Kinetics and Mechanism of Oxidation of S2O32? by a Co-Bound μ-Amido-μ-Superoxo Complex

Article abstract of DOI:10.1002/kin.20973

In acetate buffer media (pH 4.5–5.4) thiosulfate ion (S₂O₃^2?) reduces the bridged superoxo complex, [(NH₃)₄Co^III(μ-NH₂,μ-O₂)Co^III(NH₃)₄]⁴⁺ (1) to its corresponding μ-peroxo product, [(NH₃)₄Co^III(μ-NH₂,μ-O₂)Co^III(NH₃)₄]³⁺ (2) and along a parallel reaction path, simultaneously S₂O₃^2? reacts with 1 to produce the substituted μ-thiosulfato-μ-superoxo complex, [(NH₃)₄Co^III(μ-S₂O₃,μ-O₂)Co^III(NH₃)₄]³⁺ (3). The formation of μ-thiosulfato-μ-superoxo complex (3) appears as a precipitate which on being subjected to FTIR shows absorption peaks that support the presence of Co(III)-bound S-coordinated S₂O₃^2? group. In reaction media, 3 readily dissolves to further react with S₂O₃^2? to produce μ-thiosulfato-μ-peroxo product, [(NH₃)₄Co^III(μ-S₂O₃,μ-O₂)Co^III(NH₃)₄]²⁺ (4). The observed rate (k₀) increases with an increase in [T_Thio] ([T_Thio] is the analytical concentration of S₂O₃^2?) and temperature (T), but it decreases with an increase in [H⁺] and the ionic strength (I). Analysis of the log A_t versus time data (A is the absorbance of 1 at time t) reveals that overall the reaction follows a biphasic consecutive reaction path with rate constants k₁ and k₂ and the change of absorbance is equal to {a₁ exp(–k₁t) + a₂ exp(–k₂t)), where k₁ > k₂.

Full text of DOI:10.1002/kin.20973

Kinetics and Mechanism
2−
of Oxidation of S2O3
by a Co-Bound
μ-Amido-μ-Superoxo
Complex
BULA SINGH,1 RANENDU SEKHAR DAS,2 RUPENDRANATH BANERJEE,2
SUBRATA MUKHOPADHYAY2
1Department of Chemistry, Visva-Bharati, Santiniketan, 731 235, India
2Department of Chemistry, Jadavpur University, Kolkata, 700 032, India
Received 5 April 2015; revised 5 November 2015; accepted 15 November 2015
DOI 10.1002/kin.20973
Published online 21 December 2015 in Wiley Online Library (wileyonlinelibrary.com).
ABSTRACT: In acetate buffer media (pH 4.5–5.4) thiosulfate ion (S2O32−) reduces the bridged
superoxo complex, [(NH3)4CoIII(μ-NH2,μ-O2)CoIII(NH3)4]4+ (1) to its corresponding μ-peroxo
product, [(NH3)4CoIII(μ-NH2,μ-O2)CoIII(NH3)4]3+ (2) and along a parallel reaction path, simul-
taneously S2O32− reacts with 1 to produce the substituted μ-thiosulfato-μ-superoxo complex,
[(NH3)4CoIII(μ-S2O3,μ-O2)CoIII(NH3)4]3+ (3). The formation of μ-thiosulfato-μ-superoxo com-
plex (3) appears as a precipitate which on being subjected to FTIR shows absorption peaks
2−
that support the presence of Co(III)-bound S-coordinated S2O3 group. In reaction media,
2−
3 readily dissolves to further react with S2O3 to produce μ-thiosulfato-μ-peroxo product,
[(NH3)4CoIII(μ-S2O3,μ-O2)CoIII(NH3)4]2+ (4). The observed rate (k0) increases with an increase
in [TThio] ([TThio] is the analytical concentration of S2O32−) and temperature (T), but it decreases
with an increase in [H+] and the ionic strength (I). Analysis of the log At versus time data (A
is the absorbance of 1 at time t) reveals that overall the reaction follows a biphasic consec-
utive reaction path with rate constants k1 anꢀdC k2 and the change of absorbance is equal to
{a1 exp(–k1t) + a2 exp(–k2t)}, where k1 > k2.
2015 Wiley Periodicals, Inc. Int J Chem Kinet
48: 88–97, 2015
INTRODUCTION
Metal–dioxygen complexes have long been the sub-
ject of intensive investigations due to their relevance to
the physiological metabolism of dioxygen (O2) [1–5].
During the metabolic process, stepwise reduction of
Correspondence to: Bula Singh; e-mail: singhbula@rediffmail.
com, R. S. Das; ranendu147@gmail.com.
Supporting Information is available in the online issue at
www.wileyonlinelibrary.com.
ꢀC
2015 Wiley Periodicals, Inc.

2−
OXIDATION OF S2O3 BY A Co-BOUND μ-AMIDO-μ-SUPEROXO COMPLEX
89
O2 through electron-transfer reactions leads to the pro-
duction of highly reactive oxygen species (ROS) like
of the title work will offer an insight to the under-
standing of the prospective roles of sulfur-containing
small molecules, such as S2O32−, thiosalicylic acid, N-
acetyl cysteine, etc. as an effective antioxidant against
the toxic ROS species.
•−
2−
O2 and O2 [6–9]. In living systems, different en-
zymes, such as xanthine oxidase [10,11], lipoxyge-
nase, and cyclooxygenase [12,13], spontaneously but
in limited extent produce ROS species in electron-
rich environment. However, when the production of
ROS species exceeds the limit, a state called oxida-
tive stress arises which threats the integrity of various
functional biomolecules, such as proteins [14], lipids
and lipoproteins [15], and DNA [16]. The same ROS
which can be a “good Samaritan” when they are strate-
gically directed to destroy cancer cells [17–20], under
the condition of chronic oxidative stress, becomes re-
sponsible for diseases, such as diabetes [21,22], cardio-
vascular [23] and neurodegenerative diseases [24,25],
tumorigenesis [19], or the damage of mitochondrial
DNA causing cellular aging [23,26–28].
EXPERIMENTAL
Materials, Instrumentation, and
Preparation of Complex 1
The metallo-superoxide complex, viz. μ-amido-
μ-superoxo-bis[tetraaminecobalt(III)]
tetranitrate,
[(NH3)4CoIII(μ-NH2,μ-O2)CoIII(NH3)4](NO3)4 (1),
and the aqueous solution of NaS2O3 were prepared fol-
lowing the literature processes [42,43]. The details of
the materials used for the title study is provided in the
Supporting Information. Complex 1 was characterized
from its FTIR spectrum (Fig. S1 in the Supporting
Information; details are given in the Supporting
Information), and the purity of 1 was determined from
its characteristic UV–vis absorption at 700 nm (ε in
In pursuit of understanding the ROS chemistry, nu-
merous metal-bound-oxygen complexes were synthe-
sized, among which cobalt-bound superoxo complexes
are renowned since they are easy to study at ambient
conditions [29]. Taube, Sykes, and others have reported
the detailed studies on the reduction of these complexes
with different metal ions [30,31] and reducing agents
such as sulfite, arsenite, nitrite, etc. [32–34]. In this
present work, we have chosen such a complex ion, viz.
μ-amido-μ-superoxo-bis[tetraaminecobalt(III)]4+ (1)
as a model for ROS species and have studied the redox
reaction between 1 and thiosulfate anion (S2O32−).
The thiosulfate ion, S2O32−, in which one of the
two sulfur exists in the zero oxidation state can act as a
M
−1 cm−1, found 300; reported 306 [42]). The details
of the instrumental techniques used for the title study
is also provided in the Supporting Information.
Kinetics
Kinetic experiments were performed in acetic acid
buffer media in the pH range of 4.5–5.4 at ambient tem-
perature. Under this reaction condition, both complex
2−
2−
1 and S2O3 are stable and specially, S2O3 is not
2−
reducing agent and thus also as an antioxidant. S2O3
decomposed to other sulfurous products [44,45]. On
2−
finds several medicinal applications, like in treatment
of cis-platin induced nephrotoxicity, as an antidote for
extravasation of chemotherapy agents [35], in treat-
ment of ototoxicity and cochlear damage caused by
the action of aminoglycoside antibiotics [36] and in
the addition of S2O3 solution to the aqueous acidic
solution of 1 at constant ionic strength (I) condition, the
absorbance of 1 at 700 nm is gradually decreased and
such a representative UV–vis spectral series of change
in the absorbance of 1 with time is shown in Fig. 1. In
general, kinetics were studied in acidic buffer media at
25.0 0.1°C and at a fixed ionic strength, I = 0.50 M,
maintained with NaNO3, unless mentioned otherwise.
A large excess of [TThio] ([TThio] is the analytical con-
2−
calciphylaxis [41]. The mechanism, by which S2O3
reduces nephrotoxicity and ototoxicity, has been pro-
posed to involve the process of free-radical scaveng-
ing and/or covalent binding to the platinum compound
[37–40].
2−
centration of S2O3 added in reaction solution) is
The potential role of S2O32− as a radical scavenger
to the best of our knowledge is yet to be studied in
detail. Our present work that describes the interaction
maintained over [1] in all the kinetic runs to ensure
pseudo–first-order conditions. The reacting solutions
were prepared by the addition of measured volume of
buffered solution of 1, and sodium nitrate followed
by the addition of freshly prepared aqueous solution of
S2O32−. Generally, adventitious presence of metal ions
2−
of S2O3 with a superoxo complex (1) interestingly
brings forefront both the capacities of S2O32− anion as
a radical scavenger as well as a bridging ligand. The de-
tailed analyses of the experimental observation reveal
that S2O32− not only reduces the bound-superoxo moi-
ety of 1 but also through a consecutive reaction path
forms an intermediate bridging thiosulfato complex.
We believe that the proposed mechanistic elucidation
like Cu2+ is known to oxidize S2O3 [46] and thus,
2−
to prevent such interference, a well-known sequester-
ing agent, dipicolinic acid (dpa; 2.0 mM) was added in
the reaction media to completely suppress the effects
of such metal ions. Here it is noteworthy that none of
International Journal of Chemical Kinetics DOI 10.1002/kin.20973

90
SINGH ET AL.
= 10.0–40.0 mM, no perceptible precipitate is formed
during the reaction and it is evident from the evenness
of the spectral decay series shown in Fig. 1. However,
if [1] ࣙ 5.0 mM, then on additions of S2O32− a green-
ish precipitate is quickly formed in the reaction media,
which within very short time spontaneously redissolves
on standing. The precipitate was quickly separated,
washed with cold water–ethanol mixture and vacuum
dried and then it was subjected to FTIR studies. The
details of IR studies and the characterization of the
precipitate are further discussed in the next section.
0.5
0.4
0.3
0.2
0.1
0.0
1
12
F
400
500
600
Wavelength (nm)
700
800
Stoichiometry
Stoichiometric determinations were carried out by es-
Figure 1 Observed spectral decay of 1 with a reaction of
S2O32−. Spectral series of 1–12 were recorded at an interval
of 20 s. The final spectrum F was recorded after 24 h. Con-
dition: [1] = 1.0 mM; [TThio] = 10.0 mM [dpa] = 2.0 mM;
[H+] = 3.2 × 10−5 M; I = 0.5 M; T = 25.0 0.1°C.
timating the excess [1] or excess [S2O32−] from the
2−
complete reaction of complex 1 and S2O3 in differ-
ent molar ratios. Excess [1] was determined from the
difference in absorption values of pure 1 (A0) and the
equilibrium absorption values of the reaction mixtures
(Ae) at 700 nm. Remaining [S2O32−] was quantified
titrimetrically with potassium iodide following stan-
dard procedures [43]. The titrimetric results are pro-
vided in Table S1 in the Supporting Information.
Any gas evolved during the reaction was collected
by the downward displacement of water, and the mea-
sured volume of the gas is corrected to NTP (1.0 atm,
273.15 K).
-0.4
-0.8
-1.2
-1.6
-2.0
2−
The oxidized products from S2O32−, mainly S4O6
and other polythionates, are unstable in acid media
2−
and usually are converted to SO4 in part [44]. The
2−
presence of SO4 in the product mixtures from the
reactions of either excess [1] over [S2O32−] or excess
[S2O32−] over [1] was verified from the precipitation
of BaSO4 on treatment with HCl and BaCl2 [43]. Both
polythionates, present in the former case, and excess
0
50
100
150
200
250
time (s)
Figure 2 Plot of log10 At with time (where A is the ab-
sorbance at 700 nm). Condition: [1] = 1.0 mM; [TThio] =
10.0 mM; [dpa] = 2.0 mM; [H+] = 3.2 × 10−5 M; I =
0.5 M; T = 25.0 0.1°C.
S2O32−, present in the latter case, can produce SO4
under acidic condition. Thus the above test for S4O6
2−
2−
is very much qualitative. For the same reason, gravi-
2−
metric quantification of SO4 was unsuccessful and
2−
the complex 1 and S2O3 reacts with dpa. This was
the obtained results show randomness. In the product
mixture, though the presence of S4O62− was confirmed
by the spot test with SbCl3 [47] but again the interfer-
verified from by monitoring the constancy of ab-
sorbance of 1 in the presence of dpa over at least 2 h
2−
2−
and from the titrimetric determination of S2O3 con-
ences from the traces of remaining S2O3 cannot be
centration of the dpa-S2O32− mixture.
completely ignored [47].
The reaction kinetics was followed by the decrease
in absorbance at 700 nm (Fig. 2), the characteristic
visible absorption peak of 1, where all other species in
the reaction mixture are transparent. The observed rate
constants (k0) were derived using linear standard least-
square fit to the corresponding log10At versus time data
(where A is the absorbance at time t, vide infra).
Here it is important to mention that under the ex-
RESULTS AND DISCUSSION
Stoichiometry and Reaction Products
Titrimetric results shown in Table S1 in the Sup-
porting Information suggest that 1 mol of 1 is re-
2−
perimental conditions when [1] = 1.0 mM and [TThio
]
duced by 1 mol of S2O32−. The reductant S2O3 is
International Journal of Chemical Kinetics DOI 10.1002/kin.20973

2−
OXIDATION OF S2O3 BY A Co-BOUND μ-AMIDO-μ-SUPEROXO COMPLEX
91
oxidized to S4O62−, which is further decomposed to
2−
80
60
40
20
0
produce mainly SO4 and other products [44,45].
2−
Since S2O3 is an one electron donor, the super-
2−
oxo complex 1 is directly reduced by S2O3 to
1630
its corresponding peroxo complex [(NH3)4CoIII(μ-
NH2,μ-O2)CoIII(NH3)4]3+ (2). Moreover, S2O3
2−
1342
1094
also through the substitution reaction produces
μ-thiosulfato-μ-superoxo complex, [(NH3)4CoIII(μ-
S2O3,μ-O2)CoIII(NH3)4]3+ (3), which is further re-
993
653
3261 3150
2−
duced by S2O3 to μ-thiosulfato-μ-peroxo complex,
[(NH3)4CoIII(μ-S2O3,μ-O2)CoIII(NH3)4]2+ (4) (vide
infra for details). The peroxo complexes 2 and 4, how-
ever in acidic media, fast isomerize to the correspond-
ing hydroperoxo complexes [(NH3)4CoIII(μ-NH2,μ
-HO2)CoIII(NH3)4]4+ (5) and [(NH3)4CoIII(μ-S2O3,μ-
HO2)CoIII(NH3)4]3+ (6), respectively. The near-zero
absorbance at 700 nm of the spectrum number 12 in
Fig. 1 suggests that complex 1 is completely reduced to
its corresponding peroxo products 2 and 4. Spectrum
12 also bears a resemblance with the final spectrum
reported by Mandal et al. [48]. The similarity sug-
gests the formation of the μ-peroxo products 2 and 4
from 1.
4000
3000
2000
cm-1
1000
Figure 3 FTIR spectrum of the precipitate.
the bridging NH2 group by S2O32−. It is not a redox
change.
The assumption behind the formation of complex
3 is supported by the fact that when S2O3 reacts
−
2−
with the similar Co-bound peroxo complex [51,52] or
2−
SO3 reacts with Co-bound superoxo complex [53],
corresponding μ-thiosulfato and μ-sulfato complexes
are produced. Crystal structures of μ-thiosulfato com-
plex [51,52] also reveals that two of the cobalt ions
are bridged by the sulfur atom of the thiosulfato group.
Further supports are provided by the FTIR spectrum of
the precipitate (Fig. 4). Apart from the regular peaks
for the superoxo group (1094 cm−1 in contrast to 1086
cm−1 for pure 1, Fig. S1 in the Supporting Information)
and NH3 groups (3261–3150, 1630, and 1342 cm−1 in
contrast to 3279–3152, 1628, and 1370 cm−1 for pure
1, Fig. S1 in the Supporting Information), it presents
two new peaks at 993 and 653 cm−1. These two char-
acteristic peaks account for the symmetric stretching
In acid media, 5 readily mineralizes to produce
Co2+, O2, and NH3 [31,49]. As NH4 is a weak acid
+
(pKa = 9.2) [50], in the experimental pH range (4.5–
5.4) most of the NH3 produced will remain in solution
+
as solvated NH4 ion, and only O2 will evolve as the
gaseous product. The results from several stoichiomet-
ric runs show that the volume of the evolved O2 gas
is approximately molar corrected to all other factors.
Similar to 5, we assume that the other μ-hydroperoxo
complex 6 was also mineralized to Co2+, O2, S2O3
,
2−
and NH4+ in acid media. The complete mineralization
of 5 and 6 is further supported by the final spectrum F
(Fig. 1) of the reaction mixture taken after 24 h.
Thus, the stoichiometric equation representing the
main redox reaction is given as follows:
2−
of the SO3 group of the S-coordinated S2O3 group
[54], and these are in close similarity with reported
peaks of the Co(III)-bound S2O32− group at ꢀ990 and
635 cm−1 [55]. The bridging S2O32− group also shows
an important peak [56] at the wavelength range of
1175–1130 cm−1, but unfortunately that peak cannot
be identified since complex 1 also shows the peak in
this region. However, it is noteworthy that the relatively
small IR peak at 1086 cm−1 for pure 1 (Fig. S1 in
the Supporting Information) becomes much stronger
in Fig. 3 and has shifted to 1094 cm−1. We sus-
pect that the IR absorption for the bridging S2O3
group reinforces the IR peak of 1 at this wavelength
range.
The simultaneous presence of peaks for the super-
oxo group and S-coordinated S2O3 group firmly
1
4+ + 33+ 2S2O32− = 23+ + 42+ + S4O62− (1)
Characterization of the Precipitate
The buffered solutions of mere 1 ([1] ࣙ 5.0 mM) or
2−
S2O3 in experimental pH are highly stable and do
2−
not form any precipitate on long standing (at least 4 h).
Thus we suspect that the appearance of the precipitate
during the reaction actually arises due to the formation
of the μ-thiosulfato product, [(NH3)4CoIII(μ-S2O3,μ-
O2)CoIII(NH3)4]3+ (3) on substitution of the bridging
NH2 group of 1 by S2O32−. Here it is important to note
that the charge of the complex changes from +4 in 1 to
+3 in 3. The change occurs due to the substitution of
2−
supports the assumption that during the reaction the
μ-thiosulfato-μ-superoxo complex 3 was formed,
which on further reaction redissolves.
International Journal of Chemical Kinetics DOI 10.1002/kin.20973

92
SINGH ET AL.
0.4
0.0
0.4
0.0
log10A'
-0.4
-0.8
-1.2
-1.6
-2.0
-0.4
-0.8
-1.2
0
20
40
60
80
100
0
50
100
150
200
250
time (s)
time(s)
Figure 5 Plot of the log10(Aꢁ – At) versus time. Condition:
[1] = 1.0 mM; [TThio] = 10.0 mM; [dpa] = 2.0 mM; [H+]
= 3.2 × 10−5 M; I = 0.5 M; T = 25.0 0.1°C.
Figure 4 Plot of log10 At versus time and the extrapo-
lated straight line, log10 A´ at longer time. Condition: [1] =
1.0 mM; [TThio] = 10.0 mM; [dpa] = 2.0 mM; [H+] =
3.2 × 10−5 M; I = 0.5 M; T = 25.0 0.1°C.
time in Fig. 3, which clearly consists of two linear
parts: an initial part and a second part at longer time.
It can be supposed that one slow and one compara-
tively fast reaction is occurring simultaneously, and as
a result such a curved plot is yielded. One can ex-
trapolate the second linear part at longer time to zero
and can set the 2.303 times of the slope value of this
line to the rate constant for the slow reaction [58]. In
Fig. 4, we have denoted this extrapolated straight
line as log10Aꢁ. Then the logarithm of the difference,
(Aꢁ – At) is plotted in Fig. 5 from which the value
of the rate constant for the fast reaction can be ob-
tained as the 2.303 times of the slope value [58].
Though the constants k1 and k2 in Eq. (3) are inter-
changeable and they cannot be certainly associated
with any particular reaction step in the reaction se-
quence (Eq. (2)) [60], for our convenience we asso-
ciate k1 and k2 with the faster and slower reaction,
respectively.
Analysis of the Kinetic Data
From Fig. 2, it is evident that the reaction between 1
and S2O32− follows a complex kinetic trace. This type
of biphasic kinetic profile has been simply described
by a two-step consecutive reaction scheme as shown in
Eq. (2) [57–60]
k1
k2
A → B → C
(2)
where k1 and k2 are the pseudo–first-order rate con-
stants. The concentration of the species A, B, and C,
i.e. [A], [B] and [C], at any time (t) can be expressed in
terms of the initial concentration of A as [A]0, and the
overall rate expression can be expressed in terms of the
absorption (A) of the reaction mixture at any suitably
chosen wavelength (for derivation, see the Supporting
Information). The integrated expression for a consec-
utive reaction (Eq. (2)) is given by Eq. (3):
Following this analytical method described herein,
the rate constants k1 and k2 have been extracted from
the data set associated with the kinetic runs.
At − A = a1 exp(−k1t) + a2 exp(−k2t) (3)
∞
where At and Aꢁ are the absorbance values at time t
and time infinity [58].
Variation of Rate Constants with [TThio
and [1]
]
With the help of Eq. (3), one can extract the rate
constants from the plot of log10 (At – Aꢁ) versus time
data [57–60]. In our present case, the reaction was
monitored at 700 nm and the absorbance value at in-
finite time at this wavelength can be supposed to be
zero since Fig. 2 shows that just after 220 s of time
the absorbance at 700 nm has reached a near-zero
value (absorbance 0.006 in comparison to the initial
value of 0.271, Fig. 1). Thus we have simply plotted
the common logarithm of absorbance, log10At against
Plots of k1 and k2 versus [TThio] at constant [1] (Figs. 6
and 7, respectively; Table S2 in the Supporting In-
formation) are linear. This is in accordance with the
trend of the first-order reactions where rate constants
increase proportionally with the increase in [TThio].
Moreover, each of the plots passes through the origin,
which suggests that complex 1 does not autodecom-
pose under the experimental conditions.
International Journal of Chemical Kinetics DOI 10.1002/kin.20973

2−
OXIDATION OF S2O3 BY A Co-BOUND μ-AMIDO-μ-SUPEROXO COMPLEX
93
120
80
40
0
120
80
40
0
0
10
20
30
40
0
1
2
3
105[H+] (M)
[TThio] (mM)
Figure 8 Variation of rate constant k1 with [H+]. Condi-
tion: [1] = 1.0 mM; [TThio] = 10.0 mM; [dpa] = 2.0 mM;
I = 0.5 M; T = 25.0 0.1°C.
Figure 6 Variation of rate constant k1 with [TThio] . Condi-
tion: [1] = 1.0 mM; [dpa] = 2.0 mM; [H+] = 3.2 × 10−5 M;
I = 0.5 M; T = 25.0 0.1°C.
Variation of Rate Constants with Media
Acidity
120
80
40
0
Figures 8 and 9 present the plots of k1 and k2 against
the media proton concentration, [H+]. It is evident
that with the increasing proton concentration, both
the rate constants k1 and k2 decrease (Table S3 in
the Supporting Information). Since the μ-superoxo
group of 1 does not take part in any protic equilib-
rium [29,61], the observation can only be explained
by considering the dissociation constants of thiosulfu-
ric acid (H2S2O3) [62]. As the concentration of [H+]
2−
increases, the free concentration of S2O3 decreases
0
1
2
3
and subsequently the reaction rates also decrease. Plots
105[H+] (M)
Figure 9 Variation of rate constant k2 with [H+]. Condi-
tion: [1] = 1.0 mM; [TThio] = 10.0 mM; [dpa] = 2.0 mM;
I = 0.5 M; T = 25.0 0.1°C.
120
80
40
0
in Figs. 8 and 9 thus suggest that the reactive reduc-
tant is the free S2O3 anion instead of its protonated
forms.
2−
However, though the above explanation seems satis-
factory but it cannot exactly explain that why a nominal
decrease in the concentration of free S2O32− results in
the substantial lowering of the observed rate constant
values (as much as 50%). We presume that the drop
in the concentration of free S2O32−, however small it
is, actually affects the rate law rather more intricately.
As the rate law cannot be further simplified, the exact
quantitative relation responsible for this observation
cannot be realized.
0
10
20
30
40
[TThio] (mM)
Figure 7 Variation of rate constant k2 with [TThio]. Condi-
tion: [1] = 1.0 mM; [dpa] = 2.0 mM; [H+] = 3.2 × 10−5 M;
I = 0.5 M; T = 25.0 0.1°C.
International Journal of Chemical Kinetics DOI 10.1002/kin.20973

94
SINGH ET AL.
oxidants, i.e. complexes 1, 2, and 3 are cationic, the
active reductant must then bear negative charges and
in the present case it is the S2O32− anion.
-1.0
-1.2
-1.4
-1.6
-1.8
Probable Reaction Mechanism
The results of the kinetic analyses suggest that
though the title reaction involves only the oxidant,
the μ-superoxo complex, [(NH3)4CoIII(μ-NH2,μ-
O2)CoIII(NH3)4]4+ (1), and the reductant, S2O3
,
2−
the reaction proceeds through a biphasic consecu-
2−
tive reaction path. The reductant, S2O3 is an one-
electron donor and thus it reduces complex 1 to its
corresponding peroxo complex, [(NH3)4CoIII(μ-NH2,
μ-O2)CoIII(NH3)4]3+ (2) [30,31,53], which in acid me-
dia fast isomerizes to the corresponding hydroperoxo
complex [(NH3)4CoIII(μ-NH2,μ-HO2)CoIII(NH3)4]4+
(5). Now the formation of the precipitate indi-
0.5
0.6
0.7
I1/2 (M1/2
0.8
0.9
1.0
)
Figure 10 Variation of log10 k1 with I1/2. Condition: [1]
= 1.0 mM; [TThio] = 10.0 mM; [dpa] = 2.0 mM; [H+] =
3.2 × 10−5 M; T = 25.0 0.1°C.
2−
cates that along with reduction, S2O3 simultane-
ously substitutes the bridging NH2 group to produce
the intermediate complex, [(NH3)4CoIII(μ-S2O3,μ-
O2)CoIII(NH3)4]3+ (3). These two simultaneous par-
allel reactions comprise the first phase of the biphasic
reaction.
Variation of Rate Constants k1 and k2 with
Ionic Strength (I)
The effects of the change of media ionic strength (I) on
the observed rate constants were investigated by vary-
ing the concentration of NaNO3 in the reaction media
keeping all other reaction conditions otherwise same.
Changes in I have significant effects on the rates of the
reaction (Table S4 in the Supporting Information), and
the plots of log10k1 and log10k2 versus I1/2 (Figs. 10
and 11, respectively) are linear with a negative slope
(–1.1 0.2; r = 0.96 and –1.1 0.1; r = 0.98, respec-
tively). The outcome suggests a reaction between two
species with opposite charges. Since all the reactive
The μ-amido-μ-hydroperoxo complex 5 is un-
stable in acid media, and it is mineralized to
Co2+, NH4+, and O2 [49]. Thus for the second
phase of the reaction we look forward to the
reaction of the μ-thiosulfato-μ-superoxo complex,
[(NH3)4CoIII(μ-S2O3,μ-O2)CoIII(NH3)4]3+ (3), with
2−
S2O3 by which the precipitate redissolves in the so-
2−
lution. S2O3 reduces the intermediate complex 3 to
the corresponding μ-peroxo product, [(NH3)4CoIII(μ-
S2O3,μ-O2)CoIII(NH3)4]2+ (4), which isomerizes to its
corresponding μ-hydroperoxo product, [(NH3)4CoIII
(μ-S2O3,μ-HO2)CoIII(NH3)4]3+ (6), and ultimately
like complex 5 mineralizes in the acidic media.
Here it is noteworthy that the spectral feature of the
title reaction (Fig. 1) strongly supports the formation
of the hydroperoxo complexes 5 and 6. Figure 1 shows
that during the reaction a peak builds up at ꢀ500 nm,
which according to Mori and Weil [63] is due to the
isomerization of the peroxo complexes 2 and 4 to 5 and
6, respectively.
-1.2
-1.4
-1.6
-1.8
-2.0
The probable reaction mechanism describing all
these reactions is shown in Scheme 1. The respective
rate constants for the reactions 1 to 2, 1 to 3, and 3 to
4 (Eqs. (4), (5), and (6), respectively) are kꢁ, kꢁꢁ, and
kꢁꢁꢁ. The μ-peroxo complexes 2 and 4 isomerize in acid
media to 5 and 6 (Eqs. (7) and (8), respectively), which
finally decompose (Eqs. (9) and (10), respectively) to
Co2+, NH4+, O2, and S2O32−. The reductant, S2O32− is
oxidized to S4O62−, which in acid media decomposes
mainly to SO42− (Eq. (11)) along with other sulfurous
0.5
0.6
0.7
0.8
0.9
1.0
I1/2 (M1/2
)
Figure 11 Variation of log10 k2 with I1/2. Condition: [1]
= 1.0 mM; [TThio] = 10.0 mM; [dpa] = 2.0 mM; [H+] =
3.2 × 10−5 M; T = 25.0 0.1°C.
International Journal of Chemical Kinetics DOI 10.1002/kin.20973

2−
OXIDATION OF S2O3 BY A Co-BOUND μ-AMIDO-μ-SUPEROXO COMPLEX
95
2-
[(NH3)4CoIII(µ-NH2,µ-O2)CoIII(NH3)4]3+ + 1/2S4O6
k'
[(NH3)4CoIII(µ-NH2,µ-O2)CoIII(NH3)4]4+ + S2O32- + H+
(4)
(2)
(1)
k''
[(NH3)4CoIII(µ-S2O3,µ-O2)CoIII(NH3)4]3+ + NH3
(3)
[(NH3)4CoIII(µ-NH2,µ-O2)CoIII(NH3)4]4+ + S2O3
(1)
2-
(5)
(6)
k'''
[(NH3)4CoIII(µ-S2O3,µ-O2)CoIII(NH3)4]2+ + 1/2S4O6
(4)
2-
[(NH3)4CoIII(µ-S2O3,µ-O2)CoIII(NH3)4]3+ + S2O3
(3)
2-
[(NH3)4CoIII(µ-NH2,µ-HO2)CoIII(NH3)4]4+
(5)
(7)
(8)
[(NH3)4CoIII(µ-NH2,µ-O2)CoIII(NH3)4]3+ + H+
(2)
[(NH3)4CoIII(µ-S2O3,µ-HO2)CoIII(NH3)4]3+
(6)
[(NH3)4CoIII(µ-S2O3,µ-O2)CoIII(NH3)4]2+ + H+
(4)
(9)
[(NH3)4CoIII(µ-NH2,µ-HO2)CoIII(NH3)4]4+ + 9H+
(5)
2Co2+ + 9NH4+ + O2
(10)
(11)
2Co2+ + 8NH4+ + O2 +S2O3
2-
[(NH3)4CoIII(µ-S2O3,µ-O2)CoIII(NH3)4]3+ + 7H+
(6)
3S4O62- + 2H2O
7S + 5SO42- + 4H+
Scheme 1
O
O
3+
k'
CoIII(NH3)4
(NH3)4CoIII
mineralization
2-
NH2
S2O3
2 (B)
O
O
4+
(NH3)4CoIII
CoIII(NH3)4
NH2
(A)
1
O
O
k'''
2+
O
O
3+
CoIII(NH3)4
(NH3)4CoIII
CoIII(NH3)4
(NH3)4CoIII
k''
S2O3
2-
S2O3
S2O3
2-
S2O3
4
(C)
(B')
3
(k' + k'')
k'''
(B (2), B' (3))
C (4)
A (1)
Scheme 2
2−
prehend. As described earlier, the first phase consists
of two parallel reactions, 1 to 2 and 1 to 3 and the
second phase consists of one reaction, 3 to 4. If we
compare the phases with the steps of the consecutive
reaction shown in Eq. (2), the rate constant of the first
step, k1, is actually a sum of k´ and k´´ and k2 is simply
products [44]. The presence of SO4 as mentioned
earlier has been qualitatively detected but could not be
quantified (vide supra).
The reactions shown in Scheme 1 are further ar-
ranged in Scheme 2 in such a way that the phases of
the overall consecutive reaction become easier to com-
International Journal of Chemical Kinetics DOI 10.1002/kin.20973

96
SINGH ET AL.
(Fig. 5) we can obtain the values of log10a1 and log10a2
as 0.434 0.051 and 0.399 0.005, respectively.
-2
-3
-4
-5
-6
Effect of Temperature on the Rate Constant
The title reaction was also studied at different tem-
peratures and the variations in rate constant values are
summarized in Table S5 in the Supporting Information.
The Arrhenius plots of lnk1 and lnk2 versus 1/T are lin-
ear (Figs. 12 and 13, respectively). The corresponding
activation energies derived from the k1 and k2 values
are 89.5 7.3 and 85.4 5.3 kJ mol−1, respectively.
3.2
3.3
3.4
3.5
103/T (K-1)
CONCLUSION
Figure 12 Plot of lnk1 versus 1/T. Condition: [1] = 1.0
mM; [TThio] = 10.0 mM; [dpa] = 2.0 mM; [H+] = 3.2 ×
10−5 M; I = 0.5 M.
2−
In acetate buffer media, the S2O3
ion reacts
with the μ-superoxo complex, [(NH3)4CoIII(μ-NH2,μ-
O2)CoIII(NH3)4]4+ (1), through a biphasic consecu-
2−
tive reaction path. S2O3 reduces complex 1 by one-
electron transfer to the corresponding μ-peroxo com-
-4
-5
-6
plex, [(NH3)4CoIII(μ-NH2,μ-O2)CoIII(NH3)4]3+ (2)
2−
and simultaneously S2O3
substitutes the μ-NH2
group of 1 to produce the μ-thiosulfato-μ-superoxo
complex, [(NH3)4CoIII(μ-S2O3,μ-O2) CoIII(NH3)4]3+
(3). Intermediate 3 can be isolated as a precipitate,
and FTIR data of 3 confirm the presence of the S-
2−
coordinated S2O3 group. Complex 3 further reacts
2−
with S2O3 to produce the μ-thiosulfato-μ-peroxo
complex, [(NH3)4CoIII(μ-S2O3,μ-O2)CoIII(NH3)4]2+
(4). Complexes 3 and 4 belong to a novel series of
Co-bound superoxo and peroxo compounds with the
3.2
3.3
3.4
3.5
103/T (K-1)
bridging S2O32− group, and the title work offers a new
2−
insight in the area of coordination chemistry of S2O3
.
Figure 13 Plot of lnk2 versus 1/T. Condition: [1] = 1.0 mM;
[TThio] = 10.0 mM; [dpa] = 2.0 mM; [H+] = 3.2 × 10−5
M; I = 0.5 M.
The present work also describes the antioxidant prop-
2−
erty of S2O3 against the harmful ROS, which is
of considerable importance in the light of its medical
use.
kꢁꢁꢁ. Thus Eq. (3) can also be rewritten as Eq. (12).
Authors thankfullyacknowledge financial help receivedfrom
the UPE II Fund of Jadavpur University sanctioned by the
University Grants Commission, New Delhi, India.
At − A = a1 exp[−(kꢁ + kꢁꢁ)t] + a2 exp(−kꢁꢁꢁt)
∞
(12)
BIBLIOGRAPHY
Equation (12) cannot be further explored, and hence
the rate constants kꢁ and kꢁꢁ cannot be further separated
since the molar absorption coefficients for the inter-
mediates 2, 3, and 4 (for details, see the Supporting
Information) are not surely known. For this same rea-
son, it is also difficult to reconstruct the kinetic trace
from the rate constants and molar absorption coeffi-
cients values. However, from the intercept values of the
plots of log10At versus time (Fig. 4) and log10 (Aꢁ - At)
1. Klotz, I. M.; Kurtz, D. M., Jr. Chem Rev 1994, 94,
567.
2. Martell, A. E.; Sawyer, D. T. Oxygen Complexes and
Oxygen Activations by Transition Metals; Plenum Press,
New York, 1988.
3. Niederhoffer, E. C.; Timmons, J. H.; Martell, A. E.
Chem Rev 1984, 84, 137.
4. Gong, Y.; Zhou, M. Chem Rev 2009, 109, 6765.
International Journal of Chemical Kinetics DOI 10.1002/kin.20973

2−
OXIDATION OF S2O3 BY A Co-BOUND μ-AMIDO-μ-SUPEROXO COMPLEX
97
5. Tine´, M. R. Coord Chem Rev 2012, 256, 316.
6. Sharma, P.; Jha, A. B.; Dubey, R. S.; Pessarakli, M. J
Bot 2012, doi:10.1155/2012/217037.
7. Fridovich, I. Ann N Y Acad Sci 1999, 893, 13.
8. Betteridge, D. J. Metabolism 2000, 49, 3.
9. Halliwell, B. Free Radic Res 1999, 31, 261.
10. Halliwell, B. Annu Rev Nutr 1996, 16, 33.
11. Kuppusamy, P.; Zweier, J. L. J Biol Chem 1989, 264,
9880.
39. Dickey, D. T.; Wu, Y. J.; Muldoon, L. L.; Neuwelt, E.
A. J Pharmacol Exp Ther 2005, 314, 1052.
40. Muldoon, L. L.; Pagel, M. A.; Kroll, R. A.; Brummett,
R. E.; Doolittle, N. D.; Zuhowski, E. G.; Egorin, M. J.;
Neuwelt, E. A. Clin Cancer Res 2000, 6, 309.
41. Leitao, D. J.; Blakley, B.W. J Otolaryngol 2003, 32,
146.
42. Davies, R.; Mori, M.; Sykes, A. G.; Weil, J. A. Inorg
Synth 1982, 12, 206.
43. Vogel, A. I. In Vogel’s Textbook of Quantitative Inor-
ganic Analysis, 5th ed; Jeffery, G. H.; Basett, J.; Mend-
ham, J.; Denny, R. C., Eds.; Longman Scientific & Tech-
nical: London, 1989, 390–393.
44. Mizoguchi, T.; Takei, Y.; Okabe, T. Bull Chem Soc Jpn
1976, 49, 70.
45. Dinegar, R. H.; Smellie, R. T.; La Mer, V. K. J Am Chem
Soc 1951, 73, 2050.
46. Byerley, J. J.; Fouda, S. A.; Rempel, G. L. J Chem Soc,
Dalton Trans 1973, 889.
12. Zimmerman, B. J.; Granger, D. N. Am J Med Sci 1994,
307, 284.
13. Kontos, H. A.; Wei, E. P.; Ellis, E. F.; Jenkins, L. W.;
Povlishock, J. T.; Rowe, G. T.; Hess, M. L. Circ Res
1985, 57, 142.
14. Stadtman, E. R.; Levine, R. L. Ann N Y Acad Sci 2000,
899, 191.
15. Yla-Herttuala, S. Ann N Y Acad Sci 1999, 874, 134.
16. Marnett, L. J. Carcinogenesis 2000, 21, 361.
17. Watson, J. Open Biol 2013, 3, 120.
47. Feigl, F.; Anger, V. Spot Tests in Inorganic Analysis,
6th Ed.; Elsevier: Amsterdam, the Netherlands, 2012.
48. Mandal, K.; Banerjee, R.; Mukhopadhyay, S. Inorg
Chim Acta 2013, 394, 691.
49. Parry, R. W. Inorganic synthesis, Vol. 12; McGraw-Hill:
New York, 1970.
50. Seybold, P. G. Int J Quantum Chem 2008, 108, 2849.
51. Yamanari, K.; Mori, M.; Dogi, S.; Fuyuhiro, A. Chem
Lett 1993, 1855.
52. Yamanari, K.; Mori, M.; Dogi, S.; Fuyuhiro, A. Inorg
Chem 1994, 33, 4807.
18. Trachootham, D.; Alexandre, J.; Huang, P. Nat Rev Drug
Discovery 2009, 8, 579.
19. Schumacker, P. T. Cancer Cell 2006, 10, 175.
20. Luo, J.; Solimini, N. L.; Elledge, S. J. Cell 2009, 136,
823.
21. Ma, Z. A. Curr Diabetes Rev 2012, 8, 69.
22. Victor, V. M.; Rocha, M.; Herance, R.; Hernandez-
Mijares, A. Curr Pharm Des 2011, 17, 3947.
23. Haigis, M. C.; Yankner, B. A. Mol Cell 2010, 40,
333.
24. Andersen, J. K. Nat Med 2004, 10 (Suppl), S18.
25. Li, Q.; Spencer, N. Y.; Pantazis, N. J.; Engelhardt, J. F.
J Biol Chem 2011, 286, 40151.
53. Edwards, J. D.; Yang, C.; Sykes, A. G. J Chem Soc,
Dalton Trans 1974, 1561.
54. Fischmann, A. J.; Warden, A. C.; Black, J.; Spiccia, L.
Inorg Chem 2004, 43, 6568.
26. Cadenas, E.; Davies, K. J. Free Radic Biol Med 2000,
29, 222.
55. Murdock, A. R.; Tyree, T.; Otterbein, W.; Kinney, L.;
Carreras, M.; Cooper, J. N.; Elder, R. C. Inorg Chem
1985, 24, 3674.
56. Nakamoto, K. Infrared and Raman Spectra of Inorganic
and Coordination Compounds, Part B, 6th ed.; Wiley:
New York, 1997.
57. Gray, E. T., Jr.; Taylor, R. W.; Margerum, D. W. Inorg
Chem 1977, 16, 3047.
58. Ashley, K. R.; Hamm, R. E. Inorg Chem 1966, 5, 1645.
59. Espenson, J. H.; Chao, T. Inorg Chem 1977, 16, 2553.
60. Maji, D.; Banerjee, R. Transition Met Chem 2001, 26,
544.
27. Finkel, T.; Holbrook, N. J. Nature 2000, 408, 239.
28. Harman, D. J Gerontol 1956, 11, 298.
29. Sykes, A. G.; Weil, J. A. Prog Inorg Chem 1970, 13, 1.
30. Sykes, A. G. Trans Faraday Soc 1963, 59, 1325.
31. Hoffman, A.; Taube, H. Inorg Chem 1968, 7, 1971.
32. Davies, K. M.; Sykes, A. G. J Chem Soc 1971, 1418.
33. Davies, R.; Sykes, A. G. J Chem Soc 1968, 2831.
34. Sykes, A. G. Chem Br 1974, 10, 170.
35. Valentine, J. S.; Gralla, E. B. PNAS 2008, 105, 8178.
36. Hochman, J.; Blakley, B. W.; Wellman, M.; Blakley, L.
J Otolaryngol 2006, 35, 153.
37. Schlieper, G.; Brandenburg, V.; Ketteler, M.; Floege, J.
Nat Rev Nephrol 2009, 5, 539.
38. Neuwelt, E. A.; Brummett, R. E.; Remsen, L. G.; Kroll,
R. A.; Pagel, M. A.; McCormick, C. I.; Guitjens, S.;
Muldoon, L. L. Cancer Res 1996, 56, 706.
61. Richens, D. T.; Sykes, A. G. J Chem Soc, Dalton Trans
1982, 1621.
62. Page, F. M. J Chem Soc 1953, 1719.
63. Mori, M.; Weil, J. A. J Am Chem Soc 89, 6364, 1967.
International Journal of Chemical Kinetics DOI 10.1002/kin.20973