Electrocatalytic water oxidation by a water-soluble nickel porphyrin complex at neutral ph with low overpotential

Article abstract of DOI:10.1021/acs.inorgchem.5b00924

The water-soluble cationic nickel(II) complex of meso-tetrakis(4-N-methylpyridyl)porphyrin (1) can electrocatalyze water oxidation to O₂ in neutral aqueous solution (pH 7.0) with the onset of the catalytic wave appearing at 14;1.0 V (vs NHE). The homogeneous catalysis with 1 was verified. Catalyst 1 exhibited water oxidation activity in a pH range 2.0-8.0 and had a strict linear dependence of catalytic current on its concentration. After 10 h of constant potential electrolysis at 1.32 V (vs NHE), a negligible difference of the solution was observed by UV-vis. In addition, inspection of the working electrode by electrochemistry, scanning electron microscope (SEM), and energy dispersive X-ray spectroscopy (EDX) showed no sign of deposition of NiO_x films. These results strongly argued that 1 is a real molecular electrocatalyst for water oxidation. The turnover frequency (TOF) for this process was 0.67 s^-1 at 20 °C. On the basis of results from the kinetic isotope effect (KIE) and inhibition experiments, electrochemical studies in various buffer solutions with different anions and pHs, and DFT calculations, a catalytic cycle of 1 for water oxidation via a formally Ni(IV) species was proposed.

Full text of DOI:10.1021/acs.inorgchem.5b00924

Article
pubs.acs.org/IC
Electrocatalytic Water Oxidation by a Water-Soluble Nickel Porphyrin
Complex at Neutral pH with Low Overpotential
Yongzhen Han,† Yizhen Wu,† Wenzhen Lai,*,† and Rui Cao*,†,‡
†Department of Chemistry, Renmin University of China, Beijing 100872, China
‡School of Chemistry and Chemical Engineering, Shaanxi Normal University, Xi’an 710119, China
S
* Supporting Information
ABSTRACT: The water-soluble cationic nickel(II) complex
of meso-tetrakis(4-N-methylpyridyl)porphyrin (1) can electro-
catalyze water oxidation to O2 in neutral aqueous solution (pH
7.0) with the onset of the catalytic wave appearing at ∼1.0 V
(vs NHE). The homogeneous catalysis with 1 was verified.
Catalyst 1 exhibited water oxidation activity in a pH range
2.0−8.0 and had a strict linear dependence of catalytic current
on its concentration. After 10 h of constant potential
electrolysis at 1.32 V (vs NHE), a negligible difference of
the solution was observed by UV−vis. In addition, inspection
of the working electrode by electrochemistry, scanning
electron microscope (SEM), and energy dispersive X-ray
spectroscopy (EDX) showed no sign of deposition of NiOx
films. These results strongly argued that 1 is a real molecular electrocatalyst for water oxidation. The turnover frequency (TOF)
for this process was 0.67 s−1 at 20 °C. On the basis of results from the kinetic isotope effect (KIE) and inhibition experiments,
electrochemical studies in various buffer solutions with different anions and pHs, and DFT calculations, a catalytic cycle of 1 for
water oxidation via a formally Ni(IV) species was proposed.
Recently, a macrocyclic Ni(II) amine complex was reported
as the first molecular catalyst for water oxidation during the
preparation of our Article.51 However, more evidence is
required to confirm the homogeneous nature of this catalysis,
as macrocyclic Ni(II) amine complexes of this kind have
already been proven to decompose and thus to form anodic
deposition as active NiOx films for water oxidation.48
Metalloporphyrins have been studied as electrocatalysts for
the reduction of oxygen and proton,52−56 but examples for the
oxidation of water are rare. Recently, Wang and Groves
reported on a series of water-soluble Co porphyrin catalysts for
water oxidation at neutral pH with high efficiency and
stability.13 The use of cationic porphyrins inspired us to
examine the catalytic performance of Ni analogues. We
rationalized that the redox potential of high-valent Ni atoms
is considerably higher than that of Co in the same valence state,
which should be favored for the oxidation of water. Herein, we
report a water-soluble cationic Ni(II) porphyrin 1 that
functions as a single-site electrocatalyst for water oxidation in
neutral aqueous solution with low overpotential. A combination
of experiments and techniques, including pH-dependence and
control experiments, KIE measurements, buffer anion and
acetonitrile inhibition effects, electrochemistry, and spectro-
scopic and surface analysis, leads to a strong argument that
INTRODUCTION
■
Catalytic water oxidation is a subject of fundamental interest
and significance because it is the half reaction to split water and
is the key step in natural and artificial photosynthesis to convert
and store solar energy.1−9 As water oxidation that produces O2
is challenging in both thermodynamic and kinetic aspects, the
development of catalysts for oxygen evolution reaction (OER)
that operate with low energy cost and with high efficiency and
stability has attracted extensive interest.10−13 Recent progress in
homogeneous and heterogeneous OER catalysts has led to the
discovery of active transition metals, including Ru,14−20 Ir,21−25
V,26 Mn,27 Fe,28−31 Co,8,13,32−36 and Cu.11,37−42 Among these
systems, mononuclear catalysts consisting of earth-abundant
metals have many advantages, including low cost and easy
syntheses, controllable redox features, and simple assemblies to
other devices.43−45 However, OER catalysts of first-row
transition metals usually require relatively large overpotential
to initiate water oxidation even in basic media,11,34,37−40 and
few have been shown to have low overpotential.46 Therefore,
improvements in catalysts to lower the energy inputs for OER
are still needed.
Nickel is a first-row transition metal and has received
particular interest for water oxidation because of its diverse
redox properties and the strong oxidizing power of its high
valence states. NiOx and Ni(OH)2 have been shown to be
highly active for electrocatalytic water oxidation,10,47−50 yet Ni-
based molecular OER catalysts are much less explored.
Received: April 24, 2015
© XXXX American Chemical Society
A
DOI: 10.1021/acs.inorgchem.5b00924
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
catalyst 1 has homogeneous catalysis features for water
oxidation. On the basis of results from a series of electro-
chemical studies and theoretical calculations, a catalytic cycle of
1 for water oxidation is proposed, in which a two-electron
oxidized 1 is considered to be the catalytic active species to
oxidize water.
RESULTS AND DISCUSSION
Synthesis and Characterization. Water-soluble nickel(II)
porphyrin complex 1 (Figure 1) was prepared as hexafluor-
■
Figure 1. (Left) Molecular structure of Ni(II) porphyrin 1. (Right)
Thermal ellipsoid plot of the X-ray structure of 1 (50% probability).
Hydrogen atoms and counteranions are omitted for clarity.
Figure 2. (a) CVs of 1, Ni(NO3)2, and buffer-only solution at 50 mV
ophosphate 1(PF6)4 and triflate 1(OTf)4 salts. Crystals of both
salts were obtained and used for characterization. Crystals of
1(OTf)4 were used in all electrochemical and catalytic
experiments due to its better solubility. Two Soret bands of 1
in water indicated an equilibrium of Ni atoms with and without
axial aqua ligands57 (Supporting Information Figure S1). The
high-resolution mass spectrum showed an ion at a mass-to-
charge ratio of 512.0845 (Supporting Information Figure S2),
matching the calculated value of 512.0850 for the divalent ion
[1(PF6)2]2+ with the expected isotopic distribution (0.5 unit
separation between each peak). In the X-ray structures of
1(OTf)4 and 1(PF6)4 (Supporting Information Figure S4), the
two cationic Ni porphyrin units matched: each Ni atom has a
four-coordinate square-planar geometry and is located perfectly
at the center of the porphyrin ring. All four counteranions were
successfully located in the X-ray structure of 1(OTf)4 and
1(PF6)4. This result suggested the Ni(II) oxidation state.
Complex 1 was further characterized by NMR spectroscopy.
Consistent with its C4 symmetry, there are four sharp signals in
1H NMR (Supporting Information Figure S3). The peaks
shown in this spectrum are sharp and clean and have the
expected chemical shifts for a diamagnetic porphyrin species.
This result was consistent with the d8 Ni(II) electronic
structure and indicated that the demetalation of Ni(II)
porphyrin did not take place, as such a process would result
in free porphyrin ligands and paramagnetic Ni(II) ions.
Electrochemical Water Oxidation. The cyclic voltammo-
gram (CV) of 0.50 mM 1 in 0.10 M pH 7.0 phosphate buffer
exhibited a pronounced catalytic activity, which was absent in
CVs of buffer-only solutions (Figure 2a). The onset for this
catalytic wave appeared at ∼1.0 V versus the normal hydrogen
electrode (NHE, all potentials reported in this paper are
referenced to NHE). This value is smaller than that observed
for its Co porphyrin analogue by ∼200 mV,13 and is even
smaller than that for the biomimetic Cu catalyst recently
reported by Lin and co-workers.40 This feature is significant to
meet the requirement for an OER catalyst to operate with low
s
−1 scan rate. (b) Normalized CVs of 0.50 mM 1 at different scan rates
in mA cm−2 V−1/2 s1/2. Conditions: ITO working electrode (S = 0.25
cm−2), 0.10 M pH 7.0 phosphate buffer, 20 °C.
energy cost. It is necessary to point out that a shoulder anodic
wave existed at 1.07 V, which corresponded to the one-electron
redox couple of [Por•+−NiII]5+/[Por−NiII]4+ (see below). The
normalized catalytic current (icat/ν1/2) decreased with increas-
ing scan rates (Figure 2b), suggesting a catalytic process with a
chemical rate-limiting step, which is likely O−O bond
formation in this system (see below).13,39
In 20 successive CV cycles of 1, no increase of catalytic
current was observed (Supporting Information Figure S6), and
the working electrode after 20 CV scans displayed no catalytic
current in a buffer-only solution (Figure 3d). Importantly, 0.50
mM of Ni(NO3)2 under identical conditions showed no
catalysis up to 1.45 V (Figure 2a). At higher potentials, the CV
of free Ni(II) ions in 0.10 M pH 7.0 phosphate buffer displayed
a crossover and a pronounced catalytic wave (Supporting
Information Figure S5), which is indicative of the formation of
active NiOx films on the electrode. The onset potential for this
catalytic wave was 1.60 V. This value is 600 mV larger than the
onset with complex 1. These results suggested that the
observed catalytic activity of 1 was not caused by NiOx films
or free Ni(II) ions that may possibly be formed during catalysis.
Constant potential electrolysis (CPE) of 1 was carried out at
1.32 V in 0.1 M pH 7.0 phosphate buffer solution in a three-
compartment gastight electrochemical cell with a Schlenk
connection on each compartment. The evolved O2 was
analyzed by using a calibrated Ocean Optics FOXY probe
(Model NeoFox), which was immersed into the solution having
the ITO working electrode (Supporting Information Figure
S19, details are described in Experimental Section). During 10
h of CPE, substantial current maintained with no current
increase (Figure 3a). Faradaic efficiency of >90% for O2
B
DOI: 10.1021/acs.inorgchem.5b00924
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
Figure 3. (a) Catalytic currents in CPE with and without 1 in 0.10 M
pH 7.0 phosphate buffer at 1.32 V. (b) UV−vis spectra of 1 before and
after 10 h of CPE. Insert: Plots of the absorption at 535 and 564 nm as
a function of time during CPE. (c) The electric charge curve with and
without 1 during 10 h of CPE. Insert: Detection of evolved O2 during
CPE and the theoretical amount of O2 produced. (d) CVs of a freshly
cleaned ITO electrode, an ITO after successive 20 CV cycles with 1,
and an ITO after 10 h of CPE with 1 in 0.10 M pH 7.0 phosphate
buffer. Scan rate is 50 mV s−1.
Figure 4. (a) CVs of 0.5 mM 1 in 0.10 M pH 2.0−7.0 phosphate
buffers. (b) CVs of different concentrations of 1 in 0.10 M pH 7.0
phosphate buffer. Insert: Plot of catalytic current at E = 1.32 V vs [1].
Conditions: ITO working electrode, scan rate 50 mV s−1, 20 °C.
evolution was obtained (Figure 3c and Supporting Information
Figure S7).
As shown in Figure 3b and Supporting Information Figure
S8, the UV−vis spectra of the aqueous solution of 1 during and
after electrolysis showed negligible change throughout the CPE,
which strongly argued for the molecular nature of this catalysis
and the catalyst stability under turnover conditions for water
oxidation. In addition, Tyndall scattering analysis of the
aqueous solution of 1 after CPE revealed that no NiOx
particles were formed (Supporting Information Figure S9).
Graphene oxide was used as a comparison because its aqueous
solution has a similar color and has apparent Tyndall scattering
effect. After 10 h of CPE, the ITO working electrode was
removed from the solution of 1 and gently washed with
deionized water, and was reexamined in a buffer-only solution.
The CV of this ITO electrode displayed no catalytic current
(Figure 3d). The surface morphology and composition of this
ITO electrode were further inspected by SEM and EDX
(Supporting Information Figures S10 and S11), which all
clearly excluded the formation of any heterogeneous Ni phase
on the working electrode.
pH 12.0, and this current retained over hundreds of CV scans
at this pH (Supporting Information Figure S15b). However, if
this NiOx-deposited working electrode was examined in a
neutral phosphate buffer, the catalytic current disappeared
stepwise over several CV scans (Supporting Information Figure
S15c). Similar electrochemical behaviors were observed in
phosphate buffer solutions with pH > 9.0, and Supporting
Information Figure S14 showed another example at pH 10.5.
These results suggested that NiOx was unstable by acting as the
OER catalyst at neutral pH, which is consistent with the factor
that NiOx materials typically catalyze water oxidation in basic
solutions.10,48,49 In addition, the catalytic current had a strict
linear dependence on the concentration of 1 (Figure 4b), which
suggested that it is a single-site, molecular catalysis according to
eq 1.
icat = ncatFAccat kcatD
(1)
cat
Significantly, 1 was catalytically active in the pH range 2.0−
8.0 (Figure 4a and Supporting Information Figure S13a). This
result strongly argues against NiOx materials that may possibly
be generated from decomposed 1 acting as the real catalyst for
water oxidation, as such species are unstable in acidic
solutions.32 In contrast to CVs of 1 in acidic and neutral
solutions, when the pH was more than 9.0, another catalytic
process appeared (Supporting Information Figure S13). For
example, successive CV cycles of 1 in a pH 12.0 phosphate
buffer exhibited a continuous increase of the catalytic current as
well as the anodic current at ∼0.78 V and cathodic current at
∼0.67 V (Supporting Information Figure S15a), implying the
formation and the deposition of active NiOx films on the ITO
electrode.10,48,49 After 20 CV scans at pH 12.0, this working
electrode displayed catalytic current in a buffer-only solution at
Here, ncat = 4 is the number of electrons for water oxidation to
produce O2, F is the Faraday constant, A is the electrode surface
area, ccat is the concentration of catalyst, kcat is the apparent
first-order rate constant, and Dcat is the diffusion coefficient of
the catalyst (Dcat = 2.33 × 10−10 m2 s−1 was calculated in sulfate
solution at 20 °C according to eq 2 and Figure 5b, see below).
Mechanism Studies. In order to get more insights into the
water oxidation mechanism, we examined the electrochemical
behaviors of 1 in aqueous solution with different electrolytes.
Unlike in phosphate buffer, the CV of 1 in 0.10 M sulfate
solution showed no apparent catalytic current up to 1.35 V but
displayed two oxidation waves at E1/2 = 1.07 V and Epa = 1.29 V
(Figure 5a), which were assigned to the two one-electron
oxidation processes of 1. We rationalized that, in comparison to
phosphate, sulfate is a much weaker base to assist O−O bond
C
DOI: 10.1021/acs.inorgchem.5b00924
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
Figure 5. (a) CVs of 1 in 0.10 M various anion solutions with a scan
rate of 50 mV s−1. Dependence of the (b) first and (c) second
oxidation peak current of 1 recorded in sulfate solution on the square
root of scan rates. The peak currents of the first (in the absence of
catalysis) and the second oxidation processes vary linearly with the
square root of scan rates. (d) Plot of icat/id vs ν−1/2 with a slope of 1.12
V1/2 s−1/2. Conditions: ITO working electrode, 20 °C.
formation via a concerted O atom-proton transfer path-
way.58−60 This buffer anion effect was further supported by
the CV of 1 in pH 7.0 acetate buffer (Figure 5a). CV
comparison of 1 in phosphate, sulfate, and acetate solutions is
significant, as it indicated that a formally Ni(IV) species could
be the active catalytic species for water oxidation.
As shown in Figure 5b, the peak current of the first oxidation
process in sulfate solution (in the absence of catalysis) varies
linearly with the square root of scan rates, which is consistent
with the Randles−Sevcik equation (eq 2).11,39
nFυD
RT
cat
id = 0.4633nFAccat
(2)
Figure 6. CVs of 0.50 mM of 1 in phosphate buffer at pH 7.0 with
different phosphate concentrations. The ionic strength was maintained
as constant with addition of sodium sulfate during these measure-
ments. (a) The catalytic current increased rapidly with the increase of
phosphate concentration. (b) The catalytic current decreased with
further increase of phosphate concentration. (c) Plot of catalytic
current at E = 1.32 V vs [HPO42−]. Conditions: ITO working
electrode, scan rate 50 mV s−1, 20 °C.
Here n = 1 is the number of electrons transferred in this
process, and ν is the scan rate. The ratio of eq 1 to eq 2 gives eq
3
icat
id
kcatT
υ
= 0.080
(3)
Here icat/id is consistent with the trend from pure diffusion
behavior to pure kinetic behavior as scan rates decrease. With
the molecular catalysis of 1 established, according to eq 3, the
TOF value of 0.67 s−1 at room temperature can be determined
from the slope of the linear plot of icat/id versus ν−1/2 in Figure
5d. It is worth noting that, as water oxidation is complicated,
the calculated TOF is only an estimation of the catalytic
performance.11,61
Similar to its Co porphyrin analogues, an inhibition effect
was also observed with 1 at high phosphate concentrations. As
shown in Figure 6, if the concentration of 1 was kept constant
at 0.50 mM, the catalytic current initially increased rapidly and
considerably with the increase of phosphate concentration up
to 33 mM (Figure 6a), which was due to the phosphate-assisted
O−O bond formation. In the range of 35−65 mM phosphate
concentration, the change of catalytic currents was insignificant,
and further increase of phosphate concentration caused the
decrease of the catalytic currents (Figure 6b).
These results from inhibition experiments can be summar-
ized in Figure 6c, which shows a maximum catalytic current at a
phosphate concentration of ca. 67 mM. It is necessary to
mention that the ionic strength was maintained as constant
during these measurements with addition of sodium sulfate.
Wang and Groves studied this inhibition and proposed that it
was caused by the coordination of phosphate anions on the
metal center, which prevented the binding and therefore
activation of water on the metal center.13 This anion-assisted
O−O bond formation step and buffer inhibition at high
phosphate concentration are consistent with homogeneous
catalysis.13,58 In addition, such competitive coordination-caused
inhibition was verified by inhibition experiments with
acetonitrile. As shown in the electronic absorption spectra of
1 (Supporting Information Figure S17), the absorbance at 420
nm increased, while the shoulder at 445 nm decreased with an
D
DOI: 10.1021/acs.inorgchem.5b00924
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
Figure 7. (a) CV and DPV of 1 in 0.10 M pH 7.0 phosphate buffer. (b) Dependence of the first and second oxidation peak potentials of 1 on pH
from 4.58 to 8.07. DPVs of 1 in 0.10 M phosphate buffers with various pH values: (c) from 4.58 to 5.37, (d) from 5.47 to 6.35, (e) from 6.54 to 8.07.
Conditions: ITO working electrode, 20 °C. DPV experimental parameters: increase potential, 0.004 V; amplitude, 0.05 V; pulse width, 0.05 s;
sampling width, 0.0167 s; sample period, 0.5 s.
Figure 8. Proposed catalytic cycle for water oxidation with 1.
k
cat,D O = (icat,H O/icat,D O)2, Supporting Information Figure S18).
2 2 2
isosbestic point at 431 nm upon addition of acetonitrile. This
result suggested the coordination of acetonitrile at the Ni(II)
site of 1. Along with this change in UV−vis, the CV of 1
showed obvious inhibition in catalysis in the presence of
acetonitrile (Supporting Information Figure S16), which
proved that competitive coordination at the Ni(II) site could
inhibit the water oxidation catalysis.
Kinetic isotope effect experiments were carried out, and a
KIE value of 1.55 was obtained from the CVs of 1 recorded in
D2O and H2O phosphate buffers under identical conditions
This result is consistent with a rate-limiting O−O bond
formation step. As we addressed above, careful analysis of CVs
of 1 in sulfate, acetate, and phosphate solutions suggested that
water oxidation could be realized via a two-electron oxidized 1
(Figure 5a). In order to further investigate the redox process of
1 in phosphate buffer, differential pulse voltammetry (DPV)
was used because of its higher sensitivity compared to CV and
its ability to probe the redox processes under catalytic
conditions.62 The DPV of 1 displayed two oxidation waves at
Epa = 1.09 V and Epa = 1.25 V in pH 7.0 phosphate buffer
(the KIE value was calculated according to KIE = kcat,H O
/
2
E
DOI: 10.1021/acs.inorgchem.5b00924
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
Figure 9. Energy diagram for water oxidation catalyzed by complex 1. TSOO is the TS of the base-assisted WNA reaction (see Figure 10b for its
structure).
(Figure 7a), which were in excellent agreement with the two
oxidation events observed in the CV of 1 in sulfate solution
(E1/2 = 1.07 V and Epa = 1.29 V). The results from both CV and
DPV measurements suggested that a formally Ni(IV) species
generated by two one-electron oxidations of 1 was able to
oxidize water.
In order to illuminate the water oxidation mechanism with 1,
DPVs were carried out in 0.1 M phosphate buffers with
different pH values. As shown in Figure 7b−e, the first
oxidation peak potential of 1 stayed almost unchanged in the
pH range 4.58−8.07. For the second oxidation process, it is
pH-dependent in the pH range 5.37−6.54, and the plot of the
oxidation peak potential versus pH exhibited a linear
dependence with a slope of −55 mV per pH unit, which
indicated a 1H+/1e− procedure. However, the second oxidation
is 1.63 V, which corresponds to the oxidation of [Por−NiII]4+
to [Por•+−NiII]5+. This result is consistent with the fact that the
first oxidation peak potential is independent of pH. Subsequent
coordination of a water molecule on Ni atom of [Por•+−NiII]5+
gives [Por•+−NiII−OH2]5+. The further one-electron oxidation
of [Por•+−NiII−OH2]5+ gives [Por•+−NiIII−OH]5+, which has
three different pathways, namely, electron transfer followed by
proton transfer (ETPT), proton transfer followed by electron
transfer (PTET), and proton-coupled electron transfer
(PCET). The calculated oxidation potentials at pH 7.0 are
1.89 and 1.81 V for the [Por•+−NiII−OH2]5+/[Por•+−NiIII−
OH]5+ and [Por−NiIII−OH]4+/[Por•+−NiIII−OH]5+ couples,
respectively. The pH-dependence of the second oxidation in
the pH range 5.37−6.54 suggested a PCET process. At pH 7.0,
the second oxidation peak was assigned to the redox couple of
[Por−NiIII−OH]4+/[Por•+−NiIII−OH]5+ due to its pH-inde-
pendence in the pH range 6.54−8.07. Significantly, the
calculated separation between the first and second oxidation
of 1 at pH 7.0 is 0.18 V, which agrees very well with the value of
0.16 V from experimental DPV measurements.
It should be noted that, in our calculations, counteranions
were omitted for simplicity. It is known that the nature and
concentration of counteranions have substantial effects on the
redox potential of positively charged nickel porphyrin species.65
We tried to include counteranions in the calculation and found
that, with addition of four OTf− anions (the position of the
OTf− anions were determined from the X-ray structure of
1(OTf)4), the calculated potential for the first oxidation of 1
was 1.48 V. Although this value is 0.15 V lower than that
without counteranions, it is still larger than the experimental
value. Given the challenges of accurate prediction of redox
potentials using DFT methods,40 the calculated redox potential
separation is more meaningful. As we discussed above, our
calculated separation between the first and second oxidation of
1 is in excellent agreement with the experimental value.
exhibited pH-independence in the pH range 4.58−5.37 (Epa
=
1.31 V) and in the pH range 6.54−8.07 (Epa = 1.25 V). Detailed
parameters for DPV measurements are summarized in Figure 7.
The oxidation of water to generate hydrogen peroxide
(H2O2) was not likely with 1 under our experimental
conditions. First, our CPE and O2 evolution experiments
were carried out at 1.32 V in pH 7.0 phosphate buffer. Under
this condition, water could not be oxidized to H2O2 from a
thermodynamic point of view (2H2O → H2O2 + 2H+ + 2e−, Eθ
= 1.78 V; at pH = 7.0, it is 1.37 V).63 Second, no H2O2 was
detected in the aqueous solution of 1 after 10 h of CPE by
using an established method.64 Furthermore, electrochemical
measurements of 1 at rotating ring (Pt)-disk (GC) electrodes
showed that no H2O2 was produced during the oxidation of
water. The results from linear sweep voltammetry (LSV) of
disk electrode with a scan rate of 10 mV s−1 and CPE of ring
electrode at 0.70 V were shown in Supporting Information
Figure S12. On the basis of these results, no current was
observed at ring electrode due to the oxidation of H2O2.53,54
The small current change on the ring electrode was due to the
reduction of oxidized species of 1 at the ring electrode.
After two one-electron oxidation steps, the release of one
proton from [Por•+−NiIII−OH]5+ occurred to give [Por−NiIII−
O•]4+, which subsequently reacted with a water molecule to
form the O−O bond via a concerted O atom-proton transfer
pathway. Since previous studies suggested that the first and
perhaps other solvation shells may play a decisive role in proton
transfer as part of the O−O bond formation step,66−74 we
therefore investigated the water nucleophilic attack (WNA)
reaction by including a cluster of four water molecules. We
Computational Studies. Density functional theory (DFT)
calculations were carried out to better understand the
mechanism of water oxidation with 1. The single crystal X-
ray structure of the cationic Ni porphyrin (Por) unit of 1 was
used as the starting point for geometry optimization. On the
basis of our results, the catalytic cycle for water oxidation with 1
was proposed in Figure 8. The energy profile was summarized
in Figure 9. The calculated potential for the first oxidation of 1
F
DOI: 10.1021/acs.inorgchem.5b00924
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
found that the O−O bond formation is concerted with a proton
transfer process, in which one proton from the attacking water
molecule transferred to the Ni-bounded O atom via the water
cluster, leading to [Por−NiII−HOOH]4+. This reaction was
found to have an energy barrier of 24.9 kcal mol−1. The
transition state (TS) was showed in Figure 10a. On the other
and its Co porphyrin analogues. For Co porphyrin, Wang and
Groves proposed that a formally Co(V) porphyrin is the active
catalytic species for water oxidation,13 but for Ni porphyrin, a
formally Ni(IV) porphyrin, [Por−NiIII−O•]4+, is suggested to
be able to oxidize water, which caused the onset of catalytic
wave for water oxidation with 1 to lower by ∼200 mV as
compared to its Co porphyrin analogues. This improvement
was likely due to the fact that the redox potential of high-valent
Ni atoms is considerably higher than that of Co in the same
valence state.
CONCLUSION
■
As NiOx is known to be highly active for water oxidation, great
care has been made to ensure that 1 is the real molecular
catalyst. Several lines of evidence have been provided in our
work, including the following: (1) catalyst 1 is active in neutral
and even acidic solutions; (2) 1 has different electrochemical
behaviors in neutral vs basic solutions, and NiOx films
produced in basic solution of 1 are unstable at pH 7.0 under
catalytic conditions; (3) there is no crossover, which is typical
for metal oxide film deposition and catalysis, in the CV of 1;
(4) catalytic current has a linear relationship with the
concentration of 1; (5) free Ni(II) ions display no catalytic
activity under identical conditions; (6) no current increase is
observed in successive cycles of CVs and in CPE; (7) UV−vis
shows almost identical spectra during and after 10 h of CPE;
(8) the working electrode after CPE displays no catalytic
current in a buffer-only solution, and no sign of any
heterogeneous Ni phase is observed on the electrode as
inspected by SEM and EDX or in the solution by Tyndall
scattering analysis; (9) buffer and acetonitrile inhibition
observed with 1 are consistent with a homogeneous, single-
site water oxidation.
In summary, the water-soluble cationic nickel(II) porphyrin
1 was synthesized and examined as a single-site OER catalyst.
Our results confirmed that 1 was able to electrocatalyze water
oxidation to O2 at neutral pH with low onset potential at ∼1.0
V. Mechanism studies revealed that the oxidation of water by a
two-electron oxidized 1 is feasible. The onset of the catalytic
wave for water oxidation with 1 is ∼200 mV lower compared to
its Co porphyrin analogues and is the lowest among existing
molecular OER catalysts in neutral aqueous solutions. Although
the TOF of 0.67 s−1 is moderate, the verification of a single-site
Ni-based molecular OER catalyst, low overpotential, homoge-
neous catalysis in neutral aqueous solutions, and catalyst
stability are notable.
Figure 10. Transition state structure of O−O bond formation step via
water nucleophilic attack. (a) [Por−NiIII−O•]4+ with four water
molecules and (b) [Por−NiIII−O•]4+ in the presence of one water
molecule and one OAc− anion. The hydrogen atoms in the porphyrin
ring were not shown for clarity.
hand, we also studied the WNA reaction in the presence of an
external base to investigate the possible base-assisted
mechanism. Due to a large self-interaction error, the calculation
with inclusion of a phosphate ion as an external base was not
successful, as we also found in previous study of WNA reaction
of a cobalt corrole system.36 Alternatively, an acetate anion was
used as an added base in our calculations. The located TS
structure shows an increased O−O bond length (2.15 vs 1.96
Å, Figure 10), indicating a much earlier TS, as compared to that
for [Por−NiIII−O•]4+ with four water molecules. The calculated
energy barrier of 18.0 kcal mol−1 is significantly lower than that
for [Por−NiIII−O•]4+ with four water molecules. In addition,
the calculated H2O/D2O KIE value is 1.44. As shown in Figure
8, the base-assisted WNA reaction leads to the formation of
[Por−NiII−OOH]3+. A subsequent PCET reaction generates
[Por−NiII−OO•]3+ at a potential of 0.01 V. The further one-
electron oxidation of [Por−NiII−OO•]3+ results in the
regeneration of 1 and the release of O2 simultaneously.
EXPERIMENTAL SECTION
■
General Materials and Methods. All reagents were purchased
from commercial suppliers and used as received unless otherwise
noted. Deionized water was used in all experiments. Porphyrins and
nickel(II) porphyrins were prepared according to modified literature
procedures.75 1H NMR spectra were acquired on a Bruker
̈
spectrometer operating at 400 MHz. High-resolution mass spectro-
metric measurements were made on a Bruker Fourier transform ion
̈
cyclotron resonance mass spectrometer APEX IV at Peking University.
Electronic absorption spectra were acquired on a Cary 50
spectrophotometer. The morphologies of the ITO surface before
and after electrolysis were examined using a JSM-6700F field emission
scanning electron microscope (FE-SEM). The EDX spectra were
collected from three randomly selected areas of each sample. In
addition, the materials were analyzed at several local spots to ensure
their chemical homogeneity. Images were obtained with an
acceleration voltage of 5 or 10 kV.
Our calculations demonstrated that water oxidation with 1
via water nucleophilic attack to [Por−NiIII−O•]4+ is feasible.
The O−O bond formation step was suggested to be the rate-
limiting step in electrocatalytic water oxidation with 1. It is
interesting to compare the catalytic mechanism of Ni porphyrin
G
DOI: 10.1021/acs.inorgchem.5b00924
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
Syntheses. Tetrakis(4-pyridyl)porphyrin. Freshly distilled pyrrole
(5.04 g, 75 mmol) and 4-pyridinecarboxaldehyde (8.04 g, 75 mmol)
were dissolved in propionic acid (150 mL), and the resulting solution
was refluxed at 140 °C for 1.5 h in the dark. After the solvent was
removed under reduced pressure, the residue was redissolved in DMF
(220 mL) and was refluxed for 30 min. After being cooled to room
temperature, the mixture was kept at −20 °C overnight during which
time the porphyrin precipitated. The precipitate was filtered off, and
the solid residue was carefully washed with ether following being dried
under vacuum to give 1.77 g of purple solid (2.9 mmol, yield 15%).
Tetrakis(4-N-methylpyridyl)porphyrin Tetratosylate. Tetrakis(4-
pyridyl)porphyrin (1.24 g, 2.0 mmol) and methyl tosylate (18.6 g,
100 mmol) were added into DMF (200 mL), and the resulting
solution was refluxed for 16 h in the dark. The volume of DMF was
reduced to ∼10 mL in vacuum, and the porphyrin was precipitated by
addition of acetone (200 mL). The precipitate was filtered off and
redissolved in water (200 mL). The aqueous phase washed by CH2Cl2
(200 mL × 3) was reduced to ∼10 mL. The porphyrin was
precipitated again by addition of acetone (200 mL), and the precipitate
was filtered off and dried under vacuum to give 2.04 g of purple solid
(1.50 mmol, 75%). 1H NMR (400 MHz, DMSO-d6): δ 9.49 (d, J = 6.3
Hz, 8H), 9.19 (s, 8H), 8.98 (d, J = 6.4 Hz, 8H), 7.46 (d, J = 7.9 Hz,
8H), 7.05 (d, J = 7.9 Hz, 8H), 4.73 (s, 12H), 2.22 (s, 12H), −3.10 (s,
2H).
Nickel(II) Tetrakis(4-N-methylpyridyl)porphyrin Tetrahexafluoro-
phosphate (1(PF6)4). Tetrakis(4-N-methylpyridyl)porphyrin tetrato-
sylate (682 mg, 0.50 mmol) and Ni(OAc)2·4H2O (871 mg, 3.5 mmol)
were added into DMF/H2O (1:1, 30 mL), and then the resulting
solution was heated at 100 °C for 6 h in the dark. After being cooled to
room temperature, the residue from removing the solvents under
reduced pressure was dissolved in water (200 mL). Nickel(II)
porphyrin was precipitated by addition of a saturated aqueous
NH4PF6 solution (ca. 50 mL), and then the precipitate was filtered
off, washed with ether, and dried in vacuum to yield 0.56 g of red solid
(0.43 mmol, 85%). Red needle-like crystals of 1(PF6)4 were obtained
by slow evaporation from CH3CN and H2O solution at room
temperature. High-resolution ESI-MS for [NiII(C44H36N8)(PF6)2]2+:
calcd 512.0850; found 512.0845.
solvent water molecules, we did not include them in the refinement.
Details of the data quality and a summary of the residual values of the
refinements are listed in Supporting Information Table S1.
Electrochemical Studies. All electrochemical experiments were
carried out using a CH Instruments (model CHI660E electrochemical
analyzer) at 20 °C unless otherwise noted. Cyclic voltammograms
(CVs) and differential pulse voltammetries (DPVs) were performed in
6 mL of aqueous solutions and used a three-compartment cell with a
0.25 cm2 tin-doped indium oxide (ITO) electrode as the working
electrode, saturated Ag/AgCl as the reference electrode, and Pt wire as
the auxiliary electrode. All potentials are reported versus the normal
hydrogen electrode (NHE) through the addition of 0.199 V to the
measured potentials. The working electrode was ultrasonically washed
with deionized water. Constant potential electrolysis recorded in 15
mL of 0.10 M pH 7.0 phosphate buffer containing 0.75 mM of
1(OTf)4 was carried out at 1.32 V in a fritted cell with a 2.0 cm2 ITO
working electrode.
The possible H2O2 formation was examined by electrochemical
measurements using rotating ring-disk electrode. Electrochemistry of 1
was conducted at a rotating ring (Pt)-disk (GC) electrode in 0.10 M of
pH 7.0 phosphate buffer. Linear sweep voltammetry (LSV) of the disk
electrode with a scan rate of 10 mV s−1 and CPE of the ring electrode
at 0.70 V were shown in Supporting Information Figure S12. The
results showed that no oxidation of H2O2 was observed at the ring
electrode, and the decline of ring current was due to the reduction of
oxidized species of 1 at the ring electrode.
Characterization of Oxygen Evolution. CPE of 0.75 mM of 1 in
15 mL of 0.10 M pH 7.0 phosphate buffer was carried out at 1.32 V in
a three-compartment gastight electrochemical cell with a Schlenk
connection on each compartment (Supporting Information Figure
S19). One compartment contained a Pt auxiliary electrode, one
contained the 2.0 cm2 ITO working electrode, and one contained the
Ag/AgCl reference electrode. Analysis of O2 produced in CPE
experiments was conducted by using a calibrated Ocean Optics FOXY
probe (Model NeoFox) which was immersed into the solution having
the ITO working electrode. The solution was bubbled with argon for 1
h with vigorous stirring to remove O2. The phase shift of the O2 sensor
on the FOXY probe, recorded at 10 s intervals, was converted into the
concentration of O2 in the solution using a two-point calibration curve
(air, 20.9% O2; high purity argon, 0% O2). Background from the
buffer-only solution was measured in the same method and was
deducted.
Theoretical Methods. All calculations were done with density
functional theory (DFT) as implemented in the Gaussian 09 software
package.80 Optimization and frequency calculations were performed
using the B3LYP functional in combination with the def-SVP basis
set.81 The energy was then corrected by using a larger basis set, def-
TZVP.82 All optimizations and single point calculations were
performed with solvation included using conductor-like polarizable
continuum model (CPCM)83 in water. For redox potential
calculations, reference potential of normal hydrogen electrode
(NHE) was set at −4.28 eV,36,84 and the standard free energy of
proton was set at −272.2 kcal mol−1.36 The H2O/D2O KIE for the
base-assisted WNA reaction was calculated using Gaussian frequency
data based on the semiclassical Eyring equation85
Nickel(II) Tetrakis(4-N-methylpyridyl)porphyrin Tetratriflate (1-
(OTf)4). The red needle-like crystals of 1(PF6)4 were taken in a mixture
−
of CH3CN/H2O (1:1), and PF6 anions were exchanged by OTf− on
anion exchange resin by slow stirring at ambient temperature for 48 h.
The solution was filtered, and the resin was washed with methanol.
The solvents were removed under reduced pressure to give a red solid.
Red needle-like crystals of 1(OTf)4 were obtained by slow evaporation
1
from CH3OH and H2O at room temperature. H NMR (400 MHz,
CD3CN): δ 9.07 (d, J = 6.3 Hz, 8H), 9.02 (s, 8H), 8.67 (d, J = 6.3 Hz,
8H), 4.66 (s, 12H) (Supporting Information Figure S3).
Crystallographic Studies. The complete data sets for red needle-
like crystals of 1(PF6)4 and 1(OTf)4 were collected. Single crystals
suitable for X-ray analysis were each coated with Paratone-N oil,
suspended in a small fiber loop, and placed in a cooled gas stream at
150(2) K on a Bruker D8 Venture X-ray diffractometer. Diffraction
intensities were measured using graphite monochromated Mo Kα
radiation (λ = 0.710 73 Å). Data collection, indexing, data reduction,
and final unit cell refinements were carried out using APEX2;76
absorption corrections were applied using the program SADABS.77
The structure was solved by direct methods using SHELXS78 and
refined against F2 on all data by full-matrix least-squares with
SHELXL,79 following established refinement strategies. Crystallo-
graphic studies revealed that both 1(PF6)4 and 1(OTf)4 crystallized in
the monoclinic space group C2/c. All non-hydrogen atoms were
refined anisotropically. All hydrogen atoms binding to carbon were
included in the model at geometrically calculated positions and refined
using a riding model. The isotropic displacement parameters of all
hydrogen atoms were fixed to 1.2 times the U value of the atoms they
are linked to (1.5 times for methyl groups). The CIF of 1(OTf)4
contains level B alerts (isolated oxygen atom for O1S and O2S).
Atoms O1S and O2S are oxygen atoms of cocrystallized solvent water
molecules. Due to the positional disorder of the hydrogen atoms of
kH
kD
(GH# − GHR) − (GD# − GD)
R
⎡
⎤
⎥
⎢
= exp −
⎢
⎣
⎥
⎦
RT
where G is the Gibbs free energy, superscript R denotes reagent, and #
represents TS.
ASSOCIATED CONTENT
* Supporting Information
■
S
Figures S1−S19, Table S1, crystallographic data in CIF format
for 1(OTf)4 and 1(PF6)4, and Cartesian coordinates of DFT-
optimized structures. The Supporting Information is available
free of charge on the ACS Publications website at DOI:
10.1021/acs.inorgchem.5b00924.
H
DOI: 10.1021/acs.inorgchem.5b00924
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
(25) Moore, G. F.; Blakemore, J. D.; Milot, R. L.; Hull, J. F.; Song, H.
E.; Cai, L.; Schmuttenmaer, C. A.; Crabtree, R. H.; Brudvig, G. W.
Energy Environ. Sci. 2011, 4, 2389−2392.
AUTHOR INFORMATION
Corresponding Authors
*E-mail: wenzhenlai@ruc.edu.cn.
*E-mail: ruicao@ruc.edu.cn.
■
(26) Santoni, M. P.; La Ganga, G.; Nardo, V. M.; Natali, M.;
Puntoriero, F.; Scandola, F.; Campagna, S. J. Am. Chem. Soc. 2014,
136, 8189−8192.
Notes
The authors declare no competing financial interest.
(27) Shimazaki, Y.; Nagano, T.; Takesue, H.; Ye, B. H.; Tani, F.;
Naruta, Y. Angew. Chem., Int. Ed. 2004, 43, 98−100.
(28) Ellis, W. C.; McDaniel, N. D.; Bernhard, S.; Collins, T. J. J. Am.
Chem. Soc. 2010, 132, 10990−10991.
ACKNOWLEDGMENTS
■
We thank Prof. Ming-Tian Zhang and Xiao-Jun Su for helpful
discussions and assistance in O2 detection. We are grateful to
the support from the “Thousand Talents Program” in China,
the National Natural Science Foundation of China under
Grants 21101170 and 21203245, the Fundamental Research
Funds for the Central Universities, and the Research Funds of
Renmin University of China.
(29) Fillol, J. L.; Codola,
Costas, M. Nat. Chem. 2011, 3, 807−813.
̀ ́
Z.; Garcia-Bosch, I.; Gomez, L.; Pla, J. J.;
(30) Coggins, M. K.; Zhang, M.-T.; Vannucci, A. K.; Dares, C. J.;
Meyer, T. J. J. Am. Chem. Soc. 2014, 136, 5531−5534.
(31) Wu, Y. Z.; Chen, M. X.; Han, Y. Z.; Luo, H. X.; Su, X. J.; Zhang,
M.-T.; Lin, X. H.; Sun, J. L.; Wang, L.; Deng, L.; Zhang, W.; Cao, R.
Angew. Chem., Int. Ed. 2015, 54, 4870−4875.
(32) Dogutan, D. K.; McGuire, R., Jr.; Nocera, D. G. J. Am. Chem.
Soc. 2011, 133, 9178−9180.
(33) Gerken, J. B.; McAlpin, J. G.; Chen, J. Y. C.; Rigsby, M. L.;
Casey, W. H.; Britt, R. D.; Stahl, S. S. J. Am. Chem. Soc. 2011, 133,
14431−14442.
(34) Wasylenko, D. J.; Ganesamoorthy, C.; Borau-Garcia, J.;
Berlinguette, C. P. Chem. Commun. 2011, 47, 4249−4251.
(35) Nakazono, T.; Parent, A. R.; Sakai, K. Chem. Commun. 2013, 49,
6325−6327.
(36) Lei, H. T.; Han, A. L.; Li, F. W.; Zhang, M. N.; Han, Y. Z.; Du,
P. W.; Lai, W. Z.; Cao, R. Phys. Chem. Chem. Phys. 2014, 16, 1883−
1893.
REFERENCES
■
(1) Concepcion, J. J.; Jurss, J. W.; Brennaman, M. K.; Hoertz, P. G.;
Patrocinio, A. O. T.; Iha, N. Y. M.; Templeton, J. L.; Meyer, T. J. Acc.
Chem. Res. 2009, 42, 1954−1965.
(2) Dau, H.; Zaharieva, I. Acc. Chem. Res. 2009, 42, 1861−1870.
(3) Nocera, D. G. Acc. Chem. Res. 2012, 45, 767−776.
(4) Ferreira, K. N.; Iverson, T. M.; Maghlaoui, K.; Barber, J.; Iwata, S.
Science 2004, 303, 1831−1838.
(5) Loll, B.; Kern, J.; Saenger, W.; Zouni, A.; Biesiadka, J. Nature
2005, 438, 1040−1044.
(6) Kohl, S. W.; Weiner, L.; Schwartsburd, L.; Konstantinovski, L.;
Shimon, L. J. W.; Ben-David, Y.; Iron, M. A.; Milstein, D. Science 2009,
324, 74−77.
(37) Chen, Z. F.; Meyer, T. J. Angew. Chem., Int. Ed. 2013, 52, 700−
703.
(7) Eisenberg, R. Science 2009, 324, 44−45.
(38) Chen, Z. F.; Rathmell, A. R.; Ye, S. R.; Wilson, A. R.; Wiley, B. J.
Angew. Chem., Int. Ed. 2013, 52, 13708−13711.
(39) Zhang, M.-T.; Chen, Z. F.; Kang, P.; Meyer, T. J. J. Am. Chem.
Soc. 2013, 135, 2048−2051.
(40) Zhang, T.; Wang, C.; Liu, S. B.; Wang, J. L.; Lin, W. B. J. Am.
Chem. Soc. 2014, 136, 273−281.
(41) Coggins, M. K.; Zhang, M.-T.; Chen, Z. F.; Song, N.; Meyer, T.
J. Angew. Chem., Int. Ed. 2014, 53, 12226−12230.
(42) Su, X.-J.; Gao, M.; Jiao, L.; Liao, R.-Z.; Siegbahn, P. E. M.;
Cheng, J.-P.; Zhang, M.-T. Angew. Chem., Int. Ed. 2015, 54, 4909−
4914.
(8) Kanan, M. W.; Nocera, D. G. Science 2008, 321, 1072−1075.
́
(9) Smith, R. D. L.; Prevot, M. S.; Fagan, R. D.; Zhang, Z. P.; Sedach,
P. A.; Siu, M. K. J.; Trudel, S.; Berlinguette, C. P. Science 2013, 340,
60−63.
̆
(10) Dinca, M.; Surendranath, Y.; Nocera, D. G. Proc. Natl. Acad. Sci.
U.S.A. 2010, 107, 10337−10341.
(11) Barnett, S. M.; Goldberg, K. I.; Mayer, J. M. Nat. Chem. 2012, 4,
498−502.
(12) Bediako, D. K.; Costentin, C.; Jones, E. C.; Nocera, D. G.;
́
Saveant, J. M. J. Am. Chem. Soc. 2013, 135, 10492−10502.
(13) Wang, D.; Groves, J. T. Proc. Natl. Acad. Sci. U.S.A. 2013, 110,
15579−15584.
(43) Cao, R.; Lai, W. Z.; Du, P. W. Energy Environ. Sci. 2012, 5,
8134−8157.
(14) Zong, R. F.; Thummel, R. P. J. Am. Chem. Soc. 2005, 127,
12802−12803.
(44) Hetterscheid, D. G. H.; Reek, J. N. H. Angew. Chem., Int. Ed.
2012, 51, 9740−9747.
(15) Wasylenko, D. J.; Ganesamoorthy, C.; Henderson, M. A.;
Koivisto, B. D.; Osthoff, H. D.; Berlinguette, C. P. J. Am. Chem. Soc.
2010, 132, 16094−16106.
(45) Wasylenko, D. J.; Palmer, R. D.; Berlinguette, C. P. Chem.
Commun. 2013, 49, 218−227.
(46) Dey, S.; Mondal, B.; Dey, A. Phys. Chem. Chem. Phys. 2014, 16,
12221−12227.
(16) Concepcion, J. J.; Jurss, J. W.; Templeton, J. L.; Meyer, T. J. J.
Am. Chem. Soc. 2008, 130, 16462−16463.
(47) Bediako, D. K.; Lassalle-Kaiser, B.; Surendranath, Y.; Yano, J.;
Yachandra, V. K.; Nocera, D. G. J. Am. Chem. Soc. 2012, 134, 6801−
6809.
(17) Chen, Z. F.; Concepcion, J. J.; Jurss, J. W.; Meyer, T. J. J. Am.
Chem. Soc. 2009, 131, 15580−15581.
(18) Concepcion, J. J.; Tsai, M. K.; Muckerman, J. T.; Meyer, T. J. J.
Am. Chem. Soc. 2010, 132, 1545−1557.
(48) Singh, A.; Chang, S. L. Y.; Hocking, R. K.; Bach, U.; Spiccia, L.
Catal. Sci. Technol. 2013, 3, 1725−1732.
(49) Singh, A.; Chang, S. L. Y.; Hocking, R. K.; Bach, U.; Spiccia, L.
Energy Environ. Sci. 2013, 6, 579−586.
(19) Duan, L. L.; Fischer, A.; Xu, Y. H.; Sun, L. C. J. Am. Chem. Soc.
2009, 131, 10397−10399.
(20) Duan, L. L.; Bozoglian, F.; Mandal, S.; Stewart, B.; Privalov, T.;
Llobet, A.; Sun, L. C. Nat. Chem. 2012, 4, 418−423.
(21) McDaniel, N. D.; Coughlin, F. J.; Tinker, L. L.; Bernhard, S. J.
Am. Chem. Soc. 2008, 130, 210−217.
(50) Gao, M. R.; Sheng, W. C.; Zhuang, Z. B.; Fang, Q. R.; Gu, S.;
Jiang, J.; Yan, Y. S. J. Am. Chem. Soc. 2014, 136, 7077−7084.
(51) Zhang, M.; Zhang, M.-T.; Hou, C.; Ke, Z. F.; Lu, T. B. Angew.
Chem., Int. Ed. 2014, 53, 13042−13048.
(22) Cao, R.; Ma, H. Y.; Geletii, Y. V.; Hardcastle, K. I.; Hill, C. L.
Inorg. Chem. 2009, 48, 5596−5598.
́
(52) Collman, J. P.; Devaraj, N. K.; Decreau, R. A.; Yang, Y.; Yan, Y.
L.; Ebina, W.; Eberspacher, T. A.; Chidsey, C. E. D. Science 2007, 315,
1565−1568.
(23) Hull, J. F.; Balcells, D.; Blakemore, J. D.; Incarvito, C. D.;
Eisenstein, O.; Brudvig, G. W.; Crabtree, R. H. J. Am. Chem. Soc. 2009,
131, 8730−8731.
(53) Wasylenko, D. J.; Rodríguez, C.; Pegis, M. L.; Mayer, J. M. J.
Am. Chem. Soc. 2014, 136, 12544−12547.
(24) Blakemore, J. D.; Schley, N. D.; Balcells, D.; Hull, J. F.; Olack,
G. W.; Incarvito, C. D.; Eisenstein, O.; Brudvig, G. W.; Crabtree, R. H.
J. Am. Chem. Soc. 2010, 132, 16017−16029.
(54) McGuire, R., Jr.; Dogutan, D. K.; Teets, T. S.; Suntivich, J.;
Shao-Horn, Y.; Nocera, D. G. Chem. Sci. 2010, 1, 411−414.
I
DOI: 10.1021/acs.inorgchem.5b00924
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
(55) Kleingardner, J. G.; Kandemir, B.; Bren, K. L. J. Am. Chem. Soc.
2014, 136, 4−7.
(56) Lee, C. H.; Dogutan, D. K.; Nocera, D. G. J. Am. Chem. Soc.
2011, 133, 8775−8777.
(57) Pasternack, R. F.; Sutin, N.; Turner, D. H. J. Am. Chem. Soc.
1976, 98, 1908−1913.
(58) Chen, Z. F.; Concepcion, J. J.; Hu, X. Q.; Yang, W. T.; Hoertz,
P. G.; Meyer, T. J. Proc. Natl. Acad. Sci. U.S.A. 2010, 107, 7225−7229.
(59) Lai, W. Z.; Cao, R.; Dong, G.; Shaik, S.; Yao, J. N.; Chen, H. J.
Phys. Chem. Lett. 2012, 3, 2315−2319.
(60) Baran, J. D.; Gronbeck, H.; Hellman, A. J. Am. Chem. Soc. 2014,
̈
136, 1320−1326.
́
(61) Costentin, C.; Drouet, S.; Robert, M.; Saveant, J. M. J. Am.
Chem. Soc. 2012, 134, 11235−11242.
(62) Bard, A. J.; Faulkner, L. R. Electrochemical Methods; Wiley: New
York, 1980.
(63) Izgorodin, A.; Izgorodina, E.; MacFarlane, D. R. Energy Environ.
Sci. 2012, 5, 9496−9501.
(64) Mittra, K.; Chatterjee, S.; Samanta, S.; Dey, A. Inorg. Chem.
2013, 52, 14317−14325.
(65) Kadish, K. M.; Lin, M.; Van Caemelbecke, E.; De Stefano, G.;
Medforth, C. J.; Nurco, D. J.; Nelson, N. Y.; Krattinger, B.; Muzzi, C.
M.; Jaquinod, L.; Xu, Y.; Shyr, D. C.; Smith, K. M.; Shelnutt, J. A.
Inorg. Chem. 2002, 41, 6673−6687.
(66) Ertem, M. Z.; Gagliardi, L.; Cramer, C. J. Chem. Sci. 2012, 3,
1293−1299.
(67) Ertem, M. Z.; Cramer, C. J. Dalton Trans. 2012, 41, 12213−
12219.
(68) Bozoglian, F.; Romain, S.; Ertem, M. Z.; Todorova, T. K.; Sens,
C.; Mola, J.; Rodríguez, M.; Romero, I.; Benet-Buchholz, J.;
Fontrodona, X.; Cramer, C. J.; Gagliardi, L.; Llobet, A. J. Am. Chem.
Soc. 2009, 131, 15176−15187.
́
(69) Sartorel, A.; Miro, P.; Salvadori, E.; Romain, S.; Carraro, M.;
Scorrano, G.; Di Valentin, M.; Llobet, A.; Bo, C.; Bonchio, M. J. Am.
Chem. Soc. 2009, 131, 16051−16053.
(70) Sala, X.; Ertem, M. Z.; Vigara, L.; Todorova, T. K.; Chen, W. Z.;
Rocha, R. C.; Aquilante, F.; Cramer, C. J.; Gagliardi, L.; Llobet, A.
Angew. Chem., Int. Ed. 2010, 49, 7745−7747.
(71) Wang, L. P.; Wu, Q.; Van Voorhis, T. Inorg. Chem. 2010, 49,
4543−4553.
(72) Bianco, R.; Hay, P. J.; Hynes, J. T. J. Phys. Chem. A 2011, 115,
8003−8016.
(73) Li, X. C.; Chen, G. J.; Schinzel, S.; Siegbahn, P. E. M. Dalton
Trans. 2011, 40, 11296−11307.
(74) Vigara, L.; Ertem, M. Z.; Planas, N.; Bozoglian, F.; Leidel, N.;
Dau, H.; Haumann, M.; Gagliardi, L.; Cramer, C. J.; Llobet, A. Chem.
Sci. 2012, 3, 2576−2586.
(75) Romera, C.; Sabater, L.; Garofalo, A.; Dixon, I. M.; Pratviel, G.
Inorg. Chem. 2010, 49, 8558−8567.
(76) Bruker AXS. APEX2 v2009; Madison, WI, 2009.
(77) Sheldrick, G. M. SADABS, 2008/1; University of Gottingen:
̈
Gottingen, Germany, 2008.
̈
(78) Sheldrick, G. M. Acta Crystallogr. 1990, A46, 467−473.
(79) Sheldrick, G. M. Acta Crystallogr. 2008, A64, 112−122.
(80) Frisch, M. J.; et al. Gaussian 09, Revision C.01, Gaussian, Inc.:
Wallingford, CT, 2010.
(81) Schafer, A.; Horn, H.; Ahlrichs, R. J. Chem. Phys. 1992, 97,
̈
2571−2577.
(82) Schafer, A.; Huber, C.; Ahlrichs, R. J. Chem. Phys. 1994, 100,
̈
5829−5835.
(83) Barone, V.; Cossi, M. J. Phys. Chem. A 1998, 102, 1995−2001.
(84) Kelly, C. P.; Cramer, C. J.; Truhlar, D. G. J. Phys. Chem. B 2006,
110, 16066−16081.
(85) Wang, Y.; Kumar, D.; Yang, C. L.; Han, K. L.; Shaik, S. J. Phys.
Chem. B 2007, 111, 7700−7710.
J
DOI: 10.1021/acs.inorgchem.5b00924
Inorg. Chem. XXXX, XXX, XXX−XXX