Detailed kinetics and mechanism of the oxidation of thiocyanate ion (SCN-) by peroxomonosulfate ion (HSO5-). Formation and subsequent oxidation of hypothiocyanite ion (OSCN-)

Article abstract of DOI:10.1021/ic302544y

The haloperoxidase-catalyzed in vivo oxidation of thiocyanate ion (SCN ^-) by H₂O₂ is important for generation of the antimicrobial hypothiocyanite ion (OSCN^-), which is also susceptible to oxidation by strong in vivo oxidizing agents (i.e., H₂O ₂, OCl^-, OBr^-). We report a detailed mechanistic investigation on the multistep oxidation of excess SCN^- with peroxomonosulfate ion (HSO₅^- in the form of Oxone) in the range from pH 6.5 to 13.5. OSCN^- was detected to be the intermediate of this reaction under the above conditions, and a kinetic model is proposed. Furthermore, by kinetic separation of the consecutive reaction steps, the rate constant of the direct oxidation of OSCN^- by HSO ₅^- was determined: k₂ = (1.6 ± 0.1) × 10² M^-1 s^-1 at pH 13.5 and k ₂^H = (3.3 ± 0.1) × 10³ M ^-1 s^-1 at pH 6.89. A critical evaluation of the estimated activation parameters of the elementary steps revealed that the oxidations of SCN^- as well as the consecutive OSCN^- by HSO ₅^- are more likely to proceed via 2e^--transfer steps rather than 1e^- transfer.

Full text of DOI:10.1021/ic302544y

Article
pubs.acs.org/IC
Detailed Kinetics and Mechanism of the Oxidation of Thiocyanate
Ion (SCN−) by Peroxomonosulfate Ion (HSO5 ). Formation and
−
Subsequent Oxidation of Hypothiocyanite Ion (OSCN−)
Jozsef Kalmar,* Gabor Lente, and Istvan Fabian
́
́
́
́
́
́
Department of Inorganic and Analytical Chemistry, University of Debrecen, P.O. Box 21, Debrecen H-4010, Hungary
S
* Supporting Information
ABSTRACT: The haloperoxidase-catalyzed in vivo oxidation
of thiocyanate ion (SCN−) by H2O2 is important for
generation of the antimicrobial hypothiocyanite ion
(OSCN−), which is also susceptible to oxidation by strong
in vivo oxidizing agents (i.e., H2O2, OCl−, OBr−). We report a
detailed mechanistic investigation on the multistep oxidation
of excess SCN− with peroxomonosulfate ion (HSO5 in the
−
form of Oxone) in the range from pH 6.5 to 13.5. OSCN− was
detected to be the intermediate of this reaction under the
above conditions, and a kinetic model is proposed.
Furthermore, by kinetic separation of the consecutive reaction
steps, the rate constant of the direct oxidation of OSCN− by HSO5− was determined: k2 = (1.6 0.1) × 102 M−1 s−1 at pH 13.5
H
and k2 = (3.3 0.1) × 103 M−1 s−1 at pH 6.89. A critical evaluation of the estimated activation parameters of the elementary
steps revealed that the oxidations of SCN− as well as the consecutive OSCN− by HSO5 are more likely to proceed via 2e−-
−
transfer steps rather than 1e− transfer.
and Cl−/OCl− or Br−/OBr− are simultaneously present. The
INTRODUCTION
■
understanding of the in vivo redox chemistry of SCN− is further
The diverse, and often catalytic oxidation reactions of
complicated by the fact that the 2e− oxidation product of
thiocyanate anion (SCN−) are intensively studied (vide infra)
SCN−, OSCN−, can be subsequently oxidized by the same
because the biological importance of this anion is connected to
oxidizing agents as SCN−, i.e., H2O2, HOCl, or HOBr.3−5 The
its vivid in vivo redox chemistry.1 In addition, SCN− is an
reactions of OSCN− with these oxidizing agents are difficult to
environmental pollutant, and its oxidative removal from
study because OSCN− cannot be synthesized in the absence of
industrial wastewater is essential. Furthermore, the in situ
excess SCN−. This problem is attributed to the “overoxidation”
oxidation of SCN− is of fundamental importance during
of OSCN−, which is assumed to be of a similar order of
thiocyanation of various compounds in synthetic organic
chemistry.
magnitude as its production from SCN−.3,5,18 The half-life of
the decomposition of OSCN− is on the order of several
Thiocyanate ion is enzymatically oxidized by H2O2 to the
minutes at pH 7,19 further limiting the applicable experimental
antimicrobial hypothiocyanite ion2 (OSCN−) in vivo.3−5 The
rate of the direct SCN− + H2O2 reaction is slow,6 but defensive
conditions.
From environmental aspects, SCN− and cyanide ion (CN−)
are considered to be the main inorganic anionic pollutants of
wastewater and recycled water from metallurgical, mostly gold
extraction plants. Although the toxicity of CN− is undoubtedly
higher than that of SCN−, the latter can be found in larger
quantities in metallurgical wastewater and is more toxic to fish.
The removal of SCN− is important for the detoxification of
industrial water because molecular oxygen can oxidize SCN− to
sulfate (SO42−) and CN− on a longer time scale, especially
under UV radiation. For this reason, the admissible limit of
SCN− in metallurgical wastewater is usually as low as 0.05 mg/
L. Technologically feasible processes for SCN− removal almost
exclusively utilize oxidative degradation approaches.20 The
peroxidase enzymes found in neutrophils (myeloperoxidase7)
and eosinophil granulocites (eosinophil peroxidase8) catalyze
this reaction effectively. Extracellular enzymes in milk
(lactoperoxidase9), in airways (lactoperoxidase10), and in saliva
(salivary peroxidase11) are also efficient catalysts for the
oxidation of SCN− by H2O2. Furthermore, OSCN− is a potent
bactericide and plays an important role in multiple stages of the
nonimmune defense system3,4 at several physiological locations
(e.g., in the mouth12 and in airways13,14). Thus, the
insufficiently low in vivo local concentration of its precursor
SCN− is assumed to play an important role in the development
of several chronic diseases, like cystic fibrosis.1,15 SCN− can be
oxidized to OSCN− in a noncatalytic way by hypochlorous acid
(HOCl)16 and by hypobromous acid (HOBr).17 These
reactions may be important in controlling the redox activities
of those defensive peroxidase systems where SCN−/OSCN−
Received: November 20, 2012
Published: January 30, 2013
© 2013 American Chemical Society
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fundamental reactions of these processes are catalytic,21−23
photocatalytic,24,25 and also biological oxidations.26,27 The
universal goal is to convert SCN− quantitatively into environ-
mentally more tolerable cyanate ion (OCN−) [or carbonate ion
(CO32−) and nitrogen (N2)] and SO42−. The oxidizing agents
of choice are usually hydrogen peroxide (H2O2)21 and ozone
(O3)23 in order to avoid the formation of any hazardous side
products.
0.03 ms, allowing determination of the first-order rate constants (kobs
)
up to 500 s−1. The temperature was maintained within 0.1 °C during
all stopped-flow measurements with a Julabo F-12 refrigerated/heated
circulator.
Electrospray Ionization Mass Spectrometry (ESI-MS). ESI-MS
measurements were performed with a Bruker micrOTOFQ mass
spectrometer equipped with a quadrupole and time-of-flight (Q-TOF)
analyzer. The detailed description of the methodology is given in our
recent study.19 The aim of the MS experiments was to identify the
intermediates and products of the HSO5− + SCN− reaction in a quick
and sensitive, but qualitative way.
Software. The raw data sets of the measurements were processed
with instrument-controlling software. Typically, Micromath Scientist
2.042 was used for Levenberg−Marquard least-squares fitting
procedures. Multiple sets of kinetic curves recorded under different
initial condititions were simultaneously evaluated with the program
package ZiTa43 using the GEAR algorithm.44
The oxidation of SCN− plays a crucial role in organic
synthetic strategies aiming at the preparation of thiocyanated
compounds. To create a reactive intermediate for thiocyanation
of both aromatic28−30 and aliphatic31−33 compounds, a mixture
of various oxidizing agents is used in situ with thiocyanate salts.
−
Some of these synthetic methods utilize HSO5 in the form of
Oxone as the oxidizing agent in order to achieve milder
reaction conditions with higher yields.30,32,33
−
HSO5 is a peroxo compound and resembles H2O2 in its
RESULTS AND DISCUSSION
■
redox behavior, although the reactions of HSO5− with inorganic
−
Oxidation of SCN− with HSO5 in the Range from pH
reducing agents34,35 are characteristically faster than those of
7 to 13. For the sake of simplicity, we studied the reaction of
−
H2O2. The oxygen-transfer reactions of HSO5 to halide ions
HSO5 with SCN− at pH 13.5 first, where HSO5 is fully
−
−
(Cl−, Br−, and I−) are described in detail.34 Stoichiometric36
34
−
deprotonated [pKa(HSO5 ) = 8.35 at I = 1.0 M ], and
hypothiocyanite (OSCN−), the expected product of an oxygen
transfer to SCN−, is stable for hours.45 The reaction was
studied under high excess of SCN− ([SCN−]0 = 50.0−500
and kinetic results37,38 indicate that the oxidation of SCN− by
−
HSO5 is a complicated multistep process under acidic
conditions, from pH 2.2 to 6.3. The end products were
36
identified as OCN−, CN−, and SO4
.
On the basis of
2−
2−
−
mM) over SO5 ([HSO5 ]0 = 5.0 mM). The reaction is
complete in a few seconds under the conditions applied, and
additional kinetic effects were not observed on longer time
scales. Kinetic traces were recorded at various wavelengths in
the range of 350−500 nm (Figure 1) and analyzed by singular-
calculations by the initial rate method and using numerical
integration for simulating kinetic observations, Smith and
Wilson concluded that the key initial step in the overall reaction
37,38
is the equilibrium (R0H) between SCN− and HSO5 ,
−
where “HOSCN” is an uncharacterized intermediate.
HSO5− + SCN− ⇌ SO 2− + HOSCN
(R0H)
4
The formation of a second intermediate [thiocyanogen:
(SCN)2] is proposed in the rapid comproportionation of
SCN− and HOSCN, which inhibits accumulation of HOSCN
under acidic conditions. This comproportionation is well
characterized in the literature.19,39 In our present study, the
kinetics and mechanism of the oxidation of SCN− by HSO5
−
are characterized in a wide range of pH values from the
physiologically relevant neutral to the highly basic, where the
comproportionation of SCN− + OSCN− to give (SCN)2 is
negligible;19 thus, OSCN− accumulates.
2−
Figure 1. UV−vis spectral change recorded in the SO5 + SCN−
EXPERIMENTAL SECTION
reaction system at pH 13.5. The absorbance maximum at 376 nm is
characteristic of OSCN−. The first (blue) and last (red) spectra are
highlighted. Inset: Kinetic curve recorded at 376 nm. [KHSO5]0 = 5.0
mM; [NaSCN]0 = 500 mM; pH 13.5; I = 1.0 M; T = 25.0 °C.
■
Reagents and Solutions. NaSCN (99.99+ %) and Oxone were
purchased from Sigma-Aldrich and used without further purification.
The KHSO5 content of Oxone was regularly checked by iodometric
titration. All other chemicals were purchased in the highest available
purity and used without further purification. Ultrafiltered (Milli-Q,
Millipore) water was used for preparing solutions. The pH of the
reaction mixture was maintained either by phosphate buffer (from pH
6.5 to 7.8), by borate buffer (from pH 8.0 to 10), or by NaOH (from
pH 11 to 13). The pH was measured by a Metrohm 6.0234.110
combined glass electrode attached to a Metrohm 721 NET Titrino
titrating unit. The electrode was calibrated by two buffers according to
IUPAC recommendations.40 The ionic strength was adjusted by
NaClO4.
value decomposition analysis (Table S1 in the Supporting
Information) in order to determine the number of absorbing
species in the system.46,47 It was found that the increase in
absorbance during the reaction at around 375 nm can be
attributed to the formation of a single species; no absorbing
intermediates were detected. The observed product was
unambiguously identified as OSCN− with an absorbance
maximum at 376 nm.48 OSCN− was also identified with an
ESI-MS method (see the Supporting Information) described in
our recent study.19 In subsequent experiments, the kinetic
curves were exclusively recorded at 376 nm, where neither the
reactants nor other products contribute to the absorbance. All
kinetic traces (Figure 2, right panel) could be fitted with a
single-exponential function within the precision of the
experimental data (R2 > 0.98) and show the following
Kinetic Measurements. Fast kinetic experiments were carried out
in an Applied Photophysics DX-17 MV stopped-flow instrument with
either a photomultiplier tube or a photodiode array installed as the
detector. Absorbance traces were collected using an optical cell of 10.0
mm path length. The dead time of the stopped-flow instrument was
determined by the reaction of 2,6-dichlorophenol−indophenol with
ascorbic acid (AA) under pseudo-first-order conditions with AA in
excess.41 The dead time of the instrument was measured to be 1.51
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Inorganic Chemistry
Article
Figure 2. Reaction of SCN− with SO52− at pH 13.5. Left panel: experimental pseudo-first-order kinetic curves; only 5% of the recorded points are
shown for clarity. Right panel: pseudo-first-order rate constant as a function of the initial SCN− concentration. Inset: Final absorbance at 376 nm as a
function of the initial SCN− concentration. The symbols are the same for the two panels. Circles: experimental points with SCN− in high excess.
Triangles: comparable [SCN−]0 and [SO52−]0. Lines: result of a global fit to the kinetic model of Scheme 1. [KHSO5]0 = 5.0 mM; [NaSCN]0 = 5.0−
500 mM; pH 13.5; I = 1.0 M; T = 25.0 °C.
Scheme 1. Kinetic Model for the Oxidation of SCN− by
SO5 at around pH 13
characteristics: the final absorbance readings reach a limiting
value by increasing [SCN−]0, and the experimental pseudo-first-
order rate constant (kobs) is linearly dependent on [SCN−]0
with a nonzero intercept (Figure 2, left panel).
a
2−
At high SCN− excess, OSCN− is produced following a 1:1
OSCN−-to-SO5
stoichiometry. From the highest final
2−
absorbance, [OSCN−]max = 5.1 mM was calculated using the
literature value of the molar absorbance of OSCN− at 376 nm
(26.5 M−1 cm−1).48 Phenomenologically, such observations can
be explained by an equilibrium reaction between the reactants
(R0), which is pseudo-first-order in both directions49 (these
2−
conditions are fulfilled because SCN− and SO4 are in large
a
Although k1 ≪ k2, at high concentrations of SCN−, the first oxidation
step (R1) can compete with the faster second step (R2); thus, OSCN−
can accumulate. Reaction steps R3, R4, R5, and R6 in italics are
adopted from the literature52,39 and assumed to be much faster than
R1 and R2.
excess50 over the limiting reagents):
SO52− + SCN− ⇌ SO 2− + OSCN−
(R0)
4
A similar equilibrium (R0H) was proposed previously in the
kinetic model of the reaction of SCN− + HSO5 under acidic
−
conditions.37,38 However, considering that the backward
reaction of R0 is hardly feasible thermodynamically51 and
to the two-step kinetic model utilizing the ZiTa software
package and assuming that both the R1 and R2 steps are first-
order for each species involved. The actual rates of reactions R3
and R4 were considered to be comparable and much faster than
those of the rate-determining steps (R1 and R2) over the entire
course of the reaction, thus not affecting the stoichiometry of
the depletion of OSCN−. The fit was excellent even at
2−
that the kinetics of the SO5 + SCN− reaction system proved
to be indifferent to the change in [SO42−]0 (Figure S1 in the
Supporting Information), the existence of equilibrium R0 was
ruled out. An alternative kinetic model was constructed,
keeping in mind that the rate of oxidation of OSCN− by
peroxo compounds is expected to be about the same order of
magnitude as the rate of oxidation of SCN−.3,5,18 In the kinetic
model shown in Scheme 1, SCN− and its first-step oxidation
2−
comparable initial concentrations of SCN− and SO5 (Figure
2). The calculated rate constants are k1 = 8.08 0.02 M−1 s−1
and k2 = 164.8 0.5 M−1 s−1.
product OSCN− are competing for the remaining SO5 in
2−
A derivation on the basis of the system of simultaneous
differential equations defined by Scheme 1 yields the following
approximate expression for the concentration change of
OSCN− as a function of time (see the Supporting
Information):
reactions R1 and R2, respectively. With an increasing excess of
SCN−, the relative rate of the oxidation of OSCN− decreases;
thus, the surviving amount of OSCN− increases. The
stoichiometry of the production of OSCN− becomes 1:1
(OSCN−-to-SO52−) at high [SCN−]0. The oxidation product of
OSCN− (i.e., O2SCN−) is known to decompose in fast
reactions (R3−R6, adopted from the literature39,52) to yield the
final mixture of stable anions: SO42−, SO32−, CN−, and
OCN−.19,39,52 The production of SO42−, CN−, and OCN− is
k1[SCN−][SO52−
]
0
[OSCN−] =
k1[SCN−] + k2[SO52−
]
0
× {1 − exp[−(k1[SCN−] + k2[SO52−]0)t]}
(1)
−
observed during the oxidation of SCN− by HSO5 at acidic
pH.36 The formation of SO3
is detected during the
Although the derivation shows that Scheme 1 does not predict
a simple first-order kinetic trace in a strict mathematical sense,
eq 1 implies that the actual solution is close to an exponential
2−
decomposition of OSCN− at neutral pH.19 The presence of
CN− and OCN− in the basic SCN− + SO5 reaction mixture
2−
−
was verified by ESI-MS in this study (see the Supporting
Information) with a method detailed earlier.19 The exper-
imental kinetic data set recorded at pH 13.5 was fitted globally
curve with a pseudo-first-order rate constant kobs = k2[HSO5 ]0
+ k1[SCN−]0. This explains why the kinetic traces could be
fitted to a simple first-order expression with reasonable
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CN− + OSCN− → → OCN− + SCN−
accuracy. The kobs versus [SCN−]0 plot in Figure 2 (right
panel) can be fitted to kobs = m[SCN−]0 + a acceptably, yielding
(R7)
CN− is formed in reaction R5 (Scheme 1) after the oxidation of
m = 8.351
0.001 M−1 s−1 and a = 0.601
0.001 s−1 at
OSCN− by HSO5 (R2H). With increasing [SCN−]0, the
−
constant [SO52−]0 = 5.0 mM. A comparison with eq 1 gives k1
= 8.35 M−1 s−1 and k2 = 120 M−1 s−1. These rate constants are
in good agreement with those estimated by globally fitting the
experimental data set to the kinetic model of Scheme 1 without
approximations. These results lend further support to the
kinetic model proposed in Scheme 1.
produced amount of CN− decreases as the relative rate of
oxidation of OSCN− (R2H) decreases. Similar observations
were made previously regarding the role of CN− in the reaction
system of SCN− + HSO5− at highly acidic pH.36,37 The CN− +
OSCN− reaction (R7) was already investigated in detail in a
wide range of pH values by directly mixing the two reactants53
and is known to be reasonably fast only at pH < 7, which
The title reaction was also studied in detail at pH 6.89, where
−
explains the absence of this process in the SO5 + SCN−
2−
the protonated form of the oxidant, HSO5 , is predominant
over SO52−, but OSCN− is still deprotonated [pKa(HOSCN) =
4.93 at I = 1.0 M19]. In the first 100 ms after mixing the
reactants, the kinetic behavior of the reaction is similar to that
observed at pH 13.5 (Figures 3 and S2 in the Supporting
reaction system at pH 13.5. The inclusion of R7 in the kinetic
model of Scheme 1 is not necessary because the SCN−
+
HSO5 and OSCN− + CN− reactions are kinetically
independent in the applied pH range. The third and slowest
process observed at pH 6.89 resulted in the complete depletion
of OSCN− in minutes. This last and slowest process was
identified as the hydrolytic decomposition of OSCN− to give
simple inorganic anions such as CN−, OCN−, SO42−, and
SO32−, which were already described in our earlier work.19
−
The pH dependence of the (H)SO5 + SCN− reaction was
−
studied in the range from pH 6.9 to 13.5. The useful pH range
was limited because HOSCN rapidly decomposes at 2 < pH <
619 and HOSCN and SCN− rapidly comproportionate to give
(SCN)2 at pH < 2.39 In order to determine the apparent rate
constants for the oxidation of both SCN− (k1app) and OSCN−
(ka2pp) at a given pH, the kinetic experiments were run in
duplicate at two different [SCN−]0 values. The two rate
constants were computed by simultaneously fitting these pairs
of experimental kinetic curves to the kinetic model of Scheme
1. The values of both ka1pp and ka2pp decrease in a sigmoidal
fashion with increasing pH, as seen in Figure 4. An excellent
Figure 3. Experimental kinetic curves recorded in the reaction of
−
SCN− + HSO5 at pH 6.89. Two kinetically separated phases are
shown (a third slower step is not included in the figure). The faster
process can be described with the kinetic model of Scheme 1. Open
circles: experimental curves. Only 5% of the recorded points are shown
for clarity. Lines: results of the global fit of the first part of the kinetic
traces to the kinetic model of Scheme 1. [KHSO5]0 = 5.0 mM;
[NaSCN]0 = 25−500 mM; pH 6.89; I = 1.0 M; T = 25.0 °C.
Information). Thus, the findings can be interpreted using the
2−
−
kinetic model of Scheme 1 by replacing SO5 with HSO5 :
H
HSO5− + SCN− → SO 2− + OSCN− + H+
k1
4
(R1H)
H
HSO5− + OSCN− → SO 2− + O2SCN− + H+ k2
4
(R2H)
−
Figure 4. pH dependence of the SCN− + (H)SO5 reaction system.
The values of the rate constants were estimated by globally
fitting the first part of the kinetic curves with ZiTa: k1H = 203
Closed circles (right y scale): apparent rate constants for the oxidation
H
1 M−1 s−1 and k2 = (3.44 0.03) × 103 M−1 s−1 (Figures 3
of SCN− with (H)SO5 . Triangles (left y scale): apparent rate
−
−
constants for the oxidation of OSCN− with (H)SO5 . Line: result of a
and S2 in the Supporting Information). On a longer time scale
(>100 ms), two additional kinetically independent reaction
steps were observed. The second process (Figure 3) is
characterized by a small decrease in the absorbance at 376
nm in a few seconds, and according to the matrix rank analysis
of the multivawelength measurements, it does not involve the
formation of any new absorbing species (Table S1 in the
Supporting Information). A semiquantitative observation was
made that the amplitude of the absorbance change during the
second process significantly decreases with increasing [SCN−]0
and even vanishes at high [SCN−]0. This phenomenon makes
the quantitative description of this reaction step challenging
but, on the other hand, leads to the assumption that the second
process is probably attributed to the reaction of OSCN− with
CN−:
simultaneous fit to eqs 2 and 3. The y scales are normalized to each
other. [KHSO5]0 = 5.0 mM; [NaSCN]0 = 50.0 or 500 mM; I = 1.0 M;
T = 25.0 °C.
description of these pH profiles can be given by taking into
account the reactivity difference between HSO5 and SO5
−
2−
:
[H+]
Ka
k1app = k1H
+ k1
+ k2
Ka + [H+]
Ka + [H+]
(2)
(3)
[H+]
Ka
k2app = k2H
Ka + [H+]
Ka + [H+]
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−
where Ka is the acid dissociation constant of HSO5 and rate
constants k1, kH1 , k2, and k2H are defined in previous paragraphs.
In the final calculations, all kinetic traces recorded at various
reactant concentrations and pH values were fitted simulta-
neously to the kinetic model of Scheme 1, taking into account
the pH dependence of the system. These estimated values for
k1, kH1 , k2, and kH2 are listed in Table 1. The excellent fit between
Table 1. Calculated Rate Constants and Activation
−
Parameters in the Reaction System (H)SO5 + SCN−
studied at
pH
ΔH⧧
ΔS⧧
298 K
reaction
k
(M−1 s−1
)
(kJ mol−1
)
(J mol−1 K−1
)
Figure 5. Temperature dependence of the reaction between SCN−
R1
13.5
13.5
6.9
(1.0 0.2) × 101
(1.6 0.1) × 102
(2.0 0.2) × 102
(3.3 0.1) × 103
35.2 0.5
−108
2
−
and HSO5 at pH 6.54. Circles (right y scale): rate constant of the
R2
−
oxidation of SCN− with HSO5 (R1H). Triangles (left y scale): rate
R1H
R2H
26
30
1
2
−112
−77
3
5
−
constant of the oxidation of OSCN− with HSO5 (R2H). Lines:
6.9
results of nonlinear fits to the Eyring equation. [KHSO5]0 = 5.0 mM;
[NaSCN]0 = 50.0 or 500 mM; I = 1.0 M; pH 6.54 (measured at 25.0
°C).
the experimental results and the calculated data is illustrated in
−
Figure 4. The estimate for pKa(HSO5 ) = 8.43 0.05 (I = 1.0
M) is in good agreement with previously published values.34
The apparent rate constants ka1pp and k2app were found to
decrease linearly with decreasing [H+] in the range from pH
12.3 to 13.5 (Figure S3 in the Supporting Information), where
measured directly. With this approach, we lost the possibility of
determining the activation parameters of R2 but excluded the
possibility of any systematic error arising from the complicated
simultaneous evaluation of two sets of kinetic data. The Eyring
equation was used again to evaluate the activation parameters
from the temperature-dependent study (Figure S4 in the
Supporting Information), and the results are listed in Table 1.
The entropies of activation of reactions R1H, R1, and also R2H
are large negative values close to −100 J mol−1 K−1. Similar
entropies of activation were measured in the reactions of
[H+] ≪ Ka(HSO5 ) and both eqs 2 and 3 are simplified to
−
linear correlations. This observation, together with the results
of the full-scale pH-dependence study, is strong evidence
against the presence of any acid/base-catalyzed processes in the
(H)SO5 + SCN− reaction system.
−
The temperature dependence of the HSO5 + SCN−
−
reaction was studied in phosphate buffer at pH 6.54 (measured
at 25.0 °C), where the rates of the reactions R1 and R2 are
negligible compared to those of R1H and R2H, respectively.
The pH was chosen to be far (at least 1.5 units away) from the
34
−
HSO5 with halide ions (Cl−, Br−, and I−). The analogy
between the oxidation of halides and SCN− is not surprising,
taking into account the well-established classification of SCN−
as a pseudohalide ion. With the methodology presented in ref
34, one can conclude that the rate-determining step in the
−
pKa values of both HSO5 and HOSCN at 25.0 °C, and
phosphate buffer was chosen because its pH shows less than a
−0.05 unit/+10 K dependence on the temperature.54 Thus, all
complications arising from the possible temperature-dependent
shifting of the protolytic equilibria in the reaction system were
averted by carefully choosing the experimental conditions. In
order to obtain information on the temperature dependence of
the rate of reactions R1H (kH1 ) and R2H (k2H), the same
approach as that in the pH-dependence study was used: two
duplicate kinetic experiments were run at two different
[SCN−]0 values at a given temperature. The calculated rate
constants kH1 and kH2 are shown in Figure 5 in the temperature
range from 5.0 to 45.0 °C. Because both R1H and R2H are
elementary reactions, the Eyring equation was used to calculate
the activation parameters by utilizing a nonlinear least-squares
fitting algorithm:55
HSO5 + SCN− (R1H) elementary reaction cannot be a 1e−-
−
−
transfer step. At pH 7, the formal potential of the HSO5 /
SO4•− couple is (1.19 + 0.0591 × log 10−7) = 0.78 V,34
whereas the standard reduction potential of the SCN•/SCN−
couple is 1.63 V.56 Thus, the free energy of the possible 1e−
HSO5 + SCN− (R1H) elementary step would be [(1.63 −
−
0.78) × 96.485] = 82 kJ mol−1, which is much higher than the
measured free energy of activation at 298 K: (26 + 298 ×
0.112) = 59 kJ mol−1. Therefore, similarly to the reactions of
halides, the HSO5 + SCN− (R1H) elementary step is more
−
likely to proceed through an initial 2e− oxidation, probably
involving concerted heterolytic O−O bond cleavage and S−O
bond formation. This mechanism is expected to involve much
more geometric reorganization and characterized by higher
activation entropy than a 1e−-transfer rate-determining step,
which does not include atom transfer before the transition state.
In accordance with the concluding thoughts of ref 34, a
generalization is risky without a sufficiently high number of
reliably measured activation parameters in the literature for the
exp⎜ ⧧ ⎟ exp⎜ −
⎛
⎞
⎠
⎛
⎝
⎞
⎟
⎠
kbT
h
ΔS
R
ΔH⧧
RT
k =
⎝
(4)
where kb is the Boltzmann constant, h is the Planck constant, R
is the universal gas constant, and ΔS⧧ and ΔH⧧ are the
activation entropy and enthalpy. The resulting activation
parameters are listed in Table 1. The temperature dependence
reactions of HSO5 and SO52−. Still the −80 to −110 J mol−1
−
K−1 activation entropies of R1H, R1, and R2H seem to suggest
that all of these processes proceed through 2e−-transfer rate-
determining steps rather than 1e−-transfer steps.
of the oxidation of SCN− with SO5 (R1) was also studied at
2−
pH 13.5. The excess of SCN− used in these kinetic experiments
([SO52−]0 = 5.0 mM and [SCN−]0 = 750 mM) was chosen to
CONCLUSIONS
■
be high to avoid the oxidation of OSCN− by SO5 (R2). In
We showed that peroxomonosulfate ion (HSO5 ) oxidizes
2−
−
other words, the rate of reaction R2 was negligible compared to
R1 under the applied experimental conditions, and k1 could be
thiocyanate ion (SCN−) to hypothiocyanite ion (OSCN−) in
the range from pH 6.9 to 13.5 (kH1 = 2.0 × 102 M−1 s−1 at pH
2154
dx.doi.org/10.1021/ic302544y | Inorg. Chem. 2013, 52, 2150−2156

Inorganic Chemistry
Article
6.89). The oxidation of OSCN− by HSO5− (k2H = 3.3 × 103 M−1
to the kinetic measurements are greatly acknowledged. The
authors are thankful for the Hungarian Scientific Research Fund
[OTKA: K68668 and T49044 (IN64289)], TAMOP 4.2.1./B-
09/1/KONV-2010-007, and 4.2.2-08/12 CHEMIKUT. I.F. is
grateful for support by the Hungarian Academy of Sciences
through the Research Group of Homogeneous Catalysis and
Reaction Kinetics.
s−1 at pH 6.89) also takes place in the reaction system of
HSO5 + SCN−. On the basis of the estimated activation
−
parameters, we concluded that the rate-determining steps in the
above oxidation processes are more likely to be 2e−- than 1e−-
transfer steps. Because the rate constant of the oxidation of
OSCN− is an order of magnitude higher than that of SCN−,
−
these two compounds compete for the oxidant HSO5 , and the
quantitative conversion of SCN− to OSCN− is possible only by
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