**Preface to "Photocatalytic Processes for Environmental Applications"**

Photocatalysis, especially heterogeneous photocatalysis, is one of the most widely investigated processes for environmental applications. Indeed, the activation of photocatalytic materials by solar light, which is an infinite source of energy, is considered as a sustainable technology. The conversion of solar light into chemical energy via heterogeneous photocatalysts is a powerful tool to decontaminate different environment compartments (water, air, and soil). In addition to photocatalysis, other photochemical processes can be used for environmental purposes, including the photo-Fenton process and photolysis. This Special Issue on "Photocatalytic Processes for Environmental Applications" summarizes these crucial challenges faced by current and future generations.

> **Olivier Monfort and Yanlin Wu** *Editors*

### *Editorial* **Photocatalytic Processes for Environmental Applications**

**Olivier Monfort 1,\* and Yanlin Wu <sup>2</sup>**


Photocatalysis, especially heterogeneous photocatalysis, is one of the most investigated processes for environmental remediation. Indeed, the use of solar light, which is an infinite source of energy, to activate photocatalysts is a sustainable technology [1]. The conversion of solar light into chemical energy via a heterogeneous photocatalyst is a powerful tool to decontaminate the natural environment (water, air, and soil), but also to produce green energy such as hydrogen, which is considered as the most promising alternative to fossil energy [1]. Beside photocatalysis, other photochemical processes can be used for environmental applications, including the photo-Fenton process and photolysis. The Special Issue on "Photocatalytic Processes For Environmental Applications" summarizes all these crucial challenges that the future generation will have to solve. The Special Issue is available online at https://www.mdpi.com/journal/processes/special\_issues/photocatalytic\_processes (access date: 19 November 2021).

A well-known family of heterogeneous photocatalysts are transition metal chalcogenides. Among them, the most popular one is titanium dioxide (TiO2), which is still the most investigated owing to its non-toxicity, low cost, and relatively high efficiency under UVA light. While several methods including the design of composite and doping are often found to successfully increase the photocatalytic activity of TiO2 under visible light, innovative methods could also be used, such as photocatalytic activation through electron beam [2]. In this special issue, Gallegos et al. have demonstrated that microparticulate TiO2 activated with e-beam exhibited a higher degradation rate constant for Direct Blue 1 removal [2].

It is worth reminding the photocatalytic mechanism during the degradation of organic pollutants in water. After the generation of electron/hole (*e*−*/h+*) pairs under suitable irradiation (*hν > Eg*), highly reactive inorganic radicals, especially reactive oxygen species (ROS), are produced:

$$\text{H}^+ + \text{OH}^- \rightarrow \text{HO}^\bullet \tag{1}$$

$$
\epsilon^- + \mathcal{O}\_2 \to \mathcal{O}\_2^{\bullet -} \tag{2}
$$

These "primary" radicals are responsible of the efficient degradation of water contaminants.

On the other hand, as UVA light represents only 5% of the solar spectrum, intense research is also focused on visible light-driven photocatalysts like, for example, vanadiumand iron-based oxides (V2O5, BiVO4, Fe2O3, CoFe2O4, and so on), as well as metal sulfides (CdS, ZnS, CuS, and so on) [1,3,4]. Some of these materials are ternary oxides. Indeed, it helps to tune the energy bandgap, which is an essential feature in heterogeneous photocatalysis [1]. In addition, another important physical factor that should be considered in heterogeneous photocatalysis is the morphology of materials. Morphology can influence the specific surface area as well as the transport properties of charge carriers. In this way, Ullah et al. have published in this Special Issue work related to the preparation of CdS nanorods, which have excellent photocatalytic properties in the degradation of Rhodamine

**Citation:** Monfort, O.; Wu, Y. Photocatalytic Processes for Environmental Applications. *Processes* **2021**, *9*, 2080. https:// doi.org/10.3390/pr9112080

Received: 11 November 2021 Accepted: 12 November 2021 Published: 20 November 2021

**Publisher's Note:** MDPI stays neutral with regard to jurisdictional claims in published maps and institutional affiliations.

**Copyright:** © 2021 by the authors. Licensee MDPI, Basel, Switzerland. This article is an open access article distributed under the terms and conditions of the Creative Commons Attribution (CC BY) license (https:// creativecommons.org/licenses/by/ 4.0/).

B [3]. Moreover, photocatalysts can be composed of chemical elements that are active in Fenton-based reactions (Equations (3)–(5)). It is the case of CoFe2O4 that has been investigated by To Loan et al. [4]. Indeed, this visible light-driven photocatalyst in combination with Fenton-based processes can almost completely degrade Rhodamine B [4].

$$\text{Fe}^{2+} + \text{H}\_2\text{O}\_2 \rightarrow \text{Fe}^{3+} + \text{OH}^- + \text{HO}^\bullet \tag{3}$$

$$\text{Fe}^{3+} + \text{H}\_2\text{O}\_2 \rightarrow \text{Fe}^{2+} + \text{H}^+ + \text{HO}\_2^\bullet \tag{4}$$

$$\text{Fe}^{3+} + \text{H}\_2\text{O} + h\nu \rightarrow \text{Fe}^{2+} + \text{H}^+ + \text{HO}^\bullet \tag{5}$$

Concerning these Fenton-based processes, they are usually performed in homogeneous systems using dissolved iron species, thus exhibiting higher kinetic rates compared with heterogeneous systems. By looking at these reactions (Equations (3) and (4)), it appears as a catalytic process owing to the iron cycle (Fe2+ ↔ Fe3+). However, in an aqueous solution, Fenton-based reactions face two main drawbacks, including (i) the stability of iron species and (ii) the reduction of Fe3+ into Fe2+ (which is the limiting step). Therefore, using non-toxic and biodegradable ligands, stable iron complexes can be used under visible light, thus allowing their use at environmental pH and favoring the reduction process via iron photolysis (Equation (5)). That has been investigated in the work of Huang et al., where 2,4-dichlorophenol has been successfully degraded using the Fe(III)–EDDS complex in both synthetic and real water bodies [5].

Another interesting photochemical process is the photolysis of radical precursors. In such a configuration, no material is needed as the action of light of suitable energy can generate ROS like HO•. Usually, hydroxyl radicals can be generated by photolysis of H2O2 under UVB light, while strong UVC is required to photolyze water molecules. For the latter, a VUV/UV photoreactor is necessary. In this Special Issue, Luo et al. have investigated the fluid dynamics in a VUV/UV photoreactor by correlating simulations with experiments for the degradation of organic pollutants [6]. Therefore, their work brings interesting knowledge for the development of such a technology for water treatments.

The improvement of photochemical processes can also be reached using physical methods. An example is the application of a voltage to a photocatalyst in the form of a photoelectrode. Such a photoelectrocatalytic process can be employed for various applications including the production of hydrogen by water splitting, the enhanced degradation of water contaminants, and the production of hydrogen peroxide. In this Special Issue, Papagiannis et al. have studied the latter system using a photoanode [7]. Indeed, the production of H2O2 by WO3 photoanode is made sustainable by using organic fuels from biomass derivatives such as glycerol and ethanol. H2O2 is a well-known oxidant that is used in both conventional oxidation and advanced oxidation processes. It is a radical precursor used in the production of hydroxyl radicals.

This Special Issue on "Photocatalytic Processes For Environmental Applications" offers an overview (through original full-length and review articles) of the different photochemical processes (photocatalysis, photo-Fenton, and photolysis) triggered by different inorganic compounds that can be used for environmental applications, not for only water treatment, but also for hydrogen production. As guest editors, we thank all the contributors and the editorial team of *Processes* for their work and the successful publication of this Special Issue.

**Funding:** There is no external funding.

**Conflicts of Interest:** The authors declare no conflict of interest.

#### **References**


## *Review* **Binary and Ternary Vanadium Oxides: General Overview, Physical Properties, and Photochemical Processes for Environmental Applications**

**Olivier Monfort 1,\* and Patrícia Petrisková <sup>2</sup>**


**Abstract:** This review article is a comprehensive report on vanadium oxides which are interesting materials for environmental applications. Therefore, a general overview of vanadium and its related oxides are presented in the first two parts. Afterwards, the physical properties of binary and ternary vanadium oxides in single and mixed valence states are described such as their structural, optical, and electronic properties. Finally, the use of these vanadium oxides in photochemical processes for environmental applications is detailed, especially for the production of hydrogen by water splitting and the degradation of organic pollutants in water using photocatalytic and photo-Fenton processes. The scientific aim of such a review is to bring a comprehensive tool to understand the photochemical processes triggered by vanadium oxide based materials where the photo-induced properties are thoroughly discussed based on the detailed description of their intrinsic properties.

**Keywords:** vanadium oxide; photochemistry; environment; materials; pollutant; hydrogen

#### **1. Introduction**

Vanadium (V) is a chemical element discovered by Andres Manuel Del Rio in 1801 in the form of an ore, called today vanadinite (Pb5(VO4)3Cl) [1,2]. At that time, vanadium was named "panchromium" since its different salts exhibit a wide variety of colors. In 1831, Nils Gabriel Sefstrom (re)discovered this chemical element in the form of oxide [3]. He called this chemical element "vanadium" after Vanadis, a goddess in the Norse mythology. Vanadium is an early first-row transition metal (Group 5; *Z* = 23), with the electronic configuration of [Ar]3d34s2 [4]. In the earth's crust, vanadium is 5th (0.019%) among all transitional metals and 22nd among all discovered elements [5,6], that corresponds to an average amount of 159 g per ton [7]. In the soil, the average abundance of vanadium is > 100 mg·kg<sup>−</sup>1, but this value depends strongly on the geology and the human activity [2,7]. Vanadium can be found in various minerals in the form of either vanadate, silicate, sulfide, sulfate, phosphate, or oxide, such as vanadinite, carnotite, roscoelite, patronite, bravoite and davidite [2,8]. Concerning the physico-chemical properties of this transition metal, vanadium has a melting point at 1910 ◦C and forms oxide with different oxidation states (from +II to +V) and with various crystalline structures [9].

Vanadium is a transition metal of high environmental and biological importance. Indeed, vanadium is ubiquitously distributed in soil, water, air and living organisms. For instance, in sea water, vanadium (mainly in the form of H2VO4 −) is the second most abundant transition metal element with a mean concentration of 2 <sup>μ</sup>g·L−<sup>1</sup> [2,10] while in rivers, lakes and groundwater, the maximum concentration of vanadium is up to <sup>30</sup> <sup>μ</sup>g·L−<sup>1</sup> [2,11]. The presence of vanadium in the environment is due to both natural and human activities. The main natural sources for vanadium are marine aerosols, dust

5

**Citation:** Monfort, O.; Petrisková, P. Binary and Ternary Vanadium Oxides: General Overview, Physical Properties, and Photochemical Processes for Environmental Applications. *Processes* **2021**, *9*, 214. https://doi.org/10.3390/pr9020214

Academic Editor: Monika Wawrzkiewicz Received: 30 December 2020 Accepted: 21 January 2021 Published: 24 January 2021

**Publisher's Note:** MDPI stays neutral with regard to jurisdictional claims in published maps and institutional affiliations.

**Copyright:** © 2021 by the authors. Licensee MDPI, Basel, Switzerland. This article is an open access article distributed under the terms and conditions of the Creative Commons Attribution (CC BY) license (https:// creativecommons.org/licenses/by/ 4.0/).

from soils and rocks weathering, and volcanic emissions [12–15]. From anthropogenic activity, vanadium loads are mainly from mining and burning of crude oil, thus resulting in atmospheric pollution. Vanadium from human contamination is also observed in water resources (rivers, lakes and seas). The quantity of vanadium released into the atmosphere is estimated at more than 60,000 tons each year, and mostly from combustion of fossil fuels [16]. In addition, V has the ability to stay in the air, water and soil for long period of time [17].

In biochemistry, vanadium is an essential trace element for several mammals. Indeed, V has a key role in some enzymes such as nitrogenase and haloperoxidase [2]. Vanadium enters in the blood through gastro-intestinal and respiratory systems and then is transported into the other parts of the body in the form of citrates, lactates, or phosphates [18]. Most of V accumulates in vital organs such as kidneys, spleen and liver [19]. In the human body, about 100 μg of vanadium is present. It is important to notice that vanadium in the form of vanadate (VO4 <sup>3</sup>−) is isostructural to phosphate. It is therefore a competitive element with PO4 <sup>3</sup><sup>−</sup> and can inhibit and/or motivate several phosphate-metabolizing enzymes [20]. Therefore, at high levels, vanadates can be fatal for human [2]. For plants, it is known that V can affect the metabolism including mineral uptake, enzymatic activities, photosynthetic activity, and the biological yield of plants.

This review article starts with an overview of the applications of various vanadium compounds (Section 2). Afterwards binary and ternary vanadium oxides are introduced (Section 3) with an emphasis on their structure and physical properties before a discussion on their use in photochemical processes for environmental applications (Section 4). Usually, transition metal oxides (TMOs) and especially vanadium oxides have unique physical, chemical, optical, electronic, thermal, and magnetic properties. Therefore, vanadium oxides can find promising applications in various fields such as energy conversion and storage [21]. The particularity of vanadium oxide based materials is that most of them exhibit an insulator-to-metal transition (IMT). Indeed, below a critical temperature (*Tc*), the properties of the material are insulating or semiconducting while at temperature above they have a metallic behavior. There are different vanadium oxides, therefore the IMT are observed over a wide range of temperatures depending on the oxide, i.e., O/V ratio [22,23]. Due tof this IMT, vanadium oxides are known to be chromogenic materials: they can change their optical properties by applying some external stimuli in the form of photon radiation, temperature change and voltage pulse, thus defining vanadium oxides as photochromic, thermochromic, and electrochromic materials, respectively [22].

In addition to the potential use of vanadium oxides in new technologies, the challenge of environmental remediation is one of the most crucial issues for the present and future generations. The deterioration and contamination of the environment combined to the threat of energy shortage have led the society to explore alternatives to fossil fuels and to pay more attention to our actual way of life. One of the main goals of national governments is to reduce their energy consumption and greenhouse gas emission and to save the natural environment [24]. Indeed, international agreements, such as Kyoto protocol (1995), Stockholm convention (2001), and Paris agreements (2015) have been signed, thus limiting our impact on the planet. Although these attempts are unprecedented, the development of green and efficient remediation treatments is necessary. Our natural environment has been considerably damaged, especially from the last industrial revolution, and is still being deteriorated due to the human activity, in particular factories in the field of textile, chemistry, agriculture, and pharmacy [25–27]. For instance, the quality of water which is a rare and precious resource (around 0.5% of water on the Earth is drinkable and accessible) is alarming because the different types of industry release many organic and inorganic pollutants such as pesticides and fertilizers, sulfur derivatives, active pharmaceutical ingredients, etc. [25–28]. Among them, many pollutants are persistent and of emerging concern, i.e., they are known compounds which are detected by accumulation effect and whose chemistry and interaction with the environment are unknown. Such pollutants may have harmful effects because they include endocrine disruptor compounds, pharmaceutical

and personal care products, etc. [29–32]. These contaminants are potentially toxic for humans, animals and plants, thus pushing the national environmental agencies to impose strict measures to limit the pollution [26]. However, limiting the pollution does not remediate the damages already caused to the planet [33]. In addition, the water resources are subjected to an imbalance between availability and demand in water which is caused by its overuse for domestic, industrial, and agricultural purposes, but also by the climate disorder (longer, more intense and more frequent droughts).

Regarding these issues, vanadium oxides are promising materials. As semiconductors, vanadium oxides can be employed in photocatalysis which is a viable alternative to conventional biological, chemical, and physical technologies for the removal of pollutants, especially in water [25,26]. In addition, these photocatalytic (and other photo-induced processes) can be triggered by solar energy for environmental applications. In addition to decontamination treatments (Section 4.2), another task is the supply of sustainable and clean energy (Section 4.1). Indeed, vanadium oxide based materials are also able to convert solar energy for the production of hydrogen. To resume, binary and ternary vanadium oxides are attractive for a wide range of environmental applications [9]. Despite the numerous scientific contributions on vanadium oxides, the novelty of this review article is to (i) introduce this versatile element in a global overview, followed by (ii) a strong emphasis on the physical properties of vanadium oxides based materials and their use in (iii) photochemical processes for environmental applications which are one of the biggest issues in the modern time.

#### **2. Vanadium: A Versatile Element for Various Applications**

#### *2.1. Vanadium Compounds in Homogeneous Systems*

The common oxidation states of vanadium are from +II to +V. The chemistry of this transition element in aqueous solution is a key point to understand its versatile properties. In water, vanadium with oxidation states from +II to +V form either aqua-complexes or oxo-cations with typical colors. The hexa-aqua-complexes [VII(H2O)6] 2+ and [VIII(H2O)6] 3+ are violet and green, while vanadyl (VIVO2+) and pervanadyl (VVO2 +) cations are blue and yellow, respectively. Regarding the aqueous redox reactions (Equations (1)–(3)), vanadium (II) is a reducing agent while vanadium (V) is an oxidant [34,35]. In the environment, V (IV) and V (V) are in the form of vanadyl and vanadate, respectively, while V (II) is thermodynamically unstable and V (III) is stable only under strongly anoxic conditions [2].

$$\text{VO}\_2\text{}^+ + 2\text{H}^+ + e^- = \text{VO}^{2+} + \text{H}\_2\text{O}; E^0 = 1.0\text{ V} \tag{1}$$

$$\rm{VO^{2+}} + 5\rm{H\_2O} + 2\rm{H^+} + e^- = [\rm{V(F\_2O)\_6}]^{3+}; E^0 = 0.24\text{ V} \tag{2}$$

$$\left[\text{V(H}\_{2}\text{O)}\_{6}\right]^{3+} + e^{-} = \left[\text{V(H}\_{2}\text{O)}\_{6}\right]^{2+}; E^{0} = -0.26\text{ V} \tag{3}$$

In this section, the chemistry of vanadium (V) in aqueous solution is briefly discussed since many molecular species have been discovered [36]. The speciation of these V (V) compounds depends mainly on vanadium concentration, temperature, and solution pH. There are basically two main types of reaction which control the formation of vanadium (V) species: hydrolysis and condensation [37]. At room temperature, dissolved V (V) salts are solvated in aqueous solution by water molecules, thus giving rise to hydrated species [V(H2O)*n*] 5+ [37]. However, the water molecule ligands are partially deprotonated (Equation (4)) due to the strong polarizing power of highly charged vanadium (V). This hydrolysis process results in a decrease in pH.

$$\left[\mathrm{V(H\_2O)\_6}\right]^{5+} + h\mathrm{H\_2O} \rightarrow \left[\mathrm{V(H\_2O)\_{6-h}(OH)\_h}\right]^{(5-h)+} + h\mathrm{H\_3O}^+\tag{4}$$

The term "*h*" is the hydrolysis ratio and it increases with increasing pH, thus leading to the formation of aqua, hydroxo, and oxo species. Therefore, at low pH (pH < 2), the hydrolysis ration *h* = 4 and [V(OH)4(H2O)2] <sup>+</sup> is formed. Since V (V) is highly charged, internal proton transfer occurs to decrease this high positive charge and it leads to the

formation of pervanadyl species [VO2(H2O)4] <sup>+</sup> or [VO2] <sup>+</sup> with V=O double bonds (vanadyl groups) in *cis* position [38]. At pH ≈ 6, the coordination number of V (V) decreases from 6 to 4 since V-O bonds are more covalent, thus leading to the formation of four-fold coordinated vanadate [H*n*VO4] (3−*n*)−. Above pH = 12, vanadate species are fully deprotonated. It is worth noting that monomeric species are observed only in dilute solutions. At higher vanadium concentration, condensation process dominates and occurs through olation and oxolation (Equations (5) and (6)) [37]:

$$\text{Olation:} \text{ >V-OH} + \text{>V-OH}\_2 \rightarrow \text{>V-O(H)} - \text{V} < + \text{H}\_2\text{O} \tag{5}$$

$$\text{Oxidation:}\ > \text{V}-\text{OH} + > \text{V}-\text{OH} \rightarrow > \text{V}-\text{O}-\text{V}< + \text{H}\_2\text{O}\tag{6}$$

The temperature has also strong effect on the molecular structure of vanadium species [37]. Once vanadium precursors have been condensed to decavanadate species, they can transform upon heating into cyclic metavanadates (Equation (7)) [39]. This dissociation process is a reversible reaction:

$$\text{2[H}\_2\text{V}\_{10}\text{O}\_{28}\text{]}^{6-} + 4\text{H}\_2\text{O} = 5\text{[V}\_4\text{O}\_{12}\text{]}^{4-} + 12\text{H}^+\tag{7}$$

In addition to the chemistry of aqueous vanadium species, their biological relevance (Figure 1) is also a point which needs to be briefly developed. Only higher oxidation states of vanadium are considered in the physiological pH (2 < pH < 8), i.e., V (IV) and V (V) in cationic and anionic forms. On the other hand, V (III) is mainly present in primitive organisms such as ascidians and worms and it is not present in complex organisms [40,41]. Vanadium at higher oxidation states is present in most of mammal tissues (around 20 nM). In the human body, the average vanadium concentration is approximately 0.3 μM [42–48]. Vanadium in the form of vanadate (H2VO4 −), which can be present in drinking water, is partially reduced in the stomach before being precipitated in the intestines in the form of VO(OH)2 [10]. Vanadium can also enter directly in the blood stream by intravenously injection [49]. Once in the blood, vanadium undergoes redox conversion between V (V) and V (IV) according to the level of oxygen and the presence of reductants/oxidants. Thus, vanadium could form either anionic VO4 <sup>3</sup>−, cationic VO2+ or neutral/charged vanadium complex. Concerning the toxicity of vanadium, the limit values to observe harmful effect on human health during one-time exposure is around 7 mg by intravenous application and 35 mg by inhalation [10,50]. Vanadium contents (in the blood) can decrease to about 30% within 24 h [49,50] since it is eliminated via urine or distributed in tissues. In bones, the residence time is of about one month since vanadium can replace phosphorus in hydroxyapatite [51]. The vanadium path in the body is summarized in Figure 1.

On the other hand, vanadium has high potential as metallodrugs, especially for the treatment of diabetes which was the first use of vanadium in pharmacology. Indeed, in the form of salts including sodium metavanadate, sodium orthovanadate, and vanadyl sulphate, the effect on patient whom needed insulin is positive. This is due to the ability of these vanadium species to activate insulin receptor [42]. Vanadium derivatives can also inhibit/activate several ATPases and phosphatases enzymes since vanadate is a phosphate mimic anion [34,42]. Vanadium is also used for the treatment of several pathologies such as hyper-lipidemia, obesity, and hyper-tension [42]. It is worth mentioning that V has anti-cancer properties. Indeed, vanadium compounds can be used to prevent wear and tear of essential biomolecules such as DNA, thus protecting the genomic stability [52,53]. Vanadium derivatives can be also used in anti-bacterial and anti-viral applications [34,42].

**Figure 1.** The pathway of vanadium in the human metabolism.

In addition to medicinal applications, vanadium can be used in the field of energy production as, for instance, the redox flow battery. A redox flow battery is similar to an electrochemical cell composed of two compartments and two electrolyte reservoirs (Figure 2). The electrolytes circulate through the cell by a pump while the compartments are separated by an ion exchange membrane to allow the diffusion of ions (Figure 2) [54]. The main disadvantage in typical redox flow batteries such as iron/chromium or iron/titanium is the cross contamination of the electrolytes due to different redox couple species in each half-cell compartment [55,56]. In vanadium redox flow batteries, this problem is overcome by employing the same electrolyte in both compartments, thus eliminating the cross-contamination as well as the problems from the maintenance of electrolyte solution. Indeed, a vanadium electrolyte is used in both half-cells where the V2+/V3+ redox couple operates in the negative one while the VO2+/VO2 <sup>+</sup> redox couple operates in the positive half-cell [54]. Therefore, such pioneered vanadium-based batteries are considered as the 1st generation of such a type of battery.

**Figure 2.** Scheme of a redox flow battery.

All the above applications of vanadium in homogeneous systems are very different from each other due to the chemical versatility of vanadium. Depending on its oxidation states, this element promotes a wide variety of chemistries, especially in catalysis (Figure 3). Concerning the chemistry of coordination compounds, vanadium is very flexible, especially V (V). Vanadium (V) has not rigid stereochemical requirements and, thus, can form a variety of complexes, including tetrahedral and octahedral geometries, as well as trigonal and pentagonal bipyramids [57]. On the other hand, vanadium (IV) is less flexible and forms mainly square pyramidal complexes or distorted octahedral geometries [57]. Vanadium complexes like peroxido-vanadium complexes can be involved in biological and catalytic applications [34,58]. Indeed, catalytic properties of peroxido-vanadium complexes are focused on the oxidation of organic compounds. In these oxidative reactions, the complex is formed in-situ from either ammonium vanadate or V2O5 by reaction with hydrogen peroxide and the resulting complex can oxidize aliphatic and aromatic hydrocarbon structures [58]. Other catalytic applications involving vanadium complexes exist, such as polymerization, C-C bond cleavage and activation, hydrogenation, etc. (Figure 3) [59–61].

**Figure 3.** Some of the possible uses of vanadium in catalytic processes.

#### *2.2. Vanadium Materials in Heterogeneous Systems*

In this section, other applications than environmental ones (Section 4) are discussed. In the form of solid materials, vanadium is mainly used in metallurgy for the formation of ferrovanadium alloy (FeV) or as steel additive [34]. Mixed with aluminum in titanium alloys, vanadium is also used in jet engines, high-speed airframes, and dental implants [34]. The metallurgy accounts for 85% of the vanadium use. Moreover, vanadium is also used for the storage of hydrogen [62]. Hydrogen is considered as the alternative source of energy to fossil fuels. Even if it is an eco-friendly energy which is relatively easy to produce, its storage and transportation are crucial for the development of hydrogen-based energy system since hydrogen is an explosive compound [62,63]. Usually, the storage and transportation of hydrogen is based on high-pressure gaseous H2 and cryogenicallycooled liquid hydrogen. In such conditions, the storage and transportation is associated with safety issues and unpractical heavy tanks [64]. Therefore, to avoid these problems, hydrogen can be stored in the form of metal hydrides like those of vanadium which are considered as a viable method [63]. In addition to the high hydrogen storage capacity of vanadium hydrides including vanadium based alloys also exhibit good hydrogenationdehydrogenation kinetics at room temperature [62].

On the other hand, vanadium can be used in the form of chalcogenides (oxides, sulfides, etc.), nitrides, carbides, and a number of halides which are stable materials [35]. Vanadium sulfide (VS2) which belongs to the group of transition metal dichalcogenide (TMD) has a structure composed of V layers sandwiched between two layers of S [65]. These layers are stacked together by Van der Walls interactions. Vanadium sulfide has interesting electrical properties, especially high 2D metallic behavior. This 2D conducting material is of high interest for the design of planar supercapacitors [66,67]. Indeed, VS2 in the form of nanosheets can be used in next-generation energy storage devices [68]. However, the most important vanadium chalcogenide are the oxides which are the most commercially available. Among vanadium oxides, V2O5 is the most used. This yellowbrown oxide can be obtained easily either by calcination of vanadium metal powder in an excess of oxygen or by thermal decomposition of ammonium metavanadate (Equation (8)):

$$2\text{NH}\_4\text{VO}\_3 \rightarrow \text{V}\_2\text{O}\_5 + 2\text{NH}\_3 + \text{H}\_2\text{O} \tag{8}$$

Nowadays, vanadium pentoxide is a key component in the production of sulfuric acid since it catalyzes the oxidation of sulfur dioxide to sulfur trioxide. In addition, V2O5 and materials containing vanadium pentoxide are used in many other catalytic reactions [34,69]. Due to its catalytic activity and selectivity, V2O5 is a known catalyst for the selective oxidation of various organic compounds, mainly hydrocarbons and alcohols, e.g., butane to maleic anhydride, propene to acrolein and methanol to formaldehyde [34,70]. These latter cases are examples of commercial products synthesized using vanadium pentoxide. V2O5 is also used in other oxidative processes such as the oxidative dehydrogenation of alkanes [34,71]. However, V2O5 and generally vanadium oxide based materials are also important catalysts for environmental applications which are developed in Section 4. Besides the catalytic properties, the layered structure of vanadium pentoxide is promising for the design of electrodes in lithium ion batteries (LIB) [72]. Indeed, V2O5 has high ionic storage capacity and as a cathode material, vanadium pentoxide has a theoretical energy density of about 1.1 kW·h·Kg−<sup>1</sup> which corresponds to 440 mA·h·g−<sup>1</sup> for fully-lithiated V2O5 (Li3V2O5) [73]. It is far more than classical cathode material like coarse-grained LiCoO2 which has a practical energy density < 0.5 kW·h·Kg−<sup>1</sup> (274 mA·h·g−1). However, charging more than one Li ion per V2O5 unit cell leads to irreversible structural changes [74]. When one lithium is stored (LiV2O5), the lattice structure is extremely reversible with over 1000 cycles without capacity loss and the energy density is still higher than classical CoO2 cathode (294 mA·h·g−1) [75]. It is worth noting that Panasonic has commercialized vanadium pentoxide as cathode material in rechargeable Li-ion batteries. Such a cathode can be used for low-energy application only because of kinetic problems, thus other layered vanadium oxides should be designed to improve the LIB performance [76]. That is the case of V6O13 which is obtained from V2O5 (Equation (9)). Indeed, the theoretical specific capacity of V6O13 reaches up to 420 mA·h·g−1. In addition, 8 mol Li+ can be inserted in the layered structure of 1 mol V6O13. However, the poor cycling performance of this vanadium oxide is one of the most serious drawbacks for a use as cathode materials in LIBs [77]. Therefore, another strategy is to substitute V cations or vanadium vacancies with other cations such as Ag+, Mn2+, and Ti4+, thus improving the electrochemical properties of V6O13 cathode.

$$\rm{^2\theta V\_2O\_5 + 4NH\_3 \to 3V\_6O\_{13} + 2N\_2 + 6H\_2O} \tag{9}$$

Vanadium dioxide is also an intensively studied vanadium oxide phase, especially for its thermochromic properties. Indeed, in the form of coating, VO2 can change its appearance from transmissive to opaque due to a reversible insulator-to-metal transition (IMT) [78]. At a critical temperature (*TC* = 68 ◦C) and without any additional external stimuli, phase transition between IR-transparent VO2(M) and IR-translucent VO2(R) occurs (Figure 4). The insulator-to-metal transition is not only accompanied by optical changes, but also by changes in electrical properties. For instance, the electrical resistivity of VO2 varies at the critical temperature from 20 Ω·cm in the insulating phase to about 0.1 Ω·cm in the metallic phase, which is accounts to a factor 200 [79]. This reversible change in electrical properties occurs within 100 fs [80]. Regarding the ultrafast switching time and the sharp change in resistivity, VO2 is an excellent temperature sensor. Regarding all these properties, the VO2 IMT is an excellent feature for the design of smart windows since energy consumption can be saved and greenhouse gas emission reduced [9]. This is an actual issue since, according to a survey, buildings are one of the main problems in this field with an energy consumption of 40% along with 30% of anthropogenic greenhouse gas emission [81]. These negative statistics are due to the excessive use of lighting, airconditioning and heating [81]. In addition, the heat exchange between building indoor and

the outdoor environment are occurring through the windows and, thus, amplifying the energy consumption.

**Figure 4.** Principle of VO2 thermochromic coating deposited on transparent support.

The vanadium oxides are also widely used in recent technologies such as memory devices and sensors [82]. Vanadium oxide based materials are attractive due to their easy way of fabrication which is inexpensive, non-toxic, and using scalable wet chemical techniques [83]. It is important to remind that vanadium can form a large number of oxides but only few of them provide stability needed in microelectronic devices. That is the case of VO2, V2O3, and V2O5 which are oxides containing vanadium in a single oxidation state. Because of their reversible changes of physical properties, these thermoand electro-chromic vanadium oxides exhibit hysteresis loop. The width of this loop is a crucial parameter for applications in microelectronics. Indeed, the memory aspect of vanadium oxides materials lies in the hysteresis of the insulator-to-metal transition. One of the most known applications involving vanadium oxides, such as V2O3, VO2, and V2O5 are bolometers. A bolometer is a thermal infrared detector. Therefore, bolometers are used in infrared imaging applications such as thermal camera, night vision camera, mine detection, early fire detection, medical imaging and gas leakage detection [84]. Materials which are used in bolometers require high temperature coefficient of resistance (TCR) and a small noise constant (1/f) [84]. Vanadium oxides belong to this category. Although pure and stoichiometric binary oxides such as VO2, V2O3, and V2O5 have high TCR, they are difficult to synthesize as single crystals. V2O3 is an excellent candidate due to the critical temperature of its IMT transition which is far below the room temperature and, thus, the level of noise is low. However, the actual strategy is to use systems mixing different single-valence vanadium oxides in order to design stable microbolometric materials with high TCR and resistivity [85]. The use of vanadium oxides for such bolometric applications appeared after 1992 since, fprior to this date, it was classified by the US government.

#### **3. Structures and Physical Properties of Binary and Ternary Vanadium Oxides**

#### *3.1. Case of Binary Vanadium Oxides*

As most of the d-transition metals, vanadium has different oxidation states which are present in oxides as either single or mixed valence states. The vanadium oxides existing in a single valence state are typically VIIO, VIII2O3, VIVO2, and VV2O5 while couples of mixed valence oxides exist such as V3O5, V4O7, V5O9, V6O11, V6O13, V7O13, and V8O15. These latter oxides mix mainly two valence states of vanadium [22]. The formation of vanadium oxide of particular oxidation state(s) is highly dependent on the preparation method, i.e., the experimental parameters such as the temperature, the partial pressure of the oxidant/reducer [73]. For instance, the following phases are formed and transformed subsequently in an oxidative atmosphere in the order: VO2, V6O13, V4O9, V3O7, and V2O5. This order of appearance during the synthesis is closely related to the increasing oxidation

state. For these oxides, the oxidation states for the vanadium is 4.0, 4.3, 4.5, 4.7, and 5.0, respectively [86,87]. Such vanadium oxides ranging from V2O5 to VO2 belong to the Wadsley series, i.e., oxides of general formula V*n*O2*n*+1 [72]. Another important category of vanadium oxides is the so-called Magnéli phases with a general formula V*n*O2*<sup>n</sup>*−<sup>1</sup> [72].

The structure of the most important vanadium oxides are described below along with their properties. The rich coordination-style configurations for the vanadium ions, which is also one reason why vanadium has various valence states, leads to the formation of numerous and varied frameworks [88]. The control of vanadium oxide structures is not an easy task. Usually in transition metal oxides, the electronic properties are related to the sband associated with the transition metal ion and the p-band associated with the oxygen ion. As a consequence, these orbitals are pushed away from the Fermi level while the remaining d-orbitals are close to this level [89]. Thus, the d-orbitals of the metal are of significant importance in the explanation of electronic properties using either crystal field theory or conventional band theory. In the crystal field theory, the splitting energy for 3d-series ions in oxides has a value of 1–2 eV and this value increases as cation oxidation state increases. In conventional band theory, which treats the electrons as extended plane wave, the metal d-orbitals form the conduction band. According to this theory, the semiconductors are defined as materials having completely filled valence band and an empty conduction band separated by a short energy band gap (*Eg*). The structure and the morphology of solid state materials are mainly responsible of the observed properties [37]. In other words, the physical properties of a material depend strongly on its structure. The various atom-stacking style form the numerous crystal structures of vanadium oxides [90,91], thus providing the foundation to understand the interplay between the properties and the structures. In this section, binary vanadium oxides with both single and mixed valence states are discussed.

#### 3.1.1. Structure of VO2, V2O5, and V2O3

The oxides of formula VO2, V2O5, and V2O3 are formed of vanadium in one valence state. Their melting points are 1967, 690, and 1970 ◦C, respectively, while metallic vanadium melts at 1910 ◦C. The two ends of the Wadsley series, that are VO2 and V2O5, are known layered vanadium oxides [73]. For the highest valence state of vanadium (V5+), there is only one V2O5 phase which is the orthorhombic crystal system (space group of *Pmmn*) [92]. Indeed, as long as the reaction temperature is high enough and the oxygen is sufficiently provided, orthorhombic V2O5 would be exclusively formed [93]. The orthorhombic lattice of vanadium pentoxide is composed of *[VO5]* pyramids which form alternating double chains (pyramids-up/pyramids-down) along the *b*-axis. These double chains are connected laterally by bridging oxygens to form a sheet or a ribbon in the *a*–*b* plane. The planes themselves are connected by Van der Waals bonds [83]. The Van der Waals bonds are weak and create easy cleavage along these planes [83]. Since V2O5 has a layered structure it is, therefore, possible to extract few layers (of nanometer thickness) or even a monolayer of vanadium pentoxide [22]. Although the sheets/ribbons along the *b*-axis direction are very robust, the V2O5 lattice is fairly open and permeable to smaller intercalating cations. In addition, the surface of the Van der Waals planes is polar, thus allowing hydration phenomenon [94]. Theoretically, in the crystal structure of V2O5, the nominal valence state of vanadium atoms is +V, but thermal agitation and crystal defects create a small concentration of V (IV) in the solid. Therefore, electronic conduction occurs by means of electron hopping. This is equivalent to the motion of the +IV valent vanadium states. The presence of V (IV) in V2O5 also creates local lattice distortion [83].

For VO2, more than 10 types of V (IV) oxide phases have been reported and the preparation of the desired structure is hard since the structures are very similar. Among these vanadium dioxides, VO2(M), VO2(R) and VO2(B) are the three main phases, from which VO2(B) is metastable. The other vanadium dioxide polymorphs are metastable phases including VO2(A), VO2(B), VO2(C), and VO2(D), but also mineral phases such as paramontroseite and nsutite-type structures [95,96]. Compared to VO2(M) and VO2(R),

the synthesis of metastable phases usually requires a lower temperature. The VO2(M/R) exhibits an IMT where VO2(M) and VO2(R) are the insulating and the metallic phases, respectively. Below the phase transition temperature (*TC*), vanadium dioxide has a monoclinic structure (VO2(M)) [97] while above *TC*, VO2 is in rutile form (VO2(R)) [22]. In rutile phase, each vanadium is coordinated to six nearest oxygens, thus forming *[VO6]* octahedra [83]. These *[VO6]* octahedra are symmetrical and form edge-sharing chains along the *c*-axis of the crystal lattice. This tetragonal lattice has a space group *P42/mnm* while the monoclinic phase, VO2(M), has a space group *P21/c*. At room temperature, the metastable phases also exist and VO2(B) is the most investigated one. Eventually, VO2(B) can be converted to metastable VO2(A) or VO2(M) by heat treatment [98]. VO2(B) is sometimes named as VO2(M2) since it has a monoclinic structure. The main difference between the VO2(M) and VO2(B) monoclinic structures is based on the *β* angle of the lattice structure which is 122.62◦ in VO2(M) and 92.88◦ in VO2(B). In addition, VO2(B) does not exhibit an insulating-to-metal transition, neither the tetragonal VO2(A) phase. The IMT in VO2(M/R) is induced by temperature gradient leading to a huge change in resistivity. The structural changes involve small distortions of the infinitely linear V–V chains (with a V–V distance of 0.288 nm along the *c*-axis) in VO2(R), thus giving rise to the dimerization of V–V pairs (with two values of 0.262 nm and 0.316 nm between V–V) in VO2(M) [72]. In the rutile phase, the electrons from d-orbitals are shared by all of the vanadium atoms along the *c*-direction of the lattice, thus leading to the metallic behavior. In the monoclinic phase, the vanadium atoms form dimers, thus resulting in the insulating behavior (i.e., semiconducting VO2(M)) [99]. Furthermore, during the phase transition to VO2(R), the V–V pairs in the chains in VO2(M) do not only undergo dimerization, but also a twist from linear to zigzag-type chains [72] (Figure 5). The local structural rearrangement around the V atoms during the insulator-to-metal transition provides a key factor to modify the band–gap value [72]. It is important to note that monoclinic and rutile VO2 phases have different crystal systems, but their crystal lattices are of great similarity. The only difference between these two structures is the displacement of the vanadium atoms in the chains. Therefore, this structural similarities between the metallic and insulating phases makes the understanding of the IMT more difficult.

**Figure 5.** Changes in the crystalline structure during VO2 phase transition. The schematized sequence occurs in the picoseconds scale.

Concerning V2O3, there are very limited numbers of phases. This is mainly related to the chemical instability of the low valence state vanadium ions, here V (III). However, there is still intense research in searching for new V2O3 crystal polymorphs. Similar to iron(III) oxides (Fe2O3) which is known as either hematite (α-Fe2O3), beta phase (β-Fe2O3), gamma phase (γ-Fe2O3), maghemite (δ-Fe2O3), or epsilon phase (ε-Fe2O3) [100], only δ-V2O3 and ε-V2O3 have not been discovered up to now. Vanadium sesquioxide (V2O3) also exhibits an IMT where V2O3 has a monoclinic/rhombohedral structure below *TC*, and a corundum structure above *TC* [22]. In its insulating phase (below *TC*), V2O3 has body

centered monoclinic structure with space group *I2*/*a*. The metallic phase of V2O3 (above *TC*) has a trigonal corundum structure with space group *R*3*c*.

#### 3.1.2. Properties of VO2, V2O5, and V2O3

The control of the synthesis of binary vanadium oxides of particular structure is a key parameter to tailor the desired electronic properties [72]. It is well known that the charge carrier concentration in the prepared materials is crucial to understand and explain the performance of different vanadium oxides in various applications such as energy conversion, in particular photocatalysis and water photosplitting (Section 4). Indeed, the electronic properties are usually discussed using the band structure of the semiconductor materials, which is the plot of electron energies against the wave vector. Prior to further process in this section on the properties (especially electronic ones) of binary vanadium oxides with single vanadium valence, it is important to introduce some notions of the band structure theory. The number of bands (in a band structure diagram) is equal to the number of atomic orbitals in the unit cell of the compound. The overlap integral, which is the overlap between the interacting orbitals, determines the bandwidth (or the band dispersion), i.e., it is the difference in energy between the lowest and highest levels in a band. The greater is the overlap between neighboring orbitals, the greater is the bandwidth (and vice versa). Therefore, the behavior of electrons in a solid can be studied microscopically from its electronic band structure (Figure 6) [101]. From these electronic properties, and more particularly from the band structure diagram, it is possible to extract additional important information about the material such as optical properties. Indeed, in a material, optical properties originate from the response of electrons to a stimulus which is the transition between energy levels under an incident radiation. The energy band gap calculation and the absorption edge estimation are of high interest in many research fields due to potential applications in optical, electronic, and optoelectronic devices. The electronic band structure of a material is dependent on the crystal structure of the material.

**Figure 6.** The band structures of insulator, semiconductor and metal.

Concerning the main single valence state of binary vanadium oxides (VO2, V2O5, and V2O3), their electronic structure are all characterized by a strong hybridization between the 2p oxygen and 3d vanadium bands, which are the valence orbitals. Their optical properties can be explained using their electronic properties. Since most of the binary vanadium oxides exhibit an insulator-to-metal transition which leads to electrochromic and thermochromic properties, such oxides are very interesting for the design of advanced materials. Indeed, the IMT can be triggered using different stimuli including voltage and heat. Since the vanadium–oxygen bonds are highly ionic in vanadium oxides, the value of transition temperature of the IMT strongly depends on the oxidation state of the vanadium ion, i.e., *TC* increases with oxidation states of the vanadium atom [22]. The resulting phase transitions in vanadium oxides are reversible and are accompanied by changes in structural, magnetic, optical and electrical properties. The structural change is a key factor since it leads to a displacement of the atoms in the crystal lattice and to a redistribution of electronic charges and, hence, the observed physico-chemical properties change [102].

In V2O5, the Fermi energy, which corresponds to the highest occupied electronic band, is calculated at 1.62 eV. This Fermi level is separated from higher conduction band by about 0.6 eV [103] while the valence band is at about 2.7 eV. Vanadium pentoxide has an indirect bandgap of 1.7 eV calculated by DFT (Density Functional Theory) but this value is smaller than the experimental one (2.3 eV). This underestimation of the bandgap is a common problem using the DFT. Concerning the indirect *Eg*, it is the result of the split-off of oxygen 2p band up to the vanadium 3d band [83]. In addition, there are two discrete intermediate bands (situated within the bandgap) with narrow bandwidth of about 0.7 eV [103], and they can make significant contribution to charge transport phenomena or mediate optical transitions from valence to conduction bands. Concerning the optical and electrical properties, V2O5 exhibits significant optical anisotropy in the visible region [22]. This anisotropy is due to the bands which are dispersive to various extents, and this is a clear indication of the crystal structure anisotropies. For instance, the electronic conduction within the *a*−*b* planes of V2O5 crystal lattice is considerably higher than that perpendicular to these planes [83]. The electron hopping (resulting from small concentration of V4+ in V2O5) is much easier within the *a–b* planes due to the shorter distances between vanadium atoms compared to the perpendicular direction to these planes. This leads to in-plane/out-of-plane conduction anisotropy [83]. In the form of films, the electrical conductivity of V2O5 remains low (from 10−<sup>2</sup> to 10−<sup>3</sup> <sup>S</sup>·cm−1) [104]. Apart from the anisotropic behavior observed in vanadium pentoxide, the presence of dispersive bands is also an indication of strong hybridization, since the more bands are dispersive, the stronger is the hybridization, and vice versa. In addition, in the semiconducting state, vanadium pentoxide exhibits high transmittance in near-UV and blue part of the visible spectrum, and low transmittance in near-IR and red part of the visible spectrum. The most known/popular property associated with V2O5 is its electrochromic effect, which can produce series of different colours of the material. Concerning the electrochemical properties of V2O5, the voltammograms of V2O5 films exhibit only one couple of redox peaks at around 0.70 V and −0.4 V, and these peaks are ascribed to the V5+/V4+ pair (the voltammograms were obtained at a scan rate of 0.030 V·s−<sup>1</sup> over a potential range of ±1 V). As the number of deposited layers increases, the produced current linearly increases as well as the optical modulation. Indeed, a voltage of −1 V is required to color the films while the color is subsequently bleached under a voltage of 1 V. The electrochromic transitions in V2O5 films have an excellent reversibility [105]. However, it is difficult to classify V2O5 in the cathodic or anodic electrochromism category. The cathodic electrochromism refers to a decrease in transmittance caused by charge injection (the case of Ti and W oxides), while anodic electrochromism refers to an increase in the transmittance (for instance Cr, Mn, and Fe oxides). Vanadium pentoxide is unique since it exhibits a transmittance increase in the near-UV and blue part of the visible spectrum and, simultaneously, a transmittance decrease in the near-IR and red part of the visible by charge injection [106]. The charge injection in V2O5 can be performed by intercalation into the van der Waals planes of small cation such as H+, Li+, and Na+, which cause electrochemical charge transfer reactions. The electrochromic properties of V2O5 have been developed for the design smart windows. On the other hand, vanadium pentoxide is also a thermochromic material with a transition at 530 K [107]. The resulting IMT is only observed when V2O5 is deposited in the form of thin films (nanostructure effect). However, the observed transition does not involve structural change. In the bulk phase, there is not IMT since vanadium pentoxide remains semiconducting at all temperatures [108].

In VO2, the metallic rutile and the monoclinic insulating phases have a Fermi level at almost the same energy level, i.e., 7.47 eV, and 7.44 eV, respectively [22]. Vanadium dioxide is transparent in the IR region above *Tc* and the optical transmittance changes by about two orders of magnitude during the transition [109]. The IR transmittance of VO2 exhibits hysteresis loops around the transition temperature: upon heating the optical switching occurs around 80 ◦C while upon cooling, it is at about 60 ◦C. Both phases of VO2 (metallic and insulating) absorb in the visible. Indeed, the material in insulating phase has a brown-yellow color, which turns into reddish-brown when deposited in the form of thin films, while the metallic state exhibits a greenish blue color [110]. The optical and electronic properties of VO2 can be explained by the band theory, where the number of band groups in the insulating phase is exactly the same as in the metallic phase, but in the latter, each band group is doubled. Therefore, in the rutile metallic structure, there is a finite overlapping of valence and conduction bands, which is a characteristic of metallic behavior [22]. In the monoclinic insulating structure, there is a gap of energy between the bands resulting from the oxygen 2p orbitals and the vanadium 3d orbitals. As a consequence, the observed energy band gap (*Eg*) in the rutile structure is 0.6 eV [111], while for the monoclinic structure, *Eg* is 2.5 eV. These properties of VO2 (M/R) can be also described (and, thus, better understood) using the crystal field theory [22]. In the metallic rutile phase (*T* > *TC*), vanadium (IV) is at the center of an oxygen octahedron. Therefore, in a symmetry point of view, the 3d orbitals are split into doubly-degenerated *eg* and triply-degenerated *t2g* levels. The *eg* orbitals are directed towards the 2p orbitals of the oxygen ligand and they form the *eg <sup>σ</sup>* bands in Figure 7. Regarding the monoclinic and rutile symmetries, the *t2g* orbitals, which are not directed toward the 2p orbitals of oxygens but between them, form the *eg <sup>π</sup>* and *a1g* bands (Figure 7). The *t2g* orbitals are also responsible for pairing between vanadium (IV) creating dimer resulting in the splitting of *a1g* into two bands. However, this dimerization effect in the metallic rutile phase is weak and pairing of vanadium (IV) does not happen. That is why the overlapping between *a1g* and *eg <sup>π</sup>* bands occurs at the Fermi level and it results in the metallic properties of rutile VO2 (Figure 7). In the insulating monoclinic phase, the distortion of the lattice results in stronger dimerization effect between vanadium (IV) pairs, thus splitting the *a1g* into two bands without overlap between *a1g* and *eg <sup>π</sup>* at Fermi level (Figure 7). In addition, in the monoclinic phase, the *eg <sup>π</sup>* band is destabilized by V-O overlapping and thus pushes this band upwards from the Fermi level [112,113].

**Figure 7.** Energy diagrams of monoclinic insulating and metallic rutile phases of VO2.

Concerning the optical properties, VO2 shows significant anisotropy in the infrared region [22]. Similar to V2O5, VO2 can also change its optical properties in a persistent and reversible way in response to a voltage [114]. Although VO2 is electrochromic, it is rather used as thermochormic materials with an IMT at 340 K [115], since V2O5 exhibits better electrochromic performance. In addition, the optical excitation through laser pulses is also an effective method (compared to excitation using heat gradient, i.e., by change of temperature) to control the phase transition in VO2(M/R) [116]. Indeed, laser pulses of sufficiently high energy could trigger a non-thermal sub-picosecond structural phase transition [117]. During this optical excitation, electrons are promoted from the VB to the CB of VO2. Subsequently, the collapse of the band gap occurs, thus leading to the transition into metallic phase. After depopulation of the *eg <sup>σ</sup>* bonding orbitals, the reversible transition occurs.

On the other hand, the strains present in VO2 is also a way to regulate the physical and chemical properties of the VO2(M/R) material. Indeed, deposited in the form of films, the source of strain is from the lattice and elastic mismatches between the VO2 material and the substrate. According to the relative size of the mismatch, the strain can be divided into tensile stress and compressive stress. For VO2 thin films, the induced strain can effectively adjust the IMT due to the presence of lattice mismatch between the sample and the substrate, thus causing a change of the transition point of VO2 IMT [118]. In addition to the different physical excitation detailed just above, the doping is also an effective method to affect the IMT in VO2(M/R) by either decreasing or increasing *TC* [117]. Doping is a typical chemical method to control the carrier concentration in semiconducting materials by insertion of discrete energy levels within the energy band gap. Depending on the nature of the doped elements (i.e., p- or n-type dopants), electrons injection (in the conduction band) or holes injection (in the valence band) in the host material is performed. In the case of W (VI) doping which has a much larger size than vanadium, injection of electrons in the parent materials occurs as well as an additional effect: the induction of lattice strain. This structural change plays an important role in the decrease of the phase-transition temperature [119]. Indeed, the insertion of W atoms not only breaks the lattice symmetry but also increases the spacing between the metal cations [120]. On the other hand, lowvalence metal ions such as Cr3+, Ga3+, and Al3+ are used to increase the temperature of the VO2 phase transition [121,122].

Another possibility to induce IMT transition is the use of an electric field through the Joule heating effect (from leakage/loss of current) [117]. However, it is important to rule out the electric and thermal effects from other processes. By applying an electric field transverse to the material's interface with a gate electrode, i.e., metal oxide semiconductor field effect transistor (MOSFET) (Figure 8), it is observed that the IMT is suppressed at high bias [73].

**Figure 8.** Scheme of a MOSFET using thermochromic VO2.

In V2O3, a striking IMT occurs at 150 K which is marked by a jump in the resistivity by seven orders of magnitude and an antiferromagnetic ordering of spins. V2O3 also shows electrochromism with a transition at about 160 K. The Fermi level of V2O3 is at 9.2 eV while the observed energy gap is about 0.66 eV. At temperature below 150 K, vanadium trioxide exists in its insulating phase with a monoclinic structure [72]. At a temperature above 150 K, the high-temperature paramagnetic V2O3 phase exists with a corundum crystal structure that exhibits metallic behavior. Therefore, at room temperature, there is a finite overlapping of valence and conduction bands, thus giving its metallic behavior of the corundum structure. Similar to VO2, the crystal field, which is generated by the oxygen on vanadium in the octahedra *[VO6]*, leads to the splitting of the d-orbitals into two sets of bands: *t2g* (lower band) and *eg <sup>σ</sup>* (upper band). In V2O3, the upper band is doubly degenerated and it consists in fact of *dxz* and *dyz* bands. Concerning the lower band, the noncubic arrangement of vanadium ions in the lattice combined with the trigonal distortion leads to the degeneration of *t2g* orbitals into *a1g* singlet (*d3z2*−*r2*) and doubly-degenerated *eg <sup>π</sup>* bands (*dxy* and *dx2*−*y2*) as shown in Figure <sup>9</sup> [123]. In the insulating monoclinic phase, the symmetry of the crystal field is lower than in corundum phase. Therefore, the *a1g* and *eg <sup>π</sup>* bands are split, thus creating an energy gap between *a1g* and *eg <sup>π</sup>* bands while, in the metallic corundum phase, the *a1g* and *eg <sup>π</sup>* bands overlap at the Fermi level (Figure 9) [124].

**Figure 9.** Energy diagrams of monoclinic insulating and metallic corundum phases of V2O3.

It is worth noting that the degree of anisotropy in binary vanadium oxides of single vanadium valence increases in the following order: V2O3 < VO2 < V2O5. Finally, vanadium monoxide (VO) also exists but it is usually a non-stoichiometric oxide. Vanadium monoxide crystallizes in cubic rock salt structure (NaCl) where all the *[VO6]* distorted octahedra are connected by the edges [125].

#### 3.1.3. Binary Vanadium Oxides with Mixed Valence States

The oxides of the general formula V*n*O(2*n*+1), i.e., the Wadsley phases, exist between V2O5 and VO2. This means that these phases have a mixture of vanadium (IV) and (V) because of oxygen vacancies. On the other hand, the Magnéli phases, i.e., oxides of the general formula V*n*O(2*<sup>n</sup>−***<sup>1</sup>**), point out the vanadium oxides which exist between VO2 and V2O3. This means that these phases have a mixture of vanadium (IV) and (III). The compounds of Magnéli and Wadsley series are intermediate oxides with vanadium in a mixed valence state formed from slabs of rutile-like structure sharing faces [126]. This is also known as the "shear structure" concept. These binary vanadium oxides have interesting electromagnetic properties due to strong electron correlation. It includes the insulator-to-metal transition (IMT), which is one of the most important parameters in strongly correlated electron systems [126]. In the Magnéli phases, i.e., oxides between V2O3 and VO2, the arrangement of atomic planes in the rutile VO2 structure is *[ABAB* ... *]*, where *A* = O and *B* = VO. By applying the shear operations which are (i) to periodically eliminate one *A*-plane every *n B*-planes, and (ii) to close the resulting gaps by operating a shear vector, the formula *n*VO + *(n*−1*)*O=V*n*O(2*n*+1) is obtained [126]. The structure of Magnéli phases includes *n* VO2 units between shear planes. Until now, six compounds (*n* = from 3 to 8) have been isolated. Almost all V*n*O*(*2*n*+1*)* compounds (*n* = 4, 5, 6, 8) show an IMT accompanied by sharp reduction of the magnetic susceptibility due to the formation of spin-singlet V4+–V4+ pairs in the low temperature insulating phase [126]. For *n* = 3 and 7, the insulating compound V3O5 has an antiferromagnetic state below *TN* = 75 K, while the metallic compound V7O13 has a rare antiferromagnetic groundstate [126]. On the other hand, in the Wadsley series which contain oxides between VO2 and V2O5, only two compounds, V3O7 and V6O13, have been obtained in both powder and single crystal [126]. In addition, the V*n*O(2*n*+1) compounds exhibit interesting electronic properties like V2O5, i.e., electron hopping [126].

The structure of V6O13 can be explained by the shear structure concept and this vanadium oxide undergoes insulator-to-metal transition at −123 ◦C. V6O13 contains a mixture of vanadium at a valence of +IV and +V, which is beneficial to the electronic conductivity. At room temperature, i.e., high temperature phase, V6O13 is in its metallic state with a layered structure belonging to space group *C2*/*m*. In the low temperature insulating state, the structure of V6O13 belongs to space group *P21*/*a*. In addition to its IMT, V6O13 also exhibits an antiferromagnetic transition [126]. As mentioned above, the IMT in V6O13 is accompanied by a sharp reduction of magnetic susceptibility, which highlights the formation of spin-singlet V4+–V4+ pairs in the insulating phase. V6O13, as most of the Wadsley compounds (including V2O5, V6O13, V3O7, and VO2(B)), are promising materials for cathodes in rechargeable Li-ion batteries (LIBs) mainly due to their stable layered structure but also to their high specific capacity [126,127]. Furthermore, V6O13 has additional advantages such as higher rate capability, better reversible capacity, smaller volume expansion and safer lithiation potential. Indeed, at room temperature, V6O13 could manage high rate of charge and discharge because of the metallic character of this oxide at ambient temperatures. In term of molar ratio of inserted lithium ions into the layered material, the theoretical specific capacity of the cathode material is 8 mol Li+ in 1 mol V6O13. However, the preparation of V6O13 is complicated to achieve owing to the valence fluctuation of vanadium atoms, i.e., it is hard to control the ratio between +IV and +V valence states [77].

Concerning V3O7, it is also an intermediate phase between VO2 and V2O5 which is often reported in the literature as hydrate compound (V3O7·H2O) at room temperature, i.e., in its insulating phase [128]. V3O7 is a mixed valence oxide with a ratio of V4+/V5+ = 0.5. This vanadium oxide has a bandgap of 2.5 eV at room temperature, and it is rather considered as a semiconductor as most of the vanadium oxides in their insulating phase [129]. The structure of crystalline V3O7·H2O is orthorhombic with a space group *Pnam* [127,128]. This phase has a layered structure similar to orthorhombic V2O5 since V3O7·H2O is mostly prepared from vanadium (V) oxide [127]. V3O7·H2O is composed of layers of *[VO6]* octahedra and *[VO5]* trigonal bipyramids sharing corners [126,128,129]. In the structure of V3O7·H2O, the water molecule is bound to vanadium by replacing one of the oxygen atoms in the *[VO6]* octahedron, thus forming hydrogen bonds with the octahedron in the next layer which gives a 3D structure [128]. Like V6O13, V3O7·H2O is also a promising cathode material in LIBs since these materials have high initial capacity which is closely related to their layered structure together with the presence of mixed valence states of vanadium ions [127]. In addition, V3O7·H2O exhibits excellent electrochromic performance because of optical modulation in both anodic and cathodic reactions, thus making V3O7·H2O an excellent candidation cathode material for electrochromic applications [129]. For instance, in the form of films, V3O7·H2O can switch reversibly between a reduced blue state (at −1.2 V) and an oxidized orange state (at +1.9 V) [130]. These reversible changes are accompanied by changes of reflectance, especially at 590 nm. Therefore, similar to V2O5, V3O7·H2O is characterized by a typical muli-electrochromism but with faster kinetics. The fast kinetics observed in the V3O7·H2O film can be explained by its unique structure combined with mixed valence states of V (V) and V (IV). That contributes to higher electronic conductivity and more active redox sites during the reduction and oxidation processes [130]. On the other hand, non-hydrated V3O7 has received less attention than its hydrate form. This fact is mainly due to its metastable behavior. Indeed, the non-hydrated phase can be formed by heat treatment of V3O7·H2O which leads to removal of the water molecule starting at 75 ◦C and finishing at about 305 ◦C. However, the formed V3O7 is metastable and plays the role of growing seeds for the formation of V2O5. In orther words, although V3O7·H2O

can be prepared from V2O5, its heating in air gives rise to both dehydration and oxidation into V2O5. It is possible to synthesize directly V3O7 (with a stable structure) by using spray pyrolysis technique on substrate heated at 400 ◦C [129]. The crystalline phase of V3O7 has a monoclinic structure belonging to space group *Pmmn* [129]. In addition, V3O7 exhibits higher intercalation capability, fast switching intercalation/extraction rate and better reversibility than V2O5 [129].

Another interesting phase in the Wadsley series is V4O9. The synthesis of this V*n*O(2*n*+1) with *n* = 4 is a real challenge since it has a formal valence ratio V4+/V5+ = 1. In addition, V4O9 is a special compound in the Wadsley series in the sense that its structure and properties are still not clear (although its existence has been recognized), thus, this phase is often called "the missing link" in binary vanadium oxides [126]. Usually, V4O9 is synthesized by a solid state reaction in a reducing atmosphere of an appropriate mixture of binary vanadium oxides with single valence state: V2O5 and V2O3 (or VO2). As reducing agents, carbon and SO2 are the most used and the synthesis is carried out under ultra-high vacuum. However, the best method to prepare single phase V4O9 is by using sulfur (S) as a reducing agent (Equation (10)), but an excess of S leads to the formation of VO2(B) phase [126]. The structure of V4O9 is not a plane fault type but a tunnel defect type. The crystal system of V4O9 is orthorhombic and belongs to the space group *Cmcm*. Since V4O9 is prepared from V2O5 reduction where an excess of reductant gives VO2(B), this suggests that V4O9 has a layered structure similar to those from V2O5 and VO2(B).

$$2\text{V}\_2\text{O}\_5 + \text{S} \to 2\text{V}\_4\text{O}\_9 + \text{SO}\_2\tag{10}$$

To better explain the structure of V4O9, the shear structure concept can be used. The lattice of V2O5 is formed of layers of *[V2O3]* (called "*a*") which are stacked and sharing *[O2]* units in the *c*-direction (layers called "*b*"), thus V2O5 has a layered structure of *[VO5]* pyramids-up/pyramids-down along the *ab*-plane (Section 3.1.1). Hence, the structure of V2O5 has the stacking of layers as *[ababa* ... *]* in the c-direction (or parallel to the *ab*plane). On the other hand, the crystal structure of VO2(B) has a *[abaaba* ... *]* stacking which can be regarded as a daughter structure of V2O5. Indeed, this stacking manner of *[abaaba* ... *]* is obtained from V2O5 stacking by using the following shear operation: to periodically eliminate one *b*-layer every two *a*-layers, so the arrangement is *[abaxabaxa* ... *]* (*x* = gap) where the gaps are closed by connecting each *[aba]*-block, thus giving the stacking manner of *[abaabaa* ... *]*. For V4O9, it is supposed that its stacking structure has a cleavage plane. This suggests that the *ab*-plane of V4O9 is a cleavage plane with a lattice similar to that of V2O5 or VO2(B). If the structure of V4O9 were a shear structure derived from V2O5, it would have a stacking manner of *[abababaabababaab* ... *]*. This stacking manner is derived by two operations: the periodic elimination of one *b*-layer every four *a*-layers, so the arrangement is *[abababaxabababaxab* ... *]* with closing the gaps and thus obtaining the composition of 8V2O3 + 6O2 = 4V4O9. However, the observed *c*-parameter is much shorter than the calculated value for the stacking periodicity of *[abababaabababaab* ... *]*. To obtain the composition of V4O9, half of oxygen atoms should be removed from half of *b*-layers in this stacking periodicity. This model has stacking periodicity of *[abab'abaab'a* ... *]* in which half of oxygen atoms in *b'*-layers are deficient in an ordered manner to form tunnel defects. Hence, the composition is 4V2O3 (*a*-layers) + 2O2 (*b*-layers) + 2O (*b'*-layers) = 2V4O9 [126]. In other words, the obtained structure is composed of *[VO5]* pyramids, *[VO6]* distorted octahedra and *[VO4]* tetrahedra that is very similar to the structure of vanadyl pyrophosphate (VO)2P2O7 (by replacing *[PO4]* tetrahedra by *[VO4]* tetrahedra). The similarities of V4O9 to (VO)2P2O7 are even lager, especially in magnetic properties. Indeed, the excitation gap between the singlet groundstate and the excited triplet state is approximately 73 K. Therefore, the resemblance of magnetic properties between V4O9 and (VO)2P2O7 strongly supports the reliability of the structure [126].

#### *3.2. Case of Ternary Vanadium Oxides*

Ternary vanadium oxides are interesting materials since their properties can be tuned according to the third element composing them. In terms of modifying the photoinduced properties of materials, it was observed that metal impurities in homometallic oxide can modify the energy bandgap by the overlapping of outer s- and d-orbitals of the p- and d-block metals ions, respectively, thus giving rise to smaller energy band gap (*Eg*) [131]. Therefore, it was assumed that p- and d-block metals could be used to design single phase heterometallic and nowadays, many visible light-responsive photoactive materials are heterometallic oxides, e.g., ternary and quaternary oxides [131]. Concerning the oxides based on vanadium, one promising class of material is the metal orthovanadates which are composed of VO4 <sup>3</sup><sup>−</sup> tetrahedral. Several orthovanadates including BiVO4, FeVO4, Mn3(VO4)2, CeVO4, NdVO4, and GdVO4 have been widely studied during the last decades due to their promising properties and ease of preparation. Other metal vanadates such as, for instance, Cu2V2O7, Cu3V2O8, Co3V2O8, FeV3O8, Ag4V2O7 have received also a strong interest from the scientific community. Indeed, these vanadium based oxides can be used in (photo)catalysis, lithium-ion batteries, solar cells, gas sensors, water splitting technologies, optoelectronics, etc. [132–134]. Such an interest for transition metal vanadates is due to their interesting optical, electric and magnetic properties, thus being useful in the above mentioned applications [135]. Although most of the ternary vanadium oxides are not thermochromic and/or electrochromic as the binary vanadium oxides, these transition metal vanadates have attracted considerable attention for environmental applications since they are semiconductors which can be activated under UV or visible light radiations. Therefore, ternary vanadium oxide based materials can trigger photoinduced chemical reactions such as photocatalysis and photo-Fenton process for contaminant removal or photosplitting of water (for O2 and/or H2 production) as discussed in Section 4 [135].

#### 3.2.1. Bismuth Vanadates

Bismuth vanadates are ternary oxides which exhibit a wide variety of structures [136]. The most promising one is BiVO4. This oxide is a n-type semiconductor that exists in three polymorphs: monoclinic scheelite-like, tetragonal scheelite-like, and tetragonal zircon-like structures with an energy band gap (*Eg*) of 2.40, 2.34, and 2.90 eV, respectively [137–139]. All these crystal structures are built up with *[VO4]* tetrahedra and *[BiO8]* polyhedra where V (V) and Bi (III) are in the center of these units. In both scheelite phases, each *[BiO8]* is surrounded by eight vanadium units, whereas in zircon structure, Bi units are surrounded by only six *[VO4]* [139]. Comparing the scheelite structures, the difference between monoclinic and tetragonal phases is based on the local environment, where vanadium and bismuth units are more distorted in the monoclinic structure [139]. In addition, a reversible phase transition between monoclinic scheelite-like (low temperature phase) and tetragonal scheelite-like (high temperature phase) structures occurs at 255 ◦C [138,139]. However, between tetragonal zircon-like and tetragonal scheelite-like structure, an irreversible phase transition is observed at about 400–500 ◦C [138,139]. Therefore, at room temperature, there are two BiVO4 polymorphs which are stable. Among them, the monoclinic scheelite-like structure exhibits the best photoactivity under solar light, e.g., longer lifetime of photogenerated charge carriers and higher photocurrent. This is due to better photon absorption and charge carrier transport resulting from smaller *Eg* and distorted *[VO4]* and *[BiO8]* units, respectively [138,139].

The crystalline structure and the energy band gap of BiVO4 have an influence on the efficiency of photo-induced processes triggered by this semiconductor, e.g., photocatalytic reactions. However, the exposed crystal facets are maybe more important since they can affect the thermodynamic and the kinetic of a chemical reaction [137]. By designing a particular crystal, the crucial parameters of photo-induced chemical reactions in heterogeneous phase can be tuned such as the preferential adsorption of a reactant, the surface transfer of photo-generated charge carriers and the desorption of products [137,140]. In the case of bismuth vanadate where the monoclinic structure is the most photoactive one, the *{010}*

and *{110}* crystal facets has redox functions since they provide reduction and oxidation sites, respectively, in which photogenerated electrons and holes are available [141,142]. The effect of these crystal facets on the efficiency of BiVO4 in photo-induced processes is often investigated by photoelectrochemical (PEC) and transient absorption spectroscopy (TAS) analyses [143]. PEC measurements could explain the charge carriers' behavior at the surface of BiVO4, especially the surface recombination which could be evaluated by amperometry, i.e., by analyzing the shape of photogenerated current curve at potential near the onset of steady state photocurrent (larger is the current "spike", higher is the recombination). TAS is another powerful tool (which is detailed hereafter) to understand the charge carriers' dynamics including relaxation rate, recombination rate, and mechanism of recombination [143,144]. The operation of this fast spectroscopy is based on the excitation of the sample by a pulse followed by the optical measurement (using a spectrophotometer) and the analysis of spectral differences between excited and the ground state levels of the sample as a function of time [144]. Since the delay between excitation and measurement sets the time scales that can be studied, TAS is considered as a fast spectroscopic technique [144]. The advantage of TAS studies resolved in time (compared to PEC studies) is the possibility to investigate the dynamics of photogenerated holes. Indeed, TAS allows the observation of these dynamics where *h<sup>+</sup>* is responsible, for instance, to water oxidation or generation of hydroxyl radicals while in PEC analysis, the hole dynamics is indirectly deduced from the current signal which is monitored via external circuit. In the future, studies based on transient absorption spectroscopy will surely developed since currently, TAS studies of transition metal oxides are limited [143,144]. To date, the dynamics of photogenerated holes in BiVO4 was studied using transient absorption spectroscopy coupled to photoelectrochemical measurement. In such a study (in aqueous medium), the constant of decay time of spectroscopic signal of *h<sup>+</sup>* is dependent to the applied bias due to kinetic competition between water oxidation and charge carrier recombination. In addition, the time constant for charge carriers' recombination (measured with transient absorption spectroscopy) is different to the time constant of back electron transfer (measured with transient photocurrent measurements using chopped light), thus indicating two distinct recombination processes which limit the photocurrent generation in BiVO4 photoanodes. One recombination process is fast (occurring within microseconds) and assigned to direct bulk electron/hole recombination, while the other one is slow (occurring within milliseconds) and identified as recombination of surface-accumulated holes with bulk electrons. The latter process, also called back electron transfer, can be avoided at strong anodic bias which provides a large energetic barrier to prevent this recombination process [143]. With regard to TAS and PEC analyses of BiVO4 photoanaode, a model of carrier dynamics can be schematized (Figure 10). After the electron/hole separation under simulated solar light, hole trapping process in energy levels closed to VB occurs at 5 picoseconds. At 40 ps and 2.5 ns, electron relaxation and trapping, respectively, are in competition with direct recombination with trapped holes. The relaxation and trapping processes are expected to be dominant over recombination. After 10 ns to 10 μs, recombination of trapped charge carriers can occur (Figure 10). In addition, at such time scales, no spectral or kinetic differences are observed between systems under a bias voltage or not, thus indicating that the effects of bias voltage occur on longer time scales [144].

**Figure 10.** Scheme of charge carriers dynamics within the first stages of BiVO4 photoactivation.

The main advantages of BiVO4 are its non-toxicity and its energy band gap located in the visible region. However, several drawbacks are responsible of the low efficiency of bismuth vanadate in photo-induced chemical reactions such as its poor electron mobility and high *e*−*/h*<sup>+</sup> recombination rate [137,138]. In addition, BiVO4 has a short hole diffusion length (70−100 nm) that compromises the optimization of film thickness and light harvesting according to the optical penetration depth principle [138]. All these disadvantages are mainly due to the structure of BiVO4 and the vanadium 3d orbitals which constitute the conduction band, since *[VO4]* tetrahedra are not connected to each other and the V 3d orbitals are strongly localized, respectively [138].

On the other hand, there are many other bismuth vanadates in the ternary system Bi–V–O. These bismuth vanadates exhibit different crystalline structures, from which the most known are the Aurivillius-type Bi4V2O11 phase (and also, its reduced form: Bi4V2O10) [145–149]. Beside Bi4V2O11, there are also hollandite-type Bi8V2O17 structure, Bi3.33(VO4)2O2, Bi2VO5.5, etc. [136]. However, Bi4V2O11 and Bi3.33(VO4)2O2 compounds are the most studied materials in photo-induced processes.

Concerning Bi3.33(VO4)2O2, it crystallizes in the triclinic system (space group *P*1) with the following lattice parameters: a = 7.114 Å, b = 7.844 Å, c = 9.372 Å, α = 106.090◦, β = 94.468◦, and γ = 112.506◦ [136]. This crystal structure could be described by the alternative stacking along the b-axis of two types of layer (layers I and II). The layer I has four distinct Bi sites and *[VO4]* tetrahedra while the layer II has three distinct Bi sites and *[VO4]* tetrahedra. It is worth noting the deformation of coordination environments for all Bi atoms (due to lone pair of Bi3+), but the *[VO4]* tetrahedra remain regular. In the two types of tetrahedral (in Bi3.33(VO4)2O2), the mean V–O distance is 1.715 Å which is slightly shorter than 1.74 Å for tetrahedra in BiVO4. About the four Bi atoms in layer I, the coordination number is 8 while for the three Bi atoms in layer II, it is 6. Therefore, the mean interatomic distances Bi–O is between 2.39 Å and 2.61 Å and it is consistent to those in BiVO4 (2.48 Å). Bi3.33(VO4)2O2 has a yellow color and its optical band gap is estimated at 2.36 eV, which is similar to the monoclinic scheelite BiVO4 [136].

Concerning Bi4V2O11, it is a material which was discovered in the 1980s. It is defined as an intermediate compound in the system Bi2O3-V2O5. This bismuth vanadate material is particularly interesting due to its physical properties, e.g., dielectric properties, ferroelectricity and pyroelectricity at room temperature. The crystalline structure of Bi4V2O11 is the parent structure of the *BIMEVOX* family materials (*BI* = bismuth, *ME* = metal subtitution, *VOX* = vanadates) which have attractive conduction properties at moderate temperature [146]. Bi4V2O11 has also attracted interest due to its photo-induced properties for organic pollutants removal in water (Section Bismuth Vanadates in Section 4.2.3) [148].

Bi4V2O11 exhibits two reversible phase transitions at 430 ◦C and at 570 ◦C while the melting temperature is 870 ◦C [145,146,148]. Therefore, there are three phases of Bi4V2O11 which are the low-temperature α-form, the intermediate β-form, and the high-temperature γ-form which crystallize in the monoclinic, orthorhombic and tetragonal system, respectively [147]. The tetragonal gamma phase belongs to the *I4/mmm* space group while the orthorhombic beta polymorph has *Amam* space group [146]. For the alpha Bi4V2O11, it is more complicated to assign a space group since its monoclinic structure is a three-fold supercell [147]. However, the sructure was refined recently in the *A2* space group [147]. All these three modifications can be described from a mean orthorhombic subcell having am ≈ 5.53 Å, bm ≈ 5.61 Å and cm ≈ 15.29 Å (c being the stacking direction). The relationships between the unit cell parameters of the three phases and those of the subcell are: aγ = b<sup>γ</sup> = am/(2)1/2, c<sup>γ</sup> = cm; a<sup>β</sup> = 2am, b<sup>β</sup> = bm, c<sup>β</sup> = cm; a<sup>α</sup> = 6am, b<sup>α</sup> = bm and c<sup>α</sup> = cm [146]. From a structural point of view, the structure of the three Bi4V2O11 phases is built from *[Bi*2*O*2*]* <sup>2</sup>*<sup>+</sup>* layers separated by vanadium–oxygen deficient perovskite slabs *[VO*3.5*]* [145]. Therefore, Bi4V2O11 has the general formula *[Bi*2*O*2*][Am*−1*BmO*3*m+*1*]* consisting of an intergrowth between *[Bi*2*O*2*]* <sup>2</sup>*<sup>+</sup>* sheets and *[Am*−1*BmO*3*m*+1*]* <sup>2</sup><sup>−</sup> perovskite-like slabs with *m* as the number of *[BO*6*]* octahedra stacked along the direction perpendicular to the sheets. Therefore, Bi4V2O11 (or Bi2VO5.5) is the corresponding oxygen deficient structure with *m* = 1 i.e., *[Bi*2*O*2*][VO*3.5܆0.5") *[*- "symbolizes the oxygen deficiency) [146]. A deepened comparison between the polymorphs, based on their structure and their conductivity, can be also discussed. Structurally, the alpha structure mainly differs from the two high temperature polymorphs by a specific vanadium-oxygen surroundings. Indeed, the vanadium-oxygen slabs exhibit vanadium with three different oxygen environments [147]. About conductive properties, the gamma polymorph has the highest conductivity compared to the two other phases, i.e., 0.2 S·cm−<sup>1</sup> is observed at 600 ◦C with a strong drop during the phase transition from gamma to beta and beta to alpha [145]. The conductivity of beta and alpha phases is 0.01 S·cm−<sup>1</sup> at 500 ◦C and 10–5 S·cm−<sup>1</sup> at 300 ◦C, respectively. The observed drop in electrical performance is obviously due to the crystal structure, and particularly to the ordering of the oxygen vacancies in the V-O slabs [145]. These slabs are the components that mainly differ (strucrurally and electrically) from one phase to another [145]. For example, in the gamma phase, there is a large disorder of the oxygen atoms in V-O slabs which leads to fast diffusion properties. In addition to the electrical conductivity, Bi4V2O11 has interesting optical properties with an energy band gap of 2.15 eV. This low *Eg* makes Bi4V2O11 a promising visible-light-active material in photocatalysis. In addition, since Bi4V2O11 has a layered structure (Aurivillius-related phase), the separation of photogenerated carriers is favored in the stacking direction, thus improving its photocatalytic efficiency. However, within the layers, the photogenerated electrons and holes are easily recombined because of the short band gap (2.15 eV) [149]. Another characteristic of Bi4V2O11 is its ability to be easily reduced and re-oxidized as, for example, through Bi4V2O10. In this reduced phase, there are also high and low temperature polymorphs where the structure of the low temperature one (α-Bi4V2O10) is similar to α-Bi4V2O11: *[Bi*2*O*2*]* and *[VO*3*]* layers are identified but the latter one, *[VO*3*]*, is built up by *[VO*5*]* square pyramids sharing basal corners.

#### 3.2.2. Iron Vanadates

In the last decade, many scientists have focused on iron vanadates due to their versatile properties (photocatalysis, Fenton-based process, etc.). Among these ternary oxides, FeVO4, FeV3O8, Fe2VO4, FeV2O4, Fe2V4O13, Fe4(VO4)4 have been successfully synthesized using different routes [134]. Usually, vanadates exhibit higher melting point and lower decomposition rate (at high temperatures) than V2O5. This is particularly the case of iron vanadates where the incorporation of Fe leads to stabilization of the material, thus decreasing the volatility of V. The most studied iron vanadate is FeVO4 (also called iron orthovanadate) which has a stoichiometric ratio V:Fe = 1. FeVO4 is also paramagnetic at room temperature with a saturation magnetization of 0.15 emu·g−<sup>1</sup> [132]. This stoichiometric iron vanadate exhibits four different phases usually labeled as FeVO4-I, FeVO4-II, FeVO4-III, and FeVO4-IV. The FeVO4-I polymorph is the only stable phase formed at room temperature while other ones are all metastable phases formed at high pressures and/or high temperatures. Pure FeVO4-I crystallizes in triclinic system with space group *P*1 [133]. Structurally, this polymorph is built from distorted *[FeO*6*]* octahedra and distorted *[FeO*5*]* trigonal bipyramids linked together according the following sequence: two octahedra and one trigonal bypiramid. These Fe–O polyhedra are arranged in doubly-bent chains which are joined together by *[VO*4*]* tetrahedral in order to form a 3D framework. The other forms of iron orthovanadates, i.e., FeVO4-II, -III, and -IV, crystallize in the orthorhombic CrVO4, orthorhombic α-PbO2 and monoclinic wolframite NiWO4 structures, respectively [150,151]. It is worth noting that hydrated FeVO4 is amorphous, but after prolonged reaction time at normal pressure and low temperature, crystalline monoclinic FeVO4·*x*H2O can be obtained. Concerning the physico-chemical properties of crystalline FeVO4-I, it is a stable n-type semiconductor with a narrow energy band gap (*Eg*). The *Eg* has been estimated between 2.70 (direct) and 2.05 (indirect) eV depending on the synthesis parameters and the size and morphology of the prepared material [150–154]. The energetic position of the conduction band is above zero and more precisely at −0.4 eV [155]. This means that FeVO4 semiconductor can theoretically produce hydrogen (*E*<sup>0</sup>*(H+/H2)* = 0 V) under appropriate irradiation (*hν* > *Eg*), and the theoretical STH (solar-to-hydrogen) efficiency is relatively high (16%). Concerning the photophysical properties, the photocurrent density in undoped FeVO4 is low (less than 0.1 mA·cm−2) at 1.23 V vs. RHE (reversible hydrogen electrode) [154]. However, to improve the photophysical properties of FeVO4, synthesis of nanostructured materials and their modification using dopants can increase the specific surface area and improve the electronic properties (by modifying the band structure of FeVO4), respectively [156]. In the case of Mo-doped FeVO4, the photocurrent under simulated solar irradiation increases with increasing the dopant concentration with a maximum value of 0.2 mA·cm−<sup>2</sup> for 2% of Mo at 1.6 V vs. RHE [154]. The produced photocurrent is improved of about 90% compared to undoped FeVO4. At higher percentage of Mo, the photocurrent decreases since the presence of high amount of intrinsic defects act as recombination centers. Regarding the IPCE, 2% Mo-doped FeVO4 exhibits an efficiency of 7% at 400 nm, while it is only 1% for undoped FeVO4 at the same experimental conditions. Since the IPCE remains low at longer wavelength, the direct band gap appears to contribute more to the photocurrent than the indirect *Eg* [154]. In order to enhance the electronic properties without the use of dopants, quaternary oxides can be designed. Indeed, bismuth-substituted iron vanadate (Fe1*-x*Bi*x*VO4) has improved charge carrier mobility and exhibits a shift of *Eg* towards higher energies with an increase of Bi content [154]. However, during the synthesis of either pristine or modified iron vanadates, particular attention is required to limit the formation of secondary phases like Fe2V4O13 [154]. Concerning metastable FeVO4 phases, few works are reported but FeVO4-II has shown better properties than stable FeVO4-I in gas sensing and photocatalytic activity. Indeed, metastable phases usually exhibit unique properties which are better than the pure phase. Such a fact is also observed in the case of vanadium oxides with metasable VO2(B) which has better photo-induced properties than stable VO2(M), especially in the degradation of organic pollutants. Finally, although there are numerous PEC studies of FeVO4, TAS has not been performed yet, in contrary to one of the most popular vanadate, BiVO4 (Section 3.2.1).

Another widely studied iron vanadate is Fe2V4O13. This compound has a monoclinic structure which consists of tetrahedrally coordinated *[VO4]* and octahedrally coordinated *[Fe2O10]* units, where tetrahedra are linked to octahedra through edge sharing, thus forming an unusual horseshoe-like chain structure [157,158]. It is quite difficult to prepare pure Fe2V4O13 since FeVO4 is always observed as an impurity. In addition, during the crystallization of Fe2V4O13 (from the amorphous material), the specific surface area of Fe2V4O13 calculated by BET was found to significantly decrease from 52 to 22 m2·g−<sup>1</sup> while the color of the material changed from brown to light green [159]. In the structure of Fe2V4O13, there is the presence of holes and channels which are interesting for many applications. Indeed, this iron vanadate has been investigated as an efficient cathode for lithium ion batteries as

well as catalyst for either photooxidation of organic pollutants or photoreduction of CO2. For LIB, Fe2V4O13 can deliver a discharge capacity of about 154 mA·h·g−<sup>1</sup> between 1.5 and 4.0 V after 50 cycles [157]. However, under irradiation, its photocurrent density is about <sup>16</sup> <sup>μ</sup>A·cm−<sup>2</sup> at 1.23 V vs. RHE in sodium borate electrolyte. This low value is due to the poor charge transport of this iron vanadate [158]. As a consequence, the IPCE is only 1.5% under 1.4 V at 350 nm [158]. Fe2V4O13 is n-type semiconductor with an optical band gap of 2.3 eV with an absorption edge extending toward the visible light region. The conduction band minimum (CBM) of Fe2V4O13 is determined to be −0.55 eV (vs. RHE) [160] while the valence band maximum (VBM) is located at 1.69 eV.

A less known iron vanadate is FeV2O4. This spinel compound crystallizes in cubic face-centered structure and belongs to the space group *Fd*3*m*. The lattice parameters are as the following: a = 8.32 Å, b = 8.32 Å, c = 8.32 Å and α = β = γ = 90◦ with a cell volume of 588.77 Å<sup>3</sup> [161]. The iron cations are located in eight tetrahedral sites *[FeO4]* while vanadium ions are situated in 16 octahedral *[VO6]* (the cubic-closed packing structure is based on oxygen anions) [161]. FeV2O4 is a semiconductor which has a yellow color and it exhibits a direct energy band gap at about 1.9 eV and an indirect one at approximately 2.6 eV [162]. Comparing to iron (III) oxide (Fe2O3) and vanadium (V) oxide (V2O5), FeV2O4 is a material exhibiting much better electronic properties, i.e., with effective separation of photo-generated charge carriers, thus leading to better and faster charge transfer process at the interface with the liquid or gaseous medium [162]. In addition, the charge transfer resistance (RCT) is 23 KΩ which is a relatively low value. Therefore, FeV2O4 is considered as a suitable photo-electrode material for PEC cell [162]. The photocurrent of this iron vanadate under UV-visible illumination (100 mW·cm<sup>−</sup>2) can reach 0.18 mA·cm−<sup>2</sup> at 1.2 V with an IPCE of 22% at 320 nm (in sulfate electrolyte) [162]. Finally, another underrated iron vanadate is FeV3O8. It has a monoclinic structure with lattice parameters of a = 12.13 Å, b = 3.679 Å and c = 6.547 Å and a direct optical band gap of 2.23 eV, thus suggesting a potential and promising photoactivity under visible light [163].

#### 3.2.3. Copper Vanadates

Different polymorphs of copper vanadate (CVO) have been studied including CuV2O6, Cu2V2O7, Cu3V2O8, etc. The preparation of different types of copper vanadate with different stoichiometries can be achieved using different synthesis techniques and different experimental conditions [164,165]. The value of stoichiometric ratio strongly influences both physical and chemical properties. This stoichiometry is often determined by the ratio between copper (II) and vanadium (V) oxides, since the crystal structures of many copper vanadates are considered to be composed of CuO and V2O5. For example, CuV2O6, Cu2V2O7, and Cu3V2O8 exhibit molar ratio of CuO and V2O5 of 1:1, 2:1, and 3:1, respectively, and other combinations of copper oxide with V2O5 can lead to various phases of copper vanadate. In addition to the pure phase, surface modifications can be performed using metals and nonmetals, thus improving electrical, optical, and structural properties of the materials by enhancement of charge carrier transport, tuning of the band gap energy, etc. Initially, copper vanadates attracted the interest of scientists due to their high lithiated characteristics and their potential application as cathode in LIB [164]. This is particularly the case of Cu3V2O8 and its composites and doped forms [164]. Crystalline copper vanadates have interesting layered structures which are composed of Cu-O layers in octahedral coordination joined by V-O tetrahedral layers. In addition, copper vanadates are also semidonductors with interesting optical properties, i.e., the absorption of a large part of natural solar light [164].

Cu3V2O8 is one of the most investigated CVO. It can crystallized in either monoclinic, triclinic, or orthorhombic system, all with a low symmetry. Usually, most of the CVO structures offers a large range of phases [164]. The monoclinic gamma phase (γ-Cu3V2O8) consists of *[VO4]* tetrahedra and two distinct Cu units: square-planar *[CuO4]* and squarepyramidal *[CuO5]*. It is the most stable phase of this type of copper vanadate. The oxidation states for copper, vanadium and oxygen in Cu3V2O8 are +2, +5 and −2, respectively, thus

the electronic band structure is mainly composed of Cu 3d9, V 3d0, and O 2p6 orbitals [166]. Indeed, γ-Cu3V2O8 is a n-type semiconductor with an indirect band gap of about 1.8 eV and a direct band gap at approximately 2.7 eV [135]. The predicted *Eg* transitions are from O 2p to Cu 3d orbitals for direct *Eg* and from O 2p to V 3d orbitals for indirect energy band gap (Figure 11) [166]. It is worth noting that the valence band exhibits a high degree of hybridization with strong involvement of O 2p orbitals. Furthermore, the Fermi level (*EF*) is located at 4.57 eV while VBM and CBM lay at 6.03 and 4.23 eV, respectively [166]. These values are relative to the vacuum level (e.g., *E*0*(H+/H2)* is situated at 4.5 eV). Regarding these properties, γ-Cu3V2O8 is attractive for photochemical and electrochemical applications such as lithium ion batteries and photocatalysis [164,165].

**Figure 11.** The electronic band structure of γ-Cu3V2O8 (related to the vacuum energy).

Concerning the photoelectrochemical properties of γ-Cu3V2O8, the photocurrent density is measured at 62 <sup>μ</sup>A·cm−<sup>2</sup> at 1.23 V vs. RHE under simulated solar irradiation. In addition, the IPCE of γ-Cu3V2O8 reaches only 3% at 360 nm [166]. Due to a relatively low energy band gap, there is a fast recombination between the photogenerated charge carriers, thus leading to the weak photoactivity of γ-Cu3V2O8. The recombination process, which dominates this copper vanadate, is worsened by short diffusion length. Therefore, although the energy band gap and the physico-chemical stability of γ-Cu3V2O8 appear promising, the performance of this material as a photoelectrode in PEC applications is not yet suitable. Further work is required to improve the properties and performance of γ-Cu3V2O8 in order to design a next-generation photoanode [166].

Another famous copper vanadate which is also an n-type semiconductor is Cu2V2O7. Indeed, it has been already used as a cathode material for rechargeable LIB, photoelectrode in PEC technologies (for water oxidation) and as a photocatlyst [167]. Cu2V2O7 has a monoclinic crystal structure where V5+ ion is surrounded by four oxygen while Cu2+ ion is coordinated with five oxygens [168]. This semiconductor copper vanadate is also a promising candidate for visible light-driven photocatalysis because of its narrow energy band-gap (2.0 eV) which is an indirect *Eg* [167,168]. The photocurrent density achieved by Cu2V2O7 reaches about 120 <sup>μ</sup>A·cm−<sup>2</sup> at 1.58 V vs. RHE while the IPCE is 4% at 400 nm under 1.23 V vs. RHE [168]. However, as in the case of Cu3V2O8, the photocatalytic efficiency of Cu2V2O7 is limited because of the fast recombination of photogenerated holes and electrons. Therefore, composites with a Z-scheme structure have been designed in order to prevent this recombination i.e., to efficiently separate the *e*−*/h+* pairs in the composite [167]. Indeed, a direct Z-scheme system is similar to a type-II heterojunction (Figure 12), but the electrons of the holes of semiconductor No1 (SC1) recombine by physical

contact with the electrons of semiconductor No2 (SC2), thus the remaining photogenerated electrons and holes of SC1 and SC2, respectively, are spatially separated.

**Figure 12.** Electronic band structure of type-II heterojunction and direct Z-scheme with the different charge carriers migrations.

#### 3.2.4. Silver Vanadates

Among the ternary vanadium oxides that are investigated for photochemical applications, silver vanadates have attracted much attention, mainly because of their narrow band gap and highly dispersed and hybridized valence band. The valence band in silver vanadates is composed of 2p orbitals of oxygen and 4d orbitals of silver and the hydbridization between the 2p6 orbitals with completely filled 4d<sup>10</sup> orbitals leads to a VB located at a more positive energy level than VB in most of the other ternary vanadium oxides. That feature increases the oxidation power of the material. On the other hand, the conduction band consists of hybridized 5s orbital of silver with 3d orbitals of vanadium. Therefore, silver vanadates exhibit good absorption of low energy photon (e.g., visible light) and high mobility of photogenerated holes, which are crucial characteristics for application in photochemical processes, especially the degradation of organic pollutants (Section 4.2.2) [169].

A variety of ternary oxides based on silver, like AgSbO3, AgInO2, Ag2Ti4O9, etc., have been recently developed as visible-light responsive photoactive materials [169]. In addition, ternary oxides based on vanadium and with a perovskite structure, such as SrVO3 and CaVO3, are highly stable and transparent oxides. Therefore, silver vanadates have attracted much interest due to their potential applications ranging from photocatalysis to batteries and sensors, and passing through antibacterial agents [170]. Silver vanadates have excellent electrochemical properties making them suitable for primary lithium ion batteries. Indeed, in the form of electrodes, Ag2V4O11 nanowires and α-AgVO3 microrods have exhibited high discharge capacities of 366 mA·h·g−<sup>1</sup> and 324.1 mA·h·g−1, respectively [135]. Silver vanadates have also interesting optical properties with intense absorption bands in the visible region, especially for α-AgVO3, β-AgVO3, Ag4V2O7, Ag2V4O11, and Ag3VO4 [171].

Concerning AgVO3, this ternary vanadium oxide exists in two stable phases, both having a monoclinic structure: α-AgVO3 and β-AgVO3. The alpha and beta polymorph belong to space group *C2/c* and *Cm*, respectively [135,170,172]. In addition, α-AgVO3 phase can be irreversibly transformed to β-AgVO3 at around 200 ◦C [170]. The structure of crystalline AgVO3 exhibits silver ions with three different local coordinations corresponding to *[AgOx]* clusters (*x* = 5, 6, and 7) and vanadium ions coordinated to four oxygen, thus forming tetrahedron *[VO4]* [170]. AgVO3 materials are p-type semiconductors with a direct energy band gap. The *Eg* of the most stable polymorph, i.e., β-AgVO3 is estimated to be about 2.2 eV [169,173]. That clearly indicates that β-AgVO3 is a visible-light driven photoactive material with promising applications under natural sunlight. As reported in the previous paragraph, the narrow energy band gap of silver vanadates is mainly due to hydridization of O 2p and Ag 4d orbitals which compose the valence band [169]. In addition, in β-AgVO3, there are intermediate energy levels within the band gap which arise from the structural disorder of the building blocks, i.e., *[VO4]* and *[AgOx]* clusters. This disorder also enhances the separation of photogenerated electron/hole pairs and implies highly mobile charge carriers [170]. Therefore, the electrochemical performance of β-AgVO3 is better than α-AgVO3. This better performance is further improved by the tunnel structure of the beta polymorph, thus allowing better cation insertion, i.e., higher specific capacity values [172].

Another silver vanadate, Ag2V4O11, is widely used as efficient cathode in LIB for advanced biomedical devices such as implantable cardiac defibrillators [172]. Indeed, Ag2V4O11 is a silver vanadate which been successfully commercialized for use as cathode material in LIB [174]. Ag2V4O11 has a monoclinic structure with the space group *C2/m* [135]. The crystal structure of this silver vanadate is made of *[V4O16]* units which are composed of distorted *[VO6]* octahedra sharing corners in order to form an infinite *[V4O12]n* quadruple strings [135]. These chains are linked by corner-shared oxygen to provide continuous *[V4O11]n* layers which are separated by silver atoms [135]. Ag2V4O11 is a semiconductor with an indirect band of about 2.0 eV. Indeed, the absorption band of Ag2V4O11 almost covers the region from UV to about 600 nm which corresponds to the visible portion of the natural sunlight [175]. Therefore, Ag2V4O11 is not only an excellent electrochemical material, but also a promising photoactive material. The conduction band of Ag2V4O11 is composed of the V 3d and Ag 5s5p orbitals while the valence band is the result of O 2p and Ag 4d hydribization as in most of silver vanadates [175]. The energy position of VBM and CBM can be determined theoretically by using the equation related to Mulliken electronegativity and the band gap value of the semiconductor oxide (Equation (11). As a result, the conduction band minimum is determined to be 4.4 eV relative to the vacuum level (which is close to *E(H+/H2)* = 4.5 eV) while the valence band maximum is calculated to be 6.4 eV (using *Eg* = 2.0 eV) [175].

$$\rm{E\_{CBM}} = X - 0.5E\_{\%} \tag{11}$$

where ECBM and *Eg* are the CB edge and band gap energies relative to the vacuum level, respectively, and *X* is the mean of the Mulliken electronegativity of the constituent atoms in the semiconductor [175].

Concerning the other silver vanadates, pure Ag4V2O7 is difficult to prepare since this phase often coexists with Ag3VO4 during the synthesis. Ag4V2O7 crystallizes in the orthorhombic system and it exhibits a band gap of about 2.5 eV. On the other hand, Ag3VO4 has also a band gap in the visible region with a value estimated to be 2.2 eV [176]. This phase crystallizes in the monoclinic system [176]. Compared to Ag4V2O7, the energy position of the conduction band in Ag3VO4 is less negative, thus leading to less efficient reduction ability [177]. Despite most of silver vanadates have narrow energy band gaps, the separation rate of photo-generated electrons and holes remain low, that is the main drawback of these ternary vanadium oxides [176].

#### 3.2.5. Manganese and Other Transition Metal Vanadates

There are many other transition metal vanadates, of which manganese ones are among the most documented. Most of manganese vanadates have been firstly studied for their electrochemical properties in LIBs [178]. However, Mn2V2O7 has attracted attention not only for its electrochemical performance but also for its optical properties. There are two phases of this manganese vanadate: monoclinic β-Mn2V2O7 and triclinic α-Mn2V2O7 [178,179]. β-Mn2V2O7 exhibits an energy band gap of about 1.2 eV while α-Mn2V2O7 has a larger *Eg*

(1.8 eV) [178,180]. The structure of β-Mn2V2O7 is composed of *[MnO6]* octahedra sharing edges with *[V2O7]* interlayer while for α-Mn2V2O7, it is often reported as impurity [178]. Another interesting manganese vanadate is Mn(VO3)2. It crystallizes in monoclinic structure and exhibits a direct energy band gap of about 3.0 eV [181].

For the other transition metal vanadates, the Table 1 summarizes the most important optical characteristics of such materials, i.e., their photocurrent and band gap energy values, since such properties are crucial for photo-induced chemical processes [180].

**Table 1.** Measured and calculated energy band gap for different transition metal vanadates along with their produced photocurrent.


#### **4. Vanadium Oxides in Photochemical Processes for Environmental Applications** *4.1. Production of Sustainable Energy*

One of the major concerns for humanity in the 21st century is the environmental changes [164,182]. Therefore, the production of renewable and sustainable energy is a crucial problematic throughout the world. Renewable energy sources can fulfil the requirements of energy without any production of pollution (toxic air pollutants and greenhouse gases), thus it is the best alternative to overcome the currently existing problems [183,184]. One of the main sustainable and renewable sources of energy is solar light with its utilization in photocatalysis, photovoltaic cell, etc. [164]. Therefore, besides the degradation of pollutant which is developed in Section 4.2, binary and ternary vanadium oxides can be used in the field of sustainable energy. The two main applications using vanadium oxide based materials in this field are the production of hydrogen by water splitting and the production of electricity using photovoltaic devices. Hydrogen (H2), as a fuel, has a great energy yield (122 KJ·g−1) compared to fossil fuels (40 KJ·g−1) which are known to produce greenhouse gases and particulate matter (PM). A green and sustainable way to produce hydrogen is from the splitting of water molecule for two main reasons: (i) water is abundant on the planet (70% of Earth area is covered by water); and (ii) such a technology does not produce wastes (only hydrogen and oxygen). Using vanadium oxides materials as semiconductor photocatalysts, water splitting can be photo-assisted. The photosplitting of water has attracted the interest of many environmental chemists since it is considered as the "Holy Grail" in the field of energy: conversion of solar energy into chemical energy [9]. However, since the overall reaction of water splitting (Equations (12)–(14)) has relatively high Gibbs energy (ΔG<sup>0</sup> = 237.2 KJ·mol−1), the use of only solar energy is not enough to trigger the production of both hydrogen and oxygen using the actual commercial photocatalysts. Therefore, unless the discovery of an extraordinary green, efficient and inexpensive material, additional energy should be supplied to efficiently break H2O molecules into molecular oxygen and hydrogen [185–188]. Therefore, electrical energy coupled to an irradiation (especially solar light) can be used in a photoelectrochemical (PEC) system containing a photoactive electrode (Figure 13). This strategy appears the best solution (so far) to produce green energy by water splitting. PEC is also called electrochemically assisted photocatalysis (electrophotocatalysis). To produce H2 in an economically viable way and to reduce production cost, the electrical energy can derive from renewable sources, such as photovoltaic systems, so the produced hydrogen is entirely "green". In other words, photoelectrochemical (PEC) cell is the most advanced way to produce chemical energy (hydrogen) by converting solar energy.

$$2\text{H}^+ + 2\text{e}^- \rightarrow \text{H}\_2\text{:} \qquad E^0 = 0\text{ V} \tag{12}$$

$$2\text{H}\_2\text{O} + 4\text{h}^+ \rightarrow \text{O}\_2 + 4\text{H}^+; \quad E^0 = 1.23\text{ V} \tag{13}$$

$$2\text{H}\_2\text{O} + h\nu \rightarrow 2\text{H}\_2 + \text{O}\_2; \quad \Delta E = 1.23\text{ V}\tag{14}$$

**Figure 13.** Two-compartment photoelectrochemical cell: the case of photoactive anode.

A PEC system is basically a cell coupled to a potentiostat. The PEC cell is composed of an anode and a cathode in contact with an electrolyte. An electrical voltage is applied between the electrodes. There are several cell configurations with either electrodes in the same compartment or separated by an ion-transfer membrane. For water splitting, the two-compartment PEC cell is the most suitable since O2 and H2 evolve separately, thus avoiding potential explosive mixture and water recombination (Figure 13) [189]. The electrophotocatalytic process for the production of hydrogen by water splitting using a PEC cell and a semiconductor photocatalyst such as vanadium oxides requires that the redox potentials for both water oxidation (Equation (13)) and reduction (Equation (12)) should lie within the energy band gap of the photoactive material (Figure 14) [186–190]. This photo-active material is often an n-type semiconductor photocatalyst like vanadium oxides and it constitutes the anode which is irradiated (photoanode). Since ΔE = 1.23 V (for overall water splitting), the energy band gap of a photocatalyst candidate should be theoretically 1.23 eV in order to photogenerate *e*<sup>−</sup> and *h+* species of enough energy to trigger the reactions (Equations (13) and (14)) [190]. However, due to overpotential losses, the *Eg* value should be at least 2 eV [185–188]. After the generation of *e*−*/h<sup>+</sup>* pair under suitable irradiation, the PEC water splitting process using a photoanode is composed of three main steps: (i) water oxidation in the photoanode compartment, (ii) *e*− transfer through external circuit toward the cathode compartment, where (iii) reduction of water occurs [186,189].

**Figure 14.** Mechansim of overall water splitting using heterogeneous semiconductor photocatalysis.

The first report on PEC water splitting dates 1972 [191]. In this work, Fujishima and Honda used rutile TiO2 photoanode which is irradiated by UV light in a two-compartment cell. Unlike titanium dioxide which is the most popular photocatalytic material, vanadium oxides are promising candidates for PEC water splitting. One reason is that many vanadium oxides can be photoactivated by visible light. Although Fe2O3, WO3, or CdS are also visible light driven photocatalyst, they cannot be used in this sustainable process due to inappropriate band structure, low photostability, high recombination rate, or high toxicity [185,186]. Another reason why vanadium oxides are interesting is their band structure. Indeed, the energy band gap of many vanadium oxides has the suitable energetic position, thus enclosing the redox potentials for water splitting, especially for the production of H2 [192,193] while most of semiconductor transition metal oxides are efficient in partial water splitting, especially water oxidation (oxygen production). Therefore, vanadium oxide photocatalysts can be used without any electrical voltage for the production of hydrogen where V6O13, VO2(M), and VO2(B) could produce until 35 μmol of hydrogen per gram of catalyst after 12 h [192]. Supported VO2 nanorods with energy band gap of 2.7 eV are also able to photocatalytically produce hydrogen by water splitting under UV light at room temperature. Indeed, in the presence of ethanol, a maximum value of H2 production rate of 800 mmol·m−2·h−<sup>1</sup> is obtained when the light is incident to the axis of the nanorods (due to vertical alignment of the nanorods) [194]. V2O5 has also shown interesting photo-induced properties with the generation of 8.9 mm3·h−<sup>1</sup> of hydrogen by water splitting under visible light and in the presence of methanol. The role of alcohols is to trap photogenerated holes, thus improving the electron-hole pair separation where electrons can be efficiently used in the production of hydrogen (Figure 15A). The high efficiency of V2O5 in photocatalytic water splitting is also explained by the high specific surface area which is controlled by the sol-gel technique where vanadium triethoxide precursor and structure directing agent are employed [195]. In the case of using photo-sensitization process, i.e., using Eosin Y dye, the production of hydrogen can reach almost 20 mm3·h−<sup>1</sup> [195]. Indeed, under irradiation, the excited dye can transfer its electrons to the conduction band of V2O5 photocatalyst where reduction reaction for the formation of hydrogen occurs (Figure 15B).

**Figure 15.** Improvement of hydrogen production using semiconductor photocatalyst by (**A**) sacrificial agent method (here, methanol) and (**B**) by dye photosensitization (here, Eosin Y).

However, the production rate of hydrogen using only light energy (without electrical bias) is low. This poor efficiency is mainly due to kinetic issues, limited solar energy conversion (photon to electron), and important charge recombination [196,197]. In addition, some binary and ternary vanadium oxides have a conduction band minimum which is not suitable for hydrogen production via the water splitting. For instance, BiVO4 has a CBM at 0 V vs. SHE which makes the generation of H2 experimentally impossible due to overpotential losses. However, BiVO4 is an efficient photocatalyst for water oxidation which is a more challenging reaction than water reduction (four photogenerated holes are required compared to two photogenerated electrons for hydrogen production) [198]. There are potential alternatives to overcome these drawbacks by, for instance, the surface modification of photocatalysts. Indeed, the use of a synergistic co-catalyst is able to decrease the activation energy for H2 production or to improve the charge carriers separation. Another example is the use of a suitable composite photocatalyst (especially a type-II heterojunction) since the photogenerated electrons are accumulated in one semiconductor while the photogenerated holes are collected in the other semiconductor, thus improving the *e*<sup>−</sup>*/h+* pair separation under irradiation. This is the case of FeVO4/CdS where 400 μM of hydrogen H2 is produced after 5 h under visible light [197] and of BiVO4/CdS nanoparticles with a hydrogen production rate of 0.57 mmol·h−<sup>1</sup> under visible light [199]. However, the solution which appears the most efficient (so far) is to set up a PEC system (i.e., the use of electric bias) since relatively simple materials with low production cost can be used while the electric bias improves the *e*<sup>−</sup>*/h+* pair separation by enforcing the movement of electrons oppositely to that of holes.

In such a PEC configuration, vanadium oxide photoanodes are more efficient, especially BiVO4 due to its appropriate electronic band structure, excellent chemical and thermal stabilities, and environmental friendliness [193,200,201]. Its band gap of 2.4–2.5 eV (absorption edge: 500–520 nm) allows the production of 6.2–7.5 mA·cm−<sup>2</sup> photocurrent under simulated solar light. In addition, its valence band maximum which is located at 2.5 V vs. RHE provides a strong driving force for water oxidation by photogenerated holes [143]. Concerning the production of hydrogen by water splitting, bismuth vanadate photoanode can produce up to 2.5 <sup>μ</sup>mol·min−<sup>1</sup> of hydrogen depending on the experimental conditions [202]. Since BiVO4 has poor separation ability and slow mobility of carriers, it has limited efficiency in practical applications [200]. Therefore, modifications such as doping, coupling to oxygen evolution cocatalysts (OECs) and fabricating heterostructure composites (using for example FeVO4 [203,204]) are considered to be the most promising techniques for improving the efficiency of vanadium oxide based materials [200]. Indeed, co-doping of metal and non-metal into ternary iron vanadium oxide could increase the carriers' concentration and decrease the energy band gap, thus enhancing the PEC activities [205]. BiVO4/FeVO4 heterostructured photoanode exhibits a photocurrent density of 2.5 mA·cm<sup>2</sup> at 1.23 V vs. RHE, which might be attributed to excellent charge transfer/transport phenomena (Figure 16) [206]. In addition to thermodynamic parameters, the physical factors including specific surface area, exposed crystal facets, etc., are also crucial to obtain the best photo-induced efficiency. The design of nanostructured vanadium oxides in the form of ultrafine powder, hierarchical structure and porous material is an efficient approach to reach high rates in hydrogen production [207,208].

**Figure 16.** Charge transport of photogenerated holes and electrons in layered BiVO4/FeVO4 photoanode.

On the other hand, PEC processes using photoanode are also interesting for the degradation of pollutants [185,186]. The main advantage in the use of a PEC cell is the reaching of higher mineralization rates [140,185]. Since the charge carrier recombination in the photoanode is limited by the vector charge transport along the electrochemical potential gradient, the activation energy for full mineralization is easily achieved [140,185]. However, vanadium oxides have high oxidation power, thus, non-assisted photocatalysis is sufficient for efficient degradation of pollutants (Section 4.2).

Another environmental application using the photoinduced properties of vanadium oxides under solar light is the photovoltaic technology. Such devices, especially in the form of organic solar cells, are composed of an anode, a cathode, an active organic layer and interfacial layers [209]. In these latter ones, vanadium oxides are promising. The interfacial layers cover the electrode and are in contact to the organic layer. Their function is to support the transport of photogenerated charges [209]. Usually, as interfacial layer, poly(3,4-ethylenedioxythiophene) poly(styrenesulfonate) (PEDOT:PSS) is employed but its poor long-term stability has pushed the scientists to consider other possibilities to replace PEDOT:PSS such as V2O5. Vanadium pentoxide can be deposited in ambient conditions, thus avoiding high temperature annealing and other post-treatments [210]. Recently, it is in perovskite solar cells that vanadium oxides exhibited outstanding performance as hole-transport layers [211].

#### *4.2. Photo-Induced Degradation of Organic Contaminants*

The organic contamination is a critical issue for the natural environment, especially for the hydrosphere. Indeed, many pollutants are accumulated in in the aqueous media, especially the persistent organic pollutants (POPs) and the contaminants of emerging concern (CECs), i.e., pharmaceuticals and personal care products (PPCPs), endocrine-disrupting compounds (EDCs), flame retardants (FRs), pesticides, herbicides, and fungicides, etc. These CECs are frequently detected in wastewater treatment plant (WWTP) effluents. Indeed, the use of conventional treatments are not enough efficient to remove the CECs while these contaminants have a well-known harmful impact not only on the environment, but

also on the human health. Therefore, the discharge of WWTP effluents (i.e., treated wastewaters) is the main source of contamination of CECs and POPs into the natural environment, thus being a real threat for drinking water source (ground and surface water) [212].

The properties of vanadium oxides are versatile mainly due to the multivalent character of vanadium. Thank to these properties, it is possible to design efficient photo-induced processes (photocatalysis, photo-Fenton, etc.) for environmental remediation, including the removal of organic pollution [83].

#### 4.2.1. Generalities of Photo-Induced Processes

One of the most efficient techniques to remediate the environment, especially for the removal of organic contaminants from water, is the advanced oxidation process (AOP). The AOPs are based on the generation of reactive oxygen species (ROS) including inorganic radicals, such as HO•. AOPs usually enhance the biodegradation of organic pollutants but it could also lead to degradation intermediates which are more toxic than the original targeted molecules [212]. The activation of the AOPs can be performed either by catalysis, heat, light, or a combination of these means (e.g., photocatalysis and photo-Fenton). It is the technology based on the use of coupled catalysis and solar radiation which is highlighted in this section since the resulting process can reach high remediation efficiency along with sustainable development (i.e., efficient and environmentally friendly process).

In AOPs, the most important active species are hydroxyl radicals (HO•). HO• has two main characteristics. First, it is a strong oxidant with a redox potential *E*0*(HO*•*/H2O)* of 2.8 V vs. NHE. Second, it is a nonselective oxidant which is able to react with most of the organic contaminants with high bimolecular kinetic constants (108–1011 <sup>M</sup>−1·s−1). These characteristics are responsible for the high efficiency of AOPs in the removal of POPs and CECs [213]. The reaction between hydroxyl radicals and organic molecules proceeds mainly by hydrogen abstraction (from C–H, N–H, or O–H bonds) and by hydroxyl addition to unsaturated bonds. It is worth noting that the involvement of hydroxyl radicals in organic decontamination process is not easy to demonstrate since the half-life of HO• is in the order of nanosecond. Therefore, most of their identification are performed using indirect methods such as fluorescence spectroscopy or electron paramagnetic resonance in the presence of suitable probe molecules. In addition to hydroxyl radicals, the photo-induced degradation of organic pollutants involve other radicals like superoxide anion (O2 •−) and its conjugated form, hydroperoxyl radical (HO2 •). These radicals are, however, less reactive than HO• [213].

These photo-induced AOPs are suitable techniques for the degradation of organic contaminants. Among them, direct photolysis, photocatalysis and photo-Fenton reaction are the most used processes. Direct photolysis is based on the sole use of light to break the chemical bonds in the contaminant molecule. Since UVB and UVC irradiations are the most used lights in direct photolysis, this method cannot be considered as sustainable process (and it is not discussed in this section) while photocatalytic and photo-Fenton processes are developed hereafter.

#### Photocatalysis

The photocatalysis is a photo-induced process combining the use light irradiation and catalysis. Therefore, the compounds which are employed in such a process are called photocatalysts like vanadium oxide based materials. In this case, the process is referred as heterogeneous photocatalysis since solid semiconductors are used. Prior to describe how heterogeneous photocatalysis is operating, a brief historical background, and some basic principles should be introduced. Although the definition of a semiconductor is in Section 3.1.2, the term semiconducting was first used by Alessandro Volta in 1782 and it is at the beginning of the 20th century that Johan Koenigsberger classified materials into conductor, insulator, and semiconductor. On the other hand, the concept of photoconductivity appeared in 1874 [164]. Indeed, an electron can move from valence band (VB) to conduction band (CB) when the supplied energy to the semiconductor is higher than the

energy band gap (*Eg*)—also known as forbidden energy band—which is the energy gap between the VB and the CB. The value of the energy band gap is an intrinsic value, thus varying from one material to another. In other words, under incident light of sufficient energy, the electrons move from VB to CB while holes have an opposite motion, thus generating a current [164].

A key parameters in the selection of a semiconductor photocatalyst is the value of the energy band gap along with the stability of the material. The principle of heterogeneous photocatalysis using semiconductors is based on the generation of electrons and holes, i.e., the *e*−*/h+* pairs (Equation (15)). Indeed, once the *e*−*/h<sup>+</sup>* pairs are photogenerated under an incident irradiation of energy (*hν*) equal or larger than the energy band gap, these charge carriers are responsible in the formation of reactive oxygen species (Equations (16)–(22)) which, subsequently, degrade organic contaminants. The primary ROS are formed by the reaction of electron and hole with oxygen and water, respectively. However, the mechanism of hydroxyl radical generation is still the matter of discussions since the photogenerated holes can react also with surface hydroxyl function present at the surface of the photocatalyst. The degradation of pollutants can also occur directly through photogenerated holes (which are oxidant), but this reaction is not kinetically favored due to the strong reactivity of ROS with organic molecules. The formation of ROS does not only include the generation of superoxide and hydroxyl radicals, but also other species such as hydroperoxyl radicals and hydrogen peroxide, which is an important precursor of HO• (Equations (16)–(22)) [33]. H2O2 is also a key precursor in Fenton-based processes (Section Photo-Fenton and Fenton-Based Processes).

$$Semiconductor + h\nu \rightarrow \text{e}^- + h^+ \tag{15}$$

$$\mathrm{OH}^- + h^+ \to \mathrm{HO}^\bullet \tag{16}$$

$$\mathrm{H\_2O} + h^+ \to \mathrm{HO^{\bullet}} + \mathrm{H^+} \tag{17}$$

$$\bullet \text{ O}\_2 + \text{e}^- \rightarrow \text{O}\_2 \\ \bullet \text{"} \tag{18}$$

$$\rm O\_2^{\bullet-} + H^+ \to HO\_2^{\bullet} \tag{19}$$

$$\text{H}\_2\text{HO}^\bullet \rightarrow \text{H}\_2\text{O}\_2\tag{20}$$

$$\text{CH}\_2\text{O}\_2 + e^- \rightarrow \text{HO}^\bullet + \text{OH}^- \tag{21}$$

$$\rm{H\_2O + O\_2^{\bullet-} \rightarrow HO\_2^{\bullet} + OH^-} \tag{22}$$

However, several conditions should be fulfilled to observe the above mentioned photocatalytic reactions. In addition to the fact that the *e*−*/h<sup>+</sup>* pair generation is in competition with the recombination process (releasing the absorbed energy in the form of either light or heat), the photogeneation of the charge carriers (electrons and holes) should reach the surface of the photocatalyst where they react with adsorbed reactants, since the photocatalysis is a surface dependent process (Langmuir–Hinshelwood model). In addition, the redox potential of the photocatalytic reaction (e.g., the generation of ROS) has to lie within the energy band gap (Figure 17). Therefore, the feasibility of a photocatalytic reaction depends mainly on the energetic position of the valence band maximum (VBM) and the conduction band minimum (CBM) of the semiconductor photocatalyst [33].

**Figure 17.** Mechanism of primary ROS generation by heterogeneous semiconductor photocatalysis.

Photo-Fenton and Fenton-Based Processes

Another important photo-induced process is the photo-Fenton reaction. Usually, this process describes the generation of ROS in the presence of iron. Indeed, the classical Fenton reaction was discovered in 1894 by H. J. H. Fenton whom observed the oxidation of tartaric acid in the presence of iron and hydrogen peroxide [214]. About 100 years later, the heterogeneous Fenton reaction was observed (Equations (23) and (24)). Compared to the homogeneous system, heterogeneous Fenton reaction can be used over a wider range of pH (e.g., environmental pH) since there is no need to use acidic condition to avoid the production of sludge (from iron coagulation/precipitation). However, the mechanism of heterogeneous Fenton is still unresolved. This complexity is due to various interactions between the surface of the material with hydrogen peroxide, organic molecules (and their degradation byproducts), and the generated ROS [213]. It is admitted the main ROS are HO• and HO2 •/O2 •− which are either generated at the surface of the materials or at the close surroundings of the surface through possibly leached iron (i.e., homogeneous Fenton reaction) [213].

$$\equiv \text{Fe}^{\text{II}} + \text{H}\_2\text{O}\_2 \rightarrow \equiv \text{Fe}^{\text{III}}\text{OH} + \text{HO}^\bullet \tag{23}$$

$$\equiv \text{Fe}^{\text{III}} + \text{H}\_2\text{O}\_2 \rightarrow \equiv \text{Fe}^{\text{II}} + \text{H}^+ + \text{HCO}^\bullet \tag{24}$$

On the other hand, heterogeneous photo-Fenton reaction is considered as viable alternative to classical heterogeneous Fenton, since there is no need to add hydrogen peroxide to the system (Equation (25)). Indeed, the generation of inorganic radicals such as hydroxyl and superoxide radicals occurs via the photolysis of iron (III) where electron transfers occur with adsorbed water molecules, i.e., a ligand to metal charge transfer (LMCT). The AOPs generated by heterogeneous photo-Fenton reaction are efficient in the degradation of organic contaminants including pharmaceuticals, pesticides/herbicides, etc. [215,216]. Furthermore, it allows the possibility of using natural solar light, thus increasing the sustainibility of such a process.

$$\equiv \text{Fe}^{\text{III}} + \text{H}\_2\text{O} + hv \rightarrow \equiv \text{Fe}^{\text{II}} + \text{H}^+ + \text{HO}^\bullet \tag{25}$$

However, other transition metals can also trigger Fenton-type reactions. To increase the environmentally friendliness of these processes, it is possible to use elements that are abundant in the Earth's crust. Intense research is actually focused on these iron-free Fenton processes [217]. One important element is aluminum which has a stronger ability to convert H2O2 to HO• through *e*− transfer due to more negative redox potential (*E*0*(Al*<sup>3</sup>*+/Al*<sup>0</sup>*)* <sup>=</sup> −1.66 V) than iron (*E*0*(Fe*<sup>3</sup>*+/Fe*<sup>2</sup>*+)* = +0.776 V). However, the main disadvantage of using zero-valent aluminum is the formation of Al2O3 layer which is hard to remove, thus requiring strong acidic conditions to avoid the inhibition of the Fenton process. The generation of ROS in such a system mainly occurs through the in situ production of hydrogen peroxide by electron transfer to dissolved oxygen. The hydrogen peroxide is then subsequently decomposed to generate hydroxyl radicals (Figure 18) [217].

**Figure 18.** Generation of primary ROS at the surface of aluminum.

Another interesting element with high natural occurrence is manganese due to its versatility in oxidation states ranging from 0 to +VII. Among them, only Mn (II), Mn (III), and Mn (IV) can be used in Fenton-based processes. Mn (II) species are predominantly present in soluble form while Mn (III) and Mn (IV) exist in the form of oxides and hydroxides, such as Mn3O4, Mn2O3, MnOOH, and MnO2. The latter is already known as oxidizing material for direct degradation (neither light nor radical precursor assistance) of adsorbed organic contaminants. In this iron-free Fenton process involving manganese, the mechanism is complex and it is not yet resolved. However, it is assumed that hydrogen peroxide is activated by interconversion between Mn2+ and Mn4+ via intermediate Mn3+. According to the type and the form of the oxide used, the main ROS are either hydroxyl radicals or superoxide radicals. The main advantage of this Mn-based Fenton process is a significant efficiency at circum-neutral conditions, especially for pH ranging from 4 to 8 [217].

Copper is the element which is the most analogeous to iron in Fenton-based processes. Indeed, like Fe (II) and Fe (III), monovalent and divalent copper can activate hydrogen peroxide to generate reactive oxygen species, especially HO• and HO2 • (Equations (26) and (27)). The main advantage of copper oxides over iron ones is their stability (Cu leaching is limited), thus allowing this iron-free Fenton reaction to run out at circum-neutral pH. However, Cu (I) is readily oxidized into Cu (II) by oxygen, thus, its reaction with hydrogen peroxide is inhibited. This is a strong drawback which limits practical application of this system. Excess of H2O2 has to be used which increases the cost of this iron-free Fenton process but decrease the oxidation efficiency (by scavenging the generated HO•) [217].

$$\equiv \text{Cu}^{\text{II}} + \text{H}\_2\text{O}\_2 \rightarrow \equiv \text{Cu}^{\text{I}} + \text{H}^+ + \text{HCO}^\bullet \tag{26}$$

$$\mathbf{u} \equiv \mathbf{C} \mathbf{u}^{\mathrm{I}} + \mathbf{H}\_{2} \mathbf{O}\_{2} \rightarrow \equiv \mathbf{C} \mathbf{u}^{\mathrm{II}} + \mathbf{OH}^{-} + \mathbf{HO}^{\bullet} \tag{27}$$

Finally, vanadium is also an element which is active in Fenton-type reactions. Among the multiple oxidation states of vanadium, only V (V) to V (IV) can trigger a Fenton-like reaction (Equation (28)), but in the form of oxo-cations. Such reactions are observed in biochemical process involving vanadium [134], but also at the surface of vanadium oxide based materials [218]. It is worth noting that cerium, a lanthanide element, in soluble or solid form (Ce3+/4+ or CeO2) can trigger Fenton-based processes on the same scheme as in the case of iron (Equations (29) and (30)). However, considering the toxicity of cerium, it cannot be implemented in applications for the treatment of water from organic contaminants [217].

$$\rm{VO}^{2+} + \rm{H}\_2\rm{O}\_2 \rightarrow \rm{VO}\_2{}^+ + \rm{H}^+ + \rm{HO}^\bullet \tag{28}$$

$$\equiv \text{Ce}^{\text{III}} + \text{H}\_2\text{O}\_2 \rightarrow \equiv \text{Ce}^{\text{IV}} + \text{OH}^- + \text{HO}^\bullet \tag{29}$$

$$\equiv \text{Ce}^{\text{IV}} + \text{H}\_2\text{O}\_2 \rightarrow \equiv \text{Ce}^{\text{III}} + \text{H}^+ + \text{HCO}^\bullet \tag{30}$$

The mechanism using these transition metals in the photo-Fenton process has not been yet clearly demonstrated. However, since hydrogen peroxide can be in situ photochemically generated (e.g., by photocatalysis), the iron-free Fenton-based processes appear to have a significant contribution in photocatalytic systems based, for instance, on binary and ternary vanadium oxides.

#### 4.2.2. Binary Vanadium Oxides V2O5

As discussed in Section 3.1, V2O5 materials are usually studied for applications in electrochemical devices, especially for Li-ion batteries. However, during the last decades, there is a growing interest of the use of vanadium pentoxide in photocatalytic applications, especially for the degradation of organic pollutants and the production of hydrogen by water splitting. In order to be thermodynamically feasible, the valence band maximum should be more positive than the redox potential of hydroxyl radical generation (*E*0*(HO*•*/H2O)* = 2.8 V) and the conduction band minimum should be more negative than the redox potential of hydrogen formation (*E*<sup>0</sup>*(H+/H2)* = 0 V), respectively [219]. Indeed, the VBM and CBM of V2O5 is about 2.6–3.0 V and from −0.1 to +1.0 V, respectively. Therefore, the band structure of V2O5 is suitable for these photocatalytic processes. It is worth noting that the generation of superoxide radical is theoretically not feasible (*E*<sup>0</sup>*(O2/O2* •*- )* <sup>=</sup> −0.33 V), but V2O5 presents surface defects in the form of V4+ which allows the generation of superoxide radicals [219].

Vanadium pentoxide can be synthesized using different experimental methods and prepared in different forms (powders, films, etc.). Most of the preparation techniques involved wet chemical methods which allow the use of structure directing agent (Triton X-100, polyvinylpyrrolidone, polyethylen glycol, etc.) in order to design a particular morphology. Whatever the method employed, all the prepared V2O5 photocatalysts require an annealing process in order to get crystalline materials, usually at temperature about 500 ◦C.

Supported photocatalysts can be prepared using the sol-gel method along with the vanadium triisopropoxide oxide as precursor [220,221]. Another preparation procedure of supported V2O5 is the hydrothermal method [222]. Such a solvothermal method often implies ammonium metavanadate (NH4VO3) as precursor, but bulk V2O5 powder can be also directly used as precursor. Indeed, in the presence of a reducing agent such as oxalic acid, VO2(B) film is prepared (Equations (31) and (32)) which is subsequently converted into V2O5 with the appropriate morphology [222].

$$\rm{V\_2O\_5 + 3H\_2C\_2O\_4 \to 2VOC\_2O\_4 + 3H\_2O + 2CO\_2} \tag{31}$$

$$\text{VOC}\_2\text{O}\_4 \rightarrow \text{VO}\_2 + \text{CO}\_2 + \text{CO} \tag{32}$$

By this mean, V2O5 films with different morphologies (including hierarchical nanostructures, nanorods, nanowires, etc.) can be deposited on different substrates such as sapphire, indium tin oxide (ITO) coated glass, etc. [220,222]. The thickness of the films can be regulated by the number of deposited layers on the substrate, and that can have an impact on the quality and the efficiency of the supported photocatalyst [220]. The calculated energy band gap (*Eg*) is about 2.1 eV, so the V2O5 photocatalytic films can be used in the degradation of organic pollutants in aqueous or gaseous phases under simulated sunlight [220–222]. For instance, V2O5 nanowires films can degrade around 98% of toluidine blue O, an organic dye, after 50 min under UV light [220]. Flower like nonastructure films are able to remove up to 74% and 63% of methylene blue under UV and visible light, respectively, after 200 min [222]. In addition, 51% of this organic dye is mineralized into CO2 and H2O in the case of UV irradiation. Gaseous pollutants can be also efficiently degraded by V2O5 photocatalysis, such as for example 1,2-dichlobenzene [221]. Under visible light, about 45% of this organic contaminants can be degraded after 7 h

irradiation in presence of V2O5 [221]. The excellent photocatalytic properties of V2O5 films are ascribed to large specific surface area and efficient visible light utilization [221,222].

On the other hand, V2O5 powders can be also preprared. As in the case of supported photocatalysts, sol-gel, and hydrothermal methods are often chosen to synthesize V2O5 powders using NH4VO3 and H2O2 as precursors [223,224]. In addition, co-precipitation can be also used in the presence of NH4VO3 and HNO3 [225]. The energy band gap of the synthesized V2O5 powders is about 2.3 eV while the specific surface area is high for nanostructured powders [223–225]. Nanostructured powders can degrade efficiently rhodamine 6G, congo red, and methyl orange under visible light with removal extents of 78%, 99%, and 82%, respectively, after 180 min irradiation time [223,224]. Phenolic contaminants such as phenol, pyrocatechol (2-HP), 2-chlorophenol (2-CP) 2-nitrophenol (2-NP), and 2-aminophenol (2-AP) are also degraded by V2O5 photocatalysis under natural sunlight [225]. Although the degradation extents for phenol, 2-HP, 2-CP, 2-NP, and 2-AP are 13%, 40%, 68%, 81%, and 37%, the mineralization is lower with extents about 8%, 16%, 42%, 29%, and 27%, respectively [225]. The high efficiency of V2O5 powder photocatalysts is due mainly to the high specific surface area, relatively good utilization of visible light and excellent electron-hole pair separation [223,224]. Indeed, based on the produced photocurrent and compared to commercial micro-sized V2O5 powder, the number of charge carriers reaching the surface of the nanostructured materials (where reactions occur) is higher [224]. In addition, the exposition of V2O5 to irradiation induces the formation of surface defects which are composed of V4+ states, thus improving the lifetime of charge carriers by electron trapping phenomenon [225]. Moreover, the reproducibility of the photocatalytic experiments is also excellent with almost no change in removal rate after repeated runs [223]. Furthermore, V2O5 nanopowders are also efficient antimicrobial agents with improved removal of *Escherichia coli* and *Staphylococcus aureus* bacteria [223]. Recently, the effect of various extra-drying process (e.g., vacuum, freeze, oven) during the preparation of V2O5 nanoparticles through wet chemical methods has shown a limitation in the particle aggregation, thus increasing the reactivity of the photocatalyst [226]. Indeed, the nanoparticles size are in the range of 30–40 nm and the specific surface area is in the range of 125–205 m2·g−<sup>1</sup> [226]. Therefore, the synthesis procedure is crucial in the preparation of efficient photocatalysts.

Concerning the mechanism of any photo-induced degradation of organic contaminants, the blank and dark experiments should be always performed, i.e., direct photolysis and adsorption, respectively. V2O5 cannot efficiently remove organic pollution via adsorption processes while solar and visible light has not enough energy to photolyze most of the organic contaminants [220]. The photocatalytic mechanism is a surface-dependent process following the model of Langmuir–Hinshelwood, thus requiring the adsorption of the target molecule onto the surface of the photocatalyst. Many studies have shown that the photocatalytic degradation of organic contaminants using V2O5 occurs mainly via the primary ROS which are superoxide and hydroxyl radicals [221,223–225]. Indeed, hydroxyl radicals are form by reaction of photogenerated holes with surface hydroxyl group or adsorbed water molecule [221,225]. However, the involvement of O2 •− species is contradictory among the published works. Some scientists claim the reaction of photogenerated electrons with dissolved oxygen and subsequently, hydrogen peroxide and hydroxyl radicals can be formed by a cascade of reactions (Equations (19), (33)–(35)) according to the pH [221,224,225]. However, regarding the band structure of V2O5, especially the conduction band minimum, the formation of O2 •− is not thermodynamically feasible [225]. A viable tool for the identification of ROS, particularly for HO• and O2 •−, is the electron paramagnetic resonance (EPR). Indeed, the signals of DMPO-HO• and DMPO-O2 − spin adducts have a characteristics hyperfine structure [224]. By this mean, the photocatalytic degradation of organic contaminant in the presence of V2O5 suggests a negligible contribution of the superoxide radicals [224].

$$\text{HO}\_2\text{}^\bullet + \text{HO}\_2\text{}^\bullet \rightarrow \text{H}\_2\text{O}\_2 + \text{O}\_2\tag{33}$$

$$\text{O}\_2\text{O}\_2^{\bullet-} + \text{H}\_2\text{O}\_2 \rightarrow \text{OH}^\bullet + \text{OH}^- + \text{O}\_2 \tag{34}$$

$$\text{H}\_2\text{O}\_2 + hv \rightarrow 2\text{OH}^\bullet\tag{35}$$

The degradation of aromatic compounds results in ring opening through the oxidation by HO•, and ultimately leads to smaller species such as acetate and maleate [221]. In the case of organic dye, degradation occurs via *N*-deethylation, thus giving rise to smaller intermediates such as benzoic acid, hydroxybenzoic acid, etc. [223]. It is important to mention that the pH of the system is a crucial parameter [223,225]. Indeed, the ionic form of the organic pollutants (which is related to its *pKa*) and the charge surface properties of the material are dependent to the pH, thus impacting the interactions pollutant/photocatalyst. In addition, the pH affect also the generation of ROS [225]. For instance, at low pH, superoxide radicals form hydroperoxyl radicals (Equation (19)) and subsequently hydrogen peroxide (Equations (33) and (37)) which is a precursor for the generation of hydroxyl radicals (Equations (21), (34), and (35)) [225].

To resume, vanadium pentoxide nanostructured materials are promising solar light driven photocatalysts for the remediation of water from organic contaminants. So far, many studies have been focused on the removal of organic dyes, thus, the self-sensitization of the photocatalyst for the generation of ROS is possibly involved. Therefore, further works on colorless POPs and CECs are necessary in order to draw a better overview on the efficiency of V2O5 photocatalysis.

#### VO2

Vanadium dioxide is thoroughly investigated for applications as smart windows due to its thermochromic properties (Section 2.2). The study of its photocatalytic properties is not yet well documented although it is a promising photoactive material. Among the different polymorphs of vanadium dioxide, VO2(B) exhibits interesting photocatalytic properties although it is a semiconductor which is not thermochromic and which has a slightly larger energy band gap than that of V2O5, with a value of 2.8 eV [227]. Supported VO2(B) can be prepared by an aqueous sol-gel method using V2O5 and hydrogen peroxide as precursors. The deposition of the sol-gel should be performed on substrates that will not react with the vanadium oxide layer (such as Si/SiO2 and sapphire) since on lime glass, contamination by ion diffusion could occur during the annealing process [228]. Layered VO2(B) is obtained by annealing in reducing conditions, e.g., in H2/Ar atmosphere or in vacuum in order to reduce V2O5 to mixed valence V4O9 and ultimately to VO2. The photocatalytic activity of VO2(B) films is considerably better than that of V2O5 with a more efficient removal of Rhodamine B under UVA light [227]. The higher photocatalytic efficiency could be ascribe to the defect at the surface of vanadium oxide since the reductive annealing leads to more oxygen vacancies which are beneficial for the *e*−*/h<sup>+</sup>* pairs separation.

Since the literature on single phase VO2 photocatalysts is scarce, composites containing VO2 is also discussed in this section. Usually, the synthesis procedure of such materials involves hydrothermal methods using V2O5 and a reducing agent such as citric and oxalic acid [229–233]. By this mean, VO2(B) and VO2(M) can be prepared according to the experimental conditions. Then heterojunctions with another semiconductor metal oxide can be formed using, for instance, dispersion of VO2 powder into a metal alkoxide sol-gel [230,231,233]. VO2(B)/CeO2 photocatalysts in the form of nanoparticles exhibit efficient photo-induced properties for the removal of Rhodamine B with a removal extent reaching up to 66% in 180 min under UV irradiation. Similarly, VO2 alone can remove only 29% of this dye under the same experimental conditions [229]. The better photocatalytic properties of VO2(B)/CeO2 is explained by a smaller energy band gap (*Eg* = 3.25 eV) compared to that of pure VO2 (*Eg* = 3.5 eV). This shift toward lower values is due to the presence of CeO2 which absorbs larger extent of the visible light [229]. VO2(M)/ZnO composite is a photocatalyst with high specific surface area and energy band gap of 2.67 eV [231]. Indeed, in the form of a hierarchical structure, VO2(M)/ZnO gives rise to a specific surface area of about 70 m2·g−<sup>1</sup> which is a two-fold increase compared to the single component, thus, increasing the number of photocatalytic sites. In such a composite, ZnO is responsible to the widening of *Eg* due its energy band gap of 2.7 eV

compared to 2.3 eV for pure VO2. VO2(M)/ZnO photocatalyst can degrade up to 80% Rhodamine B under solar like irradiation after 20 min, which is significantly better than the single components [231]. In addition, the reproducibility after repeated photocatalytic runs is excellent while the composite remains stable without any dissolution of the material. Another efficient composite photocatalyst is VO2(M)/TiO2 which is able to degrade completely methylene blue within 10 min under UV light and about 70% of Rhodamine B after 2.5 h under solar irradiation [230,231,233]. However, the purpose of this latter composite is to design thermochromic and self-cleaning coating by combining the properties of TiO2 and VO2. Therefore, the photocatalytic properties are ascribed only to TiO2, thus, assuming a negligible contribution of VO2 in the photocatalytic mechanism [230,233].

The mechanism of photocatalytic degradation of organic pollutant in the presence of VO2 based photocatalysts involves the generation of primary ROS (i.e., HO• and O2 •−) by reaction of water and oxygen with photogenerated holes and electrons, respectively [229]. These strong oxidizing radicals are the main reactive species in the mineralization of organic contaminants. In addition, in the case of VO2(B)/CeO2, CeO2 plays the role of an electron trap, thus, the photogenerated *e*<sup>−</sup> are accumulated in there, that improves the *e*<sup>−</sup>*/h+* pairs separation [229]. It is also plausible that CeO2 induces Fenton-type reactions which would explain the excellent efficiency of this composite photocatalyst. For VO2(M)/ZnO photocatalyst, photocorrosion of ZnO was expected (Equation (36)) but in this type-II hetrojunction, the photogenerated holes are accumulated in VO2 while electrons are transferred to ZnO, thus, reducing the photocorrosion and limiting the recombination of *e*−*/h<sup>+</sup>* pairs [231].

$$\text{ZnO} + 2\text{li}^\* \rightarrow \text{Zn}^{2+} + 1/2\text{O}\_2 \tag{36}$$

Vanadium dioxide can be also combined with ternary oxides such as silver vanadate (Ag3VO4) [232]. VO2/Ag3VO4 composite prepared by hydrothermal method exhibits efficient utilization of both UVA and visible light (at λ < 420 nm) since about 73% of Rhodamine B is removed after 90 min irradiation time. The reproducibility and stability of this composite is also stable. To explain the excellent photocatalytic efficiency of VO2/Ag3VO4, VO2 is ascribed as the key component, since in this composite, it acts as electron trap, thus improving the *e*<sup>−</sup>*/h+* pairs separation [232].

#### Mixed Valence Vanadium Oxides

Mixed valence vanadium oxides are attractive materials for many applications. Among these oxides, V4O9 and V6O13 are the most investigated in photocatalytic applications. One of the first works about the photocatalytic properties of such bynary vanadium oxides appeared in 1990 where the photo-induced degradation of isopropanol was 2.5-fold enhanced under UV-visible light [234]. Indeed, V2O5 is already a good photocatalyst in the degradation of organic contaminants, but the amount of electronic defects (i.e., the presence of V4+ at V2O5 surface) is assumed to play a key role in the photocatalytic efficiency. Thus, by extending this hypothesis, vanadium oxides of mixed V (IV) and V (V) should exhibit improved photocatalytic properties [234,235]. Although such oxides are promising, they are much less studied than V2O5 and VO2 [236]. Usually V4O9 and V6O13 are synthesized from V (V) precursors such as V2O5 and NH4VO3 which are reduced by heat treatment in either vacuum or reducing atmosphere [234–236].

V6O13 is probably the most documented mixed valence vanadium oxides [234–236]. This photocatalyst can oxidize completely 3-hexanol under visible light (λ = 460 nm) while in the dark or using photo-activated VO2 and V2O5, the degradation of this alcohol is negligible [236]. In addition, the photocatalytic degradation of primary alcohols including 1-pentanol, 1-hexanol, 1-octanol, 1-nonanol, and 1-decanol is also efficient using V6O13 photocatalyst. It is worth noting that this photocatalytic degradation is selective since about 70% of the alcohols is oxidized in aldehyde [236]. Furthermore, the photocatalytic degradation of secondary alcohols using V6O13 leads to more than 70% selectivity in the formation of ketones. V6O13 photocatalyst can also efficiently degrade saturated hydrocarbons, especially aromatics, such as xylene and toluene [236]. In the form of a composite, V6O13 is also an efficient photocatalyst [235]. Indeed, VO2/V6O13 can remove organic contaminants, such as atrazine and methylene blue under visible light with an excellent reproducibility after repeated runs, thus indicating the stability of the photocatalyst. This photoactive nanostructured VO2/V6O13 is able to completely mineralize atrazine after 2 h irradiation while 75% of methylene blue can be degraded within 1 h, which is significantly better than V2O5 and VO2 photocatalysis [235]. The VO2/V6O13 exhibits high specific surface area (about 25 m2·g−1) and an energy band gap of 2.4 eV. It is worth noting that, in the VO2/V6O13, the band structure is modified due to the formation of heterojunction, thus extending the light absorption in the visible part of the spectrum [235].

Concerning the degradation mechanisms, the presence of V (IV) leads to the formation of oxygen vacancies which (i) enhanced the adsorption of the organic molecule at the surface of the material and (ii) improved the *e*−*/h<sup>+</sup>* pairs separation [234]. The photoinduced generation of primary ROS such as hydroxyl and superoxide radicals operates more efficiently due to the presence of V (IV) at the surface of the photocatalyst [234]. An alternative mechanism of photocatalytic degradation of organic molecules proceeds via chemisorption of the contaminant onto the surface of the material followed by excitation of the V-O-C bonding [236]. This phenomenon causes the elimination of protons in alpha position, and ultimately to the production of byproducts. This mechanistic scheme is different from the conventional photocatalytic mechanism [236]. In composite photocatalyst like VO2/V6O13, the adsorption of organic pollutants at the surface of the material is improved, thus being one reason of the excellent photocatalytic efficiency. In addition, the presence of both V (IV) and V (IV) improves the separation of charge carriers, namely the photogenerated electrons and holes, thus giving rise to higher production of primary reactive oxygen species that lead to the mineralization of organic contaminants into CO2 and H2O [235].

Another popular mixed valence vanadium oxide is V4O9 [237]. This compound has shown efficient but smaller photo-induced properties than V6O13. Indeed, V4O9 photocatalyst can degrade Rhodamine B under visible light with a removal extent of 45% after 3 h irradiation. The energy band gap of V4O9 is about 2.13 eV. It is also interesting to note that this mixed valence oxide is not often mentioned in the literature as a photocatalyst for several reasons. The first one is that V4O9 has similar photocatalytic efficiency as V2O5 which is easier to prepare. The second one is the similar crystalline structure between V2O5 and V4O9 which both have an orthorhombic system, thus resulting in very similar diffractions during the characterization by XRD [237]. Consequently, it is plausible that V4O9 has been confused with V2O5. The mechanism of photocatalytic degradation of organic pollutants using V4O9 exhibits predominant role of superoxide radicals and photogenerated holes [237].

#### 4.2.3. Ternary Vanadium Oxides Bismuth Vanadates

Among ternary vanadium oxides based on vanadium, BiVO4 is probably the most studied one due to bismuth which broadens the response of the material in the visible light, thus increasing its applicability for photocatalysis under sunlight. In addition, BiVO4 is stable and non-toxic which highlight its environmental compatibility [149,238]. Bismuth vanadate exists in three polymoprhs (Section 2.1) where the monoclinic scheelite phase is the most photocatalytically efficient [33]. Indeed, it is a promising visible light driven semiconductor photocatalyst which is mainly investigated in suspended systems and supported materials. The latter is more convenient for environmental applications since the immobilization of the photocatalyst on supports such as glass and polymer is beneficial for post-treatment separation [33]. BiVO4 with different morphologies can be designed using wet chemical methods, especially hydrothermal techniques [238]. By this mean, different hierarchical structures can be prepared by varying experimental conditions including pH, temperature, type of precursor, and presence of structure directing agent [238,239]. Prior to develop the applications of BiVO4 in photo-induced processes for the degradation

of organic pollutants, it is worth mentioning that the photocatalytic efficiency in the degradation of organic pollutants differs from one work to another. These differences are due to the way of synthesis of BiVO4, the type of incident radiation and the kind of targeted pollutant [33].

Most of the literature is devoted to the degradation of organic dyes. Although they are not considered as suitable model pollutants, organic dyes are environmentally relevant contaminants and they provide an overview on the feasibility of the tested material in the removal of organic pollutants. Nanostructured BiVO4 with 1D and 2D morphology can degrade up to 93% of methylene blue and 86% of methyl orange after 60 and 150 min visible light irradiation, respectively. In addition, BiVO4 nanotubes are able to completely remove Rhodamine B after only 15 min irradiation. These excellent removal extents are mainly ascribed to the high specific surface area which is ranging from 5 to 45 m2·g−<sup>1</sup> for such nanostructures. These values are considerably higher than those of bulk BiVO4 (about 0.5 m2·g<sup>−</sup>1) [238]. The high specific surface area is associated with the increase of active site number and better adsorption properties. Furthermore, the recombination rate between *e*−*/h<sup>+</sup>* pairs is smaller than in bulk structure due to better transfer of charge carriers at the surface (i.e., shorter pathway resulted from smaller particle size) [238]. Other structures such as hierarchical ones are also efficient in the degradation of organic pollutants. For example, denditric, hollow sphere and olive-like BiVO4 can degrade under visible light more than 80% of a wide range of organic dyes (e.g., Rhodamine B and methylene blue) after short irradiation times [238]. Like 1D and 2D nanostructured BiVO4, the main advantage of hierarchical structures is their high specific surface area (higher than 2 m2·g<sup>−</sup>1) [238].

Since the decolorization of organic dyes provides only basic knowledge of the feasibility of the photocatalytic reactions using recent materials, deeper insights into BiVO4 mediated photo-induced processes for the removal of organic contaminants are required [33]. To this end, two strategies are possible: (i) the study of the degradation pathway and mineralization extent of organic dyes; and (ii) the use of uncolored model pollutants in order to avoid self-sensitization process (i.e., in order to study the intrinsic properties of the photoactive material). The mineralization of the contaminants is crucial since incomplete removal could lead to more toxic intermediate byproducts [33]. Concerning the first strategy, total organic carbon and mass spectrometry studies have shown that BiVO4 is able to mineralize 65% of methylene blue under visible light while the intermediate degradation products are composed of hydroxylated aromatics [33]. In addition, the complete removal of Rhodamine B by the visible light driven BiVO4 photocatalyst is associated to only 40% mineralization extent after 100 h irradiation time [240]. The degradation of Rhodamine B follows two main pathways which are the deethylation process and the removal of the chromophore since the identified intermediates are ethylbenzene, xylene, etc. [240]. Concerning the second strategy, pharmaceutical compounds are often used such as ibuprofen (IBP) and ciprofloxacin (CIP) [241,242]. Indeed, under UVA light, BiVO4 can degrade about 90% of IBP and 65% of CIP. BiVO4 photocatalyst is also efficient in the degradation of phenolic compounds such as nonyl- and octyl-phenols which are well known endocrine disruptor compounds [243]. In this latter case, the alkyl chain of the phenolic compounds plays the role of an anchor on the surface of BiVO4, thus improving its adsorption onto the surface of the photocatalyst and improving the degradation reaction. Apart from aqueous pollutants, BiVO4 can also degrade contaminants in gaseous phase including isopropyl alcohol and nitrogen oxides [244,245]. For instance, under visible light, BiVO4 is able to convert 25% of N2O after 12 h irradiation time. To improve the photocatalytic efficiency of BiVO4, besides doping and preparation of composite, an interesting strategy is to use simultaneously different contaminants which are removed by reduction and oxidation processes. Indeed, the case of dual phenol and chromium (VI) pollution can be efficiently treated using BiVO4 since phenol can be oxidized and mineralized into CO2 and H2O while Cr (VI) can be reduced in Cr (III) (which is much less toxic) [246]. Concretely, after 3 h irrdiation time under visible light, complete removal of Cr (VI) is achieved and 60% of phenol is mineralized by the BiVO4 photocatalysis. For comparison, at similar experimental condition and using

phenol alone, only 10% of the pollutant is completely removed. Another interesting strategy to enhance the photocatalytic ability of BiVO4 is to prepare or anneal the material in reducing atmosphere (H2) in order to generate oxygen vacancies which improve the charge carrier transport, thus resulting in better photocatalytic efficiency [247]. Under visible light, hydrogenated BiVO4 can degrade up to 98% of tetracycline which is significantly higher than the 52% achieved using pristine BiVO4.

The mechanism of degradation of organic pollutants using BiVO4 is crucial and several contradictions are present in the literature. Indeed, the identification of the main reactive oxygen species gives rise to different results from one work to another. Some researchers have stated the main ROS are hydroxyl radicals while other ones have found the main contribution in organic pollutant degradation is from superoxide radicals. These differences are due to misinterpretation of results that arises from unreliable employed method to identify the primary ROS such as the scavenging method. Therefore, more sophisticated methods should be employed like for instance the electron paramagnetic resonance (EPR). By EPR, it has been proven that BiVO4 photocatalytically generates neither hydroxyl radicals nor superoxide radicals [248]. Indeed, the theoretical band structure of monoclinic scheelite BiVO4 (at normal conditions) exhibits a conduction band minimum and valence band maximum of 0 V and 2.4 V vs. NHE, respectively (Figure 19). Therefore, the redox potential for the formation of primary ROS, i.e., *E*<sup>0</sup>*(O2/O2* •*- )* and *E*0*(OH*•*/H2O)* are not within the energy band gap of BiVO4, thus, the production of these radicals is thermodynamically not favored [33,198,239,248].

**Figure 19.** Energy diagram of the band structure of BiVO4 photocatalyst.

However, why HO• is often detected in BiVO4 photocatalytic systems? One reason is the pH value of the reaction system. The theoretical band structure of BiVO4 is for pH = 0 while in practice, the pH is so-called natural. Therefore, an increase of pH leads to a shift of the bands toward negative values, thus allowing the formation of superoxide radicals. These O2 •- can be converted into HO• via the generation of H2O2 during the photocatalytic process according to the chemical reaction (Equations (21) and (37)) [240,241,249]. It is worth noting that even at natural pH, the generation of hydroxyl radicals by photogenerated holes from BiVO4 is even less favored [241].

$$\rm O\_2^{\bullet-} + 2H^+ + e^- \rightarrow H\_2O\_2 \tag{37}$$

The second reason is when organic dyes are used. In this case, sensitization of BiVO4 occurs by transfer of electrons from the excited state of the dye to the conduction band of the photocatalyst (Equations (38) and (39)). Therefore, the resulted improvement of the photocatalytic process is not due to the intrinsic photo-induced properties of BiVO4. Indeed, the injected electrons in the BiVO4 react with oxygen to produce superoxide radicals, thus following the reaction scheme above (Equations (18), (19), (33)–(35) and (37)), i.e., the generation of H2O2 and subsequently HO• [240].

$$\text{RhB} + h\nu \rightarrow \text{RhB}^\* \tag{38}$$

$$\text{BiVO}\_4 + \text{RhB}^\* \rightarrow \text{RhB}^{\bullet+} + \text{BiVO}\_4(\varepsilon^-) \tag{39}$$

The production of hydroxyl radicals is strongly dependent of the pH which is the key factor. Some researchers stated that the generation of hydroxyl radicals from superoxide radicals is favored at low pH while other scientists discussed that the alkaline pH might be more favorable due to the presence of OH−, which is easily convertible into HO• by reaction with photogenerated holes. However, the latter case has been demonstrated to be thermodynamically not feasible. Although BiVO4 is a promising photocatalyst for environmental remediation, further studies are required to develop its application for degradation of a wider range of contaminants and to clear out the uncertainties of its degradation mechanism.

Other bismuth vanadates such as Bi4V2O11 and Bi8V2O17 are also used as visible light driven photocatalysts [250,251]. The preparation methods of these materials are similar to that for BiVO4, thus BiVO4 is often present as impurity, i.e., pure mixed valence bismuth vanadate is not easily prepared. However, their use in the removal of organic pollutants is efficient under visible light. For instance, BiVO4/Bi4V2O11 can be efficiently degrade and mineralize sulfamethazine with degradation and mineralization extents of 95% and 50%, respectively [250]. The main primary ROS are identified as superoxide radicals.

#### Iron Vanadates

Iron vanadates, especially FeVO4, are interesting materials for the photo-induced degradation of organic pollutants. Indeed, due to their semiconducting properties and the presence of iron, both photocatalytic and photo-Fenton processes can be triggered using these photoactive materials. Although iron vanadates are gaining stronger interest within the photochemical community, few works on pure single phase are reported, thus, composite materials based on iron vanadates are also discussed in this section.

FeVO4 is a semiconductor with an energy band gap of about 2.15 eV, thus absorbing visible light [156]. The conduction band minimum is positioned at −0.65 eV while the valence band maximum is at 1.50 eV. FeVO4 can be synthesized by wet chemical methods using iron nitrate and ammonium metavanadate as precursors, and structure-directing agents (e.g., PEG and PVP) to design a particular morphology [132,156]. According to the morphology, the *Eg* varies from 2.1 to 3.0 eV where nanostructures exhibit higher energy band gap than bulk ones [132,156,252–254]. The specific surface area is found to vary between 15 and 30 m2·g−<sup>1</sup> [132,254]. FeVO4 photocatalysts are efficient for the removal of organic dyes under UVA light with removal extents of about 90% for Rhodamine B and 50% for phenol red after 300 min and 120 min irradiation time, respectively [156,252,254]. However, under visible light, the photocatalytic efficiency is lower [156,254]. Indeed, about 70% of Rhdomaine B can be degraded after 5 h irradiation time with less than 10% mineralization rate [151,156,253,254].

Due to the presence of iron in the material, FeVO4 can also triggered Fenton-type reactions. For instance, in the presence of H2O2, the degaradtion of Rhodmaine B is almost completed under visible light [254]. Compared to classical heterogeneous Fenton catalysts such as hematite (α-Fe2O3), magnetite (Fe3O4), and goethite (γ-FeOOH), FeVO4 is able to efficiently activate H2O2 into hydroxyl radicals with the removal of more than 90% of methyl orange after1hreaction [152]. Under visible light, photo-Fenton is also triggered by FeVO4 with subsequent activation of peroxymonosulfate (HSO5 −) into sulfate radicals (SO4 •−), thus leading, after 1 h of irradiation time, to the degradation of 96% of sulfamethoxazole, an antibiotic substance often detected in wastewater effluents [255]. However, the performance of this photo-induced process is rather low due to weak photolysis rates for iron and vanadium and unstable cycle between high and low valences species, i.e., Fe(III)/Fe(II) and V(V)/V(IV) [255].

Therefore, composite materials based on FeVO4 can be prepared to increase the efficiency of the photo-induced processes. ZnO/FeVO4 material triggers simultaneously, under UVA light and in the presence of H2O2, photocatalytic and Fenton processes with synergetic efficiency, thus achieving 93% degradation extent of sodium dodecyl sulfate within 1 h [256]. Similarly, FeVO4/CeO2 is also a material combining photocatalytic and photo-Fenton processes which are efficient in the removal of 4-nitrophenol [257]. In this case, the contribution of Fenton reaction, i.e., measured in the dark with the presence of hydrogen peroxide, has led to a 9% degradation extent after 1 h reaction while with both visible light and H2O2, the removal extent reached 94% [257]. In addition, such photoactive materials can be reused since efficiency after repeated runs is stable while the photo-induced processes, especially the Fenton reaction, do not required highly acidic conditions to be efficient. These multi-photoactive materials are efficient in the treatment of real wastewaters, thus highlighting their potential applications at industrial level [256]. On the other hand, many FeVO4 based composites are studied regarding only the photocatalytic process without taking into account the potential contribution of the photo-Fenton process. FeVO4/V2O5, FeVO4/BiVO4, and FeVO4/Fe2O3 are interesting composites which have exhibited excellent photocatalytic activity in the removal of many organic dyes, including Rhodamine B, phenol red, methyl violet, and methylene blue [133,258,259]. In such composites, direct Z-scheme or type-II heterojunction (Figure 12 in Section 3.2.3) improves the *e*<sup>−</sup>*/h+* pair separation, thus the photogenerated charge carriers react more efficiently (for the formation of ROS), but also enhances the absorption of visible light (i.e., a shift of the *Eg* toward shorter wavelengths). Compared to the single components, the use of composite photocatalysts under either UVA or visible light irradiation leads to an increase in the removal extents of organic pollutants from +10% to +80% according to the experimental conditions [133,258,259].

Concerning the photo-induced mechanisms for the degradation of organic pollutants, the photocatalytic and photo-Fenton are both involved. For photocatalysis using pure FeVO4, the band structure of this semiconductor allows the formation of superoxide radicals from the reaction of dissolved oxygen with the photogenerated electron since *E*<sup>0</sup>*(O2* •−*/O2)* is situated within the energy band gap of FeVO4. The photogenerated holes cannot produce hydroxyl radicals since *E*<sup>0</sup>*(H2O/HO*•*)* is more positive than the valence band maximum [132]. However, the *h<sup>+</sup>* can oxidize water to produce hydrogen peroxide (Equation (40)) which, subsequently, reacts with *e*− to produce hydroxyl radicals (Equation (21)) [156]. The production of H2O2 following this reaction path is not significant in the presence of an organic dye due to the photo-sensitization process (Equations (38) and (39)). Therefore, O2 •− is the main primary ROS in the degradation of organic contaminants.

$$\text{H}\_2\text{H}\_2\text{O} + 2\text{h}^+ \rightarrow \text{H}\_2\text{O}\_2 + 2\text{H}^+ \tag{40}$$

In FeVO4 composites, the choice of the second component is crucial. In FeVO4/CeO2, the CBM of FeVO4 and CeO2 are −0.45 eV and −0.32 eV, respectively, and their VBM are at 1.55 eV and 2.45 eV [257]. Therefore, under UVA, the photogenerated *e*− are accumulated in the CB of CeO2 while *h<sup>+</sup>* are transferred in the VB of FeVO4, thus defining the principle of the type-II heterojunction with increase of charge carriers' lifetime (Figure 12 in Section 3.2.3). If FeVO4 is decorated by CeO2, the latter acts rather as an electron scavenger, thus also enhancing the *e*−*/h<sup>+</sup>* pair separation. Once the surface Ce (IV) trapped an electron, the reduced Ce (III) is an efficient ion for the generation of superoxide radicals from dissolved O2 [257]. Therefore, the use of CeO2 with its Ce3+/Ce4+ cycle is an efficient method to improve the photocatalytic efficiency of FeVO4 for the degradation of organic compounds. Additionally, a type-II heterojunction, direct Z-scheme like FeVO4/Fe2O3 works on the same principle: to limit the charge carrier recombination. In a direct Zscheme, the mechanism could be viewed as an "extreme" type-II heterojunction but with

the energetic level of CBM of the first component close to that of the VBM of the second component (Figure 12 in Section 3.2.3). In FeVO4/Fe2O3, under incident light of suitable energy, the photogenerated electrons and holes in the CB of FeVO4 and VB of Fe2O3 can react rapidly to generate primary ROS since the *e*<sup>−</sup> and *h<sup>+</sup>* in the CB of Fe2O3 and VB of FeVO4 are combined fastly. This direct Z-scheme mechanism was confirmed by scavenging experiments, where the main active species for the degradation of organic pollutants are identified as HO• and *h<sup>+</sup>* (Figure 20). Indeed, such results could not be observed in the case of a type-II heterojunction mechanism, where the configuration would lead to the inability to generate any primary ROS (Figure 20) [259].

**Figure 20.** The band structure of FeVO4/Fe2O3 composite and the generation of primary ROS following either a type-II heterojunction or a direct Z-scheme mechanism.

The photo-Fenton mechanism, in presence of H2O2 and FeVO4, occurs by following a two-way reaction scheme [152,257]. Indeed, the activation of hydrogen peroxide into peroxyl radicals is performed using surface iron (III) and vanadium (V) according to the Fenton-like reactions (Equations (24) and (42)). Subsequently, the reduced species, i.e., Fe (II) and V (IV) generate hydroxyl radicals (Equations (23) and (41)). These Fenton-type pathways are still the matter of discussion within the photochemical community since the use of light irradiation could play a preponderant role [152]. Indeed, since the Fenton-like reaction are slower than Fenton reactions, the recycling of iron and vanadium species is slow, thus requiring the presence of incident light which could increase the photolysis of Fe (III) and V (V) into reduced species [152,255]. The main active species during the degradation of organic contaminants using FeVO4/H2O2/light system are HO• [152,255]. In addition, the generation of hydroxyl radicals increases by increasing the amount of surface vanadium (V) [152]. It is worth noting that instead of H2O2, peroxymonosulfate (HSO5 −) could be used to generate other efficient radicals for the degradation of organic pollutants: the sulfate radicals (SO4 •−) [255]. The production of the sulfate radicals (Equation (43)) could also lead to the generation of hydroxyl radicals by reaction with hydroxide anions (Equation (44)). In addition, singlet oxygen which is an excellent oxidant could be formed in such a photochemical system (Equations (45) and (46)) [255]. From these reactions, it is clear that there is a synergy with the photocatalytic process since photogenerated electrons are involved in several crucial reaction steps. In the form of a composite, especially FeVO4/CeO2, a third reaction pathway appears [152,257]. Indeed, surface Ce (IV) in CeO2 can activate Fenton-type processes (Equations (29) and (30)), thus generating hydroxyl radicals [257].

$$\mathbf{V} \equiv \mathbf{V}^{\mathrm{IV}} + \mathbf{H}\_2 \mathbf{O}\_2 \to \equiv \mathbf{V}^{\mathrm{V}} + \mathbf{O} \mathbf{H}^- + \mathbf{H} \mathbf{O}^\bullet \tag{41}$$

$$\equiv \text{V}^{\text{V}} + \text{H}\_{2}\text{O}\_{2} \rightarrow \equiv \text{V}^{\text{IV}} + \text{H}^{+} + \text{HCOO}^{\bullet} \tag{42}$$

$$\rm{Fe^{2+}/V^{4+} + HSO\_5^- \to SO\_4^{\bullet-} + OH^- + Fe^{3+}/V^{5+}} \tag{43}$$

$$\rm SO\_4{}^{\bullet-} + \rm OH^- \rightarrow SO\_4{}^{2-} + \rm HO^\bullet \tag{44}$$

$$\text{HSO}\_5^- + \text{SO}\_5^{\bullet -} + e^- \rightarrow \text{HSO}\_4^- + \text{SO}\_4^{2-} + ^1\text{O}\_2 \tag{45}$$

$$2\text{SO}\_5^{\bullet-} + \text{H}\_2\text{O} \rightarrow 3/2^1\text{O}\_2 + 2\text{HSO}\_4^- \tag{46}$$

Regarding the combined photo-induced processes in FeVO4 based systems, most of the published works referred to addition of H2O2 which requires the use of chemical and potentially increases the cost of implementation (or scale up) of such a photo-induced process for the removal of organic contaminants. Therefore, a key point in the research of simultaneously combined photocatalytic and photo-Fenton processes is to focus on in situ generation of H2O2 and to study the contribution of each involved process and their potential synergy.

On the other hand, other iron vanadates are gaining stronger interest in the photochemical community. It is the case of Fe2V4O13 which is a visible light driven photocatalyst with an energy band gap in the range of 2.2–2.8 eV [260,261]. The energy levels of conduction band minimum and valence band maximum are −2.05 eV and 0.15 eV, respectively [261]. Therefore, the photoctalytic reduction ability of Fe2V4O13 is relatively strong. Indeed, this iron vanadate is able to remove Cr (VI) by reduction reaction using photogenerated electrons (Equation (47)) [260]. The conversion extent of Cr (VI) into non-toxic Cr (III) reaches up to 98% which is significantly higher than that in presence of TiO2 (about 40%). The performance of Fe2V4O13 can be also ascribed to its higher specific surface area [260,261].

$$\rm{Cr\_2O\_7}^{2-} + 14H^+ + 6e^- \rightarrow 2Cr^{3+} + 7H\_2O \tag{47}$$

In the form of composite, Fe2V4O13/ZnO can degrade 90% of methyl orange under simulated solar light [261]. This enhanced removal extent compared to pristine Fe2V4O13 is due to the formation of a type-II heterojunction where, under light irradiation, the photogenerated electrons are accumulated in the conduction band of ZnO while *h+* are transferred into the valence band of Fe2V4O13. Since the CBM of ZnO in this system is at −0.35 eV, the production of superoxide radicals is favored. However, the VBM are not enough positive to allow the production of hydroxyl radical by reaction with the photogenerated holes. Furthermore, the *h<sup>+</sup>* has weak oxidation ability to directly degrade the organic dye and, therefore, the main primary ROS are O2 •−.

Fe2V4O13 can also act as a multi-Fenton catalyst. Similar to FeVO4, a two way reaction scheme occurs in the presence of H2O2 which is activated into reactive oxygen species as described above (Equations (23), (24), (41), and (42)). Therefore, Fe2V4O13 can degrade more than 90% of methyl orange associated with a mineralization extent of 40% [159]. The mechanism involves the cycle of redox couples Fe(III)/Fe(II) and V(V)/V(IV). By comparison with FeVO4/V2O5 and Fe2O3/V2O5 systems where two compounds are combined, the best results are achieved using Fe2V4O13, thus highlighting a synergy in the Fenton activity induced by iron and vanadium species [159]. Fe2V4O13 has been also tested in the degradation of real wastewater effluents spiked by organic dyes. As expected, a decrease of efficiency is observed compared to synthetic wastewater (removal extent < 50%) due to the consumption of H2O2 by the organic matter which is present in higher extents [159].

Finally, another interesting iron vanadate is FeV3O8. This materials is a semiconductor which is a promising visible light driven photocatalyst (*Eg* = 2.23 eV) [163]. Indeed, depending on its morphology, the specific surface area can reach more than 150 m2·g−<sup>1</sup> (for ultrathin nanostructured material), thus giving rise to good adsorption properties. In addition, under visible light, FeV3O8 can remove up to 92% of methyl orange after only 30 min irradiation time [163]. In such a photoactive system, the main primary ROS are identified as both superoxide and hydroxyl radicals, but further studies are required to confirm this observation since the exact role of photogenerated holes are not clearly elucidated.

Other Transition Metal Vanadates

Among the other transition metal vanadates, the most promising ones are summarized in the Table 2. These ternary oxides are semiconductor photocatalysts which could be activated either under visible or UVA light. Although they are promising materials, their use in photo-induced processes for the removal of organic contaminants are far less documented than bismuth and iron vanadates. In addition, these atypic vanadates are often coupled with another oxide to form a composite. Therefore, deeper research is necessary to elucidated the photocatalytic mechanism (of both pure and composite materials) and to highlight potentially other photo-induced processes like the Fenton-type reactions which probably occur in metal vanadates based on Ce, Cu, and Mn.

**Table 2.** Overview of different transition metal vanadate in the photo-induced degradation of organic pollutants.


#### **5. Conclusions**

Binary and ternary vanadium oxides in single and mixed valence states are promising materials for environmental applications. Based on the versatile behaviour of vanadium, the electronic and optical properties of semiconducting vanadium oxides can be tuned, thus giving rise to interesting photo-induced properties. Therefore, vanadium oxide-based materials can trigger photochemical processes for the environmental purposes including the splitting of water for hydrogen production and the photocatalytic and photo-Fenton processes for the removal of organic pollutants in water. Such problematics are highly relevant for the current and next generations. This review article provides a useful tool for scientific research and pedagogical knowledge where the comprehensive description of the structure and intrinsic properties of vanadium oxides are related to their photochemical efficiency in environmental processes.

**Author Contributions:** Conceptualization, writing: original draft preparation, writing: review and editing, supervision, O.M.; writing: original draft preparation, writing: review and editing, P.P. All authors have read and agreed to the published version of the manuscript.

**Funding:** This research received no external funding.

**Institutional Review Board Statement:** Not applicable.

**Informed Consent Statement:** Not applicable.

**Data Availability Statement:** No new data were created or analyzed in this study. Data sharing is not applicable to this article.

**Conflicts of Interest:** The authors declare no conflict of interest.

#### **References**


## *Article* **Degradation of Direct Blue 1 through Heterogeneous Photocatalysis with TiO2 Irradiated with E-Beam**

#### **Elvia Gallegos 1,2, Florinella Muñoz Bisesti 3, Katherine Vaca-Escobar 1, Cristian Santacruz 4, Lenys Fernández 1, Alexis Debut <sup>5</sup> and Patricio J. Espinoza-Montero 1,\***


Received: 1 April 2020; Accepted: 7 September 2020; Published: 18 September 2020

**Abstract:** Most dyes used in the textile industry are chemically stable and poorly biodegradable, therefore, they are persistent in the environment and difficult to degrade by conventional methods. An alternative treatment for this kind of substance is heterogeneous photocatalysis using TiO2, so, in this work, it is proposed to degrade Direct Blue 1 (DB1) using microparticulate TiO2 irradiated with e-beam at three different doses: 5, 10 and 20 kGy (J/kg). The DB1 degradation was implemented in a batch reactor (DB1 initial concentration = 50 mg L−1, pH 2.5, TiO2 concentration = 200 mg L−1). We have demonstrated that the photocatalytic power of TiO2, when irradiated with e-beam (5, 10, 20 kGy), varies slightly, with minor effects on photodegradation performance. However, the dose of 10 kGy showed a slightly better result, according to the DB1 photodegradation rate constant. Adsorption process was not affected by irradiation; its isotherm was fitted to Freundlich's mathematical model. The DB1 photodegradation rate constants, after one hour of treatment, were: 0.0661 and 0.0742 min−<sup>1</sup> for irradiated (10 kGy) and nonirradiated TiO2, respectively. The degradation rate constant has an increase of 12.3% for irradiated TiO2. Finally, there was no evidence of mineralization in the degradation process after 60 min of treatment. According to the results, the irradiation of microparticulate TiO2 with e-beam (10 kGy) slightly improves the photodegradation rate constant of DB1.

**Keywords:** Direct Blue 1; irradiated TiO2; heterogeneous photocatalysis; electron-beam; adsorption

#### **1. Introduction**

Effluents from the textile industry are considered an important environmental problem because it is difficult to treat due to their high nonbiodegradable load, partially attributed to the dye presence. At present, 100,000 different types of dyes are produced, of which about 36,000 tons dye/year are consumed by the textile industries [1]. The dyes are used in the dyeing stage, where up to 30% of them come off, becoming part of the industrial discharge [2].

Among the most used dyes are those of the azo type, which constitute around 70% of world production [3], due to their ease of preparation and the wide range of colors in which they are offered. These dyes have chromophore groups formed by double bonds –N=N– and have high stability to light

and washing. Most of them are sulfonic acids or their salts [4]. Its high stability determines properties of difficult degradation by conventional treatment methods; therefore, they are an important source of chemical and esthetic contamination [2]. Among the environmental problems caused by azo dyes are: color in effluents, visual pollution and eutrophication of water bodies, effects on gases solubility in water [5], generation of toxic, carcinogenic or mutagenic amines under specific conditions [6] and alterations in soil microbial biological activity of populations when they have been subjected to irrigation [7], among others.

One of the azo dyes used in the staining of denim fabrics in Ecuador, particularly in Pelileo (city located in the country center), is Direct Blue 1 (DB1), whose structure is shown in Figure 1, and it has two azo groups in its composition, it has a molecular weight of 992.82 g mol−1, which resists temperatures of up to 300 ◦C, and is difficult to degrade by conventional methods.

**Figure 1.** Direct Blue 1 (DB1) structure [8].

As an alternative treatment for compounds of difficult degradation, there are Advanced Oxidation Processes (AOPs), which are characterized by the oxidation reactions promoted by the hydroxyl radical (•OH) [9,10]. AOPs have been extensively studied for the oxidation of persistent compounds in water [11], such as dyes. Its advantages are simple handling, the relatively low cost of the reagents [9], as well as the oxidation of organic matter, even reaching its total mineralization (CO2, H2O and inorganic salts) [1]. A brief summary of some AOPs and other methodologies used for DB1 degradation are shown in Table 1:



#### *Processes* **2020**, *8*, 1181

A widely used AOP is heterogeneous photocatalysis with TiO2, which consists of an electron-hole pair (e−-h+) generation, by excitation of the electron from a semiconductor through UV radiation. The holes oxidize the hydroxyl ion in water to generate hydroxyl radicals (•OH), and electrons are transferred to adsorbed molecular oxygen to generate a superoxide ion. The •OH generation allows the degradation of the target pollutant [10].

TiO2 is a cheap, chemically stable, and nontoxic semiconductor. Several studies have been developed to improve its photocatalytic activity [28], for instance: doping with metal ions, doping with nonmetals, formation of compounds with TiO2 and with less energetic band gap semiconductors or with noble gases, dye sensitization, or rare earth modifications [10]. Some studies have applied ionizing radiation, particularly accelerated electrons (e-beam), to TiO2 nanoparticles, improving their photocatalytic activity [8,9]. There are few studies on this last modification, therefore, in this work the photocatalytic degradation of the DB1 dye is studied using a TiO2 microparticulate irradiated with accelerated electrons.

#### **2. Materials and Methods**

#### *2.1. Reagents and Equipment*

TiO2 Hombitan (commercial grade, 325 μm microparticulate) previously sieved, DB1 Dyetex, Quito-Ecuador and H2SO4 1 N (Merck, Quito-Ecuador) were used. The equipment used was: ELU6U linear electron accelerator with 8 MeV power, EMPIREAN model PANanalytical diffractometer, Shimadzu TOC-5 total organic carbon meter, Hitachi U 1900, Tokio-Japan spectrophotometer, HACH 2800, Loveland-USA spectrophotometer, LABEC magnetic stirrer VELP, 15 W Sylvania Mercury.

#### *2.2. TiO2 Irradiation*

A total of 2 g of TiO2 Hombitan by Venator Materials was placed in LDPE low-density polyethylene bags, and those bags were irradiated in a linear electron accelerator (ELU6U) at three doses: 5, 10 and 20 kGy (J/kg) [29,30].

#### *2.3. TiO2 Characterization*

TiO2 was characterized by X-Ray Diffraction (XRD), operating in a θ-2θ configuration (Bragg–Brentano geometry) equipped with a copper X-ray tube (Kα-λ = 1.54056 Å) at 45 kV and 40 mA. The diffractograms analysis was performed on the average of four measurements between the angles 5◦ to 90◦ (θ-2θ) using the High Score Plus software.

Diffuse Reflectance Spectroscopy (DRS) was obtained at room temperature using a UV–VIS spectrometer equipped with an integrating sphere (Perkin Elmer, Lambda 365). The energy band gap of TiO2 samples was determined by extrapolating the slope to R→0 in the plot (*R*·*h*υ) *<sup>n</sup>* versus *h*υ, where R is the reflectance, *h*υ is the photon energy and n = 2 corresponds to direct allowed transitions [31–34]. The best fit to a straight line near the absorption edge was used to determine the energy band gap.

Infrared absorption of TiO2 powder samples was obtained by measuring the Fourier transform infrared (FTIR) in a Jasco FT/IR 4700 spectrometer. Samples were analyzed at room temperature in a region of 4000–400 cm−1. Spectral resolution was 1 cm−1. Adsorbed water was removed from TiO2 powder by heating the samples to 50 ◦C for 24 h. TiO2 powder was mixed with KBr and pressed into pellets.

Finally, elemental analysis by Energy-Dispersive X-ray Spectroscopy (EDS) was performed in a RESCAN MIRA 3 scanning electron microscope.

#### *2.4. DB1 Degradation Tests with Irradiated TiO2*

The photocatalysis process was implemented in a batch reactor equipped with two Sylvania 15 W mercury lamps, with a working wavelength of 254 nm (UV region), two acrylic chambers and a water-cooling system (Figure 2), which allowed for keeping the temperature constant.

**Figure 2.** Diagram of the photoreactor used in the degradation of the DB1 dye. (1) Water inlet, (2) UV lamps, (3) Samples, (4) Magnetic stirrers, (5) Water outlet.

The effluent used was water with an initial concentration of 50 mg L−<sup>1</sup> of DB1 dye, and 200 mg L−<sup>1</sup> of both irradiated and nonirradiated TiO2 [16,17]. Then, ultrasound was applied for 10 min to the mixture.

Photodegradation was performed for 60 min, with stirring at 200 rpm [35] using a magnetic stirrer, Figure 2, and the change in pH was also followed by taking measurements at the beginning and at the end of the process. To verify degradation, aliquots were taken at different times, which were centrifuged (6000 rpm) to separate the catalyst from the sample. The sample was filtered through a Whatman filter with a pore diameter of 0.45 μm. The DB1 concentration decrease was followed by spectrophotometry, at a wavelength of 565 nm [18]. The concentrations were obtained from a previously prepared calibration curve.

The DB1 degradation reaction constant was determined after one hour of photodegradation by fitting the experimental data to a kinetic model.

The irradiation dose effect on reaction rate was analyzed using a completely randomized design (CRD), the independent variable was irradiation dose and the response variable was the reaction constant. Statistical analyzes were developed in STATGRAPHICS Centurion XVIII software. The experiments were made in triplicate.

#### *2.5. Adsorption Kinetics*

DB1 adsorption in nonirradiated and irradiated TiO2 was studied for different doses (5, 10 and 20 kGy). This study was made using the Figure 2 reactor, in the same photodegradation conditions (Co, DB1 = 50 mg L<sup>−</sup>1; catalyst dose = 200 mg L<sup>−</sup>1; pHo = 2.5; Tamb) in light absence. Samples were taken every 10 min, during 50 min and were analyzed by spectrophotometry. With the experimental data of DB1 concentration, DB1 adsorption kinetics on TiO2 were determined, with which the adsorption equilibrium time was established.

#### *2.6. Adsorption Isotherms*

Solutions with different DB1 concentrations, in a range of 25 to 250 mg L−1, were prepared. The conditions established in the previous section were used (Catalyst dose = 200 mg L<sup>−</sup>1; pH = 2.5; Tamb). Samples were taken after reaching the equilibrium time, previously determined, these samples were filtered through 0.45 μm filters and measured by UV-vis spectrophotometry.

Experimental data were analyzed to identify if they fit Langmuir or Freundlich mathematical models, and the irradiation dose influence in adsorption onto TiO2 was studied using a CRD, in which the design variable was the irradiation dose (0 kGy and best dose) and the response variable was the best adjusted kinetic constant. Statistical analyzes were done in STATGRAPHICS Centurion XVIII software. The samples were taken in triplicate.

#### *2.7. Contact Angles*

The static contact angles were measured following the procedure described by [31–35]. Dried titanium dioxide samples were stored at ambient conditions before measurement. TiO2 powder was sprinkled uniformly on a 2 × 2 cm clean flat surface. A glass microscope slide (20 × 10 × 2 mm) was covered with double-sided adhesive tape. The tape was pressed to the powder with a 500 g weight for 10 s. The excess of powder was removed by using a razor blade. Finally, the slide was shaken to remove nonadhered material. This preparation technique produces a uniform powder surface. Measurements were taken immediately at ambient conditions, 20 ◦C and 60% relative humidity. For the static contact angle, an 8 μL droplet of deionized water was placed on the slide, and within 30 s three contact angle readings were measured (Dataphysics, OCA 25 equipped with electronic syringe module ESr-N). Each measurement was taken on three different positions of the slide and repeated on three slides. The average value is reported.

#### *2.8. Mineralization Levels through TOC*

Dye mineralization levels were measured at the beginning and ending of the process, through the determination of Total Organic Carbon (TOC), using method 5310 of the Standard Methods for the Examination of Water and Wastewater.

#### **3. Results**

#### *3.1. X-ray Di*ff*raction, Band Gap by DRS and FTIR Spectroscopy*

Figure 3 show the two samples analyzed patterns (TiO2 nonirradiated and TiO2 irradiated).

**Figure 3.** TiO2 X-ray diffraction with and without irradiation.

Figure 4 shows the FTIR spectra of TiO2 samples, with and without e-beam treatment measured under ambient pressure conditions at the same humidity.

**Figure 4.** FT-IR spectra of TiO2 with and without e-beam treatment measured under ambient conditions. Inset, a magnified view of the spectrum of the samples.

Band gap measurements were performed using DRS to determine if there were any changes in the optical properties of TiO2 caused by e-beam treatment (5, 10, and 20 kGy) (Figure 5).

**Figure 5.** Diffuse Reflectance Spectroscopy (DRS) spectra of TiO2 with and without e-beam treatment are displayed. Inset, the band gaps estimated from the DRS spectra are summarized.

In order to verify the presence of C in the samples of TiO2, EDS analysis was performed. Figure 6 shows an analysis based on mapping that identifies the elements that conform the sample. It is noted that the elements of the sample are O, Ti and C. The presence of C is attributed to the impurities in the sample. The averages of each element were 48.74%, 48.26%, and 3.00%, respectively.

**Figure 6.** Energy-Dispersive X-ray Spectroscopy (EDS) mapping of composition of sample of TiO2 without irradiation, and EDS spectrum of irradiated and unirradiated TiO2.

#### *3.2. Heterogeneous Photocatalysis*

The curves obtained from DB1 dye degradation by heterogeneous photocatalysis with TiO2 irradiated at various doses and nonirradiated are shown in Figure 7.

**Figure 7.** DB1 photodegradation on TiO2 with and without e-beam treatment is displayed. Conditions: Co, DB1 <sup>=</sup> 50 mg L<sup>−</sup>1; catalyst dose <sup>=</sup> 200 mg L<sup>−</sup>1; pHo <sup>=</sup> 2,5; Tamb <sup>≈</sup> <sup>20</sup> ◦C.

To determine the dye degradation reaction constants, a linear fit was performed as shown in Figure 8. The experimental data were adapted to a pseudo first order kinetic model which is consistent with what was reported in other studies [16,36]. Results are shown in Table 2. Subsequently, an ANOVA was performed (Figure 9) in order to analyze dye degradation constant rates.

**Figure 8.** DB1 linear adjustment according to pseudo first order kinetics, data from Figure 7.

**Table 2.** Rate constants for pseudo first order reaction at different irradiation doses.

**Figure 9.** Statistical analysis of irradiation dose effect on the photodegradation of DB1. Conditions: Co, DB1 <sup>=</sup> 50 mg L<sup>−</sup>1; catalyst dose <sup>=</sup> 200 mg L<sup>−</sup>1; pHo <sup>=</sup> 2.5; Tamb <sup>≈</sup> <sup>20</sup> ◦C.

*3.3. Isotherms of Adsorption*

Adsorption isotherms obtained were "S" type, they can be observed in Figure 10:

**Figure 10.** Adsorption DB1 dye isotherms of unirradiated and irradiated TiO2 at different doses (5, 10 y 20 kGy).

With increasing concentrations, isotherms took a trend that was later adapted to the Freundlich mathematical model (Equation (1)) with high correlations (Table 3). With the experimental data from this linearization, adsorption constant Kf was obtained for each adsorption dose. Their average values are detailed in Table 3. In addition, a statistical analysis of the Kf was performed using CRD as shown in Figure 11.

$$\log(\mathcal{Q}\_{\mathbf{c}}) = \log(\mathcal{K}\_{\mathbf{f}}) + \frac{1}{\mathbf{n}} \log(\mathcal{C}\_{\mathbf{c}}) \tag{1}$$


**Table 3.** Freundlich isotherm parameters.

**Figure 11.** Statistical analysis of irradiation dose effect on DB1 adsorption. (CO, DB1 = 50 mg L<sup>−</sup>1; catalyst dose = 200 mg L<sup>−</sup>1; pHo = 2.5; Tamb).

#### *3.4. Mineralization*

In Table 4, the percentages of DB1 mineralization levels achieved after one hour of treatment are presented.

**Table 4.** % Total Organic Carbon (TOC) removal during DB1 degradation after 60 min of treatment.


#### *3.5. Contact Angles*

The contact angle by sessile drop of the nonirradiated and irradiated TiO2 samples are shown in Figure 12 (insert).

**Figure 12.** Contact angle by sessile drop without drying the sample (ambient), of the nonirradiated and irradiated TiO2. Insert, contact angle.

#### **4. Discussion**

#### *4.1. X-ray Di*ff*raction, Band Gap by DRS and FTIR Spectroscopy*

Figure 3 shows TiO2 X-ray diffraction patterns with and without irradiation. The main features of the XRD pattern of the two samples are similar, peaks at 25 and 48 degrees, which are anatase phase characteristic peaks [37], and it shows that the crystalline phase used for the photocatalyst when it was irradiated with accelerated electrons did not change. This result is in agreement with the ones found in other studies. A study has pointed out that there were no significant changes in the TiO2 phase composition and crystal lattice when it was subjected to gamma irradiation at different doses (30–150 kGy) [38]. Another study shows that there were no changes in TiO2 when it was irradiated with e-beam with an energy of 5 MeV and a current of 200 μA [39]. Also, the peaks in Figure 3 are in good agreement with the standard spectrum and XRD patterns were similar, indicating that there was no change in the crystallinity of TiO2 caused by the e-beam treatment [39,40]. However, reports that changes in structure and improvement in photocatalytic activity were obtained at highest doses (5 MGy) [41].

The FTIR analysis results of the four TiO2 samples with and without e-beam treatment, obtained under similar environmental conditions (same humidity), are shown in Figure 4. The samples treated at 5 and 10 kGy clearly show an increased signal corresponding to water. Signals at 3300, 1600 and 900 cm−<sup>1</sup> can be attributed to the OH-stretching, bending and vibrational water modes, respectively [42–44]. It is difficult to discriminate OH on C groups in IR spectra, because the C-O-H stretch overlaps the OH stretch of water adsorbed on the surface. However, in the electron beam-treated samples, especially at 5 kGy, the peak of the C–O stretching at 1400 cm−<sup>1</sup> is more intense, and this does not happen with other samples. This result was confirmed by the contact angle (Figure 12), where the sample treated at 5 kGy has the lowest contact angle (that is the highest wettability). In line with these results, it can be confirmed that the amount of water adsorbed on the TiO2 surfaces increased after e-beam treatment, i.e., the TiO2 surfaces became more hydrophilic [29]. Therefore, partial oxidation of C impurities on TiO2 could be responsible for the slight improvement of photocatalytic activity. While, the TiO2 reticular vibration signal can be assigned to the band between 600–700 cm−<sup>1</sup> [44]. The e-beam treatment had almost no influence on the geometric and electronic structures of TiO2.

To corroborate the presence of residual carbon in TiO2, from its synthesis, studies of EDS of the samples treated and not treated with e-beam were carried out. According to the results found, the percentage of carbon decreased from 3.00% for the untreated sample to 0.50 % for the sample treated at 10 kGy (Figure 6). This leads us to presume that part of the impure C is released from the surface of TiO2 with radiation, however this does not occur with the carbon that is within the TiO2 crystalline structure. Apparently, the TiO2 treatment with e-beam causes the carbon to acquire different oxidation states, which influences the photocatalytic behavior of TiO2.

Diffuse Reflectance Spectroscopy (DRS) was used to determine band gap, which was obtained by the Kubelka–Munk function. It was plotted as a wavelength function, where the intercept on the x-axis (energy) of the linear extrapolation of the Kubelka–Munk function corresponds to the material's band-gap (Figure 5) [45]. The energy band gap for all TiO2 samples, irradiated and unirradiated, was around 3.25 eV (Insert of the Figure 5) which corresponds to reported value of anatase [34]. The value of the band gap does not change significantly up to irradiation, indicating that the change in optical properties cannot be the reason for the slight variation in photocatalytic activity. These results are consistent with those reported in the literature [29].

Furthermore, XPS studies reported by Kim M. J. et al., for the TiO2 surface, show that Ti 2p spectra for untreated and e-beam-treated TiO2 samples (5, 10, and 15 kGy) show no structural changes. In contrast, the C 1s spectrum was significantly influenced by the e-beam treatment, i.e., the existing C as an impurity underwent a noticeable structural change in front of the e-beam. Finally, it has been shown that when 5 and 10 kGy e-beam doses were used to irradiate TiO2 samples, the absolute intensity of the C 1s state increases with respect to the untreated sample. The center of the C 1s peak shifts to a higher binding energy after e-beam treatment. In the case of e-beam exposure at 15 kGy, the intensity of the C 1s peak decreases. The change in C 1s state after e-beam treatment is due to the formation of C-O bonds, causing the TiO2 surface to become more hydrophilic [29,46]. The literature is complemented by the results shown in this paper, with regard to contact angle, band gap, EDS and FTIR.

In the case of TiO2 irradiated at 20 kGy, C 1s is removed, provoking a slightly decrease in the FTIR signals intensity and an increase in the contact angle (reduces hydrophilicity). The decrease of C 1s is due to the elimination of C by oxidizing agents created in the atmosphere by the e-beams and may become more dominant than the C deposition from the atmosphere, or oxidation of existing ones [29].

#### *4.2. E*ff*ect of Irradiation on Heterogeneous Photocatalysis*

According to Figure 7, the degradation of the DB1 in the first 5 min is slightly faster in the case of TiO2 irradiated at 10 kGy, however, after this time the degradation of the dye is slightly faster with nonirradiated TiO2. Unexpectedly, after 50 min of treatment, there is again a breaking point, where the irradiated TiO2 becomes lighter faster. Therefore, a statistical analysis of the degradation rate constant is convenient, and it is shown later. A degradation of around 98% is reached for all cases at 60 min of treatment.

A linear fit was performed to determine the dye degradation rate constants, as shown in Figure 8. The experimental data were adapted to a pseudo first order kinetic model. The rate constant averages for three replications are detailed in Table 2. Subsequently, an ANOVA was performed, Figure 9, in which we determined that there is a statistically significant difference (*p* < 0.05) in the reaction rate constants. This analysis indicates that the 10 kGy dose is the one that presents an improvement in the reaction rate constant, which is due an increase in the concentration of C-O bonds coming from impurities on the surface of TiO2 [29].

#### *4.3. E*ff*ect of Irradiation Dose on Adsorption*

The adsorption isotherms obtained were "S" type; that is, at low concentrations of adsorbate (DB1), the surface (TiO2) presents collateral associations, a process known as "cooperative adsorption" [47]. This kind of adsorption behavior implies a moderate intermolecular attraction between the adsorbate and the adsorbent, where there is a strong competition for the adsorption centers with solvent molecules or other adsorbates [48]. This process is identified in the isotherm as a turning point in the low concentration zone (Figure 10), as the adsorbate concentration increases, it presents better affinity

towards the surface, and can be seen as an increase in the curve slope until saturation of available sites is reached (maximum adsorption).

With increasing concentrations, isotherms take a trend that is later adapted to the Freundlich mathematical model with high correlations, Table 3. With the experimental data from this linearization, the adsorption constant Kf was obtained for each adsorption dose; this constant is an approximate indicator of the maximum adsorption capacity on TiO2 [49]. Their average values are detailed in Table 3.

Using a CRD, in which the independent variable was the irradiation dose and the response variable was the adsorption constant, it was determined that there was a statistically significant difference between these values (*p* < 0.05). Figure 11 reflects that there is no a statistically significant difference between 0 and 10 kGy doses, nor between 5 and 20 kGy doses, but there is a statistically significant difference between 0 and 5 kGy doses, 0 and 20 kGy doses, 10 and 5 kGy doses, and between 10 and 20 kGy doses.

Particle size is a determining factor for the adsorption process, since it is related to the specific surface area, while the particle size decreases, the specific surface area increases and therefore, the number of active sites [50]. The TiO2 particle size used in this work was around 325 μm (considered a micrometric size), which means that it could improve the adsorption capacity if the particle size decreases.

Nevertheless, an excessive adsorption might block the way of the light towards the nanoparticle and provoke a decrease in the photocatalytic efficiency, especially when there are big molecules such as DB1 [2]. This can be confirmed with the experimental results of Kf, for TiO2 without irradiation and TiO2 irradiated at 10 kGy, where the lowest value of Kf is obtained.

#### *4.4. Contac Angle*

Contact angle refers to the angle that the surface of a liquid makes when it encounters a solid. The value of the contact angle depends mainly on the relationship between the adhesive forces between the liquid and the solid and the cohesive forces of the liquid. When the adhesive forces with the surface of the solid are very large in relation to the cohesive forces, the contact angle is less than 90 degrees, resulting in the liquid wetting the surface. Contact angle measurement is useful in the evaluation of surface macroscopic properties, such as surface energy and wettability [51].

The variation of the contact angle between irradiated and nonirradiated TiO2 samples is small. The e-beam produces a change on the surface that affects the contact angle variation, which goes down from 131.6◦ (TiO2, 0 kGy) to 130.8◦ (TiO2, 10 kGy), Figure 12; this slight decrease can be attributed to the increase in C-OH, as discussed above, which causes a slight increase of the wettability. That is likely related to the slight increase of the photocatalytic activity of the TiO2 irradiated at 10 kGy. However, studies show that to achieve a change in the oxidation state of Ti, an increase in the Ti3+/Ti4<sup>+</sup> ratio, and irradiation with an MeV electron beam is required [52]. There are some studies that compares the photocatalytic activity and the super-hydrophilicity of TiO2 induced by UV irradiation. The results are inconclusive, however, the composition and the processing of the material increases the photocatalytic activity and reduces the super-hydrophilic character, or vice versa [53–55]. According to Yu et al., the mechanism of photoinduced super-hydrophilicity of the films is different from that of photocatalytic oxidation, super-hydrophilicity is attributed to formation of surface trapped electrons and Ti3<sup>+</sup> defective sites [56]. All these studies show that the photoinduced super-hydrophilicity is not permanent and reduces when the sample is stored in the dark for a couple of days.

#### *4.5. Mineralization*

According to Table 4, the mineralization after 60 min of treatment with TiO2, not irradiated and irradiated at 10 kGy, is very low. However, it is important to note that the interest of this research was not to achieve total mineralization, but to study the effect of irradiation on the degradation of the DB1. However, in a study developed with azo dyes, it was determined that total mineralization occurs at longer times, around 5 h [57], which means that this low percentage could be caused by the time taken

to mineralize the dye. Other studies attribute the low TOC to structure cleavage and to the production of organic molecular small fragments that are not fully mineralized under the prevailing oxidative conditions [58,59].

#### **5. Conclusions**

TiO2 irradiated at 10 kGy allows a slightly faster degradation of DB1, compared to the degradation with TiO2 irradiated at 0, 5, and 20 kGy; in this sense, the rate constant of degradation of DB1, on TiO2 irradiated at 10 kGy, is slightly higher statistically, according to ANOVA analysis. The results of XRD and band gap are identical for the irradiated and unirradiated TiO2 with e-beam, showing that e-beam irradiation show no changes in the crystalline structure of the TiO2 samples. However, EDS and XPS results from similar e-beam irradiation show that its surface properties suffer slight modifications, and the residual carbon on the TiO2 structure is oxidized by irradiation, forming oxygenated species (C-O), which causes an increase in the TiO2 hydrophilicity. This was corroborated experimentally by decreasing the contact angle, increasing the FT-IR (signals to the C-OH). In the case of TiO2 irradiated at 20 kGy, C-OH groups are lost; causing a decrease in the intensity of the corresponding signals in FTIR intensity and an increase in the contact angle (reduces wettability). Regarding the absorption, the TiO2 irradiated at 10 kGy absorbs less than the other materials used (TiO2 irradiated at 0, 5 and 20 kGy); this apparently slightly favors the degradation of DB1, since an excessive adsorption would prevent the adequate passage of light towards the photocatalyst, causing a decrease in the speed of the degradation.

**Author Contributions:** Conceptualization, P.J.E.-M. and F.M.B.; methodology, P.J.E.-M. and F.M.B.; software, K.V.-E., A.D. and C.S.; validation, E.G., F.M. and P.J.E.-M.; formal analysis, E.G., C.S. and A.D.; investigation, E.G.; resources, F.M.B. and P.J.E.-M.; data curation, K.V.-E.; writing—original draft preparation, E.G., P.J.E.-M., L.F. and F.M.B.; writing—review and editing, E.G., P.J.E.-M., K.V.-E., L.F. and F.M.B.; visualization, P.J.E.-M. and F.M.B.; supervision, P.J.E.-M. and F.M.B.; project administration, P.J.E.-M. and F.M.B.; funding acquisition, P.J.E.-M. and F.M.B. All authors have read and agreed to the published version of the manuscript.

**Funding:** This research was funded by ESCUELA POLITÉCNICA NACIONAL, grant number PIMI-14-13.

**Acknowledgments:** The authors wish to thank the Laboratorio de Investigaciones Aplicadas de la Escuela Politécnica Nacional for the support receive during this research and the Centro de Nanociencia y Nanotecnología, Universidad de las Fuerzas Armadas ESPE for the X-ray and EDS analysis.

**Conflicts of Interest:** The authors declare no conflict of interest.

#### **References**


© 2020 by the authors. Licensee MDPI, Basel, Switzerland. This article is an open access article distributed under the terms and conditions of the Creative Commons Attribution (CC BY) license (http://creativecommons.org/licenses/by/4.0/).

## *Article* **Visible Light-Driven Photocatalytic Rhodamine B Degradation Using CdS Nanorods**

**Haseeb Ullah 1,\*, Eva Viglašová <sup>2</sup> and Michal Galamboš 2,\***


**Abstract:** In this work, highly crystalline CdS nanorods (NRs) were successfully synthesized by a facile, one-step solvothermal method. The as-prepared CdS NRs powder was characterized by XRD, FESEM, Raman, PL, XPS, BET, and UV-visible techniques to evaluate the structural, morphological, and optical properties. The photocatalytic performance of the as-synthesized CdS NRs was investigated for the photodegradation of RhB dye under visible light irradiations. It has been found that CdS NRs show maximum RhB degradation efficiency of 88.4% in 120 min. The excellent photodegradation ability of the CdS NRs can be attributed to their rod-like structure together with their large surface area and surface state. The kinetic study indicated that the photodegradation process was best described by the pseudo-first-order kinetic model. The possible mechanism for the photodegradation of RhB dye over CdS NRs was proposed in this paper.

**Keywords:** CdS; rhodamine B; photocatalysis; pseudo-first order kinetics

#### **Citation:** Ullah, H.; Viglašová, E.; Galamboš, M. Visible Light-Driven Photocatalytic Rhodamine B Degradation Using CdS Nanorods. *Processes* **2021**, *9*, 263. https:// doi.org/10.3390/pr9020263

Academic Editor: Olivier Monfort Received: 30 November 2020 Accepted: 24 January 2021 Published: 29 January 2021

**Publisher's Note:** MDPI stays neutral with regard to jurisdictional claims in published maps and institutional affiliations.

**Copyright:** © 2021 by the authors. Licensee MDPI, Basel, Switzerland. This article is an open access article distributed under the terms and conditions of the Creative Commons Attribution (CC BY) license (https:// creativecommons.org/licenses/by/ 4.0/).

#### **1. Introduction**

In recent years, the ever-increasing global population, unplanned urbanization, and rapid industrialization have caused enormous environmental pollution, becoming an increasingly serious threat to the survival of human beings [1,2]. Organic dyes are considered a major class of environmental hazards, extensively produced by some industries such as the manufacture of plastics, papers, textile, food, cosmetic, drugs, and printing industries [3–5]. Most of the dyes have xenobiotic properties and complex aromatic structures and are resistant to degradation. Inhalation of organic dyes can cause eye burns, profuse sweating, vomiting, nausea, and mental confusion [6,7]. Nowadays, around 100,000 different types of synthetic dyes are produced with a production rate of 7 × 105 tons per year, and around 15% of the dyes are released into the natural ambiance [8]. The presence of dyes in waterbodies, even in a small amount, is a very concerning problem, as it not only affects the aesthetic nature but also reduces the photosynthetic activity of aquatic plants and algae in aqueous environments [9].

Various treatment methods, e.g., advanced oxidation processes [10], ion flotation [11], ozonation [12], membrane separation [13], anaerobic and aerobic treatment [14], biological treatment [15], and adsorption [16] have been used to treat wastewater containing organic dyes. However, these methods have several limitations, such as high cost, slow kinetics, and low performance. Therefore, it is important to develop an efficient and cost-effective technique for the decontamination of wastewater containing organic dyes. Recently, semiconductor-based photocatalysis has shown great potential in the remediation of the degradation of organic dyes as it can partially mineralize macromolecular organic dyes into small molecules such as H2O and CO2 [17]. During the photocatalysis process under visible light irradiation, the electrons excites from the filled valence band (VB) of a semiconductor to the empty conduction band (CB), generates excited electron-hole pairs, which reacts with the oxygen and hydroxyl ions to generate active radicals that can degrade

water contaminants readily and quickly [18]. Among the known semiconductor photocatalysts, titanium oxide (TiO2) has been extensively studied and used [19]. The traditional photocatalysts such as ZnO andTiO2 have been extensively used due to their low-cost, reusability, stability, and availability. However, their practical application as photocatalyst has been severely restricted due to their high bandgap (TiO2 ~ 3.2 eV, ZnO ~ 3.37 eV) and ultraviolet (UV) light activity that accounts only for 4% of the total UV-visible light [20–22]. Therefore, there is an imminent need to develop an efficient, low-cost visible-light active materials to tackle environmental pollution.

Among the semiconductor photocatalysts, cadmium sulfide (CdS) is the most extensively used photocatalyst in environmental remediation and one of the most important semiconductor materials (bandgap ~ 2.4 eV), absorbs in the visible region [23,24]. However, the practical usage of CdS in photocatalysis is limited due to some serious shortcomings, including particle aggregation, photo corrosion, and a short lifetime of photo-generated carriers [25,26]. In this context, various efforts have been made to improve CdS photocatalytic activity and stability, such as morphological and structural modification, ion doping, and semiconductors hybridization, etc. It is widely known that the photocatalytic properties of the semiconductor nanomaterials significantly depend on their specific properties such as size, bandgap, morphology, spatial architectures, and especially the specific surface area [27,28]. Therefore various synthetic methods have been investigated for the synthesis of 1D to 3D CdS nanostructures photocatalysts such as hydrothermal/solvothermal method, colloidal method, template method, thermal evaporation, electrochemically induced deposition method, and laser-assisted catalytic growth process [29,30]. Among all these methods, a solvothermal method is the most promising and cost-effective technique to direct and control the nanocrystalline properties of the resultant nanomaterials such as morphology, crystal phase, luminescence, and size [31]. It has been reported that CdS nanorods (NRs) with a hexagonal crystal structure, good crystallinity, and large surface area have proved higher photocatalytic activity compared to the pure cubic phases [32]. CdS NRs with a high aspect ratio showed higher photocatalytic activity due to the restrained electrons in a specific direction that influences the electronic and optical properties of the photocatalyst [33]. Liu et al. found the improved CdS photocatalytic performance due to the increase of exposed 002 facets [34].

Herein, by taking the above consideration and discussion, the main aim of our work was established. The morphology-controlled synthesis of hexagonal phase CdS NRs with high-aspect ratio, high surface area, and exposed 002 facets was performed. The photocatalytic activity of the CdS NRs for the photodegradation of rhodamine B (RhB) was investigated using visible light irradiation. Besides, the effect of several experimental parameters, viz., the catalyst dosage and the initial RhB concentration, were also examined.

#### **2. Materials and Methods**

#### *2.1. Chemicals*

Cadmium acetate dihydrate (Cd(CH3COO)2·2H2O), sodium diethyldithiocarbamate tetrahydrate (C5H10NNaS2·4H2O), diethylenetriamine (HN(CH2CH2NH2)2), and ethanol (C2H5OH) were purchased from Sigma-Aldrich (Islamabad, Pakistan). Rhodamine B (RhB) dye was chosen as a model for the environmental pollutant. All chemicals were of reagent grade and were used as received.

#### *2.2. Synthesis of CdS NRs*

In a typical synthesis procedure, Cd(CH3COO)2·2H2O (7.379 g), C5H10NNaS2·4H2O (5.25 g), and diethylenetriamine (120 mL) were added to a clean Teflon-lined autoclave. The mixture was vigorously stirred for 30 min to form a homogenous dispersion. The autoclave was then sealed and heated at 180 ◦C for 24 h in an oven and subsequently air cooled naturally. The resulting yellow product of CdS NRs was collected after centrifugation and washed with ethanol and distilled water, followed by drying at 80 ◦C for 10 h.

#### *2.3. Characterization*

The crystal phase identification of the photocatalyst was performed by X-ray diffraction (XRD) analysis. The XRD measurement was done using Bruker D2 Phaser diffractometer with Cu-Kα radiation (λ = 1.5406 Å) in the 2θ range between 10–80◦. The structure and morphology were examined by a field emission scanning electron microscope (FESEM-EDS, Zeiss Ultra Plus), and the photoluminescence (PL) measurement was performed on a PL lifetime spectrometer (Edinburgh Instruments). The optical property was determined by using a UV-visible spectrophotometer (Shimadzu UV-1800). The electron binding energies and surface chemical composition were determined by X-ray photoelectron spectroscopy (XPS, Thermo K alpha). The pore size and surface area of the sample were measured by Brunauer–Emmett–Teller (BET) nitrogen adsorption measurement (BET Micromeritics ASAP 2020). Raman spectrum of the sample was recorded on an inVia Renishaw spectrometer.

#### *2.4. Photocatalytic Degradation of RhB*

The photocatalytic ability of the as-synthesized CdS NRs was assessed by the photocatalytic degradation of RhB under visible light irradiation using a 300 W incandescent lamp (Philips) placed at a working distance of 30 cm above the surface of the reaction medium. For RhB degradation, 10 mg sample of the CdS NRs catalyst was added into 80 mL of a 10 mg/L RhB solution. Before light illumination, the mixture was stirred continuously in the dark for 30 min to ensure desorption-adsorption equilibrium. During the subsequent irradiations, 4-mL of suspension was drawn at regular interval of time and centrifuged to obtain a clear solution. The remaining concentration of RhB dye was measured by a UV-vis spectrophotometer.

#### **3. Results and Discussions**

#### *3.1. Structural, Chemical, and Optical Analysis of CdS NRs*

The phase purity and crystallographic structure of the as-synthesized CdS NRs catalyst were investigated by using the XRD technique. The XRD pattern of CdS NRs is shown in Figure 1a. The observed diffraction peaks at 2θ 25.06◦, 26.67◦, 28.40◦, 36.96◦, 43.92◦, 48.10◦, 52.23◦, 54.94◦, and 67.31◦ and were assigned to the (100), (002), (101), (102), (110), (103), (112), (004), and (210) crystal planes of the hexagonal CdS, respectively (ICCD, PDF, 00-001-0780). The sharp diffraction peaks indicate the higher crystallinity of the as-synthesized CdS photocatalyst. The XRD peak intensities of the CdS NRs differ from the bulk material. It is noted that XRD peak intensity at 26.67◦ corresponding to (002) facet of CdS, is much more intense with a small width than all the other peaks, indicating the orientation growth of the nanorods along the c-axis [35], in agreement with FESEM results. Moreover, the XRD pattern of the as-synthesized CdS NRs did not show any peaks except for CdS, which confirms the high purity of the catalyst. The crystalline phase, homogeneity, and surface condition of the CdS NRs was further confirmed by Raman spectroscopy, as shown in Figure 1b. The CdS shows two characteristic peaks at 300 and 603 cm−<sup>1</sup> corresponded to the 1LO and 2LO vibrational modes, respectively. The sharp Raman peaks indicate the high crystallinity of the synthesized photocatalyst, which is consistent with the XRD observations.

**Figure 1.** XRD patterns of CdS (**a**), Raman spectrum (**b**).

Field emission scanning electron microscopy (FESEM) was performed to determine the surface morphology and shape of the CdS nanostructures. The FESEM images of CdS nanostructures are shown in Figure 2, which provides clear evidence for the formation of a rod-like morphology of CdS.

**Figure 2.** FESEM images of CdS NRs (**a**,**b**).

UV-visible spectrophotometric analysis (UV-vis) was performed to analyze the optical absorption properties of the as-prepared CdS NRs. Figure 3a shows the typical UV-vis spectrum of the CdS NRs in the 350–700 nm wavelength range with absorption edges at around 480 and 510 nm. The energy bandgap *E*g was calculated from the following formula

$$
abla \nu = A \left( \hbar \nu - E\_{\%} \right)^{1/2} \tag{1}$$

where *α* is the absorption coefficient, *A* is a constant, *hν* is the photon energy, and *Eg* is the direct bandgap. The bandgap of the synthesized CdS NRs was calculated to be ∼2.35 eV from the (α.E)2 vs. energy (eV) plot, and therefore, this nanomaterial can be used as an absorption photocatalyst over a wide range of the visible spectrum.

**Figure 3.** UV-vis DRS (**a**) and estimated bandgap energy (**b**) of CdS NRs.

The specific surface area and pore structure of CdS NRs were investigated. As shown in Figure 4, the observed isotherms can be assigned as type-IV sorption isotherms with an H3-type of hysteresis loop over a pressure range between 0.8 to 1.0 suggests the presence of slit-like pores [36]. The BET surface area of the CdS NRs was calculated as 51 m2/g. The higher surface area of the as-synthesized CdS NRs photocatalyst compared to the reported CdS photocatalyst [37] makes the transfer of photo-generated charges easier due to the higher number of available active sites on the surface, resulting in improved photocatalytic performance. The average pore size was calculated as 32.7 nm.

**Figure 4.** Nitrogen adsorption–desorption isotherm curve of CdS NRs.

The photoluminescence (PL) measurements were performed using an excitation wavelength of 380 nm to investigate the lifetime and separation of photo-induced electron-hole pairs in CdS NRs. The PL spectrum was recorded in the 425–600 nm spectral range (Figure 5a). It is well known that there is a close link between the PL data and the recombination process of photo-generated electron-hole pairs of the photocatalyst. Generally, a higher value of PL peak intensity indicates the rapid recombination of the photo-generated electron-hole pairs of the catalyst, while a lower value indicates the higher separation of photo-generated electron-hole pairs [38]. The PL emission spectrum of CdS NRs exhibits a broad emission peak around 510 nm due to the near-band-edge emission, which is consistent with the UV-vis spectroscopy results. Time-resolved PL spectrum was captured to calculate the lifetime of the photo-generated electron-hole pairs of the photocatalyst; this was found to be 1.83 ns (Figure 5b) [39].

**Figure 5.** PL spectrum of CdS NRs (**a**), time-resolved PL spectrum (**b**).

The valence state and surface chemical composition of the CdS NRs were determined via XPS analysis. The XPS survey spectrum given in Figure 6a, show the presence of Cd and S in the majority, together with trace amount of C and O, which are probably due to contamination and adsorbed oxygen during sample exposure to the air. The two strong peaks located at 404.2 and 411.2 eV (Figure 6b) corresponds to the Cd 3d5/2 and Cd 3d3/2 states of Cd2+, respectively [40]. The high-resolution S 2p spectrum in Figure 6c shows two strong peaks at 162.2 eV and 160.9 eV; and one weak peak at 168.3 eV. The strong peaks at 162.2 eV and 160.9 eV corresponds to the S 2p3/2 and S 2p1/2, respectively, which can be attributed to S2<sup>−</sup> ions, while the weak peak at 168.3 eV is due to the higher valence state of S, which came from the hydrolysis of sulfur source. Figure 6d shows the binding energy of the O 1s peak of the surface adsorbed oxygen atoms on CdS.

**Figure 6.** XPS survey spectrum of CdS NRs (**a**), high-resolution XPS spectra of Cd 3d (**b**), S 2p (**c**), and O 1s (**d**).

#### *3.2. Photocatalytic Degradation of RhB*

The photocatalytic ability of the as-prepared CdS NRs on the photodegradation of RhB in an aqueous solution under visible light illumination was investigated, and the results are shown in Figure 7. Before the photocatalytic reaction, a blank experiment without adding photocatalyst was carried out to examine the self-degradation of RhB dye. It can be seen in Figure 7b, there is no obvious change in the concentration of dye in the absence of CdS, which suggested that the self-degradation of RhB dye during photocatalytic degradation reaction is negligible. However, by adding the photocatalyst in the solution, the absorption intensity of RhB decreases over time under visible light illumination, verifying that CdS NRs photocatalyst exhibits photocatalytic activity (Figure 7b). The photocatalytic decomposition of RhB reaches 88.4% in 120 min. However, it is very difficult to achieve the maximum degradation value during the photocatalysis process.

**Figure 7.** Time-dependent absorption spectra of RhB solution in visible light (**a**), and the degradation of RB in visible light (**b**).

The effect of photocatalyst dosage (5 mg, 7.5 mg, and 10 mg) on the photocatalytic activity was examined. As can be seen from Figure 8a, increasing the amount of CdS from 5 mg to 10 mg results in an increase in the photodegradation of RhB from 57 to 88.4%, respectively. This increase in photodegradation can be attributed to increases in the number of surface active sites and the total surface area of the catalyst available for the photocatalytic reaction [41].

**Figure 8.** Effect of photocatalyst loading (**a**), Effect of initial RhB concentration (**b**), and Plot of ln (Co/C) vs. time (**c**) (10 mg photocatalyst, 10 mg/L RhB dye, Room temperature, 300 W incandescent lamp).

The effect of the initial RhB concentration on the photocatalytic activity was investigated by varying the concentration from 5 mg/L to 10 mg/L and 15 mg/L (Figure 8b), as it is an important parameter in the photodegradation of organic pollutants. The photocatalytic degradation efficiency decreases from 99 to 55% with the increase in the concentration of RhB from 5 mg/L to 15 mg/L. The reduced rate of degradation efficiency at higher concentrations is due to the unavailability of photoactive sites by increased adsorption of dye molecules and decreased rate of light penetration into the suspension [42,43]. At a high dye concentration, large portion of light is absorbed by the dye molecules, and only small portion of light is absorbed by the catalyst. As a result, the reduced photo flux reaching the catalyst surface causes a reduction in degradation efficiency.

The chemical kinetics of the photocatalytic reaction were studied to obtain some insight into the mechanism of the photodegradation process and to assess the feasibility and efficiency of treating the dye this way. The Langmuir–Hinshelwood kinetics and its pseudofirst-order reaction can be represented by the following equation as (Equations (2) and (3));

$$\frac{\text{dC}}{\text{d}t} = -\mathbb{K}\_{app} \text{ C} \tag{2}$$

$$\ln(\frac{\text{Co}}{\text{C}}) = K\_{app} \text{ t} \tag{3}$$

where *Kapp* is the constant of degradation rate and *t* is the irradiation time. The *ln(Co/C)* versus *t* plot is shown in Figure 8c. A linear fit between *ln(Co/C)* and *t* was obtained for each plot supporting the conclusion that the photodegradation reaction follows pseudofirst-order kinetics with respect to dye concentration. The *Kapp* values were calculated from the slopes of *ln(C*0*/C)* vs. *t* plots. For visible-light irradiation, the correlation coefficient was *<sup>R</sup>*<sup>2</sup> = 0.945 and the *Kapp* value was 0.15 × <sup>10</sup>−<sup>2</sup> min<sup>−</sup>1.

In addition, to study the photostability and reusability of the as-synthesized CdS NRs photocatalyst, an experiment of five consecutive cycles for the photodegradation of RhB dye was carried out. The photodegradation activity of the CdS NRs remained almost constant over five-cycles as shown in Figure 9a, suggests excellent cycling stability photocatalyst. The XRD pattern (Figure 9b) of the used photocatalyst after five cycles shows that no additional peak was observed and the crystal structure of the photocatalyst has been preserved, indicating the high stability and sustainability of the photocatalyst.

**Figure 9.** Stability test of CdS NRs for five cycles (**a**), XRD pattern for CdS NRs before the activity and after recycling five times for RhB photodegradation (**b**).

Figure 10, depicts the plausible mechanism for RhB dye degradation over CdS NRs. Upon visible light illumination, the electrons (e) are excited from the filled valence band ¯

(VB) of CdS to the conduction band (CB), creating a hole (*h*+) in the VB. The generated electron from the conduction band is taken up by the adsorbed O2 at the surface of the photocatalyst and produces superoxide radical ion (•O2 −). The •O2 − further reacts with H2O and produces hydroxyl radical (•OH) or can directly oxidize RhB dye. Simultaneously, some of the photo-induced holes (*h*+) react with OH<sup>−</sup> or H2O to generate active specie •OH, which oxidize the adsorbed RhB dye.

$$\text{CdS NRs} + h\nu \rightarrow \text{CdS (\bar{e} + h^{+})} \tag{4}$$

$$\text{O}\_2 + \text{\color{red}{e}} \rightarrow \text{O}\_2\text{\color{red}{\cdot}}\tag{5}$$

$$\text{O}\_2\text{}^- + \text{ë} + 2\text{H}^+ \rightarrow \text{H}\_2\text{O}\_2 \tag{6}$$

$$\mathrm{H\_2O\_2 + \bar{e} \to OH + OH^-} \tag{7}$$

$$\rm{H}^{+} + \rm{OH}^{-} / \rm{H}\_{2}\rm{O} \rightarrow \rm{OH} \tag{8}$$


$$\text{O}\_2\text{ }^- + \text{MB dye} \rightarrow \text{Decolourization} \tag{11}$$

**Figure 10.** Proposed mechanism for the photocatalytic degradation of RB under visible light illumination.

#### **4. Conclusions**

In the present work, the highly crystalline CdS NRs were successfully synthesized by a one-step solvothermal method. The CdS NRs were characterized in terms of structure, morphological, and optical properties by XRD, Raman, FESEM, BET, PEL, UV-vis, and XPS techniques. The photocatalytic activity investigation was carried out by performing the photodegradation of RhB under visible light illumination over CdS NRs. Moreover, the photocatalytic activity of the CdS NRs depends on the various parameters such as initial dye concertation and the photocatalyst amount. It can be concluded that as-prepared CdS NRs show higher photocatalytic activity at a high dosage of photocatalyst and the low initial concentration of the RhB solution. The photodegradation of RhB follows a pseudo-first-order kinetic model.

**Author Contributions:** Conceptualization, H.U. and M.G.; investigation, H.U. and E.V. All authors have read and agreed to the published version of the manuscript.

**Funding:** This research was funded by the Higher Education Commission of Pakistan.

**Institutional Review Board Statement:** Not applicable.

**Informed Consent Statement:** Not applicable.

**Data Availability Statement:** The data presented in this study are available in the main text of the article.

**Conflicts of Interest:** The authors declare no conflict of interest.

#### **References**

