A biomimetic copper water oxidation catalyst with low overpotential

Article abstract of DOI:10.1021/ja409267p

Simply mixing a Cu(II) salt and 6,6′-dihydroxy-2,2′-bipyridine (H₂L) in a basic aqueous solution afforded a highly active water oxidation catalyst (WOC). Cyclic voltammetry of the solution at pH = 12-14 shows irreversible catalytic current with an onset potential of ～0.8 V versus NHE. Catalytic oxygen evolution takes place in controlled potential electrolysis at a relatively low overpotential of 640 mV. Experimental and computational studies suggest that the L ligand participates in electron transfer processes to facilitate the oxidation of the Cu center to lead to an active WOC with low overpotential, akin to the use of the tyrosine radical by Photosystem II to oxidize the CaMn₄ center for water oxidation.

Full text of DOI:10.1021/ja409267p

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Article
A Biomimetic Copper Water Oxidation Catalyst with Low Overpotential
Teng Zhang, Cheng Wang, Shubin Liu, Jin-Liang Wang, and Wenbin Lin
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A Biomimetic Copper Water Oxidation Catalyst with Low
Overpotential
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Teng Zhang,†,1 Cheng Wang,†,1,2, Shubin Liu,3 Jin-Liang Wang,2 Wenbin Lin*,1
1Department of Chemistry, University of Chicago, 929 E. 57th Street, Chicago, IL 60637, USA
2Department of Chemistry, University of North Carolina, Chapel Hill, NC 27599, USA
3Research Computing Center, University of North Carolina, Chapel Hill, NC 27599, USA
ABSTRACT: Simply mixing a Cu(II) salt and 6,6’-dihydroxy-2,2’-bipyridine (H2L) in a basic aqueous solution afforded a
highly active water oxidation catalyst (WOC). Cyclic voltammetry of the solution at pH = 12-14 shows irreversible catalytic
current with an onset potential of ~0.8 V vs. NHE. Catalytic oxygen evolution takes place in controlled potential electroly-
sis at a relatively low overpotential of 640 mV. Experimental and computational studies suggest that the L ligand partici-
pates in electron transfer processes to facilitate the oxidation of the Cu center to lead to an active WOC with low overpo-
tential, akin to the use of the tyrosine radical by Photosystem II to oxidize the CaMn4 center for water oxidation.
Introduction
between the catalytic center, the CaMn4 cluster, and the
oxidant, the photochemically generated P680+. Mean-
while, tyrosine Z as well as the adjacent histidine residue
His-190 also participates in the “proton-rocking” process,
enabling the proton-coupled-electron-transfer (PCET)
mechanism for water oxidation reaction.
Splitting water into hydrogen and oxygen is one of the
most attractive scenarios for solar energy harvesting and
sustainable energy production.1-7 One of the key challeng-
es to water splitting is the development of efficient cata-
lysts for the water oxidation half reaction with low over-
potentials, good stability, and high turnover rates.8,9 Dis-
covering efficient catalysts for the water oxidation reac-
tion has rekindled great interest in recent years, and vari-
ous molecular and non-molecular catalysts based on no-
ble10-17 and earth-abundant18-28 metals have been reported.
Despite much progress in water oxidation catalysis, major
improvements in several areas, including lowering over-
potentials, increasing catalyst durability, and using earth-
abundant elements, are needed before efficient photo-
catalytic water splitting can be realized.
H+
HN
N
O
O
H H+
e-
e-
O
N
N
Ca
Cu
O
O
O
OH2
Mn
Mn
Mn
O
Mn
O
O
Recently, copper-based water oxidation catalysts
(WOCs) such as the Cu-bipy system,25 the Cu-carbonate
system28 and the Cu-peptide system29 have attracted great
interest due to the high abundance and low cost of Cu
and the simplicity of these systems. Although these cop-
per-based WOCs show high current densities and good
stabilities under basic conditions, difficulties in accessing
high-oxidation-state copper species (CuIII and/or CuIV)
lead to high overpotentials for water oxidation and se-
verely limit their practical utility. It is thus highly desira-
ble to identify new strategies to lower overpotentials of
these promising Cu-based WOCs.
Scheme 1. A typical PCET process in Photosystem II
(left) and the proposed ligand-assisted PCET in the
Cu-L system (right).
We report here the use of a ligand containing suitable
pendant groups to mimic the functions of tyrosine Z in
facilitating the oxidation of the Cu center to lead to a
more active WOC. We designed a copper-based WOC
with 6,6’-dihydroxy-2,2’-bipyridine (H2L) as the ligand,
trying to mimic the role of tyrosine Z in Photosystem II by
not only providing a redox-accessible ligand but also hav-
ing the hydroxyl groups participating in the PCET pro-
cesses, to lower the overpotential and enhance the WOC
activity (Scheme 1). The H2L ligand has been used to syn-
thesize rhodium, ruthenium, and iridium complexes as
catalysts for organic transformations,38-41 CO2 reduction,42
and, during the preparation of this manuscript, water
oxidation.43 The pendant hydroxyl groups were proposed
to assist hydride transfer in the catalytic cycles,42 howev-
er, the redox activity of the H2L ligand has not yet been
As in-depth knowledge of the natural water oxidation
reaction becomes available,30-37 more efficient molecular
WOCs with lower overpotentials can be designed by
mimicking the key features of the water oxidation process
in natural photosynthesis. In Photosystem II, a redox-
active tyrosine residue, usually referred as tyrosine Z or
YZ, serves as a mediator in the electron transfer process
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Page 2 of 9
Figure 2. CVs of the Cu-L complex (red) and the parent Cu-
utilized. The idea of using redox active, non-innocent
ligands to affect metal reactivities has recently been ex-
ploited to enhance catalytic activities,11,16,44-46 and exam-
ples of ligand-assisted proton transfer have also been re-
ported.47,48
1
2
3
4
bipy complex (CuLa, blue) in basic aqueous solution (pH =
12.4, 0.1 M NaOH/NaOAc electrolyte) on a glassy carbon
electrode (a, S = 0.07 cm2) and an ITO electrode (b, S = 1
cm2).
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Results and Discussion
To verify and quantify the electrochemical oxygen gen-
eration, control potential electrolysis (CPE) experiments
were carried out at a potential of 0.90V vs. Ag/AgCl refer-
ence electrode (1.135V vs. NHE, η = 640mV) with an ITO
electrode in a 0.1 M NaOAc/NaOH solution at pH = 12.4
(Figure 3). A substantial current of >0.15 mA/cm2 was
maintained during the whole period of electrolysis (Fig-
ure S5, SI). The amount of oxygen was determined by gas
chromatography, and the Faradaic efficiency was calcu-
lated to be (85±5)% based on total charge passing
through. CPE experiments under the same pH and the
same potential with a catalyst-free solution only gave a
current density of 0.003 mA/cm2 and a negligible amount
of oxygen, showing that the copper complex does serve as
an effective WOC under such conditions.
Electrochemical water oxidation of the CuL com-
plex. H2L was dissolved in a basic aqueous solution, and
a Cu(II) salt (such as CuSO4) was added to form a blue-
green solution. Cyclic voltammograms (CVs) were ob-
tained at pH = 11.0-14.0. The CVs showed large, irreversi-
ble oxidative waves that correspond to catalytic water
oxidation (Figure 1 and S1, supporting information [SI]).
In the pH range of 11.0-12.2, complicated behaviors were
observed, which are attributed to oligmerization or
polymerization of the CuL complex. The catalytic cur-
rents were well above the background level when the so-
lution contained no catalyst (Figure S2, SI), showing that
the CuL complex is responsible for the catalytic activity.
CV of a solution containing only ligand L showed the lig-
and oxidation peak at a similar position (Figure S2, SI).
The water oxidation occurs at a relatively low overpoten-
tial (η) of 510-560 mV based on the half-peak potential of
the CVs in pH 12.0-14.0, significantly lower than the pre-
viously-reported parent Cu-bipy compound (also denoted
as CuLa in Scheme 2) by about 200 mV (Figure 2).25 This
trend is consistent on both glassy carbon and ITO elec-
trodes, though on ITO electrodes the current shows a
linear relationship to the potential. The CVs and water
oxidation catalytic currents observed for the Cu-L system
did not depend on the choices of Cu(II) salts or inert elec-
trolyte (Figure S3-4, SI), indicating that the counterions
do not coordinate to the Cu centers during the water oxi-
dation reaction.
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2
1
0
a)
b)
O2 N2
CH4 (internal standard)
0
50
100
Time (min)
150
0
1
2
3
4
5
6
time / min
Figure 3. a) Control potential electrolysis of a solution con-
taining 1 mM catalyst (black solid line) and the buffer only
(red dash line) on an ITO electrode (S = 1 cm2). b) GC traces
from a representative CPE experiment. Blue line shows the
air background. Oxygen amount is determined by comparing
relative intensity of oxygen and internal standard peaks.
0.0
CPE experiments were also performed with a large-area
glassy carbon electrode. A current density of 0.2 mA/cm2
was achieved and a Faradaic efficiency of 59% was calcu-
lated based on total charge passing through. The lower
Faradaic efficiency is due to oxidation of the carbon elec-
trode (Figure S6, SI).
-0.1
-0.2
pH 14.0
pH 13.6
pH 13.2
pH 12.8
pH 12.4
pH 12.0
pH 11.0
The catalyst stability was examined by taking UV-Vis
spectra before and after electrolysis. When CPE experi-
ments were performed in a fritted cell as in oxygen quan-
tification experiments, little to no depletion of UV-Vis
signal was observed after 3 h of electrolysis (Figure S7, SI).
An average turnover number of ~1 O2/Cu was reached
based on the total CuL amount in the solution, but a
more realistic TON of ~400 was estimated based on the
amount of CuL involving in electrolysis (SI).49 We also
recovered the free ligand H2L by first acidifying the solu-
tion and then precipitating H2L by removing Cu(II) ions
with silica-supported triamine tetraacetate. NMR spec-
troscopy (Figure S8, SI) indicated the recovery of >90% of
the H2L ligand, which not only supports the catalytic na-
ture of the electrolysis and stability of the complex but
also proves water (hydroxide) as the source of oxygen.
However, when electrolysis was performed in a spectroe-
0.2
0.4
0.6
0.8
1.0
1.2
E vs. Ag/AgCl (V)
Figure 1. CVs of solutions containing 1 mM of CuSO4 and L2-
at various pH’s on a glassy carbon electrode (S = 0.07 cm2).
a)
0
-2
-4
b)
0.0
-0.2
-0.4
-0.6
0.4
0.6
0.8
1.0
1.2
1.4
0.4
0.6
0.8
1.0
1.2
1.4
E vs. Ag/AgCl (V)
E vs. Ag/AgCl (V)
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(≤ 20 mV/s). The ratio of the catalytic current over the
lectrochemical cell with no frit or membrane to separate
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the electrode compartments, the ligand absorption max-
ima in the UV-Vis spectra decreased by ~30% after the
same operating time and a similar turnover number based
on bulk concentration (Figure S9, SI). The catalyst loss is
thus attributed to reduction of Cu(II) ions on the counter
electrode, releasing free ligand which is readily oxidized.
Similar Cu(II) reduction was also observed for the Cu-La
system on the counter electrode.25
diffusive peak current, icat/id, matches the trend from a
pure diffusion behavior (DP) region to a pure kinetic be-
havior (KP) region as scan rates decrease (Figure 4b).50 In
the KP region, the catalytic peak current in cyclic volt-
ammetry is given by equation 2:50
ꢀꢎꢏꢐ ꢂ ꢃꢎꢏꢐꢄꢅꢆ
ꢎꢏꢐꢑ
ꢒ
ꢎꢏꢐꢓ
(Eq 2)
where ncat = 4 is the number of electrons transferred in
the catalytic reaction (water oxidation), F is Faraday’s
constant, A is the area of the electrode, ccat is the bulk
catalyst concentration, kcat is the apparent first-order rate
constant, and D is the diffusion coefficient of the catalyst.
9
During electrolysis, a deposit was observed on the elec-
trode. Further studies prove that the film is oligomer or
polymer of the CuL complex instead of metal oxide. First,
the freshly deposited film showed catalytic water oxida-
tion activity in a fresh pH 12.4 buffer without the catalyst,
but dissolved back into the solution after 30 min of elec-
trolysis. By taking UV-Vis spectra (Figure S10, SI) and
ICP-MS analyses of the solution of the dissolved film, a
Cu/L ratio of 1.12±0.15 was determined, indicating that the
film is likely a coordination polymer containing the ligand
as well as Cu, but not copper oxides or hydroxides. Se-
cond, the film formation was not observed at higher pH’s.
If copper oxide or hydroxide presents, its formation
should be facilitated by a higher base concentration,
which is opposite to our observation.
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By taking the ratio of equation 2 over equation 1 and
plugging in all of the constants, we obtain equation 3:
ꢔ
ꢙ
ꢕꢖꢗ
ꢕꢖꢗ
ꢊ
ꢂ 1.424ꢇ
(Eq 3)
ꢔ
ꢘ
The pseudo-first order rate constant of the catalytic wa-
ter oxidation (or hydroxide oxidation at high pH), kcat, is
usually referred as turnover frequency (TOF) of the cata-
lyst in the literature. From the slope of the plot of icat/id vs.
v-1/2, an apparent kcat = 0.4 s-1 was calculated. This value is
comparable to most reported molecular water oxidation
catalysts,14,22,51 yet much lower than the recently reported
copper based ones.25,29 However, as the 4-electron water
oxidation reaction is much more complicated than the
simplified model of ErCcat reaction, the calculated kcat val-
ue can only serve as an estimate of the catalytic rate.
The electrochemical deposit behaves differently from a
direct deposit of the CuL complex on an ITO electrode
prepared by drying a methanol solution of Cu-L via sol-
vent evaporation. The direct deposit readily dissolves in a
basic solution, while the electrochemical deposit only
dissolves when a potential bias high enough to drive wa-
ter oxidation is applied as described above. This different
behavior can be explained by partial oxidation of the elec-
trochemical deposit which would make the coordination
polymer less soluble. Additional catalytic cycles are need-
ed to return the coordination polymer to its “original”
state (i.e., the reduced state) before it can dissolve in a
basic solution.
a)
0
100mV/s
50mV/s
20mV/s
15mV/s
10mV/s
8mV/s
-50
-100
-150
5mV/s
2mV/s
At longer electrolysis time the film grows thicker to
form an insulating layer on the electrode and, leading to a
decrease of electrolysis current. Such “thick” films have
Cu/L ratios larger than 1, indicating ligand decomposition
in the deposit. This at least partially explains the lower
Faradaic efficiency for the Cu-L system when compared to
other copper-based water oxidation systems.
-0.8
-0.4
0.0
0.4
0.8
1.2
E vs. Ag/AgCl (V)
b) 24
20
CVs were recorded at different scan rates in order to
obtain kinetic information of the Cu-L system (Figure 4a).
It is convenient to compare the catalytic current to a re-
versible diffusive current, which is given by equation 1:50
16
12
8
ꢀꢁ ꢂ 0.4633ꢃꢄꢅꢆꢇꢈꢉꢊꢋ
(Eq 1)
ꢌꢍ
where n is the electron transferred in the non-catalytic
reaction, v is the scan rate, R is the gas constant, and T is
the absolute temperature.
0
6
12
v-1/2 (s1/2
18
24
·
V-1/2
)
Figure 4. a) CVs of a 1 mM Cu/L solution under various scan
rates on a glassy carbon electrode (S = 0.07 cm2, pH 12.4, 0.1
M NaOH/NaOAc, background corrected). b) Linear fitting
plot of icat/id vs. v-1/2 for TOF calculations.
The diffusive peak current in the absence of catalytic
processes was estimated using the diffusion-controlled
Cu(II)/Cu(I) peak at ~-0.4 V vs. Ag/AgCl (Figure 4a). The
Cu(II)/Cu(I) peak is quasi-reversible at high scan rates
(100 mV/s and 50 mV/s) while reversible at low scan rates
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Characterization of aqueous species in Cu-L solu- ior at high Cu/L concentrations (Figure 6). This equilibri-
1
2
3
4
5
6
7
8
tions. Small needle-like crystals were observed by adjust-
ing the pH of an aqueous CuL solution from highly basic
conditions (pH > 12) to ~11.0. These crystals are of the
same phase as single crystals of a coordination polymer
grown from a mixed methanol/water solution as indicat-
ed by their identical PXRD patterns (Figure S14, SI). A
single crystal X-ray diffraction study shows that in the
crystal, the Cu center and the L ligand form 1-D coordina-
tion polymer with the formula of [Cu(HL)(ꢀ2-OH)]n. In
the crystal structure, each Cu center adopts a slightly dis-
torted square planar coordination environment with a
mono-deprotonated ligand HL- and two bridging hydroxo
groups. Two oxygen atoms from two adjacent HL ligands
sit at the axial positions with Cu-O distances of 2.690(3)
and 3.076(4) Å, respectively, showing very weak interac-
tions with the Cu center. The O···O distances between a
µ-hydroxo group and its nearest neighbors from the HL-
ligands are 2.452(5) and 2.497(4) Å, respectively, indicat-
ing the existence of hydrogen bonding (Figure 5 and Fig-
ure S15, SI).
um is also proposed to be responsible for the film for-
mation, presumably owing to the local pH decrease dur-
ing the control potential electrolysis.
a)
0
1mM
-20
800
500
400
200
100
ꢀ
ꢀ
ꢀ
ꢀ
ꢀ
M
M
M
M
M
-40
-60
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80ꢀ
50ꢀ
40ꢀ
20ꢀ
10ꢀ
M
M
M
M
M
-80
-100
0.2
0.4
0.6
0.8
1.0
1.2
E vs. Ag/AgCl (V)
b) 120
80
40
0
0
200
400
600
800
1000
concentration (
ꢀ
M)
Figure 6. a) CVs of Cu/L solutions at various concentrations
on a glassy carbon electrode (S = 0.07 cm2, pH 12.4, 0.1 M
NaOH/NaOAc, v = 100 mV/s, background corrected). b) The
relationship between the concentration and CV peak current
of the CuL complex.
12
9
Figure 5. X-ray structure of 1-D polymer [Cu(HL)(OH)]n.
Atoms shown at 50% probability. Red: O; blue: N; yellow:
Cu; white: C.
6
Precipitate appeared
We observed precipitate formation at pH 12.4 when the
concentration of the CuL complex exceeded 2 mM, sug-
gesting the presence of an equilibrium between monomer
and oligomer/polymer species in solution. Hydroxo-
bridged dinuclear species are well-known for Cu-bipy or
Cu-diamine complexes in neutral to weakly basic aqueous
solutions.52-55 However, in the Cu-L system, formation of
the Cu2(µ2-OH)2 dimer is less favorable due to the steric
effect of the 6,6’-substituents. Formation of oligomer or
polymer is thus favored because less steric hindrance is
experienced in the oligomer/polymer species. The pres-
ence of polynuclear species was also supported by ESI-MS
studies (Figure S16, SI). An equilibrium between mono-
mer and oligomer/polymer species also accounts for an
unusual trend of concentration-dependent CVs, in which
the catalytic peak currents showed a curving-over behav-
H2L dissolved
3
0
20
40
ꢀ
60
based added (
mol)
Figure 7. Titration curve of 10 mL of 1 mM Cu/H2L solution.
X
Y
N
N
L
a: X = H, Y =H
H2Lb: X = OH, Y = H
Lc: X = OMe, Y = H
L
H2 d: X = COOH, Y = H
X
Y
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Scheme 2. A list of 2,2’-bipy derivatives used in this
work.
better identifies the positions of redox peaks due to its
higher sensitivity, and it also gives redox potentials under
catalytic conditions.56,57
1
2
3
4
5
6
7
8
Titration curve of the Cu-L system was also obtained
SWVs under the same conditions showed a pH-
dependent peak between 1.01-1.07 V vs. NHE (Figure 8a).
As indicated by CVs and CPE experiments, this peak is
responsible for catalytic water oxidation. Peaks at higher
potentials were observed at lower pH’s, which are at-
tributed to protonation and polymerization of the copper
species.
(Figure 7). At the initial state, the H2L ligand did not dis-
solve in water. It slowly dissolved upon addition of base
and began to form a blue-green complex with the Cu(II)
ion. However, a blue-green precipitate formed after ~1.5
equiv. of base was added, and a titration jump was ob-
served. This is attributed to dimerization or polymeriza-
tion of the Cu(HL)+ and CuL complexes. A second titra-
tion jump was observed upon addition of ~2 equiv. of
base, indicating a complete transformation to the neutral
[Cu(HL)(OH)]n(OH2)x polymer. The precipitate, presum-
ably the neutral polymer, did not fully dissolve until ex-
cess (> 20 eq) of base was added. The titration experiment
thus indicates that an anionic species, CuL(OH)(OH2)-, is
the dominant species in the catalytically active solution,
and possibly the true catalyst.
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Interestingly, CuL serves as the only example with wa-
ter oxidation observed at a lower overpotential than the
parent CuLa complex. SWVs of the parent complex Cu-
La(OH)2 in basic aqueous solution show a small, pH-
dependent peak between 0.93-1.03 V vs. NHE (Figure 8b)
which is attributed to the Cu(III)/Cu(II) redox pair. A
second oxidation occurs at around ~1.23 V vs. NHE and
leads to catalytic water oxidation. Similar behaviors were
observed in 4,4’-substituted systems Cu-Lc and Cu-Ld
(Figure 8d-e). Both the Cu(III)/Cu(II) peak and the water
oxidation peak are only slightly effected by substituents at
4,4’ positions, consistent with the previous report by
Mayer and coworkers.25
Electrochemistry of analogous systems. In order to
elucidate the water oxidation mechanism and the reason
for the lower overpotential observed for the Cu-L system,
we carried out electrochemical studies on the M-L sys-
tems as well as related M-bipy systems. Scheme 2 shows a
list of the 2,2’-bipy derivatives we examined. Square wave
voltammetry (SWV) was used instead of CV because it
0
0
0
b)
a)
c)
-80
-160
-240
-50
pH 14.0
pH 13.8
pH 13.6
pH 13.4
pH 13.2
pH 13.0
pH 12.5
pH 14.0
pH 13.6
pH 13.2
pH 14.0
pH 13.6
pH 13.2
pH 12.8
pH 12.4
pH 12.0
-50
pH 12.8
pH 12.4
pH 12.0
-100
-100
-150
0.4
0.6
0.8
1.0
1.2
0.4
0.6
0.8
1.0
1.2
0.2
0.4
0.6
0.8
1.0
1.2
1.4
E vs. Ag/AgCl (V)
E vs. Ag/AgCl (V)
E vs. Ag/AgCl (V)
d) 0
0
-10
-20
-30
-40
-50
0
f)
e)
-10
-20
-30
-40
-50
pH 14.0
pH 13.6
pH 13.2
pH 12.8
pH 12.4
pH 12.0
pH 11.4
pH 10.9
pH 8.9
pH 14.0
pH 13.5
pH 13.0
pH 12.5
pH 12.0
pH 14.0
pH 13.6
pH 13.2
pH 12.9
pH 12.0
-50
pH 8.2
-100
0.4
0.6
0.8
1.0
1.2
0.2
0.4
0.6
0.8
1.0
1.2
0.4
0.6
0.8
1.0
1.2
E vs. Ag/AgCl (V)
E vs. Ag/AgCl (V)
E vs. Ag/AgCl (V)
h) 0
0
g)
-10
-20
-30
-40
-8
pH 14.0
pH 13.6
pH 13.2
pH 12.8
pH 12.4
pH 12.0
Cu-L
Cu-Lb
-16
Zn-L
-24
0.4
-50
0.2
0.8
1.2
1.6
0.4
0.6
0.8
1.0
1.2
E vs. Ag/AgCl (V)
E vs. Ag/AgCl (V)
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Figure 8. SWVs of aqueous solutions containing 1 mM of (a) Cu2+ and L2-; (b) Cu2+ and La; (c) Cu2+ and Lb2-; (d) Cu2+ and Lc; (e)
Cu2+ and Ld2-; (f) L2-; (g) Zn2+ and L2- under various pH’s (glassy carbon electrode, S = 0.07 cm2, 1 M KOH/KNO3). (h) SWVs of 1
mM CuL, CuLb, and ZnL in mix DMF/H2O solution with 2 eq. of KOH (glassy carbon electrode, S = 0.07 cm2, 0.1 M NMe4BF4).
1
2
3
4
5
6
7
8
Surprisingly, the 4,4’-substituted analog of H2L, 4,4’-
dihydroxy-2,2’-bipyridine (H2Lb), showed an even higher
overpotential for water oxidation than the parent complex
CuLa(OH)2 as indicated by SWVs (Figure 8b-c). The dif-
ferences between CuL and CuLb complexes indicate that
the lower overpotential observed in the CuL complex is
not simply due to the electron-donating nature of the
substituents or introduction of additional negative
charge.
dependent UV-Vis spectra of the free ligand (Figure 9)
show a blue shift of the absorption maximum upon pro-
tonation from L2- to HL-, and a red shift upon protonation
from HL- to H2L. The UV-Vis spectra of CuL and ZnL
complexes show absorption peaks at a similar position as
HL- (Figure S21, SI). The similarity of the UV-Vis spectra
of HL- and CuL or ZnL complexes suggests similarity of
9
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45
46
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48
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53
54
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58
59
60
their structures.
A planar structure of the mono-
protonated form HL- is proposed (Scheme 3), which is
corroborated by DFT calculations (Figure S22, SI). Given
similar structures of HL- and CuL or ZnL complexes and
similar electronic effects of H+, Cu(OH)(OH2)+, and
Zn(OH)(OH2)+, it is not surprising to see similar ligand
oxidation potentials of the three.
To further investigate the roles of 6,6’-substituted hy-
droxyl groups in water oxidation, SWVs were also record-
ed in a DMF/H2O (v/v 9:1) solution to reduce the catalytic
signal. One equivalent of copper trifluorosulfonate or zinc
trifluorosulfonate, one equivalent of H2L or H2Lb ligand,
and two equivalent of KOH were dissolved in the
DMF/H2O mixed solvent to afford a solution with a metal
concentration of 1 mM. Two sequential redox peaks were
observed in the Cu-L solution at 1.18 and 1.35 V vs.
Ag/AgCl, while only the first peak was observed in the
Cu-Lb solution (Figure 8h). Thus, the first peak was as-
signed to the Cu(III)/Cu(II) redox pair. An additional re-
dox peak was also observed in the Cu-Lb solution at 0.82
V vs. Ag/AgCl, which could be assigned to some DMF-
coordinated copper species. The second peak was as-
signed to ligand oxidation by comparing to the Zn-L sys-
tem, for which a redox peak was observed at 1.33 V vs.
Ag/AgCl (Figure 8h). As the ligand oxidation potential is
very close to the Cu(III)/Cu(II) potential, we propose that
ligand oxidation is involved in the catalytic cycle and
plays a key role in lowering the overpotential for water
oxidation.
16000
12000
pH 7
pH 8
pH 9
pH 10
pH 11
pH 12
pH 13
8000
4000
0
300
350
400
450
500
λ
/ nm
Figure 9. UV-Vis spectra of free L2- ligand in aqueous solu-
tion at various pH’s.
We also compared SWVs of the CuL complex with free
ligand L2-/HL- and analogous ZnL complex in aqueous
solution to confirm the ligand oxidation in catalytic cycle.
Free ligand shows pH-independent oxidation peak at ~1.01
V vs. NHE between pH 14.0-12.8, and shifts to ~1.12 V vs.
NHE between pH 12.8-9.0. The shift on peak position is
attributed to the protonation from L2- to HL-. A second
protonation was observed around pH 8.5 as indicated by
another peak position shift, corresponding to the proto-
nation from HL- to H2L (Figure 8f). The second protona-
tion is consistent with the titration curve of L2- (Figure
S18, SI), and both protonation steps are consistent with a
change on UV-Vis spectra (Figure 9). The ZnL complex
shows a pH-independent redox peak at ~1.16 V vs. NHE
(Figure 8g). We propose that the dominant species of the
ZnL complex is ZnL(OH)(OH2)- as indicated by a titration
curve (Figure S20, SI), and assign the redox peak to the
ligand oxidation of the species.
Scheme 3. Proposed structures for mono-
deprotonated and di-deprotonated forms of H2L.
Computational studies on the catalysis mecha-
nism. The catalytic cycle for water oxidation is proposed
in Scheme 4b. DFT calculations were performed on the
Cu-L complex for a better understanding of the water
oxidation mechanism. As shown in Scheme 4a, the redox
active L2- ligand plays an important role in the second
oxidation step from a Cu(III) species to an apparent
“Cu(IV)” species. As shown by calculation results, in the
parent Cu-bipy complex the second oxidation can only
take place at the hydroxo group with a calculated oxida-
tion potential of 1.4 V vs NHE due to the difficulty in oxi-
dizing the bipy ligand (Scheme S1, SI). However, in the
Cu-L complex, the second oxidation takes place at the
ligand as indicated by the ligand-centered singly occupied
As shown in Figure 8f and 8g, the oxidation of the L lig-
and is little affected by metal coordination when the met-
al complexes bear the same charge as the protonated free
ligand. It could be explained by the structural similarity
between HL- and CuL or ZnL complexes. The pH-
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Journal of the American Chemical Society
pKa’s and redox potentials are not so informative. Fortu-
molecular orbital (SOMO) of 9 (Figure 10). The calculated
1
2
3
4
5
6
7
8
spin density (Figure 10) of 9 mainly locates at the ligand,
indicating that it is more like a Cu(III) species with the
ligand having the radical character. By involving the lig-
and in the oxidation process, the second redox potential
is lowered by ~0.3 V as indicated by the DFT calculations
compared to the corresponding oxidation step for the
parent Cu-bipy complex (Scheme 4 and Scheme S1, SI),
leading to stabilization of the highly oxidized intermedi-
ate. Thus, the redox active L2- ligand serves as an electron
transfer mediator in the catalytic cycle and is responsible
for the low overpotential observed, and water oxidation is
found to be driven by a Cu(III) center other than a hard-
er-to-access Cu(IV) one. However, as the oxyl-centered
radical character of a Cu(IV)-oxo species58 (i.e., 6’ in the
Cu-bipy system, Scheme S1, SI) makes it more electro-
philic and reactive than the Cu(III)-hydroxyl species 9,
the formation of peroxo species 11 (Scheme 4b), which is
believed to be the rate determine step of the water oxida-
tion process, is probably disfavored in the case of the Cu-
L complex. This could in part explain the lower current
density and the slower TOF value for the Cu-L complex
compared to the parent Cu-bipy system (Figure 2a).
nately, the trends we obtained from calculations on the
Cu-L, Cu-bipy, and Cu-Lb systems match with our exper-
imental results, lending support to our conclusions.
Concluding Remarks
The introduction of 6,6’-dihydroxyl groups on the bipy
moiety allows the L ligand to be intimately involved in
the water oxidation catalytic cycle via ligand oxidation
and significantly lowers the water oxidation overpoten-
tial. This work thus uncovers an effective biomimetic
strategy toward ligand design for preparing highly effi-
cient WOCs.
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60
Due to the difficulty in characterizing the intermediates
in electrocatalysis, we were not able to experimentally
prove or disprove the intramolecular PCET process. Fur-
ther experimental and theoretical studies are under way
to fully unravel the detailed mechanisms and to design
more efficient WOCs based on our knowledge of this bi-
omimetic strategy.
Experimental Section
Electrochemistry and Gas Analysis. Cyclic voltam-
mograms (CVs), square wave voltammograms (SWVs),
and controlled potential electrolysis (CPE) were recorded
on a CHI420 electrochemistry work station. Regular 3-
electrode systems were used. CVs and SWVs were record-
ed using a glassy carbon disk working electrode (S = 0.07
cm2) unless noted. Bulk electrolysis was performed in a
gas-tight fritted cell to separate the cathode and the an-
ode. No further calibration on iR drop was applied. A
Ag/AgCl/1M KCl electrode was used as the reference elec-
trode in all experiments and its potential (0.235 V vs.
Figure 10. Calculated SOMO (left) and spin density (right) of
structure 9.
4-
NHE) was calibrated with the Fe(CN)63-/ Fe(CN)6 cou-
ple.59 For CVs and SWVs in aqueous solution, the electro-
lyte (1 M) was prepared by titrating 1 M KOH solution
with HNO3 to the desired pH. 0.1 M solution of NMe4BF4
in DMF/H2O (90:10, v/v) was used for non-aqueous CVs
and SWVs. The pH 12.4 buffer was prepared by dissolving
NaOH (0.025 M) and the inert salt (NaOAc, NaNO3, or
NaClO4, 0.075M) in the same solution then adjusting to
the desired pH with the corresponding acid. Analysis of
the gas product in CPE experiments was conducted by gas
chromatography (Varian 450-GC, molecular sieve col-
umns, pulsed discharge helium ionization detector,
PDHID). Background from air was calibrated with respect
to the nitrogen signal, and the amount of oxygen was de-
termined by comparing to a methane internal standard.
Calculation on the Cu-Lb system uncovers why the
CuLb complex does not serve as an effective water oxida-
tion catalyst as CuL. Although CuLb(OH)(OH2)- exhibits a
lower Cu(III)/Cu(II) redox potential, the second oxidation
on ligand is ~0.2 V higher than the 6,6’-substituted analog
(Scheme S2, SI), consistent with the electrochemical be-
haviors observed in both aqueous and non-aqueous solu-
tions. The higher ligand oxidation potential for Lb thus
explains why CuLb exhibits a higher overpotential for wa-
ter oxidation.
However, we should note that accurate prediction of
pKa’s and redox potentials is still a challenging problem
with calculated values significantly deviating from exper-
imental ones. Thus, the absolute values of the calculated
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0
-
2-
O-
O-
O-
-
a)
b)
1
2
3
4
5
6
7
8
O-
H
H
H
N
N
N
N
N
N
OH
CuII
O
CuII
O
H
pKa=10.9
pKa=21.7
CuII
N
N
O
CuII
-e-, -H+
OH
H
OH
H
O
H
OH
O-
O-
O-
H
-
O-
2
3
1
O-
H
O
CuIII
2
1.4V
0.4V
2.4V
N
N
+
0
-
2-
O-
O-
O-
O-
-2e-, -2H+
-O2, +H2O
O
H
H
H
9
N
N
N
N
N
N
N
N
H
OH
CuIII
O
CuIII
O
CuIII
O
CuIII
pKa=-7.1
pKa=36.5
pKa=6.1
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
O-
6
OH
H
OH
H
O
H
O
H
O-
O-
O-
O-
-e-
5
6
7
4
-
O-
O
H
1.6V
1.1V
-0.2V
H
O
0
-
+
O
O
O
N
N
N
OH
O
CuIII
H
O
CuII
H
O
CuIII
N
N
N
N
N
N
O
CuIII
O
O
H
pKa=-2.4
pKa=13.1
N
CuIII
H
O
O
H
OH
H
O-
O
H
H2O, -H+
11
9
O-
O
O
10
9
8
Scheme 4. a) Calculated thermodynamic pathways for the first and second oxidation of the Cu-L complex. Struc-
tures in red indicate the proposed most probable pathway based on DFT calculations. Redox potential values
are reported vs. NHE. b) Proposed mechanism for water oxidation.
Computational Details. All quantum chemical calcu-
ACKNOWLEDGMENT
lations were performed using the density functional theo-
ry (DFT) functional B3LYP/6-311G+g(d) as implemented
in the Gaussian 09 software suite.60 Open-shell and
closed-shell electronic structure complexes were opti-
mized at the unrestricted and restricted level, respective-
ly. Solvation was considered in the calculations using the
conductor-like polarizable continuum model (CPCM).
We thank the NSF-DMR for funding support. C.W. gratefully
acknowledges the UNC Department of Chemistry for an Ern-
est L. Eliel fellowship and a Venable award. We thank Mr.
Christopher Poon and Ms. Stephanie A. Kramer for ICP-MS
experiments.
References
1. FUJISHIMA, A.; HONDA, K. Nature 1972, 238, 37.
2. Youngblood, W. J.; Lee, S. A.; Kobayashi, Y.; Hernandez-Pagan,
E.; Hoertz, P. G.; Moore, T. A.; Moore, A. L.; Gust, D.;
Mallouk, T. E. J. Am. Chem. Soc. 2009, 131, 926.
3. Ohno, T.; Bai, L.; Hisatomi, T.; Maeda, K.; Domen, K. J. Am.
Chem. Soc. 2012, 134, 8254.
4. Reece, S. Y.; Hamel, J. A.; Sung, K.; Jarvi, T. D.; Esswein, A. J.;
Pijpers, J. J. H.; Nocera, D. G. Science 2011, 334, 645.
5. Wang, X.; Xu, Q.; Li, M.; Shen, S.; Wang, X.; Wang, Y.; Feng, Z.;
Shi, J.; Han, H.; Li, C. Angewandte Chemie International
Edition 2012, 51, 13089.
6. Gao, Y.; Ding, X.; Liu, J.; Wang, L.; Lu, Z.; Li, L.; Sun, L. J. Am.
Chem. Soc. 2013, 135, 4219.
7. Swierk, J. R.; Mallouk, T. E. Chem. Soc. Rev. 2013, 42, 2357.
8. Eisenberg, R.; Gray, H. B. Inorg. Chem. 2008, 47, 1697.
9. Barber, J. Chem. Soc. Rev. 2009, 38, 185.
10. Hara, M.; Waraksa, C. C.; Lean, J. T.; Lewis, B. A.; Mallouk, T.
E. J Phys Chem A 2000, 104, 5275.
11. Wada, T.; Tsuge, K.; Tanaka, K. Inorg. Chem. 2001, 40, 329.
12. McDaniel, N. D.; Coughlin, F. J.; Tinker, L. L.; Bernhard, S. J.
Am. Chem. Soc. 2008, 130, 210.
13. Concepcion, J. J.; Jurss, J. W.; Brennaman, M. K.; Hoertz, P. G.;
Patrocinio, A. O. T.; Murakami Iha, N. Y.; Templeton, J. L.;
Meyer, T. J. Acc. Chem. Res. 2009, 42, 1954.
14. Hull, J. F.; Balcells, D.; Blakemore, J. D.; Incarvito, C. D.; Ei-
senstein, O.; Brudvig, G. W.; Crabtree, R. H. J. Am. Chem.
Soc. 2009, 131, 8730.
pKa’s and redox potentials were calculated from free
energy values obtained from frequency calculations. Ref-
erence potential of NHE was set at 4.28V for redox poten-
tial calculation, while standard free energy of proton was
set at -11.72 eV for pKa calculation as reported in the liter-
ature.61
ASSOCIATED CONTENT
Supporting Information. Crystallographic data (cif file),
detailed experimental procedure, synthesis of compound
H2L, calculation of TON in electrolysis, additional CVs and
electrolysis data of Cu-L solutions, ESI-MS spectrum of Cu-L,
ICP-MS results and UV-vis spectra of the film, table of redox
potentials for SWVs in Fig. 8, titration curves of L2-, Cu-Lb
and Zn-L solutions, UV-Vis spectra of Cu-L and Zn-L solu-
tions, and computational results of Cu-La and Cu-Lb systems.
This material is available free of charge via the Internet at
http://pubs.acs.org.
AUTHOR INFORMATION
Corresponding Author
*wenbinlin@uchicago.edu.
Author Contributions
15. Duan, L.; Bozoglian, F.; Mandal, S.; Stewart, B.; Privalov, T.;
Llobet, A.; Sun, L. Nat Chem 2012, 4, 418.
†These authors contributed equally.
ACS Paragon Plus Environment

Page 9 of 9
Journal of the American Chemical Society
16. Kärkäs, M. D.; Åkermark, T.; Johnston, E. V.; Karim, S. R.;
47. Wilson, A. D.; Newell, R. H.; McNevin, M. J.; Muckerman, J.
1
2
3
4
5
6
7
8
Laine, T. M.; Lee, B.; Åkermark, T.; Privalov, T.; Åkermark,
B. Angewandte Chemie International Edition 2012, 51, 11589.
17. Wang, C.; Wang, J.; Lin, W. J. Am. Chem. Soc. 2012, 134, 19895.
18. Kanan, M. W.; Nocera, D. G. Science 2008, 321, 1072.
19. Dismukes, G. C.; Brimblecombe, R.; Felton, G. A. N.; Pryadun,
R. S.; Sheats, J. E.; Spiccia, L.; Swiegers, G. F. Acc. Chem.
Res. 2009, 42, 1935.
20. Jiao, F.; Frei, H. Angewandte Chemie International Edition
2009, 48, 1841.
21. Yin, Q.; Tan, J. M.; Besson, C.; Geletii, Y. V.; Musaev, D. G.;
Kuznetsov, A. E.; Luo, Z.; Hardcastle, K. I.; Hill, C. L. Sci-
ence 2010, 328, 342.
T.; Rakowski DuBois, M.; DuBois, D. L. J. Am. Chem. Soc.
2006, 128, 358.
48. Dogutan, D. K.; McGuire, R.; Nocera, D. G. J. Am. Chem. Soc.
2011, 133, 9178.
49. Costentin, C.; Drouet, S.; Robert, M.; Savéant, J. J. Am. Chem.
Soc. 2012, 134, 11235.
50. Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Funda-
mentals and Applications; John Wiley & Sons, Inc.: 2001.
51. Concepcion, J.; Jurss, J.; Hoertz, P.; Meyer, T. Angewandte
Chemie International Edition 2009, 48, 9473.
52. Ryl, L. B.; Ronay, G. S.; Fowkes, F. M. J. Phys. Chem. 1958, 62,
798.
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
22. Ellis, W. C.; McDaniel, N. D.; Bernhard, S.; Collins, T. J. J. Am.
Chem. Soc. 2010, 132, 10990.
23. Wasylenko, D. J.; Ganesamoorthy, C.; Borau-Garcia, J.; Ber-
linguette, C. P. Chem. Commun. 2011, 47, 4249.
24. Fillol, J. L.; Codolà, Z.; Garcia-Bosch, I.; Gómez, L.; Pla, J. J.;
Costas, M. Nat Chem 2011, 3, 807.
25. Barnett, S. M.; Goldberg, K. I.; Mayer, J. M. Nat Chem 2012, 4,
498.
26. Zhang, F.; Yamakata, A.; Maeda, K.; Moriya, Y.; Takata, T.;
Kubota, J.; Teshima, K.; Oishi, S.; Domen, K. J. Am. Chem.
Soc. 2012, 134, 8348.
27. Trotochaud, L.; Ranney, J. K.; Williams, K. N.; Boettcher, S.
W. J. Am. Chem. Soc. 2012, 134, 17253.
28. Chen, Z.; Meyer, T. J. Angewandte Chemie International Edi-
tion 2013, 52, 700.
29. Zhang, M.; Chen, Z.; Kang, P.; Meyer, T. J. J. Am. Chem. Soc.
2013, 135, 2048.
30. KOK, B.; FORBUSH, B.; McGLOIN, M. Photochem. Photobiol.
1970, 11, 457.
31. Zouni, A.; Witt, H.; Kern, J.; Fromme, P.; Krauss, N.; Saenger,
W.; Orth, P. Nature 2001, 409, 739.
32. Ferreira, K. N.; Iverson, T. M.; Maghlaoui, K.; Barber, J.; Iwata,
S. Science 2004, 303, 1831.
33. Loll, B.; Kern, J.; Saenger, W.; Zouni, A.; Biesiadka, J. Nature
2005, 438, 1040.
34. McEvoy, J. P.; Brudvig, G. W. Chem. Rev. 2006, 106, 4455.
35. Umena, Y.; Kawakami, K.; Shen, J.; Kamiya, N. Nature 2011,
473, 55.
36. Klauss, A.; Haumann, M.; Dau, H. Proceedings of the National
Academy of Sciences 2012, 109, 16035.
37. Cox, N.; Pantazis, D. A.; Neese, F.; Lubitz, W. Acc. Chem. Res.
2013, 46, 1588.
38. Conifer, C. M.; Taylor, R. A.; Law, D. J.; Sunley, G. J.; White, A.
J. P.; Britovsek, G. J. P. Dalton Trans. 2011, 40, 1031.
39. Conifer, C. M.; Law, D. J.; Sunley, G. J.; Haynes, A.; Wells, J.
R.; White, A. J. P.; Britovsek, G. J. P. European Journal of
Inorganic Chemistry 2011, 2011, 3511.
40. Nieto, I.; Livings, M. S.; Sacci, J. B.; Reuther, L. E.; Zeller, M.;
Papish, E. T. Organometallics 2011, 30, 6339.
41. Kawahara, R.; Fujita, K.; Yamaguchi, R. J. Am. Chem. Soc. 2012,
134, 3643.
53. Gustafson, R. L.; Martell, A. E. J. Am. Chem. Soc. 1959, 81, 525.
54. Fabian, I. Inorg. Chem. 1989, 28, 3805.
55. Garribba, E.; Micera, G.; Sanna, D.; Strinna-Erre, L. Inorg.
Chim. Acta 2000, 299, 253.
56. Mirceski, V.; Komorsky-Lovric, S.; Lovric, M. Square-Wave
Voltammetry: Theory and Application; Springer-Verlag:
Berlin, Heidelberg, 2007.
57. Osteryoung, J.; O'Dea, J. J. In Bard, A. J., Ed.; Electroanalytical
Chemistry: A series of Advances; Marcel Dekker, Inc.: New
York, 1986; Vol. 14, pp 209-308.
58. Winkler, J. R.; Gray, H. B. Struct. Bond. 2012, 142, 17.
59. Kolthoff, I. M.; Tomsicek, W. J. J. Phys. Chem. 1934, 39, 945.
60. Frisch, M. J., et al. Gaussian 09, Revision C. 01; Gaussian, Inc.:
Wallingford CT, 2010.
61. Marenich, A. V.; Majumdar, A.; Lenz, M.; Cramer, C. J.;
Truhlar, D. G. Angewandte Chemie International Edition
2012, 51, 12810.
TOC Graphic
42. Wang, W.; Hull, J. F.; Muckerman, J. T.; Fujita, E.; Himeda, Y.
Energy Environ. Sci. 2012, 5, 7923.
43. DePasquale, J.; Nieto, I.; Reuther, L. E.; Herbst-Gervasoni, C.;
Paul, J. J.; Mochalin, V.; Zeller, M.; Thomas, C. M.; Addison,
A. W.; Papish, E. T. Inorg. Chem. 2013, 52, 9175.
44. Ward, M. D.; McCleverty, J. A. J. Chem. Soc. , Dalton Trans.
2002, 275.
45. Ringenberg, M. R.; Kokatam, S. L.; Heiden, Z. M.; Rauchfuss,
T. B. J. Am. Chem. Soc. 2008, 130, 788.
46. Nippe, M.; Khnayzer, R. S.; Panetier, J. A.; Zee, D. Z.; Olaiya,
B. S.; Head-Gordon, M.; Chang, C. J.; Castellano, F. N.;
Long, J. R. Chem. Sci. 2013, 4, 3934.
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