Oxidation of thiocyanate by iron(V) in alkaline medium

Article abstract of DOI:10.1016/j.ica.2004.07.001

The oxidation of thiocyanate (SCN^-) by iron(V) (Fe(V)) was studied as a function of pH in alkaline solutions by a premix pulse radiolysis technique. The reaction precedes via a two-electron oxidation, which converts Fe(V) to Fe(III). The rates decrease with an increase in pH. The reactivity of diprotonated Fe(V) species, H2FeO4- with SCN^- could account for the decease in rate with increasing pH. The oxidation of thiocyanate by iron(V) (Fe(V)) was studied as a function of pH in alkaline solutions by a premix pulse radiolysis technique. The rates decrease with an increase in pH. The rate law for the oxidation of SCN^- by Fe(V) was obtained as -d[Fe(V)]/dt = k₁₀{[H⁺]²/([H⁺]² + K ₂[H⁺] + K₂K₃)}[Fe(V)][SCN ^-], where k₁₀ = 5.72 ± 0.19 × 10⁶ M^-1 s^-1, pK₂ = 7.2, and pK₃ = 10.1. The reaction precedes via a two-electron oxidation, which converts Fe(V) to Fe(III). Thiocyanate reacts approximately 10³× faster with iron(V) than does with iron(VI).

Full text of DOI:10.1016/j.ica.2004.07.001

Inorganica Chimica Acta 357 (2004) 4587–4591
www.elsevier.com/locate/ica
Note
Oxidation of thiocyanate by iron(V) in alkaline medium
a,
b
c
Virender K. Sharma *, Donald B. OÕConnor , Diane Cabelli
a
Department of Chemistry, Florida Institute of Technology, 150 West University Boulevard, Melbourne, FL 32901, USA
b
Department of Chemistry, University of Texas at Austin, Austin, TX 78712, USA
Chemistry Department, Brookhaven National Laboratory, Upton, Long Island, NY 11973, USA
c
Received 16 December 2003; accepted 5 July 2004
Available online 11 August 2004
Abstract
The oxidation of thiocyanate by iron(V) (Fe(V)) was studied as a function of pH in alkalinꢀe solutions by a premix pulse radiolysis
technique. The rates decrease with an increase in pH. The rate law for the oxidation of SCN by Fe(V) was obtained as ꢀd[Fe(V)]/
dt = k10{[H+]2/([H+]2 + K2[H+] + K2K3)}[Fe(V)][SCNꢀ], where k10 = 5.72 0.19 · 106 Mꢀ1 sꢀ1, pK2 = 7.2, and pK3 = 10.1. The re-
action precedes via a two-electron oxidation, which converts Fe(V) to Fe(III). Thiocyanate reacts approximately 103· faster with
iron(V) than does with iron(VI).
ꢀ 2004 Elsevier B.V. All rights reserved.
Keywords: Iron(V); Oxidation; Thiocyanate; Pulse radiolysis
1. Introduction
ous reactions, sporadic information is available on the
fundamental properties of Fe(V) and Fe(IV) in the ab-
There has been increasing interest in higher oxidation
states of iron (Fe(IV), Fe(V), and Fe(VI)) because of
their unusual solid-state structures and electronic prop-
erties [1–4]. In aqueous solutions, the highly oxidized ir-
ons have shown a potential as an alternate for battery
cathodes, oxidant for organic synthesis, and an interme-
diate in Fenton reactions and biological transfer proc-
esses [5–8]. Ferryl (FeIV = O) and perferryl (FeV = O)
units have been postulated in several iron–dioxygen sys-
tems [9]. Recently, structural evidences for the reactive
intermediates having an FeIV = O unit in the activation
of dioxygen by nonheme iron enzymes were given [10].
Fe(V) and Fe(IV) are frequently postulated as interme-
diates in the oxidation of inorganic and organic sub-
strates by Fe(VI) [11–20]. Despite the proposed
involvement of these oxidation states of iron in numer-
sence
of
a
specialized
ligand
system
ðFeVO4 and FeIVO44ꢀÞ in aqueous solution [21,22].
A limited number of studies on the oxidations by
Fe(V) in its simple state as an oxyanion ðFeVO43ꢀÞ in
aqueous solution have been conducted [23–27]. It has
been suggested that Fe(V) oxidizes amino acids and
carboxylic acids by a two-electron step (Eq. (1)).
3ꢀ
FeðVÞþamino acids ! FeðIIIÞþNH3 þaꢀketo acids
ð1Þ
The rates for the oxidation of carboxylic acids by
Fe(V) depend on the nature of substituent group at
the a-carbon atom of the acids and decrease in the order
of a-C–NH2 > a-C–OH > a-C–H [23–25].
We have used a premix pulse radiolysis technique to
study the reaction of Fe(V) with thiourea and cyanide
[26,27]. Similar to previous studies on Fe(V), the reac-
tion of Fe(V) and thiourea appear to occur via a con-
certed two-electron oxidation. However, in studying a
reaction of cyanide with Fe(V) in alkaline solution, we
*
Corresponding author. Tel.: +1 321 674 7310; fax: +1 321 674
8951.
E-mail address: vsharma@fit.edu (V.K. Sharma).
0020-1693/$ - see front matter ꢀ 2004 Elsevier B.V. All rights reserved.
doi:10.1016/j.ica.2004.07.001

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V.K. Sharma et al. / Inorganica Chimica Acta 357 (2004) 4587–4591
found for the first time the spectroscopic evidence for
sequential one-electron reductions of Fe(V) to Fe(III)
by cyanide; Fe(V) is one electron reduced to Fe(IV),
which is subsequently one-electron reduced to Fe(III) by
cyanide. The results also showed that Fe(V) and Fe(IV)
are much more reactive than Fe(VI) (k(Fe(V) + cya-
nide)/k(Fe(VI) + cyanide) ꢁ 104; k(Fe(IV) + cyanide)/
k(Fe(VI) + cyanide) ꢁ 102).
The Fe(VI) solution (0.001 M borate/0.005 M phos-
phate; pH 9.0) in one syringe was mixed with the thiocy-
anate solution (0.10 M phosphate buffer; 0.7 M
methanol) of the other syringe. Both solutions were sat-
urated with nitrous oxide; [N2O] = 0.026 M. The mixed
solution was promptly injected into the optical cell (2.5
cm pathlength) and exposed to the ionizing pulse. Typ-
ically, the time between mixing and pulsing was <2 s,
sufficiently short for performing experiments. The se-
quence of reactions, following the electron pulse (Eq.
(I)) occurring within a fraction of a microsecond
[35,36] that lead to formation of Fe(V) are shown in
Eqs. (2)–(5). The number in parentheses in Eq. (I) are
G values, that is the number of radicals formed per
100 eV of energy dissipated in the aqueous solution.
The solutions were saturated with N2O and contained
methanol. The H atoms are converted to hydrated elec-
trons (Eq. (3)), which react with nitrous oxide to form
hydroxyl radicals (Eq. (2)). The OH radicals react with
methanol to form reducing radicals (Eq. (4)), which in
turn reduces Fe(VI) to Fe(V) (Eq. (5))
Recently, we published a very detailed study of the re-
action of Fe(VI) with thiocyanate [28], determining the
stoichiometry of the overall reduction of Fe(VI) to
Fe(III) reaction and showing the effect of pH on the
mechanism. In that study we measured the rate constant
for Fe(V) reaction with SCNꢀ at a single pH and
showed it was three orders of magnitude faster than that
of Fe(VI) with SCNꢀ. In the present paper, we have in-
vestigated the oxidation of thiocyanate (SCNꢀ) by
Fe(V) in alkaline medium as a function of pH using a
premix pulse radiolysis technique. The substrate SCNꢀ
was chosen because it is a simple sulfur compound
and is of interest in oscillating reactions, catalytic perox-
idation by lactoperoxidase and other peroxidases, phys-
iology, and industrial processes [29–32]. Additionally,
the results of the present study can be compared with
an earlier work on the oxidation of cyanide by Fe(V).
Å
H2O,Hð0:55Þ; eꢀaqð2:65Þ; OH ð2:75Þ; H3Oþð2:65Þ;
H2O2ð0:72Þ; H2ð0:45Þ
ðIÞ
ð2Þ
ð3Þ
ð4Þ
Å
N2O þ eꢀ þ H2O ! OH þ OHꢀ þ N2
aq
H þ OHꢀ () eaqꢀ þ H2O
2. Experimental
ꢀÅ
ꢀ
Å
Å
2.1. Materials
OH=O þ CH3OH ! H2O=OH þ CH2OH
Å
All chemicals used (Sigma, Aldrich) were of reagent
grade or better and were used without further purifica-
tion. Solutions were prepared with water that had been
distilled and then passed through an 18 MX Milli-Q wa-
ter purification system. Potassium ferrate (K2FeO4) of
high purity (98% plus) was prepared by the method of
Thompson et al. [33]. The Fe(VI) solutions were pre-
pared by addition of solid samples of K2FeO4 to 0.005
M Na2HPO4/0.001 M borate, pH 9.0, where solutions
are most stable [33]. Phosphate serves as a complexing
agent for Fe(III), which otherwise precipitates rapidly
as a hydroxide that instantly interferes with the optical
monitoring of the reaction and also accelerates the spon-
taneous decomposition of Fe(VI). A molar absorption
coefficient e510 nm = 1150 Mꢀ1 cmꢀ1 was used for the
calculation of ½FeO42ꢀꢂ at pH 9.0 [34].
FeðVIÞ þ CH2OH ! FeðVÞ þ Product
k ¼ 9 ꢃ 109Mꢀ1 sꢀ1
ð5Þ
2.3. Equipment
A 4 MeV Van de Graff accelerator was used to per-
form kinetic studies. This produced electron pulses
(<500 ns) resulting in irradiation doses of 2–15 gray in
the solutions. Details of the accelerator and absorption
spectrometer have been reported elsewhere [37]. The thi-
ocyanate dosimeter (0.01 M KSCN, 0.026 M N2O, pH
5.5) was used as a calibrant taking ðGðSCNÞ ꢀÞ ¼ 6:13
2
(radicals/100 eV) and e472 nm = (7950 3%).
3. Results and discussion
2.2. Fe(V) Kinetics
Initially, spectral studies of thiocyanate reaction with
The reaction of Fe(VI) with thiocyanate occurred so
rapidly under our experimental conditions [28] as to ne-
cessitate the use of a premix apparatus to study the reac-
tion of Fe(V) with thiocyanate using pulse radiolysis.
The premix pulse radiolysis apparatus consists of two
30 cm3 glass syringes mounted in a double syringe drive.
Fe(V) were performed at pH 10.67. 1.00 · 10ꢀ4
M
Fe(VI) and 1 · 10ꢀ3 M thiocyanate solutions were
mixed and pulse irradiated. Kinetic traces and spectra
obtained following the pulse are shown in Fig. 1. These
spectra manifest the decay of Fe(V) (kmax = 380 nm)
during the reaction. The characteristic spectrum of

V.K. Sharma et al. / Inorganica Chimica Acta 357 (2004) 4587–4591
4589
0.030
0.025
0.020
0.015
0.010
0.005
0.000
-0.005
decay rates of Fe(V) with methanol at different pH were
used to correct rates for reactions of Fe(V) with thiocy-
anate [26]. The experiments at different pH showed no
change in the observed rate law. The spontaneous decay
rate of Fe(V) with itself (2kobs ꢁ 107 Mꢀ1 sꢀ1) was so
fast below pH 10.1 that it contributes significantly to
the studied reaction, therefore, rates were limited to
pH P 10.1.
The rate of the reaction increases with a decrease in
pH (Table 1). The protonation of thiocyanate does not
occur in the pH range studied (pK = ꢀ1.28 0.03
[38]), therefore, the dependence is related to protona-
tions of Fe(V) (Eqs. (8) and (9))
0.010
0.005
0.000
-0.005
-0.010
-0.015
-0.020
1.18 ms
4.72
11.24
300 350 400 450 500 550 600
λ, nm
H2FeO4 () Hþ þ HFeO4
pK2 ¼ 7:2 [22]
pK3 ¼ 10:1 [22]
ð8Þ
ð9Þ
ꢀ
2ꢀ
þ
2ꢀ
3ꢀ
HFeO4 () H þ FeO4
0.000
0.005
0.010
0.015
0.020
time, sec
The three forms of Fe(V) react with thiocyanate (Eqs.
(10)–(12)).
H2FeO4ꢀ þ SCNꢀ ! Products
Fig. 1. Kinetic traces and spectra obtained upon pulse radiolysis of
solutions (pH 10.67) during Fe(V) oxidation of SCNꢀ.
ð10Þ
ð11Þ
ð12Þ
HFeO42ꢀ þ SCNꢀ ! Products
FeO43ꢀ þ SCNꢀ ! Products
Fe(IV) (kmax = 430 nm) [21] was not observed in the ex-
periments. As discussed in the earlier publication [34],
the OD of the Fe(V) reaction with thiocyanate was
The rate of disappearance of Fe(V) can be given by
ꢀd½FeðVÞꢂ=dt ¼ k10½H2FeO4ꢀꢂ½SCNꢀꢂ
þ k11½HFeO42ꢀꢂ½SCNꢀꢂ
1
the same as that of Fe(V) spontaneous decay, indicating
that reaction between Fe(V) and thiocyanate was stoi-
chiometric and there was no chain decomposition of
Fe(VI) arising from the possible free radical intermedi-
ates (Eq. (7)). The SCNÅ could have possibly reduced
Fe(VI) to Fe(V) at a faster rate, k ꢁ 109 Mꢀ1 sꢀ1
Å
þ k12½FeO43ꢀꢂ½SCNꢀꢂ
ð13Þ
The substitutions of equilibria of Fe(V) (Eqs. (8) and
(9)) gives the following equation:
FeðVÞ þ SCNꢀ ! FeðIVÞ þ SCN
ð6Þ
ð7Þ
2
ꢀd½FeðVÞꢂ=dt ¼ fðk10½Hþꢂ þ k11K2½Hþꢂ
Å
FeðVIÞ þ SCN ! FeðVÞ þ SCNðoxidizedÞ
þ k12K2K3Þ=Zg½FeðVÞꢂ½SCNꢀꢂ;
ð14Þ
The absence of reaction 6 suggests that the reaction
of Fe(V) with thiocyanate proceeds via a concerted
two-electron oxidation, which converts Fe(V) to Fe(III).
The kinetics of Fe(V) reaction with thiocyanate was
studied under pseudo-order conditions with thiocyanate
in excess. Reactions were monitored by measuring the
absorbance of Fe(V) at 380 nm wavelength as a function
of time. The disappearance of Fe(V) occurred by a first-
order decay that was also first order in thiocyanate
concentration. The second-order rate constants were de-
termined as a function of pH (Table 1). The first-order
where Z = [H+]2 + K2[H+] + K2K3
The empirical fit of the observed second-order rate
constants suggest that second and third terms of Eq.
(14) are not needed to fit the data. This is supported
by a plot of second-order rate constants with respect
2ꢀ
to HFeO4
i.e. {[H+]/([H+] + K3), which was found
non-linear (inset Fig. 2). The rate constants can be
ꢀ
explained by considering H2FeO4 is the only reactive
species of Fe(V). Moreover, a plot of the observed sec-
ond-order rate constants versus {[H+]2/([H+]2 + K2
[H+] + K2K3)} was also found to be linear (Fig. 2). This
further suggests rates are dependent on concentration of
ꢀ
Table 1
Oxidation of thiocyanate by Fe(V) at different pH values at 22 ꢁC
H2FeO4 in the reaction solution. The observed rate law
may be written as
pH
k (Mꢀ1 sꢀ1
)
2
2
ꢀd½FeðVÞꢂ=dt ¼ k10f½Hþꢂ =ð½Hþꢂ þ K2½Hþꢂ
10.10
10.46
10.67
10.95
11.20
3630 121
þ K2K3Þg½FeðVÞꢂ½SCNꢀꢂ
ð15Þ
value of k10
1162 43.3
513 10.8
104 2.14
63.7 1.29
The slope in Fig.
5.7 0.2 · 106 Mꢀ1 sꢀ1
2
gave
a
=
.

4590
V.K. Sharma et al. / Inorganica Chimica Acta 357 (2004) 4587–4591
imately three orders of magnitude more reactive than
Fe(VI). The bimolecular decomposition rates of Fe(V),
4000
3000
2000
1000
0
k ꢁ 107
M
ꢀ1 sꢀ1, suggest the radical nature of Fe(V)
r2 = 0.99
species [22]. This radical character of Fe(V)
(FeV = O M FeIV–OÅ) may be responsible for higher re-
activity of Fe(V).
The oxidations of thiocyanate and cyanide by higher
oxidation states of iron occur faster at a lower pH (Fig.
3). Recent spin density calculations suggested that the
protonated Fe(VI) species ðHFeO4ꢀÞ has a larger spin
density on the oxo ligands than the deprotonated
ðHFeO42ꢀÞ species [39]. This would result in higher oxi-
dation power of protonated Fe(VI) and thus higher re-
activity at a lower pH.
4000
3000
2000
1000
0
0.0
0.2
+
0.4
0.6
+
[H ]([H ]+K
)
3
The rates for reactivity of Fe(VI) and Fe(V) with thi-
ocyanate are lower than those of cyanide (Fig. 3). Cya-
nide has two species (HCN () H++CNꢀ; pK = 9.3
[40]) in the pH range of this study. The protonated cya-
nide (HCN) would decrease the electrorepulsion be-
tween negatively charged Fe(VI) and Fe(V) species
(e.g. Eq. (8)) and contributes to higher reactivity of cya-
nide. Moreover, activation energy for the oxidation of
cyanide by Fe(VI), DEà = 38.9 1.0 kJ molꢀ1 was found
to be higher than DEà = 32.7 2.1 kJ molꢀ1 for thiocy-
anate [34,41]. Less favorable activation energies of the
reactions may also contribute to the slower rates for
the oxidation of SCNꢀ.
0.0000 0.0002 0.0004 0.0006 0.0008
[H+]2/([H+]2 + K2[H+] + K2K3)
Fig. 2. Hydrogen ion dependence on the rate of Fe(V) oxidation of
thiocyanate in 0.1 M sodium phosphate at 22 ꢁC.
A comparison of the reactivity of Fe(V) and Fe(VI)
with thiocyanate is shown in Fig. 3. Fig. 3 also gives
comparison of rate constants for oxidations of thiocya-
nate and cyanide by Fe(V) and Fe(VI). Fe(V) is approx-
Fe(V)
107
106
Cyanide
4. Conclusions
105
104
103
The rate law for the reaction of Fe(V) with thiocya-
nate showed a first-order dependence on Fe(V) and
SCNꢀ. Reduction of Fe(V) to Fe(III) by SCNꢀ proceeds
by a two-electron process. The reactivity of diprotonated
102
ꢀ
Fe(V) species, H2FeO4 with thiocyanate could account
Thiocyanate
101
for the increase in rate with decreasing pH. Fe(V) reacts
approximately three orders of magnitude faster with thi-
ocyanate than does Fe(VI). The observed rates for oxi-
dation of thiocyanate by Fe(V) are lower than that of
cyanide reaction with Fe(V).
10.0
10.5
11.0
11.5
12.0
12.5
pH
Fe(VI)
103
102
101
100
10-1
Acknowledgement
A part of this research is supported by under contract
DE-AC02-98CH10886 with the USDOE and its Divi-
sion of Chemical Sciences, Office of Basic Energy
Sciences.
Cyanide
Thiocyanate
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10.0
11.0
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