Electrochemical Synthesis of Photoelectrodes and Catalysts for Use in

Review pubs.acs.org/CR

Electrochemical Synthesis of Photoelectrodes and Catalysts for Use in Solar Water Splitting Donghyeon Kang, Tae Woo Kim, Stephen R. Kubota, Allison C. Cardiel, Hyun Gil Cha, and Kyoung-Shin Choi* Department of Chemistry, University of WisconsinMadison, Madison, Wisconsin 53706, United States ABSTRACT: This review focuses on introducing and explaining electrodepostion mechanisms and electrodeposition-based synthesis strategies used for the production of catalysts and semiconductor electrodes for use in water-splitting photoelectrochemical cells (PECs). It is composed of three main sections: electrochemical synthesis of hydrogen evolution catalysts, oxygen evolution catalysts, and semiconductor electrodes. The semiconductor section is divided into two parts: photoanodes and photocathodes. Photoanodes include n-type semiconductor electrodes that can perform water oxidation to O2 using photogenerated holes, while photocathodes include p-type semiconductor electrodes that can reduce water to H2 using photoexcited electrons. For each material type, deposition mechanisms were reviewed first followed by a brief discussion on its properties relevant to electrochemical and photoelectrochemical water splitting. Electrodeposition or electrochemical synthesis is an ideal method to produce individual components and integrated systems for PECs due to its various intrinsic advantages. This review will serve as a good resource or guideline for researchers who are currently utilizing electrochemical synthesis as well as for those who are interested in beginning to employ electrochemical synthesis for the construction of more efficient PECs.

CONTENTS 1. Introduction 2. Electrochemical Synthesis of Hydrogen Evolution Catalysts 2.1. Pure Metals 2.1.1. Platinum 2.1.2. Non-Noble Metals 2.2. Metal Alloys and Mixtures 2.3. Compounds or Composites Containing Metal and Nonmetal Elements 2.3.1. Metal Oxides 2.3.2. Metal Sulfides and Selenides 2.3.3. Metal−Phosphorus Composites 3. Electrochemical Synthesis of Oxygen Evolution Catalysts 3.1. IrO2 3.2. RuO2 3.3. Manganese-Based Catalysts (MnOOH and MnxOy) 3.4. CoOOH-Based Catalysts 3.5. NiOOH- and FeOOH-Based Catalysts 3.6. Spinel-Type Oxides 4. Electrochemical Synthesis of Photoanodes 4.1. Binary Oxide Photoanodes 4.1.1. α-Fe2O3 4.1.2. WO3 4.1.3. ZnO 4.1.4. In2O3 and Tin-Doped In2O3 (ITO) 4.1.5. SnO2 © XXXX American Chemical Society

4.1.6. TiO2 4.2. Ternary Oxide Photoanodes 4.2.1. BiVO4 4.2.2. CuWO4 4.2.3. Bi2WO6 4.2.4. Bismuth Molybdates 5. Electrochemical Synthesis of Photocathodes 5.1. Binary Oxide Photocathodes 5.1.1. Cu2O 5.1.2. CuO 5.2. Ternary Oxide Photocathodes 5.2.1. CuFeO2 5.2.2. CuBi2O4 5.3. Chalcopyrite Photocathodes 5.3.1. CuInS2 5.3.2. CuGaS2 and CuGaSe2 5.3.3. Cu(InxGa1−x)Se2 5.3.4. Cu2ZnSnS4 5.4. Silicon 5.5. III−V Semiconductors 6. Summary Author Information Corresponding Author Notes Biographies

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Special Issue: Solar Energy Conversion Received: August 23, 2015

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and the thickness of the films can be easily controlled by monitoring charge passed to induce electrochemical reactions. These are great advantages over other solutions-based synthesis methods where colloidal particles are first produced and then cast onto the conducting substrates. Fifth, since electrodeposition allows for conformal deposition even when the WE has complex surface morphologies, it is an optimal technique for producing multijunction electrodes with uniform and intimate junctions via consecutive depositions of multiple layers. This feature will be particularly beneficial for integration of a semiconductor electrode and a catalyst, multiple semiconductor layers, or a semiconductor and a passivation layer. Due to these reasons, many recent innovations in the development of photoelectrodes and catalysts for use in PECs have been achieved by electrodeposition or electochemical synthesis, which has drawn more interest than ever to electrodeposition. This review will focus on introducing, organizing, and explaining electrodepostion mechanisms and strategies used for the synthesis of catalysts and semiconductor electrodes for use in water-splitting PECs. It is composed of three main sections: electrochemical synthesis of hydrogen evolution catalysts, oxygen evolution catalysts, and semiconductor electrodes. The semiconductor section is divided into two parts: photoanodes and photocathodes. Photoanodes include n-type semiconductor electrodes that can perform water oxidation to O2 using photogenerated holes, while photocathodes include ptype semiconductor electrodes that can reduce water to H2 using photoexcited electrons. Since the primary focus of this review is electrochemical synthesis methods, it will not include fundamental principles or detailed mechanistic understanding of photoelectrodes and catalysts, although notable performances of the resulting materials are discussed briefly. We hope that this review will serve as a good resource or guideline for researchers who are currently utilizing electrochemical synthesis as well as for those who are interested in beginning to employ electrochemical synthesis for the construction of more efficient PECs.

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1. INTRODUCTION Solar water splitting provides a sustainable and environmentally benign route for the production of H2, which can be used as a clean fuel. Since Honda and Fujishima reported the first photoelectrochemical cell (PEC) for water splitting using TiO2 as a photoelectrode in 1972,1 various material systems have been investigated to construct more efficient PECs. The key components of the water-splitting PEC are semiconductor electrodes that harvest the solar spectrum to generate, separate, and transport charge carriers to the semiconductor/aqueous interfaces for use in water reduction and oxidation reactions.2,3 Since the surfaces of semiconductor electrodes are not particularly catalytic for water reduction or oxidation, they need to be coupled with hydrogen evolution catalysts (HEC) and oxygen evolution catalysts (OEC) to facilitate interfacial charge transfer reactions.2,3 PEC hydrogen production has been successfully demonstrated on the laboratory scale. However, the commercial viability of PECs depends critically on the cost of H2 produced by the PECs, which is affected by the cost of PEC construction.4 Therefore, a true breakthrough for PEC development will be possible only when efficiency enhancement is achieved at low cost. The production cost of PECs depends on not only the cost of raw materials used for the photoelectrode and catalyst syntheses but also the cost required for the synthesis procedures. If the synthesis involves complicated steps and extreme conditions, the cost for PEC construction will inevitably increase, even if inexpensive raw materials are selected. In addition to being simple, practical, and inexpensive, the synthesis methods must also have high degrees of freedom in tuning compositions and morphologies of photoelectrodes and catalysts because subtle modifications of morphologies and compositions can markedly affect the performances of photoelectrodes and catalysts. Electrodeposition or electrochemical synthesis is an ideal method to produce individual components and integrated systems for PECs due to its unique intrinsic advantages.5,6 First, it is a solution-based synthesis compatible with ambient conditions, which makes it inexpensive and practical. Electrodeposition is already used in painting automobile bodies, proving that industrial level scale-up of electrodeposition is possible. Second, the material types that can be produced by electrochemical synthesis are broad and include almost all of the semiconductors and catalysts that have been used for solar water splitting (e.g., metals, alloys, oxides, chalcogenides, Si, and III−V semiconductors). Third, the solution-based nature of electrochemical synthesis allows for the easy manipulation of various synthetic variables (e.g., pH, additives, types of solvents, temperature) that markedly affect morphologies (e.g., surface areas, nanostructures, orientations). In addition, deposition potential and current can be used as additional synthesis parameters to finely control the nucleation and growth processes of desired materials. At the same time, uniform doping and solid solutions can also be easily achieved by modifying the compositions of the plating solutions. As a result, an exceptional level of morphology and composition control is possible. Fourth, in electrochemical synthesis, materials grow directly from a conducting substrate (working electrode, WE),

2. ELECTROCHEMICAL SYNTHESIS OF HYDROGEN EVOLUTION CATALYSTS HECs can be divided into three broad categories: pure metals, alloys or mixtures of metals, and compounds or composites containing both metal and nonmetal elements. In this section, electrodeposition methods and mechanisms used to prepare a few exemplary HECs of each category are summarized. Most studies reviewed here are those that utilized electrodeposition for the specific purpose of preparing and testing HECs. However, some studies that did not contain hydrogen evolution reaction (HER) properties were included when necessary for elucidating deposition mechanisms. Since the main focus of this study is to provide a review for general electrodeposition mechanisms used for the synthesis of HECs, for detailed performance comparisons and mechanistic understanding of HER reactions of the catalysts included here, the reader is directed to the available review literature.2,7,8 2.1. Pure Metals

2.1.1. Platinum. The classic model used for the HER on metal catalysts first involves the reduction of a proton to form an adsorbed hydrogen (Hads) on the metal catalyst surface (Volmer step). Release of hydrogen gas is accomplished by either the combination of two adsorbed hydrogen atoms (Tafel B

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step) or the reduction of an additional solution proton species (Heyrovsky step). These steps in acidic and basic conditions can be described as follows, where M represents the metal catalyst.9,10 Acidic Conditions

In-depth reviews of metal electrodeposition processes are available in various review articles published previously.19−21 In simplified terms, typical metal deposition occurs by nucleation onto the WE (typically at defect sites) followed by the growth of nuclei. Nucleation generally requires a larger overpotential than growth. Many factors including the type of WE, type of metal deposited, concentration of metal precursor, and applied potential affect the nucleation and growth processes. Some of these basic principles are illustrated with Pt deposition discussed below. In general, slow deposition using a small deposition overpotential results in 2D layer growth or growth of larger particles (Figure 2a).18,22,23 When moderate or large over-

M + e− + H+ → M−Hads (Volmer step) M−Hads + M−Hads → 2M + H 2 (Tafel step) M−Hads + H+ + e− → M + H 2 (Heyrovsky step)

Basic conditions M + e− + H 2O → M−Hads + OH− (Volmer step) M−Hads + M−Hads → 2M + H 2 (Tafel step) M−Hads + H 2O + e− (Heyrovsky step) → M + H 2 + OH−

Platinum is the best performing catalyst discovered to date for HER. This is because Pt can form Pt−H bonds with ideal bond strength, strong enough for facile adsorption and reduction of H+ yet weak enough to allow for easy release of H2 when reduction is complete. This feature places Pt at the top of the volcano plot presented in Figure 1, which shows the relationship between the exchange current for H2 evolution and the M−H bond strength.11

Figure 2. SEM images of Pt deposited on a Au substrate at (a) 500, (b) 100, and (c) 25 mV vs RHE. Applying less positive potential is equivalent to applying more overpotential for Pt deposition. Adapted with permission from ref 25. Copyright 2006 Elsevier.

potentials are used, 3D clusters or islands of Pt are deposited (Figure 2b).16,24 As the deposition potential increases further, more nuclei can form, which results in rougher surfaces (Figure 2c). When nucleation occurs throughout deposition (i.e., progressive nucleation), a wide range of particle sizes can result. Uniformly sized particles can be obtained when deposition conditions are optimized to induce instantaneous nucleation (opposed to progressive nucleation mentioned above) where nucleation occurs primarily at the beginning of a deposition followed only by the growth of already formed nuclei without inducing further nucleation.22 It is also possible to electrodeposit monolayer Pt films by using underpotentially deposited hydrogen (Hupd) to control the growth of Pt as demonstrated by Liu et al. They used a 3 mM K2PtCl4 solution (pH 4) as the plating solution and applied a potential that was sufficient to deposit Pt metal while simultaneously forming Pt−H bonds by the reduction of H+ to surface-adsorbed H atoms but not enough for complete reduction to H2. These Pt−H bonds passivated deposited Pt, resulting in the formation of 2D monolayers and suppression of 3D growth. Switching between pulses of reduction (Pt monolayer deposition/Hupd) and oxidation (removal of Hupd) potentials allowed for controlled sequential growth of Pt monolayers in a manner analogous to atomic layer deposition (Figure 3).26 Since Pt is an expensive element, many efforts have been made to decrease the amount of Pt needed by increasing the surface area of the Pt deposits. One approach is to use high surface area substrates, such as fibrous high surface area conducting polymer electrodes (Figure 4a), as the WE to deposit Pt.27−30 Another approach is to deposit Pt films with meso- or macroscale porous structures. For example, mesoporous Pt films have been deposited using self-assembly

Figure 1. Volcano plot showing the relationship between the exchange current and the M−H bond strength. Adapted with permission from ref 11. Copyright 1972 Elsevier.

Platinum electrodeposition typically involves reduction of Pt2+ or Pt4+ species in acidic aqueous solutions using PtCl42− or PtCl62− as Pt sources (eqs 1−3).12 For the case of Pt4+ reduction, deposition has been proposed to involve either the formation of Pt2+ as an intermediate13−15 (eqs 1 and 2) or a direct reduction to Pt (eq 3).16−18 Pt(IV)Cl 6 2 − + 2e− → Pt(II)Cl4 2 − + 2Cl− E° = 0.77 V (1)

Pt(II)Cl4 2 − + 2e− → Pt + 4Cl− E° = 0.75 V

(2)

Pt(IV)Cl 6 2 − + 4e− → Pt + 6Cl− E° = 0.76 V

(3) C

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Despite the high catalytic activity of Pt, finding a less expensive and more earth-abundant material is an active area of research due to the prohibitive cost of Pt. Pt is often used as a material for demonstrating an upper limit of hydrogen evolution when developing less expensive HECs. 2.1.2. Non-Noble Metals. Moving down the volcano plot (Figure 1) from Pt, the metals have decreased catalytic activity, but many are much cheaper than Pt. Therefore, if their performance and surface area can be optimized, they may achieve a preferable efficiency-to-cost ratio. Ni is one such metal that has received extensive research for HER. Ni has been typically deposited cathodically from a solution containing Ni salts (e.g., sulfate or chloride) as a Ni2+ source (eq 4).35

Figure 3. (a) Schematic of Hupd-terminated Pt deposition on Au (111), and (b) STEM image of 2D Pt layer deposited on Au(111). Adapted with permission from ref 26. Copyright 2012 American Association for the Advancement of Science.

Ni 2 + + 2e− → Ni E° = −0.250 V

(4) 36−38

Ni2+ reduction to Ni involves a multistep process. For a reduction mechanism that involves two consecutive singleelectron transfer processes, it was proposed that the initial reduction leads to adsorbed Ni+ species (eq 5), which likely remain solvated or complexed with solution species (e.g., Ni(Cl)ads or Ni(OH)ads). Further reduction of adsorbed Ni+ species leads to the formation of metallic Ni (eqs 6 and 7). Ni 2 + + e− → Ni+ads

(5)

Ni+ads + Ni 2 + + 2e− → Ni + Ni+ads

(6)

Ni+ads + e− → Ni

(7)

Similar to Pt, morphological control is possible by altering deposition parameters such as deposition potential as well as composition of the plating solution (e.g., concentration, pH, and additives). For example, Torabi and Dolati demonstrated the morphology change of Ni films (e.g., nucleation density, particle size, surface roughness) by altering concentration of Ni precursor, deposition potential, and addition of additives (Figure 5a and 5b).39 Templates have also been used for Ni synthesis. Simultaneous HER during deposition can produce bubbles on the electrode that act as templates that Ni deposits around (Figure 5c).40 Additionally, hard templates such as anodized aluminum oxide (AAO) membranes have been used to synthesize Ni nanowire arrays. The insulating membrane is not electrochemically active, so Ni can only grow in the pores that contact the WE. After deposition, the membrane can be dissolved away, leaving Ni wires (Figure 5d).41,42 Other non-noble metal catalysts such as Fe and Co can likewise be deposited from reduction of the corresponding ions in the plating solution prepared with metal salts (eqs 8 and 9).35

Figure 4. (a) SEM image of Pt particles deposited in a fibrous polyaniline film on a glassy carbon substrate. Adapted with permission from ref 28. Copyright 1988 American Chemical Society. (b) TEM image of Pt deposited from a hexagonal liquid crystalline solution. Adapted with permission from ref 31. Copyright 1997 American Association for the Advancement of Science. (c) Cross-sectional SEM image of a Pt film deposited through a self-assembled polystyrene sphere template. Adapted with permission from ref 33. Copyright 2002 American Chemical Society. FE-SEM image of a (d) Pt/Cu codeposit on a glassy carbon substrate and (e) porous Pt film made by removing the Cu through electrochemical oxidation. Adapted with permission from ref 34. Copyright 2006 IOP Publishing.

of surfactants or block copolymers as soft templates for electrodeposition (Figure 4b).31,32 Macroporous Pt films have been produced using larger-scale templates such as selfassembled polystyrene spheres (Figure 4c).33 Additionally, codepositing Pt with a less electrochemically stable metal followed by chemical or electrochemical removal of the less stable metal can also be used to achieve higher surface area Pt. For example, Liu et al. codeposited Pt and Cu films onto a glassy carbon substrate (Figure 4d). Cu was then selectively removed by applying a potential to anodically dissolve Cu as Cu2+, which was not sufficient to oxidize Pt. This resulted in a porous Pt film, with pores ranging from 100 to 300 nm in diameter (Figure 4e).34

Fe2 + + 2e− → Fe E° = − 0.440 V

(8)

Co2 + + 2e− → Co E° = − 0.277 V

(9)

However, non-noble metals are often not stable in acidic environments,35 so most HER tests for these metals are performed in alkaline conditions.39,40,43−46 They can require an overpotential of 200 mV or greater to achieve a current density of 10 mA/cm2.39,43,44 (The overpotential required to achieve 10 mA/cm2 will be used as a benchmark for HER activity and will be denoted as η10 mA.) Therefore, it is unlikely that a non-noble metal alone can compete with Pt, which typically shows η10 mA < 50 mV in terms of HER activity. One strategy to improve the D

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they can be present in various phases. They may exist as a mixture of two phases (metal A and metal B), or they may form solid solutions with A’s or B’s pristine crystal structure (e.g., A1−xBx (x < 0.5) with crystal structure of A, AxB1−x with crystal structure of B, or a mixture of A1−xBx and AxB1−x). Formation of solid solutions will most likely involve a change in lattice constants and, therefore, can be examined by X-ray diffraction studies. Another possibility is A and B forming intermetallic alloys, which have crystal structures different from those of A and B. In intermetallic alloys atoms A and B are not statistically mixed, but they have distinctive sites in the lattice. Alloys that are superior in HER to their pure metal counterparts generally owe their improved activity to three possible factors. One is simply increased surface area due to codeposition altering the morphology significantly.50 Another possible factor is an electronic improvement of the material for HER. For example, mixing metals with different d-orbital character (i.e., early and late transition metals) changes the strength of the M−H bond, thus affecting the HER activity.51 Finally, a possible improvement is a synergistic effect. For reactions that require multiple steps, such as HER, the ratelimiting step can vary for different catalysts. Combining metals with optimal activity for each different step can result in a catalyst that is superior to each constituent metal. For example, Highfield et al. proposed for a NiMo system they studied that Ni and Mo could be considered two separate pure phases. The Ni sites can reduce protons and then transfer adsorbed hydrogen to Mo sites, where further reduction to molecular hydrogen is faster.52 While there exists a large number of literature examples of various metal alloy HECs, NiMo alloys in particular have received a great deal of attention owing to their relatively high activity and stability. NiMo is typically deposited in a plating solution containing Ni salts and MoO42− (e.g., Na2MoO4). Podlaha and Landolt proposed that NiMo codeposition (as well as FeMo and CoMo) is possible because Ni species catalyze the reduction of Mo by forming a NiMo intermediate, either through the adsorbed intermediate (Ni+(ads)) or directly with solution-based Ni species.53,54 Like their individual constituent metals, mixed-metal catalysts are often tested for HER activity in basic conditions.47−49,55−57 However, if the alloy contains a metal with superior chemical resistance, then operation in acidic conditions is also possible (e.g., Mo-containing films).50,58,59 Mixed-metal catalysts typically display improved activity compared to their corresponding single-metal films. For example, Navarro-Flores et al. synthesized Ni and various Ni alloys (Mo, W, Fe), where Ni activity was η10 mA = 510 mV and NiMo showed η10 mA = 100 mV in 0.5 M H2SO4. Wang et al. deposited NiMo onto a high surface area Cu foam substrate and reported an activity of η10 mA = 10 mV in 1 M NaOH which even surpassed Pt at higher overpotentials.60

Figure 5. AFM images of Ni deposits synthesized by Torabi and Dolati. (a) with and (b) without cumarine as an additive to affect morphology. Adapted with permission from ref 39. Copyright 2010 Springer Science and Business Media. (c) Optical micrograph of a Ni deposit. Morphology is due to simultaneous HER during deposition. Adapted with permission from ref 40. Copyright 2000 Elsevier. (d) Ni nanowires deposited through an AAO membrane onto ITO-coated glass. The AAO membrane was etched in a 10% NaOH solution to leave only the Ni nanowire array, but some residual Al2O3 was indicated by the authors in the figure. Adapted with permission from ref 41. Copyright 2002 American Chemical Society.

activity and stability of such films is to investigate mixed-metal or alloy catalysts. 2.2. Metal Alloys and Mixtures

Forming mixed-metal or metal alloy films is straightforward if the constituent metals can be deposited using similar deposition mechanisms and conditions. When the plating solution contains all of the desired precursor ions, applying a voltage with a sufficient enough potential for reduction of all of the desired species can result in simultaneous deposition. Making alloys of Ni, Fe, and/or Co is a good example.47−50 The metal composition in the deposited films can be easily tuned by varying the concentration of the metal sources in the plating solution. However, codepositing two or more metals (elements) can exhibit behavior different from depositing single metals alone. One metal may suppress or promote the deposition of one of the other elements in the solution. This effect may even enable the deposition of a metal or element that cannot be electrodeposited alone. Deposition of Ni−Mo alloys (Ni−Mo alloys with various compositions are hereafter abbreviated as NiMo) is a good example; electrochemical deposition of Mo metal is not possible in aqueous media, but Mo can be codeposited with Ni, which will be discussed further below. After alloy depositions, elemental analysis is used to identify the compositions of the deposits, but identifying the phases (i.e., crystal structure) of the codeposited metals is also critical in understanding and reproducing their HER catalytic properties. When metal A and metal B are codeposited, for example,

2.3. Compounds or Composites Containing Metal and Nonmetal Elements

Recently, various compounds and composites of metals with nonmetal elements have been investigated for HER in search of a high-activity, low-cost catalyst. They include oxides, chalcogenides (sulfides and selenides), and phosphides. Compared to metal and metal alloys, electrodeposition has not been widely used to prepare these HECs. However, it is possible to develop electrodeposition conditions to produce many, if not all, of them. In this section some of these catalysts E

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broad Ru peaks from X-ray photoelectron spectroscopy (XPS).63 Electrochemically synthesized RuO2 has not been used much as a HEC. However, some examples have been reported.61,66−69 Ni−RuO2 composite electrodes, where RuO2 nanoparticles are embedded in a Ni matrix, showed values of η10 mA that ranged from 50 to 150 mV in basic conditions. When tested in 0.5 M H2SO4, these films showed comparable performances, indicating that RuO2 is a good HEC in both acidic and basic media.61 These values are similar to RuO2 made through more conventional thermal decomposition routes.70,71 2.3.2. Metal Sulfides and Selenides. Molybdenum disulfide, MoS2, was originally investigated as a hydrodesulfurization catalyst,72 but it has recently received a great deal of attention as a HEC as well.73 MoS2 films can be prepared by cathodic deposition using a MoS42− precursor as both the sulfur and the Mo precursor by eq 11.74 When deposited in aqueous media these films are amorphous, but they can be converted to crystalline MoS2 by annealing in an inert atmosphere (e.g., argon) or vacuum at temperatures > 500 °C.74−76

that have been prepared by electrodeposition to date are summarized. 2.3.1. Metal Oxides. Some transition metal oxides can be stable under the cathodic conditions of HER and can be very catalytically active. RuO2 in particular is a promising oxide HEC material studied in both acidic and basic conditions. RuO2composite electrodes can be made from preprepared RuO2 particles suspended in a solution with metal ions (e.g., Ni2+). Vigorous mixing of the solution while cathodically depositing the metal ion can result in RuO2 particles embedded in a metal matrix (Figure 6).61

MoS4 2 − + 2e− + 2H 2O → MoS2 + 2SH− + 2OH− (11)

Figure 6. SEM images of (a) Ni film deposited from a solution of 0.2 M NiCl2 and 2 M NH4Cl and (b) Ni−RuO2 film deposited from the same solution with a 50 g/L RuO2 suspension. Adapted with permission from ref 61. Copyright 2007 Elsevier.

Murugesan et al. reported that crystalline MoS2 can be obtained as deposited when ionic liquids are used as plating solutions.77 They used molybdenum glycolate and 1,4butanedithiol as the Mo and S precursors, respectively, in the ionic liquid N-methyl-N-propylpiperidinium (cation) and bis(trifluoromethanesulfonyl)imide (anion). Ionic liquids offer a useful route for electrochemical synthesis by providing a large potential and temperature window where the solvent is stable, providing the means to synthesize compounds that may not be deposited in aqueous environments.78 The stability of the ionic liquid allowed for deposition at an elevated temperature (100 °C) as well as a very negative applied potential, which resulted in the deposition of crystalline MoS2 as deposited (Figure 7a).77

RuOx can also be electrodeposited directly using aqueous solutions of RuCl3 as plating solutions by cathodic, anodic, or potential sweeping (i.e., cyclic voltammetry (CV)) methods where both cathodic and anodic reactions occur alternately.62−65 However, as-deposited films are amorphous, and the oxidation states of Ru in the as-deposited films are generally not well defined due to the variety of possible Ru oxidation states. Even oxidation states and coordination of Ru species present in the plating solution can be quite complex.62 Due to this complexity, elucidation of deposition mechanisms is difficult. Additionally, as-deposited films are often composed of a variety of phases, and so they are typically described using a general formula of hydrous ruthenium oxide, RuOx·nH2O. For cathodic deposition, water can be reduced to generate OH− at the WE, resulting in a local increase in pH (eq 10).35 This pH change can result in the precipitation of Ru ions as various oxide and hydroxide species. Under these cathodic conditions, Ru metal can also be simultaneously deposited. Films can then be converted to RuO2 by annealing in air at >200 °C.64 −

2H 2O + 2e → H 2 + 2OH +

−

(or 2H + 2e → H 2

−

Figure 7. (a) Crystalline MoS2 deposited from an ionic liquid solution onto a glassy carbon substrate by Murugesan et al. Adapted with permission from ref 77. Copyright 2013 American Chemical Society. (b) Amorphous MoS3 deposited on ITO-coated glass by Hu and coworkers. Adapted with permission from ref 81. Copyright 2011 The Royal Society of Chemistry.

E° = − 0.83 V E° = 0.00 V)

(10)

Alternatively, if deposition was performed by CV, Ru metal and Ru−oxide/hydroxides deposited during the cathodic portion of the cycle can be oxidized to RuO2 during the anodic portion of the cycle. For anodic deposition, it was assumed that the oxidation of Ru3+ ions to Ru4+ ions resulted in the deposition of RuO2·nH2O.65 However, Jow et al. examined films deposited from the RuCl3 solution using three different conditions (cathodic, anodic, and CV) and reported that all films, even films deposited under strictly anodic conditions, contained Ru with mixed oxidation states, which were shown as

For crystalline MoS2, both theory and experimental results based on crystals synthesized by physical vapor deposition indicate that bulk MoS2 is not catalytic, rather Mo-edge sites are the active sites for HER.79,80 Therefore, producing MoS2 with morphologies that can maximize catalytically active edge sites has been of great interest. Recent studies by Hu and co-workers suggest that crystalline MoS2 is not necessarily required to achieve HER activity.81,82 F

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They investigated electrochemically synthesized, HER active amorphous MoSx (Figure 7b), and proposed that active sites similar to those present in crystalline MoS2 are present in amorphous MoSx based on the appreciable HER activity observed by the amorphous films. They found that MoSx films made cathodically, anodically, or through CV exhibited approximately equivalent HER activity when used as deposited. The phase deposited from the MoS42− solution under cathodic bias or during cathodic scan should be amorphous MoS2 (eq 11), while the phase deposited under anodic bias or during anodic scan should be amorphous MoS3 formed by oxidation of S2− in the MoS42− species to sulfur (eq 12).83 Hu and coworkers propose that, regardless of the as-deposited composition, anodically and cathodically made films are activated by converting to MoS2+x under HER conditions, where the Mo:S ratio is close to 1:2, although they are not crystalline.82 MoS4 2 − → MoS3 + 1/8S8 + 2e−

Although there have been multiple studies reporting electrodeposition of MSe films, including CoSe, their purposes were not to investigate them as HECs. However, Carim et al. recently reported electrochemically synthesized CoSe films with promising HER activity. The as-deposited films were amorphous but considered to be CoSe species in a Se matrix. The films, without any postdeposition annealing procedure, demonstrated an activity of η10 mA = 135 mV and appeared to be stable in acidic conditions for hours.93 This activity is comparable to CoSe2 nanoparticles made through other methods.95 2.3.3. Metal−Phosphorus Composites. Metal−phosphides are another class of compounds that demonstrates very promising HER activity in acidic conditions. M−P films (M = Fe, Ni, and Co) have been cathodically electrodeposited using solutions containing salts of the desired metals with NaH2PO2 or H3PO3 added as the P source.96−98 A review of Ni−P electrodeposition is provided by Daly and Barry.99 As deposited, films are typically amorphous or show the crystallinity of only the metal component.100,101 Due to this amorphous nature, the exact chemical nature of films has not always been completely elucidated. Studies that included XPS indicated that the as-deposited films include metal−phosphides, but they also include metal and phosphorus species with various oxidation states.98,102,103 Since we were unsure of the exact nature of the films prepared by various groups, we refer to these films as M−P films instead of metal−phosphides in this section. Two mechanisms have been proposed for the cathodic synthesis of M−P films. The first is a direct codeposition of M and P (eqs 18 and 19).104

(12)

The onset potential and activity of MoSx films can be improved by incorporating transition metal cations as promotors.84 When MoSx films are prepared by electrodeposition, this can be easily accomplished by adding metal salt precursors to the deposition solution.85 For M−MoSx films (M = Mn, Fe, Co, Ni, Cu, Zn) improved HER activity was observed for Fe, Ni, and Co incorporation. In 1 M H2SO4, Fe incorporation resulted in the largest η10 mA improvement, from ca. 180 mV for unpromoted MoSx to 130 mV for Fe-promoted films. The improved activity was proposed to be from increased surface area, as seen in SEM, and also possibly from an improvement in intrinsic activity caused by the metal ions interacting with the active sites for HER. These MoSx films were tested in acidic conditions and appeared to be stable for hours.81,85 More recently, HER activities of transition metal selenides have been reported.86,87 Metal selenide films can be electrodeposited from solutions of metal ion precursors with H2SeO3 as a Se precursor. With a Se precursor, MSe deposition can be observed at a potential more positive than the potential normally required to deposit the M without Se.88,89 Se has a relatively positive reduction potential, so it can be deposited first (eq 13),35 which can then increase the ease of metal deposition due to favorable MSe formation (eq 14).90,91

M2 + + 2e− → M

H3PO2 + e− + H+ → P + 2H 2O (or H3PO3 + 3H+ + 3e− → P + 3H 2O)

(13)

Se + M

−

+ 2e → MSe

(14)

At sufficiently negative potentials, Se can be further reduced to H2Se (eq 15).35 It is then possible that H2Se can react with solution metal species and precipitate as MSe (eq 16). Alternatively, it can react with solution Se species, resulting in the deposition of excess Se (eq 17), which is not desirable.90 Se + 2H+ + 2e− → H 2Se E° = −0.399 V

(15)

H 2Se + M2 + → MSe + 2H+

(16)

2H 2Se + H 2SeO3 → 3Se + 3H 2O

(17)

(19)

It should be noted that P cannot be electrodeposited by itself but can be codeposited with metals.104 For example, for the preparation of Ni−P films, deposition begins with Ni nucleation and then P can be codeposited due to strong Ni− P interaction.105 The second deposition mechanism involves an indirect pathway, where first PH3 is produced through electrochemical reduction of P precursors (e.g., H3PO2 or H3PO3) (eq 20).35 The PH3 can then chemically react with metal ions, resulting in M0 and P0 deposition (eq 21).106,107 It is also possible that the reaction of PH3 with metal ions can directly form metal phosphides. This deposition mechanism has been supported by detection of PH3 during deposition.106,107

H 2SeO3 + 4H+ + 4e− → Se + 3H 2O E° = 0.741 V 2+

(18)

H3PO2 + 4H+ + 4e− → PH3 + 2H 2O E° = − 0.174 V (or H3PO3 + 6H+ + 6e− → PH3 + 3H 2O)

(20)

2PH3 + 3M2 + → 3M + 2P + 6H+

(21)

Electrodeposited Fe−P, Co−P, and Ni−P films have been tested as HECs in basic solutions and shown η10 mA in the range of 100−150 mV.96,97,101,103,108 The best performance was observed by Saadi et al. using Co−P films. As deposited, the films had an excess of Co, attributed to cobalt oxide species. These films demonstrated an activity of η10 mA = 85 mV in a 0.5 M H2SO4 solution.98 After HER, the Co:P ratio decreased to

Deposited films can range from amorphous to crystalline phases depending on the deposition conditions such as temperature or pH.88,89,92−94 For example, Zhang et al. showed the crystallinity of as-deposited films improved by decreasing the solution pH from 4 to 2.92 G

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1:1 accompanied by a morphology change, indicating that impurity phases were dissolved during HER (Figure 8).

Figure 8. SEM images of Co−P films deposited on a Cu substrate by Saadi et al. (a) as deposited and (b) after voltammetry in a 0.5 M H2SO4 solution. Adapted with permission from ref 98. Copyright 2014 American Chemical Society. Figure 9. SEM images of (a) IrO2 nanoparticles deposited on ITO, (b) ZnO nanorod arrays used as WE, (c) IrO2 nanoparticles deposited on ZnO nanorod arrays, and (d) IrO2 nanotube arrays obtained after dissolution of ZnO nanorods prepared by Zhao et al. Adapted with permission from ref 113. Copyright 2012 Elsevier.

3. ELECTROCHEMICAL SYNTHESIS OF OXYGEN EVOLUTION CATALYSTS The most intensively studied OECs can be categorized into three classes. The first class includes binary oxides of precious metals, such as RuO2 and IrO2, which have shown the best oxygen evolution reaction (OER) performances to date. The second class includes oxides or oxyhydroxides of more earthabundant, inexpensive transition metals such as Mn, Co, Ni, and Fe. The third class includes more complex ternary oxides that have a specific crystal structure, such as spinel and perovskite. In this section, electrodeposition methods developed to date that produce well-known OECs are reviewed. Since the focus of this review lies in explaining mechanisms and strategies used for electrochemical synthesis, detailed OER mechanisms or performance comparisons, which can be found elsewhere, will not be covered.2,109−111

electrode was tested as an OEC in 1 M KOH and showed a stable performance for nearly 2 days at 100 mA/cm2. Ir 3 + + 2H 2O → IrO2 + 4H+ + e− E o = 0.233 V

(22)

Blakemore et al. used a similar mechanism by oxidizing organometallic complexes of Ir3+ (i.e., Cp*Ir(H2O)3]SO4 and [(Cp*Ir)2(OH)3]OH, where Cp* = pentamethylcyclopentadienyl) to form amorphous iridium oxide films.115,116 The authors noted that using the molecular precursors resulted in the formation of catalysts that are significantly more active than crystalline IrO2. Specifically, this catalyst has a very low overpotential for water oxidation, approximately 200 mV at 0.5 mA/cm2 when the performance was tested in 0.1 M KNO3 (pH 2.9).115 A study by Yagi et al. deposited IrO2 films using K3IrCl6 as the Ir3+ source and oxalate as a complexing agent in a pH 10 solution where Ir3+ is otherwise insoluble. In this case, both Ir3+ and oxalate appeared to be oxidized during anodic deposition to form IrO2. The resulting films were composed of amorphous particles with a diameter of 100−250 nm, and their OER performance was tested in 0.1 M KNO3 (pH 6.3).117

3.1. IrO2

IrO2 is one of the most widely studied materials for OER and has shown superior performance over most other OECs. IrO2 has been mainly electrodeposited by two different anodic deposition mechanisms. The first involves the use of a complexing agent, which stabilizes otherwise insoluble Ir4+ ions. When the complexing agents or the bonds between the complexing agent and the Ir4+ ions are oxidized, the resulting free Ir4+ ions can be deposited as IrO2 on the WE.112 For example, Zhao et al. deposited nanoparticulate IrO2 films (diameter, 5−10 nm) using K2IrCl6 as the Ir4+ source and oxalate as the complexing agent (Figure 9a).113 They also demonstrated the deposition of IrO2 nanoparticles onto ZnO nanorods that were used as the WE (Figure 9b and 9c). The IrO2-modified ZnO nanorods were then immersed in acid to etch away the ZnO rods, resulting in IrO2 nanotubes (Figure 9d). The IrO2 nanotube arrays were reported to show better OER performances and a markedly enhanced stability compared to IrO2 nanoparticles when tested in a 0.1 M phosphate buffer solution (pH 7). The second mechanism to deposit IrO2 involves the use of soluble Ir3+ species in the plating solution. When Ir3+ species are oxidized to insoluble Ir4+ ions, IrO2 can be deposited (eq 22). For example, Battaglia et al. used this mechanism to deposit amorphous nanoparticles of IrO2 from a deaerated pH 9 solution of Ir(OH2)2Cl4−.114 Ni nanowires were used as the WE, which provided high surface area. The resulting IrO2

3.2. RuO2

Similar to IrO2, RuO2 is another precious metal binary oxide that has been widely studied and shown to have great OER performance. RuO2 has been electrodeposited anodically mainly by two different mechanisms. The first method is to form composite films that contain RuO2 particles. In this case already prepared RuO2 particles (0.1−0.2 μm) were suspended in the plating solution, and these particles were entrapped in the matrix films during the anodic electrodeposition of the matrix films. For example, Musiani et al. showed the formation of PbO2/RuO2 films by incorporating RuO2 particles into anodically deposited PbO2 films. A solution composed of lead acetate and lead nitrate (pH 4.4), which also contained 0.1−0.8 wt % RuO2 in suspension, was used as the plating solution.118 By oxidizing soluble Pb2+ ions to insoluble Pb4+ ions, the matrix film PbO2 was deposited (eq 23) with RuO2 particles included within the PbO2 film. Optimum RuO2 content was investigated by testing the OER performance of the resulting films in a 0.5 H

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M H2SO4 (pH 0.3) solution.118,119 A similar deposition strategy was introduced in the HEC section when discussing RuO2. The only difference is that when preparing RuO2 as a HEC, a matrix that is electrochemically stable against reduction such as Ni was electrodeposited, while when preparing RuO2 as an OEC, a matrix that is electrochemically stable against oxidation such as PbO2 was electrodeposited. The morphologies of the resulting RuO2 films were similar to the RuO2 films shown in the HEC section (Figure 6) prepared using the same deposition technique.

3Mn 2 + + 4H 2O → Mn3O4 + 8H+ + 2e− E° = 1.824 V (26)

Zaharieva et al. deposited Mn oxides using an aqueous solution containing Mn(CH3COO)2·4H2O and MgSO4 by anodic deposition and by CV. In a 0.1 M phosphate buffer (pH 7), the authors found that only the catalyst deposited by CV was active for OER.124 The authors suggested that when the deposition was performed under a constant anodic potential, MnO2 was formed, which limited the OER performance. However, the CV method resulted in the formation of mixedvalent Mn3+/4+ ions in the deposits, which were more catalytic for OER in pH 7. The authors noted that unless the Mn film was covered in Nafion, the activity of the catalysts decreased significantly within minutes. Zhou et al. deposited various manganese oxide-based OECs from an ionic liquid, ethylammonium nitrate, containing 10% water and Mn(CH3COO)2 at 120 °C at various potentials depending on the pH of the solution.125 The potential needed for electrodeposition decreased with an increase in pH of the electrolyte. The authors determined that the film deposited from acidic electrolyte was mainly composed of a birnessite-like MnOx phase (i.e., a phase closely related to MnO2 having mainly Mn4+ ions) (Figure 11a), and the film deposited from

Pb2 + + 2H 2O → PbO2 + 4H+ + 2e− E o = 1.449 V (23) 3+

The second method involves the oxidation of Ru ions to insoluble Ru4+ ions, which resulted in the deposition of amorphous RuO2 films, as demonstrated by Tsuji et al.120 The resulting films were composed of densely packed particles creating a rough surface (Figure 10a). A solution containing

Figure 10. SEM images of (a) as-deposited and (b) crystalline (after 300 °C annealing) RuO2 prepared by Tsuji et al. Adapted with permission from ref 120. Copyright 2010 Elsevier. Figure 11. SEM images of manganese oxide films deposited from (a) acidic, (b) basic, and (c) neutral electrolytes by Zhou et al. Adapted with permission from ref 125. Copyright 2012 Wiley-VCH.

RuCl3·xH2O was used as the plating solution. The amorphous as-deposited films could be converted to crystalline RuO2 films by annealing at ≥300 °C (Figure 10b), but the authors noted that amorphous RuO2 films required less overpotential to initiate OER than the crystalline films.120 RuO2 or IrO2 OECs can also be prepared by electrodeposition of Ru or Ir metals followed by oxidation. Even if no intentional oxidation step is added, these metals are expected to be oxidized to form oxides, oxyhydroxides, or hydroxides under the anodic condition used for OER.121

basic electrolyte mainly consisted of Mn3O4 with a spinel structure with an overall trend of increasing oxidation state of Mn with lower pH (Figure 11b). Neutral electrolytes yielded material with an intermediate oxidation state, such as Mn2O3 or a mixture of the birnessite-like phase and Mn3O4 (Figure 11c). The catalytic activities of the resulting films were evaluated in 1 M NaOH. The authors observed a higher catalytic activity for films deposited from acidic and neutral electrolytes, suggesting that Mn2O3 and birnessite-like MnOx are more active for OER than Mn3O4 in 1 M NaOH. ́ et al. anodically deposited amorphous manganese Ramirez oxide films and converted them to various crystalline manganese oxide films by changing the annealing conditions.126 For example, annealing at 773 K in air and 873 K under N2 led to a conversion into the crystallized phases of α-Mn2O3 and Mn3O4, respectively. The OER performances and stabilities of the resulting films were examined in both neutral and alkaline solutions. The authors noted that more stable performances were observed in alkaline solutions.

3.3. Manganese-Based Catalysts (MnOOH and MnxOy)

Precious metals like Ir and Ru are expensive and limited in supply. Therefore, developing OECs containing inexpensive and abundant materials like Mn have been of great interest. Mn-based OECs, which include MnOOH and MnxOy, have been mainly electrodeposited by anodic deposition, where soluble Mn2+ is oxidized to insoluble Mn3+ or Mn4+ (eqs 24−26).122−126 For example, Ohsaka and co-workers deposited γ-MnOOH from an aqueous solution containing Na2SO4 and Mn(CH3COO)2 by CV and reported that the catalytic performance of the resulting electrodes was better in alkaline medium than in acidic or neutral media.122,123

3.4. CoOOH-Based Catalysts

2Mn 2 + + 3H 2O → Mn2O3 + 6H+ + 2e− E° = 1.443 V

Cobalt is another abundant and relatively inexpensive transition metal, and Co-based OECs such as CoOOH and Co3O4 have been extensively investigated. Electrodeposition of CoOOH has been mainly achieved by oxidation of Co2+ ions in the plating solution to insoluble Co3+ ions, which precipitate as CoOOH

(24)

Mn 2 + + 2H 2O → MnO2 + 4H+ + 2e− E° = 1.228 V (25) I

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on the WE.127,128 For example, Pauporté et al. anodically electrodeposited CoOOH from a Co(NO3)2 solution (pH 7.4) at varying temperatures. They found that crystalline CoOOH can be deposited above 60 °C, while deposition at room temperature results in the deposition of amorphous CoOOH.128 More recently, developed methods by Koza et al. and Liu et al. utilized reduction of Co3+ to Co2+ in an alkaline NaOH solution (pH 14) to produce Co(OH)2 films first and then oxidized the Co(OH)2 films to CoOOH. Since Co3+ ions are not soluble in basic media, ethylenediamine (en) was added as a complexing agent to form [Co(en)3]3+. When a cathodic bias was applied, [Co(en)3]3+ species were reduced to [Co(en)3]2+ (eq 27). However, since Co2+ ions bind with a higher affinity to O ligands than to N ligands, en cannot serve as an effective complexing agent for Co2+ in strong basic conditions, leading to the precipitation of Co2+ as crystalline β-Co(OH)2 films (eq 28) (Figure 12a). [Co(en)3 ]3 + + e− → [Co(en)3 ]2 +

(27)

[Co(en)3 ]2 + + 2OH− → Co(OH)2 + 3en

(28)

Figure 13. (a) SEM image of Co−Pi catalyst prepared by Kanan et al. Adapted with permission from ref 131. Copyright 2008 American Association for the Advancement of Science. SEM images of Co films prepared in (b) methylphosphonate and (c) borate by Surendranath et al. Adapted with permission from ref 135. Copyright 2009 American Chemical Society.

Co−O clusters with phosphate as a disordered component.133,134 Surendranath et al. replaced the phosphate buffer with methylphosphonate (pH 8) and borate (pH 9.2) buffers to deposit Co-based OECs (Figure 13b and 13c) and found a similar activity to that observed when phosphate was used.135 Similarly, Risch et al. deposited Co-based catalyst films from weakly buffering acetate electrolytes, nonbuffering chloride electrolytes, and phosphate to study the effect of electrolyte variation on the structure and function of Co-based OECs. They also found that the Co catalyst deposited in phosphate electrolyte had the best catalytic activity.136 Following up these studies, various mechanistic and structural studies to further investigate the Co−Pi catalyst were published.137,138 The Co−Pi catalyst has also been generated by the anodization of cobalt−metal thin films or cobalt metal nanoparticles in a phosphate buffer (pH 7). The term anodization is used when the WE is not inert, and the substance used as the WE is oxidized to form the desired product. On the other hand, the term anodic deposition is used when the WE is inert, and the desired product is deposited by the oxidation of the species in the solution at the WE. Young et al. used sputter-coated Co films,139 while Cobo et al. used electrodeposited nanoparticulate Co films for anodization.140 The chemical composition, morphology, and catalytic activity of the resulting films were comparable to those of the electrochemically deposited Co−Pi catalyst. Cobo et al. also showed that electrodeposited Co nanoparticles can be used as a catalyst for HER and that the conversion between Co nanoparticles and Co−Pi catalyst can be achieved reversibly by applying oxidative and reductive potentials as necessary in a phosphate buffer solution (pH 7).140

Figure 12. SEM images of (a) Co(OH)2, (b) CoOOH, and (c) Co3O4 prepared by Liu et al. Adapted with permission from ref 130. Copyright 2014 Elsevier.

Koza et al. investigated the catalytic ability of Co(OH)2 for OER in 1 M KOH and noted that the surface of the film was oxidized to CoOOH by the anodic bias used for OER.129 Liu et al. continued with this work and fully converted the electrodeposited Co(OH)2 film to CoOOH (Figure 12b) by electrochemical oxidation at 95 °C in 1 M KOH. Liu et al. also showed the conversion of electrodeposited Co(OH)2 films to Co3O4 by thermal decomposition at 300 °C in air (Figure 12c). The OER catalytic activities of CoOOH and Co3O4 were compared in 1 M KOH. Co3O4 had an earlier onset and higher current density when geometric area was considered. However, when electrochemically active surface area was considered, their catalytic performances were comparable with Tafel slopes of about 60 mV/dec.130 Kanan et al. showed that when CoOOH was deposited by oxidation of Co2+ ions in a very dilute Co(NO3)2 (0.5 mM) solution containing 0.1 M potassium phosphate (pH 7.0), a substantial amount of phosphate ions was incorporated into the films (Figure 13a). The resulting amorphous OEC was referred to as Co−Pi catalyst by the authors, and its catalytic performance and stability were investigated in a phosphate buffer solution (pH 7).131,132 The as-deposited film was amorphous, but later structural analysis using X-ray absorption spectroscopy (XAS) and X-ray pair distribution function analysis suggested that it contains domains of cubane-type

3.5. NiOOH- and FeOOH-Based Catalysts

Ni- and Fe-based oxyhydroxides are more examples of earthabundant elements that have been used to develop OECs. As in the case of CoOOH, NiOOH and FeOOH can also be deposited by oxidizing Ni2+ or Fe2+ ions in the plating solution to insoluble Ni3+ or Fe3+ ions, which precipitate as NiOOH or FeOOH on the WE. As in the case of CoOOH, when NiOOH was deposited from a buffer solution containing a small amount of Ni2+ ions, the resulting film contained a significant amount of buffer anions. For example, Dincă et al. anodically prepared NiOOH by oxidation of Ni2+ in a borate buffer (pH 9.2) containing 0.2 M borate and 2 mM Ni(NO3)2 (Figure 14a). The resulting NiOOH film contained borate ions in the NiOOH deposits and was referred to as a Ni−Bi catalyst by the authors. The Ni−Bi catalyst was tested in borate buffer (pH J

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oxidation states of Ni. Trotochaud et al. also electrodeposited Ni(OH)2 films to prepare NiOOH catalysts to demonstrate the intrinsic catalytic properties of a pure NiOOH phase and the effect of intentional and incidental Fe incorporation into the NiOOH structure on the OER performances of NiOOH in 1 M KOH solution.148 3.6. Spinel-Type Oxides

Crystalline spinel-type Co3O4 can be anodically deposited by oxidation of Co2+ to Co3+ in basic aqueous media. In order to solubilize Co2+ ions in a basic solution, complexing agents such as tartrate and glycine are added.149,150 For example, Koza et al. used a 2 M NaOH solution containing 5 mM Co2+ and 6 mM tartrate (tart). Upon application of anodic bias, Co2+ ions were oxidized to Co3+ ions (eq 31), which were coprecipitated with Co2+ to form Co3O4 (eq 32).150 Koza et al. noted that the crystallinity of the electrodeposited Co3O4 was a strong function of the deposition temperature and current density.

Figure 14. SEM image of (a) Ni−Bi film prepared by Dincă et al. Adapted with permission from ref 141. Copyright 2010 National Academy of Sciences. SEM images of NiOOH deposited using (b) [Ni(NH3)6]Cl2 and (c) [Ni(en)3]Cl2 by Singh et al. Adapted with permission from ref 144. Copyright 2013 The Royal Society of Chemistry.

9.2) for 8 h and showed a stable performance.141 The asdeposited film was amorphous, but using XAS studies, Risch et al. elucidated that its structure is related to that of γ-NiOOH with water and borate molecules between the oxide sheets.142 Bediako et al. followed up on this work and studied the oxidation state and structural changes associated with the Ni− Bi catalyst activity. The Ni−Bi films required an oxidative pretreatment (i.e., applying 1.64 V vs RHE in 1 M borate solution (pH 9.2) before using as an OEC) in order to achieve maximum catalytic activity, and the authors determined that this anodic pretreatment increased the average oxidation state of the Ni from +3.16 to +3.6 and improved the catalytic activity by 3 orders of magnitude.143 Singh et al. used the same deposition mechanism to prepare Ni−Bi films but used three different Ni precursors, [Ni(en)3]Cl2, [Ni(OH2)6](NO3)2, and [Ni(NH3)6]Cl2, and investigated the morphologies and performances of the resulting films (Figure 14b and 14c). The catalyst films were amorphous, but XAS studies indicated that the γ-NiOOH phase was present in all films.144 The authors noted that the films deposited using [Ni(en)3]Cl2 were found to be more homogeneous and to have a higher electroactive surface area. The authors reported stability and robustness of the catalyst in a 0.6 M borate buffer (pH 9.2).144 Another frequently used method to prepare NiOOH films is to deposit and then oxidize Ni(OH)2 to form NiOOH. The deposition of Ni(OH)2 is typically achieved in a plating solution containing Ni2+ by electrochemical generation of OH−, which increases the local pH near the WE, leading to the precipitation of Ni2+ as Ni(OH)2. Electrochemical generation of OH− is commonly achieved by reduction of water or reduction of nitrate (eqs 10 and 29).

2Co2 +(tart) → 2Co3 + + 2(tart) + 2e− 2Co3 + + Co2 +(tart) + 8OH− → Co3O4 + 4H 2O + (tart)

(32)

The resulting Co3O4 was tested for OER in 1 M NaOH or KOH solution. The films commonly showed two anodic peaks before water oxidation, which were assumed to be due to the oxidation of Co2+ to Co3+ and Co3+ to Co4+.149,150 The stability of the film was tested in 1 M KOH at 10 mA/cm2, and the Co3O4 film showed a stable performance for 48 h.150 Electrochemical synthesis has also been used to prepare ternary spinel-type oxides as OECs, such as NiCo2O4 and ZnCo2O4.151,152 For example, Kim et al. electrochemically prepared ZnCo2O4 and Co3O4 to compare their OER catalytic performances. ZnCo2O4 has a regular spinel structure where Zn2+ only replaces Co2+ in the tetrahedral (Td) sites in Co3O4. Therefore, the authors postulated that comparing OER performances of ZnCo2O4 and Co3O4 could enable them to investigate whether the Co2+ ions in the tetrahedral sites have a major catalytic role for the OER reaction of Co3O4. If ZnCo2O4 shows comparable catalytic performances to Co3O4, it would mean that Co2+ in the Co3O4 is not catalytically important, and Co2+ can be replaced by a much cheaper Zn2+. In order to prepare Co3O4 films, Kim et al. cathodically deposited Co(OH)2 first by nitrate reduction (eq 29) and converted Co(OH)2 to Co3O4 by annealing at 500 °C in air. For the synthesis of ZnCo2O4, ZnO and Co(OH)2 were first codeposited using nitrate reduction and annealed to form ZnCo2O4. (Mechanisms of ZnO deposition are discussed in detail in the ZnO section.) The authors intentionally deposited excess ZnO to ensure uniform conversion to ZnCo2O4 throughout the film during annealing and removed remaining ZnO by dissolution in strong base where ZnCo3O4 is stable (Figure 15).152 The resulting ZnCo2O4 and Co3O4 electrodes were tested in both 1 M KOH and pH 7 phosphate solutions and showed comparable performances. Since these two electrodes were prepared from the same synthesis conditions with comparable film thicknesses and morphologies, these results could be used to suggest that the Co2+ ions in the Td sites in the Co3O4 spinel structure are not catalytically critical for the enhancement of OER, and ZnCo 2 O 4 can be a more economical and environmentally benign replacement for Co3O4 as an OEC.152

NO3− + H 2O + 2e− → NO2− + 2OH− E o = − 0.24 V (29)

The oxidation of Ni(OH)2 to NiOOH can be achieved either by an intentional postdeposition electrochemical oxidation process (e.g., applying an anodic bias in an electrolyte that does not contain metal ions) before the OER test or by the anodic bias applied during the OER test (eq 30).145,146 Ni(OH)2 + OH− → NiOOH + H 2O + e−

(31)

(30)

Louie and Bell used this mechanism to prepare Ni 1 − x Fe x OOH via cathodic electrodeposition of Ni1−xFex(OH)2 and identified an optimum Fe content to achieve the best OER performance in 1 M KOH.147 They also studied how the presence of Fe affected the local structure and K

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4.1. Binary Oxide Photoanodes

4.1.1. α-Fe2O3. Hematite (α-Fe2O3) is one of the most extensively studied n-type binary oxides for use in solar water splitting. Its band gap (Eg = 2.0−2.2 eV) allows for the utilization of a significant portion of visible light; a theoretical maximum of its solar-to-hydrogen efficiency (STH) exceeds 12%.154−156 It is also chemically stable in neutral and basic media, and it does not suffer from photocorrosion. Additionally, it is environmentally benign and inexpensive. The conduction band minimum (CBM) of α-Fe2O3 is too positive (ca. 0.2 V vs RHE) to utilize the photoexcited electrons to reduce H+ at the counter electrode without applying an external bias. However, it can be coupled with a photocathode (i.e., p-type semiconductor) suited for solar H2 production in the form of a photoelectrochemical diode or a tandem device to achieve solar water splitting. The very short hole diffusion length (2−4 nm), poor electrical conductivity, and facile surface recombination of Fe2O3 have been identified as its major limitations for use as a photoanode.157−160 In order to overcome these limitations, various synthesis methods and strategies have been developed, which include composition and morphology optimizations.161,162 The goal of composition tuning is to increase charge transport properties or to reduce surface states, while the goal of morphology optimization is to reduce the distance that the holes need to travel to reach the interface by incorporating nanostructures. Fe2O3 has been prepared by both anodic and cathodic deposition. The anodic deposition involves electrochemical oxidation of Fe2+ to Fe3+.35 In slightly acidic aqueous media (3 < pH < 7), Fe2+ is soluble (eq 33) but Fe3+ is insoluble (eq 34). Therefore, the electrochemical oxidation of Fe2+ to Fe3+ (eq 35) can result in precipitation of Fe3+ as FeOOH on the WE (eq 36).

Figure 15. (a) Schematic illustration of the synthesis of ZnCo2O4 using electrochemical codeposition of Co(OH)2 and ZnO. (b) SEM image of ZnCo2O4 film after ZnO removal. (c) Side view SEM image of ZnCo2O4 film. (d) Top-view and (e) side-view SEM images of Co3O4 prepared by Kim et al. Adapted with permission from ref 152. Copyright 2014 American Chemical Society.

Although the use of electrodeposition to prepare ternary oxide-based OECs has been limited to date, it should be possible to electrochemically prepare more diverse oxide-based OECs (e.g., NiLa2O4, LaCoO3) using the various mechanisms and methods discussed in this review as well as those previously published.5 More studies on electrodeposited OECs are expected in the future due to the simplicity and versatility of electrochemical synthesis methods.

log[Fe2 +] = 13.29 − 2pH

(33)

log[Fe3 +] = 4.84 − 3pH

(34)

Fe2 + → Fe3 + + e− E° = 0.771 V

(35)

Fe3 + + 2H 2O → FeOOH + 3H+

(36)

Spray and Choi used this mechanism to anodically deposit transparent FeOOH films in a slightly acidic solution (pH 4.1) (Figure 16a).163 The as-deposited FeOOH films were X-ray amorphous, but Raman spectroscopy suggested that the local structure of Fe3+ resembles that of γ-FeOOH. The as-deposited films were converted to transparent, nanocrystalline Fe2O3 films by annealing at 520 °C in air (Figure 16b). SEM studies showed that the resulting Fe2O3 films were composed of spherical particles with sizes of approximately 10−100 nm. Spray et al. also demonstrated the incorporation of Al3+ and Sn4+ ions into nanoparticulate Fe2O3 electrodes using a simple postdeposition surface modification method, which reduced surface states responsible for rapid surface recombination.164 After the surface treatment with Al3+ and Sn4+ ions, a significantly enhanced photocurrent and a shift of photocurrent onset potential to the negative direction were observed in 1 M NaOH solution. Further photocurrent enhancement for water oxidation was observed when Co2+ ions were adsorbed on the Fe2O3 surface as OECs. Unlike anodic deposition, cathodic deposition of FeOOH uses Fe3+ ions as the Fe source, and the oxidation state of Fe3+ ions does not change during deposition. For example, Schrebler

4. ELECTROCHEMICAL SYNTHESIS OF PHOTOANODES This section aims to give an overview of electrochemically deposited n-type semiconductors that have been investigated for use as photoanodes in a water-splitting PEC. They include various binary and ternary oxides. While chalcogenide (sulfides, selenides, and tellurides) semiconductors have also been studied as photoanodes in PECs, they were utilized not for photo-oxidation of water but for photo-oxidation of hole acceptors such as sulfide and/or sulfite. Some studies implied water oxidation by chalcogenide photoanodes, but no O2 detection was demonstrated to support their claims.153 Since the main goal of this review is to examine deposition methods to prepare photoanodes that have been used for solar water oxidation, electrochemical synthesis of n-type chalcogenide photoanodes was not covered in this review. For each of the photoanodes in this section, various electrodeposition conditions and mechanisms are reviewed first, followed by discussion of any notable photoelectrochemical properties of the resulting electrodes. L

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deposition in acidic aqueous media can result in poor quality films because the deposition potential required to reduce Fe2+ to Fe is sufficient to initiate H2 evolution (eq 39). H2 evolution during Fe deposition can affect the uniformity and adhesion of Fe deposits. In particular, if Fe needs to be deposited on a transparent conducting oxide substrate like fluorine-doped tin oxide (FTO) to allow back-side illumination of the resulting Fe2O3/FTO photoanodes, obtaining a good quality Fe film can be more difficult than the cases where a metal WE that is poorly catalytic for H2 evolution and has a good affinity to Fe metal can be used. In order to overcome this issue and deposit uniform Fe films, McDonald and Choi used a nonaqueous solvent, dimethyl sulfoxide (DMSO), containing Fe(ClO4)2 as the plating solution. Prior to deposition, the plating solution was purged with argon to remove dissolved O2. The reduction of O2 can compete with the reduction of Fe2+ ions, lower the deposition efficiency, and alter the solution composition (e.g., OH− generation) (eq 40). (39)

O2 + 2H 2O + 4e− → 4OH− E° = 0.401 V

(40)

The as-deposited nanocrystalline Fe films were converted to α-Fe2O3 films by annealing at 500 °C (Figure 16c and 16d). The resulting nanoparticulate Fe2O3 films were uniform, transparent, and adherent. McDonald and Choi used these αFe2O3 electrodes to investigate various photodeposition conditions where photogenerated holes in α-Fe2O3 were directly utilized to form Co−Pi OECs on the α-Fe2O3 surface. Enache et al. used similar deposition conditions to prepare Sidoped Fe2O3 electrodes by adding tetraethyl-orthosilicate as the Si source to the DMSO plating solution.175 The best photoelectrochemical water oxidation performance of electrochemically prepared Fe2O3 photoanodes was reported by Zeng et al. (Figure 17d).176 They first electrodeposited Fe films from a basic aqueous solution (pH 10.8) where Fe2+ ions

Figure 16. SEM images of (a) anodically synthesized FeOOH as deposited and (b) α-Fe2O3 prepared by annealing by Spray and Choi. Adapted with permission from ref 163. Copyright 2009 American Chemical Society. SEM images of (c) cathodically as-deposited Fe and (d) α-Fe2O3 obtained by annealing by McDonald and Choi. Adapted with permission from ref 174. Copyright 2011 American Chemical Society.

et al. used a plating solution composed of FeCl3, KF, KCl, and H2O2.165 F− was used as a complexing agent to form soluble FeF2+ species since Fe3+ is insoluble. During cathodic deposition, H2O2 was reduced to generate OH− ions (eq 37). The increase in OH− concentration on the WE decreased the solubility of the FeF2+ species and resulted in the deposition of FeOOH (eq 38). In general, the reduction of Fe3+ can compete with the reduction of H2O2, but in this study the reduction of Fe3+ was effectively suppressed by the presence of F− forming FeF2+ species. The as-deposited films were converted to crystalline Fe2O3 films by annealing at 500 °C in air. H 2O2 + 2e− → 2OH−

2H+ + 2e− → H 2 E° = 0.000 V

(37)

3H 2O2 + 2FeF2 + + 6e− → 2FeOOH + 2F− + 2H 2O (38)

The study by Schrebler et al. did not investigate photoelectrochemical properties of the resulting Fe2O3 films, but their deposition method was adopted by several groups to prepare α-Fe2O3 photoanodes or M-doped α-Fe2O3 photoanodes (M = Al, Ti, Cr, Ni, Sn, Zr, Zn, Mo, Pt, etc.) for use in solar water oxidation.166−173 For example, Hu et al. prepared αFe2O3 and Pt-doped α-Fe2O3 using this mechanism because Pt4+ ions in Pt-doped α-Fe2O3 are substitutionally doped at the Fe3+ sites, which can result in an increase in carrier density.166 In order to prepare Pt-doped Fe2O3 electrodes, H2PtCl6 was added to the plating solution as the Pt source with the Pt content in the electrolyte ([Pt]/([Pt] + [Fe]) adjusted to be 2−10 atom %. Another cathodic deposition method used to prepare αFe2O3 photoanodes involved the deposition of nanocrystalline Fe metal films (eq 8), followed by thermal oxidation of Fe to αFe2O3, as demonstrated by McDonald and Choi.174 Fe

Figure 17. (a) J−V plots and (b) J−t plots at 1.23 V vs RHE of photooxidation of water by α-Fe2O3 and Co−Pi/α-Fe2O3 films reported by Zeng et al. (1 M KOH, AM 1.5 illumination). Adapted with permission from ref 176. Copyright 2015 The Royal Society of Chemistry. (c) SEM image and (d) J−t plots at 1.32 V vs RHE of Co− Pi/α-Fe2O3 electrode with an inversed opal structure, reported by Shi et al. Adapted with permission from ref 178. Copyright 2013 Elsevier. M

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were stabilized with ammonia and converted the Fe films to αFe2O3 by thermal oxidation. When the resulting film was combined with Co−Pi OECs, a photocurrent density of 1.9 mA/cm2 was achieved at 1.23 V vs RHE in a 1 M KOH solution (Figure 17a and 17b). Shi et al. produced mesoporous Fe2O3 electrodes with an inverse opal structure by depositing Fe metals around polystyrene beads assembled on the WE (Figure 17c).177 During the annealing step, Fe was converted to Fe2O3 and the polystyrene beads were removed. When combined with Co−Pi OECs, the resulting Fe2O3 electrodes achieved a photocurrent density of ca. 2.4 mA/cm2 at 1.32 V vs RHE in a 1 M NaOH solution,178 which is comparable to the performance reported by Zeng et al.176 4.1.2. WO3. Tungsten oxide (WO3) is another n-type semiconductor that has been extensively studied as a photoanode for use in solar water splitting.179 WO3 has an indirect band-gap energy of 2.7−2.8 eV, and, therefore, it can utilize only a limited portion of visible light. Also, while its valence band maximum (VBM) position can provide sufficient overpotential for photogenerated holes to oxidize water, its CBM position (0.3 V vs RHE) cannot allow photoexcited electrons to reduce water to hydrogen without an applied bias.180−183 However, it is one of the few oxides that are chemically stable in acidic aqueous media. WO3 is also nontoxic and inexpensive. WO3 films can be prepared by cathodic deposition. The most commonly used plating solution is prepared by dissolution of tungsten powder in 30% H2O2, where H2O2 acts as both an oxidizing and a complexing agent to form a predominately dimeric tetraperoxoditungstate species (W 2 O 11 2− or [W2(O)3(O2)4(H2O)2]2−) (eq 41).184

Figure 18. SEM images of (a) WO3 film and (b) WO3/Co−Pi OEC film prepared by Seabold and Choi. Adapted with permission from ref 187. Copyright 2011 American Chemical Society. (c) SEM image of WO3 films prepared by Kwong et al. using oxalic acid as an additive. Adapted with permission from ref 190. Copyright 2013 American Chemical Society. (d) SEM image of porous WO3 film prepared by Hill et al. using ethylene glycol as a stabilizer. Adapted with permission from ref 191. Copyright 2012 American Chemical Society.

was tested in a pH 7 solution, complete suppression of peroxo formation was achieved and all photogenerated holes were used for O2 evolution. As a result, the WO3/Co−Pi OEC photoelectrode also showed long-term photostability. Mi et al. tested photoelectrochemical properties of electrodeposited WO3 photoanodes in a series of acidic aqueous electrolytes (i.e., HCl, H2SO4, and HClO4) without adding an OEC.188 They reported that in these acidic solutions, photooxidation of the respective acid anions (Cl−, SO42−, ClO4−) instead of water was the predominant process at WO3 photoanodes, which also confirmed the poor catalytic nature of the bare WO3 surface for O2 evolution. Cathodic deposition of WO3 in aqueous media has also been used by Baeck et al. to combinatorially prepare tungsten oxidebased composite photoanodes.189 In this study, plating solutions were prepared by adding varying amounts of metal (M) chloride salts (M = Ni, Co, Cu, Zn, Pt, Ru, Rh, Pd, and Ag) to a peroxotungstate solution. From this solution, WO3 was codeposited with M or MxOy depending on the reduction potential of the M ions. Since the as-deposited films were amorphous, compositional analysis of the films was performed after the films were annealed at 600 °C. The resulting films were (i) a mixture of WO3 and MxOy (when M = Co, Cu, Zn), (ii) a mixture of WO3, MxOy, and M (when M = Pt, Ru, Rh, Pd, and Ag), or (iii) a mixture of WO3 and MxWyOz (e.g., NiWO4 when M = Ni). For the films that contained MxOy or NiWO4, it was not clear whether these oxide phases were formed during the deposition or during the annealing step since the asdeposited films were reported to be amorphous. Morphology tuning of WO3 films could be achieved by simply adjusting the compositions of the plating solution. For example, Kwong et al. used the same deposition mechanism and studied the effect of various carboxylic acid additives (e.g formic acid, oxalic acid, and citric acid) on the growth of WO3

2W(s) + 10H 2O2 → [W2(O)3 (O2 )4 (H 2O)2 ]2 − + 2H3O+ + 5H 2O

(41)

During cathodic deposition, peroxo bonds of W2O112−, which are critical for the solubility of tungsten species, are cleaved, resulting in the deposition of tungsten oxide films (eq 42). W2O112 − + (2 + x)H+ + x e− + ((2 + x)/2)/H 2O + ((8 − x)/4)O2

(42)

The as-deposited tungsten oxides are usually in an amorphous hydrated form (WO3·nH2O) but can be converted to crystalline WO3 by annealing at >400 °C. In general, the peroxotungstate solutions are not stable and can go through hydrolysis and condensation reactions, resulting in the bulk precipitation of tungsten oxides in the plating solution. To increase the stability of peroxotungstate species, alcohol, like isopropanol or ethanol, is often added to the solution as a stabilizer.185,186 Seabold and Choi prepared WO3 photoanodes using the aforementioned mechanism.187 The resulting WO3 electrode had a flat and featureless surface (Figure 18a). In aqueous solutions with no hole acceptors other than water, WO3 photoanodes are known to produce peroxo species as well as O2 during photo-oxidation of water because the surface of WO3 is not particularly catalytic for O2 evolution. The accumulation of peroxo species on the surface was reported to cause a gradual loss of photoactivity of WO3.187 The authors demonstrated that when a Co−Pi OEC layer was electrodeposited on WO3 (Figure 18b) and the resulting WO3/Co−Pi OEC photoanode N

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Figure 19. (a) Tandem junction microwire array with a buried homojunction (n−p+-Si) coated by ITO and n-WO3. (b) Two-dimensional crosssection of an individual tandem junction array unit cell. (c) Electronic structure of the tandem device in the dark showing the buried n−p+-Si junction, ohmic Si/ITO/n-WO3 junction, and n-WO3/liquid junction. The device is shown equilibrated with the oxygen-evolution potential. (d) Steady-state electronic structure of the tandem device under illumination with the carrier movement directions shown. Adapted with permission from ref 193. Copyright 2014 The Royal Society of Chemistry.

films and investigated the relationship between the morphologies (e.g., grain size) and the photocurrent of the resulting WO3 films (Figure 18c).190 Hill and Choi reported a modified cathodic deposition condition for WO3 that can be used to prepare high surface area porous WO3 electrodes.191 The modification is to use ethylene glycol (EG) as a stabilizer instead of isopropanol (Figure 18d). The authors observed that the amount of EG in the plating solution and deposition temperature directly affected the porosity of WO3. The authors postulated that addition of EG makes the plating medium more viscous and hindered facile ion diffusion, which triggered diffusion-limited growth and created high surface area films. As expected, the surface area increase directly resulted in the increase in photocurrent generation observed in their study. In addition to the peroxotungstate solutions, solutions prepared by the dissolution of Na2WO4 in 2 M H2SO4 have also been used as plating solutions for the deposition of WO3 films, although the resulting WO3 films were not tested for photoelectrochemical properties.192 In this case, the cathodic deposition processes were believed to involve the reduction of W6+ ions to form either hydrogen tungsten oxide bronzes, HxWO3 (0 < x < 1), or nonstoichiometric tungsten oxides with the formula, WO3−y (0 < y < 1), where the oxidation state of a portion of the W ions is lower than +6. The resulting films can be converted to WO3 by thermal or electrochemical oxidation. Electrodeposition of WO3 using a peroxotungstate solution has been used for the construction of a tandem device. Shaner et al. electrodeposited a thin layer of n-type WO3 on ITOcoated n−p+-Si core−shell microwire arrays to create a tandem junction photoanode (Figure 19a and 19b).193 Since the buried n−p+-Si junction provided an additional photovoltage to the WO3 photoanode, unassisted H2 production at the cathode was demonstrated in a 1 M H2SO4 solution (Figure 19c and 19d). This result clearly showed the ability of electrodeposition to conformally deposit semiconductor films on complex surfaces. 4.1.3. ZnO. Zinc oxide is an n-type oxide that is environmentally friendly and inexpensive.194,195 Its band-gap (3.2 eV) and band-edge positions are similar to those of TiO2, but it was reported that ZnO has a higher electron mobility

than TiO2.194−199 Since its photon absorption is quite limited, due to its wide band gap, it has mainly been used as a substrate for dye-sensitized solar cells and for the deposition of smaller band-gap photon absorbers or as a buffer layer for solid-state solar cells.194,198−200 Since ZnO can be easily produced by various methods with diverse morphologies, it has also been used as a photoanode for water-splitting PECs to demonstrate new concepts regardless of its low efficiency. Electrochemical synthesis of ZnO is particularly useful because ZnO can be deposited as a highly crystalline phase when the deposition temperature is >50 °C.201,202 Other oxides are generally deposited as hydroxide or amorphous hydrated oxides, which require annealing treatments to obtain crystalline films. To date, three major mechanisms have been established for electrochemical synthesis of ZnO. The first is by electrochemical generation of base. In this case, slightly acidic plating solutions containing Zn2+ ions and species that can generate OH− by electrochemical reduction were generally used. NO3−, O2, and H2O2 are typical species used for electrochemical generation of base (eqs 29, 40, and 37).202−204 Under cathodic bias, electrochemically generated OH− at the WE increases the local pH, hence decreasing the solubility of Zn2+(eq 43). As a result, Zn2+ precipitates as Zn(OH)2 on the WE (eq 44). At elevated temperatures (>50 °C), simultaneous dehydration occurs and crystalline ZnO is deposited on the WE (eq 45). Since ZnO is deposited as a crystalline phase, crystal habit control (hexagonal plates vs hexagonal rods) is also possible by adding additives that affect the relative growth rates along different crystallographic directions.6 log[Zn 2 +] = 10.96 − 2pH

(43)

Zn 2 + + 2OH− → Zn(OH)2

(44)

Zn(OH)2 → ZnO + H 2O

(45)

The same mechanism can be used to form solid solutions, Zn1−xMxO, if M2+ ions can be stabilized in a tetrahedral site of Zn2+ in the wurtzite structure. For example, Jaramillo et al. O

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Figure 20. Oxidation reaction of the L-ascorbate dianion in alkaline solutions (pH > 11.7). Adapted with permission from ref 206. Copyright 2006 American Chemical Society.

template to deposit ZnO films with a lamellar structure (d001 = 3.5 with SDS and d001 = 2.8 nm with LA) (Figure 21b). Finally, ZnO electrodes can also be produced by electrodeposition of Zn metal by reduction of Zn2+ followed by thermal conversion to ZnO. In general, the growth mechanism and growth pattern of metals are quite different from that of metal oxides. Therefore, metal oxide electrodes prepared by thermal oxidation of electrodeposited Zn metal can have significantly different morphologies that may not be obtained by direct electrodeposition of ZnO. For example, López and Choi prepared Zn metal electrodes using cathodic deposition in a DMSO solution containing Zn(ClO4)2·6H2O as the zinc source and LiClO4·3H2O as the supporting electrolyte.208 When the deposition potential and temperature were optimized, Zn films were deposited with a uniform fibrous morphology (Figure 21c), which could be converted to fibrous ZnO films by annealing at 450 °C in the air. In addition, López and Choi electrodeposited Zn metal films with fern-shaped dendritic frameworks from nonaqueous formamide media containing Zn(ClO4)2·6H2O and LiClO4·3H2O by changing the interplay between the growth rate and the mass transport rate (Figure 21d).209 Nonaqueous media such as DMSO and formamide offer a wider range of possible deposition temperatures and deposition potentials than aqueous solutions, thus allowing more freedom in altering the deposition rate and temperature, which in turn affects the crystal growth rate and mass transport-limited growth. Although photocurrent generated by ZnO for water oxidation is not notable, as expected from its wide band gap, ZnO photoanodes have been used to demonstrate new concepts. For example, the feasibility of photodepositing Co− Pi OECs on a photoanode using photogenerated holes in the photoanode was first demonstrated by Steinmiller and Choi using electrodeposited ZnO photoanodes.210 During their attempt to electrodeposit Co−Pi OEC on the ZnO surface following the original electrochemical method reported by Kanan and Nocera, Steinmiller and Choi noted that the oxidation potential of Co2+ to Co3+ is similar to that of water and, therefore, any photoanodes that can photo-oxidize water should also be able to utilize their photogenerated holes to photo-oxidize Co2+ to Co3+ to produce Co−Pi OECs (Figure 22). They also noted that photodeposition of OECs automatically ensures the selective placement of the OECs at locations where they can best utilize photogenerated holes for water oxidation because the formation of OECs also occurs preferentially where holes are readily available. This makes photodeposition of OECs the most efficient way to couple OECs with photoanodes if the OECs are deposited by oxidation reactions. The photodeposition method of OECs is now widely used in preparing integrated photoanode systems. 4.1.4. In2O3 and Tin-Doped In2O3 (ITO). Due to its wide band gap (2.8−3.5 eV), it is unlikely that n-type In2O3 will be used as a photoanode for water oxidation by itself, although its

demonstrated electrodeposition of Zn1−xCoxO films by adding Co(NO3)2 as the Co2+ source in a plating solution used for ZnO deposition.205 The second method is an anodic deposition that involves electrochemical generation of acids. ZnO is amphoteric and soluble in both strong acids and bases. (The solubility of ZnO is the lowest in pH 10.3.) Therefore, electrochemically lowering the local pH at the WE in a strong basic solution containing soluble Zn(OH)42− and Zn(OH)3− species can also lead to the deposition of ZnO films, as demonstrated by Limmer et al.206 During anodic deposition, ascorbate dianion was oxidized to diketogulonic acid (Figure 20), lowering the local pH at the WE, resulting in the deposition of ZnO (Figure 21a).

Figure 21. SEM image of ZnO film deposited anodically on Au(111) by Limmer et al. Adapted with permission from ref 206. Copyright 2006 American Chemical Society. (b) SEM and TEM (inset) images of mesoporous ZnO films with a lamellar structure deposited with 5 wt % lauric acid by Steinmiller and Choi. Adapted with permission from ref 207. Copyright 2007 American Chemical Society. (c) SEM image of fibrous Zn film deposited from a DMSO solution by López and Choi. Adapted with permission from ref 208. Copyright 2005 The Royal Society of Chemistry. (d) Fern-shaped Zn film deposited from a formamide solution by López and Choi. Adapted with permission from ref 209. Copyright 2006 American Chemical Society.

Steinmiller and Choi adopted this method and deposited mesoporous ZnO films with ordered lamellar structures using an interfacial surfactant templating method.207 In this study surfactants such as sodium dodecyl sulfate (SDS) and lauric acid (LA) were added to the plating solution, and their selfassembly on the WE during ZnO deposition served as a P

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Figure 22. (a) Schematic representation of Co−Pi OECs on ZnO. (b) Band diagram of ZnO and redox potentials of water and Co2+ showing the feasibility of anodically photodepositing Co−Pi PECs on ZnO. SEM images of Co−Pi OECs on ZnO rods when (c) photodeposition and (d) electrodeposition were used. Adapted with permission from ref 210. Copyright 2014 National Academy of Sciences.

Figure 23. (a) SEM images of In(OH)3 nanorods and (b) In(OH)3 nanocubes prepared by Gan et al. Adapted with permission from ref 215. Copyright 2014 The Royal Society of Chemistry. SEM images of InSnOH films deposited (c) without and (d) with KNO3 as an additive by Kovtyukhova and Mallouk. Adapted with permission from ref 216. Copyright 2010 American Chemical Society.

CBM (ca. −0.2 V vs RHE) is favorably located for water splitting.211−213 However, In2O3 may be used to form a composite structure with a narrower band-gap photon absorber. For example, Li et al. utilized an electrodeposition method to prepare a photoelectrode consisting of In2O3/In2S3 core−shell nanocubes.214 For this structure, In(OH)3 was first prepared by electrochemical reduction of nitrate in a In(NO3)3 solution at 70 °C (eq 29). As nitrate reduction increased the local pH at the WE, In3+ ions were deposited as In(OH)3 nanocubes with an average size of 300 nm. The In(OH)3 nanocubes were converted to In2O3 by annealing at 350 °C in air and retained the original morphology. The In2O3 nanocubes showed a relatively narrow direct optical band gap of 2.83 eV. The resulting In2O3 was converted to an In2O3@In2S3 heterojunction by hydrothermally growing In2S3 (Eg = 2.24 eV) on the In2O3 substrate, which was then studied for photocurrent generation. Gan et al. showed that In(OH)3 crystals can be deposited with different shapes, such as nanorods, nanosheets, and nanocubes, by varying the concentration of In3+ and NH4Cl used as an additive and other deposition conditions (Figure 23a and 23b).215 The In(OH)3 crystals with varying shapes could be converted to In2O3 crystals, preserving the original shapes by a heat treatment at 400 °C. Tin-doped indium oxide (ITO) with a Sn/In atomic ratio of 0−0.1 has also been prepared by electrodeposition by Kovtyukhova and Mallouk.216 ITO is a transparent conducting oxide that is often used as a substrate to deposit photoelectrodes. Therefore, although ITO itself will not be used as a photoelectrode, electrodeposition methods to produce high surface area ITO films can be very useful. In general, 2D thin films of ITO are most commonly prepared by physical vapor deposition, electron beam evaporation, or various sputtering deposition methods. ITO films by Kovtyukhova and Mallouk were prepared by cathodic deposition of crystalline Sn-doped In(OH)3 films followed by thermal oxidation at 300 °C. The

acidic aqueous plating solution (pH 2.3) contained In(NO3)3· xH2O, SnCl4·xH2O, and HNO3 with KNO3 or NaCl added as a supporting electrolyte to affect the morphology of the films. The authors proposed that when SnCl4 was added to a solution containing [In(H2O)6]3+, bi- or polynuclear indium−tin complexes are formed in solution as shown in eq 46 where 1 ≤ n ≤ 4m and m ≥ 1. SnCl4(s) + H 2O + [In(H 2O)6 ]3 + → [[(HO)4 − n Sn(μ‐O(OH))]m In n(H 2O)5n ]2n + + H+ + Cl−

(46)

Cathodic deposition was carried out at 79−80 °C. When the local pH at the WE was raised either by reduction of nitrate (eq 29) or by dissolved O2 (eq 40) in the plating solution, Sndoped In(OH)3 was deposited. Without the presence of KNO3 or NaCl, Sn-doped In(OH)3 was deposited as 2D sheets, but with the presence of KNO3 or NaCl, bundles of rods were obtained (Figure 23c and 23d). These films were converted to ITO by annealing and maintained their original morphologies. The resistivity of the ITO films decreased gradually as the tin content increased and reached 0.9 Ω·cm when the Sn/In atomic ratio was ca. 0.1. The transmittance of this film was up to 84%, which is comparable to that of typical ITO films prepared by vapor deposition methods. 4.1.5. SnO2. SnO2 is an n-type semiconductor that has a band gap of 3.5−3.8 eV and a CBM position of 0.19 V vs NHE.217−219 Like In2O3, it may be used as a substrate to deposit narrower band-gap semiconductors. Crystalline SnO2 can be directly deposited by nitrate reduction (eq 29) using a plating solution composed of tin dichloride, sodium nitrate, and nitric acid at 85 °C, as demonstrated by Chang et al.220 The asdeposited film was composed of crystalline tetragonal SnO2 particle aggregates as confirmed by XRD and SEM. The fact that SnO2 was cathodically deposited from a solution Q

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conformally deposit high-quality protection layers in a more practical manner, would be highly beneficial. Electrodeposition of TiO2 films has been achieved either by anodic deposition or by cathodic deposition. (Electrochemical production of TiO2 by anodization of Ti metal is not included in this review because many reviews are available for this subject.) The first anodic deposition condition for TiO2 reported by Kavan et al.230 used acidic aqueous TiCl3 solutions (pH 1.5−3.1). The Ti3+ ions were present as TiOH2+ in the solution (eq 47) and oxidized to Ti4+ during anodic deposition at room temperature to form an amorphous hydrous film, which was noted as Ti4+ polymers by the authors (eq 48).230 The as-deposited film could be converted to crystalline TiO2 films by heating at ≥400 °C (Figure 25a).

containing SnCl2 suggests that the dissolution of SnCl2 in the nitric acid solution generated Sn4+ ions. The as-deposited film was annealed at 200−400 °C to improve its crystallinity. Spray and Choi used the same deposition mechanism to prepare mesoporous SnO2 films. They added 2 wt % of anionic surfactant SDS to the plating solution to utilize self-assembly of SDS formed on the WE as a template to deposit mesoporous SnO2 films (Figure 24a).221 The resulting SnO2 film possessed

Ti 3 +(aq) + H 2O → TiOH2 + + H+

(47)

TiOH2 + + e− → Ti4 + polymers

(48)

Figure 24. (a) Schematic representation for using surface selfassembly of SDS−Sn4+ to deposit mesoporous SnO2 films. (b) HRTEM image of the mesoporous SnO2 film prepared by Spray and Choi. Adapted with permission from ref 221. Copyright 2007 The Royal Society of Chemistry. SEM images of Sn films with (c) cauliflower and (d) wire morphology prepared by Santato et al. Adapted with permission from ref 222. Copyright 2007 Elsevier.

a 3D worm-like mesoporous structure (Figure 24b). The surface area and the mean pore size of the resulting SnO2 films were 163 m2/g and 2.8 nm, respectively, which were analyzed by N2 adsorption−desorption measurements. SnO2 electrodes have also been prepared by cathodic deposition of Sn metal followed by thermal oxidation, which was reported by Santato et al.222 In this study, high surface area Sn metal films with cauliflower-like and wire morphologies were deposited by reduction of Sn2+ to Sn0 in a nonaqueous DMSO solution containing SnCl2 and NaNO3 (Figure 24c and 24d). These films were converted to crystalline SnO2 electrodes by annealing at 550 °C and retained the original morphologies. 4.1.6. TiO2. Titanium dioxide, which is an n-type semiconductor with a band gap of 3.0−3.2 eV, is undoubtedly one of the most important oxides used in photoelectrochemistry.223,224 It has been investigated for use in photocatalysis, dyesensitized solar cells, and water-splitting PECs.3,223−226 In fact, the first water-splitting PEC cell reported by Honda and Fujishima in 1972 used TiO2 as the photoanode.1 Due to its excellent photostability and chemical stability, both in acidic and in basic solutions, it has also been utilized as protection or passivation layers for photoelectrodes (e.g., Cu2O, Si, GaAs, GaP, etc.) that are not chemically or photochemically stable in aqueous media.227−229 TiO2 films used as protection layers, which need to be thin and conformally coated, were typically deposited using atomic layer deposition (ALD).227−229 Developing inexpensive solution-based methods, which can

Figure 25. (a) SEM image of anodically deposited TiO2 by Kavan et al. Adapted with permission from ref 230. Copyright 1993 Elsevier. (b) Side-view SEM image of W:BiVO4 after an amorphous TiO2 layer was electrodeposited by Eisenberg et al. Adapted with permission from ref 234. Copyright 2014 American Chemical Society. (c) SEM image of anodically deposited crystalline TiO2 by Wessels et al. Adapted with permission from ref 235. Copyright 2006 The Electrochemical Society. (d) SEM photograph of cathodically deposited TiO2 by Natarajan et al. after a heat treatment at 400 °C. Adapted with permission from ref 236. Copyright 2014 The Electrochemical Society.

When porous alumina membranes were used as the WE, TiO2 nanowires could be obtained after annealing the asdeposited films to form crystalline TiO2 (anatase), followed by the removal of the alumina template in an alkaline solution.231−233 Eisenberg et al. used the same deposition method to place a thin layer of amorphous TiO2 (80−120 nm thick deposited for 15−30 s) on W-doped BiVO4 films prepared as photoanodes (Figure 25b), which resulted in a significant photocurrent enhancement as well as a shift of the photocurrent onset potential to the negative direction.234 The authors explained that the enhancement was partly due to the TiO2 layer passivating the FTO surface not completely covered by Wdoped BiVO4, which minimized back reduction of photoR

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4.2. Ternary Oxide Photoanodes

oxidized intermediates on the conducting substrate. It was also partly due to the TiO2 layer passivating defect sites on the surface of W-doped BiVO4 films, which can serve as recombination centers. The authors also noted that another possible contribution might be improved charge collection owing to band bending at the TiO2−BiVO4 interface. While the aforementioned anodic condition resulted in amorphous Ti(IV)−O films as deposited, Wessels et al. reported an anodic deposition condition that could produce crystalline TiO2 films as deposited.235 An acidic aqueous solution of TiCl3 (pH 2.5) was used as the plating solution with SDS added to control the morphology of the deposits, and anodic deposition was carried out at 80 °C. XRD study showed the presence of multiple crystalline phases, a rutile phase as a major phase and an anatase as a minor phase. After calcination at 450 °C, anatase peaks disappeared, while the rutile peaks became sharper (Figure 25c). In addition to anodic conditions, a few cathodic deposition conditions have been reported for the preparation of TiO2 films. The first method reported by Natarajan et al. dissolved Ti powder in a solution containing H2O2 and ammonia. In this basic solution, H2O2 was deprotonated to form 2HO2−, which reacted with Ti to form TiO2+ species (eq 49). Further reaction of TiO2+ with OH− and water led to the formation of monomeric or oligomeric forms of Ti4+ species observed as a yellow-colored gel (eq 50).236 Ti + 2HO2− + H 2O → TiO2 + + 4OH−

While the studies on binary compounds provided us with various important insights and understanding of water-splitting PECs, most of these compounds have critical limitations for further improving their efficiencies of solar water splitting. Therefore, more recent studies have focused on examining various ternary oxides to identify more promising photoanodes for solar water splitting. Most of these oxides had been already tested as powder-type photocatalysts, but their potentials as photoanodes for solar water oxidation have not been accurately evaluated. The challenge is that producing high-quality photoelectrodes (e.g., high purity, good contact) of ternary compounds can be more difficult than simpler binary compounds because more thermodynamically stable binary compounds can easily be formed as impurities when hightemperature methods are used. When ternary oxides are first synthesized as powder-type samples by low-temperature solution methods and then prepared as electrodes by casting them on a conducting substrate, the contact between the particles or between the particles and the conducting substrates may not be sufficient to properly evaluate charge transport properties or charge collection efficiencies. Electrodeposition can produce ternary oxides as high-quality electrodes directly on the conducting substrate with good electrical continuity using mild solution synthesis conditions. When electrodeposition is coupled with various postdeposition treatments, the number of compounds that can be electrochemically produced can be further extended. In this section, various electrodeposition-based synthesis strategies for ternary oxide photoanodes are reviewed. The methods introduced here can easily be modified to produce more diverse ternary oxides or even more complex quaternary oxide electrodes. 4.2.1. BiVO4. BiVO4 is an n-type semiconductor that has recently emerged as one of the most promising photoanodes for use in water-splitting PECs.239,240 It can absorb a substantial portion of the visible spectrum (Eg = 2.4 eV) and has a favorable CBM position very near the thermodynamic H2 evolution potential, which allows BiVO4 to achieve more than 1 V of photovoltage for water oxidation (i.e., the potential difference between thermodynamic water oxidation potential, E° = 1.23 V, and photocurrent onset for water oxidation).239−241 Due to these favorable features, the most recent exciting developments in water-splitting PECs utilized BiVO4 as the photoanode.242−247 BiVO4 films have been produced electrochemically by anodization, anodic deposition, and cathodic deposition. The anodization method used to produce BiVO4 involved oxidation of a Bi metal WE in an aqueous solution (pH 7) containing NH4VO3 as the vanadium precursor. Upon oxidation of Bi metal to Bi3+, Bi3+ reacted with the VO3− in the solution to form an amorphous Bi−V−O layer, which was converted to crystalline BiVO4 films by annealing at 450 °C.248,249 Anodic electrodeposition of BiVO4 was first reported by Seabold and Choi in 2012, who used a plating solution containing Bi(NO3)3 and VOSO4 as Bi3+ and V5+ sources, respectively.250 During anodic deposition, VO2+ ions were oxidized to V5+ ions (eq 54), which coprecipitated with Bi3+ to form an amorphous Bi−V−O film (eq 55) (Figure 26a).

(49)

TiO2 + + 2OH− + x H 2O → TiO(OH)2 ·x H 2O (yellow gel)

(50)

TiO(OH)2 ·x H2O (yellow gel) + H 2SO4 → TiO(SO4 ) + (2 + x)H 2O

(51)

When the yellow gel was dissolved in concentrated H2SO4 solution, a red-colored solution containing TiO(SO4) was formed (eq 51), which was used as the plating solution after an aqueous solution of KNO3 was added. During cathodic deposition, nitrate was reduced to form OH− ions (eq 29), and the local pH increase at the WE surface resulted in the deposition of a TiO(OH)2·xH2O gel film. The as-deposited film was converted to nanocrystalline TiO2 films after heat treatment at ≥400 °C (Figure 25d). The cathodic deposition method reported by Matsumoto et al. used a (NH4)2[TiO(C2O4)2] solution (pH 4) and utilized reduction of water to OH− (eq 10) to induce deposition of amorphous TiO2 films (eq 52).237 TiO(C2O4 )2 2 − + 2OH− → TiO2 + H 2O + 2(C2O4 )2 − (52)

Alternatively, Lokhande et al. reported cathodic deposition of amorphous TiO2 films from an aqueous alkaline solution of Ti4+ where bulk precipitated TiO2 powders were in equilibrium with solution Ti(OH)4.238 The authors claimed that during cathodic deposition, OH− in Ti(OH)4 was reduced, resulting in the deposition of TiO2 (eq 53). However, it should be also possible that water, instead of OH− in Ti(OH)4, was reduced to H2, generating OH− (eq 10), which decreased the solubility of Ti(OH)4 and resulted in the deposition of TiO2. Ti(OH)4 + 2e− → TiO2 + H 2 + 2OH−

(53) S

VO2 + + H 2O → VO2+ + e− + 2H+

(54)

Bi 3 + + VO2+ + 2H 2O → BiVO4 + 4H+

(55)

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The plating solution used for the synthesis of BiOI films contained Bi(NO3)3 and KI as the Bi and I source, respectively, which formed [BiI4]− species in solution. From this solution BiOI films can be deposited when OH− is generated electrochemically at the WE by nitrate reduction for example (eqs 29 and 56). [BiI4]− + 2OH− → BiOI + H 2O + 3I−

(56)

However, the authors explained that nitrate reduction and other typically used reduction reactions that can generate OH− could not be used in this case because Bi3+ is more easily reducible to Bi (eq 57). McDonald and Choi overcame this issue by exploiting electrochemical reduction of p-benzoquinone (eq 58), which has a much more positive reduction potential than Bi3+,253 to increase the local pH at the WE. As a result, BiOI could be deposited without codepositing Bi metal as impurities. Figure 26. SEM images of (a) anodically as-deposited Bi−V−O film and (b) crystalline BiVO4 film obtained by annealing by Seabold and Choi (scale bar, 250 nm). Adapted with permission from ref 250. Copyright 2012 American Chemical Society. SEM images of (c) asdeposited BiOI and (d) top-view and (e) side view of BiVO4 electrode obtained from BiOI. Adapted with permission from refs 252 and 255. Copyright 2012 and 2014 The Royal Society of Chemistry and American Association for the Advancement of Science.

Bi 3 + + 3e− → Bi E° = 0.308 V

(57)

In order to convert BiOI to BiVO4, excess V2O5 was dissolved in a NH4OH solution and placed on the BiOI film before annealing at 520 °C in air. During the annealing step, each 2D plate of BiOI was converted to a round particle of BiVO4. (Unlike BiOI, BiVO4 has a 3D crystal structure.) Since the morphology of the BiOI film possessed ample space between the extremely thin 2D BiOI plates (Figure 26c), the grain growth of BiVO4 could be effectively suppressed, resulting in the formation of a porous film composed of submicrometersized BiVO4 particles. After BiVO4 was formed, the unreacted V2O5 was removed by dissolution in 1 M NaOH, leaving a pure BiVO4 film. Using excess V2O5 and removing unreacted V2O5 ensured more uniform formation of BiVO4 than using a stoichiometric amount of V2O5, since the conversion had to occur by a solid-state reaction where the diffusion of ions is not as easy as in solution reactions. The resulting BiVO4 film showed a slight increase in photocurrent generation compared to anodically prepared undoped BiVO4, but a significant increase in electron−hole separation was not observed. This is because although the films were porous, the size of individual BiVO4 particles in the film were too large compared to the hole diffusion length of BiVO4, which is about 100 nm (Figure 26d and 26e).254 On the basis of this result, Kim and Choi made a few modifications to the method developed by McDonald and Choi, which enabled the production of nanoporous BiVO4 that is composed of much smaller particles (average diameter, ca. 76 nm), achieving a surface area of 31.8 m2/g.255 The major change was the replacement of the solution used as the V source for the conversion of BiOI to BiVO4. Kim et al. observed that the surface of BiOI film is highly hydrophobic; thus, the aqueous vanadium solution used in the study by McDonald and Choi could not easily penetrate into the BiVO4 films. As a result, the conversion of BiOI to BiVO4 was initiated only at the top of the film surface, and the grains of BiVO4 grew toward the bottom of the film by slow solid-state diffusion, resulting in the formation of large particles of BiVO4 and no porosity in that direction (i.e., perpendicular to the substrate), limiting the

However, based on EDS composition analysis, the authors stated that the as-deposited phase was most likely not BiVO4 but Bi4V6O21, which is reported to be a member of the Bi−V− O family composed of Bi3+ and V5+ ions. The as-deposited film was annealed at 500 °C to form crystalline BiVO4 and amorphous V2O5. The amorphous V2O5 could be dissolved in 1 M KOH to leave a pure BiVO4 film (Figure 26b). When combined with FeOOH as an OEC, the resulting BiVO4/FeOOH achieved 1 mA/cm2 at ca. 0.6 V vs RHE with a photocurrent onset potential as positive as ca. 0.2 V.250 This result demonstrated the unique ability of BiVO4 to generate more than 1 V of photovoltage for water oxidation while being able to utilize visible light to generate sufficient photocurrent in the low-bias region (650 °C). Second, electrodeposited silicon did not reach the high purity required for the electronic or PV industry.

Figure 44. (a) Schematic depiction of an electrochemical liquid− liquid−solid electrodeposition process yielding crystalline silicon from liquid Ga electrode proposed by Gu et al. (b) SEM image of crystalline Si obtained from liquid Ga electrode. Adapted with permission from ref 420. Copyright 2013 American Chemical Society. (c) SEM image of crystalline Si electrodeposited from CaCl2 molten salts by Cho et al. (d) J−V plots of the electrodeposited p-Si electrode (solid line) and commercial silicon wafer (dashed line) obtained for photoreduction of ethyl viologen in acetonitrile under 1 sun illumination. Adapted with permission from ref 421. Copyright 2012 Wiley-VCH.

pool as a WE, crystalline silicon was successfully electrodeposited by reduction of SiCl4 in a propylene carbonate solution at 80 °C. In another study, Cho et al. reported the first photoactive silicon produced by electrodeposition.421 In this study, they used SiO2 nanoparticles in a CaCl2 molten salt (850 °C) as the Si precursor. The as-deposited Si was crystalline (Figure 44c), as expected for those obtained from hightemperature molten salts systems, and it showed p-type photocurrent for the reduction of ethyl viologen in an acetonitrile solution, although the amount of photocurrent generated was not as high as those obtained from commercially available Si electrodes (Figure 44d). These results encourage further development in Si electrodeposition to produce more photoactive, crystalline Si photoelectrodes. 5.5. III−V Semiconductors

III−V semiconductors, such as GaAs, GaP, InP, and their solid solutions (e.g., GaInP2), have played a vital role in applications involving solar energy conversion. Like Si, the doping type (n or p) of III−V semiconductors can be systematically controlled. This makes it possible to use these compounds as photocathodes when p-doped and as photoanodes when n-doped. The band-edge positions of III−V compounds as well as Si in pH 0 solution along with their band gaps are shown in Figure 45.422,423 This figure shows that when p-doped, Si and III−V semiconductors can serve as photocathodes in pH 0 solution for water reduction. In general, the band-edge positions of oxides against the redox potentials of water (i.e., H2 evolution and O2 evolution) are fixed regardless of the pH since both the redox potentials and the band-edge positions of oxides shift AG

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conditions; several studies reported that as-deposited films are crystalline,431−435 while the rest reported that as-deposited films were amorphous or poorly crystalline and needed a postdeposition annealing step at 300 °C under N2 atmosphere to obtain crystalline GaAs.428,436 Instead of codepositing Ga and As, sequential deposition of Ga and As layers in aqueous solutions has also been used at room temperature.437,438 In these studies, Ga was deposited in 5 M KOH, and As was deposited in 7 M HCl sequentially. The as-deposited Ga−As bilayer films were annealed in a vacuum furnace at 80−190 °C437 or under N2 atmosphere at 450 °C438 to obtain crystalline GaAs films. In addition to bulk GaAs films, a high-quality monolayer thick GaAs has also been deposited by Villegas and Stickney using electrochemical atomic layer epitaxy (ECALE) methods.439 In this study, specifically oriented monolayers of GaAs could be obtained depending on the orientation of the substrate used (e.g., GaAs (100) on Au (100) and GaAs (110) on Au (110). More recently, Fahrenkrug et al. demonstrated the preparation of crystalline GaAs in an alkaline aqueous solution containing As2O3 using a liquid gallium WE at 90 °C, which is a relatively mild condition to produce crystalline GaAs films.440 In their electrochemical liquid−liquid−solid process, HAsO32−, which was obtained by dissolution of As2O3 in an aqueous alkaline solution (pH 13), was reduced at the interface of Ga(l)/electrolyte to form GaAs crystals (Figure 46). In this system, the liquid gallium served as both the WE and the Ga precursor.

Figure 45. Conduction band and valence band positions for GaP, GaAs, InP, and Si at pH 0.

equal amounts, following Nernstian behavior (−59 mV/pH) when pH increases. However, unlike metal oxides, the bandedge positions of III−V semiconductors and Si do not show a systematic Nernstian shift when the pH changes. As a result, their relative band-edge positions against water oxidation and reduction potentials vary depending on the pH of the solutions. For this reason, when n-doped, Si and III−V semiconductors can serve as photoanodes for water oxidation in alkaline solutions where the water oxidation potential is raised above their VBM.229 There have been various attempts to electrodeposit III−V electrodes, which can significantly reduce production cost and simplify preparation processes. The fact that toxic gas phases, such as arsenic compounds (e.g., AsH3(g) and AsCl3(g)), which are precursors for traditional III−V synthesis, are not used for solution-based electrodeposition methods can be an additional advantage for using electrodeposition to produce III−V semiconductors. Electrodeposited III−V electrodes have not yet demonstrated notable performances for photoelectrochemical water splitting. However, considering that III−V electrodes produced by other methods have been extensively studied for photoelectrochemical water splitting and that there are clear advantages to producing III−V electrodes by electrodeposition, a brief review on the electrodeposition methods developed to produce III−V electrodes to date is provided in this section. GaAs has been deposited by reduction of Ga3+ ions and As3+ species (e.g., H3AsO3, HAsO32−)424 using various types of plating solutions. Earlier studies used molten salts of B2O3/ NaF/Ga2O3/NaAsO2 as the plating solution to deposit thin films of crystalline GaAs at 720−760 °C.425 When molten salts of KCl/GaCl3/AsI3 were used, the deposition temperature could be lowered to 300 °C.426 Room-temperature deposition of GaAs has been achieved when ionic liquids or aqueous media are used as plating solutions. Carpenter and Verbrugge used AlCl3/1-methyl-3ethylimidazolium chloride (ImCl) ionic liquid containing GaCl3 and AsCl3 as the plating solution.427 During the cathodic deposition, Ga and As metals were deposited from solution GaCl4− and AsCl4− species, leading to the formation of GaAs, which was confirmed by XPS. Reports on aqueous deposition mostly used strongly acidic plating solutions (0.7 ≤ pH < 3) where Ga metal or GaCl3 was dissolved as the Ga source and AsCl3, As2O3, or As2O5 was used as the As source.428−436 A basic plating solution containing GaCl3 and As2O3 has also been used, but poor adhesion of GaAs to the WE was reported.428 The crystallinity of GaAs films deposited from aqueous media appears to vary significantly depending on the deposition

Figure 46. Photos showing the formation of GaAs from a liquid Ga pool over time demonstrated by Fahrenkrug et al. Adapted with permission from ref 440. Copyright 2012 American Chemical Society.

Several authors reported electrochemical and photoelectrochemical studies of electrodeposited GaAs.430,432,434,435 In these reports, electrodeposited GaAs showed only n-type conductivity, and their photoelectrochemical properties were tested using hole acceptors such as sulfides. Further development will be necessary to electrodeposit p-GaAs relevant for solar hydrogen production. Unlike GaAs, only a few studies have been reported for the electrodeposition of GaP, which has an indirect band gap of 2.2−2.3 eV. All reported methods involved the use of molten salt systems such as NaF/NaPO3/Ga2O3 at 750−900 °C.441,442 Two possible deposition mechanisms to form GaP suggested by Mattei et al. are shown in eqs 89 and 90.442 Electrochemical or photoelectrochemical properties of the resulting GaP films were not investigated. Ga 3 + + 4PO3− + 8e− → GaP + 3PO4 3 − AH

(89)

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6. SUMMARY Electrodeposition combines electrochemistry and materials synthesis in a way that allows for the functionality of deposited materials to be controlled by atomic level electrochemical reactions. While electrodeposition has long been known as a synthesis method for metal plating, various systems reviewed in this study clearly demonstrate the versatility of electrodeposition in materials synthesis. While utilizing solutionbased electrochemical reactions that are fundamentally very simple, electrodeposition allows for the production of numerous semiconductors and catalysts as electrode-type materials with desired morphologies. The synthesis versatility together with the practicality of electrodeposition makes this method optimal for the construction of electrodes for use in energy-related applications where low cost for large-scale production is critical. For most of the material types discussed here, the use of electrodeposition is still in its infancy. For example, the types of ternary or quaternary oxides and chalcogenides that have been electrodeposited have been very limited. However, combination or modification of the methods discussed in this review can enable electrodeposition of more diverse and complex oxide and chalcogenide-based photoelectrodes and catalysts. Also, the recent development on the electrodeposition of Si and III−V semiconductors suggests that electrodeposition may be a viable synthesis method to produce these materials as photoactive semiconductor electrodes. While this review discusses electrodeposition methods used for the synthesis of materials relevant to PEC constructions, we believe that the methods included here will also be useful and applicable when producing electrode materials for other applications. Electrodeposition as a synthetic technique will continue to grow as the types and functionality of materials that can be electrodeposited expand.

(90)

InP (Eg = 1.35 eV) and its solid solution with GaP, GaInP2 (Eg = 1.83 eV), have much narrower band gaps than GaP. As a result, p-type InP and p-type GaInP2 have been intensively studied as photocathodes for water-splitting PECs.443−448 ptype InP can generate a substantial amount of photocurrent (25−30 mA/cm2 at 0 V vs RHE) for water reduction in a strongly acidic solution (e.g., 1 M of HClO4). Although p-type InP is also known to suffer from photocorrosion in acidic solution (eqs 91 and 92),449,450 recent studies demonstrated enhanced stability of InP after it was combined with a protection layer of TiO2 and HECs such as Ru, Pt, and MoS3.451−454 The InP photoelectrodes used in these studies were either commercially available InP wafers, which are produced by the Czochralski process, or InP films prepared by a metal organic vapor-phase epitaxy system (MOVPE)453 or electron beam evaporation, followed by phosphorization.454 InP + 3e− + 3H+ → In 0 + PH3

(91)

InP + 3e− + 3H+ + 4H 2O → In 0 + H3PO4 + 4H 2 (92)

Electrodeposition of InP has been also studied by several groups. In early studies, a eutectic melt of In2O3/NaPO3/ KPO3/NaF/KF salts was used at 600 °C.455 The proposed reaction mechanism is as given in eqs 93 and 94. Cathode: In 3 + + 4PO3− + 8e− → InP + 3PO4 3 −

(93)

Anode: 4PO4 3 − → 4PO3− + 2O2 + 8e−

(94)

Sahu et al. demonstrated InP electrodeposition in both aqueous and nonaqueous media using InCl3 and NH4PF6 as In and P precursors, respectively.456−458 In both aqueous (pH 2) or nonaqueous dimethylformamide media, InP films were cathodically electrodeposited at room temperature. XRD studies revealed that as-deposited films from both solutions were crystalline InP. Vacuum annealing at 300 °C could be used to improve their crystallinity further. The resulting films were reported to show an n-type conductivity, but no photoelectrochemical properties were reported. Another established method involves two steps: electrodeposition of In metal followed by postdeposition phosphorization.459−461 After In metal was cathodically electrodeposited (eq 76) on various substrates, phosphorization was achieved by annealing the In films with red phosphorus (vacuum, 400 °C),460 PH3 (H2 flowing, 450 °C),461 or 10% of PH3 in H2 (H2 flowing, 750 °C).459 These two-step phosphorization methods typically produced n-type InP.460,461 However, Musiani et al. showed that when a thin layer of Zn is deposited on the surface of In films before phosphorization, p-type InP could also be generated.461 In the same study, the authors tested the photoelectrochemical properties of electrochemically prepared n-InP and p-InP using a I−/I3− redox couple in acidic solution to find that both n-InP and p-InP generated a significant amount of photocurrent (i.e., up to 20 mA/cm2 for n-InP and ca. −10 mA/cm2 for p-InP). These highly photoactive electrodeposited InP films encourage further studies on developing electrochemical methods to produce high-quality III−V photoelectrodes for use in water-splitting PECs in a more cost-effective manner.

AUTHOR INFORMATION Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest. Biographies

Donghyeon Kang received his B.S. degree from Sogang University (South Korea) in 2008, summa cum laude with honors in Chemisty, and his M.S. degree from the same university in 2010, receiving the Outstanding Inorganic Chemistry Graduate Student Award. He has been a Ph.D. student under the supervision of Professor Kyoung-Shin Choi in the Department of Chemistry at the University of AI

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WisconsinMadison since 2011. His research focuses on electrochemical synthesis of semiconductor electrodes and electrocatalysts and their integrations for use in water-splitting photoelectrochemical cells.

Allison C. Cardiel received her B.A. degree in Chemistry from Carleton College in 2013, graduating summa cum laude and with distinction in her major. She received her M.S. degrees in Materials Chemistry and Environmental Studies from the University of WisconsinMadison in 2015. She is currently in her third year of graduate studies in the Materials Chemistry program at the University of WisconsinMadison, working in Dr. Kyoung-Shin Choi’s group. Her current research focuses on electrochemical synthesis of oxidebased electrocatalysts for various oxidation reactions.

Tae Woo Kim earned his Ph.D. degree in the Department of Materials Science and Engineering at Yonsei University (South Korea) in 2011 (Supervisor Prof. Sang-Hoon Hyun). He then worked as a postdoctoral researcher at Ewha Womans University in Seoul during 2011−2012 (Supervisor Prof. Seong-Ju Hwang). In 2012 he joined Prof. Kyoung-Shin Choi’s lab at the University of Wisconsin Madison as a postdoctoral researcher and worked on the development of photoanodes for use in photoelectrochemical cells. He is currently a senior researcher in the Korea Institute of Energy Research (KIER). His research interests include the development of environmentally friendly metal oxide-based nanocomposites for use in Li secondary batteries, photocatalysts, gas adsorbent, and solar cells.

Hyun Gil Cha received his B.S. degree in Polymer Engineering from Pukyoung National University (South Korea) in 2005 and Ph.D. degree in Physical Chemistry from Sogang University (South Korea) under the supervision of Prof. Young Soo Kang in 2011. He joined Prof. Kyoung-Shin Choi’s lab at the University of Wisconsin Madison as a postdoctoral researcher and has been working on the design and construction of electrochemical and photoelectrochemical cells for biomass conversion combined with fuel production.

Stephen R. Kubota received his B.A. degree in Chemistry from Lewis & Clark College in 2013 with honors. He is currently a third year materials chemistry graduate student in Dr. Kyoung-Shin Choi’s research group at the University of WisconsinMadison. His current research involves the development of electrocatalysts for various reduction reactions including electrochemical conversion of biomassderived compounds into commercially and industrially useful

Kyoung-Shin Choi received her Ph.D. degree in Chemistry from Michigan State University (2000). After spending 2 years at the

chemicals. AJ

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WE XAS XPS XRD η10 mA

University of California, Santa Barbara, as a postdoctoral researcher, she joined the chemistry faculty at Purdue University as an assistant professor in 2002 and was promoted to an associate professor in 2008. She was a visiting scholar at the National Renewable Energy Laboratory (NREL) during the fall of 2008. In 2012, she joined the chemistry faculty at the University of WisconsinMadison as a full professor. Her research combines solid-state chemistry, electrochemistry, and materials chemistry in order to address materialsrelated issues of electrode materials for use in photoelectrochemical and electrochemical devices. Her research group develops electrodeposition mechanisms/conditions to produce a variety of electrode and catalyst materials for use in electrochemical and photoelectrochemical applications. She was a recipient of a 2006 Alfred P. Sloan Research Fellowship, the 2007 ACS ExxonMobil Faculty Fellowship in Solid-State Chemistry, and the 2010 Iota Sigma Pi Agnes Fay Morgan Research Award. She served as the 2014 Chair of the Gordon Research Conference: Electrodeposition and is currently an associate editor for Chemistry of Materials.

Φsep

working electrode X-ray absorption spectroscopy X-ray photoelectron spectroscopy X-ray diffraction overpotential required to achieve a current of 10 mA/ cm2 electron−hole separation efficiency

REFERENCES (1) Fujishima, A.; Honda, K. Electrochemical Photolysis of Water at a Semiconductor Electrode. Nature 1972, 238, 37−38. (2) Walter, M. G.; Warren, E. L.; McKone, J. R.; Boettcher, S. W.; Mi, Q.; Santori, E. A.; Lewis, N. S. Solar Water Splitting Cells. Chem. Rev. 2010, 110, 6446−6473. (3) Grätzel, M. Photoelectrochemical Cells. Nature 2001, 414, 338− 344. (4) U.S. Department of Energy, Hydrogen and Fuel Cells Program (2005): Hydrogen and Fuel Cell Activities, Progress, and Plans (www. hydrogen.energy.gov/congress_reports.html). (5) Therese, G. H. A.; Kamath, P. V. Electrochemical Synthesis of Metal Oxides and Hydroxides. Chem. Mater. 2000, 12, 1195−1204. (6) Choi, K.-S.; Jang, H. S.; McShane, C. M.; Read, C. G.; Seabold, J. A. Electrochemical Synthesis of Inorganic Polycrystalline Electrodes with Controlled Architectures. MRS Bull. 2010, 35, 753−760. (7) Zeng, M.; Li, Y. Recent Advances in Heterogeneous Electrocatalysts for the Hydrogen Evolution Reaction. J. Mater. Chem. A 2015, 3, 14942−14962. (8) Zou, X.; Zhang, Y. Noble Metal-Free Hydrogen Evolution Catalysts for Water Splitting. Chem. Soc. Rev. 2015, 44, 5148−5180. (9) Conway, B. E.; Salomon, M. Electrochemical Reaction Orders: Applications to the Hydrogen- and Oxygen-Evolution Reactions. Electrochim. Acta 1964, 9, 1599−1615. (10) Bockris, J. O. M.; Potter, E. C. The Mechanism of the Cathodic Hydrogen Evolution Reaction. J. Electrochem. Soc. 1952, 99, 169−186. (11) Trasatti, S. Work Function, Electronegativity, and Electrochemical Behaviour of Metals III. Electrolytic Hydrogen Evolution in Acid Solutions. J. Electroanal. Chem. Interfacial Electrochem. 1972, 39, 163−184. (12) Feltham, A. M.; Spiro, M. Platinized Platinum Electrodes. Chem. Rev. 1971, 71, 177−193. (13) Simonov, A. N.; Cherstiouk, O. V.; Vassiliev, S. Y.; Zaikovskii, V. I.; Filatov, A. Y.; Rudina, N. A.; Savinova, E. R.; Tsirlina, G. A. Potentiostatic Electrodeposition of Pt on GC and on HOPG at Low Loadings: Analysis of the Deposition Transients and the Structure of Pt Deposits. Electrochim. Acta 2014, 150, 279−289. (14) Shimazu, K.; Weisshaar, D.; Kuwana, T. Electrochemical Dispersion of Pt Microparticles on Glassy Carbon Electrodes. J. Electroanal. Chem. Interfacial Electrochem. 1987, 223, 223−234. (15) Lau, A. L. Y.; Hubbard, A. T. Study of the Kinetics of Electrochemical Reactions by Thin-Layer Voltammetry: III. Electroreduction of the Chloride Complexes of Platinum(II) and (IV). J. Electroanal. Chem. Interfacial Electrochem. 1970, 24, 237−249. (16) Yasin, H. M.; Denuault, G.; Pletcher, D. Studies of the Electrodeposition of Platinum Metal from a Hexachloroplatinic Acid Bath. J. Electroanal. Chem. 2009, 633, 327−332. (17) El Sawy, E. N.; Birss, V. I. A Comparative Study of the Electrodeposition of Nanoporous Ir and Pt Thin Films. J. Electrochem. Soc. 2013, 160, D386−D393. (18) Uosaki, K.; Ye, S.; Naohara, H.; Oda, Y.; Haba, T.; Kondo, T. Electrochemical Epitaxial Growth of a Pt(111) Phase on an Au(111) Electrode. J. Phys. Chem. B 1997, 101, 7566−7572. (19) Bicelli, L. P.; Bozzini, B.; Mele, C.; D’Urzo, L. A Review of Nanostructural Aspects of Metal Electrodeposition. Int. J. Electrochem. Sci. 2008, 3, 356−408. (20) Budevski, E.; Staikov, G.; Lorenz, W. J. Electrocrystallization: Nucleation and Growth Phenomena. Electrochim. Acta 2000, 45, 2559−2574.

ACKNOWLEDGMENTS This work was supported by the Division of Chemical Sciences, Geosciences, and Biosciences, Office of Basic Energy Sciences of the U.S. Department of Energy through Grant DESC0008707. ABBREVIATIONS AAO anodized aluminum oxide AFM atomic force microscopy ALD atomic layer deposition AM air mass AZO aluminum-doped zinc oxide CBM conduction band minimum CIGSe Cu(InxGa1−x)Se2 CV cyclic voltammetry CZTS Cu2ZnSnS4 DMSO dimethyl sulfoxide E° standard reduction potential Eg band gap EG ethylene glycol FE faradaic efficiency FTO fluorine-doped tin oxide h+ photogenerated holes Hads adsorbed hydrogen HEC hydrogen evolution catalyst HER hydrogen evolution reaction Hupd underpotentially deposited H ITO indium tin oxide J−V current density-potential J−t current density−time LA lactic acid NHE normal hydrogen electrode OEC oxygen evolution catalyst OER oxygen evolution reaction PEC photoelectrochemical cell PV photovoltaic RHE reversible hydrogen electrode SDS sodium dodecyl sulfate SEM scanning electron microscopy STH solar-to-hydrogen efficiency STM scanning tunneling microscopy TEM transmission electron microscopy VBM valence band maximum AK

DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX

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and Direct Current Electrodeposition. Chem. Mater. 2002, 14, 4595− 4602. (42) Yin, A. J.; Li, J.; Jian, W.; Bennett, A. J.; Xu, J. M. Fabrication of Highly Ordered Metallic Nanowire Arrays by Electrodeposition. Appl. Phys. Lett. 2001, 79, 1039−1041. (43) Marozzi, C. A.; Chialvo, A. C. Development of Electrode Morphologies of Interest in Electrocatalysis Part 2: Hydrogen Evolution Reaction on Macroporous Nickel Electrodes. Electrochim. Acta 2001, 46, 861−866. (44) Savadogo, O.; Carrier, F. The Hydrogen Evolution Reaction in Basic Medium on Iron Electrodeposited with Heteropolyacids. J. Appl. Electrochem. 1992, 22, 437−442. (45) Rojas, M.; Fan, C. L.; Miao, H. J.; Piron, D. L. Characterization of Hydrogen Evolution on Cobalt Electrodeposits in Water Electrolysis. J. Appl. Electrochem. 1992, 22, 1135−1141. (46) DeCarvalho, J.; Filho, G. T.; Avaca, L. A.; Gonzalez, E. R. Electrodeposits of Iron and Nickel-Iron for Hydrogen Evolution in Alkaline Solutions. Int. J. Hydrogen Energy 1989, 14, 161−165. (47) Lupi, C.; Dell’Era, A.; Pasquali, M. Nickel-Cobalt Electrodeposited Alloys for Hydrogen Evolution in Alkaline Media. Int. J. Hydrogen Energy 2009, 34, 2101−2106. (48) Correia, A. N.; Machado, S. A. S.; Avaca, L. A. Studies of the Hydrogen Evolution Reaction on Smooth Co and Electrodeposited Ni-Co Ultramicroelectrodes. Electrochem. Commun. 1999, 1, 600−604. (49) Solmaz, R.; Kardaş, G. Electrochemical Deposition and Characterization of NiFe Coatings as Electrocatalytic Materials for Alkaline Water Electrolysis. Electrochim. Acta 2009, 54, 3726−3734. (50) Navarro-Flores, E.; Chong, Z.; Omanovic, S. Characterization of Ni, NiMo, NiW and NiFe Electroactive Coatings as Electrocatalysts for Hydrogen Evolution in an Acidic Medium. J. Mol. Catal. A: Chem. 2005, 226, 179−197. (51) Jaksic, M. M. Hypo-Hyper-d-Electronic Interactive Nature of Interionic Synergism in Catalysis and Electrocatalysis for Hydrogen Reactions. Int. J. Hydrogen Energy 2001, 26, 559−578. (52) Highfield, J. G.; Claude, E.; Oguro, K. Electrocatalytic Synergism in Ni/Mo Cathodes for Hydrogen Evolution in Acidic Medium: A New Model. Electrochim. Acta 1999, 44, 2805−2814. (53) Podlaha, E. J.; Landolt, D. Induced Codeposition I. An Experimental Investigation of Ni-Mo Alloys. J. Electrochem. Soc. 1996, 143, 885−892. (54) Podlaha, E. J.; Landolt, D. Induced Codeposition III. Molybdenum Alloys with Nickel, Cobalt, and Iron. J. Electrochem. Soc. 1997, 144, 1672−1680. (55) Krstajić, N. V.; Jović, V. D.; Gajić-Krstajić, L.; Jović, B. M.; Antozzi, A. L.; Martelli, G. N. Electrodeposition of Ni-Mo Alloy Coatings and Their Characterization as Cathodes for Hydrogen Evolution in Sodium Hydroxide Solution. Int. J. Hydrogen Energy 2008, 33, 3676−3687. (56) Hong, S. H.; Ahn, S. H.; Choi, J.; Kim, J. Y.; Kim, H. Y.; Kim, H.-J.; Jang, J. H.; Kim, H.; Kim, S.-K. High-Activity Electrodeposited NiW Catalysts for Hydrogen Evolution in Alkaline Water Electrolysis. Appl. Surf. Sci. 2015, 349, 629−635. (57) Sanches, L. S.; Domingues, S. H.; Marino, C. E. B.; Mascaro, L. H. Characterisation of Electrochemically Deposited Ni-Mo Alloy Coatings. Electrochem. Commun. 2004, 6, 543−548. (58) Lucatero, S.; Tamayo, G.; Crespo, D.; Mariño, E.; Videa, M. NiMo Nanoparticles Electrodeposited by Pulsed Current and Their Catalytic Properties for Hydrogen Production. J. New Mater. Electrochem. Syst. 2013, 16, 177−182. (59) Videa, M.; Crespo, D.; Casillas, G.; Zavala, G. Electrodeposition of Nickel-Molybdenum Nanoparticles for Their use as Electrocatalyst for the Hydrogen Evolution Reaction. J. New Mater. Electrochem. Syst. 2010, 13, 239−244. (60) Wang, Y.; Zhang, G.; Xu, W.; Wan, P.; Lu, Z.; Li, Y.; Sun, X. A 3D Nanoporous Ni-Mo Electrocatalyst with Negligible Overpotential for Alkaline Hydrogen Evolution. ChemElectroChem 2014, 1, 1138− 1144. (61) Vázquez-Gómez, L.; Cattarin, S.; Guerriero, P.; Musiani, M. Preparation and Electrochemical Characterization of Ni + RuO2

(21) Penner, R. M. Mesoscopic Metal Particles and Wires by Electrodeposition. J. Phys. Chem. B 2002, 106, 3339−3353. (22) Zoval, J. V.; Lee, J.; Gorer, S.; Penner, R. M. Electrochemical Preparation of Platinum Nanocrystallites with Size Selectivity on Basal Plane Oriented Graphite Surfaces. J. Phys. Chem. B 1998, 102, 1166− 1175. (23) Waibel, H.-F.; Kleinert, M.; Kibler, L. A.; Kolb, D. M. Initial Stages of Pt Deposition on Au(111) and Au(100). Electrochim. Acta 2002, 47, 1461−1467. (24) Zubimendi, J. L.; Vázquez, L.; Ocón, P.; Vara, J. M.; Triaca, W. E.; Salvarezza, R. C.; Arvia, A. J. Early Stages of Platinum Electrodeposition on Highly Oriented Pyrolitic Graphite: Scanning Tunneling Microscopy Imaging and Reaction Pathway. J. Phys. Chem. 1993, 97, 5095−5102. (25) Plyasova, L. M.; Molina, I. Y.; Gavrilov, A. N.; Cherepanova, S. V.; Cherstiouk, O. V.; Rudina, N. A.; Savinova, E. R.; Tsirlina, G. A. Electrodeposited Platinum Revisited: Tuning Nanostructure via the Deposition Potential. Electrochim. Acta 2006, 51, 4477−4488. (26) Liu, Y.; Gokcen, D.; Bertocci, U.; Moffat, T. P. Self-Terminating Growth of Platinum Films by Electrochemical Deposition. Science 2012, 338, 1327−1330. (27) Herrasti, P.; Diaz, R.; Ocón, P. Characterization of Poly-oToluidine. Electrodeposition of Platinum and Palladium into Polymer Films for Electrocatalysis. New J. Chem. 1993, 17, 279−285. (28) Kost, K. M.; Bartak, D. E.; Kazee, B.; Kuwana, T. Electrodeposition of Platinum Microparticles into Polyaniline Films with Electrocatalytic Applications. Anal. Chem. 1988, 60, 2379−2384. (29) Bartak, D. E.; Kazee, B.; Shimazu, K.; Kuwana, T. Electrodeposition and Characterization of Platinum Microparticles in Poly(4vinylpyridine) Film Electrodes. Anal. Chem. 1986, 58, 2756−2761. (30) Trueba, M.; Trasatti, S. P.; Trasatti, S. Electrocatalytic Activity for Hydrogen Evolution of Polypyrrole Films Modified with Noble Metal Particles. Mater. Chem. Phys. 2006, 98, 165−171. (31) Attard, G. S.; Bartlett, P. N.; Coleman, N. R. B.; Elliott, J. M.; Owen, J. R.; Wang, J. H. Mesoporous Platinum Films from Lyotropic Liquid Crystalline Phases. Science 1997, 278, 838−840. (32) Choi, K.-S.; McFarland, E. W.; Stucky, G. D. Electrocatalytic Properties of Thin Mesoporous Platinum Films Synthesized Utilizing Potential-Controlled Surfactant Assembly. Adv. Mater. 2003, 15, 2018−2021. (33) Bartlett, P. N.; Baumberg, J. J.; Birkin, P. R.; Ghanem, M. A.; Netti, M. C. Highly Ordered Macroporous Gold and Platinum Films Formed by Electrochemical Deposition through Templates Assembled from Submicron Diameter Monodisperse Polystyrene Spheres. Chem. Mater. 2002, 14, 2199−2208. (34) Liu, H.; He, P.; Li, Z.; Li, J. High Surface Area Nanoporous Platinum: Facile Fabrication and Electrocatalytic Activity. Nanotechnology 2006, 17, 2167−2173. (35) Pourbaix, M. Atlas of electrochemical equilibria in aqueous solutions, 2nd ed.; National Association of Corrosion Engineers: Houston, TX, 1974. (36) Epelboin, I.; Joussellin, M.; Wiart, R. Impedance Measurements for Nickel Deposition in Sulfate and Chloride Electrolytes. J. Electroanal. Chem. Interfacial Electrochem. 1981, 119, 61−71. (37) Epelboin, I.; Wiart, R. Mechanism of the Electrocrystallization of Nickel and Cobalt in Acidic Solution. J. Electrochem. Soc. 1971, 118, 1577−1582. (38) Saraby-Reintjes, A.; Fleischmann, M. Kinetics of Electrodeposition of Nickel from Watts Baths. Electrochim. Acta 1984, 29, 557−566. (39) Torabi, M.; Dolati, A. A Kinetic Study on the Electrodeposition of Nickel Nanostructure and Its Electrocatalytic Activity for Hydrogen Evolution Reaction. J. Appl. Electrochem. 2010, 40, 1941−1947. (40) Marozzi, C. A.; Chialvo, A. C. Development of Electrode Morphologies of Interest in Electrocatalysis. Part 1: Electrodeposited Porous Nickel Electrodes. Electrochim. Acta 2000, 45, 2111−2120. (41) Chu, S.-Z.; Wada, K.; Inoue, S.; Todoroki, S.-I. Fabrication and Characteristics of Ordered Ni Nanostructures on Glass by Anodization AL

DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

Composite Cathodes of Large Effective Area. Electrochim. Acta 2007, 52, 8055−8063. (62) Hu, C.-C.; Huang, Y.-H. Cyclic Voltammetric Deposition of Hydrous Ruthenium Oxide for Electrochemical Capacitors. J. Electrochem. Soc. 1999, 146, 2465−2471. (63) Jow, J.-J.; Lee, H.-J.; Chen, H.-R.; Wu, M.-S.; Wei, T.-Y. Anodic, Cathodic and Cyclic Voltammetric Deposition of Ruthenium Oxides from Aqeous RuCl3 Solution. Electrochim. Acta 2007, 52, 2625−2653. (64) Zhitomirsky, I.; Gal-Or, L. Ruthenium Oxide Deposits Prepared by Cathodic Electrosynthesis. Mater. Lett. 1997, 31, 155−159. (65) Hu, C.-C.; Liu, M.-J.; Chang, K.-H. Anodic Deposition of Hydrous Ruthenium Oxide for Supercapacitors. J. Power Sources 2007, 163, 1126−1131. (66) Tavares, A. C.; Tarasatti, S. Ni+RuO2 Co-Deposited Electrodes for Hydrogen Evolution. Electrochim. Acta 2000, 45, 4195−4202. (67) Iwakura, C.; Tanaka, M.; Nakamatsu, S.; Inoue, H.; Matsuoka, M.; Furukawa, N. Electrochemical Properties of Ni/(Ni + RuO2) Active Cathodes for Hydrogen Evolution in Chlor-Alkali Electrolysis. Electrochim. Acta 1995, 40, 977−982. (68) Miousse, D.; Lasia, A. Hydrogen Evolution Reaction on RuO2 Electrodes in Alkaline Solutions. J. New Mater. Electrochem. Syst. 1999, 2, 71−78. (69) Xiong, K.; Li, L.; Deng, Z.; Xia, M.; Chen, S.; Tan, S.; Peng, X.; Duan, C.; Wei, Z. RuO2 Loaded into Porous Ni as a Synergistic Catalyst for Hydrogen Production. RSC Adv. 2014, 4, 20521−20526. (70) Kodintsev, I. M.; Trasatti, S. Electrocatalysis of H2 Evolution on RuO2 + IrO2 Mixed Oxide Electrodes. Electrochim. Acta 1994, 39, 1803−1808. (71) Fachinotti, E.; Guerrini, E.; Tavares, A. C.; Trasatti, S. Electrocatalysis of H2 Evolution by Thermally Prepared Ruthenium Oxide. Effect of Precursors: Nitrate vs. Chloride. J. Electroanal. Chem. 2007, 600, 103−112. (72) Pecoraro, T. A.; Chianelli, R. R. Hydrodesulfurization Catalysis by Transition Metal Sulfides. J. Catal. 1981, 67, 430−445. (73) Laursen, A. B.; Kegnaes, S.; Dahl, S.; Chorkendorff, I. Molybdenum Sulfides-Efficient and Viable Materials for Electro- and Photoelctrocatalytic Hydrogen Evolution. Energy Environ. Sci. 2012, 5, 5577−5591. (74) Ponomarev, E. A.; Neumann-Spallart, M.; Hodes, G.; LévyClément, C. Electrochemical Deposition of MoS2 Thin Flms by Reduction of Tetrathiomolybdate. Thin Solid Films 1996, 280, 86−89. (75) Ghosh, S. K.; Bera, T.; Karacasu, O.; Swarnakar, A.; Buijnsters, J. G.; Celis, J. P. Nanostructured MoSx-Based Thin Films Obtained by Electrochemical Reduction. Electrochim. Acta 2011, 56, 2433−2442. (76) Albu-Yaron, A.; Lévy-Clément, C.; Katty, A.; Bastide, S.; Tenne, R. Influence of the Electrochemical Deposition Parameters on the Microstructure of MoS2 Thin Films. Thin Solid Films 2000, 361−362, 223−228. (77) Murugesan, S.; Akkineni, A.; Chou, B. P.; Glaz, M. S.; Bout, D. A. V.; Stevenson, K. J. Room Temperature Electrodeposition of Molybdenum Sulfide for Catalytic and Photoluminesence Applications. ACS Nano 2013, 7, 8199−8205. (78) Abbott, A. P.; McKenzie, K. J. Application of Ionic Liquids to the Electrodeposition of Metals. Phys. Chem. Chem. Phys. 2006, 8, 4265−4279. (79) Hinnemann, B.; Moses, P. G.; Bonde, J.; Jørgensen, K. P.; Nielsen, J. H.; Horch, S.; Chorkendorff, I.; Nørskov, J. K. Biomimetic Hydrogen Evolution: MoS2 Nanoparticles as Catalyst for Hydrogen Evolution. J. Am. Chem. Soc. 2005, 127, 5308−5309. (80) Jaramillo, T. F.; Jørgensen, K. P.; Bonde, J.; Nielsen, J. H.; Horch, S.; Chorkendoff, I. Identification of Active Edge Sites for Electrochemical H2 Evolution from MoS2 Nanocatalysts. Science 2007, 317, 100−102. (81) Merki, D.; Fierro, S.; Vrubel, H.; Hu, X. Amorphous Molybdenum Sulfide Films as Catalysts for Electrochemical Hydrogen Production in Water. Chem. Sci. 2011, 2, 1262−1267. (82) Vrubel, H.; Hu, X. Growth and Activation of an Amorphous Molybdenum Sulfide Hydrogen Evolving Catalyst. ACS Catal. 2013, 3, 2002−2011.

(83) Bélanger, D.; Laperrière, G.; Marsan, B. The Electrodeposition of Amorphous Molybdenum Sulfide. J. Electroanal. Chem. 1993, 347, 165−183. (84) Byskov, L. S.; Nørskov, J. K.; Clausen, B. S.; Topsøe, H. DFT Calculations of Unpromoted and Promoted MoS2-Based Hydrodesulfurization Catalysts. J. Catal. 1999, 187, 109−122. (85) Merki, D.; Vrubel, H.; Rovelli, L.; Fierro, S.; Hu, X. Fe, Co, and Ni Ions Promote the Catalytic Activity of Amorphous Molybdenum Sulfide Films for Hydrogen Evolution. Chem. Sci. 2012, 3, 2515−2525. (86) Gao, M.-R.; Lin, Z.-Y.; Zhuang, T.-T.; Jiang, J.; Xu, Y.-F.; Zheng, Y.-R.; Yu, S.-H. Mixed-Solution Synthesis of Sea Urchin-Like NiSe Nanofiber Assemblies as Economical Pt-Free Catalysts for Electrochemical H2 Production. J. Mater. Chem. 2012, 22, 13662−13668. (87) Kong, D.; Cha, J. J.; Wang, H.; Lee, H. R.; Cui, Y. First-Row Transition Metal Dichalcogenide Catalysts for Hydrogen Evolution Reaction. Energy Environ. Sci. 2013, 6, 3553−3558. (88) Kowalik, R.; Ż abiński, P.; Fitzner, K. Electrodeposition of ZnSe. Electrochim. Acta 2008, 53, 6184−6190. (89) Thanikaikarasan, S.; Mahalingam, T.; Sundaram, K.; Kathalingam, A.; Kim, Y. D.; Kim, T. Growth and Characterization of Electrosynthesized Iron Selenide Thin Films. Vacuum 2009, 83, 1066−1072. (90) Natarajan, C.; Sharon, M.; Lévy-Clément, C.; NeumannSpallart, M. Electrodeposition of Zinc Selenide. Thin Solid Films 1994, 237, 118−123. (91) Liu, F.; Wang, B.; Lai, Y.; Li, J.; Zhang, Z.; Liu, Y. Electrodeposition of Cobalt Selenide Thin Films. J. Electrochem. Soc. 2010, 157, D523−D527. (92) Zhang, Z.; Pang, S.; Xu, H.; Yang, Z.; Zhang, X.; Liu, Z.; Wang, X.; Zhou, X.; Dong, S.; Chen, X.; et al. Electrodeposition of Nanostructure Cobalt Selenide Films towards High Perfomance Counter Electrodes in Dye-Sensitized Solar Cells. RSC Adv. 2013, 3, 16528−16533. (93) Carim, A. I.; Saadi, F. H.; Soriaga, M. P.; Lewis, N. S. Electrocatalysis of the Hydrogen-Evolution Reaction by Electrodeposited Amorphous Cobalt Selenide Films. J. Mater. Chem. A 2014, 2, 13835−13839. (94) Zainal, Z.; Saravanan, N.; Mien, H. L. Electrodeposition of Nickel Selenide Thin Films in the Presence of Triethanolamine as a Complexing Agent. J. Mater. Sci. 2005, 16, 111−117. (95) Kong, D.; Wang, H.; Lu, Z.; Cui, Y. CoSe2 Nanoparticles Grown on Carbon Fiber Paper: An Efficient and Stable Electrocatalyst for Hydrogen Evolution Reaction. J. Am. Chem. Soc. 2014, 136, 4897− 4900. (96) Paseka, I.; Velicka, J. Hydrogen Evolution and Hydrogen Sorption on Amorphous Smooth Me-P(x) (Me = Ni, Co and Fe-Ni) Electrodes. Electrochim. Acta 1997, 42, 237−242. (97) Burchardt, T. The Hydrogen Evolution Reaction on NiPx Alloys. Int. J. Hydrogen Energy 2000, 25, 627−634. (98) Saadi, F. H.; Carim, A. I.; Verlage, E.; Hemminger, J. C.; Lewis, N. S.; Soriaga, M. P. CoP as an Acid-Stable Active Electrocatalyst for the Hydrogen-Evolution Reaction: Electrochemical Synthesis, Interfacial Characterization and Performance Evaluation. J. Phys. Chem. C 2014, 118, 29294−29300. (99) Daly, B. P.; Barry, F. J. Electrochemical Nickel−phosphorus Alloy Formation. Int. Mater. Rev. 2003, 48, 326−338. (100) Pillai, A. M.; Rajendra, A.; Sharma, A. K. Electrodeposited Nickel-Phosphorous (Ni-P) Alloy Coating: An In-Depth Study of Its Preparation, Properties, and Structural Transitions. J. Coat. Technol. Res. 2012, 9, 785−797. (101) Sequeira, C. A. C.; Santos, D. M. F.; Brito, P. S. D. Electrocatalytic Activity of Simple and Modified Fe-P Electrodeposits for Hydrogen Evolution from Alkaline Media. Energy 2011, 36, 847− 853. (102) Morikawa, T.; Nakade, T.; Yokoi, M.; Fukumoto, Y.; Iwakura, C. Electrodeposition of Ni-P Alloys from Ni-Citrate Bath. Electrochim. Acta 1997, 42, 115−118. AM

DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

(103) Jiang, N.; You, B.; Sheng, M.; Sun, Y. Electrodeposited CobaltPhosphorous-Derived Films as Competent Bifunctional Catlaysts for Overall Water Splitting. Angew. Chem., Int. Ed. 2015, 54, 6251−6254. (104) Brenner, A. Electrodeposition of Alloys Principles and Practice; Academic Press: New York, 1963. (105) Kurowski, A.; Schultze, J. W.; Staikov, G. Initial Stages of Ni-P Electrodeposition: Growth Morphology and Composition of Deposits. Electrochem. Commun. 2002, 4, 565−569. (106) Zeller, R. L.; Landau, U. Electrodeposition of Ni-P Amorphous Alloys. J. Electrochem. Soc. 1992, 139, 3464−3469. (107) Harris, T. M.; Dang, Q. D. The Mechanism of Phosphorus Incorporation during the Electrodeposition of Nickel-Phosphorus Alloys. J. Electrochem. Soc. 1993, 140, 81−83. (108) Popczyk, M.; Budniok, A.; Lasia, A. Electrochemical Properties of Ni-P Electrode Materials Modified with Nickel Oxide and Metallic Cobalt Powders. Int. J. Hydrogen Energy 2005, 30, 265−271. (109) McCrory, C. C. L.; Jung, S. H.; Peters, J. C.; Jaramillo, T. F. Benchmarking Heterogeneous Electrocatalysts for the Oxygen Evolution Reaction. J. Am. Chem. Soc. 2013, 135, 16977−16987. (110) Trassati, S. Electrochemistry of Novel Materials; VCH Publishers: New York, 1994. (111) Trassati, S. Interfacial Electrochemistry: Theory, Experiment, and Applications; Marcel Dekker: New York, 1999. (112) Yamanaka, K. Anodically Electrodeposited Iridium Oxide Films (AEIROF) from Alkaline Solutions for Electrochromic Display Devices. Jpn. J. Appl. Phys., Part 1 1989, 28, 632−637. (113) Zhao, C. X.; Yu, H. T.; Li, Y. C.; Li, X. H.; Ding, L.; Fan, L. Z. Electrochemical Controlled Synthesis and Characterization of WellAligned IrO2 Nanotube Arrays with Enhanced Electrocatalytic Activity toward Oxygen Evolution Reaction. J. Electroanal. Chem. 2013, 688, 269−274. (114) Battaglia, M.; Inguanta, R.; Piazza, S.; Sunseri, C. Fabrication and Characterization of Nanostructured Ni-IrO2 Electrodes for Water Electrolysis. Int. J. Hydrogen Energy 2014, 39, 16797−16805. (115) Blakemore, J. D.; Mara, M. W.; Kushner-Lenhoff, M. N.; Schley, N. D.; Konezny, S. J.; Rivalta, I.; Negre, C. F. A.; Snoeberger, R. C.; Kokhan, O.; Huang, J.; et al. Characterization of an Amorphous Iridium Water-Oxidation Catalyst Electrodeposited from Organometallic Precursors. Inorg. Chem. 2013, 52, 1860−1871. (116) Blakemore, J. D.; Schley, N. D.; Olack, G. W.; Incarvito, C. D.; Brudvig, G. W.; Crabtree, R. H. Anodic Deposition of a Robust Iridium-Based Water-Oxidation Catalyst from Organometallic Precursors. Chem. Sci. 2011, 2, 94−98. (117) Yagi, M.; Tomita, E.; Kuwabara, T. Remarkably High Activity of Electrodeposited IrO2 Film for Electrocatalytic Water Oxidation. J. Electroanal. Chem. 2005, 579, 83−88. (118) Musiani, M.; Furlanetto, F.; Bertoncello, R. Electrodeposited PbO2+RuO2: A Composite Anode for Oxygen Evolution from Sulphuric Acid Solution. J. Electroanal. Chem. 1999, 465, 160−167. (119) Bertoncello, R.; Cattarin, S.; Frateur, I.; Musiani, M. Preparation of Anodes for Oxygen Evolution by Electrodeposition of Composite Oxides of Pb and Ru on Ti. J. Electroanal. Chem. 2000, 492, 145−149. (120) Tsuji, E.; Imanishi, A.; Fukui, K.; Nakato, Y. Electrocatalytic Activity of Amorphous RuO2 Electrode for Oxygen Evolution in an Aqueous Solution. Electrochim. Acta 2011, 56, 2009−2016. (121) Rivera-Vélez, N. E.; Valdez, T.; González, I.; Fachini, E.; Manzo, M.; Cabrera, C. R. Iridium and Ruthenium Electrodeposition at Platinum Nanopowder Using the Rotating Disc Slurry Electrode Technique. ECS Trans. 2011, 35, 47−55. (122) Mohammad, A. M.; Awad, M. I.; El-Deab, M. S.; Okajima, T.; Ohsaka, T. Electrocatalysis by Nanoparticles: Optimization of the Loading Level and Operating pH for the Oxygen Evolution at Crystallographically Oriented Manganese Oxide Nanorods Modified Electrodes. Electrochim. Acta 2008, 53, 4351−4358. (123) El-Deab, M. S.; Awad, M. I.; Mohammad, A. M.; Ohsaka, T. Enhanced Water Electrolysis: Electrocatalytic Generation of Oxygen Gas at Manganese Oxide Nanorods Modified Electrodes. Electrochem. Commun. 2007, 9, 2082−2087.

(124) Zaharieva, I.; Chernev, P.; Risch, M.; Klingan, K.; Kohlhoff, M.; Fischer, A.; Dau, H. Electrosynthesis, Functional, and Structural Characterization of a Water-Oxidizing Manganese Oxide. Energy Environ. Sci. 2012, 5, 7081−7089. (125) Zhou, F. L.; Izgorodin, A.; Hocking, R. K.; Spiccia, L.; MacFarlane, D. R. Electrodeposited MnOx Films from Ionic Liquid for Electrocatalytic Water Oxidation. Adv. Energy Mater. 2012, 2, 1013− 1021. (126) Ramírez, A.; Hillebrand, P.; Stellmach, D.; May, M. M.; Bogdanoff, P.; Fiechter, S. Evaluation of MnOx, Mn2O3, and Mn3O4 Electrodeposited Films for the Oxygen Evolution Reaction of Water. J. Phys. Chem. C 2014, 118, 14073−14081. (127) Friebel, D.; Bajdich, M.; Yeo, B. S.; Louie, M. W.; Miller, D. J.; Casalongue, H. S.; Mbuga, F.; Weng, T. C.; Nordlund, D.; Sokaras, D.; et al. On the Chemical State of Co Oxide Electrocatalysts during Alkaline Water Splitting. Phys. Chem. Chem. Phys. 2013, 15, 17460− 17467. (128) Pauporté, T.; Mendoza, L.; Cassir, M.; Bernard, M. C.; Chivot, J. Direct Low-Temperature Deposition of Crystallized CoOOH Films by Potentiostatic Electrolysis. J. Electrochem. Soc. 2005, 152, C49−C53. (129) Koza, J. A.; Hull, C. M.; Liu, Y. C.; Switzer, J. A. Deposition of β-Co(OH) 2 Films by Electrochemical Reduction of Tris(ethylenediamine)cobalt(III) in Alkaline Solution. Chem. Mater. 2013, 25, 1922−1926. (130) Liu, Y. C.; Koza, J. A.; Switzer, J. A. Conversion of Electrodeposited Co(OH)2 to CoOOH and Co3O4, and Comparison of Their Catalytic Activity for the Oxygen Evolution Reaction. Electrochim. Acta 2014, 140, 359−365. (131) 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. (132) Kanan, M. W.; Surendranath, Y.; Nocera, D. G. CobaltPhosphate Oxygen-Evolving Compound. Chem. Soc. Rev. 2009, 38, 109−114. (133) Risch, M.; Khare, V.; Zaharieva, I.; Gerencser, L.; Chernev, P.; Dau, H. Cobalt-Oxo Core of a Water-Oxidizing Catalyst Film. J. Am. Chem. Soc. 2009, 131, 6936−6937. (134) Du, P.; Kokhan, O.; Chapman, K. W.; Chupas, P. J.; Tiede, D. M. Elucidating the Domain Structure of the Cobalt Oxide Water Splitting Catalyst by X-ray Pair Distribution Function Analysis. J. Am. Chem. Soc. 2012, 134, 11096−11099. (135) Surendranath, Y.; Dincă, M.; Nocera, D. G. ElectrolyteDependent Electrosynthesis and Activity of Cobalt-Based Water Oxidation Catalysts. J. Am. Chem. Soc. 2009, 131, 2615−2620. (136) Risch, M.; Klingan, K.; Ringleb, F.; Chernev, P.; Zaharieva, I.; Fischer, A.; Dau, H. Water Oxidation by Electrodeposited Cobalt Oxides- Role of Anions and Redox-Inert Cations in Structure and Function of the Amorphous Catalyst. ChemSusChem 2012, 5, 542− 549. (137) Surendranath, Y.; Kanan, M. W.; Nocera, D. G. Mechanistic Studies of the Oxygen Evolution Reaction by a Cobalt-Phosphate Catalyst at Neutral pH. J. Am. Chem. Soc. 2010, 132, 16501−16509. (138) Khnayzer, R. S.; Mara, M. W.; Huang, J.; Shelby, M. L.; Chen, L. X.; Castellano, F. N. Structure and Activity of Photochemically Deposited ″CoPi″ Oxygen Evolving Catalyst on Titania. ACS Catal. 2012, 2, 2150−2160. (139) Young, E. R.; Nocera, D. G.; Bulović, V. Direct Formation of a Water Oxidation Catalyst from Thin-Film Cobalt. Energy Environ. Sci. 2010, 3, 1726−1728. (140) Cobo, S.; Heidkamp, J.; Jacques, P. A.; Fize, J.; Fourmond, V.; Guetaz, L.; Jousselme, B.; Ivanova, V.; Dau, H.; Palacin, S.; et al. A Janus Cobalt-Based Catalytic Material for Electro-Splitting of Water. Nat. Mater. 2012, 11, 802−807. (141) Dincă, 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. (142) Risch, M.; Klingan, K.; Heidkamp, J.; Ehrenberg, D.; Chernev, P.; Zaharieva, I.; Dau, H. Nickel-Oxido Structure of a Water-Oxidizing Catalyst Film. Chem. Commun. 2011, 47, 11912−11914. AN

DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

(143) 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. (144) Singh, A.; Chang, S. L. Y.; Hocking, R. K.; Bach, U.; Spiccia, L. Highly Active Nickel Oxide Water Oxidation Catalysts Deposited from Molecular Complexes. Energy Environ. Sci. 2013, 6, 579−586. (145) Kostecki, R.; McLarnon, F. Electrochemical and In Situ Raman Spectroscopic Characterization of Nickel Hydroxide Electrodes: I. Pure Nickel Hydroxide. J. Electrochem. Soc. 1997, 144, 485−493. (146) Lyons, M. E. G.; Brandon, M. P. The Oxygen Evolution Reaction on Passive Oxide Covered Transition Metal Electrodes in Aqueous Alkaline Solution. Part 1-Nickel. Int. J. Electrochem. Sci. 2008, 3, 1386−1424. (147) Louie, M. W.; Bell, A. T. An Investigation of Thin-Film Ni-Fe Oxide Catalysts for the Electrochemical Evolution of Oxygen. J. Am. Chem. Soc. 2013, 135, 12329−12337. (148) Trotochaud, L.; Young, S. L.; Ranney, J. K.; Boettcher, S. W. Nickel-Iron Oxyhydroxide Oxygen-Evolution Electrocatalysts: The Role of Intentional and Incidental Iron Incorporation. J. Am. Chem. Soc. 2014, 136, 6744−6753. (149) Nakaoka, K.; Nakayama, M.; Ogura, K. Electrochemical Deposition of Spinel-Type Cobalt Oxide from Alkaline Solution of Co2+ with Glycine. J. Electrochem. Soc. 2002, 149, C159−C163. (150) Koza, J. A.; He, Z.; Miller, A. S.; Switzer, J. A. Electrodeposition of Crystalline Co3O4-A Catalyst for the Oxygen Evolution Reaction. Chem. Mater. 2012, 24, 3567−3573. (151) Wu, G.; Li, N.; Zhou, D. R.; Mitsuo, K.; Xu, B. Q. Anodically Electrodeposited Co + Ni Mixed Oxide Electrode: Preparation and Electrocatalytic Activity for Oxygen Evolution in Alkaline Media. J. Solid State Chem. 2004, 177, 3682−3692. (152) Kim, T. W.; Woo, M. A.; Regis, M.; Choi, K. S. Electrochemical Synthesis of Spinel Type ZnCo2O4 Electrodes for Use as Oxygen Evolution Reaction Catalysts. J. Phys. Chem. Lett. 2014, 5, 2370−2374. (153) Miao, J.; Yang, H. B.; Khoo, S. Y.; Liu, B. Electrochemical Fabrication of ZnO-CdSe Core-Shell Nanorod Arrays for Efficient Photoelectrochemical Water Splitting. Nanoscale 2013, 5, 11118− 11124. (154) Li, J.; Wu, N. Semiconductor-Based Photocatalysts and Photoelectrochemical Cells for Solar Fuel Generation: A Review. Catal. Sci. Technol. 2015, 5, 1360−1384. (155) van de Krol, R.; Liang, Y.; Schoonman, J. Solar Hydrogen Production with Nanostructured Metal Oxides. J. Mater. Chem. 2008, 18, 2311−2320. (156) Bak, T.; Nowotny, J.; Rekas, M.; Sorrell, C. C. PhotoElectrochemical Hydrogen Generation from Water Using Solar Energy. Materials-Related Aspects. Int. J. Hydrogen Energy 2002, 27, 991−1022. (157) Katz, M. J.; Riha, S. C.; Jeong, N. C.; Martinson, A. B. F.; Farha, O. K.; Hupp, J. T. Toward Solar Fuels: Water Splitting with Sunlight and ″Rust”? Coord. Chem. Rev. 2012, 256, 2521−2529. (158) Dotan, H.; Sivula, K.; Grätzel, M.; Rothschild, A.; Warren, S. C. Probing the Photoelectrochemical Properties of Hematite (α-Fe2O3) Electrodes Using Hydrogen Peroxide as a Hole Scavenger. Energy Environ. Sci. 2011, 4, 958−964. (159) Hamann, T. W. Splitting Water with Rust: Hematite Photoelectrochemistry. Dalton Trans. 2012, 41, 7830−7834. (160) Dare-Edwards, M. P.; Goodenough, J. B.; Hamnett, A.; Trevellick, P. R. Electrochemistry and Photoelectrochemistry of Iron(III) Oxide. J. Chem. Soc., Faraday Trans. 1 1983, 79, 2027−2041. (161) Kay, A.; Cesar, I.; Grätzel, M. New Benchmark for Water Photooxidation by Nanostructured α-Fe2O3 Films. J. Am. Chem. Soc. 2006, 128, 15714−15721. (162) Mayer, M. T.; Lin, Y.; Yuan, G.; Wang, D. Forming Heterojunctions at the Nanoscale for Improved Photoelectrochemical Water Splitting by Semiconductor Materials: Case Studies on Hematite. Acc. Chem. Res. 2013, 46, 1558−1566.

(163) Spray, R. L.; Choi, K.-S. Photoactivity of Transparent Nanocrystalline Fe2O3 Electrodes Prepared via Anodic Electrodeposition. Chem. Mater. 2009, 21, 3701−3709. (164) Spray, R. L.; McDonald, K. J.; Choi, K.-S. Enhancing Photoresponse of Nanoparticulate α-Fe2O3 Electrodes by Surface Composition Tuning. J. Phys. Chem. C 2011, 115, 3497−3506. (165) Schrebler, R.; Bello, K.; Vera, F.; Cury, P.; Muñoz, E.; del Río, R.; Gómez Meier, H.; Córdova, R.; Dalchiele, E. A. An Electrochemical Deposition Route for Obtaining α-Fe2O3 Thin Films. Electrochem. Solid-State Lett. 2006, 9, C110−C113. (166) Hu, Y.-S.; Kleiman-Shwarsctein, A.; Forman, A. J.; Hazen, D.; Park, J.-N.; McFarland, E. W. Pt-Doped α-Fe2O3 Thin Films Active for Photoelectrochemical Water Splitting. Chem. Mater. 2008, 20, 3803− 3805. (167) Kleiman-Shwarsctein, A.; Hu, Y.-S.; Forman, A. J.; Stucky, G. D.; McFarland, E. W. Electrodeposition of α-Fe2O3 Doped with Mo or Cr as Photoanodes for Photocatalytic Water Splitting. J. Phys. Chem. C 2008, 112, 15900−15907. (168) Hu, Y. S.; Kleiman-Shwarsctein, A.; Stucky, G. D.; McFarland, E. W. Improved Photoelectrochemical Performance of Ti-Doped αFe2O3 Thin Films by Surface Modification with Fluoride. Chem. Commun. 2009, 2652−2654. (169) Kleiman-Shwarsctein, A.; Huda, M. N.; Walsh, A.; Yan, Y.; Stucky, G. D.; Hu, Y.-S.; Al-Jassim, M. M.; McFarland, E. W. Electrodeposited Aluminum-Doped α-Fe2O3 Photoelectrodes: Experiment and Theory. Chem. Mater. 2010, 22, 510−517. (170) Kumar, P.; Sharma, P.; Shrivastav, R.; Dass, S.; Satsangi, V. R. Electrodeposited Zirconium-Doped α-Fe2O3 Thin Film for Photoelectrochemical Water Splitting. Int. J. Hydrogen Energy 2011, 36, 2777−2784. (171) Liu, Y.; Yu, Y.-X.; Zhang, W.-D. Photoelectrochemical Properties of Ni-Doped Fe2O3 Thin Films Prepared by Electrodeposition. Electrochim. Acta 2012, 59, 121−127. (172) Cai, J.; Li, S.; Li, Z.; Wang, J.; Ren, Y.; Qin, G. Electrodeposition of Sn-Doped Hollow α-Fe2O3 Nanostructures for Photoelectrochemical Water Splitting. J. Alloys Compd. 2013, 574, 421−426. (173) Mirbagheri, N.; Wang, D.; Peng, C.; Wang, J.; Huang, Q.; Fan, C.; Ferapontova, E. E. Visible Light Driven Photoelectrochemical Water Oxidation by Zn- and Ti-Doped Hematite Nanostructures. ACS Catal. 2014, 4, 2006−2015. (174) McDonald, K. J.; Choi, K.-S. Photodeposition of Co-Based Oxygen Evolution Catalysts on α-Fe2O3 Photoanodes. Chem. Mater. 2011, 23, 1686−1693. (175) Enache, C. S.; Liang, Y. Q.; van de Krol, R. Characterization of Structured α-Fe2O3 Photoanodes Prepared via Electrodeposition and Thermal Oxidation of Iron. Thin Solid Films 2011, 520, 1034−1040. (176) Zeng, Q.; Bai, J.; Li, J.; Xia, L.; Huang, K.; Li, X.; Zhou, B. A Novel In Situ Preparation Method for Nanostructured α-Fe2O3 Films from Electrodeposited Fe Films for Efficient Photoelectrocatalytic Water Splitting and the Degradation of Organic Pollutants. J. Mater. Chem. A 2015, 3, 4345−4353. (177) Shi, X.; Zhang, K.; Shin, K.; Moon, J. H.; Lee, T. W.; Park, J. H. Constructing Inverse Opal Structured Hematite Photoanodes via Electrochemical Process and Their Application to Photoelectrochemical Water Splitting. Phys. Chem. Chem. Phys. 2013, 15, 11717−11722. (178) Shi, X.; Zhang, K.; Park, J. H. Understanding the Positive Effects of (Co−Pi) Co-Catalyst Modification in Inverse-Opal Structured α-Fe2O3-Based Photoelectrochemical Cells. Int. J. Hydrogen Energy 2013, 38, 12725−12732. (179) Liu, X.; Wang, F.; Wang, Q. Nanostructure-Based WO3 Photoanodes for Photoelectrochemical Water Splitting. Phys. Chem. Chem. Phys. 2012, 14, 7894−7818. (180) Butler, M. A. Photoelectrolysis and Physical Properties of the Semiconducting Electrode WO3. J. Appl. Phys. 1977, 48, 1914−1918. (181) Koffyberg, F. P.; Dwight, K.; Wold, A. Interband-Transitions of Semiconducting Oxides Determined from Photoelectrolysis Spectra. Solid State Commun. 1979, 30, 433−437. AO

DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

from Oxygenated Qqueous Zinc Chloride Solutions. J. Electrochem. Soc. 1998, 145, 864−874. (203) Izaki, M.; Omi, T. Transparent Zinc Oxide Films Prepared by Electrochemical Reaction. Appl. Phys. Lett. 1996, 68, 2439−2440. (204) Pauporté, T.; Lincot, D. Hydrogen Peroxide Oxygen Precursor for Zinc Oxide Electrodeposition I. Deposition in Perchlorate Medium. J. Electrochem. Soc. 2001, 148, C310−C314. (205) Jaramillo, T. F.; Baeck, S.-H.; Kleiman-Shwarsctein, A.; Choi, K.-S.; Stucky, G. D.; McFarland, E. W. Automated Electrochemical Synthesis and Photoelectrochemical Characterization of Zn1‑xCoxO Thin Films for Solar Hydrogen Production. J. Comb. Chem. 2005, 7, 264−271. (206) Limmer, S. J.; Kulp, E. A.; Switzer, J. A. Epitaxial Electrodeposition of ZnO on Au(111) from Alkaline Solution: Exploiting Amphoterism in Zn(II). Langmuir 2006, 22, 10535−10539. (207) Steinmiller, E. M. P.; Choi, K.-S. Anodic Construction of Lamellar Structured ZnO Films Using Basic Media via Interfacial Surfactant Templating. Langmuir 2007, 23, 12710−12715. (208) López, C. M.; Choi, K.-S. Enhancement of Electrochemical and Photoelectrochemical Properties of Fibrous Zn and ZnO Electrodes. Chem. Commun. 2005, 3328−3323. (209) López, C. M.; Choi, K.-S. Electrochemical Synthesis of Dendritic Zinc Films Composed of Systematically Varying Motif Crystals. Langmuir 2006, 22, 10625−10629. (210) Steinmiller, E. M. P.; Choi, K.-S. Photochemical Deposition of Cobalt-Based Oxygen Evolving Catalyst on a Semiconductor Photoanode for Solar Oxygen Production. Proc. Natl. Acad. Sci. U. S. A. 2009, 106, 20633−20636. (211) Babu, V. J.; Vempati, S.; Uyar, T.; Ramakrishna, S. Review of One-Dimensional and Two-Dimensional Nanostructured Materials for Hydrogen Generation. Phys. Chem. Chem. Phys. 2015, 17, 2960−2986. (212) Sharma, R.; Mane, R. S.; Min, S. K.; Han, S.-H. Optimization of Growth of In2O3 Nano-Spheres Thin Films by Electrodeposition for Dye-Sensitized Solar Cells. J. Alloys Compd. 2009, 479, 840−843. (213) Yang, X.; Xu, J.; Wong, T.; Yang, Q.; Lee, C.-S. Synthesis of In2O3−In2S3 Core−Shell Nanorods with Inverted Type-I Structure for Photocatalytic H2 Generation. Phys. Chem. Chem. Phys. 2013, 15, 12688−12686. (214) Li, H.; Chen, C.; Huang, X.; Leng, Y.; Hou, M.; Xiao, X.; Bao, J.; You, J.; Zhang, W.; Wang, Y.; et al. Fabrication of In2O3@In2S3 Core−Shell Nanocubes for Enhanced Photoelectrochemical Performance. J. Power Sources 2014, 247, 915−919. (215) Gan, J.; Lu, X.; Zhai, T.; Zhao, Y.; Xie, S.; Mao, Y.; Zhang, Y.; Yang, Y.; Tong, Y. Vertically Aligned In2O3 Nanorods on FTO Substrates for Photoelectrochemical Applications. J. Mater. Chem. 2011, 21, 14685−14688. (216) Kovtyukhova, N. I.; Mallouk, T. E. Electrochemically Assisted Deposition as a New Route to Transparent Conductive Indium Tin Oxide Films. Chem. Mater. 2010, 22, 4939−4949. (217) Xu, Y.; Schoonen, M. A. A. The Absolute Energy Positions of Conduction and Valence Bands of Selected Semiconducting Minerals. Am. Mineral. 2000, 85, 543−556. (218) Bedja, I.; Hotchandani, S.; Kamat, P. V. Preparation and Photoelectrochemical Characterization of Thin SnO2 Nanocrystalline Semiconductor Films and Their Sensitization with Bis(2,2′bipyridine)(2,2′-bipyridine-4,4′-bicarboxylic acid)ruthenium(II) Complex. J. Phys. Chem. 1994, 98, 4133−4140. (219) Vinodgopal, K.; Bedja, I.; Kamat, P. V. Nanostructured Semiconductor Films for Photocatalysis. Photoelectrochemical Behavior of SnO2/TiO2 Composite Systems and Its Role in Photocatalytic Degradation of a Textile Azo Dye. Chem. Mater. 1996, 8, 2180−2187. (220) Chang, S. T.; Leu, I. C.; Hon, M. H. Preparation and Characterization of Nanostructured Tin Oxide Films by Electrochemical Deposition. Electrochem. Solid-State Lett. 2002, 5, C71−C74. (221) Spray, R. L.; Choi, K.-S. Electrochemical Synthesis of SnO2 Films Containing Three-Dimensionally Organized Uniform Mesopores via Interfacial Surfactant Templating. Chem. Commun. 2007, 3655−3653.

(182) Iwai, T. Temperature Dependence of the Optical Absorption Edge of Tungsten Trioxide Single Crystal. J. Phys. Soc. Jpn. 1960, 15, 1596−1600. (183) Mi, Q.; Ping, Y.; Li, Y.; Cao, B.; Brunschwig, B. S.; Khalifah, P. G.; Galli, G. A.; Gray, H. B.; Lewis, N. S. Thermally Stable N2Intercalated WO3 Photoanodes for Water Oxidation. J. Am. Chem. Soc. 2012, 134, 18318−18324. (184) Meulenkamp, E. A. Mechanism of WO3 Electrodeposition from Peroxy-Tungstate Solution. J. Electrochem. Soc. 1997, 144, 1664− 1671. (185) Baeck, S. H.; Jaramillo, T.; Stucky, G. D.; McFarland, E. W. Controlled Electrodeposition of Nanoparticulate Tungsten Oxide. Nano Lett. 2002, 2, 831−834. (186) Monllor-Satoca, D.; Borja, L.; Rodes, A.; Gómez, R.; Salvador, P. Photoelectrochemical Behavior of Nanostructured WO3 Thin-Film Electrodes: The Oxidation of Formic Acid. ChemPhysChem 2006, 7, 2540−2551. (187) Seabold, J. A.; Choi, K.-S. Effect of a Cobalt-Based Oxygen Evolution Catalyst on the Stability and the Selectivity of PhotoOxidation Reactions of a WO3 Photoanode. Chem. Mater. 2011, 23, 1105−1112. (188) Mi, Q.; Zhanaidarova, A.; Brunschwig, B. S.; Gray, H. B.; Lewis, N. S. A Quantitative Assessment of the Competition Between Water and Anion Oxidation at WO3 Photoanodes in Acidic Aqueous Electrolytes. Energy Environ. Sci. 2012, 5, 5694−5697. (189) Baeck, S. H.; Jaramillo, T. F.; Brändli, C.; McFarland, E. W. Combinatorial Electrochemical Synthesis and Characterization of Tungsten-Based Mixed-Metal Oxides. J. Comb. Chem. 2002, 4, 563− 568. (190) Kwong, W. L.; Nakaruk, A.; Koshy, P.; Sorrell, C. C. Tunable Photoelectrochemical Properties by Nanostructural Control in WO3 Thin Films Prepared by Carboxylic Acid-Assisted Electrodeposition. J. Phys. Chem. C 2013, 117, 17766−17776. (191) Hill, J. C.; Choi, K.-S. Effect of Electrolytes on the Selectivity and Stability of n-type WO3 Photoelectrodes for Use in Solar Water Oxidation. J. Phys. Chem. C 2012, 116, 7612−7620. (192) Kulesza, P. J.; Faulkner, L. R. Electrocatalysis at a Novel Electrode Coating of Nonstoichiometric Tungsten(VI,V) Oxide Aggregates. J. Am. Chem. Soc. 1988, 110, 4905−4913. (193) Shaner, M. R.; Fountaine, K. T.; Ardo, S.; Coridan, R. H.; Atwater, H. A.; Lewis, N. S. Photoelectrochemistry of Core−Shell Tandem Junction n−p+-Si/n-WO3 Microwire Array Photoelectrodes. Energy Environ. Sci. 2014, 7, 779−790. (194) Djurišić, A. B.; Chen, X.; Leung, Y. H.; Man Ching Ng, A. ZnO Nanostructures: Growth, Properties and Applications. J. Mater. Chem. 2012, 22, 6526−6510. (195) Kołodziejczak- Radzimska, A.; Jesionowski, T. Zinc Oxide From Synthesis to Application: A Review. Materials 2014, 7, 2833− 2881. (196) Ö zgür, Ü .; Alivov, Y. I.; Liu, C.; Teke, A.; Reshchikov, M. A.; Doğan, S.; Avrutin, V.; Cho, S. J.; Morkoç, H. A Comprehensive Review of ZnO Materials and Devices. J. Appl. Phys. 2005, 98, 041301. (197) Tang, H.; Prasad, K.; Sanjinès, R.; Schmid, P. E.; Lévy, F. Electrical and Optical Properties of TiO2 Anatase Thin Films. J. Appl. Phys. 1994, 75, 2042−2047. (198) Wang, Z. L. Zinc Oxide Nanostructures: Growth, Properties and Applications. J. Phys.: Condens. Matter 2004, 16, R829−R858. (199) Chandiran, A. K.; Abdi-Jalebi, M.; Nazeeruddin, M. K.; Grätzel, M. Analysis of Electron Transfer Properties of ZnO and TiO2 Photoanodes for Dye-Sensitized Solar Cells. ACS Nano 2014, 8, 2261−2268. (200) Lévy-Clément, C.; Elias, J. Optimization of the Design of Extremely Thin Absorber Solar Cells Based on Electrodeposited ZnO Nanowires. ChemPhysChem 2013, 14, 2321−2330. (201) Oekermann, T. On Solar Hydrogen and Nanotechnology; John Wiley & Sons, Ltd.: Chichester, UK, 2010. (202) Peulon, S.; Lincot, D. Mechanistic Study of Cathodic Electrodeposition of Zinc Oxide and Zinc Hydroxychloride Films AP

DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

(222) Santato, C.; López, C. M.; Choi, K.-S. Synthesis and Characterization of Polycrystalline Sn and SnO2 Films with Wire Morphologies. Electrochem. Commun. 2007, 9, 1519−1524. (223) Tian, J.; Zhao, Z.; Kumar, A.; Boughton, R. I.; Liu, H. Recent Progress in Design, Synthesis, and Applications of One-Dimensional TiO2 Nanostructured Surface Heterostructures: A Review. Chem. Soc. Rev. 2014, 43, 6920−6937. (224) Chen, X.; Mao, S. S. Titanium Dioxide Nanomaterials: Synthesis, Properties, Modifications, and Applications. Chem. Rev. 2007, 107, 2891−2959. (225) Schneider, J.; Matsuoka, M.; Takeuchi, M.; Zhang, J.; Horiuchi, Y.; Anpo, M.; Bahnemann, D. W. Understanding TiO2 Photocatalysis: Mechanisms and Materials. Chem. Rev. 2014, 114, 9919−9986. (226) O’Regan, B.; Grätzel, M. A Low-Cost, High-Efficiency Solar Cell Based on Dye-Sensitized Colloidal TiO2 Films. Nature 1991, 353, 737−740. (227) Paracchino, A.; Laporte, V.; Sivula, K.; Grätzel, M.; Thimsen, E. Highly Active Oxide Photocathode for Photoelectrochemical Water Reduction. Nat. Mater. 2011, 10, 456−461. (228) Zhao, J.; Minegishi, T.; Zhang, L.; Zhong, M.; Gunawan; Nakabayashi, M.; Ma, G.; Hisatomi, T.; Katayama, M.; Ikeda, S.; et al. Enhancement of Solar Hydrogen Evolution from Water by Surface Modification with CdS and TiO2 on Porous CuInS2 Photocathodes Prepared by an Electrodeposition−Sulfurization Method. Angew. Chem., Int. Ed. 2014, 53, 11808−11812. (229) Hu, S.; Shaner, M. R.; Beardslee, J. A.; Lichterman, M.; Brunschwig, B. S.; Lewis, N. S. Amorphous TiO2 Coatings Stabilize Si, GaAs, and GaP Photoanodes for Efficient Water Oxidation. Science 2014, 344, 1005−1009. (230) Kavan, L.; O’Regan, B.; Kay, A.; Grätzel, M. Preparation of TiO2 (Anatase) Films on Electrodes by Anodic Oxidative Hydrolysis of TiCl3. J. Electroanal. Chem. 1993, 346, 291−307. (231) Lei, Y.; Zhang, L. D.; Fan, J. C. Fabrication, Characterization and Raman Study of TiO2 Nanowire Arrays Prepared by Anodic Oxidative Hydrolysis of TiCl3. Chem. Phys. Lett. 2001, 338, 231−236. (232) Liu, S.; Huang, K. Straightforward Fabrication of Highly Ordered TiO2 Nanowire Arrays in AAM on Aluminum Substrate. Sol. Energy Mater. Sol. Cells 2005, 85, 125−131. (233) Lei, Y.; Zhang, L. D.; Meng, G. W.; Li, G. H.; Zhang, X. Y.; Liang, C. H.; Chen, W.; Wang, S. X. Preparation and Photoluminescence of Highly Ordered TiO2 Nanowire Arrays. Appl. Phys. Lett. 2001, 78, 1125−1124. (234) Eisenberg, D.; Ahn, H. S.; Bard, A. J. Enhanced Photoelectrochemical Water Oxidation on Bismuth Vanadate by Electrodeposition of Amorphous Titanium Dioxide. J. Am. Chem. Soc. 2014, 136, 14011−14014. (235) Wessels, K.; Feldhoff, A.; Wark, M.; Rathousky, J.; Oekermann, T. Low-Temperature Preparation of Crystalline Nanoporous TiO2 Films by Surfactant-Assisted Anodic Electrodeposition. Electrochem. Solid-State Lett. 2006, 9, C93−C96. (236) Natarajan, C.; Nogami, G. Cathodic Electrodeposition of Nanocrystalline Titanium Dioxide Thin Films. J. Electrochem. Soc. 1996, 143, 1547−1550. (237) Matsumoto, Y.; Ishikawa, Y.; Nishida, M.; Ii, S. A New Electrochemical Method To Prepare Mesoporous Titanium(IV) Oxide Photocatalyst Fixed on Alumite Substrate. J. Phys. Chem. B 2000, 104, 4204−4209. (238) Lokhande, C. D.; Min, S. K.; Jung, K. D.; Joo, O. S. Cathodic Electrodeposition of Amorphous Titanium Oxide Films from an Alkaline Solution Bath. J. Mater. Sci. 2004, 39, 6607−6610. (239) Kudo, A.; Omori, K.; Kato, H. A Novel Aqueous Process for Preparation of Crystal Form-Controlled and Highly Crystalline BiVO4 Powder from Layered Vanadates at Room Temperature and Its Photocatalytic and Photophysical Properties. J. Am. Chem. Soc. 1999, 121, 11459−11467. (240) Park, Y.; McDonald, K. J.; Choi, K.-S. Progress in Bismuth Vanadate Photoanodes for Use in Solar Water Oxidation. Chem. Soc. Rev. 2013, 42, 2321−2337.

(241) Li, Z.; Luo, W.; Zhang, M.; Feng, J.; Zou, Z. Photoelectrochemical Cells for Solar Hydrogen Production: Current State of Promising Photoelectrodes, Methods to improve Their Properties, and Outlook. Energy Environ. Sci. 2013, 6, 347−370. (242) Abdi, F. F.; Han, L.; Smets, A. H. M.; Zeman, M.; Dam, B.; van de Krol, R. Efficient Solar Water Splitting by Enhanced Charge Separation in a Bismuth Vanadate-Silicon Tandem Photoelectrode. Nat. Commun. 2013, 4, 2195. (243) Kim, J. H.; Lee, J. S. BiVO4-Based Heterostructured Photocatalysts for Solar Water Splitting: A Review. Energy Environ. Focus 2014, 3, 339−353. (244) Huang, Z.-F.; Pan, L.; Zou, J.-J.; Zhang, X.; Wang, L. Nanostructured Bismuth Vanadate-Based Materials for Solar-EnergyDriven Water Oxidation: A Review on Recent Progress. Nanoscale 2014, 6, 14044−14063. (245) Chen, Y.-S.; Manser, J. S.; Kamat, P. V. All Solution-Processed Lead Halide Perovskite-BiVO4 Tandem Assembly for Photolytic Solar Fuels Production. J. Am. Chem. Soc. 2015, 137, 974−981. (246) Zhong, M.; Hisatomi, T.; Kuang, Y.; Zhao, J.; Liu, M.; Iwase, A.; Jia, Q.; Nishiyama, H.; Minegishi, T.; Nakabayashi, M.; et al. Surface Modification of CoOx Loaded BiVO4 Photoanodes with Ultrathin p-Type NiO Layers for Improved Solar Water Oxidation. J. Am. Chem. Soc. 2015, 137, 5053−5060. (247) Pihosh, Y.; Turkevych, I.; Mawatari, K.; Uemura, J.; Kazoe, Y.; Kosar, S.; Makita, K.; Sugaya, T.; Matsui, T.; Fujita, D.; et al. Photocatalytic Generation of Hydrogen by Core-Shell WO3/BiVO4 Nanorods with Ultimate Water Splitting Efficiency. Sci. Rep. 2015, 5, 11141. (248) Dall’Antonia, L. H.; de Tacconi, N. R.; Chanmanee, W.; Timmaji, H.; Myung, N.; Rajeshwar, K. Electrosynthesis of Bismuth Vanadate Photoelectrodes. Electrochem. Solid-State Lett. 2010, 13, D29−D32. (249) Myung, N.; Ham, S.; Choi, S.; Chae, Y.; Kim, W.-G.; Jeon, Y. J.; Paeng, K.-J.; Chanmanee, W.; de Tacconi, N. R.; Rajeshwar, K. Tailoring Interfaces for Electrochemical Synthesis of Semiconductor Films: BiVO4, Bi2O3, or Composites. J. Phys. Chem. C 2011, 115, 7793−7800. (250) Seabold, J. A.; Choi, K.-S. Efficient and Stable Photo-Oxidation of Water by a Bismuth Vanadate Photoanode Coupled with an Iron Oxyhydroxide Oxygen Evolution Catalyst. J. Am. Chem. Soc. 2012, 134, 2186−2192. (251) Park, Y.; Kang, D.; Choi, K.-S. Marked Enhancement in Electron−Hole Separation Achieved in the Low Bias Region using Electrochemically Prepared Mo-Doped BiVO4 Photoanodes. Phys. Chem. Chem. Phys. 2014, 16, 1238−1246. (252) McDonald, K. J.; Choi, K.-S. A new Electrochemical Synthesis Route for a BiOI Electrode and Its Conversion to a Highly Efficient Porous BiVO4 Photoanode for Solar Water Oxidation. Energy Environ. Sci. 2012, 5, 8553−8557. (253) Zhou, Q.; Xie, Q.; Fu, Y.; Su, Z.; Jia, X. e.; Yao, S. Electrodeposition of Carbon Nanotubes−Chitosan−Glucose Oxidase Biosensing Composite Films Triggered by Reduction of pBenzoquinone or H2O2. J. Phys. Chem. B 2007, 111, 11276−11284. (254) Rettie, A. J. E.; Lee, H. C.; Marshall, L. G.; Lin, J.-F.; Capan, C.; Lindemuth, J.; McCloy, J. S.; Zhou, J.; Bard, A. J.; Mullins, C. B. Combined Charge Carrier Transport and Photoelectrochemical Characterization of BiVO4 Single Crystals: Intrinsic Behavior of a Complex Metal Oxide. J. Am. Chem. Soc. 2013, 135, 11389−11396. (255) Kim, T. W.; Choi, K.-S. Nanoporous BiVO4 Photoanodes with Dual-Layer Oxygen Evolution Catalysts for Solar Water Splitting. Science 2014, 343, 990−994. (256) Kang, D.; Park, Y.; Hill, J. C.; Choi, K.-S. Preparation of BiBased Ternary Oxide Photoanodes BiVO4, Bi2WO6, and Bi2Mo3O12 Using Dendritic Bi Metal Electrodes. J. Phys. Chem. Lett. 2014, 5, 2994−2999. (257) Benko, F. A.; MacLaurin, C. L.; Koffyberg, F. P. CuWO4 and Cu3WO6 as Anodes for the Photoelectrolysis of Water. Mater. Res. Bull. 1982, 17, 133−136. AQ

DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

(258) Arora, S. K.; Mathew, T.; Batra, N. M. Electrochemical Characteristics of Copper Tungstate Single Crystals. J. Phys. D: Appl. Phys. 1990, 23, 460−464. (259) Chang, Y.; Braun, A.; Deangelis, A.; Kaneshiro, J.; Gaillard, N. Effect of Thermal Treatment on the Crystallographic, Surface Energetics, and Photoelectrochemical Properties of Reactively Cosputtered Copper Tungstate for Water Splitting. J. Phys. Chem. C 2011, 115, 25490−25495. (260) Yourey, J. E.; Bartlett, B. M. Electrochemical Deposition and Photoelectrochemistry of CuWO4, a Promising Photoanode for Water Oxidation. J. Mater. Chem. 2011, 21, 7651−7610. (261) Yourey, J. E.; Kurtz, J. B.; Bartlett, B. M. J. Phys. Chem. C 2012, 116, 3200−3205. (262) Hill, J. C.; Choi, K.-S. Synthesis and Characterization of High Surface Area CuWO4 and Bi2WO6 Electrodes for Use as Photoanodes for Solar Water Oxidation. J. Mater. Chem. A 2013, 1, 5006−5014. (263) Hill, J. C.; Ping, Y.; Galli, G. A.; Choi, K.-S. Synthesis, Photoelectrochemical Properties, and First Principles Study of n-Type CuW1−xMoxO4 Electrodes Showing Enhanced Visible Light Absorption. Energy Environ. Sci. 2013, 6, 2440−2446. (264) Zhang, L.; Wang, H.; Chen, Z.; Wong, P.-K.; Liu, J. Bi2WO6 Micro/nano-Structures: Synthesis, Modifications and Visible-LightDriven Photocatalytic Applications. Appl. Catal., B 2011, 106, 1−13. (265) Finlayson, A. P.; Tsaneva, V. N.; Lyons, L.; Clark, M.; Glowacki, B. A. Evaluation of Bi-W-Oxides for Visible Light Photocatalysis. Phys. Status Solidi A 2006, 203, 327−335. (266) Kudo, A.; Hijii, S. H2 or O2 Evolution from Aqueous Solutions on Layered Oxide Photocatalysts Consisting of Bi 3+ with 6s2 Configuration and d0 Transition Metal Ions. Chem. Lett. 1999, 28, 1103−1104. (267) Zhang, L.; Zhu, Y. A Review of Controllable Synthesis and Enhancement of Performances of Bismuth Tungstate Visible-LightDriven Photocatalysts. Catal. Sci. Technol. 2012, 2, 694−706. (268) Yu, J.; Kudo, A. Hydrothermal Synthesis and Photocatalytic Property of 2-Dimensional Bismuth Molybdate Nanoplates. Chem. Lett. 2005, 34, 1528−1529. (269) Zheng, F.-L.; Li, G.-R.; Yu, X.-L.; Tong, Y.-X. Synthesis of Bismuth Molybdate Nanowires via Electrodeposition−Heat-Treatment Method. Electrochem. Solid-State Lett. 2009, 12, K56−K58. (270) Hara, M.; Kondo, T.; Komoda, M.; Ikeda, S.; Kondo, J. N.; Domen, K.; Shinohara, K.; Tanaka, A. Cu2O as a Photocatalyst for Overall Water Splitting under Visible Light Irradiation. Chem. Commun. 1998, 357−358. (271) de Jongh, P. E.; Kelly, J. J. Cu2O: A Catalyst for the Photochemical Decomposition of Water? Chem. Commun. 1999, 1069−1070. (272) Gerischer, H. On the Stability of Semiconductor Electrodes against Photodecomposition. J. Electroanal. Chem. Interfacial Electrochem. 1977, 82, 133−143. (273) Wang, L.; Tao, M. Fabrication and Characterization of p-n Homojunctions in Cuprous Oxide by Electrochemical Deposition. Electrochem. Solid-State Lett. 2007, 10, H248−H250. (274) McShane, C. M.; Choi, K.-S. Photocurrent Enhancement of ntype Cu2O Electrodes Achieved by Controlling Dendritic Branching Growth. J. Am. Chem. Soc. 2009, 131, 2561−2569. (275) Wang, W.; Wu, D.; Zhang, Q.; Wang, L.; Tao, M. pHDependence of Conduction Type in Cuprous Oxide Synthesized from Solution. J. Appl. Phys. 2010, 107, 123717. (276) Grez, P.; Herrera, F.; Riveros, G.; Ramírez, A.; Henríquez, R.; Dalchiele, E.; Schrebler, R. Morphological, Structural, and Photoelectrochemical Characterization of n-Type Cu2O Thin Films Obtained by Electrodeposition. Phys. Status Solidi A 2012, 209, 2470−2475. (277) de Jongh, P. E.; Vanmaekelbergh, D.; Kelly, J. J. Cu2O: Electrodeposition and Characterization. Chem. Mater. 1999, 11, 3512− 3517. (278) Paracchino, A.; Brauer, J. C.; Moser, J.-E.; Thimsen, E.; Grätzel, M. Synthesis and Characterization of High-Photoactivity Electro-

deposited Cu2O Solar Absorber by Photoelectrochemistry and Ultrafast Spectroscopy. J. Phys. Chem. C 2012, 116, 7341−7350. (279) Nian, J.; Hu, C.; Teng, H. Electrodeposited p-Type Cu2O for H2 Evolution from Photoelectrolysis of Water under Visible Light Illumination. Int. J. Hydrogen Energy 2008, 33, 2897−2903. (280) Laidoudi, S.; Bioud, A. Y.; Azizi, A.; Schmerber, G.; Bartringer, J.; Barre, S.; Dinia, A. Growth and Characterization of Electrodeposited Cu2O Thin Films. Semicond. Sci. Technol. 2013, 28, 115005. (281) Zhou, Y. C.; Switzer, J. A. Electrochemical Deposition and Microstructure of Copper (I) Oxide Films. Scr. Mater. 1998, 38, 1731−1738. (282) Rakhshani, A. E.; Varghese, J. Galvanostatic Deposition of Thin Films of Cuprous Oxide. Sol. Energy Mater. 1987, 15, 237−248. (283) Golden, T. D.; Shumsky, M. G.; Zhou, Y. C.; VanderWerf, R. A.; VanLeeuwen, R. A.; Switzer, J. A. Electrochemical Deposition of Copper(I) Oxide Films. Chem. Mater. 1996, 8, 2499−2504. (284) Wang, L. C.; de Tacconi, N. R.; Chenthamarakshan, C. R.; Rajeshwar, K.; Tao, M. Electrodeposited Copper Oxide Films: Effect of Bath pH on Grain Orientation and Orientation-Dependent Interfacial Behavior. Thin Solid Films 2007, 515, 3090−3095. (285) Mizuno, K.; Izaki, M.; Murase, K.; Shinagawa, T.; Chigane, M.; Inaba, M.; Tasaka, A.; Awakura, Y. Structural and Electrical Characterizations of Electrodeposited p-Type Semiconductor Cu2O Films. J. Electrochem. Soc. 2005, 152, C179−C182. (286) Sowers, K. L.; Fillinger, A. Crystal Face Dependence of pCu2O Stability as Photocathode. J. Electrochem. Soc. 2009, 156, F80− F85. (287) Kim, M.; Yoon, S.; Jung, H.; Lee, K.-J.; Lim, D.-C.; Kim, I.-S.; Yoo, B.; Lim, J.-H. The Influence of Polarity of Electrodeposited Cu2O Thin Films on the Photoelectrochemical Performance. Jpn. J. Appl. Phys. 2014, 53, 08NJ01. (288) Septina, W.; Ikeda, S.; Khan, M. A.; Hirai, T.; Harada, T.; Matsumura, M.; Peter, L. M. Potentiostatic Electrodeposition of Cuprous Oxide Thin Films for Photovoltaic Applications. Electrochim. Acta 2011, 56, 4882−4888. (289) McShane, C. M.; Choi, K.-S. Junction Studies on Electrochemically Fabricated p−n Cu2O Homojunction Solar Cells for Efficiency Enhancement. Phys. Chem. Chem. Phys. 2012, 14, 6112− 6118. (290) Brown, K. E. R.; Choi, K.-S. Electrochemical Synthesis and Characterization of Transparent Nanocrystalline Cu2O Films and Their Conversion to CuO Films. Chem. Commun. 2006, 3311−3313. (291) Zhang, Z.; Wang, P. Highly Stable Copper Oxide Composite as an Effective Photocathode for Water Splitting via a Facile Electrochemical Synthesis Strategy. J. Mater. Chem. 2012, 22, 2456−2464. (292) Zhang, Z.; Dua, R.; Zhang, L.; Zhu, H.; Zhang, H.; Wang, P. Carbon-Layer-Protected Cuprous Oxide Nanowire Arrays for Efficient Water Reduction. ACS Nano 2013, 7, 1709−1717. (293) Dubale, A. A.; Su, W.-N.; Tamirat, A. G.; Pan, C.-J.; Aragaw, B. A.; Chen, H.-M.; Chen, C.-H.; Hwang, B.-J. The Synergetic Effect of Graphene on Cu2O Nanowire Arrays as a Highly Efficient Hydrogen Evolution Photocathode in Water Splitting. J. Mater. Chem. A 2014, 2, 18383−18397. (294) Paracchino, A.; Mathews, N.; Hisatomi, T.; Stefik, M.; Tilley, S. D.; Grätzel, M. Ultrathin Films on Copper(I) Oxide Water Splitting Photocathodes: A Study on Performance and Stability. Energy Environ. Sci. 2012, 5, 8673−8681. (295) Tilley, S. D.; Schreier, M.; Azevedo, J.; Stefik, M.; Grätzel, M. Ruthenium Oxide Hydrogen Evolution Catalysis on Composite Cuprous Oxide Water-Splitting Photocathodes. Adv. Funct. Mater. 2014, 24, 303−311. (296) Morales-Guio, C. G.; Tilley, S. D.; Vrubel, H.; Grätzel, M.; Hu, X. Hydrogen Evolution from a Copper(I) Oxide Photocathode Coated with an Amorphous Molybdenum Sulphide Catalyst. Nat. Commun. 2014, 5, 1−7. (297) Morales-Guio, C. G.; Liardet, L.; Mayer, M. T.; Tilley, S. D.; Grätzel, M.; Hu, X. Photoelectrochemical Hydrogen Production in Alkaline Solutions Using Cu2O Coated with Earth-Abundant Hydrogen Evolution Catalysts. Angew. Chem., Int. Ed. 2014, 54, 664−667. AR

DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

Bandgap p-Copper (II) Oxide Film With Excellent Photoactivity. J. Electrochem. Soc. 2011, 158, D578−D584. (317) Nakaoka, K.; Ogura, K. Electrochemical Preparation of p-Type Cupric and Cuprous Oxides on Platinum and Gold Substrates from Copper(II) Solutions with Various Amino Acids. J. Electrochem. Soc. 2002, 149, C579−C585. (318) Nakaoka, K.; Ueyama, J.; Ogura, K. Photoelectrochemical Behavior of Electrodeposited CuO and Cu2O Thin Films on Conducting Substrates. J. Electrochem. Soc. 2004, 151, C661−C665. (319) Siegfried, M. J.; Choi, K.-S. Conditions and Mechanism for the Anodic Deposition of Cupric Oxide Films in Slightly Acidic Aqueous Media. J. Electrochem. Soc. 2007, 154, D674−D677. (320) Switzer, J. A.; Kothari, H. M.; Poizot, P.; Nakanishi, S.; Bohannan, E. W. Enantiospecific Electrodeposition of a Chiral Catalyst. Nature 2003, 425, 490−493. (321) Bohannan, E. W.; Kothari, H. M.; Nicic, I. M.; Switzer, J. A. Enantiospecific Electrodeposition of Chiral CuO Films on SingleCrystal Cu(111). J. Am. Chem. Soc. 2004, 126, 488−489. (322) Kothari, H. M.; Kulp, E. A.; Boonsalee, S.; Nikiforov, M. P.; Bohannan, E. W.; Poizot, P.; Nakanishi, S.; Switzer, J. A. Enantiospecific Electrodeposition of Chiral CuO Films from Copper(II) Complexes of Tartaric and Amino Acids on Single-Crystal Au(001). Chem. Mater. 2004, 16, 4232−4244. (323) Joseph, S.; Kamath, P. V.; Upadhya, S. Electrochemical Synthesis of Oriented CuO Coatings on Stainless Steel Substrates: Solution-Mediated Control over Orientation. J. Electrochem. Soc. 2009, 156, E18−E22. (324) Zhang, R.; Liu, J.; Guo, H.; Tong, X. Synthesis of CuO Nanowire Arrays as High-Performance Electrode for Lithium Ion Batteries. Mater. Lett. 2015, 139, 55−58. (325) Sagu, J. S.; Peiris, T. A. N.; Wijayantha, K. G. U. Rapid and Simple Potentiostatic Deposition of Copper (II) Oxide Thin Films. Electrochem. Commun. 2014, 42, 68−71. (326) Zhao, X.; Wang, P.; Yan, Z.; Ren, N. Ag Nanoparticles Decorated CuO Nanowire Arrays for Efficient Plasmon Enhanced Photoelectrochemical Water Splitting. Chem. Phys. Lett. 2014, 609, 59−64. (327) Read, C. G.; Park, Y.; Choi, K.-S. Electrochemical Synthesis of p-Type CuFeO2 Electrodes for Use in a Photoelectrochemical Cell. J. Phys. Chem. Lett. 2012, 3, 1872−1876. (328) Hahn, N. T.; Holmberg, V. C.; Korgel, B. A.; Mullins, C. B. Electrochemical Synthesis and Characterization of p-CuBi2O4 Thin Film Photocathodes. J. Phys. Chem. C 2012, 116, 6459−6466. (329) Nakabayashi, Y.; Nishikawa, M.; Nosaka, Y. Fabrication of CuBi2O4 Photocathode through Novel Anodic Electrodeposition for Solar Hydrogen Production. Electrochim. Acta 2014, 125, 191−198. (330) Bohannan, E. W.; Jaynes, C. C.; Shumsky, M. G.; Barton, J. K.; Switzer, J. A. Low-Temperature Electrodeposition of the HighTemperature Cubic Polymorph of Bismuth(III) Oxide. Solid State Ionics 2000, 131, 97−107. (331) Laurent, K.; Wang, G.; Tusseau-Nenez, S.; Leprince-Wang, Y. Structure and Conductivity Studies of Electrodeposited δ-Bi2O3. Solid State Ionics 2008, 178, 1735−1739. (332) Zhang, S. B.; Wei, S.-H.; Zunger, A. A Phenomenological Model for Systematization and Prediction of Doping Limits in II−VI and I−III−VI2 Compounds. J. Appl. Phys. 1998, 83, 3192−3196. (333) Chen, S.; Gong, X. G.; Wei, S.-H. Band-Structure Anomalies of the Chalcopyrite Semiconductors CuGaX2 versus AgGaX2 (X = S and Se) and Their Alloys. Phys. Rev. B: Condens. Matter Mater. Phys. 2007, 75, 205209. (334) Wu, W.; Jiang, C.; Roy, V. A. L. Recent Progress in Magnetic Iron Oxide−Semiconductor Composite Nanomaterials as Promising Photocatalysts. Nanoscale 2015, 7, 38−58. (335) Martinez, A.; Fernández, A.; Arriaga, L.; Cano, U. Preparation and Characterization of Cu−In−S Thin Films by Electrodeposition. Mater. Chem. Phys. 2006, 95, 270−274. (336) Asenjo, B.; Chaparro, A.; Gutiérrez, M.; Herrero, J. Electrochemical Growth and Properties of CuInS2 Thin Films for Solar Energy Conversion. Thin Solid Films 2006, 511, 117−120.

(298) Azevedo, J.; Steier, L.; Dias, P.; Stefik, M.; Sousa, C. T.; Araújo, J. P.; Mendes, A.; Grätzel, M.; Tilley, S. D. On the Stability Enhancement of Cuprous Oxide Water Splitting Photocathodes by Low Temperature Steam Annealing. Energy Environ. Sci. 2014, 7, 4044−4052. (299) Engel, C. J.; Polson, T. A.; Spado, J. R.; Bell, J. M.; Fillinger, A. Photoelectrochemistry of Porous p-Cu2O Films. J. Electrochem. Soc. 2008, 155, F37−F42. (300) Wu, L.; Tsui, L.-k.; Swami, N.; Zangari, G. Photoelectrochemical Stability of Electrodeposited Cu2O Films. J. Phys. Chem. C 2010, 114, 11551−11556. (301) Shyamal, S.; Hajra, P.; Mandal, H.; Singh, J. K.; Satpati, A. K.; Pande, S.; Bhattacharya, C. Effect of Substrates on the Photoelectrochemical Reduction of Water over Cathodically Electrodeposited p-Type Cu2O Thin Films. ACS Appl. Mater. Interfaces 2015, 7, 18344−18352. (302) Chauhan, D.; Satsangi, V. R.; Dass, S.; Shrivastav, R. Preparation and Characterization of Nanostructured CuO Thin Films for Photoelectrochemical Splitting of Water. Bull. Mater. Sci. 2014, 29, 709−716. (303) Koffyberg, F. P.; Benko, F. A. A Photoelectrochemical Determination of the Position of the Conduction and Valence band Edges of p-Type CuO. J. Appl. Phys. 1982, 53, 1173−1177. (304) Hsu, Y.-K.; Yu, C.-H.; Lin, H.-H.; Chen, Y.-C.; Lin, Y.-G. Template Synthesis of Copper Oxide Nanowires for Photoelectrochemical Hydrogen Generation. J. Electroanal. Chem. 2013, 704, 19− 23. (305) Chiang, C.-Y.; Shin, Y.; Aroh, K.; Ehrman, S. Copper Oxide Photocathodes Prepared by a Solution Based Process. Int. J. Hydrogen Energy 2012, 37, 8232−8239. (306) Chiang, C.-Y.; Aroh, K.; Franson, N.; Satsangi, V. R.; Dass, S.; Ehrman, S. Copper Oxide Nanoparticle Made by Flame Spray Pyrolysis for Photoelectrochemical Water Splitting - Part II. Photoelectrochemical Study. Int. J. Hydrogen Energy 2011, 36, 15519−15526. (307) Lim, Y.-F.; Chua, C. S.; Lee, C. J. J.; Chi, D. Sol−Gel Deposited Cu2O and CuO Thin Films for Photocatalytic Water Splitting. Phys. Chem. Chem. Phys. 2014, 16, 25928−25934. (308) Emin, S.; Abdi, F. F.; Fanetti, M.; Peng, W.; Smith, W.; Sivula, K.; Dam, B.; Valant, M. A novel Approach for the Preparation of Textured CuO Thin Films from Electrodeposited CuCl and CuBr. J. Electroanal. Chem. 2014, 717−718, 243−249. (309) Barreca, D.; Fornasiero, P.; Gasparotto, A.; Gombac, V.; Maccato, C.; Montini, T.; Tondello, E. The Potential of Supported Cu2O and CuO Nanosystems in Photocatalytic H2 Production. ChemSusChem 2009, 2, 230−233. (310) Jin, Z.; Li, P.; Liu, G.; Zheng, B.; Yuan, H.; Xiao, D. Enhancing Catalytic Formaldehyde Oxidation on CuO−Ag2O Nanowires for Gas Sensing and Hydrogen Evolution. J. Mater. Chem. A 2013, 1, 14736− 14738. (311) Zhang, L.; Liu, Y.-N.; Zhou, M.; Yan, J. Improving Photocatalytic Hydrogen Evolution over CuO/Al2O3 by PlatinumDepositing and CuS-Loading. Appl. Surf. Sci. 2013, 282, 531−537. (312) Xu, S.; Du, A. J.; Liu, J.; Ng, J.; Sun, D. D. Highly Efficient CuO Incorporated TiO2 Nanotube Photocatalyst for Hydrogen Production from Water. Int. J. Hydrogen Energy 2011, 36, 6560−6568. (313) Guo, X.; Diao, P.; Di, X.; Huang, S.; Yang, Y.; Jin, T.; Wu, Q.; Xiang, M.; Zhang, M. CuO/Pd Composite Photocathodes for Photoelectrochemical Hydrogen Evolution Reaction. Int. J. Hydrogen Energy 2014, 39, 7686−7696. (314) Ng, K. H.; Kadir, H. A.; Minggu, L. J.; Kassim, M. B. Stability of WO3/CuO Heterojunction Photoelectrodes in PEC System. Mater. Sci. Forum 2013, 756, 219−224. (315) Poizot, P.; Hung, C.-J.; Nikiforov, M. P.; Bohannan, E. W.; Switzer, J. A. An Electrochemical Method for CuO Thin Film Deposition from Aqueous Solution. Electrochem. Solid-State Lett. 2003, 6, C21−C25. (316) Izaki, M.; Nagai, M.; Maeda, K.; Mohamad, F. B.; Motomura, K.; Sasano, J.; Shinagawa, T.; Watase, S. Electrodeposition of 1.4-eVAS

DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

(356) Chirilă, A.; Buecheler, S.; Pianezzi, F.; Bloesch, P.; Gretener, C.; Uhl, A. R.; Fella, C.; Kranz, L.; Perrenoud, J.; Seyrling, S.; et al. Highly Efficient Cu(In,Ga)Se2 Solar Cells Grown on Flexible Polymer Films. Nat. Mater. 2011, 10, 857−861. (357) Dergacheva, M. B.; Urazov, K. A.; Pen’kova, N. V. Electrodeposition of CuInSe2 Films onto a Molybdenum Electrode. Russ. J. Appl. Chem. 2010, 83, 653−658. (358) Lai, Y.; Liu, J.; Yang, J.; Wang, B.; Liu, F.; Zhang, Z.; Li, J.; Liu, Y. Incorporation Mechanism of Indium and Gallium during Electrodeposition of Cu(In,Ga)Se2 Thin Film. J. Electrochem. Soc. 2011, 158, D704−D709. (359) Kemell, M.; Saloniemi, H.; Ritala, M.; Leskelä, M. Electrochemical Quartz Crystal Microbalance Study of the Electrodeposition Mechanisms of CuInSe2 Thin Films. J. Electrochem. Soc. 2001, 148, C110−C118. (360) Chiu, Y.-S.; Hsieh, M.-T.; Chang, C.-M.; Chen, C.-S.; Whang, T.-J. Single-Step Electrodeposition of CIS Thin Films with the Complexing Agent Triethanolamine. Appl. Surf. Sci. 2014, 299, 52−57. (361) Huang, C.-J.; Su, Y.-K.; Chen, K.-L.; Lai, M. Y. Characteristics of Copper Indium Diselenide Thin Films Formed on Flexible Substrates by Electrodeposition. Jpn. J. Appl. Phys. 2005, 44, 7795− 7800. (362) Fahoume, M.; Boudraine, H.; Aggour, M.; Chraïbi, F.; Ennaoui, A.; Delplancke, J. L. One, Step Electrodeposition of Cu(Ga,In)Se2 Thin Films from Aqueous Solution. J. Phys. IV 2005, 123, 75−80. (363) Yang, J.; Huang, C.; Jiang, L.; Liu, F.; Lai, Y.; Li, J.; Liu, Y. Effects of Hydrogen Peroxide on Electrodeposition of Cu(In,Ga)Se2 Thin Films and Band Gap Controlling. Electrochim. Acta 2014, 142, 208−214. (364) Li, M.; Zheng, M.; Zhou, T.; Li, C.; Ma, L.; Shen, W. Fabrication and Characterization of Ordered CuIn(1‑x)GaxSe2 Nanopore Films via Template-Based Electrodeposition. Nanoscale Res. Lett. 2012, 7, 675. (365) de S Lucas, F. W.; Lima, A. R. F.; Mascaro, L. H. The Electrodeposition of Ga-Doped CuInSe2 Thin Film in the Presence of Triton 100-X. Electrochim. Acta 2014, 147, 47−53. (366) Sebastian, P. J.; Calixto, M. E.; Bhattacharya, R. N.; Noufi, R. A 10% Cu(In,Ga)Se2 Based Photovoltaic Structure Formed by Electrodeposition and Subsequent Thermal Processing. J. Electrochem. Soc. 1998, 145, 3613−3615. (367) Sebastian, P. J.; Calixto, M. E.; Bhattacharya, R. N.; Noufi, R. CIS and CIGS Based Photovoltaic Structures Developed from Electrodeposited Precursors. Sol. Energy Mater. Sol. Cells 1999, 59, 125−135. (368) Bhattacharya, R. N.; Hiltner, J. F.; Batchelor, W.; Contreras, M. A.; Noufi, R. N.; Sites, J. R. 15.4% CuIn1−xGaxSe2-Based Photovoltaic Cells from Solution-Based Precursor Films. Thin Solid Films 2000, 361−362, 396−399. (369) Jacobsson, T. J.; Platzer- Björkman, C.; Edoff, M.; Edvinsson, T. CuInxGa1−xSe2 as an Efficient Photocathode for Solar Hydrogen Generation. Int. J. Hydrogen Energy 2013, 38, 15027−15035. (370) Kumagai, H.; Minegishi, T.; Sato, N.; Yamada, T.; Kubota, J.; Domen, K. Efficient Solar Hydrogen Production from Neutral Electrolytes using Surface-Modified Cu(In,Ga)Se2 Photocathodes. J. Mater. Chem. A 2015, 3, 8300−8307. (371) Yokoyama, D.; Minegishi, T.; Maeda, K.; Katayama, M.; Kubota, J.; Yamada, A.; Konagai, M.; Domen, K. Photoelectrochemical Water Splitting Using a Cu(In,Ga)Se2 Thin Film. Electrochem. Commun. 2010, 12, 851−853. (372) Contreras, M. A.; Ramanathan, K.; AbuShama, J.; Hasoon, F.; Young, D. L.; Egaas, B.; Noufi, R. Diode Characteristics in State-ofthe-Art ZnO/CdS/Cu(In1‑xGax)Se2 Solar Cells. Prog. Photovoltaics 2005, 13, 209−216. (373) Green, M. A.; Emery, K.; Hishikawa, Y.; Warta, W.; Dunlop, E. D. Solar Cell Efficiency Tables (Version 39). Prog. Photovoltaics 2012, 20, 12−20.

(337) Yuan, J.; Shao, C.; Zheng, L.; Fan, M.; Lu, H.; Hao, C.; Tao, D. Fabrication of CuInS2 Thin Film by Electrodeposition of Cu−In Alloy. Vacuum 2014, 99, 196−203. (338) Garuthara, R.; Wijesundara, R.; Siripala, W. Characterization of CuInS2 Thin Films Prepared by Electrodeposition and Sulfurization with Photoluminescence Spectroscopy. Sol. Energy Mater. Sol. Cells 2003, 79, 331−338. (339) Gunawan; Septina, W.; Ikeda, S.; Harada, T.; Minegishi, T.; Domen, K.; Matsumura, M. Platinum and Indium Sulfide-Modified CuInS2 as Efficient Photocathodes for Photoelectrochemical Water Splitting. Chem. Commun. 2014, 50, 8941−8943. (340) Luo, J.; Tilley, S. D.; Steier, L.; Schreier, M.; Mayer, M. T.; Fan, H. J.; Grätzel, M. Solution Transformation of Cu2O into CuInS2 for Solar Water Splitting. Nano Lett. 2015, 15, 1395−1402. (341) Iwase, A.; Ng, Y. H.; Amal, R.; Kudo, A. Solar Hydrogen Evolution using a CuGaS2 Photocathode Improved by Incorporating Reduced Graphene Oxide. J. Mater. Chem. A 2015, 3, 8566−8570. (342) Kaga, H.; Kudo, A. Cosubstituting Effects of Copper(I) and Gallium(III) for ZnGa2S4 with Defect Chalcopyrite Structure on Photocatalytic Activity for Hydrogen Evolution. J. Catal. 2014, 310, 31−36. (343) Quintans, C. S.; Kato, H.; Kobayashi, M.; Kaga, H.; Iwase, A.; Kudo, A.; Kakihana, M. Improvement of Hydrogen Evolution under Visible Light over Zn1−2x(CuGa)xGa2S4 Photocatalysts by Synthesis Utilizing a Polymerizable Complex Method. J. Mater. Chem. A 2015, 3, 14239−14244. (344) Wang, M.; Yang, S.; Li, H.; Yi, J.; Li, M.; Lv, X.; Niu, G.; Zhong, J. Double Pulse Electrodeposition of Cu-Ga Precursor Layer for CuGaS2 Solar Energy Thin Film. J. Electrochem. Soc. 2013, 160, D459−D464. (345) Niu, G.; Yang, S.; Li, H.; Yi, J.; Wang, M.; Lv, X.; Zhong, J. Electrodeposition of Cu-Ga Precursor Layer from Deep Eutectic Solvent for CuGaS2 Solar Energy Thin Film. J. Electrochem. Soc. 2014, 161, D333−D338. (346) Zhang, L.; Minegishi, T.; Nakabayashi, M.; Suzuki, Y.; Seki, K.; Shibata, N.; Kubota, J.; Domen, K. Durable Hydrogen Evolution from Water Driven by Sunlight Using (Ag,Cu)GaSe2 Photocathodes Modified with CdS and CuGa3Se5. Chem. Sci. 2015, 6, 894−901. (347) Zhang, L.; Minegishi, T.; Kubota, J.; Domen, K. Hydrogen Evolution from Water Using AgxCu1−xGaSe2 Photocathodes under Visible Light. Phys. Chem. Chem. Phys. 2014, 16, 6167−6168. (348) Kumagai, H.; Minegishi, T.; Moriya, Y.; Kubota, J.; Domen, K. Photoelectrochemical Hydrogen Evolution from Water Using Copper Gallium Selenide Electrodes Prepared by a Particle Transfer Method. J. Phys. Chem. C 2014, 118, 16386−16392. (349) Marsen, B.; Cole, B.; Miller, E. L. Photoelectrolysis of Water Using Thin Copper Gallium Diselenide Electrodes. Sol. Energy Mater. Sol. Cells 2008, 92, 1054−1058. (350) Moriya, M.; Minegishi, T.; Kumagai, H.; Katayama, M.; Kubota, J.; Domen, K. Stable Hydrogen Evolution from CdS-Modified CuGaSe2 Photoelectrode under Visible-Light Irradiation. J. Am. Chem. Soc. 2013, 135, 3733−3735. (351) Steichen, M.; Thomassey, M.; Siebentritt, S.; Dale, P. J. Controlled Electrodeposition of Cu−Ga from a Deep Eutectic Solvent for Low Cost Fabrication of CuGaSe2 Thin Film Solar Cells. Phys. Chem. Chem. Phys. 2011, 13, 4292−4211. (352) Oda, Y.; Minemoto, T.; Takakura, H. Electrodeposition of Crack-Free CuGaSe2 Thin Films from Single Bath. J. Electrochem. Soc. 2008, 155, H292−H295. (353) Muthuraj, J. J. J.; Rasmussen, D. H.; Suni, I. I. Electrodeposition of CuGaSe2 from Thiocyanate-Containing Electrolytes. J. Electrochem. Soc. 2011, 158, D54−D56. (354) Liu, F.; Yang, J.; Zhou, J.; Lai, Y.; Jia, M.; Li, J.; Liu, Y. OneStep Electrodeposition of CuGaSe2 Thin Films. Thin Solid Films 2012, 520, 2781−2784. (355) Wei, S.-H.; Zhang, S. B.; Zunger, A. Effects of Ga Addition to CuInSe2 on Its Electronic, Structural, and Defect Properties. Appl. Phys. Lett. 1998, 72, 3199−3201. AT

DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

(374) Ito, K.; Nakazawa, T. Electrical and Optical Properties of Stannite-Type Quaternary Semiconductor Thin Films. Jpn. J. Appl. Phys. 1988, 27, 2094−2097. (375) Ramasamy, K.; Malik, M. A.; O’Brien, P. Routes to Copper Zinc Tin Sulfide Cu2ZnSnS4 a Potential Material for Solar Cells. Chem. Commun. 2012, 48, 5703−5714. (376) Pawar, S. M.; Pawar, B. S.; Moholkar, A. V.; Choi, D. S.; Yun, J. H.; Moon, J. H.; Kolekar, S. S.; Kim, J. H. Single Step Electrosynthesis of Cu2ZnSnS4 (CZTS) Thin Films for Solar Cell Application. Electrochim. Acta 2010, 55, 4057−4061. (377) Jeon, M.; Shimizu, T.; Shingubara, S. Cu2ZnSnS4 Thin Films and Nanowires Prepared by Different Single-Step Electrodeposition Method in Quaternary Electrolyte. Mater. Lett. 2011, 65, 2364−2367. (378) Valdes, M.; Modibedi, M.; Mathe, M.; Hillie, T.; Vazquez, M. Electrodeposited Cu2ZnSnS4 thin films. Electrochim. Acta 2014, 128, 393−399. (379) Gurav, K. V.; Kim, Y. K.; Shin, S. W.; Suryawanshi, M. P.; Tarwal, N. L.; Ghorpade, U. V.; Pawar, S. M.; Vanalakar, S. A.; Kim, I. Y.; Yun, J. H.; et al. Pulsed Electrodeposition of Cu2ZnSnS4 Thin Films: Effect of Pulse Potentials. Appl. Surf. Sci. 2015, 334, 192−196. (380) Tao, J.; Liu, J.; He, J.; Zhang, K.; Jiang, J.; Sun, L.; Yang, P.; Chu, J. Synthesis and Characterization of Cu2ZnSnS4 Thin Films by the Sulfurization of Co-Electrodeposited Cu-Zn-Sn-S Precursor Layers for Solar Cell Applications. RSC Adv. 2014, 4, 23977−23984. (381) Araki, H.; Kubo, Y.; Jimbo, K.; Maw, W. S.; Katagiri, H.; Yamazaki, M.; Oishi, K.; Takeuchi, A. Preparation of Cu2ZnSnS4 Thin Films by Sulfurization of Co-Electroplated Cu-Zn-Sn Precursors. Phys. Status Solidi C 2009, 6, 1266−1268. (382) He, X.; Shen, H.; Pi, J.; Zhang, C.; Hao, Y. Synthesis of Cu2ZnSnS4 Films from Sequentially Electrodeposited Cu−Sn−Zn Precursors and Their Structural and Optical Properties. J. Mater. Sci.: Mater. Electron. 2013, 24, 4578−4584. (383) Scragg, J. J.; Dale, P. J.; Peter, L. M.; Zoppi, G.; Forbes, I. New Routes to Sustainable Photovoltaics: Evaluation of Cu2ZnSnS4 as an Alternative Absorber Material. Phys. Status Solidi B 2008, 245, 1772− 1778. (384) Rovelli, L.; Tilley, S. D.; Sivula, K. Optimization and Stabilization of Electrodeposited Cu2ZnSnS4 Photocathodes for Solar Water Reduction. ACS Appl. Mater. Interfaces 2013, 5, 8018−8024. (385) Stephen, R. G.; Riley, F. L. Oxidation of Silicon by Water. J. Eur. Ceram. Soc. 1989, 5, 219−222. (386) Bao, X.-Q.; Liu, L. Improved Photo-Stability of Silicon Nanobelt Arrays by Atomic Layer Deposition for Efficient Photocatalytic Hydrogen Evolution. J. Power Sources 2014, 268, 677−682. (387) Seger, B.; Tilley, D. S.; Pedersen, T.; Vesborg, P. C. K.; Hansen, O.; Grätzel, M.; Chorkendorff, I. Silicon Protected with Atomic Layer Deposited TiO2: Durability Studies of Photocathodic H2 Evolution. RSC Adv. 2013, 3, 25902. (388) Seger, B.; Pedersen, T.; Laursen, A. B.; Vesborg, P. C. K.; Hansen, O.; Chorkendorff, I. Using TiO2 as a Conductive Protective Layer for Photocathodic H2 Evolution. J. Am. Chem. Soc. 2013, 135, 1057−1064. (389) Morales-Guio, C. G.; Thorwarth, K.; Niesen, B.; Liardet, L.; Patscheider, J.; Ballif, C.; Hu, X. Solar Hydrogen Production by Amorphous Silicon Photocathodes Coated with a Magnetron Sputter Deposited Mo2C Catalyst. J. Am. Chem. Soc. 2015, 137, 7035−7038. (390) Fan, R.; Dong, W.; Fang, L.; Zheng, F.; Su, X.; Zou, S.; Huang, J.; Wang, X.; Shen, M. Stable and Efficient Multi-Crystalline n+p Silicon Photocathode for H2 Production with Pyramid-Like Surface Nanostructure and Thin Al2O3 Protective Layer. Appl. Phys. Lett. 2015, 106, 013902−013905. (391) Bao, X.-Q.; Petrovykh, D. Y.; Alpuim, P.; Stroppa, D. G.; Guldris, N.; Fonseca, H.; Costa, M.; Gaspar, J.; Jin, C.; Liu, L. Amorphous Oxygen-Rich Molybdenum Oxysulfide Decorated p-type Silicon Microwire Arrays for Efficient Photoelectrochemical Water Reduction. Nano Energy 2015, 16, 130−142. (392) Benck, J. D.; Lee, S. C.; Fong, K. D.; Kibsgaard, J.; Sinclair, R.; Jaramillo, T. F. Designing Active and Stable Silicon Photocathodes for

Solar Hydrogen Production Using Molybdenum Sulfide Nanomaterials. Adv. Energy Mater. 2014, 4, 1400739. (393) Yang, J.; Walczak, K.; Anzenberg, E.; Toma, F. M.; Yuan, G.; Beeman, J.; Schwartzberg, A.; Lin, Y.; Hettick, M.; Javey, A.; et al. Efficient and Sustained Photoelectrochemical Water Oxidation by Cobalt Oxide/Silicon Photoanodes with Nanotextured Interfaces. J. Am. Chem. Soc. 2014, 136, 6191−6194. (394) Kenney, M. J.; Gong, M.; Li, Y.; Wu, J. Z.; Feng, J.; Lanza, M.; Dai, H. High-Performance Silicon Photoanodes Passivated with Ultrathin Nickel Films for Water Oxidation. Science 2013, 342, 836− 840. (395) Chen, Y.; Tran, P. D.; Boix, P.; Ren, Y.; Chiam, S. Y.; Li, Z.; Fu, K.; Wong, L. H.; Barber, J. Silicon Decorated with Amorphous Cobalt Molybdenum Sulfide Catalyst as an Efficient Photocathode for Solar Hydrogen Generation. ACS Nano 2015, 9, 3829−3836. (396) Czochralski, J. Ein neues Verfahren zur Messung der Kristallisationsgeschwindigheit der Metalle. Z. Phys. Chem. 1918, 92, 219−221. (397) Cohen, U.; Huggins, R. A. Silicon Epitaxial Growth by Electrodeposition from Molten Fluorides. J. Electrochem. Soc. 1976, 123, 381−383. (398) Cohen, U. Some Prospective Applications of Silicon Electrodeposition from Molten Fluorides to Solar Cell Fabrication. J. Electron. Mater. 1977, 6, 607−643. (399) Rao, G. M.; Elwell, D.; Feigelson, R. S. Electrowinning of Silicon from K2SiF6-Molten Fluoride Systems. J. Electrochem. Soc. 1980, 127, 1940−1944. (400) Rao, G. M.; Elwell, D.; Feigelson, R. S. Electrodeposition of Silicon onto Graphite. J. Electrochem. Soc. 1981, 128, 1708−1711. (401) Elwell, D.; Rao, G. M. Mechanism of Electrodeposition of Silicon from K2SiF6-Flinak. Electrochim. Acta 1982, 27, 673−676. (402) Boen, R.; Bouteillon, J. The Electrodeposition of Silicon in Fluoride Melts. J. Appl. Electrochem. 1983, 13, 277−288. (403) Frazer, E. J.; Welch, B. J. A galvanostatic Study of the Electrolytic Reduction Silica in Molten Cryolite. Electrochim. Acta 1977, 22, 1179−1182. (404) De Mattei, R. C.; Elwell, D.; Feigelson, R. S. Electrodeposition of Silicon at Temperatures above Its Melting Point. J. Electrochem. Soc. 1981, 128, 1712−1714. (405) Takeda, Y.; Kanno, R.; Yamamoto, O.; Mohan, T. R. R.; Lee, C.-H.; Kröger, F. A. Cathodic Deposition of Amorphous Silicon from Tetraethylorthosilicate in Organic Solvents. J. Electrochem. Soc. 1981, 128, 1221−1224. (406) Lee, C. H.; Kröger, F. A. Cathodic Deposition of Amorphous Alloys of Silicon, Carbon, and Fluorine. J. Electrochem. Soc. 1982, 129, 936−942. (407) Rama Mohan, T. R.; Kröger, F. A. Cathodic Deposition of Amorphous Silicon from Solutions of Silicic Acid and Tetraethyl Ortho-Silicate in Ethylene Glycol and Formamide Containing HF. Electrochim. Acta 1982, 27, 371−377. (408) Munisamy, T.; Bard, A. J. Electrodeposition of Si from Organic Solvents and Studies Related to Initial Stages of Si Growth. Electrochim. Acta 2010, 55, 3797−3803. (409) Agrawal, A. K.; Austin, A. E. Electrodeposition of Silicon from Solutions of Silicon Halides in Aprotic Solvents. J. Electrochem. Soc. 1981, 128, 2292−2296. (410) Nishimura, Y.; Fukunaka, Y. Electrochemical Reduction of Silicon Chloride in a Non-Aqueous Solvent. Electrochim. Acta 2007, 53, 111−116. (411) Nicholson, J. P. Electrodeposition of Silicon from Nonaqueous Solvents. J. Electrochem. Soc. 2005, 152, C795−C802. (412) Gobet, J.; Tannenberger, H. Electrodeposition of Silicon from a Nonaqueous Solvent. J. Electrochem. Soc. 1988, 135, 109−112. (413) Vichery, C.; Le Nader, V.; Frantz, C.; Zhang, Y.; Michler, J.; Philippe, L. Stabilization Mechanism of Electrodeposited Silicon Thin Films. Phys. Chem. Chem. Phys. 2014, 16, 22222−22228. (414) Zein El Abedin, S.; Borissenko, N.; Endres, F. Electrodeposition of Nanoscale Silicon in a Room Temperature Ionic Liquid. Electrochem. Commun. 2004, 6, 510−514. AU

DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

(415) Liu, X.; Zhang, Y.; Ge, D.; Zhao, J.; Li, Y.; Endres, F. ThreeDimensionally Ordered Macroporous Silicon Films Made by Electrodeposition from an Ionic Liquid. Phys. Chem. Chem. Phys. 2012, 14, 5100−5105. (416) Borisenko, N.; Zein El Abedin, S.; Endres, F. In Situ STM Investigation of Gold Reconstruction and of Silicon Electrodeposition on Au(111) in the Room Temperature Ionic Liquid 1-Butyl-1methylpyrrolidinium Bis(trifluoromethylsulfonyl)imide. J. Phys. Chem. B 2006, 110, 6250−6256. (417) Al-Salman, R.; Mallet, J.; Molinari, M.; Fricoteaux, P.; Martineau, F.; Troyon, M.; El Abedin, S. Z.; Endres, F. Template assisted Electrodeposition of Germanium and Silicon Nanowires in an Ionic Liquid. Phys. Chem. Chem. Phys. 2008, 10, 6233−6237. (418) Mallet, J.; Molinari, M.; Martineau, F.; Delavoie, F.; Fricoteaux, P.; Troyon, M. Growth of Silicon Nanowires of Controlled Diameters by Electrodeposition in Ionic Liquid at Room Temperature. Nano Lett. 2008, 8, 3468−3474. (419) Mallet, J.; Martineau, F.; Namur, K.; Molinari, M. Electrodeposition of Silicon Nanotubes at Room Temperature using Ionic Liquid. Phys. Chem. Chem. Phys. 2013, 15, 16446−16449. (420) Gu, J.; Fahrenkrug, E.; Maldonado, S. Direct Electrodeposition of Crystalline Silicon at Low Temperatures. J. Am. Chem. Soc. 2013, 135, 1684−1687. (421) Cho, S. K.; Fan, F.-R. F.; Bard, A. J. Electrodeposition of Crystalline and Photoactive Silicon Directly from Silicon Dioxide Nanoparticles in Molten CaCl2. Angew. Chem., Int. Ed. 2012, 51, 12740−12744. (422) Tran, P. D.; Artero, V.; Fontecave, M. Water Electrolysis and Photoelectrolysis on Electrodes Engineered Using Biological and BioInspired Molecular Systems. Energy Environ. Sci. 2010, 3, 727−747. (423) Krawicz, A.; Cedeno, D.; Moore, G. F. Energetics and Efficiency Analysis of a Cobaloxime-Modified Semiconductor under Simulated Air Mass 1.5 Illumination. Phys. Chem. Chem. Phys. 2014, 16, 15818−15824. (424) Tomilov, A. P.; Smetanin, A. V.; Chernykh, I. N.; Smirnov, M. K. Electrode Reactions Involving Arsenic and Its Inorganic Compounds. Russ. J. Electrochem. 2001, 37, 997−1011. (425) de Mattei, R. C.; Elwell, D.; Feigelson, R. S. The Synthesis of GaAs by Molten Salt Electrolysis. J. Cryst. Growth 1978, 43, 643−644. (426) Dioum, I. G.; Vedel, J.; Tremillon, B. Properties of Arsenic in Molten Potassium Tetrachlorogallate at 300°C. J. Electroanal. Chem. Interfacial Electrochem. 1982, 139, 329−333. (427) Carpenter, M. K.; Verbrugge, M. W. Electrochemical Codeposition of Gallium and Arsenic from a Room Temperature Chlorogallate Melt. J. Electrochem. Soc. 1990, 137, 123−129. (428) Yang, M. C.; Landau, U.; Angus, J. C. Electrodeposition of GaAs from Aqueous Electrolytes. J. Electrochem. Soc. 1992, 139, 3480− 3488. (429) Gheorghies, L.; Gheorghies, C. Preparation of GaAs Thin Films from Acid Aqueous Solution. J. Optoelectron. Adv. Mater. 2002, 4, 979−982. (430) Mahalingam, T.; Lee, S.; Lim, H.; Moon, H.; Kim, Y. D. Electrosynthesis and Characterization of GaAs in Acid Solutions by Potentiostatic Method. Sol. Energy Mater. Sol. Cells 2006, 90, 2456− 2463. (431) Murali, K. R.; Trivedi, D. C. Preparation and Characterization of Pulse Electrodeposited GaAs Films. Phys. Status Solidi A 2006, 203, 874−877. (432) Chandra, S.; Khare, N.; Upadhyaya, H. M. Photoelectrochemical Solar Cells Using Electrodeposited GaAs and AlSb Semiconductor Films. Bull. Mater. Sci. 1988, 10, 323−332. (433) Murali, K. R.; Subramanian, V.; Rangarajan, N.; Lakshmanan, A. S.; Rangarajan, S. K. Preparation of GaAs Films by the Pulse Plating Technique. J. Mater. Sci.: Mater. Electron. 1991, 2, 149−151. (434) Gao, Y. K.; Han, A. Z.; Lin, Y. Q.; Zhao, Y. C.; Zhang, J. D. A Study of Electrodeposited GaAs Films. Thin Solid Films 1993, 232, 278−281.

(435) Gao, Y. K.; Han, A. Z.; Lin, Y. Q.; Zhao, Y. C.; Zhang, J. D. Electrodeposition and Characterization of GaAs Polycrystalline Thin Films. J. Appl. Phys. 1994, 75, 549−552. (436) Lajnef, M.; Chtourou, R.; Ezzaouia, H. Electric Characterization of GaAs Deposited on Porous Silicon by Electrodeposition Technique. Appl. Surf. Sci. 2010, 256, 3058−3062. (437) Kozlov, V. M.; Bozzini, B.; Bicelli, L. P. Formation of GaAs by Annealing of Two-Layer Ga-As Electrodeposits. J. Alloys Compd. 2004, 379, 209−215. (438) Andreoli, P.; Cattarin, S.; Musiani, M.; Paolucci, F. Electrochemical Approaches to GaAs1−xSbx Thin Films. J. Electroanal. Chem. 1995, 385, 265−268. (439) Villegas, I.; Stickney, J. L. Preliminary Studies of GaAs Deposition on Au (100), (110), and (111) Surfaces by Electrochemical Atomic Layer Epitaxy. J. Electrochem. Soc. 1992, 139, 686− 694. (440) Fahrenkrug, E.; Gu, J.; Maldonado, S. Electrodeposition of Crystalline GaAs on Liquid Gallium Electrodes in Aqueous Electrolytes. J. Am. Chem. Soc. 2013, 135, 330−339. (441) Cuomo, J. J.; Gambino, R. J. The Synthesis and Epitaxial Growth of GaP by Fused Salt Electrolysis. J. Electrochem. Soc. 1968, 115, 755−759. (442) De Mattei, R. C.; Elwell, D.; Feigelson, R. S. Conditions for Stable Growth of Epitaxial GaP Layers by Molten Salt Electrodeposition. J. Cryst. Growth 1978, 44, 545−552. (443) Heller, A.; Shalom, E. A.; Bonner, W. A.; Miller, B. HydrogenEvolving Semiconductor Photocathodes. Nature of the Junction and Function of the Platinum Group Metal Catalyst. J. Am. Chem. Soc. 1982, 104, 6942−6948. (444) Shalom, E. A.; Heller, A. Efficient p-InP(Rh-H alloy) and pInP(Re-H alloy) Hydrogen Evolving Photocathodes. J. Electrochem. Soc. 1982, 129, 2865−2866. (445) Ang, P. G. P.; Sammells, A. F. Hydrogen Evolution at p-InP Photocathodes in Alkaline Electrolyte. J. Electrochem. Soc. 1984, 131, 1462−1464. (446) Goodman, C. E.; Wessels, B. W. Function of Cobalt and Platinum on p-InP in the Photoevolution of Hydrogen from Alkaline Solutions. Appl. Phys. Lett. 1986, 49, 829−830. (447) Gorochov, O.; Stoicoviciu, L. Remarks on the Hydrogen Photoevolution at p-InP and Metallized p-InP Electrodes. J. Electrochem. Soc. 1988, 135, 1159. (448) Khaselev, O.; Turner, J. A. Electrochemical Stability of pGaInP2 in Aqueous Electrolytes toward Photoelectrochemical Water Splitting. J. Electrochem. Soc. 1998, 145, 3335−3339. (449) Uosaki, K.; Kita, H. Photocathodic Reactions at p-InP. Sol. Energy Mater. 1983, 7, 421−429. (450) Döscher, H.; Supplie, O.; May, M. M.; Sippel, P.; Heine, C.; Muñoz, A. G.; Eichberger, R.; Lewerenz, H.-J.; Hannappel, T. Epitaxial III-V Films and Surfaces for Photoelectrocatalysis. ChemPhysChem 2012, 13, 2899−2909. (451) Lin, Y.; Kapadia, R.; Yang, J.; Zheng, M.; Chen, K.; Hettick, M.; Yin, X.; Battaglia, C.; Sharp, I. D.; Ager, J. W.; et al. Role of TiO2 Surface Passivation on Improving the Performance of p-InP Photocathodes. J. Phys. Chem. C 2015, 119, 2308−2313. (452) Lee, M. H.; Takei, K.; Zhang, J.; Kapadia, R.; Zheng, M.; Chen, Y.-Z.; Nah, J.; Matthews, T. S.; Chueh, Y.-L.; Ager, J. W.; et al. p-Type InP Nanopillar Photocathodes for Efficient Solar-Driven Hydrogen Production. Angew. Chem., Int. Ed. 2012, 51, 10760−10764. (453) Gao, L.; Cui, Y.; Wang, J.; Cavalli, A.; Standing, A.; Vu, T. T. T.; Verheijen, M. A.; Haverkort, J. E. M.; Bakkers, E. P. A. M.; Notten, P. H. L. Photoelectrochemical Hydrogen Production on InP Nanowire Arrays with Molybdenum Sulfide Electrocatalysts. Nano Lett. 2014, 14, 3715−3719. (454) Hettick, M.; Zheng, M.; Lin, Y.; Sutter-Fella, C. M.; Ager, J. W.; Javey, A. Nonepitaxial Thin-Film InP for Scalable and Efficient Photocathodes. J. Phys. Chem. Lett. 2015, 6, 2177−2182. (455) Elwell, D.; Feigelson, R. S.; Simkins, M. M. Electrodeposition of Indium Phosphide. J. Cryst. Growth 1981, 51, 171−177. AV

DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

(456) Sahu, S. N. Aqueous Electrodeposition of InP Semiconductor Films. J. Mater. Sci. Lett. 1989, 8, 533−534. (457) Sahu, S. N. Structural, Optical and Electrical Properties of First Aqueous Electrodeposited InP Semiconductor Films. Sol. Energy Mater. 1990, 20, 349−358. (458) Sahu, S. N. Some Properties of the First Non-Aqueous ElectroCodeposited InP, In and P Thin Films. J. Mater. Sci.: Mater. Electron. 1992, 3, 102−106. (459) Lobaccaro, P.; Raygani, A.; Oriani, A.; Miani, N.; Piotto, A.; Kapadia, R.; Zheng, M.; Yu, Z.; Magagnin, L.; Chrzan, D. C.; et al. Electrodeposition of High-Purity Indium Thin Films and Its Application to Indium Phosphide Solar Cells. J. Electrochem. Soc. 2014, 161, D794−D800. (460) Ortega, J.; Herrero, J. Preparation of In X (X = P, As, Sb) Thin Films by Electrochemical Methods. J. Electrochem. Soc. 1989, 136, 3388−3391. (461) Cattarin, S.; Musiani, M.; Casellato, U.; Rossetto, G.; Razzini, G.; Decker, F.; Scrosati, B. Preparation of n- and p-InP Films by PH3 Treatment of Electrodeposited In Layers. J. Electrochem. Soc. 1995, 142, 1267−1272.

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DOI: 10.1021/acs.chemrev.5b00498 Chem. Rev. XXXX, XXX, XXX−XXX