Kinetics and mechanism of the comproportionation of hypothiocyanous acid and thiocyanate to give thiocyanogen in acidic aqueous solution

Article abstract of DOI:10.1021/ic061470i

The kinetics of comproportionation of hypothiocyanous acid (HOSCN) and thiocyanate (SCN^-) to give thiocyanogen ((SCN)₂) in acidic aqueous solutions have been determined by double-mixing stopped-flow UV spectroscopy. Hypothiocyanite (OSCN^-) was generated at pH 13 by oxidation of excess SCN^- with hypobromite (OBr), followed by a pH jump to acidic conditions ([H⁺] = 0.20-0.46 M). The observed pseudo-first-order rate constants exhibit first-order dependencies on [H ⁺] and [SCN^-] with overall third-order kinetics. The corresponding kinetics of hydrolysis of (SCN)₂ have also been examined. Under conditions of high (and constant) [H⁺] and [SCN ^-], the kinetics exhibit second-order behavior with respect to [(SCN)₂] and complex inverse dependences on [H⁺] and [SCN^-]. Under conditions of low [H⁺] and [SCN ^-], the kinetics exhibit first-order behavior with respect to [(SCN)²] and independence with respect to [H⁺] and [SCN^-]. We attribute this behavior to a shift in the rate-limiting step from disproportionation of HOSCN (second-order dependency on [(SCN) ₂]) to rate-limiting hydrolysis (first-order dependency on [(SCN)₂]). Thus, we have determined the following equilibrium constant by the kinetic method: (SCN)₂ + H₂O ? HOSCN + SCN^- + H⁺; K_hyd = [HOSCN][SCN ^-][H⁺]/[(SCN)₂] = K_hyd/k _comp = 19.8(±0.7) s^-1/ 5.14(±0.07) × 10³ M^-2 s^-1 = 3.9 × 10^-3 M².

Full text of DOI:10.1021/ic061470i

Inorg. Chem. 2007, 46, 285−292
Kinetics and Mechanism of the Comproportionation of Hypothiocyanous
Acid and Thiocyanate to Give Thiocyanogen in Acidic Aqueous Solution
Pe´ter Nagy, Kelemu Lemma, and Michael T. Ashby*
Department of Chemistry and Biochemistry, UniVersity of Oklahoma, Norman, Oklahoma 73019
Received August 4, 2006
The kinetics of comproportionation of hypothiocyanous acid (HOSCN) and thiocyanate (SCN-) to give thiocyanogen
((SCN)2) in acidic aqueous solutions have been determined by double-mixing stopped-flow UV spectroscopy.
Hypothiocyanite (OSCN-) was generated at pH 13 by oxidation of excess SCN- with hypobromite (OBr-), followed
by a pH jump to acidic conditions ([H+]
) 0.20−0.46 M). The observed pseudo-first-order rate constants exhibit
first-order dependencies on [H+] and [SCN-] with overall third-order kinetics. The corresponding kinetics of hydrolysis
of (SCN)2 have also been examined. Under conditions of high (and constant) [H+] and [SCN-], the kinetics exhibit
second-order behavior with respect to [(SCN)2] and complex inverse dependences on [H+] and [SCN-]. Under
conditions of low [H+] and [SCN-], the kinetics exhibit first-order behavior with respect to [(SCN)2] and independence
with respect to [H+] and [SCN-]. We attribute this behavior to a shift in the rate-limiting step from disproportionation
of HOSCN (second-order dependency on [(SCN)2]) to rate-limiting hydrolysis (first-order dependency on [(SCN)2]).
Thus, we have determined the following equilibrium constant by the kinetic method: (SCN)2
SCN- H+; Khyd [HOSCN][SCN-][H+]/[(SCN)2] 0.7) s-1/ 5.14(
khyd/kcomp 19.8( 0.07)
10- M2.
+
×
H2O
h
HOSCN
+
2
1
+
3
)
)
)
±
±
103 M- s-
)
3.9
×
Introduction
very acidic conditions by oxidation of SCN-.6-9 Stanbury’s
first studies of the oxidation of SCN- by ClO2 and H2O2
were hampered by the sluggish kinetics of the oxidation
reactions and subsequent decomposition reactions of (SCN)2
that competed with its formation.7,8 This problem was
overcome by employing Cl2/HOCl as an oxidant mixture,
which apparently produces (SCN)2 within the mixing time
of a stopped-flow experiment.9 Stanbury’s investigations have
provided evidence that (SCN)2, in the presence of excess
Relatively little is known about the reaction chemistry of
the human defense factor hypothiocyanite (OSCN-),1 in part
because it is unstable, but also because until recently the
only recognized means of preparing it has been the enzyme-
catalyzed oxidation of SCN- by H2O2 and hydrolysis of
(SCN)2. The enzymatic process is limited to near-neutral pH,
which is a condition for which OSCN- is relatively unstable.2
Early studies of OSCN- have included hydrolysis of (SCN)2,3
a reaction that is usually initiated by the heterogeneous
extraction of (SCN)2 from organic solvents (in which it was
prepared by reaction of lead(II) thiocyanate with bromine)
into aqueous solutions, a technique that is not conducive to
studying fast reactions.4,5 Stanbury has recently investigated
the aqueous chemistry of (SCN)2, which was generated under
-
SCN-, is in rapid equilibrium with trithiocyanate (SCN)3
-
(cf. the corresponding equilibrium between I2 and I3 ):6
-
(SCN)2 + SCN- h (SCN)3 K1 ) 0.43 M-1
(1)
More recently, Stanbury et al. have investigated the decom-
position of (SCN)2 under acidic conditions.9 These studies
have yielded an equilibrium constant (K2) for the hydrolysis
of (SCN)2 to give hypothiocyanous acid (HOSCN) and
SCN-, as well as a rate constant for the disproportionation
* To whom correspondence should be addressed. E-mail: MAshby@
ou.edu.
(1) Ihalin, R.; Loimaranta, V.; Tenovuo, J. Arch. Biochem. Biophys. 2006,
445, 261-268.
(2) Pruitt, K. M., Tenovuo, J. O., Eds. Immunology Series: The Lactop-
eroxidase System: Chemistry and Biological Significance; Dekker:
New York, 1985; Vol. 27, p 257.
(3) Wood, J. L. Org. React. 1946, 3, 240-266.
(4) Aune, T. M.; Thomas, E. L. Eur. J. Biochem. 1977, 80, 209-214.
(5) Hoogendoorn, H.; Piessens, J. P.; Scholtes, W.; Stoddard, L. A. Caries
Res. 1977, 11, 77-84.
(6) Barnett, J. J.; Stanbury, D. M. Inorg. Chem. 2002, 41, 164-166.
(7) Figlar, J. N.; Stanbury, D. M. Inorg. Chem. 2000, 39, 5089-5094.
(8) Figlar, J. N.; Stanbury, D. M. J. Phys. Chem. A 1999, 103, 5732-
5741.
(9) Barnett, J. J.; McKee, M. L.; Stanbury, D. M. Inorg. Chem. 2004, 43,
5021-5033.
10.1021/ic061470i CCC: $37.00
Published on Web 12/13/2006
© 2007 American Chemical Society
Inorganic Chemistry, Vol. 46, No. 1, 2007 285

Nagy et al.
standardized HCl solution using phenolphthalein as an indicator.
The buffer solutions were prepared from the solids NaH2PO4‚H2O,
Na2HPO4, and Na3PO4‚12H2O, the ionic strength was adjusted with
NaClO4, and the pH was adjusted with NaOH or HClO4. Stock
solutions of NaSCN were prepared after drying in a 150 °C oven
to constant weight. Stock solutions of NaOCl were prepared by
sparging Cl2 into a 0.3 M solution of NaOH. The sparging was
stopped when the [OCl-] achieved ca. 100 mM, as determined
spectrophometrically (ꢀ(OCl-)292nm ) 350 M-1 cm-1). Solutions
of NaOBr were prepared by adding Br2 to ice-cold solutions of
NaOH.19 Solutions of OBr- were standardized spectrophotometri-
cally at 329 nm (ꢀ329 ) 332 M-1 cm-1).19 The solutions of OCl-
and OBr- were used within 2 h of the preparations to minimize
errors due to decomposition.
UV/Visible Spectroscopy. Electronic spectra were measured
using a HP 8452A diode-array spectrophotometer or the mono-
chromator of a HI-TECH SF-61 DX2 stopped-flow instrument.
pH Measurements. The [OH-] for the unbuffered solutions was
determined by acid-base titration against standardized HCl solu-
tions. HClO4 and HCl were standardized against bicarbonate. The
[H+] of the buffered solutions were determined with an Orion Ion
Analyzer EA920 using an Ag/AgCl combination pH electrode. The
ionic strength was kept constant at 1.0 M for all solutions (NaClO4
+ HClO4 + NaSCN + NaH2PO4). All pH measurements were
corrected for the “Irving factor” of the working medium. A value
of pKw of 13.79 was used for the [H+] or [OH-] calculations
according to Martell and Smith.20
General Description of the Stopped-Flow Studies. Kinetic
measurements for the reaction of HOSCN with SCN- and the
hydrolysis of (SCN)2 were made with a HI-TECH SF-61 DX2
stopped-flow spectrophotometer using a Xe arc lamp and a PMT
detector. A double-mixing mode was used for all of the experiments.
All monochromatic traces were collected at λ ) 300 nm using a 1
cm optical path length. During all of the stopped-flow measure-
ments, a temperature of 291 K was maintained in the observation
cell with a Lauda RC-20 circulator. The adiabatic temperature
increases that were associated with the pH-jump experiments were
determined experimentally to be less than 1 °C.
Calibration of the HI-TECH SF-61 DX2 Stopped-Flow
Instrument. The instrument was calibrated in both single- and
double-mixing modes using the reaction of 2,6-dichlorophenol-
indophenol (DCIP) and ascorbic acid (AA) reaction at pH 1.80
under pseudo-first-order conditions (excess [AA]).21 The plot of
[AA] vs kobs started to deviate from linearity when kobs was larger
than 400 s-1, indicating that the highest limit of kobs that could be
measured for a (pseudo) first-order reaction on our HI-TECH SF-
61 DX2 stopped-flow instrument is 400 s-1. The dead time and
the pretriggering time of the instrument were determined on the
basis of the following equations:
of HOSCN (k3):
(SCN)2 + H2O h HOSCN + SCN- + H+
SCN
k2
k
hyd
K2 )
)
) 5.7 × 10-4 M2 (2)
SCN
comp
k-2
k
2HOSCN f O2SCN- + SCN- + 2H+
k3 ) 6.9 × 104 M-1 s-1 (3)
We have recently investigated the kinetics of the reaction
of hypochlorite (OCl-)10 and hypobromite (OBr-)11 with
SCN- under alkaline conditions and learned that reactions
between the corresponding hypohalous acids and SCN- occur
very fast and with near diffusion-controlled kinetics in the
case of HOBr:
HOX h OX- + H+ pKa(HOCl) ) 7.4, pKa(HOBr) ) 8.6
(4)
HOX + SCN- f OSCN- + X- + H+
k(X ) Cl) ) 2.3 × 107 M-1 s-1, k(X ) Br) ) 2.3 ×
109 M-1 s-1 (5)
In contrast to OCl-, OBr- also reacts with SCN- with a
measurable rate, albeit with a rate constant (3.8 × 104 M-1
s-1) that is 5 orders of magnitude smaller than that of its
conjugate acid.11 Through the comparison of time-resolved
electronic spectra, time-resolved 13C NMR spectra, and time-
resolved 15N NMR spectra, we have concluded that the
enzyme system (lactoperoxidase/H2O2/SCN-), hydrolysis of
(SCN)2, and oxidation of SCN- by HOX (X ) Cl, Br) all
yield the same initial species.12 The experimental12-17 and
computed18 NMR shielding constants are consistent with a
formulation of OSCN-,12 and the NMR spectra appear to
rule out alternative formulations of hypothiocyanite (e.g.,
OCNS-, SOCN-, ONCS-, and OSNC-).12 The development
of a facile in situ synthesis of OSCN- by oxidation of SCN-
with OX- (X ) Cl, Br)10,11 affords the opportunity to
investigate some of its chemical reactions, including the
comproportionation of HOSCN with SCN- (k-2 ) k
the reverse of eq 2), which we report here together with a
reinvestigation of the hydrolysis of (SCN)2 that yields a direct
measurement of the rate constant for eq 2 (k2 ) kShyCdN).
SCN
comp
,
Experimental Section
Reagents. All chemicals were ACS certified grade or better.
Water was doubly distilled in glass. Solutions of NaOH, mostly
free of CO2 contamination, were quantified by titration with a
obstd
Dead time: Aobs ) Atot e-k
(10) Ashby, M. T.; Carlson, A. C.; Scott, M. J. J. Am. Chem. Soc. 2004,
126, 15976-15977.
(11) Nagy, P.; Beal, J. L.; Ashby, M. T. Chem. Res. Toxicol. 2006, 19,
587-593.
obsts
Pretriggering time: Atot ) Afit e-k
(12) Nagy, P.; Alguindigue, S. S.; Ashby, M. T. Biochem. 2006, 45, 12610-
12616.
where Aobs is the observed absorbance maximum; Atot is the effective
absorbance change associated with the reaction; Afit is the fitted
absorbance change (extrapolated to t ) 0); td is the dead time; and
ts is the pretriggering time of the instrument. It has been previously
(13) Arlandson, M.; Decker, T.; Roongta, V. A.; Bonilla, L.; Mayo, K.
H.; MacPherson, J. C.; Hazen, S. L.; Slungaard, A. J. Biol. Chem.
2001, 276, 215-24.
(14) Modi, S.; Behere, D. V.; Mitra, S. Biochim. Biophys. Acta 1991, 1080,
45-50.
(15) Modi, S.; Deodhar, S. S.; Behere, D. V.; Mitra, S. Biochem. 1991,
30, 118-124.
(16) Pollock, J. R.; Goff, H. M. Biochim. Biophys. Acta 1992, 1159, 279-
85.
(19) Troy, R. C.; Margerum, D. W. Inorg. Chem. 1991, 30, 3538-43.
(20) Martell, A. E.; Smith, R. M. Critical Stability Constants, Vol. 4:
Inorganic Complexes; Plenum Press: New York, 1976.
(17) Walker, J. V.; Butler, A. Inorg. Chim. Acta 1996, 243, 201-206.
(18) Sundholm, D. J. Am. Chem. Soc. 1995, 117, 11523-11528.
(21) Tonomura, B.; Nakatani, H.; Ohnishi, M.; Yamaguchi-Ito, J.; Hiromi,
K. Anal. Biochem. 1978, 84, 370-83.
286 Inorganic Chemistry, Vol. 46, No. 1, 2007

Comproportionation of HOSCN and SCN-
Figure 1. Time-absorbance trace at λ ) 300 nm for the hydrolysis of
thiocyanogen under relatively high [H+] and high [SCN-] (30% of the data
points are illustrated). This measurement was achieved using a double-
mixing stopped-flow sequence, where Cl2 was generated in situ by mixing
a solution of 20 mM HOCl (pH 7.91 in 1 M NaClO4) and 0.93 M H+ in
the first mixing cycle, which results in a reaction mixture of 10 mM Cl2 in
0.47 M H+ (I ) 1 M). This reaction mixture was then mixed with 100 mM
SCN- in 0.47 M H+ (I ) 1 M). The concentrations at t ) 0 are [(SCN)2]o
) 5 mM, [SCN-] ) 40 mM, [H+] ) 0.47 M, I ) 1 M (NaClO4 + HClO4
+ NaSCN). A homogeneous second-order fit is illustrated together with
its residual.
shown mathematically that it is not necessary to correct for a
concentration gradient within the stopped-flow observation cell for
a (pseudo) first-order reaction (as opposed to fast second-order
reactions, where the aging of the reaction mixture has to be taken
into account and the observed rate constants has to be corrected
using the geometrical parameters of the observation cell).22 We note
that the effect of the mixing rate constant (kmix) on kobs was
investigated, and we have found that no correction was necessary
for the data collected with our instrument.
Reaction of HOSCN with SCN-. The H+ and SCN- concentra-
tion dependencies of the reaction rates were investigated under
pseudo-first-order conditions with at least a 20-fold excess of SCN-
over HOSCN. The species OSCN- was generated in situ by the
reaction of SCN- and OBr- at pH 13 in the first mixing cycle11
followed by a pH jump to the desired pH in the second mixing
cycle. The following concentrations were achieved after the second
mixing cycle:
Figure 2. Time-absorbance traces at λ ) 300 nm for the hydrolysis of
thiocyanogen as a function of the [(SCN)2]o (0.5, 1.0, 2.0, and 3.0 mM) at
pH 1.88 and [SCN-] ) 0.1 M (20% of the data points are illustrated).
Equation 13 would be required to fit the entire absorption-time traces, but
there are too many parameters in eq 13 to resolve them with the available
data. However, the data can be modeled with single-exponential functions
after the induction periods. First-order fits of the indicated portions of the
data (solid lines) are illustrated together with their residuals. The rate
constants that are computed for [(SCN)2]o ) 0.5-3.0 mM are 18.4 ((0.2),
20.9 ((0.1), 20.42 ((0.06), and 22.27 ((0.04) s-1, respectively.
Hydrolysis of (SCN)2. The species (SCN)2 was generated in
situ by the reaction of Cl2/HOCl with excess SCN- at [H+] ) 0.2
M in the first mixing cycle. The effects of [H+], [SCN-], and
[(SCN)2]o on the reaction rates were investigated under pseudo-
first-order conditions with at least a 25-fold excess of SCN- over
(SCN)2. A pH jump to the desired values in the second mixing
cycle was achieved using a 0.2 M phosphate buffer of ca. pH 6.
The following concentrations were achieved after the second mixing
cycle:
Effect of [SCN-]: [H+] ) 0.46 M; [OSCN-]o ) 1.25 mM;
[SCN-] ) 23.8-123.8 mM
Effect of [H+]: [SCN-] ) 0.10 M; [(SCN)2]o ) 2.4 mM;
Effect of [H+]: [SCN-] ) 73.8 mM; [OSCN-]o ) 1.25 mM;
[H+] ) 0.20-0.46 M
[H+] ) 5.0-37.0 mM.
Effect of [SCN-]: [H+] ) 12.9 mM; [(SCN)2]o ) 2.0 mM;
[SCN-] ) 50-130 mM.
Effect of [(SCN)2]o: [H+] ) 13.2 mM; [SCN-] ) 0.10 M;
[(SCN)2]o ) 0.5-3.0 mM.
We have previously reported that a 400-fold excess of SCN- was
required to produce stoichiometric amounts of OSCN- with OCl-
at pH 13 when employing a Bio-Logic SFM-400/Q mixer.12 We
have determined that a 50-fold excess over OCl- is required to
produce >95% yield of HOSCN using the HI-TECH SF-61 DX2
stopped-flow instrument. Under similar conditions, OBr- yields
∼80% yield of HOSCN with the HI-TECH SF-61 DX2 stopped-
flow instrument. For the pseudo-first-order conditions (excess SCN-
and constant [H+]) that were employed in the present study, we
have determined that the over-oxidation products do not influence
the kinetics of the reaction of HOSCN with SCN-.
Kinetic Data Analysis. The monochromatic kinetic traces were
fit with HI-TECH KinetAsyst 3.14 software (Hi-Tech, UK). The
concentration dependencies of the pseudo-first-order rate constants
were obtained by linear least-squares fits of the data with Kaleida-
Graph 3.6 (Synergy Software).
Results
Kinetics of Hydrolysis of Thiocyanogen. Our efforts to
measure the kinetics of the hydrolysis of (SCN)2 by the initial
(22) Dunn, B. C.; Meagher, N. E.; Rorabacher, D. B. J. Phys. Chem. 1996,
100, 16925-16933.
Inorganic Chemistry, Vol. 46, No. 1, 2007 287

Nagy et al.
Figure 5. Observed pseudo-first-order rate constants for the compropor-
tionation of HOSCN as a function of [H+]. This measurement was achieved
using a double-mixing stopped-flow sequence, where OSCN- was generated
in the first mixing cycle by the reaction of SCN- and OBr- at pH ) 13
followed by a pH jump. Concentrations are for after the second mixing
cycle: [SCN-] ) 73.8 mM; [OSCN-]o ) 1.25 mM; I ) 1.0 M (NaClO4
+ HClO4 + NaSCN); [H+] ) 0.20-0.46 M. The rate constant that is
Figure 3. Change in absorption at λ ) 300 nm (20% of the data illustrated)
during the hydrolysis of (SCN)2 (2.0 mM) at pH 1.89 in the presence of 50
(O), 80 (0), 110 (]), and 130 (×) mM SCN-. Linear fits are illustrated
with slopes of 20.2, 20.1, 19.3, and 18.8 s-1 for [SCN-] ) 50, 80, 110,
and 130 mM.
computed from these data is kcomp ) 5.14 ((0.07) × 103 M-2 s-1
.
Table 1. Observed First-Order Rate Constants of the Hydrolysis of
Thiocyanogen as a Function of pH
a
a
pH
k
obs (s-1
)
pH
k )
obs (s-1
1.43
1.62
20.31 ( 0.06
20.70 ( 0.06
1.84
2.30
23.80 ( 0.05
27.37 ( 0.04
a The estimated error is for a least-squares fit of an average of seven
mixing cycles.
Figure 6. Observed pseudo-first-order rate constants for the compropor-
tionation of HOSCN as a function of [SCN-]. This measurement was
achieved using the procedure that is described in the caption of Figure 4.
Concentrations are for after the second mixing cycle: [H+] ) 0.46 M;
[OSCN-]o ) 1.25 mM; I ) 1.0 M (NaClO4 + HClO4 + NaSCN); [SCN-]
) 23.8-123.8 mM. The rate constant that is computed from these data is
kcomp ) 5.15 ((0.08) × 103 M-2 s-1
.
dependency on [H+] was observed for a given [SCN-], but
the first-order rate constant appears to converge on a discrete
value as the pH was raised (Table 1). There is an induction
period at lower pH (e.g., Figure 2) which disappears at higher
pH (e.g., Figure 4).
Kinetics of Comproportionation of Hypothiocyanous
Acid and Thiocyanate. This reaction was investigated using
a double-mixing stopped-flow sequence in which OSCN-
was generated in the first mixing cycle by oxidation of SCN-
with OBr- and in the second mixing cycle a pH jump was
induced.11 In the presence of excess SCN- and at constant
pH, pseudo-first-order kinetics were observed. The observed
pseudo-first-order rate constants exhibit first-order depen-
dency on [H+] (Figure 5) and [SCN-] (Figure 6). Thus, the
reaction exhibits overall third-order kinetics.
Figure 4. Time-absorbance traces at λ ) 300 nm for the hydrolysis of
thiocyanogen at pH 2.30 (30% of the data points are illustrated). Condi-
tions: [(SCN)2]o ) 2.4 mM, [SCN-] ) 0.1 M, I ) 1.0 M. A first-order fit
is illustrated together with its residual.
rate method for [H+] ) 0.5 M and [SCN-] ) 0.1 M were
unsuccessful. The pseudo-zero-order rate constants exhibited
a second-order dependency on [(SCN)2]o (unpublished
results), which suggested that disproportionation of HOSCN
is the rate-limiting step under such conditions. Second-order
kinetics, as previously observed by Stanbury et al.,9 were
confirmed for more acidic conditions by monitoring the entire
reaction and fitting the absorption-time trace to a homo-
geneous second-order function (Figure 1). However, under
less acidic conditions when the [SCN-] was relatively low,
first-order kinetics were observed. The first-order dependency
on [(SCN)2] under such conditions was confirmed by varying
the [(SCN)2]o (Figure 2). The rate law does not depend upon
the [SCN-] (Figure 3). Within the first-order regime, a slight
Discussion
Rate Law for Hydrolysis of Thiocyanogen. In deriving
a rate law for the decomposition of (SCN)2, Stanbury et al.
-
assumed HOSCN, (SCN)2, and (SCN)3 are in rapid equi-
librium.9 Through [SCN-]-jump stopped-flow experiments,
we have shown that the equilibrium between (SCN)2 and
288 Inorganic Chemistry, Vol. 46, No. 1, 2007

Comproportionation of HOSCN and SCN-
Scheme 1. Overall Stoichiometries of the Hydrolysis and Subsequent
Disproportionation of (SCN)2 for the Two Proposed Mechanisms
For Mechanism II, assuming once more that the dispro-
portionation of HOSCN is rate-limiting and that the subse-
quent fast reaction of HO2SCN consumes an equivalent of
HOSCN, we can define the rate as:
3 × {(SCN)2 + H2O h HOSCN + SCN- + H+}
2HOSCN f HO2SCN + SCN- + H+
HO2SCN + HOSCN f O3SCN- + SCN- + 2H+
3(SCN)2 + 3H2O f O3SCN- + 5SCN- + 6H+
d[(SCN)2]
-1
3
Rate )
×
) k3[HOSCN]2
(7)
dt
The two mechanisms differ in the definition of k3 (the
relationship between the definition of the rate in terms of
the observable (SCN)2 and the unobservable intermediate
HOSCN). Since the next step of the derivation involves
setting d[HOSCN]/dt to zero, it does not matter whether its
proportionality constant is 1/3 or 2/3. However, the propor-
tionality constant with respect to (SCN)2 does matter, in that
it defines k3 in terms of the observable. Since the present
study does not differentiate between the two aforementioned
mechanisms, we will define k3′ ) Pk3, where P ) 3/2 and
3 for Mechanisms I and II, respectively. From this point
forward, the derivation of the rate law is the same for the
two mechanisms, beginning with the steady-state approxima-
tion:
(SCN)3- occurs within the mixing time (unpublished results).
However, the first-order kinetics we observed at low [H+]
and [SCN-] suggest that hydrolysis can become rate-limiting
under certain circumstances. Thus, in deriving our rate law
we will assume that the equilibrium between (SCN)2 and
-
(SCN)3 is rapid, that the [HOSCN] achieves the steady-
state, and that the disproportionation of HOSCN is rate-
limiting. The mechanism that was previously proposed9 is
illustrated as Mechanism I of Scheme 1. The rapid equilib-
-
rium between (SCN)2 and (SCN)3 will be taken into
consideration later. In addition to Mechanism I, Mechanism
II is also conceivable (Scheme 1). The two mechanisms differ
in the third step, namely disproportionation of HO2SCN for
the former mechanism and oxidation of HO2SCN by HOSCN
in the latter mechanism. We will derive a general rate law
next, so that the reader may compare our results to those
that have been previously reported by Stanbury et al.9 The
derivation will also prove useful when we discuss the
induction period that occurs at lower pH (e.g., Figure 2).
Since neither our data nor that of Stanbury et al.9 are able to
differentiate between these alternative reaction pathways, we
consider both here, beginning with Mechanism I.
-d[HOSCN]
) k3′[HOSCN]2 +
dt
k-2[H+][SCN-][HOSCN] - k2[(SCN)2] ) 0 (8)
[HOSCN]SS
)
-k [H+][SCN-] (
k
x
2[H+]2[SCN-]2 + 4k2k3′[(SCN)2]
-2
-2
2k3′
For Mechanism I, assuming that the disproportionation of
HOSCN is rate-limiting and that the subsequent fast dispro-
portionation of HO2SCN yields an equivalent of HOSCN,
we can define the rate in terms of the observable (SCN)2 as:
(9)
It is useful at this point to define a constant:
3
2
× {(SCN)2 + H2O h HOSCN + SCN- + H+}
4k2k3′
Let R )
(10)
2HOSCN f HO2SCN + SCN- + H+
k-22[H+]2[SCN-]2
1
2
× {2HO2SCN f O3SCN- + HOSCN + H+}
which when incorporated into eq 9 yields:24
3
2
3
1
5
(SCN)2 + H2O f O3SCN- + SCN- + 3H+
2
2
2
2k2[(SCN)2]
d[(SCN)2]
[HOSCN]SS
)
(11)
-2
3
Rate )
×
) k3[HOSCN]2
k [H+][SCN-]( 1 + R[(SCN) ] + 1)
(6)
x
-2
2
dt
where the rates in eq 6 are defined in terms of the elementary
reactions above.23
Substituting [HOSCN]SS into the rate law (eq 12) gives eq
13:
(23) The proportionality constants are defined for the elementary reaction
(in some cases in curly brackets) using the convention for: aA f
bB; -d[A]/(a‚dt) ) d[B]/(b‚dt) ) k[A].
(24) Espenson, J. H. Chemical Kinetics and Reaction Mechanisms;
McGraw-Hill: New York, 1981; pp 79-80.
Inorganic Chemistry, Vol. 46, No. 1, 2007 289

Nagy et al.
increasing pH that is evident in Table 1 is due to an
increasing contribution of OH-, even under the very acidic
conditions of our measurements. There is a discernible
induction period in the kinetic traces that are collected below
ca. pH 2. We attribute this induction period to the formation
of a steady-state concentration of HOSCN (eq 2). Thus, the
kinetic traces illustrated in Figure 2, which were collected
at pH 1.88, are best described by eq 13, which is somewhere
between the limiting extremes of eqs 14 and 15. However,
it is clear that the rate law that describes the kinetic traces
in Figure 2 is much closer to eq 14 than eq 15. Thus, the
traces after the induction periods are described well by a
single exponential (Figure 2), the rate constants that are
obtained when [(SCN)2]o is varied are essentially the same
(Figure 2, kavg ) 20.5 ((1.6) s-1), and the (pseudo)-first-
order rate constants that are obtained when [SCN-] is varied
are essentially the same (Figure 3, kavg ) 19.7 ((0.7) s-1).
The induction period disappears when the pH is greater than
2 and the resulting kinetics exhibit clean first-order behavior
(Figure 4). When the [H+] is varied, the rate constant appears
to converge on a value of ca. 20 s-1 (Table 1). Thus, the
observed kinetics of hydrolysis is completely consistent with
the model that has been employed to derive eq 13. Accord-
ingly, these results suggest a fairly narrow window of [H+]
and [SCN-] in which overall first-order behavior is observed.
It is noteworthy that an intermediate forms when I2 reacts
with H2O that is best described as H-O-I-H+(vide infra).
A corresponding intermediate forms when I2 reacts with OH-
-d[(SCN)2]
dt
Rate )
-d[(SCN)2]
) k3′[HOSCN]2
(12)
) k3′ ×
dt
2
2k2[(SCN)2]
4k2k3′
k-22[H+]2[SCN-]2
k-2[H+][SCN-]
1 +
[(SCN)2] + 1
[
]
(
)
x
(13)
Equation 13 is useful for describing limiting extremes. If 1
, R[(SCN)2] (which is expected to occur for high pH and
low [SCN-]), the k2 step becomes rate-limiting, and first-
order kinetics are expected:
2
-d[(SCN)2]
dt
2k2[(SCN)2]
) k3′
)
+
-
[
]
k [H ][SCN ] R[(SCN) ]
x
-2
2
2
4K2 k3′[(SCN)2]
) k2[(SCN)2] (14)
R[H+]2[SCN-]2
In contrast, if 1 . R[(SCN)2] (which can be expected for
low pH and high [SCN-]), the k3 step becomes rate-limiting
and second-order kinetics can be expected:
K2 k3′[(SCN)2]2
2
2
-d[(SCN)2]
dt
k2[(SCN)2]
) k3′
)
)
2
+
-
2
k-2[H ][SCN ]
[H+]o [SCN-]o
[
]
-
that is best described as I-I-OH (vide infra). While we
keff[(SCN)2]2 (15)
find it difficult to conceive of analogous intermediates for
the reaction of H2O and OH- with (SCN)2, our measurements
do not preclude the involvement of such species. In contrast
to I2, no such adduct between H2O or OH- and X2 (X ) Cl,
Br) has been detected in previous studies.26,27
Rate Law for Comproportionation of Hypothiocyanous
Acid and Thiocyanate. The rate law we employed for the
determination of k-2 is more readily derived. Since under
very acidic conditions the equilibrium of eq 2 lies far to the
side of (SCN)2, the rate of approach to equilibrium beginning
with HOSCN is effectively:
Much of the data that were analyzed by Stanbury et al.
exhibit clean second-order kinetics,9 although we conclude
that some of the data at lower [H+] should be essentially
first-order. In contrast, we have focused on the regime of
[H+] and [SCN-] where the observed kinetics are essentially
first-order. Accordingly, while the previously reported
second-order kinetic data that were based upon changes in
absorbance had to be converted to molar pseudo-second-
order rate constants by applying a conversion factor,9 k2 (the
first-order rate constant for hydrolysis of (SCN)2) can be
obtained directly from the spectral kinetic traces we have
measured. Since the equilibrium between (SCN)2 and
+d[(SCN)2]
Rate )
) k-2[HOSCN][SCN-][H+] (16)
dt
-
(SCN)3 is fast with respect to the rate of hydrolysis of
(SCN)2 under the conditions that were employed during our
measurements, the first-order decay of the absorbance (which
is actually a composite of the absorption of (SCN)2 and
As expected, the rate of comproportionation of HOSCN and
SCN- exhibits overall third-order kinetics with first-order
dependencies on [HOSCN] (i.e., first-order kinetic traces in
the presence of excess SCN- and H+), [H+] (Figure 5),and
[SCN-] (Figure 6).
Equilibrium Constant for the Reversible Hydrolysis of
Thiocyanogen. Stanbury et al. have taken two approaches
to evaluating K2. In the first approach, the initial absorption
at λ ) 300 nm was fit to an equation that was derived using
the assumption that the resulting absorption is the composite
of the equilibria of eqs 1 and 2:6
-
(SCN)3 ) is proportional to -d[(SCN)2]/dt.
The data in Table 1 suggest that the observed first-order
rate constants (kobs) approach a pH-independent value of ca.
20 s-1, which is consistent with the hydrolysis of (SCN)2
by H2O (eq 2). This reaction can be considered to be a
nucleophilic attack of H2O on (SCN)2. A 7-orders-of-
magnitude difference exists between the second-order rate
constants for nucleophilic attack of H2O versus OH- on I2,25
so it is conceivable that the gradual increase of kobs with
(26) Beckwith, R. C.; Wang, T. X.; Margerum, D. W. Inorg. Chem. 1996,
35, 995-1000.
(27) Wang, T. X.; Margerum, D. W. Inorg. Chem. 1994, 33, 1050-5.
(25) Lengyel, I.; Epstein, I. R.; Kustin, K. Inorg. Chem. 1993, 32, 5880-
2.
290 Inorganic Chemistry, Vol. 46, No. 1, 2007

Comproportionation of HOSCN and SCN-
ꢀ(SCN) + ꢀ(SCN) - - K [SCN-]
Table 2. Rate Constants for the Hydrolysis of X2 (X2 + H2O), Rate
Constants for Comproportionation of HOX and X- (HOX + X- + H+),
and the Corresponding Equilibrium Constants
1
2
3
ꢀeff
)
(17)
K2
1 + K1[SCN-] +
kXhyd
kXcomp
KXhyd ) kXhyd kX
/
comp
[SCN-][H+]
X
(s-1
)
(M-2 s-1
)
(M2)
ref
Cl
Br
I
15
98
1.8 × 104
1.6 × 1010
2.0 × 1020
5.1 × 103
8.3 × 10-4
6.1 × 10-9
5.4 × 10-13
3.9 × 10-3
25a
24b
23c
More recently, beginning with an equilibrium mixture of
-
(SCN)2 and (SCN)3 , Stanbury et al. have investigated the
1.1 × 108
reaction sequence of eqs 1-3.9 Their pseudo-second-order
kinetic traces were fit to the rate law (where P was assigned
a value of 3/2 by Stanbury et al.):9
SCN
20
this studyd
a These data were determined at 20 °C and µ ) 0.5 M. For comparison,
Cl
hyd
K
) 8.3 × 10-4 M2 at 25 °C and µ ) 0.1 M. b These data were
Br
hyd
determined at 25 °C and µ ) 0.5 M. For comparison, K ) 6.2 × 10-9
-d[(SCN)2]
M2 at 25 °C and µ ) 1.0 M. c Consult the discussion of eqs 20 and 21
herein and ref 23 for an explanation of the origin of these rate constants.
d These data were determined at 18 °C, µ ) 1.0 M.
) P ×
dt
2
K2 k3[(SCN)2]2
+ K1[(SCN)2]3[H+]2 + K2[SCN-][H+]
- 2
constants, KXhyd ) khXyd/kcXomp. Despite the fact that the redox
potential of SCN- lies somewhere between that of Br- and
I-,28 the equilibrium constant for (SCN)2 is more similar to
Cl2 than either Br2 or I2. Given that khyd is very similar for
X ) Cl, Br, and SCN, the similarity of K for X ) Cl and
SCN is due to their closely matched values of kXcomp, which
at ca. 104 M-2 s-1 are substantially smaller than the
corresponding values that have been determined for Br and
I. Before continuing this discussion, it is important to note
that the aqueous chemistry of I2 and HOI is substantially
more complicated than the corresponding chemistry for X
) Cl, Br, and apparently for SCN. At high pH (above ca.
pH 7), the major pathway of I2 hydrolysis proceeds through
nucleophilic attack that produces the intermediate I2OH-.
However, in the pH regime of this study (below pH 2), the
major hydrolysis pathway for I2 is via addition of H2O and
the intermediate H2OI+. Thus, the equilibrium and rate
+ 2
{
}
[H ] [SCN ]
(18)
The complexity of eq 18 (with respect to eq 15) stems from
the fact that for constant [SCN-] and [H+] the observed rate
data are pseudo-second-order (kobs′). Since homogeneous
second-order kinetics depend upon the initial concentration
of the reactant, kobs′ must be written in terms of the
-
partitioning between (SCN)2 and (SCN)3 (i.e., by taking
into account K1):
2
K2 k3
[H+]2[SCN-]2
kobs′ ) P ×
×
-2
K2
1 + K1[SCN-] +
(19)
-
+
(
)
[SCN ][H ]
In fitting their data to eq 19, K1 was held constant (0.43
constants for X ) I in Table 2 are given by (KIhyd) K20K21):
2
M-1), and thus two parameters were obtained, K2 k3 and K2.
25,29,30
The value for K2 that was obtained using eq 17 (K2 ) 5.4
((1.9) × 10-4 M2) and the value that was obtained using
eq 19 (K2 ) 5.66 ((0.77) × 10-4 M2) were self-consistent.
By comparison, we obtained a value of K2 ) [HOSCN]-
[SCN-][H+]/[(SCN)2] ) k2/k-2 ) khSyCdN/kcSoCmNp ) 19.8 ((0.7)
s-1/ 5.14 ((0.07) × 103 M-2 s-1 ) 3.8 × 10-3 M2, which
differs from the values that were determined by Stanbury et
al. by roughly an order of magnitude. However, given the
potential sources of error (including a large estimated error
for K1 ) 0.43 ((0.29) M-1 and the likelihood that at least
part of the data that were analyzed by Stanbury et al. were
probably in the regime where eq 14 was more appropriate),
we find the value of K2 to be quite comparable to the value
that was determined by largely independent means by
Stanbury et al. This latter conclusion regarding the equilib-
rium constant of eq 2 adds credence to our conclusion that
we have measured the rate constants of eq 2.
Comparisons with the Halogens. The rate constants for
hydrolysis of the halogens (kXhyd) are compared with the rate
constant for hydrolysis of (SCN)2 (i.e., k2) in Table 2. It is
remarkable how similar kXhyd is for X ) Cl, Br, and SCN
(i.e., ca. 10-100 s-1). In sharp contrast, kIhyd for I2 is 5
orders of magnitude larger than the other (pseudo)halogens.
Table 2 also summarizes the rate constants for compropor-
tionation of HOX and X- (kcXomp), as well as the equilibrium
constants that are derived from the ratio of the two rate
I2 + H2O h H2OI+ + I-
K20 ) k20/k-20 ) 0.12 s-1/1 × 1010 M-1 s-1
)
1.2 × 10-11 M (20)
H2OI+ h HOI + H+
K21 ) k21/k-21 ) 9 × 108 s-1/2 × 1010 M-1 s-1
)
4.5 × 10-2 M (21)
Accordingly, the hydrolysis of I2 is not an elementary
reaction.
The hydrolysis of I2 aside, it is apparent that the rates of
hydrolysis of X2 for X ) Cl, Br, and SCN are comparable
and that the differences between the positions of the equilibria
lie in the relative rates of comproportionation. From a
thermodynamic perspective, since HX are fully ionized, HOX
are completely un-ionized, and the values of kXhyd are
comparable, the trend of kScoCmNp< kcColmp , kcBormp must relate to
the relative stabilities of X2 versus HOX and X-. These
(28) Hughes, M. N. General chemistry [of thiocyanic acid and its
derivatives]. In Chemistry and Biochemistry of Thiocyanic Acid Its
DeriVatiVes; Academic Press: New York, 1975; pp 1-67.
(29) Sebok-Nagy, K.; Koertvelyesi, T. Int. J. Chem. Kin. 2004, 36, 596-
602.
(30) Schmitz, G. Int. J. Chem. Kin. 2004, 36, 480-493.
Inorganic Chemistry, Vol. 46, No. 1, 2007 291

Nagy et al.
relative stabilities are the result of a combination of bond
enthalpies (e.g., X-X vs O-X) and solvation effects.
Solvation energies are expected to be particularly important
for the anions and for the species that can form strong
hydrogen bonds. From a kinetic perspective, the trend that
is observed for kXhyd must reflect the relative nucleophilic-
ity31 of X- versus the electrophilicity of HOδ-X δ+. Regard-
ing the nucleophilicity, we note that SCN- is an ambidentate
ligand and that it is known to produce products that are the
result of both S and N nucleophilic attack.32 This naturally
raises the question of whether the reactions we have studied
involve HOSCN and/or HONCS. As noted earlier, we have
previously investigated the structure of the initial product
that is produced by the reaction of HOBr and SCN-, the
reaction we have employed herein to produce “HOSCN”,
and we are convinced that the product is indeed HOSCN
and not HONCS.12 Thus, we are confident that we are
measuring the rate of nucleophilic attack on HOSCN.
Regarding the attacking nucleophile, the simplest mecha-
nisms would be nucleophilic attack by the S of SCN-. Since
the Swain-Scott parameters suggest that Br- and SCN- are
comparable nucleophiles and that both are better nucleophiles
than Cl- (n ) 2.99, 4.02, and 4.93 for Cl-, Br-, and SCN-,
respectively),33 the origin of the trend in kXhyd may be due to
the relative electrophilicities of HOX. It is noteworthy that
HOBr generally reacts several orders of magnitude faster
with nucleophiles than HOCl does.34,35 Thus, HOBr appears
to be the better electrophile. Unfortunately, there is a dearth
of information regarding the reactivity of HOSCN, and in
fact the comproportionation of HOSCN and SCN- is the first
reaction we are aware for which the rate has been measured.
Fortunately, the corresponding rates of reaction of HOCl and
HOBr with SCN- are known. The second-order rate con-
stants for the reaction of HOX with SCN- are 2.3 × 107,
2.3 × 109, and 5.1 × 103 M-1 s-1 for X ) Br, Cl, and SCN,
respectively.10,11 Since these reactions involve a common
reactant, we conclude that the relative electrophilicities are
HOBr > HOCl . HOSCN, at least with respect to the
nucleophile SCN-. However, we note that the relative
reactivities of HOCl and HOBr toward nucleophiles is largely
independent of the nature of the nucleophile itself.34,35 It
would appear likely that HOSCN is generally a relatively
poor electrophile. Accordingly, it is not surprising that the
only nucleophiles (other that SCN-) that we have found that
it reacts with are thiolates.36 Since thiolates are very good
nucleophiles, it is not surprising that they are among the most
reactive species toward HOCl and HOBr as well.34,35
Conclusion
We have measured the first-order rate constant for the
hydrolysis of (SCN)2 by H2O (kShyCdN) and the corresponding
third-order rate constant for the comproportionation of
HOSCN + SCN- + H+ (kcSoCmNp). The values of khXyd for X2
are within an order of magnitude for X ) Cl, Br, and SCN.
In contrast, while the values of kXcomp are comparable for
HOX + X- + H+ for X ) Cl and SCN, kcBormp for X ) Br
is 6 orders of magnitude larger. We propose that the
mechanism of comproportionation involves nucleophilic
attack on HOX by X-. While the relative nucleophilicities
of X- are SCN- > Br- . Cl-, the electrophilicities of HOX
are HOBr > HOCl . HOSCN. In addition to explaining
the trend in kXcomp, the ordering of nucleophilicities and
electrophilicities rationalize the trend in the reactions of HOX
with SCN-: HOBr > HOCl . HOSCN.
Acknowledgment. We appreciate the National Science
Foundation (CHE-0503984), the American Heart Association
(0555677Z), and the Petroleum Research Fund (42850-AC4)
for their financial support. We are also grateful to reviewers
of this manuscript who provided insightful comments.
(31) Phan, T. B.; Breugst, M.; Mayr, H. Angew. Chem., Int. Ed. 2006, 45,
3869-3874.
(32) Loos, R.; Kobayashi, S.; Mayr, H. J. Am. Chem. Soc. 2003, 125,
14126-14132.
IC061470I
(33) Koskikallio, J. Acta Chem. Scand. 1969, 23, 1477-89.
(34) Pattison, D. I.; Davies, M. J. Chem. Res. Toxicol. 2001, 14, 1453-
64.
(36) Ashby, M. T.; Aneetha, H. J. Am. Chem. Soc. 2004, 126, 10216-
(35) Pattison, D. I.; Davies, M. J. Biochem. 2004, 43, 4799-4809.
10217.
292 Inorganic Chemistry, Vol. 46, No. 1, 2007