A family of Tri- and dimetallic pyridine dicarboxamide cryptates: Unusual O, N, O -coordination and facile access to secondary coordination sphere hydrogen bonding interactions

Article abstract of DOI:10.1021/ic502873d

A series of tri- and dimetallic metal complexes of pyridine dicarboxamide cryptates are reported in which changes to the base and metal source result in diverse structure types. Addition of strong bases, such as KH or KN(SiMe₃)₂, followed by divalent metal halides allows direct access to trinuclear complexes in which each metal center is coordinated by a dianionic N,N,N-chelate of each arm. These complexes bind a guest K⁺ cation within the central cavity in a trigonal planar coordination environment. Minor changes to the solvent and equivalents of base used in the syntheses of the triiron(II) and tricobalt(II) complexes affords two trinuclear clusters with atypical O,N,O-coordination by each pyridine dicarboxamide arm; the amide carbonyl O atoms are oriented toward the interior of the cavity to coordinate to each metal center. Finally, varying the base enables the selective synthesis of dinuclear nickel(II) and copper(II) complexes in which one pyridine dicarboxamide arm remains protonated. These amide protons are at one end of a hydrogen bonding network that extends throughout the internal cavity and terminates at a metal bound hydroxide, carbonate, or bicarbonate donor. In the dinickel complex, the bicarbonate cannot be liberated as CO₂ either thermally or upon sparging with N₂, which differs from previously reported monometallic complexes. The carbonate or bicarbonate ligands likely arise from sequestration of atmospheric CO₂ based on the observed reaction of the di(hydroxonickel) analog.

Full text of DOI:10.1021/ic502873d

Article
pubs.acs.org/IC
A Family of Tri- and Dimetallic Pyridine Dicarboxamide Cryptates:
Unusual O,N,O‑Coordination and Facile Access to Secondary
Coordination Sphere Hydrogen Bonding Interactions
†
†
‡
§
Gary L. Guillet,†,⊥ Jesse B. Gordon, Gianna N. Di Francesco, Matthew W. Calkins, Erik Cizmar,
̌
̌
́
Khalil A. Abboud,† Mark W. Meisel,‡ Ricardo García-Serres,∥ and Leslie J. Murray*,†
†Center for Catalysis, Department of Chemistry, University of Florida, Gainesville, Florida 32611-7200, United States
‡Department of Physics and the National High Magnetic Field Laboratory, University of Florida, Gainesville, Florida 32611-8440,
United States
§
̌
Institute of Physics, Faculty of Science, P.J. Safar
́
ik University, 04154 Kosi
̌
ce, Slovakia
∥Laboratoire de Chimie de Biologie des Met
Martyrs, 38054 Grenoble Cedex 9, France
́
aux, UMR 5249, Universite
́
Joseph Fourier, Grenoble-1, CNRS-CEA, 17 Rue des
S
* Supporting Information
ABSTRACT: A series of tri- and dimetallic metal complexes of
pyridine dicarboxamide cryptates are reported in which changes
to the base and metal source result in diverse structure types.
Addition of strong bases, such as KH or KN(SiMe3)2, followed
by divalent metal halides allows direct access to trinuclear
complexes in which each metal center is coordinated by a
dianionic N,N,N-chelate of each arm. These complexes bind a
guest K+ cation within the central cavity in a trigonal planar
coordination environment. Minor changes to the solvent and
equivalents of base used in the syntheses of the triiron(II) and
tricobalt(II) complexes affords two trinuclear clusters with
atypical O,N,O-coordination by each pyridine dicarboxamide
arm; the amide carbonyl O atoms are oriented toward the
interior of the cavity to coordinate to each metal center. Finally, varying the base enables the selective synthesis of dinuclear
nickel(II) and copper(II) complexes in which one pyridine dicarboxamide arm remains protonated. These amide protons are at
one end of a hydrogen bonding network that extends throughout the internal cavity and terminates at a metal bound hydroxide,
carbonate, or bicarbonate donor. In the dinickel complex, the bicarbonate cannot be liberated as CO2 either thermally or upon
sparging with N2, which differs from previously reported monometallic complexes. The carbonate or bicarbonate ligands likely
arise from sequestration of atmospheric CO2 based on the observed reaction of the di(hydroxonickel) analog.
nuclearity or the coordination environment of each metal
center, whereas the latter can suffer from poor tunability or low
synthetic yields for the ligand.5a,b
INTRODUCTION
■
Biological activation of small molecule substrates, such as O2,
CO2, and N2, typically occurs at metal clusters housed within
the active site of metalloenzymes.1 In these systems, redox
cooperativity between metal ions and the high or intermediate
spin states of the metal centers within the clusters are key
requirements for the observed reactivity.1b−d,p,2 Coupled to the
properties inherent to the metal clusters, hydrogen bonding
interactions within the secondary coordination sphere interact
with and orient substrates to favor bond scission, redox events,
or proton transfers.1d,2c,3 Significant effort has focused on
synthetic modeling of chalcogenide-bridged biological clusters
(e.g., ferredoxins, oxygen-evolving complex) using either self-
assembly,4 in which the cluster is generated by the reaction of a
low molecular weight ligand with a metal precursor, or by
metalation of a ligand which preorganizes cluster assembly.5 In
the former approach, there is limited control of the cluster
In contrast to the sulfide- or oxide-bridged clusters, the
diiron, tricopper, and heme-iron−copper active sites of the
bacterial monooxygenases, multicopper oxidases, and cyto-
chrome c oxidase, respectively, display weak interactions
between the metal centers in the reduced state, but substrate
binding activates cooperativity between the metal
centers.1a−e,p,2b,c,6 A number of groups have developed self-
assembled and designed ligands to model these active sites;6b,7
however, these systems display limited a priori control of the
metal−metal distance or are unable to probe the role of
proximal hydrogen bonding interactions. In addition, few
Received: December 1, 2014
© XXXX American Chemical Society
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examples have been reported in which hydrogen bonding
interactions are incorporated into dinucleating ligands.8
held within the N,N,N-chelate of each pyDCA arm. In addition,
we report the characterization of two unexpected families of
complexes derived from minor changes to our synthetic
protocol for the trinuclear N,N,N-coordinated series. The first
are the triiron and tricobalt complexes, in which each metal
center is coordinated by the two amide O atoms and the
pyridyl N atom of each arm, and these compounds represent
uncommon examples of O,N,O-coordination by pyDCA
donors. For the second, we observed that selective dimetalation
of H6L can be accomplished with copper(II) or nickel(II) to
afford complexes containing a proximal hydrogen bonding
network within the internal cavity.
Our focus has been on the use of macrobicyclic ligands to
template the assembly of multimetallic complexes, wherein
substrate binding within the internal void space facilitates
electronic communication between the metal centers and leads
to cooperative activation of substrates.9 Here, we employ our
previously reported cryptand, H6L, composed of three pyridine
dicarboxamide donor arms and two tri(5-ethoxyphenyl)-
methane caps (Figure 1).9a In contrast to the monometallic
RESULTS AND DISCUSSION
■
Synthesis, Characterization, and Reactivity of N,N,N-
Coordinated Trinuclear Cryptates. As anticipated from our
prior report,9a deprotonation of H6L with a strong base
followed by addition of a divalent metal halide indeed affords
the corresponding trimetallic complexes in yields of 39%, 78%,
92%, and 62% for the tricopper(II) (1), triiron(II) (2),
tricobalt(II) (3), and trinickel(II) (4) compounds. Upon
metalation of the dianionic N3-chelate of each pyDCA arm,
we observe distinctive changes in the infrared spectra relative to
the protonated ligand: there is a decrease in the intensity of the
absorptions at 1670 and 1683 cm−1 and an increase in intensity
at ∼1600 cm−1 in the complexes relative to the free ligand, as
well as strong absorption bands between 1590 and 1570 cm−1
for 1−3. Minor differences from this general trend are observed
in spectra of the trinickel(II) complex 4; specifically, the
presence of absorption bands at 1610, 1630, and 1590 cm−1
(Figures S1−S4, Supporting Information). These spectral
changes are indicative of N,N,N-coordination of the bound
metal centers and provide a facile method to aid in identifying
the cryptand donor atoms in the reaction product (vide infra).
In the solid state structure of 1 (Figure 2), one of the two
[K(18-crown-6)]+ countercations stacks within a cleft between
two pyDCA arms of the dianionic complex [(CuCl)3KL]2−
with the other two clefts occupied by solvent molecules
(Supporting Information Figure S8). Each copper(II) center is
held in a distorted square planar coordination environment (τ4
= 0.14−0.20) composed of three pyDCA N-atoms and one
Figure 1. Structure of ligand H6L.
pyridine-2,6-dicarboxamide (pyDCA) complexes in which
bulky aryl substituents on the amide N atoms preclude
formation of bis(pyDCA) compounds, the cryptand structure
of H6L provides the steric protection for each metal site,
simultaneously positions the metal ions in close proximity, and,
critically, ensures that the kinetically labile coordination sites
are directed toward the internal cavity. Inspired by recent
reactivity studies,10 we aimed to expand our library of
complexes to include divalent late 3d metal centers. Herein,
we report the synthesis of trimetallic complexes of iron(II),
cobalt(II), nickel(II), and copper(II), in which each metal is
Figure 2. Solid state structure of 1 collected at 100 K (a) and cut away of the central cavity (b). C, N, O, Cu, K, and Cl atoms are represented as
gray, blue, red, yellow-green, pink, and green ellipsoids (15%), respectively. Two K[18-crown-6]+ ions, solvents of crystallization, H atoms, and OEt
and iPr substituents have been removed for clarity.
B
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Figure 3. Solid state structure of 4 collected at 100 K (a) and cut away of the central cavity (b). C, N, O, Ni, and Cl atoms are represented as gray,
blue, red, aqua, and green ellipsoids (70%), respectively. A K[18-crown-6]+ cation, solvents of crystallization, H atoms, and OEt and iPr substituents
have been removed for clarity.
chloride donor, and the copper(II) ions lie nominally out of the
N3Cl plane (<0.04 Å). As expected from the constraints placed
on the complex by the cryptand ligand, there are a number of
similarities between the structure of 1 and that reported for the
triiron(III) complex, [(FeCl2)3L]3−.9a First, each metalated
ligand arm is overall monoanionic as a chloride coordinates to
each metal center. Whereas a number of monocopper(II)
pyDCA complexes are solvent-adducts, 1 incorporates chloride
donors instead of solvent in the fourth coordination site,
presumably because the steric constraints enforced by the
cryptand. Second, the plane of the pyridyl ring in each arm of 1
is twisted out of the plane of the aryl rings in each
triphenylmethane cap (60.4−71.1°). This configuration results
in the Npyr−Cu vector of each arm oriented slightly away from
the center of the internal cavity. Finally, the metal−metal
distances in 1 (6.525−6.694 Å) are similar to those in the ferric
complex (6.723 Å).9a
We were surprised that the tricopper(II) cryptate accom-
modates a guest K+ ion, which is coordinated to only the Cl−
donors, instead of isolating the trianionic [(CuCl)3L]3−
complex. The K+−arenecap (4.865−5.128 Å) and K+−Hcap
(3.121−3.385 Å) distances are significantly longer than those
usually observed for cation−π interactions,11 which supports
our assignment of a nominally three-coordinate K+ ion. On the
basis of prior reports of the nucleophilicity of the fourth (non
N-atom) donor in mononuclear pyDCA compounds,10a−g the
three chloride ligands likely create a favorable electrostatic
pocket to interact with the alkali cation to stabilize the observed
and unusual low coordination number. The synthetic approach
for 1 also suggests that this alkali cation binding site is selective
for K+ over other cations and also has a higher affinity for K+
than 18-crown-6. Specifically, sufficient equivalents of 18-
crown-6 are present in the reaction mixture to account for each
charge balancing K+ for the expected complex [(CuCl)3L]3−;
yet, mass spectrometry, crystallography, and combustion
analysis data provide evidence for formation of only
[(CuCl)3KL]2−. Moreover, our preliminary attempts to test
this hypothesis by replacing KN(SiMe3)2 with NaN(SiMe3)2 in
the synthesis of 1 do not appear to afford a complex with guest
Na+ cation, but rather [(CuCl)3L]3− is isolated. To our
knowledge, 1 represents a unique example of a three-coordinate
K+ center and hints at a different design strategy for selective
alkali cation receptors. Although our attempts to obtain single
crystals of 2 and 3 of sufficient quality for X-ray diffraction
analysis are ongoing, HRMS-ESI(−) data and combustion
analyses of 1−3 support a parent ion formula consistent with
the doubly charged anionic complex containing one K+ cation,
three transition metal ions, and three chloride ions (Figure S5−
S7, Supporting Information). Given the similarities between the
IR spectra of these three compounds and the apparent stability
of the parent ion [(MCl)3KL]2−, we tentatively assign 2 and 3
as isostructural complexes to 1.
In contrast to data collected on complexes 1−3, the trinickel
complex was isolated as the monoanionic species [(Ni-
(THF))2(NiCl)L]− in the solid state (Figure 3). With respect
to 1, the [K(18-crown-6)]+ counteraction is displaced from the
cleft and stacks between the isopropyl groups in the cap, and
two nickel centers coordinate solvent molecules rather than
halide donors (Figures S8 and S9, Supporting Information).
Two structural changes allow 4 to accommodate the THF
donors and point to the unexpected flexibility of the ligand
(Table S2, Supporting Information). First, the distance
between the two central carbon atoms of the triphenylmethane
caps contracts slightly in 4 (9.23 Å) as compared to 1 (9.59 Å),
leading to a widening of the clefts between the cryptand arms
and an average increase in the metal−metal distances (6.9 0.3
Å in 4 vs 6.6
0.1 Å in 1). This increase in metal−metal
separation is more evident in the ∼7 Å distance between the
two THF-coordinated NiII centers in 4. Second, the angle
between the planes of the pyridyl rings coordinated to Ni1 and
Ni3 and those of the adjacent aryl groups on the triphenyl-
methane caps are larger (73.4−86.5°) than the comparable
angles in 1, which results in the Npyr−Ni vector being oriented
further away from the central cavity of the ligand and toward
the cleft. We cannot exclude the possibility that minor
differences in the synthetic procedures of 1 and 4 might result
in changes to product speciation or that one structure type
selectively crystallizes depending on the metal identity.
However, the subtle differences in the metal−ligand bond
distances that are expected between the nickel and copper ions
and each pyDCA arm could be amplified within the confines of
the cryptand cavity. Specifically, the Ni−Cl bond distances are
reported to be shorter than the comparable Cu−Cl distances in
the monometallic complexes,10a,h which would afford a larger
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central cavity lined by the halide donors for a hypothetical
[(NiCl)3L]3− complex. In addition, the tetragonal coordination
environment around each copper center in 1 allows the
chloride donor to be positioned slightly within the internal
cavity (cf., Npyr−Cu−Cl bond angles, Table S3, Supporting
Information), which further contracts the chloride−chloride
distances to likely favor K+ binding. Regardless, the structures
of 1 and 4 indicate that ligand flexibility can provide a pathway
for halide exchange with substrates and product dissociation
during reactions of the N,N,N family of complexes.
In cyclic voltammograms on solutions of 1 in DMF, we
observe a quasireversible wave at E1/2 = −0.043 V vs Fc+/Fc
(Figure 4), which is at a comparable potential to the redox
Figure 5. Cyclic voltammograms for 2 and 3 in MeCN at a scan rate of
150 mV/s. 0.1 M TBAPF6 used as supporting electrolyte. Working
electrode: 1 mm Pt button. Auxiliary electrode: Pt wire. Reference
electrode: Ag/AgNO3 in MeCN.
oriented toward the internal cavity, we explored other metal ion
sources and, specifically, the hexamethyldisilazide and trifluor-
omethanesulfonate (−OTf) salts. For the former, products were
ill-defined and difficult to isolate and characterize. In the course
of our studies using iron(II) and cobalt(II) triflates (triflate =
trifluoromethanesulfonate) as metal precursors, we synthesized
two complexes as the major product in which each pyridine
dicarboxamide adopts an atypical O,N,O-coordination mode to
each metal center.
Reaction of 3.3 equiv of KH with H6L followed by addition
of the bis(acetonitrile) adduct of either Fe(OTf)2 or Co(OTf)2
yielded the trinuclear O,N,O-coordinated compounds [(Fe-
(MeCN))3(μ3-O)LH5][OTf]4 (5) and [Co3(μ-OH)(μ3-
OH)(μ-OTf)(OTf)(MeCN)LH6][OTf]2 (6), respectively.
Both complexes are synthesized in good yield at 79% and
51% for the iron and cobalt species, respectively. Two aspects
of the syntheses are worth noting: first, the optimized syntheses
require half of the equivalents of KH as compared to the
syntheses of 1−4, and second, acetonitrile is an essential
cosolvent in the reaction as the products are isolated in poor
yield if this solvent is omitted. Independent of the intense
infrared absorptions arising from the triflate anions in 5 and 6
(viz. 1028 and 637 cm−1), there are other systematic differences
between the IR spectra of 5 and 6, and of 1−4: the relative
intensity for the absorptions at 1100 and 960 cm−1 are
significantly reduced, and new intense features arise at 1631,
1556, 1233, 1203, and 1170 cm−1 in spectra of 5 and 6 (Figures
S11 and S12, Supporting Information).
As alluded to above, each pyDCA arm coordinates to the
metals in an O,N,O-fashion, and a μ3-oxide bridge is installed
within the center of the cluster in the structure of 5 (Figure 6).
To adopt this unique coordination mode for pyDCA to iron,
the donor arms rotate the appended carbonyl oxygen atoms
toward the interior of the scaffold and dramatically alter the
aspect ratio of the complex relative to the N,N,N-chelated
compounds 1−4. The distance between the central C atoms in
each cap increases by ∼3.45 Å, and the pyridyl N-atoms move
inward toward the internal cavity by ∼3.21 Å as compared to 1.
The coordination geometry around each iron atom is best
described as distorted square pyramidal with O1, N2, O2, and
Figure 4. Cyclic voltammogram of 1 in DMF over a limited potential
window at a scan rate of 100 mV/s. 0.1 M TBAPF6 used as supporting
electrolyte. Working electrode: 1 mm Pt button. Auxiliary electrode: Pt
wire. Reference electrode: Ag/AgNO3 in MeCN.
process assigned as the CuII/III couple by Tolman and co-
workers (−0.076 V vs Fc+/Fc) for a monocopper(II)-hydroxide
pyDCA complex.10f The wave observed here is broad, and no
other redox events are present over the scanned potential
window, indicative of minimal electronic communication
between the metal centers. In contrast, two closely spaced
but discernible redox events are present in voltammograms of 2
at E1/2 = −0.152 and −0.422 V, and only ill-defined anodic and
cathodic events are observed for 3 at 0.047 and −0.444 V
(Figure 5). The redox processes for 2 are analogous to those
for mononuclear ferrous (−0.490 V vs SCE) and ferric (−0.570
V vs SCE) pyDCA complexes12 and are consistent with the
pyridine-diimine13 and pyridine-diamine14 iron(II) compounds
if the trend in donor strengths is considered. The richer redox
chemistry of 2 is surprising as compared to the other
compounds in this family, and could suggest either ligand-
based redox events [the tris(alkoxyphenyl)methane caps are
oxidatively sensitive15] or access to unique structures that afford
improved electronic coupling (vide infra).
O,N,O-Ligated Trimetallic Complexes Containing a μ3-
Oxide or μ3-Hydroxide. In three of the four trinuclear
complexes reported above, metalation using metal chlorides
afforded compounds in which a chloride coordinates to each
metal center in the complex. In an effort to use the steric
constraints of the ligand to enforce an open-coordination site
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Figure 6. Solid state structure of 5 collected at 100 K (a) and cut away of the central cavity (b). C, H, N, O, and Fe atoms are represented as gray,
light gray, blue, red, and orange ellipsoids (60%), respectively. All H atoms have been removed (except those on amide moieties which were freely
refined) along with solvents of crystallization, four triflate anions, and the OEt and iPr substituents for clarity. Symmetry operations were used to
generate equivalent atoms: y − x, 1 − x, z; 1 − y, 1 + x − y, z.
O5 forming the basal plane and N4 of the coordinated MeCN
occupying the axial position (τ = 0.19) and Fe1 held 0.616 Å
out of the basal plane and shifted toward N4. The four outer-
sphere triflate counteranions in the structure of 5 and the amide
protons, which were refined freely, suggest that the complex is
presumably tetracationic with the formula [Fe3(MeCN)3(μ3-
O)H6L]4+ and an oxo-bridged triiron(II) cluster housed within
the ligand; however, this formulation is not congruent with all
spectroscopic and magnetic data. The planarity observed for
the {Fe3O} cluster in 5 is similar to that reported for other oxo-
bridged triiron compounds,16 whereas hydroxide bridged
complexes are typically more pyramidalized. In addition, the
Fe−Ooxo bond length of 1.8721(6) Å in 5 is comparable to
ferrous pentanuclear helicates reported by Kawata16a (1.868 Å,
1.929 Å) and Oshio16b (1.882 Å, 1.904 Å) with all metal ions
adopting a high-spin configuration. Two examples of mixed-
valent, helicate polynuclear iron compounds are known in
which monometallic pyDCA complexes with N,N,N-coordina-
tion nucleate the cluster.17,18 One example shows clear
shortening of the Fe−Ooxo bond length (av 1.856 Å) and loss
of C3 symmetry in the solid state17 as expected with the
incorporation of a ferric ion, whereas the other is assigned as an
FeIII2/FeII cluster and is C3 symmetric in the solid state with
longer Fe−Ooxo bonds (1.887 Å).18 In 5, however, the 3-fold
crystallographic symmetry could allow one of the six amide
protons to be disordered over either the upper or lower
hemispheres of the structure, which in turn would afford a
mixed valent species. At present, this argument seems the most
reasonable considering the magnetic susceptibility data, the
Given that oxo-bridged all-ferrous clusters are unusual, we
first attempted to validate the protonation state of the amide N
1
atoms by reacting 5 with MeOTf. H NMR spectra of the
organics isolated from acid hydrolysis of that product isolated
after treating 5 with MeOTf were indistinguishable from that of
H6L. To probe the oxidation state of the iron centers within the
cluster, Mossbauer spectra and variable temperature magneto-
̈
metry data were recorded on 5 (Figures 7 and 8, respectively).
The zero-field Mossbauer spectrum measured at 80 K (Figure
̈
Figure 7. Zero-field, 80 K Mossbauer spectra of (a) a powder sample
̈
of 5 and (b) an acetonitrile solution of 5 synthesized with 57Fe-
enriched iron. The blue lines are quadrupole doublet simulations of
the two components assigned to high-spin FeII centers with the
parameters stated in the text.
Mossbauer data, and the combustion analysis measurements
̈
(vide infra).
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Figure 8. Variable temperature magnetometry in 10 mT (left) and field dependence of magnetization data collected at 5 K (right) of 5.
Figure 9. Solid state structure of 6 collected at 100 K (a) and cut away from the central cavity (b). C, H, N, O, S, F, and Co atoms are represented as
gray, light gray, blue, red, yellow, teal, and pink ellipsoids (30%), respectively. All hydrogens have been removed (except calculated H positions on
O13 and O14) along with solvents of crystallization, and two triflate anions for clarity.
7a) comprises a very broad doublet with parameters character-
istic of high-spin Fe(II) and a magnetically split component
accounting for roughly one-third of the total iron. Each peak of
the doublet shows shoulders that clearly indicate that it is
composed of at least two separate species. Indeed, it was
successfully simulated with two symmetric doublets of equal
Oshio.16b Our values here are larger than those for the iron(II)
centers in clusters reported by Betley and co-workers as might
be expected, given the significant structural and electronic
differences in their complexes as compared to 5 (e.g., absence
of metal−metal interactions).20 The magnetic component is a
broad-line sextet centered on 0.64 mm/s that does not change
shape with increasing temperature (up to 210 K). This
observation is indicative of a half-integer spin system with a
long spin−lattice relaxation time. The overall splitting and the
isomer shift of the centroid are consistent with high-spin FeIII.
The line width and relaxation behavior may be attributed to
magnetic ordering in the solid phase. In such a case, a solution
spectrum should display different behavior as the paramagnets
would be separated by solvent molecules. However, a sample of
the complex synthesized from 57Fe-enriched iron was
intensities and parameters δ1 = δ2 = 1.16 mm/s, and ΔEQ1
=
2.04 mm/s and ΔEQ2 = 3.08 mm/s. The individual peaks are
broad (Γ1(FWHL) = 0.83 mm/s and Γ2 = 0.64 mm/s) and are
best fitted with Gaussian lines. These doublets are of equal
intensity, and the simulated parameters are consistent with
iron(II) centers in 5. The δ and ΔEQ values are comparable to
those reported by Agapie and co-workers for the iron(II)
centers in the mixed valent heterometallic oxo-triiron cluster19
and in the oxo-triiron(II) complexes reported by Kawata16a and
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subsequently dissolved in acetonitrile and yielded an identical
spectrum with again a relative integration of the two signals of
2:1 (Figure 7b). This surprising invariability between solution
and solid-state spectra could arise from the presence of
insoluble nanoparticles of, for example, ferric oxide. We have
however discarded this hypothesis based on other analyses (viz.
elemental analyses and X-ray crystallography) that indicate that
the complex is pure to 2% with the formula determined by X-
ray crystallography. Additionally, the fact that the two
Co centers and the central O atom donor are 2.117(5) Å
(Co1), 2.057(5) Å (Co2), and 1.940(5) Å (Co3), and the
central O atom lies 0.4 Å below the Co3 plane. These bond
distances are slightly longer than the values reported for a
hydroxide-bridged tricobalt cluster, and the hydroxide donor
lies similarly out-of-plane of the three metal centers in that
reported complex.25 Finally, bond valence sum analysis
supports the oxidation assignment of each metal center as
CoII, leading to an assignment of the bridging ligands as
hydroxides (Table S4, Supporting Information).
The redox properties of the trinuclear clusters in 5 and 6
were investigated by cyclic voltammetry. Two broad,
irreversible processes at −1.614 and −2.504 V are present in
voltammograms of 5 as well as two pseudoreversible processes
at 1.846 and 2.136 V. The broadness of the peaks is similar to
our reported data for [(FeCl2)3L]3−, and agrees with the weak
coupling between the metal centers from our simulations of the
magnetometry data. A similar scenario was observed for 6 with
the oxidative (2.003 and 1.683 V) and reductive (−1.227 and
−1.597 V) events shifted to lower absolute values for the
potentials (Figures S13−S14, Supporting Information).
The absence of reversible one-electron redox events in the
data collected on both 5 and 6 was surprising given that the
structural support offered by the macrobicyclic ligand was
expected to provide stability to redox cycling. These results
contrast the reversible couples observed for the heterometallic
oxo-bridged triiron clusters reported by Herbert et al.19 and the
rich redox chemistry reported for the strongly interacting
triiron clusters.20 This difference can likely arise from the
significantly weaker coupling between the metal centers in 5 as
compared to the aforementioned ones. One possibility is that
the perceived irreversibility could arise from ligand rearrange-
ment upon reduction or oxidation of the cluster. Such structural
changes are supported in our reactivity results (vide infra) in
which treatment of N3-ligated trimetallic compound with
TMSOTf causes demetalation and rotation of one ligand arm
from the N,N,N- to the O,N,O-arrangement.
Mossbauer components are in fixed, integer proportions
̈
support their assignment as different iron centers of the same
multimeric complex, namely an FeIIFeIIFeIII mixed-valent trimer
with localized valences.
On the basis of the oxidation states determined from the
Mossbauer data, the spin system for 5 was defined with S = 5/
̈
1
2
and S2 = S3 = 2 corresponding to one iron(III) and two iron(II)
centers in high spin configurations, respectively. From prior
reports for (μ3-oxo)triiron(II) and the mixed valent ana-
logs,16,21 EasySpin22 simulations were constrained to single-ion
anisotropy with the iron(III) and the iron(II) sites restricted to
|D| < 3 cm−1 and |D| < 30 cm−1, respectively. In addition, the
two iron(II) ions (S2 and S3) were treated as equivalent without
including a rhombic anisotropy term to reduce the parameter
space for the simulation and to avoid overparameterization.
From our simulations, reasonable agreement to the temper-
ature-dependent susceptibility data was only achieved for both
J12 and J13 < J23. Small negative value of D for iron(III) sites was
required to model the low temperature regime of the
susceptibility as well as the curvature of the field dependence
of the magnetization. The best agreement, as shown by the
solid lines in Figure 8, was obtained with values of g1 = 2, g2 =
g3 = 2.03, D1 = −1.45 cm−1, D2 = D3 = 3.47 cm−1, and coupling
constants of J12 = J13 = −3.05 cm−1 and J23 = −6.35 cm−1.
3
Alternate values for S1 of 2, /2, and 1 were also used in these
simulations, but a value of 5/2 afforded the best agreement with
the high temperature χT data.
In contrast to 5, three unique cobalt centers are present in 6,
and each metal center is held in a pseudo-octahedral
coordination environment (Figure 9). As observed in 5, each
pyDCA arm coordinates to a metal center through the amide
carbonyl oxygen atoms and the pyridyl N-atom with a similar
change in the aspect ratio of the complex as compared to 1.
Here, however, the observed electron density for the amide
protons refines poorly, and the protons were therefore
calculated. Further support of the assignment of each arm as
protonated comes from the CO and CN bond distances,
which fall within the range typical for O-bound neutral amides
(OCNHR) as opposed to the deprotonated variant
(−OCNR).23 We were unable to locate electron density
corresponding to protons on the bridging O atom donors, and
relied on the CoO bond distances to assign the oxidation
state of the metal ions and protonation state of the bridging
ligands. The cationic complex is balanced by two outer-sphere
triflate anions and contains two coordinated triflate ligands, for
which the μ-1,1 bridging mode has not been reported
previously. Therefore, the trimetallic core of 6 is likely
tetracationic, suggesting that four possible assignments for the
complex in which protonation of the bridging O atom donor is
We were interested in the source of oxide or hydroxide
ligands in 5 and 6 as the reaction was carried out under
anaerobic and anhydrous conditions. The isolated yields of
both 5 and 6 suggest that neither ligand decomposition nor
trace contaminants in the reagents is the source of these O
atoms. Organic products consistent with solvent deoxygena-
tion26 were not observed by GC in the reaction headspace
during the synthesis of 5. The most likely candidate remains
adventitious water, possibly from one of the reagents, solvents,
and/or the glass surface. We attempted to synthesize 5 in the
absence of KH, but this route was unsuccessful. Solvent,
specifically acetonitrile, is the probable proton source because
addition of acetonitrile to solutions of L6− regenerates the free-
base ligand. Moreover, the products are neither isolated nor
observed from reactions in which acetonitrile is omitted. Access
to N,N,N- versus O,N,O-coordination modes here could be
determined by the ease of reprotonation and the size of the
counteranion as compared to the internal cavity; that is, the
smaller chloride is accommodated within the cavity whereas
triflate is not, and the complex rearranges to generate
coordinatively saturated metal centers. The effect of proto-
nation state of the amide N atoms and the resultant effect on
metal binding and the initial coordination mode are currently
being explored to provide insights into the design of complexes
with dynamic stimuli-driven structural properties. By way of a
preliminary attempt to probe such chemical stimuli to effect
balanced by a CoII instead of CoIII ion (e.g., [CoII (μ3-OH)(μ-
3
OH)]4+ and [CoII CoIII(μ3-O)(μ-OH)]4+). First, the Co1
2
O13 and Co2O13 bond distances are 1.930(5) Å and
1.958(5) Å and are within the range reported for high-spin
CoIIOH bonds.24 Second, the distances between the three
G
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Figure 10. Solid state structure of 7 collected at 100 K (a) and cut away of the central cavity (b). C, H, N, O, and Ni represented as gray, light gray,
blue, red, and aqua ellipsoids (60%), respectively. All H atoms have been removed except those which constitute the hydrogen bonding network in
the cavity along with solvents of crystallization, three tetraethylammonium cations, one triflate anion, and the OEt and iPr substituents.
structural changes, we were intrigued as to whether the two
complex types (i.e., N,N,N vs O,N,O) could be interconverted.
Our first hypothesis was that halide abstraction from 2 would
lead to coordinatively unsaturated metal centers frustrated by
the steric constraints imposed by the ligand, and ultimately
result in complex rearrangement. Preliminary work in which 2
is treated with 3 equiv of Me3SiOTf yields a dimetallic complex,
albeit in poor yield, for which the coordination mode changes
from N,N,N to O,N,O for the two retained iron centers (Figure
S16, Supporting Information). This preliminary result agrees
with the proposed synergy between the electronic and steric
effects in these complexes.
Preliminary Studies on O Atom Transfer Reactivity of
2. The ability of 2 to act as an oxygen atom transfer agent was
probed using iodosobenzene as the O atom source and
triphenylphosphine as the substrate. Although the size of the
internal cavity would not accommodate PPh3, its extensive use
in O atom transfer studies and the flexibility of the ligand arms
led to this choice of substrate. A series of conditions were
surveyed in which the ratio of 2:PPh3:PhIO was varied, and we
observed only Ph3PO and no unreacted starting phosphine in
reactions with the ratios of 1:1:3, 1:2:4, and 1:3:5; control
reactions of PhIO with PPh3 under similar conditions did not
yield Ph3PO (Figures S19−20 Supporting Information).
Reactions with fewer equivalents of PhIO, such as 1:1:2 and
1:1:1, showed incomplete substrate oxidation with both Ph3PO
and PPh3 present in the NMR sample. The number of
equivalents of PhIO required to effect complete oxidation of
triphenylphosphine suggests a possible activation step where 2
accepts an O atom from iodosobenzene, followed by structural
rearrangements to afford the reactive O atom transfer complex.
Such a rearrangement could traverse compounds in which
O,N,O- or N,N,O-coordination modes are observed as the
complex collapses to an oxo-centered trinuclear cluster. To
probe whether O atom incorporation in 2 leads to other ligand
coordination modes, we reacted 2 with iodosobenzene in the
absence of substrates, but we were unable to isolate crystalline
products from the reaction mixture, and the characteristic
changes to the IR spectrum for O,N,O-coordination were not
observed. Using a similar procedure as for Ph3P oxidation, we
evaluated O atom transfer by 2 from PhIO to styrene; however,
we observed only styrene and not expected oxidation products
in the product mixtures by gas-chromatography (data not
shown). These results suggest a limit for the O atom transfer
reactivity of the product of 2 and iodosobenzene, although
steric effects cannot be excluded here. Ongoing work aims to
define the substrate scope both in terms of sterics and
electronic factors as well as the mechanism of O atom transfer.
Incorporation of Hydrogen Bonding Networks in
Dimetalated Complexes. Prior reports on nickel(II) and
copper(II) pyDCA complexes used hydroxide sources instead
of stronger bases to simultaneously metalate and install a
hydroxide ligand on the resultant complex.10a,d,j These
mononuclear compounds sequester carbon dioxide reversibly
(nickel)10a−c and support reactive metal−oxygen species
(copper).10e−h,j In light of these types of reactivity, we pursued
a similar synthetic approach using H6L to define the
interactions between the metal-hydroxides within the internal
cavity. As a first step toward synthesizing these types of
complexes, we aimed to test whether alkylammonium
hydroxides were competent for deprotonation of the ligand
and subsequent metalation using either copper or nickel triflate.
Therefore, these reactions were performed on the benchtop,
and we anticipated carbonate or bicarbonate complexes
resulting from sequestration of atmospheric CO2 based on
reports by others on the monometallic pyDCA complexes.
Reaction of H6L with 3 equiv of either Ni(OTf)2 or
Cu(OTf)2 followed by excess tetraethylammonium hydroxide
as a methanolic solution in the presence of ambient atmosphere
led to an immediate color change from pale green or yellow to
dark red or dark green, respectively, and ultimately afforded
[Et4N]2[Ni(OCOOH)Ni(OH)H2L] (7) or [Et4N]2[Cu3(μ-
η1:η2-OCOO)(μ-OH)(OH)H2L] (9) in good yield. Even
though we were able to isolate the dinickel complex from the
product mixture (Figure S17, Supporting Information), the
reaction residues were quickly washed with aqueous sodium or
potassium hydroxide solutions to remove the [Et4N][OTf]
byproduct. Surprisingly, this seemingly nominal difference
between alkali cations afforded unique products for both the
Ni and Cu complex syntheses. Reaction of 3 equiv of Ni(OTf)2
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and 9 equiv of Et4NOH followed by a KOH wash reproducibly
afforded 7 whereas a NaOH wash led to mixtures from which
we were able to isolate a sodium inclusion complex,
[Et4N]2[Ni2(η1-OH)(μ-η1:η2-CO3)LH2Na(THF)], in poor
yield (Figure S18, Supporting Information).
ing similar arrays in extended solids or molecular systems will
lead to highly effective CO2 capture reagents.
The bicarbonate ligand could arise from reaction of Et4NOH
with carbon dioxide prior to complex formation. As noted
above, however, we suspected that metalation under conditions
used to synthesize 7 initially affords a complex containing either
one or two nickel-hydroxide centers, in which a NiII−OH
moiety rapidly sequesters an equivalent of atmospheric carbon
dioxide. Then, the steric constraints of the internal cavity,
hydrogen bonding interactions, or other electrostatic effects
would disfavor capture of a second equivalent of CO2 to install
a second bicarbonate donor. To test this hypothesis, the
di(hydroxonickel(II)) complex 8 was synthesized using a
procedure analogous to that for 7 except that the reaction was
sparged with nitrogen prior to adding Et4NOH and maintained
under an inert atmosphere thereafter. ESI(−)/MS data
collected on solutions of 8 contain an isotopic pattern
consistent with the expected [C89H99N9Ni2O14]2− at 817.8005
m/z as well as an ion envelope of much lower intensity
corresponding to 7, which could be present as a minor product
or arise from reaction with adventitious carbon dioxide during
MS data collection (Figure S22, Supporting Information). We
were unable to remove the EtN4OTf byproduct from the
reaction for satisfactory combustion analysis, and further
studies were carried out on the complex as-synthesized. To
avoid exposure to CO2, NMR spectra were recorded on a
rigorously dried sample of 8 in anhydrous solvent. From NMR
spectra of both 7 and 8 under air-free and anhydrous conditions
in d8-THF, it was readily apparent that, whereas spectra of 7 in
solvents containing trace water were well-resolved, those
collected in anhydrous solvent indicated that multiple species
are present upon exclusion of water. We postulated that the
complex might lock into multiple conformations of comparable
energies in the absence of water, likely enforced by the
hydrogen bonding interactions on the interior of the
compound. Thus, exposure of an anhydrous sample of 7 in
CDCl3 to air for a short period of time results in broadening of
the spectrum and ultimately reversion to spectra recorded in as-
received CDCl3. Although broad resonances are present in
spectra of the di(NiOH) complex 8, we observe resonances in
the range −3.5 to −5 ppm, which are typical for terminal
nickel-hydroxides.10a,d,30 In support of our hypothesis that CO2
can be sequestered by a transient di(NiOH) complex in the
synthesis of the bicarbonate-containing complex 7, exposure of
an anhydrous sample of the di(NiOH) complex 8 in THF to
CO2 for 1 h results in changes in the spectrum that are
consistent with formation of 7 as well as free ligand
(Supporting Information Figure S24). Similarly, exposure of a
sample of the di(NiOH) complex 8 as a solid to air for 12 h
afforded the bicarbonate complex 7, albeit with a similar
X-ray quality crystals of 7 were isolated as red blocks from
diffusions of diethyl ether into tetrahydrofuran solutions of the
reaction mixture (Figure 10). As expected, the complex
incorporates a bicarbonate ligand on one nickel center;
however, we were surprised that the ligand here is only
partially metalated, with two nickel cations held within an
N,N,N chelate, and the third pyDCA arm remains unmetalated.
The bicarbonate ligand is hydrogen bonded to the hydroxide
ligand on the neighboring nickel ion as well as to a water
molecule within the central cavity. Second coordination sphere
hydrogen bonding interactions have been previously incorpo-
rated into monometallic copper pyDCA complexes. In these
systems, hydrogen bond donors were appended to the pyDCA
amide groups and interact with the terminal chloride
coordinated to the metal center.27 Unlike these monometallic
complexes, hydrogen bonding interactions in 7 result from
partial metalation of the cryptate, incorporates solvent
molecules, and extends throughout the internal cavity,
beginning with the metal-bound bicarbonate and hydroxide
and terminating with the amide protons H9 and H7 on the
unmetalated pyDCA arm. We were unable to locate the
bicarbonate proton in the difference map, but it was instead
calculated on O14 on the basis of the following observations.
First, the C90−O bond distances are comparable to those
reported for the bicarbonatonickel(II) pyDCA compound
reported by Holm and co-workers, and the C90−O14 distance
is longer than that for C90−O15.10a Second, the protons on the
water molecule within the cavity and on the hydroxide ligand
were located and freely refined, and the orientations of these H
atoms and the distance between O14 and O18 are consistent
with the presence of a hydrogen bond. Third, charge balance
considerations favor a bicarbonate donor. Although there are
numerous prior examples of polynuclear nickel(II) complexes
2−
−
in which either CO3 or HCO3 coordinate in a bridging
fashion,28 this complex represents a unique case in which both a
terminal Ni−OH and an Ni−OC2OH coexist in the same
complex. It is evident that the ligand precludes a bridging
coordination mode for the bicarbonate donor as the cryptand
enforces a limited range of possible metal−metal distances in
the N,N,N-coordination mode.
A number of mononuclear hydroxide complexes sequester
carbon dioxide as bicarbonate or carbonate with subsequent
release of CO2 either by sparging with an inert gas or by
degassing a solution of the bicarbonate complex.10a−c,28h,29
These compounds parallel the reactivity of carbonic anhydrase,
in which a terminal hydroxide bound to the active site ZnII
captures CO2, which then dissociates as bicarbonate.1l We
therefore investigated if release of CO2 could be accomplished
from 7. To our surprise, sparging solutions of 7 with N2,
repeated cycles of freeze−pump−thaw, or heating a solid
sample under dynamic vacuum (<10 mT) resulted in no
1
amount of the free ligand as determined by H NMR spectra
(Supporting Information Figure S25). As anticipated then, one
of the terminal nickel-hydroxide sites within 8 is competent for
CO2 capture. This result also suggests that whereas the
bicarbonate complex 7 is stable in wet NMR solvents, the
di(NiOH) complex is more sensitive and undergoes
demetalation more readily.
Demonstrating the generality of this synthetic approach to
the dimetalated compounds, reaction of 3 equiv of copper(II)
triflate with H6L followed by excess Et4NOH afforded
[Et4N]2[Cu3(μ-η1:η2-CO3)(μ-OH)(OH) H2L], 9, in good
yield (40%). The molecular structure of this copper complex
is similar to that of the bicarbonate-containing compound 7 as
1
change in the H NMR or IR spectra (Figure S21, Supporting
Information). Evidently then, CO2 release in our multimetallic
system is more energetically demanding than the comparable
monometallic complexes, and we attribute this difference to the
additional energy required to break the hydrogen bonding
network. This result provides strong evidence that incorporat-
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Figure 11. Solid state structure of 9 collected at 100 K (a) and cut away of the central cavity (b). C, H, N, O, and Cu atoms are represented as gray,
light gray, blue, red, and yellow-green ellipsoids (30%), respectively. All hydrogens have been removed except amide and hydroxide hydrogens (all
shown H positions are calculated) along with solvents of crystallization, two tetraethylammonium cations, and the OEt and iPr substituents.
there are two Cu2+ ions coordinated in an N,N,N-chelate of two
arms of the cryptand (Figure 11). However, a third metal ion is
present within the internal cavity of 9 and is held within a
distorted square planar coordination environment comprising
the hydroxide ligated to Cu2, η2-coordination of the carbonate
on Cu1, and an additional hydroxide donor. This structure was
not of the appropriate quality to refine the positions of H atoms
on O16, O17, N7, or N9, so these positions were calculated
and included on the basis of charge balance and the observed
bond lengths. In particular, the Cu2−O17 distance is 1.971(2)
Å, the Cu3−O17 distance is 1.852(2) Å, and the Cu3−O16
distance is 1.9773(17) Å. Cu−O bonds lengths (∼1.90 Å) in
the hydroxo-bridged copper-pyDCA complex reported by
Tolman10j are comparable to those of 9. Elongation of the
Cu2−O17 (1.971(2) Å) and Cu3−O16 (1.9773(17) Å) is
observed; however, only 73% occupancy of Cu2 is observed as
this metal center is disordered over two positions in the
structure solution. Square planar copper(II) complexes for
which a carbonate coordinates in an η2 manner have only been
reported previously for N atom donor ligands.28b,31 The subtle
differences that generate the trimetallic species 9 as compared
with the dimetallic 7 are unclear, but the greater basicity of the
terminal hydroxides of copper-pyDCA complexes as compared
to nickel-pyDCA compounds may favor coordination of the
excess Cu(II) in solution.
We initially expected from reports on mono-pyDCA
complexes that substituting Et4NOH for stronger bases would
yield complexes with each doubly deprotonated pyDCA arm
coordinating a metal center. For both copper(II) and nickel(II),
however, only two of the three pyDCA donor arms are
metalated. One hypothesis is that specificity for partial
metalation arises from a balance between the increasing pKa
values for each subsequent amide proton and that of water.31d
In particular for these hydroxide-decorated complexes, an
electrostatic component could contribute to the partial
metalation as each anionic hydroxide donor is oriented toward
the internal cavity, thereby disfavoring both deprotonation and
approach of weaker bases, and the association of countercations
with the partially metalated complexes could contribute a steric
effect. Interestingly, the ability to use solvent and choice of base
to target these dimetallic complexes is analogous to recent
reports by Powers et al. in which solvent choice biases the
metalation of a hexadentate ligand to afford either trinuclear or
dinuclear complexes. Such subtle strategies may therefore be
common approaches to tuning the nuclearity of metal clusters
templated by multidentate ligands.32
OUTLOOK
■
The macrobicyclic cryptand, H6L, offers an unusual utilitarian
platform to assemble multimetallic compounds as the donor
atom types can be readily tuned, hydrogen bonding interactions
incorporated, and the extent of metalation controlled. The
results presented here point to an ability to rationally modulate
all of these parameters using base strength, choice of metal
precursor, and solvent identity. Initial results point to clear
divergence from reaction outcomes seen for monometallic
complexes with different reactivity patterns and access to
unique structural motifs.
EXPERIMENTAL SECTION
■
General Considerations. Reactions necessitating dry, air free
conditions were performed using standard Schlenk techniques or in an
Innovative Technologies glovebox. Anhydrous solvents were extracted
from an Innovative Technologies solvent purification system.
Anhydrous divalent metal halides were purchased from either Strem
Chemical or Sigma-Aldrich and dried first with SOCl2 and then heated
under vacuum. Both potassium and sodium hexamethyldisilazide were
purchased from Sigma-Aldrich and used as received. Potassium
hydride was purchased from Sigma-Aldrich and was rinsed with
anhydrous hexanes prior to use. 18-Crown-6 was gently heated under
vacuum prior to storage in the glovebox. NMR spectra were recorded
on either a 500 MHz Inova or 300 MHz Mercury spectrophotometer
with the spectra referenced to the residual protonated solvent signal,
7.27 ppm for CDCl3 and 3.58 for THF-d8. Deuterated solvents were
purchased from Cambridge Isotope Laboratories and when necessary
dried according to standard procedures. Infrared spectra were recorded
as solids on a Bruker Vertex 80v FTIR using a Pike GladiATR stage, or
were collected under air-free conditions on a Bruker Alpha FTIR
equipped with a diamond ATR stage inside the glovebox. Cyclic
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voltammetry was performed under a nitrogen atmosphere using a
standard three electrode setup. Electrodes were purchased from either
BASi, Inc., or CH Instruments, Inc. Potential sweeps were controlled
by a Princeton Applied Research Versastat II potentiostat. Elemental
analyses were performed by Complete Analysis Laboratories, Inc.
(Parsippany, NJ). The bis(acetonitrile) adducts of the metal
trifluoromethanesulfonate salts were synthesized by reacting a slurry
of metal dihalide in anhydrous air-free acetonitrile with 2 equiv of
trimethylsilyltrifluoromethanesulfonate (Me3SiOTf) for 2 h, and the
products recrystallized by diffusion of diethyl ether into a saturated
MeCN solution. Coordinated solvent molecules were removed by
heating under vacuum. Tetraethylammonium hydroxide solution (1.5
M in methanol) was purchased from Sigma-Aldrich. The solution was
transferred from the reagent bottle into a storage flask by cannula and
stored under an inert atmosphere, and aliquots for reactions were
removed using a syringe and under a positive pressure of dinitrogen.
Mass spectra were collected by direct injection using either negative or
position mode and an appropriate anhydrous solvent. Samples were
input into the instrument using gastight syringes and a syringe pump
via HPLC tubing. The tubing was prerinsed with anhydrous solvent
also in a gastight syringe before introduction of the sample.
X-ray Crystallography. X-ray intensity data were collected at 100
K on a Bruker DUO diffractometer using Mo Kα radiation (λ =
0.710 73 Å) or Cu Kα radiation (λ = 1.541 78 Å) from an ImuS power
source, and an APEXII CCD area detector. Raw data frames were read
by the SAINT3 program and integrated using 3D profiling algorithms.
The resulting data were reduced to produce hkl reflections and their
intensities and estimated standard deviations. The data were corrected
for Lorentz and polarization effects, and numerical absorption
corrections were applied on the basis of indexed and measured
faces. The structures were solved and refined in SHELXTL6.1, using
full matrix least-squares refinement. The non-H atoms were refined
with anisotropic thermal parameters, and all of the H atoms were
calculated in idealized positions and refined riding on their parent
atoms unless otherwise stated. Details specific to data refinement and
structure solutions for each compound are reported in the Supporting
Information.
Reaction of Ph3P with Iodosobenzene and 2. A portion of 2
(20 mg, 0.0083 mmol) was dissolved in MeCN (7 mL), and
triphenylphosphine (1, 2, or 3 equiv to 2) was added. Then,
iodosobenzene (1, 2, or 3 equiv to 2) was added to the reaction, and
the mixture stirred at room temperature for 6 h. The solvent was
removed under reduced pressure and the residue redissolved with
vigorous stirring in deuterated benzene. The sample was filtered
through Celite, and 31P NMR spectra were collected using 85% H3PO4
as the reference. The ratio of the integrands for peaks corresponding
to triphenylphosphine and triphenylphosphineoxide were determined.
[K(18-Crown-6)]2[K(CoIICl)3L] (3). The tricobalt complex was
synthesized using the method described for 2 with the following
observed differences. Addition of CoCl2 to the deprotonated ligand
afforded a green reaction mixture, and the crude product was
recrystallized from diffusion of diethyl ether into a concentrated
solution of THF yielding the desired product as a green microcrystal-
line powder in 92% yield. IR: 2865 (m), 1573 (s), 1453 (s), 1351 (s)
1103 (s) cm−1. Evans method (CD3CN): μeff = 6.81(3) μB. HRMS-
ESI(−): m/z 902.1948, calcd for (C89H95N9O12CoII3Cl3K)2−
902.1907. Anal. Calcd for C113H143Cl3Co3K3N9O24: C, 56.27; H,
5.98; N, 5.23. Found: C, 56.12; H, 5.88; N, 5.25.
[18-Crown-6-K][(Ni(THF))2(NiCl)L] (4). A sample of H6L (38 mg,
0.026 mmol) was dissolved in THF (12 mL), and KH (6.7 mg, 0.17
mmol) was added to the solution. The reaction was stirred for 1 h and
then filtered through Celite. 18-Crown-6 (7 mg, 0.03 mmol) was then
added to the yellow solution followed by NiCl2·1.5THF (20 mg, 0.084
mmol). The reaction was stirred for 48 h at 50 °C to afford a red-
brown slurry, which was then filtered through Celite and the filtrate
dried under vacuum. The brown residue was recrystallized by diffusion
of diethyl ether into a THF solution of the residue affording brown
crystals. The crystals were dried for 4 h at 50 °C under reduced
pressure (34 mg, 62% yield). IR: 2961 (m), 1630 (s), 1590 (m), 1449
(s), 1028 (s), 901 (m) cm−1. HRMS-ESI(−): m/z 873.7324, calcd. for
[C89H97N9Cl2Ni3O13]2− 873.7311. Anal. Calcd for C109H135N9Ni3O20:
C, 61.12; H, 6.35; N, 5.89. Found: C, 61.01; H, 6.37; N, 5.84.
[(Fe(MeCN))3(μ3-O)LH5][OTf]4 (5). A 100 mL round-bottom flask
was charged with H6L (0.882 g, 0.592 mmol) and THF (40 mL). A
50% dispersion of KH (0.157 g, 0.195 mmol) was added at room
temperature causing the solution to change from colorless to yellow
with the release of hydrogen gas. The reaction was stirred for 1 h at
ambient temperature after which Fe(OTf)2·2MeCN (0.853 g, 1.95
mmol) was added in one portion. The dark intensely colored reaction
mixture was stirred for 12 h after which the solvent was removed
under reduced pressure. The residue was dissolved in a minimum of
acetonitrile (10 mL) and the solution filtered through Celite. Toluene
diffusion into the acetonitrile solution over the course of 2 days
resulted in crystallization of KOTf (263 mg), which was removed by
filtration. 18-Crown-6 (150 mg, 0.568 mmol) was added to the
solution, which was then diluted with toluene and placed in a freezer at
−40 °C for 2 h. The flocculent, purple precipitate was collected by
filtration, redissolved in a minimum of MeCN (8 mL), and diffusion of
toluene into this solution afforded the target complex as black crystals
(1.111 g, 79%). IR: 2964 (m), 1631 (s), 1556 (s), 1456 (s), 1209 (s)
1028 (s), 636 (s) cm−1. Anal. Calcd for C99H110F12Fe3N12O25S4: C,
49.71; H, 4.64; N, 7.03. Found: C, 49.72; H, 4.72; N, 6.92.
[K(18-Crown-6)]2[K(CuIICl)3L] (1). A portion of H6L (200 mg,
0.134 mmol) was dissolved in THF (∼8 mL) and cooled to −78 °C. A
solution of KHMDS (177 mg, 0.887 mmol) was made in THF (2 mL)
and also cooled to −78 °C. The KHMDS solution was then added
dropwise to the solution of H6L causing an immediate change from a
colorless solution to yellow. The reaction was stirred for approximately
30−60 min followed by the addition of a solution of 18-crown-6 (107
mg, 0.403 mmol) in 0.5 mL of THF. CuCl2 (60 mg, 0.44 mmol) was
then added to the reaction. The reaction immediately took on a deep
green color. The reaction was warmed to room temperature and
stirred overnight. The reaction was then filtered, and the solvents were
removed under reduced pressure. The residue was taken up in a
minimum of THF (2 mL) and the complex precipitated with diethyl
ether (15 mL). The residue was collected by filtration, suspended, and
stirred in diethyl ether for 8 h, and recollected by filtration to afford a
green powder (127 mg, 39% yield). X-ray quality crystals were grown
by slow diffusion of dimethoxyethane into a tetrahydrofuran solution.
IR: 2866 (m), 1585 (s), 1453 (s), 1208 (s), 1102 (s), 898 (m) cm−1.
HRMS-ESI(−): m/z 909.1862, calcd for [C89H95N9O12CuII Cl3K]2−
Reaction of 5 with Methyl Trifluoromethylsulfonate.
Complex 5 (54 mg, 0.023 mmol) was dissolved in THF (8 mL),
and methyl trifluoromethanesulfonate (2.8 μL, 0.025 mmol) was
added. The reaction was stirred overnight after which the solvent was
removed under reduced pressure to yield a tacky black solid. This
residue was removed from the anaerobic chamber, redissolved in
dichloromethane (8 mL), and hydrolyzed and demetalated by washing
with an aqueous solution containing a large excess of Na2EDTA (8
mL). The organic layer was separated, dried with magnesium sulfate,
and the solvent removed under reduced pressure yielding 25 mg of off-
white powder. NMR spectra collected on the residue in CDCl3
indicated that the isolated compound was the nonalkylated, free
ligand H6L. Isolated mass corresponds to 46% recovered yield.
[Co3(μ-OH)(μ3-OH)(μ-OTf)(OTf)LH6][OTf]2 (6). H6L (203 mg,
0.136 mmol) was dissolved in THF (∼10 mL), and KH (18 mg, 0.43
3
909.1845. Anal. Calcd for C113H143Cl3Cu3K3N9O24: C, 55.95; H, 5.94;
N, 5.20; Cu, 7.86; Cl, 4.38; K 4.84. Found: C, 55.98; H, 6.01; N, 5.04;
Cu, 8.02; Cl, 4.28; K, 4.53.
[K(18-Crown-6)]2[K(FeIICl)3L] (2). This compound was synthe-
sized using the method described for 1 with the following exceptions.
Addition of FeCl2 to the ligand deprotonated with 6.6 equiv of KH
afforded a red/brown reaction mixture, and the crude product was
recrystallized from diffusion of diethyl ether into a concentrated
solution of THF yielding the desired product as a red/brown powder
in 78% yield. IR: 2866 (m), 1575 (s), 1456 (s), 1351 (s) 1103 (s)
cm−1. Evans’ method (CD3CN): μeff = 7.83(4). HRMS-ESI(−): m/z
897.6907, calcd for [C89H95N9O12FeII Cl3K]2− 897.6931. Anal. Calcd
3
for C113H143Cl3Fe3K3N9O24: C, 56.49; H, 6.00; N, 5.25. Found: C,
56.32; H, 6.35; N, 5.25.
K
DOI: 10.1021/ic502873d
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
mmol) was added in one portion. Once effervescence ceased,
Co(O3SCF3)2·2MeCN (190 mg, 0.432 mmol) was added to the
pale yellow reaction mixture, which immediately became a mahogany
colored mixture. The reaction was stirred overnight, and then was
filtered, and the solvent removed from the filtrate under reduced
pressure. The residue was dissolved in a minimum of MeCN and
toluene diffused into the solution to afford colorless crystals of KOTf,
which were separated from the mother liquor by filtration. This
method of removing KOTf from the crude product was repeated a
second time. 18-Crown-6 (10 mg, 0.38 mmol) was then added to the
mother liquor, followed by toluene. The mixture was filtered and the
desired product isolated as a green solid. Storing the filtrate at −32 °C
yielded a second crop of the product. The isolated compound was
dried under reduced pressure yielding a green solid (162 mg, 51%).
IR: 2964 (m), 1633 (s), 1556 (s), 1461 (s), 1209 (s) 1028 (s), 634 (s)
cm−1. Evans method (CD3CN): μeff = 7.54(3) μB. Anal. Calcd for
C95H106F12Co3N10O26S4: C, 48.82; H, 4.57; N, 5.99. Found: C, 48.82;
H, 4.39; N, 5.93.
traps immersed in a slurry of liquid nitrogen and chloroform. THF was
removed under reduced pressure, and the red solid was then exposed
to the atmosphere and dissolved in CDCl3. NMR was collected on the
residue indicating the presence of free ligand and 7.
[Et4N]2[Cu3(μ-η1:η2-OCOO)(μ-OH)(OH)H2L] (9). Synthesis of the
target complex was carried out on the benchtop with exposure to
ambient air. Cu(OTf)2 (79 mg, 0.22 mmol) was added to a solution of
H6L (102 mg, 0.0683 mmol) in THF (10−15 mL), after which
tetraethylammonium hydroxide (1.5 M in methanol, 0.424 mL, 0.636
mmol) was added dropwise to the reaction mixture. The dark greenish
blue reaction mixture was stirred overnight and then filtered through
Celite followed by solvent removal from the filtrate under vacuum.
The residue was redissolved in dichloromethane (10−15 mL) and
extracted with 1 M NaOH (2 × 10 mL). The organic layer was dried
with magnesium sulfate, filtered through Celite, and then dried under
reduced pressure. The solid was redissolved in THF, and n-hexane was
diffused into the solution to afford green crystals after 24 h (54.8 mg,
40% yield). IR: 2964 (m), 1672 (m), 1625 (m) 1579 (s), 1450 (s),
1031 (s) cm−1. Evans method (CDCl3): μeff = 2.80(1) μB. HRMS-ESI:
m/z 893.2361, calcd for [C90H99N9Cu3O17, −OH, +Cl]2− 893.2360.
Anal. Calcd for C107H142N11Cl3Cu3O18: C, 59.30; H, 6.60; N, 7.11.
Found: C, 58.97; H, 6.60; N, 6.97.
[Et4N]2[Ni(OCOOH)Ni(OH)H2L] (7). Synthesis of the target
complex was carried out on the benchtop and exposed to ambient
air. To a solution of H6L (51 mg, 0.034 mmol) in THF (10 mL) was
added Ni(OTf)2·2MeCN (45 mg, 0.10 mmol). Tetraethylammonium
hydroxide in methanol (1.5 M, 0.210 mL, 0.316 mmol) was added
dropwise to the pale green reaction mixture. The reaction flask was
capped with a septum, and the resulting red solution was stirred
overnight at room temperature and subsequently filtered through
Celite, and the filtrate dried under reduced pressure. The red solid was
dissolved in dichloromethane (10−15 mL) and washed with 1 M
KOH (2 × 10 mL). The combined organics were dried with
magnesium sulfate, filtered through Celite, and dried overnight under
vacuum at ambient temperature. The solid was redissolved in THF,
and diethyl ether was diffused into the solution to afford red crystals,
which were suitable for X-ray diffraction experiments. Analytically pure
solid was isolated after heating the solid at 75 °C under vacuum for 4 h
(24 mg, 36% yield). Compound 7 could be prepared by an alternate
method using 2 equiv of Ni(OTf)2·2MeCN and 6 equiv of Et4NOH,
and washing the product with 1 M NaOH instead of 1 M M KOH
(40% yield). IR: 2962 (m), 1675 (m), 1617 (s) 1586 (s), 1453 (s),
[Et4N]2[Cu(OCOOH)Cu(OH)H2L] (10). Synthesis of the target
complex was carried out on the benchtop with exposure to ambient
air. Cu(OTf)2 (49 mg, 0.14 mmol) was added to a solution of H6L
(101 mg, 0.0676 mmol) in THF (12 mL) followed by
tetraethylammonium hydroxide (1.5 M in methanol, 0.413 mL,
0.636 mmol) to afford a dark green mixture. The reaction flask was
capped with a septum, and the mixture was stirred overnight and then
filtered through Celite, and the filtrate dried under vacuum. The
residue was dissolved in dichloromethane (15 mL) and the deep-green
solution washed with 1 M KOH (2 × 10 mL). The organics were then
dried with magnesium sulfate, the solution filtered through Celite, and
filtrate dried under vacuum overnight. The solid was dissolved in THF,
and n-hexane was diffused into the solution to afford needle-like green
crystals (69.5 mg, 52% yield). IR: 2963 (m), 1681 (s), 1586 (s), 1454
(s), 1033 (s) cm−1. Evans method (CDCl3): μeff = 2.89(1). HRMS-
ESI(−): m/z 844.7916, calcd for [C90H99N9Cu2O16]2− 844.7913. Anal.
Calcd for C106H143N11Cu2O18: C, 64.09; H, 7.26; N, 7.76. Found: C,
63.88; H, 7.20; N, 7.82.
1
1032 (s) cm−1. H NMR (500 MHz, CDCl3) δ ppm 0.11 (m, 57 H),
1.21 (m, 22 H), 1.27 (d, J = 6.73 Hz, 8 H), 2.39 (br s, 12 H), 2.84 (m,
3 H), 2.98 (br q, J = 7.10 Hz, 16 H), 3.13 (m, 4 H), 3.27 (m, 6 H),
3.51 (spt, J = 7.50 Hz, 2 H), 3.66 (spt, J = 7.60 Hz, 2 H), 3.74 (m, 2
H), 6.56 (s, 2 H), 6.59 (d, J = 1.92 Hz, 2 H), 6.79 (d, J = 2.33 Hz, 2
H), 6.81 (d, J = 2.33 Hz, 2 H), 6.91 (d, J = 2.33 Hz, 2 H), 7.00 (d, J =
2.33 Hz, 2 H), 7.51 (d, J = 7.69 Hz, 2 H), 7.54 (d, J = 7.69 Hz, 2 H),
7.70 (d, J = 2.47 Hz, 2 H), 7.78 (t, J = 7.69 Hz, 1 H), 7.84 (t, J = 7.69
Hz, 1 H), 7.99 (t, J = 7.76 Hz, 1 H), 8.27 (d, J = 7.82 Hz, 2 H), 11.32
(s, 2 H). 13C NMR (126 MHz, CDCl3) δ ppm 7.38, 15.38−15.88,
23.57, 23.71, 24.20, 24.43, 24.64, 26.33−26.60, 30.92, 40.19, 53.04,
68.70, 69.73, 69.92, 119.05, 121.45, 121.94, 122.33, 123.86, 125.20,
125.57, 125.67−125.87, 125.91−126.19, 126.69, 132.88, 137.83,
138.03, 138.27, 138.72, 139.09, 140.83, 140.99, 141.38, 141.68,
141.90, 142.17, 150.36−150.52, 151.84−151.94, 152.08, 152.34,
152.57, 159.35, 163.66, 167.72. HRMS-ESI(−): m/z 839.7996, calcd
Magnetic Measurements and Data Simulations. Sample
preparation, data collection, and simulations were carried out as
described previously.9b In these simulations, which employed codes
written to work with EasySpin22 running in Matlab, a correction that
accounts for underlying temperature independent contributions χdia
and χTIP was included with value −0.001 38 emu/mol, according to χdia
estimated from Pascal’s constants and χTIP based on reports
elsewhere.33 In addition, a paramagnetic impurity contribution was
included in susceptibility and magnetization with Sfree = 2, gfree = 2.03
in form of (1 − A)χ5 + Aχfree with a value of 0.01 for A.
Mossbauer Spectroscopy. Mossbauer spectra were recorded at
̈
4.2 K, either on a low field Mossbauer spectrometer equipped with a
̈
Janis SVT-400 cryostat or on a strong field Mossbauer spectrometer
̈
equipped with an Oxford Instruments Spectromag 4000 cryostat
containing an 8 T split-pair superconducting magnete4. Both
spectrometers were operated in a constant acceleration mode in
transmission geometry. The isomer shifts were referenced against that
of a room temperature metallic iron foil. Analysis of the data was
performed with the program WMOSS (WEB Research).
for [C90H99N9Ni2O16]2− 839.7964. Anal. Calcd for C106H143
-
N11Ni2O18: C, 64.41; H, 7.29; N, 7.79. Found: C, 64.29; H, 7.26; N,
7.69.
[Et4N]2[(NiOH)2H2L] (8). A solution of H6L (55 mg, 0.037 mmol)
in THF (10 mL) was sparged with N2 for 30 min, and Ni(OTf)2·
2MeCN (36 mg, 0.081 mmol) was added, forming a pale green slurry.
Tetraethylammonium hydroxide in methanol (1.5 M, 0.153 mL, 0.229
mmol) was added dropwise to the pale green reaction mixture. Then,
the resulting red solution was stirred overnight at room temperature
and was dried under reduced pressure for 18 h. The red solid was
brought into the glovebox where it was dissolved in THF, filtered
through Celite, and dried under reduced pressure.
ASSOCIATED CONTENT
* Supporting Information
Experimental details, supporting figures, and X-ray crystallo-
graphic data in CIF format. This material is available free of
charge via the Internet at http://pubs.acs.org.
■
S
Reaction of 8 with CO2. Complex 8 was dissolved in THF (30
mL) in a Schlenk flask to afford a red solution and was freeze−pump−
thawed. The solution was stirred for 1 h under a constant pressure of
research grade CO2 that was passed through an O2 scrubber and two
AUTHOR INFORMATION
Corresponding Author
*E-mail: murray@chem.ufl.edu.
■
L
DOI: 10.1021/ic502873d
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
3981. (e) Vela, J.; Stoian, S.; Flaschenriem, C. J.; Munck, E.; Holland,
Present Address
̈
⊥Department of Chemistry and Physics, Armstrong State
University, 11935 Abercorn Street, Savannah, GA 31419.
P. L. J. Am. Chem. Soc. 2004, 126, 4522. (f) Venkateswara Rao, P.;
Holm, R. H. Chem. Rev. 2004, 104, 527. (g) Cady, C. W.; Crabtree, R.
H.; Brudvig, G. W. Coord. Chem. Rev. 2008, 252, 444. (h) Sproviero, E.
Author Contributions
́
M.; Gascon, J. A.; McEvoy, J. P.; Brudvig, G. W..; Batista, V. S. Coord.
The manuscript was written through contributions of all
authors. All authors have given approval to the final version of
the manuscript.
Chem. Rev. 2008, 252, 395. (i) Mukherjee, S.; Stull, J. A.; Yano, J.;
Stamatatos, T. C.; Pringouri, K.; Stich, T. A.; Abboud, K. A.; Britt, R.
D.; Yachandra, V. K.; Christou, G. Proc. Natl. Acad. Sci. U.S.A. 2012,
109, 2257. (j) Lee, S. C.; Lo, W.; Holm, R. H. Chem. Rev. 2014, 114,
3579. (k) Brown, E. C.; York, J. T.; Antholine, W. E.; Ruiz, E.; Alvarez,
S.; Tolman, W. B. J. Am. Chem. Soc. 2005, 127, 13752. (l) York, J. T.;
Bar-Nahum, I.; Tolman, W. B. Inorg. Chem. 2007, 46, 8105.
Notes
The authors declare no competing financial interest.
ACKNOWLEDGMENTS
■
(5) (a) Stack, T. D. P.; Holm, R. H. J. Am. Chem. Soc. 1987, 109,
2546. (b) Zhou, J.; Raebiger, J. W.; Crawford, C. A.; Holm, R. H. J.
Am. Chem. Soc. 1997, 119, 6242. (c) Kanady, J. S.; Lin, P.-H.; Carsch,
K. M.; Nielsen, R. J.; Takase, M. K.; Goddard, W. A., III; Agapie, T. J.
Am. Chem. Soc. 2014, 136, 14373. (d) Tsui, E. Y.; Tran, R.; Yano, J.;
Agapie, T. Nat. Chem. 2013, 5, 293. (e) Tsui, E. Y.; Kanady, J. S.;
Agapie, T. Inorg. Chem. 2013, 52, 13833.
(6) (a) Feig, A. L.; Lippard, S. J. Chem. Rev. 1994, 94, 759.
(b) Wallar, B. J.; Lipscomb, J. D. Chem. Rev. 1996, 96, 2625.
(c) Proshlyakov, D. A.; Pressler, M. A.; DeMaso, C.; Leykam, J. F.;
DeWitt, D. L.; Babcock, G. T. Science 2000, 290, 1588. (d) Jensen, K.
P.; Bell, C. B.; Clay, M. D.; Solomon, E. I. J. Am. Chem. Soc. 2009, 131,
The following institutions are acknowledged: NSF CRIF award
to the University of Florida (UF) Chemistry Department
(CHE-1048604) (L.J.M.); NSF (CHE-0821346) and UF for
funding X-ray equipment purchase (K.A.A.); The University of
Florida (L.J.M., K.A.A.); ACS Petroleum Research Fund 52704-
DNI3 (L.J.M.); Labex ARCANE ANR-11-LABX-0003−01
(R.G.-S.); Fulbright Commission of the Slovak Republic
̌
(M.W.M.); APVV-0132-11 and VEGA 1/0145/13 (E.C.);
National Science Foundation (NSF) awards DMR-1202033
(M.W.M.) and DMR-1157490 (NHMFL). Contributions from
M.K. Peprah and E.S. Knowles during the early stages of the
magnetometry studies are gratefully acknowledged.
12155. (e) Wikstrom, M. Biochim. Biophys. Acta 2012, 1817, 468.
̈
(f) Pinakoulaki, E.; Daskalakis, V.; Ohta, T.; Richter, O.-M. H.;
Budiman, K.; Kitagawa, T.; Ludwig, B.; Varotsis, C. J. Biol. Chem. 2013,
288, 20261.
REFERENCES
■
(7) (a) Cole, A. P.; Root, D. E.; Mukherjee, P.; Solomon, E. I.; Stack,
T. D. P. Science 1996, 273, 1848. (b) Wallar, B. J.; Lipscomb, J. D.
Chem. Rev. 1996, 96, 2625. (c) Du Bois, J.; Mizoguchi, T. J.; Lippard,
(1) (a) Tinberg, C. E.; Lippard, S. J. Acc. Chem. Res. 2011, 44, 280.
(b) Solomon, E. I.; Augustine, A. J.; Yoon, J. Dalton Trans. 2008, 3921.
(c) Solomon, E. I.; Heppner, D. E.; Johnston, E. M.; Ginsbach, J. W.;
Cirera, J.; Qayyum, M.; Kieber-Emmons, M. T.; Kjaergaard, C. H.;
Hadt, R. G.; Tian, L. Chem. Rev. 2014, 114, 3659. (d) Blomberg, M. R.
A.; Borowski, T.; Himo, F.; Liao, R.-Z.; Siegbahn, P. E. M. Chem. Rev.
2014, 114, 3601. (e) Murray, L. J.; Lippard, S. J. Acc. Chem. Res. 2007,
40, 466. (f) Baik, M.-H.; Newcomb, M.; Friesner, R. A.; Lippard, S. J.
Chem. Rev. 2003, 103, 2385. (g) Dey, A.; Jiang, Y.; Ortiz de
Montellano, P.; Hodgson, K. O.; Hedman, B.; Solomon, E. I. J. Am.
Chem. Soc. 2009, 131, 7869. (h) Gherman, B. F.; Baik, M.-H.; Lippard,
S. J.; Friesner, R. A. J. Am. Chem. Soc. 2004, 126, 2978. (i) Denisov, I.
G.; Makris, T. M.; Sligar, S. G.; Schlichting, I. Chem. Rev. 2005, 105,
2253. (j) Gourlay, C.; Nielsen, D. J.; White, J. M.; Knottenbelt, S. Z.;
Kirk, M. L.; Young, C. G. J. Am. Chem. Soc. 2006, 128, 2164.
(k) Jeoung, J.-H.; Dobbek, H. Science 2007, 318, 1461. (l) Domsic, J.
F.; Avvaru, B. S.; Kim, C. U.; Gruner, S. M.; Agbandje-McKenna, M.;
Silverman, D. N.; McKenna, R. J. Biol. Chem. 2008, 283, 30766.
(m) Appel, A. M.; Bercaw, J. E.; Bocarsly, A. B.; Dobbek, H.; DuBois,
D. L.; Dupuis, M.; Ferry, J. G.; Fujita, E.; Hille, R.; Kenis, P. J. A.;
Kerfeld, C. A.; Morris, R. H.; Peden, C. H. F.; Portis, A. R.; Ragsdale,
S. W.; Rauchfuss, T. B.; Reek, J. N. H.; Seefeldt, L. C.; Thauer, R. K.;
Waldrop, G. L. Chem. Rev. 2013, 113, 6621. (n) Can, M.; Armstrong,
F. A.; Ragsdale, S. W. Chem. Rev. 2014, 114, 4149. (o) Burgess, B. K.;
Lowe, D. J. Chem. Rev. 1996, 96, 2983. (p) Holm, R. H.; Kennepohl,
P.; Solomon, E. I. Chem. Rev. 1996, 96, 2239. (q) Hoffman, B. M.;
Lukoyanov, D.; Yang, Z.-Y.; Dean, D. R.; Seefeldt, L. C. Chem. Rev.
2014, 114, 4041.
́
S. J. Coord. Chem. Rev. 2000, 200−202, 443. (d) Kim, E.; Chufan, E.
E.; Kamaraj, K.; Karlin, K. D. Chem. Rev. 2004, 104, 1077. (e) Lewis, E.
A.; Tolman, W. B. Chem. Rev. 2004, 104, 1047. (f) Mirica, L. M.;
Ottenwaelder, X.; Stack, T. D. P. Chem. Rev. 2004, 104, 1013.
(g) Tshuva, E. Y.; Lippard, S. J. Chem. Rev. 2004, 104, 987. (h) Ohi,
H.; Tachi, Y.; Itoh, S. Inorg. Chem. 2006, 45, 10825. (i) Brown, E. C.;
Johnson, B.; Palavicini, S.; Kucera, B. E.; Casella, L.; Tolman, W. B.
Dalton Trans. 2007, 3035. (j) Do, L. H.; Hayashi, T.; Moenne-Loccoz,
̈
P.; Lippard, S. J. J. Am. Chem. Soc. 2010, 132, 1273. (k) Friedle, S.;
Reisner, E.; Lippard, S. J. Chem. Soc. Rev. 2010, 39, 2768. (l) Citek, C.;
Lyons, C. T.; Wasinger, E. C.; Stack, T. D. P. Nat. Chem. 2012, 4, 317.
(m) Gupta, A. K.; Tolman, W. B. Inorg. Chem. 2012, 51, 1881. (n) Li,
Y.; Soe, C. M. M.; Wilson, J. J.; Tuang, S. L.; Apfel, U.-P.; Lippard, S. J.
Eur. J. Inorg. Chem. 2013, 2013, 2011.
(8) (a) Zinn, P. J.; Powell, D. R.; Day, V. W.; Hendrich, M. P.;
Sorrell, T. N.; Borovik, A. S. Inorg. Chem. 2006, 45, 3484. (b) Zinn, P.
J.; Sorrell, T. N.; Powell, D. R.; Day, V. W.; Borovik, A. S. Inorg. Chem.
2007, 46, 10120. (c) Shook, R. L.; Borovik, A. S. Chem. Commun.
2008, 6095. (d) Graef, T.; Galezowska, J.; Dechert, S.; Meyer, F. Eur. J.
Inorg. Chem. 2011, 2011, 4161. (e) Ng, G. K.-Y.; Ziller, J. W.; Borovik,
A. S. Inorg. Chem. 2011, 50, 7922. (f) Ng, G. K.-Y.; Ziller, J. W.;
Borovik, A. S. Chem. Commun. 2012, 48, 2546.
(9) (a) Guillet, G. L.; Sloane, F. T.; Dumont, M. F.; Abboud, K. A.;
Murray, L. J. Dalton Trans. 2012, 41, 7866. (b) Guillet, G. L.; Sloane,
F. T.; Ermert, D. M.; Calkins, M. W.; Peprah, M. K.; Knowles, E. S.;
̌
̌
́
Cizmar, E.; Abboud, K. A.; Meisel, M. W.; Murray, L. J. Chem.
(2) (a) Solomon, E. I.; Sundaram, U. M.; Machonkin, T. E. Chem.
Rev. 1996, 96, 2563. (b) Liu, J.; Chakraborty, S.; Hosseinzadeh, P.; Yu,
Y.; Tian, S.; Petrik, I.; Bhagi, A.; Lu, Y. Chem. Rev. 2014, 114, 4366.
(c) Siegbahn, P. E. M.; Blomberg, M. R. A. Chem. Rev. 2010, 110,
7040.
(3) (a) Nagano, S.; Poulos, T. L. J. Biol. Chem. 2005, 280, 31659.
(b) Dempsey, J. L.; Winkler, J. R.; Gray, H. B. Chem. Rev. 2010, 110,
7024. (c) Kaila, V. R. I.; Verkhovsky, M. I.; Wikstrom, M. Chem. Rev.
Commun. 2013, 49, 49.
(10) (a) Huang, D.; Holm, R. H. J. Am. Chem. Soc. 2010, 132, 4693.
(b) Huang, D.; Makhlynets, O. V.; Tan, L. L.; Lee, S. C.; Rybak-
Akimova, E. V.; Holm, R. H. Proc. Natl. Acad. Sci. U.S.A. 2011, 108,
1222. (c) Huang, D.; Makhlynets, O. V.; Tan, L. L.; Lee, S. C.; Rybak-
Akimova, E. V.; Holm, R. H. Inorg. Chem. 2011, 50, 10070. (d) Zhang,
X.; Huang, D.; Chen, Y.-S.; Holm, R. H. Inorg. Chem. 2012, 51, 11017.
(e) Donoghue, P. J.; Gupta, A. K.; Boyce, D. W.; Cramer, C. J.;
Tolman, W. B. J. Am. Chem. Soc. 2010, 132, 15869. (f) Donoghue, P.
J.; Tehranchi, J.; Cramer, C. J.; Sarangi, R.; Solomon, E. I.; Tolman, W.
B. J. Am. Chem. Soc. 2011, 133, 17602. (g) Halvagar, M. R.; Tolman,
W. B. Inorg. Chem. 2013, 52, 8306. (h) Tehranchi, J.; Donoghue, P. J.;
̈
2010, 110, 7062.
(4) (a) Christou, G. Acc. Chem. Res. 1989, 22, 328. (b) Yachandra, V.
K.; Sauer, K.; Klein, M. P. Chem. Rev. 1996, 96, 2927. (c) Lee, S. C.;
Holm, R. H. Chem. Rev. 2004, 104, 1135. (d) Mukhopadhyay, S.;
Mandal, S. K.; Bhaduri, S.; Armstrong, W. H. Chem. Rev. 2004, 104,
M
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Cramer, C. J.; Tolman, W. B. Eur. J. Inorg. Chem. 2013, 2013, 4077.
(i) Halvagar, M. R.; Neisen, B.; Tolman, W. B. Inorg. Chem. 2013, 52,
793. (j) Halvagar, M. R.; Solntsev, P. V.; Lim, H.; Hedman, B.;
Hodgson, K. O.; Solomon, E. I.; Cramer, C. J.; Tolman, W. B. J. Am.
Chem. Soc. 2014, 136, 7269. (k) Noveron, J. C.; Olmstead, M. M.;
Mascharak, P. K. J. Am. Chem. Soc. 2001, 123, 3247. (l) Marlin, D. S.;
Olmstead, M. M.; Mascharak, P. K. Angew. Chem., Int. Ed. 2001, 40,
4752.
Gambarotta, S.; Yap, G. P. A.; Thompson, L. K.; Herbst-Irmer, R.
Chem.Eur. J. 1999, 5, 577. (e) Korobkov, I.; Gambarotta, S.; Yap, G.
P. A. Organometallics 2001, 20, 2552.
(27) (a) Shirin, Z.; Thompson, J.; Liable-Sands, L.; Yap, G. P. A.;
Rheingold, A. L.; Borovik, A. S. J. Chem. Soc., Dalton Trans. 2002,
1714. (b) Wang, D.; Lindeman, S. V.; Fiedler, A. T. Eur. J. Inorg. Chem.
2013, 2013, 4473.
(28) (a) Rawle, S. C.; Harding, C. J.; Moore, P.; Alcock, N. W. J.
Chem. Soc., Chem. Commun. 1992, 1701. (b) Kitajima, N.; Hikichi, S.;
Tanaka, M.; Morooka, Y. J. Am. Chem. Soc. 1993, 115, 5496.
(c) Escuer, A.; Vicente, R.; Kumar, S. B.; Solans, X.; Font-Bardía, M.;
Caneschi, A. Inorg. Chem. 1996, 35, 3094. (d) Escuer, A.; Fallah, M. S.;
Kumar, S. B.; Mautner, F.; Vicente, R. Polyhedron 1998, 18, 377.
(e) Escuer, A.; Vicente, R.; Kumar, S. B.; Mautner, F. A. J. Chem. Soc.,
(11) (a) Barrett, A. G. M.; Crimmin, M. R.; Hill, M. S.; Hitchcock, P.
B.; Kociok-Kohn, G.; Procopiou, P. A. Inorg. Chem. 2008, 47, 7366.
̈
(b) Appelt, C.; Slootweg, J. C.; Lammertsma, K.; Uhl, W. Angew.
Chem., Int. Ed. 2012, 51, 5911. (c) Horn, B.; Limberg, C.; Herwig, C.;
Feist, M.; Mebs, S. Chem. Commun. 2012, 48, 8243. (d) Thirumoorthi,
R.; Chivers, T. Eur. J. Inorg. Chem. 2012, 2012, 3061. (e) Moore, M.;
Gambarotta, S.; Bensimon, C. Organometallics 1997, 16, 1086.
(f) Chisholm, M. H.; Iyer, S. S.; Streib, W. E. New J. Chem. 2000,
24, 393. (g) Andrews, P. C.; Armstrong, D. R.; Raston, C. L.; Roberts,
B. A.; Skelton, B. W.; White, A. H. J. Chem. Soc., Dalton Trans. 2001,
Dalton Trans.. 1998, 3473. (f) Lozano, A. A.; Sae
́ ́
z, M.; Perez, J.;
García, L.; Lezama, L.; Rojo, T.; Lopez, G.; García, G.; Santana, M. D.
́
Dalton Trans. 2006, 3906. (g) Wikstrom, J. P.; Filatov, A. S.;
Mikhalyova, E. A.; Shatruk, M.; Foxman, B. M.; Rybak-Akimova, E. V.
Dalton Trans. 2010, 39, 2504. (h) Martínez-Prieto, L. M.; Real, C.;
́
996. (h) Chen, X.; Lim, S.; Plecnik, C. E.; Liu, S.; Du, B.; Meyers, E.
́
́
Avila, E.; Alvarez, E.; Palma, P.; Cam
́
pora, J. Eur. J. Inorg. Chem. 2013,
A.; Shore, S. G. Inorg. Chem. 2005, 44, 6052. (i) Layfield, R. A.; García,
F.; Hannauer, J.; Humphrey, S. M. Chem. Commun. 2007, 5081.
(12) (a) Korendovych, I. V.; Staples, R. J.; Reiff, W. M.; Rybak-
Akimova, E. V. Inorg. Chem. 2004, 43, 3930. (b) Korendovych, I. V.;
Kryatova, O. P.; Reiff, W. M.; Rybak-Akimova, E. V. Inorg. Chem.
2007, 46, 4197.
2013, 5555.
(29) (a) Looney, A.; Han, R.; McNeill, K.; Parkin, G. J. Am. Chem.
Soc. 1993, 115, 4690. (b) Company, A.; Jee, J.-E.; Ribas, X.; Lopez-
Valbuena, J. M.; Gom
Buchholz, J.; Costas, M.; van Eldik, R. Inorg. Chem. 2007, 46, 9098.
(c) Troeppner, O.; Huang, D.; Holm, R. H.; Ivanovic-Burmazovic, I.
́
ez, L.; Corbella, M.; Llobet, A.; Mahía, J.; Benet-
́
́
(13) (a) de Bruin, B.; Bill, E.; Bothe, E.; Weyhermuller, T.;
̈
Dalton Trans. 2014, 43, 5274. (d) Looney, A.; Parkin, G.; Alsfasser, R.;
Ruf, M.; Vahrenkamp, H. Angew. Chem., Int. Ed. 1992, 31, 92.
(e) Zhang, X.; van Eldik, R.; Koike, T.; Kimura, E. Inorg. Chem. 1993,
32, 5749. (f) Bazzicalupi, C.; Bencini, A.; Bencini, A.; Bianchi, A.;
Corana, F.; Fusi, V.; Giorgi, C.; Paoli, P.; Paoletti, P.; Valtancoli, B.;
Zanchini, C. Inorg. Chem. 1996, 35, 5540.
Wieghardt, K. Inorg. Chem. 2000, 39, 2936. (b) Gibson, V. C.;
Redshaw, C.; Solan, G. A. Chem. Rev. 2007, 107, 1745.
(14) Taktak, S.; Ye, W.; Herrera, A. M.; Rybak-Akimova, E. V. Inorg.
Chem. 2007, 46, 2929.
(15) Dinger, M. B.; Scott, M. J. J. Chem. Soc., Perkin Trans. 1 2000,
1741.
(30) Schmeier, T. J.; Nova, A.; Hazari, N.; Maseras, F. Chem.Eur. J.
(16) (a) Yoneda, K.; Adachi, K.; Nishio, K.; Yamasaki, M.; Fuyuhiro,
A.; Katada, M.; Kaizaki, S.; Kawata, S. Angew. Chem., Int. Ed. 2006, 45,
5459. (b) Bao, X.; Leng, J.-D.; Meng, Z.-S.; Lin, Z.; Tong, M.-L.;
Nihei, M.; Oshio, H. Chem.Eur. J. 2010, 16, 6169.
(17) Burkill, H. A.; Robertson, N.; Vilar, R.; White, A. J. P.; Williams,
D. J. Inorg. Chem. 2005, 44, 3337.
2012, 18, 6915.
(31) (a) Sletten, J.; Hope, H.; Julve, M.; Kahn, O.; Verdaguer, M.;
Dworkin, A. Inorg. Chem. 1988, 27, 542. (b) Kitajima, N.; Fujisawa, K.;
Koda, T.; Hikichi, S.; Moro-oka, Y. J. Chem. Soc., Chem. Commun.
1990, 1357. (c) Mao, Z.-W.; Heinemann, F. W.; Liehr, G.; van Eldik,
R. J. Chem. Soc., Dalton Trans. 2001, 3652. (d) Mateus, P.; Delgado, R.;
(18) Saalfrank, R. W.; Trummer, S.; Krautscheid, H.; Schunemann,
̈
́
Lloret, F.; Cano, J.; Brandao, P.; Felix, V. Chem. Eur. J. 2011, 17,
̃
V.; Trautwein, A. X.; Hien, S.; Stadler, C.; Daub, J. Angew. Chem., Int.
Ed. 1996, 35, 2206.
11193.
(32) Powers, T. M.; Gu, N. X.; Fout, A. R.; Baldwin, A. M.;
Hernandez Sanchez, R.; Alfonso, D. M.; Chen, Y.-S.; Zheng, S.-L.;
Betley, T. A. J. Am. Chem. Soc. 2013, 135, 14448.
(33) Boca, R. In Magnetic Functions Beyond the Spin-Hamiltonian;
(19) Herbert, D. E.; Lionetti, D.; Rittle, J.; Agapie, T. J. Am. Chem.
Soc. 2013, 135, 19075.
́
́
(20) (a) Eames, E. V.; Betley, T. A. Inorg. Chem. 2012, 51, 10274.
(b) Powers, T. M.; Betley, T. A. J. Am. Chem. Soc. 2013, 135, 12289.
(21) (a) Boudalis, A. K.; Clemente-Juan, J.-M.; Dahan, F.;
Tuchagues, J.-P. Inorg. Chem. 2004, 43, 1574. (b) Kushch, L. A.;
Nikolaev, A. V.; Yagubskii, E. B.; Simonov, S. V.; Shibaeva, R. P.;
Sadakov, A. V.; Omel’yanovskii, O. E.; Mironov, V. S. Inorg. Chem.
̌
Mingos, D. M. P., Ed.; Springer-Verlag: Berlin/Heidelberg, 2006; Vol.
117, pp 1−264.
Commun. 2012, 21, 57. (c) Mereacre, V.; Klower, F.; Lan, Y.; Cler
́
ac,
̈
R.; Wolny, J. A.; Schunemann, V.; Anson, C. E.; Powell, A. K. Beilstein
̈
J. Nanotechnol. 2013, 4, 807.
(22) Stoll, S.; Schweiger, A. J. Magn. Reson. 2006, 178, 42.
(23) (a) Kapoor, R.; Pathak, A.; Kapoor, P.; Venugopalan, P.
Polyhedron 2006, 25, 31. (b) Kapoor, P.; Pannu, A. P. S.; Sharma, M.;
Hundal, G.; Kapoor, R.; Hundal, M. S. J. Coord. Chem. 2011, 64, 256.
(c) Ali, A.; Hundal, G.; Gupta, R. Cryst. Growth Des. 2012, 12, 1308.
(24) (a) Reynolds, R. A., III; Yu, W. O.; Dunham, W. R.;
Coucouvanis, D. Inorg. Chem. 1996, 35, 2721. (b) Higgs, T. C.;
Carrano, C. J. Inorg. Chem. 1997, 36, 291. (c) Handley, D. A.;
Hitchcock, P. B.; Leigh, G. J. Inorg. Chim. Acta 2001, 314, 1.
(d) Schweinfurth, D.; Klein, J.; Hohloch, S.; Dechert, S.; Demeshko,
S.; Meyer, F.; Sarkar, B. Dalton Trans. 2013, 42, 6944.
(25) Nefedov, S. E.; Denisova, T. O. Russ. J. Inorg. Chem. 2006, 51,
1404.
(26) (a) Jubb, J.; Gambarotta, S. Inorg. Chem. 1994, 33, 2503.
(b) Aspinall, H. C.; Tillotson, M. R. Inorg. Chem. 1996, 35, 2163.
(c) Jubb, J.; Scoles, L.; Jenkins, H.; Gambarotta, S. Chem.Eur. J.
1996, 2, 767. (d) Kasani, A.; Kamalesh Babu, R. P.; Feghali, K.;
N
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