Iron(II) cage complexes of N-heterocyclic amide and bis(trimethylsilyl) amide ligands: Synthesis, structure, and magnetic properties

Article abstract of DOI:10.1021/ic102341a

Metallation of hexahydropyrimidopyrimidine (hppH) by [Fe{N(SiMe ₃)₂}₂] (1) produces the trimetallic iron(II) amide cage complex [{(Me₃Si)₂NFe}₂(hpp) ₄Fe] (2), which contains three iron(II) centers, each of which resides in a distorted tetrahedral environment. An alternative, one-pot route that avoids use of the highly air-sensitive complex 1 is described for the synthesis of the iron(II)-lithium complex [{(Me₃Si) ₂N}₂Fe{Li(bta)}]₂ (3) (where btaH = benzotriazole), in which both iron(II) centers reside in 3-coordinated pyramidal environments. The structure of 3 is also interpreted in terms of the ring laddering principle developed for alkali metal amides. Magnetic susceptibility measurements reveal that both compounds display very weak antiferromagnetic exchange between the iron(II) centers, and that the iron(II) centers in 2 and 3 possess large negative axial zero-field splittings.

Full text of DOI:10.1021/ic102341a

ARTICLE
pubs.acs.org/IC
Iron(II) Cage Complexes of N-Heterocyclic Amide and
Bis(trimethylsilyl)amide Ligands: Synthesis, Structure, and Magnetic
Properties
Scott A. Sulway, David Collison, Joseph J. W. McDouall, Floriana Tuna, and Richard A. Layfield*
School of Chemistry, The University of Manchester, Oxford Road, Manchester M13 9PL, U.K.
S Supporting Information
b
ABSTRACT: Metallation of hexahydropyrimidopyrimidine (hppH) by [Fe{N(SiMe3)2}2] (1) produces the trimetallic iron(II)
amide cage complex [{(Me3Si)2NFe}2(hpp)4Fe] (2), which contains three iron(II) centers, each of which resides in a distorted
tetrahedral environment. An alternative, one-pot route that avoids use of the highly air-sensitive complex 1 is described for the
synthesis of the iron(II)-lithium complex [{(Me3Si)2N}2Fe{Li(bta)}]2 (3) (where btaH = benzotriazole), in which both iron(II)
centers reside in 3-coordinated pyramidal environments. The structure of 3 is also interpreted in terms of the ring laddering principle
developed for alkali metal amides. Magnetic susceptibility measurements reveal that both compounds display very weak
antiferromagnetic exchange between the iron(II) centers, and that the iron(II) centers in 2 and 3 possess large negative axial
zero-field splittings.
’ INTRODUCTION
We now report the trimetallic iron(II) amide cage complex
[{(Me3Si)2NFe}2(hpp)4Fe] (2) and the tetrametallic iron(II)-
lithium complex [{(Me3Si)2N}2Fe{Li(bta)}]2 (3) (where
The iron(II) amide [Fe{N(SiMe3)2}2] (1)1 is a landmark
compound that has stimulated extensive investigations into low-
coordinate iron chemistry.2,3 In addition to the fundamental
interest in the coordination chemistry of iron complexes with
bulky amide ligands, the reactions of 1 with a broad range of E-
H acidic substrates have enabled systematic studies of low-
coordinate iron complexes containing nitrogen,4-10 oxygen,11,12
phosphorus,13 sulfur,9,10,14-22 or selenium23,24 donor ligands.
Chaudret et al. have recently shown that 1 can also be thermally
decomposed in an atmosphere of dihydrogen in the presence of
long-chain acids and long-chain amines to produce monodis-
perse iron nanoparticles.25,26
The reactions of 1 with thiols and elemental sulfur have
attracted particular interest, because of the range of structurally
diverse iron-sulfur cage compounds that can be accessed via this
synthetic route.17,20,22 However, synthetic routes to polymetallic
iron cage compounds in which the iron centers are bridged by
amido nitrogens have hitherto not been reported. Previous
studies on the synthesis of manganese(II) amide cages from
the deprotonation of functionalized aromatic primary amines
by manganocene27 suggested that iron(II) amide cage com-
pounds should be accessible using 1 as the precursor. Thus,
our interest in 1 stems from the Brønsted basic reactivity of the
[(Me3Si)2N]- ligands toward heterocycles that contain acidic
N-H groups, and the potential for this reactivity to be applied in
the synthesis of iron(II) cages. Furthermore, we recently re-
ported that deprotonation of the N-heterocycle benzotriazole
(btaH) by tris(cyclopentadienyl)dysprosium forms the dimeric
complex [Cp2Dy(μ-bta)]2, a complex which shows single-mo-
lecule magnet (SMM) behavior,28,29 hence given the current
interest in SMMs based on 3d transition metals, it is conceivable
that the chemistry of 1 could be developed to provide an
alternative synthetic method in the SMM field.
[hpp]-
=
hexahydropyrimidopyrimidide, [bta]-
=
benzotriazolide). Compounds 2 and 3 were characterized by
X-ray crystallography, SQUID magnetometry, and NMR spec-
troscopy, and DFT (density functional theory) calculations were
carried out on compound 2.
’ RESULTS AND DISCUSSION
Synthesis and Structural Characterization. The reaction of
1 with (hpp)H produced a brown solution, which, after being
concentrated and stored at -30 ꢀC, yielded yellow-green crystal-
line needles. X-ray crystallography revealed the product to be the
trimetallic iron(II) cage complex 2 (see Scheme 1 and Figure 1).
In the structure of 2, the central Fe(2) atom resides in a
distorted tetrahedral environment through complexation by four
N-donor atoms of the [hpp]- ligands. Of the four N-donor
Scheme 1. Synthesis of Compound 2
Received: November 23, 2010
Published: February 11, 2011
r
2011 American Chemical Society
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Inorganic Chemistry
ARTICLE
Scheme 2. Synthesis of Compound 3
type ligands related to
2 are [{hpp(H)}2FeCl2] and
[{H2C(hpp)}2FeCl2], each of which contains an iron(II) center
in a tetrahedral environment, with the hpp ligands coordinated
through imino nitrogens.32 The [hpp]- ligand is well-known in
transition-metal chemistry for its ability to encourage the forma-
tion of metal-metal bonds, with compounds having the general
formula of [M2(hpp)4] being particularly prevalent.33 A search of
the Cambridge Structural Database reveals that more than half of
the transition-metal complexes of the [hpp]- contain a metal-
metal bond.34
Figure 1. Molecular structure of compound 2. Hydrogen atoms are
omitted for the sake of clarity.
However, the Fe Fe separation of 3.127 Å in 2 is too large
3 3 3
for an Fe-Fe bond to be feasible. Indeed, a Mayer bond order
analysis confirms that the iron centers in 2 only bond to nitrogen
(see Figure 2), while the Fe Fe interactions show bond orders
3 3 3
of only 0.06.
A drawback with 1 is its extremely high sensitivity to air. We
have found that an alternative, “one-pot” approach has potential
to be developed as a general synthetic route to iron(II) cages.
With this method, we combined the N-H acidic heterocycle
benzotriazole (bta)H, which is known to encourage cage
formation,35 with solid [FeBr2(thf)2] and the somewhat less
air-sensitive lithium bis(trimethylsilyl)amide (although care was
still taken to exclude oxygen and moisture from the reaction).
The solids were cooled to -78 ꢀC and then a toluene was added.
The reaction developed a green color upon warming to room
temperature, which then became brown and was accompanied by
the formation of a precipitate. Filtering the reaction and storing
the concentrated solution at -30 ꢀC produced green crystals,
which X-ray crystallography revealed to be [{(Me3Si)2N}2-
Fe{Li(bta)}]2 (3) (see Scheme 2 and Figure 3).
Figure 2. Mayer bond order (B3LYP/Def2-SVP) analysis for com-
pound 2.
atoms coordinated to Fe(2), N(5) and N(5A) engage in end-on
bonding, with the Fe(2)-N(5/5A) distance being 2.0409(19)
Å, whereas N(3) and N(3A) also μ-bridge to Fe(1) and Fe(1A)
to result in Fe(2)-N(3/3A) and Fe(1)-N(3/3A) distances of
2.1064(18) and 2.3066(18) Å, respectively. The coordination
geometry of Fe(1) is also a distorted tetrahedron, consisting of a
κ2-[hpp]- ligand that coordinates through N(2) and N(3), a
κ1-[hpp]- ligand that coordinates through N(4), and a terminal
[(Me3Si)2N]- ligand that coordinates to Fe(1) through N(7).
The resulting Fe-N distances are as follows: Fe(1)-N(2),
2.0645(19) Å; Fe(1)-N(4), 1.9901(19) Å; and Fe(1)-N(7),
1.9880(19) Å. The distorted nature of the tetrahedral iron
coordination environments in 2 is apparent from the range of
N-Fe-N bond angles. For Fe(1), the N-Fe-N bond angles
are 61.95(7)ꢀ and 113.98(7)-122.69(8) Å; for Fe(2), the angles
The molecular structure of 3 is a centrosymmetric dimer that
consists of a central {Li(bta)}2 core capped on each end by one
[Fe{N(SiMe3)2}2] unit. Combining the two {Li(bta)} units in
the center of 3 generates a 6-membered Li2N4 ring with the
Li(1)-N(2) and Li(1)-N(3A) distances being 2.044(5) and
1.992(5) Å, respectively. A μ-{N(SiMe3)2} ligand links the
{Li(bta)} core to the iron amide capping unit to result in an
Li(1)-N(4) bond distance of 2.109(5) Å, and an Li(1)-N(4)-
Fe(1) angle of 93.73(15)ꢀ. The Fe(1) atom in 3 resides in a
distorted pyramidal environment, with the Fe(1)-N(1), Fe(1)-
N(4), and Fe(1)-N(5) distances being 2.052(2), 1.9764(19),
and 1.909(2) Å, respectively, and the N(1)-Fe-N(4), N(4)-
Fe-N(5) and N(5)-Fe-N(1) angles being 104.50(8)ꢀ,
141.31(9)ꢀ, and 112.55(9)ꢀ. The Fe(1/1A) centers are posi-
tioned 0.13 Å out of the plane of their three respective nitrogen
are in the range of 101.58(7)ꢀ-121.72(7)ꢀ. The Fe(1) Fe(2)
3 3 3
separation is 3.127 Å and the Fe(1)-Fe(2)-Fe(1A) angle is
153.86(1)ꢀ, which indicates a “bending” of the molecule and
leads to C2 point symmetry. The 1H NMR spectrum of 2 shows a
sharp singlet at 2.39 ppm corresponding to the SiMe3 substitu-
ents, and resonances at 7.24 and 7.20 ppm due to the hydrogens
of the [hpp]- ligands, suggesting that the solid-state structure is
preserved in benzene solution.
1
Although the coordination chemistry of guanidinate ligands is
well established,30 surprisingly few iron complexes of these
ligands are known,31 and complex 2 is the first iron complex of
the anionic [hpp]- ligand. Two iron complexes of neutral hpp-
donor atoms. A sharp singlet at 0.11 ppm in the H NMR
spectrum of 3 can be assigned to the SiMe3 substituents, but
the resonances due to the [bta]- protons are less distinct, with a
weak, broad signal being observed at -11.2 ppm (ω1/2 = 380Hz).
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ARTICLE
Scheme 3
Figure 3. Molecular structure of compound 3. Hydrogen atoms are
omitted for the sake of clarity.
Alkali metal amido ferrate complexes have recently been
shown to display interesting applications in synthesis,36 although
crystallographically authenticated structures such as that of 3 are
scarce.36b The {Li(bta)}2 core of 3 is isostructural with the
{Li(bta)}2 units found in the polymeric ladder structure of
[(dmso)Li(bta)]¥ (4) (where dmso = dimethylsulfoxide) (see
Scheme 3).37 While crystalline 4 could only be obtained by
dissolving [Li(bta) thf] in hot dmso, the SiMe3 substituents in 3
3
result in good solubility in toluene at room temperature. Because
the polymeric structure of [(dmso)Li(bta)]¥ can be interpreted
in terms of the ring laddering principle developed for alkali-metal
amides,38 the central {Li(bta)}2 core in 3 can be regarded as an
intercepted dimeric unit of the [Li(bta)]¥ ladder. However,
whereas structural studies on intercepted alkali-metal amide
ladders normally deploy tertiary amines as the interceptors, in
3, the intercepting units are molecules of [Fe{N(SiMe3)2}2].
This interpretation gives insight into the mechanism through
which 3 is formed. The first component to form is probably
[Fe{N(SiMe3)2}2], through transmetallation of [FeBr2(thf)2]
by LiN(SiMe3)2, which also explains the initial green color of the
reaction mixture. The second step is the deprotonation of btaH
by the remaining LiN(SiMe3)2 to generate [Li(bta)], which then
dimerizes but cannot polymerize, because of the presence of the
[Fe{N(SiMe3)2}2] traps, leading to a third step in which
molecules of 3 form (see Scheme 3). The formation of 3 could
be of general significance, because it may be possible to trap units
of other alkali-metal amide ladders with other low-coordinate
metal silyl-amides.
Magnetochemistry and DFT Calculations. Magnetic sus-
ceptibility measurements were carried out on polycrystalline
samples of compounds 2 and 3 in an applied field of 0.1 T
(see Figure 4). For compound 2, the room-temperature χMT
value of 9.6 cm3 K mol-1 corresponds to three uncoupled
tetrahedral Fe(II) d6 centers in 2. Upon cooling, χMT decreases
smoothly to 8.94 cm3 K mol-1 at 50 K and then more rapidly to
reach 3.87 cm3 K mol-1 at 2 K. In addition, the temperature
dependence of the susceptibility of 2 shows a continuous increase
between 300 K (χM = 0.031 cm3 mol-1) and 2 K (χM = 0.193
cm3 mol-1), without showing a maximum above 1.8 K, suggest-
ing a very weak antiferromagnetic exchange between Fe(1/1A)
and Fe(2). Fitting the susceptibility data with the Curie-Weiss
law, χM = C/(T - θ), gave values of C = 9.6 cm3 K mol-1 and θ =
-3.4 K (or -2.4 cm-1), where C and θ represent the Curie
constant and the Weiss temperature, respectively. The value of C
Figure 4. Plots of χMT vs temperature and of magnetization versus field
at 2 K (inset) for 2 and 3.
is in agreement with the presence of three tetrahedral Fe(II) ions
with S = 2, and the negative θ value indicates some antiferro-
magnetic exchange in 2.
The magnetization of 2 increases with increasing field at 2 K to
reach 7.28 NμB at 70 kG, but does not saturate (see Figure 4,
inset), and the data could only be modeled by taking into account
axial (D) and rhombic (E) zero-field splitting (ZFS) terms,
according to the following equation (eq 1):39
ꢀ
ꢁ
X
X
ꢂ
1
H ¼ - 2J
SiSj þ D
Siz2 - SiðSi þ 1Þ
3
i < j
X
ꢃ
2
þ E
Six2 - Siy
ð1Þ
Applying this equation to the susceptibility and magnetization
data for 2 produced values of J = -0.08 cm-1, D = -10.2 cm-1
,
|E| = 0.06 cm-1, and g = 2.05. The red traces in Figure 4 represent
the data as modeled with MAGPACK,39 using these parameters.
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ARTICLE
’ CONCLUSION
In summary, [Fe{N(SiMe3)2}2] (1) has been used as a
precursor to the trimetallic iron(II) amide cage compound
[{(Me3Si)2NFe}2(hpp)4Fe] (2). An alternative, one-pot syn-
thetic route has been used to synthesize the iron(II)-lithium
complex [{(Me3Si)2N}2Fe{Li(bta)}]2 (3). Despite the differing
iron coordination geometries in 2 and 3, their ZFS values were
found to be similar, which is an observation that was rationalized
in terms of their respective Fe(II) centers possessing similar
d-orbital structures. Complexes 2 and 3 contain one and two
potentially reactive [N(SiMe3)2]- ligands, respectively, which
suggests that they themselves could be developed as precursors
to iron(II) amide cages of even higher nuclearity. Our ongoing
research will develop this chemistry.
Figure 5. B3LYP/Def2-SVP spin density plot for compound 2 (isosur-
face value = 0.004 a.u.).
’ EXPERIMENTAL SECTION
Synthetic Methods, Physical Measurements, and Calcula-
tions: General Considerations. All reactions were performed with
rigorous exclusion of oxygen and moisture, using standard Schlenk
techniques. [Fe{N(SiMe3)2}2] (1) was prepared according to a litera-
ture procedure.1 Solvents were dried and degassed using an Innovative
Technologies Solvent Purification System. 1H NMR spectra were
recorded on a Bruker Avance III 400 MHz spectrometer, and UV/vis
spectra were recorded on a Shimadzu UV-2401 instrument. Elemental
analysis results were obtained using the elemental analysis service of
London Metropolitan University.
[{(Me3Si)2NFe}2(hpp)4Fe] (2). A solution of [Fe{N(SiMe3)2}2] (1)
(0.77 g, 2.0 mmol) in toluene (10 mL) was added to a solution of
hexahydropyrimidopyrimidine (0.28 g, 2.0 mmol) in toluene (15 mL)
at -78 ꢀC. The reaction mixture was stirred for 30 min before being
slowly warmed to room temperature. After stirring for 4 h, the resulting
deep brown solution was filtered (Celite, porosity 3) and the solvent
reduced in volume to ∼10 mL. Storage of the solution at -30 ꢀC
overnight resulted in the formation of air-sensitive yellow-green
crystals of 2 (0.57 g, 28%). Analysis: calcd: C 46.15, H 8.14, N
18.85%; found: C 46.16, H 8.00, N 18.64%. 1H NMR (C6D6, 400.23
MHz, δ/ppm): 7.24 and 7.20 ppm, 48 H, hpp; 2.39 ppm, 36 H, SiMe3.
UV/vis (toluene), λmax (nm) (ε (dm3 M-1 cm-1)): 438 (3450), 298
(9800).
The large value of D implies considerable orbital angular
momentum within the d6 configurations of the Fe(II) centers
in 2, and the nonzero value of E likely reflects the distorted
tetrahedral iron coordination geometries. The negative sign of J,
the exchange coupling constant, is consistent with antiferromag-
netic coupling between Fe(1/1A) and Fe(2). However, this
exchange interaction is small, as indicated by the small value of J.
A DFT study of 2 determined that Fe(1/1A) and Fe(2) carry
Mulliken spin densities of þ3.72 and -3.76, respectively, with
the remaining spin density being distributed over the other
atoms. Thus, each Fe(II) center in 2 carries four spins (S = 2),
and the spin density plot (Figure 5) reveals the Fe(1/1A) to be
spin up and Fe(2) to be spin down.
Of the various spin states available to 2, DFT clearly predicts
the antiferromagnetically coupled S = 2 state to be the most
stable, in agreement with the experimental susceptibility data.
The plots of χMT versus temperature and of magnetization
versus field for 3 shows a decrease in χMT from 7.2 cm3 K
mol-1 at 300 K to 4.2 cm3 K mol-1 at 2 K, and the isothermal M
versus H plot at 2 K reaches 5.3 NμB at 70 kG and shows no
saturation. Simulation of the susceptibility and magnetization
data using the spin Hamiltonian described in eq 1 gives J = -0.03
cm-1, D = -10.5 cm-1, E = 0, and g = 2.2, with the large value of
the axial ZFS once again indicating the presence of orbital angular
momentum.
It is noteworthy that compounds 2 and 3 possess very similar,
large ZFS values, despite the differing iron(II) coordination
geometries in each complex. Additional insight into this observa-
tion was not possible with Q-band EPR spectroscopy, because
the spectra of each compound consist only of weak resonances,
although the absence of a resonance at g = 2 is consistent with a
large ZFS. However, it has been shown that the large orbital
contribution to the angular momentum in three coordinate,
distorted trigonal-planar, high-spin iron(II) complexes arises
from the effect of spin-orbit coupling in the magnetic field
between the low-lying, but nondegenerate dz2 (doubly occupied)
and dyz (singly occupied) orbitals,40 which we propose to explain
the large magnitude of the ZFS in 3. Analysis of the d-orbital
structure in 2 does in fact produce a similar picture to that found
in 3-coordinated iron complexes (see Table S1 in the Supporting
Information), with one essentially doubly occupied d orbital with
z2 character and four singly occupied d orbitals with large dyz
character per iron(II) center, which can ultimately be used to
rationalize the similar ZFS values in 2 and 3.
[{(Me3Si)2N}2Fe{Li(bta)}]2 (3). As solids, [FeBr2(thf)2] (0.72 g, 2.0
mmol), LiN(SiMe3)2 (0.65 g, 4.0 mmol) and 1H-benzotriazole (0.24 g,
2.0 mmol) were combined and cooled to -78 ꢀC, and toluene (30 mL)
was then added. The reaction was stirred for 30 min at -78 ꢀC before
being warmed slowly to room temperature. The initial green-brown
color of the reaction mixture changed to a deep brown color after 2 h of
stirring. After 14 h at room temperature, the solution was filtered
(porosity 3, Celite) and the solvent volume was reduced in vacuo to
∼12 mL. Storing the resulting brown solution at -30 ꢀC overnight
resulted in the formation of air-sensitive green crystals of 3 (1.42 g,
72%). Analysis: calcd: C 43.11, H 8.04, N 13.97%; found: C 43.25, H
1
8.20, N 13.62%. H NMR (C6D6, 400.23 MHz, δ/ppm): 0.11 ppm,
36 H, SiMe3; -11.2 ppm, broad, bta. UV/vis (toluene), λmax (ε (dm3
M-1 cm-1)): 286 nm (2650).
X-ray Crystallography. X-ray crystallographic data on 2 were
collected using synchrotron radiation at the Diamond Light Source
(Beamline I19), and data on 3 were collected using an Oxford Instru-
ments XCalibur 2 diffractometer (Table 2). The CCDC references
codes are 798917 (2) and 798918 (3).
Magnetic Susceptibility Measurements. Magnetic suscept-
ibility measurements on polycrystalline samples of both compounds
were carried out on a Quantum Design MPMS XL SQUID magnet-
ometer in O-ring-sealed KelF capsules. Data were corrected for
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Inorganic Chemistry
ARTICLE
Table 2. Crystal Data for Complexes 2 and 3
2
3
formula
C
47H92Fe3N14Si4
C18H40FeLiN5Si4
501.68
formula weight
T, K
1133.23
100(1)
100(2)
wavelength, Å
crystal system
space group
1.54178
0.71073
monoclinic
P21/n
orthorhombic
Pccn
a, Å
22.522(4)
18.0458(7)
8.5059(3)
18.7423(7)
90
b, Å
14.455(3)
c, Å
17.979(3)
R, deg
90
β, deg
90
101.276(4)
90
γ, deg
90
V/Å3
5853.2(19)
2821.34(19)
4
Z
4
density (calculated), g/cm3
absorption coefficient, μ/mm-1
crystal size, mm3
θ range, deg
reflections collected
independent reflections
R(int)
1.286
1.181
5.230
0.718
0.1 ꢀ 0.1 ꢀ 0.1
4.39 to 89.23
0.1 ꢀ 0.1 ꢀ 0.1
2.79-28.45
10747
26727
6621
7065
0.0309
0.0688
% completeness
data/restraints/parameters
goodness-of-fit on F2
final R indices [I > 2σ(I)]
R indices (all data)
98.8
98.1
6621/0/316
7065/0/274
0.836
1.139
R1 = 0.0386, wR2 = 0.0893
R1 = 0.0448, wR2 = 0.0972
R1 = 0.0378, wR2 = 0.0813
R1 = 0.0699, wR2 = 0.0855
diamagnetism using Pascal constants, and for the diamagnetic contribu-
tion of the sample holder by measurement.
1786. (b) Olmstead, M. M.; Power, P. P.; Shoner, S. C. Inorg. Chem.
1991, 30, 2547–2551.
(2) Lappert, M. F.; Protchenko, A.; Power, P. P.; Seeber, A. In Metal
Amide Chemistry; John Wiley and Sons, Ltd.: Chichester, U.K., 2009;
Chapter 6.
(3) Power, P. P. Comments Inorg. Chem. 1989, 8, 177–202.
(4) Bartlett, R. A.; Power, P. P. J. Am. Chem. Soc. 1987, 109, 7563–7564.
(5) Chen, H.; Bartlett, R. A.; Dias, H. V. R.; Olmstead, M. M.; Power,
P. P. J. Am. Chem. Soc. 1989, 111, 4338–4345.
(6) Panda, A.; Stender, M.; Wright, R. J.; Olmstead, M. M.; Klavins,
P.; Power, P. P. Inorg. Chem. 2002, 41, 3909–3916.
(7) Reiff, W. M.; Schulz, C. E.; Whangbo, M. H.; Seo, J. I.; Lee, Y. S.;
Potratz, G. R.; Spicer, C. W.; Girolami, G. S. J. Am. Chem. Soc. 2009, 131,
404–405.
DFT Calculations. Density functional calculations were carried
out using the Gaussian 03 suite of programs,41 with the Def2-SVP basis
set on all atoms.42 The crystallographically determined geometry was
used in conjunction with the BP8643 and B3LYP44 exchange-correlation
functionals to determine the ground state multiplicity of 2.
’ ASSOCIATED CONTENT
S
Supporting Information. X-ray crystallography data in
b
CIF format. B3LYP/Def2-SVP natural orbitals of compound 2.
This material is available free of charge via the Internet at http://
pubs.acs.org.
(8) Merrill, W. A.; Stich, T. A.; Brynda, M.; Yeagle, G. J.; Fettinger,
J. C.; De Hont, R.; Reiff, W. M.; Schulz, C. E.; Britt, R. D.; Power, P. P. J.
Am. Chem. Soc. 2009, 131, 12693–12702.
’ AUTHOR INFORMATION
(9) Ellison, J. J.; Ruhlandt-Senge, K.; Power, P. P. Angew. Chem., Int.
Ed. 1994, 22, 1178–1180.
(10) MacDonnell, F. M.; Ruhlandt-Senge, K.; Ellison, J. J.; Holm,
R. H.; Power, P. P. Inorg. Chem. 1995, 34, 1815–1822.
(11) Bartlett, R. A.; Ellison, J. J.; Power, P. P.; Shoner, S. C. Inorg.
Chem. 1991, 30, 2888–2894.
(12) Nehete, U. N.; Anantharaman, G.; Chandrasekhar, V.; Mur-
ugavel, R.; Walawalker, M. G.; Roesky, H. W.; Vidovic, D.; Magull, J.;
Samwer, K.; Sass, B. Angew. Chem., Int. Ed. 2004, 43, 3832–3835.
(13) Chen, H.; Olmstead, M. M.; Pestana, D. C.; Power, P. P. Inorg.
Chem. 1991, 30, 1783–1787.
Corresponding Author
*E-mail: Richard.Layfield@manchester.ac.uk.
’ ACKNOWLEDGMENT
The authors acknowledge the support of the EPSRC, The
EPSRC UK National Electron Paramagnetic Resonance Service
at The University of Manchester, and The Royal Society. R.A.L.
thanks the Alexander von Humboldt Foundation for the award of
a Fellowship for Experienced Researchers.
(14) Verma, A. K.; Lee, S. C. J. Am. Chem. Soc. 1999, 121, 10838–
10839.
(15) Duncan, J. S.; Nazif, T. M.; Verma, A. K.; Lee, S. C. Inorg. Chem.
2003, 42, 1211–1224.
’ REFERENCES
(1) (a) Andersen, R. A.; Faegri, K., Jr.; Green, J. C.; Haaland, A.;
Lappert, M. F.; Leung, W.-P.; Rypda, K. Inorg. Chem. 1988, 27, 1782–
2525
dx.doi.org/10.1021/ic102341a |Inorg. Chem. 2011, 50, 2521–2526

Inorganic Chemistry
ARTICLE
(16) Sydora, O. L.; Wolczanski, P. T.; Lobkovski, E. B. Angew. Chem.,
Int. Ed. 2003, 42, 2685–2687.
lvarez, P.; Kennedy, A. R.; Klett, J.; Mulvey, R. E.; Rentschler, E.; Russo,
L. Angew. Chem., Int. Ed. 2009, 48, 3317–3321.
(17) Ohki, Y.; Sunada, Y.; Honda, M.; Katsada, M.; Tatsumi, K. J.
Am. Chem. Soc. 2003, 125, 4052–4053.
(18) Ohki, Y.; Sunada, Y.; Tatsumi, K. Chem. Lett. 2005, 34, 172–
173.
(37) (a) Lambert, C.; Hampel, F.; von R. Schleyer, P. J. Organomet.
Chem. 1993, 455, 29–35. See also: (b) Andrews, P. C.; Clegg, W.;
Mulvey, R. E.; O'Neill, P. A.; Wilson, H. M. M. J. Chem. Soc., Chem.
Commun. 1993, 1142-1143.
(19) Sydora, O. L.; Henry, T. P.; Wolczanski, P. T.; Lobkovski, E. B.;
Rumbovsky, E.; Hendrickson, D. N. Inorg. Chem. 2006, 45, 609–626.
(20) Ohki, Y.; Ikagawa, Y.; Tatsumi, K. J. Am. Chem. Soc. 2007, 129,
10457–10465.
(21) Pryadun, R.; Holm, R. H. Inorg. Chem. 2008, 47, 3366–
3370.
(22) Ohki, Y.; Imada, M.; Murata, A.; Sunada, Y.; Ohta, S.; Honda,
M.; Sasamori, T.; Tokitoh, N.; Katada, M.; Tatsumi, K. J. Am. Chem. Soc.
2009, 131, 13168–13178.
(38) (a) Mulvey, R. E. Chem. Soc. Rev. 1998, 27, 339–346. (b)
Mulvey, R. E. Chem. Soc. Rev. 1991, 20, 167–209. (c) Armstrong, D. R.;
Barr, D.; Clegg, W.; Hodgson, S. M.; Mulvey, R. E.; Reed, D.; Snaith, R.;
Wright, D. S. J. Am. Chem. Soc. 1989, 111, 4719–4727.
(39) Borrꢀas-Almenar, J. J.; Clemente-Juan, J. M.; Coronado, E.;
Tsukerblat, B. S. J. Comput. Chem. 2001, 22, 985–991.
(40) Andres, H.; Bominaar, E. L.; Smith, J. M.; Eckert, N. A.;
Holland, P. L.; M€unck, E. J. Am. Chem. Soc. 2002, 124, 3012–3025.
(41) Frisch, M. J.; et al. Gaussian 03, Revision E.01; Gaussian, Inc.:
Wallingford, CT, 2004.
(23) Hauptmann, R.; Kliss, R.; Schneider, J.; Henkel, G. Z. Anorg.
Allg. Chem. 1998, 624, 1927–1936.
(24) Hauptmann, R.; Kliss, R.; Henkel, G. Angew. Chem., Int. Ed.
1999, 38, 377–379.
(25) Dumestre, F.; Chaudret, B.; Amiens, C.; Renaud, P.; Fejes, P.
Science 2004, 303, 821–823.
(42) Weigend, F.; Ahlrichs, R. Phys. Chem. Chem. Phys. 2005, 7,
3297.
(43) (a) Becke, A. D. Phys. Rev. A 1988, 38, 3098. (b) Perdew, J. P.
Phys. Rev. B 1986, 33, 8822. (b) Becke, A. D. J. Chem. Phys. 1993, 98,
5648.
(26) Lacroix, L. -M.; Lachaize, S.; Falqui, A.; Respaud, M.; Chau-
dret, B. J. Am. Chem. Soc. 2009, 131, 549–557.
(44) (a) Mayer, I. Chem. Phys. Lett. 1983, 97, 270. (b) Mayer, I. Int. J.
Quantum Chem. 1984, 26, 151.
(27) (a) Layfield, R. A. Chem. Soc. Rev. 2008, 37, 1098–1107. (b)
Alvarez, C. S.; Bond, A. D.; Harron, E. A.; Layfield, R. A.; Mosquera,
M. E. G.; McPartlin, M.; Rawson, J. M.; Wright, D. S. Dalton Trans. 2003,
3002–3008. (c) Alvarez, C. S.; Bond, A. D.; Cave, D.; Mosquera,
M. E. G.; Harron, E. A.; Layfield, R. A.; McPartlin, M.; Rawson, J. M.;
Wood, P. T.; Wright, D. S. Chem. Commun. 2002, 2980–2981. (d)
Aluarez, C. S.; Bond, A. D.; Harron, E. A.; Layfield, R. A.; McAllister,
J. A.; Pask, C. P.; Rawson, J. M.; Wright, D. S. Organometallics 2001, 20,
4135–4138.
(28) Layfield, R. A.; McDouall, J. J. W.; Sulway, S. A.; Tuna, F.;
Collison, D.; Winpenny, R. E. P. Chem.—Eur. J. 2010, 16, 4442–4446.
(29) Layfield, R. A.; Bashall, A.; McPartlin, M.; Rawson, J. M.;
Wright, D. S. Dalton Trans. 2006, 1660–1666.
(30) Bailey, P. J.; Pace, S. Coord. Chem. Rev. 2001, 214, 91–141.
(31) (a) Foley, S. R.; Yap, G. P. A.; Richeson, D. S. Chem. Commun.
2000, 1515–1516. (b) Foley, S. R.; Yap, G. P. A.; Richeson, D. S. Inorg.
Chem. 2002, 41, 4149–4157. (c) Rose, R. P.; Jones, C.; Schulten, C.;
Aldridge, S.; Stasch, A. Chem.—Eur. J. 2008, 14, 8477–8480.
(32) Oakley, S. H.; Coles, M. P.; Hitchcock, P. B. Inorg. Chem. 2004,
43, 7564–7566.
(33) (a) Cotton, F. A.; Gruhn, N. E.; Gu, J.; Huang, P.; Lichtenber-
ger, D. L.; Murillo, C. A.; Van Dorn, L. O.; Wilkinson, C. C. Science 2002,
298, 1971–1974. (b) Cotton, F. A.; Gruhn, N. E.; Lichtenberger, D. L.;
Murillo, C. A.; Timmons, D. J.; Van Dorn, L. O.; Villagrꢀan, D.; Wang, X.
Inorg. Chem. 2006, 45, 201–213. (c) Cotton, F. A.; Daniels, L. M.;
Murillo, C. A.; Timmons, D. J.; Wilkinson, C. C. J. Am. Chem. Soc. 2002,
124, 9249–9256. (d) Cotton, F. A.; Murillo, C. A.; Reibenspies, J. H.;
Villagrꢀan, D.; Wang, X.; Wilkinson, C. C. Inorg. Chem. 2004, 43, 8373–
8378. (e) Fernandez-Cortabitarte, García, F.; Morey, J. V.; McPartlin,
M.; Singh, S.; Wheatley, A. E. H.; Wright, D. S. Angew. Chem., Int. Ed.
2007, 46, 5425. (f) Chisholm, M. H.; Gallucci, J.; Hadad, C. M.; Huffma,
J. C.; Wilson, P. J. J. Am. Chem. Soc. 2003, 125, 16040–16049. (g)
Cotton, F. A.; Dalal, N. S.; Huang, P.; Ibragimov, S. A.; Murillo, C. A.;
Piccoli, P. M. B.; Ramsey, C. M.; Schultz, A. J.; Wang, X.; Zhao, Q. Inorg.
Chem. 2007, 46, 1718–1726. (h) Berry, J. F.; Cotton, F. A.; Huang, P.;
Murillo, C. A.; Wang, X. Dalton Trans. 2005, 3713–3715. (i) Irwin,
M. D.; Abdou, H. E.; Mohamed, A. A.; Fackler, J. P., Jr. Chem. Commun.
2003, 2882. (j) Cotton, F. A.; Murillo, C. A.; Wang, X.; Wilkinson, C. C.
Dalton Trans. 2007, 3943–3951.
(34) Allan, F. H. Acta Crystallogr., Sect. B: Struct. Sci. 2002, B58, 380–
388.
(35) Collison, D.; McInnes, E. J. L.; Brechin, E. K. Eur. J. Inorg. Chem.
2006, 2725–2733.
(36) (a) Wunderlich, S. H.; Knochel, P. Angew. Chem., Int. Ed. 2009,
ꢀ
48, 9717–9720. (b) Alborꢀes, P.; Carrella, L. M.; Clegg, W.; García-A
2526
dx.doi.org/10.1021/ic102341a |Inorg. Chem. 2011, 50, 2521–2526