Spectroelectrochemistry of potassium ethylxanthate, bis(ethylxanthato)nickel(II) and tetraethylammonium tris(ethylxanthato)-nickelate(II)

Article abstract of DOI:10.1039/b105683m

Electrochemical and chemical oxidation of S₂COEt^-, Ni(S₂COEt)₂, and [Ni(S₂COEt)₃]^- have been studied by CV and in situ UV-VIS spectroscopy in acetonitrile. Cyclic voltammograms of S₂COEt^- and Ni(S₂COEt)₂ display one (0.00 V) and two (0.35 and 0.80 V) irreversible oxidation peaks, respectively, referenced to an Ag/Ag⁺ (0.10 M) electrode. However, the cyclic voltammogram of [Ni(S₂COEt)₃]^- displays one reversible (-0.15 V) and two irreversible (0.35, 0.80 V) oxidation peaks, referenced to an Ag/Ag⁺ electrode. The low temperature EPR spectrum of the oxidatively electrolyzed solution of (NEt₄)[Ni(S₂COEt)₃] indicates the presence of [Ni^III(S₂COEt)₃], which disproportionates to Ni(S₂COEt)₂, and the dimer of the oxidized ethylxanthate ligand, (S₂COEt)₂ ((S₂COEt)₂ = C₂H₅OC(S)SS(S)COC₂H₅), with a second order rate law. The final products of constant potential electrolysis at the first oxidation peak potentials of S₂COEt^-, Ni(S₂COEt)₂, and [Ni(S₂COEt)₃]^- are (S₂COEt)₂; Ni²⁺(sol) and (S₂COEt)₂; and Ni(S₂COEt)₂ and (S₂COEt)₂, respectively. The chemical oxidation of S₂COEt^- to (S₂COEt)₂, and [Ni(S₂COEt)₃]^- to (S₂COEt)₂ and Ni(S₂COEt)₂ were also achieved with iodine. The oxidized ligand in the dimer form can be reduced to S₂COEt^- with CN^- in solution.

Full text of DOI:10.1039/b105683m

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Spectroelectrochemistry of potassium ethylxanthate, bis(ethyl-
xanthato)nickel(II) and tetraethylammonium tris(ethylxanthato)-
nickelate(II)†
Ömer Dag,‡ Seniz Özalp Yaman,§ Ahmet M. Önal* and Huseyin Isci*
Department of Chemistry, Middle East Technical University, 06531 Ankara, Turkey
Received 28th June 2001, Accepted 31st July 2001
First published as an Advance Article on the web 12th September 2001
Electrochemical and chemical oxidation of S2COEtϪ, Ni(S2COEt)2, and [Ni(S2COEt)3]Ϫ have been studied by CV
and in situ UV-VIS spectroscopy in acetonitrile. Cyclic voltammograms of S2COEtϪ and Ni(S2COEt)2 display one
(0.00 V) and two (0.35 and 0.80 V) irreversible oxidation peaks, respectively, referenced to an Ag/Agϩ (0.10 M)
electrode. However, the cyclic voltammogram of [Ni(S2COEt)3]Ϫ displays one reversible (Ϫ0.15 V) and two
irreversible (0.35, 0.80 V) oxidation peaks, referenced to an Ag/Agϩ electrode. The low temperature EPR spectrum
of the oxidatively electrolyzed solution of (NEt4)[Ni(S2COEt)3] indicates the presence of [NiIII(S2COEt)3], which
disproportionates to Ni(S2COEt)2, and the dimer of the oxidized ethylxanthate ligand, (S2COEt)2 ((S2COEt)2 =
C2H5OC(S)SS(S)COC2H5), with a second order rate law. The final products of constant potential electrolysis at
the first oxidation peak potentials of S2COEtϪ, Ni(S2COEt)2, and [Ni(S2COEt)3]Ϫ are (S2COEt)2; Ni2ϩ(sol) and
(S2COEt)2; and Ni(S2COEt)2 and (S2COEt)2, respectively. The chemical oxidation of S2COEtϪ to (S2COEt)2, and
[Ni(S2COEt)3]Ϫ to (S2COEt)2 and Ni(S2COEt)2 were also achieved with iodine. The oxidized ligand in the dimer form
can be reduced to S2COEtϪ with CNϪ in solution.
ligands are reported to be effective in stabilizing Ni(). Studies
of the nickel complexes with sulfur donor ligands are also
important in view of the fact that nickel is coordinated to
sulfur atoms in hydrogenase enzymes. Up to now the redox
Introduction
The chemistry of nickel is mostly centered on its “ϩ2” oxid-
ation state with d8 electron configuration. Octahedral com-
plexes are paramagnetic, S = 1, and square-planar complexes
are diamagnetic, S = 0. In spite of the fact that Ni(), d9, and
Ni(), d7, were reported as early as 1913 (K4[Ni2(CN)6])
and 1936 ([Ni(PEt3)2Br3]),1 the interest in the chemistry of Ni()
and Ni() started after the discovery of nickel as a bioessential
trace element in several important enzymes in the 1970s.2 Today
the presence of nickel has been demonstrated in urease, hydro-
genase, CO oxidoreductase, and methyl-coenzyme M reductase
enzymes.3–5 Detailed structural information on the geometry
and coordination of the nickel in these enzymes is very import-
ant for understanding the mechanism of catalysis. A study on
the crystal structure of nickel–iron hydrogenase from Desulfo-
vibrio gigas indicated that the active site contains nickel
coordinated to four sulfur donor ligands with a geometry
between square-planar and tetrahedral.6 EPR studies on several
nickel containing enzymes isolated from bacteria suggested that
during the catalytic turnover nickel changes its oxidation state
from NiI to NiII to NiIII.3 Thus, the activity of the nickel center
in these enzymes is attributed to its redox chemistry. Unless the
coordination environment is known, it is not possible to dis-
tinguish Ni() from Ni() by using EPR data. The possibility of
antiferromagnetic coupling of paramagnetic nickel with the
[4Fe–4S] cluster complicates the issue even further.7
behavior of Ni() complexes with 1,2-dithiolates,9–21 dithio-
carbamates,8,22 and thioether ligands,23 have been studied in
considerable detail. In the case of 1,2-dithiolate complexes
the redox is centered on the ligands and no changes took
place in the metal oxidation state. With other ligands, the
formation12–14,16 and the isolation24–26 of Ni() complexes were
reported.
The studies on redox properties of nickel complexes with
xanthates and thioxanthates, which are good ligands for nickel,
are limited. Hendrickson et al.8 reported that the electro-
chemical oxidation of [Ni(S2COEt)3]Ϫ yielded [NiIII(S2COEt)3].
More recently, Chakravorty and co-workers investigated the
electrochemical behavior of [Ni(S2COR)3]Ϫ (R = CH3, C2H5,
n-C3H7, n-C4H9 and i-C4H9) complexes.27 From the low tem-
perature EPR measurement they claimed that electrochemical
oxidation of [Ni(S2COEt)3]Ϫ indeed led to the formation of
[NiIII(S2COEt)3], in which the nickel is in its “ϩ3” oxidation
state.
In this work we studied electrochemical and chemical
oxidation of S2COEtϪ, Ni(S2COEt)2, and [Ni(S2COEt)3]Ϫ in
acetonitrile. First we measured the cyclic voltammogram of
each species using an Ag/Agϩ reference electrode, then carried
out constant potential electrolysis at their lowest oxidation
peak potentials. Since the quantitative electronic absorption
spectra are known,28 the changes in the UV-VIS absorption
spectra are followed in situ during the electrolysis to identify the
oxidation products. In situ EPR spectra were measured in order
to detect any radical intermediates.
In the last two decades a great deal of research effort has
been directed towards the synthesis of complexes of Ni() and
Ni(),1,8–16 which can serve as models to further our under-
standing of this metal in biological systems. Polypeptides,17,18
tetraaza macrocycles,11 dioximes, and deprotonated amide1
† Part of this work was presented at the NATO ASI Conference, 15–26
June, 1996, Przesieka, Poland.
‡ Present address: Department of Chemistry, Bilkent University,
Ankara, Turkey.
§ Present address: Faculty of Engineering, Atılım University, Ankara,
Turkey.
Experimental
29
The ligand K(S2COEt) and the complexes Ni(S2COEt)2 and
(C2H5)4N[Ni(S2COEt)3]19 were prepared by the literature
methods. (n-C4H9)4N(BF4) was purchased from Aldrich and
DOI: 10.1039/b105683m
J. Chem. Soc., Dalton Trans., 2001, 2819–2824
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Fig. 1 The apparatus used for measuring the in situ spectral changes
during the constant potential electrolysis at room temperature; RE: Ag-
wire reference electrode; CE: Pt-sieve counter electrode; WE: Pt-wire
working electrode.
Fig. 2 Decomposition kinetics of [NiIII(S2COEt)3] from CV measure-
ments at 273 K. (a) The decay of the reduction peak current; (b) First
order fit (R2 = 0.918); (c) Second order fit (R2 = 0.997).
time, at 273 K, is presented in Fig. 2(a). To find the order of
the decomposition reaction with respect to [NiIII(S2COEt)3],
first ln(i0 Ϫ i∞)/(it Ϫ i∞) and then 1/(it Ϫ i∞) versus time graphs
were plotted using the same data, where i0 and i∞ values are
obtained by plotting the cathodic peak heights versus time
(i.e. time intervals between the anodic and cathodic peaks on
the voltammogram). This experiment was performed at three
different temperatures, 253, 273 and 297 K. The plots for the
decomposition reaction at 273 K are shown in Fig. 2(b) and 2(c)
as examples.
used as supporting electrolyte without further purification.
Iodine and potassium cyanide used were reagent quality.
Acetonitrile, which was used as solvent, was Aldrich spectro-
scopic grade. The quantitative electronic absorption spectra28
were used for identification. The elemental analysis for each
compound was satisfactory.
Voltammetric recordings were made with a Lloyd PL3 XY/t
recorder. A Pt-bead working, Pt-wire counter and an Ag/Agϩ
(0.10 M) reference electrode were used for CV studies. These
three electrodes were positioned as close as possible to minimize
IR drop. Cyclic voltammograms were recorded under a nitro-
gen atmosphere in acetonitrile at room temperature. Controlled
potential electrolysis at the anodic peak potentials, were carried
out with an Ag-wire reference electrode using the apparatus
shown in Fig. 1. Both reference electrodes were calibrated
against a ferrocene/ferrocenium couple, the oxidation potential
was found to be 0.02 V versus Ag/Agϩ (0.10 M) and 0.52 V
versus Ag-wire. Oxidation processes have been followed in situ
on a Hewlett-Packard UV-VIS HP 8452A diode array spectro-
photometer. An FTIR Nicolet 510 spectrometer was also used
for characterization of the compounds.
EPR method
The dichloromethane solution of (C2H5)4N[Ni(S2COEt)3] in the
EPR cell is electrolyzed for 60 s using constant current (10 µA).
EPR measurements were made on a Varian E12 EPR spec-
trometer. After the electrolysis stopped, the signal height was
measured at different time intervals. The rate of the decom-
position reaction was studied by following the decay of the
EPR signal with time. This experiment was carried out at 231,
238, 245, and 258 K. The EPR signal obtained during the
electrolysis at 245 K is shown in Fig. 3(a) and a representative
example for ln((signal height)0/(signal height)t) and 1/(signal
height)t versus time plots are given in Fig. 3(b) and 3(c),
respectively.
The rate of the decomposition reaction of [NiIII(S2COEt)3]
formed by the electrochemical oxidation of [Ni(S2COEt)3]Ϫ,
was studied using CV and EPR methods.
CV method
Results and discussion
The [NiIII(S2COEt)3] complex was produced by electrolysis in a
CV cell. The anodic scan (50 mV sϪ1) started at Ϫ0.40 V and
stopped at Ϫ0.15 V, which is the reversible oxidation peak
potential for [Ni(S2COEt)3]Ϫ, and the electrolysis was carried
out for 60 s. After the electrolysis, the cycle (between 0.00 V and
Ϫ0.40 V) was completed at different scan rates. Since the scan
rate between Ϫ0.40 V and Ϫ0.15 V, electrolysis time and the
concentration of [Ni(S2COEt)3]Ϫ were kept constant, the con-
centration of [NiIII(S2COEt)3] produced each time should also
be the same. Completing the cycle at different scan rates gives
different reduction peak heights, which reflects the different
concentrations of the remaining [NiIII(S2COEt)3] species in
solution. This allowed us to study the kinetics of the decom-
position reaction. The decay of the reduction peak current with
Oxidation of K(S2COEt)
The cyclic voltammogram of K(S2COEt) in acetonitrile shows
an irreversible oxidation peak at around 0.00 V referenced to
Ag/Agϩ electrode (Fig. 4(a)). Electrochemical oxidation has
been carried out in CH3CN at this potential and the spectral
changes are followed in situ on a UV-VIS spectrophotometer.
The spectral changes during the constant potential electrolysis
are shown in Fig. 5(a). The most intense band at 308, as well as
the band at 226 nm in the spectrum28 of S2COEtϪ are dimin-
ished, while the two new bands at 286 and 246 nm are gradually
developed. These bands are characteristic of the oxidized dimer
form of the ethylxanthate ligand, (S2COEt)2. The isosbestic
points at 290 and 240 nm clearly indicate a one-step reaction,
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Fig. 3 (a) EPR spectrum of [NiIII(S2COEt)3] measured at 245 K; (b) Decay of EPR signal at 245 K, first order fit (R2 = 0.901); (c) Decay of EPR
signal at 245 K, second order fit (R2 = 0.967).
spectrum, influences the appearance of the higher energy part
of the absorption spectrum of (S2COEt)2 (IϪ has absorption
peaks at 246 (ε = 15300 cmϪ1 MϪ1) and 206 nm (ε = 20600 cmϪ1
M
Ϫ1) in acetonitrile).
Our further investigation shows that the electrochemically
generated dimer undergoes reduction with CNϪ (Fig. 5(c)) as in
eqn. (3).
(S2COEt)2 ϩ 2CNϪ
2S2COEtϪ ϩ (CN)2
(3)
Comparison of the changes in the absorption spectra given in
Fig. 5, indicate that the electrochemical and chemical oxidation
of ethyl xanthate anion yields the same product, and it is also
possible to reverse the oxidation by CNϪ to regenerate the ethyl
xanthate anion.
Oxidation of Ni(S2COEt)2
There are two irreversible oxidation peaks in the cyclic voltam-
mogram of Ni(S2COEt)2 in acetonitrile at 0.35 and 0.80 V
versus Ag/Agϩ reference electrode (Fig. 4(c)). Constant poten-
tial electrolysis has been carried out at the first peak potential.
The spectral changes during the electrolysis of about 10Ϫ5
M
Ni(S2COEt)2 are shown in Fig. 6. The bands at 414, 474, 316
and 218 nm, which are characteristic bands in the spectrum of
Ni(S2COEt)2 gradually lose their intensities, as the bands at
286 and 246 nm, originating from (S2COEt)2, intensify with
time during the electrolysis. Two sharp isosbestic points at 290
and 224 nm are observed. The final spectrum is consistent with
that of (S2COEt)2. This result suggests that oxidation is taking
place as in eqn. (4). At this concentration the presence of Ni2ϩ
28
Fig.
[Ni(S2COEt)3]; and (c) Ni(S2COEt)2, in acetonitrile–(n-C4H9)4N(BF4)
solvent–electrolyte couple at a voltage scan rate of 200 mV sϪ1
4
Cyclic voltammograms of (a) K(S2COEt); (b) (Et4N)-
.
leading to the formation of (S2COEt)2 in the electrolysis solu-
tion. This observation is consistent with one electron oxidation
of ethyl xanthate anion forming a radical followed by fast
dimerization to form disulfide(dimer) as in eqn. (1) and (2).
Ni(S2COEt)2
Ni2ϩ(sol) ϩ (S2COEt)2 ϩ 2eϪ
(4)
Ϫ
S2COEtϪ
2S2COEt
S2COEt ϩ e
(1)
(2)
ؒ
ion in solution is not detectable from the electronic absorption
spectrum, because the d–d absorption of the hexaacetonitrile-
nickel() complex is very weak.30 The presence of Ni2ϩ ion in
solution has been demonstrated by titration of the electrolyzed
solution with CNϪ, which yielded the [Ni(CN)4]2Ϫ complex,31
xanthate ligand, and cyanogen according to eqn. (5).
ؒ
(S2COEt)2
The chemical oxidation of S2COEtϪ by molecular iodine in
CH3CN also gives the same spectral changes (the isosbestic
point at 290 nm and the formation of the band at 286 nm),
Fig. 5(b). In this process S2COEtϪ is oxidized and I2 is reduced
to yield (S2COEt)2 and IϪ ion, respectively. However, the form-
ation of IϪ, which strongly absorbs in the UV region of the
Ni2ϩ(sol) ϩ (S2COEt)2 ϩ 6CNϪ
[Ni(CN)4]2Ϫ ϩ 2S2COEtϪ ϩ (CN)2 (5)
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Fig. 6 UV-VIS absorption spectral changes during the constant
potential electrolysis of a 3 × 10Ϫ5 M solution of Ni(S2COEt)2 in
acetonitrile. (There are 5 min intervals between scans.)
Fig. 5 UV-VIS absorption spectral changes during (a) constant
potential electrolysis of a 5 × 10Ϫ5 M solution of K(S2COEt) (there are
5 min intervals between scans), (b) chemical oxidation of K(S2COEt)
with I2, and (c) chemical reduction of (S2COEt)2 with CNϪ in
acetonitrile.
When the acetonitrile solution of Ni(S2COEt)2 was titrated
with saturated KCN solution (such that the volume change is
negligible during titration) and absorption spectral changes
were measured in situ, a set of sharp isosbestic points were
observed until the mole ratio of CNϪ/Ni2ϩ was 4/3. These
isosbestic points were lost upon addition of more CNϪ and a
new set developed. The spectrum of the solution at the end was
the same as the overlapped spectrum of [Ni(CN)4]2Ϫ and
K(S2COEt) in acetonitrile. These experimental results are
interpreted as two consequent reactions as in eqn. (6) and (7).
Fig. 7 UV-VIS absorption spectral changes during, (a) the constant
potential electrolysis of [Et4N][Ni(S2COEt)3] at the first oxidation peak
potential (there are 5 min intervals between scans); (b) followed, at the
second oxidation peak potential (5 min intervals between scans); and (c)
the chemical oxidation of [Et4N][Ni(S2COEt)3] by I2 in acetonitrile.
nitrile displays one reversible, at Ϫ0.15 V, and two irreversible,
at 0.35 and 0.80 V, oxidation peaks versus the Ag/Agϩ reference
electrode (Fig. 4(b)). Constant potential electrolysis of this
complex in acetonitrile was performed at the lowest oxidation
peak potential. The absorption spectral changes, which were
measured in situ during the electrolysis, are presented in
Fig. 7(a). Seven isosbestic points are observed as the bands due
to the starting complex disappeared and the characteristic
bands of Ni(S2COEt)2 developed. The final spectrum is the
same as the one obtained from the 1 : 2 mixture of (S2COEt)2
3Ni(S2COEt)2 ϩ 4CNϪ
[Ni(CN)4]2Ϫ ϩ 2[Ni(S2COEt)3]Ϫ (6)
2[Ni(S2COEt)3]Ϫ ϩ 8CNϪ
2[Ni(CN)4]2Ϫ ϩ 6S2COEtϪ (7)
Oxidation of (NEt4)[Ni(S2COEt)3]
The cyclic voltammogram of (NEt4)[Ni(S2COEt)3] in aceto-
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of the electrochemical oxidation at the first peak potential
(Fig. 7(a)). The differences around 246 and 206 nm are due to
the absorption of IϪ ion. This observation suggests that both
electrochemical and chemical oxidation yield the same final
products.
The electrochemical oxidation of [Ni(S2COEt)3]Ϫ was studied
by Chakravorty et al.27 Their experimental observations were
similar to ours. The equilibrium constant for the decomposition
reaction of [NiIII(S2COEt)3] to Ni(S2COEt)2 and (S2COEt)2 was
calculated to be 391 M at 298 K. The rate of the decomposition
reaction of nickel() complex was claimed to be first order with
a rate constant of 1.0 × 10Ϫ3 sϪ1 at 253 K; and the activation
entropy, ∆S*, was found to be about Ϫ40 eu. In fact, we have
also found a negative entropy of activation both from the CV
(≈ Ϫ46 eu) and EPR method (≈ Ϫ60 eu). If the decomposition
reaction rate is first order with respect to [NiIII(S2COEt)3], the
activation step in the proposed mechanism would be the
formation of [Ni (S2COEt)2(S2COEt )] complex radical from
[NiIII(S2COEt)3]. It is very difficult to visualize how an electron
transfer from a ligand to the nickel center results in such a large
decrease in entropy. The large negative entropy of activation
can be explained by a mechanism whereby the [NiIII(S2COEt)3]
decomposes to Ni(S2COEt)2 and (S2COEt)2 as shown in the
following scheme:
Fig. 8 EPR spectrum of [NiIII(S2COEt)3] in acetonitrile at 10 K.
(Frequency = 9.4653600 GHz.)
and Ni(S2COEt)2. This observation indicates that the oxidation
of [Ni(S2COEt)3]Ϫ at a potential corresponding to the lowest
oxidation peak potential in its cyclic voltammogram, yields
Ni(S2COEt)2, and (S2COEt)2, as the final products. Our room
temperature in situ measured UV-VIS spectral data do not
show any sign of intermediate formation during the oxidation
process. However, the in situ electrolysis of a dichloromethane
solution of (C2H5)4N[Ni(S2COEt)3] in an EPR cell, at low tem-
perature, yielded a singlet with a “g” value of 2.09 as shown in
Fig. 3(a). We have also measured the EPR spectrum of the
acetonitrile solution of tris(ethylxanthato)nickel() at 10 K,32
after constant potential electrolysis at 253 K. This spectrum is
presented in Fig. 8.
II
ؒ
The EPR spectrum clearly indicates the formation of an
Ni() species in the oxidative electrolysis process. Thus, the
reversible peak obtained on the voltammogram corresponds
to an Ni()/Ni() redox couple. Since no absorbance due
to the Ni() species24 has been recorded by in situ UV-VIS
measurements during the room temperature electrolysis, the
[NiIII(S2COEt)3] complex formed at the electrode surface
must decompose to the observed final products before it reaches
the light path in the UV cell. Therefore, the mechanism of the
electrochemical oxidation process at the first oxidation peak
potential can be represented by eqn. (8)–(10); where an electron
The above mechanistic scheme is a manifestation of eqn. (9)
and (10). The rate law expression for the decomposition of
[NiIII(S2COEt)3] to Ni(S2COEt)2 and (S2COEt)2, can be derived
using a rapid equilibrium assumption and has the following
form:
(8)
(9)
(10)
is reversibly removed from the tris(ethylxanthato)nickel()
anion forming [NiIII(S2COEt)3], which yields a radical com-
II
ؒ
plex, [Ni (S2COEt)2(S2COEt )], in a rapid equilibrium step.
The radical complex undergoes a bimolecular reaction to yield
Ni(S2COEt)2 and (S2COEt)2 as final products.
Since we have detected the presence of nickel() species by
EPR at low temperature, and by CV at room temperature, the
rapid equilibrium must mainly lie to the left. This implies that
2
Further controlled potential oxidative electrolysis of the
above electrolyzed solution was carried out at the second oxid-
ation peak potential. Spectral changes during the second
electrolysis are shown in Fig. 7(b). The spectral changes dis-
played in this figure and in Fig. 6, which displays the spectral
changes observed during the electrolysis of Ni(S2COEt)2, are
strikingly similar. Two sharp isosbestic points at 290 and 224
nm are common to both. These results leave no doubt that the
Ni(S2COEt)2 complex formed at the end of the electrolysis at
the first oxidation peak potential, is further oxidized to produce
Ni2ϩ(sol) and the dimer of the oxidized ligand, as shown in
eqn. (4).
K1 << 1, and k = 2k2K1 . If our proposed mechanism is correct
then the rate of the decomposition reaction must be second
order with respect to [NiIII(S2COEt)3]. In order to check this
hypothesis we followed the decay kinetics of nickel() species
using EPR and CV methods as described in the experimental
section. The kinetic data obtained from EPR and CV experi-
ments gave a better fit to a second order integrated rate law,
rather than the first order one (see Fig. 2 and 3). It was also very
clear from our data that the half-lives of the decay reaction
of the nickel() species were initial concentration dependent
(Fig. 2(a)). Both half-life and Wilkinson33 methods also sug-
gested a second order decay reaction. All these observations are
consistent with our proposed mechanism, in which the rate
determining step is the bimolecular reaction of the radical
complex (eqn. (10)), yielding a large negative entropy of
activation.
The chemical oxidation of [Ni(S2COEt)3]Ϫ in acetonitrile was
also carried out with iodine. The spectral changes measured
during the titration of [Ni(S2COEt)3]Ϫ with a concentrated
acetonitrile solution of I2 (Fig. 7(c)), are very similar to that
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Conclusion
The lowest oxidation peaks in the cyclic voltammograms of
S2COEtϪ, Ni(S2COEt)2, and [Ni(S2COEt)3]Ϫ appear at 0.00,
0.35, and Ϫ0.15 V, respectively, versus Ag/Agϩ reference
electrode in acetonitrile. The oxidation peak for [Ni(S2COEt)3]Ϫ
is reversible. Constant potential electrolysis carried out at
potentials corresponding to the above peak potentials for each
species, always yielded (S2COEt)2 as one of the final products.
The chemical oxidation of S2COEtϪ and [Ni(S2COEt)3]Ϫ with I2
resulted in the same final products as the electrochemical
oxidation.
The electrochemical oxidation of [Ni(S2COEt)3]Ϫ by constant
potential electrolysis at the first oxidation peak potential
resulted in the nickel() species, which then decomposed to
Ni(S2COEt)2 and (S2COEt)2 as the final products. The kinetic
data obtained in this work suggest that the decomposition
reaction is second order with respect to [NiIII(S2COEt)3].
In this study, we have demonstrated that electrochemical
and chemical oxidation–reduction processes can be followed
by in situ UV-VIS spectrophotometry to identify the final prod-
ucts of the process if the absorption spectra of the expected
products are known. However, it may not be possible to obtain
a real picture if the lifetime of the redox species are not long
enough.
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Acknowledgements
We would like to thank Middle East Technical University
Research Fund and Project TBAG-1520 for support of this
work.
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K was measured at MPI für
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J. Chem. Soc., Dalton Trans., 2001, 2819–2824