Cu-Based Materials as Photocatalysts for Solar Light Artificial Photosynthesis: Aspects of Engineering Performance, Stability, Selectivity

Abstract

Cu-oxide nanophases (CuO, Cu₂O, Cu⁰) constitute highly potent nanoplatforms for the development of efficient Artificial Photosynthesis catalysts. The highly reducing conduction band edge of the d-electrons in Cu₂O dictates its efficiency towards CO₂ reduction under sunlight excitation. In the present review, we discuss aspects interlinking the stability under photocorrosion of the (CuO/Cu₂O/Cu⁰) nanophase equilibria, and performance in H₂-production/CO₂-reduction. Converging literature evidence shows that, because of photocorrosion, single-phase Cu-oxides would not be favorable to be used as a standalone cathodic catalyst/electrode; however, their heterojunctions and the coupling with proper partner materials is an encouraging approach. Distinction between the role of various factors is required to protect the material from photocorrosion, e.g., use of hole scavengers/electron acceptors, band-gap engineering, nano-facet engineering, and selectivity of CO₂-reduction pathways, to name a few possible solutions. In this context, herein we discuss examples and synthesis efforts that aim to clarify the role of interfaces, faces, and phase stability under photocatalytic conditions.

1. Introduction

Over the last decades, the negative impact of greenhouse gas emissions on the environment is strongly correlated with air pollution, climate change, and global warming. Carbon dioxide (CO₂) is considered one of the main greenhouse gases, thus strategies need to be developed in order to decrease CO₂ levels in the atmosphere. The “Artificial Photosynthesis” approach [1] aims to exploit photocatalytic technology, that is the use of solar photons to produce hydrogen (H₂) and—ideally—couple it to CO₂ reduction towards carbon-based fuels. Economically and environmentally, this is a sustainable, circular economy approach since CO₂ reduction can result in useful products such as formic acid (HCOOH), formaldehyde (HCHO), methanol (CH₃OH), methane (CH₄), and carbon monoxide (CO), to name a few [2].

Both photocatalytic H₂ evolution from H₂O and photocatalytic CO₂ reduction reactions (CO₂RR) share some common steps, with their difference being on the specific surface reactions of the photogenerated electrons. These steps include, see Figure 1: photon-adsorption by the semiconductor, electron-hole separation and their migration to the semiconductor surface [3], and subsequent incorporation of the electron to a H⁺, leading towards H₂ or to one CO₂ radical, leading towards the HCOOH/HCOH/CH₃OH/CH₄ chain. H₂ production can be identified as a less-complex process, compared to the CO₂ reduction, the latter involving CO₂ adsorption and activation, product formation, and desorption. Thus, CO₂ reduction is a multi-proton and multi-electron transfer process and selectivity of the products remains a major challenge (see Figure 1).

The difficulty in activating the inert CO₂ molecule lies in its closed-shell electronic configuration, linear geometry, and D_∞h symmetry [2]. Thus, it requires a highly negative reduction potential of −1.9 V (vs. NHE at pH 7) to activate CO₂ and form the CO₂⁻ radical, which the vast majority of semiconductors cannot provide [2].

Fujishima and Honda in 1972, [4], paved the way for light-induced water splitting by TiO₂ and, since then, numerous studies on photocatalysts were presented, and in particular TiO₂ [5,6,7]. Moreover, research interest in photocatalytic and photo-electrochemical applications has increased, prompting successful engineering of various semiconducting photocatalysts [5,8,9,10]. So far, TiO₂ remains the most studied photocatalyst, due to its chemical stability, low cost, and availability [11]. However, the large energy gap of TiO₂ (3.2 eV) [3] limits its photoactivity exclusively in the ultraviolet radiation range, typically λ < 360 nm, which accounts for 2–5% of the sunlight [12]. Moreover, the conduction band (CB) edge of TiO₂ is E_CB = −160 mV vs. NHE (pH = 0) or E_CB = −50 mV (pH = 7) [3], which disfavors its photoexcited electrons to efficiently achieve reduction of CO₂ towards the key intermediate CO₂⁻. In this context, p-type Cu₂O is considered as highly promising and attractive since it can absorb visible-light photons, E_g in the range 2.0–2.2 eV [13,14,15,16], and mostly due to the highly reducing energy positioning of its conduction-band edge, which is approximately E_CB = −1000 mV vs NHE (pH = 0) (see Figure 2) [10].

As an element, Cu has high abundance on Earth’s crust, ease of handling and manufacturing, low cost, and low environmental toxicity. In the family of stable Cu-oxides, Cu₂O and CuO are direct-bandgap semiconductors with E_g in the range 2.0–2.2 eV (Cu₂O) and 1.3–1.7 eV (CuO) [13,14,15,16], respectively. The variations in E_g values depend mainly on the particle-morphology and nano- vs. bulk-size [17]. The small bandgap energies allow Cu₂O and CuO to absorb the majority of solar-photons while their direct-bandgaps endow the two oxides with large photon-absorption coefficients [18,19]. Additionally, it is reported that copper-based catalysts and the different Cu oxidation states can significantly improve the kinetics and thermodynamics of CO₂ adsorption, activation, and CO dimerization [20,21]. Li et al. [9] discuss the selectivity of various cocatalysts for the photocatalytic CO₂ reduction. In this context, metallic Cu⁰ is more favorable for the production of hydrocarbons, primarily CH₄, while Cu₂O favors CH₃OH production [2,9].

While CuO is known to be the most stable copper oxide, the stability of Cu₂O under photochemical processes, is a matter of concern. The usually reported lack of long-term stability can be attributed to “photocorrosion”, a term referring to the occurrence of lattice-destabilization linked to consumption of the photogenerated carriers, see Figure 2, affecting its overall photo-performance [22]. These events, together with eventual enhanced electron-hole recombination rate in Cu₂O and CuO, may be detrimental for Hydrogen-Evolution Reaction photocathodes [13,18,23,24,25]. Pertinent review articles have been published in the literature on the synthesis of Cu-oxides [18,26,27,28], their use as photocatalysts for pollutant degradation, H₂-production from water [6,29], and CO₂ reduction [10,20,30].

Herein, the aim of this mini-review is to discuss aspects on the photostability of copper oxides, CuO and Cu₂O, with focus on some specific, pertinent strategies that have been—so far—reported to improve the photostability. The first part of the review focuses on the photostability problems and the self-decomposition mechanism of CuO and Cu₂O. Then, we discuss the different material-engineering approaches such as heterojunctions, core-shell structures, strain engineering, combination with supports. Finally, we provide a targeted overview of some promising Cu-based materials for artificial photosynthesis and discuss the potential strategies employed to enhance the photocatalytic performance. For conciseness, the present review focuses on the reducing process (H₂-evolution, CO₂ reduction). The oxidizing-side (holes) optimization and mechanisms are out of the scope of the present review.

2. The Problem of Photocorrosion and Some Specific Approaches to Prevent It

From the redox point of view, photocorrosion of Cu₂O can be explained, see Figure 2, as being triggered either:

[i]

by the oxidation of Cu¹⁺ to Cu²⁺, by photogenerated holes. Formally, this tends to convert Cu₂O to CuO.

[ii]

by the reduction of Cu¹⁺ to Cu⁰ by photogenerated or electrochemically provided electrons. Formally, this tends to convert Cu₂O to metallic Cu⁰.

Many nanoengineering approaches aiming to diminish photocorrosion are based on the working hypothesis of the prevention of electron-hole recombination and promoting them out of the Cu-oxide lattice. Hereafter, see Figure 3, we summarize some approaches in this direction.

2.1. Case−Study A: Photocathode-Heterojunctions (e.g., Cu_xO/TiO₂, CuO/BiVO₄, CuO/Cu₂O)

CuO can be an effective photocatalyst and a photocathode if the self-reduction problem of the Cu²⁺ sites to Cu¹⁺ is solved via a heterojunction with a n-type semiconductor such as TiO₂ [19,34]. This TiO₂@CuO core-shell scheme can act as a protective layer, see Figure 3 (Case-Study A). The efficiency of this approach can be influenced by the preparation method. Masudy-Panah et al. [24], via control of the sputtering power, achieved to optimize the TiO₂-layer over CuO, and this in turn had a beneficial effect on the photocorrosion stability of CuO-based photocathodes [24]. Under five hours of photocorrosion stability tests, increasing the sputtering power was found to significantly suppresses the formation of Ti³⁺, which increased the Hydrogen Evolution Reaction (HER) efficiency of the photocathode [24]. The Nyquist-curves recorded by Electrochemical Impedance Spectroscopy (EIS) of thin TiO_2-film indicated easier transfer of photoinduced electrons to the electrolyte solution [24]. Xing et al. [26] reported that a protective layer of TiO₂ formed by sol-gel method, can be used to protect the CuO photocathode [26]. In this case, they had achieved improved photocurrent density by TiO₂@CuO vs. bare CuO photocathode. In [26], the long-term stability was impaired due to a less-homogeneous protective TiO₂ layer on the CuO photocathode [26]. Specifically, in the TiO₂@CuO core-shell, it was suggested [26] that photoexcited electrons of CuO and TiO₂ easily migrate to the CB of TiO₂ and the holes from VB of TiO₂ transfer to the VB of CuO, while the internal electric field further promotes the separation and transfer of photogenerated carriers in the two components of the Z-heterojunction direction, resulting to a more stable and efficient CuO photocathode [26]. Similarly, the results of [35,36] can be understood as being due to a Z-scheme configuration [37] that is considered to improve electron-hole separation and transfer them to the CB of TiO₂ and VB of Cu₂O, respectively.

In [23], precise control of Oxygen-rich CuO and Cu-rich CuO regions on the same material was investigated via sputtering on Fluorine Tin Oxide (FTO)-coated glass. It was observed that the stability of CuO electrodes was considerably influenced by the precise local balance of Cu or O elements, with O-rich CuO electrodes to be most stable against photocorrosion [23]. O-rich materials show higher stability against photocorrosion, which might be a result from slower self-reduction of CuO to Cu₂O. EIS data showed that, in O-rich CuO electrodes, charge transfer resistance was decreased [23], thus improving interfacial charge transport and photocatalytic performance.

Jeong et al. [38] had fabricated BiVO₄/CuO heterojunction electrodes by spin-coating capping layers of BiVO₄ on CuO photoelectrodes, in order to prevent self-reduction of CuO to Cu₂O. Using X-Ray Photoelectron Spectroscopy (XPS) data, they have observed that CuO photoelectrodes with poor photostability occurred when self-reduction of CuO to Cu₂O was occurring, while after repeated BiVO₄ deposition cycles, the photocorrosion was suppressed, i.e., a > 76% photostability was reported [38].

In conclusion, heterojunction of Cu-oxide semiconductors with appropriate n-type semiconductors can be a promising strategy to address the issue of Cu²⁺/Cu¹⁺ self-reduction. Properly accounting for the relative CB and VB positioning, and the ensuing band-bending, is of primary importance in this approach.

2.2. Case−Study B: The Case of e⁻ Capture/Acceptor (CuO/Pt and CuO/Pd-Au)

In another approach, to prevent Cu²⁺ to Cu¹⁺ self-reduction [23], an oxygen-rich CuO film was coated with a noble-metal nanolayer, e.g., Pt or Pd-Au, using sputtering-technology. Improved stability was confirmed by EIS, showing enhanced CuO stability to correlate with decrease of the resistance [23], i.e., photoinduced electrons had higher mobility [23]. Noticeably, in that work, electron-hole generation was improved via the interaction of plasmonic Au with the semiconductor [23]. In a similar approach, adding Pt as an electron accepting catalyst on a CuO was reported in [26], using an electro-deposition method [26] (see Figure 3 Case-Study B). By adding a metal with higher work function than CuO [Φ = 5.3 eV] e.g. such as Pt [Φ = 5.65 eV], the photogenerated electrons can migrate from CuO to the Pt layer, thus be trapped there, until thermodynamic equilibrium is established, forming an inner electric-field between CuO and Pt [39]. The benefit of this approach, is that, by decreasing electron accumulation on CuO, it minimizes the self-reductive corrosion [26].

In an alternative approach, Zhang et al. [40] had anchored single Cu-atoms (Cu-SA) on TiO₂ structure by a bottom-up approach, where a metal organic framework (MOF) MIL-125 was used as a substrate to ensure atomic dispersion of atomic cocatalyst and enabled the highest loading amount of approximately 1.5 wt.% Cu. Anchoring Cu on the TiO₂ matrix significantly decreased the charge carrier recombination. Loading single Cu-atoms on TiO₂ versus pristine TiO₂ indicated that the Cu-SA loading might effectively facilitate the transfer of photogenerated electrons from TiO₂ to the Cu active sites [40]. This can be attributed to the fact that the redox-potential of Cu²⁺/Cu¹⁺ (0.16 V vs. NHE) is more positive than CB of TiO₂ (−0.1 V vs. NHE). Electrochemical Impedance Spectroscopy Nyquist curves of Cu-SA/TiO₂ verify that the Cu-SA species can act as electron acceptors, thus facilitating the interfacial charge separation [40]. The decreased electron-hole recombination was also evident by the photoluminescence data [40], i.e., in Cu-SA TiO₂ the recombination of electrons and holes was slower. As a result, long-term stability was reported, i.e., 380 days with minor loss on the catalytic performance [40].

The aforementioned examples, indicate that fast transfer of electrons out from the Cu-oxide matrix, in a non-reversible way, is the key-condition to improve Cu²⁺ to Cu¹⁺ self-reduction. Thus, the redox positioning of the electron-acceptors vs. the CB of the Cu-atoms in Cu-oxides, is one of the key-parameters to be considered.

2.3. Case−Study C: The Use of Hole Scavengers

As exemplified in Figure 2, in Cu₂O, self-oxidation of Cu¹⁺ to Cu²⁺ by photogenerated holes, can be a primary photocorrosion source [31]. The highly negative potential CB-edge of Cu₂O, typically −1000 mV vs. NHE (pH = 0) [10] boosts the electron transfer to interfacial H⁺, towards H₂ production. Concomitantly, h⁺ are gradually accumulated and this may promote photocorrosion, oxidation of Cu¹⁺ to Cu²⁺, since the low level of valence band (VB) prevents h⁺ to participate in water oxidation. Thus, an efficient hole scavenger would be required to achieve a rapid withdrawal of h⁺ in order to suppress this Cu₂O self-oxidation (see Figure 3 Case-Study C inset). The presence of a h⁺ scavenger with a suitable oxidizing potential (see Figure 4) facilitates the h⁺ transfer, thus improving the photostability of Cu₂O. In this context, careful selection of a sacrificial hole-scavenger may improve stability by preventing self-photo-oxidation of Cu₂O. This in turn, is expected to enhance the overall photocatalytic H₂ evolution by Cu₂O, permitting an educing potential to be built-up. In this context, Toe et al. [31] demonstrated that among some hole scavengers, e.g., Na₂SO_3, methanol, and ethanol, Na₂SO₃ was the most effective in photocatalytic H₂ evolution by Cu₂O catalysts [31].

Importantly, this study showed that via this approach [31], photocatalytic H₂-production from [H₂O+scavenger] could be accomplished without any secondary components compared, for example, to the common case of use of alcohols (methanol, isopropanol) as hole scavengers. In [31], Na₂SO₃ enabled h⁺ scavenging to oxidize SO₃²⁻ into SO₄²⁻, leaving free e⁻ in the conduction band for HER. Importantly, Na₂SO₃ allowed an effect 12-fold better than ethanol [31], allowing a 15 h production of H₂ by the Cu₂O/Na₂SO₃, without any signs of photocorrosion [31]. This approach is interesting since it highlights the possibility to engineer Cu₂O without addition of metal or cocatalysts.

2.4. Case−Study D: The Effect of Carbonaceous Materials

Reduced graphene oxide (rGO) can act as a protecting agent. An et al. [32] have studied the effect of rGO interfaced with different exposed facets of Cu₂O with (100) to be better for CO₂R than (100) [32]. The proposed mechanism is described in Figure 3 Case-Study D). They measured Cu-leaching using ICP (inductively coupled plasma optical emission spectrometry) and they demonstrated a Cu-atom leaching of ~3% after 3 h of light-irradiated Cu₂O/rGO [32]. EIS data showed a smaller resistance indicating better charge-transfer in Cu₂O/rGO vs. Cu₂O, as can be seen in Figure 5a. In a Mott–Schottky analysis they had a negative slope [32], which indicates p-type semiconductor properties in Cu₂O/rGO and an increased donor-density vs. Cu₂O.

A different carbonaceous substrate, such as g-C₃N₄/Vulcan Carbon, was prepared by Hussain et al. [42], protecting the (111) facets of Cu₂O, resulting in improved stability of Cu₂O nanoparticles. A different approach was given by Sun et al. [43], who developed a heterostructure of Cu₂O quantum dots (QDs) supported on a 3D g-C₃N₄ foam. In this work, the authors suggested that the (111)Cu₂O facet, when interfaced with g-C₃N₄, allowed creation of additional DOS at the Cu₂O/g-C₃N₄ interface [43]. In such heterojunction, the photoexcited electrons are suggested to be transferred from g-C₃N₄ to Cu₂O QDs, thus avoiding Cu₂O photocorrosion and enhance photocatalytic activity. Enhanced stability was also evidenced by the reusability of this Cu₂O/g-C₃N₄ heterojunction up to five times [43].

The protective potential of interfacial carbon was also evidenced by Wang et al., who synthesized Cu₂O/graphited carbon interfaces, using a polyol-method, to produce Cu₂O nano-flowers on graphene [45]. Increased porosity of the catalyst achieved by 90 °C treatment improved the internal charge transfer, as evidenced by EIS Nyquist data and as can be seen in Figure 5c. [45].

Overall, carbon-nanomatrices can act as protecting agents for Cu-oxides. The enhanced electron-storage capacity of graphitized carbon, together with their high chemical resilience are key-beneficial parameters that renders this approach as highly promising towards industrial exploitability.

2.5. Case−Study E: Core-Shell and Alloys

Besides the conventional approach of core-shell structures, in the case of Cu_-oxides, one can create core-shell structures by starting from a fully reduced Cu state and result to a reduced metallic-core, with a stable oxidized layer. Yang et al. [18] proposed a two-step method, firstly by electrodeposition of Cu₂O on FTO film and a subsequent thermal oxidation (time was the factor that controlled the thickness of the outer layer) in which the outer film was transformed to CuO. The so-formed Cu₂O/CuO bilayer was tested for its photostability, and Electrochemical Impedance Spectroscopy data showed very small resistance under illumination, while the Mott–Schottky measurements demonstrated large carrier density and a high flat-band potential, evidencing good conductivity and a high degree of band bending [18]. This approach places the Cu₂O/CuO composite as promising photocathode candidate for HER, according to its good stability in alkaline media, high photocurrent, and the ease of method.

A self-supported Cu/Cu₂O-CuO/rGO nanowire array [44] showed good stability for more than 15 h in alkaline medium, as can be seen in Figure 5b, in accordance with low resistance evidenced by EIS-loops, which indicates fast charge transfer. This provides evidence that the presence of metallic Cu⁰ as a hetero-component—not a standalone one—can offer high-speed paths for electron transportation. In addition, the highly reductive rGO shell also provides a conductive network that may promote charge transfer, which are beneficial for the Hydrogen Evolution Reaction process.

A highly efficient and low cost Cu₂O-SnO₂ core-shell electrocatalyst has been reported in [46] via control of the shell thickness. An optimized 5 nm thick SnO₂ shell gave the faradaic efficiency, i.e., >90% of CO₂ to CO reduction [46]. These Cu₂O-SnO₂ electrocatalyst achieved a good stability over 18 h of test at −0.6 V vs. RHE in 0.5 M KHCO₃ electrolyte.

A three-dimensional nanoporous Cu-Ru alloy as a Pt-free catalyst for HER, was prepared by Wu et al. [28] using a dealloying process, as shown schematically in Figure 6a,c. They demonstrated that the so-obtained nanoporous Cu₅₃Ru₄₇ exhibits long-term operation stability for more than 24 h in neutral pH (1.0 M PbS) and, after the first 3 h, fully stable at the alkaline pH (1 M KOH), as can be seen in Figure 6b.

Shen et al. started with a fundamental basis, which relies on electronic modification of a graphitic shell [33]. As reported from other groups too, with a graphitic layer to act as a protective shell for the core, this method can improve significantly the durability of the catalyst. Therefore, this group [33] synthesized a nickel-copper alloy (core) encapsuled in graphitic layer (shell) (NiCu@C) via chemical vapor deposition method, using CH₄ as a carbon source. This allowed them to obtain a tunable C-thicknesses. In this way, they had obtained three materials: NiCu@C-1 had mostly single-layered (78.4%) C-coating, NiCu@C-2 a well-defined, 2–3 nm thickness, core-shell structure, corresponding to 5–10 graphitic layers, and NiCu@C-3 with an 8–15 nm shell. Raman Spectroscopy, see Figure 5d, shows the three characteristic peaks, D, G, and 2D, that increased upon increase the graphitic shell (C) [32]. Shift in Raman peaks could be related to the modification of electronic structures of graphitic layers arisen from the interaction of carbon shell NiCu core. The stability of the catalyst was evaluated by chronoamperometric measurements, which verify that NiCu@C catalysts show slower decay in both acidic and alkaline pH than in pristine NiCu. From EIS, they demonstrated [32] also that the charge transfer resistance follows the same rule, as the thickness is increasing, the resistance increasing too, which can be explained by slower reaction kinetics and electron transport at the thicker graphitic shells (see Figure 3 Case-Study E- inset).

2.6. Case−Study F: Lattice-Size-Shape-Facets

Lattice-strain engineering can be effective on altering the DOS near the Fermi level, thus regulating adsorption energy of catalysts [27]. This was exemplified by Kang et al. [27], who used plasma spraying (PS) to prepare the coating layer in self-supported Cu electrode with intensive strain. The electrode’s stability was significant for more than 30 h at 100 mA cm⁻² (see Figure 3 Case-Study F). EIS data show that their PS-Cu had the lowest resistance of all tested samples, and the stability was best at acid medium and the worst at alkaline medium.

In addition, lattice-strain engineering may improve interaction between catalyst and intermediates [47]. For instance, tensile strain [48] could increase the interatomic distance inside the lattice catalyst, thus decreasing the overlap of d-orbitals, and/or upshift the d-band center and intermediates [49]. In this context, so far, several methodologies have been exploited to generate lattice-strain, such as, introducing interfaces with lattice mismatch [47], creating defects [48], and preparing ultra-thin nanosheets [50].

The different crystallite-facets of metal oxide nanoparticles may affect the activity and stability of photocatalysts for HER and CO₂RR, thus controlling preferential-facet growth can be a fruitful strategy. In a commonly applied methodology, the differential affinity of certain surfactants for certain Cu₂O-facets may allow control of surface oxidation extent, oxophilicity, hydrophilicity, and their activity to HER [51]. In a refined work, Gao et al. [52] studied Cu₂O nanoparticles (NPs) with (100) facets, octahedral Cu₂O NPs with exposed (111) and t-Cu₂O with both (111) and (100) facets. Studies on these Cu₂O NPs showed that the types of interfaced crystal facets exhibited different stabilities and different catalytic activities [53]. Shape of Cu₂O particles has been shown to be correlated with stability during the water splitting process [54]. More specifically, when comparing Cu₂O cubes with (100) facets, Cu₂O octahedra with (111) facets and Cu₂O rhombic dodecahedra with (110) facets, it was concluded that—practically—standalone Cu₂O, irrespective of the faced exposure, could not be very stable, no matter its facet engineering [53].

From the above-mentioned examples, it may be anticipated that, after finding which facet is best for the catalyst’s activity, we should protect that facet, by an appropriate method, e.g., p-n heterojunctions or coating or by mixing it with carbon materials as graphene.

3. Hydrogen Evolution by Cu-Oxide Based Materials

Both Cu₂O and CuO are p-type semiconductors [13,14,15,16] with band gaps 2.0–2.5 eV and 1.3–1.7 eV, respectively, see Figure 2, thus, absorbing a significant proportion of the solar-spectrum. However, their stability issues, when combine with fast e⁻-h⁺ recombination phenomena, make them rather unfavorable photocatalysts for H₂ evolution from H₂O [54]. Nonetheless, recent studies demonstrate that we can minimize these disadvantages by coupling copper oxides with other, more stable and suitable semiconductors. Doing so, the electrical and optical properties of Cu₂O and CuO are significantly changed by the material they are coupled with [24,28,33].

Hydrogen Evolution Reaction (HER) [55] is considered an extraordinary reaction in electrochemistry for two reasons: firstly, for its simplicity, because it is always a 2e⁻ process, no matter the pH [55], and secondly because it is a direct process to produce high-purity H₂, considered to be the key fuel for the future. Mechanistically, the basis of HER is described by the reactions in

Table 1. Overall reaction pathways for HER in acidic/alkaline mediums.

Overall Reaction (Conditions)	Reaction Pathway	Equations
Acidic Media $2\left\{{\mathrm{H}}^{+}\right\}+2\left\{{\mathrm{e}}^{-}\right\}\to {\mathrm{H}}_2$	$\left\{{\mathrm{H}}^{+}\right\}+\left\{{\mathrm{e}}^{-}\right\}+ * \to {\mathrm{H}}^{*} \left(\mathrm{Volmer}\right)$	1.i
	$\left\{{\mathrm{H}}^{+}\right\}+\left\{{\mathrm{e}}^{-}\right\}+{ \mathrm{H}}^{*}\to {\mathrm{H}}_2 \left(\mathrm{Heyrovsky}\right)$	1.ii
	${\mathrm{or} \mathrm{H}}^{*}+{ \mathrm{H}}^{*}\to {\mathrm{H}}_2 \left(\mathrm{Tafel}\right)$	1.iii
Alkaline Media $2{\mathrm{H}}_2\mathrm{O}+2\left\{{\mathrm{e}}^{-}\right\} \to {\mathrm{H}}_2+2{\mathrm{OH}}^{-}$	${\mathrm{H}}_2+\left\{{\mathrm{e}}^{-}\right\}\to {\mathrm{H}}^{*}+{\mathrm{OH}}^{-} \left(\mathrm{Volmer}\right)$	2.i
	${\mathrm{H}}_2\mathrm{O}+\left\{{\mathrm{e}}^{-}\right\}+{ \mathrm{H}}^{*}\to {\mathrm{H}}_2+{\mathrm{OH}}^{-} \left(\mathrm{Heyrovsky}\right)$	2.ii
	${\mathrm{or} \mathrm{H}}^{*}+{ \mathrm{H}}^{*}\to {\mathrm{H}}_2 \left(\mathrm{Tafel}\right)$	2.iii

[56]. HER in electrochemistry is a 2e⁻ process [55], which is divided mainly in two steps: first is the Volmer step (1.i in Table 1) that includes the adsorption of one H-atom on the catalyst surface, by transferring one proton H⁺ from the acid electrolyte which combines with an electron e⁻ transferred from the catalyst surface to form an adsorbed hydrogen atom (H* = H⁺ + e⁻); second is the Heyrovsky step (Equation (1.ii)) in which this H* combines with one electron e⁻ and one H⁺ to form a H₂ molecule, see also Figure 7. In addition, there is possibility to form one H₂ molecule through the Tafel step (1.iii) i.e., the combination of two H* [55].

In acidic electrolytes, it has been generally accepted that the differences in reaction rates and activity are closely related to the differences of free-energy of hydrogen-adsorption on different materials [26]. In alkaline media, due to the lack of H⁺, the reaction must begin from dissociation of H₂O molecules, see Volmer–Heyrovsky steps (Equations (2.i) and (2.ii)). In addition, in alkaline media, as in the case of acidic media, there is the possibility to form a H₂ molecule through the Tafel step (2.iii), i.e., the combination of two H*. This additional step of proton-generation from H₂O in alkaline media introduces an additional energy barrier and affects the whole reaction kinetics and that is the main reason why acid medium is preferred in this reaction. In alkaline environments, there are two approaches to start with: [i] by the single descriptor of hydrogen binding energy [57], or [ii] the bifunctional nature of HER, a good catalyst needs both a beneficial OH_ad energetics (as it is slower when it has to do with metal; thus, is the rate-limiting step) and beneficial H_ad energetics [51].

A summary of pertinent Cu-based semiconducting materials is listed in Table 2 (photocatalytic H₂ production), Table 3 (electrocatalytic H₂ production) and Table 4 (photoelectrocatalytic H₂ production). Hereafter, we highlight some engineering aspects related to the control/improvement of hydrogen evolution by such catalysts.

Table 2. Summary of pertinent Cu-based heterojunctions and photocatalytic H₂ evolution rates.

Photocatalytic Hydrogen Production
Catalyst	Hole Scavenger	Irradiation Source	H2 Evolution Rate	Ref.
Cubic-Cu2O	Pure H2O (Water Splitting)	300 W Xe-Lamp	Not detected	[54]
Octahedra-Cu2O			>0.4 μmol g−1 h−1
Rhombic Dodecahedra-Cu2O			~1.6 μmol g−1 h−1
Cu2O/TiO2 (C-1.5/T-2)	30%MeOH (H2O/MeOH)	Full-arc Xe Lamp 100 mW cm−2	11 mmol g−1 h−1	[58]
	30%MeOH (Seawater/MeOH)		5.1 mmol g−1 h−1
CuO (Later Cu2O)-TiO2	10% Glycerol (H2O/Glycerol)	300 W Xe-Lamp	336 μmol g−1 h−1	[59]
Cu0-TiO2			~867 μmol g−1 h−1
Cu2O/TiO2	10% MeOH (H2O/MeOH)	300 W Xe-Lamp	70 μmol g−1 h−1	[60]
3 wt.% Cu-TiO2	~8% MeOH (H2O/MeOH 11:1)	125 W Hg-Lamp (325 & 365nm)	2.07 mmol g−1 h−1	[61]
13.5 wt.% Cu-TiO2			2.48 mmol g−1 h−1
Electrodeposited Cu2O-WO3	Pure H2O (Water Splitting)	400 W Hg-Lamp	~7 μmol g−1	[29]
Cu@Cu2O/ZnO	0.75M Na2S and 1.05M Na2SO3	300 W Xe-Lamp	1.47 mmol g−1 h−1	[62]
ZnO/Cu2O-CuO	60mM Na2S	150 W Xe-Lamp	1092.5 μmol g−1 h−1	[63]
CuSA-TiO2 (~1.5 wt%)	~70% MeOH (H2O/MeOH 1:2)	500 W Xe-Lamp	101.74 mmol g−1 h−1	[40]
0.75% Cu atom-TiO2	25% MeOH (H2O/MeOH 3:1)	100 W Xe-Lamp	16.6 mmol g−1 h−1	[64]
NDs-Cu2O	20% EtOH (H2O/EtOH)	300 W Xe-Lamp (100 mW cm−2)	1597 μmol g−1 h−1	[65]
		Visible Light 420–760 nm 77.5 mW cm−2	824 μmol g−1 h−1
Cu2O@g-C3N4 (CN5 5wt.%)	Pure H2O	300 W Xe-Lamp (≥420 nm)	795 μmol g−1 h−1	[66]

Table 2. Summary of pertinent Cu-based heterojunctions and photocatalytic H₂ evolution rates.

Photocatalytic Hydrogen Production
Catalyst	Hole Scavenger	Irradiation Source	H2 Evolution Rate	Ref.
Cubic-Cu2O	Pure H2O (Water Splitting)	300 W Xe-Lamp	Not detected	[54]
Octahedra-Cu2O			>0.4 μmol g−1 h−1
Rhombic Dodecahedra-Cu2O			~1.6 μmol g−1 h−1
Cu2O/TiO2 (C-1.5/T-2)	30%MeOH (H2O/MeOH)	Full-arc Xe Lamp 100 mW cm−2	11 mmol g−1 h−1	[58]
	30%MeOH (Seawater/MeOH)		5.1 mmol g−1 h−1
CuO (Later Cu2O)-TiO2	10% Glycerol (H2O/Glycerol)	300 W Xe-Lamp	336 μmol g−1 h−1	[59]
Cu0-TiO2			~867 μmol g−1 h−1
Cu2O/TiO2	10% MeOH (H2O/MeOH)	300 W Xe-Lamp	70 μmol g−1 h−1	[60]
3 wt.% Cu-TiO2	~8% MeOH (H2O/MeOH 11:1)	125 W Hg-Lamp (325 & 365nm)	2.07 mmol g−1 h−1	[61]
13.5 wt.% Cu-TiO2			2.48 mmol g−1 h−1
Electrodeposited Cu2O-WO3	Pure H2O (Water Splitting)	400 W Hg-Lamp	~7 μmol g−1	[29]
Cu@Cu2O/ZnO	0.75M Na2S and 1.05M Na2SO3	300 W Xe-Lamp	1.47 mmol g−1 h−1	[62]
ZnO/Cu2O-CuO	60mM Na2S	150 W Xe-Lamp	1092.5 μmol g−1 h−1	[63]
CuSA-TiO2 (~1.5 wt%)	~70% MeOH (H2O/MeOH 1:2)	500 W Xe-Lamp	101.74 mmol g−1 h−1	[40]
0.75% Cu atom-TiO2	25% MeOH (H2O/MeOH 3:1)	100 W Xe-Lamp	16.6 mmol g−1 h−1	[64]
NDs-Cu2O	20% EtOH (H2O/EtOH)	300 W Xe-Lamp (100 mW cm−2)	1597 μmol g−1 h−1	[65]
		Visible Light 420–760 nm 77.5 mW cm−2	824 μmol g−1 h−1
Cu2O@g-C3N4 (CN5 5wt.%)	Pure H2O	300 W Xe-Lamp (≥420 nm)	795 μmol g−1 h−1	[66]

3.1. Photocatalytic Hydrogen Production

3.1.1. Facet Engineering

The shape of Cu₂O nanocrystals can affect the photocatalytic H₂ production performance [54]. Cu₂O nanopowders with different morphologies, i.e., cubes with (100) facets, octahedra with (111) facets, or rhombic dodecahedra with (110) facets, which were evaluated for long-term stability on hydrogen production as well as for overall water-splitting. All materials presented photocorrosion phenomena under irradiation in the form of phase transformation of Cu₂O to CuO, verified both by XPS and X-Ray Diffraction (XRD) measurements. Noteworthy, the stability trend was completely different under dark conditions. Moreover, the absence of sacrificial agent during water splitting led to an absence of O₂ production, as normally H₂ and O₂ should have been produced in a 2:1 ratio. Thus, in [54], the photogenerated holes of Cu₂O that were supposed to oxidize H₂O and produce O₂ were consumed in the oxidation of Cu₂O to CuO. Furthermore, the best performing material, rhombic dodecahedra Cu₂O, was tested for long-term stability where at 48 h the produced H₂ was 8-fold less and at 72 h no H₂ was detected. To prevent the Cu₂O oxidation, the nanocrystals were coated with a TiIrO_x layer [54]. This hybrid TiIrO_x/Cu₂O material had a better photocatalytic performance with a H₂:O₂ ratio of 2:1_, indicating that TiIrO_x facilitated the hole transfer, thus preventing photocorrosion of Cu₂O, at the same time enabling O₂-production from water oxidation.

3.1.2. Coupling with Semiconductors

• The case of Cu-TiO₂

Another strategy is the coupling of CuO or Cu₂O with other semiconductors and the establishment heterojunctions or Z-scheme mechanism [37]. Coupling copper oxides with other semiconductors, such as TiO₂ or ZnO with lower energy band positions can lead to the formation of Z-Scheme or Type-II electron transfer mechanism between the two materials. Therefore, Cu₂O/TiO₂ hybrid systems have been widely investigated in the fields of photocatalysis for both H₂ evolution [60] and CO₂ reduction [67] and degradation of organics [68]. Lv et al. [58] coupled Cu₂O nanoparticles onto the surface of TiO₂ and showed that, due to the Z-scheme mechanism, carriers with strong oxidation and reduction ability were confined to TiO₂, inhibiting the photocorrosion of the Cu₂O. XPS and EPR data, see Figure 8a–c, indicated presence of surface lattice defects, oxygen vacancies (O_V), which can act both as electron traps and simultaneously enhance the adsorption capability of H⁺ ions in the hydrogen evolution reaction [58].

Moreover, they notice an upward shift and a downward shift on the XPS spectra of Ti_2p and Cu_2p, respectively [58], indicating a directional carrier transport from TiO₂ to Cu₂O (Figure 8a,b). Regarding the photocatalytic activity, their Cu₂O/TiO₂ composite exhibited high H₂ production rates, 11 mmol g⁻¹ h⁻¹ in water and 5.1 mmol g⁻¹ h⁻¹ in seawater [57]. Compared with the pristine Cu₂O, yielding 0.5 mmol g⁻¹ h⁻¹ and 0.3 mmol g⁻¹ h⁻¹ in water and seawater, respectively, the Cu₂O/TiO₂ composite showed a clear improvement. This significant improvement of photocatalytic activity and stability is attributed on the consumption of the highly oxidizing holes of Cu₂O from the O_V’s of TiO₂ and the Z-scheme mechanism [58] (Figure 8d). Generally, stabilizing CuO or Cu₂O nanoparticles on TiO₂ is a common approach, Tian et al. have prepared a TiO₂ photocatalyst loaded with Cu NPs through an ion exchange (IE) process, where they report the co-existence of Cu²⁺, Cu¹⁺, and Cu⁰, based on XPS-data. The final composite material gave a purple hue which persisted during the photocatalytic experiments, indicating that copper was at the metallic state [60]. In another study, Jung et al. also used TiO₂ as a substrate to deposit small, finely-dispersed CuO NPs (0.5–2.5 nm) [59] (Figure 8e). During H₂ generation, they made two observations, [i] there was a color change from dark-grey to deep-violet, attributed to change of the Cu oxidation state, and [ii] there was an initial delay of the hydrogen production. XPS and AES profile at different illumination times show the slow transformation of Cu oxidation states from Cu²⁺ to Cu¹⁺. The authors [59] suggest that the photogenerated electrons are consumed to reduce CuO to Cu₂O instead of reducing H⁺ and produce H₂ (Figure 8f,g). When the reduction of copper has been accomplished, these photogenerated electrons participate in H₂ production. In the case of Cu⁰, as in the cases of noble metal cocatalysts, there is a continuum of states within the TiO₂ energy gap where Cu⁰ attracts the photogenerated electron from TiO₂, thus facilitating electron transfer to protons (see Figure 8h).

• Coupling of Cu with Non-TiO₂ Semiconductors

Besides TiO₂, other semiconductors, especially n-type, have been also used for the stabilization of Cu NPs. In this context, Hu et al. coupled p-type Cu₂O with n-type WO₃, forming a p-n heterojunction [29]. Unlike pristine Cu₂O, which did not produce H₂ during water splitting, Cu₂O/WO₃ heterojunction facilitated H₂ production, due to efficient charge separation. Lou et al. designed a Cu⁰@Cu₂O/ZnO heterostructure in order to stabilize the plasmonic Cu core, by decorating it with Cu₂O shell, and finally stabilizing this core-shell structure onto ZnO nanorods [62] (See Figure 9a). This hybrid nanostructure exhibited better stability even after four catalytic cycles (Figure 9b), and superior photocatalytic performance, i.e., 1472μmol g¹ h⁻¹ compared with the pristine Cu@Cu₂O material which yields 5 μmol g⁻¹ h^—1 (See Figure 9c). This improvement on the photocatalytic activity was attributed on the Localized Surface Plasmon Resonance (LSPR) caused by photoexcitation of metallic Cu⁰ particles [69]. Arguably, the LSPR was suggested to be linked with hot-electron transfer, as evidenced by Photoluminescence (PL) data. Specifically, the PL emission spectra indicated that Cu@Cu₂O/ZnO nanocomposites could delay the recombination rate of e⁻-h⁺ pairs. In the proposed model, see Figure 9d, surface plasmons are photoexcited upon irradiation and the hot electrons occupy the surface plasmon states above the Fermi energy. A fraction of these hot electrons are injected onto the CB of Cu₂O, by overcoming the Schottky barrier formed at the Cu₂O-Cu⁰ interface, and finally they leap on the CB of ZnO. This mechanistic path for the electrons is believed to intercept photocorrosion phenomena [62].

3.1.3. Carbon-Based Materials and Core-Shell Structures

Use of carbonaceous materials is another interesting approach towards the stabilization and the improvement of the photocatalytic performance of copper nanocatalysts. Lin et al. has decorated 50 nm Cu₂O with small ~3 nm nanodiamonds (NDs), and the optimal loading was found to be 3 wt.% enabling a H₂ production rate of 1597 μmol g⁻¹ h⁻¹ with a solar-to-hydrogen efficiency of 0.85%. This nanocomposite exhibited high photocatalytic stability attributed to the narrower band gap, compared with pristine Cu₂O, which improves the ability to harvest solar light [65] (See Figure 9e). As for the proposed photocatalytic mechanism, the photoexcited electrons from the NDs surface were injected into the Cu₂O and at the same time the photogenerated holes of Cu₂O migrated to the NDs (See Figure 9f). The excess of electrons on Cu₂O performed the reduction of H₂O to H₂ and the excess of holes on the NDs oxidized the ethanol. Another material of great interest is graphitic carbon nitride (g-C₃N₄). Liu et al. synthesized a Cu₂O@g-C₃N₄ core-shell structure onto the Cu₂O octahedra with exposed (111) facets. The efficiency of the composite material under visible light irradiation was attributed to the synergistic effect at the interface of Cu₂O and g-C₃N₄, i.e., transfer of the photogenerated electrons of Cu₂O to the g-C₃N₄ shell. In contrast, UV irradiation on the Cu₂O crystals had a negative effect, and this was attributed to a mechanism where O_V in Cu₂O act as electron traps, leading to the reduction of Cu₂O to Cu⁰ exhibiting lower photocatalytic activity [66].

3.1.4. The Case of Cu-Single Atoms

Lastly, single-atom catalysts (SACs) are gaining more and more interest owing to their ability in maximizing reaction active sites [70]. These isolated metal atoms anchored in the surface of photocatalysts may offer more reaction active sites, but their stability and tendency to aggregate during photocatalysis is still an issue to address. Very recently, Lee et al. [64] designed atomically-dispersed Cu on TiO₂ by incorporating site-specific single atoms of Cu. Interestingly, it was shown that photogenerated electrons can be transferred from the CB of TiO₂ to the d-orbitals of the isolated Cu²⁺ atoms. These trapped electrons induced a polarization field resulting in lattice distortion of TiO₂, that was suggested to be linked with the enhanced photocatalytic H₂ activity. This phenomenon was reversible when the material was exposed to O₂ without irradiation. They also report a change of color during photocatalysis, in accordance with other studies on TiO₂-based nanocatalysts [58,59]. In another very recent study, Zhang et al. managed to deposit higher amounts of CuSA (>1 wt%) on TiO₂, at this point we should note that normally the loading percentages of SACs is near 0.1–0.3 wt% [40]. Their optimized material exhibited the impressive H₂ evolution rate of 101.7 mmol g⁻¹ h⁻¹, which in their case was higher than PtSA-TiO₂ (95.3 mmol g⁻¹ h⁻¹). This impressive photocatalytic performance is attributed to the larger amount of Cu SACs as well as to the high dispersion onto the TiO₂ surface. Regarding the photocatalytic mechanism, efficient electron transfer is attributed to the reversible redox process between Cu²⁺ and Cu¹⁺ atoms. These findings are supported by Electron Paramagnetic Resonance (EPR) data, showing a strong Cu²⁺ EPR signal before irradiation, which during irradiation was converted to Cu⁺, which was re-oxidized to Cu²⁺ after exposure to air O₂ [40]. This in situ self-healing process enables CuSA-TiO₂ to achieve these remarkable H₂ production rates. Furthermore, long-term H₂ evolution experiments, i.e., 380 days later, showed no loss of the photocatalytic performance [40].

3.2. Electrocatalytic Hydrogen Production

Electrochemistry permits quantitative comparisons of activity stability, while photo-electrochemistry permits more advanced parameters to be monitored, i.e., photocurrent density, photo-response, photostability and photocorrosion. In brief, the tools used can be [56]:

[i]

exploring the intrinsic activity of catalysts by measuring the capacitance of double layer Cdl to detect the electrochemical active surface area (ESCA) [71],

[ii]

overpotential needs to be the lowest possible, at 10 mA/cm2 [72],

[iii]

the Tafel slope shows the rate of adsorption–desorption kinetics, where the electrochemical desorption step is rate-determining according to the Volmer–Heyrovsky mechanism [73].

[iv]

through EIS measurements the charge transfer resistance can easily be found [56].

[v]

Mott–Schottky plots can provide information about flat band potential and the populations of donors and acceptors [74].

The double-layer capacitance which can be determined by cyclic voltammetry measurements, the electrochemically active surface area, current density [56], and the faradaic efficiency of the sample [75]. Linear Sweep Voltammetry (LSV) provides a measure of overpotentials necessary to achieve j = 10 mA/cm², and after 2 h to check the stability, or even a long-term stability test which can last up to 24 h [76].

In

Table 3. Summary of Electrocatalytic Parameters, related to Hydrogen Evolution Reaction, for Cu-based catalysts.

Electrocatalytic Hydrogen Production
Catalyst	Electrolyte or pH	Cdl (mF/cm2)	η(mV) @ −10 mA/cm2	Tafel Slope (mV dec−1)	Ref.
Cu/Cu2O-CuO/rGO-400	1 M KOH	130	105	124	[44]
PS-Cu	0.5 M H2SO4	8.77	182	99.16	[27]
	1 M PBS	10.01	261	143.58
	1 M KOH	13.56	121	136.54
Cu-foam	0.5 M H2SO4	4.86	411	192.22
	1 M PBS	5.60	429	211.95
	1 M KOH	5.91	473	230.27
Cu-NPs	0.5 M H2SO4	6.87	530	132.17
	1 M PBS	6.47	707	285.99
	1 M KOH	9.72	454	148.86
Pt-foil	0.5 M H2SO4		78	79.09
	1 M PBS		186	194.05
	1 M KOH		178	167.84
NiCu@C-1	pH = 0		48	94.5	[33]
	pH = 7		164	94.6
	pH = 14		74	74
Graphene coated Cu	0.5 M H2SO4		252	67	[77]
Cu@NC NT/CF-500	1 M KOH	101	123	61	[78]
Cu53Ru47	1 M KOH	59	15	30	[28]
	1 M PbS	59	41	35
Cu-m/Cu-W/NiCo-LDH	0.5 M H2SO4		15	79.4	[79]
	1 M KOH	19.8	27	50.5

, we summarize pertinent values for these electrochemical parameters for selected Cu-oxide catalysts.

In Table 3, pertinent literature works are presented. In all cases, we mention works where they did not use noble metals as co-catalysts, but use only Cu-based electrocatalysts. Some of them had merged the electrolytes with graphene as a core-shell structure in order to protect them from corrosion [77], while others [44,78] used carbon nanotubes with Cu or Cu/Cu₂O/CuO structures merged in reduced graphene oxide. An important contribution is a strain-activated copper catalyst [27], operating under a wide pH range. In [27] the current density achieved was1.987 A cm⁻² at −1.2 V vs. RHE, which is 2.85 times vs. a reference Pt foil.

Moreover, a three-dimensional nanoporous Cu-Ru alloy as an outstanding Pt-free catalyst for HER has been prepared by a dealloying process by Wu et al. [28]. This catalyst had a good activity with a very low overpotential and very low Tafel slopes in alkaline and neutral environments. By alloying those two metals, Cu and Ru, the incorporation of Ru atoms in the Cu matrix enhanced the strength of the Cu-H interaction, while it weakened the Ru-H interaction, but improved the overall activity [27]. The Electrochemical impedance spectroscopy measurements of np-Cu₅₃Ru₄₇ further confirmed that the incorporation of Ru atoms into the Cu matrix brings about small internal resistance and fast charge-transfer behavior. The same group, [27], demonstrated with Density Functional Theory (DFT) that the incorporation of Ru atoms into the Cu matrix could effectively optimize the d-electron domination of Cu and Ru atoms and decrease the energy barrier and improving the HER activity.

Another approach [79] presented a complex heterostructure of NiCo layered double hydroxide wrapped around Cu nanowires grown on top of commercially available Cu mesh (Cu-m). This Cu-m/Cu-W/NiCo-LDH material outperformed the benchmark 40 wt% Pt/C at a wide pH range [79]. As shown in Figure 10a,b, the required overpotential was smaller for alkaline medium than 40%Pt/C, confirmed by the Tafel plots at Figure 10c. While in acidic environment, as shown in Figure 10d,e, the Cu-m/Cu-W/NiCo-LDH had increased overpotential versus 40%Pt/C at −10 mA cm⁻², while the Tafel slope in Figure 10f shows a decreased value for Cu-m/Cu-W/NiCo-LDH. Figure 10h shows the required potential across the pH range to reach −10 mA cm⁻² for Cu-m/Cu-W/NiCo-LDH versus 40 wt% Pt/C. In Figure 10g, high-resolution transmission electron microscopy (HRTEM) shows the (111) exposed facets of Cu-m and Cu-Ws, and the latter’s growth direction is along (111).

3.3. Photoelectrochemical Hydrogen Evolution Reaction

Photoelectrochemical cells (PEC) are widely considered potent for solar water splitting devices [80] because they combine solar energy collection with water electrolysis and because they spatially separate two half-cell reactions, HER and OER. For a PEC cell, the process depends on three steps [55]: (i) light absorption by a semiconductor material, (ii) electron-hole pair generation, and, as a result, (iii) photoinduced electrons (or holes) are driven, by space-charge field, to the semiconductor/solution interface where they reduce or oxidize water.

Usually, a PEC cell has a n-type semiconductors, such as TiO₂, ZnO, BiVO₄, WO₃, Fe₂O₃ [4,81,82,83] whose valence band edges are more positive than the potential of the H₂O/O₂ redox couple, that makes them ideal for OER photoanodes. On the other hand, p-type semiconductors such as p-CdS, p-WSe, p-InP and p-type Copper Oxides [25,84,85], whose conduction band edge is more negative than the H₂O/H₂ potential, are ideal for HER photocathodes (see Figure 11). In Table 4, we selected some interesting works from several groups, who have presented durable and stable photocathodes for a PEC cell. In addition, very important for these catalysts is to have good light response and photocurrent to benchmark their photocatalytic activity.

In order to explore the photocorrosion of CuO, Xing et al. [26] had tested heterojunctions CuO: CuO/TiO₂, CuO/Pt, and CuO/TiO₂/Pt as photocathodes. Under increased protection of the photocathodes, allowed increase of the photocurrent density at −0.55 V, as shown in Figure 10j. The composite of CuO/TiO₂/Pt had the highest photocurrent density. In addition, according to the Figure 10k, CuO/TiO₂/Pt seems to have a very good photo-response, at −0.55 V vs. Ag/AgCl applied under 30 s light on/off cycles for more than 300 s. This electrode was stable against photocorrosion in 1 M NaOH electrolyte at −0.55 V vs. Ag/AgCl under simulated sunlight illumination for 2 h, Figure 10l. The EIS Nyquist plot, shown in Figure 10m, reveals that the CuO/TiO₂/Pt has decreased resistance in charge transfer phenomena, which indicates easier transfer of photoinduced electrons to the electrolyte solution.

Table 4. Pertinent Cu-based materials for PhotoElectrochemical Hydrogen Evolution Reaction based on their catalytic activity.

Photoelectrochemical Hydrogen Production
Catalyst	Environment	Light Source	J @ Applied Potential	Ref.
CuO/TiO2	1 M KOH	500 W Xe-Lamp	−0.54 mA/cm2 @ −0.55 V vs. Ag/AgCl	[26]
CuO/Pt	1 M KOH		−0.57 mA/cm2 @ −0.55 V vs. Ag/AgCl
CuO/TiO2/Pt	1 M KOH		−0.75 mA/cm2 @ −0.55 V vs. Ag/AgCl
Cu2O/CuO Bilayered composite	0.5 M Na2SO4 + 1 M KOH	300 W Xe-Lamp (1000 mW m−2)	−3.15 mA/cm2 @ 0.40V vs. RHE	[18]
Cu2O/Ga2O3/TiO2/Rux	0.5 M Na2SO4 + 0.1 M phosphate solution	300 W Xe-Lamp (1000 mW m−2)	−10 mA/cm2 @ 0V vs. RHE	[87]
3D CuO		150 W solar simulator (1000 mW m−2)	−3.15 mA/cm2 @ 0.42V vs. RHE	[88]
CuO/thin film	0.1 M Na2SO4	Sunlight (1000 mW m−2)	−3.1 mA/cm2 @ 0V vs. RHE	[23]
Au-Pd decorated CuO thin film	0.1 M Na2SO4	Sunlight (1000 mW m−2)	−3.88 mA/cm2 @ 0V vs. RHE
CuSA-TiO2	0.2 M Na2SO4	150 W Xe-Lamp	−10 mA/cm2 @ −0.72V vs. NHE	[40]

Table 4. Pertinent Cu-based materials for PhotoElectrochemical Hydrogen Evolution Reaction based on their catalytic activity.

Photoelectrochemical Hydrogen Production
Catalyst	Environment	Light Source	J @ Applied Potential	Ref.
CuO/TiO2	1 M KOH	500 W Xe-Lamp	−0.54 mA/cm2 @ −0.55 V vs. Ag/AgCl	[26]
CuO/Pt	1 M KOH		−0.57 mA/cm2 @ −0.55 V vs. Ag/AgCl
CuO/TiO2/Pt	1 M KOH		−0.75 mA/cm2 @ −0.55 V vs. Ag/AgCl
Cu2O/CuO Bilayered composite	0.5 M Na2SO4 + 1 M KOH	300 W Xe-Lamp (1000 mW m−2)	−3.15 mA/cm2 @ 0.40V vs. RHE	[18]
Cu2O/Ga2O3/TiO2/Rux	0.5 M Na2SO4 + 0.1 M phosphate solution	300 W Xe-Lamp (1000 mW m−2)	−10 mA/cm2 @ 0V vs. RHE	[87]
3D CuO		150 W solar simulator (1000 mW m−2)	−3.15 mA/cm2 @ 0.42V vs. RHE	[88]
CuO/thin film	0.1 M Na2SO4	Sunlight (1000 mW m−2)	−3.1 mA/cm2 @ 0V vs. RHE	[23]
Au-Pd decorated CuO thin film	0.1 M Na2SO4	Sunlight (1000 mW m−2)	−3.88 mA/cm2 @ 0V vs. RHE
CuSA-TiO2	0.2 M Na2SO4	150 W Xe-Lamp	−10 mA/cm2 @ −0.72V vs. NHE	[40]

As general conclusion, photocorrosion must be scrutinized as another important phenomenon in PEC for solar hydrogen production. This aspect is considered using molecular components such as linkers, non-oxide material modifiers, or catalysts. The most common method, as mentioned in Table 4, is by using TiO₂ as a matrix where CuO, Cu₂O can be protected by photocorrosion through the Z-scheme mechanism. In the case of single-Cu atoms, there is a different approach, where Cu(II) transforms reversibly to Cu(I) during photoelectrocatalysis as described before.

4. CO₂ Reduction by Cu-Based Materials

From an economic and environmental point of view, (electro)photocatalytic CO₂ reduction is a forward-looking realm in catalytic technology. Inspired by natural-photosynthesis, scientists are trying to realize “Artificial-Photosynthesis”, convert CO₂ into useful and high value chemical fuels, using the highly abundant solar energy (see Figure 1). Inoue together with Fujishima and Honda pioneered the idea of the feasibility of catalytic conversion of CO₂ to C1-fuels using semiconductors [4,89]. Since then, many efforts were made to improve the overall performance and tune the selectivity towards a specific product [9]. In this context, Cu-oxides and metal-Cu particles emerge among the most promising photocatalysts thanks to good electron transfer properties and loosely bound d-electrons, thus having great potential for facilitating CO₂ photoreduction [20]. Herein, we focus on Cu-based semiconductors coupled with other oxide semiconductors, carbon-based materials or plain copper catalysts with facet modifications and discuss stability issues and product selectivity (see

Table 5. Pertinent Cu-based semiconductors used for the photocatalytic CO₂ conversion into various chemical fuels.

Photocatalytic CO2 Reduction
Catalyst	Hole Scavenger/ Reaction Conditions	Irradiation Source	Main Products	Ref.
(110) Cu2O	Saturated H2O in CO2	300 W Xe-Lamp	CH3OH: 1.2 mol g−1 h−1	[90]
Dodeca-Cu2O/rGO	Saturated H2O in CO2	300 W Xe-Lamp (λ ≤ 420 nm)	CH3OH: 355.3 μmol g−1 h−1	[91]
Pt-Cu2O/TiO2	Saturated H2O in CO2 (71kPa)	500 W Xe-arc Lamp (300 nm < λ < 400 nm) (20.5 mW cm−2)	CH4: 1.42 μmol g−1 h−1 CO: 0.05 μmol g−1 h−1	[92]
Octa-Cu2O/TiO2	Water vapor—1atm CO2 (g)	1 kW Hg (Xe) arc lamp (λ ≥ 305 nm)	CO: 2.11 μmol g−1 h−1	[67]
1% Cu/TiO2 (H2)	Mixed gas CO2/H2O	150 W solar simulator (90 mW cm−2, 200 ≤ λ ≤ 1000 nm)	CH4: 25 μmol g−1 h−1 CO: 4.4 μmol g−1 h−1	[93]
2%CuO-19%ZnO/TiO2	Saturated H2O in CO2, 0.2 M NaOH	18 W Hg-Lamp (λ = 25 4nm)	CH4: 184 μmol g−1 (after 24 h)	[94]
ZnO-Cu2O	Saturated H2O in CO2, 0.2 M Na2CO3	300 W Xe-Lamp	CH4: 1080 μmol g1 h−1	[95]
5wt% CuO/NaTaO3	CO2, Isopropanol	250 W Hg-Lamp (365 nm)	CH3OH: 1302.22 μmol g−1 h−1	[96]
Cu2O	Saturated H2O in CO2, NaOH, Na2SO3	500 W Xe-arc Lamp (400 nm < λ < 700nm)	CH3OH: 104 μmol g−1	[97]
Cu2O-SiC			CH3OH: 191 μmol g−1
BiVO4/C-coated Cu2O	Saturated H2O in CO2	300 W Xe-Lamp (100 mW cm−2) (λ > 420 nm)	CO: ~3 μmol g−1 h−1	[98]
3%Cu/TiO2	Saturated H2O in CO2, 1 M NaOH	UV-Lamp (254 nm, 5.4 mW cm−2) Visible: 500 W Xe-arc Lamp	HCOOH: > 4500 μmol g−1 h−1 (Visible) // CH3OH: ~300 μmol g−1 h−1 (Visible) HCOOH: >2000 μmol g−1 h−1 (UV)// CH3OH: >75 μmol g−1 h−1 (UV)	[99]
3% Cu/g-C3N4			HCOOH: >3500 μmol g−1 h−1 (Visible)//CH3OH: ~350 μmol g−1 h−1 (Visible) HCOOH: >3750 μmol g−1 h−1 (UV)// CH3OH: >200 μmol g−1 h−1 (UV)
(g-C3N4)/(3%Cu/TiO2) (30:70)			HCOOH: 5069 μmol g−1 (Visible)// CH3OH: 2574 μmol g−1 (Visible) HCOOH: 6709 μmol g−1 (UV)// CH3OH: 614 μmol g−1 (UV)
rGO-CuO	DMF, Saturated H2O in CO2	20 W LED (85 W m−2)	CH3OH: 1228 μmol g−1	[100]
rGO-Cu2O			CH3OH: 862 μmol g−1
c-Cu2O_gCN	~1 bar CO2 with moisture	8 W LED Lamp	CO: 0.002 μmol g1 h−1	[101]
g-C3N4/CuO@MIL-125(Ti)	H2O 0.3%CO2 (1 MPa)	300 W Xe-Lamp (326.1 W m−2) (λ = 420 nm)	CO: 180.1 μmol g−1 CH3OH: 997.2 μmol g−1 C2H5OH: 531.5 μmol g−1 CH3CHO: 1505.7 μmol g−1	[102]

).

4.1. Photocatalytic CO₂ Reduction

4.1.1. Facet Dependency of CO₂ Reduction

As in the case of photocatalytic H₂ evolution, see Section 3.1.1, facet engineering of Cu oxides shows high promise towards increased performance in photocatalytic CO₂ reduction [10]. Recently, Wu et al. reported the photocatalytic reduction of CO₂ on facet specific active sites of Cu₂O [90]. They demonstrated that the (110) facet of a single Cu₂O particle was the photoactive site towards CO₂ reduction to CH₃OH, while the (100) facet remained inert [90]. This observation can be attributed to electron-density difference on Cu active sites on the (110) facet and a shift from Cu(I) to Cu(II) due to CO₂ and H₂O adsorption [90]. In this way, during CO₂ reduction, the Cu₂O catalyst manages to oxidize H₂O accompanied with a lattice expansion due to the CO₂ adsorption. This process showed a selectivity towards methanol yielding 1.2 mol CH₃OH g⁻¹ h⁻¹ and reaching an internal quantum yield of ~72% [90].

4.1.2. Coupling with Semiconductors

• The Case of Cu-TiO₂

Cu₂O coupling with TiO₂ has been explored as an efficient strategy to enhance photocatalytic performance. More specifically, Aguirre et al. synthesized a p-n Cu₂O/TiO₂ heterojunction, where the Z-scheme mechanism of electron transfer [37] enhances the stability of Cu₂O [67] (see Figure 12a,b). The energetics of this example are educative. TiO₂ alone exhibits very low CO₂-to-fuel-conversion efficiency under UV excitation, while Cu₂O possesses a more favorable E_CB = −1.4 eV vs. NHE (pH = 7) [103]; however, it suffers from photostability issues [13]. In the p-n Cu₂O/TiO₂ scheme, Cu(0) was not observed during photocatalytic experiments. In contrast, compared with pure Cu₂O, an increase of Cu(II)/Cu(I) ratio [67] was observed (see Figure 12c). Liu et al. had prepared a Cu/TiO₂ catalyst tailoring Cu valence and oxygen vacancies of the composite [93]. By thermal treatment under reducing atmosphere (H₂ and He) they exhibit the formation of defect sites, i.e., O_V’s and Ti³⁺ centers which affected the CO production. Moreover, during the reduction process through calcination, there was a change on the Cu oxidation states from Cu²⁺ to Cu¹⁺ or Cu⁰. These Cu¹⁺ species can effectively trap electrons due to the more positive reduction potential of Cu⁺/Cu⁰ (+0.52eV) couple vs. that of Cu²⁺/Cu⁰ (+0.34 eV) [104]. The Cu⁺/Cu⁰ couple can play a dual role where Cu¹⁺ species trap electrons, and Cu⁰ species can effectively trap holes [93]. Xiong et al. had used TiO₂ crystals as a matrix to deposit Pt and Cu₂O NPs where Cu₂O promoted CH₄ production but suppressed H₂ evolution [92]. In that case, Pt favored the activation of H₂O while Cu₂O performed the CO₂ reduction. After the photocatalytic reaction, most of the Cu₂O phase was reduced to Cu⁰, indicating that Pt promotes e⁻-transfer to Cu₂O during photocatalysis. Afterwards, the so-formed metallic Cu⁰ enhanced the selective CH₄ production, a well-known property of metallic Cu [92].

• The Case of Cu Coupling with Non-TiO₂ Semiconductors

As we have already stated, p-n junctions of Cu₂O with appropriate n-type metal-oxides can exhibit better charge separation and enhanced photocatalytic activity for CO₂ reduction. ZnO as an n-type semiconductor possesses a large energy gap of 3.2–3.3 eV with high electron mobility and low dielectric constant, e.g., compared to TiO₂. Bae et al. [95] used ZnO-Cu₂O nanohybrids to reduce CO₂ to CH₄ in an CO₂-saturated aqueous medium without use of hole scavenger [95]. Once again, the Z-scheme was the proposed mechanism and when compared to TiO₂(P25)-Cu₂O, the ZnO-Cu₂O exhibited superior reaction activity and selectivity. Another work [96] reported the synthesis of a CuO-NaTaO₃ hybrid with the ability to reduce CO₂ to CH₃OH with a maximum yield of 1302 μmol CH₃OH g⁻¹ h⁻¹. As expected, smaller, and more uniform distribution of CuO NPs exhibited better catalytic performance. Noteworthy, in that work [96], CuO was suggested to be the CO₂ reduction site, while NaTaO₃ was the hole-scavenging site, and oxidized isopropanol to acetone. XPS analysis indicated the presence of Cu²⁺ species only, i.e., by CuO; however, a post catalytic analysis was missing.

4.1.3. Carbon-Based Materials and Core-Shell Cu-Oxide Structures

Incorporating copper oxide in carbon-based materials such as rGO, g-C₃N₄, or carbon coating, is a widely used strategy in photocatalytic CO₂ reduction. Graphene-based materials are gaining attention, due to their characteristics, high specific area and adsorption capacity, as well as excellent electron mobility and chemical stability [105]. Gusain et al. prepared rGO-CuO/Cu₂O nanohybrids and used them for photocatalytic conversion of CO₂ to CH₃OH [100] (see Figure 12d). rGO interfaced with semiconducting materials can serve as an electron bridge and enhance the electron transfer process, a step crucial for the photocatalytic reactions [106]. In that work [106], it was suggested that the photogenerated electrons of CuO can be efficiently transferred to rGO, inhibiting in this way the electron hole recombination (see Figure 12e). Moreover, the high surface area and the defects of rGO can enhance CO₂ adsorption [100]. Noticeably, CuO (Cu²⁺) nanorods grafted on rGO exhibited better photocatalytic activity than a rGO/Cu₂O (Cu¹⁺) nanocomposite. Stability experiments indicated that after six catalytic runs, there were no significant changes in the morphological and chemical characteristics of the nanocomposite [100] (see Figure 12f). In another work, Chang et al. fabricated Cu₂O/gCN nanocomposites with different Cu₂O morphologies [101]. The main gaseous product of CO₂ reduction was CO with c-Cu₂O/gCN being the best performing photocatalyst. The high activity of c-Cu₂O/gCN was suggested to be linked with the improved CO₂ adsorption as well as the formation of CuO [101].

5. Conclusions

Herein, we discuss cases of pertinent Cu-based materials, that under appropriate coupling and engineering (facet and strain) aimed to resolve stability problems (corrosion and photocorrosion). There are many factors to be taken into consideration, such as the catalytic environment, the wavelength and intensity of irradiation source, and primarily the Cu oxidation state, which is proven to be a crucial factor determining the catalytic performance. Fundamental steps towards the stabilization and durability of Cu-based materials include the selection of a proper semiconductor coupling, suitable sacrificial agents, surface protection through core-shell structures and stabilization on defect sites of the substrate matrix. Z-scheme mechanisms could lead to the stabilization of Cu₂O through enhanced charge-separation of the photogenerated carriers. In addition, the proper selection of Cu oxidation states can tune the selectivity towards specific CO₂ reduction products. It is important to mention that the vast majority of Cu-based nanocatalysts discussed herein were noble-metal free. To conclude, the reviewed works indicate that the Cu-oxide nanophases constitute highly potent nanoplatforms for development of efficient artificial photosynthesis catalysts. There is converging evidence that—probably—pure Cu-phases would be difficult to be used as standalone cathodic catalysts or electrodes; however, their heterojunctions with proper partner materials are an encouraging approach. Distinction between the role of numerous factors is required, protection from photocorrosion vs. CO₂-reduction pathways, band-gap engineering, nano-facet engineering.