Solid State Nanostructured Metal Oxides as Photocatalysts and Their Application in Pollutant Degradation: A Review

Abstract

Most dyes used in various industries are toxic and carcinogenic, thus posing a serious hazard to humans as well as to the marine ecosystem. Therefore, the impact of dyes released into the environment has been studied extensively in the last few years. Heterogeneous photocatalysis has proved to be an efficient tool for degrading both atmospheric and aquatic organic contaminants. It uses the sunlight in the presence of a semiconductor photocatalyst to accelerate the remediation of environmental contaminants and the destruction of highly toxic molecules. To date, photocatalysis has been considered one of the most appealing options for wastewater treatment due to its great potential and high efficiency by using sunlight to remove organic pollutants and harmful bacteria with the aid of a solid photocatalyst. Among the photocatalysts currently used, nanostructured metal oxide semiconductors have been among the most effective. This review paper presents an overview of the recent research improvements on the degradation of dyes by using nanostructured metal oxide semiconductors obtained by a solid-state method. Metal oxides obtained by this method exhibited better photocatalytic efficiency than nanostructured metal oxides obtained using other solution methods in several cases. The present review discusses examples of various nanostructured transition metal oxides—such as TiO₂, Fe₂O₃, NiO, ReO₃, IrO₂, Rh₂O₃, Rh/RhO₂, and the actinide ThO₂—used as photocatalysts on methylene blue. It was found that photocatalytic efficiency depends not only on the bandgap of the metal oxide but also on its morphology. Porous nanostructured metal oxides tend to present higher photocatalytic efficiency than metal oxides with a similar band gap.

1. Introduction

Currently, researchers and scientists are becoming increasingly interested in metal oxide nanostructures due to their important technological applications in electronic and optoelectronic devices, sensors, medicines, and renewable energy sources. Reducing the size of materials (metal oxides) to nano-level imparts properties different from the bulk or crystalline forms; these nanoparticles demonstrate behaviors characteristic of isolated atoms and molecules [1,2,3,4,5].

Industrial plants generate increasing amounts of wastewater, which often causes severe environmental problems [6,7]. Wastewater produced in many industrial processes typically contains organic compounds that are toxic and not amenable to biological treatments [7]. There are several different types of organic pollutants, including dyes, phenols, biphenyls, pesticides, fertilizers, hydrocarbons, plasticizers, detergents, oils, greases, pharmaceuticals, proteins, carbohydrates, and so on [8].

Among the various physical, chemical, and biological technologies used in pollution control, advanced oxidation processes such as photocatalysis are being increasingly adopted in the destruction of organic contaminants due to their high efficiency, simplicity, good reproducibility, and ease of handling [8]. Heterogeneous photocatalysis possesses some critical advantages that have feasible applications in wastewater treatments, including: ambient operating temperatures and pressures, and the complete mineralization of contaminants and their intermediate compounds without leaving secondary pollutants. However, for large scale municipal applications, current photocatalytic water treatment systems are less attractive because they are more time consuming and have higher costs than other existing advanced oxidation techniques such as UV/H₂O₂, O₃/H₂O₂, and UV/O₃ technologies [9].

Therefore, semiconducting metal oxides have high potential due to their capability to generate charge carriers when stimulated with the required amount of energy, and to their applications in environmental remediation and electronics. Some semiconducting metal oxide synthesis methods are the chemical vapor deposition technique, the hydrothermal method, the laser ablation technique, and the electro-deposition method [10,11]. Out of these techniques, the hydrothermal method is the most user-friendly because of its reduced cost and ease of handling as well as being chemically reactive at low temperatures. However, other methods including the thermal decomposition of solids such as molecular complexes and macromolecular complexes [12] are gaining relevance due to their ease of preparation and final product purity.

An example of environmental remediation processes is the degradation of organic dyes [13] and the generation of clean energies such as “green hydrogen” [14].

In this context, one of the main applications of nanostructured metal oxides is the photocatalytic degradation of organic pollutants, which is in the field of environmental remediation.

The main characteristics that a suitable metal oxide photocatalytic system must have are [15]:

• An adequate bandgap;
• Suitable morphology;
• High surface area;
• Stability and reusability.

Semiconductor metal oxides with a bandgap near 3.2 eV are UV-light active, while semiconductor metal oxides with a bandgap near 2.7 eV are Visible-light active [16,17]. Metal oxides exhibiting these features—such as oxides of vanadium, chromium, titanium, zinc, tin, and cerium—follow similar primary photocatalytic processes such as light absorption, which induces a charge separation process with the formation of positive holes that can oxidize organic substrates. In this process, a metal oxide is activated with either UV light, visible light, or a combination of both, and photoexcited electrons are promoted from the valence band to the conduction band, forming an electron/hole pair (e⁻/h⁺). The photogenerated pair (e⁻/h⁺) can reduce and/or oxidize a compound adsorbed on the photocatalyst surface. A schematic representation of this process is shown in Figure 1.

As Figure 1 shows, the bandgap of the metal oxides dictates their photocatalytic activity; crucially, it establishes whether their energy matches or exceeds the band gap energy of the semiconductor. In

Table 1. Bandgap values for the main metal oxides along the periodic table.

Metal Oxides	Compounds	Bandgap (eV) a,b	Ref.
Metal transition	Cr2O3	3.2	[18]
	V2O5	2.3	[18]
	Co3O4	1.5	[18]
	TiO2	3.3 (anatase)	[18]
	TiO2	3.0 (rutile)	[18]
	Mn2O3	3.27	[19]
	WO3	2.8	[19]
	MoO3	3.14	[18]
	NiO	3.5	[19]
	Fe2O3	2.1	[18]
	Fe3O4	2.61	[19]
	Cu2O	2.2	[18]
	CuO	1.6	[18]
Metal representative	ZnO	3.2,34	[19,20]
	SnO2	4.2	[19]
	Ga2O3	4.85	[19]
	Sb2O3	4.49	[19]
Lanthanides	CeO2	3.0–3.6	[19]
	La2O3	4.3	[19]
Actinides	ThO2	3.1	[21]

^a Values from references: [17,18,19,20,21,22]. ^b Other similar values are reported for some of the metal oxides.

, a summary of the bandgap values for the most common metal oxides along the periodic table is displayed. It is important to note that the bandgap values vary with the size of the nanoparticle, preparation method, doping, etc.

The photocatalytic activity of metal oxides comes from two sources [13]:

(i)

Generation of ^•OH radicals by oxidation of OH⁻ anions;

(ii)

Generation of O₂⁻ radicals by reduction of O₂.

Nanostructured metal oxide semiconductors have been widely used in photocatalytic redox processes because of their filled valence band (VB) and empty conduction band (CB) electronic configurations. When exposed to a photon with energy exceeding the bandgap (hν > E_g), it generates an electron-hole pair with one electron in VB pumped into CB, leaving a hole behind in VB. The holes generated in VB have great oxidation capability, while the electrons in the CB have high reduction potential.

These highly-reactive electrons and holes participate in the photocatalytic organic degradation.

As mentioned above, the factors relevant for a photocatalyst to be efficient include adequate bandgap, suitable morphology, high surface area, stability, and reusability.

Whether a given nanostructured metal oxide meets these characteristics depends on its preparation method. For instance, TiO₂ is one of the most used and efficient metal oxides for the photocatalytic degradation of several organic dye pollutants [23]. However, its relative efficiency depends on the preparation method, which in turn determines its bandgap, morphology, surface area, stability and reusability.

For instance, the dependence of the catalytic effectiveness of TiO₂ on the preparation method is illustrated in

Table 2. Photocatalytic Performances and Key Factors (Crystal Phase, Size and Morphology) of Different TiO₂ Materials in the Degradation of Organic Water Pollutants [24].

Organic Water Pollutant	Photocatalyst (Concentration)	Irradiation Light	Reaction Kinetic	Removal %	Irradiation Time
MB (1 × 10−5 M)	TiO2 anatase. 1 g × Lt−1 with 340 nm cut-off filter, 330 nm > λ > 680 nm	Xe lamp (150 W)	Pseudo first-order kinetic	86.5%	25 min
MB (10 ppm)	TiO2 nanoparticles (2.5 g L−1)	UV lamp (40 W)	Langmuir–Hinshelwood Pseudo first-order kinetic	71%	60 min
MB (50 ppm)	TiO2 nanofibers	Xenon lamp UV-vis (150 W) with AM 1.5 G filter λ > 400 nm	NM	100%	180 min
MB(10 mgL−1)	TiO2 nanotubes (0.16 g L−1)	xenon visible light (500 W)	NM	99.1%	40 min
MB (0.01 mM)	mesoporous TiO2 0.17 g L−1)	UV light irradiation using eight tubes with a power source of 6 W, λ = 365 nm	NM	85%	60 min
MB (0.75 × 10−5 M)	Hollow titania micro-spheres (HTS)	UV lamp 15 W	Pseudo first-order kinetic	53%	90 min
MB (1 mM)	TiO2 nanoparticles	Blacklight lamp (1 mW)	NM	45%	20 min
Other similar organic water pollutants malachite green, MG (10 ppm)	TiO2nanoparticles (2.5 g L−1)	UV lamp (40 W)	Langmuir–Hinshelwood Pseudo first-order kinetic	78%	60 min
Rhodamine B, RhB, and methyl orange MO (0.01 mM)	TiO2 hollow tetragonal nanocone (0.1 g L−1)	full-arc Xe lamp (300 W) with a cutoff filter, λ > 420 nm	NM	95.0% for RhB, 90.7% for MO	30 min
Bisphenol A BPA, (200 μM)	TiO2 powder, Degussa P25 (0.5 g L−1)	Xe arc lamp (300 W), IR water filter and cutoff filter, λ > 420 nm	Pseudo first-order kinetic	75% for BPA	4 h
acetaminophen, Ace (1.3 μM)	TiO2 powder, Degussa P25 (NM)	UVA/LED λmax = 366 nm	NM	100%	8 min
RhB (NM)	TiO2 nanopowder (0.3 g L−1)	visible light λ > 420 nm	NM	90%	5 h
Methyl orange 4 × 10−5 M	TiO2 powder (10 mg)	Xe lamp 300 W	NM	90%	3 h

.

The E_g bandgap also exhibits an inverse dependence on particle size. For instance, theoretical studies have explained this relationship with the following expression [25]:

$$E_g^{*}\cong E_g^{bulk}+\frac{{\hslash }^2{\pi }^2}{2e{r}^2}\left(\frac{1}{m_e}+\frac{1}{m_h}\right)-\frac{1.8e^2}{4{\pi \epsilon \epsilon }_{0}r}$$

(1)

where E_g bulk is the bulk energy gap, r is the particle radius, m_e is the effective mass of the electrons, m_h is the effective mass of the holes, ε is the relative permittivity, ε₀ is the permittivity of free space, $\hslash$ is Plancks constant divided by 2π, and e is the charge of the electron.

The experimental inverse, E_g vs. particle size, has been observed for several nanostructured metal oxide nanoparticles [25,26,27], including ZnO, Fe₂O₃, Co₃O₄, NiO and SnO₂.

On the other hand, the bandgap can also be modified by doping. In general, doping decreases the bandgap, as is reported in the case of ZnO/Mn²⁺, ZnO/Cu and TiO₂/doping, among others [24]. Some of the most studied are the TiO₂/dopant and ZnO/dopant.

Table 3. Photocatalytic improvement parameter of various ZnO nano-photocatalysts (Adapted from reference [24]).

Photocatalyst	Bandgap	Surface Area
Bi3+-ZnO	3.15–2.6	Increase
Co-ZnO	3.34–3.06	Increase
N-ZnO	3.15–2.86	Increase
F-ZnO	3.35–2.51	Increase
Fe-ZnO	3.24–3.16	Increase
Ag/ZnO	3.30–3.21	Increase
B/ZnO	3.2–3.1	Increase
Bi-TiO2 Ni-TiO2	2.99–3.08 3.02, 2.99–3.03	Decrease
Ag-TiO2	3.0–2.6	Increase
Fe-TiO2	3.2–2.98	Increase
N-TiO2	3.1–2.7	Increase
Ti/WO3	3.4–3.31	Decrease
Zn-WO3	3.2–3.12	Slight decrease
Fe-SnO2	3.8–1.65	Increase
Zn-SnO2	3.50–3.17	Increase
Cu-SnO2	3.02–2.2	Increase

and

Table 4. Photocatalytic improvement performance of various ZnO nano-photocatalysts (Adapted from reference [24]).

N°	Composite	Synthesis Method	Pollutant for Degradation	PC Performance	Irradiation
1	Co-ZnO	Sol-gel	MB	3 at.% Co-ZnO exhibited 92% degradation in 60 min.	Visible light
2	La-doped ZnO	Hydrothermal oxidation	MO and MB	Best PCA observed by S0.005 due to defects.	UV light
3	Pd/ZnO	Microwave hydrothermal, borohydride and photoreduction	CR	Pd/ZnO synthesized by borohydride method has the highest PCA compared to other routes.	UV light
4	C, N co-doped ZnO	Two-step pyrolysis	MB	6C25 showed the best degradation due to its larger number of active sites.	Solar stimulated light
5	X-ZnO (X = Li, Al, N, P)	Mass production technology	RhB	PCA decreased as follows: N > Li > P > Al	Sunlight
6	Fe-doped ZnO	Combustion	BPA	FexZn1−xO (where x = 0.03) showed noticeable efficacy in the series.	Sunlight
7	Gd-ZnO films	RF magnetron sputtering	MB	0.7 at.% Gd-ZnO exhibited higher PCA than ZnO.	UV light
8	Au-ZnO	Combustion	MB	Ag-ZnO demonstrated better activity than Au-ZnO.	UV light
	Ag-ZnO
9	Eu3+-doped ZnO	Coprecipitation	RhB	Doped ZnO (100%) degraded dye faster than ZnO.	UV light
10	Ce-ZnO	Hydrothermal	Pharmaceutica	Nizatidine, levofloxacin and acetaminophen degraded around 95% within 4 h.	UV light
11	Cu-ZnO	Chemical growth	MO and MB	Increase in PC efficiency was 57.5% for MO and 60% for MB in 180 min.	UV light
12	B-ZnO	Sol-gel	CN−	Doped sample with 1.5 wt% exhibited 89% degradation whereas pure ZnO exhibited 75%.	Solar stimulated light
13	In-ZnO	Plasma assisted chemical vapors	MB	4 at.% In-ZnO showed improved PCA compared to ZnO and 8 at.% In-ZnO.	Solar stimulated light
14	Sm-ZnO	Chemical precipitation	MB	Zn1−xSmxO x = 0.04 expressed maximum PC degradation (94.94%).	Visible light
15	WO3-doped ZnO	hydrothermal	Diazinon	10 mg/L diazinon, 10 g/cm—2 2% O3-ZnO exhibited 99% degradation in 180 min.	UV light

summarize some of the results of M_xO_y/dopant systems. One of the explanations for this decrease in bandgap after doping is the increase in surface area. The increased surface area of the photocatalyst also increases the photodegradation of dye pollutants. This is because materials with a large surface area have more active sites than materials with a low surface area. This is illustrated in Table 3.

As mentioned before, ZnO is one of the most common and effective photocatalysts for the photodegradation of organic dye pollutants. ZnO is insoluble in water, existing in the form of a white powder, and it is extracted from the mineral Zincite through synthetic methods. Plastic, rubber, glass, ceramic, and cement industries use ZnO in various processes. ZnO is an n-type semiconductor with high bandgap energy that belongs to group II–IV of the periodic table [26,27]. ZnO is chemically stable, nontoxic, and inexpensive. It possesses high radiation hardness, effective transparency, remarkable optical absorption in the UV range, and outstanding thermal properties. Due to these characteristics, ZnO is used in a number of applications, such as optoelectronics, solar cells, medicine, catalysis, photocatalysis, and as an antibacterial agent. Doping in ZnO produces a decrease in the bandgap, which enhances photocatalytic efficiency. In Table 4, some doped ZnO are shown.

Titanium dioxide, TiO₂, is also known as Titania or titanium (IV) oxide. TiO₂ is the most promising photocatalyst since it possesses long-term chemical and physical stability along with low production costs [28]. It has a wide bandgap (3.0–3.2 eV) in the UV light range and exhibits three polymorphs, namely rutile, anatase, and brookite. Among these three phases, rutile is the most stable phase; the others are metastable, but they can transform into rutile phase irreversibly at elevated temperatures. Furthermore, the TiO₂ photocatalyst has important uses in water treatment, lithium-ion batteries, sensors, catalysis, antibacterial, and anticancer applications [29,30,31].

Table 5. Summary of various TiO₂ nano-photocatalysts, pollutants, and irradiation sources along with their measured Photocatalysis (Adapted from reference [24]).

N°	Doped	Pollutants	PCA	Irradiation Source
1	Ce-TiO2 La-TiO2 V-TiO2	RhB	1.0%-Ce-TiO2 > 1.0%-V-TiO2 > undoped TiO2 > 1.0%-La-TiO2 showed degradation (%) 83.43 > 53.74 > 21.56 > 11.09, respectively.	Solar light
2	PF co-doped anatase TiO2	MO	It demonstrated excellent PCA compared to undoped TiO2, F-TiO2, P-TiO2 and Degussa P25.	Full-spectrum light
3	Ga-doped TiO2nanopowder	MO	0.6 mol% Ga-TiO2 demonstrated up to 82% degradation.	Visible light
4	F, N co-doped TiO2	MB	97.31% degradation was achieved within 5 h.	Visible light
5	Nd-TiO2 film	MB	0.1% Nd-doped TiO2 showed maximum degradation (92%).	UV light
6	Moroccan natural P-TiO2	IC	Its degradation increased at high values of pH, initial concentration and amount of catalyst.	UV light
7	Nb2O5-TiO2	MB	5 mol% expressed the highest PCA under visible light and equal efficiency as TiO2 under UV.	UV and Visible light
8	Mesoporous Ag-TiO2	MO	TiO2 with the lowest content of Ag exhibited higher PCA.	UV and solar light
9	Fe3+-TiO2	CV	Degradation kinetics rate increased with an increase in iron content.	UV
10	Ru/TiO2	2-CP	0.4 wt% Ru/TiO2 showed high PCA using UV light and 0.2% Ru/TiO2 using visible light.	UV and visible light
11	Gd-TiO2	RhB	0.3% Gd-TiO2 demonstrated the best PCA.	Visible light
12	C-TiO2	RhB	More decolorization compared to fluorescence spectroscopy.	Visible light
13	Pd-TiO2	MB and MO	Maximum degradation shown by 0.75 wt% Pd-TiO2 for mixture of dyes and 0.5 wt% Pd-TiO2 for single dye.	Visible or solar light
14	Graphene/TiO2	BPA	It showed remarkable PCA compared to pure TiO2.	Solar light
15	Ni/TiO2	Malathion	94% degradation achieved.	UV light

shows some examples of TiO₂/dopants and their photocatalytic activity toward different organic dye pollutants.

2. Materials and Methods

Nanostructured Metal Oxides Preparation Methods

Several solution methods have been developed for synthesizing metal or metal oxide nanoparticles, such as the solvothermal, hydrothermal, and sol-gel methods [10,11]. Several of these nanostructured metal oxides have been suitable for photocatalytic degradation of organic dye pollutants. Usually, the preparation of nanostructured metal oxides for photocatalysis involves some solution method. However, for some of these methods, drawbacks like the aggregation of the NPs during the adsorption stage cause a reduction in the efficiency of the desired application [32,33,34]. When the metal or metal oxide is obtained as a colloid, as is commonly observed, their separation by centrifugation often produces agglomeration of the NPS [34]. To solve this type of problem, a solid-state method is used in which the metal nanoparticles are generally obtained as solid pure phase material. Therefore, the solventless synthesis of nanostructures is highly significant due to its economical, eco-friendly, and industrially viable nature. However, the development of new solid-state methods to prepare metallic nanostructured materials is a constant challenge. We have previously described a new solid-state method to synthesize metallic nanostructured nanomaterials from the pyrolysis of metallic and organometallic derivatives of poly- and oligo-phosphazene under air and at 800 °C [35,36,37,38]. Metallic nanostructured materials of the type M, M_xO_y, and M_xP_yO_z are obtained depending on the nature of the metal. Another method—used when the respective metallic or organometallic derivative is not possible to prepare—uses MLn/[NP(O₂C₁₂H₈)]₃ mixtures [39,40,41]. However, in several of these systems, the M or M_xO_y phase is accompanied by a phosphate phase. These methods have been discussed in detail in several papers [35,36,37,38,39,40,41] and books [32,33,42,43,44]. When we used a polymer that does not contain phosphorus in its polymeric chain—such as the chitosan MXn and PS-co-4-PVP MXn macromolecular precursors (MXn = metallic salt)—and then subjected it to solid-state pyrolysis under air and at 800 °C (see Figure 2), pure phase nanostructured M_xO_y were obtained [45,46,47,48,49,50,51,52,53,54,55,56,57]. Now we will discuss the use of these nanostructured metal oxides M_xO_y obtained by this method in the degradation of organic dye pollutants.

3. Results

3.1. TiO₂

The titania to be used as photocatalyst was obtained by pyrolysis of the (Cp₂TiCl₂)(Quitosano) (I),(Cp₂TiCl₂)(PS-co-4-PVP) (II), (TiOSO₄)(Chitosan) (III), (TiOSO₄)(PS-co-4-PVP) (IV), (TiO(acac)₂)(Chitosan) (V), and (TiO(acac)₂)(PS-co-4-PVP) (VI) macromolecular precursors under air and at several temperatures (500 °C, 600 °C, 700 °C, and 800 °C [23]. The bandgaps of the TiO₂ obtained using these precursors are displayed in

Table 6. Bandgap for TiO₂ obtained using different precursors and temperatures (Adapted from reference [23]).

TiO2-(I)p			TiO2-(II)p
Temperature (°C)	Phase	Bandgap (eV)	Temperature (°C)	Phase	Bandgap (eV)
500	Anatase	3.66	500	Anatase	3.43
600	Mixture	3.43	600	Anatase	3.30
700	Mixture	3.63	700	Anatase	3.22
800	Rutile	3.40	800	Mixture	3.27
TiO2-(III)p			TiO2-(IV)p
Temperature (°C)	Phase	Bandgap (eV)	Temperature (°C)	Phase	Bandgap (eV)
500	Anatase	3.32	500	Anatase	3.53
600	Anatase	3.43	600	Anatase	3.45
700	Anatase	3.28	700	Mixture	3.37
800	Anatase	3.65	800	Mixture	3.38
TiO2-(V)p			TiO2-(VI)p
Temperature (°C)	Phase	Bandgap (eV)	Temperature (°C)	Phase	Bandgap (eV)
500	Anatase	3.47	500	Anatase	3.72
600	Anatase	3.42	600	Mixture	3.57
700	Mixture	3.21	700	Mixture	3.36
800	Mixture	3.33	800	Rutile	3.24

. The values range from 3.21 to 3.72 eV, which are adequate for photocatalytic degradation of organic dyes using UV-visible irradiation.

The results of the photocatalytic degradation of methylene blue are shown in

Table 7. Kinetic data for degradation of methylene blue with nanostructured TiO₂ (Adapted from reference [23]).

Sample	Apparent Photodegradation Rate Constant k (10−2 min−1)	Degradation η (%)	R2 (%)
TiO2-Anatase-(I)p 500 °C	0.40 ± 0.04	11	93.9
TiO2-Mixture-(I)p 600 °C	0.80 ± 0.04	20	98.4
TiO2-Mixture-(I)p 700 °C	0.06 ± 0.03	14	98.0
TiO2-Rutile-(I)p 800 °C	1.30 ± 0.10	30	93.8
TiO2-Anatase-(II)p 500 °C	0.40 ± 0.04	11	91.9
TiO2-Anatase-(II)p 600 °C	0.33 ± 0.03	7	88.8
TiO2-Anatase-(II)p 700 °C	0.40 ± 0.02	10	97.6
TiO2-Mixture-(II)p 800 °C	1.00 ± 0.06	23	97.2
TiO2-Anatase-(III)p 500 °C	3.90 ± 0.20	65	96.9
TiO2-Anatase-(III)p 600 °C	4.10 ± 0.20	65	97.4
TiO2-Anatase-(III)p 700 °C	2.00 ± 0.06	39	99.2
TiO2-Anatase-(III)p 800 °C	7.13 ± 0.01	87	99.8
TiO2-Anatase-(IV)p 500 °C	2.40 ± 0.01	45	97.8
TiO2-Anatase-(IV)p 600 °C	3.10 ± 0.01	55	98.0
TiO2-Mixture-(IV)p 700 °C	0.20 ± 0.02	5	91.1
TiO2-Mixture-(IV)p 800 °C	4.00 ± 0.20	63	97.1
TiO2-Anatase-(V)p 500 °C	3.40 ± 0.10	59	99.1
TiO2-Anatase-(V)p 600 °C	1.30 ± 0.04	27	99.2
TiO2-Mixture-(V)p 700 °C	1.30 ± 0.09	27	96.4
TiO2-Mixture-(V)p 800 °C	0.60 ± 0.03	14	98.0
TiO2-Anatase-(VI)p 500 °C	1.70 ± 0.03	34	99.7
TiO2-Mixture-(VI)p 600 °C	2.40 ± 0.02	48	95.9
TiO2-Mixture-(VI)p 700 °C	2.60 ± 0.02	51	94.6
TiO2-Rutile-(VI)p 800 °C	0.50 ± 0.02	14	98.5

.

As can be seen in Table 7, the most efficient catalytic degradation arises from the TiO₂ produced from the TiO₂-A-(III)p(800) system precursors, with 87% degradation after 25 min. Based on this table, the photocatalytic efficiency could be expressed by the following multifactorial relationship:

PE = a bandgap + b particle morphology + c particle size + d Titania phase + e pyrolysis temperature

(2)

where PE is photocatalytic efficacy and the Titania phase can be anatase, rutile, or brookita. In the case of TiO₂ obtained by our solid-state method using TiO₂-A-(III)p(800) as the precursor, factor b is highly porous (see Figure 3) since, as it is known, the anatase form is the most efficient phase as a photocatalyst and, in solid-state, the temperature of the thermal treatment is crucial to determine morphology and particle size. In this case, factor a does not appear to be very important since, as shown in Table 7, the values are similar.

The results shown in Table 7 were obtained at pH 6.5; however, a pH effect study showed that, at alkaline pH 9.5, the degradation was 98% [23].

3.2. Fe₂O₃

Fe₂O₃ was prepared by pyrolysis of the Chitosan(FeCl₂)_y, Chitosan(FeCl₃)_y, PS-b-4-PVP(FeCl₂)_y, and PS-b-4-PVP(FeCl₃)_y macromolecular complexes under air and at 800 °C [47]. The product was Fe₂O₃ in the hematite phase in all the cases. The bandgap of the as-prepared Fe₂O₃ ranged from 1.83 eV to 2.12 eV, which indicated an effective photocatalyst behavior under UV-visible irradiation (see

Table 8. Kinetic data for the degradation of MB with α-Fe₂O₃ obtained from PS-co-4-PVP and chitosan macromolecular precursors (Adapted from reference [47]).

Precursor	Eg (eV)
Chitosan(FeCl3) 1:1	1.83
Chitosan(FeCl3) 1:5	2.15
PS-co-4-PVP(FeCl3) 1:1	2.12
Fe+3-PS-co-4-PVP(FeCl3) 1:5	2.12
Chitosan(FeCl2) 1:1	2.15
Chitosan(FeCl2) 1:5	2.15
PS-co-4-PVP(FeCl2) 1:1	1.90
PS-co-4-PVP(FeCl2) 1:5	2.09

).

Some parameters for the degradation experiments of MB with α-Fe₂O₃ obtained from PS-co-4-PVP and chitosan macromolecular precursors are summarized in

Table 9. Kinetic data for the degradation of MB with α-Fe₂O₃ obtained from PS-co-4-PVP (1:1) and chitosan (1:1) macromolecular precursors (Adapted from reference [47]).

Photocatalyst	Apparent Photodegradation Rate Constant k (10−2 min−1)	Discoloration Rate η(%) at 60 min	Discoloration Rate η(%) at 150 min
α-Fe2O3 (PS-co-4-PVP)	1.2 ± 0.04	62.6	86.9
α-Fe2O3 (Chitosan)	2.1 ± 0.1	73.4	94.6

. The highest extent of degradation of MB (98.6 %) at 150 min of irradiation time was achieved for α-Fe₂O₃ obtained from the PS-co-4-PVP(FeCl₂)y precursor in a 1:1 ratio.

The size and morphology of the two systems were determined by electron microscopy studies, showing fused nanoparticles and 3D networks with average sizes of 150–200 nm for α-Fe₂O₃ from the precursor Chitosan(FeCl₂) 1:1 and 55–100 nm for α-Fe₂O₃ from the precursor PS-co-4-PVP(FeCl₂) 1:1. These morphologies show a random network of the surface area, quasi-linear nanoparticle chains that fold into larger porous powder particles [47]. It appears to be that for these Fe systems, the multifactorial relationship that holds this behavior could be the following:

PE= a bandgap + b particle morphology + c particle size + d (Fe precursor)

(3)

where the (Fe precursor) factor refers to the FeCl₂ or FeCl₃ metal salt source linked to the polymer. Once again, the most important factor dictating the photocatalyst efficiency of α-Fe₂O₃ toward the degradation of methylene blue is the b factor, more specifically, the material porosity. Table 9 shows that the bandgaps of α-Fe₂O₃ are similar; however, factor a could be less important. On the other hand, factors c and d could be somewhat significant.

3.3. NiO

NiO is a p-type semiconductor with E_g = 3.5 eV with multiple practical applications [57,58,59,60]. However, its bandgap can be modified by doping with other metal oxide semiconductors or by formation of nanocomposites in which one of the components would be in less quantity (mainly its photocatalyst properties) [59,60]. NiO has been used in catalysis, battery cathodes, fuel-cell electrodes, electrochromic films, electrochemical supercapacitors, and magnetic materials [59,60]. As these applications depend on its bandgap value and, in turn, the bandgap properties of superconductors depend on the environment of the metal oxide materials [61,62], there are no systematic, detailed studies on the effect of the medium on the bandgap. Although some M_xO_y/M’_xO’_y have been prepared—where M_xO_y is a nanostructured metal oxide and M’_xO’_y is a metal oxide matrix—no systematic effect of the matrix on the photocatalytic efficiency toward the degradation of organic dyes has been reported. By preparing the NiO/SiO₂, NiO/TiO₂, NiO/Al₂O₃, and NiO/glasses nanocomposites through a solid-state method, we have investigated the effect of the SiO₂, TiO₂, Al₂O₃, and Na_4.2Ca_2.8(Si₆O₁₈)(glass) matrices on the photocatalytic properties of NiO toward the degradation of methylene blue [57]. The composites were prepared by solid-state thermal treatment of the Chitosan(NiCL₂)_x/M’_xO’and PS-co-4-PVP(NiCL₂)_x//M’_xO’_y precursors, where M’_xO’_y could be SiO₂, TiO₂, and Al₂O₃ glasses. Kinetic data for the photodegradation process of MB with NiO and with the NiO/SiO₂, NiO/TiO₂, NiO/Al₂O₃, and NiO/glasses composites are displayed in

Table 10. Kinetic data for the photodegradation process of MB with NiO and with the NiO/SiO_2, NiO/TiO₂, NiO/Al₂O₃, and NiO/glasses composites (Adapted from reference [57]).

Photocatalyst	Apparent Photodegradation Rate Constant k (10−2 M·min−1)	Discoloration Rate (%)	R2 Linear Fit (%)
NiO-chitosan a	2.4	71%	0.998
NiO-PS-4-PVP	2.2	68%	0.991
NiO/SiO2-chitosan	2.3	69%	0.999
NiO/SiO2-PS-4-PVP	1.6	48%	0.996
NiO/TiO2-chitosan	2.9	91%	0.992
NiO/TiO2-PS-4-PVP	2.6	81%	0.980
NiO/Al2O3-chitosan	1.5	45%	0.990
NiO/Na4.2Ca2.8(Si6O18)	2.6	75%	0.990

^a Nomenclature: Composite, hyphen (-), name of the precursor polymer.

.

As can be observed in Table 6, the NiO from the precursor containing the chitosan as a solid-state template produces much higher activity than the NiO sample arising from the PSP-4-PVP(NiO)_x precursor. The polymer precursor clearly influences the photocatalytic activity. In this sense, using chitosan as a solid-state template produces much higher activity than the NiO sample arising from the PSP-4-PVP(NiO)_x precursor. On the other hand, the most efficient photocatalytic activity was observed for the NiO/TiO₂ composite, with 90% degradation of methylene blue in 5 h.

In general, the photocatalytic activity shown in Table 10 is probably related to a matrix effect of NiO inside the different SiO₂, TiO₂, and Al₂O₃ matrices. Different matrices can influence NiO in different ways. For example, when the matrix is TiO₂, a p-n junction p-NiO/TiO₂ can be formed; this reduces the recombination rate of the photogenerated electron-hole pairs, which is known to enhance the photocatalytic activity of TiO₂ (see Figure 4) [60]. In the case of the NiO/TiO₂ composite, which has the highest catalytic behavior, the matrix effect may arise from the enhanced activity of TiO₂ by NiO. Therefore, in this case, NiO acts as a matrix in the active semiconductor for the degradation of methylene blue. On the other hand, the NiO/Al₂O₃ compound is one of the least efficient photocatalysts for the degradation of methylene blue, which is probably due to an insulating effect caused by Al₂O₃ that cuts off the communication between p-NiO and methylene blue to start the degradation of the dye. This is also in agreement with the low photocatalyst efficiency of the insulating SiO₂ matrix. An explanation for the results of the photocatalytic activity of the NiO/Na_4.2Ca_2.8(Si₆O₁₈) composite obtained from the Chitosan(NiCl₂·6H₂O)_x//Na₂O·CaO·SiO₂ precursor is that Na_4.2Ca_2.8(Si₆O₁₈) has similar characteristics to those of SiO₂ as a matrix.

3.4. Precious Metal Oxides Ir, Rh, Re: IrO₂, Rh/RhO₂, Rh₂O₃, ReO₃

Among the metals in the periodic table, the so called precious—such as Ir, Rh, and Re, among others—are some of the most catalytically active [63]. Their activity is hugely enhanced at the nano-level [64,65]. These metals, as well as their metal oxides, exhibit high catalytic activity [63,65].

3.5. Ir

IrO₂ is a promising conducting oxide used, for example, as an electrode material in ferroelectric capacitors for nonvolatile memory applications [66]. However, to the best of our knowledge, there are no reports on its photocatalytic activity. The only related studies are those about the catalytic reduction of 4-Nitrophenol [67] and its electrochemical catalytic activity toward oxygen evolution [68]. As with the previous metal oxides, the iridium oxide for the catalytic experiments was prepared by thermal treatment of the Chitosan(IrCl₃)_x and PSP-4-PVP(IrCl₃)_x macromolecular precursors [55]. The as-prepared IrO₂has an adequate bandgap for the degradation of methylene blue under UV-visible irradiation (see

Table 11. Kinetic data for the photodegradation process of MB with IrO₂, Rh/RhO₂, Rh/Rh₂O₃, and ReO₃.

Photocatalyst	Photodegradation Rate Constant k (10−3 M·min−1)	Discoloration Rate (%)	R2 Linear Fit (%)	Ref.
IrO2-PS-4-PVP	1.7	53%	0.995	[55]
IrO2-chitosan	2.4	38%	0.991	[55]
Rh/RhO2	a	78%	b	[56]]
Rh2O3	a	70%	b	[56]
ReO3-PS-4-PVP	2.8	64%	0.977	[54]
ReO3-chitosan	2.8	53%	0.997	[54]

^a Photodegradation rate constant k not informed. ^b Linear fit not informed.

). As previously mentioned, we have used this dye as a model dye to assay the photocatalytic properties of IrO₂. Kinetic data for the photodegradation process of MB with IrO₂ are displayed in Table 11. IrO₂ from the PVP precursor exhibits better photocatalytic activity (57% photodegradation in 300 min) than the IrO₂ from the chitosan precursor (38% photodegradation in 300 min). This effect was associated with the more porous morphology of IrO₂ from the PVP precursor [55].

3.6. Rh

In this case, the Rhodium oxides Rh/RhO₂ and Rh/Rh₂O₃ were obtained by pyrolysis of the PSP-4-PVP(RhCl₃)_x and Chitosan(RhCl₃)_x macromolecular precursors, respectively [56]. The bandgap values (see

Table 12. Bandgap for the Ir, Rh, and Re oxides.

Photocatalyst	Bandgap (eV)	Ref.
ReO3	4.36	[54]
IrO2	2.4–2.6	[55]
Rh/RhO2	3.7	[56]
Rh2O3	3.0	[56]

) are adequate for photocatalytic activity under UV-visible irradiation. The results (see Table 11) indicate that the Rh/RhO₂ sample exhibits better photocatalytic activity (78% photodegradation in 300 min) than the Rh/Rh₂O₃ one (70% photodegradation in 300 min). As the porosity of both materials is similar, the higher photocatalytic activity of Rh/RhO₂ could be due to its small particle size in the range 10–20 nm [56].

3.7. Re

ReO₃ for the photocatalytic essays was obtained by solid-state thermal treatment of the Chitosan(ReCl₃)_x and PSP-4-PVP(ReCl₃)_x macromolecular precursors [54]. The bandgap for ReO₃ indicates an adequate value for UV-visible photoactivation (see Table 12). ReO₃ was also found to catalyze the photodegradation of MB with an efficiency of 53% and 64% in 300 min for the chitosan and PVP precursors, respectively (see Table 11). ReO₃ is an unusual transition metal oxide as it presents a metallic behavior with conductivity close to that of copper [69]. Therefore, nanoparticles of rhenium trioxide produce a SERS (surface-enhanced Raman spectra) effect on some organic compounds, such as pyridine [70]. Despite this, their catalytic activity has only been proved in the catalytic degradation of methyl orange [71].

3.8. Th

Among the actinides oxides, thoria is an important and promising material used in ceramic catalyst sensor solid electrolytes, catalysis, optical materials, and the traditional nuclear industry [72,73,74,75]. The thorium oxide was prepared by thermal treatment of the Chitosan Th(NO₃)₄ and PS-co-4-PVPTh(NO₃)₄ macromolecular precursors at 800 °C. The ThO₂/SiO₂ and ThO₂/TiO₂ composites were also synthesized by pyrolysis of the Chitosan Th(NO₃)₄//MO₂ andPS-co-4-PVP Th(NO₃)₄//MO₂ macromolecular composites, where MO₂ is SiO₂ or TiO₂.

To the best of our knowledge, no photocatalytic studies on ThO_2, ThO₂/SiO₂, or ThO₂/TiO₂ have been reported. Regarding thoria, scarce literature data have been reported. As for ThO₂ nanoparticles, Aller et al. [76] report values ranging from 5.13 eV to 4.5 eV, while for ThO₂ thin films, Buono-Core et al. [77] report values from 4.5 eV to 4.61 eV. However, some of these values are theoretical calculations. Using the solid-state UV-visible absorption and the Tauc plot for ThO₂, a value of 5.66 eV is estimated for the chitosan polymer precursor and 5.76 eV for the PS-4-co-PVP polymer precursor.

With this in mind, using UV-visible radiation, some catalytic activity for ThO₂ could be expected. In fact, the values for the catalytic results are shown in

Table 13. Kinetic data for the photodegradation process of MB with ThO₂ and with the ThO₂/SiO_2, ThO₂/TiO₂ composites (Adapted from reference [21]).

Photocatalyst	Eg (eV)	Apparent Photodegradation	Discoloration Rate (%)	R2 Linear Fit (%)
ThO2(chitosan precursor)	5.66	3.7 × 10−3	67	0.992
ThO2 (PS-4-PVP precursor)	5.75	2.2 × 10−3	66	0.967
ThO2/SiO2 (chitosan precursor)	5.50	7.7 × 10−4	24	0.979
ThO2/SiO2 (PS-4-PVP precursor)	5.6	8.5 × 10−4	25	0.923
ThO2/TiO2 (chitosan precursor)	3.14	1.4 × 10−3	39	0.815
ThO2/TiO2 (PS-4-PVP precursor)	3.14	8.7 × 10−4	27	0.941

. The values for the ThO₂/SiO₂ and ThO₂/TiO₂ nanocomposites are also shown for comparison. On the other hand, the bandgap values for thoria included in TiO₂ as matrix exhibited values of 3.14 eV and 3.15 eV for the ThO₂/TiO₂ composites from chitosan and PS-4-co-PVP polymer precursors, respectively. These values are lower than those of ThO₂ and ThO₂/SiO₂; however, as Buono-Core et al. point out [77], the optical bandgap energy is very sensitive to the preparation method and the experimental parameters applied in the synthesis [77]. In fact, Mahmoud reported a value of 3.82 eV for ThO₂ prepared by a spray pyrolysis technique [78]. The estimated rate constant for the degradation of methylene blue in the presence of thorium prepared without the SiO₂ and TiO₂ matrices is greater than that of the pristine compounds, suggesting that the structural modification and synergy of the inorganic components play a fundamental role. This is related to the increase in the photocatalytic efficiency of the semiconductor as a result of the greater number of active sites provided by the (ThO₂) PS-4-co-PVP and (ThO₂) chitosan precursors.

In general and according to the examples shown here, the efficiency in photocatalysis using photocatalyst semiconductor oxides obtained by a solid method is higher than that exhibited by the same photocatalysts obtained by solution methods. This may be due to the high porosity of the catalysts obtained by the solid state method.

3.9. Heterojunction Structure

These structures are formed by coupling two semiconductors whose band structures are aligned in such a way that they allow for an enhanced charge separation compared to the single system. A maximum of two alignments are feasible with this binary combination, viz., type I and type II [79,80,81,82]. Figure 5 shows a schematic of the band alignment of the nanocomposites for both heterostructures. In the type I alignment, both the VB and CB edge potential of either semiconductor lies within the bandgap of the other semiconductor, as the latter possesses a wider bandgap. Since one has lower VB and CB potentials than the other, both bands act as hole and electron collecting sites, as is shown in Figure 5. In the type II heterostructure, the CB position of the first semiconductor lies above that of the second, whereas its VB lies within the bandgap of the second. Therefore, the photogenerated holes tend to migrate to the VB of the first one, whereas the excited electrons tend to migrate to the CB of the second. The mentioned band position and migration of the charge carriers are depicted in Figure 5b. In general, the type II heterostructure is widely preferred because it allows the migration of electrons and holes in the opposite direction (see Figure 5b) [79,80,81,82].

Some examples of both type I and type II band alignments are presented and discussed in reference [81].

3.10. Photocatalyst Mechanism of Dye Degradation by Nanostructures

In general, the system of heterogeneous photocatalysis using semiconductor materials consists of a light-harvesting antenna and several active species to facilitate pollutant degradation [83]. The series of chain oxidative-reductive reactions that occur at the photon-activated surface has been broadly proposed as:

Photocatalyst + hν → h⁺ + e⁻

h⁺ + H₂O → *OH + H⁺

h⁺ + OH⁻ → *OH

h⁺ + Pollutant → (Pollutant)⁺

e⁻ + O₂ → *O₂⁻

*O₂⁻ + H⁺ → *OOH⁻

2 *OOH → O₂ + H₂O₂

H₂O₂ + *O₂⁻ → *OH + OH⁻ + O₂

H₂O₂ + hν → 2 *OH

Hence, the final reaction is:

Pollutant + (^•OH, h⁺, ^•OOH or O₂⁻) → degradation products

which produces the degradation of the pollutant.

When the semiconductor is irradiated by an input light possessing ultra-band-gap energy (hν > E_g), a valence band (VB) electron (e⁻) is excited to the conduction band (CB), leaving behind a photogenerated hole (h⁺) at the VB. Accordingly, the produced e⁻/h⁺ pairs are able to migrate to the surface of the semiconductor and participate in redox reactions.

The photocatalytic reaction usually involves three main active species: a hydroxyl radical (^•OH), h⁺, and a superoxide radical (^•O₂⁻), where ^•OH is the primary oxidant in the photocatalytic degradation of the pollutant in the aqueous solution. ^•OH radicals are normally generated via two routes: (i) H₂O and OH⁻ in a water environment are readily oxidized by photogenerated h⁺ to form ^•OH radicals; (ii) O₂ present in an aqueous solution is reduced by photogenerated e⁻ to form O₂ radicals, which subsequently react with h⁺ forming ^•OOH radicals, whose further decomposition produces ^•OH radicals.

Complementary to what was discussed above, two interesting reviews [84,85] analyze the photocatalytic properties of the TiO₂ system incorporated in a carbon gel matrix and covalent organic frameworks (COFs). In Dongge et al., although still in its incipient stage in comparison with other traditional inorganic metal oxides, titania is proposed as having some advantages such as having a high surface due to the high porosity of the carbon gel matrix. In the second case, Yuhang et al. point out that even when COFs are highly porous and crystalline polymeric materials, they are photocatalytic systems different from nanostructured metal oxide semiconductors, and are a valid alternative in environmental remediation. Meanwhile, recent investigations [86,87] have proven that photocatalysis of nanostructured metal oxides is being increasingly adopted.

4. Conclusions

Nanostructured metal oxides play an important role in the environmental decontamination of organic dyes. An alternative route that often improves photocatalytic efficiency is the use of photocatalysts based on metal oxides obtained by a solid-state method. Specifically, the solid-state method based on the thermal treatment of Chitosan·MXn and PS-co-4-PVP MXn precursors affords pure nanostructured metal oxides M_xO_y, which are often more efficient photocatalysts than the respective M_xO_y obtained by other solution methods. It was found that the photocatalytic efficiency depends on the nature of the MXn precursor salts and the calcination temperature, as in the case of TiO₂, where the most efficient photocatalyst was the one obtained from the Chitosan TiSO₄ precursor pyrolyzed at 800 °C. The efficiency of most of the studied photocatalysts can be described by the following relationship:

PE = a band gap + b particle morphology + c particle size + d crystalline phase + e pyrolysis temperature

(4)

where PE is the photocatalyst efficiency. In particular, for a nanostructured metal oxide semiconductor photocatalyst obtained using our solid-state method, morphology is more important than the bandgap. In turn, when using the solid-state preparation method of the metal oxide photocatalyst, the temperature often modulates the morphology. This can, for example, generate a highly porous morphology, thus causing a very efficient photocatalytic activity.

However, research on this topic is a challenge since the parameters that induce a more efficient photocatalysis in solid-state synthesis are not fully known. An achievement in this matter could be a significant advance for environmental decontamination.