-
ing [S2O42−] in excess over [MnIII] at pH 5.63 and at
room temperature (298 K). After completion of the re-
actions, the excess substrate was determined. The con-
sumption ratio (ꢀ[MnIII] /ꢀ[S2O42−]) was found to be
2.0 (Table I). The observed consumption ratios support
the formation of bisulfite as a product of oxidation of
S2O42−. The formation of bisulfite was confirmed by
the tests described below [13].
The polymerization test was carried out by adding
5 cm3 of 40% (w/v) acrylamide solution to a mix-
ture containing MnIII and dithionite in concentrations
5.0 × 10−4 and 5.0 × 10−3 mol dm−3, respectively,
at pH 5.63. The onset of polymerization could not be
identified owing to the appearance of turbidity in the
initially clear solution. However, when a large excess of
methanol was added, a polymeric suspension formed,
indicating the intervention of free radicals during the
Table I Stoichiometric Results of the Oxidation of S2O42− by Mn(III)
Initial [MnIII
]
Initial [S2O42−
]
)
Final [MnIII
]
Final [S2O42−
]
)
ꢀ [MnIII] /
(103 mol dm−3
)
(103 mol dm−3
(103 mol dm−3
)
(103 mol dm−3
ꢀ[S2O42−
]
0.30
0.40
0.50
1.00
2.00
5.00
0.00
0.00
0.00
0.85
1.81
4.76
1.99
2.13
2.08
International Journal of Chemical Kinetics DOI 10.1002/kin.21022

4
GUPTA, BHATTACHARJEE, AND PAL
reaction [24]. Cationic and anionic polymerizations
have been ruled out under the conditions at which the
reactions were studied. Blank experiments from which
either the S2O42− or MnIII was excluded gave no poly-
meric product.
RESULTS
Effect of [MnIII(C5H4NCO2)3] and [S2O42−
]
on kobs
The reactions were studied at various initial concentra-
tions of MnIII in the range (2.0–8.0) × 10−5 mol dm−3
but at constant [S2O42−], pH, and temperature of
1.0 × 10−3 mol dm−3, 5.63, and 288 K, respectively.
The average pseudo–first-order rate constant was found
to be (20.3 0.5) × 10−4 s−1. The reaction is thus first
order with respect to [MnIII].
Figure 1 Variation−of1the reaction rate with S2O42− concen-
tration. Plots of kobs versus [S2O42− −1
at four different
]
temperatures. [MnIII] = 5.0 × 10−5 mol dm−3 and pH 5.63.
Table III Values of Equilibrium Constant for the Fast
Step and Disproportionation Constant for the Slow Step
at Different Temperatures. [MnIII] = 5.0 ×
10−5 mol dm−3 and pH 5.63
The pseudo–first-order rate constants were also de-
termined (Table II) for different S2O42− concentrations
in the range (0.2–1.2) × 10−3 mol dm−3 but at con-
stant [MnIII] and pH of 4.0 × 10−5 mol dm−3 and 5.63,
respectively. When 1/kobs values were plotted against
1/[S2O42−], straight lines were obtained making inter-
cepts on the y-axis at different temperatures (Fig. 1), in-
dicating that the reaction between MnIII and S2O42− oc-
curs via an intermediate 1:1 complex formation. Again,
the abrupt increase in absorbance at 350 nm on addition
of S2O42− to MnIII in a buffer medium is also suggestive
of the formation of a complex between the reactants.
From the slope and intercept of each plot of 1/kobs
versus 1/[S2O42−], the values of Ke (the equilibrium
constant for complex formation) and kd (the dispro-
portionation constant) were calculated at four different
temperatures. The values of Ke and kd at different tem-
peratures are presented in Table III. The value of Ke
was also determined spectrophotometrically at 350 nm
Temperature (K)
Ke (dm3 mol−1
)
102 kd (s−1
)
288
293
298
303
474
502
528
585
0.621
0.854
1.18
1.49
using the relation [25–27]:
ꢀε/Ca = KeꢀεꢁIC0 − Keꢀε
(2)
where Ke is the equilibrium constant, l is optical path
length in centimeters, Ca is the concentration of the
substrate, C0 is the concentration of manganese(III),
and ꢀε is difference between the extinction coeffi-
cient of the complex and that of manganese(III). From
the slope of the plot of ꢀε/Ca versus ꢀε, the spec-
trophotometric value of Ke was determined to be 4.3 ×
102 dm3 mol−1 at 288 K.
Table II Effect of Variation of [S2O42−] on the
Pseudo–First-Order Rate Constant at Different
Temperatures. [MnIII] = 5.0 × 10−5 mol dm−3 and pH
5.63
103 kobs (s−1
)
Effect of Variation of Na(pic)-picH
103 [S2O42−
]
The effect of changing [H+] (done by varying the con-
centration ratio of picolinate–picolinic acid) on kobs
was also studied in the pH range 5.22–6.10 and at
[MnIII], [S2O42−] and temperature of 5.0 × 10−5 mol
dm−3, 5.0 × 10−4 mol dm−3, and 293 K, respectively.
Ionic strength was not maintained constant since it has
no influence on the reaction rate. The values of kobs
were found to be independent at different hydrogen
(mol dm−3
)
288 K
293 K
298 K
303 K
0.4
0.5
0.6
0.8
1.0
1.2
0.99
1.20
1.35
1.73
2.03
2.22
1.43
1.72
1.96
2.44
2.86
3.23
2.08
2.44
2.78
3.57
4.10
4.55
2.77
3.45
4.03
4.55
5.41
6.25
International Journal of Chemical Kinetics DOI 10.1002/kin.21022

REDUCTION OF TRIS(PYRIDINE-2-CARBOXYLATO)MANGANESE(III) BY DITHIONITE
5
ion concentrations. The absence of the change of rate
DISCUSSION
with change in pH of the buffer system indicates that
the concentrations of the ingredients do not have any
influence on the reaction.
The rates of decomposition of 5.0 × 10−3 mol dm−3
Na2S2O4 solution (in terms of the slopes of the
concentration–time curves) at 296 K are reported to
be 1.85 × 10−6, 0.7 × 10−6, 0.28 × 10−6, and 0.12
× 10−6 mol dm−3 s−1 at pH values of 3.5, 4.0, 4.5,
and 5.0, respectively [10]. At higher pH, the rate of de-
composition is even slower. The rate of decomposition
of 1.0 × 10−3 mol dm−3 Na2S2O4 at 296 K and pH
4.5 in the absence of any stabilizing cation was found
Activation and Thermodynamic Parameters
The reactions have been studied at different tempera-
tures in the 288–303 K regions. The pseudo–first-order
rate constant was found to increase with an increase in
temperature (Fig. 1). The thermodynamic parameters
(ꢀS° and ꢀH°) for the equilibrium step and activation
to be 0.087 × 10−6 mol dm−3
s
−1. The above results
ࣔ
ࣔ
suggest that the hydrolysis of dithionite is too slow to
be significant under the studied pH range (5.22–6.10).
Dithionous acid is a strong acid, which ionizes ac-
cording to Eqs. (5) and (6).
parameters (ꢀH and ꢀS ) for the slowest step in the
oxidation of dithionite were calculated from Eqs. (3)
and (4) as shown below.
log Ke = [ꢀS◦ − ꢀH ◦/T ] 2.303 R
(3)
K1
H2S2O4 ꢀ HS2O−4 + H+
(5)
log (k/T) = [log (kB/h) + ꢀS=/2.303R]
−ꢀH=/2.303RT
K2
(4)
H2S2O−4 ꢀ S2O42− + H+
(6)
where kB is the Boltzmann constant, h is Planck’s con-
stant, and the other terms have their usual meaning.
The values of K1 and K2 have been reported [3] to
be 4.5 × 10−1 and 3.5 × 10−3 at 298 K. Since
2-
S
S
N
O-
N
N
O
Ke
O-
N
O
O
MnIII
N
S
S
+
(7)
(8)
O
O
O
MnIII
N
O-
S
O-
O
O
O
O
2-
-
S
N
kd
slow
N
O
O
O-
N
O-
N
-
MnII
N
O
O
+
SO2 + SO2
O
O
MnIII
N
O
O
-
N
fast
[MnII(C5H4NCO2)2 (H2O)2]
N
O
+
(9)
+
2H2O
MnII
N
-
N
CO2
O
O
fast
-
H+
(10)
(11)
HSO3
+
SO2 + H2O
fast
-
+
[MnIII(C5H4NCO2)3 ]
SO2
+
[MnII(C5H4NCO2)3 ]-
SO2
Scheme 1 Overall reaction mechanism involving oxidation of dithionite by MnIII
International Journal of Chemical Kinetics DOI 10.1002/kin.21022
.

6
GUPTA, BHATTACHARJEE, AND PAL
Scheme 2 Alternative reaction mechanism involving oxidation of dithionite by MnIII
.
the reaction has been carried out in the pH range
5.22–6.10, it is likely that the S2O42− ion is the pre-
dominant reductant species, which reduces MnIII. The
oxidation of S2O42− proceeds via a 1:1 intermediate
complex formed between [MnIII(C5H4NCO2)3] and
An alternative mechanistic pathway (Scheme 2)
shows a chelate ring opening in the MnIII complex
to give a penta-coordinated species, which binds to
the S2O42− moiety. After the electron transfer, the pi-
colinate ion is expected to rechelate. This alternative
mechanistic pathway should not affect the rate expres-
sion shown below.
S2O42−
.
The MnIII complex has been reported to be a dis-
torted octahedron with the axial positions occupied by
N atoms [12]. In the intermediate complex, it is quite
plausible that the S2O42− moiety occupies the seventh
coordination site; the coordination takes place through
one of the negatively charged O atoms. The intermedi-
ate complex breaks down in the rate-determining step
to give SO2 and the radical anion SO2•−− [11] thereby re-
ducing MnIII to [MnII(C5H4NCO2)3] . The MnII com-
plex in the presence of water is likely to form the more
stable aquated form [MnII(C5H4NCO2)2(H2O)2] with
the liberation of a picolinate ion. The SO2•− under-
goes further one-electron oxidation by another MnIII
molecule to give a second molecule of SO2. Similar ob-
servation has also been reported in the kinetic studies
of the reduction of Fe(CN)63− by the dithionite ion [12]
under the stoichiometric ratio of [FeIII]:[S2O42−] =
2:1 that suggests the the dissociation of the reductant
in the rate-determining step with a subsequent rapid
reduction of another Fe(CN)63− by the SO2•− radical
ion.
The reaction obeys the following rate expression:
ꢀ
ꢁ
ꢀ
ꢁ ꢀ
ꢁ
d MnIII
kdKe MnIII S2O24−
−
=
ꢀ
ꢁ
(15)
1 + Ke S2O24−
dt
ꢀ
ꢁ
ꢀ
ꢁ
d MnIII
dt
kdKe S2O24−
1 + Ke S2O24−
1
ꢀ
ꢁ
kobs = −
=
ꢀ
ꢁ
(16)
(17)
MnIII
1
kobs
1
1
=
+
ꢀ
ꢁ
kdKe S2O24−
kd
In Scheme 2, the reaction has been shown to take
place between two negatively charged ions. The re-
actions that take place between two similarly charged
ions [28,29] are characterized by negative entropies of
ࣔ
activation. The present value of ꢀS , which is –137
13 J K−1 mol−1, supports the suggested step (13)
and mechanism. Again oxygen-free dithionite solution
shows a strong ESR signal, which indicates that S2O42−
The mechanism suggested for the oxidation of
dithionite by [MnIII(C5H4NCO2)3] is shown in
Scheme 1.
International Journal of Chemical Kinetics DOI 10.1002/kin.21022

REDUCTION OF TRIS(PYRIDINE-2-CARBOXYLATO)MANGANESE(III) BY DITHIONITE
7
dissociates to give SO2•−
:
S2O24− ꢀ 2 SO•2−
(18)
The equilibrium was used [30] to study the sul-
fur dioxide radical ion SO2•− during the oxidation
of dithionite by Co(III), Fe(III), and Mn(III) com-
plexes. Since [SO2•−] = K1/2 [S2O42− 1/2
] , the abso-
lute rate constant was calculated using the relation
k = kobs/K1/2[S2O42− 1/2
] . On the other hand, during
the reduction of electron transfer proteins and other
compounds of biological interest by dithionite, the
equilibrium constant of S2O42− ꢀ 2SO2− has been
determined [31] by electron paramagnetic resonance
spectroscopy. The equilibrium constant between the
monomer form and the dimer form of dithionite has
been calculated using different buffers such as sodium
pyrophosphate, sodium phosphate, and glycinate in the
pH range 7.25–10.4 at constant ionic strength (0.8 M).
The values of K have been found to be in the range
(0.85–1.5) × 10−19 M. A reasonable value of the equi-
librium constant (K) of the reaction in pyrophosphate
buffer at pH 8.0 is (1.4 0.4) × 10−9 M, which has
been utilized in calculating the second-order rate con-
stant for the reaction between the oxidant and SO2−.
The authors [30,31] have calculated rate constants in
alkaline medium unlike the present investigation which
has been carried out in acid medium.
Figure 3 Influence of temperature on the disproportiona-
tion constant. Plot of log (kd/T) versus 1/T. [MnIII] = 5.0 ×
10−5 mol dm−3 and pH 5.63.
Table IV Experimental and Calculated
Pseudo–First-Order Rate Constant Values at 288 K
104 [S2O42−
]
(mol dm−3
)
103 kobs (expt) (s−1 103 kobs (Calcd) (s−1 a
) )
4.0
5.0
6.0
8.0
10.0
12.0
0.99
1.20
1.35
1.73
2.03
2.22
0.91
1.10
1.27
1.59
1.87
2.11
The plots as shown in Fig. 1 justify the derived rate
equation (15). The enthalpy change (ꢀH°) associated
with the equilibrium step was calculated from the slope
of the least-square plot of log Ke versus 1/T (Fig. 2).
The corresponding entropy change ꢀS° was calculated
from Eq. (3). The values of ꢀH° and ꢀS° are 9.0
aCalculated kobs values are obtained using the spectrophotometric
value of Ke of 430 dm−3 mol−1 at 288 K.
2 kJ mol−1 and –20 7 J K−1 mol−1, respectively. The
ࣔ
enthalpy of activation (ꢀH ) was calculated to be 43
4 kJ mol−1 from the slope of the least-squares plot of
log(kd/T) versus 1/T (Fig. 3). The entropy of activation
ࣔ
(ꢀS ) was found to be –137 13 J K−1 mol−1 using
Eq. (4).
The value of Ke obtained from spectrophotometric
measurement has been utilized to calculate kobs us-
ing the deduced rate expression. The experimental and
calculated values are presented in Table IV. It is to be
mentioned that freshly prepared solutions of dithionite
were used at the beginning of the spectrophotometric
experiment to minimize decomposition of dithionite
solution. Despite this precaution some decomposition
may also occur. However, the observed and calculated
values of kobs are not widely different thereby justifying
the suggested mechanism and rate expression.
We have earlier investigated the oxidation of some
inorganic reducing substrates by this oxidant [32–35].
Figure 2 Influence of temperature on the equilibrium con-
stant. Plot of log (Ke/T) versus 1/T. [MnIII] = 5.0 ×
10−5 mol dm−3 and pH 5.63.
ࣔ
ࣔ
The values of ꢀH and ꢀS for these reactions have
been mentioned in our recent report [7]. The values of
International Journal of Chemical Kinetics DOI 10.1002/kin.21022

8
GUPTA, BHATTACHARJEE, AND PAL
Table V Values of Thermodynamic Parameters for the Oxidation of Different Inorganic Reductants by Mn(III) in Acid
Medium
ࣔ
Reductant
N3− [32]
Reaction
E0 (V) [36]
ꢀG0 (kJ mol−1
)
ꢀG (kJ mol−1
)
3N2 + 2H+ + 2e ꢁ 2HN3
−3.1
598.3
87.8
3N2 + 2e ꢁ 2N3−
−3.4
NH2OH [33]
H3PO2 [34]
N2 + 2H2O+ 4H+ + 2e ꢁ 2NH2OH·H+
H3PO3 + 2H+ + 2e ꢁ H3PO2 + H2O
H3PO4 + 2H+ + 2e ꢁ H3PO3 + H2O
SO42− + 4H+ + 2e ꢁ H2SO3 + H2O
S4O62− + 2e ꢁ 2S2O32−
−1.87
−0.50
−0.276
0.172
0.08
360.9
96.5
53.3
−33.2
−15.4
10.8
37.2
92.3
98.3
61.7
72.5
82.5
H3PO3 [34]
SO32− [35]
S2O32− [7]
S2O42− (present study)
2H2SO3 + H+ + 2e ꢁ HS2O4−
−0.06
tetrathionate and sulfite, respectively. It has been re-
ported [41] that the ferrate ion (HFeO4−), which is a
powerful two electron oxidant (E0 = 1.9 V) and an oxy-
gen transfer agent, rapidly oxidizes S2O32− to SO32− in
acid medium when the reductant is in excess to that of
oxidant. This is because S2O32− is oxidized initially to
give OSSO32− and Fe(IV) followed by the reaction of
OSSO32− with another molecule of the oxidant to give
Fe(IV) and S2O52−. The latter finally gives SO32− by
reacting rapidly with water as mentioned below.
HFeO−4 + S2O23− → Fe (IV) + OSSO23− + H+ (19)
HFeO−4 + OSSO32− → Fe (IV) + S2O25− + H+ (20)
Figure 4 Correlation plot of the thermodynamic parame-
ters in the oxidation of different inorganic compounds by
MnIII. Plot of ꢀG# versus ꢀG0.
S2O52− + H2O → 2SO23−
(21)
ࣔ
ꢀG have been calculated (Table V) using the relation
Such a possibility that SO32− is formed in the oxida-
tion of S2O32− by manganese(III) has been ruled out
since Mn(III) is a less powerful oxidant than HFeO4−.
It is pertinent to compare the oxidation of S2O32− and
S2O42− by MnIII as an oxidant. There are few similar-
ities in these two reactions. Both the reactions occur
through intermediate formation of free radicals and
also through the inner sphere reduction process. The
formation of MnII in both these reactions has been
confirmed by the appearance of typical six-line EPR
spectrum. On the other hand, there is no evidence that
the S–S bond is broken in the oxidation of thiosulfate
where the substrate is oxidized to give tetrathionate and
no complex formation occurs prior to electron transfer
unlike the oxidation of dithionite where the reaction
occurs through intermediate formation of complex be-
tween the two reactants followed by the decomposition
to give products. The activation parameters obtained in
the oxidation of dithionite are different from thiosul-
fate oxidation there by supporting different mechanism
and S–S bond cleavage.
ࣔ
ࣔ
ࣔ
ꢀG = ꢀH – TꢀS at 288 K. If ꢀG0 is the standard
free energy of the elementary electron transfer step in
the prevailing electrolyte medium and 2 is the number
of electrons transferred, the values of ꢀG0 (ꢀG0 =
–2FE0) have been calculated using the known value
[36] of E0 and are presented in Table V. For weak in-
teraction (neglible overlap between the orbitals in the
activated state), Marcus theory predicts [37–39] a slope
of ½ for the ꢀG versus ꢀG0 plot. This was later on
ࣔ
verified by Nikolov [40] for a number of redox reac-
tions. We have also plotted ꢀG versus ꢀG0 (Fig. 4).
ࣔ
The plot is linear for the reactions when the redox
potentials are in the lower region (–0.06 to 0.5 V),
whereas for higher E0 values, e.g., 2N3− → 3N2 + 2e
(–3.4 V) and NH2OH → N2 + 2e (–1.87 V) the plot
is nonlinear (Fig. 4). However, theoretically predicted
slope of ½ has been observed from the linear portion
of the plot.
Both S2O32− and S2O42−, which contain S–S link-
age, have been oxidized by different oxidants to give
International Journal of Chemical Kinetics DOI 10.1002/kin.21022

REDUCTION OF TRIS(PYRIDINE-2-CARBOXYLATO)MANGANESE(III) BY DITHIONITE
9
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International Journal of Chemical Kinetics DOI 10.1002/kin.21022