Electrosynthesis of highly transparent cobalt oxide water oxidation catalyst films from cobalt aminopolycarboxylate complexes.

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DOI: 10.1002/cssc.201403188

Electrosynthesis of Highly Transparent Cobalt Oxide Water Oxidation Catalyst Films from Cobalt Aminopolycarboxylate Complexes Shannon A. Bonke,[a, b] Mathias Wiechen,[a, b] Rosalie K. Hocking,[c] Xi-Ya Fang,[d] David W. Lupton,[a] Douglas R. MacFarlane,[a, b] and Leone Spiccia*[a, b] and morphology to enable the production of thin, catalytic CoOx films on a conductive substrate, which can be exploited in integrated photoelectrochemical devices. Notably, under a bias of 1.0 V (vs. Ag/AgCl), the film deposited from [Co(NTA)(OH2)2] (NTA = nitrilotriacetate) decreased the transmission by only 10 % at l = 500 nm, but still generated > 80 % of the water oxidation current produced by a [Co(OH2)6]2 + -derived oxide film whose transmission was only 40 % at l = 500 nm.

Efficient catalysis of water oxidation represents one of the major challenges en route to efficient sunlight-driven water splitting. Cobalt oxides (CoOx) have been widely investigated as water oxidation catalysts, although the incorporation of these materials into photoelectrochemical devices has been hindered by a lack of transparency. Herein, the electrosynthesis of transparent CoOx catalyst films is described by utilizing cobalt(II) aminopolycarboxylate complexes as precursors to the oxide. These complexes allow control over the deposition rate

Introduction Solar radiation striking the earth’s surface is both, renewable, and sufficient to satisfy human energy demands. However, without the development of suitable energy conversion and storage systems, it is unable to surpass fossil fuels as the world’s primary energy source.[1] Although there are challenges associated with the harvesting of solar energy,[2] significant issues related to its conversion and storage are equally prevalent.[1d,e, 3] The conversion of solar energy into a fuel that can be easily stored and utilized has the potential to address these challenges. One intensely discussed approach is to use solar energy to power electrolytic water splitting, thereby producing hydrogen as a fuel. Such an approach could, ideally, utilize sunlight directly in an integrated light-harvesting/water-splitting device, or indirectly with photovoltaics to capture solar energy and then generate electricity, which would power the process. The combustion of hydrogen as a fuel produces water as the only waste product to create a sustainable energy cycle.[1b,c–e, 3a]

The overall water-splitting reaction [Reaction (1)] is divided into two redox half-reactions: the oxidation of water to evolve oxygen [Reaction (2)] and the reduction of protons to evolve hydrogen [Reaction (3)]. The thermodynamic requirement for the overall process is 237 kJ per mole of hydrogen produced, which corresponds to an equilibrium potential difference of 1.23 V between two electrodes under standard conditions.[1c] However, the mechanistically complex oxygen evolution reaction, in particular, does not proceed readily near the equilibrium potential. Thus, effective catalysis is required to avoid significant energy losses. The effectiveness of a catalyst is often discussed in terms of the overpotential required for the reaction to proceed at an appreciable rate. It is defined as the difference between the actual applied potential and the thermodynamically required potential.[4] 2 H2 O Ð O 2 þ 2 H 2 DE 0 ¼ 1:23 V=474 kJ mol1

[a] S. A. Bonke, Dr. M. Wiechen, Prof. Dr. D. W. Lupton, Prof. Dr. D. R. MacFarlane, Prof. Dr. L. Spiccia School of Chemistry, Monash University Victoria, 3800 (Australia) E-mail: [email protected]

2 H 2 O Ð O 2 þ 4 H þ þ 4 e E 0 ¼ 1:23 V 2 H þ þ 2 e Ð H 2

ð1Þ

ð2Þ

ð3Þ

[b] S. A. Bonke, Dr. M. Wiechen, Prof. Dr. D. R. MacFarlane, Prof. Dr. L. Spiccia ARC Centre of Excellence for Electromaterials Science (ACES) Monash University, Victoria, 3800 (Australia)

E 0 ¼ 0:00 V

[c] Dr. R. K. Hocking Discipline of Chemistry, School of Pharmacy and Molecular Sciences James Cook University, Queensland, 4811 (Australia)

For practical water splitting to be realized, robust, cheap, and easily manufactured catalysts must be deployed.[1d,e, 3c, 4] Although there has been a great deal of research on catalysts based on precious metals (e.g., Ru or Ir),[5] in recent years water oxidation catalysts based on the oxides of more abundant transition metals, such as manganese (MnOx),[6] nickel (NiOx),[7] iron (Fe2O3),[8] and cobalt (CoOx),[6b, 9] has been devel-

[d] Dr. X.-Y. Fang Monash Centre for Electron Microscopy Monash University, Victoria, 3800 (Australia) Supporting Information for this article is available on the WWW under http://dx.doi.org/10.1002/cssc.201403188.

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Full Papers oped to address these requirements. Films of CoOx, MnOx, and NiOx can be readily formed as active water oxidation catalysts through oxidative electrodeposition from electrolytes containing simple salts of the respective transition-metal ions.[6f,g, 7a, 9a,e,l] For the cobalt oxide of interest herein, electrodeposition generally utilizes [Co(OH2)6]2 + in either phosphate (pH 7) or borate (pH 9.2) buffered solutions.[9] Systematic studies of the catalytic activity of CoOx catalyst films have shown a region of linear dependence between the amount of Co deposited and the catalytic current density.[9j,p] This direct proportionality indicates that catalytically active sites are accessible throughout the oxide material and not only at the interface between CoOx and the bulk electrolyte. However, this linear interdependence breaks down at higher overpotentials,[9j,p] at which mass transport, specifically proton transfer and abstraction, has been shown to become a limiting process.[9o,p] Nevertheless, increasing the CoOx layer thickness has been suggested as a means to improve the catalytic currents,[9p] and an ideal film thickness has been predicted.[9o] Other research has focused on functionalizing light harvesters with CoOx films to directly achieve the desired solar-to-fuel energy conversion.[10] For example, composite metal oxide structures with the CoOx water oxidation catalyst coupled to an a-Fe2O3 photoanode significantly lowers the bias required for water oxidation under illumination.[10a,c,d, 11] Solar water splitting has also been demonstrated by integrating the CoOx catalyst into an amorphous silicon-based photovoltaic device.[9m, 10g,h] However, the efficiency of both the a-Fe2O3 photoanodes and the silicon-based integrated device were reported to suffer from the limited transparency of the CoOx films.[10a,g] As mentioned previously, increasing the CoOx film thickness results in faster water oxidation,[9j,p] but the increasingly opaque CoOx film also limits the amount of solar energy that reaches the light harvester. Thus, to facilitate efficient and direct solar to fuel energy conversion, it is important to develop methods for the preparation of transparent, but highly active, CoOx water oxidation catalysts. Metal complexes have been used extensively as precursors in nanoparticle synthesis,[12] and the use of complexing agents, such as aminopolycarboxylates, is well established in electroplating/coating to alter the morphology of the final surface.[13] These complexing agents alter the deposition process, but are typically not incorporated into the film. Recent work to apply metal complexes as precursor compounds in the preparation of water oxidation catalysts has led to improved catalytic activity of both nickel[7c] and manganese[6c,k] oxide films. Initial work has also shown that cobaloximes can be effective as CoOx precursors,[14] which indicates that cobalt complexes can form an important strategy for the optimization of electrodeposited CoOx water oxidation catalysts. Building upon these previous findings, we herein report a systematic examination of the electrodeposition of CoOx films from mononuclear CoII complexes as a method for the preparation of highly active, yet highly transparent, water oxidation catalysts. Chelating bi- (glycine (GLY)), tri- (iminodiacetate (IDA)), tetra- (nitrilotriacetate (NTA)), and hexadentate (ethylenediaminotetraacetate (EDTA)) aminocarboxylate ligands

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Figure 1. Structures of [Co(IDA)2]2 (left), [Co(NTA)(OH2)2] (middle), and [Co(EDTA)]2 (right) used as precursors for the electrodeposition of CoOx films as water oxidation catalysts.

were selected as cobalt(II) coordinating agents (Figure 1). These ligands are inexpensive and easily sourced/produced, which is advantageous for the development of functional materials on a commercial scale. To enable comparison with the literature,[9a,j,m] CoOx films were deposited from [Co(OH2)6]2 + under the same conditions as those reported.

Results and Discussion Electrodeposition of CoOx Films The cobalt(II) aminopolycarboxylate complexes were prepared in situ prior to the electrodeposition of CoOx films. The UV/Vis spectra of [Co(OH2)6]2 + , [Co(IDA)2]2, [Co(NTA)(OH2)2] , and [Co(EDTA)]2 exhibited an absorption band between l = 400 and 600 nm with extinction coefficients (e) at lmax of 4, 8, 18, and 12 m1 cm1, respectively (Figure S1 in the Supporting Information); these values are consistent with literature data.[15] The spectra of [Co(IDA)2]2, [Co(NTA)(OH2)2] , and [Co(EDTA)]2 were unchanged in an aqueous solution of NaOH or 0.10 m borate buffer (both pH 9.2), which indicated that the complexes remained intact in these solutions. Salts of [Co(OH2)6]2 + in borate buffer were insoluble at concentrations required for UV/Vis spectrum acquisition. In the case of GLY, precipitates formed regardless of the ligand (L)/Co ratio used (range: 1–3) and, consequently, these complexes were excluded from further examination. To carry out the electrodeposition from [Co(IDA)2]2, [Co(NTA)(OH2)2] , [Co(EDTA)]2, and [Co(OH2)6]2 + , the in situ prepared solutions of the respective complex ([CoLx]) were diluted with borate buffer (Bi, pH 9.2) to give 1.0 mm [CoLx] solutions in 0.10 m Bi ([CoLx]/Bi). Additionally, the deposition of [Co(OH2)6]2 + was examined to facilitate comparison with the literature. Data for the deposition of [Co(OH2)6]2 + from 0.50 mm precursor solutions are reported herein and comparative data for a 1.0 mm solution of [Co(OH2)6]2 + (close to the solubility limit in 0.10 Bi) is given in Figure S2 in the Supporting Information. The oxidation of the CoII complexes from each [CoLx]/Bi solution was examined by cyclic voltammetry (CV). Each CV trace displayed an oxidative wave that increased in magnitude with subsequent cycles (Figure 2); this is in agreement with observations reported in the literature for the deposition of CoOx from a solution of [Co(OH2)6]2 + /Bi and is consistent with the 2


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Full Papers the CV results above, the deposition of CoOx at a lower potential to minimize water oxidation is not possible from the CoII IDA, NTA, and EDTA complexes, although it can be performed from [Co(OH2)6]2 + .[9j,p] By comparing the CV traces of the complexes with those of the deposited films, it can be seen that the onset of water oxidation from the deposited film occurs at a lower potential than that required for effective film deposition from the complexes (cf. Figure 2 with Figure 7 A, which is discussed later); this confirms that water oxidation will make up a significant fraction of the charge passed in the electrodeposition process.

General characterization To enable a detailed comparison of the CoOx films electrodeposited from the different CoII complexes, comprehensive characterization was carried out by means of a variety of techniques. The morphology of the CoOx films was examined by SEM (Figure 3). The blank substrate shows FTO crystals (Fig-

Figure 2. CV scans of [Co(OH2)6]2 + (0.50 mm, black), [Co(IDA)2]2 (1.0 mm, red), [Co(NTA)(OH2)2] (1.0 mm, green), and [Co(EDTA)]2 (1.0 mm, blue) in 0.10 m Bi buffer solutions. Subsequent cycles generated greater current densities.

deposition of an electrocatalytic species on an electrode.[9e] Increasing the denticity of the ligands resulted in the onset of electrodeposition occurring at sequentially higher potentials. Oxidative and reduction peaks can be seen for [Co(OH2)6]2 + between 500 and 900 mV (Figure 2). As the denticity of the ligands in the complex increased, this process was observed at lower intensities. This is likely to be due to less material being deposited on the electrode. It is comparable to the peaks observed during electrochemical characterization of the films (Figure 7 A, below). The CoOx films were electrodeposited by chronoamperometry (CA; the term controlled potential electrolysis (CPE) is also used in the literature) of each [CoLx]/Bi solution at 1.4 V versus Ag/AgCl. This deposition potential was selected because it enabled deposition from each complex under identical conditions and at suitable rates. CA was performed by using a freshly cleaned fluorine-doped tin oxide (FTO) electrode with an electroactive area of 4 mm by 4 mm (see the Experimental Section for details). To facilitate comparison between the CoOx films, each deposition was allowed to proceed until a maximum current density was reached (Figure S3 in the Supporting Information). In all cases, an initially increasing current density could be observed; this approached an asymptotic limit, which mirrored that reported in the literature.[9e] The interval after which the current density ceased to increase was strongly dependent on the respective precursor complex used for the deposition. In the case of [Co(OH2)6]2 + /Bi, a rapidly increasing current density was observed, which peaked after 4 min and for [Co(IDA)2]2/Bi after 10 min. For the [Co(NTA)(OH2)2]/Bi and [Co(EDTA)]2/Bi systems, longer intervals of 60 and 120 min, respectively, were required to reach peak currents. A comparison of the current density during the deposition process is not a useful indication of catalytic activity because both water and cobalt oxidation processes take place simultaneously at the anode at 1.4 V versus Ag/AgCl. Thus, it is not possible to electrochemically determine the current arising from each of these processes. As shown by the discussion of ChemSusChem 0000, 00, 0 – 0

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Figure 3. Scanning electron micrographs of blank FTO (A) and CoOx films from [Co(OH2)6]2 + (B), [Co(IDA)2]2 (C), and [Co(NTA)(OH2)2] (D).

ure 3 A), whereas the film electrodeposited from [Co(OH2)6]2 + is evident in the form of a forest of tall nodules (Figure 3 B). The film does not appear to have fully coated the substrate surface because FTO crystals remain visible between the CoOx nodules. In contrast to the larger (1–5 mm) particles reported in the literature,[9a,e] the shorter deposition time used in this study has resulted in smaller nodules. The CoOx particles found here are in the size range of about 100–250 nm and they do not appear to connect separate FTO crystals. The [Co(IDA)2]2 electrodeposit has formed as a thin layer that conformally coats the substrate surface and grows between the FTO crystals (Figure 3 C). No nodule-like particles in the style of the [Co(OH2)6]2 + -derived CoOx film were observed. This was the case even when a sixfold longer (60 min) deposition time was used to deposit a significantly thicker film (Figure S4 in the Supporting Information). However, the possibility that increasing the deposition time may result in the formation of nodules as the coating becomes 3


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Full Papers increasingly thick cannot be totally excluded. The [Co(NTA)(OH2)2]-derived CoOx electrode did not show clear indications that oxide material was present on the surface. Only at high magnification are small sections of a deposited CoOx film visible (Figure 3 D). This is most apparent where the CoOx film bridges separate FTO crystals (see Figure S5 in the Supporting Information for a higher magnification SEM image). By comparing Figure 3 D to the blank substrate shown in Figure 3 A, it is apparent that there may also be a CoOx coating present on the substrate surface. No evidence of CoOx could be found in the case of the deposition from [Co(EDTA)]2 (Figure S6 in the Supporting Information). Preliminary analysis of the amount of deposited CoOx was carried out by using point and large-area energy-dispersive Xray (EDX) analysis (see Figure S7 in the Supporting Information). This indicated the presence of Co on the surface for the films deposited from [Co(OH2)6]2 + , [Co(IDA)2]2, and [Co(NTA)(OH2)2] . In the case of CoOx derived from [Co(EDTA)]2, no increase in the amount of Co relative to blank FTO was found. A comparison of the ratio of EDX counts from Sn (from FTO substrate) to Co indicated a correlation between the deposited amount of Co and the CoOx precursor in the order of [Co(OH2)6]2 + > [Co(IDA)2]2 > [Co(NTA)(OH2)2] (Table S1 in the Supporting Information). Inductively coupled plasma mass spectrometry (ICP-MS) was used to more accurately determine the amount of cobalt present in the deposited films (Table 1

CoOx deposited was negligible. Because this film also showed little improvement in current above the FTO baseline, it was excluded from further study. For [Co(OH2)6]2 + , the CoOx films electrodeposited from 0.5 mm versus 1 mm solutions were analyzed by ICP-MS and electrochemically. The greater cobalt content of the film from the 1 mm solution (Table S2 in the Supporting Information) does not result in significant improvements in catalytic activity (Figure S19—S21 in the Supporting Information), and a much lower average catalytic activity per cobalt center (in mA mmol1) was determined for the film from the more concentrated solution (Tables S2 and S3 in the Supporting Information). Significant interest in coupling CoOx water oxidation catalysts with light harvesters to enable direct light-driven water splitting in a photoelectrochemical cell (PEC)[10] led us to investigate the transparency of the prepared CoOx films. The transparency measurements were performed by utilizing in situ UV/ Vis spectrophotometry, with and without an applied potential. The spectroelectrochemical data revealed notable differences in the transparency of the films derived from different precursors (Table 1, Figure 4, and Figure S8 in the Supporting Infor-

Table 1. ICP-MS determination of the amount of cobalt (mmol cm2 1 standard deviation) deposited onto FTO from each precursor and the transparency of each film at l = 500 nm calculated from the UV/Vis spectra (see Figure 4 and Figure S9 in the Supporting Information). Precursor

Co deposited[a] [mmol cm2]

Transmission at l = 500 nm[c] [%]

[Co(OH2)6]2 + [Co(IDA)2]2 [Co(NTA)(OH2)2] [Co(EDTA)]2 FTO substrate

0.06 0.01 0.04 0.01 0.019 0.008 0.00[b] (used as blank)[a]

57 71 72 75 78

(40) (57) (68) (75) (78)

Figure 4. UV/Vis transmission spectra of CoOx films deposited from [Co(OH2)6]2 + (black) and [Co(NTA)(OH2)2] (green) measured by using in situ spectroelectrochemistry, along with the spectrum of the FTO substrate (gray). Spectra were collected at open-circuit voltage (Eoc, dash–dot lines), as well as 650 (dashed lines) and 1000 mV bias versus Ag/AgCl (solid lines).

[a] CoOx film deposition and analysis was performed in triplicate for each precursor. The FTO blank was subjected to the same treatment as the slides used for CoOx deposition, excluding the deposition step. The background amount of Co found for the FTO blank ((0.054 0.004) mmol cm2) was subtracted from the other Co values given in the table. [b] A slightly negative value was obtained after subtraction of the FTO blank, which implied that no or a negligible amount of Co was deposited. [c] The transparency was determined from UV/Vis spectra of the prepared CoOx films for l = 500 nm at Eoc and 1.0 V versus Ag/AgCl. Values in parentheses are those with a bias of 1.0 V. Errors associated with transparency measurements are detailed in the Experimental Section.

mation). The spectra measured with and without applied potential display the same transparency dependence on the denticity of the ligands, that is, increasing denticity leads to more transparent films in the order of Co(OH2)6]2 + < [Co(IDA)2]2 < [Co(NTA)(OH2)2] . Indeed, the films made by using the two complexes only decreased the total transmission by 6–7 % at l = 500 nm relative to the FTO control. The CoOx films exhibited an electrochromic effect, that is, the application of a potential decreased their transparency. The magnitude of this effect is also correlated with the initial (open-circuit) transparency and, hence, with the amount of CoOx deposited in each case (Table 1). A comparison of the open-circuit transmission with that measured at 1.0 V (vs. Ag/AgCl) for each film revealed that, at l = 500 nm, the highly transparent [Co(NTA)(OH2)2]-de-

and Table S2 in the Supporting Information). The results exhibit a similar trend to those from EDX analysis; the sample deposited from [Co(OH2)6]2 + has the highest amount of Co per unit area and smaller amounts are present for those deposited from [Co(IDA)2]2 and [Co(NTA)(OH2)2] . For the film made from [Co(EDTA)]2, the ICP-MS data was within the error of the FTO baseline value (Table 1), which indicated that the amount of

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Full Papers rived CoOx film only showed a decrease of about 4 % to 68 % transmission, whereas the more opaque film derived from [Co(OH2)6]2 + decreased by 17 % to 40 %. It is important to note here that the films contain different amounts of CoOx and, therefore, the results should not necessarily be taken to suggest a difference in the molar absorptivity of the deposited material. In summary, analysis of the CoOx films prepared from [Co(OH2)6]2 + , [Co(IDA)2]2, [Co(NTA)(OH2)2] , and [Co(EDTA)]2 clearly shows that Co has been deposited in decreasing amounts according to the increasing denticity of the ligands in the precursor. Connected to the lower amount of cobalt, the transparency of the films increases with increasing ligand denticity. Furthermore, the morphology of the cobalt oxide films formed on the substrate can be controlled through selection of the precursors. It should also be noted, however, that the film morphology and transparency may not be independent because smoother surfaces are likely to reduce scattering effects.

the influence of the buffer solution on the disorder of the CoOx phase,[4, 9b,g,l,m] the possibility exists that different precursors influence the structure. The XANES data (Figure S9 in the Supporting Information), measured ex situ, were shifted to slightly lower energies relative to previous in situ measurements.[9g] A possible explanation for this observation is that the films deposited in this study may be more prone to reduction. The thin nature of the CoOx films prepared from the cobalt complexes also had an effect on the EXAFS resolution because the FTO coating on the glass substrate affects the fluorescence intensity measurements. The lack of long-range order in these materials excludes XRD as a characterization technique (Figure S10 in the Supporting Information for Fourier-transformed (FT) EXAFS).[4, 6c, 16]

Electrochemical testing To identify suitable testing conditions, initial experiments were carried out in 0.10 and 0.60 m Bi with CoOx films deposited from [Co(NTA)(OH2)2] . Furthermore, Na2SO4 was added to the solutions of Bi to allow effects from the increased buffering capacity to be separated from ionic strength changes. The CV results revealed that the moderate activity of the CoOx films in 0.10 m Bi could be substantially improved by increasing the borate concentration to 0.60 m (cf. black and green traces in Figure 6). The addition of 0.32 m Na2SO4 to the 0.10 m Bi buffer

Characterization by X-ray absorption spectroscopy Detailed structural analysis of the CoOx films is complicated due to the low amount of Co deposited, the poor crystallinity of the formed oxides, and the high crystallinity of the FTO substrate. Thus, X-ray absorption spectroscopy (XAS) was performed to gain an insight into the structure of the electrodeposited catalysts. Both the X-ray absorption near-edge structure (XANES; Figure S9 in the Supporting Information) and extended X-ray absorption fine structure (EXAFS; Figure 5) meas-

Figure 5. EXAFS data for electrodeposited and reference cobalt oxides (CoOx). Spectra are shown for two reference cobalt oxide phases, the cobalt(II,III) oxide spinel phase (Co3O4, blue) and heterogenite (CoO(OH), gray), and for CoOx prepared from [Co(OH2)6]2 + (black), [Co(IDA)2]2 (red), and [Co(NTA)(OH2)2] (green).

Figure 6. CV results for CoOx deposited onto FTO recorded in Bi electrolytes of different concentrations and ionic strengths. The 0.10 (black) and 0.60 m (green) Bi buffers were used to identify suitable conditions for catalyst testing. The buffer ionic strength was adjusted to 1.0 m by adding 0.32 (red) and 0.23 m (blue) Na2SO4 to the 0.10 and 0.60 m Bi buffers, respectively. A scan rate of 20 mV s1 was used with stirring at 650 rpm and the electrolyte temperature maintained at 35 8C. The same trends were observed for CoOx films deposited from each precursor with data from [Co(NTA)(OH2)2]-derived CoOx shown herein.

urements indicate that a similar material is formed from each precursor. In all cases, the material closely resembles the heterogenite (CoO(OH)) cobalt oxide phase; this is in agreement with previous results on CoOx catalysts prepared from [Co(OH2)6]2 + /Bi solutions.[9b,g,l] Thus, a cobalt(III) oxide is formed regardless of the precursor used and different cobalt(II) aminocarboxylate complexes do not appear to influence the structure (phase) of the CoOx formed. However, differences in the structural disorder of the materials on a scale below the sensitivity of the technique cannot be excluded. By considering previous results reported in the literature regarding, for instance, ChemSusChem 0000, 00, 0 – 0

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increases the ionic strength of the electrolyte to about 1 and greatly increases the measured current (cf. black and red traces in Figure 6). The addition of 0.23 m Na2SO4 to the 0.60 m Bi solution also adjusts the ionic strength of this buffer to about 1 M (cf. green and blue trace in Figure 6). A comparison of the two curves for electrolytes with an ionic strength of about 1 M 5


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Full Papers (cf. red and blue traces in Figure 6) indicates that below 10 mA cm2 there is little effect of using a higher buffer concentration. The increased buffering capacity of the electrolyte becomes a factor only above this threshold, at which the lower buffer concentration appears to be limiting the current (red trace, Figure 6). This testing was performed with CoOx from each of the precursors; the trends are the same in each case. Given that a stable current was obtained with 0.60 m Bi + 0.23 m Na2SO4 electrolyte, it was selected for the detailed studies discussed below. Furthermore, a constant temperature of 35 8C was selected for testing to ensure that the electrolyte remained in solution; marginally higher currents were measured at this temperature compared with 25 8C (Figure S11 in the Supporting Information). The literature does not define a standard scan speed when analyzing/approximating catalyst activity by potentiodynamic measurements (e.g., CV or linear scan voltammetry (LSV)). To investigate the effect of the scan rate on the observed current, measurements were recorded at different scan rates (Figure S12 in the Supporting Information). No differences were observed in the catalytic currents produced by the films, but variations in the magnitude of precatalytic features were evi-

dent at different scan speeds. The faster scan speed of 20 mV s1 revealed the oxidative activation of the film much more clearly than the 1 mV s1 scan rate (Figure S12, inset, in the Supporting Information). For this reason, the CV data shown herein were produced at a scan rate of 20 mV s1. At this rate, inspection of the reserve CV scan (i.e., from 1.4 to 0 V) indicates that the water oxidation current is independent of film oxidation. Analysis of all films was completed at both 1 and 20 mV s1; results at the former scan rate are given in Figure S13 in the Supporting Information). The CV analysis of the films in the optimized electrolyte indicated two oxidative processes (Figure 7 A). The first is an oxidation peak observed at about 500–800 mV. In this range, the [Co(OH2)6]2 + -derived CoOx film gave rise to the highest current, followed by that from [Co(IDA)2]2, whereas an almost negligible current was displayed by the [Co(NTA)(OH2)2]-derived film (Figure 7 A, inset). It has been shown by electron paramagnetic resonance (EPR) and XAS that this first anodic process is related to the oxidation of cobalt(III) to cobalt(IV) within the film.[9f–h,p] The magnitude of this oxidative peak has also been shown to correlate with the amount of cobalt present within the film.[9p] Accordingly, the different currents observed herein

Figure 7. A) CV results for CoOx derived from [Co(OH2)6]2 + (black), [Co(IDA)2]2 (red), [Co(NTA)(OH2)2] (green), and [Co(EDTA)]2 (blue) precursors on FTO substrates and an FTO blank (gray) recorded at a scan rate of 20 mV s1 in 0.60 m Bi with 0.23 m Na2SO4, pH 9.2, stirring at 650 rpm and T = 35 8C. The thermodynamic requirement for water oxidation under these conditions is 470 mV versus Ag/AgCl. B) The current from each film normalized to the amount of cobalt deposited per cm2 of the FTO substrate, as determined by ICP-MS. C) CA of [Co(OH2)6]2 + (black), [Co(IDA)2]2 (red), and [Co(NTA)(OH2)2] (green) CoOx films on FTO carried out at 1.0 V versus Ag/AgCl in 0.60 m Bi, pH 9.2, with 0.23 m Na2SO4, stirring at 650 rpm and T = 35 8C. D) Current densities normalized to the amount of cobalt present.

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Full Papers correspond well with the amounts of cobalt deposited from each precursor, which is consistent with the results obtained by ICP-MS measurements (Table 1). A second redox process was identified at potentials > 800 mV that was consistent with water oxidation (Figure 7 A).[9a] The [Co(IDA)2]2- and [Co(NTA)(OH2)2]-derived CoOx films produce a lower current density than that of the [Co(OH2)6]2 + -derived CoOx (as also confirmed by the steadystate current experiments described below).

a 24 h period under the conditions used for the CA measurements detailed above (Figure S16 in the Supporting Information). The current obtained was stable for the experiment duration, which indicated that these thin oxide films were extremely stable during catalytic testing. Referring to previous studies, cobalt oxide films derived from [Co(OH2)6]2 + have not shown decomposition over 12 h of measurements.[9a] Furthermore, with appropriate engineering controls (a flow-through electrochemical cell with temperature-controlled electrolyte), the catalyst was shown to be stable for 90–100 h of operation.[9j] The 24 h test did appear to result in some film morphology changes that were visible by SEM (cf. Figure 3 D and Figure S17 in the Supporting Information); however, this did not greatly affect the film transparency (Figure S18 in the Supporting Information). The steady-state currents of 4.0 and 3.8 mA cm2, for the films made from [Co(IDA)2]2 and [Co(NTA)(OH2)2] , respectively, are only 13–20 % smaller than those measured for the CoOx films from [Co(OH2)6]2 + , despite the smaller amount of CoOx involved. Thus, the current densities normalized to the amount of Co deposited per cm2 (from ICP-MS data) plotted in Figure 7 B clearly show that both the [Co(IDA)2]2- and [Co(NTA)(OH2)2]-derived CoOx films exhibit a much higher relative catalytic activity. As also noted earlier, the CA and CV (or LSV) measurements showed that the CoOx films deposited from 0.5 mm solutions of [Co(OH2)6]2 + exhibited much higher activity per amount of cobalt deposited than those deposited from 1.0 mm solutions (Figures S19–S21 and Tables S2 and S3 in the Supporting Information).

Tafel region analysis A detailed comparison of the electrocatalytic activity over a range of current densities can be obtained through a series of CA measurements. A Tafel region of the data (i.e., linear dependence of overpotential vs. log current density) can then be identified. Analysis of the different CoOx films revealed that a Tafel region existed for each film in the 800–900 mV range versus Ag/AgCl (Figure S14 in the Supporting Information). Analysis of the data yielded Tafel slopes of (61 3) mV per decade at 35 8C in all cases. More detail is provided in the Supporting Information, along with the Tafel plots (Tables S4 and S5 and Figure S14). This is in agreement with previous reports of 59 mV (25 8C) for [Co(OH2)6]2 + CoOx films.[9a,h,j] The data used for comparison with results reported in the literature (Tables S6 and S7 in the Supporting Information) and comparative tables and plots (Table S8 and Figure S15 in the Supporting Information) are also provided.

CA analysis

Conclusions

Steady-state CA measurements were performed on each CoOx film at 1.0 V versus Ag/AgCl in the optimized electrolyte (0.60 m Bi + Na2SO4) for 60 min (Figure 7 C) to examine whether the activity of the catalyst varied with time. The applied potential of 1.0 V versus Ag/AgCl corresponds to an overpotential of 530 mV. A comparison of the current densities at 60 min (Table 2) again highlights the higher relative activity of the films produced from some of the Co complexes (Figure 7 D and Table 2). Importantly, this measurement also provided a preliminary indication of film stability during catalytic testing. In addition, the thinnest film produced in this study, the [Co(NTA)(OH2)2] CoOx film was tested at 1.0 V versus Ag/AgCl over

Catalytically active cobalt oxide films were deposited from simple and inexpensive aminopolycarboxylate complexes, specifically [Co(IDA)2]2 and [Co(NTA)(OH2)2] . The use of these complexes facilitated the formation of thin films. Under the same conditions, higher denticity ligands resulted in lower amounts of cobalt being deposited. XAS analysis of the films showed that the heterogenite (CoO(OH)) phase was formed in each case. SEM examination of the film morphology indicated the formation of films that conformally coated the substrate rather than the more particle-like material formed under the same conditions from [Co(OH2)6]2 + . The thin films formed from the cobalt complexes were highly transparent, but retained a high catalytic activity, which made them suitable for direct coupling to light-harvesting materials. The use of precursor complexes could be particularly useful in the assembly of photoelectrochemical devices, in which it should be possible to tailor the photoelectrodeposition of transparent catalyst films to a specific light harvester/sensitizer by selecting a precursor that deposited in a controlled manner under irradiation and/or application of a suitable potential bias. The formation of highly transparent cobalt catalyst films addressed transparency issues previously encountered in the literature with triple-junction silicon cells and hematite photoanodes.[10a,g] The application of this methodology to suitable photoanodes can be identified as a promising future research objective. Such work will assist in

Table 2. Steady-state current densities for the CoOx films measured after 60 min at 1.0 V versus Ag/AgCl in an aqueous electrolyte containing 0.6 m Bi and 0.23 m Na2SO4 (pH 9.2), stirring at 650 rpm and T = 35 8C. Precursor

Current density[a] [mA cm2]

Relative current[b] [mA mmol1]

[Co(OH2)6]2 + [Co(IDA)2]2 [Co(NTA)(OH2)2]

4.56 0.06 4.05 0.04 3.8 0.2

76 13 101 25 200 85

[a] Values are the average of two measurements. [b] Relative currents (mA mmol1) were calculated by using the current density values (mA cm2) in this table, and the ICP-MS results (mmol cm2 ; Table 1). The error displayed is 1 standard deviation.


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Full Papers ber, the counter electrode was immersed in 0.10 m Bi (20 mL; pH 9.2). In the working electrode compartment, a FTO-coated glass electrode, with 0.16 cm2 electroactive area, was immersed in a 1.0 mm solution of CoLx (20 mL) containing 0.10 m Bi (pH 9.2). The electroactive area of the FTO-coated glass was defined by laser engraving the conductive layer on three sides of the desired area and by using polyimide (Kapton) tape on the fourth side to allow an electrical connection. A glass-bodied Ag/AgCl RE was immersed in a glass chamber (RE chamber) containing a 0.10 m solution of Bi (2 mL) with a P4 ceramic frit at the bottom. This RE chamber was then immersed in the working electrode compartment. The RE chamber provided an additional separating barrier to prevent the RE from being contaminated with the cobalt solution. This RE chamber was used when the cobalt salts or complexes were present in solution, but not for catalyst film testing. The potential drop across this additional frit was < 1 mV. Deposition times and applied potential differences were those specified throughout the study.

advancing the development of photoelectrochemical watersplitting devices that use only abundant elements.

Experimental Section Materials Reagent- or analytical-grade chemicals were used as received from commercial suppliers. FTO (F:SnO2, Dyesol TEC8 Glass Plates) coated glass was used with a sheet resistance of 8 W/sq. Reverseosmosis purified water with a quoted resistivity of 1 MW cm1 at 25 8C was used in the preparation of all electrolyte solutions.

Preparation of electrolytes Borate buffer solutions (Bi) were prepared by first dissolving boric acid in water, followed by the adjustment of the solution pH to 9.2 with a freshly prepared solution of NaOH. The pH of the stirred solution was monitored with a glass pH electrode. After pH adjustment, the solution was made up to the desired volume to reach the desired borate concentration. Any addition of Na2SO4 was performed prior to the pH adjustment.

CV and CA analyses CV analysis of electrodeposited CoOx films was performed at 1 and 20 mV s1 by using a two-compartment electrochemical cell, equivalent to that used for film deposition, with both compartments charged with cobalt-free buffered electrolyte (20 mL). The CoOxcoated FTO glass and Ag/AgCl RE were placed within the working compartment, whereas the Ti auxiliary electrode was immersed in the counter electrode compartment. These conditions were maintained for CA (steady-state) measurements. Tafel slopes were generated by measuring the current at 25 mV applied potential steps within the region of exponentially increasing oxidative current, as determined by CV. Each potential was applied for 10 min. The potential steps were applied in the descending (1000–675 mV) then ascending (675–1000 mV) directions. The solution temperature was controlled with the use of a temperature-controlled bath connected to a custom jacketed electrochemical cell. Unless otherwise indicated, all water oxidation experiments were carried out at 35 1 8C. For the 24 h CA measurement, the soldered electrode connection (FTO to wire) was covered with a hot-melt adhesive to insulate the connection. This prevented any electrical short-circuits due to the evaporation and condensation of water over time.

Formation of cobalt(II) aminocarboxylate complexes All cobalt complexes were formed in situ by adding aqueous solutions of ligand (10 mL of 50 mm for NTA and EDTA; 100 mm for IDA) to aqueous solutions of CoSO4 (10 mL, 50 mm) in molar proportions corresponding to the Co/ligand stoichiometry of the final complex. Adjustment of volume to 100 mL with H2O, and the pH to 9.2 with an aqueous solution of NaOH afforded a solution with a 10 mm complex concentration for use in electrodeposition. The pH was monitored with a glass pH electrode. The formation of the cobalt(II) complexes was confirmed by means of UV/Vis spectrophotometry (Figure S1 in the Supporting Information). Immediately prior to deposition, the solutions were diluted into solutions of Bi, pH 9.2, to form 1.0 mm CoLx in a 0.10 m solution of Bi (CoLx/Bi).

Electrochemical methods All electrochemical experiments were performed with a Bio-Logic VSP potentiostat and BASi Ag/AgCl (3 m NaCl) reference electrode (RE) under standard laboratory conditions, unless otherwise specified. The potential of the RE was checked against a master electrode periodically throughout the study. The variation of the reference was < 15 mV. The solution was stirred at 650 rpm with a 15 mm magnetic bar during all measurements and no compensation for iR drop was made in any of the reported results. The separation between the working and REs was maintained at (8 0.5) mm. Coiled titanium wire was used as the counter electrode and FTO-coated glass for the working electrode was used for all experiments. The FTO slides were cleaned prior to use by sonication for 20 min in surfactant (Hellmanex, Sigma–Aldrich), 20 min in water, followed by a further 20 min in ethanol (96 %). The slides were then cut to size, sonicated once more in ethanol for 5 min, and dried with either compressed air or at 110 8C.

UV/Vis spectrophotometry Solution UV/Vis spectra of the complexes were measured on a Varian 300 UV/Vis spectrophotometer by using 1 cm cuvettes. The cobalt(II) complexes were diluted in either aqueous NaOH (pH 9.2) or Bi (pH 9.2) media for analysis. The cuvettes were then left in air for one week, after which time the spectra were remeasured. The UV/Vis spectra of the solid cobalt oxide films were measured on a PerkinElmer LAMBDA 950 UV/Vis/near-IR spectrophotometer. The electrodes were mounted in 1 cm cuvettes containing 0.1 m Bi. For measurements with in situ electrochemistry, the RE and counter electrode were placed above the electroactive area to avoid blocking the beam. The incident beam was allowed to pass through the back-side of the electrode (i.e., glass, FTO, then the cobalt oxide surface). A potential was applied and spectra were collected once the current stabilized. Collection was performed in triplicate to ensure that the film was no longer changing under the electrochemical bias. The error in transparency for three repeat measurements on the same film was 1—2 % (one standard deviation), whereas the error for different samples was up to 6 % of the absolute transmission due to variation in the lamp intensity.

Electrodeposition of catalyst films All electrochemical depositions were performed at room temperature in a two-compartment electrochemical cell fitted with a P4 ceramic frit to separate the two compartments. In the auxiliary cham-

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ria, Australia. We also acknowledge the operational support of the High Energy Accelerator Research Organisation (KEK) in Tsukuba, Japan, and access to the Australian National Beamline Facility. We thank Prof. A. M. Bond for helpful discussions.

Cobalt oxide films were dissolved in 70 % HNO3 (10 mL) with the solution subjected to ultrasonication and allowed to rest for > 12 h to ensure complete dissolution of the film. These solutions were then diluted (1:20) and analyzed by using an Optimass 9500 ICPTOF MS instrument. Raw count rates from the analyses were externally standardized by means of a calibration curve constructed by using a commercially available stock solution of cobalt chloride. Instrumental drift was monitored and corrected through the use of scandium ions as an internal standard. The error, shown in Table 1 as 1 standard deviation in cobalt content, was determined by performing triplicate film depositions and subsequent cobalt analysis.

Keywords: cobalt · electrochemistry · energy conversion · heterogeneous catalysis · water splitting [1] a) G. Ciamician, Science 1912, 36, 385 – 394; b) M. Grtzel, Acc. Chem. Res. 1981, 14, 376 – 384; c) A. J. Bard, M. A. Fox, Acc. Chem. Res. 1995, 28, 141 – 145; d) N. S. Lewis, D. G. Nocera, Proc. Natl. Acad. Sci. USA 2006, 103, 15729 – 15735; e) N. Armaroli, V. Balzani, Angew. Chem. Int. Ed. 2007, 46, 52 – 66; Angew. Chem. 2007, 119, 52 – 67; f) W. Lubitz, E. J. Reijerse, J. Messinger, Energy Environ. Sci. 2008, 1, 15 – 31; g) J. Barber, Chem. Soc. Rev. 2009, 38, 185 – 196; h) IPCC, Special Report on Renewable Energy Sources and Climate Change Mitigation, Cambridge University Press, Cambridge, 2011; i) J. Messinger, Aust. J. Chem. 2012, 65, 573 – 576; j) S. Styring, Faraday Discuss. 2012, 155, 357 – 376; k) T. A. Faunce, W. Lubitz, A. W. Rutherford, D. MacFarlane, G. F. Moore, P. D. Yang, D. G. Nocera, T. A. Moore, D. H. Gregory, S. Fukuzumi, K. B. Yoon, F. A. Armstrong, M. R. Wasielewski, S. Styring, Energy Environ. Sci. 2013, 6, 695 – 698. [2] a) A. Hagfeldt, M. Grtzel, Acc. Chem. Res. 2000, 33, 269 – 277; b) M. Grtzel, Acc. Chem. Res. 2009, 42, 1788 – 1798; c) A. Hagfeldt, G. Boschloo, L. C. Sun, L. Kloo, H. Pettersson, Chem. Rev. 2010, 110, 6595 – 6663; d) P. Docampo, S. Guldin, T. Leijtens, N. K. Noel, U. Steiner, H. J. Snaith, Adv. Mater. 2014, 26, 4013 – 4030. [3] a) T. R. Cook, D. K. Dogutan, S. Y. Reece, Y. Surendranath, T. S. Teets, D. G. Nocera, Chem. Rev. 2010, 110, 6474 – 6502; b) N. S. Lewis, Science 2007, 315, 798 – 801; c) R. E. Blankenship, D. M. Tiede, J. Barber, G. W. Brudvig, G. Fleming, M. Ghirardi, M. R. Gunner, W. Junge, D. M. Kramer, A. Melis, T. A. Moore, C. C. Moser, D. G. Nocera, A. J. Nozik, D. R. Ort, W. W. Parson, R. C. Prince, R. T. Sayre, Science 2011, 332, 805 – 809. [4] H. Dau, C. Limberg, T. Reier, M. Risch, S. Roggan, P. Strasser, ChemCatChem 2010, 2, 724 – 761. [5] a) S. W. Gersten, G. J. Samuels, T. J. Meyer, J. Am. Chem. Soc. 1982, 104, 4029 – 4030; b) A. Harriman, I. J. Pickering, J. M. Thomas, P. A. Christensen, J. Chem. Soc. Faraday Trans. 1 1988, 84, 2795 – 2806; c) T. J. Meyer, Acc. Chem. Res. 1989, 22, 163 – 170; d) M. Hara, J. T. Lean, T. E. Mallouk, Chem. Mater. 2001, 13, 4668 – 4675; e) C. Sens, I. Romero, M. Rodriguez, A. Llobet, T. Parella, J. Benet-Buchholz, J. Am. Chem. Soc. 2004, 126, 7798 – 7799; f) P. G. Hoertz, Y.-I. Kim, W. J. Youngblood, T. E. Mallouk, J. Phys. Chem. B 2007, 111, 6845 – 6856; g) J. J. Concepcion, J. W. Jurss, M. K. Brennaman, P. G. Hoertz, A. O. T. Patrocinio, N. Y. Murakami Iha, J. L. Templeton, T. J. Meyer, Acc. Chem. Res. 2009, 42, 1954 – 1965; h) R. Brimblecombe, G. C. Dismukes, G. F. Swiegers, L. Spiccia, Dalton Trans. 2009, 9374 – 9384; i) W. J. Youngblood, S. H. A. Lee, Y. Kobayashi, E. A. Hernandez-Pagan, P. G. Hoertz, T. A. Moore, A. L. Moore, D. Gust, T. E. Mallouk, J. Am. Chem. Soc. 2009, 131, 926 – 927; j) X. Sala, I. Romero, M. Rodriguez, L. Escriche, A. Llobet, Angew. Chem. Int. Ed. 2009, 48, 2842 – 2852; Angew. Chem. 2009, 121, 2882 – 2893; k) L. Duan, F. Bozoglian, S. Mandal, B. Stewart, T. Privalov, A. Llobet, L. Sun, Nat. Chem. 2012, 4, 418 – 423; l) Y. Gao, X. Ding, J. H. Liu, L. Wang, Z. K. Lu, L. Li, L. C. Sun, J. Am. Chem. Soc. 2013, 135, 4219 – 4222; m) K. S. Joya, Y. F. Joya, K. Ocakoglu, R. van de Krol, Angew. Chem. Int. Ed. 2013, 52, 10426 – 10437; Angew. Chem. 2013, 125, 10618 – 10630. [6] a) M. M. Najafpour, T. Ehrenberg, M. Wiechen, P. Kurz, Angew. Chem. Int. Ed. 2010, 49, 2233 – 2237; Angew. Chem. 2010, 122, 2281 – 2285; b) F. Jiao, H. Frei, Angew. Chem. Int. Ed. 2009, 48, 1841 – 1844; Angew. Chem. 2009, 121, 1873 – 1876; c) R. K. Hocking, R. Brimblecombe, L. Y. Chang, A. Singh, M. H. Cheah, C. Glover, W. H. Casey, L. Spiccia, Nat. Chem. 2011, 3, 461 – 466; d) I. Zaharieva, M. M. Najafpour, M. Wiechen, M. Haumann, P. Kurz, H. Dau, Energy Environ. Sci. 2011, 4, 2400 – 2408; e) M. Wiechen, I. Zaharieva, H. Dau, P. Kurz, Chem. Sci. 2012, 3, 2330 – 2339; f) I. Zaharieva, P. Chernev, M. Risch, K. Klingan, M. Kohlhoff, A. Fischer, H. Dau, Energy Environ. Sci. 2012, 5, 7081 – 7089; g) F. L. Zhou, A. Izgorodin, R. K. Hocking, L. Spiccia, D. R. MacFarlane, Adv. Energy Mater. 2012, 2, 1013 – 1021; h) M. Wiechen, H. M. Berends, P. Kurz, Dalton Trans. 2012,

SEM measurements Scanning electron micrographs were recorded on a FEI Nova NanoSEM 450 FEGSEM instrument. This instrument was equipped with a Bruker QUANTAX X-Flash silicon drift detector (SDD) for EDX analysis. Samples were deposited on FTO glass, rinsed with water, and allowed to dry before sputter-coating with platinum ( 1 nm Pt layer) to increase conductivity prior to imaging.

XAS measurements The Co K-edge XAS spectra were recorded on the multipole wiggler XAS beam-line 12 ID at the Australian Synchrotron, in operational mode 1, or at the Australian National Beamline Facility (ANBF; beam-line 20B at the Photon Factory, Tsukuba).[17] The beam energy was 3.0 GeV and the maximum beam current was 400 mA. The energy scale was calibrated by using a Co foil as an internal standard, for which the calibration energy of 7709.5 eV corresponds to the first inflection point of the foil. The average program[18] was used to average raw data files and exclude channels. Fluorescent-mode XAS data were taken directly on the electrodes at room temperature by placing the electrodes at 458 to the incoming beam. Solid-state reference samples were prepared by mixing with boron nitride and ground to a uniform fine powder, which was then loaded into 1 mm thick Al spacers and sealed with 63.5 mm Kapton tape windows. The program PySpline[19] was used to background subtract data, spline the data, and calculate the Fourier transform of the EXAFS. An E0 value of 7725.0 eV was used when defining the k range of the EXAFS. A four-segment spline of order 2 was fitted to the EXAFS region, and all data were normalized to the edge jump at 7770 eV.

Acknowledgements We are grateful to the Australian Research Council (ARC) for financial support provided through the ARC Centre of Excellence for Electromaterials Science and an Australian Laureate Fellowship awarded to D.R.M., the Australian Government for providing an Australian Postgraduate Award (S.B.), the Alexander von Humboldt Foundation for a Feodor Lynen Fellowship (M.W.), and a Senior Research Award (L.S.). We acknowledge the use of facilities in the Monash Centre for Electron Microscopy (MCEM). Access to equipment and assistance from Dr. M. Raveggi (School of Geosciences, Monash University) is greatly appreciated. Portions of this research were undertaken on the X-ray absorption spectroscopy (XAS) beamline at the Australian Synchrotron, VictoChemSusChem 0000, 00, 0 – 0

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Received: October 27, 2014 Revised: January 22, 2015 Published online on && &&, 0000

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FULL PAPERS Cover up! Thin, transparent cobalt oxide films, prepared from cobalt polyaminocarboxylate precursors, are shown to be excellent water oxidation catalysts with catalytic activities comparable to those of thicker films derived from Co2 + .


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S. A. Bonke, M. Wiechen, R. K. Hocking, X.-Y. Fang, D. W. Lupton, D. R. MacFarlane, L. Spiccia* && – && Electrosynthesis of Highly Transparent Cobalt Oxide Water Oxidation Catalyst Films from Cobalt Aminopolycarboxylate Complexes

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Boosting electrochemical water oxidation through replacement of Oh Co sites in cobalt oxide spinel with manganese.

Cobalt-phosphate-assisted photoelectrochemical water oxidation by arrays of molybdenum-doped zinc oxide nanorods.

silicon photoanodes with nanotextured interfaces.

Water oxidation catalyzed by a dinuclear cobalt-polypyridine complex.

Hierarchical cobalt-based hydroxide microspheres for water oxidation.

Green cobalt oxide (CoOx) film with nanoribbon structures electrodeposited from the BF₂-annulated cobaloxime precursor for efficient water oxidation.

Edge reactivity and water-assisted dissociation on cobalt oxide nanoislands.

Electrodeposited cobalt-sulfide catalyst for electrochemical and photoelectrochemical hydrogen generation from water.

What determines catalyst functionality in molecular water oxidation? Dependence on ligands and metal nuclearity in cobalt clusters.

cobalt oxide mixed metal catalysts for methane oxidation.

A Highly Active and Easily Accessible Cobalt Catalyst for Selective Hydrogenation of C═O Bonds.

cobalt(III) phenanthroline complexes. 2. Cobalt chelation of two critical SH groups.

High-performance hybrid oxide catalyst of manganese and cobalt for low-pressure methanol synthesis.

A highly active and stable hydrogen evolution catalyst based on pyrite-structured cobalt phosphosulfide.

Immobilization of a cobalt catalyst on fullerene in molecular devices for water reduction.

Selective cobalt oxidation as a means to differentiate metal-binding sites of cobalt alkaline phosphatase.