Influence of mixed thiolate/thioether versus dithiolate coordination on the accessibility of the uncommon +I and+III oxidation states for the nickel ion: An experimental and computational study

Article abstract of DOI:10.1021/ic200063d

Sulfur-rich nickel metalloenzymes are capable of stabilizing Ni^I and Ni^III oxidation states in catalytically relevant species. In an effort to better understand the structural and electronic features that allow the stabilization of s

Full text of DOI:10.1021/ic200063d

ARTICLE
pubs.acs.org/IC
Influence of Mixed Thiolate/Thioether versus Dithiolate Coordination
on the Accessibility of the Uncommon þI and þIII Oxidation States for
the Nickel Ion: An Experimental and Computational Study
||
Marcello Gennari,† Maylis Orio,‡,^ Jacques Pꢀecaut,§ Eberhard Bothe,|| Frank Neese,‡,
Marie-No€elle Collomb,*,† and Carole Duboc*,†
†Universitꢀe Joseph Fourier Grenoble 1/CNRS, Dꢀepartement de Chimie Molꢀeculaire, UMR-5250, Laboratoire de Chimie Inorganique
Redox, Institut de Chimie Molꢀeculaire de Grenoble FR-CNRS-2607, BP-53, 38041 Grenoble Cedex 9, France
‡Institute for Physical and Theoretical Chemistry Universit€at Bonn, Wegelerstrasse 12, D-53113 Bonn, Germany
§Laboratoire de Reconnaissance Ionique et Chimie de Coordination, Service de Chimie Inorganique et Biologique, (UMR E-3
CEA/UJF, FRE3200 CNRS), CEA-Grenoble, INAC, 17 rue des Martyrs 38054, Grenoble cedex 9, France
Max-Planck-Institut f€ur Bioanorganische Chemie, Stiftstr. 34-36, D-45470 M€ulheim an der Ruhr, Germany
S Supporting Information
b
ABSTRACT: Sulfur-rich nickel metalloenzymes are capable of stabilizing NiI
and NiIII oxidation states in catalytically relevant species. In an effort to better
understand the structural and electronic features that allow the stabilization
of such species, we have investigated the electrochemical properties of two
mononuclear N2S2 NiII complexes that differ in their sulfur environment.
Complex 1 features aliphatic dithiolate coordination ([NiL], 1), and complex
2I is characterized by mixed thiolate/thioether coordination ([NiLMe]I, 2I).
The latter results from the methylation of a single sulfur of 1. The X-ray
structure of 2I reveals a distorted square planar geometry around the NiII ion,
similar to what was previously reported by us for 1. The electrochemical
investigation of 1 and 2þ shows that the addition of a methyl group shifts the
potentials of both redox NiII/NiI and NiIII/NiII redox couples by about 0.7
and 0.6 V to more positive values. Through bulk electrolyses, only the
mononuclear dithiolate [NiIL]ꢀ (1-) and the mixed thiolate/thioether
[NiIIILMe 2þ (22þ) complexes were generated, and their electronic proper-
]
ties were investigated by UVꢀvis and EPR spectroscopy. For 1- (NiI, d9 configuration) the EPR data are consistent with a dx2-y2 based
singly occupied molecular orbitals (SOMOs). However, DFT calculations suggest that there is also pronounced radical character.
This is consistent with the small g-anisotropy observed in the EPR experiments. The spin population (Mulliken analysis) analysis
of 1- reveals that the main contribution to the SOMO (64%) is due to the bipyridine unit. Time dependent density functional
theory (TD-DFT) calculations attribute the most prominent features observed in the electronic absorption spectrum
of 1- to metal to ligand charge transfer (MLCT) transitions. Concerning 22þ, the EPR spectrum displays a rhombic signal with
gx = 2.236, gy = 2.180, and gz = 2.039 in CH3CN. The giso value is larger than 2.0, which is consistent with metal based oxidation. The
unpaired electron (NiIII, d7 configuration) occupies a Ni-dz2 based molecular orbital, consistent with DFT calculations. Nitrogen
hyperfine structure is observed as a triplet in the gz component of the EPR spectrum with AN = 51 MHz. This result indicates the
coordination of a CH3CN molecule in the axial position. DFT calculations confirm that the presence of a fifth ligand in the
coordination sphere of the Ni ion is required for the metal-based oxidation process. Finally, we have shown that 1 exhibits catalytic
reductive dehalogenation activity below potentials of ꢀ2.00 V versus Fc/Fcþ in CH2Cl2.
’ INTRODUCTION
acetyl-coenzyme A (acetyl-CoA) from CO, CoA, and a methyl
group.12ꢀ14 In these enzymes, the nickel sites exhibit extreme
plasticity in both metal coordination and redox chemistry. In
addition to the regular þII oxidation state of the thiolateꢀnickel
Among the eight known nickel enzymes engaged in various
redox activities,1 four involve thiolateꢀnickel bonds in their
active site. The Ni superoxide dismutase (NiSOD) catalyzes the
disproportionation of superoxide radical anion into hydrogen
peroxide and molecular dioxygen.2ꢀ5 NiFe hydrogenase rever-
sibly reduces protons to H2.6ꢀ8 CO dehydrogenase reversibly
oxidizes CO to CO2.9ꢀ11 Finally, acetyl-CoA synthase produces
Received: January 11, 2011
Published: March 23, 2011
r
2011 American Chemical Society
3707
dx.doi.org/10.1021/ic200063d Inorg. Chem. 2011, 50, 3707–3716
|

Inorganic Chemistry
ARTICLE
Scheme 1
complexes, the uncommon þI and þIII oxidation states are
involved in several catalytically relevant species.
In view of elucidating the catalytic mechanisms of these sulfur-
rich metalloenzymes, considerable efforts have been directed
toward synthesizing low molecular weight Ni complexes that
mimic some of the active site's structural, spectroscopic, and
functional properties.15ꢀ17 Importantly, a number of square
planar dithiolate NiII compounds in an N2S2 coordination sphere
have been described.15 Their structural, electronic, and redox
properties have been investigated in detail, as well as their
reactivity in S-oxygenation18ꢀ21 or/and S-alkylation22ꢀ26 reac-
tions. These studies have given insight into the nucleophilicity of
the S-thiolate ligands versus the reaction products, i.e., S-oxyge-
nated derivates or S-thioethers, respectively.
By contrast, substantially fewer studies were directed toward
NiIII thiolate complexes. To the best of our knowledge, a single
example of such a species has been structurally characterized
so far.27 The other available reports concern in situ generated
NiIII complexes, either by chemical or electrochemical oxidation
process of the corresponding NiII species.15 The exploration of
the electronic structure of these NiIII complexes by EPR spec-
troscopy and DFT calculations has suggested a highly covalent
character for the NiꢀS bond, as well as a better stabilization of
this oxidation state in a five- or six-coordinated environment
compared to a square planar geometry.19,28ꢀ33 However, the
influence of a mixed thiolate/thioether environment on the
electronic properties of the NiIII ion has not been investigated
in depth so far.33,34
Figure 1. ORTEP drawing of the two independent molecules of
[NiLMe]I 2H2O (2I 2H2O, 30% thermal ellipsoids). Hydrogen atoms,
iodide counteranions, and water molecules are omitted for clarity.
3
3
types of aliphatic sulfur donors, e.g., pure thiolate versus mixed
thiolate/thioether environments. We have recently reported the
synthesis and characterization of a mononuclear diamine alipha-
tic dithiolate NiII complex, [NiIIL] (1) [L2- = 2,20-(2,20-bipyr-
idine-6,60-diyl)bis(1,1-diphenylethanethiolate), Scheme 1], that
does yield metal-based oxidation. Interestingly, upon electro-
chemical oxidation in the presence of imidazole (im), the square
planar N2S2 NiII complex is converted into a square pyramidal
N3S2 NiIII complex, [NiL(im)]þ, thus mimicking the redox
structural changes of NiSOD during its catalytic cycle.
Synthetic NiI thiolate complexes are even more sparse than
those of NiIII because of their instability with respect to dispro-
portionation to NiII and Ni0 species.35 Examples of structurally
characterized mononuclear thiolate NiI complexes have only
been reported in organometallic chemistry.36 A few NiI thiolate
complexes have been generated in situ. However this occurred
exclusively from NiII complexes with aromatic thiolates37ꢀ40 or
with mixed sulfur environments (thiolate/thioether or thiolate/
SO2),22 both of which are known for stabilizing the þI oxidation
state. Furthermore, the investigation of the electronic properties
of these NiI species has, until now, been very limited.22,37,38,40
Consequently, to the best of our knowledge, no NiI complex in an
aliphatic thiolate environment, as found in the mimicked metal-
loenzymes, has been reported.
We report here the isolation and structural characterization of
the mixed thiolate/thioether [NiIILMe]I complex (2I), which
results from the mono-S-methylation of 1 (Scheme 1). The
redox chemistry of these two NiII complexes was investigated
either in coordinating (CH3CN or DMF)41 or in noncoordinat-
ing (CH2Cl2) solvent, with particular attention given to the
influence of the thiolate/thioether donors in the accessibility of
the þI and þIII oxidation states for the metal ion. Unexpectedly,
It is thus of crucial importance to improve our knowledge on
the [NiIIILMe 2þ
]
(22þ) complex with a mixed thiolate/thioether
the electronic properties of NiIII and NiI species with different
3708
dx.doi.org/10.1021/ic200063d |Inorg. Chem. 2011, 50, 3707–3716

Inorganic Chemistry
ARTICLE
Table 1. Selected Bond Lengths (Å) and Angles (deg) Issued from the X-ray Structures of 2I 2H2O and 1 As Well As from Their
3
Optimized Structures
2I 2H2O
1
3
exptl
calcd
1.897
exptl32
calcd
1.918
Ni(1)ꢀN(1)
1.899(8)
1.914(7)
1.918(7)
1.959(8)
2.169(3)
2.174(3)
2.172(2)
2.187(2)
84.6(3)
NiꢀN(1)
1.934(2)
1.935(2)
2.173(1)
2.176(1)
84.56(7)
98.73(5)
161.11(5)
159.97(5)
99.06(5)
84.25(2)
Ni(2)ꢀN(51)
Ni(1)ꢀN(2)
1.934
2.215
2.223
83.4
NiꢀN(2)
1.918
2.180
2.180
84.5
Ni(2)ꢀN(52)
Ni(1)ꢀS(1)
NiꢀS(1)
Ni(2)ꢀS(51)
Ni(1)ꢀS(2)
NiꢀS(2)
Ni(2)ꢀS(52)
N(1)ꢀNi(1)ꢀN(2)
N(51)ꢀNi(2)ꢀN(52)
N(1)ꢀNi(1)ꢀS(1)
N(51)ꢀNi(2)ꢀS(51)
N(1)ꢀNi(1)ꢀS(2)
N(51)ꢀNi(2)ꢀS(52)
N(2)ꢀNi(1)ꢀS(1)
N(52)ꢀNi(2)ꢀS(51)
N(2)ꢀNi(1)ꢀS(2)
N(52)ꢀNi(2)ꢀS(52)
S(1)ꢀNi(1)ꢀS(2)
S(51)ꢀNi(2)ꢀS(52)
N(1)ꢀNi-N(2)
N(1)ꢀNi-S(1)
N(1)ꢀNi-S(2)
N(2)ꢀNi-S(1)
N(2)ꢀNi-S(2)
S(1)ꢀNi-S(2)
85.0(3)
94.8(2)
94.2
100.6
157.5
157.6
100.6
83.1
98.0 (3)
173.7(2)
176.2(2)
162.3(2)
163.2(2)
98.2(2)
173.6
160.0
99.9
98.4(2)
80.78(9)
79.27(10)
80.9
coordination sphere and the aliphatic dithiolate [NiIL]ꢀ (1-)
complex were successfully obtained by exhaustive electrolysis. To
the best of our knowledge, these two compounds represent the
first examples of NiI and NiIII species in such environments.
Their electronic properties were investigated by UVꢀvis and
EPR spectroscopies combined with DFT calculations. Finally,
the capability of 1 to perform electrocatalytic reductive dehalo-
genation reactions in CH2Cl2 solutions is discussed.
’ RESULTS
1. Synthesis and Characterization of [NiLMe þ (2þ). The
]
mono S-alkylated [NiLMe]I complex (2I) was synthesized from
the reaction of 132 with 1.2 molar equiv of CH3I in CH2Cl2
accompanied by a drastic change of the color solution, from
violet to dark yellow (Scheme 1). This complex was obtained in a
good yield as expected from previous studies demonstrating that
the second S-alkylation occurs more slowly than the first
one.24,42,43 To avoid the presence of the electroactive iodide
counteranion during the electrochemical investigation, this anion
was exchanged by CF3SO3ꢀ (OTfꢀ) by adding 1 molar equiv of
AgCF3SO3 in a 2I solution, giving 2OTf (see Experimental
Section).
Figure 2. CV of 1 (1.6 mM) in DMF (—) and in CH2Cl2 ( ), 0.1 M
3 3 3
Bu4NPF6 at room temperature. Scan rate of 100 mV sꢀ1
.
3
The nearest iodide is located at a distance of 3.778 Å from the
nickel (Ni1), showing that the anion is not coordinated to
the metal.
The NiꢀS bond lengths are not very sensitive to the nature of
the S-donor ligand, with the NiꢀSthiolate bond being slightly
shorter than the NiꢀSthioether one. In contrast, the NiꢀN bond
trans to the thioether is shorter than the bond trans to the
thiolate, as previously observed in other monomethylated N2S2
NiII complexes.21,23,24 Upon S-methylation of 1 into 2þ, a
stronger sensitivity of the NiꢀS distances than the NiꢀN ones
would be expected because of the weaker σ-donor character of
the neutral thioethers than the anionic thiolates. However, the
invariance of the NiꢀS bond length upon S-alkylation is attrib-
uted to a relief of the dπꢀpπ antibonding interaction between the
metal and sulfur.22,24,25,44,45
The ORTEP drawing of the two independent molecules
present in the unit cell of 2I is shown in Figure 1, and selected
bond lengths and angles are given in Table 1 together with those
of 1. The NiII ion is in the center of a square plane in a
N2Sthioether thiolate donor environment. Slight distortions are
S
evidenced by the fact that the Ni, N, and S atoms are displaced
from the mean NiN2S2 plane from 0.07 to 0.20 Å. The six-
membered nickel-aza-thiacyclohexane ring in 2I obviously leads
to larger deviation from planarity with respect to the more common
five-membered nickel-aza-thiacyclopentane ring complexes.21,23
The S-methylation of 1 leads to the closure of the SꢀNiꢀS
angle (about 4ꢀ), while no significant changes are observed in the
3709
dx.doi.org/10.1021/ic200063d |Inorg. Chem. 2011, 50, 3707–3716

Inorganic Chemistry
ARTICLE
Table 2. Electrochemical Potentials for 1 and 2þ at Room Temperature (vs Fc/Fcþ, 100 mV sꢀ1
)
3
complex
solvent
oxidation process NiIIfNiIII
reduction process NiIIfNiI
[NiL] (1)
[NiL] (1)
DMF
Epa = þ0.05 V; Epc = ꢀ0.73 V
Epa = þ0.13 V; Epc = ꢀ0.55 V
E1/2 = þ0.60 V (ΔEp = 70 mV)
E1/2 = þ0.68 V (ΔEp = 140 mV)
E1/2 = ꢀ1.85 V (ΔEp = 70 mV)
E1/2 = ꢀ1.90 V (ΔEp = 130 mV)
E1/2 = ꢀ1.19 V (ΔEp = 60 mV)
E1/2 = ꢀ1.22 V (ΔEp = 100 mV)
CH2Cl2
CH3CN
CH2Cl2
[NiLMe þ
[NiLMe þ
]
(2þ)
(2þ)
]
Scheme 2
NꢀNiꢀN angle (less than 0.5ꢀ). The SꢀCCH3 distance of
1.858(9) Å is also typical.25
2. Electrochemical Properties of Complexes [NiL] (1). We
have previously reported the redox properties of complex 1 in
CH2Cl2.32 The cyclic voltammogram (CV) displays two one-
electron metal based processes: a poorly reversible NiII/NiI
reduction wave at E1/2 = ꢀ1.91 V versus Fc/Fcþ and an
irreversible NiIII/NiII oxidation peak at Epa = þ0.13 V.46 The
NiIII complex undergoes a fast chemical reaction leading to a NiIII
dimer, which is irreversibly reduced at Epc = ꢀ0.55 V. In the
presence of imidazole, the electrochemically oxidized species was
characterized as a mononuclear [NiIIIL(im)]þ complex.
Figure 3. Evolution of the absorption spectrum of 1 (0.16 mM, 1 cm
optical path) in DMF, 0.1 M Bu4NPF6 during the exhaustive electrolysis
at ꢀ1.96 V vs Fc/Fcþ at room temperature (time between two spectra:
Δt = 2 min).
In this work, the reduction properties of 1 were further
investigated in DMF and CH2Cl2. The CV of 1 in DMF is
analogous to that observed in CH2Cl2 (Figure 2). We can
however notice slight potential shifts of the redox systems, and
a better reversibility of the NiII/NiI reduction process in DMF
(Table 2). The poor reversibility of this wave in CH2Cl2 arises
from catalytic reductive dehalogenation of the solvent. This has
been evidenced when a potential of ꢀ2.00 V versus Fc/Fcþ is
applied on a CH2Cl2 solution of 1. First, a lack of current
decrease is observed upon this applied potential. Second, when
total charge (Q) approached a calculated value of 8 electrons per
initial 1 complex, the CV of the solution is similar to the initial
one except for the appearance of an irreversible oxidation peak at
0.8 V corresponding to the electroactivity of the released chloride
anions. Furthermore, the amount of involved electrons exactly
corresponds to the produced amount of Clꢀ. On the basis of
these data together with those of the literature, the mechanism
displayed in Scheme 2 may be proposed.
Figure 4. Experimental (—) and simulated (---) X-band EPR spectra
(30 K) of 1-, electrogenerated from a solution of 1 (0.15 mM) at ꢀ25 ꢀC
in CH2Cl2, 0.1 M Bu4NPF6. Parameters used for the simulation: g =
2.110, g^ = 2.034, A = 212 MHz, A^ = 32 MHz.
The one-electron reduced form of 1, [NiL]ꢀ (1-), can be
afforded by a bulk electrolysis carried out at room temperature
(E = ꢀ1.96 V) in DMF (Figure S2). Upon the reduction process,
the solution changes from violet to brown with the appearance of
two intense transitions at 25 000 and 21 050 cmꢀ1, as well as two
less intense transitions at 12 350 and 11 100 cmꢀ1 (Figure 3).
The intensity of these electronic transitions is maximal after the
exchange of 1.75 electrons per molecule of 1. This excess of
coulometry is proposed to be the result of the spontaneous
reoxidation of 1- to 1 in the course of the electrolysis. This occurs
without any degradation of the complex as attested by the CV
recorded at the end of the electrolysis (Figure S2). From
chronoamperometry, 1- is generated with an estimated yield of
90%. In CH2Cl2, the same NiI species can also be produced in a
comparable amount, but only at low temperature (Figure S3).
Finally, the metal based character of the process is confirmed
by EPR spectroscopy. The frozen solution EPR spectrum of an
electrogenerated solution of 1- in DMF or CH2Cl2 (Figure 4)
displays an axial signal with g > g^ > 2.0023, typical of NiI
1
complexes having an S = /2 ground state with the unpaired
electron (d9 for a NiI ion) occupying a predominantly dx2-y2
orbital.26,47,48 However, owing to the small anisotropy of g (g^ =
2.034, g = 2.110), some radical character of the species is
suggested. Five 14N hyperfine lines are observed on the g
component, consistent with the delocalization of the spin density
3710
dx.doi.org/10.1021/ic200063d |Inorg. Chem. 2011, 50, 3707–3716

Inorganic Chemistry
ARTICLE
Figure 7. Experimental (—) and simulated (---) X-band EPR spectra
(100 K) of 22þ, electrogenerated from 1 mM solutions of 2þ (CH3CN
or CH2Cl2, 0.1 M Bu4NPF6).
Figure 5. CV of 2OTf (1 mM) in CH3CN, 0.1 M Bu4NPF6 (a) at room
temperature (__) and at ∼ ꢀ30 ꢀC ( ) and (b) after exhaustive
Upon examination of first the reduction of 2þ, although the
NiII/NiI wave appears reversible in CH3CN, and the correspond-
ing redox potential that is less negative by 0.7 V compared to that
of 1, an exhaustive reduction of the solution at room temperature
(E = ꢀ1.35 V) shows that the reduced NiI species (2) is not
stable. Indeed, at the end of the electrolysis (about 1 electron
consumed), the resulting CV shows the disappearance of the
initial NiII/NiI and NiIII/NiII waves while several waves corre-
sponding to undefined electroactive species appeared. Attempts
to stabilize the reduced NiI species by performing reductive
electrolysis at low temperature have failed.
3 3 3
electrolysis at þ0.70 V at ∼ꢀ30 ꢀC. Scan rate of 100 mV sꢀ1
.
Regarding the oxidation process, when an electrolysis at þ0.70
V of a CH3CN solution of 2þ is carried out at room temperature,
one electron per initial complex is exchanged. The CV of the
resulting solution shows the nearly complete disappearance of
the NiIII/NiII wave at E1/2 = þ0.60 V, and the appearance of a
shoulder at Epc = ꢀ1.00 V that precedes the NiII/NiI reduction
wave of 2þ (Figure S6). This demonstrates the instability of
[NiIIILMe 2þ (22þ) at room temperature and its chemical
]
Figure 6. Absorption spectra of 2þ (1 mM) in CH3CN and of 22þ
,
electrogenerated from 1 mM solutions of 2þ (CH3CN or CH2Cl2, 0.1 M
conversion to another electroactive species. It is proposed that
22þ decomposes via the formation of an intermediate thiyl
radical leading to an SꢀS bond. This chemical transformation
is reversible since an exhaustive reduction of the oxidized
solution carried out at ꢀ1.10 V restores 2þ via disulfide reduc-
tion, with a yield of about 80% (Figure S6).
Bu4NPF6) (1 mm optical path).
of the unpaired electron over the bipyridine units (two nitrogen
atoms).
3. Electrochemical Properties of Complexes [NiLMe]þ (2þ).
The redox properties of 2þ have been investigated in CH3CN
and CH2Cl2. The CV of 2þ in CH3CN (Figure 5a) displays two
reversible one-electron metal based processes: a NiII/NiI reduc-
tion wave at E1/2 = ꢀ1.19 V, and a NiIII/NiII oxidation wave at
E1/2 = þ0.60 V. By cycling at lower potential values, a second
poorly reversible reduction wave at Epc = ꢀ2.00 V is observed
that might be attributed to a metal or ligand based reduction
process (Figure S4). As for 1, the nature of the solvent (CH3CN
versus CH2Cl2) slightly affects the potentials of the redox
systems of 2þ (Table 2), and the NiII/NiI reduction wave is also
less reversible in CH2Cl2 (Figure S5).
However, if exhaustive oxidations of 2þ are carried out at low
temperature (∼ꢀ30 ꢀC), the mononuclear 22þ complex is almost
quantitatively generated either in CH3CN (E = þ0.70 V) or in
CH2Cl2 (E = þ0.89 V) (exchange of one electron per 2þ)
(Figures 5b and S5b). Upon oxidation, the color of the solution
turns from dark yellow to red. The resulting absorbance and EPR
spectra are slightly different in both solvents, possibly due to the
coordination of either a CH3CN molecule or a triflate counter-
anion (in CH2Cl2) to 22þ. In CH2Cl2, the spectrum exhibits two
main transitions at 21 350 and 18 870 cmꢀ1 (shoulder), while
in CH3CN additional transitions are observed at 26 320 and
13 700 cmꢀ1 (Figure 6). The EPR spectra of 22þ recorded in
both solvents are consistent with a mononuclear NiIII species
(Figure 7). The frozen solution spectra were simulated with the
following parameters: gx =2.238, gy = 2.178, gz = 2.039 in CH3CN
and gx =2.231, gy = 2.172, gz = 2.039 in CH2Cl2. The giso values
significantly larger than 2.0 are consistent with an S = 1/2 ground
statefortheNiIII species with the unpaired electron (d7 for a NiIII ion)
A comparison with the potential values of 1 shows that the
addition of a methyl group induces a positive potential shift for
both redox NiII/NiI and NiIII/NiII systems of about 0.7 and 0.6 V,
respectively. This potential shift is consistent with the decreased
donor ability of the sulfur upon S-alkylation and with the
development of positive charge on the mono-S-methylated
complex.22,24,25
3711
dx.doi.org/10.1021/ic200063d |Inorg. Chem. 2011, 50, 3707–3716

Inorganic Chemistry
ARTICLE
Figure 8. Localized SOMOs for 1-, 22þ, and 2(CH3CN)2þ complexes.
occupying a predominantly Ni-dz2 orbital.49 It has been shown
that the NiIII state is stabilized by interactions with donor axial
ligand(s). This is confirmed by the EPR spectrum recorded in
CH3CN. The overall EPR spectral envelope is similar to that in
CH2Cl2, with the exception of the three 14N hyperfine lines
imposed in the gz component with AN = 51 MHz. The triplet
is compatible with the coordination of one CH3CN molecule
in the axial coordination site of a z-axis elongated square
pyramid.27,29,33 In CH2Cl2, the assumption that the fifth ligand
(such as a water molecule or a triflate anion) is sufficiently
coordinating to interact with the NiIII ion would account for the
identical position of the gz signal in both solvents.
Figure 9. TDDFT assignment of the calculated transitions for 1- in
CH2Cl2. The population of the relevant MOs (HOMO: highest
occupied molecular orbital; LUMO, lowest unoccupied molecular) is
indicated in parentheses as well as the wavelength of the optical
transitions. Legend: Ni represents the metal, S the sulfur atoms, and L
the rest of the ligand.
4. Electronic Structure of [NiIL]ꢀ (1-) and [NiIIILMe 2þ (22þ
)
]
and Interpretation of Their Spectroscopic Properties. To
obtain some insight into the geometrical and electronic structure
of the NiI complex, 1-, DFT calculations were carried out.
Comparison between the optimized structures of 1 and 1- reveals
a comparable environment around the nickel ion. Nevertheless,
the reduction of the Ni ion engenders an overall expected
expansion of its coordination sphere with longer NiꢀN/S
distances of about 0.01 and 0.04 Å, respectively (Table S2).
Upon examination of the electronic structure of this NiI species,
the singly occupied molecular orbital (SOMO) of 1- indicates a
strong delocalization of the unpaired electron over the metal and
the ligand, and especially over the bipyridine unit (Figure 8). The
Mulliken population analysis further supports that the spin
density is spread all over the complex since positive spin
populations are found at the Ni ion (0.31), and the bipyridine
unit (0.69 with 0.14 on each nitrogen atom) (Table S3). The
dominant contribution of the bipyridine unit to the total spin
density provides evidence for partial radical character for the 1-
species, which is consistent with the experimental EPR data.
Interestingly, a non-negligible part of the unpaired electron is
located in the Ni dx2-y2 metal-based MO. On the other hand, the
low sulfur contribution to the SOMO indicates a limited degree
of covalency for the NiꢀS bond in contrast to what is observed in
the previously described thiolate NiII and NiIII complexes.30,31,33
The EPR parameters calculated on the optimized structure of
1- (gx = 2.000, gy = 2.025, and gz = 2.107) are in very good
agreement with experiment and confirm the validity of the DFT
results for this NiI species (Table S4).
mainly from metal-to-ligand charge transfer (MLCT) transitions
since the donor orbital is mainly metal-centered. For the transi-
tions at 18 657 and 20 964 cmꢀ1, the acceptor orbitals are ligand-
based MOs that are fully delocalized over the four phenyl groups
of L. In the case of the 24 752 cmꢀ1 transition, the acceptor
orbital is also ligand-based but features a dominant contribution
from the carbon atoms of the bipyridine unit.
DFT calculations were also undertaken to model the electro-
nic and geometrical structure of the NiIII complex, 22þ. When
looking at the electronic structure of the optimized NiIII species
initiated from the X-ray structure of 2þ, we observe that DFT
calculations predict the single unpaired electron to mainly reside
in a S(π)-based MO (Figure 8). In fact, the SOMO of 22þ
features a large contribution from the thiolate unit (46%), which
dominates over the Ni character (19%). This is further confirmed
by the Mulliken population analysis showing that the thiolate
bears most of the spin density with positive spin populations
found at the S(1) atom (0.54), and the Ni ion (0.18) (Table S6).
Consequently, this electronic structure may best be described as
consisting of a NiIIꢀthiyl radical species. This result is in
complete disagreement with the EPR spectroscopic data re-
corded on the one-electron oxidized solution of 2þ (Table S4
and Figure 7). Moreover, it is worth noting that a similar result
was previously obtained from DFT calculations performed on a
diamine dithiolate NiIII complex which showed that the addition
of a fifth ligand was necessary to reproduce the spectroscopic
properties of this complex.31
For a better understanding of the optical properties of 1-, TD-
DFT calculations were undertaken on its optimized structure. By
employing this computational method, we can provide the
assignment of its salient features by calculating the main transi-
tion energies and their relative intensities. Our calculations
adequately reproduce the key features of the spectrum in terms
of energy and intensity, namely, the two intense bands at 25 000
and 21 050 cmꢀ1 (calc: λmax(f) = 24 752 (0.019) and 20 964
(0.044) cmꢀ1, respectively) and the shoulder at 18 800 cmꢀ1
(calc: λmax(f) = 18 657 (0.011) cmꢀ1) (Table S5). Regarding the
nature of these absorptions, Figure 9 shows that they all arise
To be consistent with the spectroscopic data, CH3CN solvent
molecules in our subsequent computational models of 22þ were
thus explicitly included. The best geometry optimized model,
namely 2(CH3CN)2þ, is generated with only one CH3CN
ligand coordinated to an axial position of the Ni ion featuring a
NiꢀNCH3CN distance of 2.108 Å (Table S6). The computed
electronic structure of this latter species correctly places the
3712
dx.doi.org/10.1021/ic200063d |Inorg. Chem. 2011, 50, 3707–3716

Inorganic Chemistry
ARTICLE
It has been previously observed that the NiI oxidation state is
stabilized upon conversion of thiolate(s) to thioether(s), while
the NiIII oxidation state is destabilized in the cationic thioether
derivatives.22ꢀ26 This behavior has been explained by the loss of a
SpπꢀNidπ antibonding interaction in the NiꢀSR HOMO, and
by the increase of the positive charge on the complex.23,24,44,45
Unexpectedly, in this work, we were able, with the help of the
ligands L and LMe, to stabilize a mononuclear dithiolate [NiIL]ꢀ
(1-) complex and a mixed thiolate/thioether [NiIIILMe 2þ (22þ
)
]
complex. All of our attempts to generate either the dithiolate
[NiIIIL]þ (1þ) complex in the absence of a strong donor ligand
such as imidazole or the mixed thiolate/thioether [NiILMe] (2)
species have failed.
The surprising stability of 1-, which represents the first
example of a NiI with a purely aliphatic dithiolate N2S2 environ-
ment, can be explained by the presence of the bipyridine ligand,
which allows for the partial delocalization of the electron density
onto coordinated ligand. In fact, the significant radical character
of 1- was attested to by the small g-anisotropy, and the composi-
tion of the redox active molecular orbital of the complex (64% of
the SOMO is localized on the bipyridine unit). The stability of 1-
can also be related to the higher reduction potential of 1 (about
ꢀ1.25 V vs NHE for the NiII/NiI redox system) with respect to
the other previously reported N2S2 dithiolate NiII complexes
(around ꢀ2.0 V vs NHE).24 Furthermore, the NiII/NiI potential
value of 1 falls near those of the mixed thiolate/thioether NiII
complexes (between ꢀ1.18 and ꢀ1.37 V vs NHE), for which the
corresponding NiI species were generated.50,51
Figure 10. TDDFT assignment of the calculated transitions for
2(CH3CN)2þ in CH2Cl2. The population of the relevant MOs
(HOMO, highest occupied molecular orbital; LUMO, lowest unoccu-
pied molecular orbital) is indicated in parentheses as well as the
wavelength of the optical transitions. Legend: Ni represents the metal,
S the sulfur atoms, and L the rest of the ligand.
unpaired electron in the Ni(dz2)-based MO (Figure 8). The
SOMO of 2(CH3CN)2þ is thus primarily metal-based (67%)
with a non-negligible contribution from the thiolate (20%),
consistent with the EPR data (Table S3). As previously observed,
the conversion of a square planar geometry to a square pyramidal
environment drastically stabilizes the NiIII state.32
Regarding 22þ, this is the first example of a mononuclear NiIII
complex in a mixed thiolate/thioether environment. An earlier study
of Brunold et al.30 reports the investigation of the electronic
structures of NiIII complexes with pure thiolate S4N2 or S2N2
environments by several spectroscopic techniques and DFT calcula-
tions. They have also observed that addition of solvent is required to
get optimized structures for thiolate NiIII species consistent with the
experimental data. Interestingly, we observed that the conversion of
one thiolate by a thioether apparently leads to an increased
participation of the remaining thiolate in the SOMO, and conse-
quently to a decrease of its metal character (20% of S and 67% of Ni
in 2(CH3CN)2þ). Indeed, in the case of the dithiolate N2S2 NiIII
complex studied by Brunold et al.30 and of the dithiolate N3S2
[NiIIIL(im)]þ previously described by us,32 the metal character of
the SOMO is more pronounced (80% and 76%, respectively) and
the thiolate participation is smaller (4% and 15%, respectively).
Finally, we have also shown that complex 1 presents catalytic
reductive dehalogenation activity. Such activity is of great interest
since it can be used for decontaminating soil and water from
chlorinated solvents. Indeed, even if a huge number of methodolo-
gies has been developed for this purpose in recent years, electro-
chemical reductive methods of organic halides represent a promising
approach as they are intrinsically milder, more selective, and easier to
implement that most of the other methods.52 The ability of Ni
complexes to mediate homogeneous reductive dehalogenation of
chlorinated products has already been reported, but, up to now,
mainly by a chemical agent as sodium borohydride.53 The capability
of 1 to catalytically reduce halogenated compounds, via an electro-
chemical method, is therefore promising. However, the mechanism
should be investigated in detail and the efficiency of the complex
evaluated. Is complex 1 active at low concentration of halogenated
derivatives? Is it capable of reducing other organic halides besides
CH2Cl2? Works are in progress in our laboratory to answer these
questions.
The computed EPR parameters for 2(CH3CN)2þ (gx =2.237,
gy = 2.162, gz = 2.070, and AN = 44 MHz) are in good agreement
with the experimental data, thus validating 2(CH3CN)2þ as the
one-electron oxidized species of 2þ. We note in passing that the
EPR parameters calculated for 22þ without any axial ligand
display very small g anisotropy, consistent with a radical based
system (Table S4). This provides further evidence that this
species does not represent the situation present in solution. .
TD-DFT calculations were undertaken on 2(CH3CN)2þ in
order to assign the two main absorptions of its electronic spectra,
i.e., the two bands at 21 350 and 18 870 cmꢀ1. The calculations
nicely reproduce these key features in the spectrum in terms of
energy and intensity (calc: λmax(f) = 21 786 (0.017) and 17 762
(0.012) cmꢀ1, respectively) (Table S7). Regarding the nature of
these absorptions, Figure 10 shows that they all arise mainly from
ligand-to-metal charge transfer (LMCT) transitions since the
acceptor orbital is mainly metal-centered (44% Ni contribution).
The donor orbitals of both absorptions are delocalized π-orbitals
but feature completely different compositions. While the main
contribution to the donor orbital of the 21 786 cmꢀ1 transition
comes from one pyridine ring of the bipyridine unit, the donor
orbital of the 17 762 cmꢀ1 transition is delocalized over the two
phenyl rings of the thioether arm.
’ FINAL REMARKS AND CONCLUSIONS
Given the relevance of the uncommon thiolate NiI and NiIII
species in the catalytic cycle of sulfur rich Ni enzymes, we have
sought to improve our understanding on the stability of such
complexes through the examination of their electronic structure
with different chemical environments, namely pure aliphatic
thiolate versus mixed thiolate/thioether.
3713
dx.doi.org/10.1021/ic200063d |Inorg. Chem. 2011, 50, 3707–3716

Inorganic Chemistry
ARTICLE
’ EXPERIMENTAL SECTION
methods and refined on F2 by full matrix least-squares techniques using
the SHELXTL package.57 All non-hydrogen atoms were refined aniso-
tropically, and hydrogen atoms were placed at their calculated position.
The structure contains four cocrystallized water molecules per two
complexes, 1.5 being on symmetric elements and 2.5 disordered.
Electrochemistry. Electrochemical measurements were carried
out under an argon atmosphere (in glovebox at room temperature or
by argon purging at low temperature). Cyclic voltammetry and con-
trolled potential electrolysis experiments were performed by using an
EG&G model 173 potentiostat/galvanostat equipped with a PAR model
universal programmer and a PAR model 179 digital coulometer. A
standard three-electrode electrochemical cell was used. Potentials were
referred to an Ag/0.01 M AgNO3 reference electrode in CH3CN þ 0.1
M Bu4NClO4. Measured potentials were calibrated through the use of an
internal Fc/Fcþ standard. The working electrode was a vitreous carbon
disk (3 mm in diameter) polished with 2 μm diamond paste (Mecaprex
Presi) for cyclic voltammetry (Epa, anodic peak potential; Epc, cathodic
peak potential; E1/2 = (Epa þ Epc)/2; ΔEp = Epa ꢀ Epc). Exhaustive
electrolyses were carried out on reticulated vitreous carbon electrode 45
PPI (the electrosynthesis Co. Inc.) (1 cm3). The auxiliary electrode was
a Pt wire in DMF, CH3CN, or CH2Cl2 þ 0.1 M Bu4NPF6 solution.
Computational Details. Theoretical calculations were based on
density functional theory (DFT) and have been performed with the
ORCA program package.58 Full geometry optimizations were carried
out for all complexes using the GGA functional BP8659ꢀ61 in combina-
tion with the TZV/P62 basis set for all atoms and by taking advantage of
the resolution of the identity (RI) approximation in the Split-RI-J
variant63 with the appropriate Coulomb fitting sets.64 The structure
optimizations of the NiI and NiIII complexes were both initiated from the
X-ray structure of 1 and 2I, respectively. Increased integration grids
(Grid4 in ORCA convention) and tight SCF convergence criteria were
used. For all molecular property calculations, solvent effects were
accounted for according to the experimental conditions. For that
purpose, we used the DMF (ε = 38.3), CH2Cl2 (ε = 9.08), and
acetonitrile (ε = 36.6) solvents within the framework of the conductor
like screening (COSMO) dielectric continuum approach.65 EPR para-
meters were obtained from additional single-point calculations using the
hybrid functional B3LYP.66,67 For these calculations, we employed three
types of basis sets: (i) Kutzelnigg’s NMR/EPR (IGLO-III) basis set68 for
all ligating N atoms, (ii) Ahlrichs’ valence triple-basis set69 with two sets
of polarization functions (TZVPP)70 for the Ni and S atoms, and (iii)
Ahlrichs’ valence triple-basis set69 with one set of polarization functions
(TZVP)70 for all remaining atoms. Special care was also taken to ensure
accurate results by increasing the size of the integration grid to 7 (ORCA
convention) for the metal center.71 Optical properties were also
obtained from additional single-point calculations using the
B3LYP66,67 functional with the TZVP62 basis set. Electronic transition
energies and dipole moments for all models were calculated using time-
dependent DFT (TD-DFT)72ꢀ74 within the TammꢀDancoff
approximation.75,76 To increase computational efficiency, the RI
approximation77 was used in calculating the Coulomb term and at least
30 excited states were calculated in
General. The ligand L54 and the complex [NiL] (1)32 were prepared
according to reported methods. The synthesis of [NiLMe]I (2I) was
performed under argon (Schlenk techniques). The solvents used for
electrochemical and spectroelectrochemical analyses (acetonitrile, Rath-
burn, HPLC grade, dichloromethane, Carlo Erba reactifs-SDS, anhy-
drous, and DMF, Sigma Aldrich, anhydrous, 99.8%) and the electrolyte
tetrabutylammonium hexafluorophosphate (Bu4NPF6) were used as
received and stored in a glovebox.
Synthesis of [NiLMe]I (2I). Iodomethane (24 μL, 0.385 mmol) was
added to a suspension of 1 (200 mg, 0.314 mmol) in dichloromethane
(30 mL). The mixture was stirred overnight, giving a brown solution. The
solvent was removed under vacuum, and the product was extracted from the
residual solid with an excess of acetonitrile. The solvent was removed in
vacuo, and THF (5 mL) was added to the residue. After a few minutes of
stirring, a brown precipitate was formed. It was filtered, washed with THF
(3 mL), and redissolved in a minimum volume of dichloromethane. Diethyl
ether was allowed to slowly diffuse into this solution to afford a brown-orange
microcrystalline solid after 24 h. This was filtered, washed with diethyl ether,
and dried (2I, 181 mg, 0.232 mmol, 74%). IR (cmꢀ1): 3053 m, 2922w,
1956w, 1886w, 1808w, 1603s, 1569 m, 1489s, 1443s, 1265w, 1186w, 1126w,
1034w, 980 m, 748s, 698vs, 602 m. 1H NMR (400 MHz, CD3CN): δ 3.56
(s, 2H, CH2), 3.77 (s, 2H, CH2), 4.64 (s br, 3H, CH3), 7.10 (t, J = 7.1 Hz,
2H, CH Ph), 7.28 (m, 6H, CH Ph), 7.39 (d,J =7.4 Hz, 4H, CHPh), 7.44 (t,
J = 7.0 Hz, 4H, CH Ph), 7.58 (d, J = 7.2 Hz, 4H, CH Ph), 8.00 (t, J = 7.4 Hz,
1H, CH bpy), 8.07 (t, J = 7.7 Hz, 1H, CH bpy), 8.25 (br, 1H, CH bpy), 8.34
(br, 1H, CH bpy), 9.31 (br, 1H, CH bpy), 10.16 (br, 1H, CH bpy). ESI-MS
(CH3CN, m/z, I%): 651.2, 100 2þ. Anal. Calcd for C39H33N2S2NiI H2O
(797.43): C, 58.74; H, 4.42; N, 3.51. Found: C, 58.73; H, 4.24; N, 3.33.
3
Brown single crystals of 2I 2H2O suitable for X-ray diffraction were obtained
3
by slow evaporation of a THF/CH2Cl2 (1/1) solution of the product.
Synthesis of [NiLMe](CF3SO3) (2OTf). 2OTf was obtained
quantitatively by reacting 2I with an equimolar amount of Ag(CF3SO3)
in CH3CN. The AgI precipitate that formed was filtered off, and the
resulting solution was evaporated to dryness to give a brown powder
corresponding to 2OTf (yield: 95%). 1H NMR (400 MHz, CD3CN): δ
2.94 (s br, 3H, CH3), 4.00 (s, 2H, CH2), 4.11 (s, 2H, CH2), 7.16 (t, J =
7.3 Hz, 2H, CH Ph), 7.27 (t, J = 7.6 Hz, 4H, CH Ph), 7.32 (t, J = 7.4 Hz,
2H, CH Ph), 7.42 (m, 5H, 4 CH Ph þ 1 CH bpy), 7.49 (m, 5H, 4 CH Ph
þ 1 CH bpy), 7.57 (d, J = 7.8 Hz, 4H, CH Ph), 7.80 (t, J = 7.9 Hz, 1H,
CH bpy), 7.91 (t, J = 7.9 Hz, 1H, CH bpy), 8.32 (d br, 1H, CH bpy), 8.47
(br, 1H, CH bpy).
Physical Measurements. The infrared spectra were recorded on a
Magna-IR TM 550 Nicolet spectrometer as KBr pellet. 1H NMR spectra
were recorded on a Bruker Avance 400 MHz spectrometer. Resonance
assignments were achieved from the analysis of 1Hꢀ H COSY experi-
1
ments. Electronic absorption spectra were recorded on a Varian Cary
100 absorption spectrophotometer in quartz cells. X-band EPR spectra
were recorded on a Bruker EMX, equipped with the ER-4192 ST Bruker
cavity and an ER-4131 VT at 100 K. The elemental analyses were carried
out with a C, H, N analyzer (SCA, CNRS). ESI-MS experiments were
performed on a Bruker Esquire 3000 Plus ion trap spectrometer
equipped with an electrospray ion source (ESI). The simulations of
the EPR spectra were performed by the EasySpin software.55 The sample
was analyzed in positive ionization mode by direct perfusion in the ESI-
MS interface (ESI capillary voltage 2 kV, sampling cone voltage 40 V).
X-ray Crystallography. A summary of data collection and struc-
ture refinement for 2I is reported in Table S1. Single-crystal diffraction
data were taken using an Oxford-Diffraction Xcallibur S Kappa geometry
diffractometer (Mo K radiation, graphite monochromator, λ 0.710 73 Å)
at 150 K. An absorption correction was applied using the ABSPACK
Oxford-diffraction program56 with transmission factors in the range
0.623ꢀ0.940. The molecular structure of 2I was solved by direct
’ ASSOCIATED CONTENT
Supporting Information. Additional structural (2þ),
S
b
electrochemical (1 and 2þ), and theoretical data.This material
is available free of charge via the Internet at http://pubs.acs.org.
’ AUTHOR INFORMATION
Corresponding Author
*E-mail: marie-noelle.collomb@ujf-grenoble.fr (M.-N.C.),
carole.duboc@ujf-grenoble.fr (C.D.).
3714
dx.doi.org/10.1021/ic200063d |Inorg. Chem. 2011, 50, 3707–3716

Inorganic Chemistry
ARTICLE
Present Addresses
(25) Grapperhaus, C. A.; Mullins, C. S.; Mashuta, M. S. Inorg. Chim.
Acta 2005, 358, 623–632.
(26) Musie, G.; Reibenspies, J. H.; Darensbourg, M. Y. Inorg. Chem.
1998, 37, 302–310.
(27) Hanss, J.; Kruger, H. J. Angew. Chem., Int. Ed. 1998,
37, 360–363.
^Laboratoire de Spectrochimie Infrarouge et Raman, CNRS
LASIR-UMR 8516, Bat C4 F-59655 Villeneuve d’Ascq Cedex,
France.
(28) Mathrubootham, V.; Thomas, J.; Staples, R.; McCraken, J.;
Shearer, J.; Hegg, E. L. Inorg. Chem. 2010, 49, 5393–5406.
(29) Kruger, H. J.; Peng, G.; Holm, R. H. Inorg. Chem. 1991,
30, 734–742.
(30) Fiedler, A. T.; Brunold, T. C. Inorg. Chem. 2007, 46,
8511–8523.
(31) Fiedler, A. T.; Bryngelson, P. A.; Maroney, M. J.; Brunold, T. C.
J. Am. Chem. Soc. 2005, 127, 5449–5462.
(32) Gennari, M.; Orio, M.; Pꢀecaut, J.; Neese, F.; Collomb, M.-N.;
’ ACKNOWLEDGMENT
M.G. thanks the CNRS for his postdoctoral fellowship. Frank
Reikowski and Eckhard Bill are acknowledged for EPR experi-
ments. F.N. gratefully acknowledges the University of Bonn and
the Max-Planck society (via a Max-Planck fellowship) for finan-
cial support of this work. Furthermore, the special research unit
SFB 813 (Chemistry at Spin-Centers, University of Bonn) is
gratefully acknowledged.
Duboc, C. Inorg. Chem. 2010, 49, 6399–6401.
(33) Green, K. N.; Brothers, S. M.; Jenkins, R. M.; Carson, C. E.;
Grapperhaus, C. A.; Darensbourg, M. Y. Inorg. Chem. 2007,
46, 7536–7544.
(34) Nakane, D.; Kuwasako, S.-i.; Tsuge, M.; Kubo, M.; Funahashi,
Y.; Ozawa, T.; Ogura, T.; Masuda, H. Chem. Commun. 2010,
46, 2142–2144.
(35) Rosen, W.; Busch, D. H. J. Am. Chem. Soc. 1969, 91, 4694–
4697.
(36) Ito, M.; Matsumoto, T.; Tatsumi, K. Inorg. Chem. 2009,
48, 2215–2223.
’ REFERENCES
(1) Ragsdale, S. W. J. Biol. Chem. 2009, 284, 18571–18575.
(2) Miller, A. F. Curr. Opin. Chem. Biol. 2004, 8, 162–168.
(3) Bryngelson, P. A.; Maroney, M. J. In Nickel and Its Surprising
Impact in Nature; Sigel, A., Sigel, H., Sigel, R. K. O., Eds.; John Wiley &
Sons Ltd.: West Sussex, U.K., 2007; pp 417ꢀ444.
(4) Barondeau, D. P.; Kassmann, C. J.; Bruns, C. K.; Tainer, J. A.;
Getzoff, E. D. Biochemistry 2004, 43, 8038–8047.
(5) Wuerges, J.; Lee, J.-W.; Yim, Y.-I.; Yim, H.-S.; Kang, S.-O.;
Djinovic Carugo, K. Proc. Natl. Acad. Sci. U.S.A. 2004, 101, 8569–8574.
(6) Lubitz, W.; Gastel, M. V.; G€artner, W. In Nickel and Its Surprising
Impact in Nature; Sigel, A., Sigel, H., Sigel, R. K. O., Eds.; John Wiley &
Sons Ltd.: West Sussex, U.K., 2007; pp 279ꢀ322.
(37) Baidya, N.; Olmstead, M.; Mascharak, P. K. Inorg. Chem. 1991,
30, 929–937.
(38) Baidya, N.; Olmstead, M. M.; Mascharak, P. K. J. Am. Chem. Soc.
1992, 114, 9666–9668.
(39) Baidya, N.; Olmstead, M. M.; Whitehead, J. P.; Bagyinka, C.;
Maroney, M. J.; Mascharak, P. K. Inorg. Chem. 1992, 31, 3612–3619.
(40) Gomez, L.; Pereira, E.; de Castro, B. J. Chem. Soc., Dalton Trans.
2000, 1373–1379.
(7) Vignais, P. M.; Billoud, B. Chem. Rev. 2007, 107, 4206–4272.
(8) Garcin, E.; Vernede, X.; Hatchikian, E. C.; Volbeda, A.; Frey, M.;
Fontecilla-Camps, J. C. Structure 1999, 7, 557–566.
(9) Ragsdale, S. W.; Pierce, E. Biochim. Biophys. Acta 2008,
1784, 1873–1898.
(41) In contrast to 2þ, the neutral 1 complex has been studied in
DMF since it is insoluble in CH3CN.
(10) Dobbek, H.; Svetlitchnyi, V.; Gremer, L.; Huber, R.; Meyer, O.
(42) Burke, J. A.; Brink, E. C. Inorg. Chem. 1969, 8, 386–387.
(43) Shoup, J. C.; Burke, J. A. Inorg. Chem. 1973, 12, 1851–1855.
(44) Ashby, M. T.; Enemark, J. H.; Lichtenberger, D. L. Inorg. Chem.
1988, 27, 191–197.
(45) Maroney, M. J.; Choudhury, S. B.; Bryngelson, P. A.; Mirza,
S. A.; Sherrod, M. J. Inorg. Chem. 1996, 35, 1073–1076.
(46) In this manuscript, all redox potential are referred to Fc/Fcþ.
(47) Lovecchio, F. V.; Gore, E. S.; Busch, D. H. J. Am. Chem. Soc.
1974, 96, 3109–3118.
(48) Gagne, R. R.; Ingle, D. M. Inorg. Chem. 1981, 20, 420–425.
(49) Lappin, A. G.; McAuley, A. Adv. Inorg. Chem. 1988,
32, 241–295.
(50) Ge, P.; Riordan, C. G.; Yap, G. P. A.; Rheingold, A. L. Inorg.
Chem. 1996, 35, 5408–5409.
(51) Yamamura, T.; Sakurai, S.; Arai, H.; Miyamae, H. J. Chem. Soc.,
Chem. Commun. 1993, 1656–1658.
(52) Laine, D. F.; Cheng, I. F. Microchem. J. 2007, 85, 183–193.
(53) Alonso, F.; Beletskaya, I. P.; Yus, M. Chem. Rev. 2002,
102, 4009–4092.
(54) Kopf, M. A.; Varech, D.; Tuchagues, J. P.; Mansuy, D.; Artaud, I.
J. Chem. Soc., Dalton Trans. 1998, 991–998.
(55) Stoll, S.; Schweiger, A. J. Magn. Reson. 2006, 178, 42–55.
(56) ABSPACK CrysAlis PRO RED; Oxford Diffraction Ltd: Abing-
don, Oxfordshire, U.K., 2010.
Science 2001, 293, 1281–1285.
(11) Drennan, C. L.; Heo, J. Y.; Sintchak, M. D.; Schreiter, E.;
Ludden, P. W. Proc. Natl. Acad. Sci. U.S.A. 2001, 98, 11973–11978.
(12) Doukov, T. I.; Blasiak, L. C.; Seravalli, J.; Ragsdale, S. W.;
Drennan, C. L. Biochemistry 2008, 47, 3474–3483.
(13) Doukov, T. I.; Iverson, T. M.; Seravalli, J.; Ragsdale, S. W.;
Drennan, C. L. Science 2002, 298, 567–572.
(14) Darnault, C.; Volbeda, A.; Kim, E. J.; Legrand, P.; Vernede, X.;
Lindahl, P. A.; Fontecilla-Camps, J. C. Nat. Struct. Biol. 2003, 10, 271–
279.
(15) Bouwman, E.; Reedijk, J. Coord. Chem. Rev. 2005, 249,
1555–1581.
(16) Evans, D. J. Coord. Chem. Rev. 2005, 249, 1582–1595.
(17) Harrop, T. C.; Mascharak, P. K. Coord. Chem. Rev. 2005,
249, 3007–3024.
(18) Grapperhaus, C. A.; Darensbourg, M. Y. Acc. Chem. Res. 1998,
31, 451–459.
(19) Mullins, C. S.; Grapperhaus, C. A.; Frye, B. C.; Wood, L. H.;
Hay, A. J.; Buchanan, R. M.; Mashuta, M. S. Inorg. Chem. 2009,
48, 9974–9976.
(20) Dey, A.; Green, K. N.; Jenkins, R. M.; Jeffrey, S. P.; Darens-
bourg, M.; Hodgson, K. O.; Hedman, B.; Solomon, E. I. Inorg. Chem.
2007, 46, 9655–9660.
(21) Grapperhaus, C. A.; Mullins, C. S.; Kozlowski, P. M.; Mashuta,
(57) Sheldrick, G. M. SHELXTL-Plus, Structure Determination Soft-
ware Programs (Version 5.1.); Madison, WI, 1998.
(58) Neese, F. ORCA—An Ab Initio, Density Functional and Semi-
empirical Program Package (v. 2.6ꢀ35); Universit€at Bonn: Bonn,
Germany, 2007.
(59) Perdew, J. P. Phys. Rev. B 1986, 33, 8822–8824.
(60) Perdew, J. P. Phys. Rev. B 1986, 34, 7406–7406.
(61) Becke, A. D. Phys. Rev. A 1988, 38, 3098–3100.
M. S. Inorg. Chem. 2004, 43, 2859–2866.
(22) Farmer, P. J.; Reibenspies, J. H.; Lindahl, P. A.; Darensbourg,
M. Y. J. Am. Chem. Soc. 1993, 115, 4665–4674.
(23) Smee, J. J.; Miller, M. L.; Grapperhaus, C. A.; Reibenspies, J. H.;
Darensbourg, M. Y. Inorg. Chem. 2001, 40, 3601–3605.
(24) Bellefeuille, J. A.; Grapperhaus, C. A.; Derecskei-Kovacs, A.;
Reibenspies, J. H.; Darensbourg, M. Y. Inorg. Chim. Acta 2000,
300ꢀ302, 73–81.
3715
dx.doi.org/10.1021/ic200063d |Inorg. Chem. 2011, 50, 3707–3716

Inorganic Chemistry
ARTICLE
(62) Sch€afer, A.; Huber, C.; Ahlrichs, R. J. Chem. Phys. 1994,
100, 5829–5835.
(63) Neese, F. J. Comput. Chem. 2003, 24, 1740–1747.
(64) Weigend, F. Phys. Chem. Chem. Phys. 2006, 8, 1057–1065.
(65) Klamt, A.; Sch€urmann, G. J. Chem. Soc., Perkin Trans. 2
1993, 799–805.
(66) Becke, A. D. J. Chem. Phys. 1993, 98, 5648–5652.
(67) Lee, C. T.; Yang, W. T.; Parr, R. G. Phys. Rev. B 1988,
37, 785–789.
(68) Kutzelnigg, W.; Fleischer, U.; Schindler, M. In The IGLO
Method: Ab Initio Calculation and Interpretation of NMR Chemical Shifts
and Magnetic Susceptibilities; Springer-Verlag: Heidelberg, 1990; Vol. 23.
(69) Sch€afer, A.; Horn, H.; Ahlrichs, R. J. Chem. Phys. 1992,
2571–2577.
(70) Ahlrichs, R. Unpublished results.
(71) Neese, F. Inorg. Chim. Acta 2002, 337, 181–192.
(72) Casida, M. E. In Recent Advances in Density Functional Methods;
Chong, D. P., Ed.; World Scientific: Singapore, 1995.
(73) Stratmann, R. E.; Scuseria, G. E.; Frisch, M. J. J. Chem. Phys.
1998, 109, 8218–8224.
(74) Bauernschmitt, R.; Ahlrichs, R. Chem. Phys. Lett. 1996, 454
464.
(75) Hirata, S.; Head-Gordon, M. Chem. Phys. Lett. 1999, 314, 291–
299.
(76) Hirata, S.; Head-Gordon, M. Chem. Phys. Lett. 1999, 302, 375–
382.
(77) Neese, F.; Olbrich, G. Chem. Phys. Lett. 2002, 170–178.
3716
dx.doi.org/10.1021/ic200063d |Inorg. Chem. 2011, 50, 3707–3716