Copper(I)-Based p-Type Oxides for Photoelectrochemical and

Review pubs.acs.org/cm

Copper(I)-Based p‑Type Oxides for Photoelectrochemical and Photovoltaic Solar Energy Conversion Ian Sullivan, Brandon Zoellner, and Paul A. Maggard* Department of Chemistry, North Carolina State University, Raleigh, North Carolina 27695-8204, United States ABSTRACT: Recent research efforts have been growing into p-type copper(I) based oxides for development of their use in solar energy applications. The oxides of interest include the binary Cu2O and a number of new ternary CuxMyOz oxides. Both the binary and ternary Cu(I)-based oxides have many advantages when compared to other well-known p-type oxides such as NiO, III−V, and II−VI semiconductors. The benefits found within the diverse group of Cu(I)-containing oxides include bandgap sizes that can be tuned from ∼1.2 to >3.0 eV, high charge carrier mobility, and favorable band energies relative to fuel-producing redox couples. These properties give them potential utility in a variety of different solar applications, such as in dye-sensitized solar cells and suspended powder photocatalysis. Research efforts into surface modifications and changes in their chemical compositions and structures have allowed for greater stability and greater efficiency in aqueous solutions, both of which have represented two key barriers for this class of materials. Presented in this review are the currently known binary and ternary Cu(I)-oxides and relationships of their syntheses and structures with their visible-light and ultraviolet bandgap sizes, band energies, and photoelectrochemical properties. As their constituent elements are relatively abundant and nontoxic, they represent an attractive class of materials that can be used in the conversion of sunlight to electricity or solar fuels.

I. INTRODUCTION I.A. Motivation. The world’s energy consumption was estimated to be 13.5 TW in 2001 and has been predicted to double to 27 TW by 2050.1 A majority of this energy is fossilfuel based and results in the emission of greenhouse gases. In 2013, ∼81% of the U.S. energy consumption was fossil-fuel based, including petroleum-based products, coal, and natural gas.2 While the demand for energy is growing, an attractive alternate source of energy is the Sun. The energy from sunlight incident upon the Earth’s surface in 1 h can theoretically supply all of the energy needed by the world in a single year.3 To harness this intermittent source of energy, sunlight needs to be absorbed and converted into either electrical power or chemical energy for storage. Semiconducting materials are excellent candidates for this application as they have the ability to absorb light and directly convert it to electrical power which can be either stored in batteries or used to split water or reduce carbon dioxide to chemical fuels. Copper(I)-based metal oxides are promising materials as ptype semiconductors for solar energy applications. These materials have a wide range of bandgap sizes, ranging from as low as ∼1.2 eV to greater than 3.0 eV, enabling their potential use in various types of system configurations for harvesting solar energy. They also can exhibit relatively high charge carrier mobilities4 (100 cm2 V−1 s−1) as compared to other p-type materials, such as NiO (0.53 cm2 V−1 s−1),5 as well as relatively better stability in aqueous media compared to III−V or II−VI semiconductors, such as p-GaAs or CdS. The p-type nature of these materials also facilitates photon-driven redox chemistry © 2016 American Chemical Society

that n-type semiconductors are not suitable for as a working electrode under irradiation, such as H2O or CO2 reduction.5,6 I.B. Background. Much research into solar energy and artificial photosynthesis has been focused on the use of metal oxides such as TiO2, ZnO, Fe2O3, and ZrO2.7 There are many different configurations in which these catalysts can be used to produce chemical or electrical energy from sunlight, such as suspended particle photocatalysis, as photoelectrodes for photoelectrochemical (PEC) water splitting, in photovoltaics, or as a component in dye-sensitized solar cells (DSSCs).8,9 These metal oxides are examples of n-type semiconductors and are typically used as photoanodes in PEC cells or for transport of electrons in DSSCs and PV cells. Many commonly studied ntype semiconductors are also wide bandgap materials, with bandgap sizes greater than 3.0 eV, which limits their light absorption to the UV region.10 A common limitation for smaller bandgap metal oxides is that they can suffer from low charge carrier mobility, as is the case for α-Fe2O3.11 As with all electrochemical systems, both oxidation and reduction reactions must take place. The n-type semiconductors are typically used for oxidation of water under irradiation, whereas p-type semiconductors are used for the reduction of water or CO2 under irradiation.12 For instance, in water splitting, water is oxidized to O2 at the photoanode, and the protons from this reaction are reduced to H2 at the photocathode. While much research has been reported in the area of oxidation/n-type Received: March 4, 2016 Revised: August 2, 2016 Published: August 4, 2016 5999

DOI: 10.1021/acs.chemmater.6b00926 Chem. Mater. 2016, 28, 5999−6016

Review

Chemistry of Materials

contact. The minority carriers, holes (h+), will travel to the surface/electrolyte interface when under irradiation.12 For an oxidation reaction to occur, the valence band of n-type photoelectrodes must be lower in energy (i.e., more positive potential) than the redox potential of interest, such as the oxidation of water to O2 (1.23 V vs SHE). Well-known examples of n-type semiconductors are TiO2 and α-Fe2O3.10,17 In p-type semiconductors, the opposite is true; downward band bending is present, and electrons, which are the minority charge carriers, will diffuse to the surface−electrolyte interface. Concomitantly, the majority carriers, i.e., holes, will diffuse to the back contact and be transported to the counter electrode. For a p-type metal oxide to be effective, the conduction band must be higher in energy (i.e., more negative potential) than the redox potential of interest, for example, H+/H2 in proton reduction (0.0 V vs SHE).18 For further details regarding band bending, extensive reviews can be found in the literature.19−21 While a semiconductor can absorb light that excites an electron from the valence band to the conduction band, the charge carriers must be separated and able to diffuse to the surface/electrolyte interface and to the back contact, allowing a current to flow. Thus, charge carrier mobility is another important characteristic in semiconductors. The higher the charge carrier mobility, the higher the conductivity that can be achieved in the material, allowing for diffusion of charge carriers to the surfaces and driving the redox chemistry. These are related in eq 1:

metal oxides, the reduction/p-type side has shown relatively less progress. While there are several related Cu(I)-containing chalcogenides such as CIS and CIGS that show excellent potential for solar energy conversion as photovoltaics, the scope of these is outside the current review.13,14 Herein, we report on binary and ternary Cu(I)-based oxides which show promise as p-type materials for solar energy conversion through H2O or CO2 reduction and for use as photovoltaic materials. I.C. Photoelectrochemical and Photovoltaic Overview. There are several requirements necessary for a material to be an effective semiconductor for solar energy capture and conversion.15 The ability to absorb a large fraction of visible light in the solar spectrum, stability against degradation, low cost, and nontoxicity are some of the required characteristics. Electronic characteristics of the semiconductor determine how it will function and how efficient it will be, including its space-charge width, absorption coefficient, electron−hole separation and lifetime, and charge carrier mobility.16 Shown in Figure 1 is an

σ = neμ

(1)

where σ is the conductivity (S·cm−1), n is the number of charge carriers per unit volume (cm−3), e is the elementary charge (C), and μ is the mobility of the charge carriers (cm2 V−1 s−1). The mobility is also related to the effective mass, m* (kg), shown in eq 2:

m* =

eτs μ

(2)

where τs is the mean scattering time (s). In this relationship the effective mass is inversely proportional to the mobility, i.e., a larger effective mass results in lower mobility. The effective mass is related to the band dispersion in the electronic structure in k-space, shown in eq 3: ⎛ ∂ 2E ⎞−1 m* = ℏ2⎜ 2 ⎟ ⎝ ∂k ⎠

(3)

where ℏ is the reduced Planck’s constant, E is energy, and k is a wave vector in k-space. Thus, semiconductors with a large band dispersion will exhibit a reduced effective mass and a higher conductivity. The physical thickness of a semiconducting film is another important factor, as there is a trade-off between the absorption coefficient and the charge carrier diffusion length. The charge carrier diffusion coefficient is given by eq 4:

Figure 1. Valence and conduction bands before equilibration for a ptype (a) and an n-type (c) semiconductor. After equilibration, downward band bending is present for p-type semiconductors (b) and upward band bending for n-type semiconductors (d). The diagrams from left to right are drawn to the same relative potential scale. CB, conduction band; VB, valence band; Ef, Fermi level; ERedox, redox potential of the electrolyte; VBB, band bending potential.19

illustration of the electronic energy levels of a p-type semiconductor (a) and an n-type semiconductor (c). As with all semiconductors, the conduction and valence bands are in a flat-band state prior to the equilibration of the Fermi levels at the surface−electrolyte interface. After the Fermi levels are equilibrated, the bands now exhibit a potential energy gradient near the surfaces, i.e., band bending, as shown in Figure 1b,d. The majority carriers for n-type semiconductors are electrons (e−) that diffuse from the space charge layer to the back

D=

kBT μ e

(4)

where D is the diffusion coefficient (cm2 s−1) of the charge carrier in n- or p-type semiconductors and kB is Boltzmann’s constant (eV·K−1).12 The diffusion coefficient is a relative measure of the distance the charge carriers can diffuse. The absorption coefficient is given by eq 5: 6000

DOI: 10.1021/acs.chemmater.6b00926 Chem. Mater. 2016, 28, 5999−6016

Review

Chemistry of Materials α=

I 1 ln 0 d I

states are thus regenerated. Ideally, the excited electron in the dye does not relax back to the ground state before hole injection (step 5) or recombine with the hole in the valence band (step 6). The valence band should be lower in energy than the LUMO of the dye molecule but higher in energy than the HOMO, allowing hole injection into the semiconductor. The conduction band should be sufficiently high in energy to keep it from participating in any of these photoelectrochemical steps.24 The reduction potential of the redox mediator should be lower in energy than the LUMO of the dye molecule, allowing electrons to reduce the mediator. Research into new combinations of metal-oxide semiconductors, dyes, and redox mediators can be found in the recent literature.22−29 Photovoltaic solar cells (PV cells) are based on heterojunctions of p-type and n-type materials (p-n junctions). These can be composed of a single element (i.e., p-Si and n-Si), or metal chalcogenides, metal oxides, or a combination thereof. As the Fermi levels of the p- and n-type semiconductors equilibrate, this results in a built-in potential gradient at their interfaces. This determines the maximum photovoltage of the device and the efficiency of the cell. Many different parameters govern the efficiency of the PV cell, including bandgap size, band alignment, and resistivity.30 For both DSSCs and PV cells, the power conversion efficiency is given by eq 6:

(5)

where α is the absorption coefficient (cm−1), d is the depth of light penetration into the semiconductor (cm), I is the incident light intensity (W), and I0 is the transmitted light intensity. Ideally the carrier diffusion length is on the order of the absorption coefficient, in order to enable all photon-generated charge carriers to reach the surfaces. If there is a large difference between the two, higher recombination rates will most likely occur. For instance, α-Fe2O3 is an n-type semiconductor with a diffusion coefficient of ∼0.01 cm2 s−1 and a hole diffusion length of ∼2−4 nm. Owing to its low diffusion coefficient and short hole diffusion length, there is a relatively higher probability that the excited charge carriers will not reach the surface/electrolyte interface before recombining. This has been cited as the main reason for the limited performance of α-Fe2O3 as a viable photoanode for water oxidation.11 In DSSC photovoltaics, a wide band gap semiconductor (Eg > 3.0 eV) acts as a charge transport layer that has been sensitized, such as by a molecular dye or by quantum dots,22,23 in order to extend light absorption into the visible-light range. The valence and conduction band positions relative to the energy levels of the dye molecule and redox mediator are of great importance, as this will determine (thermodynamically) the maximum efficiency of the solar cell.8 Figure 2 illustrates an

η = VocIsc

FF × 100 Ic

(6)

where η is efficiency, Voc is the open circuit voltage (V), Isc is the short circuit current (A), FF is the fill factor, and Ic is the power of incident light (W/cm2). The fill factor is given by eq 7:

FF =

IscVoc ImpVmp

(7)

where Imp and Vmp are the current and voltage at maximum power, respectively. The efficiency can be calculated from values found in current−density potential curves (j−v curves).30,31 In-depth reviews on semiconductor electrochemistry can be found in the literature.12,32−34 Faradaic efficiency is also another important aspect in solar energy harvesting and is determined by the amount of charge passed in order to form moles of a product during an electrochemical reaction. For instance, the reduction of protons is a two electron process, and the Faradaic efficiency (η) is given by eq 8:35

Figure 2. Energy level diagram for a p-DSSC. The dye absorbs light, and an electron is promoted to an excited state (1); hole injection then occurs between the dye and the p-type semiconductor (2); the dye is now in a reduced state and should transfer an electron to the redox mediator (3); the redox mediator then diffuses to the counter electrode and is oxidized by holes from the semiconductor (4). Other processes, such as relaxation of the electron from the LUMO to the HOMO (5) or recombination of the excited dye (6) or redox mediator (7), represent competitive processes that can diminish the efficiency.24

η=

mol H2·2e−·F × 100 Q

(8)

where F is Faraday’s constant (C/mol) and Q is the total charge passed (C). Total charge can be calculated by integration of the current produced during the electrochemical reaction over time. The Faradaic efficiency can be determined by measuring the product amount and comparing to the theoretical amount that should have been produced. If there are other electrochemical reactions that take place (i.e., selfreduction or corrosion) or a large resistance causing current loss, then the Faradaic efficiency will be less than 100%. The Faradaic efficiency is also necessary to determine the solar-tofuel efficiency (SFE), as given by eq 9:35

energy level schematic of the conduction and valence bands, as compared with the dye molecule and redox mediator energy levels, for a p-type DSSC. In step 1, light excites an electron from the HOMO of the dye to the LUMO. In step 2, hole injection into the valence band of the semiconductor occurs, leaving a reduced dye. In step 3, the dye should then transfer the electron to the redox mediator (M). The positive charge from the initial hole-injection is collected at the counter electrode and oxidizes the redox mediator in step 4. The initial 6001

DOI: 10.1021/acs.chemmater.6b00926 Chem. Mater. 2016, 28, 5999−6016

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Chemistry of Materials

and ascorbic acid for the decomposition of methyl orange (MO).45 The Cu2+ ions are converted to Cu(OH)2 upon addition of the NaOH solution and are then reduced by the ascorbic acid to Cu2O nanoparticles of ∼100 nm in diameter. Through UV−vis/DRS measurements, the absorption was found to center at ∼470−490 nm, resulting in a band gap of ∼2.5−2.6 eV. While the authors report that the blue shift is expected from quantum size effects,46 it cannot be overlooked that the first allowed transition in Cu2O dominates at smaller particle sizes.41 The concentration of NaOH was found to be an important aspect of the synthesis, as high concentrations of NaOH yielded cubic shaped particles of ∼1 μm in size. These exhibited a band gap of ∼2.1 eV, consistent with bulk Cu2O.47 The cubic morphology of the larger particles had a larger decomposition rate of the methyl orange dye, as compared to the spherical nanoparticles. While the nanoparticles had a larger surface area for the decomposition, there is no well-defined facet surface owing to its spherical shape, thereby inhibiting a more efficient diffusion of electrons to the surface/electrolyte interface. Susman et al. found that the particle morphology growth of Cu2O nanocrystals on Au-seeded glass substrates using citrate, tartrate, and EDTA solutions could be controlled based on the chelating salt.48 Depending on the complexing salt and time used in the synthesis, octahedra, truncated octahedra, cuboctahedra, and cubes varying in sizes from ∼197 nm to ∼73 nm were observed. Bandgap sizes were also shown to be blueshifted as the particle size decreased, e.g., as compared to bulk Cu2O (Eg: ∼ 2.1 eV), ranging from ∼2.3 eV to ∼2.5 eV. Exposure of only the most active facets through precise morphology control allowed for surface effects to yield the most efficient photocatalysis. Films of Cu2O have also been intensely investigated for the conversion of solar energy to chemical fuels by means of the water splitting reaction. A tandem Cu2O/BiVO4 photoelectrochemical cell was reported by Bornoz et al. for overall water splitting.49 The backside of the cell with BiVO4 was irradiated first because of its larger band gap (∼2.5 eV) compared to Cu2O. This allowed for the lower energy light to pass through to the frontside of the Cu2O electrode. A film thickness of 50 nm was found to be optimal for the backside irradiated BiVO4, allowing more photons to pass through to the frontside Cu2O electrode. While there was a modest conversion efficiency for this cell, it nonetheless demonstrated that unassisted water splitting can be performed using the Cu2O/ BiVO4 combination. The ease with which Cu2O is electrodeposited or thermally grown on Cu substrates is advantageous in making controllable film thicknesses for photovoltaic applications.50 This has allowed Cu2O to be used as a p-type layer in photovoltaics in many prior studies over the years.51−56 Septina et al. investigated the electrochemical deposition of Cu2O on fluorine-doped tin oxide (FTO) electrodes under potentiostatic conditions.57 Potentials used ranged from −0.50 V to −0.65 V vs Ag/AgCl with copper metal being formed at −0.65 V. The films exhibit a cubic morphology on the surface of the FTO substrate with a preferential orientation of the (111) plane. The thicknesses of the Cu2O films were controlled by adjusting the applied potential, during which a total current of ∼1.4 C/cm2 was passed through the cell. Films of ∼2.0 μm were developed under a potential range of −0.5 V to −0.55 V. As the applied potential was increased to −0.6 V and −0.65 V, the thickness of the films decreased to 1.6 and 1.5 μm, respectively. Photocurrents of up to ∼7.1 mA/cm2 were observed under simulated AM 1.5 irradiation using films made

(1.23 V) ·J ·η (9) I where 1.23 V is the thermodynamic potential of water splitting, J is the current density (mA/cm2), and I is the power of incident irradiation (mW/cm2). While photocurrents may be produced in a photoelectrochemical cell, it is critical to always verify that the current is a result of the desired electrochemical reactions (i.e., H2 and O2 formation) in order to determine the Faradaic efficiency of the system. SFE =

II. COPPER(I) OXIDES II.A. Cuprous Oxide (Cu2O). Research into the p-type semiconductor Cu2O can be found throughout the literature owing to its use as an electrode, catalyst, photoelectrode, and a p-type layer in PV cells.36−40 In its pure form, Cu2O is a cubic material crystallizing in the space group Pn3̅m. Copper is linearly coordinated between tetrahedrally coordinated oxygen atoms within the unit cell, shown in Figure 3.41 The band gap

Figure 3. Crystal structure of Cu2O, with the unit cell outlined; labeled atom types are Cu (blue) and O (red).

size of Cu2O is ∼2.1 eV, with the conduction and valence bands straddling the redox potentials of water oxidation and proton reduction. It therefore meets theoretical requirements to thermodynamically split water into hydrogen and oxygen gas under illumination,42 as illustrated alongside several other Cu(I)-based oxides in Figure 14. There are several reasons why Cu2O has been such a widely studied material; it is relatively abundant, is easy to deposit as a film, is nontoxic, and has a relatively high charge carrier mobility among p-type oxides. High-purity single crystals of Cu2O grown thermally on sheets of Cu have a reported carrier mobility of up to ∼1.8 × 105 cm2 V−1 s−1,4 while films of Cu2O with grain sizes of ∼1−100 mm2 have exhibited mobilities of up to ∼100 cm2 V−1 s−1.43 Control of particle morphology and size has been shown to be a critical component of maximizing the efficiency of photocatalytic surface reactions.44 For example, specific crystallographic directions of a crystallite may have a higher or lower mobility for electrons and/or holes owing to anisotropies in band dispersion. Further, smaller sized particles may have a larger surface area and increase the amount of available active sites for oxidation or reduction reactions. Nanoparticles of Cu2O are easily synthesized with various particle morphologies. Zhang et al. have synthesized spherical Cu2O nanoparticles using aqueous solutions of CuSO4, NaOH, 6002

DOI: 10.1021/acs.chemmater.6b00926 Chem. Mater. 2016, 28, 5999−6016

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Chemistry of Materials with the deposition potential of −0.60 V, along with an additional 200 nm layer of sputtered aluminum zinc oxide (AZO) in a solar cell configuration. These conditions resulted in a compact, planar Cu2O layer with no voids in the film, lowering the leakage current and resulting in the highest efficiency. The impact of AZO and other coating layers will be covered in more detail below. While p-type Cu2O meets most requirements for an effective photocathode, it is not thermodynamically stable in aqueous solutions. The redox couple for the reduction of Cu2O to Cu lies within the band gap of the material, and it will self-reduce under aqueous conditions, as represented in eq 10.58 Cu 2O + 2H+ + 2e− → 2Cu + H 2O

Heterojunctions of Cu2O with other wide bandgap n-type oxides have achieved high power conversion efficiencies, especially with AZO and Ga2O3 as reported by Minami, Nishi, and Tanaka.43,53,63 Among the highest power conversion efficiencies (∼6.1%) for a Cu2O-based photovoltaic is that recently reported by Minami et al.64 Copper sheets were thermally oxidized to form Cu2O on the surface, followed by a treatment with sodium which produced Na-doped Cu2O in order to lower the overall resistivity to as low as 10−1 Ω cm. The composited and transparent n-type layer of AZO/ Al0.025Ga0.975O exhibited the lowest series resistance and highest shunt resistance, and which correlated with the atomic percentage of Al in the Al−Ga−O layer. Similarly, Nishi et al. have reported a power conversion efficiency of 4.12% for a Cu2O based photovoltaic.43 Again, Cu metal was thermally oxidized to Cu2O with a hole mobility of ∼100 cm2 V−1 s−1. Thin layers of ZnO were deposited by pulsed laser deposition (PLD), ranging from ∼0−150 nm in thickness. Next, Al-doped ZnO (AZO) was deposited on the ZnO layer, completing the cell. An optimal thickness of the ZnO layer was ∼50 nm that gave the largest power conversion efficiencies for the cell. Other thicknesses showed similar fill factors (FF) and open circuit voltages (Voc). However, the 50 nm layer had the highest transmittance, enabling larger shortcircuit currents (Isc). Much recent work from Minami and coworkers can be found regarding the optimization of Cu2O based photovoltaics.65,66 While great progress has been made in the generation of relatively larger and more stable photocurrents and hydrogen production using p-Cu2O by several groups, the largest photocurrents produced are around 10 mA/cm2, or about two-thirds of its maximum theoretical photocurrent obtainable, based on 100% efficiency of light absorption and conversion to photocurrent.67,68 II.B. Advantages of Cu(I)-Based Mixed-Metal Oxides. While Cu2O shows growing promise as a photoelectrode, the band gap of the material (∼2.1 eV) only absorbs a portion of the solar spectrum, limiting the theoretical maximum efficiency to ∼20% with a maximum possible photocurrent of ∼15 mA/ cm2.67,69 Recently, many ternary metal oxides have been investigated as alternative p-type semiconductors that are similarly comprised of a Cu(I)-based valence band, but which are also combined with an early transition metal (e.g., Nb(V), Ta(V), and V(V)) that serves to form the conduction band. Depending upon the selection of the early transition metal and the Cu(I) coordination environment, as well as the extended structural connectivity, bandgap sizes as low as ∼1.2 eV and up to >3.0 eV can be attained. The tunability of this class of p-type oxides can be used to tailor them to the specific needs of DSSCs, as Z-scheme tandem solar cells,70 in suspended powder photocatalysts, or within multijunction photovoltaic cells. By allowing for the absorption of a wider range of wavelengths and favorable band energies with respect to mediator or fuelproducing redox couples, larger photocurrents and higher efficiencies can thus be obtained. Another advantage of mixed-metal oxides is the potential for better stabilization of the Cu(I) cation within the material. As noted before, Cu2O will self-reduce under irradiation in aqueous solution. This is owing to the excitation of an electron from the copper 3d10 valence band to the primarily copper 4s conduction band.41 Adding another metal with unfilled d orbitals allows for electrons to be excited into its conduction band and helps to inhibit the reduction of Cu(I) at the surfaces.

(10)

The formation of Cu0 at the surface of Cu2O photoelectrodes has been observed through XPS after irradiation in aqueous solution.36 To inhibit the self-reduction process, several groups have investigated protecting Cu2O by adding cocatalysts to the surface or by depositing thin layers of other metal oxides/ chalcogenides on the surface of the Cu2O photoelectrodes.59−61 Paracchino et al. deposited layers of Al-doped ZnO (AZO), TiO2, and Pt nanoparticles on Cu2O films.42 This results in a staggered type-II band offset, allowing the excited electrons from Cu2O to travel through the layers to the electrolyte and drive the water reduction reaction. Current densities of nearly 8 mA/cm2 were observed at a 0 V vs RHE (pH = 4.9), as well as a Faradaic efficiency of nearly 100% for H2 production.42 Morales-Guio et al. investigated a similar surface-protected Cu2O electrode using amorphous molybdenum sulfide.61 Again, layers of AZO and TiO2 were deposited on the surface of Cu2O, followed by the MoS2+x layer which acts as a hydrogen evolution reaction (HER) surface. A Na2SO4 solution was used as the electrolyte, and H2 evolution was performed over a range of different pH conditions. The MoS2+x layer was reported to protect both the Cu2O and the TiO2 layers from the acidic solution longer than a similar electrode with Pt on the surface. Chronoamperometric measurements indicated a stable current density of ∼5.0 mA/cm2 at a pH of 4.0 for 10 h. The layering of various oxides and chalcogenides functioned to protect the Cu2O from aqueous conditions and provided favorable band offsets, allowing for the efficient transfer of electrons to the surface/electrolyte interface while preventing the reduction of Cu2O. Photocathodes of Cu2O nanowires have been reported by Luo et al. exhibiting current densities of up to 10 mA/cm2 in aqueous media at −0.3 V vs RHE.62 Nanowires with an average thickness of ∼100−300 nm and length of ∼1−3 μm were grown by heating anodized Cu substrates under an Ar atmosphere at different times and temperatures to produce Cu2O. To protect the Cu2O nanowires from self-reduction in aqueous media, layers of n-type AZO and TiO2 were deposited by ALD to form a buried p−n junction, followed by the photoelectrochemical deposition of RuOx to enable efficient hydrogen evolution. Photoelectrochemical measurements of the layered films were made at an applied potential of 0.0 V vs RHE at a pH = 5 in aqueous solution. The photocathode exhibited 100% Faradaic efficiency for over 55 h, as well as a current density of over 4 mA/cm2. These large photocurrents were attributed to the larger surface area of the nanowires, as well as efficient charge collection at their surfaces owing to the thickness of the nanowires being on the order of the minority carrier diffusion length. 6003

DOI: 10.1021/acs.chemmater.6b00926 Chem. Mater. 2016, 28, 5999−6016

Review

Chemistry of Materials

shown to produce a current of ∼0.02 mA/cm2 under simulated solar conditions in an aqueous environment. The photocurrents were increased to ∼0.05 and 0.22 mA/cm2 as the polycrystalline films of Cu3VO4 were heated in air at 300 and 350 °C for 15 min. When heated in air, Cu(I) is oxidized to Cu(II) which increased the amount of Cu(I) vacancies within its structure. Scanning electron microscopy and energy dispersive spectroscopy indicated the formation of CuO and nanorods of Cu3V2O8 covering all sides of the particles. Linear sweep voltammetry experiments, e.g., comparing bare Cu3VO4 to when CuO and Cu3V2O8 are present at the surface, illustrated the key role of these Cu(II) oxides at the surfaces in increasing the photocurrents. The photocurrent was similarly reproduced when CuO was chemically deposited onto the surfaces of Cu3VO4. Greater surface coverage led to higher cathodic photocurrents that were measured for over 4000 s of irradiation at a potential of 0 V vs SCE. II.C.ii. Copper Niobates. Research into the Cu(I)-containing niobate system has currently yielded the synthesis of CuNbO3,79−81 CuNb3O8,82 Cu2Nb8O21,77 and CuNb13O33.83 From this set of compounds, CuNbO3, Cu2Nb8O21, and CuNb3O8 have been studied for their photocatalytic and photoelectrochemical properties. The first reported synthesis of CuNbO3 by Sleight and Prewitt was carried out under 65 kbar in a temperature range of 1000 to 1200 °C, which produced black crystals with a monoclinic cell.79 It was determined that this initial synthesis produced a mixture of Cu(I) and Cu(II) cations within the crystals. Later investigations by Wahlstrom and Marinder resulted in the successful synthesis of CuNbO3 containing only the Cu(I) cation.80,81 The CuNbO3 phase was found to crystallize in the space group C2/m, as shown in the structure illustrated in Figure 5. Linearly bonded Cu(I) cations bridge between two corrugated layers of corner-sharing Nb4O16 clusters that are composed of edge-sharing NbO6 octahedra.80 Two layers of Nb4O16 clusters stack along the top and bottom of the Cu(I) layer, giving the stoichiometry of CuNbO3 for the unit cell. In recent research by the Maggard group,84,85 the CuNbO3 phase was prepared by solid-state methods and found to exhibit an indirect bandgap size of ∼2.0 eV, consistent with its cherryred color. Electronic band-structure calculations showed that electrons are excited from the filled Cu 3d10 orbitals to the empty Nb 4d0 orbitals that make up the majority of the valence and conduction band states, respectively. Polycrystalline electrodes of CuNbO3 were prepared on FTO and irradiated under solar-simulated light (AM 1.5 G, ∼100 mW/cm2) in an aqueous 0.5 M Na2SO4 solution adjusted to a pH of ∼12. The polycrystalline films had been heated in air from 250 to 500 °C for 3 h each. As this heating temperature increased, larger cathodic photocurrents were produced, of up to ∼1.5 mA/cm2 after heating to an optimal temperature of 350 °C. After heating to temperatures greater than 350 °C, the CuO impurity phase was observed by powder XRD and the photocurrent decreased. Recently, a Nb/Ta solid solution, i.e., Cu(Nb1−xTax)O3, was synthesized by the solid state route and found to result in decreased dark currents in the polycrystalline films while maintaining similar photocurrent densities.85 However, photocatalytic rates for hydrogen production during suspended particle photocatalysis were observed to decrease with an increasing substitution of Ta into the structure. Under visible-light irradiation, photocatalysis by suspended CuNbO3 particles produced ∼142 μmol/g of hydrogen. However, the amount of hydrogen production dropped to ∼90 μmol/g when

II.C. Group-V Copper(I) Oxides. In contrast to Cu2O, group-V copper oxides have varying structural features and a wide range of bandgap sizes. While Cu(I) ions are typically linearly coordinated in metal oxides, the group-V copper oxides contain Cu coordinated in distorted tetrahedral and octahedral environments. These materials exhibit a wide range of visiblelight absorption, with bandgap sizes ranging from ∼1.2 to 2.7 eV. Solid solutions of these materials can show dramatic shifts in their absorption edges, e.g., allowing one semiconducting oxide system to span the entire visible range of the solar spectrum, while also maintaining favorable band energies. The latter chemical tunability can allow a concomitant tuning of the bandgap size and band energies around specific redox couples. II.C.i. Copper Vanadates. There are few known coppervanadate compounds containing the Cu(I) cation, e.g., Cu3VO4 and CuVO3. Other known examples of copper vanadates contain either Cu(II) or a mixture of Cu(I) and Cu(II). Initial synthesis of Cu3VO4 produced black crystals with a space group of I4̅2m symmetry.71,72 A unique feature within the crystal structure is the tetrahedral coordination environment of the Cu(I) cation, together with vanadium that is also in a tetrahedral coordination geometry, as illustrated in Figure 4.

Figure 4. Crystal structure of Cu3VO4 with the unit cell outlined; labeled atom types are Cu (blue), O (red), and V (orange polyhedra).

This places Cu3VO4 into a relatively small group of known cuprous oxides that do not have linearly coordinated Cu(I) cations, including CuNb3O8 and Cu2Nb8O21.73−77 Notably, this change in coordination geometry results in the highest Cu(I)based valence band energies (i.e., all more positive than the O2/ H2O redox couple) and smallest bandgap sizes (i.e., ∼1.2 to 1.4 eV) within these systems, as shown in Figure 14 and listed in Table 1. Recently, Sahoo et al. from the Maggard group investigated the photoelectrochemical properties of Cu3VO4 as a visiblelight active photocathode for the reduction of water to H2.78 The small bandgap size of ∼1.2 eV allowed for most of the visible-light spectrum to be absorbed (∼70%), as a result of excitation of electrons from the filled Cu 3d10 orbitals to the empty V 3d0 orbitals. Electronic structure calculations confirmed that copper and vanadium make up the majority of the states at the top of the valence band and bottom of the conduction band, respectively. Calculations also show a surprisingly wide conduction band dispersion (∼2 eV), especially given the lack of an extended vanadate connectivity. Despite the low band gap of Cu3VO4, polycrystalline films were 6004

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Chemistry of Materials Table 1. Bandgap Sizes, Photocurrents, Stability, and Cu(I)-Coordination Environments in Cu(I)-Containing Oxides

material

band gap (eV)

maximum photocurrent (mA/cm2);a potential (V)b

pH

Cu coordination environment

2.00

7.6; 0.0

4.9

Linear

5.7; 0.0

5.0

10.0; −0.3

5.0

5; 0.0

5.0

4.3; 0.0

5.0

1.54; 0.0

6.8

0.18; −0.138

11 5.8

Tetrahedral

PEC cell and photocatalysis measurement conditions

stability J/J0 (%); time

H2 Faradaic efficiency (%)/H2 evolution rate

78%; 20 min

100%

99%; 10 h

100%

70%; 55 h

100%

94%; 8 h

100%

87.7%; 20 min 74.4%; 20 min N/A

84%

N/A

N/A

20%; 4000 s

N/A

80%; 5000 s

62% 12.3 μmol/h

N/A

N/A

95%; 11000 s

142 μmol/g

50%; 1000 s

N/A

47%; 1000 s

N/A

N/A

0.1 μmol/h

25−50%; 1000 s

0.6 μmol/h

N/A

Cu3VO4

1.20

2.08−10; 0.0 (two electrode) 0.25; −0.042

CuNb3O8

1.26

0.85; −0.108

5.8

Octahedral

Cu2Nb8O21

1.43

1.3; −0.358

12

Tetrahedral

CuNbO3

2.00

0.45; −0.358

6.3

Linear

α-Cu2Ta4O11

2.65

0.5; −0.358

6.5

Linear

β-Cu2Ta4O11

2.65

1.5; −0.358

6.5

Linear

Cu3Ta7O19

2.60

2.6; −0.358

6.5

Linear

Cu5Ta11O30

2.55

2.5; −0.358

6.5

Linear

CuFeO2

1.36

1.5; −1.26

6.8

Linear

from LSV measurements; layered film with AZO, TiO2 and Pt; ref 42 from LSV measurements; layered film with AZO and MoS2−x; ref 61 from LSV measurements; Cu anodized to CuO and reduced to Cu2O; ref 62 from LSV measurements; layered film with AZO, TiO2 and RuO2; ref 59 from LSV measurements; layered with CuO and Ni nanoparticles; ref 60 from LSV measurements; Cu/Cu2O/CuO layered; ref 50 from LSV measurements; thin film in aqueous media; ref 39 from Isc condition; photovoltaic cells Cu2O/ZnO, highest efficiency 6.1%; refs 43, 52, 53, 56, 66 from LSV measurements; oxidized film electrode; ref 78 from LSV measurements; oxidized film electrode; ref 74;suspended particle catalysis in methanol with a 1% Pt cocatalyst; ref 75 from LSV measurements; oxidized film electrode; ref 77 from LSV measurements; oxidized film electrode; ref 84; suspended particle catalysis in methanol with 1% Pt cocatalyst; ref 85 from LSV measurements; oxidized film electrode; ref 100 from LSV measurements; film electrode; oxidized; ref 100 from LSV measurements; oxidized film electrode; ref 95; suspended particle catalysis in methanol with 0.5% Pt; ref 96 from LSV measurements; oxidized film electrode; refs 95, 97; suspended particle catalysis in methanol with 0.5% Pt; ref 96 from LSV measurements; pellet electrode; ref 94

CuRhO2 CuAlO2

1.90 3.00

7.0; −1.26 0.3; 0.0 (two electrode) 0.954; 0.0 (two electrode) 0.384; 0.0 (two electrode) 0.415; 0.0 (two electrode) 2.05; 0.0 (two electrode) N/A

14

Linear Linear

from LSV measurements; pellet electrode; ref 115 from Isc condition; 0.04% efficiency; DSSC; ref 120

60%; 10 h N/A

N/A (10% CO2 reduction) 80% N/A

from Isc condition; 0.04% efficiency; DSSC; lref 121

N/A

N/A

from Isc condition; 0.026% efficiency; DSSC; ref 122 N/A

N/A

from Isc condition; 0.045% efficiency; 1% Mg doped; N/A DSSC; ref 123 from Isc condition; 0.182% efficiency; DSSC; ref 124 N/A

N/A

Wurtzite structure; refs 118, 119

N/A

Cu2O

α-CuGaO2

β-CuGaO2

3.00

1.50

-

Linear

-

Tetrahedral

N/A

N/A N/A

N/A

a

All values are changed to positive values but may be reported as negative for convention. bAll values are shifted to reflect the potentials vs RHE (Ag/AgCl + 0.197; SCE + 0.242).

Cu0.81ICu0.17IINb2.97O8.82 While the structure type is the same as that reported for LiNb3O8, there are significant differences in the Li and Cu positions that lead to the highly distorted tetrahedral coordination environment of the Cu(I) cations. Despite the distortion at the copper sites, the NbO6 octahedra remain nearly identical between the two phases.82 Their similar structures suggested that a solid solution could be formed between the two phases across all possible compositions, as was found and reported for Li1−xCuxNb3O8.76 As lithium is replaced with copper, the bandgap size shows a dramatic decrease in size from ∼3.89 eV (x = 0) to ∼1.45 eV (x = 1) owing to the introduction of the higher lying Cu 3d10 orbitals that form a

6% Ta was added, i.e., CuNb0.94Ta0.06O3. When the maximum percentage of tantalum was included in the structure (x = 0.25), the hydrogen production dropped to ∼67 μmol/g. The relatively stable photocurrents with increasing Ta substitution, but decreasing photocatalytic rates, would be consistent with a decreasing Faradaic efficiency of the polycrystalline films. Another Cu(I) niobate that has been investigated as a photoelectrode and photocatalyst is CuNb3O8, shown with the unit cell illustrated in Figure 6. The CuNb3O8 phase was initially found during phase composition studies in the copper− niobium−oxygen system and suggested to have copper deficiencies that made the stoichiometry closer to 6005

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within a CuCl flux.77 Single crystal X-ray diffraction was used to determine the new composition and crystal structure, as shown in Figure 7. Its structure is unlike any of the previously

Figure 5. Crystal structure of CuNbO3 as viewed along the [010] direction, with the unit cell outlined; labeled atom types are Cu (blue), O (red), and Nb (green polyhedra). Figure 7. Crystal structure of Cu2Nb8O21 as viewed along the [010] direction, with the unit cell outlined; labeled atom types are as in Figure 5.

described phases in this system and consists of chains of distorted CuO4 tetrahedra and layers of NbO7 pentagonal bipyramids and monocapped trigonal prisms. The pentagonal bipyramidal coordination environment is predominantly observed within the Cu(I) tantalate system,86 such as found for Cu5Ta11O3087 and Cu3Ta7O19.88 One other example of NbO7 polyhedra can be found in the solid solution NaCu(Ta1−xNbx)4O11 phase, wherein niobium partially substitutes for tantalum that is in a pentagonal bipyramidal coordination environment within the structure of NaCuTa4O11.89 However, the Cu2Nb8O21 phase thermally decomposes beginning at only ∼250 °C owing to the high mobility of Cu(I) cations within the structure and their disproportionation at the surfaces. The Cu2Nb8O21 phase exhibits a small bandgap size of ∼1.43 eV, the second smallest bandgap size in this system. Mott− Schottky measurements locate the potentials of the valence and conduction bands at ∼0.12 V and ∼−1.5 V vs RHE (pH = 12), respectively, as shown in Figure 14. The position of the conduction band provides an overpotential of ∼0.8 V to drive the reduction of water, while its valence band is less positive than the oxidation potential of water. Polycrystalline electrodes of Cu2Nb8O21 were prepared and their photoelectrochemical properties investigated when irradiated with visible light. No appreciable photocurrent was produced by the electrode over a scanned potential from 0.2 to −0.6 V (vs SCE). However, when the polycrystalline film was heated at 350 and 450 °C for 3 h, photocurrents of ∼0.25 and 0.75 mA/cm2 were observed together with a relatively large dark current. The increase in photocurrent arises from the formation of CuO at the surfaces of the particles as the electrode material was heated in air. If the islands are removed from the surface, the photocurrents decrease by up to 10−50%. If the electrode was heated at temperatures of 500 °C or higher, significant decomposition of Cu2Nb8O21 into CuNb2O6 and Nb2O5 occurs with diminishing photocurrents.77 II.C.iii. Copper Tantalates. The family of p-type copper tantalates includes Cu2Ta4O11, Cu3Ta7O19, and Cu5Ta11O30. A representative crystal structure of Cu3Ta7O19 is shown in Figure 8. Compounds in this family all contain the local α-U3O8related structural feature, shown in Figure 9, and share the general formula AxM3n+1O8n+3 (e.g., A = Na, Ca, Ag, Cu; M = Nb or Ta). Their structures are composed of single (n = 1) and/or double layers (n = 2) of edge-shared pentagonal

Figure 6. Crystal structure of CuNb3O8 as viewed along the [010] direction, with the unit cell outlined; labeled atom types are as in Figure 5.

new valence band. The decreased band gap allowed for the materials to function as visible-light active photocatalysts that generated H2 from aqueous solutions. The most active compound for H2 production was found to be a 50:50 mixture of lithium and copper (Li0.5Cu0.5Nb3O8), yielding a total of 18.29 μmoles of H2 gas after visible-light irradiation for 1 h with a 1% Pt cocatalyst. The p-type CuNb3O8 has also been investigated by the Maggard group as a polycrystalline photoelectrode. Through a combination of diffuse reflectance spectroscopy and Mott− Schottky analysis, the positions of the valence and conduction bands were found to be located at ∼0.55 V and ∼−0.71 V vs RHE (pH 6.3).74 Current densities of up to ∼0.4 mA/cm2 were found after heating the films to 350 and 450 °C for 3 h. A maximum incident photon-to-current efficiency of ∼7% was achieved under monochromatic irradiation of ∼352 nm. In a separate study by King et al., nanoislands of CuO were found to form at the surface of CuNb3O8 when heated in air at temperatures of 450 °C and greater, as described below in Section IV and shown in Figure 13.75 This heat treatment causes Cu(I) to migrate from the interior of the particles to the surface and react with O2 in air, forming CuO. These CuO surface islands have been postulated to help to form a type-II band offset that drives more efficient change separation at the surfaces. The recently discovered copper niobate, Cu2Nb8O21, was synthesized by a solvothermal reaction of CuCl and Li3NbO4 nanoparticles, or by reacting nanoparticles of Cu2O and Nb2O5 6006

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of visible light down to ∼2.5 to 2.6 eV, similarly for all members of this system. For example, the substitution of Cu for Na in the solid solution Na2−xCuxTa4O11 reduces the bandgap size by ∼1.4 eV, from ∼4.0 eV (x = 0) to ∼2.6 eV (x = 0.78).93 Electronic structure calculations of the x = 0 and x = 0.78 compositions showed that as the Cu(I) percentage increases, an increasing contribution of Cu 3d10 orbitals comprises a new higher-energy valence band. This red shift arises from the introduction of Cu 3d10 orbitals, which lie higher in energy than the O 2p in Na2Ta4O11. The bandgap sizes for these materials are closely similar owing to the relatively constant Cu 3d10 valence band energy levels and high lying Ta 5d0 based conduction band.90 As a p-type material, this high energy conduction band yields a highly negative reduction potential, which, under band gap excitation, can potentially drive the reduction of H+ and/or CO2. The excitation of an electron from the Cu 3d10 based valence band to the Ta 5d0 based conduction band can also help to stabilize the material from self-reduction as compared to Cu2O.94 The photoelectrochemical properties of the Cu5Ta11O30 and Cu3Ta7O19 phases were first investigated as p-type polycrystalline films by Fuoco et al.95 Increasing cathodic photocurrents were observed when polycrystalline films of these materials were heated in air from 350 to 550 °C for 3 h each, ranging from ∼0.2 mA/cm2 when not heated in air, to ∼1.5 mA/cm2 at 350 °C and 2.6 mA/cm2 after heating in air to 550 °C. Mott− Schottky measurements were used to determine the valence and conduction band potentials as well as to estimate the charge carrier density (∼1016 cm−3). These materials exhibited a valence band potential of 1.06 V for Cu5Ta11O30 and 1.19 V for Cu3Ta7O19 at pH 6.3 vs NHE and conduction band energies of −1.54 V and −1.28 V, respectively. As illustrated in Figure 14, these band potentials straddle both the oxidation and the reduction potentials for water (−0.372 V H+/H2, 0.858 V O2/H2O), making it a suitable material for total water splitting. However, their photocatalytic activities for hydrogen production as suspended powders show relatively low rates as compared to the Cu(I) niobates. Work by Kato et al. focused on the related solid solution of Cu3xLa1−xTa7O19, wherein the lanthanum was replaced by Cu(I) in the structure.96 A full substitution from a pure lanthanum to a pure Cu(I)-containing phase was achieved owing to the similarities between the LaTa7O19 and Cu3Ta7O19 structures. As a result, the bandgap size was red-shifted by ∼1.6 eV. These compounds containing copper were found to reduce water to hydrogen in a methanol/ water solution after depositing a 0.5% Pt (by weight) cocatalyst on the surface. The most active composition, Cu1.8La0.4Ta7O19, was found to produce ∼2.6 μmol/h of H2 while the pure copper compound produced ∼0.1 μmol/h. Further detailed studies of the effects of both heating time and temperature on the Cu5Ta11O30 films were reported by Sullivan et al.97 After heating polycrystalline films in air at 350− 550 °C for 15, 30, and 60 min each, CuO nanoparticles were observed via SEM to increasingly grow in size and surface coverage on the crystallites. The presence of Cu(II) was confirmed through XPS measurements of the surfaces. Powder X-ray diffraction (XRD) measurements showed a decreased unit cell volume and a copper-deficient formula of Cu3.2(1)Ta11O30 from Rietveld refinements, indicating a significant loss of copper from the compound with heating in air. It was found that Cu(I) preferentially migrates out of the ab planes of the crystallites and is oxidized to CuO nanoparticles at the surfaces, as illustrated in Figure 10. These CuO

Figure 8. Crystal structure of Cu3Ta7O19 along the [010] direction, with the unit cell outlined; labeled atom types are Cu (blue), O (red), and Ta (green polyhedra).

Figure 9. Polyhedral view of the local α-U3O8 structural feature common to all members of the copper(I) tantalate family; labeled atom types are as in Figure 8.

bipyramids, as known for Cu3Ta7O19 and Cu5Ta11O30,90 respectively. The value of n may also be nonintegral, and then represents a mix of single and double layers as found for Cu2Ta4O11.91 These pentagonal bipyramidal layers alternate with layers of isolated TaO6 octahedra and A site cations that can be 2, 7, or 8-coordinated. The TaO6 octahedra are constructed from the apical oxygen atoms from two adjacent layers of edge-sharing TaO7 pentagonal bipyramids.92 Owing to close similarities in their crystalline structures, these materials exhibit similar optical, electronic, and photoelectrochemical properties. The incorporation of Cu(I) into these tantalate phases yields significantly lowered bandgap sizes that facilitate the absorption 6007

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Figure 10. (a) Polyhedral model of the hexagonal crystal morphology of Cu5Ta11O30 with arrows indicating the preferential directions of copper migration within the ab plane and (b) a local structural view of the Cu(I)/TaO6 layer with the three symmetry-equivalent directions for copper migration labeled with arrows. Reprinted with permission from ref 97. Copyright 2014 American Chemical Society.

Figure 11. Two main polymorphic modifications of the delafossite structure that crystallize in the trigonal R3̅m (a) and hexagonal P63/mmc (b) space groups.

the visible-light range. Nanocrystalline films were prepared and sensitized with two different Zn-porphyrin based dyes, with light absorption at ∼2.0 eV and sensitizing them to visible-light wavelengths. Spectroelectrochemical measurements of the dyes confirmed that the HOMO and LUMO were at the proper energy levels for hole injection. The photophysical analysis of this system found fluorescence quenching of the dyes (an indication of hole injection), as well as hole injection lifetimes of 8 ps and Nb > V, yielding the decreasing bandgap size. The delafossites have a more complex electronic structure owing to the different dn configurations of the trivalent metal cations, with several computational investigations focused on the orbital mixing between the conduction and valence bands.109,113 Generally, however, the Cu(I) delafossites follow a similar trend. The sbased conduction bands of CuBO2, CuAlO2, CuGaO2, and CuInO2 (≥3.0 eV band gaps) all are higher in energy than the lower lying d-based bands of CuFeO2 and CuRhO2 (∼1.4 eV and ∼1.9 eV band gaps, respectively).94,115 The Cu(I) coordination environment also has a significant impact on their bandgap sizes as well, with a higher coordination number leading to relatively smaller band gaps, as listed in Table 1. For example, the Cu(I) niobates show a large variation in bandgap size, ranging from ∼1.26 to ∼2.0 eV, owing in part to the change in Cu(I) coordination in the structures. The smallest bandgap size of 1.26 eV occurs for CuNb3O8 with the relatively uncommon octahedrally-coordinated Cu(I) cation. Next, Cu2Nb8O21 follows with a band gap 6010

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Figure 13. SEM images of the particle morphologies and surfaces of Cu(I)-containing oxides, including Cu3VO4 (a, b), CuNb3O8 (c, d), and Cu5Ta11O30 (e, f). The formation of Cu(II)-containing oxides at the surfaces are shown in b, d, and f, after heating each in air. Parts (a) and (b) adapted with permission from ref 78. Copyright 2015 The Royal Society of Chemistry. Parts (c) and (d) and (e) and (f) reprinted with permission from refs 75 and 97. Copyright 2014 American Chemical Society.

of ∼1.6 eV and contains tetrahedrally coordinated Cu(I) cations. Finally, CuNbO3 has the largest band gap of ∼2.0 eV and has linearly coordinated Cu(I). This trend is also observed for the two different CuGaO2 polymorphs. The α-CuGaO2 phase has a band gap of ∼3.6 eV owing to higher lying Ga 4s conduction band states;122 however, Omata et al. have reported a β-CuGaO2 phase which has a band gap of 1.47 eV with a mainly Cu 3d-based valence band and a Ga 4s-based conduction band.118 Again, the coordination environment of Cu(I) is a key factor, as the delafossite type α-CuGaO2 has a linearly coordinated Cu(I), while the wurtzite type β-CuGaO2 phase has tetrahedrally coordinated Cu(I). Similarly, the copper tantalates all have very similar structure types and linearly coordinated Cu(I), all with relatively larger and similar bandgap sizes of ∼2.5−2.6 eV.91,95

movement of Cu(I) cations was found to occur along the [100] and [010] directions, allowing for the preferential migration to specific crystal facets, as illustrated in Figure 10b. In each of these studies, it has been found that photocurrent densities were increased after heating polycrystalline films of the group-V copper oxides in air. This is attributed to the process of Cu(I) being oxidized to Cu(II), inducing p-type defects and increasing the conductivity of the semiconductors. However, investigations into the addition CuO to the surfaces of Cu5Ta11O30 and Cu3VO4, with minimal oxidation of Cu(I), have shown increased photocurrents as well. This is owing to a favorable type-I or type-II band offset in the heterojunction of the parent material and Cu(II)-oxide at the surface (i.e., for CuO or Cu2V3O8), allowing for increased charge transport from the surfaces to the electrolyte and increased photocurrents.42,97 Another possible enhancement arises from the increased visible-light absorption, such as especially for the larger band gap Cu(I) tantalates. As CuO forms on the surface, the color of the material darkens from yellow to black, and subband gap absorption is observed by UV−vis diffuse-reflectance spectroscopy. This allows for light absorption as low as ∼1.5 eV for Cu3Ta7O19, ∼1.3 eV for Cu5Ta11O30, and 3.0 eV. This is dependent upon the energetic position of the conduction band as determined by the choice of the second metal cation (i.e., Nb5+, In3+, or Fe3+), as well as structural differences, such as the coordination environment of the Cu(I) cations in the structure. The ability to tune their band gaps facilitates the absorption of a broader range of wavelengths of the solar spectrum. The graphic in Figure 14 illustrates some of the band positions relative to water redox couples for several Cu(I) oxides, clearly showing that most have band positions thermodynamically favorable for CO2 reduction and/or for overall water splitting. Future promising research directions in this area include surface protecting layers as used very effectively for Cu2O (e.g., Al-doped ZnO, TiO2), investigation of heterojunction effects with CuO, and dye-sensitization of the wide band gap compounds (e.g., CuAlO2). Several investigations have shown that heterojunctions of p-type Cu(I) oxides with other n-type metal oxides can effectively separate charge carriers and result in large photocurrents. As an additional effect, this can also

■

AUTHOR INFORMATION

Corresponding Author

*(P.A.M.) E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors acknowledge support from the Department of Chemistry at North Carolina State University. P.A.M. also thanks the Research Corporation for Science Advancement for a Scialog award.

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REFERENCES

(1) Lewis, N. S.; Nocera, D. G. Powering the Planet: Chemical Challenges in Solar Energy Utilization. Proc. Natl. Acad. Sci. U. S. A. 2006, 103, 15729−15735. (2) U.S. DOE. Annual Energy Outlook 2015; 2015; p 154. (3) Tsao, J.; Lewis, N.; Crabtree, G. Solar FAQs; U.S. Department of Energy: 2006; pp 1−24. (4) Ito, T.; Yamaguchi, H.; Okabe, K.; Masumi, T. Single-Crystal Growth and Characterization of Cu2O and CuO. J. Mater. Sci. 1998, 33, 3555−3566. (5) Flynn, C. J.; Oh, E. E.; McCullough, S. M.; Call, R. W.; Donley, C. L.; Lopez, R.; Cahoon, J. F. Hierarchically-Structured NiO Nanoplatelets as Mesoscale P-Type Photocathodes for Dye-Sensitized Solar Cells. J. Phys. Chem. C 2014, 118, 14177−14184.

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DOI: 10.1021/acs.chemmater.6b00926 Chem. Mater. 2016, 28, 5999−6016

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DOI: 10.1021/acs.chemmater.6b00926 Chem. Mater. 2016, 28, 5999−6016