Noble Metal-Free Copper Hydroxide as an Active ... - ACS Publications

Research Article pubs.acs.org/journal/ascecg

Noble Metal-Free Copper Hydroxide as an Active and Robust Electrocatalyst for Water Oxidation at Weakly Basic pH Shengsheng Cui, Xiang Liu, Zijun Sun, and Pingwu Du* CAS Key Laboratory of Materials for Energy Conversion, Department of Materials Science and Engineering, iChEM (Collaborative Innovation Center of Chemistry for Energy Materials), University of Science and Technology of China (USTC), 96 Jinzhai Road, Hefei, Anhui Province 230026, P. R. China S Supporting Information *

ABSTRACT: Copper hydroxide (Cu(OH)2) is a quite inexpensive and abundant material, but the literature contains no reports of using it as a stable water oxidation catalyst (WOC). In this study, we report for the first time that Cu(OH)2 material synthesized from a simple copper salt can be used as a WOC with good activity and stability. Under optimal conditions using Cu(OH)2 as the electrocatalyst, a catalytic current density of 0.1 mA/cm2 can be achieved under an applied potential of ∼1.05 V relative to Ag/AgCl at pH 9.2. The slope of the Tafel plot is 78 mV/dec. The Tafel plot indicates that a current density of ∼0.1 mA/cm2 requires an overpotential of 550 mV. The Faradaic efficiency was measured to be ∼95%. The as-synthesized Cu(OH)2 material was characterized by X-ray powder diffraction, scanning electron microscopy, Fourier transform infrared spectroscopy, and X-ray photoelectron spectroscopy. KEYWORDS: Water oxidation, Electrocatalysis, Copper hydroxide, Catalyst

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for water oxidation catalysis.29−34 The complex (2,2′bipyridine)Cu(OH)2 was applied for this purpose, and using it, water oxidation catalysis occurs at approximately 750 mV of overpotential.29 Homogeneous water oxidation catalysis can be conducted using simple Cu(II) salts in Na2CO3 solution, and the study provided some evidence for water oxidation catalysis by a single metal site, dicopper, and even interfacial catalysis.30 Although these homogeneous copper catalysts generally require high overpotentials to realize water oxidation, their tunability in molecular and electronic structure are advantageous for gaining insight into the catalytic mechanism. Recently, our group reported that heterogeneous CuO material electrodeposited from a molecular copper(II) 2-pyridylmethylamine complex is an active catalyst for water oxidation.10 The following studies using CuO synthesized from a simple copper salt directly showed that its catalytic activity can be tuned by size and morphology.35 Different from CuO-based materials, Cu(OH)2 is another common copper-based compound. Previous studies in molecular catalysis proposed that the structure of Cu−OH might be the active site for catalysis.31 In addition, some layered double hydroxides (LDHs) containing transition metal cations, particularly Ni, Fe, Zn, Cu, and Co, are reported for water oxidation catalysis.36−41 For Cu-containing LDHs, the catalysts usually contain not only copper but also other metal elements.42−44 For example, the (Cu/Ti)LDH material was

INTRODUCTION Hydrogen production via water splitting using solar energy has attracted great attention because sunlight is a renewable, carbon-neutral, and abundant energy source of sufficient scale to replace fossil fuels and meet increasing global energy demands.1−4 In addition to the photolysis of water, hydrogen production by electrolysis of water is also an appealing way to store energy if the electric power is obtained from renewable sources of energy, such as photovoltaics.5,6 With this aim, one of the crucial scientific challenges is to develop efficient catalytic systems for electrochemical production of hydrogen. However, the kinetic bottleneck of water splitting is the oxygen evolution reaction (OER) because of the four electron (4e−) transfer process.7,8 Thus, there is a high demand for an electrocatalyst with high activity, excellent stability, and low overpotential for OER.9−11 Electrocatalysts based on precious metals, such as IrO2 and RuO2, have been developed as water oxidation catalysts (WOCs) since the late 1970s,12−14 but their practical applications are significantly hampered due to the low abundance and high-cost of these noble metals.15,16 Therefore, great efforts have been made to develop low-cost alternatives based on metals such as cobalt,17−19 nickel,20−22 iron,23,24 and manganese.25−27 As an earth-abundant metal, copper is a well-known catalyst for oxidizing organic compounds such as phenols, alcohols, and hydrocarbons.28 In recent years, there has been increasing interest in copper-based WOCs due to their high abundance, low cost, and rich redox properties. Since 2012, a few copper(II)-based homogeneous systems have been reported © 2016 American Chemical Society

Received: January 12, 2016 Revised: March 20, 2016 Published: March 21, 2016 2593

DOI: 10.1021/acssuschemeng.6b00067 ACS Sustainable Chem. Eng. 2016, 4, 2593−2600

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ACS Sustainable Chemistry & Engineering

elemental composition of the Cu(OH)2 catalyst and the valence states of metal elements were probed with an ESCALAB 250 X-ray photoelectron spectroscopy (XPS) instrument. The high resolution XPS spectra of Cu 2p, Cu LMM, and O 1s were obtained. The binding energies were obtained with reference to the C 1s peak (285.0 eV). The crystalline phase analysis of the catalyst was measured by X-ray diffraction (XRD, D/max-TTR III) via graphite monochromatized Cu Kα radiation of 1.54178 Å, operating at 40 kV and 200 mA. The scanning rate was 5° min−1 from 10° to 80° in 2θ.

reported to produce oxygen in water under visible light in the presence of AgNO3 as a sacrificial agent.43 To date, it is not clear if copper plays a key role in this LDH material. To the best of our knowledge, no pure Cu(OH)2 material has been reported for water oxidation catalysis. In this present study, we report that pure Cu(OH)2 material directly synthesized from simple copper divalent salts behaves as an active electrocatalyst for water oxidation. The morphologies and compositions of Cu(OH)2 were characterized by X-ray powder diffraction (XRD), scanning electron microscopy (SEM), Fourier transform infrared spectroscopy (FTIR), and X-ray photoelectron spectroscopy (XPS).

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RESULTS AND DISCUSSION Crystalline Cu(OH)2 material was facilely synthesized from a simple copper salt with no template, as described in the Experimental Details. For the formation of Cu(OH)2 to be confirmed, the resulting material was characterized by powder X-ray diffraction (XRD) (Figure 1a). The diffraction peaks of

EXPERIMENTAL DETAILS

Chemicals. Commercially available Cu(NO3)2·3H2O (99%), ammonium hydroxide (NH3·H2O, 25%−28%), NaOH (96%), and Cu(OH)2 (98%) were obtained from Aldrich or Acros and used without further purification. All aqueous solutions were prepared with Millipore water (resistivity: ∼18 MΩ cm). The Cu(OH)2 material was loaded onto glass plates coated with fluorine-doped tin oxide (FTO, purchased from Zhuhai Kaivo Electronic Components Co., Ltd.; surface resistivity: 8−12 Ω sq−1), which were cleaned by sonication three times each in deionized water and ethanol for 10 min, followed by drying in air. Preparation of Cu(OH)2 Material. One gram of Cu(NO3)2·3H2O was initially dissolved in 100 mL of water followed by the addition of 30 mL of NH3·H2O (0.3 M) under constant stirring. A blue precipitate of Cu(OH)2 was produced by slowly adding the NaOH (1.0 M) solution until the pH value reached 9−10. Then, the blue Cu(OH)2 precipitate was filtered, washed three times by water to obtain a solid product, and dried at room temperature. Electrochemical Methods. Electrochemical property analyses were performed at room temperature in a three-electrode cell connected to a CHI602E potentiostat (Shanghai Chenhua Instrument Co., Ltd.). The reference electrode was a Ag/AgCl electrode (in 3 M KCl solution). Pt wire was used as the counter electrode, and a fluorine-doped tin oxide (FTO) plate was used as the working electrode. All cyclic voltammograms (CV) were measured with a scan rate of 50 mV s−1 unless otherwise noted, and iR compensations were applied during the measurements. Bulk electrolysis was carried out at variable potentials with no iR compensation and mild stirring. Cu(OH)2/FTO Electrode Preparation. The Cu(OH)2 material was loaded onto an FTO electrode via the drop-casting method. Ten milligrams of Cu(OH)2 was dispersed in 0.9 mL of ethanol and 0.1 mL of 5% Nafion solution. The resulting solution was homogenized by ultrasonic dispersion for 20 min. Then, 40 μL of the solution was drop-cast onto a clean FTO plate (0.5 cm2), and the Cu(OH)2/FTO electrode was dried at room temperature. The pH dependence of the CVs was measured for the Cu(OH)2/FTO electrode in 0.1 M boric acid electrolyte (pH 5.92), and the pH was then adjusted by adding different amounts of 0.1 M KOH. Faradaic Efficiency. For controlled potential electrolysis (CPE), the generated oxygen in the headspace was quantified by a fluorescence-based oxygen sensor (Ocean Optics). The experiment was performed in a gastight electrochemical cell, and the Cu(OH)2/ FTO electrode was used as the working electrode. The solution was degassed by bubbling with high purity N2 for 15 min. The reference electrode was positioned several millimeters (5−10 mm) from the working electrode. The O2 sensor on the FOXY probe, recorded at 2 s intervals, provided the data to calculate the partial pressure of O2 in the headspace. After recording the partial pressure of O2 for 30 min to check the stability in the absence of an applied potential, bulk electrolysis was initiated at 1.20 V without iR drop compensation. Physical Characterization. Scanning electron microscopy (SEM) images were obtained on a SIRION200 Schottky field emission scanning electron microscope (SFE-SEM). The samples were rinsed three times by Millipore water, dried in air, and then coated with Au to make the samples conductive before loading into the instrument. The

Figure 1. (a) XRD patterns of synthesized (red) and commercial (black) Cu(OH)2. The bottom lines are standard XRD patterns of Cu(OH)2 (PDF 35-0505) (blue).

the Cu(OH)2 sample were sharp, and no impurity peaks were observed, indicating high crystallinity and good purity. The main peaks are located at 16.7°, 23.8°, 34.1°, 35.9°, 39.7°, 53.3°, and 63.0° in 2θ, corresponding to the (020), (021), (002), (111), (130), (150), and (200) planes, respectively, of orthorhombic Cu(OH)2 (PDF 35-0505). The black plot in Figure 1a is the XRD data of commercial Cu(OH)2, which shows very similar diffraction peaks. For further confirmation of the Cu(OH)2 material, XPS was performed to verify the composition and valence states of the sample, as shown in Figure 2. Figure 2a is the survey data, which shows the presence of Cu 2p, O 1s, and C 1s peaks. Figure 2b is the high-resolution XPS spectra of Cu 2p. In the Cu 2p spectra, typical Cu 2p3/2 and Cu 2p1/2 peaks are observed with binding energies of 933.8 and 953.8 eV, respectively, as well as their associated shakeup peaks at 941.2 and 961.2 eV. These Cu 2p peaks are consistent with previous studies of Cu(OH)2.45−47 The presence of Cu(OH)2 can be further confirmed by the presence of a peak at 916.8 eV in the Cu LMM spectrum (Figure 2c).48 Moreover, a wellresolved O 1s peak is observed at 531.2 eV (Figure 2d), also confirming the presence of Cu(OH)2.49 The morphologies of the as-synthesized and commercial Cu(OH)2 materials were examined by SEM (Figure 3). In Figure 3a, the morphology of as-prepared Cu(OH)2 material shows uneven nanosheets with a thickness of ∼50 nm and some short nanowires. The diameter of the nanowires in the asprepared Cu(OH)2 sample is much shorter than that in commercial Cu(OH)2. The overall structure is loose and irregular. The commercial Cu(OH)2 analysis in Figure 3b has 2594


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Figure 2. X-ray photoelectron spectroscopy (XPS) survey data (a) and high-resolution XPS spectra of Cu 2p (b); high-resolution XPS spectra of Cu LMM (c) and O 1s (d) of as-synthesized Cu(OH)2 material.

Figure 3. SEM images of (a) as-synthesized and (b) commercial Cu(OH)2.

quite uniform nanowire structure with an average diameter of ∼100 nm. Although each single nanowire probably exhibits relatively high specific surface area,50 all of the nanowires are aggregated together, which may greatly decrease the total surface area. The great difference in morphologies between the commercial and as-synthesized Cu(OH)2 may be one of the major factors that affects the subsequent catalytic performance toward water oxidation. Figure 4a shows cyclic voltammograms (CVs) of the assynthesized Cu(OH)2 material in 0.1 M potassium borate (KBi) solution at pH 9.2 using different loading amounts of Cu(OH)2 sample onto FTO plates. Pt wire was used as the counter electrode, and Ag/AgCl (3 M KCl) was used as the reference electrode. From the CVs results, an obvious catalytic activity for water oxidation was observed at an onset potential of ∼1.05 V, producing a current density of 0.1 mA/cm2 (note: all potentials are vs Ag/AgCl in this paper). Gas bubbles were clearly seen on the electrode and were confirmed to be oxygen by both gas chromatography and a fluorescence-based oxygen sensor. This observation also indicates that water oxidation catalysis did take place during the anodic scan, and the catalytic activity may occur because the as-synthesized Cu(OH)2

material has some structures or defect sites similar to the Cubased catalysts in previous reports.29,30 The catalytic current density can be improved by increasing the loading amount from 0.2 to 0.8 mg/cm2 (Figure 4a). When using 0.2 mg/cm2 of Cu(OH)2, the catalytic current density reached 2.0 mA/cm2 under an applied potential of 1.5 V. The current densities reached 2.2 and 3.2 mA/cm2 when the loading of Cu(OH)2 was increased to 0.6 and 0.8 mg/cm2, respectively. However, the catalytic current density decreased with further Cu(OH)2 loading. For this result to be fully understood, SEM images and Nyquist diagrams of the as-synthesized Cu(OH)2 at various loading amounts were collected. As shown in Figure S1, there is no morphology change with different amounts because the Cu(OH)2 material was synthesized before loading onto the FTO electrode. Therefore, the steady increase of current density probably did not have a direct relationship with any morphology change but may have been influenced by the increasing amount of the catalyst. Meanwhile, the reason that the current density dropped with 1 mg/cm2 loading is probably the poor conductivity of Cu(OH)2 material, which can be demonstrated by the Nyquist diagrams of the Cu(OH)2/FTO electrode (Figure S2). The resistance of the electrode increased 2595


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formation of a high oxidation state Cu intermediate(s) during oxidation.10,31,51 Interestingly, as shown in the reverse/cathodic scan of CV, a clear cathodic peak can be observed at ∼0.90 V, which is probably due to reduction of an intermediate. According to the cathodic scan, the peak might be assigned to an oxidation/reduction wave rather than oxidative decomposition of the material. On the basis of the above results and previous reports, CuII is probably oxidized to produce high oxidation state Cu intermediate(s), which is(are) responsible for the water oxidation catalysis. Figure 5a shows the CVs of the as-synthesized Cu(OH)2 material in 0.1 M KBi solution at pH 9.2 with different scan rates. The value of the peak current is increased with a higher scan rate, and this observation might be related to a kinetic process. The peak current varies linearly with the square root of the scan rate (insert, Figure 5a). This result is consistent with the Randles−Sevcik equation, indicating that the catalytic process is a diffusion-controlled process.52,53 In addition, the current densities for catalytic water oxidation are highly dependent on pH values (Figure 5b). Higher catalytic current intensities were observed when the pH was increased to 13.03. For example, the current density was 0.6 mA/cm2 under an applied potential of 1.5 V at pH 7.92. When pH values were increased to 13.03, the current density reached 6.4 mA/cm2 under the same potential. The black plot in Figure 5b shows no appreciable catalytic current when the pH is 5.92, indicating that the Cu(OH)2 catalyst has a weak catalytic activity in acidic solution. Bulk electrolysis experiments were performed in 0.1 M KBi solution at pH 9.2 (Figure 6a). Under applied potentials of 1.05 and 1.20 V, stable catalytic current densities for water oxidation reached 0.2 and 0.9 mA/cm2, respectively. Obvious gas bubbles were clearly observed on the surface of the Cu(OH)2/FTO electrode during electrolysis under 1.05 V (inset, Figure 6a). Interestingly, the current density gradually increased in the initial hour, and then both catalytic current plots showed good stability over a period of >8 h. We hypothesized that this initial period may be the activation process of Cu(OH)2 material to produce high oxidation intermediates. To verify this supposition, we performed bulk electrolysis under a potential of 1.20 V for 6 h and then stopped the catalytic reaction. The same Cu(OH)2/FTO electrode was used for the second and third runs of bulk electrolysis experiments under the same potential. As shown in Figure 6b, the first run (6 h) did show evidence of an activation process. The current density increased in the first run, and it stayed stable in the following two runs.

Figure 4. (a) CVs obtained by loading different amounts of synthesized Cu(OH)2 as working electrode for the water oxidation reaction; (b) comparison of water oxidation catalysis using commercial (red) and as-synthesized (blue) Cu(OH)2 materials. The black line data are obtained from bare FTO. Inset: Magnification of the blue plot between 0.6 and 1.15 V.

with a higher loading amount. When the amount of catalyst is high enough, the increasing resistance will result in a decrease of the catalytic activity. The results indicate that an optimal loading of Cu(OH)2 is around 0.8 mg/cm2, which was used in all subsequent experiments. For comparison, a bare FTO plate and commercial Cu(OH)2 were used for water oxidation (Figure 4b). No obvious catalytic activity was observed using a bare FTO plate as the working electrode for an anodic scan. In contrast, the commercial Cu(OH)2 showed catalytic activity for water oxidation, but it was lower than that of the as-synthesized Cu(OH)2 material with a catalytic potential at ∼1.15 V reaching 0.1 mA/cm2. In Figure 4b, the CV curves of commercial and as-synthesized Cu(OH)2 have a peak/plateau at ∼1.40 V, which is consistent with the CV curves of as-synthesized Cu(OH)2 under different scan rates (Figure 5a). This peak could be assigned to the

Figure 5. (a) CVs in 0.1 M KBi solution at pH 9.2 using synthesized Cu(OH)2 as working electrode at different scan rates. Inset: dependence of the peak current at Eappl = 1.4 V (vs Ag/AgCl) vs ν1/2. (b) CVs in 0.1 M KBi solution at different pH levels by adding 0.1 M KOH electrolytes. 2596


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To further check the stability of the Cu(OH)2 catalyst, we collected XRD, FTIR, and SEM data after electrocatalytic water oxidation for 6 h. The black plot in Figure 7a shows the diffraction peaks of as-synthesized Cu(OH)2 before and after electrolysis. The typical Cu(OH)2 diffraction peaks remained the same as those before electrolysis. In addition, no obvious color change can be observed in the optical pictures of Cu(OH)2 before and after electrolysis (Figure 7a, Inset). In addition, the FTIR spectrum was used to investigate the surface properties of the Cu(OH)2. As shown in Figure 7b, the peaks in the infrared spectra of Cu(OH)2 at low frequencies below 700 cm−1 are due to Cu−O vibrations, whereas bands centered at ∼3572 and ∼3351 cm−1 correspond to the hydroxyl ions.28,54 The broad bands centered at 3426 and 1632 cm−1 are attributed to the O−H stretching and bending modes, respectively, of water.55,56 The peak at ∼1384 cm−1 is from the C−H vibration, and the absorption peak at ∼1116 cm−1 can be assigned to Cu−O vibration corresponding to metal cations. Therefore, none of the peaks showed obvious change after bulk electrolysis. The SEM image of Cu(OH)2 after electrolysis under an applied potential of 1.20 V was obtained, as shown in Figure S3. The result shows that the electrolysis caused no significant change in the morphology of the Cu(OH)2 material. XPS was performed to check the valence states of the assynthesized Cu(OH)2 sample after electrolysis. Panels c and d in Figure 7 show the high-resolution XPS spectra of Cu 2p and Cu LMM after electrolysis. The later XPS spectra are nearly identical to the ones before electrolysis. In the Cu 2p spectra, the existence of Cu(OH)2 was confirmed by the binding energies of Cu 2p3/2 and Cu 2p1/2 located at 933.8 and 953.8 eV, respectively. The typical Cu LMM peak is located at 916.8 eV. Therefore, XRD, FTIR, SEM, and XPS results indicate that

Figure 6. (a) Profiles of bulk electrolysis obtained by using assynthesized Cu(OH)2 as the catalyst for water oxidation under 1.05 V (black) and 1.20 V (red). Inset: oxygen bubbles on the Cu(OH)2/ FTO electrode during bulk electrolysis under 1.20 V; (b) current density plot using synthesized Cu(OH)2 under 1.20 V for three successive runs of bulk electrolysis. All of these experiments were performed in 0.1 M KBi solution at pH 9.2.

Figure 7. (a) XRD patterns of as-synthesized Cu(OH)2 material before (black) and after (red) electrolysis for 6 h at 1.20 V. The bottom lines are standard Cu(OH)2 XRD pattern (Cu(OH)2, PDF 35-0505). Inset: photos of the Cu(OH)2 catalyst before and after electrolysis; (b) FT-IR spectra of the as-synthesized Cu(OH)2 before (black) and after (red) bulk electrolysis under an applied potential of 1.20 V. High-resolution XPS spectra of Cu 2p (c) and Cu LMM (d) of as-synthesized Cu(OH)2 material before and after bulk electrolysis. All these experiments were performed in 0.1 M KBi solution at pH 9.2. 2597


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single peak located at 917.8 eV, confirming the existence of CuO instead of Cu(OH)2.48,58 The high-resolution O 1s spectrum is also provided (Figure S4d). One peak at 529.4 eV is probably from the oxygen element in CuO, and another peak located at 532.1 eV can be assigned to the chemisorbed water, as is also reported in previous studies.48 To understand the electrocatalytic properties of Cu(OH)2, we obtained the Tafel plot by measuring the stable current density (j) of the catalyst under various potentials as a function of the overpotential (η) based on the Nernstian potential for the H2O/O2 half-reaction. The experiment was run in 0.1 M KBi solution at pH 9.2 (Figure 9a). The overpotential is defined as η = Vappl − iR − EpH, where Vappl is the applied potential vs NHE, i is the stable current, R is the uncompensated resistance, and EpH is the thermodynamic potential for water oxidation (EpH = 1.23−0.059 pH vs NHE). From the result in Figure 9a, appreciable catalytic current is observed starting at η = 470 mV (onset), and a current density of ∼0.1 mA/cm2 required η = 550 mV. A nearly linear relationship with the slope of the line being ∼78 mV/decade was observed from 0.47 to 0.63 V, demonstrating efficient kinetics of water oxidation catalysis. Table S1 summarizes the catalytic activity and Tafel slopes for the reported Cu-based catalysts for water oxidation. The Faradaic efficiency of the Cu(OH)2 catalyst was measured by a fluorescence-based oxygen sensor, and the results are shown in Figure 9b. Bulk electrolysis was performed in 0.1 M KBi solution at pH 9.2 in a gastight electrochemical cell under an inert atmosphere. The sensor was placed in the headspace of the cell, and the working electrode was the Cu(OH)2/FTO electrode. After initiating electrolysis, oxygen bubbles were rapidly generated accompanied by a rise of the oxygen percentage in the headspace. The theoretical yield of oxygen was calculated by assuming that all of the currents were caused by 4e− oxidation of water to produce oxygen. Comparing the experimental with the theoretical data, a Faradaic yield of ∼95% for oxygen production was achieved during electrolysis for 1 h.

the Cu(OH)2 material can be used as a robust catalyst for the water oxidation reaction. The catalytic performance of the Cu(OH)2 material was further studied under different pH values in various electrolytes (0.1 M potassium phosphate (KPi) at pH 7.0, 0.1 M KBi at pH 9.2, 0.1 M KOH at pH 13.0, and 1.0 M KOH at pH 13.6), and the results are shown in Figure 8. The CVs results show that

Figure 8. Cyclic voltammograms (CVs) obtained using the synthesized Cu(OH)2 as the catalyst for water oxidation reactions in 0.1 M KPi (black), 0.1 M KBi (red), 0.1 M KOH (blue), and 1.0 M KOH (cyan) solutions at different pH values.

Cu(OH)2 can catalyze water oxidation in all four electrolytes. The catalytic current densities increased when the pH values were higher, which is consistent with the shifts in the thermodynamic requirements for water oxidation. These results further demonstrated that the Cu(OH)2 material is an active catalyst in both neutral and basic solutions. Interestingly, the Cu(OH)2 rapidly changed color from blue to black during electrolysis in 0.1 and 1.0 M KOH electrolytes, indicating that the Cu(OH)2 catalyst is not very robust under strong basic conditions. The black color suggested the formation of active CuO as the real catalyst, as suggested by our previous study.57 Further XPS characterizations confirmed the formation of CuO in strong basic solution (Figure S4). In the Cu 2p spectrum (Figure S4b), two obvious peaks with binding energies at 933.9 and 953.9 eV were assigned to Cu 2p3/2 and Cu 2p1/2, respectively, matching the typical character of Cu(II) in CuO.10,35,48 Because the binding energy of Cu 2p3/2 in both CuO and Cu(OH)2 is located at a similar position, the Cu LMM spectrum was measured to identify the formation of CuO (Figure S4c). The Cu LMM spectrum showed an obvious

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CONCLUSIONS In summary, for the first time, pure crystalline Cu(OH)2 material was studied as an active water oxidation catalyst. The results showed that Cu(OH)2 can serve as a good and robust catalyst in 0.1 M KBi electrolyte at pH 9.2 and exhibits an obvious catalytic current density of 0.1 mA/cm2 with the onset potential at ∼1.05 V. A Faradaic efficiency of ∼95% for

Figure 9. (a) Tafel plot, η = (Vappl − iR − EpH), of the synthesized Cu(OH)2 in a 0.1 M KBi buffer at pH 9.2. (b) Theoretical and experimental oxygen evolution from water by bulk electrolysis at 1.20 V using the as-synthesized Cu(OH)2 catalyst at pH 9.2. 2598


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efficient water oxidation. ACS Appl. Mater. Interfaces 2014, 6, 10929− 10934. (12) Horkans, J.; Shafer, M. W. Investigation of Electrochemistry of a Series of Metal Dioxides with Rutile-Type Structure - MoO2, WO2, ReO2, RuO2, OsO2, and IrO2. J. Electrochem. Soc. 1977, 124, 1202− 1207. (13) Galizzioli, D.; Tantardini, F.; Trasatti, S. Ruthenium dioxidenew electrode material. 1. behavior in acid solutions of inert electrolytes. J. Appl. Electrochem. 1974, 4, 57−67. (14) Harriman, A.; Pickering, I. J.; Thomas, J. M.; Christensen, P. A. Metal-oxides as heterogeneous Catalysts for oxygen evolution under photochemical conditions. J. Chem. Soc., Faraday Trans. 1 1988, 84, 2795−2806. (15) Fang, Y. H.; Liu, Z. P. Mechanism and tafel lines of electrooxidation of water to oxygen on RuO2(110). J. Am. Chem. Soc. 2010, 132, 18214−18222. (16) Gust, D.; Moore, T. A.; Moore, A. L. Solar Fuels via Artificial Photosynthesis. Acc. Chem. Res. 2009, 42, 1890−1898. (17) Kanan, M. W.; Nocera, D. G. In situ formation of an oxygenevolving catalyst in neutral water containing phosphate and Co2+. Science 2008, 321, 1072−1075. (18) McAlpin, J. G.; Surendranath, Y.; Dinca, M.; Stich, T. A.; Stoian, S. A.; Casey, W. H.; Nocera, D. G.; Britt, R. D. EPR Evidence for Co(IV) Species Produced During Water Oxidation at Neutral pH. J. Am. Chem. Soc. 2010, 132, 6882−6883. (19) Nocera, D. G. The Artificial Leaf. Acc. Chem. Res. 2012, 45, 767−776. (20) Dinca, M.; Surendranath, Y.; Nocera, D. G. Nickel-borate oxygen-evolving catalyst that functions under benign conditions. Proc. Natl. Acad. Sci. U. S. A. 2010, 107, 10337−10341. (21) Bediako, D. K.; Lassalle-Kaiser, B.; Surendranath, Y.; Yano, J.; Yachandra, V. K.; Nocera, D. G. Structure-Activity Correlations in a Nickel-Borate Oxygen Evolution Catalyst. J. Am. Chem. Soc. 2012, 134, 6801−6809. (22) Yu, X.; Hua, T.; Liu, X.; Yan, Z.; Xu, P.; Du, P. Nickel-based thin film on multi-walled carbon nanotubes as an efficient bifunctional electrocatalyst for water splitting. ACS Appl. Mater. Interfaces 2014, 6, 15395−15402. (23) Fillol, J. L.; Codola, Z.; Garcia-Bosch, I.; Gomez, L.; Pla, J. J.; Costas, M. Efficient water oxidation catalysts based on readily available iron coordination complexes. Nat. Chem. 2011, 3, 807−813. (24) Ellis, W. C.; McDaniel, N. D.; Bernhard, S.; Collins, T. J. Fast Water Oxidation Using Iron. J. Am. Chem. Soc. 2010, 132, 10990− 10991. (25) Dismukes, G. C.; Brimblecombe, R.; Felton, G. A. N.; Pryadun, R. S.; Sheats, J. E.; Spiccia, L.; Swiegers, G. F. Development of Bioinspired Mn4O4-Cubane Water Oxidation Catalysts: Lessons from Photosynthesis. Acc. Chem. Res. 2009, 42, 1935−1943. (26) McEvoy, J. P.; Brudvig, G. W. Water-Splitting Chemistry of Photosystem II. Chem. Rev. 2006, 106, 4455−4483. (27) Tagore, R.; Crabtree, R. H.; Brudvig, G. W. Oxygen Evolution Catalysis by a Dimanganese Complex and its Relation to Photosynthetic Water Oxidation. Inorg. Chem. 2008, 47, 1815−1823. (28) Prathap, M. U. A.; Kaur, B.; Srivastava, R. Hydrothermal synthesis of CuO micro-/nanostructures and their applications in the oxidative degradation of methylene blue and non-enzymatic sensing of glucose/H2O2. J. Colloid Interface Sci. 2012, 370, 144−154. (29) Barnett, S. M.; Goldberg, K. I.; Mayer, J. M. A soluble copperbipyridine water-oxidation electrocatalyst. Nat. Chem. 2012, 4, 498− 502. (30) Chen, Z. F.; Meyer, T. J. Copper(II) Catalysis of Water Oxidation. Angew. Chem., Int. Ed. 2013, 52, 700−703. (31) Zhang, M. T.; Chen, Z. F.; Kang, P.; Meyer, T. J. Electrocatalytic Water Oxidation with a Copper(II) Polypeptide Complex. J. Am. Chem. Soc. 2013, 135, 2048−2051. (32) Coggins, M. K.; Zhang, M. T.; Chen, Z. F.; Song, N.; Meyer, T. J. Single-Site Copper(II) Water Oxidation Electrocatalysis: Rate Enhancements with HPO42‑ as a Proton Acceptor at pH 8. Angew. Chem., Int. Ed. 2014, 53, 12226−12230.

water oxidation was achieved using Cu(OH)2/FTO as the working electrode under an applied potential of 1.20 V vs Ag/ AgCl. The Tafel plot was determined at pH 9.2, and its slope was ∼78 mV/decade. Because Cu(OH)2 catalyst is based on a low-cost and earth-abundant transition metal (Cu), the present results highlight a new exploration for the development of easily prepared, earth-abundant, heterogeneous copper-based catalysts that oxidize water to produce oxygen.

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S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acssuschemeng.6b00067. SEM images, XPS data after water oxidation catalysis, and Nyquist diagrams (PDF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel/Fax: 86-551-63606207. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was financially supported by the National Natural Science Foundation of China (21271166, 21473170), the Fundamental Research Funds for the Central Universities (WK2060140015, WK2060190026), the Program for New Century Excellent Talents in University (NCET), and the Thousand Young Talents Program.

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