Kinetics and mechanism of the decomposition of tetrathionate ion in alkaline medium

Article abstract of DOI:10.1021/ic700992u

The kinetics of the alkaline decomposition of tetrathionate has been studied spectrophotometrically in the pH range of 9.2-12.2 using both phosphate/hydrogen phosphate and carbonate/hydrogen carbonate as buffer systems and by adjusting the ionic strength to I = 0.5 M with sodium perchlorate at T = 25.0 ± 0.1°C. Matrix rank analysis of the spectra recorded between 265 and 330 nm shows the presence of three independent absorbing species. Besides tetrathionate, thiosulfate and trithionate are identified as absorbing products of the decomposition, but sulfite and a trace amount of sulfate are also formed during the alkaline degradation process. With pentathionate, sulfoxylic acid (S(OH)₂), S₂O₃OH^-, and S₃O₃OH^- as key intermediates, a 10-step kinetic model is proposed with six fitted kinetic parameters to take all the important characteristics of the experimental curves into account. On the basis of the stoichiometric measurements and the model proposed, it is also enlightened that the product distribution of the reaction continuously varies with pH; thus the kinetic traces cannot be evaluated by assuming a single stoichiometry in contrast to recent studies. Buffer dependence of the decomposition is also discussed.

Full text of DOI:10.1021/ic700992u

Inorg. Chem. 2007, 46, 7654−7661
Kinetics and Mechanism of the Decomposition of Tetrathionate Ion in
Alkaline Medium
De´nes Varga and Attila K. Horva´th*
Department of Physical Chemistry, UniVersity of Szeged, P.O. Box 105, Szeged H-6701, Hungary
Received May 22, 2007
The kinetics of the alkaline decomposition of tetrathionate has been studied spectrophotometrically in the pH range
of 9.2 12.2 using both phosphate/hydrogen phosphate and carbonate/hydrogen carbonate as buffer systems and
by adjusting the ionic strength to I 0.5 M with sodium perchlorate at T 25.0 0.1 C. Matrix rank analysis
−
)
)
±
°
of the spectra recorded between 265 and 330 nm shows the presence of three independent absorbing species.
Besides tetrathionate, thiosulfate and trithionate are identified as absorbing products of the decomposition, but
sulfite and a trace amount of sulfate are also formed during the alkaline degradation process. With pentathionate,
sulfoxylic acid (S(OH)2), S2O3OH-, and S3O3OH- as key intermediates, a 10-step kinetic model is proposed with
six fitted kinetic parameters to take all the important characteristics of the experimental curves into account. On the
basis of the stoichiometric measurements and the model proposed, it is also enlightened that the product distribution
of the reaction continuously varies with pH; thus the kinetic traces cannot be evaluated by assuming a single
stoichiometry in contrast to recent studies. Buffer dependence of the decomposition is also discussed.
Introduction
Degradation of polythionates in alkaline medium seems to
be a promising solution to this problem.
The leaching of copper, silver, and gold from their ores
with thiosulfate seems to be a less harmful alternative method
than the traditional cyanidation process. As a result, this
promising method has been extensively studied by several
research groups.1-3 It is already known that metal ions can
be readily brought into solution from their ores as metal-
thiosulfate complexes, but recovery of the metal from the
complexes continues to be a problem. As a possible solution,
usage of strong base ion exchange resins has been suggested
very recently.4 It has been found, however, that polythionates
(mainly tri- and tetrathionate) formed in the oxidation of
thiosulfate by metal ions5,6 may easily cause4 the exchange
resin to foul. It is therefore crucial to remove the polythion-
ates from the leach prior to the ion exchange process.
The alkaline decomposition of tetrathionate has been
studied since the beginning of the last century. Gutman7,8
and Riesenfeld9 have independently reported that if tetrathion-
ate was left for a long time in a strong alkaline medium,
sulfite and thiosulfate were finally formed through the
following equation.
2S4O62- + 6OH- f 3S2O32- + 2SO32- + 3H2O (1)
Kuertenacker et al.10-12 and Goehring13 have later shown that
the decomposition strongly depends on the alkalinity of the
solution. Initially, tetrathionate decomposes into trithionate
and pentathionate according to the following equation.
2-
2S4O62- f S3O62- + S5O6
(2)
* To whom correspondence should be addressed. E-mail: horvatha@
chem.u-szeged.hu.
(1) Zipperian, D.; Raghavan, S. Hydrometallurgy 1988, 19, 361.
(2) Tao, J.; Jin, C.; Shi, X. XVIII International Mineral Processing
Congress, Sydney, May 23-28, 1993; p 1141,
(3) Abbruzzese, C.; Fornari, P.; Massidda, R.; Veglio, F.; Ubaldini, S.
Hydrometallurgy 1995, 39, 265.
(7) Gutman, A. Ber. 1906, 39, 509.
(8) Gutman, A. Ber. 1907, 40, 3614.
(9) Riesenfeld, E. H. Z. Anorg. Allg. Chem. 1924, 141, 109.
(10) Kurtenacker, A.; Kaufmann, M. Z. Anorg. Allg. Chem. 1925, 148,
43.
(4) Zhang, H.; Dreisinger, D. B. Hydrometallurgy 2002, 66, 59.
(5) Naito, K.; Yoshida, M.; Shieh, M.; Okabe, T. Bull. Chem. Soc. Jpn.
1970, 43, 1365.
(6) Byerley, J. J.; Fouda, S. A.; Rempel, G. L. J. Chem. Soc., Dalton
Trans. 1972, 889.
(11) Kurtenacker, A.; Kaufmann, M. Z. Anorg. Allg. Chem. 1925, 148,
369.
(12) Kurtenacker, A.; Mutschin, A.; Stastny, F. Z. Anorg. Allg. Chem. 1935,
224, 339.
(13) Goehring, M. Fortschr. Chem. Forsch. 1952, 2, 444.
7654 Inorganic Chemistry, Vol. 46, No. 18, 2007
10.1021/ic700992u CCC: $37.00
© 2007 American Chemical Society
Published on Web 07/31/2007

Decomposition of Tetrathionate Ion in Alkaline Medium
Table 1. Initial Composition of the Kinetic Runsa
no.
[PO34-
]
0
[HPO24-
]
0
[S4O2-
]
0
[S2O2-
]
0
3
6
1-7
52.5
52.5
50.6
38.5
23.1
16.5
16.5
10.5
10.5
16.5
38.5
69.3
82.5
82.5
0.3, 0.5, 0.7, 1.0, 1.4, 2.0, 3.0
1.0
0
8-13
0.3, 0.5, 0.7, 1.0, 2.0, 3.0
0
0
0
14-20
21-27
31-34
35-41
42-47
0.3, 0.5, 0.7, 1.0, 1.4, 2.0, 3.0
0.3, 0.5, 0.7, 1.0, 1.4, 2.0, 3.0
0.3, 0.5, 0.7, 1.0, 1.4, 2.0, 3.0
0.3, 0.5, 0.7, 1.0, 1.4, 2.0, 3.0
1.0
0
0.3, 0.5, 0.7, 1.0, 2.0, 3.0
No.
[CO23-
]
0
[HCO-]0
[S4O2-
]
[S2O2-
]
0
3
0
3
6
48
49
50-52
53-55
56
67.5
63.0
50.25
50.25
35.0
7.5
21.0
50.25
50.25
105
0.6
0.6
0.3, 0.6, 1.0
0.6
0.6
0
0
0
0.3, 0.6, 1.0
0
a All concentrations are given in mM.
In very weak alkaline solutions (pH < 9.0), both trithionate
and pentathionate are fairly stable toward further degradation.
If, however, the pH slightly increases, pentathionate is no
longer stable and decomposes according to the following
equation14
One of the latest studies on the alkaline decomposition of
tetrathionate has shown4 that the stoichiometry of the reaction
is even more complicated than was thought because sulfide
ion is also formed during the degradation, but this result is
argued by a recent study of Breuer and Jeffrey.19
The aim of our paper is to set up a consistent kinetic model
for the alkaline decomposition of tetrathionate that is able
to take the changes of the stoichiometry into account properly
and to provide a reliable starting point for the design of
maximizing the recovery of thiosulfate from the leach after
the alkaline treatment of the polythionates. To date no
comprehensive kinetic and mechanistic studies are yet
available in the literature.
2S5O62- + 6OH- f 5S2O32- + 3H2O
(3)
resulting in an overall process shown below.
4S4O62- + 6OH- f 2S3O62- + 5S2O32- + 3H2O (4)
If the pH increases further (>12.0), disproportionation of
trithionate starts to rise5 according to equation
S3O62- + 2OH- f S2O32- + SO42- + H2O
(5)
Experimental Section
Materials and Buffers. Potassium trithionate was prepared by
the reaction of potassium thiosulfate (Riedel-de Hae¨n) with sulfur
dioxide.20 Its purity was checked21 and found to be better than
98.5%. All other chemicals were of the highest purity, commercially
available (potassium tetrathionate (Fluka), sodium thiosulfate,
sodium phosphate, sodium hydrogen phosphate, sodium hydrogen
carbonate, and sodium carbonate (Reanal), sodium perchlorate
(Sigma-Aldrich)), and used without further purification. Four times
distilled water was used to prepare all the stock solutions.
Two series of experiments were conducted in different buffer
compositions. In the first series of experiments the pH was kept
constant by using phosphate/hydrogen phosphate buffer (pKa )
11.75), while in the other series carbonate/hydrogen carbonate buffer
(pKa ) 9.7) was used.35 The ionic strength was also kept constant
(I ) 0.5 M) by adding the necessary amount of sodium perchlorate
to the buffer components. Table 1 contains the composition of the
56 kinetic runs.
or at pH)13 per equation.
2S3O62- + 6OH- f S2O32- + 4SO32- + 3H2O (6)
It means that the overall degradation can be described by
either of the following processes
2S4O62- + 6OH- f 3S2O32- + 2SO32- + 3H2O (7)
4S4O62- + 10OH- f 7S2O32- + 2SO42- + 5H2O (8)
It has also been found that thiosulfate ion catalyzes the
decomposition of tetrathionate in neutral solution15-17 through
the following steps.
S4O62- + S2O32- h S5O62- + SO3
S4O62- + SO32- f S3O62- + S2O3
(9)
2-
Methods and Instrumentation. The kinetic measurements were
carried out in a standard 1 cm quartz cuvette equipped with a Teflon
cap and a magnetic stirrer. First, the buffer solution (I ) 0.5 M)
2-
(10)
Later Rolia and Chakrabarti18 have shown that thiosulfate
catalysis also occurs at higher pH (10-12), where eq 4 is
the dominating stoichiometry. The rate equation for the
consumption of tetrathionate at T ) 35 °C was found
to be
(14) Wagner, H.; Schreier, H. Phosphorus Sulf. Relat. Elem. 1978, 4, 281.
(15) Goehring, M.; Helbing, W.; Appel, I. Z. Anorg. Chem. 1947, 254,
185.
(16) Foss, O. Acta Chem. Scan. 1949, 3, 1385.
(17) Fava, A.; Bresadola, S. J. Am. Chem. Soc. 1955, 77, 5792.
(18) Rolia, E.; Chakrabarti, C. L. EnViron. Sci. Technol. 1982, 16, 852.
(19) Breuer, P. L.; Jeffrey, M. I. Hydrometallurgy 2004, 72, 335.
(20) Brauer, G. Handbook of PreparatiVe Inorganic Chemistry, 2nd ed.;
Academic Press: New York, 1963; Vol. 1, p 398.
2-
d[S4O6
]
-
) (0.1 + 12.6[S2O32-])[S4O62-][OH-] (11)
dt
(21) Koh, T. Anal. Sci. 1990, 6, 3.
Inorganic Chemistry, Vol. 46, No. 18, 2007 7655

Varga and Horva´th
Table 2. Determination of the Formed Thiosulfate and Trithionate
Concentration Versus the Initial Tetrathionate Concentration after the
Reaction Is Completed in Phosphate Buffera
pH
[S4O2-
]
SR2
SR3
SI
SA
0
6
10.8
10.8
10.8
10.8
10.8
10.8
10.8
12.2
12.2
12.2
12.2
12.2
12.2
12.2
11.0
11.5
12.0
0.003
0.002
0.0014
0.001
0.0007
0.0005
0.0003
0.003
0.002
0.0014
0.001
0.0007
0.0005
0.0003
0.002
1.285
1.285
1.286
1.290
1.297
1.305
1.325
1.365
1.390
1.410
1.425
1.440
1.454
1.470
1.285
1.310
1.360
0.460
0.455
0.450
0.440
0.425
0.405
0.360
0.280
0.230
0.190
0.157
0.126
0.097
0.065
0.450
0.390
0.285
0.020
0.040
0.056
0.080
0.112
0.160
0.260
0.420
0.520
0.600
0.672
0.736
0.796
0.860
0.060
0.200
0.420
0.030
0.025
0.022
0.020
0.019
0.015
0.010
0.010
0.010
0.010
0.007
0.006
0.005
0.005
0.020
0.010
0.005
Figure 1. Molar absorbances of tetrathionate (dotted line), trithionate (solid
line), and thiosulfate (dashed line) used in the calculation procedure.
0.002
0.002
was delivered into the cuvette followed by the thiosulfate solution
(I ) 0.5 M adjusted by sodium perchlorate if necessary). Finally,
the tetrathionate solution (I ) 0.5 M adjusted by sodium perchlorate)
was added from a fast delivery pipet to start the reaction. The
reaction was followed with a Zeiss Specord S10 diode array
spectrophotometer within the 240-500 nm wavelength range. Each
solution was illuminated only 300 times so as to minimize the
photochemical decomposition of tetrathionate.23 A separate experi-
ment showed that the final spectra of the illuminated solution agreed
well, within the experimental error, with the final spectra of the
unilluminated solution of the same composition. Since the spectra
of tetrathionate, thiosulfate, and trithionate are overlapping, it is
crucial to determine the molar absorptivity of these species precisely
for the simultaneous curve-fitting method. Thus, the molar absor-
bance of tetrathionate as well as that of trithionate and thiosulfate
was calculated for each wavelength from three to eight different
reagent concentrations. The calculated spectra of tetrathionate,
thiosulfate, and trithionate depicted in Figure 1 agree satisfactorily
with the data obtained by Lorenz and Samuel.24 Sulfite and sulfate
do not absorb the light over 265 nm at all.
For qualitative determination of the end products, a Bio-Rad
Digilab Division dedicated FT-Raman spectrometer was used with
laser-type Nd:YVO4 at 1064 nm. After the decomposition had been
completed, the solvent was evaporated under vacuum at 35 °C.
The Raman spectrum of the solid material was registered for
identification of the end products.
Data Treatment. For the Zeiss Specord S10 diode-array
spectrophotometer, the constant uncertainty in the absorbance
measurement is (0.003 AU. As the absorbance increases beyond
unity, the transmitted light intensity decreases and circuit noise
begins to limit the precision with which the absorbance can be
measured. For this reason we set an upper limit of 1.0 AU beyond
which no kinetic data was used for further data analysis.
a SR2 and SR3 are defined as [S2O2-]∞/[S4O2-]0 and [S3O2-]∞/[S4O2-]0,
3
6
6
6
respectively. SI and SA, defined as [SO23-]∞/[S4O2-
]
0
and [SO24-]∞/[S4
6
O26-]0, respectively, are Calculated with the help of eqs 12 and 13.
are 265, 270, 275, 280, 285, 290, and 295 nm. The reason for
selecting these wavelengths was to gain as much information as
possible about tetrathionate and the products since no wavelength
could be found where any of the absorbing species had exclusive
contribution. Since each kinetic curve contained 300 data pairs,
the number of points in each run was reduced to 50 to avoid
unnecessary and time-consuming calculations. The essence of the
reduction method has already been described elsewhere.27 Alto-
gether, almost 19600 experimental points from 56 absorbance-
time series (each containing 50 points) were used for simultaneous
fitting. The sum of squares of the deviations between the measured
and calculated absorbances was selected as the parameter to be
minimized. Our quantitative criterion for an acceptable fit was that
the average deviation for the absolute fit approach 0.003 AU, which
is close to the uncertainty of the spectrophotometer.
Results
Stoichiometry. Identification of the end products, includ-
ing classical analytical procedures (gravimetry, iodometry)
and Raman spectroscopy, has clearly shown that thiosulfate,
trithionate, and sulfite were the main products of the
decomposition but trace amount of sulfate could also be
detected in a relatively wide pH range of 9.2-12.2. Since
trithionate and thiosulfate absorb light in the wavelength
range of 265-330 nm, the ratio of the formed trithionate
over consumed tetrathionate concentration (SR3) as well as
that of the formed thiosulfate over consumed tetrathionate
concentration ratio (SR2) could be determined from the final
spectra of the solutions. After the reaction had been
completed, the thiosulfate and the trithionate concentrations
were determined from the final spectrum of the solution at
the wavelength range of 265-295 nm. With the final
concentration values of thiosulfate and trithionate, one can
easily calculate the final concentrations of sulfite and sulfate
taking the mass and oxidation number balance into consid-
eration according to the following equations.
Matrix rank analysis studies25 have been executed on those parts
of the wavelength range of 240-500 nm where the absorbance
did not exceed 1.0 AU. The experimental curves were analyzed
with the program package ZiTa.26 Seven different wavelengths have
been chosen for simultaneous fitting of the kinetic data and they
(22) IUPAC Stability Constant Database; Royal Society of Chemistry:
London, 1992-1997.
(23) Horva´th, A. K.; Nagypa´l, I.; Epstein, I. R. J. Am. Chem. Soc. 2002,
124, 10956.
(24) Lorenz, L.; Samuel, R. Z. Phys. Chem. 1932, B14, 218.
(25) Peintler, G.; Nagypa´l, I.; Jancso´, A.; Epstein, I. R.; Kustin, K. J. Phys.
Chem. A 1997, 103, 8013.
(26) Peintler, G. ZiTa, version 5.0; a comprehensive program package for
fitting parameters of chemical reaction mechanism; Attila Jo´zsef
University: Szeged, Hungary, 1989-1998.
(27) Horva´th, A. K.; Nagypa´l, I.; Epstein, I. R. Inorg. Chem. 2006, 45,
9877.
7656 Inorganic Chemistry, Vol. 46, No. 18, 2007

Decomposition of Tetrathionate Ion in Alkaline Medium
[SO42-]∞ ) 2[S2O32-]∞ + [S3O62-]∞ - 3[S4O62-]0 (12)
2-
[SO32-]∞ ) 7[S4O62-]0 - 4[S2O32-]∞ - 4[S3O6
] (13)
∞
The corresponding results can be seen in Table 2 together
with the ratio of each of final concentration of the sulfate
and sulfite over the initial tetrathionate concentration deter-
mined by eqs 12 and 13, respectively. It is clearly indicated
that SR2 slightly increases with decreasing initial tetrathionate
concentration and with increasing pH. The opposite effect
is observed in the dependence of SR3 against the initial
tetrathionate concentration and the pH. It is also important
to note that the relative amount of sulfite formed during the
decomposition increases with increasing pH and with
decreasing initial tetrathionate concentration. The amount of
sulfate formed seems to be subtle, but later we shall see that
the route, which produces sulfate, plays an important role
in determining the kinetic behavior of the system. From Table
2, it is clear that the stoichiometry of the reaction is very
complex and must be augmented with the following equa-
tions.
Figure 2. Typical absorbance change during the alkaline decomposition
of tetrathionate in the UV range. [S4O2-
] ) 0.002 M, pH ) 11.5
0
6
(adjusted by phosphateEnDash/hydrogen phosphate buffer), [S2O2-]0 ) 0
3
M. Lines with boldface type represent the seven selected wavelengths, where
the fitting procedure was executed.
applied in several kinetic systems studied previously.27,29-32
The essence of the procedure is to take all of the chemically
plausible mono- and bimolecular reactions and start the fitting
with rate coefficients assigned to the reactions in such a way
that the initial mean square deviation between the measured
and calculated absorbances is sensitive to each of them.
Parameters that become insensitive, meaning that the cor-
responding process does not affect the fitted curves, are
omitted at each step. Finally, only the most sensitive
parameters remain in the model. The most critical part of
this procedure is to choose the set of species involved in the
mono- and bimolecular reactions. In fact the reactants, the
end products, and the solvent must be included. This group
consists of S4O62-, S2O32-, S3O62-, SO32-, SO42-, OH-, and
H2O. Additional intermediates S5O62-, S2O3OH-, S3O3OH-,
HS(OH), and S(OH)2 were chosen for the following reasons:
(i) The role of S2O23OH- has already been well-
established27,33,34 in the redox reactions of tetrathionate. Since
S2O3OH- plays a central role in the tetrathionate-hydrogen
peroxide and the tetrathionate-chlorite reactions, it looks
reasonable to postulate it as an intermediate.
4S4O62- + 6OH- f 2S3O62- + 5S2O32- + 3H2O (S1)
2S4O62- + 6OH- f 3S2O32- + 2SO32- + 3H2O (S2)
4S4O62- + 10OH- f 7S2O32- + 2SO42- + 5H2O (S3)
As we shall see later, stoichiometry S1, with slight contribu-
tion of S3, governs the reaction at slightly alkaline pHs where
pentathionate is not a stable end product due to its fast
alkaline degradation.14 The importance of eq S3 increases
with increasing tetrathionate concentration. As the pH
increases, however, the contribution of eq S2 quickly rises
and starts to dominate the stoichiometry in a strong alkaline
solution. These stoichiometries can be easily understood by
the proposed kinetic model of the degradation process
(explained later). We also would like to stress that trithionate
can be treated as an end product only in our kinetic time-
scale. It is well-known18,28 that trithionate hydrolyzes at a
much longer time-scale meaning that this process cannot have
a significant effect in our experimental circumstances.
MRA Studies. Figure 2 shows a typical absorbance
change during the course of the reaction within the 260-
330 nm wavelength range. Since there is no explicit sign in
the absorbance-time sections for formation of a long-lived
intermediate, detailed matrix rank analysis25 has been
executed in the wavelength range of 265-330 nm on all the
measured series simultaneously where the measured values
did not exceed 1.0 AU. This study has clearly revealed three
absorbing species (tetrathionate, thiosulfate, and trithionate)
in the wavelength range in question meaning that there is
no absorbing intermediate formed in significant concentra-
tion.
(ii) Pentathionate is known to participate in catalytic
rearrangement of tetrathionate;15-17 therefore this species
must also be included as an intermediate.
(iii) S3O3OH-, HS(OH), and S(OH)2 have already been
considered in the degradation process of pentathionate;14
therefore these species may also play important roles in the
alkaline decomposition of tetrationate.
This set of reactants, intermediates, and end products leads
to 25 plausible mono- or bimolecular steps. Because the
(29) Horva´th, A. K.; Nagypa´l, I. J. Phys. Chem. A 1998, 102, 7267.
(30) Horva´th, A. K.; Nagypa´l, I.; Peintler, G.; Epstein, I. R.; Kustin, K. J.
Phys. Chem. A 2003, 107, 6966.
(31) Horva´th, A. K.; Nagypa´l, I.; Epstein, I. R. J. Phys. Chem. A 2003,
107, 10063.
(32) Horva´th, A. K.; Nagypa´l, I. J. Phys. Chem. A 2006, 110, 4753.
(33) Kurin-Cso¨rgei, K.; Orba´n, M.; Ra´bai, G.; Epstein, I. R. J. Chem. Soc.,
Faraday Trans. 1996, 92, 2851.
Proposed Kinetic Model. The method, with which the
best fitting model was found, has already been successfully
(28) Naito, K.; Hayata, H.; Mochizuki, M. J. Inorg. Nucl. Chem. 1975,
37, 889.
(34) Voslar, M.; Matejka, P.; Schreiber, I. Inorg. Chem. 2006, 45, 2624.
Inorganic Chemistry, Vol. 46, No. 18, 2007 7657

Varga and Horva´th
decomposition is strongly affected by the pH, hydroxide-
catalyzed pathways of all the 25 reactions were also
considered initially. After a long but straightforward reduc-
tion procedure, the following kinetic model emerged.
1.5 M) where the rate coefficient of the decomposition was
obtained. Thus, taking the positive kinetic salt effect into
consideration, the difference between the value determined
in the present and the previous studies can easily be
explained.
Step R2 is the second-order decomposition of S2O3OH-
producing thiosulfate and sulfite. Evidently, this reaction is
not elementary, and it probably proceeds via the following
consecutive processes.
PO43- + H+ h HPO4
(E1)
(E2)
(R1)
2-
-
CO32- + H+ h HCO3
S4O62- + OH- f S2O32- + S2O3OH-
2S2O3OH- f S2O32- + S2O52- + 2H+
S2O52- + H2O f 2SO32- + 2H+
(14)
(15)
2S2O3OH- + 4OH- f S2O32- + 2SO32- + 3H2O (R2)
S4O62- + S2O3OH- f S5O62- + SO42- + H+ (R3)
The calculations have revealed that the rate coefficient kR2
is in total correlation with kR3; thus we have fixed kR2 ) 107
M-1 s-1 that is fast enough to treat S2O3OH- as a steady-
state intermediate in accordance with MRA studies. This
pathway directly leads to stoichiometry S2. The role of this
step was confirmed by further calculation in which this step
was omitted from the proposed model. The result showed
on the one hand an unacceptably high average deviation and
on the other hand a significant loss of source of sulfite. We,
therefore, concluded that this step is strongly supported by
the experiments.
Step R3 is the other route for conversion of the intermedi-
ate S2O3OH-. Evidently, its role increases with increasing
initial tetrathionate concentration. As one may notice, this
step converts the reactants into sulfate and pentathionate. This
pathway is responsible for the increasing amount of sulfate
with increasing initial tetrathionate formation (see Table 2).
Since the proposed kinetic model does not contain any other
process than step R3 yielding sulfate, its role is essential.
One might argue that the easily conceivable alkaline
decomposition of trithionate
2-
S4O62- + S2O32- h S5O62- + SO3
S4O62- + SO32- f S3O62- + S2O3
(R4)
(R5)
2-
S5O62- + OH- f + S2O32- + S3O3OH- (R6)
S3O3OH- + OH- f S2O32- + S(OH)2
(R7)
S4O62- + S3O3OH- f S5O62- + S2O3OH- (R8)
2S(OH)2 + 2OH- f SO32- + HS(OH) + 2H2O (R9)
S2O32- + HS(OH) f S2O32- + H2O
(R10)
The rapid de- and reprotonation processes (E1 and E2) were
taken into account with known equilibrium constants35 to
follow the small change in pH during the decomposition
(pKE1 ) 11.5, pKE2 ) 9.7). These acid dissociation reactions
may be regarded as auxiliary processes, necessary only for
detailed calculations, but they are not the central part of the
proposed model.
S3O62- + H2O f SO42- + S2O32- + 2H+
(16)
Discussion
might also be a possible alternative for producing sulfate as
an end product. If, however, step R3 was substituted with
eq 16, then not only would the average deviation increase
to an unacceptably high 0.0141 AU but also k16 would be
0.018 ( 0.003 M-1 s-1. It is, however, orders of magnitude
higher than that obtained from independent measurements
by Rolia and Chakrabarti.18 Therefore, we have concluded
that step R3 cannot be replaced by eq 16. Further support of
step R3 is in our latest results on thiosulfate-hypochlorous
acid system36 that explains simultaneous formation of higher
polythionates and sulfate. It should also be noted that
stoichiometry S3 can be generated from 2*(R1) + 2*(R3)
+ 2*(R6) + 2*(R7) + (R9) + (R10).
Step R1 is the initiating rate-determining step of the
alkaline decomposition of tetrathionate that proceeds via
scission of the inner S-S bond yielding thiosulfate and
S2O3OH-. The formation of the adduct S4O6OH3- is probably
followed by the heterolytic inner S-S bond cleavage
resulting in the nucleophilic displacement of thiosulfate by
hydroxide ion. Our fitting procedure has yielded a value of
0.044 ( 0.004 M-1 s-1 for kR1 that is approximately an order
of magnitude higher than that could be calculated from Rolia
and Chakrabarti’s study18 but an order of magnitude lower
than that determined by Zhang and Dreisinger.4 They have
obtained respective values of 0.0051 and 0.38 M-1 s-1 for
the rate-determining step of the alkaline decomposition. This
difference is, however, due to the effect of ionic strength as
noted by Breuer and Jeffrey.19 Rolia and Chakrabarti have
conducted their experiments without buffer and supporting
electrolyte resulting in an ionic strength (approximately 0.02
M) as determined by tetrathionate and the products. In
contrast, Zhang and Dreisinger have used high buffer
concentration yielding higher ionic strength (approximately
Step R4 is the well-known thiosulfate-assisted rearrange-
ment of tetrathionate studied by several research groups.15-17
Recent result on the dynamic instability of the electrochemi-
cal oxidation of thiosulfate is also explained by the help of
(35) IUPAC Stability Constant Database; Royal Society of Chemistry:
London, 1992-1997.
(36) Varga, D.; Horva´th, A. K.; Nagypa´l, I. J. Phys. Chem. B 2006, 110,
2467.
7658 Inorganic Chemistry, Vol. 46, No. 18, 2007

Decomposition of Tetrathionate Ion in Alkaline Medium
this step.37 Fava and Bresadola have determined17 a value
of 3.9 × 10-3 M-1 s-1 for the forward reaction at 50 °C and
at an I ) 0.94 M ionic strength in neutral solution, but they
were not able to determine the individual rate coefficient of
the reverse reaction. Foss and Kringlebotn have obtained38
1.3 × 10-3 M-1 s-1 for the rate coefficient of the forward
reaction at 25 °C and at an I ) 1.15 M ionic strength. Our
calculation has yielded kR4 ) (1.94 ( 0.08) × 10-4 M-1
s-1, which is approximately an order of magnitude lower
than the value determined by Foss and Kringlebotn. This
difference, however, can easily be explained by the kinetic
salt effect of reaction between two like-charged ions.
Nevertheless, the exact value of k-R4 cannot be determined
from our measurements. The back reaction must be rapid
enough to provide sufficiently low level of pentathionate,
since it does not accumulate in detectable amount. Further
details about the reactions of pentathionate will be discussed
later (see paragraphs below), but we were able to calculate
only the ratio of kR6/k-R4 that was found to be 0.82 ( 0.01
with k-R4 ) 200 M-1 s-1. As one may notice thiosulfate
acts as an autocatalyst of the decomposition of the tetrathion-
ate, but the characteristic sigmoidal-shaped curve is not
observed under our experimental circumstances. It is,
however, due to the fact that kR4 is low compared to the
value of kR1, which means that the S-shaped kinetic curve
cannot be manifested. The low value of kR4 also indicates
that the forward reaction may be neglected from the pro-
posed model. By use of kR4 ) 0, we arrive to a somewhat
higher, but acceptable, 0.0036 average deviation with slight
changes (<20%) of the remaining fitted parameters. Al-
though the effect of this process is not striking under our
experimental conditions, we have kept it in the final model
for two reasons: first, it decreases the average deviation by
15% approximately; second, more importantly, the value
for kR4 obtained by the present work is in complete coin-
cidence with that of determined previously by independent
works.17,38
Step R6 is the attack of hydroxide ion at the γ-sulfur atom
of pentathionate yielding thiosulfate and S3O3OH-. As it was
already discussed, only the ratio of kR6/k-R4 could be
calculated from our measurements. Since pentathionate was
not accumulated in detectable amount in our experimental
circumstances, k-R4 must be higher than 200 M-1 s-1 that
results in kR6 g 164 M-1 s-1. The latter value is, however,
2 orders of magnitude higher than that obtained by Chris-
tiansen et al.43 for the alkaline degradation of pentathionate
but is in complete coincidence with the qualitative observa-
tion published by Wagner and Schreier,14 who have stated
2-
that “The alkaline degradation of the S5O6 proceeds very
quickly”. If we had decreased kR6 to 2 M-1 s-1 resulting in
k-R4 to be 2.44 M-1 s-1, then pentathionate would have been
built up to such an extent (>10-5 M) that the absorbance-
time curves would have shown a maximum since pentathion-
ate absorbs the UV light much stronger24 than tetrathionate
does. Nevertheless, MRA studies should have also indicated
the presence of an absorbing intermediate. Thus, we have
concluded that k-R4 must be higher than 200 M-1 s-1 to keep
the pentathionate concentration under 10-7 M. Further
consequence of the high rate coefficient of the alkaline
decomposition of pentathionate is that nucleophilic displace-
ment of the thiosulfate with sulfite group on pentathionate
is more than 2 orders of magnitude higher compared to that
of tetrathionate ion in our experimental circumstances. Our
calculation has yielded k-R4/kR5 g 400 that is at least 1 order
of magnitude higher than obtained by Fava and Bresadola17
or Foerster and Centner.42 In both cases, however, the
experiments were conducted at around neutral pHs, where
significant amount of HSO3- was present beside SO32- that
may have influence on the apparent rate coefficient kR5 to
-
be decreased by the different reactivities of SO32- and HSO3
toward tetrathionate.
Step R7 is one of the possible routes for disappearance of
S3O3OH- formed in step R6 and was already suggested by
Wagner and Schreier.14 Our calculation has revealed total
correlation between parameters kR7 and kR8; therefore we
were able to determine only kR8/kR7 ) 374 ( 12 with kR7 g
103 M-1 s-1.
Step R8, the other pathway responsible for keeping the
concentration of S3O3OH- low, has not been proposed so
far. The role of this step was also confirmed by further
calculations by omission this step from the proposed model.
The results indicated an increase of the average deviation to
0.0139 from 0.003, meaning that the role of this step is
strongly supported.
Steps R9 and R10 are fast processes and were already
proposed by Wagner and Schreier.14 Their individual rate
coefficients could not be determined from our measurements,
and only lower limits could be assigned to these processes.
Although recent studies44-46 consider sulfoxylate ion to have
longer lifetime in 0.5 M NaOH solution, the protonated
Step R5 is the well-known sulfitolysis of tetrathionate. The
reaction was first studied by Kurtenacker and Goldbach39
and was reported to be applied for successful determination
of polythionates.40 The procedure was later modified by
Iwasaki and Suzuki41 to make it suitable for analyzing
microamounts of polythionates. We have determined kR5 to
be 0.547 ( 0.005 M-1 s-1 that agrees well with the value
that could be estimated from the latter study. This value is
also in fairly good agreement with the value of 0.172 M-1
s-1 determined by Foerster and Centner at 0 °C without
adjusting the ionic strength by supporting electrolyte.42 As
one may notice stoichiometry S1 may be generated from 2*-
(R1) + (R2) + 2*(R5).
(37) Du, Z.; Gao, Q.; Feng, J.; Lu, Y.; Wang, J. J. Phys. Chem. B 2006,
110, 26098.
(38) Foss, O.; Kringlebotn, I. Acta Chem. Scand. 1961, 15, 1608.
(39) Kurtenacker, A.; Goldbach, E. Z. Anorg. Allgem. Chem. 1927, 166,
177.
(40) Goehring, M.; Feldmann, U.; Helbing, W. Z. Anal. Chem. 1949, 129,
346.
(43) Christiansen, J. A.; Drost-Hansen, W.; Nielsen, A. E. Acta Chem. Scan.
1952, 6, 333.
(44) Svaroksky, S. A.; Simoyi, R. H.; Makarov, S. V. J. Chem. Soc., Dalton
Trans. 2000, 511.
(41) Iwasaki, I.; Suzuki, S. Bull. Chem. Soc. Jpn. 1966, 39, 576.
(42) Foerster, F.; Centner, K. Z. Anorg. Chem. 1926, 157, 45.
(45) Svaroksky, S. A.; Simoyi, R. H.; Makarov, S. V. J. Phys. Chem. B
2001, 105, 12634.
Inorganic Chemistry, Vol. 46, No. 18, 2007 7659

Varga and Horva´th
Table 3. Fitted and Fixed Rate Coefficients Used in the Fitting
Procedurea
step
rate equation
parameter (M-1 s-1
)
kR1[S4O2-][OH-]
(R1)
(R2)
(R3)
(R4)
(-R4)
(R5)
(R6)
(R7)
(R8)
(R9)
(R10)
0.044 ( 0.004
107
57.3 ( 1.6
k
k
R2[S2O63OH-]2
R3[S4O2-][S2O3OH-]
6
kR4[S4O2-][S2O2-
]
]
(1.94 ( 0.08)×10-4
1220 ( 40
0.547 ( 0.005
1000
6
3
k-R4[S5O2-][SO23-
kR5[S4O2-6][SO23-
]
6
k
R6[S5O2-][OH-]
6
k
R7 [S3O3OH-][OH-]
1000
kR8 [S4O2-][S3O3OH-]
(3.74 ( 0.12)×105
6
kR9 [S(OH)2]2
g103
kR10 [SO2-][HS(OH)]
g106
3
Figure 5. Measured (symbols) and calculated kinetic (solid lines) curves
in the absence of thiosulfate at pH ) 11.50 using phosphate buffer. [S4
O26-]0 ) 3.0 (O), 2.0 (9), 1.4 (4), 1.0 ([), 0.7 (3), 0.5 (b), 0.3 (]) mM.
a Note: the actual calculation has yielded the ratios of kR3/kR2 ) (5.73
( 0.16) × 10-6, kR6/k-R4 ) 0.82 ( 0.1, and kR8/kR7 ) 374 ( 12.
Figure 3. Measured (symbols) and calculated kinetic (solid lines) curves
Figure 6. Measured (symbols) and calculated kinetic (solid lines) curves
in the absence of thiosulfate at [S4O2-
] ) 3.0 mM using phosphate
0
in the presence of thiosulfate at [S4O2-
] ) 1.0 mM, pH ) 12.2 using
0
buffer. pH ) 10.80 (O), 11.02 (9), 11.650 (4), 11.99 ([), 12.20 (0).
6
phosphate buffer. [S2O2-]0 ) 0.0 (O), 0.3 (9), 0.5 (4), 0.7 ([), 1.0 (0),
3
2.0 (b), 3.0 (]) mM.
Figure 4. Measured (symbols) and calculated kinetic (solid lines) curves
in the absence of thiosulfate at [S4O2-]0 ) 0.6 mM using carbonate buffer.
6
Figure 7. Measured (symbols) and calculated kinetic (solid lines) curves
pH ) 10.65 (O), 10.18 (9), 9.70 (4), 9.22 ([).
in the presence of thiosulfate at [S4O2-
] ) 0.6 mM using carbonate
0
6
buffer. [S2O2-]0 ) 0.0 (O), 0.3 (9), 0.6 (4), 1.0 ([) mM.
3
-
sulfoxylate ion (HSO2 ) and especially sulfoxylic acid can
disproportionate rapidly in weakly alkaline solution.47 There-
fore, we have set lower limits for the values of kR9 and kR10
as 103 and 106 M-1 s-1, respectively, which keep the
concentration levels of S(OH)2 and HS(OH) at sufficiently
low values. One might argue that step R9 can also be written
with elimination of HS(OH) into the following form.
Since the alkaline decomposition of tetrathionate yields
enough sulfite in our experimental conditions, it could be
done without any change in the final results. If, however,
the sulfite concentration decreases to a sufficiently low level,
then spontaneous decomposition of HS(OH) gives colloidal
sulfur precipitation,
2S(OH)2 + 2OH- f S2O32- + 3H2O
(17)
HS(OH) f S + H2O
(18)
which is known to frequently accompany the later stages of
the alkaline degradation of polythionates.12,14,43 We have also
tried to insert other reactions of HS(OH); for instance
(46) Kudrik, E. V.; Makarov, S. V.; Zahl, A.; van Eldik, R. Inorg. Chem.
2005, 44, 6470.
(47) Makarov, S. V. Russ. Chem. ReV. 2001, 70, 885.
7660 Inorganic Chemistry, Vol. 46, No. 18, 2007

Decomposition of Tetrathionate Ion in Alkaline Medium
that thiosulfate catalyzes the degradation process in agree-
ment with Rolia and Chakrabarti.18 Furthermore, it is also
shown that the rate of decomposition does not depend on
the buffer applied in the case of phosphate and carbonate
buffer components. It follows from the indirect fact that the
kinetic curves measured in the presence of different buffer
components (phosphate and carbonate) can readily be treated
together in the simultaneous curve-fitting procedure. It,
however, does not mean that no specific buffer dependence
of the alkaline decomposition of tetrathionate could be
observed at all. Although it was not investigated in the
present work, it seems firmly established that the decomposi-
tion is accelerated by ammonia6 most probably through the
Figure 8. Calculated concentration-time series of tetrathionate (b),
thiosulfate (0), and trithionate ([) compared to the concentration ranges
(dashed lines) obtained from the decomposition of the absorbance-time
series of 265-295 nm wavelength range supposing that all the absorbances
-
formation5 of SO3NH2 . Certainly it is not a unique explana-
tion, and therefore further investigations are needed to clarify
the effect of ammonia on the decomposition.
have (0.003 AU uncertainty. [S4O2-]0 ) 0.002 M, pH ) 10.80.
6
The present study may also be helpful in explaining the
experimentally observed multicomponent convection, in non-
buffered medium, in the chlorite-tetrathionate reaction at
high initial hydroxide concentration48 where significant
amounts of thiosulfate, trithionate, and sulfite may also be
present in a short period of time as a result of the alkaline
decomposition of tetrathionate. In fact, at a usual composi-
tion,48 our proposed model in unbuffered medium predicts
that the concentration of thiosulfate becomes commensurable
with that of tetrathionate within 15 min. Since formation of
the thiosulfate is accompanied by consumption of hydroxide
ions, the pH of the solution reaches neutrality opening up
the possibility of igniting the chlorite-thiosulfate reaction.49
The contributory effect of thiosulfate, trithionate, or sulfite
by their reactions with chlorite49,50 to the mother reaction
may, therefore, not be negligible.
reactions with tetrathionate, pentathionate, thiosulfate, and
the intermediates, but none of them proved to be necessary
for quantitative description of the kinetic curves.
Table 3 contains the final parameter set obtained from our
experiments. The average deviation was found to be 0.003
AU, which we believe to be close to the experimentally
achievable limit of error. The final results can be illustrated
in Figures 3-7. They clearly show that the model works
nicely in wide concentration ranges of the reactants and pHs
as well as using different buffer components.
It should also be mentioned that the absence of accumulat-
ing intermediates allows us to decompose the absorbance-
time series into concentration-time series of tetrathionate,
thiosulfate, and trithionate. As one may notice, this algebraic
decomposition might be seriously affected by the error of
the absorbance measurement especially in case of trithionate,
since its molar absorbance is much lower than that of the
tetrathionate. Therefore, we have also checked if the proposed
model is able to predict properly the concentration-time
series of tetrathionate, thiosulfate, and trithionate simulta-
neously that are obtained from the decomposition of the
absorbance-time series. Figure 8 also provides further
support for the proposed model, since the concentration-
time series are well within the concentration range deter-
mined from the absorbance-time series supposing that all
the measured absorbance have (0.003 uncertainty.
We strongly believe that the model may also serve as a
starting point to design lowering of polythionate concentra-
tion in the leaching processes of silver or gold ores as well
as the maximum thiosulfate recovery for the recycling
process.
Acknowledgment. This work was supported by the
Hungarian Research Fund (Grant No. T047031). The authors
are grateful to Professor Istva´n Nagypa´l for his valuable
advice after reading the manuscript. The authors are also
thankful to Dr. Otto´ Berkesi for his assistance in the Raman
measurements.
Conclusion
Although the alkaline decomposition of the tetrathionate
is thought to be a complex process for a long time, no
comprehensive mechanism has yet been proposed to incor-
porate this feature. Our suggested model clearly shows that
the product distribution of the reaction is strongly affected
by pH, the initial concentration of tetrathionate, and the fact
IC700992U
(48) Ba´nsa´gi, T., Jr.; Horva´th, D.; To´th, AÄ . Chem. Phys. Lett. 2004, 384,
153.
(49) Nagypa´l, I.; Epstein, I. R. J. Phys. Chem. 1986, 90, 6285.
(50) Huff-Hartz, K. E.; Nicoson, J. S.; Wang, L.; Margerum, D. W. Inorg.
Chem. 2003, 42, 78.
Inorganic Chemistry, Vol. 46, No. 18, 2007 7661