Catalytic Water Oxidation by a Bio-inspired Nickel Complex with a Redox-Active Ligand

Article abstract of DOI:10.1021/acs.inorgchem.7b02166

The oxidation of water (H₂O) to dioxygen (O₂) is important in natural photosynthesis. One of nature's strategies for managing such multi-electron transfer reactions is to employ redox-active metal-organic cofactor arrays. One prototype example is the copper tyrosinate active site found in galactose oxidase. In this work, we have implemented such a strategy to develop a bio-inspired nickel phenolate complex capable of catalyzing the oxidation of H₂O to O₂ electrochemically at neutral pH with a modest overpotential. Employment of the redox-active ligand turned out to be a useful strategy to avoid the formation of high-valent nickel intermediates while a reasonable turnover rate (0.15 s^-1) is retained.

Full text of DOI:10.1021/acs.inorgchem.7b02166

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Cite This: Inorg. Chem. XXXX, XXX, XXX-XXX
Catalytic Water Oxidation by a Bio-inspired Nickel Complex with a
Redox-Active Ligand
Dong Wang* and Charlie O. Bruner†
Department of Chemistry and Biochemistry and Center for Biomolecular Structure and Dynamics, University of Montana, Missoula,
Montana 59812, United States
S
* Supporting Information
Scheme 1. Key Redox-Active Species Involved in the Catalytic
Cycle of (A) GAO and (B) the Nickel-Phenolate Complex 1
Studied in This Work
ABSTRACT: The oxidation of water (H2O) to dioxygen
(O2) is important in natural photosynthesis. One of
nature’s strategies for managing such multi-electron
transfer reactions is to employ redox-active metal−organic
cofactor arrays. One prototype example is the copper
tyrosinate active site found in galactose oxidase. In this
work, we have implemented such a strategy to develop a
bio-inspired nickel phenolate complex capable of catalyz-
ing the oxidation of H2O to O2 electrochemically at neutral
pH with a modest overpotential. Employment of the
redox-active ligand turned out to be a useful strategy to
avoid the formation of high-valent nickel intermediates
while a reasonable turnover rate (0.15 s−1) is retained.
electrochemically at neutral pH and a moderate overpotential
and demonstrated that both redox-active sites were utilized to
store and transfer oxidizing equivalents in the catalytic cycle
(Scheme 1B). Our findings have thus provided new insights into
the rational design of molecular water oxidation electrocatalysts.
The complex NiII-L1 (1) can be synthesized by the reaction of
1:1 nickel(II) salt and ligand L1 (Scheme S1 and Figure S1) in
wet tetrahydrofuran in the presence of 1 equiv of tetrabuty-
lammonium hydroxide to deprotonate the phenol oxygen-donor.
The crystal structure of 1 (Figure 1) shows that the tetradentate
ligand L1 supports the formation of a mononuclear octahedral
nickel(II) complex with two H2O molecules bound in a cis
configuration. The Ni−O bond distance of 2.061 Å for the
phenolate oxygen-donor compares favorably with the ones
reported for other phenolate ligands.7 Notably, we found that a
slight decrease of the steric hindrance from mesityl to 4-
fluorophenyl (Figure S2) led to the formation of a dinuclear
nickel complex (NiII-L2)2 (2), in which each phenolate acts as a
bridging instead of a terminal ligand (Figure S3), even when the
synthesis was carried out under identical conditions. Such a
bis(phenolate)-bridged dinuclear structure has been widely
observed in copper-phenolate complexes.6c,8
The optical spectrum of 1 in CH3CN exhibited two primary
absorptions at 309 nm (ε = 4200 M−1 cm−1) and 385 nm (sh, ε =
650 M−1 cm−1) and two weak peaks at 596 nm (ε = 14 M−1
cm−1) and 950 nm (ε = 22 M−1 cm−1) (Figure S4). These
features resemble the ones reported for other nickel-phenolate
complexes.7,9 The absorption at 309 nm is assignable to the
phenolate-to-nickel charge transfer transition.9 When the
spectrum was measured in buffered H2O at pH ≥ 7, this band
underwent a blue shift and became a shoulder at 296 nm (ε =
he oxidation of water (H2O) to dioxygen (O2) is an
Tenergetically uphill reaction (E° = +1.23 V vs NHE at pH 0)
responsible for providing 99% of the driving force for natural
photosynthesis.1 This process occurs at the CaMn4O4 oxygen-
evolving complex upon the accumulation of four oxidizing
equivalents in a single turnover,2 a strategy that minimizes extra
energy input (overpotential)3 and avoids the formation of
partially oxidized product(s) [e.g., hydrogen peroxide (H2O2)].
One of the great challenges in the development of efficient
artificial water oxidation electrocatalysts is to manipulate the
driving force of the catalytic cycle close to the thermodynamic
surface of water oxidation while retaining a reasonable turnover
rate.
For single-site, homogeneous transition-metal-based electro-
catalysts, most efforts have focused on systems where the metal is
the only site for the storage of oxidizing equivalents,4 while a few
examples are available with both the metal and a redox-active
ligand involved in the catalytic cycle.5 To date, a system where
the redox potentials of the metal and ligand could be modulated
independently has yet to be reported.
We took inspiration from galactose oxidase (GAO), a
mononuclear copper enzyme that catalyzes the two-electron
oxidation of a range of primary alcohols to their corresponding
aldehydes with the concomitant reduction of O2 to H2O2.6 The
key redox-active component is a copper(I) tyrosinate, which is
oxidized by two electrons to generate a copper(II) tyrosyl radical
as the reactive oxidant (Scheme 1A). This example of synergy
between a metal and an organic radical provides a useful strategy
for mediating multi-electron transfer chemistry. In this work, we
implemented this strategy to develop a bio-inspired nickel-
phenolate complex capable of catalyzing water oxidation
Received: August 30, 2017
© XXXX American Chemical Society
A
DOI: 10.1021/acs.inorgchem.7b02166
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Communication
according to eq 1.5d,11 When CV of the clean buffer was
performed using the working electrode removed from the
catalyst-containing solution after gentle washing (see the
Supporting Information for details),12 no catalytic current was
observed (Figure 2, purple), indicating that there is no deposit of
a heterogeneous layer on the surface of the working electrode.
On the other hand, complex 2 showed a lower catalytic activity
than 1.
2
KIE = k
/k
= (i
/i
)
cat,H2O cat,D2O
cat,H2O cat,D2O
(1)
When normalized on the basis of the square root of the scan
rate (ν−1/2), the catalytic current decreased upon an increase of
the scan rate from 50 to 500 mV s−1 (Figure S6). Together with
the notable KIE, these results suggested that the attack of H2O to
form an O−O bond by an active oxidant is the rate-determining
step in the catalytic cycle. The turnover frequency (kobs) was
determined to be 0.15 s−1 by a widely used method described
previously5d based on eq 2 (Figure S6, inset)
icat/idiff = 2.24n(RTkobs/Fν)1/2
(2)
Figure 1. Crystal structure of 1 with thermal ellipsoids drawn at the 50%
probability level. Hydrogen atoms are omitted for clarity. Selected bond
distances (Å): Ni−O1, 2.061(1); Ni−O2, 2.048(1); Ni−O3, 2.125(2);
Ni−N1, 2.086(2); Ni−N2, 2.066(2); Ni−N3, 2.156(2).
where n = 4 and icat and idiff represent the catalytic current of
water oxidation and the diffusional current of the PhO•/PhO−
couple (see below), respectively. This value compares favorably
with those reported for other nickel-based water oxidation
catalysts at neutral pH.5b,13
3200 M−1 cm−1; Figure S5). This feature began to lose its
intensity at pH < 7, suggesting that the phenolate was protonated
at low pH. A pKa of 6.2 was obtained for the phenolate oxygen-
donor by a pH titration experiment (Figure S5).
The cyclic voltammogram (CV) of 1 in a sodium phosphate
(Na-Pi) aqueous buffer at pH 7 using a glassy carbon working
electrode (surface area ∼0.07 cm2) showed a strong, irreversible
catalytic current with an onset potential of ∼1200 mV vs Ag/
AgCl reference (Figure 2, red). The catalytic current has a linear
A solution containing 0.5 mM 1 in a Na-Pi buffer at pH 7 was
then subjected to bulk electrolysis using an indium−tin oxide
(ITO) working electrode (surface area ∼7 cm2). We applied a
potential that is more positive than the onset potential to ensure
fast production of O2. As shown in Figure S7, the current
obtained in the presence of 1 was much larger than that in the
solution containing the buffer alone. No lag period was observed,
indicating that the catalytic activity was not a result of the
decomposed species of 1. The Faradaic yield of O2 formation was
determined using a Clark electrode to be 80−90% (Figure S8).
Importantly, we performed another electrolysis of the clean
buffer solution using the working electrode (Figure S9) removed
from the catalyst-containing solution after 30 min of electrolysis.
No significant current was observed relative to a freshly prepared
electrode (Figure S7, blue). These observations confirm that
there is no formation of a heterogeneous nickel oxide layer at the
surface of the ITO electrode.
We also studied the electrochemical properties of Ni(NO3)2,
which is known to form heterogeneous nickel oxides active in
catalyzing water oxidation at alkaline pHs.14 As shown in Figure
S10, CV of 0.5 mM Ni(NO3)2 in a Na-Pi buffer at pH 7 exhibited
distinctly different behavior compared to that of 1. In the forward
scan, a catalytic current was not observed until the potential
reached ∼1500 mV, which is ∼300 mV higher than the onset
potential found for 1. Moreover, a crossover catalytic wave and a
broad cathodic feature at ∼950 mV were observed on the reverse
scan. In addition, we carried out bulk electrolysis of 0.5 mM
Ni(NO3)2 at pH 7 using an ITO electrode. There is clear
formation of a black nickel oxide layer on the electrode surface
after electrolysis (Figure S9). Taken together, these results
demonstrated that 1 is a homogeneous catalyst but not the
precursor of a heterogeneous nickel oxide species.
Figure 2. CVs of a 0.1 M Na-Pi buffer at pH 7 (black), 0.5 mM 1 in H2O
(red) and D2O (blue), and a clean buffer using the working electrode
removed from a 0.5 mM 1 containing solution (purple). Scan rate = 100
mV s−1. Inset: Plot of the catalytic current at 1500 mV vs [1]. The red
line represents the best linear fit up to [1] = 0.5 mM.
dependence on the concentration of 1 in the range of 0.05−0.5
mM (Figure 2, inset), suggesting that 1 is a single-site catalyst in
an aqueous solution. The catalytic current became saturated at a
catalyst loading higher than 0.5 mM. This observation is similar
to that of a copper bipyridine water oxidation catalyst recently
reported.10 CV carried out in D2O shows a much smaller catalytic
current (Figure 2, blue), suggesting that the transfer of proton(s)
from H2O is a significant component of the rate-determining
step. A solvent kinetic isotope effect (KIE) of 3.3 was determined
At low potential, a non-catalytic anodic peak was observed at
Ep,a = 780 mV, which could be assigned to the oxidation of
phenolate to the phenoxyl radical (Figure S11).6c,7 This process
is partially reversible, as revealed by the reverse scan even when
the forward scan returned at 1000 mV, showing first a cathodic
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DOI: 10.1021/acs.inorgchem.7b02166
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peak at Ep,c(1) = 720 mV. The peak separation of 60 mV is
indicative of a one-electron process; however, the reduced peak
current compared to that of the anodic peak suggested that there
is another process following the oxidation of phenolate.
Additionally, there is a second cathodic peak at Ep,c(2) = 350
mV. It is well-known that the phenoxyl radical could be stabilized
through delocalization over the entire phenol ring.6c We thus
attribute Ep,c(2) to the one-electron reduction of the delocalized
radical species.
In order to probe the mechanism of water oxidation catalysis,
we carried out more CV studies in aqueous solutions buffered at a
range of pHs from 7 to 11. As shown in Figures S12 and S13, the
non-catalytic anodic peak is insensitive to the pH of the solution,
consistent with our assignment of this feature to the PhO•/PhO−
couple. In contrast, the onset potential for the catalytic process
decreased with increasing pH. A plot of the potential measured at
a constant current (e.g., −40 μA) as a function of the pH (Figure
S12, inset) is linear with a slope of −60 mV/pH, indicating that
generation of the reactive species for attacking H2O is a metal-
based one-electron process coupled with the transfer of one
proton.
phenolate ligand occurs prior to that of the metal, which grants
sufficient time for decomposition of the phenoxyl radical to take
place. Further studies are aimed at exploring possible
decomposition pathways of 116 and enhancing its stability
through modification of the ligand L1.
In conclusion, we have provided a proof-of-concept
demonstration that a nickel-phenolate complex, 1, could act as
a single-site, molecular catalyst for the oxidation of H2O to O2 at
neutral pH and a modest overpotential. Several lines of evidence
have confirmed the homogeneous nature of 1 during catalysis.
Both the metal and phenolate ligand were utilized to store and
transfer oxidizing equivalents in the catalytic cycle, an advanta-
geous strategy to avoid the formation of high-valent nickel
intermediates. Furthermore, the flexibility of the ligand scaffold
has provided unprecedented opportunities to modulate the
redox properties of the metal and ligand individually to improve
the catalytic activity, energy efficiency, and catalyst stability in
future studies.
ASSOCIATED CONTENT
■
S
* Supporting Information
We thus propose a mechanism to summarize the key
experimental findings (Scheme 2). Overall, the starting nickel-
The Supporting Information is available free of charge on the
ACS Publications website at DOI: 10.1021/acs.inorg-
chem.7b02166.
Scheme 2. Proposed Water Oxidation Mechanism Catalyzed
by 1
Experimental methods, Scheme S1, and Figures S1−S13
(PDF)
Accession Codes
CCDC 1571238−1571239 contain the supplementary crystallo-
graphic data for this paper. These data can be obtained free of
charge via www.ccdc.cam.ac.uk/data_request/cif, or by emailing
data_request@ccdc.cam.ac.uk, or by contacting The Cambridge
Crystallographic Data Centre, 12 Union Road, Cambridge CB2
1EZ, UK; fax: +44 1223 336033.
AUTHOR INFORMATION
■
Corresponding Author
*E-mail: dong1.wang@umontana.edu.
ORCID
Dong Wang: 0000-0002-6370-1942
Present Address
†Victor High School, Victor, MT 59812.
Notes
The authors declare no competing financial interest.
ACKNOWLEDGMENTS
■
(II)-phenolate complex was oxidized by two one-electron steps
to generate the reactive oxidant, a nickel(III) phenoxyl radical
species. Both the solvent H/D KIE and the scan rate dependence
studies demonstrated that the rate-determining step is the O−O
bond formation. The resulting hydroperoxide intermediate likely
continues to be oxidized on the electrode surface to produce O2.
Notably, 1 takes advantage of the phenolate ligand as the second
redox-active site besides the metal center. Therefore, formation
of a putative high-valent nickel−oxo species as the reactive
intermediate15 could be avoided, while a reasonable turnover rate
is retained. On the other hand, 1 has limited stability under
catalytic conditions, as revealed in Figure S7. This is likely due to
the facts that (1) the phenolate moiety in ligand L1 has
unsubstituted ortho and para positions susceptible to decom-
position of the phenoxyl radical and (2) oxidation of the
Support of this work was provided by the Center for
Biomolecular Structure and Dynamics CoBRE (Grant NIGMS
P20GM103546) and the University of Montana. C.O.B. thanks
the M. J. Murdock Charitable Trust for a Partnership in Science
award (2015409: JAT: 2/25/2016). We thank Daniel A. Decato
for his assistance in solving the crystal structures.
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