Thiocyanogen as an intermediate in the oxidation of thiocyanate by hydrogen peroxide in acidic aqueous solution

Article abstract of DOI:10.1021/ic000561r

The kinetics of the reaction of H₂O₂ with excess SCN^- in acidic media was studied by use of Ti(IV) as an indicator for the concentration of H₂O₂. Pseudo-first-order behavior was realized by this method, and these data confirm the acid-catalyzed rate law and rate constant reported some 40 years ago for this reaction under conditions of excess H₂O₂. Under the same conditions except without Ti(IV), repetitive-scan spectra reveal the formation and decay of an intermediate that absorbs in the UV. In the proposed mechanism, HOSCN is produced in the first step and it is converted rapidly to (SCN)₂ through its equilibrium reaction with SCN^-. The observed intermediate is believed to be (SCN)₂, which decays on a longer time scale. Excellent global fits of this mechanism to the repetitive-scan data are obtained with rate constants constrained by the Ti(IV) data and published previously in our study of the ClO₂/SCN^- reaction. These fits yield a spectrum for (SCN)₂ that is characterized by λ(max) = 297 nm and ε₂₉₇ = 147.M^-1 cm^-1, in fine agreement with our prior report.

Full text of DOI:10.1021/ic000561r

Inorg. Chem. 2000, 39, 5089-5094
5089
Thiocyanogen as an Intermediate in the Oxidation of Thiocyanate by Hydrogen Peroxide in
Acidic Aqueous Solution
James N. Figlar and David M. Stanbury*
Department of Chemistry, Auburn University, Auburn, Alabama 36849
ReceiVed May 23, 2000
The kinetics of the reaction of H2O2 with excess SCN- in acidic media was studied by use of Ti(IV) as an
indicator for the concentration of H2O2. Pseudo-first-order behavior was realized by this method, and these data
confirm the acid-catalyzed rate law and rate constant reported some 40 years ago for this reaction under conditions
of excess H2O2. Under the same conditions except without Ti(IV), repetitive-scan spectra reveal the formation
and decay of an intermediate that absorbs in the UV. In the proposed mechanism, HOSCN is produced in the first
step and it is converted rapidly to (SCN)2 through its equilibrium reaction with SCN-. The observed intermediate
is believed to be (SCN)2, which decays on a longer time scale. Excellent global fits of this mechanism to the
repetitive-scan data are obtained with rate constants constrained by the Ti(IV) data and published previously in
our study of the ClO2/SCN- reaction. These fits yield a spectrum for (SCN)2 that is characterized by λmax ) 297
nm and ꢀ297 ) 147 M-1 cm-1, in fine agreement with our prior report.
Introduction
of the spectrum of (SCN)2 and its decay rate law would serve
as an important test of the mechanism proposed for the ClO2/
SCN- reaction and could make a valuable contribution to the
general field of thiocyanate oxidative chemistry.
The oxidative chemistry of thiocyanate has enjoyed continued
interest in a variety of fields including, but not limited to, the
following: oligooscillating reactions,1-3 outer-sphere electron
transfer,4-6 and catalytic peroxidation by lactoperoxidase7,8 and
other peroxidases.9-13 The mechanistic pathways associated with
the evolution of the oxidized products of thiocyanate have been
an area of strong debate and a source of often contradictory
information.14 In a recent report on the reaction of ClO2 and
SCN-, it was noted that thiocyanogen (SCN)2 is a likely
intermediate in that reaction in 1 M HClO4 and an aqueous
spectrum of this species was extracted.14 The complex mech-
anism proposed for that reaction was intimately tied to the
decomposition of (SCN)2 and offered reasonable rate constants
for a number of steps, most importantly for the hydrolytic
decomposition of (SCN)2. The extracted spectrum for this
species was found to overlap significantly with the broad
absorbance band of the reactant ClO2, which may have adversely
affected the parameters of the fit. An independent confirmation
The present work discloses information on aqueous (SCN)2
based on the reaction of H2O2 with SCN-. This reaction was
investigated in considerable detail about 40 years ago by Wilson
and Harris, although without the benefit of modern spectro-
photometric methods.15,16 It was found that the reaction has a
rate law that is first order in each of the reactants between pH
12 and 4,15 but at higher acidities it is acid-catalyzed.16 A recent
capillary electrophoresis study of the acid-independent reaction
has called the rate law into question, but no data were obtained
bearing on the acid-catalyzed reaction.17 Detailed studies by
Wilson and Harris of the acid-catalyzed kinetics were performed
under conditions of excess H2O2, and under these conditions
the reaction shows mild inhibition by HCN, which is also a
product. When the HCN inhibition can be neglected, the rate
law simplifies to
-d[H2O2]/dt ) k[H2O2][SCN-][H+]
The overall reaction in acidic media is
3H2O2 + SCN- f HSO4- + HCN + 2H2O
(1)
(1) Chinake, C. R.; Mambo, E.; Simoyi, R. H. J. Phys. Chem. 1994, 98,
2908-2916.
(2) Doona, C. J. J. Phys. Chem. 1995, 99, 17059-17060.
(3) Alamgir, M.; Epstein, I. R. J. Phys. Chem. 1985, 89, 3611-3614.
(4) Alfassi, Z. B. In N-Centered Radicals; Alfassi, Z. B., Ed.; John Wiley
& Sons: New York, 1998; pp 549-562.
(5) Nord, G. Comments Inorg. Chem. 1992, 13, 221-239.
(6) Hung, M.-L.; Stanbury, D. M. Inorg. Chem. 1994, 33, 4062-4069.
(7) Pollock, J. R.; Goff, H. M. Biochim. Biophys. Acta 1992, 1159, 279-
285.
(2)
Deviations from this ideal stoichiometry can occur through the
formation of S(CN)2 and its hydrolysis products. S(CN)2 was
proposed to arise through the reaction of HCN with an
intermediate, HOSCN. The overall proposed mechanism had
as initial steps
(8) The Lactoperoxidase System: Chemistry and Biological Significance;
Pruitt, K. M., Tenovuo, J. O., Eds.; Marcel Dekker: New York, 1985.
(9) Slungaard, A.; Mahoney, J. R. J. Biol. Chem. 1991, 266, 4903-4910.
(10) van Dalen, C. J.; Whitehouse, M. W.; Winterbourn, C. C.; Kettle, A.
J. Biochem. J. 1997, 327, 487-492.
(11) Walker, J. V.; Butler, A. Inorg. Chim. Acta 1996, 243, 201-206.
(12) Pruitt, K. M.; Mansson-Rahemtulla, B.; Baldone, D. C.; Rahemtulla,
F. Biochemistry 1988, 27, 240-245.
H2O2 + SCN- + H+ f HOSCN + H2O
(slow) (3)
(13) Adak, S.; Mazumdar, A.; Banerjee, R. K. J. Biol. Chem. 1997, 272,
11049-11056.
(15) Wilson, I. R.; Harris, G. M. J. Am. Chem. Soc. 1960, 82, 4515-4517.
(16) Wilson, I. R.; Harris, G. M. J. Am. Chem. Soc. 1961, 83, 286-289.
(17) Christy, A. A.; Egeberg, P. K. Talanta 2000, 51, 1049-1058.
(14) Figlar, J. N.; Stanbury, D. M. J. Phys. Chem. A 1999, 103, 5732-
5741.
10.1021/ic000561r CCC: $19.00 © 2000 American Chemical Society
Published on Web 10/12/2000

5090 Inorganic Chemistry, Vol. 39, No. 22, 2000
Figlar and Stanbury
H2O2 + HOSCN f HO2SCN + H2O
(fast) (4)
Further reactions of HO2SCN would lead to final products. Thus,
the reaction of HCN with HOSCN could diminish the rate of
consumption of H2O2 by as much as a factor of 2. More recent
evidence for HOSCN as the product in the first step is the
adherence of the third-order rate constant to a LFER (linear
free energy relationship) correlating rates of reactions of
nucleophiles with H2O2, all of which are believed to have
analogous rate-limiting steps.18 A potential complication in the
reaction of H2O2 with SCN- is copper catalysis, which can lead
to oscillations and bistability.19-21 Apparently this catalysis
occurs only at high pH such that catalysis by adventitious copper
ions can be neglected under the acidic conditions relevant to
the present study.
In the course of our study of the reaction of ClO2 with SCN-
in highly acidic media, we developed a mechanism in which
the equilibrium between HOSCN and (SCN)2 was central:
(SCN)2 + H2O h HOSCN + SCN- + H+
(5)
Figure 1. Multiwavelength rise and fall profile of the H2O2/SCN-
reaction. The raw kinetic data were taken on the HP 8453 spectropho-
tometer at 25 °C in 1 M HClO4, path length 1 cm, [H2O2]o ) 15 mM,
[SCN-]o ) 180 mM. The data displayed are sparsed to highlight the
rise-fall behavior in the UV.
This equilibrium is as expected given that SCN- is generally
regarded as a pseudohalide. In view of equilibrium 5, it is
reasonable to anticipate that the reaction of H2O2 with SCN-
would yield (SCN)2 as an intermediate under conditions of high
concentrations of H+ and SCN-. Moreover, through the study
of the reaction with low concentrations of H2O2, the yield of
HCN would be low enough that the complicating reaction of
HOSCN with HCN could be avoided.
Accordingly, the present paper describes spectrophotometric
studies of the reaction of H2O2 with excess SCN- in acidic
media. These studies afford an opportunity to test the ideas
described above, to obtain a spectrum of aqueous (SCN)2
without the spectral interferences encountered in our study of
the ClO2/SCN- reaction, and to verify the kinetics of decom-
position of (SCN)2.
Kinetic Methods. Kinetic measurements were conducted on a HP
8453 diode array spectrophotometer with a 1 cm cell. The temperature
was maintained at 25 °C with a Brinkmann RM6 circulating water bath.
In most experiments, the ionic strength (µ) was maintained at 1.25 M
through the addition of appropriate amounts of HClO4 and LiClO4.
Equal volumes of reactants were delivered to the cell which was then
capped and vigorously stirred prior to measurements. Initial experiments
indicated a photochemical reaction which deposited a white precipitate
(presumably sulfur arising from SCN- photolysis23) onto the cell walls
at the beam point of entry and exit. This interference was negated by
placing a UV optical cutoff filter (HP 08451-60302, 50% transmittance
at 265 nm) in the light path prior to the sample. Spectral changes were
monitored in the UV-vis region (265-400 nm) and analyzed addition-
ally by using the Specfit software package.24 Mechanistic simulations
were also performed by utilizing the Specfit package.
Experimental Section
Reagents. Aqueous solutions of hydrogen peroxide were prepared
from sodium stannate stabilized and nonstabilized 30% stocks purchased
from Fischer. The stock solutions were standardized iodometrically with
freshly prepared thiosulfate (Fischer). Distilled deionized water was
used for all dilutions and experiments. NaSCN (Baker and Fischer)
was recrystallized from water, and thiocyanate concentrations were
determined idometrically. Fischer 75% perchloric acid was used in these
experiments without further purification. Titanium(IV) chloride (99.9%)
was purchased from Aldrich and used without further purification.
Results
Detection of an Intermediate. Monitoring the direct reduc-
tion of H2O2 in excess thiocyanate (280 mM) in 1 M HClO4
produces spectroscopic traces indicative of simple rise-fall
behavior in the UV as shown by the overlay spectra in Figure
1. An alternative view is given by the single-wavelength trace
at 410 nm shown in Figure 2. Qualitatively similar behavior
was observed for a series of reactions with [SCN-]0 decreasing
to 10 mM, although flooding conditions do not apply for the
latter. Despite this seemingly uncomplicated biphasic kinetics,
fitting the data at high [SCN-] was not achieved to any degree
of success by employing a routine double-exponential function
corresponding to a A f B f C mechanism. Nevertheless, it is
clear that an intermediate is produced, which absorbs in the
UV. Its maximum absorbance at 295 nm is comparable to the
absorbance reported to arise from (SCN)2 produced in the
reaction of ClO2 with SCN- under similar conditions.14 The
intermediate is produced on a time scale comparable to that for
the consumption of H2O2 as predicted by Wilson and Harris’
rate law for the reaction of H2O2 with SCN-.16 Qualitatively,
2+
Synthesis of Peroxotitanium(IV). TiO2 was prepared by the
methods described by Thompson22 with few alterations. Fresh TiCl4
(5 mL) was dispensed in a N2 purged glovebox and immediately
transferred to an exhaust hood and added dropwise to a 20-50 mL
portion of water (note this reaction is extremely exothermic and must
be performed with caution). The resultant solution was passed through
a Dowex 50W-X8 H+ cation exchange column which was pretreated
with three passes of strong HCl followed by three passes of pH 2.0
HClO4. The TiO2+ was eluted with 2 M HClO4 and collected in 10
equivalent volume fractions. Concentrations of each fraction were
determined spectroscopically as described previously22 by adding an
excess of peroxide to the TiO2+ solutions and measuring the absorbance
of the resultant TiO2 complex at 410 nm (ꢀ410 ) 717 M-1 cm-1).
2+
(18) Hoffmann, M.; Edwards, J. O. Inorg. Chem. 1977, 16, 3333-3338.
(19) Orba´n, M. J. Am. Chem. Soc. 1986, 108, 6893-6898.
(20) Luo, Y.; Orba´n, M.; Kustin, K.; Epstein, I. R. J. Am. Chem. Soc. 1989,
111, 4541-4548.
(21) Orba´n, M.; Kurin-Cso¨rgei; Ra¨bai, G.; Epstein, I. R. Chem. Eng. Sci.
2000, 55, 267-273.
(22) Thompson, R. C. Inorg. Chem. 1984, 23, 1794-1798.
(23) Dogliotti, L.; Hayon, E. J. Phys. Chem. 1968, 72, 1800-1807.
(24) Binstead, R. A.; Zuberbu¨hler, A. D. SPECFIT; Spectrum Software
Associates: Chapel Hill, NC, 1995.

Thiocyanogen as an Intermediate
Inorganic Chemistry, Vol. 39, No. 22, 2000 5091
Table 1. Mechanism for the Reaction of Hydrogen Peroxide and Thiocyanate
rate
no.
reaction
constant
SR1
SR2
SR3
SR4
SR5
H+ + SCN- + H2O2 f HOSCN + H2O
2HOSCN f HO2SCN + SCN- + H+
2HO2SCN f O3SCN- + HOSCN + H+
O3SCN- + H2O f HSO4- + HCN
k1
kdisp
3.3 × 10-2 M-2 s-1
8.0 × 102 M-1 s-1
2 × 108 M-1 s-1
5 × 108 s-1
(SCN)2 + H2O h HOSCN + H+ + SCN-
Khyd
4.4 × 103 s-1
3.05 × 105 M-2 s-1
Table 2. Kinetic Data for the Reaction of TiO2 and SCN- in
2+
High Acid Concentrationsa
104 × kobs, s-1b replicates [H2O2]tot,o, mM [Ti]tot, mM [H+], M
1.12(5)
1.74(19)
4.70(24)
9.03(41)
13.8(7)
3
4
3
3
4
3
3
3
4
4
3
3
3
0.3
0.3
0.3
0.3
0.3
0.3
0.3
0.3
4.78
3.00
1.20
1.0
1.0
1.0
1.0
1.0
1.0
0.1
0.5
1.0
1.0
1.0
1.0
1.0
0.598
0.239
0.120
0.239
0.239
0.239
0.624
0.624
0.624
0.624
22.5(2)
2.76(72)
9.49(39)
13.8(7)
6.86(161)
5.31(61)
5.21(70)
5.35(36)
0.3
0.01
0.025
0.05
0.1
a [SCN-] ) 250 mM, 25 °C, µ ) 1.25 M LiClO4 where applicable.
b Measured the consumption of TiO22+ at 410 nm. Parenthetical values
represent the error as 3 standard deviations.
the recent report that the rate law for the acid-independent
reaction given by Wilson and Harris is incorrect.17
Peroxotitanium as an Indicator in the Oxidation of SCN-
by H2O2. Preliminary studies of the reaction of H2O2 with
excess SCN- revealed that there is no simple direct spectro-
photometric method to monitor the loss of H2O2 because of the
high background absorbance from SCN- in the UV, the
photochemical complications arising from monitoring in the far-
UV, and the absorbance of the intermediate in the nearer regions
of the UV. These difficulties can be avoided by employing an
appropriate UV cutoff filter and by using peroxotitanium(IV)
(TiO22+) as a nonparticipant chemical indicator that absorbs in
the visible region. This use of peroxotitanium(IV) is suggested
by the work of Lydon and Thompson, which shows that the
complex is formed rapidly and reversibly and that it is generally
much less reactive than H2O2 itself.25 This complex forms as
given by
Figure 2. Single-wavelength (295 nm) decay profile of the H2O2/SCN-
reaction and results of the numerical integration of the multiwavelength
data to the mechanism presented in Table 1. Best fit results associated
with the rate constants are displayed therein. Open circles represent
the experimental data points, and the solid line represents the fit. Top
box displays the residuals of the fit. Data taken at 25 °C in 1 M HClO4,
path length 1 cm, [H2O2]o ) 15 mM, [SCN-]o ) 180 mM.
the intermediate decays on the same time scale as we found for
the decay of (SCN)2 in the reaction of ClO2 with SCN-.14
On the basis of these observations, a reasonable hypothesis
for the mechanism (Table 1) can be developed by combining
portions of the mechanisms previously proposed for the reactions
of SCN- with H2O2 and ClO2: SCN- in acid reacts with H2O2
to give HOSCN, which then forms (SCN)2 through an acid-
driven equilibrium comproportionation with SCN-. Decay of
(SCN)2 obeys second-order kinetics because of the second-order
decomposition of HOSCN. There appears to be no analytical
solution to the system of differential equations arising from this
mechanism. Thus, as is described below, we use a numerical
ODE (ordinary differential equation) solver in order to fit the
parameters of the mechanism to the data. The task is greatly
simplified as the number of adjustable parameters is reduced,
and the parameter most directly accessible is the rate constant
for the consumption of H2O2. Although the rate law for the
acidic reaction of SCN- with H2O2 appears to be well
established, the data upon which it is based were obtained under
conditions of a large excess of H2O2.15,16 As the intermediate
reported in the present study is generated with excess SCN-, it
is necessary to verify the rate law for H2O2 consumption under
these conditions. Further motivation for this reinvestigation is
TiO2+ + H2O2 h TiO22+ + H2O
KTi
(6)
The stability constant for this equilibrium was reported by
Thompson in 1 M perchloric acid and was found to be 8.7 ×
103 M-1 22
Under conditions similar to those suggested by
.
Lydon and Thompson,25 pseudo-first-order behavior can be
observed for the reaction of H2O2 and SCN- by monitoring the
decay in absorbance of TiO22+ at 410 nm over time. This method
achieves pseudo-first-order kinetics for a variety of conditions,
and a typical reaction trace is shown in Figure S1 of the
Supporting Information. The pseudo-first-order rate constants
can be evaluated by a single-exponential fit, and the results are
shown in Table 2. Note that under these conditions most of the
2+
peroxide is bound as TiO2
.
Interpretation of these pseudo-first-order rate constants is
achieved by considering the following derivation. Under condi-
tions applicable to these experiments, the rate of decay of
(25) Lydon, J. D.; Thompson, R. C. Inorg. Chem. 1986, 25, 3694-3697.

5092 Inorganic Chemistry, Vol. 39, No. 22, 2000
Figlar and Stanbury
value for k1. On the other hand, the slope has a well-defined
value of (2.2 ( 0.05) × 10-6 M-1 s-1 for k1/KTi. From the
literature value for KTi given above, this ratio leads to a value
of 1.9 × 10-2 M-2 s-1 for k1. This result is in fine agreement
with that reported by Wilson and Harris,16 and it shows that
rate law 1 applies under conditions of both excess SCN- and
excess H2O2. Moreover, our value for k1 implies an intercept
of 210 s for Figure 3, which is in agreement with the observed
result. The agreement between our raw data and eq 10, along
with the agreement between our value for k1 and that reported
by Wilson and Harris, gives strong support to the assumption
that Ti(IV) acts only as an indicator and, if anything, deactivates
H2O2 as suggested by Lydon and Thompson.25
Discussion
Table 1 presents a five-step mechanism that accounts for our
observations and is consistent with the chemistry reported in
the literature. This mechanism consists of two sets of reactions
which account for the rise-fall behavior noted in the UV for
the SCN-/H2O2 reaction. The first set, consisting of reactions
SR1 and SR5, describes the direct two-electron oxidative
conversion of SCN- to hypothiocyanous acid, which produces
a discernible intermediate, (SCN)2, in 1 M H+ by equilibrium
comproportionation. The second set of reactions, SR2-SR4 and
SR5 again, is widely accepted and accounts for the second-
order decay of (SCN)2 to the likely and well-documented
products of hydrolysis.26,27
Our proposed mechanism differs from that proposed by
Wilson and Harris16 both by the number of steps and by the
proposed end products of the reaction. Most notably, the
mechanism of Wilson and Harris omits the formation of (SCN)2
through equilibrium SR5. Another difference is that Wilson and
Harris propose the reaction of HOSCN with H2O2 as the major
fate of HOSCN. While this reaction with H2O2 may be
appropriate under the conditions of excess H2O2 used by Wilson
and Harris, we find it necessary to use a pathway second-order
in HOSCN under our conditions of excess SCN-. Our mech-
anism omits the reaction of HOSCN with HCN, as we believe
that it plays an insignificant role with the low concentrations
of HCN generated in the present study. Finally, Wilson and
Harris propose hydrolysis of HO2SCN where we propose its
disproportionation; in fact, our data do not allow us to
distinguish between these two possibilities.
Success of the Model. The mechanism in Table 1 provides
for the net stoichiometry given by reaction 2 as well as adhering
to the rate law given in eq 1. The numerical integration of the
system of differential equations arising from this mechanism
provides compelling evidence of its applicability for this reaction
system. These simulations were performed with the Bulirsch-
Stoer stiff integration routine available in the Specfit software
package, making use of the reactions in Table 1 with the
associated rate constants. The spectra of H2O2, (SCN)2, and
products were treated as variables in these fits. The results at
295 nm, shown in Figure 2, provide quantitative agreement
between the observed kinetics and simulations as demonstrated
by the fine overlap of the traces in 1 M H+ with 180 mM SCN-.
The spectra of H2O2, (SCN)2, and products generated by these
fits are shown in Figure 4. The spectra obtained are entirely
consistent with expectations for H2O2 and products, while the
spectral trace for thiocyanogen is fairly consistent with expecta-
tions in terms of λmax and molar absorptivity from that previously
Figure 3. Plot of 1/kobs vs [Ti]tot at [H2O2]tot,0 ) 0.3 mM, [SCN-] )
250 mM, [H+] ) 1 M, µ ) 1.25 M (HClO4). The linear plot has a
slope of (1.84 ( 0.04) × 106 s/M and an intercept of 130 ( 90 s with
R ) 0.9991.
hydrogen peroxide in the absence of Ti(IV) according to Wilson
and Harris reduces to eq 1.16 Because of the rapid equilibria
that are established in the presence of Ti(IV), we define the
following two quantities:
[H2O2]tot ) [TiO22+] + [H2O2]
[Ti]tot ) [TiO22+] + [TiO2+]
(7)
(8)
In the presence of excess Ti(IV), the concentration of free H2O2
is essentially buffered through reaction 6. Moreover, it is
assumed as suggested by Lydon and Thompson that SCN-
2+ 25
reacts with H2O2 but not TiO2
.
As described in the
Supporting Information, when [Ti]tot . [TiO22+] these consid-
erations lead to rate law 9:
2+
2+
d[TiO2
]
k1[SCN-][H+][TiO2
KTi[Ti]tot + 1
]
-
)
(9)
dt
2+
Equation 9 leads to pseudo-first-order kinetics for loss of TiO2
with kobs defined by
k1[SCN-][H+]
kobs
)
(10)
KTi[Ti]tot + 1
Equation 10 predicts a first-order dependence of kobs on [H+]
and a zero-order dependence on [H2O2]tot,0, both of which are
demonstrated by the data in Table 2. The inverse of eq 10 is
KTi[Ti]tot
k1[SCN-][H+] k1[SCN-][H+]
1
kobs
1
)
+
(11)
which implies a linear dependence of 1/kobs on [Ti]tot. Figure 3
displays such a plot for data obtained in 1 M HClO4 with 250
mM SCN-. This linear graph has a slope of (1.84 ( 0.04) ×
106 s/M and an intercept of 133 ( 90 s. Note that the intercept
is not well defined by the data, and its value is highly sensitive
to the results obtained at low [Ti]tot. The data at low [Ti]tot,
however, were obtained under conditions that do not comply
strictly with the approximation that [Ti]tot . [TiO22+]. These
same reactions also displayed small deviations from pseudo-
first-order behavior in the initial portions of the traces. Accord-
ingly, the intercept in Figure 3 does not lead directly to a reliable
(26) Bjerrum, N.; Kirschner, A. Chem. Abstr. 1919, 13, 1057-1060.
(27) Stedman, G.; Whincup, P. A. E. J. Chem. Soc. A 1969, 1145-1148.

Thiocyanogen as an Intermediate
Inorganic Chemistry, Vol. 39, No. 22, 2000 5093
changes in the fits, and therefore, they were removed. Addition-
ally, the pseudo-first-order decay pattern established in the
2+
TiO2 reactions suggested no significant role for OSCN-, as
it had in the autocatalytic reaction of ClO2.14
The rate constant used in SR1, (3.3 ( 0.3) × 10-2 M-2 s-1
,
was established by Wilson and Harris16 and confirmed through
our work as described above. Further evidence for its validity
is demonstrated by the numerical fitting of this mechanism. All
data sets can be fit by this mechanism, and when this rate
constant is allowed to vary while the rate constants related to
the decay of thiocyanogen remain constant, the stiff integrator
converges on an answer which produces this value within
reasonable limits ((10%). This was an important result because
it allowed for the further refinement of the second phase of the
mechanism for which a good deal of conflicting evidence has
been reported in the literature.
2. Hydrolysis of (SCN)2, SR2-SR5. As was the case in the
reduction of ClO2 by thiocyanate under strongly acidic condi-
tions,14 the hydrolytic equilibrium of (SCN)2 represented by Khyd
is a critical step in these simulations. Similarly, the rate constant
Figure 4. Spectra generated by a three-component fit of the mechanism
in Table 1 with the rate constants as given. Data taken at 25 °C in 1 M
HClO4, path length 1 cm, [H2O2]o ) 15 mM, [SCN-]o ) 180 mM.
Circles represent the data associated with (SCN)2, squares represent
the proxy spectra of products, and triangles represent the spectrum for
H2O2. All data are sparsed at an interval of 5 nm for clarity. The aqueous
spectrum for (SCN)2 extracted from these kinetic fits is characterized
kdisp for the disproportionation of HOSCN is also crucial, while
SR3 and SR4 are projected to be rapid steps that merely account
for the formation of products, not for affecting the quality of
the fit. As we have discussed for the ClO2/SCN- reaction, the
decomposition kinetics of (SCN)2 in strong acid obeys a rate
2
law that permits determination of the value for kdispKhyd but
by λmax ) 297 nm and ꢀ297 ) 147 M-1 cm-1
.
not the independent values for kdisp and Khyd. Despite the less
complex mechanism invoked and the smaller degree of spectral
interferences associated with the H2O2/SCN- system, the values
of kdisp and Khyd are still inseparable. As a consequence, fits
were performed with the rate constants held constant, the
component spectra being treated as variables. By trial and error
with manual adjustment of the rate constants, it was found that
when Khyd was increased by a factor of 1.5 or kdisp was decreased
by roughly the same magnitude from the previously reported
values, excellent fits to the raw data were obtained. In Table 1,
we arbitrarily show the altered value of Khyd while retaining
the prior value for khyd, recognizing that equally good fits can
be obtained for other combinations of rate constants within
reported.14 We note that step SR5 causes the ratio [HOSCN]/
[(SCN)2] to be a constant during any one experiment, and thus,
the spectra of HOSCN and (SCN)2 cannot be deconvoluted in
these fits. Hence, the (SCN)2 spectrum returned by the fits is
actually the sum of the contributions of these two species. On
the other hand, step SR5 causes the ratio [HOSCN]/[(SCN)2]
to be dependent on the concentration of SCN-. Simulations
performed with the same parameters gave excellent fits for
reactions with [SCN-] as low as 10 mM. In these latter
reactions, the concentration of SCN- is not in large excess and
hence decreases significantly during the reaction; this decrease
does not pose a problem in the integrations as it is accounted
for explicitly in the model. These simulations yielded spectra
of (SCN)2 that were essentially identical, which implies that
HOSCN does not absorb significantly in this spectral region.
Details of the Mechanism. 1. Direct Reaction of H2O2 and
SCN-, SR1. The initial step of the acidic reaction of H2O2 and
SCN- is accounted for by the simple two-electron reduction of
H2O2 to H2O and HOSCN. This type of reduction (nuclephilic
displacement at peroxide) is commonly invoked for peroxide
reactions where there is a lack of species present to stabilize
“Fenton-like” radicals produced by one-electron reductions.16,18,28
Similarly, the two-electron oxidation product of SCN- is a
commonly reported intermediate for these types of processes
and accounts, through widely accepted decomposition pathways,
for the products reported for this reaction. A value of 5.3 has
been reported for the pKa of HOSCN,29 while the pKa of HSCN
is -0.9.30 Under the acidity conditions of the present study,
the concentration of OSCN- is, therefore, completely negligible
and the concentration of HSCN is never very significant.
Inclusion of these equilibria in the mechanism produced no
2
reasonable limits, so long as the value of kdispKhyd is unper-
turbed.
In our prior study of the ClO2/SCN- reaction, a 14-step
mechanism was required to explain the complex behavior.14 The
hydrolytic decomposition of (SCN)2 played an important role
in that study, although the spectral data upon which it was based
were somewhat obscured by the strong UV absorbance of ClO2.
An important result from the current study is thus the strong
support it gives for the mechanism proposed for the ClO2/SCN-
reaction. Despite this corroboration of our previous work, we
believe independent studies of (SCN)2 hydrolysis are required
to gain a more complete understanding of this process. Direct
determination of the magnitude of Khyd remains a goal, and the
question of whether HO2SCN decays through hydrolysis or
disproportionation needs to be answered. A conflict that remains
to be resolved is the rapid decomposition of HOSCN that our
model requires and the relative stability of HOSCN reported
by others in less acidic media.7,17,29,31 Efforts to probe these
issues in the current study were hampered by slowness of SR1,
which leads to formation of (SCN)2 occurring on a time scale
comparable to that of its decay.
(28) Masarwa, M.; Cohen, H.; Meyerstein, D.; Hickman, D. L.; Bakac,
A.; Espenson, J. H. J. Am. Chem. Soc. 1988, 110, 4293-4297.
(29) Thomas, E. L. Biochemistry 1981, 20, 3273-3280.
(30) Smith, R. M.; Martell, A. E.; Motekaitis, R. J. NIST Critically Selected
Stability Constants of Metal Complexes Database, 4.0; U.S. Depart-
ment of Commerce: Gaithersburg, MD, 1997.
(SCN)2 UV Spectrum. The spectrum of (SCN)2 derived by
kinetic fitting is shown in Figure 4 and is similar to our
(31) Aune, T. M.; Thomas, E. L. Eur. J. Biochem. 1977, 80, 209-214.

5094 Inorganic Chemistry, Vol. 39, No. 22, 2000
Figlar and Stanbury
previously reported spectrum of this species in aqueous solu-
tion.14 The importance of this second finding involves the clarity
of this broad band absorbance from which a more definitive
λmax (297 nm) is extracted. The changes in UV absorbance
throughout these experiments are almost entirely due to (SCN)2
and do not suffer from alternative species overlap, making this
evaluation more reliable. The spectrum is characterized by λmax
) 297 nm and ꢀ297 ) 147 M-1 cm-1 and is in good agreement
with our previous work14 as well as that shown by Bacon and
Irwin in CCl4.32 Dolbear and Taube cite a report on the spectrum
of (SCN)2 in aqueous solution in the Ph.D. thesis of Long;33
the λmax (302 nm) is consistent with our findings, but the molar
absorptivity (840 M-1 cm-1) is roughly 6 times larger. We
attribute this disagreement to the preliminary nature of the cited
work.
Through the use of peroxotitanium(IV) as an indicator for this
reaction, pseudo-first-order behavior for the loss of H2O2 was
achieved, confirming the rate constant previously reported, while
at the same time limiting the variables for nonlinear fitting
routines. By the fitting of the multiwavelength data to a logical
mechanism for the rise and fall of thiocyanogen, an excellent
spectrum for this species in aqueous solution was extracted. This
work continues the elucidation of often contradictory informa-
tion regarding the decay of thiocyanogen in aqueous solution.
Furthermore, this research supports the importance of thiocy-
anogen in the oxidative reactions of thiocyanate in acidic media.
Acknowledgment. The NSF is gratefully acknowledged for
funding this research. Professor Fred Anson (Caltech) is thanked
for graciously acting as a sabbatical host to D.M.S. during the
preparation of this manuscript. Helpful conversations with Prof.
R. C. Thompson (Missouri) are gratefully acknowledged.
Conclusions
The UV-monitored kinetics of the H2O2/SCN- reaction under
strongly acidic conditions provide definitive evidence for the
appearance and decay of an oxidative thiocyanate intermediate.
Supporting Information Available: A single-wavelength (410 nm)
decay profile of the reaction of TiO22+ and SCN- in 1 M HClO4 at 25
°C (Figure S1) and the derivation of eq 9. This material is available
free of charge via the Internet at http://pubs.acs.org.
(32) Bacon, R. G. R.; Irwin, R. S. J. Chem. Soc. 1958, 778-784.
(33) Dolbear, G. E.; Taube, H. Inorg. Chem. 1967, 6, 60-64.
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