Trivalent Lewis Acidic Cations Govern the Electronic Properties and Stability of Heterobimetallic Complexes of Nickel

Article abstract of DOI:10.1002/chem.201704006

Assembly of heterobimetallic complexes is synthetically challenging due to the propensity of ditopic ligands to bind metals unselectively. Here, we employ a novel divergent approach for selective preparation of a variety of bimetallic complexes within a ditopic macrocyclic ligand platform. In our approach, nickel is readily coordinated to a Schiff base cavity, and then a range of redox-inactive cations (M=Na⁺, Ca²⁺, Nd³⁺, and Y³⁺) are installed in a pendant crown-ether-like site. This modular strategy allows access to complexes with the highly Lewis acidic trivalent cations Nd³⁺ and Y³⁺, a class of compounds that were previously inaccessible. Spectroscopic and electrochemical studies reveal wide variations in properties that are governed most strongly by the trivalent cations. Exposure to dimethylformamide drives loss of Nd³⁺ and Y³⁺ from the pendant crown-ether site, suggesting solvent effects must be carefully considered in future applications involving use of highly Lewis acidic metals.

Full text of DOI:10.1002/chem.201704006

A Journal of
Accepted Article
Title: Trivalent Lewis Acidic Cations Govern the Electronic Properties
and Stability of Heterobimetallic Complexes of Nickel
Authors: Amit Kumar, Davide Lionetti, Victor Day, and James
Blakemore
This manuscript has been accepted after peer review and appears as an
Accepted Article online prior to editing, proofing, and formal publication
of the final Version of Record (VoR). This work is currently citable by
using the Digital Object Identifier (DOI) given below. The VoR will be
published online in Early View as soon as possible and may be different
to this Accepted Article as a result of editing. Readers should obtain
the VoR from the journal website shown below when it is published
to ensure accuracy of information. The authors are responsible for the
content of this Accepted Article.
To be cited as: Chem. Eur. J. 10.1002/chem.201704006
Link to VoR: http://dx.doi.org/10.1002/chem.201704006
Supported by

10.1002/chem.201704006
Chemistry - A European Journal
FULL PAPER
Trivalent Lewis Acidic Cations Govern the Electronic Properties
and Stability of Heterobimetallic Complexes of Nickel
Amit Kumar,[a] Davide Lionetti,[a] Victor W. Day,[a] and James D. Blakemore*[a]
Abstract: Assembly of heterobimetallic complexes is synthetically
challenging due to the propensity of ditopic ligands to bind metals
unselectively. Here, we employ a novel divergent approach for
selective preparation of a variety of bimetallic complexes within a
ditopic macrocyclic ligand platform. In our approach, nickel is readily
coordinated to a Schiff-base cavity, and then a range of redox-
inactive cations (M = Na+, Ca2+, Nd3+, and Y3+) are installed in a
pendant crown ether-like site. This modular strategy allows access
to complexes with the highly Lewis-acidic trivalent cations Nd3+ and
Y3+, a class of compounds that were previously inaccessible.
Spectroscopic and electrochemical studies reveal wide variations in
properties that are governed most strongly by the trivalent cations.
Exposure to dimethylformamide drives loss of Nd3+ and Y3+ from the
pendant crown-ether site, suggesting solvent effects must be
carefully considered in future applications involving use of highly
Lewis acidic metals.
a descriptor of binding and activity.[8][9]
In part inspired by these observations from natural systems,
synthetic efforts to understand and exploit the interactions of
transition metals and Lewis acidic redox-inactive cations is an
active research area. For example, model compounds for the
OEC in PSII that contain [Mn] and [Ca] have been
assembled.[ 10 ][ 11 ][ 12 ][ 13 ] Some compounds rely on multitopic
ligands that can coordinate the metals and hold them near to
each other. More broadly, a number of groups have studied
novel reactivity patterns obtained with systems including Lewis
acidic cations. Seminal work from Lau and co-workers has
demonstrated significant rate accelerations in metal-catalyzed
oxidations of organic substrates in the presence of Lewis
acids.[14] Collins also observed enhanced O-atom transfer by
binding of secondary cations to
a high-valent metal oxo
complex.[15] Early electrochemical work from Horwitz examined
interactions between oxo-bridged manganese complexes and
added Lewis acids.[16] More recently, Borovik has investigated
the effects of redox-inactive metal ions on the properties of
manganese and iron complexes. [17] [18]
Introduction
In several interesting cases, the highly Lewis-acidic metals
Sc3+ and Y3+ promote peroxide O–O bond cleavage in FeIII–O2–
M3+ species, leading to generation of reactive [FeIV=O]
intermediates.[19 ][20 ][21 ] In this chemistry, and in some of the
cases discussed above, the heterobimetallic complexes are
often not isolated, but formed in situ by combination of multiple
metal-containing precursors. This straightforward approach is
especially useful for catalytic applications, where multiple
additives can be mixed to assemble an active catalyst. Notably,
the Lewis acidities of Y3+ and other trivalent ions are significantly
greater than other ions such as Ca2+ or Sr2+; consequently, there
is a privileged position for Y3+ and Sc3+ in modulating the
reactivity of the iron complexes described above.[19]
However, despite the relevance of these heterometallic
effects to possible applications in catalysis[2] and energy
science,[ 22 ] chemical preparation of compounds that contain
trivalent Lewis-acidic ions with a transition metal has received
less attention than it deserves. Tsui et al. and Herbert et al. have
quantified Lewis acidity effects in multimetallic manganese[23]
and iron[ 24 ] oxo-bridged clusters, respectively. They find a
distinctive shift in reduction potential of their clusters as a
function of the Lewis acidity of a bound redox-inactive metal,
including the trivalent rare-earth cations Y3+ and Nd3+.
Heterometallic sites that have two or more metals in close
proximity (< 5 Å) often show unique behavior that is inaccessible
with simpler homometallic analogues.[ 1 ][ 2 ] The presence of
multiple metal centers can give rise to synergistic properties,
including improved substrate binding,[ 3 ] enhanced rates of
catalysis, and unique stability.[4] These features contribute to the
high activity of multimetallic active sites in enzymes that catalyze
complex reactions, for example N2 fixation by nitrogenase[5] or
O2 reduction by cytochrome c oxidase.[6]
Redox-inactive metals can also modulate the properties of
multimetallic active sites. For example, the oxygen-evolving
complex (OEC), a tetramanganese oxo-bridged cluster in the
enzyme Photosystem II contains a crucial Ca2+ ion that assists in
oxidation of water to generate oxygen gas.[7] The Lewis acidity of
this ion, which is redox-inactive at biologically-relevant potentials,
is postulated to be important in this role; Sr2+-substituted PSII is
also active for O2 evolution, but while some other cations
(including dications such as Cd2+) do bind in the OEC, these
analogues do not display O2 evolution activity. This highlights
that charge is not the sole feature that regulates oxygen
evolution, and suggests that Lewis acidity may be implicated as
One possible strategy for preparation of heterobimetallic
complexes involves synthesis of ditopic frameworks with binding
motifs that display distinct ligand properties and that can
therefore facilitate orthogonal metallations.[1] With this approach,
the precise tuning of the properties and reactivity of
heterobimetallic complexes can be understood by comparison of
the monometallic and the isolated heterobimetallic species.
Several ligands of this variety are known in the literature.[25]
[a]
A. Kumar, D. Lionetti, V. W. Day, and Prof. Dr. J. D. Blakemore
Department of Chemistry
University of Kansas
1251 Wescoe Hall Drive, Lawrence, Kansas 66045 USA
E-mail: blakemore@ku.edu
Supporting information for this article is given via a link at the end of
the document.
This article is protected by copyright. All rights reserved.

10.1002/chem.201704006
Chemistry - A European Journal
FULL PAPER
Reinhoudt and co-workers introduced a ditopic ligand[26] based
on a metal salophen core with a nearby crown-ether-like
moiety[27] for binding a second metal (see Scheme 1). Related
macrocyclic salen or salophen type complexes have been
28 ][ 29
]
extensively used in the fields of coordination[
and
supramolecular[30][31] chemistry; several studies have examined
complexes of this general type incorporating multiple
metals.[16][32][33] Along this line, Yang and co-workers recently
reported heterobimetallic complexes of Co2+ with Na+, K+, Ca2+,
and Sr2+; the CoII/I reduction potentials and spectroscopic
properties of these compounds were found to vary as a function
of the Lewis acidity of the redox-inactive cation.[34] This study
demonstrates the usefulness of ditopic ligands in assembling
well-defined heterobimetallic compounds, as the compounds are
quite stable and amenable to solution studies.
We have recently been investigating multidentate,
polyphenoxide ligand scaffolds that support trivalent cations,[35]
and our interest was piqued by the possibility of introducing
analogous fragments into
a ditopic framework to access
heterobimetallic species. In particular, we were interested in
testing how trivalent cations (specifically, Y3+ and Nd3+) would
modulate the structure, electronic properties, and reduction
potentials of a nearby nickel center. We anticipated that such
studies would provide important structural and spectroscopic
insights into the use of highly Lewis acidic metals in catalytic
contexts.[19]
Along this line, we now report a divergent synthetic strategy
for generation of stable heterobimetallic complexes based on
metallation of the macrocyclic, ditopic platform L-Ni (see
Scheme 1)[26] with a variety of redox-inactive metals (Na+, Ca2+,
Nd3+, Y3+). The Nd3+ and Y3+ complexes are the first examples of
stable heterobimetallic complexes of these elements with a first-
row transition metal in a macrocyclic ligand environment. They
complement unsupported analogues that feature bridging
ligands but without a stabilizing macrocyclic environment.[36][37]
Our macrocyclic complexes are stable in acetonitrile solution
and amenable to study by solution spectroscopic methods and
electrochemistry. However, electrochemical experiments reveal
that the complexes of Nd3+ and Y3+ are unstable in
dimethylformamide (DMF); conversely, they can survive for
hours in acetonitrile (MeCN), similar to the complexes of Na+
and Ca2+. We conclude that the high Lewis acidity of these
elements drives association with DMF, and thus we propose that
future studies relying on these elements could usefully include
studies of solvent-dependent behavior.
Scheme 1. Divergent synthetic strategy for preparation of
heterobimetallic nickel complexes.
Results
In this work, we have employed a divergent synthetic strategy,
in which preparation of a common nickel-containing material (L-
Ni; see Scheme 1) precedes metallation by suitable reagents of
redox-inactive metal cations. This approach has the advantage
that once a supply of the monometallic precursor complex is
prepared, a variety of heterobimetallic complexes can be readily
accessed. Our strategy relies on the ditopic ligand platform that
was originally developed by Reinhoudt and co-workers.[26] In this
chemistry, 2,3-dihydroxybenzaldehyde and tetraethylene glycol
ditosylate are reacted together to provide a linear dialdehyde
precusor.
This
dialdehyde
undergoes
Ba2+-templated
macrocyclization with o-phenylenediamine to yield complex H2L-
Ba (Scheme 1) in reasonable yields. Treatment of the barium
complex with guanidinium sulfate in a biphasic mixture of
chloroform (CHCl3) and water results in precipitation of BaSO4;
subsequent addition of nickel(II) acetate in methanol (MeOH)
generates clean L-Ni. From this point, L-Ni is a suitable synthon
for preparation of heterobimetallic complexes.
Although L-Ni and related analogues have been deployed in
a variety of contexts, no solid-state structure of L-Ni has been
available. For comparison to our heterobimetallic structures
(vide infra), single crystals of L-Ni suitable for X-ray diffraction
(XRD) studies were grown by slow cooling of a concentrated
solution of L-Ni in MeCN. As predicted from nuclear magnetic
This article is protected by copyright. All rights reserved.

10.1002/chem.201704006
Chemistry - A European Journal
FULL PAPER
resonance (NMR) experiments, the nickel center in L-Ni is
formally in the +2 oxidation state, and is coordinated to the
square-planar Schiff base cavity of the macrocyclic ligand
framework (Figure 1). This cavity provides an L2X2-type
environment by coordination of two imine nitrogens and two
phenolate oxygens to the nickel(II) center. The complex is
overall neutral, with co-crystallized MeCN as the only other
component found in the crystal data (see SI for details).
incorporated into the crown cavity has a CN of eight, with
coordination of the six macrocyclic O-atoms, plus two triflate
anions bound in the κ1 mode (Figure 2, lower). The higher CN
for Ca2+ is not unexpected, due to the higher charge and the
presence of the more coordinating triflate counteranion.
Figure 1. Solid-state structure (XRD) of L-Ni. Hydrogen atoms and two
molecules of co-crystallized MeCN solvent omitted for clarity.
Displacement ellipsoids shown at the 50% probability level.
With L-Ni in hand, we targeted incorporation of a range of
Lewis acidic cations for comparisons of structural, spectroscopic,
and electrochemical properties across the range of
heterobimetallic compounds. In these studies we elected to use
the pKa of the redox-inactive metal–aqua complex as a measure
of Lewis acidity,[38 ] an approach that has successfully been
employed in several literature reports.[23][24] Specifically, we
chose sodium (pKa = 14.7), calcium (12.7), neodymium (9.0),
and yttrium (8.3) as redox-inactive metals for the present work.
Yttrium and neodymium are especially attractive, as they are
relevant to enhancements in oxidative reactivity with iron
Figure 2. Solid-state structures (XRD) of L-Ni-Na (upper structure) and
L-Ni-Ca (lower structure). Displacement ellipsoids are shown at the 50%
probability level. Hydrogen atoms and one hexafluoroophosphate
counteranion in the structure of L-Ni-Na are omitted for clarity.
compounds
(vide
supra).[19][20][21]
Notably,
structurally
characterized heterobimetallic nickel compounds of Nd3+ or Y3+
in macrocyclic ligands have not been available in prior synthetic
and spectroscopic work.[36][37] Furthermore, the broad range of
Lewis acidities displayed by the chosen redox-inactive metals
facilitates interpretation of any observed trends in structural or
spectroscopic features.
Treatment of L-Ni with sodium hexafluorophosphate or
calcium trifluoromethanesulfonate (triflate, –OTf) in MeCN results
in generation of heterobimetallic complexes in which sodium and
calcium, respectively, are bound in the crown ether-like cavity of
L-Ni (see Experimental Section for details on synthetic
procedures). Single crystals of L-Ni-Na and L-Ni-Ca were grown
by vapor diffusion of Et2O into concentrated solutions of the
complexes in MeCN and MeOH, respectively. The solid-state
structure of L-Ni-Na reveals incorporation of the sodium ion in
the crown-like cavity with a coordination number (CN) of seven
To our delight, treatment of L-Ni with neodymium triflate or
yttrium triflate results in a distinctive color change from deep red
to yellow. Spectroscopic evidence (vide infra) is consistent with
coordination of the trivalent ions into the crown-like cavity.
Single-crystal XRD confirms the successful preparation of both
L-Ni-Nd and L-Nd-Y (Figure 3). L-Ni-Nd features the Nd3+ ion
fully coordinated by the six crown-like O-atoms, plus three triflate
anions bound in the κ1 mode, giving a CN of 9 (Figure 3, upper).
In accord with the similar Lewis acidities of Nd3+ and Y3+,[38] the
yttrium center in L-Nd-Y is also 9-coordinate with three κ1
triflates (Figure 3, lower). In L-Ni-Y, one triflate serves as a
bridging ligand between the Ni and Y centers; intriguingly, there
is a weak interaction (ca. 2.79 Å) between the Ni metal center
and an O-atom (O12) of a Y-bound triflate counterion.
(Figure
2,
upper).
As
expected,
an
outer-sphere
hexafluorophosphate ion is present, but does not interact
covalently with the heterobimetallic fragment. Conversely, the
solid-state structure of L-Ni-Ca reveals that the calcium ion
This article is protected by copyright. All rights reserved.

10.1002/chem.201704006
Chemistry - A European Journal
FULL PAPER
(2.306(5) Å), due to the unique bridging nature of the triflate
anion containing O11 in L-Ni-Y. On average, the Nd–Otriflate
distances are significantly longer (2.428 Å), while the Ca–Otriflate
distances are similar on average (2.331 Å)
Table 1. Selected distances (Å) computed from the solid-state (XRD)
structures of the heterobimetallic complexes.
Ni–M[a]
pKa of
M(H2O)n
Ni–N
Ni–O
M–O
Complex
(avg)[b]
(avg)[c]
(avg)[d]
L-Ni
-
-
1.854
1.858
1.859
1.854
1.847
1.854
1.844
1.857
1.858
1.844
-
L-Ni-Na
L-Ni-Ca
L-Ni-Nd
L-Ni-Y
14.7
12.7
9.0
3.504(1)
3.475(1)
3.505(1)
3.398(1)
2.466
2.433
2.458
2.348
8.3
[a] M is the redox-inactive metal in the crown-ether-like pocket. [b] Average of
the distances between Ni and N1 and N2 of the imine groups. [c] Average of
the distances between Ni and O1 and O2 of the phenoxide ligands. [d]
Average of the distances between M and O1 and O2 of the bridging phenoxide
ligands.
With these results in hand, we next moved to interrogate the
electronic structure of the heterobimetallic compounds by
spectroscopy (data summarized in Table 2). Electronic
absorption spectra collected for the parent monometallic
complex L-Ni reveal a response typical of Schiff base nickel(II)
complexes: a lowest energy band with λmax at 486 nm with molar
absorptivity
of
6800
M–1
cm–1. Based on its position and intensity, and in agreement with
prior literature, this transition can be assigned to a metal-to-
ligand charge transfer.[39] Upon installation of the redox-inactive
cations in the crown-ether-like cavity, this MLCT band shifts to
higher energy. This is consistent with a shift in the position of the
metal-centered HOMO of these compounds to lower energies by
interaction of the ligand with the redox-inactive cations via the
phenoxide groups. Specifically, attenuated donation from the
phenoxide groups to the nickel center should lower the HOMO,
resulting in widening of the HOMO-LUMO gap. According to this
model, the effect of the redox-inactive cations on the ligand π-
system (LUMO) must be minor in comparison to the effect on the
nickel center. Plotting the energy of the MLCT band (in eV)
versus the pKa of the corresponding metal–aqua ion reveals a
uniform, linear trend in the shift in the HOMO-LUMO gap of 46
meV per pKa unit. While the electronic structure of the metal
complexes is most strongly perturbed by the trivalent cations,
the effects engendered by these metals are in line with those
observed for divalent or monovalent ions. This is in agreement
with work examining the effect of redox-inactive cations on the
reduction potentials of manganese[23] or iron[24] oxo clusters.
The chemical shift for the signal corresponding to the imine
proton in the 1H NMR spectra for the bimetallic complexes
provides a useful readout of the effects of incorporation of Lewis
acids. A minimal shift in the position of the imine proton
Figure 3. Solid-state structures (XRD) of L-Ni-Nd (upper structure) and
L-Ni-Y (lower structure). Displacement ellipsoids shown at the 50%
probability level. Hydrogen atoms omitted for clarity.
Inspection of this series of analogous structures reveals no
discernible trends in bond distances around the Ni metal center
as a function of the Lewis acidity of the redox-inactive cation M.
Likewise, we observe no significant trend in distances between
the redox inactive ion and the bridging O-atoms of the
phenoxide ligands (See Table 1 for select bond metrics).
However, we do observe that the structure of L-Ni-Y displays a
significantly shorter intermetallic Ni–M distance (by nearly 0.1 Å)
than all the other structures. This is in accord with the shortest
M–Ophenoxide distance of the series, and is likely due to the
presence of the unique bridging triflate ligand (vide supra).
Interestingly, this bridging triflate motif was only observed for
yttrium, the most Lewis acidic metal in our study. We note that
the average Ni-Nimine distance in L-Ni-Y is also the shortest of
our series, in accord with pronounced changes in the electronic
structure of this compound due to the close proximity of the
trivalent cation (vide infra). Comparing the distances between
the redox-inactive M and the triflate O-atoms, the only distance
of note is the rather long Y–O11 distance of 2.350(4) Å. This is
significantly longer than Y–O21 (2.291(5) Å) and Y–O31
This article is protected by copyright. All rights reserved.

10.1002/chem.201704006
Chemistry - A European Journal
FULL PAPER
resonance from 8.6 for L-Ni-Na to 8.7 ppm for L-Ni-Ca.
However, this resonance is shifted to 16.5 ppm in L-Ni-Nd, and
to 19.0 ppm for L-Ni-Y. These results indicate a marked
acidification of the imine protons upon coordination of the
trivalent ions, which results in the observed deshielding effect.
This observation may form the basis for an important design rule
for future work, in that acidification of vulnerable moieties on
organic ligands may activate decomposition mechanisms that
are normally not accessible without the action of strongly Lewis
acidic metal ions.
upon coordination of the Lewis acidic cations. This can be
assigned to arise from strengthening of the C–N interaction as
the Lewis acidity of the associated ion increases. Presumably,
this bond strengthens due to decreased back donation into the
imine group upon coordination of the Lewis acidic cations. As
observed in the 1H NMR resonances of the imine protons, the
largest shift in the imine stretch was encountered for L-Ni-Nd
and L-Ni-Y. As in the case of the MLCT energy, the imine
stretch was found to vary linearly over the range of Lewis
acidities investigated here, with a shift of 8.7 cm–1 per pKa unit
(see SI for plot). This uniform trend is consistent with a similar
mode of interaction between the imine moieties and the Lewis
acidic cations in the compounds; the cations are tightly bound in
their macrocyclic host ligand environment, even in coordinating
acetonitrile solvent.
Table 2. Selected spectroscopic features of the heterobimetallic complexes in
acetonitrile solvent.
[b]
[c]
pKa of
M(H2O)n
λmax (nm)[a]
δ(Himine
(ppm)
)
νimine
(cm–1
)
Complex
[ε (M–1 cm–1)]
L-Ni
-
486 [6800]
475 [7400]
457 [12900]
422 [2800]
430 [3900]
-
1610
1613
1616
1659
1664
L-Ni-Na
L-Ni-Ca
L-Ni-Nd
L-Ni-Y
14.7
12.7
9.0
8.6
8.7
16.5
19.0
8.3
Figure 4. Stacked electronic absorption spectra of the heterobimetallic
complexes. Conditions: [L-Ni-M] = 10–5 M in MeCN.
[a] UV-vis spectra shown in Figure 4. [b] NMR spectra shown in Supporting
Information. Solvent: CD3CN. [c] Infrared spectra shown in Supporting
Information.
Considering the high stability exhibited by all of our
heterobimetallic complexes in acetonitrile, we were encouraged
to interrogate their electrochemical properties. Electrochemical
data has not been available for heterobimetallic nickel
complexes with similar Schiff-base ligands in the past; this may
be due in part to poor stability of non-macrocyclic ligand
environments. Here, we began by examining the electrochemical
properties of the compounds in dimethylformamide (DMF)
solution. (L-Ni-Ca is only sparingly soluble in MeCN, making
experiments more challenging.) Details regarding the
electrochemical studies are given in the Experimental Section.
Cyclic voltammograms collected for the compounds in
dimethylformamide reveal a single, irreversible reduction (see SI
for voltammograms; electrochemical data summarized in Table
3). An analogous reduction has been observed previously for
monometallic nickel derivatives with unsubstituted salophen
ligands.[ 40 ] L-Ni is reduced near –2 V vs. the ferrocenium/
ferrocene couple (denoted hereafter as Fc+/0) (Ep,c = –1.89 V
where Ep,c denotes the cathodic peak current of the irreversible
process). The potentials required for reduction of L-Ni-Na and L-
Ni-Ca, Ep,c = –1.73 V and –1.53 V, respectively, are consistent
with coordination of the Lewis acidic ion resulting in a shift of the
reduction potentials to less negative values. The reductive
electrochemistry can be assigned to undergo an EC-type
mechanism,[41] as the initial electrochemical reduction is followed
Figure 5. Dependence of the lowest-energy absorption band (MLCT
energy) of the L-Ni-M complexes on the Lewis acidity (pKa) of the
corresponding M-aqua complexes.
Consistent with this observation of acidification of the imine
protons by 1H NMR, we also observe a shift of the imine C=N
stretch to higher energies in infrared spectra (see SI for spectra)
This article is protected by copyright. All rights reserved.

10.1002/chem.201704006
Chemistry - A European Journal
FULL PAPER
by a chemical reaction. At least one product of this chemical
reaction is susceptible to re-oxidation at potentials shifted
of L-Ni-Nd and L-Ni-Y did not show similar clearly resolved
reduction processes in DMF. The reduction potential measured
in both of these latter cases was near –1.8 V (see SI for
voltammograms). This value of –1.8 V does not fit with the trend
set by the other compounds; significantly, the peak potential
measured for L-Ni-Nd and L-Ni-Y (near –1.8 V) is very similar to
L-Ni, which is reduced at –1.89 V. This suggested to us that the
highly Lewis acidic Nd3+ and Y3+ ions might be lost from the
macrocyclic ligand in the presence of highly coordinating DMF
solvent.
positive of Ep,c
.
Table 3. Reduction potentials of the heterobimetallic complexes in
dimethylformamide and acetonitrile.
Complex
L-Ni
Ep,c in DMF (V)[a]
–1.89
Ep,c in MeCN (V)[b]
–1.84
The electrochemical response of the heterobimetallic
complexes in MeCN is much simpler (see Figure 7). The single,
irreversible reduction event observed with L-Ni in DMF solvent is
retained in MeCN albeit with a minor shift (Ep,c = –1.84 V in
MeCN vs. –1.89 V in DMF). The heterobimetallic complexes
show similar irreversible reduction processes (again consistent
with EC behavior) that shift to less negative potentials as a
function of the Lewis acidity of the paired cation. Specifically,
examination of Ep,c as a function of the Lewis acidity of the
added cation gives an average shift in the reduction potential of
the resulting heterobimetallic complex of 70 mV per pKa unit (see
Figure 8). As in the spectroscopic data (vide supra), the trivalent
and highly Lewis acidic cations resulted in the most pronounced
shift of Ep,c. Although the Lewis acidity of yttrium is greater than
that of neodymium (pKa = 8.3 vs. 9.0), the reduction peak
potential measured for L-Ni-Nd is more positive than that of L-
Ni-Y. This can be ascribed to uncertainties in the
electrochemical data, as interpretations regarding peak
potentials for irreversible processes can be complicated by
secondary effects. Similar uncertainty was encountered by
Agapie and co-workers,[23][24] and we have quantified our error
here as 70±12 mV/pKa which corresponds to ±17% error. Thus,
there is a clear and uniform trend of modulated reduction
potential upon coordination of Lewis acidic metals to L-Ni, even
for the unusual trivalent analogues that are the focus of this
study.
L-Ni-Na
L-Ni-Ca
L-Ni-Nd
L-Ni-Y
–1.73
–1.65
–1.53
–1.42
N/A [c]
–1.18
N/A [c]
–1.22
[a] Reported versus the ferrocenium/ferrocene couple (Fc+/0). Conditions: 0.1
M [nBu4N]+[PF6]– in DMF; 100 mV/s. [b] Reported versus Fc+/0. Conditions: 0.1
M [nBu4N]+[PF6]– in MeCN; 100 mV/s. [c] Compound not stable as gauged from
electronic absorption data and electrochemical measurements (see main text).
To confirm that the redox events observed by cyclic
voltammetry indeed correspond to electron transfer processes
involving the desired compounds, we titrated a sample of L-Ni
(in the electrochemical cell) with increasing quantities of
Ca(OTf)2. Addition of 0.5 equiv. of Ca2+ results in appearance of
a new reduction wave, Ep,c, at –1.51 V. The original reduction
wave for L-Ni (at –1.89 V) is diminished upon Ca2+ addition,
confirming conversion of L-Ni to L-Ni-Ca by rapid metalation of
the crown-ether-like cavity of the macrocyclic ligand (see Figure
6). A second, minor reduction process is observed with Ep,c ≈ –
1.75 V, perhaps consistent with interaction of the Ca2+ ion with
multiple Ni centers.
Figure 6. Electrochemical response of L-Ni in DMF upon addition of
various amounts of Ca(OTf)2. Growth of the reduction process with Ep,c
=
–1.51 V indicates coordination of Ca2+ to L-Ni, giving in-situ generation
of L-Ni-Ca.
Figure 7. Overlay of cyclic voltammetry data for L-Ni and L-Ni-M
complexes in MeCN (0.1 M [nBu4N]+[PF6]–, 100 mV/s).
We were surprised, however, to observe that voltammograms
This article is protected by copyright. All rights reserved.

10.1002/chem.201704006
Chemistry - A European Journal
FULL PAPER
Figure 9. Electrochemical response of L-Ni in DMF upon addition of
various amounts of Y(OTf)3. The initial reduction of L-Ni occurs near –
1.8 V, and this signal undergoes only minor changes as Y3+ is added.
Figure 8. Plot of Ep,c(NiII/I) vs. pKa of M(H2O)n.
A second set of electronic absorption spectra confirms DMF-
driven loss of Nd3+ and Y3+ from our macrocyclic ligand
environment (see SI for spectra). There are only minor
differences in the spectra of L-Ni-Nd and L-Ni-Y versus L-Ni in
DMF, consistent with virtually complete loss of the trivalent ion
from the isolated compounds. Consistent with this assignment,
there were only minor changes in cyclic voltammograms of L-Ni
upon titration with up to 13 equiv. of Y(OTf)3. Rather than
appearance of a new, unique reduction wave corresponding to
the heterobimetallic compound, we observe only a single, broad
reduction that gradually shifts to less negative potentials as a
function of added yttrium (see Figure 9). In this experiment, Ep,c
shifts over a narrow potential range from –1.88 to –1.77 V. The
original, reasonably sharp peak for reduction of L-Ni is located at
–1.89 V, and thus addition of even 13 equiv. of Y3+ does not
result in a significant shift in the reduction potential. This is
consistent with an equilibrium exchange process that results in
only minor occupation of the macrocyclic ligand environment by
Y3+.[16a] The broadened peak position in the titration experiment
is further suggestive of a rapid equilibrium—the various forms of
the metal complex present in solution are rapidly interconverting
under these conditions and thus are not stable on the timescale
of the cyclic voltammetry experiment.
Discussion
Tsui et al. recently demonstrated that heterometallic
manganese–oxido clusters containing a redox-inactive cation
show a reduction potential shift of 100 mV per pKa unit.[23] In
these clusters, the redox-active component of the system, a
nearby manganese center, is closely associated with the redox-
inactive metal, giving rise to the distinctive shifts in reduction
potential. Similar results have been demonstrated for iron
analogues of these compounds, with a measured dependence
on the Lewis acidity of the redox-active metal center of 70 mV
per pKa unit.[24] The attenuated coupling in these two cases can
be ascribed to differences between iron and manganese,
although exact comparisons are difficult.
In our new work, the coupling between the nickel center and
the redox-inactive metal is apparently not significantly different
from these cases. We have determined an average reduction
potential shift of 70(±12) mV per pKa unit in our heterobimetallic
compounds. This is consistent with coupling driven by the
bridging phenolate groups in our compounds. Appealingly, our
results here span the same Lewis-acidity range (from Na+ (14.7)
to Y3+ (8.3)) used in the manganese and iron cluster work,
allowing this comparison of coupling to be made under
reasonable conditions. Additionally, our results align well with
those from a narrower Lewis acidity range studied by Yang and
co-workers in heterobimetallic complexes of cobalt.[34]
Our compounds and electrochemical studies extend the
range of available data on the modulation of reduction potentials
by trivalent redox-inactive metals like Y3+. First, in terms of
useful trends, we find a linear relationship between most
spectroscopic properties and Lewis acidity. Second, we find that
the heterobimetallic compounds containing trivalent Lewis acidic
cations may be especially prone to speciation or decomposition
in coordinating solvents like DMF. We were able to clearly
assign loss of both Y3+ and Nd3+ from our ligand framework; the
This article is protected by copyright. All rights reserved.

10.1002/chem.201704006
Chemistry - A European Journal
FULL PAPER
loss occurs despite the use of a hexadentate and macrocyclic
ligand that we expected would impart stability to the
heterobimetallic compounds.[26] Conversely, L-Ni, L-Ni-Na, and
L-Ni-Ca are stable in both DMF and MeCN, allowing their
reduction peak potentials (Ep,c) to be measured in both solvents.
The potentials are not significantly different, suggesting that
DMF and MeCN do not significantly perturb the coordination
environment around Ni or M in these compounds.
As noted in the introduction, addition of the highly Lewis
acidic Y3+ and Sc3+ ions results in enhanced reactivity of certain
iron peroxo complexes.[20][21] The enhanced activity is readily
achieved by addition of multiple equivalents of M3+ to the
reaction mixture of interest. Our electrochemical studies show
that (even in a macrocyclic environment designed to favor
retention of a heterobimetallic structure) redox-inactive cations
like Y3+ can engage in rapid equilibrium processes. Thus,
speciation processes under catalytic or electrochemical
conditions could contribute to reactivity, or even afford a range
of structures that differ from the starting materials. In the current
work, we encountered loss of defined structure in the presence
of DMF. Although the stability of cation-crown complexes has
been of interest for some years,[27] more data on the stability of
trivalent cations in crown-ether-like cavities is needed in this
area.[ 42 ][ 43 ] We suggest that future work on multimetallic
complexes could usefully include such related stability studies,
as the information gained on the speciation of catalysts could
lead to improvements in their activity by rational design.[44]
All manipulations were carried out in dry N2-filled gloveboxes
(Vacuum Atmospheres Co., Hawthorne, CA) or under N2 atmosphere
using standard Schlenk techniques unless otherwise noted. All solvents
were of commercial grade and dried over activated alumina using a PPT
Glass Contour (Nashua, NH) solvent purification system prior to use, and
were stored over molecular sieves. All chemicals were from major
commercial suppliers and used after extensive drying. 2,3-
dihydroxybenzaldehyde was sublimed in vacuo. Deuterated NMR
solvents were purchased from Cambridge Isotope Laboratories; CD3CN
was dried over molecular sieves. 1H, 13C, 19F, and 31P NMR spectra were
collected on 400 and 500 MHz Bruker spectrometers and referenced to
the residual protio-solvent signal in the case of 1H and 13C, and to the
deuterium lock signal in the case of 19F and 31P unless otherwise noted.
Chemical shifts (δ) are reported in units of ppm and coupling constants
(J) are reported in Hz. NMR spectra are given in the SI (Figures S1 to
S13). Elemental analyses were performed by Midwest Microlab, Inc.
(Indianapolis, IN).
Synthesis
General procedure for preparation of L-Ni-M. Under inert
atmosphere, a heterogenous solution of L-Ni in MeCN was added to 1
equiv. of corresponding metal salt solution in MeCN and stirred for 30
minutes. The color changed from red to yellow, orange or deep red. The
solvent was removed in vacuo to give the desired product in good yield.
Yields were in the range of 85-95%. Additional details are given in the
Supporting Information.
Mass spectrometry data collected on solutions containing the
heterobimetallic compounds revealed only the presence of L-Ni with an
associated potassium cation in each case. Therefore, we conclude that
the compounds are not stable to the conditions present in our
electrospray ionization (ESI) mass spectrometer. For L-Ni: ESI-MS
(positive) m/z: 557.1 (100%) (L-Ni + K+), 558.1 (31%), 559.1 (54%).
Conclusions
We have synthesized and characterized a small family of
heterobimetallic complexes in a macrocyclic, ditopic ligand
framework. These complexes pair a nickel salophen moiety with
a variety of redox-inactive Lewis acidic cations (Na+, Ca2+, Nd3+,
and Y3+) bound in a crown ether-like environment. We find that
the energy of the MLCT band, the acidity of the imine protons,
and the imine stretching frequency are all correlated linearly with
the Lewis acidity of the redox-inactive metal as judged by the
pKa of the corresponding M(H2O)n species. At the extreme end
of the pKa scale, for the ions Y3+ and Nd3+, the heterobimetallic
complexes are unstable in DMF. Exposure of our compounds to
DMF results in demetallation of the crown ether-like cavity of the
ligand framework, as judged by both electronic absorption
spectra and electrochemical studies. However, the macrocyclic
framework successfully imparts sufficient stability to the
L-Ni-Na. Yield: 85%. 1H NMR (400 MHz, CD3CN) δ 8.45 (s, 2H),
7.85 (dd, 3JH,H = 6.3 Hz, 3JH,H = 3.3 Hz, 2H), 7.29 (dd, 3JH,H = 6.3 Hz, 3JH,H
= 3.3 Hz, 2H), 7.15 (dd, 3JH,H = 8.3 Hz, 3JH,H = 1.5 Hz, 2H), 6.93 (dd, 3JH,H
3
= 7.7 Hz, JH,H = 1.5 Hz, 2H), 6.69 (dd, J = 8.1, 7.7 Hz, 2H), 4.12 – 4.05
(m, 4H), 3.89 – 3.82 (m, 4H), 3.74 (s, 4H). 13C{1H} NMR (126 Hz,
CD3CN) δ 157.27, 155.16, 149.48, 143.18, 128.78, 126.38, 120.63,
116.42, 115.13, 69.49, 68.04, 67.62. 19F NMR (376 Hz, CD3CN) δ –
72.02, –73.90. 31P NMR (162 Hz, CD3CN) δ –131.55, –135.91, –140.27,
–144.64,
–149.00,
–153.36,
–157.72.
Anal.
Calcd.
for
C28H27N3NiNaO6PF6 (L-Ni-Na): C 46.18, H 3.74, N 5.77; Found: C 46.24,
H 3.79, N 5.77. Electronic absorption spectrum (MeCN): 249 (43000),
294 (18200), 305 (17400), 373 (23700), 475 nm (7400 M–1 cm–1).
heterobimetallic compounds in MeCN to enable
a
full
L-Ni-Ca. Yield: 93%. 1H NMR (500 MHz, CD3OD) δ 10.47 (bs, 2H),
8.13 (dd, 3JH,H = 6.0 Hz, 3JH,H = 2.8 Hz, 2H), 7.49 (dd, 3JH,H = 5.9 Hz, 3JH,H
= 2.8 Hz, 2H), 7.43 (dd, 3JH,H = 8.0 Hz, 3JH,H = 1.7 Hz, 2H), 7.35 (dd, 3JH,H
= 8.0 Hz, 3JH,H = 2.2 Hz, 2H), 6.93 (dd, J = 7.9, 2.3 Hz, 2H), 4.41-4.35 (m,
4H), 3.89-3.82 (m, 4H), 4.02 (s, 4H). 13C{1H} NMR (126 Hz, CD3OD) δ
158.13, 150.91, 150.28, 141.11, 129.22, 127.03, 121.01, 117.51, 116.69,
116.26, 69.05, 67.73, 66.83. 19F NMR (376 Hz, CD3OD) δ –80.11.
Electronic absorption spectrum (MeCN): 241 (70100), 295 (32300), 306
(34600), 354 (41000), 367 (40400), 457 nm (12900 M–1 cm–1). Anal.
electrochemical study in this solvent. The results, taken together,
demonstrate that heterobimetallic complexes of nickel and
trivalent redox-inactive cations are highly Lewis acidic and could
find future applications, if conditions are carefully selected to
maintain their unique structures.
Experimental Section
General Considerations
This article is protected by copyright. All rights reserved.

10.1002/chem.201704006
Chemistry - A European Journal
FULL PAPER
Calcd. for C28H24N2NiCaO12F6S2 (L-Ni-Ca): C 39.22, H 2.82, N 3.27;
Infrared spectra were recorded on a PerkinElmer Spectrum 100 FTIR
spectrometer at room temperature; solution samples were prepared
inside a dry N2-atmosphere glovebox and sealed in 0.1 mm NaCl cells.
All measurements were collected in acetonitrile (MeCN) solution in order
to access the spectral regions of interest. Electronic absorption spectra
were collected with an Ocean Optics Flame spectrometer, in a 1-cm path
length quartz cuvette.
Found: C 39.15, H 2.86, N 3.26.
L-Ni-Nd. Yield: 89%. 1H NMR (400 MHz, CD3CN) δ 16.48, 11.00,
9.51, 8.62, 6.92, 3.67, 5.29, 2.40, 1.84. 19F NMR (376 Hz, CD3OD) δ –
80.11. Electronic absorption spectrum (MeCN): 238 (9200), 307 (8200),
332 (8400), 347 (8700), 422 nm (2800 M–1 cm–1). Anal. Calcd. for
C29H24N2NiNdO15F9S3 (L-Ni-Nd): C 31.36, H 2.18, N 2.52; Found: C
27.44, H 2.08, N 2.65. Satisfactory analysis (±0.4%) could not be
obtained for L-Ni-Nd; this is likely due to high sensitivity of the compound
to water.
Acknowledgements
The authors thank David Sconyers for assistance with
electrochemistry, and Dr. Justin Douglas and Sarah
Neuenswander for assistance with NMR spectroscopy. This
work was supported by the University of Kansas, and by the US
National Science Foundation through award CBET-1605524.
L-Ni-Y. Yield: 88%. 1H NMR (400 MHz, CD3CN) δ 18.8-19.2 (bs,
3
2H), 7.96-8.04 (m, 2H), 7.74-7.79 (m, 4H), 7.56 (d, JH,H = 7.9 Hz, 2H),
7.07 (t, 3JH,H = 8.0 Hz, 2H), 6.69 (dd, J = 8.1, 7.7 Hz, 2H), 4.64 (t, 3JH,H
=
5.0 Hz, 4H), 4.39 (t, JH,H = 5.0 Hz, 4H), 4.30 (s, 4H). 19F NMR (376 Hz,
CD3CN) δ –79.24. Electronic absorption spectrum (MeCN): 231 (17400),
305 (11500), 341 (12000), 351 (12400), 429 nm (3900 M–1 cm–1). Anal.
Calcd. for C29H24N2NiYO15F9S3 (L-Ni-Y): C 33.01, H 2.29, N 2.65; Found:
C 33.16, H 2.44, N 2.66.
3
Keywords: Lewis acids • redox chemistry • macrocyclic ligands
• molecular electrochemistry • nickel
X-ray crystallography
References
Single crystals of the complex L-Ni were obtained by slow cooling of a
concentrated solution of the species in MeCN. Single crystals of the
complexes L-Ni-Na, L-Ni-Nd, L-Ni-Y were obtained by vapor diffusion of
Et2O in a concentrated solution of each species in MeCN. Single crystals
of the complex L-Ni-Ca were obtained by vapor diffusion of Et2O in a
concentrated solution of the species in MeOH. For additional details for
collection and refining of data see the ESI.
[1] A. L. Gavrilova, B. Bosnich, Chem. Rev. 2004, 104, 349-384.
[2] P. Buchwalter, J. Rosé, P. Braunstein, Chem. Rev. 2015, 115, 28-126.
[3] S. Ogo, R. Kabe, K. Uehara, B. Kure, T. Nishimura, S. C. Menon, R.
Harada, S. Fukuzumi, Y. Higuchi, T. Ohhara, T. Tamada, R. Kuroki,
Science 2007, 316, 585-587.
[4] B. Kure, M. Sano, T. Nakajima, T. Tanase, Organometallics 2014, 33,
3950-3965.
CCDC entries 1569883–1569887 contain the supplementary
crystallographic data for this paper. These data are provided free of
charge by the Cambridge Crystallographic Data Centre.
[5] J. B. Howard, D. C. Rees, Proc. Natl. Acad. Sci. U.S.A. 2006, 103, 17088-
17093.
[6] B. E. Schultz, S. I. Chan, Annu. Rev. Biophys. Biomol. Struct. 2001, 30, 23-
65.
Electrochemistry
[7] a) J. P. McEvoy, G. W. Brudvig, Chem. Rev. 2006, 106, 4455-4483. b) J.
Yano, V. Yachandra, Chem. Rev. 2014, 114, 4175-4205.
[8] G. Hendry, T. Wydrzynski, Biochemistry 2003, 42, 6209-6217.
[9] J. S. Vrettos, D. A. Stone, G. W. Brudvig, Biochemistry 2001, 40, 7937-
7945.
Electrochemical experiments were carried out in a N2-filled glovebox
in dry, degassed MeCN or DMF. 0.10
M tetra(n-butylammonium)
hexafluorophosphate ([nBu4N]+[PF6]–); Sigma-Aldrich, electrochemical
grade) served as the supporting electrolyte. Measurements were made
with a Gamry Reference 600+ Potentiostat/Galvanostat using a standard
three-electrode configuration. The working electrode was the basal plane
of highly oriented pyrolytic graphite (HOPG) (GraphiteStore.com, Buffalo
Grove, Ill.; surface area: 0.09 cm2), the counter electrode was a platinum
wire (Kurt J. Lesker, Jefferson Hills, PA; 99.99%, 0.5 mm diameter), and
a silver wire immersed in electrolyte served as a pseudo-reference
electrode (CH Instruments). The reference was separated from the
working solution by a Vycor frit (Bioanalytical Systems, Inc.). Ferrocene
(Sigma Aldrich; twice-sublimed) was added to electrolyte solution prior to
the beginning of each experiment; the midpoint potential of the
ferrocenium/ferrocene couple (denoted as Fc+/0) served as an external
standard for comparison of the recorded potentials. Concentrations of
analyte for cyclic voltammetry were ca. 10–2 to 10–3 mM unless otherwise
noted.
[10] J. Limburg, J. S. Vrettos, L. M. Liable-Sands, A. L. Rheingold, R. H.
Crabtree, G. W. Brudvig, Science 1999, 283, 1524-1527.
[11] a) J. S. Kanady, E. Y. Tsui, M. W. Day, T. Agapie, Science 2011, 333,
733-736. b) E. Y. Tsui, J. S. Kanady, T. Agapie, Inorg. Chem. 2013, 52,
13833-13848.
[12] S. Mukherjee, J. A. Stull, J. Yano, T. C. Stamatatos, K. Pringouri, T. A.
Stich, K. A. Abboud, R. D. Britt, V. K. Yachandra, G. Christou, Proc. Nat.
Acad. Sci. U.S.A. 2012, 109, 2257-2262.
[13] C. Zhang, C. Chen, H. Dong, J.-R. Shen, H. Dau, J. Zhao, Science 2015,
348, 690.
[14] a) T. C. Lau, C. K. Mak, J. Chem. Soc. Chem. Commun. 1993, 766-767.
b) T.-C. Lau, Z.-B. Wu, Z.-L. Bai, C.-K. Mak, J. Chem. Soc. Dalton
Trans. 1995, 695. c) C. M. Ho, T. C. Lau, New J. Chem. 2000, 24, 587-
590. d) S. M. Yiu, Z. B. Wu, C. K. Mak, T. C. Lau, J Am Chem Soc
2004, 126, 14921-14929. e) H. Du, P. K. Lo, Z. Hu, H. Liang, K. C. Lau,
Y. N. Wang, W. W. Lam, T. C. Lau, Chem. Commun. 2011, 47, 7143-
7145.
[15] C. G. Miller, S. W. Gordon-Wylie, C. P. Horwitz, S. A. Strazisar, D. K.
Peraino, G. R. Clark, S. T. Weintraub, T. J. Collins, J. Am. Chem. Soc.
1998, 120, 11540-11541.
Spectroscopy
This article is protected by copyright. All rights reserved.

10.1002/chem.201704006
Chemistry - A European Journal
FULL PAPER
[16] a) C. P. Horwitz, Y. Ciringh, Inorg. Chim. Acta 1994, 225, 191-200. b) C. P.
Horwitz, Y. Ciringh, S. T. Weintraub, Inorg. Chim. Acta 1999, 294, 133-
139.
[33] S. Tamburini, S. Sitran, V. Peruzzo, P. A. Vigato, Eur. J. Inorg. Chem.
2009, 2009, 155-167.
[34] A. H. Reath, J. W. Ziller, C. Tsay, A. J. Ryan and J. Y. Yang, Inorg. Chem.,
2017, 56, 3713-3718.
[17] Y. J. Park, J. W. Ziller, A. S. Borovik, J. Am. Chem. Soc. 2011, 133, 9258-
9261.
[35] D. Lionetti, V. W. Day, J. D. Blakemore, Dalton Trans. 2017, Article ASAP,
DOI: 10.1039/C7DT02577G
[18] a) Y. J. Park, S. A. Cook, N. S. Sickerman, Y. Sano, J. W. Ziller, A. S.
Borovik, Chem. Sci. 2013, 4, 717-726. b) Y. Sano, A. C. Weitz, J. W.
Ziller, M. P. Hendrich, A. S. Borovik, Inorg Chem 2013, 52, 10229-
10231.
[36] a) A. Gleizes, M. Julve, N. Kuzmina, A. Alikhanyan, F. Lloret, I. Malkerova,
J. L. Sanz, F. Senocq, Eur. J. Inorg. Chem. 1998, 1998, 1169-1174. b)
J.-P. Sutter, S. Dhers, J.-P. Costes, C. Duhayon, C. R. Chim. 2008, 11,
1200-1206. c) R. Maurice, L. Vendier, J.-P. Costes, Inorg. Chem. 2011,
50, 11075-11081. d) E. Colacio, J. Ruiz, A. J. Mota, M. A. Palacios, E.
Cremades, E. Ruiz, F. J. White, E. K. Brechin, Inorg. Chem. 2012, 51,
5857-5868. e) G. Cosquer, F. Pointillart, B. Le Guennic, Y. Le Gal, S.
Golhen, O. Cador, L. Ouahab, Inorg. Chem. 2012, 51, 8488-8501.
[37] a) R. Koner, H.-H. Lin, H.-H. Wei, S. Mohanta, Inorg. Chem. 2005, 44,
3524-3536. b) T. D. Pasatoiu, J.-P. Sutter, A. M. Madalan, F. Z. C.
Fellah, C. Duhayon, M. Andruh, Inorg. Chem. 2011, 50, 5890-5898. c)
L. Ding, W. Jin, Z. Chu, L. Chen, X. Lü, G. Yuan, J. Song, D. Fan, F.
Bao, Inorg. Chem. Commun. 2011, 14, 1274-1278. d) B. Cristóvão, J.
Kłak, B. Miroslaw, Polyhedron 2012, 43, 47-54. e) H. J. Im, S. W. Lee,
Polyhedron 2015, 101, 48-55. f) S. Y. Im, S. J. Park, H. J. Im, S. W. Lee,
Polyhedron 2016, 117, 231-243.
[19] a) W. Nam, Acc. Chem. Res. 2015, 48, 2415-2423. b) S. Fukuzumi, K.
Ohkubo, Y.-M. Lee, W. Nam, Chem. – Eur. J. 2015, 21, 17548-17559.
[20] F. Li, K. M. Van Heuvelen, K. K. Meier, E. Münck, L. Que, J. Am. Chem.
Soc. 2013, 135, 10198-10201.
[21] Y.-M. Lee, S. Bang, Y. M. Kim, J. Cho, S. Hong, T. Nomura, T. Ogura, O.
Troeppner, I. Ivanovic-Burmazovic, R. Sarangi, S. Fukuzumi, W. Nam,
Chem. Sci. 2013, 4, 3917-3923.
[22] M. Wiechen, I. Zaharieva, H. Dau, P. Kurz, Chem. Sci. 2012, 3, 2330-
2339.
[23] a) E. Y. Tsui, R. Tran, J. Yano, T. Agapie, Nat. Chem. 2013, 5, 293-299.
b) E. Y. Tsui, T. Agapie, Proc. Nat. Acad. Sci. U.S.A. 2013, 110, 10084-
10088. c) P.-H. Lin, M. K. Takase, T. Agapie, Inorg. Chem. 2015, 54,
59-64.
[24] D. E. Herbert, D. Lionetti, J. Rittle, T. Agapie, J. Am. Chem. Soc. 2013,
135, 19075-19078.
[38] Perrin, D. D. Ionisation Constants of Inorganic Acids and Bases in
Aqueous Solution (Pergamon, 1982).
[25] H. H. Dam, D. N. Reinhoudt, W. Verboom, Chem. Soc. Rev. 2007, 36,
367-377.
[39] a) S. M. Crawford, Spectrochim. Acta 1963, 19, 255-270. b) S. Di Bella, I.
Fragalà, I. Ledoux, M. A. Diaz-Garcia, T. J. Marks, J. Am. Chem. Soc.
1997, 119, 9550-9557.
[26] C. J. Van Staveren, J. Van Eerden, F. C. J. M. Van Veggel, S. Harkema,
D. N. Reinhoudt, J. Am. Chem. Soc. 1988, 110, 4994-5008.
[27] a) G. W. Gokel, W. M. Leevy, M. E. Weber, Chem. Rev. 2004, 104, 2723-
2750. b) J. S. Bradshaw, R. M. Izatt, Acc. Chem. Res. 1997, 30, 338-
345.
[40] A. A. Isse, A. Gennaro, E. Vianello, Electrochim. Acta 1992, 37, 113-118.
[41] J.-M. Saveant, Elements of Molecular and Biomolecular Electrochemistry,
Wiley, Hoboken, NJ, 2006.
[42] a) G. H. Rounaghi, M. Mohajeri, F. Soruri, R. M. Kakhki, J. Chem. Eng.
Data 2011, 56, 2836-2840. b) J.-C. G. Bünzli, B. Klein, D. Wessner, K.
J. Schenk, G. Chapuis, G. Bombieri, G. De Paoli, Inorg. Chim. Acta
1981, 54, L43-L46.
[28] P. A. Vigato, S. Tamburini, Coord. Chem. Rev. 2004, 248, 1717-2128.
[29] S. Akine, F. Utsuno, S. Piao, H. Orita, S. Tsuzuki, T. Nabeshima, Inorg.
Chem. 2016, 55, 810-821.
[30] T. Nabeshima, M. Yamamura, Pure Appl. Chem. 2013, 85, 763-776.
[31] C. Kremer, A. Luetzen, Chem. - Eur. J. 2013, 19, 6162-6196.
[32] S. Tamburini, P. A. Vigato, M. Gatos, L. Bertolo, U. Casellato, Inorg.
Chim. Acta 2006, 359, 183-196.
[43] a) R. D. Rogers, L. K. Kurihara, E. J. Voss, Inorg. Chem. 1987, 26, 2360-
2365. b) G. R. Willey, M. T. Lakin, N. W. Alcock, J. Chem. Soc., Dalton
Trans. 1993, 3407-3411.
[44] D. J. Sconyers, J. D. Blakemore, Chem. Commun. 2017, 53, 7286-7289.
This article is protected by copyright. All rights reserved.

10.1002/chem.201704006
Chemistry - A European Journal
FULL PAPER
Entry for the Table of Contents
FULL PAPER
Two’s Company: Incorporation of
highly Lewis-acidic Nd3+ or Y3+ to form
heterobimetallic complexes results in
modulated reduction potentials and
spectroscopic properties. Use of a
macrocyclic ligand framework enables
study of these complexes in MeCN.
However, the complexes of the
trivalent ions are unstable in
Amit Kumar, Davide Lionetti, Victor W.
Day, and James D. Blakemore*
Page No. – Page No.
Trivalent Lewis Acidic Cations
Govern the Electronic Properties and
Stability of Heterobimetallic
Complexes of Nickel
dimethylformamide, suggesting future
catalytic studies could usefully
examine environmental effects on
stability and speciation.
This article is protected by copyright. All rights reserved.