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An Investigation of Thin-Film Ni
−Fe Oxide Catalysts for the
Electrochemical Evolution of Oxygen Mary W. Louie and Alexis T. Bell *
Joint Center for Arti
ﬁcial Photosynthesis, Materials Science Division, Lawrence Berkeley National Laboratory, California 94720,
United States
Department of Chemical and Biomolecular Engineering, University of California , Berkeley, Berkeley, California 94720, United States
*S Supporting Information
ABSTRACT:
A detailed investigation has been carried out of
the structure and electrochemical activity of electrodeposited Ni
−Fe ﬁlms for the oxygen evolution reaction (OER) in
alkaline electrolytes. Ni
−Fe ﬁlms with a bulk and surface
composition of 40% Fe exhibit OER activities that are roughly 2 orders of magnitude higher than that of a freshly deposited Ni
ﬁlm and about 3 orders of magnitude higher than that of an Fe
ﬁlm. The freshly deposited Ni ﬁlm increases in activity by as much as 20-fold during exposure to the electrolyte (KOH); however, all
ﬁlms containing Fe are stable as deposited. The
oxidation of Ni(OH)2 to NiOOH in Ni ﬁlms occurs at potentials below the onset of the OER. Incorporation of Fe into the
ﬁlm increases the potential at which Ni(OH)
2/NiOOH
redox occurs and decreases the average oxidation state of Ni in NiOOH. The Tafel slope (40 mV dec−
1) and reaction order in
OH− (1) for the mixed Ni
−Fe ﬁlms (containing up to 95% Fe) are the same as those for aged Ni ﬁlms. In situ Raman spectra
acquired in 0.1 M KOH at OER potentials show two bands characteristic of NiOOH. The relative intensities of these bands vary with Fe content, indicating a change in the local environment of Ni
−O. Similar changes in the relative intensities of the bands
and an increase in OER activity are observed when pure Ni
ﬁlms are aged. These observations suggest that the OER is catalyzed
by Ni in Ni
−Fe ﬁlms and that the presence of Fe alters the redox properties of Ni, causing a positive shift in the potential at
which Ni(OH)2/NiOOH redox occurs, a decrease in the average oxidation state of the Ni sites, and a concurrent increase in the activity of Ni cations for the OER.
1. INTRODUCTION
The electrolysis of water to form hydrogen and oxygen o
ﬀers a
possible means for storing energy obtained from intermittent sources such as the sun.
1,2 However, the electrolysis of water
requires voltages in substantial excess of the thermodynamic potential for water-splitting (H2O → H2 + 1/2O2), 1.23 V, due primarily to the slow kinetics of the oxygen evolution reaction (OER).
1 At present, the most active catalysts for the OER are
RuO2 and IrO2, but even these operate with overpotentials in excess of 200 mV (at a current density of 10 mA cm−
2). 1−3
Moreover, the scarcity of Ru and Ir makes it impractical to use the metals on a large scale. For these reasons, there has been considerable interest in the discovery and development of OER catalysts based on earth-abundant metals.
A review of the literature suggests that Ni
−Fe catalysts oﬀer
a promising alternative to catalysts based on precious metals.
4
− 12 The lowest overpotential reported, ∼230 mV at
10 mA cm−
2 for electrodeposited Ni
−Fe ﬁlms, 4 is comparable
to overpotentials of 280 and 220 mV for chemically synthesized IrO2 and RuO2 ﬁlms.
3 It is also notable that Ni
−Fe catalysts are
signi
ﬁcantly more active for oxygen evolution than either Ni or
Fe alone, which exhibit overpotentials of 350
−450 and ∼500
mV, respectively, at 10 mA cm−
2. 13,14 While the bene
ﬁcial
e
ﬀects of Fe on the OER activity of Ni have been reported in a
number of studies,
7,10
−15 little is known about the structural or
chemical characteristics of Ni
−Fe catalysts, particularly under
conditions where the OER occurs. A small number of in situ spectroscopic studies have been carried out but yield contradictory results. Ni
−Fe catalysts (9−20 at % Fe)
characterized by either in situ X-ray absorption or Mossbauer spectroscopy, at potentials relevant for oxygen evolution, have been found to contain Ni(III), but the oxidation state of Fe is not clearly de
ﬁned; some authors conclude that it is Fe(III) 6,15
and others, that it is Fe(IV).
16
−18
We report here a structural and electrochemical investigation
of Ni
−Fe catalysts used for the OER in alkaline electrolyte. The
structure of the electrodeposited Ni
−Fe ﬁlm was characterized
in situ by Raman spectroscopy as a function of the applied potential, and the surface compositions were determined by ex situ X-ray photoelectron spectroscopy (XPS). The observed Raman characteristics of the Ni
−Fe series combined with the
Received:
May 28, 2013
Published:
July 16, 2013
Article
pubs.acs.org/JACS
© 2013 American Chemical Society
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Ni redox behavior and the kinetic parameters for the OER are used to correlate structural characteristics of the catalyst under reaction conditions with the catalyst composition and the catalytic activity.
2. EXPERIMENTAL SECTION
2.1. Electrocatalyst Preparation. Electrocatalyst
ﬁlms containing
Ni and/or Fe were prepared by electrodeposition onto gold electrodes. These gold substrates (99.95%, DOE Business Center for Precious Metals Sales and Recovery (BCPMSR)) were sheathed in Te
ﬂon,
leaving exposed surfaces which were either 4 or 5 mm in diameter. Prior to electrodeposition, the gold substrates were polished with alumina slurries, from 5
μm down to 50 nm in particle size, with
15 min of sonication in water after each polish. Electrodes which were used for Raman measurements were additionally roughened, in KCl (Sigma-Aldrich, P3911), using a previously reported electrochemical cycling procedure in order to produce substrates that are ideal for surface-enhanced Raman measurements.
19 (For details see section
S1.1, Supporting Information [SI].) Both smooth and roughened gold substrates were used for electrochemical measurements; the OER activity of these substrates was veri
ﬁed to be negligible compared to
that of the
ﬁlms measured in this work (Figure S5, SI).
Electrolyte solutions for electrodeposition were obtained by
dissolving nickel sulfate hexahydrate (
≥99.99% trace metals basis,
Sigma-Aldrich 467901) and/or iron sulfate heptahydrate (ACS Reagent
≥99.0%, Sigma-Aldrich 215422) in ultrapure water
(18.2 M
Ω, EMD Millipore). The metal concentrations were varied
between 0 and 0.01 M such that deposition baths with less than 50% Fe contained 0.01 M Ni and a varying Fe concentration, and baths with greater than 50% Fe contained 0.01 M Fe and a varying Ni concentration. Prior to dissolution of the appropriate quantities of sulfate salts, the water was sparged with nitrogen gas for 1 h to prevent oxidation of Fe(II) to Fe(III) and enable well-controlled depositions. Films were electrodeposited galvanostatically with a cathodic current density of 50
μA cm−2 applied for 1125 s and with nitrogen gas ﬂowing
in the headspace. The potential of the working electrode was monitored using a Ag/AgCl reference electrode with 4 M KCl
ﬁlling
solution (Pine RREF0021) or a leakless Ag/AgCl electrode with 3.4 M KCl
ﬁlling solution (eDaq ET072). The counterelectrode was a coiled
Pt wire (99.95%, DOE BCPMSR) which was routinely soaked in 5 M nitric acid to remove any deposited Ni or Fe.
Aged Ni
ﬁlms were prepared by immersion of as-deposited Ni ﬁlms,
without potential cycling, in 10 M KOH solutions for over 24 h. Immersion of the
ﬁlms was carried out in Teﬂon or polypropylene
cells/containers so as to minimize contaminants from glassware. Electrochemical and in situ Raman characterization of these
ﬁlms was
carried out after the Ni redox peaks and OER currents observed in the cyclic voltammograms were stable, typically 5
−10 cycles.
2.2. Physical Characterization. The compositions and quantities
of the electrodeposited materials were determined by elemental analysis. Electrodeposited
ﬁlms were dissolved by sonication in high-
purity 5 M HNO3 (Sigma-Aldrich 84385 or EMD Millipore NX0407). These solutions were diluted and adjusted such that the
ﬁnal solutions
contained 5% w/w HNO3 and 1000 ppb of yttrium internal standard (Sigma Aldrich 01357). Samples were characterized by inductively coupled plasma optical emission spectroscopy (ICP Optima 7000 DV, Perkin-Elmer). Calibration solutions contained 5% w/w HNO3, 1000 ppb Y, and both Ni and Fe, each with concentrations between 0 and 2000 ppb (Sigma-Aldrich 28944 and 43149 for Ni and Fe sources, respectively). On the basis of these measurements, we estimate the thicknesses of Ni
−Fe ﬁlms to be ∼25 and ∼70 nm when deposited on
roughened and polished gold substrates, respectively. (For details see section S2.2, SI.)
The surface compositions of the electrodeposited Ni
−Fe ﬁlms were
determined by X-ray photoelectron spectroscopy (XPS). Both as- deposited Ni
−Fe ﬁlms and ﬁlms which were polarized under OER
conditions (335 mV overpotential in 0.1 M KOH) for 1
−2 h were
examined. The XPS measurements were carried out with a Kratos Axis Ultra spectrometer using a monochromatic Al K
α source (15 mA,
15 kV). The instrument base pressure was 10−
9 Torr, and the charge
neutralizer system was used for all measurements. High-resolution spectra for Ni 2p, Fe 2p, O 1s, and C 1s were collected using a pass energy of 20 eV, a step energy of 50 meV, and dwell times of 200
−
400 ms. Angle-resolved XPS measurements were carried out by varying the electron takeo
ﬀ angle between 0° and 75° with respect to
the sample normal. The resulting spectra were analyzed using CasaXPS (Casa Software, Ltd.). A standard Shirley baseline with no o
ﬀset was used for background correction. In the case of Fe 2p spectra,
an additional correction was necessary due to the presence of a Ni LMM Auger peak. (For details see sections S4.2
−S4.3, SI.) The C 1s
spectrum for adventitious carbon (284.8 eV) was used for charge correction.
2.3. Electrochemical Characterization. Electrochemical charac-
terization of the Ni
−Fe catalysts was carried out in KOH electrolytes
(ACS reagent
≥85%, Sigma-Aldrich 221473) with concentrations of
0.1
−4.6 M in ultrapure water. (This KOH source is speciﬁed by the
supplier to have
≤0.001% Fe and ≤0.001% Ni.) A Hg/HgO reference
electrode (CH Instruments, ET072) with 1 M KOH
ﬁlling solution
was used throughout the experiments; the
ﬁlling solution was
exchanged before each experiment and measured against a second, unused Hg/HgO reference electrode stored in 1 M KOH. The counter electrode was a coiled Pt wire, cleaned routinely by nitric acid to remove any accumulated Ni or Fe deposits. All potentials reported in this work, unless otherwise noted, are measured against this Hg/HgO (1 M KOH) reference which has a potential of 0.098 V vs the normal hydrogen electrode (NHE). The equilibrium potential for oxygen evolution at any given pH is therefore (1.23
− 0.098 − 0.059 × pH) V.
Two electrochemical cells were used to measure the current
−
voltage characteristics of the Ni
−Fe catalysts. One was a home-built
Te
ﬂon cell20 which was designed for eﬃcient collection of Raman
signals. The second was a rotating disk electrode (RDE) apparatus (Pine Instruments) employed for additional electrochemical character- ization of Ni
−Fe ﬁlms, particularly for acquiring data for analysis of
OER kinetics in the absence of mass transfer e
ﬀects. Measurements
were carried out using either a Gamry Reference 600 or a BioLogic VSP potentiostat. IR compensation was applied at 85
−95% using the
ohmic resistance determined by AC impedance methods. Speci
ﬁcally,
impedance spectra were obtained at 0
± 10 mV (vs Hg/HgO)
between 1 MHz and 10 mHz, and the ohmic contribution was estimated from the Nyquist plots. For 1 and 0.1 M KOH electrolytes, the ohmic resistances were typically
∼5 and ∼40 Ω, respectively.
When necessary, compensation of the remaining 5
−15% of the ohmic
resistance was applied manually to the current
−potential data. This
procedure, while having negligible impact on the shape of the redox features, was necessary for acquiring accurate OER currents and therefore Tafel slopes, particularly for measurements in electrolytes with lower KOH concentrations and/or for electrocatalysts with high OER activity. It should be noted that the compensated resistance is noticeably higher during Raman spectra acquisition in 0.1 M KOH (100
−200 Ω), since immersion of the objective reduces the eﬀective
cross section for ionic conduction.
2.3.1. Surface Area. The surface areas of the ﬁlms were estimated
by measuring the electrochemical capacitance of the
ﬁlm−electrolyte
interface in the double-layer regime of the voltammograms. Using 0.1 M KOH, the electrode was potentiostatically cycled, typically between and 0.1 and 0.16 V vs Hg/HgO (1 M KOH), at scan rates between 1 and 10 mV s−
1 until the measured voltammograms had
stabilized. The positive and negative capacitance currents at the center of the potential window were averaged and plotted against the scan rate to extract the measured capacitance. The surface areas reported were obtained by using a speci
ﬁc capacitance of 60 μF cm−2 for
oxides.
21 (Representative plots are provided in section S2.3 of the SI.)
The measured currents were normalized by this area to obtain the speci
ﬁc current density. The challenges associated with determining
the surface area of catalysts have been reviewed by Trasatti and Petrii.
21 Extracted values of the speci
ﬁc current density, while reliable
for comparing across the Ni
−Fe system, should be used with care
when comparing to other catalysts reported in the literature. Additional discussion is provided in the SI, section S2.3.
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2.3.2. Cyclic Voltammograms. Cyclic voltammograms were
recorded between 0 and 0.7
−0.8 V vs Hg/HgO in 0.1 M KOH,
with the high-potential limit adjusted so as to minimize the amount of current driven through the high-activity catalysts. For the same reason, the potential window was adjusted by
∼60 mV per unit pH, since
higher OER currents were observed at higher KOH concentrations. Voltammograms were recorded until the redox peaks and the oxygen evolution currents show negligible change, typically 5
−10 cycles, with
the exception of as-deposited Ni
ﬁlms which slowly age/transform in
alkaline electrolytes.
22 While low scan rates are desirable for obtaining
current
−voltage curves for kinetic analysis, we chose a scan rate of
10 mV s−
1 to limit the overall time spent under oxygen evolution
conditions. At slower scan rates, increasingly more bubbles formed under OER conditions and collected on or near the electrode surface. The accumulation of bubbles on the electrode surface caused a drop in the measured current due to coverage of the active sites and/or to the additional ohmic resistance not accounted for by the uncompensated resistance values measured under non-OER conditions. For experi- ments carried out in the RDE apparatus, the rotation rate was varied between 0 and 2400 rotations per minute (RPM).
The activity for the OER was determined from cyclic voltammo-
grams by reading the speci
ﬁc current density at either a constant
overpotential of 300 mV or the overpotential at a constant speci
ﬁc
current density of 10 mA cm−
2. The Tafel slope was obtained from
data collected by rotating the electrode at 2400 rpm to reduce mass transport e
ﬀects and to increase the potential window for Tafel
analysis. In selecting the region for the Tafel
ﬁt, we avoided high
potentials at which oxygen bubble evolution causes mass transport limitations, and low potentials at which the redox transition for Ni(II)/ Ni(III) occurs.
The dependence of the OER current density on the concentrations
of OH− and O2 was determined. KOH concentrations of 0.1, 0.22, 0.46, 1.0, 2.2, and 4.6 M were used; these correspond to pHs of 13, 13.3, 13.7, 14, 14.3 and 14.7, respectively. Cyclic voltammograms were collected for each electrolyte concentration; measurements were repeated for the high and low concentrations of 0.1 and 4.6 M (pH 13 and 14.7) at the end of the concentration series to verify stability of the catalyst
ﬁlms with changing KOH concentration. To examine
e
ﬀect of O
2 concentration, voltammograms were measured in 0.1 M
KOH electrolyte sparged for one hour with either N2 or O2 prior to measurement. It should be noted that KOH concentrations below pH 12.5 resulted in gradual changes in the nickel redox features and OER current with time.
2.3.3. Calculation of the Turnover Frequency. The turnover
frequency (TOF) based on Ni sites can be computed in two ways. Taking the Ni in the entirety of the
ﬁlm to be catalytically active, a
lower limit, TOFmin, can be calculated using the number of Ni atoms determined by elemental analysis. On the other hand, the upper limit, TOFmax, can be estimated by taking the measured surface area of a ﬁlm and computing the number of Ni atoms at the surface using a value of
6.4
× 1014 Ni atoms per cm2 area 23 and the surface fraction of Ni as
determined by XPS. Although TOFmin and TOFmax diﬀer by 1−2 orders of magnitude, the dependence of TOF on composition is not a
ﬀected. Here, we report TOF
min, but both TOFmin and TOFmax are
presented in the SI section S7.
2.4. In Situ Raman Spectroscopy. Raman spectra were acquired
under controlled electrochemical potentials using a homemade Te
ﬂon
cell, which contained a working electrode (4 mm Au, Te
ﬂon-sheathed)
oriented at the bottom of the cell.
20 We employed a water-immersion
objective (70
× mag., N. A. = 1.23, LOMO) which was protected from
the corrosive KOH electrolytes by a 0.001-in. thick
ﬂuorinated
ethylene propylene
ﬁlm (McMaster-Carr) or 0.0005-in. thick Teﬂon
ﬁlm (American Duraﬁlm); a droplet of water was placed between the objective lens and the
ﬁlm to retain the high illumination/collection
e
ﬃciencies. Additional details are provided in ref 20.
Raman spectra were collected using a confocal Raman microscope
(LabRAM HR, Horiba Yvon Jobin) with a wavelength of 633 nm and a power of 1
−3 mW at the objective. The spot size of the laser beam is
estimated to be between 1 and 2
μm. Acquisition times for Ni−Fe
ﬁlms were typically 3 s for spectral range of 1100 cm−1 window. Using a 600 g/mm grating, the spectral resolution is
∼1 cm−1. Spectral shifts
were calibrated routinely against the value of 520.7 cm−
1 for a silicon
wafer. Raman spectra were collected at selected potentials as the potential of working electrode was scanned at a rate of 1 mV s−
1.
Raman spectra were not background-corrected due to the complexity of SERS backgrounds. When quantifying relative changes in the peak heights of Raman bands (carried out for cases when surface- enhancement contributions are minimized), a linear background correction was used. The sampling depth/volume of Ni
−Fe ﬁlms by
surface-enhanced Raman spectroscopy is discussed in the SI, section S5.5.
3. RESULTS AND DISCUSSION
3.1. Electrochemical Characteristics of Ni
−Fe Films.
Representative cyclic voltammograms for Ni and Fe
ﬁlms, i.e.,
the end-members of the catalyst series, are shown Figure 1. The cyclic voltammogram for an as-deposited Ni
ﬁlm in 0.1 M
KOH (Figure 1a) exhibits two primary characteristics, a redox couple at 0.47 V vs Hg/HgO (1 M KOH) and a positive (oxidation) current visible at overpotentials greater than 0.65 V. Both are well-known features for Ni electrodes in alkaline electrolytes. The redox peaks are attributed to the trans- formation between Ni(OH)2 and NiOOH,
14,22 which proceeds
as Ni(OH)2 + OH
− ↔ NiOOH + H
2O + e
− in alkaline
electrolytes. Oxidation currents at higher potentials are due to the evolution of oxygen, 4OH−
→ O
2 + 2H2O + 4e
−.
The cyclic voltammogram for Fe in 0.1 M KOH (Figure 1b)
shows an oxidation current attributed to the OER visible at
Figure 1.
Cyclic voltammograms for (a) Ni and (b) Fe
ﬁlms deposited on an Au substrate, measured in 0.1 M KOH at room temperature with
10 mV s−
1 scan rate and 2400 rpm. The equilibrium potential for OER is 0.365 V vs Hg/HgO (1 M KOH).
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potentials positive of 0.7 V. No redox transitions for Fe are observed in this potential window, consistent with previous reports in literature, since the oxidation of metallic Fe to Fe(II)/Fe(III) oxides/hydroxides occur at potentials between
−0.5 V and −1.2 V vs Hg/HgO (1 M). 13,24 The reduction peak observed at 0.22 V is that for the underlying gold substrate (Figure S1, SI); this feature is also present in the case of Ni
ﬁlms but not apparent in Figure 1a due to the signiﬁcantly higher currents observed compared to that of Fe.
The cyclic voltammograms for mixed Ni
−Fe ﬁlms are
characterized by two primary features: one for the Ni(OH)2/ NiOOH redox couple and the other for the OER. Two measures were de
ﬁned for the OER activity, one being the
speci
ﬁc current density at a constant overpotential of 300 mV
and the other the overpotential at a speci
ﬁc current density of
10 mA cm−
2. A plot of these two parameters as a function of
the surface composition of the Ni
−Fe ﬁlms, as determined by
X-ray photoelectron spectroscopy (section S.4, SI), reveals that the OER current density varies across 3 orders of magnitude with composition (Figure 2). The maximum speci
ﬁc current
density of 20 mA cm−
2 and the minimum overpotential of
280 mV were measured at a composition of 40% Fe. The
observation of this dramatic enhancement of the OER activity indicates that Ni
−Fe catalysts do not behave as simple mixtures
of the end-members and that interaction between Ni and Fe results in an improvement in the rate of oxygen evolution. It is notable that, across a composition range of 15
−50% Fe, the
speci
ﬁc current density only varies by 2-fold (between 10 and
20 mA cm−
2) and the overpotential by 20 mV. This behavior is
consistent with the discrepancies found in literature
4
− 10
regarding the composition of highest OER activity for the Ni
−Fe system; the optimum composition has been reported to
be as low as 10% Fe
6,7 and as high as 50% Fe. 4
The voltammograms for the Ni
−Fe series diﬀer noticeably in
their Ni(OH)2/NiOOH redox characteristics (Figure 3a). As more Fe is incorporated into the Ni
−Fe ﬁlm, the Ni(OH)
2/
NiOOH redox couple shifts to higher potentials, consistent with previous reports,
4,5,7,10 and the peak area decreases (Figure
3a). At an Fe content of 41% or greater, the oxidation wave is no longer visible due to its coincidence with the rapid rise in the OER current, which increases initially with increasing Fe content (Figure 2). However,
ﬁlms for which both oxidation
and reduction peaks are visible indicate that the two shift in tandem and have comparable integrated areas (Figure S4a, SI). Therefore, we used the reduction peak to quantify changes in the redox characteristics of Ni
−Fe ﬁlms as a function of
composition. A strong linear correlation is observed between the reduction peak potential and the Fe content (Figure 3b); the reduction peak shifts to positive potentials by as much as 150 mV relative to that for pure Ni
ﬁlms when the ﬁlm contains
70% Fe. This shift in the redox potential implies that the electrochemical oxidation of Ni(OH)2 to NiOOH is suppressed by the presence of Fe. The reduction peaks were integrated in order to determine the extent of Ni reduction/oxidation in the Ni
−Fe ﬁlms. As shown in Figure 3c for pure Ni ﬁlms deposited
atop roughened Au substrates, the charge passed during redox is greater than can be accounted for by assuming a one-electron redox reaction; we
ﬁnd that 1.2 electrons are transferred per Ni
atom in a redox cycle. (The number of Ni atoms deposited was measured by elemental analysis.) This result is consistent with reports that the Ni can exist as
γ-NiOOH in which Ni has an
average oxidation state as high as 3.7,
14,25
− 28 and therefore that
the
α-Ni(OH)
2/γ-NiOOH
transformation can involve the
transfer of up to 1.7 electrons per Ni atom. Figure 3c shows that the number of electrons transferred during Ni(OH)2/ NiOOH redox depends strongly on the Fe content, decreasing
Figure 2.
Oxygen evolution activity of electrodeposited Ni
−Fe
catalysts, taken at 300 mV overpotential and 10 mA cm−
2 speci
ﬁc
current density, as a function of composition in 0.1 M KOH. Filled and open markers correspond to measurements taken using rotating and stationary Au working electrode substrates for which electro- deposited
ﬁlms are ∼70 and ∼25 nm thick, respectively.
Figure 3.
E
ﬀect of Fe on the Ni(OH)
2/NiOOH redox couple in 0.1 M KOH. (a) Selected cyclic voltammograms for Ni−Fe ﬁlms on polished gold
substrates collected at 10 mV s−
1 and 2400 rpm, (b) reduction peak potential, and (c) the number of electrons transferred during redox as a function
of Fe content, shown for
ﬁlms on roughened gold substrates. Lines in (b) and (c) indicate linear ﬁts to the data.
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from
∼1.2 for pure Ni ﬁlms to ∼0.5 for ﬁlms containing 90%
Fe. Since all the Ni atoms in the
ﬁlms undergo redox, this
implies that addition of 90% Fe causes the average oxidation state of Ni to decrease from
∼3.2 to ∼2.5. We note that for
thicker
ﬁlms deposited onto polished substrates the number of
electrons per Ni involved in redox decreased with Fe content; however, in this case not all of the Ni in the
ﬁlm underwent
redox. (Additional discussion provided in section S2.4 of the SI.)
3.2. Electrochemical Characteristics of Aged Ni Films.
Ni electrodes are known to age in alkaline electrolytes.
23,25,29
Aging is attributed to the transformation of the
α-Ni(OH)
2/
γ‑NiOOH couple to the ordered and compact β‑Ni(OH)
2/
β‑NiOOH couple. Associated with this transformation is an anodic shift in the Ni(II)/Ni(III) redox potential, a decrease in the Ni oxidation state, and an increase in the OER activity.
14,22
− 25,29,30 This phenomenon was also observed in
our Ni
ﬁlms. The cyclic voltammograms for as-deposited and
aged Ni
ﬁlms are compared in Figure 4. Here, the redox peaks
for aged Ni
ﬁlms occur at potentials 40 mV higher than those
for as-deposited Ni
ﬁlms, and the OER activity is ∼4 mA cm−2
(at 300 mV overpotential), 20-fold higher than that for as- deposited
ﬁlms. As we will show later, we use the characteristics
of both types of Ni
ﬁlms to elucidate the behavior of the Ni−Fe
system.
It has been reported that potential cycling of nominally pure
Ni
ﬁlms for extended periods of time can result in the
incorporation of Fe present as impurities in the KOH electrolyte. In our measurements, the concentrations of Fe in freshly deposited and aged Ni
ﬁlms are comparable and close to
the detection limit for elemental analysis. Therefore, although the possibility of Fe incorporation cannot be eliminated, the dramatic change in the current
−voltage characteristics during
aging is likely not due to a change in Fe content. We attribute the negligible change in the Fe content to the absence of an applied potential during the aging process.
3.3. In Situ Raman Spectroscopy of Ni
−Fe Catalysts.
Figure 5 shows a series of Raman spectra for a Ni
ﬁlm
deposited onto a Au substrate and immersed in 0.1 M KOH, acquired as a function of the applied potential (vs Hg/HgO) during an oxidation sweep. The spectral features and their changes with potential are consistent with the transformation of Ni(OH)2 to NiOOH which starts at a potential of 0.47 V vs
Hg/HgO (as shown in Figure 1a). The band at 449 cm−
1 is
attributable to
α-Ni(OH)
2 and β-Ni(OH)2, both of which
exhibit Ni
−O bands in the range 445−465 cm−1 .22,26,31−33 The
band at 494 cm−
1 and the broad, underlying feature centered at
∼530 cm−1 are best assigned to defective or disordered Ni(OH)2 .
22,26,32,34
−37 Previous studies have reported that
disordered or doped Ni(OH)2 exhibits a Ni−O vibration that is shifted positively by as much as
∼65 mV of the band at 445−
465 cm−
1. (Additional discussion of this feature is given in
sections S5.3 and S5.4.1 of the SI.) The latter feature indicates that the Ni
ﬁlms produced in this study are not perfectly
crystalline but, instead, are disordered or defective. At high wavenumbers (Figure 5b), the bands at 3581
−3668 cm−1 can
be attributed to O
−H vibrations of Ni(OH)
2. β-Ni(OH)2
exhibits a single O
−H band at 3580 cm−1, whereas α-Ni(OH)
2
displays broader bands at 3625
−3670 cm−1. 22,26,31,38 A band at
3580 cm−
1 has also been reported for the so-called
α-Ni(OH)
2
phases, a consequence of the variability in the structure of the more disordered
α-Ni(OH)
2. The work of Bernard et al.
38
clearly shows that phases between completely disordered
α‑Ni(OH)
2 (with no O−H band at 3580 cm
−1) and perfectly
crystalline
β-Ni(OH)
2 (with a single O−H band at 3580 cm
−1)
exist. Therefore, on the basis of the features observed in the O
−H regime (Figure 5b), we conclude that our electro-
deposited
ﬁlms are not composed of perfectly crystalline β-
Ni(OH)2. Rather α-Ni(OH)2 or some form of disordered β- Ni(OH)2 is present.
The pair of bands at 474 and 554 cm−
1, observed at 0.5
−
0.7 V vs Hg/HgO (Figure 5a), are attributed to Ni
−O
vibrations in NiOOH. These two vibrations are known to have high Raman cross section due to resonance e
ﬀ ects,26 consistent
with the higher signal-to-noise ratio for these bands compared to the bands for the Ni(OH)2 phase. Both γ-NiOOH and β- NiOOH exhibit a pair of bands at these wavenumbers.
23,31,33
However, the relative intensities of the two bands have been found to di
ﬀer; speciﬁcally, in β-NiOOH, the ratio of the
intensity of the 474 cm−
1 band to that of the 554 cm−1 band is
lower than that for
γ-NiOOH. 22,23,37 On the basis of the
relative intensities of these two bands, the assignments for the Ni(OH)2 phase, and the redox potential measured for these
ﬁlms (Figure 4), we attribute these bands primarily to γ- NiOOH. The presence of
γ‑NiOOH is consistent with the Ni
oxidation state of 3.2 estimated for as-deposited Ni
ﬁlms
Figure 4.
Cyclic voltammograms for as-deposited and aged Ni
ﬁlms in
0.1 M KOH. The aged Ni
ﬁlms exhibit a higher redox potential for the
Ni(OH)2/NiOOH transition and a higher activity for the OER. Voltammograms were collected at 10 mV s−
1 and 2400 rpm.
Figure 5.
In situ Raman spectra collected at (a) low and (b) high
wavenumbers for Ni
ﬁlms atop a roughened Au substrate as a function
of potential vs Hg/HgO(1 M KOH) in 0.1 M KOH, for which the equilibrium potential for the OER is 0.365 V.
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(Figure 3c). The absence of O 
 −H vibrations for NiOOH 
 (Figure 5b, 0.5 
 −0.7 V) is consistent with previous reports. The 
 broad bands present at 3100 
 −3600 cm−1 are well-known 
 characteristics of water. 
 22,39 (Figure S21, SI) 
 In the case of Fe, Figure 6, an intense band at 571 cm− 
 1 is 
 observed at 0 V. This band shows very little change with the 
 applied potential, consistent with the electrochemical character- istics (Figure 1b). An apparent decrease in the intensity of these bands at potentials above 0.35 V is a consequence of the increased background due to the electrochemical oxidation of the underlying gold substrate. (Figure S17, SI) Additional bands are observed at 410 cm− 
 1 and in the range of 650 
 − 
 720 cm− 
 1; these features change slightly in their relative 
 intensities with potential but are too low in intensity for reliable analysis. Unlike the Raman spectra for Ni 
 ﬁlms, the Fe phase 
 displays no structural O 
 −H stretches. Identiﬁcation of the 
 bands observed for Fe 
 ﬁlms is rather diﬃcult owing to the many 
 possible iron oxide and oxyhydroxide phases, many of which are structurally similar. 
 40 The primary band at 571 cm−1 is 
 closest in shift to those for 
 α-FeOOH (560 cm−1) and Fe 
 3O4 
 (553 cm− 
 1), and the low-intensity band at 410 cm−1 matches 
 that of 
 α-Fe 
 2O3 (409 cm 
 −1). 41 The bands at 650−720 cm−1 
 may originate from Fe3O4, γ-Fe2O3 and γ-FeOOH which exhibit bands at 670, 650 
 −740, and 660 cm−1, respectively.41 
 (See additional discussion in section S5.2 of the SI.) 
 We captured the main features of the Ni 
 −Fe system by 
 plotting the Raman spectra of Ni 
 −Fe ﬁlms as functions of 
 composition at two di 
 ﬀerent potentials, 0.2 and 0.6 V vs Hg/ 
 HgO, the latter of which corresponds to an OER overpotential of 0.235 V (Figures 7 and 8, respectively). These two potentials are selected to show composition e 
 ﬀects at potentials below 
 and above the Ni(OH)2/NiOOH redox potential for each catalyst. 
 At 0.2 V, the Raman features for mixed Ni 
 −Fe ﬁlms can be 
 assigned to those of the end-members. The sharp bands of Ni(OH)2 (450 and 494 cm 
 −1) are present for ﬁlms with up to 
 32% Fe (Figure 7a) and disappear at higher Fe contents. On the other hand, the broad band of Ni(OH)2 (∼530 cm 
 −1) 
 remains at these intermediate compositions but has shifted to higher wavenumbers ( 
 ∼560 cm−1). Contributions from the Fe 
 phase are also visible at 578 and 650 
 −720 cm−1. The 
 disappearance of the sharp Ni(OH)2 bands and the persistence of the broad underlying band suggests that Ni 
 −Fe is 
 disordered. Spectra collected at higher wavenumbers (Figure 7b) support this conclusion, namely, that the sharp band at 3580 cm− 
 1 corresponding to the ordered 
 β-Ni(OH) 
 2 phase 
 disappears with increasing Fe content, leaving only a broad feature at 3500 
 −3660 cm−1 for ﬁlms with 19−51% Fe. 
 Under an oxygen evolution potential of 0.6 V vs Hg/HgO 
 (Figure 8a), the two characteristic bands for NiOOH (at 475 and 555 cm− 
 1 for the Ni 
 ﬁlm) are visible for ﬁlms with as much 
 as 90% Fe, owing to the high Raman cross section of NiOOH. 
 26 These two bands appear to merge with increasing 
 Fe content, and an additional band at 494 cm− 
 1 is apparent, as a 
 shoulder on the band at 475 cm− 
 1, for 
 ﬁlms containing 30−80% 
 Fe. The most dramatic e 
 ﬀect of composition is the change in 
 relative intensities of the pair of NiOOH bands; the 555 cm− 
 1 
 band increases in intensity with Fe incorporation. This change in the relative intensities is readily apparent for 
 ﬁlms with as 
 little as 19% Fe and clearly visible for 
 ﬁlms with higher Fe 
 contents despite the appearance of the 494 cm− 
 1 band. At Fe 
 contents of 90% or higher, features associated with Fe (and Au) phases, at 573 and 650 
 −720 cm−1, can be seen. It should be 
 Figure 6. 
 In situ Raman spectra collected for Fe 
 ﬁlms deposited on a 
 roughened Au substrate, shown as a function of potential vs Hg/HgO (1 M KOH) in 0.1 M KOH. The equilibrium potential for the OER is 0.365 V. 
 Figure 7. 
 In situ Raman spectra collected for Ni 
 −Fe catalysts, as a 
 function of composition, in 0.1 M KOH at a potential of 0.2 V vs Hg/ HgO (1 M KOH), shown for the (a) low and (b) high wavenumber regions. 
 Figure 8. 
 In situ Raman spectra collected in 0.1 M KOH at an OER 
 overpotential of 0.235 V, or 0.6 V vs Hg/HgO (1 M KOH), for (a) Ni 
 −Fe ﬁlms as a function of composition and (b) an aged Ni ﬁlm, 
 shown in comparison to the as-deposited Ni 
 ﬁlm from (a). 
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noted that the high Raman cross section of the NiOOH vibrations limits the amount of information we can extract from the Fe phases in catalysts with low-to-intermediate Fe contents. However, we can reliably analyze the NiOOH bands even for Ni
−Fe ﬁlms with small quantities of Ni. We veriﬁed that the
changes in NiOOH bands and the presence of the 494 cm−
1
band are not artifacts of the Fe and/or Au phases at
∼573 cm−1
which overlap with the NiOOH band at 555 cm−
1 (Figure S23,
SI). We also show here the e
ﬀect of aging on the two NiOOH
bands (Figure 8b); as is the case with Fe incorporation, aging of a Ni
ﬁlm also results in an increase in the intensity of the
555 cm−
1 band relative to the 475 cm−1 band.
With the exception of the feature at 494 cm−
1, no bands
attributable to new phases of Ni and/or Fe were observed. In particular, NiFe2O4, which contains 67% Fe(III) and 33% Ni(II), exhibits high-intensity Raman bands at approximately 490 and 700 cm−
1. 42,43 This pair of bands was not observed at
0.2 V, at which Ni exists as Ni(II). At 0.6 V, this pair of bands, particularly the band at 700 cm−
1, was not observed even for
high Fe content (for which contribution of the high-intensity NiOOH bands are minimal). The absence of these bands is particularly noteworthy because it has been suggested that NiFe2O4 is the cause of high OER activities in mixed Ni−Fe catalysts.
6 However, our in situ Raman observations do not
show evidence for NiFe2O4 in highly active Ni−Fe catalysts.
3.4. Analysis of OER Kinetics. The Tafel slope and the
reaction order in OH− for the Ni
−Fe catalyst series are shown
in Figure 9. The Tafel slopes for aged Ni
ﬁlms and mixed
Ni
−Fe ﬁlms are close to 40 mV dec−1, while those for as-
deposited Ni
ﬁlms and pure Fe ﬁlms are ∼55 mV dec−1. The
reaction order in OH− is unity for all
ﬁlms, with the exception
of as-deposited Ni
ﬁlms for which the reaction order could not
be determined due to the gradual aging of the
ﬁlms in alkaline
electrolytes. The OER current was found to be independent of O2 concentration in the electrolyte solution. (See additional discussion in section S3.2 of the SI.) The Tafel slope of 40 mV dec−
1 and reaction order of 1 in OH− for aged Ni and
mixed Ni
−Fe has also been found by Lyons and Brandon for
nickel electrodes;
14 these two kinetic parameters have been
rationalized using a reaction pathway composed of an initial reversible discharge of an OH− ion followed by a rate-limiting electron transfer step which results in the formation of a physisorbed hydrogen peroxide intermediate. The observation
that the kinetic parameters for aged Ni and for mixed Ni
−Fe
ﬁlms are the same implies a common rate-limiting step for these catalysts. That is, the enhancement in the OER activity due to Fe incorporation into the
ﬁlm results from an increase in
the rate of the rate-limiting step rather than from a change in the reaction pathway. That the kinetic parameters for mixed Ni
−Fe are the same as those for aged Ni but not as-deposited
Ni implies that aged Ni
ﬁlms and mixed Ni−Fe ﬁlms share a
common reaction pathway.
3.5. Implications for OER Activity. Raman spectroscopy
indicates that, under OER potentials, Ni in Ni
−Fe catalysts is
present as NiOOH and that the structure of this phase is modi
ﬁed dramatically as a consequence of Fe incorporation.
The increase in the intensity of the 560 cm−
1 band relative to
the 480 cm−
1 band is clearly noticeable, even for Fe contents as
low as 20%. This change in the intensity ratio is reminiscent of that observed during the transformation of
γ-NiOOH to
β‑NiOOH, the latter of which has been widely agreed to be the preferred phase for catalyzing the OER,
14,23,44,45 also con
ﬁrmed
in this study (Figure 4). Similarities in the e
ﬀects of aging
freshly deposited Ni
ﬁlms and the incorporation Fe into Ni
ﬁlms on the physical characteristics of Ni suggests that Ni atoms present at the catalyst surface are the active centers involved in the evolution of O2. To test this hypothesis, we compute the turnover frequency based on the number of Ni atoms in the catalyst
ﬁlms, TOF
min. We note that the TOFs
reported in this work are lower limits, since the entirety of the
ﬁlm may not be accessible for oxygen evolution. (TOF
max,
based on the number of Ni atoms at the surfaces of the
ﬁlms, is
provided in section S7 of the SI.) Although the absolute value of the TOF depends on assumptions of the OER-active content of the catalyst
ﬁlms, the composition dependence of the TOF
does not. In Figure 10a, TOFmin is plotted versus the ratio of the band intensities for the NiOOH vibrations, I475/I555. This plot suggests that a NiOOH-type phase with a lower ratio of Raman peak heights yields higher OER activities. (Scatter in the plot at low I475/I555 values, or high Fe content, is attributed to decreasing NiOOH band intensities and the presence of the 494 cm−
1 band.) The data point for the aged Ni
ﬁlm is included
in this correlation, suggesting that changes in the NiOOH bands, regardless of their source, are linked to the increased OER activity. This relationship is further corroborated by the Ni(OH)2/NiOOH redox behavior. As seen in b and c of Figure 10, respectively, the TOF correlates with the reduction peak potential and the average oxidation state of Ni (from b and c of Figure 3, respectively). The TOF increases linearly with the potential for the reduction peak, for potentials up to 0.54 V vs Hg/HgO, and despite scatter in the data, the TOF decreases with an increase in the average oxidation state of Ni. Here again, we observe that the data point for the aged Ni
ﬁlm agrees
with the correlations. These observations suggest that suppression of the oxidation of Ni(OH)2 to NiOOH, regardless of the cause, results in higher OER activities. Stated another way, the OER activity of Ni cations is higher the lower the average oxidation state of Ni. This trend agrees with the consensus in literature that
β-NiOOH, for which Ni exists as
Ni
3+, exhibits a much higher OER activity than
γ-NiOOH for
which Ni exists as Ni
3.7+. 14,25−28 The results reported here
suggest that further reduction in the average oxidation state of Ni by the addition of Fe results in higher OER activities which surpass that of aged Ni.
It should be noted that the Raman spectra for Ni
−Fe ﬁlms
also show evidence of lower crystallinity and a smaller
Figure 9.
Extracted Tafel slopes and reaction orders in the
concentration of OH− for the Ni
−Fe ﬁlms, taken from measurements
in 0.1
−4.6 M KOH. Dashed lines for a Tafel slope of 40 mV dec−1 and
a reaction order of unity are shown for reference.
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contribution from
β-Ni(OH)
2 with increasing Fe content at 0.2
V (Figure 7a). Furthermore, the unidenti
ﬁed band at 494 cm−1
observed at 0.6 V for Ni
−Fe ﬁlms with 30−80% Fe (Figure 8a)
matches very well to that observed for defective Ni(OH)2 (Figure 7a).
23,25,31,33
− 36 These observations indicate that the
Ni
−O vibrations which result from structural defects in
Ni(OH)2 are retained when Ni(OH)2 is oxidized to NiOOH for
ﬁlms which contain a signiﬁcant amount of Fe. (See
additional discussion in section S5.4.1 of the SI.) The observation that these highly active Ni
−Fe ﬁlms are disordered
is consistent with previous reports of amorphous metal oxides for the OER.
12,46 Thus, although the Ni
−Fe system displays
trends in NiOOH-type vibrations similar to that for the
γ−β
transformation, the structure of the Ni phase in Ni
−Fe catalyst
may not be exactly that of
β-NiOOH. Previous in situ X-ray
absorption spectroscopy studies
15,16 have reported the parent
structure of oxidized Ni
ﬁlms with low Fe content (
to be similar to
γ-NiOOH. Therefore, it is likely that Raman
spectroscopy probes local characteristics of Ni
−O not captured
by the extended X-ray absorption
ﬁne structure (EXAFS)
which, while able to provide the local structure and coordination around Ni, does not readily provide structural information beyond the Ni
−O planes of layered nickel
(oxy)hydroxides.
15,16,47 That is, while the Raman bands at
∼475 and ∼555 cm−1 are generated by the same basic structural unit, their relative intensities re
ﬂect the local structure
around Ni
−O which is inﬂuenced by factors including the
interlayer spacing between Ni
−O sheets, the presence of
protons or other cations (such as potassium) between the sheets, structural disorder within sheets, and the metal oxidation state.
22 Similarly, the energetics of the Ni(OH)
2/
NiOOH transformation may also be in
ﬂuenced by such factors.
The similarity in the vibrational and redox characteristics of
aged Ni
ﬁlms and the NiOOH phase in mixed Ni−Fe ﬁlms
suggest that Ni
−O vibration may be an indicator of the OER
activity; this is consistent with the OER pathway which must involve adsorbed oxygen-based intermediates at the presumed Ni active site. The suggestion that Ni plays an active role in facilitating the OER in these catalysts is consistent with the kinetic parameters for these
ﬁlms. Speciﬁcally, the Tafel slope
and reaction order in OH− for mixed Ni
−Fe ﬁlms agree with
those for aged Ni
ﬁlms but not as-deposited Ni ﬁlms or Fe
ﬁlms, indicating that Ni−Fe catalysts share a common reaction
pathway with aged Ni. A further implication is that, while aging of Ni
ﬁlms can achieve a maximum of ∼4 mA cm−2, the
addition of Fe to Ni can modify the Ni
−O character so as to
surpass the OER currents of aged Ni
ﬁlms by 5-fold
(20 mA cm−
2 for Ni
−Fe containing ∼40% Fe). On the basis
of Ni as the active site for OER, we interpret the maximum in the speci
ﬁc current density for the OER (Figure 2) to be a
competition between the increasing activity of the Ni sites, as in
ﬂuenced by its Fe neighbors, and the decreasing quantity of
Ni sites as more Fe is incorporated.
4. CONCLUSION
We have examined the electrochemical and structural character- istics of Ni
−Fe OER catalysts across the full composition range.
Electrodeposited Ni
−Fe ﬁlms containing 40% Fe exhibit OER
current densities which are 2 orders of magnitude higher than freshly deposited Ni
ﬁlms and 3 orders of magnitude higher
than Fe
ﬁlms. Electrochemical measurements show that the
Ni(OH)2/NiOOH redox couple shifts monotonically to higher (anodic) potentials with increasing Fe content, indicating that Fe suppresses the electrochemical oxidation of Ni(OH)2 to NiOOH. Correspondingly, the number of electrons transferred during redox indicates that the average oxidation state of the nominally Ni(III) sites decreases with Fe incorporation. Characterization of the Ni
−Fe ﬁlms by in situ Raman
spectroscopy reveals that the catalysts, particularly those exhibiting high OER activities, display some degree of disorder. Raman spectra also reveal that, under OER potentials, Ni in the Ni
−Fe mixtures contain a structural unit similar to NiOOH.
Furthermore, the local environment of Ni
−O, as described by
the relative intensities of the two NiOOH bands at 475 and 555 cm−
1, is modi
ﬁed dramatically by the presence of Fe. Both
properties of the Ni
−Fe ﬁlms, that is, the redox behavior and
the Raman characteristics of the NiOOH-type phase, correlate with the OER activity. This correlation is consistent with the e
ﬀect of aging on Ni ﬁlms; that is, aging of Ni ﬁlms improves
the OER activity while modifying the redox and Raman characteristics of Ni in a manner similar to that observed for Fe incorporation. Similarities between aged Ni and Ni
−Fe
mixtures are also supported by the observation that the Tafel slope and the reaction order in OH− for aged Ni
ﬁlms match
those for the mixed Ni
−Fe ﬁlms, suggesting a common reaction
Figure 10.
Turnover frequency (TOFmin) based on Ni sites for the OER plotted against the corresponding (a) intensity ratio of the 475 cm
−1 band
to the 555 cm−
1 band obtained from in situ Raman spectra, (b) the NiOOH reduction peak potential, and (c) the average oxidation state of Ni.
Filled markers correspond to the as-deposited Ni
−Fe ﬁlms, and the unﬁlled marker corresponds to an aged Ni ﬁlm. Lines are ﬁts to the data. The
corresponding plots for the speci
ﬁc current densities are shown in Figure S25 of the SI.
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pathway. Taken together, our observations demonstrate that the local environment of Ni in Ni 
 −Fe ﬁlms strongly aﬀects the 
 average oxidation state and OER activity of Ni in alkaline electrolytes. 
 ■ ASSOCIATEDCONTENT 
 *S Supporting Information Additional experimental procedures, supplementary electro- chemical data, Raman spectroscopy, X-ray photoelectron spectra. This material is available free of charge via the Internet at http://pubs.acs.org. 
 ■ AUTHORINFORMATION 
 Corresponding Author bell@cchem.berkeley.edu Notes The authors declare no competing 
 ﬁnancial interest. 
 ■ ACKNOWLEDGMENTS 
 This material is based upon work performed by the Joint Center for Arti 
 ﬁcial Photosynthesis, a DOE Energy Innovation 
 Hub, supported through the O 
 ﬃce of Science of the U.S. 
 Department of Energy under Award Number DE-SC0004993. M.W.L. is supported by the University of California President 
 ’s 
 Postdoctoral Fellowship Program. We gratefully acknowledge Ian D. Sharp (Joint Center for Arti 
 ﬁcial Photosynthesis) for 
 valuable discussion and assistance with XPS analysis of the electrocatalysts used in this study, and Elena Kreimer (University of California, Berkeley, College of Chemistry) for assistance with elemental analysis. We also thank Eric Granlund (University of California, Berkeley, College of Chemistry) as well as James Wu and Doug Jamieson (Lawrence Berkeley National Laboratory, Materials Science Division) for the fabrication of our electrodes and electrochemical cells. 
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 (28) Cordoba-Torresi, S. I.; Gabrielli, C.; Hugotlegoff, A.; Torresi, R. 
 J. Electrochem. Soc. 1991, 138, 1548. 
 (29) Wehrens-Dijksma, M.; Notten, P. H. L. Electrochim. Acta 2006, 
 51, 3609. 
 (30) Godwin, I. J.; Lyons, M. E. G. Electrochem. Commun. 2013, 32, 
 39. 
 (31) Johnston, C.; Graves, P. R. Appl. Spectrosc. 1990, 44, 105. (32) Bernard, M. C.; Cortes, R.; Keddam, M.; Takenouti, H.; 
 Bernard, P.; Senyarich, S. J. Power Sources 1996, 63, 247. 
 (33) Cornilsen, B. C.; Shan, X. Y.; Loyselle, P. L. J. Power Sources 
 1990 
 , 29, 453. 
 (34) de Torresi, S. I. C.; Provazi, K.; Malta, M.; Torresi, R. M. J. 
 Electrochem. Soc. 2001, 148, A1179. 
 (35) Vidotti, M.; Salvador, R. P.; de Torresi, S. I. C. Ultrason. 
 Sonochem. 2009, 16, 35. 
 (36) Vidotti, M.; Salvador, R. P.; Ponzio, E. A.; Cordoba de Torresi, 
 S. I. J. Nanosci. Nanotechnol. 2007, 7, 3221. 
 (37) Kostecki, R.; McLarnon, F. In Proceedings of the Symposium on 
 Electrode Materials and Processes for Energy Conversion and Storage IV; McBreen, J., Mukerjee, S., Srinivasan, S., Eds.; Electrochemical Society, Inc.: Pennington, NJ, 1997; p viii. 
 (38) Bernard, M. C.; Bernard, P.; Keddam, M.; Senyarich, S.; 
 Takenouti, H. Electrochim. Acta 1996, 41, 91. 
 (39) Hibben, J. H. J. Chem. Phys. 1937, 5, 166. (40) Cornell, R. M.; Schwertmann, U. The Iron Oxides: Structure, 
 Properties, Reactions, Occurrences, and Uses, 2nd completely rev. and extended ed.; Wiley-VCH: Weinheim, 2003. 
 (41) Thierry, D.; Persson, D.; Leygraf, C.; Boucherit, N.; 
 Hugotlegoff, A. Corros. Sci. 1991, 32, 273. 
 (42) Graves, P. R.; Johnston, C.; Campaniello, J. J. Mater. Res. Bull. 
 1988 
 , 23, 1651. 
 (43) Jacob, J.; Khadar, M. A. J. Appl. Phys. 2010, 107, 114310. (44) Lu, P. W. T.; Srinivasan, S. J. Electrochem. Soc. 1978, 125, 1416. (45) Cappadonia, M.; Divisek, J.; Vonderheyden, T.; Stimming, U. 
 Electrochim. Acta 1994, 39, 1559. 
 (46) Merrill, M. D.; Dougherty, R. C. J. Phys. Chem. C 2008, 112, 
 3655. 
 (47) Pandya, K. I.; Hoffman, R. W.; McBreen, J.; O 
 ’Grady, W. E. J. 
 Electrochem. Soc. 1990, 137, 383. 
 Journal of the American Chemical Society 
 Article 
 dx.doi.org/10.1021/ja405351s 
 | J. Am. Chem. Soc. 2013, 135, 12329−12337 
 12337

pageNum = 10

pageNum = 11

pageNum = 12

pageNum = 13

pageNum = 14

pageNum = 15

(Figure 3c). The absence of O 
 −H vibrations for NiOOH 
 (Figure 5b, 0.5 
 −0.7 V) is consistent with previous reports. The 
 broad bands present at 3100 
 −3600 cm−1 are well-known 
 characteristics of water. 
 22,39 (Figure S21, SI) 
 In the case of Fe, Figure 6, an intense band at 571 cm− 
 1 is 
 observed at 0 V. This band shows very little change with the 
 applied potential, consistent with the electrochemical character- istics (Figure 1b). An apparent decrease in the intensity of these bands at potentials above 0.35 V is a consequence of the increased background due to the electrochemical oxidation of the underlying gold substrate. (Figure S17, SI) Additional bands are observed at 410 cm− 
 1 and in the range of 650 
 − 
 720 cm− 
 1; these features change slightly in their relative 
 intensities with potential but are too low in intensity for reliable analysis. Unlike the Raman spectra for Ni 
 ﬁlms, the Fe phase 
 displays no structural O 
 −H stretches. Identiﬁcation of the 
 bands observed for Fe 
 ﬁlms is rather diﬃcult owing to the many 
 possible iron oxide and oxyhydroxide phases, many of which are structurally similar. 
 40 The primary band at 571 cm−1 is 
 closest in shift to those for 
 α-FeOOH (560 cm−1) and Fe 
 3O4 
 (553 cm− 
 1), and the low-intensity band at 410 cm−1 matches 
 that of 
 α-Fe 
 2O3 (409 cm 
 −1). 41 The bands at 650−720 cm−1 
 may originate from Fe3O4, γ-Fe2O3 and γ-FeOOH which exhibit bands at 670, 650 
 −740, and 660 cm−1, respectively.41 
 (See additional discussion in section S5.2 of the SI.) 
 We captured the main features of the Ni 
 −Fe system by 
 plotting the Raman spectra of Ni 
 −Fe ﬁlms as functions of 
 composition at two di 
 ﬀerent potentials, 0.2 and 0.6 V vs Hg/ 
 HgO, the latter of which corresponds to an OER overpotential of 0.235 V (Figures 7 and 8, respectively). These two potentials are selected to show composition e 
 ﬀects at potentials below 
 and above the Ni(OH)2/NiOOH redox potential for each catalyst. 
 At 0.2 V, the Raman features for mixed Ni 
 −Fe ﬁlms can be 
 assigned to those of the end-members. The sharp bands of Ni(OH)2 (450 and 494 cm 
 −1) are present for ﬁlms with up to 
 32% Fe (Figure 7a) and disappear at higher Fe contents. On the other hand, the broad band of Ni(OH)2 (∼530 cm 
 −1) 
 remains at these intermediate compositions but has shifted to higher wavenumbers ( 
 ∼560 cm−1). Contributions from the Fe 
 phase are also visible at 578 and 650 
 −720 cm−1. The 
 disappearance of the sharp Ni(OH)2 bands and the persistence of the broad underlying band suggests that Ni 
 −Fe is 
 disordered. Spectra collected at higher wavenumbers (Figure 7b) support this conclusion, namely, that the sharp band at 3580 cm− 
 1 corresponding to the ordered 
 β-Ni(OH) 
 2 phase 
 disappears with increasing Fe content, leaving only a broad feature at 3500 
 −3660 cm−1 for ﬁlms with 19−51% Fe. 
 Under an oxygen evolution potential of 0.6 V vs Hg/HgO 
 (Figure 8a), the two characteristic bands for NiOOH (at 475 and 555 cm− 
 1 for the Ni 
 ﬁlm) are visible for ﬁlms with as much 
 as 90% Fe, owing to the high Raman cross section of NiOOH. 
 26 These two bands appear to merge with increasing 
 Fe content, and an additional band at 494 cm− 
 1 is apparent, as a 
 shoulder on the band at 475 cm− 
 1, for 
 ﬁlms containing 30−80% 
 Fe. The most dramatic e 
 ﬀect of composition is the change in 
 relative intensities of the pair of NiOOH bands; the 555 cm− 
 1 
 band increases in intensity with Fe incorporation. This change in the relative intensities is readily apparent for 
 ﬁlms with as 
 little as 19% Fe and clearly visible for 
 ﬁlms with higher Fe 
 contents despite the appearance of the 494 cm− 
 1 band. At Fe 
 contents of 90% or higher, features associated with Fe (and Au) phases, at 573 and 650 
 −720 cm−1, can be seen. It should be 
 Figure 6. 
 In situ Raman spectra collected for Fe 
 ﬁlms deposited on a 
 roughened Au substrate, shown as a function of potential vs Hg/HgO (1 M KOH) in 0.1 M KOH. The equilibrium potential for the OER is 0.365 V. 
 Figure 7. 
 In situ Raman spectra collected for Ni 
 −Fe catalysts, as a 
 function of composition, in 0.1 M KOH at a potential of 0.2 V vs Hg/ HgO (1 M KOH), shown for the (a) low and (b) high wavenumber regions. 
 Figure 8. 
 In situ Raman spectra collected in 0.1 M KOH at an OER 
 overpotential of 0.235 V, or 0.6 V vs Hg/HgO (1 M KOH), for (a) Ni 
 −Fe ﬁlms as a function of composition and (b) an aged Ni ﬁlm, 
 shown in comparison to the as-deposited Ni 
 ﬁlm from (a). 
 Journal of the American Chemical Society 
 Article 
 dx.doi.org/10.1021/ja405351s 
 | J. Am. Chem. Soc. 2013, 135, 12329−12337 
 12334

pageNum = 16

pageNum = 17

pageNum = 18

pathway. Taken together, our observations demonstrate that the local environment of Ni in Ni 
 −Fe ﬁlms strongly aﬀects the 
 average oxidation state and OER activity of Ni in alkaline electrolytes. 
 ■ ASSOCIATEDCONTENT 
 *S Supporting Information Additional experimental procedures, supplementary electro- chemical data, Raman spectroscopy, X-ray photoelectron spectra. This material is available free of charge via the Internet at http://pubs.acs.org. 
 ■ AUTHORINFORMATION 
 Corresponding Author bell@cchem.berkeley.edu Notes The authors declare no competing 
 ﬁnancial interest. 
 ■ ACKNOWLEDGMENTS 
 This material is based upon work performed by the Joint Center for Arti 
 ﬁcial Photosynthesis, a DOE Energy Innovation 
 Hub, supported through the O 
 ﬃce of Science of the U.S. 
 Department of Energy under Award Number DE-SC0004993. M.W.L. is supported by the University of California President 
 ’s 
 Postdoctoral Fellowship Program. We gratefully acknowledge Ian D. Sharp (Joint Center for Arti 
 ﬁcial Photosynthesis) for 
 valuable discussion and assistance with XPS analysis of the electrocatalysts used in this study, and Elena Kreimer (University of California, Berkeley, College of Chemistry) for assistance with elemental analysis. We also thank Eric Granlund (University of California, Berkeley, College of Chemistry) as well as James Wu and Doug Jamieson (Lawrence Berkeley National Laboratory, Materials Science Division) for the fabrication of our electrodes and electrochemical cells. 
 ■ REFERENCES 
 (1) Walter, M. G.; Warren, E. L.; McKone, J. R.; Boettcher, S. W.; Mi, 
 Q. X.; Santori, E. A.; Lewis, N. S. Chem. Rev. 2010, 110, 6446. 
 (2) Cook, T. R.; Dogutan, D. K.; Reece, S. Y.; Surendranath, Y.; 
 Teets, T. S.; Nocera, D. G. Chem. Rev. 2010, 110, 6474. 
 (3) Matsumoto, Y.; Sato, E. Mater. Chem. Phys. 1986, 14, 397. (4) Corrigan, D. A. J. Electrochem. Soc. 1987, 134, 377. (5) Hu, C. C.; Wu, Y. R. Mater. Chem. Phys. 2003, 82, 588. (6) Landon, J.; Demeter, E.; Inoglu, N.; Keturakis, C.; Wachs, I. E.; 
 Vasic, R.; Frenkel, A. I.; Kitchin, J. R. ACS Catal. 2012, 2, 1793. 
 (7) Li, X. H.; Walsh, F. C.; Pletcher, D. Phys. Chem. Chem. Phys. 
 2011 
 , 13, 1162. 
 (8) Miller, E. L.; Rocheleau, R. E. J. Electrochem. Soc. 1997, 144, 3072. (9) Potvin, E.; Brossard, L. Mater. Chem. Phys. 1992, 31, 311. (10) Singh, R. N.; Pandey, J. P.; Anitha, K. L. Int. J. Hydrogen Energy 
 1993 
 , 18, 467. 
 (11) Trotochaud, L.; Ranney, J. K.; Williams, K. N.; Boettcher, S. W. 
 J. Am. Chem. Soc. 2012, 134, 17253. 
 (12) 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. 
 (13) Lyons, M. E. G.; Brandon, M. P. Int. J. Electrochem. Soc. 2008, 3, 
 1463. 
 (14) Lyons, M. E. G.; Brandon, M. P. Int. J. Electrochem. Soc. 2008, 3, 
 1386. 
 (15) Kim, S. H.; Tryk, D. A.; Antonio, M. R.; Carr, R.; Scherson, D. J. 
 Phys. Chem. 1994, 98, 10269. 
 (16) Balasubramanian, M.; Melendres, C. A.; Mini, S. J. Phys. Chem. B 
 2000 
 , 104, 4300. 
 (17) Corrigan, D. A.; Conell, R. S.; Fierro, C. A.; Scherson, D. A. J. 
 Phys. Chem. 1987, 91, 5009. 
 (18) Guerlou-Demourgues, L.; Fournes, L.; Delmas, C. J. Electrochem. 
 Soc. 1996, 143, 3083. 
 (19) Gao, P.; Gosztola, D.; Leung, L. W. H.; Weaver, M. J. J. 
 Electroanal. Chem. 1987, 233, 211. 
 (20) Yeo, B. S.; Klaus, S. L.; Ross, P. N.; Mathies, R. A.; Bell, A. T. 
 ChemPhysChem 2010, 11, 1854. 
 (21) Trasatti, S.; Petrii, O. A. Pure Appl. Chem. 1991, 63, 711. (22) Kostecki, R.; McLarnon, F. J. Electrochem. Soc. 1997, 144, 485. (23) Yeo, B. S.; Bell, A. T. J. Phys. Chem. C 2012, 116, 8394. (24) Doyle, R. L.; Lyons, M. E. G. J. Electrochem. Soc. 2013, 160, 
 H142. 
 (25) Oliva, P.; Leonardi, J.; Laurent, J. F.; Delmas, C.; Braconnier, J. 
 J.; Figlarz, M.; Fievet, F.; Deguibert, A. J. Power Sources 1982, 8, 229. 
 (26) Desilvestro, J.; Corrigan, D. A.; Weaver, M. J. J. Electrochem. Soc. 
 1988 
 , 135, 885. 
 (27) Fu, X. Z.; Zhu, Y. J.; Xu, Q. C.; Li, J.; Pan, J. H.; Xu, J. Q.; Lin, J. 
 D.; Liao, D. W. Solid State Ionics 2007, 178, 987. 
 (28) Cordoba-Torresi, S. I.; Gabrielli, C.; Hugotlegoff, A.; Torresi, R. 
 J. Electrochem. Soc. 1991, 138, 1548. 
 (29) Wehrens-Dijksma, M.; Notten, P. H. L. Electrochim. Acta 2006, 
 51, 3609. 
 (30) Godwin, I. J.; Lyons, M. E. G. Electrochem. Commun. 2013, 32, 
 39. 
 (31) Johnston, C.; Graves, P. R. Appl. Spectrosc. 1990, 44, 105. (32) Bernard, M. C.; Cortes, R.; Keddam, M.; Takenouti, H.; 
 Bernard, P.; Senyarich, S. J. Power Sources 1996, 63, 247. 
 (33) Cornilsen, B. C.; Shan, X. Y.; Loyselle, P. L. J. Power Sources 
 1990 
 , 29, 453. 
 (34) de Torresi, S. I. C.; Provazi, K.; Malta, M.; Torresi, R. M. J. 
 Electrochem. Soc. 2001, 148, A1179. 
 (35) Vidotti, M.; Salvador, R. P.; de Torresi, S. I. C. Ultrason. 
 Sonochem. 2009, 16, 35. 
 (36) Vidotti, M.; Salvador, R. P.; Ponzio, E. A.; Cordoba de Torresi, 
 S. I. J. Nanosci. Nanotechnol. 2007, 7, 3221. 
 (37) Kostecki, R.; McLarnon, F. In Proceedings of the Symposium on 
 Electrode Materials and Processes for Energy Conversion and Storage IV; McBreen, J., Mukerjee, S., Srinivasan, S., Eds.; Electrochemical Society, Inc.: Pennington, NJ, 1997; p viii. 
 (38) Bernard, M. C.; Bernard, P.; Keddam, M.; Senyarich, S.; 
 Takenouti, H. Electrochim. Acta 1996, 41, 91. 
 (39) Hibben, J. H. J. Chem. Phys. 1937, 5, 166. (40) Cornell, R. M.; Schwertmann, U. The Iron Oxides: Structure, 
 Properties, Reactions, Occurrences, and Uses, 2nd completely rev. and extended ed.; Wiley-VCH: Weinheim, 2003. 
 (41) Thierry, D.; Persson, D.; Leygraf, C.; Boucherit, N.; 
 Hugotlegoff, A. Corros. Sci. 1991, 32, 273. 
 (42) Graves, P. R.; Johnston, C.; Campaniello, J. J. Mater. Res. Bull. 
 1988 
 , 23, 1651. 
 (43) Jacob, J.; Khadar, M. A. J. Appl. Phys. 2010, 107, 114310. (44) Lu, P. W. T.; Srinivasan, S. J. Electrochem. Soc. 1978, 125, 1416. (45) Cappadonia, M.; Divisek, J.; Vonderheyden, T.; Stimming, U. 
 Electrochim. Acta 1994, 39, 1559. 
 (46) Merrill, M. D.; Dougherty, R. C. J. Phys. Chem. C 2008, 112, 
 3655. 
 (47) Pandya, K. I.; Hoffman, R. W.; McBreen, J.; O 
 ’Grady, W. E. J. 
 Electrochem. Soc. 1990, 137, 383. 
 Journal of the American Chemical Society 
 Article 
 dx.doi.org/10.1021/ja405351s 
 | J. Am. Chem. Soc. 2013, 135, 12329−12337 
 12337

pageNum = 19

pageNum = 20

pageNum = 21

pageNum = 22

pageNum = 23

pageNum = 24

(Figure 3c). The absence of O 
 −H vibrations for NiOOH 
 (Figure 5b, 0.5 
 −0.7 V) is consistent with previous reports. The 
 broad bands present at 3100 
 −3600 cm−1 are well-known 
 characteristics of water. 
 22,39 (Figure S21, SI) 
 In the case of Fe, Figure 6, an intense band at 571 cm− 
 1 is 
 observed at 0 V. This band shows very little change with the 
 applied potential, consistent with the electrochemical character- istics (Figure 1b). An apparent decrease in the intensity of these bands at potentials above 0.35 V is a consequence of the increased background due to the electrochemical oxidation of the underlying gold substrate. (Figure S17, SI) Additional bands are observed at 410 cm− 
 1 and in the range of 650 
 − 
 720 cm− 
 1; these features change slightly in their relative 
 intensities with potential but are too low in intensity for reliable analysis. Unlike the Raman spectra for Ni 
 ﬁlms, the Fe phase 
 displays no structural O 
 −H stretches. Identiﬁcation of the 
 bands observed for Fe 
 ﬁlms is rather diﬃcult owing to the many 
 possible iron oxide and oxyhydroxide phases, many of which are structurally similar. 
 40 The primary band at 571 cm−1 is 
 closest in shift to those for 
 α-FeOOH (560 cm−1) and Fe 
 3O4 
 (553 cm− 
 1), and the low-intensity band at 410 cm−1 matches 
 that of 
 α-Fe 
 2O3 (409 cm 
 −1). 41 The bands at 650−720 cm−1 
 may originate from Fe3O4, γ-Fe2O3 and γ-FeOOH which exhibit bands at 670, 650 
 −740, and 660 cm−1, respectively.41 
 (See additional discussion in section S5.2 of the SI.) 
 We captured the main features of the Ni 
 −Fe system by 
 plotting the Raman spectra of Ni 
 −Fe ﬁlms as functions of 
 composition at two di 
 ﬀerent potentials, 0.2 and 0.6 V vs Hg/ 
 HgO, the latter of which corresponds to an OER overpotential of 0.235 V (Figures 7 and 8, respectively). These two potentials are selected to show composition e 
 ﬀects at potentials below 
 and above the Ni(OH)2/NiOOH redox potential for each catalyst. 
 At 0.2 V, the Raman features for mixed Ni 
 −Fe ﬁlms can be 
 assigned to those of the end-members. The sharp bands of Ni(OH)2 (450 and 494 cm 
 −1) are present for ﬁlms with up to 
 32% Fe (Figure 7a) and disappear at higher Fe contents. On the other hand, the broad band of Ni(OH)2 (∼530 cm 
 −1) 
 remains at these intermediate compositions but has shifted to higher wavenumbers ( 
 ∼560 cm−1). Contributions from the Fe 
 phase are also visible at 578 and 650 
 −720 cm−1. The 
 disappearance of the sharp Ni(OH)2 bands and the persistence of the broad underlying band suggests that Ni 
 −Fe is 
 disordered. Spectra collected at higher wavenumbers (Figure 7b) support this conclusion, namely, that the sharp band at 3580 cm− 
 1 corresponding to the ordered 
 β-Ni(OH) 
 2 phase 
 disappears with increasing Fe content, leaving only a broad feature at 3500 
 −3660 cm−1 for ﬁlms with 19−51% Fe. 
 Under an oxygen evolution potential of 0.6 V vs Hg/HgO 
 (Figure 8a), the two characteristic bands for NiOOH (at 475 and 555 cm− 
 1 for the Ni 
 ﬁlm) are visible for ﬁlms with as much 
 as 90% Fe, owing to the high Raman cross section of NiOOH. 
 26 These two bands appear to merge with increasing 
 Fe content, and an additional band at 494 cm− 
 1 is apparent, as a 
 shoulder on the band at 475 cm− 
 1, for 
 ﬁlms containing 30−80% 
 Fe. The most dramatic e 
 ﬀect of composition is the change in 
 relative intensities of the pair of NiOOH bands; the 555 cm− 
 1 
 band increases in intensity with Fe incorporation. This change in the relative intensities is readily apparent for 
 ﬁlms with as 
 little as 19% Fe and clearly visible for 
 ﬁlms with higher Fe 
 contents despite the appearance of the 494 cm− 
 1 band. At Fe 
 contents of 90% or higher, features associated with Fe (and Au) phases, at 573 and 650 
 −720 cm−1, can be seen. It should be 
 Figure 6. 
 In situ Raman spectra collected for Fe 
 ﬁlms deposited on a 
 roughened Au substrate, shown as a function of potential vs Hg/HgO (1 M KOH) in 0.1 M KOH. The equilibrium potential for the OER is 0.365 V. 
 Figure 7. 
 In situ Raman spectra collected for Ni 
 −Fe catalysts, as a 
 function of composition, in 0.1 M KOH at a potential of 0.2 V vs Hg/ HgO (1 M KOH), shown for the (a) low and (b) high wavenumber regions. 
 Figure 8. 
 In situ Raman spectra collected in 0.1 M KOH at an OER 
 overpotential of 0.235 V, or 0.6 V vs Hg/HgO (1 M KOH), for (a) Ni 
 −Fe ﬁlms as a function of composition and (b) an aged Ni ﬁlm, 
 shown in comparison to the as-deposited Ni 
 ﬁlm from (a). 
 Journal of the American Chemical Society 
 Article 
 dx.doi.org/10.1021/ja405351s 
 | J. Am. Chem. Soc. 2013, 135, 12329−12337 
 12334

pageNum = 25

pageNum = 26

pageNum = 27

pathway. Taken together, our observations demonstrate that the local environment of Ni in Ni 
 −Fe ﬁlms strongly aﬀects the 
 average oxidation state and OER activity of Ni in alkaline electrolytes. 
 ■ ASSOCIATEDCONTENT 
 *S Supporting Information Additional experimental procedures, supplementary electro- chemical data, Raman spectroscopy, X-ray photoelectron spectra. This material is available free of charge via the Internet at http://pubs.acs.org. 
 ■ AUTHORINFORMATION 
 Corresponding Author bell@cchem.berkeley.edu Notes The authors declare no competing 
 ﬁnancial interest. 
 ■ ACKNOWLEDGMENTS 
 This material is based upon work performed by the Joint Center for Arti 
 ﬁcial Photosynthesis, a DOE Energy Innovation 
 Hub, supported through the O 
 ﬃce of Science of the U.S. 
 Department of Energy under Award Number DE-SC0004993. M.W.L. is supported by the University of California President 
 ’s 
 Postdoctoral Fellowship Program. We gratefully acknowledge Ian D. Sharp (Joint Center for Arti 
 ﬁcial Photosynthesis) for 
 valuable discussion and assistance with XPS analysis of the electrocatalysts used in this study, and Elena Kreimer (University of California, Berkeley, College of Chemistry) for assistance with elemental analysis. We also thank Eric Granlund (University of California, Berkeley, College of Chemistry) as well as James Wu and Doug Jamieson (Lawrence Berkeley National Laboratory, Materials Science Division) for the fabrication of our electrodes and electrochemical cells. 
 ■ REFERENCES 
 (1) Walter, M. G.; Warren, E. L.; McKone, J. R.; Boettcher, S. W.; Mi, 
 Q. X.; Santori, E. A.; Lewis, N. S. Chem. Rev. 2010, 110, 6446. 
 (2) Cook, T. R.; Dogutan, D. K.; Reece, S. Y.; Surendranath, Y.; 
 Teets, T. S.; Nocera, D. G. Chem. Rev. 2010, 110, 6474. 
 (3) Matsumoto, Y.; Sato, E. Mater. Chem. Phys. 1986, 14, 397. (4) Corrigan, D. A. J. Electrochem. Soc. 1987, 134, 377. (5) Hu, C. C.; Wu, Y. R. Mater. Chem. Phys. 2003, 82, 588. (6) Landon, J.; Demeter, E.; Inoglu, N.; Keturakis, C.; Wachs, I. E.; 
 Vasic, R.; Frenkel, A. I.; Kitchin, J. R. ACS Catal. 2012, 2, 1793. 
 (7) Li, X. H.; Walsh, F. C.; Pletcher, D. Phys. Chem. Chem. Phys. 
 2011 
 , 13, 1162. 
 (8) Miller, E. L.; Rocheleau, R. E. J. Electrochem. Soc. 1997, 144, 3072. (9) Potvin, E.; Brossard, L. Mater. Chem. Phys. 1992, 31, 311. (10) Singh, R. N.; Pandey, J. P.; Anitha, K. L. Int. J. Hydrogen Energy 
 1993 
 , 18, 467. 
 (11) Trotochaud, L.; Ranney, J. K.; Williams, K. N.; Boettcher, S. W. 
 J. Am. Chem. Soc. 2012, 134, 17253. 
 (12) 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. 
 (13) Lyons, M. E. G.; Brandon, M. P. Int. J. Electrochem. Soc. 2008, 3, 
 1463. 
 (14) Lyons, M. E. G.; Brandon, M. P. Int. J. Electrochem. Soc. 2008, 3, 
 1386. 
 (15) Kim, S. H.; Tryk, D. A.; Antonio, M. R.; Carr, R.; Scherson, D. J. 
 Phys. Chem. 1994, 98, 10269. 
 (16) Balasubramanian, M.; Melendres, C. A.; Mini, S. J. Phys. Chem. B 
 2000 
 , 104, 4300. 
 (17) Corrigan, D. A.; Conell, R. S.; Fierro, C. A.; Scherson, D. A. J. 
 Phys. Chem. 1987, 91, 5009. 
 (18) Guerlou-Demourgues, L.; Fournes, L.; Delmas, C. J. Electrochem. 
 Soc. 1996, 143, 3083. 
 (19) Gao, P.; Gosztola, D.; Leung, L. W. H.; Weaver, M. J. J. 
 Electroanal. Chem. 1987, 233, 211. 
 (20) Yeo, B. S.; Klaus, S. L.; Ross, P. N.; Mathies, R. A.; Bell, A. T. 
 ChemPhysChem 2010, 11, 1854. 
 (21) Trasatti, S.; Petrii, O. A. Pure Appl. Chem. 1991, 63, 711. (22) Kostecki, R.; McLarnon, F. J. Electrochem. Soc. 1997, 144, 485. (23) Yeo, B. S.; Bell, A. T. J. Phys. Chem. C 2012, 116, 8394. (24) Doyle, R. L.; Lyons, M. E. G. J. Electrochem. Soc. 2013, 160, 
 H142. 
 (25) Oliva, P.; Leonardi, J.; Laurent, J. F.; Delmas, C.; Braconnier, J. 
 J.; Figlarz, M.; Fievet, F.; Deguibert, A. J. Power Sources 1982, 8, 229. 
 (26) Desilvestro, J.; Corrigan, D. A.; Weaver, M. J. J. Electrochem. Soc. 
 1988 
 , 135, 885. 
 (27) Fu, X. Z.; Zhu, Y. J.; Xu, Q. C.; Li, J.; Pan, J. H.; Xu, J. Q.; Lin, J. 
 D.; Liao, D. W. Solid State Ionics 2007, 178, 987. 
 (28) Cordoba-Torresi, S. I.; Gabrielli, C.; Hugotlegoff, A.; Torresi, R. 
 J. Electrochem. Soc. 1991, 138, 1548. 
 (29) Wehrens-Dijksma, M.; Notten, P. H. L. Electrochim. Acta 2006, 
 51, 3609. 
 (30) Godwin, I. J.; Lyons, M. E. G. Electrochem. Commun. 2013, 32, 
 39. 
 (31) Johnston, C.; Graves, P. R. Appl. Spectrosc. 1990, 44, 105. (32) Bernard, M. C.; Cortes, R.; Keddam, M.; Takenouti, H.; 
 Bernard, P.; Senyarich, S. J. Power Sources 1996, 63, 247. 
 (33) Cornilsen, B. C.; Shan, X. Y.; Loyselle, P. L. J. Power Sources 
 1990 
 , 29, 453. 
 (34) de Torresi, S. I. C.; Provazi, K.; Malta, M.; Torresi, R. M. J. 
 Electrochem. Soc. 2001, 148, A1179. 
 (35) Vidotti, M.; Salvador, R. P.; de Torresi, S. I. C. Ultrason. 
 Sonochem. 2009, 16, 35. 
 (36) Vidotti, M.; Salvador, R. P.; Ponzio, E. A.; Cordoba de Torresi, 
 S. I. J. Nanosci. Nanotechnol. 2007, 7, 3221. 
 (37) Kostecki, R.; McLarnon, F. In Proceedings of the Symposium on 
 Electrode Materials and Processes for Energy Conversion and Storage IV; McBreen, J., Mukerjee, S., Srinivasan, S., Eds.; Electrochemical Society, Inc.: Pennington, NJ, 1997; p viii. 
 (38) Bernard, M. C.; Bernard, P.; Keddam, M.; Senyarich, S.; 
 Takenouti, H. Electrochim. Acta 1996, 41, 91. 
 (39) Hibben, J. H. J. Chem. Phys. 1937, 5, 166. (40) Cornell, R. M.; Schwertmann, U. The Iron Oxides: Structure, 
 Properties, Reactions, Occurrences, and Uses, 2nd completely rev. and extended ed.; Wiley-VCH: Weinheim, 2003. 
 (41) Thierry, D.; Persson, D.; Leygraf, C.; Boucherit, N.; 
 Hugotlegoff, A. Corros. Sci. 1991, 32, 273. 
 (42) Graves, P. R.; Johnston, C.; Campaniello, J. J. Mater. Res. Bull. 
 1988 
 , 23, 1651. 
 (43) Jacob, J.; Khadar, M. A. J. Appl. Phys. 2010, 107, 114310. (44) Lu, P. W. T.; Srinivasan, S. J. Electrochem. Soc. 1978, 125, 1416. (45) Cappadonia, M.; Divisek, J.; Vonderheyden, T.; Stimming, U. 
 Electrochim. Acta 1994, 39, 1559. 
 (46) Merrill, M. D.; Dougherty, R. C. J. Phys. Chem. C 2008, 112, 
 3655. 
 (47) Pandya, K. I.; Hoffman, R. W.; McBreen, J.; O 
 ’Grady, W. E. J. 
 Electrochem. Soc. 1990, 137, 383. 
 Journal of the American Chemical Society 
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pageNum = 29

pageNum = 30

pageNum = 31

pageNum = 32

pageNum = 33

(Figure 3c). The absence of O 
 −H vibrations for NiOOH 
 (Figure 5b, 0.5 
 −0.7 V) is consistent with previous reports. The 
 broad bands present at 3100 
 −3600 cm−1 are well-known 
 characteristics of water. 
 22,39 (Figure S21, SI) 
 In the case of Fe, Figure 6, an intense band at 571 cm− 
 1 is 
 observed at 0 V. This band shows very little change with the 
 applied potential, consistent with the electrochemical character- istics (Figure 1b). An apparent decrease in the intensity of these bands at potentials above 0.35 V is a consequence of the increased background due to the electrochemical oxidation of the underlying gold substrate. (Figure S17, SI) Additional bands are observed at 410 cm− 
 1 and in the range of 650 
 − 
 720 cm− 
 1; these features change slightly in their relative 
 intensities with potential but are too low in intensity for reliable analysis. Unlike the Raman spectra for Ni 
 ﬁlms, the Fe phase 
 displays no structural O 
 −H stretches. Identiﬁcation of the 
 bands observed for Fe 
 ﬁlms is rather diﬃcult owing to the many 
 possible iron oxide and oxyhydroxide phases, many of which are structurally similar. 
 40 The primary band at 571 cm−1 is 
 closest in shift to those for 
 α-FeOOH (560 cm−1) and Fe 
 3O4 
 (553 cm− 
 1), and the low-intensity band at 410 cm−1 matches 
 that of 
 α-Fe 
 2O3 (409 cm 
 −1). 41 The bands at 650−720 cm−1 
 may originate from Fe3O4, γ-Fe2O3 and γ-FeOOH which exhibit bands at 670, 650 
 −740, and 660 cm−1, respectively.41 
 (See additional discussion in section S5.2 of the SI.) 
 We captured the main features of the Ni 
 −Fe system by 
 plotting the Raman spectra of Ni 
 −Fe ﬁlms as functions of 
 composition at two di 
 ﬀerent potentials, 0.2 and 0.6 V vs Hg/ 
 HgO, the latter of which corresponds to an OER overpotential of 0.235 V (Figures 7 and 8, respectively). These two potentials are selected to show composition e 
 ﬀects at potentials below 
 and above the Ni(OH)2/NiOOH redox potential for each catalyst. 
 At 0.2 V, the Raman features for mixed Ni 
 −Fe ﬁlms can be 
 assigned to those of the end-members. The sharp bands of Ni(OH)2 (450 and 494 cm 
 −1) are present for ﬁlms with up to 
 32% Fe (Figure 7a) and disappear at higher Fe contents. On the other hand, the broad band of Ni(OH)2 (∼530 cm 
 −1) 
 remains at these intermediate compositions but has shifted to higher wavenumbers ( 
 ∼560 cm−1). Contributions from the Fe 
 phase are also visible at 578 and 650 
 −720 cm−1. The 
 disappearance of the sharp Ni(OH)2 bands and the persistence of the broad underlying band suggests that Ni 
 −Fe is 
 disordered. Spectra collected at higher wavenumbers (Figure 7b) support this conclusion, namely, that the sharp band at 3580 cm− 
 1 corresponding to the ordered 
 β-Ni(OH) 
 2 phase 
 disappears with increasing Fe content, leaving only a broad feature at 3500 
 −3660 cm−1 for ﬁlms with 19−51% Fe. 
 Under an oxygen evolution potential of 0.6 V vs Hg/HgO 
 (Figure 8a), the two characteristic bands for NiOOH (at 475 and 555 cm− 
 1 for the Ni 
 ﬁlm) are visible for ﬁlms with as much 
 as 90% Fe, owing to the high Raman cross section of NiOOH. 
 26 These two bands appear to merge with increasing 
 Fe content, and an additional band at 494 cm− 
 1 is apparent, as a 
 shoulder on the band at 475 cm− 
 1, for 
 ﬁlms containing 30−80% 
 Fe. The most dramatic e 
 ﬀect of composition is the change in 
 relative intensities of the pair of NiOOH bands; the 555 cm− 
 1 
 band increases in intensity with Fe incorporation. This change in the relative intensities is readily apparent for 
 ﬁlms with as 
 little as 19% Fe and clearly visible for 
 ﬁlms with higher Fe 
 contents despite the appearance of the 494 cm− 
 1 band. At Fe 
 contents of 90% or higher, features associated with Fe (and Au) phases, at 573 and 650 
 −720 cm−1, can be seen. It should be 
 Figure 6. 
 In situ Raman spectra collected for Fe 
 ﬁlms deposited on a 
 roughened Au substrate, shown as a function of potential vs Hg/HgO (1 M KOH) in 0.1 M KOH. The equilibrium potential for the OER is 0.365 V. 
 Figure 7. 
 In situ Raman spectra collected for Ni 
 −Fe catalysts, as a 
 function of composition, in 0.1 M KOH at a potential of 0.2 V vs Hg/ HgO (1 M KOH), shown for the (a) low and (b) high wavenumber regions. 
 Figure 8. 
 In situ Raman spectra collected in 0.1 M KOH at an OER 
 overpotential of 0.235 V, or 0.6 V vs Hg/HgO (1 M KOH), for (a) Ni 
 −Fe ﬁlms as a function of composition and (b) an aged Ni ﬁlm, 
 shown in comparison to the as-deposited Ni 
 ﬁlm from (a). 
 Journal of the American Chemical Society 
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pageNum = 36

pathway. Taken together, our observations demonstrate that the local environment of Ni in Ni 
 −Fe ﬁlms strongly aﬀects the 
 average oxidation state and OER activity of Ni in alkaline electrolytes. 
 ■ ASSOCIATEDCONTENT 
 *S Supporting Information Additional experimental procedures, supplementary electro- chemical data, Raman spectroscopy, X-ray photoelectron spectra. This material is available free of charge via the Internet at http://pubs.acs.org. 
 ■ AUTHORINFORMATION 
 Corresponding Author bell@cchem.berkeley.edu Notes The authors declare no competing 
 ﬁnancial interest. 
 ■ ACKNOWLEDGMENTS 
 This material is based upon work performed by the Joint Center for Arti 
 ﬁcial Photosynthesis, a DOE Energy Innovation 
 Hub, supported through the O 
 ﬃce of Science of the U.S. 
 Department of Energy under Award Number DE-SC0004993. M.W.L. is supported by the University of California President 
 ’s 
 Postdoctoral Fellowship Program. We gratefully acknowledge Ian D. Sharp (Joint Center for Arti 
 ﬁcial Photosynthesis) for 
 valuable discussion and assistance with XPS analysis of the electrocatalysts used in this study, and Elena Kreimer (University of California, Berkeley, College of Chemistry) for assistance with elemental analysis. We also thank Eric Granlund (University of California, Berkeley, College of Chemistry) as well as James Wu and Doug Jamieson (Lawrence Berkeley National Laboratory, Materials Science Division) for the fabrication of our electrodes and electrochemical cells. 
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 Electrochem. Soc. 1990, 137, 383. 
 Journal of the American Chemical Society 
 Article 
 dx.doi.org/10.1021/ja405351s 
 | J. Am. Chem. Soc. 2013, 135, 12329−12337 
 12337

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