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1. Introduction

Pharmaceuticals (analgesics and anti-inflammatories, antibiotics, anticonvulsants, β-blockers, endocrine active pharmaceuticals, lipid regulators, psychiatric drugs [1]) have become an essential part of our life, benefiting society’s health and wellbeing. Improved life expectancy, the better prevention of diseases and significantly reduced mortality could not be imagined without appropriate medication [2]. The consumption of pharmaceuticals continues to rise because of the ongoing demographical changes: growing world population, population density and aging societies (especially in developed countries) [1,3]. A growing population also means that more food is needed, leading to the increase in number and density of livestock, especially in intensive farming [1]. As a result, the demand for pharmaceuticals used for veterinary purposes also rises. Besides, the continuous progress of medical science, improvements in research and development, high health-care investments and the availability of the global market contribute to the development of more and more new drugs and their wider availability [2,3].

However, a huge environmental concern arises when pharmaceutical substances enter natural aquatic ecosystems [1,2,4,5]. Approximately 2000 pharmaceutically active compounds are administered globally in prescription and non-prescription medicines, as well as veterinary drugs [1,6]. These active pharmaceutical ingredients (APIs) are complex molecules that belong to micropollutants, as they are found in the range of ng/L–µg/L in water bodies [7]. APIs are being found worldwide (Figure 1) in surface water, groundwater, seawater, drinking water, sediment, soil, manure and biota [3,6,8,9]. Such a widespread occurrence of pharmaceuticals is a result of their high consumption and increased production [1].

1.1. Antibiotics Occurrence in Aqueous Environment

Among all the pharmaceuticals, antibiotics pose one of the most significant concerns. They are widely used to treat human infectious diseases, as well as for the treatment and prevention of infections in animals, apart from promoting their growth [10,11]. Approximately 100,000 tons [12,13] of antibiotics are manufactured annually in the world, and 50–60% of this amount is spent on veterinary purposes [8,13,14].

Antibiotics are chemotherapeutic agents that kill or inhibit the metabolic activity of various microorganisms, such as bacteria, fungi, viruses, microalgae or protozoa [9,11,15,16]. They can be natural, synthetic or semi-synthetic compounds, which show antibacterial, antiparasitic or antifungal activities [11,16]. There are different groups of antibiotics, depending on their mechanism of action or chemical structure. They are, for example, β-lactams (amoxicillin, cefalexin), aminoglycosides (kanamycin, neomycin, streptomycin), diaminopyrimidines (trimethoprim), glycopeptides (teicoplanin, vancomycin), macrolides (azithromycin, erythromycin, tylosin), quinolones (enrofloxacin, ciprofloxacin, levofloxacin, norfloxacin, ofloxacin), sulfonamides (sulfamethoxazole, sulfathiazole, sulfamethazine), tetracyclines (tetracycline, oxytetracycline) and others [4,10,11,14]. According to the assessment of antibiotic consumption in 76 countries, global consumption grew to 42.3 billion defined daily doses (DDDs) between 2000 and 2015, which is a 39% increase [6,16] and is still growing. The reason for that is obviously growth of population and, therefore, increasing demand of the animal protein [16]. As a result, antibiotics are being found in natural water bodies across the whole world [4,9,12]—concentrations of different antibiotics from various classes found in surface waters, in areas away from wastewater treatment plants (WWTPs) and industrial production sites, are depicted in Figure 2.

Although individual antibiotics are always detected at quite low concentrations, ranging from ng/L to μg/L [10,17], the simultaneous presence of several antibiotics even at low concentrations might result in their much stronger cumulative effect [14,16]. Properties of antibiotics clearly demonstrate why their presence in natural aquatic environment is highly undesirable: antibiotics are, first of all, persistent, so that they are non-degradable and are able to withstand natural decomposition [1,6,18]. Secondly, they are highly mobile, bioaccumulative and toxic [1,2,19]. Antibiotics are deliberately designed to affect living organisms and cause pharmacological response even at low doses. This means that aquatic organisms exposed to such pharmaceuticals may develop negative chronic effects, which affect their behavior, growth and reproduction [16].

1.2. Sulfamethoxazole, Trimethoprim and Ciprofloxacin as Antibiotics of Great Concern

Among the antibiotics, sulfamethoxazole (SMX), trimethoprim (TMP) and ciprofloxacin (CIP) are the most often detected substances [3,12,15] (Table 1, Scheme 1, their physicochemical properties are presented in Table 2). These pharmaceuticals have been found at high concentrations in influents and effluents of wastewater treatment plants, surface waters (rivers and seas), groundwater, and drinking water, not only across Europe, but worldwide [15,20,21]. Apart from that, maximum concentrations of SMX, TMP and CIP detected in Africa were ~100 times, ~54 times and ~125 times, respectively, higher than those in Europe [22]. This happens because these antibiotics are widely used in human and veterinary medicine, have the largest numbers of manufacturers [22], are toxic and persistent in the aquatic environment [4,16].

Sulfonamides are one of the oldest antibiotics groups that is still widely used [6]. All substances from this group have been detected in WWTPs effluents, surface water and groundwater, as well as in drinking water [10]. Sulfamethoxazole is the most often detected sulfonamide in the environment [6,10,23]. It is usually prescribed together with trimethoprim to treat various infectious diseases in humans [10,25,26]. Fluoroquinolones are also very often found in water bodies, especially ciprofloxacin. Hydrophilic properties of this antibiotic class make them highly mobile in the aquatic environment. This, together with their high consumption both in human and veterinary medicine, contributes to wide occurrence of ciprofloxacin, from WWTP effluents to drinking water [6,27,28]. Sulfamethoxazole, trimethoprim and ciprofloxacin have been identified in hospital wastewater with concentrations up to 35.5 µg/L [17,29,30]. Besides, sulfanomides and fluoroquinolones have demonstrated very high toxicity of aquatic organisms, especially to cyanobacteria, freshwater algae and duckweeds [10]. These organisms play a crucial role in aquatic ecosystems [10], and therefore whole ecosystems can be disturbed when antibiotics are released into them [31]. Such a wide spreading of these antibiotics resulted in adding ciprofloxacin to the second EU watch list of substances for union-wide monitoring in the field of water policy (EU Decision, 2018/840 of June 5, 2018), and sulfamethoxazole and trimethoprim were selected to be included in the third watch list [32].

1.3. Antimicrobial Resistance as a Global Health Crisis

The presence of antibiotics in natural aquatic environments raises many concerns, first of all because of the development of antimicrobial resistance (AMR). If such pharmaceuticals are continuously released into water bodies, bacterial populations undergo selective pressure and therefore develop antimicrobial resistance in order to survive (Figure 3) [1,14]. This is an adaptive genetic trait of bacterial populations, enabling them to resist the effect of the drug that used to successfully kill or inhibit these microorganisms [1,2,33]. The misuse and overuse of antibiotics in human medicine, as well as inappropriate and frequent use of antibiotics as therapeutics and growth promoters in animal farming, contribute to the development and wide spread of antibiotic resistance [34].

The fast development and spread of resistance to antibiotics is a significant health issue because it affects the ability to treat various infections [5]. Antibiotic resistance to most of the common antibiotics has already been developed, and even last-generation antibiotics have become less effective to treat bacterial infections [10,14]. Thus, the shortage of treatment options develops [14]; higher medical costs are required to create new antibiotics; and mortality increases [10]. Antibiotic resistance is a global health crisis with an enormous potential for health and economic consequences [1]. Currently, 700,000 deaths per year are associated with antibiotic-resistant infections, but it is estimated that this number could increase to 10 million per year by 2050 [1,35].

Therefore, each time that antibiotics are released into the natural environment, it contributes to the development of antimicrobial resistance in microorganisms and thus the spread of antibiotic resistant bacteria (ARB) and antibiotic resistance genes (ARGs). ARGs are contaminants of emerging concern, and they are not commonly monitored in the environment. However, they have strong potential in leading to negative ecological and human health effects [16,36]. Unfortunately, there are no legal regulations that define permitted levels of antibiotics or ARB and ARGs allowed to be released into the environment [2,16,17].

1.4. Pathways of Antibiotics Release into Aquatic Environment

Discharge of antibiotics into the natural aquatic environment occurs as a result of (1) industrial production of antibiotics; (2) consumption and excretion of antibiotics by humans and animals; (3) improper disposal of unused or expired antibiotics [1]. Hotspots of antibiotic discharge in the environment include wastewater treatment plants (receiving effluents from households and hospitals), antibiotic manufacturing plants, agriculture and aquaculture (Figure 4) [1,6,9,16]. Disposed antibiotics from households and hospitals end up in landfill, where they can eventually leak to pollute ground water. The usage of farming animal manure and WWTP sludge contaminated with antibiotics as fertilizer also results in the occurrence of antibiotics in ground water [1]. This is because 30 to 90% of oral antibiotic doses taken both by humans and animals are excreted in its unmetabolized form [1,6,17]. This means that antibiotics can be released into the environment as entirely biologically active substances [37,38]. In groundwater, under anoxic conditions, antibiotics remain unchanged for a very long time or might undergo minor degradation and produce even more toxic metabolites [6]. The consumption of groundwater (most often used as a source of drinking water) contaminated with antibiotics, ARB or ARGs might lead to the development of antibiotic resistance in humans [6]. Industrial and agricultural runoff (rainfall and snowmelt runoff) also results in antibiotic occurrence in natural water bodies [2,8,13,39].

Wastewater treatment plants are often recognized as the main source of pharmaceuticals discharge into the environment [37,40]. The reason for this lies in the common design of WWTPs.

Wastewater treatment at WWTP usually comprises a combination of physical, chemical and biological treatment processes. First, physical methods are utilized: large sediments, grit and grease are removed during pre-treatment and primary treatment stages using physical barriers and gravity forces [35]. Physical processes include, for example, sedimentation, flocculation, filtration, etc. [35,41,42]. Afterwards, secondary treatment place-biological processes are employed in order to remove organic content [35,39,42] through aerobic and anaerobic processes. For this purpose, biological reactors are utilized, in which the biochemical degradation of organic substances takes place in a similar way to the one occurring in natural aquatic systems [35]. Biological processes employ various microorganisms (bacteria, fungi, algae) and plants to transform biodegradable organic substances into simple products [41,42]. Biological methods utilize biomass in different forms: suspended (activated sludge), fixed (trickling filters) or suspended in aquatic systems (aerated lagoons) [35,42]. Tertiary treatment takes place in exceptional cases, only when additional water purification is required; for example, the disinfection or removal of specific nutrients or toxic compounds [35]. For this purpose, different chemical and physical methods are used: chlorination (most often), ozonation, UV treatment, membrane filtration, adsorption processes [35,43].

Although common treatment processes are very efficient to remove suspended solids, biodegradable organic compounds, inorganic matter and pathogens, they fail to remove pharmaceuticals and antibiotics, in particular [1,35,44,45]. Biological treatment, which is the foundation of any WWTP, is simple and easy to operate, cost-effective and does not generate toxic by-products [39,41,46]. However, the degradation efficiency of non-biodegradable antibiotics by this method is low [41,42,45,46,47]. Additionally, what is worse, antibiotics present in wastewater are able to change the composition of microbial communities in bioreactors [8,35] and make bacteria there multi-resistant to antibiotics. As a result, discharge from WWTPs containing antibiotic resistant bacteria and genes, contributes to the wide spread of AMR in the aquatic environment. This poses a significant threat to human health if receiving water bodies are used as a source of drinking water or for recreational purposes [48].

In this review, different advanced water treatment techniques, which are capable of the effective removal of antibiotics when applied in WWTP, are briefly introduced and compared. Among them, various advanced oxidation processes are analyzed, and heterogeneous photocatalysis is highlighted. The fundamentals and mechanism of heterogeneous photocatalysis over TiO₂ are discussed in detail. Surface modification strategies of TiO₂ nanomaterials in order to improve their photoactivity are presented. With this in view, the focus of the present review is on the removal efficiency of SMX, TMP and CIP from the aqueous environment by photocatalysis over various recently developed unmodified and modified TiO₂ materials. The role of different operating parameters on the removal efficiency is outlined. The ecotoxicity of the final products is also considered. Finally, further research needs are presented. We expect that the present review will provide some new ideas for the design and development of TiO₂-based photocatalysts to be successfully applied in antibiotics elimination from the aquatic environment.

2. Water Treatment Techniques for Degradation of Antibiotics

It is clear from the previous section that conventional wastewater treatment plants have not been designed to eliminate pharmaceuticals from wastewater, and therefore they are unable to provide the efficient treatment of water contaminated with antibiotics [35,41,46,47,49]. Therefore, the application of advanced water treatment technologies is required in order to achieve the effective removal of pharmaceuticals in aqueous systems. Among these technologies, adsorption, membrane filtration and advanced oxidation processes gain a lot of interest. Short descriptions, advantages and drawbacks are compiled in Table 3.

As follows from the information in Table 3, physical processes, such as adsorption and membrane filtration, have one significant limitation—the production of secondary waste. The problem is that pollutants are transferred from the liquid phase of wastewater to the solid phase of adsorbents in the case of adsorption process, or are collected on the membrane surface in the case of membrane filtration without any decomposition [44]. This means that pharmaceuticals do not undergo any changes in chemical structure and therefore, still retain their harmful potential [35]. Moreover, the obtained waste should be properly disposed of, which makes a treatment technology more expensive. Thus, the main advantage of advanced oxidation processes is the obvious, efficient degradation of pollutants with the possibility of their complete mineralization.

Currently, AOPs are widely studied by different research groups, as these methods give very promising results and have high potential for broad application. Carried out studies state that the best performance is associated with the antibiotic removal strategies that involve combined processes [41,46]. They include, for example, a combination of AOPs with biological treatment [43,57] or membrane filtration [42,57]. However, further improvements are needed, mainly in the development of new materials for AOPs [46], in order to make these processes cost-effective while retaining high efficiency.

2.1. Advanced Oxidation Processes

Advanced oxidation processes (AOPs) are powerful technologies used in wastewater treatment to degrade organic pollutants, including various pharmaceuticals [58]. Most common AOPs include the Fenton process, heterogeneous photocatalysis, ozonation, electrochemical oxidation, etc., and are briefly described and compared in Table 3. As can be seen from the table, AOPs include many different processes, and the only feature that unites all of them is the production of highly reactive species, for example, hydroxyl radicals (^•OH), ozone molecules (O₃), superoxide radicals (O₂^•−), hydrogen peroxide molecules (H₂O₂), etc.

The efficient generation of hydroxyl radicals (^•OH) in situ is the purpose of AOPs, as it defines their performance. These radicals are the second highest powerful species after fluorine with oxidizing potential E₀ = 2.80 V [52,55]. ^•OH radicals attack organic molecules rapidly and non-selectively, oxidizing them and thus transforming them into more biodegradable and less toxic substances. What is more, ^•OH radicals are also capable of mineralizing organic contaminants, yielding CO₂, H₂O and inorganic ions. In this case, organic pollutants are completely destroyed and therefore no further post-treatment is required.

The combination of different AOPs is gaining lots of interest due to enhanced process efficiency. For example, the performance of homogeneous Fenton reaction, being the most widely applied AOP [55], can be significantly increased through combination with UV light (photo-Fenton [23,41]), with electrochemical process (electro-Fenton [55]), or even with both (photo-electro-Fenton [42]).

After Fenton process, heterogeneous photocatalysis is the second most widely applied AOP [55]. This technology has numerous advantages over the other types of AOPs, mainly simplicity and sustainability. The possibility of the utilization of solar energy as an abundant, clean and renewable light source makes solar photocatalysis a green and sustainable technology for wastewater treatment. It is also very cost-efficient compared to the technologies utilizing artificial sources of light or electrodes. Besides, it was found that solar photocatalysis-based AOPs have the lowest global warming potential compared to the other wastewater treatment methods [57].

2.2. Heterogeneous Photocatalysis over TiO₂

Heterogeneous photocatalysis employs various semiconducting materials as photocatalysts, for example, metal oxides (TiO₂, SnO₂, Fe₂O₃, WO₃, ZnO, Ag₃O₄ [35,44,53,59]), metal sulfides (ZnS, CdS [44,53]), as well as g-C₃N₄-based materials [56]. Among these, TiO₂ is the most widely utilized photocatalyst due to its unique properties. They include high physical and chemical stabilities, high photocatalytic activity, nontoxicity, large abundance and low cost [60,61]. Besides, TiO₂ employment is preferred in the form of nanostructures, which have improved surface chemistry and increased surface area [62]. This results in the faster and more efficient mineralization of organic pollutants.

What makes TiO₂ such an effective photocatalyst, is the favorable thermodynamic processes taking place in TiO₂ photocatalysis. It is known that the conduction band (CB) energy of TiO₂ anatase (with a band gap E_g = 3.2 eV) is –0.51 V [63] at neutral pH, while valence band (VB) energy is 2.69 V under the same conditions. This means that CB energy level is more negative than potential of oxygen reduction (E₀(O₂/O₂^•⁻) = −0.33 V), and that VB energy level is more positive than the potential of water oxidation (E₀(H₂O/^•OH) = 2.29 V). This makes it possible to obtain highly reactive oxidizing species-superoxide radicals O₂^•⁻ and hydroxyl radicals ^•OH in photocatalytic process over TiO₂, which non-selectively degrade and mineralize organic pollutants.

In general, a photocatalytic process over TiO₂ comprises the following stages [62,64]:

Adsorption of organic pollutants on TiO₂ surface;
The photocatalytic degradation of the adsorbed organic pollutants via oxidation-reduction reactions with photogenerated electrons, holes and reactive species (depicted in Figure 5);
Desorption of degradation products.

Therefore, photocatalytic process itself is initiated when TiO₂ absorbs a photon from the light source with the energy (hv), equal to or greater than the band gap (E_g) of the photocatalyst [65]. This makes the photocatalyst electrons e⁻ excited, so that they move from the valence band to the conduction band, and holes h⁺ are generated in the valence band (Equation (1)) [45,54].

TiO₂ + hv → e⁻_CB+ h⁺_VB (if hv ≥ E_g)(1)

After excitation, separated charge carriers (e⁻ and h⁺) might recombine in the bulk with an energy release (Equation (2)) [45]. Otherwise, they migrate to the surface of TiO₂, where they participate in redox reactions with organic pollutants previously adsorbed on the photocatalyst surface. The recombination of the charge carriers on the catalyst surface might also take place.

e⁻_CB + h⁺_VB → energy(2)

Photogenerated holes have high oxidizing ability: they degrade organic compounds directly or oxidize water, so that hydrogen ions (H⁺) and hydroxyl radicals (^•OH) are generated (Equation (3)) [18,53].

h⁺_VB+ H₂O → H⁺ + ^•OH(3)

Besides, holes oxidize OH⁻ at the catalyst surface, and more ^•OH radicals are produced (Equation (4)).

h⁺_VB + OH⁻ → ^•OH(4)

^•OH radicals being very powerful oxidizing species degrade organic pollutants to carbon dioxide and water (Equation (5)) [54].

^•OH + organic pollutants → CO₂ + H₂O(5)

Photogenerated electrons, in turn, have high reducing ability. Thus, electrons react with O₂ to form superoxide radicals O₂^•− (Equation (6)) [18,53].

e⁻_CB + O₂ → O₂^•⁻(6)

These superoxide radicals react with hydrogen ions (H⁺) and generate hydroperoxyl radicals ^•HO₂ (Equation (7)), which further produce hydrogen peroxide H₂O₂ (Equation (8)).

O₂^•⁻ + H⁺ → ^•HO₂(7)

^•HO₂ + ^•HO₂ → H₂O₂ + O₂(8)

Hydrogen peroxide H₂O₂ might react with superoxide radicals O₂^•⁻ (Equation (9)) or is decomposed by the light (Equation (10)), and ^•OH radicals are produced in both cases.

H₂O₂ + O₂^•− → ^•OH + OH⁻ + O₂(9)

H₂O₂ + hv → ^•OH + ^•OH(10)

Generated reactive species (O₂^•−, ^•HO₂) are capable of degrading and mineralizing any organic compounds to CO₂ and H₂O (Equations (11) and (12)) [51,66].

O₂^•⁻ + organic pollutants → CO₂ + H₂O(11)

^•HO₂ + organic pollutants → CO₂ + H₂O(12)

In TiO₂ photocatalysis, the oxidation of organic substances might take place either through aliphatic oxidation (^•OH radicals remove H⁺ from the molecule), or through aromatic oxidation (O₂ and O₂^•⁻ oxidize aromatic molecule) [62].

2.3. Doping and Modification of TiO₂

Although TiO₂ has high photocatalytic activity, there are two main drawbacks limiting its wide practical application [67,68]. These limitations include unfavorable dynamics of photogenerated charge carriers (fast recombination of electrons and holes in bulk and on the photocatalyst surface) and poor visible light harvesting [63]. Wide band gap of TiO₂ (3.0–3.2 eV [63]) limits TiO₂ usage to UV light only (wavelengths shorter than 390 nm) according to Equation (13) [63]:

(13) $λ = \frac{h \times c}{E_{g}} \approx 390 n m,$

where h is the Planck constant, c is the speed of light, E_g is the band gap of the photocatalyst (3.2 eV for TiO₂ anatase) [63].

It is known that UV light accounts for only 3–5% of the solar light, while visible light constitutes 45% of the solar spectrum [18,54]. It is very important that process is cost-effective and sustainable when implemented on an industrial scale. If heterogeneous photocatalysis over TiO₂ cannot be carried out under visible light, expensive UV light sources have to be used [18,44]. It is obvious that for successful TiO₂ employment under visible light, the wavelength adsorption of the photocatalyst should be extended into the visible region, thus making TiO₂ active in visible light [44,66].

Therefore, in order to prevent the recombination of electrons and holes and ensure the effective utilization of visible light, particular measures should be taken. The separation of charge carriers should be enhanced, their lifetime should be prolonged, photocatalyst band gap should be narrowed and TiO₂ surface area should be increased [63,69]. For this purpose, various strategies of TiO₂ surface modification have been proposed [55]. They include mainly metal and non-metal doping [18,44,54,70], metal and non-metal co-doping [44], coupling with other semiconductors [18,44,66], constructing heterojunctions [54,63] and Z-schemes [45], dye sensitization [44].

In metal doping, metal ions (transition metals, noble metals, rare earth metals) [18,69] substitute for Ti⁴⁺. This suppresses electron-hole recombination, effectively reduces the photocatalyst band gap, enhances visible light harvesting and improves photocatalyst surface morphology [18,45,71]. Nonmetal dopants (N, C, S, B, F) also reduce the band gap and are stated to be more effective in making TiO₂ active in visible light [18,71]. In co-doping, TiO₂ is doped with two or more metals, nonmetals or both metals and nonmetals, in order to overcome the limitations of a single-doped photocatalyst [18].

Coupling TiO₂ with another semiconductor (SnO₂, ZnO, SiO₂, CdS) to obtain heterojunctions also contributes to the better separation of the charge carriers and improves utilization of visible light [18,45,64]. In the case of Z-scheme photocatalysts, band structure remains similar to the one of heterojunctions. However, the transfer mechanism of electrons and holes changes. In particular, recombination of the charge carriers with weaker redox abilities take place, while electrons and holes with stronger redox abilities are separated and preserved [45]. Another way to enhance the separation of electrons and holes is to create vacancy engineered photocatalysts containing crystal defects-oxygen vacancies. Thus, a new energy level is formed between the valence band and conduction band and, as a result, band gap is reduced, and visible light response is improved [45]. Dye-sensitized photocatalysts are also used to generate more electrons and therefore enhance photocatalytic performance [70,71].

Still, there are many other parameters affecting the efficiency of the photocatalytic degradation process; for example, organic pollutant concentration, photocatalyst dosage, light intensity, pH of reaction solution, etc. [51]. Thus, process conditions should be properly designed, numerously tested and optimized, so that maximum performance can be achieved while keeping costs to a minimum.

3. Parameters Affecting Efficiency of Photocatalytic Degradation over TiO₂-Based Photocatalysts

3.1. Effect of Antibiotic Concentration

The initial concentration of a contaminant is a very important parameter in a photocatalytic process. Researchers study its effect on the photodegradation of pharmaceuticals by varying the initial concentration of target antibiotics under other similar conditions.

The effect of initial concentrations of SMX was studied by Tiwari et al. [72] in photocatalytic removal over thin film photocatalyst Ag0(NP)/TiO₂(B). Researchers reported a decrease in photodegradation efficiency from 57% to 20% when SMX concentration was increased from 0.5 mgL⁻¹ to 15.0 mgL⁻¹.

A similar inverse relationship between the photodegradation efficiency and antibiotic concentration was observed and reported by Alfred et al. [73], Ioannidou et al. [74], Cai and Hu [75], Sarafraz et al. [24], Karim and Shriwastav [76], Wu et al. [77] and Hassani et al. [78], who used different photocatalyst systems for removal of different target antibiotics.

Researchers attribute the hindered decomposition of pharmaceuticals when antibiotic concentration is increased to a number of reasons. First of all, higher contaminant concentration means that higher number of contaminant molecules are adsorbed on active sites of the photocatalyst [75,77]. As a result, generation of reactive species is suppressed [75]. Secondly, as antibiotic concentration increases, more byproducts and intermediates are produced that compete with antibiotic molecules for a limited number of active sites on photocatalyst surface and reactive species present [24,73,77,78]. In addition, solution transmittance might be reduced when pharmaceutical concentration in the solution is increased. Weaker transmittance results in longer pathways for photons to get to photocatalyst, so that lower number of photons is adsorbed on photocatalyst surface and, therefore, a lower number of reactive species are generated [77,78]. Moreover, some antibiotics (for example, CIP) adsorb photons themselves, so that a lower number of photons are available for the photocatalyst [24].

3.2. Effect of Catalyst Concentration

The influence of photocatalyst concentration on removal efficiency of different antibiotics has been studied by many researchers. For example, Xie et al. [79] performed photocatalytic experiments under visible light irradiation utilizing different concentrations of a photocatalyst (Zn-TiO₂/pBC): 0.625 gL⁻¹, 1.25 gL⁻¹ and 2.5 gL⁻¹. The optimal concentration was found to be 1.25 g, resulting in 81.21% (k_app = 0.0087 min⁻¹) removal efficiency of 10 mgL⁻¹ SMX after 3 h of irradiation. For comparison, 0.625 gL⁻¹ and 2.5 gL⁻¹ catalyst concentrations could achieve 57.73% (k_app = 0.0043 min⁻¹) and 76.28% (k_app = 0.0075 min⁻¹) SMX degradation under the same conditions, respectively. The same results were obtained previously by Zhang et al. [80] who tested TiO₂/pBC photocatalyst for SMX removal under the same conditions but utilizing UV light instead of visible.

Obviously, increased catalyst concentration results in improved photocatalytic performance only up to a point, and any catalyst dosage beyond the optimal value leads to a decrease in removal efficiency. Such results were obtained also in the works of Ioannidou et al. [74], Kim and Kan [81], Chiang and Doong [82], Sarafraz et al. [24], Karim and Shriwastav [76], Hassani et al. [78] and Abellan et al. [83].

Researchers explain this phenomenon as follows: an increase in photocatalyst concentration increases the number of active sites available on the catalyst surface [82]. A higher number of active sites absorb more photons and, as a result, a higher number of reactive species (for example, hydroxyl radicals) are produced, which take part in antibiotic mineralization [24,74,79]. As a result, higher antibiotic removal and process efficiency can be achieved. However, excessive catalyst dosages limit the utilization of light because of greater photocatalyst aggregation and increased turbidity of the reaction solution [78,79,80]. The first one leads to a decrease of the catalyst active surface area, while the latter, to the reduced penetration of light through the light scattering effect [24,76,80,82]. This results in the generation of smaller number of reactive species and, therefore, photocatalytic efficiency decreases.

However, the situation is the opposite when it comes to mineralization efficiency. Kim and Kan [81] studied the effect of biochar/TiO₂ catalyst concentration (2.5–10 gL⁻¹) on the removal efficiency of 0.1 L SMX solution (10 mgL⁻¹). Although increased catalyst dosage led to a slight decrease in photocatalytic SMX removal from 86% to 76%, COD removal efficiency improved from 35% to 65%. Authors explain this phenomenon as follows: SMX can be better degraded by photolysis in UV-C light (SMX absorbs light at 250–270 nm [81,84]), than by photocatalysis. High photocatalyst loading results in high UV light absorption on TiO₂ surface. As a result, interaction between UV light and SMX decreases, and therefore photolysis efficiency also decreases. However, higher photocatalyst dosage contributes to the enhanced generation of ^•OH radicals and, therefore, significantly improves COD removal efficiency that is impossible under photolysis only. The same conclusions were drawn by Nasuhoglu et al. [85], Gong and Chu [86] and Choi et al. [87].

3.3. Effect of pH

pH plays a very important role in the photocatalytic decomposition of organic pollutants. Its change affects the concentration of H⁺ and OH⁻ ions in a solution [88], which has a direct influence on the generation of reactive species (photogenerated holes h⁺, hydroxyl radicals ^•OH, superoxide radicals O₂^•−) and surface chemistry (surface charge of photocatalyst, reactivity of pollutants and electrostatic interaction between pollutants and photocatalyst) [76,77,88].

Photocatalysis efficiency greatly depends on the surface properties of the photocatalyst. The surface charge of TiO₂ in aqueous solutions depends on the solution pH and zero-charge point of TiO₂. The point of zero charge (pH_PZC) is pH at which catalyst surface is uncharged and is calculated as follows [63]:

pH_PZC = 1/2 (pK_a1 + pK_a2)(14)

Ti–OH₂⁺ ↔ Ti–OH + H⁺ (pK_a1)(15)

Ti–OH ↔ Ti–O⁻ + H⁺ (pK_a2)(16)

pH_PZC of TiO₂ P25 is reported to be ~6.2 (pK_a1 = 4.5 and pK_a2 = 8) [63]. Thus, TiO₂ is positively charged at pH < 6.2 and negatively charged at pH > 6.2 [75].

Alfred et al. [73], Wang et al. [89], Yuan et al. [84], Xie et al. [79], Tiwari et al. [72], Zhang et al. [80], Cai and Hu [75] studied the influence of pH on the photocatalytic degradation of SMX and observed decrease in SMX removal efficiency with an increase in initial pH. For example, the kinetic parameters of SMX photocatalytic degradation gradually decreased from k = 0.0087 min⁻¹ to k = 0.0049 min⁻¹ when pH increased from 5.03 to 10.97, respectively [79]. This can be explained when the ionic states of both pharmaceutical compound and photocatalyst are considered. The pK_a values of SMX are reported to be as follows: pK_a1 = 1.85 ± 0.30 and pK_a2 = 5.60 ± 0.04 [75,84]. It means that SMX exists in cationic form when pH < 1.85, in anionic form when pH > 5.6, and in neutral form when 1.85 < pH < 5.6. When SMX and TiO₂ surfaces are both positively or negatively charged, there is no electrostatic attraction between them. As a result, adsorption cannot contribute to SMX removal, and only photocatalytic degradation takes place [90]. Besides, when initial pH is close to pH_ZPC of photocatalyst, catalyst aggregation might occur, resulting in reduced surface area and decreased photodegradation efficiency [84]. Results suggest that the neutral form of SMX (pH 3–5) is more easily photodegraded than its charged species [73,75], as SMX shows stronger light absorption and higher photochemical reactivity in its neutral form [84].

On the contrary, Mourid et al. [90] and Dlugosz et al. [91] reported higher rates of photocatalytic SMX removal at higher pH than at lower pH. Such contradicting results of pH effect on SMX photocatalytic removal demonstrate importance of taking into account the role of reactive species. When the main reactive species in the degradation process are holes (h⁺), low pH values favor efficiency of photocatalytic decomposition [89]. On the other hand, if the main reactive species are ^•OH, then an increase in pH leads to an ^•OH content increase and, therefore, enhances photodegradation efficiency [84,89,90]. For example, in their research, Wang et al. [89] found that h⁺ played major role in the photocatalytic degradation of SMX, while de Matos Rodrigues et al. [92] reported ^•OH and O₂^•− to be the main reactive species in SMX decomposition. This clearly demonstrates that all the experimental conditions should be considered in order to evaluate the pH effect on SMX photocatalytic removal.

Cai and Hu [75] found that the photocatalytic removal rate of TMP (pK_a = 6.7) was almost similar at pH 3, 5, and 9 (~90%) and lower at pH 7 (68%). The authors attributed such a performance to the effect of pH on the absorption spectrum of the antibiotic. Luo et al. [93] also did not observe any significant influence of pH on TMP photodegradation.

The influence of pH on the decomposition rate of CIP was studied by Sarafraz et al. [24] who carried out photocatalytic degradation of CIP at initial pH of 4, 5.5, 7, 8.5 and 10 and obtained the following kinetic parameter values: 0.0237, 0.0417, 0.0763, 0.028 and 0.0196, respectively. The highest removal rate of CIP was observed at pH 7. A similar trend (increase in removal efficiency with pH increase up to a point and a subsequent decrease in efficiency with further pH increase) was also reported by Hu et al. [94] (highest efficiency at pH 6.3), Karim and Shriwastav [76] (highest efficiency at pH 7), Hassani et al. [78] (highest efficiency at pH 5), Gad-Allah et al. [95] (highest efficiency at pH 5.8). Such results can be explained by the ionic values of CIP and the surface charge of different photocatalyst used. It is known that CIP has two pK_a values: pK_a1 = 5.9 and pK_a2 = 8.9 [24,77,78]. It means that CIP is present in cationic form CIP⁺ when pH < 5.9, in zwitterionic form CIP^± when 5.9 < pH < 8.9 and in anionic form CIP⁻ when pH > 8.9 [24,77]. If TiO₂ with pH_PZC = 6.2 is used as a photocatalyst, then, in alkaline conditions when both TiO₂ surface and CIP are charged negatively (in acidic conditions, both charged positively), electrostatic repulsion occurs between them, decreasing removal efficiency [24,78,94]. Results obtained in the abovementioned studies demonstrated that photocatalytic decomposition efficiency was significantly improved at neutral pH, when CIP was present in its zwitterionic form [24,76]. Besides, predominant reactive species in the process should also be taken into account [77]. Sarafraz et al. [24], Hu et al. [94], Li and Hu [96] found that h⁺ and ^•OH played a major role in the CIP degradation process. At the same time, Li et al. [97], Huang et al. [98] and Du et al. [99] confirmed h⁺ and O₂^•− to be the main reactive species, while Gan, Zhang and Xiong et al. [100] reported all three radicals (h⁺, ^•OH and O₂^•−) to play a significant part in CIP photocatalytic decomposition.

3.4. Effect of Presence of Inorganic Ions

Inorganic anions naturally present in water bodies have a significant impact on the photocatalytic removal of organic pollutants. Their presence affects both working pH and ionic strength of a reaction solution that eventually leads to the change in concentration of active sites on the photocatalyst surface, catalyst adsorption capacity and electrostatic interaction between organic molecules and the catalyst surface [84]. Besides, inorganic anions compete with antibiotic molecules for the same active sites on the photocatalyst surface [79,84]. These anions are also able to react with reactive species-photogenerated holes and hydroxyl radicals, decreasing their amount in the system. When h⁺ are captured by inorganic anions, the production of ^•OH is hindered and, therefore, the decomposition rate of antibiotics decreases [24,78,79,89]. The presence of some inorganic ions might also result in photocatalyst agglomeration, reducing the photocatalytic efficiency of the process [79].

Yuan et al. [84] studied the impact of five inorganic anions (Cl⁻, SO₄²⁻, HCO₃⁻, H₂PO₄⁻, HPO₄²⁻) on the photocatalytic removal of SMX over TiO₂ P25. The authors reported the inhibition of SMX oxidation in the order as follows: HCO₃⁻ > SO₄²⁻ > Cl⁻ with increased inhibition when concentrations of the anions increased. Thus, k_app values decreased from 0.544 (without inorganic anions) to 0.485 (1 mM) and 0.212 (3 mM) when HCO₃⁻ anions were added; to 0.432 (1 mM) and 0.314 (3 mM) when SO₄²⁻ anions were added; and to 0.530 (1 mM) and 0.501 (3 mM) when Cl⁻ anions were added. k_app value also decreased to 0.416 (0.1 mM) and 0.305 (0.3 mM) when HPO₄²⁻ anions were added. However, when H₂PO₄⁻ anions were present, a slight improvement in SMX photocatalytic decomposition was observed. k_app value increased from 0.544 (without H₂PO₄⁻ anions) to 0.552 and 0.586 when 0.1 mM and 0.3 mM of H₂PO₄⁻ anions were added, respectively. Researchers attributed inhibited SMX decomposition to the suppression of reactive species (h⁺, ^•OH) by anions. On the contrary, the slightly accelerated removal of SMX in case of H₂PO₄⁻ addition was attributed to the effect of a phosphorus-containing anion. It was reported to be able to enhance e⁻ transfer between h⁺ and water molecules to form hydroxyl radicals ^•OH and, therefore, improve SMX decomposition. Similarly, Xie et al. [79] reported the weakened photocatalytic performance of TiO₂ in SMX removal when SO₄²⁻, Cl⁻, NO₃⁻ were present in the reaction system, and Wang et al. [89] when HCO₃⁻ anions were added.

Photocatalytic oxidation of ciprofloxacin was affected in a similar way. Sarafraz et al. [24] found that HCO₃⁻, SO₄²⁻, Cl⁻, and NO₃⁻ anions slightly decreased the efficiency of CIP photocatalytic decomposition. Among the studied anions, HCO₃⁻ anions were reported to cause the most significant inhibitory effect, while Cl⁻ anions were the least significant. Hassani et al. [78] also assessed the influence of HCO₃⁻, SO₄²⁻, I⁻, and Cl⁻ anions on CIP photocatalytic removal and obtained different results regarding the highest inhibition. Researchers reported a decrease in decomposition efficiency from 61.70% to 52.87%, 42.94%, 39.86% and 25.83%, respectively.

3.5. Effect of Presence of Natural Organic Matter

Natural organic matter (NOM) is widely present in natural water bodies in relatively high quantities (0.1–50 mgL⁻¹) [96]. It consists of different aromatic and aliphatic organic compounds, including humic substances [84,89,96] and negatively affects the efficiency of photocatalytic processes. The photocatalytic decomposition of antibiotics is reduced for a number of reasons. First of all, NOM competes with antibiotic molecules for the same active sites on the photocatalyst surface and for the same number of reactive species (^•OH, h⁺ and O₂^•−) present in reaction media [84,87,89,96,101]. Secondly, NOM can absorb light both in the UV and visible range that is also harvested by the photocatalyst, thus creating an inner filter effect [84,96,101]. As a result, photocatalyst performance is significantly reduced. However, the inhibitory impact of natural organic matter depends greatly on the type and concentration of NOM, photocatalyst properties and target antibiotics [84,96,101].

Awfa et al. [101] assessed the photocatalytic oxidation of SMX in the presence of NOM using pure TiO₂ P25 and magnetic carbon nanotube-TiO₂ P25 composites (MCNT-TiO₂). Researchers reported the significantly reduced rate constants of SMX removal from 4.7 × 10⁻² min⁻¹ to 5 × 10⁻³ min⁻¹ over TiO₂ P25; and from 4.9 × 10⁻² min⁻¹ to 7 × 10⁻³ min⁻¹ over MCNT-TiO₂ in the absence and in the presence of NOM, respectively. Yuan et al. [84] studied the effect of three different types of NOM on SMX photocatalytic decomposition and observed decreased antibiotic removal when NOM concentration was increased. Photodegradation rates of SMX oxidation dropped from 0.544 min⁻¹ in the absence of NOM to 0.081–0.141 min⁻¹ in the presence of NOM, depending on the NOM type. Similar results were obtained by Li and Hu [96], who observed a negative influence of different NOM isolates on CIP photocatalytic removal. Sarafraz et al. [24] also carried out research on the influence of NOM different concentrations on CIP photocatalytic removal. Researchers reported the inhibition of CIP decomposition rate from 91.5% to 59% when the concentration of NOM was increased from 5 to 30 mgL⁻¹.

3.6. Effect of Reaction Media

At present, antibiotics can be widely found in natural water bodies and different wastewater effluents [79,102]. These water matrices are complex systems—they contain large amounts of organic matter or various inorganic anions that significantly affect photocatalysis efficiency due to the reasons described above [80,103]. Many researchers have carried out studies on the photocatalytic degradation of pharmaceuticals in different water matrices in order to evaluate their inhibitory effect.

Wang et al. [89] studied the photocatalytic degradation of SMX in natural water bodies under solar light. Authors reported decreased removal efficiency (86% in river water and 90.6% in lake water) compared to 100% removal efficiency in deionized water. Inhibited photocatalytic performance was attributed to the presence of organic matter and high turbidity in natural water that resulted in the lower utilization efficiency of the photocatalyst. Xie et al. [79] also observed a significant decrease in photocatalytic degradation of SMX from ~81% in ultrapure water to ~54% in river water under visible light [79]. Carbajo et al. [102] used natural ground water with reduced carbonates to study the photocatalytic degradation of SMX. The values of apparent kinetic constants decreased from 0.095 min⁻¹ in deionized water to 0.061 min⁻¹ in natural ground water when the home-made TiO₂ photocatalyst was utilized. Hassani et al. [78] also observed the inhibited photocatalytic decomposition of CIP when well water was used as a matrix.

Porcar-Santos et al. [104] carried out the photocatalytic decomposition of SMX using TiO₂ P25 under simulated solar irradiation, both in deionized water and simulated seawater. As a result, SMX degradation was twice as fast in deionized water than in saline water (pseudo-first order rate constants were 0.041 min⁻¹ and 0.020 min⁻¹, respectively). The investigation of the degradation mechanisms revealed a major contribution of reactive halogen species on the photocatalytic degradation of SMX in seawater, while hydroxyl radicals (^•OH) played their main role in photocatalytic SMX decomposition in deionized water. Yang et al. [105] utilized seawater (34‰ salinity), diluted seawater (20‰ salinity) and distilled water to study the effect of salinity on the photocatalytic degradation of SMX. Authors observed a significant decrease in photocatalysis efficiency with an increase in salinity: the remaining amount of SMX increased from 7.7% in distilled water to 34% in diluted seawater and to 49% in seawater. Apart from that, the inhibitory effect of cations and anions present in seawater was studied and reported to be as follows: Ca²⁺ > Mg²⁺ > K⁺ > Na⁺ for cations, I⁻ > SO₄²⁻ > Cl⁻ > Br⁻ for anions. Interestingly, the mineralization rate of the antibiotic was also severely affected. The remaining COD concentration increased from 34% in freshwater experiments to 75% and 84% when salinity was 20‰ and 34‰, respectively.

Malesic-Eleftheriadou et al. [103] carried out photocatalytic experiments with SMX as a target antibiotic using effluent from an urban wastewater treatment plant. The increased complexity of the water matrix resulted in decreased kinetic rates from 0.015 min⁻¹ in distilled water to 0.03 min⁻¹ in wastewater effluent. Ioannidou et al. [74] studied SMX photocatalytic degradation in three different matrices: ultrapure water, secondary treated wastewater and drinking water. Researchers reported slightly decreased removal efficiency in drinking water (apparent rate constant ~0.1 min⁻¹) and significantly inhibited the photocatalytic process in wastewater (apparent rate constant ~0.015 min⁻¹) compared to the results obtained when ultrapure water was used as a matrix (apparent rate constant ~0.13 min⁻¹). Karim and Shriwastav [76] observed a decrease in CIP photocatalytic removal from 77% to 41% when secondary treated wastewater was used as a matrix instead of distilled water. Large amounts of non-target organic compounds in wastewater effluents consumed reactive species and thus interfered with the photocatalytic decomposition of target antibiotics [74,103].

3.7. Effect of Dissolved Oxygen and Oxidants

The concentration of dissolved oxygen is an important parameter that affects the efficiency of photocatalytic processes. High contents of dissolved oxygen enhance the photocatalytic oxidation of organic pollutants for a number of reasons. First of all, oxygen acts as a trap for photoexcited electrons, decreasing the recombination rate of charge carriers [106]. Secondly, the reaction between oxygen and electrons results in the generation of reactive species: superoxide radicals O₂^•−, hydroxyl radicals ^•OH and peroxide radicals ^•HO₂ that improve the oxidation rate of pharmaceuticals [107]. As the concentration of dissolved oxygen constantly changes in natural water bodies [90], many researchers have studied its influence on the photocatalytic oxidation of antibiotics.

Mourid et al. [90] studied the effect of dissolved oxygen on SMX removal by carrying out photocatalytic tests in aerated and de-aerated media. The decomposition of the antibiotic decreased from 80% in aerated medium to 45% in de-aerated medium, proving the significance of dissolved oxygen in the photocatalytic process. Diao et al. [107] obtained similar results when CIP was treated under aerobic and anoxic conditions, as removal efficiency was reduced from 97.2% to 56.1%, respectively. The purging of nitrogen (N₂) gas through the reaction system to remove dissolved oxygen was reported by Khan et al. [108] to have an inhibitory effect on the efficiency. CIP photodegradation dropped from 85.91% to 70.98%. Xekoukoulotakis et al. [106] performed photocatalytic tests under two different conditions. In the first case, an open to air reactor was used, while in the second case, additional oxygen supply was provided. Researchers did not observe a significant influence of mode on CIP photocatalytic degradation. However, mineralization rate was greatly affected, as TOC reduction increased from 76% to 96% when aeration was applied.

The presence of different chemical oxidants can also enhance the efficiency of the photocatalytic process. For example, as Hassani et al. [78] observed in their research, the addition of hydrogen peroxide, potassium iodate and potassium persulfate increased the decomposition efficiency of CIP from 61.70% to 83.36%, 86.43% and 97.88%, respectively. Improved photocatalytic performance was attributed to the increased generation of reactive radicals, mainly ^•OH, through various mechanisms.

3.8. Effect of Light Source

The efficiency of the photocatalytic decomposition of target pharmaceuticals greatly depends on the wavelength of light applied and light intensity [45,88]. An increase in light intensity increases the number of photons in a reaction system. Thus, more charge carriers are generated that take part in photocatalytic oxidation and contribute to an improved removal rate of antibiotics [45]. Taking into account that pure TiO₂ has a band gap of 3.0 eV (rutile phase) and 3.2 eV (anatase phase) [47], it can utilize light in the UV region only (wavelengths below 390 nm) [54]. However, various TiO₂-based photocatalysts can be successfully applied under visible light [109]. Many researchers have studied the performance of TiO₂-based photocatalysts under UV-light [87,110,111,112], visible light [77,79,113,114], simulated solar light irradiation [94,101,104,115] and natural sunlight [116,117,118]. Eskandarian et al. [119] studied the photocatalytic removal of SMX using LED lamps emitting light in different ranges of UV: UV-A (365 ± 10 nm), UV-B (300 ± 5 nm) and UV-C (260 ± 10 nm). Researchers reported 58%, 80% and 100% SMX decomposition efficiency after 3 h of irradiation when UV-A, UV-B and UV-C LED lamps were employed, respectively. The mineralization of the antibiotic followed the same trend, increasing from 35.1% (UV-A) to 59.9% (UV-C). An increase in CIP removal efficiency was observed in the same order (UV-A < UV-B < UV-C) by Hassani et al. [78]. Authors attributed such results to the increase in energy of photons emitted. A higher number of reactive species was produced under UV-C light compared to UV-B and UV-A. Lin et al. [120] compared the photocatalytic performance of the synthesized TiO₂-based photocatalysts under UV irradiation and visible light source. The removal efficiency of SMX was significantly improved when UV light was used due to the higher utilization rate of photons. Although higher rates of antibiotic decomposition can be achieved under UV light [120,121], TiO₂-based photocatalysts have the potential to be efficient enough under solar irradiation [116,118,122] that is a clean and sustainable source of energy [45] and, therefore, is more preferable.

4. Degradation Pathways of Antibiotics

Many researchers have studied degradation pathways of sulfamethoxazole in photocatalytic processes over TiO₂ and TiO₂-based materials, for example, Yuan et al. [84], Gong and Chu [86], Xie et al. [79], Wang et al. [89], Mourid et al. [90], Borowska et al. [123], and Yu et al. [124]. They reported three degradation reactions to be the main transformation pathways: hydroxylation, cleavage of S-N bond and isoxazole ring opening. Although different research groups obtained different amounts and types of photoproducts (from six intermediates over Bi₂O₃-TiO₂/PAC [89] and LDH-TiO₂ [90] to sixteen intermediates over CoFe₂O₄/TiO₂ [86]), the main transformation products of SMX (C₁₀H₁₁N₃O₃S, m/z 253) in the studies are the following: C₁₀H₁₂N₃O₄S, m/z 269; C₁₀H₁₃N₃O₅S, m/z 287; C₈H₉N₃O₄S, m/z 243; C₆H₇N₃O₂S, m/z 197; C₆H₇NO₃S, m/z 173; C₆H₆O₃S, m/z 158; and C₆H₇N, m/z 93. Proposed pathways of SMX photocatalytic oxidation are presented in Scheme 2 (products are labeled as P followed by the corresponding mass of the compound). Photoproducts with low molecular weight might be further oxidized to inorganic substances: SO₄²⁻, NO₃⁻, NH⁴⁺, H₂O and CO₂ [79,90].

Photodegradation routes of trimethoprim in TiO₂ photocatalysis are less studied than those of sulfamethoxazole. Sirtori et al. [125] and Samy et al. [126] investigated phototransformation products of TMP when TiO₂ P25 and S-TiO₂ were employed as catalysts, respectively. Both research groups reported hydroxylation, demethylation, and cleavage to be the main transformation pathways in the TMP (C₁₄H₁₈N₄O₃; m/z 291) degradation process. In the hydroxylation reaction, polyhydroxylated compounds with the general formula of C₁₄H₁₉N₄O_3+x (x varying from 1 to 4) are formed. For example, C₁₄H₁₉N₄O₄; m/z 307 and C₁₄H₁₉N₄O₅; m/z 323 occur in this way. In demethylation-hydroxylation reactions, compounds with the general formula of C_14-yH_19-2yN₄O_3+x (y varying from 1 to 3) are formed: C₁₃H₁₇N₄O₅, m/z 309; C₁₃H₁₇N₄O₇, m/z 341; C₁₂H₁₅N₄O₅, m/z 295; C₁₂H₁₅N₄O₄, m/z 279; C₁₁H₁₃N₄O₅, m/z 281; and C₁₃H₁₆N₄O₃, m/z 277. Cleavage of TMP results in the generation of the following products: C₅H₆N₄O, m/z 139; C₅H₆N₄O₂, m/z 155; C₈H₁₀O₄, m/z 171; and C₅H₈N₄O, m/z 141. Besides, trimethoxybenzoylpyrimidine (C₁₄H₁₇N₄O₄, m/z 305) is a product of TMP oxidation. Based on the identification of these phototransformation products, a suggestion of possible TMP degradation pathway is presented in Scheme 3 (products are labeled as P followed by the corresponding mass of the compound).

The photodegradation pathways of ciprofloxacin over TiO₂ and TiO₂-based materials have been studied by many research groups: Wang et al. [114], Sarafraz et al. [24], Manasa et al. [118], Karim and Shriwastav [76], Huerta-Aguilar et al. [109], Li and Hu [96], and Calza et al. [127]. Obtained photoproducts varied greatly and are summarized in an excellent review by Hu et al. [128]. In CIP photodegradation, four main transformation reactions can be outlined: cleavage of piperazine ring, defluorination, decarboxylation and oxidation of the cyclopropyl group. Possible photoproducts in the CIP degradation process are presented in Scheme 4 (products are labeled as P followed by the corresponding mass of the compound).

5. Ecotoxicity of Photoproducts

Considering a wide variety of transformation products occurring during the photocatalytic degradation of antibiotics, evaluation of their toxic effects on ecosystems is highly important for the practical application of the process in water treatment.

The phytotoxicity of photocatalysis products towards plant Lepidium sativum was studied by Borowska et al. [123]. After 60 min of solar irradiation of TiO₂ catalysts doped with noble metals, toxicity was eliminated: germination index increased from 43 ± 4% for initial 1 mgL⁻¹ SMX solution to 93 ± 10% and 110 ± 16% for solutions treated with Pd/TiO₂ and Pt/TiO₂, respectively. Decreased ecotoxicity was attributed to very low concentrations of remaining antibiotics. Negligible toxicity towards Vibrio fischeri and fully removed antibiotic activity against E. coli was achieved by Durán-Álvarez et al. [121], when 30 mgL⁻¹ CIP solution was treated using TiO₂ modified with bi-metallic nanoparticles. This was a result of the complete mineralization of the parent compound after 360 min of photocatalytic treatment under simulated sunlight irradiation.

The antimicrobial activity of 15 mgL⁻¹ CIP solution towards Vibrio fischeri was diminished significantly upon a photocatalytic process with graphitized mesoporous carbon-TiO₂ nanocomposite under UV-C irradiation in the study carried out by Zheng et al. [129]. Resulting non-toxicity was attributed to the almost complete mineralization of the antibiotic after 120 min of reaction. Cai and Hu [75] applied UV/TiO₂ P25 system and observed no acute toxicity (30 min of incubation) of 1 ppm SMX and 1 ppm TMP towards Vibrio fischeri. However, significant chronic toxicity (24 h of incubation) was reported. Sarafraz et al. [24] treated CIP solutions using a N-TiO₂/visible LED system that resulted in significantly decreased toxicity towards Daphnia magna compared to the parent compound. A reduction in antibacterial activity of CIP against Staphylococcus aureus was also obtained by Gan et al. [130], who applied synthesized TiO₂ under artificial sunlight.

Gong and Chu [86] also observed a significant reduction of toxic activity against aquatic organisms when a UV-C/CoFe₂O₄/TiO₂ system for 100 µM SMX degradation is utilized. In the case of green alga Chlorella vulgaris, SMX was suggested to have been transformed into nutrients (source of carbon and inorganic salts) upon photocatalytic process, favoring algae growth. In contrast, SMX and its photoproducts inhibited the feeding of the brine shrimp Artemia salina, demonstrating that although the toxic effect of transformation products was significantly decreased, it was not completely eliminated. High rates of SMX decomposition and mineralization contributed to negligible toxicity towards Daphnia magna and Escherichia coli in the study carried out by Kim and Kan [81]. Researchers utilized UV-C/biochar/TiO₂ system for the photocatalytic degradation of 10 mgL⁻¹ SMX and thus, obtained nontoxic products (nitrates, sulfates and some organic acids) after 6 h of reaction. However, when nitrate and bicarbonate were added into reaction solution, toxicity was highly increased up to 100%, suggesting the cautious application of the given process to the treatment of natural water bodies.

Ioannidou et al. [74] reported a small decrease in antimicrobial activity after photocatalytic treatment with the WO₃-doped TiO₂ P25 catalyst, attributed to the slow mineralization of by-products after 180 min of simulated solar irradiation. Ecotoxic effects of SMX and its transformation products in that case were assessed towards bacterial strains: Escherichia coli and Enterococcus faecalis. During the oxidation process, toxicity first increased from a 50% population reduction of both E. coli and E. faecalis (20 mgL⁻¹ initial SMX solution) to 64% and 57%, but eventually decreased to 45% and 37%, respectively. Murgolo et al. [131] also observed only a partial decrease in toxicity when a synthesized TiO₂ catalyst was utilized under UV-C irradiation for the photocatalytic degradation of a mixture of pharmaceuticals containing TMP. Acute toxicity tests were performed with green alga Selenastrum capricornutum, embryo of fish Danio rerio (FET) and AMES fluctuation test.

However, Nasuhoglu et al. [85] observed the formation of more toxic products towards Daphnia magna than initial 60 mgL⁻¹ SMX solution when the target antibiotic was treated with UV-C/TiO₂ P25 for 13 h. Sirtori et al. [125] also reported the increased inhibition of Vibrio fischeri by TMP transformation products compared to parent compound (20 mgL⁻¹ TMP), both in demineralized water and seawater, when the photocatalytic process with TiO₂ P25 under solar irradiation was carried out. In addition, the toxicity of CIP solutions treated by UV-A/TiO₂ P25 increased towards Vibrio fischeri compared to the original compound in the work of Silva et al. [132]. Although toxicity decreased after 15 min of the photocatalytic process, CIP solutions became very toxic (70% of luminescence inhibition) after 45 min of treatment due to higher quantities of transformation products being generated. An interesting observation was made by Yang et al. [105]: the antimicrobial activity of SMX transformation products against bacteria Vibrio fischeri increased compared to the parent compound (50 mgL⁻¹ SMX) when photocatalytic tests were carried out in distilled water (UV-A/TiO₂ P25 system, 510 min). However, no significant toxicity was observed when tests were performed in different media: seawater with 20‰ and 34‰ salinity. Explanation for such findings might be in decreased photocatalytic decomposition of antibiotic when salinity was increased. Thus, smaller quantity of transformation products resulted in decreased biotoxicity. Xing et al. [133] utilized Toxicity Estimation Software Tool (T.E.S.T.) software to assess the toxicity of initial CIP solution and its photoproducts. Although most of the transformation products were reported to be less toxic than untreated antibiotic, some degradation products were found to be more toxic.

The photocatalytic performance of various TiO₂ and TiO₂-based materials in degradation processes of sulfamethoxazole, trimethoprim and ciprofloxacin, along with photocatalysts characterization and process description, is presented in Table 4.

6. Challenges and Future Research Needs

Among the different advanced oxidation processes, heterogeneous photocatalysis over TiO₂-based materials is one of the most efficient techniques for the removal of antibiotics. However, practical application on a large scale remains a challenge. This process has great potential, but it needs to be improved and optimized to be implemented in industry. The most significant issues are outlined as follows:

Huge variety of TiO₂-based materials with unique features have been synthesized, as well as many different techniques having been reported to synthesize efficient TiO₂-based photocatalysts. However, as the operational cost is crucial for practical applications, there is still a strong need for simple and cost-effective synthesis and modification processes to decrease the cost of the photocatalytic process. This can be achieved by implementing low-cost synthesis methods or by using low-cost materials, or both. The introduction of multiple cheaper modificators might be more cost-effective than introducing one expensive dopant, while retaining high photocatalytic activity. For example, low-cost alternatives to photocatalysts modified with expensive and scarce noble metals could be TiO₂-doped with cocatalysts like Ni and Cu [154,155] or Co and Ni [156]; Mo-doped TiO₂ [157], graphene-doped TiO₂ [158], Cu nanowires decorated with TiO₂ [159].
Special emphasis should be placed on the development of TiO₂-based photocatalysts active under solar light. Considering that high energy costs are associated with the usage of visible and ultraviolet light sources, solar photocatalysis is not only a cost-efficient solution, but also a sustainable one. Besides, further cost reduction can be achieved through the development of highly stable photocatalysts that can be easily separated from aqueous solution (immobilized/supported catalysts) and successfully reused over multiple cycles;
Further research is needed in studying the potential of photocatalytic process over TiO₂-based materials for the successful removal not only of antibiotics, but also of ARB and ARGs from wastewaters;
In order to fulfill industrial application, photocatalytic tests should be carried out on real wastewater samples. The composition of wastewater is complicated: it contains a mixture of different pharmaceuticals and other organic pollutants, as well as inorganic substances. This decreases the degradation efficiency of antibiotics compared to the degradation efficiency of a single antibiotic in purified water. Moreover, there is a need to design and develop pilot plant installations that work under solar irradiation and in the continuous mode of operation. These are essential requirements for industrial treatment systems;
Further investigations should be performed on the efficiency of combined methods. Heterogeneous photocatalysis combined with conventional treatment methods, for example with biological treatment, should be studied in order to develop a technically feasible and cost-effective solution for wastewater treatment;
Assessment of eco-toxicity should be an essential part of degradation tests over TiO₂-based photocatalysts. The complete mineralization of antibiotics remains a challenge, and incomplete mineralization leads to the formation of various intermediates and by-products, sometimes more toxic than the parent compound. Degradation mechanisms require deeper studies.

7. Conclusions

The current review focuses on the recent progress in the photocatalytic removal of antibiotics over TiO₂ and TiO₂-based materials. Heterogeneous photocatalysis has shown huge potential as an efficient, cost-effective and energy-efficient advanced oxidation process that is able to utilize the green and sustainable source of light – solar irradiation. The global occurrence of antibiotics in the environment, their fate and harmful effects on ecosystems and human health have been covered by the present review. Special emphasis is given to sulfamethoxazole, trimethoprim and ciprofloxacin as the most commonly detected antibiotics in different aquatic systems, raising significant concerns. The huge need for antibiotics removal from contaminated waters has been highlighted. Since conventional wastewater treatment fails to remove such complicated pharmaceuticals pollutants, various antibiotic removal technologies have been briefly discussed. The fundamentals of photocatalytic process over TiO₂ have been presented, along with the modification methods of TiO₂ to overcome its limitations and enhance efficiency. The photocatalytic performance of various TiO₂ and TiO₂-based materials and the effect of different process parameters (amounts of pollutant and photocatalyst, pH, light, reaction media, presence of inorganic ions, natural organic matter and oxidants) have been the main focus of this review. Summarizing recent progress in the photocatalytic removal of the three widespread antibiotics, heterogeneous photocatalysis over TiO₂-based materials has demonstrated its high potential for industry applications. We believe that the present review will contribute to the further progress in the design and development of efficient and cost-effective TiO₂-based photocatalytic systems to meet the outlined challenges of large-scale application.

Author Contributions

Conceptualization, A.K. and T.D.; investigation, A.K.; resources, W.K.; data curation, A.K.; writing—original draft preparation, A.K.; writing—review and editing, T.D. and W.K.; visualization, A.K.; supervision, T.D. and W.K. All authors have read and agreed to the published version of the manuscript.

Funding

This research received no external funding.

Data Availability Statement

The data presented in this study are available in Application of TiO₂-based Photocatalysts to Antibiotics Degradation: Cases of Sulfamethoxazole, Trimethoprim and Ciprofloxacin.

Acknowledgments

Anastasiya Kutuzova would like to gratefully acknowledge scientific support by Matthias Beller and Jennifer Strunk, LIKAT, and the grant from Erasmus+ Program-International Credit Mobility project # 2019-1-IE02-KA107-000650 funded by the European Commission that made it possible to work at the University of Limerick, Ireland.

Conflicts of Interest

The authors declare no conflict of interest.

Abbreviations

AMR	antimicrobial resistance
AOPs	advanced oxidation processes
APIs	active pharmaceutical ingredients
ARB	antibiotic resistant bacteria
ARGs	antibiotic resistant genes
BET	Brunauer-Emmett-Teller
CB	conduction band
CIP	ciprofloxacin
COD	chemical oxygen demand
DOC	dissolved organic carbon
NOM	natural organic matter
NPOC	non-purgeable organic carbon
PZC	point of zero charge
SEM	scanning electron microscopy
SMX	sulfamethoxazole
TEM	transmission electron microscopy
TMP	trimethoprim
TOC	total organic carbon
UV	ultraviolet
VB	valence band
WWTP	wastewater treatment plant
XRD	X-ray diffraction

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

Figures, Schemes and Tables

Figure 1. Global occurrence of pharmaceuticals.

Figure 2. Antibiotics concentrations in surface waters across the world.

Scheme 1. Molecular structure of SMX, TMP and CIP and their 3D models.

Figure 3. Concerns associated with antibiotics presence in the aquatic environment.

Figure 4. Pathways of antibiotics release into the environment.

View Image - Figure 5. Photocatalytic process over TiO2: (1) Photon absorption; (2) Generation and separation of electron-hole pairs (e- and h+); (3) Transport of electrons and holes in the bulk to the photocatalyst surface; (4) Recombination of electrons and holes in the bulk; (5) Surface recombination of electrons and holes; (6) Reduction reaction on TiO2 surface; (7) Oxidation reactions on TiO2 surface.

Figure 5. Photocatalytic process over TiO2: (1) Photon absorption; (2) Generation and separation of electron-hole pairs (e- and h+); (3) Transport of electrons and holes in the bulk to the photocatalyst surface; (4) Recombination of electrons and holes in the bulk; (5) Surface recombination of electrons and holes; (6) Reduction reaction on TiO2 surface; (7) Oxidation reactions on TiO2 surface.

Scheme 2. Proposed degradation pathways of SMX in photocatalytic oxidation over TiO2.

Scheme 3. Proposed degradation pathways of TMP in photocatalytic oxidation over TiO2.

Scheme 4. Proposed degradation pathways of CIP in photocatalytic oxidation over TiO2.

Table 1

Concentrations of SMX, TMP and CIP detected in surface waters worldwide.

Antibiotic	Sulfamethoxazole	Trimethoprim	Ciprofloxacin
Number of countries where antibiotic was detected	47	29	20
Average concentration (µg/L)	0.095	0.037	18.99
Maximum concentration (µg/L)	53.8	13.6	6500

Table 2

Physicochemical properties of SMX, TMP and CIP.

Antibiotic	Sulfamethoxazole	Trimethoprim	Ciprofloxacin
Class	Sulfonamides	Diaminopyrimidines	Fluoroquinolones
Molecular formula	C₁₀H₁₁N₃O₃S	C₁₄H₁₈N₄O₃	C₁₇H₁₈FN₃O₃
Molecular weight (g/mol)	253.28	290.32	331.34
pK_a1, pK_a2	1.7; 5.6 [23]	3.2; 6.7 [23]	5.9; 8.9 [24]
Solubility in water (mg/L) [8]	610	400	30000

Table 3

Advanced technologies applied to the treatment of wastewater contaminated with antibiotics.

Technology	Description	Advantages	Drawbacks
Adsorption	A mass transfer process of accumulation of chemicals from the liquid phase into the solid phase of adsorbent [23,42]. A wide variety of adsorbents is utilized (clays and minerals, metal oxides, polymers, nanocomposites, carbon nanotubes; activated carbon and biochars are among the most commonly used) [23,46].	High removal efficiency [35,45,47]	Complex and expensive production of some adsorbents [46]
		Short treatment period and simple operation [43,46]	High material costs and high regeneration costs of the used adsorbent [23,43,47]
		Low operation costs [23,41]	Nondestructive process, production of secondary waste [35,41,47]
Membrane filtration	Physical process of chemicals separation using synthetic membranes [42].Only nanofiltration and reverse osmosis can be efficiently applied to antibiotics removal [46].	High removal efficiency [42,45]	High costs of operation and maintenance [35,43]
		Fast and simple process with no chemicals involved [50,51]	Membrane fouling [50,51]
			Not suitable for large volumes of wastewater [45,51]
			Nondestructive process [41,45]
Advanced oxidation processes (AOPs)	Processes based on utilization of highly reactive chemical species that are efficient in oxidizing and mineralizing organic compounds [35,41].	High efficiency of pollutants degradation and possibility of complete mineralization in a short period of time [43,46,47]	High operation costs [43,46]. In case of incomplete mineralization post-treatment is required to remove toxicity [43]
		Include environmentally friendly, safe and sustainable processes [52]
		Disinfection properties [43]
AOPs: Ozonation	Ozone molecule O₃ has high oxidation capability (E₀ = 2.07 V), and therefore is able to efficiently oxidize organic pollutants [41,42].	On-site generation of ozone [43]	High operation and energy costs [45,47,49]
		No waste production [49]	Possible formation of harmful by-products [35,46,47]
			Limited pollutant mineralization [43]
AOPs: Fenton process (homogeneous)	Fenton reagent consisting of Fe²⁺ and H₂O₂ produces highly reactive hydroxyl radicals ^•OH that oxidize organic pollutants [49,53].	Fast, effective and safe process [42,49]	Reaction is limited to acidic conditions pH 2.8–3.5 [41,43,53]
		Easy operation [53]	Large volumes of ferrous sludge produced [42,43,53]
		Possible mineralization of pollutants [43]	Complicated recovery of Fe²⁺/Fe³⁺ ions [23]
		Iron is abundant and non-toxic [42,49]
AOPs: Electrochemical processes	Processes of oxidizing organic substances using electric current [53]. Most widely employed anode materials include graphite and TiO₂, as well as Ti-based alloys, Ru and Ir oxides [41].	Easy to operate, safe and highly efficient process [23,43]	High operation costs [23,42,45]
		No chemical reagents required and no generation of secondary wastes [23,41]	Expensive electrodes [41], short electrode life time, electrode fouling [23]
		Suitable for waste waters with high concentrations of pharmaceuticals [45]	It is required that wastewater is highly conductive (otherwise electrolytes should be added) [23,41,42]
			Mass transfer resistance [41]
			Applicable to wastewater with low flow rate [45]
			Water oxidation occurs faster than oxidation of organic pollutants [42]
AOPs: Ultrasonication	Processes employing sound waves for formation, growth and collapse of bubbles in liquid media [23].	No chemical reagents required [43]	High operation costs [43]
AOPs: Ultrasonication			Destruction and corrosion of reactor metallic surface [23]
AOPs: Radiation assisted catalytic reaction	Processes employing electromagnetic radiation (for example, microwaves, x-rays, gamma rays) to form highly reactive species (^•OH, e⁻, ^•H, H₂, H₂O₂, H₃O⁺) that oxidize organic pollutants [23,53].	Fast and energy-efficient process [23]	Toxicity of intermediates in mineralization process should be considered [23]
AOPs: Radiation assisted catalytic reaction		No chemical reagents required [23]	Various factors affect degradation efficiency (dose of radiation, pH, water matrix composition) [51]
AOPs: Catalytic wet peroxide oxidation	Processes of pollutants degradation in aqueous media through catalytic H₂O₂ reduction to OH⁻ and ^•OH under extreme pressures and temperatures [42].	Fast and efficient process [23]	High operation and material costs [23]
AOPs: Catalytic wet peroxide oxidation			Some substances (containing nitro functional groups and halogens) are difficult to degrade [42]
AOPs: Photolysis	Processes employing light (artificial or natural) for the generation of reactive species and subsequent degradation of organic pollutants [41].	Cost-effective process [41]	Lowest degradation efficiency among AOPs [45]
AOPs: Photolysis			Applicable only to photo-sensitive pollutants [41]
AOPs: Photocatalysis	Degradation of organic contaminants using semiconducting materials (photocatalysts) and light (artificial or natural) [42].	Easy to operate, highly efficient and environmentally friendly process [18,42,51]	Fast recombination of photogenerated charge carriers decreases process efficiency [42]
		Performed under ambient temperature and pressure and utilizes atmospheric oxygen as oxidant [23,44]	Limited visible light response [42]
		Complete mineralization of organic pollutants is possible, no waste generation [42,43]	Laboratory scale [43]
		Efficient recovery and reuse of photocatalysts is possible [52,54]	Not applicable to water with high concentrations of organic pollutants [42,51]
			Losses of photocatalyst under long-term operation [51]
			Toxicity of by-products should be considered [42]

Table 4

Performance of various TiO₂ and TiO₂-based materials in photocatalytic degradation of SMX, TMP and CIP.

Catalyst	Characteristics	Process conditions	Performance	Ref.
Sulfamethoxazole
Kaolinite Clay-TiO₂- ZnWO₄ and agrowaste (carica papaya seeds or musa paradisiaca peels) nanocomposite	Particle size: 62–257 nm (SEM)Band gap: 2.56–2.89 eV	Medium: Ultrapure waterAnt. Conc.: 0.05 gL⁻¹Cat. Conc: 0.05 gL⁻¹pH: 6.8Light source: sunlight irradiation (10 AM–5 PM)Cat.: TZPP₅ (based on musa paradisiaca peels and calcined at 500 °C)	Photodegradation: 60% (60 min)Mineralization: 50% (60 min) CODk(1st) = 0.0227 min⁻¹t_1/2 = 30.48 minStability: over 4 cycles	[73]
LDH-TiO₂(Zn₂-Al-CO₃/TiO₂ P25)	-	Ant. Conc.: 0.02 gL⁻¹Cat. Conc.: 0.5 gL⁻¹pH: 10Light source: 300 W UV-A lamp (300–400 nm) with high pressure tungsten filamentCat.: LDH-TiO₂ (10% TiO₂)	Photodegradation: 100% (360 min)Mineralization: 100% (144 h) CODStability: 90.5% after 5 cycles	[90]
TiO₂ P25(Evonik, Germany)	Particle size: 21 nm Composition: 80% anatase, 20% rutileBET surface: 35–65 m²g⁻¹Band gap: 3.2 eV	Medium: deionized waterAnt. Conc.: 0.001 gL⁻¹Cat. Conc.: 0.1 gL⁻¹pH: 6.0Light source: solar simulation chamber with a 1.5 kW Xenon lamp (290–400 nm)Temperature: 20 °C	Photodegradation: 100% (120 min)k(1st) = 0.041 min⁻¹	[104]
F-Pd co-doped TiO₂	Particle size: 5–25 nm (TEM)Band gap: 0.54–3.26 eV	Medium: deionized waterAnt. Conc.: 0.03 gL⁻¹Cat. Conc.: 1 gL⁻¹Light source: direct sunlight irradiationTemperature: 29–31 °CCat.: FPd-TiO₂ (10 mol.% Pd)	Photodegradation: 98.4% (40 min)Mineralization: 93% (40 min) TOC	[116]
Bi₂O₃-TiO₂/PAC (powdered activated carbon)	Crystallite size: 48 nm (XRD)Band gap: 2.58 eV	Medium: deionized waterAnt. Conc.: 0.02 gL⁻¹Cat. Conc.: 0.2 gL⁻¹Light source: solar simulator (300 W Xe arc lamp)	Photodegradation: 100% (30 min)Stability: 90.4% after 5 cycles	[89]
TiO₂ P25	Particle size: 21 nm Composition: 80% anatase, 20% rutileBET surface: ~50 m²g⁻¹Band gap: 3.2 eV	Ant. Conc.: 0.04 gL⁻¹Cat. Conc.: 0.024 gL⁻¹pH: 6.0Light source: UV-C (254 nm), light intensity 15 mWcm⁻²Temperature: 25 °C	Photodegradation: 95.0% (120 min)Mineralization: 66% (120 min) TOC	[84]
TiO₂	BET surface: 52 m²g⁻¹Band gap: 3.08 eV	Ant. Conc.: 100 µgL⁻¹Cat. Conc.: 0.2 gL⁻¹Light source: solar simulator with Xenon lamp (UV irradiance 30 Wm⁻²)	Photodegradation: 100% (30 min)	[134]
Au-TNWs/TNAs (Au nanoparticle-decorated TiO₂ nanowires on TiO₂ nanotube arrays)	Composition: 100% anataseCrystallite size: 21.3–24.7 nm (XRD)	Medium: blank wastewater samples with 0.1% (v.v) formic acidAnt. Conc.: 500 ng/mL, 30mL solutionsLight source: UV-VIS 100 W Xenon lamp, 120 mWcm⁻²Temperature: 32–33 °C	Photodegradation:k(1st) = 1.05 min⁻¹	[135]
Fe/TiO₂	Composition: 100% anataseCrystallite size: 27 nm (XRD)Band gap: 3 eV	Medium: ultrapure waterAnt. Conc.: 234 µgL⁻¹Cat. Conc.: 1 gL⁻¹pH: 6Light source: solar simulator with 100 W Xenon lampCat.: 0.04 mol.% Fe	Photodegradation: 95% (90 min)k(1st) = 0.029 min⁻¹Stability: 5 cycles (55%)	[136]
TiO₂	Composition: 93% anatase, 7% rutileCrystallite size: 8.91 nm anatase; 14.7 nm rutile (XRD)BET surface: 134 m²g⁻¹	Medium: ultrapure waterAnt. Conc.: 0.03 gL⁻¹Cat. Conc.: 0.5 gL⁻¹pH: 5.1Light source: medium pressure mercury vapour lamp (UV-vis: λ > 350 nm; 50 mWcm⁻²)	Photodegradation: 100% (120 min)Mineralization: 40% (180 min) TOC	[137]
TiO₂/CNT (carbon nanotubes), 10% CNT	Composition: 97% anatase, 3% rutileCrystallite size: 14.6 nm anatase; 40.7 nm rutile (XRD)BET surface: 142 m²g⁻¹		Photodegradation: 100% (120 min)Mineralization: 70% (180 min) TOCStability: 3 cycles	[137]
Pt/TiO₂ P25(1.0 wt.% Pt)	Composition: 80% anatase, 20% rutileCrystallite size: 220 nm anatase, 216 nm rutile (XRD)BET surface: ~50 m²g⁻¹Band gap: 3.18 eV	Medium: ultrapure waterAnt. Conc.: 0.001 gL⁻¹Cat. Conc.: 0.05 gL⁻¹Light source: natural sunlight	Photodegradation: 90% (30 min)k(1st) = 0.076 min⁻¹t_1/2 = 9.1 minMineralization: 29±10% (60 min) DOC, C_cat = ~25 mgL⁻¹Ecotoxicity: Lepidium sativum, no phytotoxicity	[123]
Pd/TiO₂ P25(1.0 wt.% Pd)	Composition: 89% anatase, 11% rutileCrystallite size: 206 nm anatase, 180 nm rutile (XRD)BET surface: ~50 m²g⁻¹Band gap: 2.92 eV		Photodegradation: 100% (10 min)k(1st) = 0.521 min⁻¹t_1/2 =1.3 minMineralization: 45 ± 2% (60 min) DOC, C_cat = ~25 mgL⁻¹Ecotoxicity: Lepidium sativum, no phytotoxicity	[123]
Ce_0.8Gd_0.2O_2-_δ/TiO₂	Particle size: 7–20 nm (TEM)Composition:55.78% ceria;27.01% rutile;17.21% anataseCrystallite size: 72.94 nm TiO₂; 19.71 nmCe_0.8Gd_0.2O_2-_δ (XRD)BET surface: 5.11 m²g⁻¹Band gap: 2.84 eV	Ant. Conc.: 0.025 gL⁻¹Cat. Conc.: 0.1 gL⁻¹Light source: 15 W mercury UV lamp	Photodegradation: 97% (120 min)k(2nd) = 0.2959 mg⁻¹min⁻¹Stability: 5 cycles	[92]
Biobased-PET-TiO₂ P25 composite films	TiO₂ P25:Particle size: 20–30 nmComposition: 80% anatase; 20% rutileBET surface: 56 m²g⁻¹PZC: 6.3–6.8	Medium: ultrapure waterAnt. Conc.: 0.001 gL⁻¹ (for each antibiotic, mixture of eight)Cat. Conc.: 0.05 gL⁻¹Light source: solar simulator with xenon lamp (1.5 kW, 500 Wm⁻²)Cat.: PET-10%-TiO₂ (10 wt.% of TiO₂)	Photodegradation: 98% (6 h)k(1st) = 0.015 min⁻¹t_1/2 = 46.2 minStability: 5 cycles	[103]
Zn-TiO₂/pBC (reed straw biochar)	Crystallite size: 9.4 nm (XRD)BET surface: 169.12 m²g⁻¹	Medium: ultrapure waterAnt. Conc.: 0.01 gL⁻¹Cat. Conc.: 1.25 gL⁻¹pH: 4Light source: 50 W Xenon lamp with 420 nm cutoff filter (visible)Temperature: 25 °CCat.: Zn10-TiO₂/pBC (10 wt.% Zn)	Photodegradation: 80.81% (3 h)k(1st) = 0.0085 min⁻¹t_1/2 = 81 minMineralization: 56.13% (3 h) CODStability: 5 cycles (77.41%)	[79]
Magnetic carbon nanotube-TiO₂ P25 (MCNT-TiO₂) composites	BET surface: 151 m²g⁻¹	Medium: ultrapure waterAnt. Conc.: 150 μgL⁻¹Cat. Conc.: 0.1 gL⁻¹pH: 7.0±0.2Light source: solar simulator, 1000 Wm⁻²Temperature: 26±3 °CCat.: MCNT-TiO₂ (1:5 mass ratio)	Photodegradation: 92% (30 min)k(1st) = 0.05 min⁻¹Stability: 5 cycles	[101]
TiO₂ P25	Composition: 80% anatase, 20% rutile	Medium: secondary urban wastewaterAnt. Conc.: 100 µgL⁻¹Cat. Conc.: 1 gL⁻¹Light source: four 9 W UVA-LEDs, 381 nm	Photodegradation: 100% SMX (30 min)k(1st) = 0.2126 min⁻¹Disinfection: total heterotrophs, E. coli and enterococci	[111]
Porous titanium-titanium dioxide (PTT) substrates	Composition: anataseBand gap: 3.0 eVIsoelectric point: 6.0	Medium: ultrapure waterAnt. Conc.: 300 mL of 2 µgL⁻¹ (eighteen pharmaceuticals)pH: ~5Light source: UV-LEDs, 1.7×10⁻³ W, 365 nm	Photodegradation: 72.74% (300 min)k(1st) = 0.00435 min⁻¹	[138]
Graphene-based TiO₂ P25 (TiO₂-rGO)	BET surface: 44.761 m²g⁻¹	Medium: membrane bioreactor-treated urban wastewaterAnt. Conc.: 100 μgL⁻¹ (three compounds)Cat. Conc.: 0.1 gL⁻¹pH: 5.2–6.2Light source: solar simulator with 1 kW Xenon lamp, 63 Wm⁻²Temperature: 25 ± 1 °CCat.: TiO₂-rGO-PH	Photodegradation: 50 ± 3% (60 min)Disinfection: E. coli complete inactivation (180 min)	[139]
Ag⁰(NP)/TiO₂ thin films	Particle size: 10–15 nm (TEM)BET surface: 12.02 m²g⁻¹	Medium: purified waterAnt. Conc.: 50 mL of 1.0 mgL⁻¹pH: 6.0Light source: 9 W UVA 360 nmTemperature: 25 ± 1 °CCat.: Ag⁰(NP)/TiO₂ (B)	Photodegradation: 57% (120 min)k(1st) = 0.0067 min⁻¹Mineralization: 30.2% (120 min) NPOCStability: 6 cycles	[72]
TiO₂ supported on reed straw biochar TiO2/pBC	Composition: 100% anataseCrystallite size: 10.1 nmBET surface: 102.16 m²g⁻¹	Medium: ultrapure waterAnt. Conc.: 0.01 gL⁻¹ Cat. Conc.: 1.25 gL⁻¹pH: 4.0Light source: 50 W xenon lamp with visible light filterTemperature: 25 °CCat.: TiO₂/pBC (300)	Photodegradation: 91.27% (3 h)k(1st) = 0.0130 min⁻¹t_1/2 = 53.32 minMineralization: 57.44% (3 h) CODStability: 5 cycles	[80]
TiO₂ P25-WO₃	Composition: 77% anatase, 23% rutileCrystallite size: 25 nm anatase, 78 nm rutile (XRD)BET surface: 32 m²g⁻¹Band gap: 3.0 eV	Medium: ultrapure waterAnt. Conc.: 350 µgL⁻¹Cat. Conc.: 0.250 gL⁻¹Light source: solarsimulator with 100 W xenon lamp, 420 nm cut-off filterTemperature: 25 °CCat.: 4% W-P25(700)	Photodegradation:100% (60 min) solar irradiation25% (120 min) visible lightk(1st) = 0.133 min⁻¹Mineralization: 28% (6 h), 20 mgL⁻¹ SMX with 1 gL⁻¹ 4% W-P25(700)Ecotoxicity: Escherichia coli, Enterococcus faecalis	[74]
TiO₂ P25-Fe immobilized on optical fibers	Grain size: 7.41 nm (XRD)Composition: 54% anatase, 46% rutileBand gap: 2.40 eV	Medium: deionized waterAnt. Conc.: 0.005 gL⁻¹Cat. Dosage: 30 pieces of 10 cm photocatalyst-coated SOFspH: 6.0Light source: visible light source (halogen lamp 150 W)Temperature: 23 °C	Photodegradation: 35% (6 h)k(1st) = 0.082 min⁻¹	[120]
TiO₂ P25-reduced graphene oxide (TiO2-rGO) immobilized on optical fibers	Grain size: 6.52 nm (XRD)Composition: 69% anatase, 31% rutileBand gap: 2.85 eV		Photodegradation: 35% (6 h)k(1st) = 0.079 min⁻¹	[120]
CoFe₂O₄/TiO₂ (+TiO₂ P25)	-	Medium: ultrapure waterAnt. Conc.: 100 µMCat. Conc.: 0.5 gL⁻¹Light source: photo-chemical reactor with twelve mercury lamps (350 nm)	Photodegradation: 100% (5 h)Mineralization: 50% (5 h) TOCEcotoxicity: green alga Chlorella vulgaris, brine shrimp Artemia salina	[86]
Biochar/TiO₂	Composition: anatase	Medium: pure waterAnt. Conc.: 0.01 gL⁻¹Cat. Conc.: 5 gL⁻¹pH: 4.0Light source: UV-C, 15 W, 254 nmTemperature: 293 K	Photodegradation: 75% (3 h)Mineralization: 65% (3 h) CODEcotoxicity: Daphnia magna, E.Coli	[81]
TiO₂ (TiEt-450)	Hydrodynamic particle size: 3.0 µmComposition: 100% anataseBET surface: 43 m²g⁻¹Isoelectric point: 7.1Band gap: 3.22 eV	Medium: deionized waterAnt. Conc.: 100 µgL⁻¹ of each contaminant (mixture of five, 500 µgL⁻¹ in total)Cat. Conc.: 0.5 gL⁻¹pH: naturalLight source: solar radiation pilot plantTemperature: ambient	Photodegradation: 100% (30 min)k(1st) = 0.095 min⁻¹	[102]
TiO₂ P25	Particle size: 30 nmComposition: 70% anatase, 30% rutileBET surface: 50 ± 15 m²g⁻¹Band gap: 3.15 eV	Medium: high purity waterAnt. Conc.: 0.02 gL⁻¹Cat. Conc.: 0.5 gL⁻¹pH: naturalLight source: UV-LEDs: 365 ± 10 nm, 300 ± 5 nm, and 260 ± 10 nmTemperature: 25 °C	Photodegradation:58% (3 h) UV-A85% (3 h) UV-B100% (3 h) UV-CMineralization:35.1% (3 h) UV-A40.5% (3 h) UV-B59.9% (3 h) UV-C	[119]
TiO₂ immobilized on porous supports: quartz fiber filters (QFT) or porous titanium sheets (PTT)	Composition: anataseIsoelectric point: 4 (QFT), 6 (PTT)Band gap: 3.18 eV (QFT), 3.0 eV (PTT)	Medium: ultrapure waterAnt. Conc.: 300 mL of 2.0 µgL⁻¹ pH: 4.5–5Light source: UV-LED 365 nm, 1.7 mW, 0.13 mW/cm²Temperature: 24 ± 2 °CCat.: TiO₂ PTT	Photodegradation:k(1st) = 0.0069 min⁻¹Stability: 2 cycles	[140]
TiO₂ P25 immobilized on PVDF-coated steel mesh (SM-TiO₂)	-	Ant. Conc.: 10 µMCat. Conc.: 0.02 gL⁻¹pH: 6.8–7.0Light source: six blacklight blue lamps, 4 W, 350–400 nm	Photodegradation: 100% (120 min)k(1st) = 0.0568 min⁻¹Mineralization: 15% (180 min) TOCStability: 20 cycles	[141]
Floating TiO₂-expanded perlite (EP-TiO₂-773)	-	Medium: deionized waterAnt. Conc.: 0.1 gL⁻¹ Cat. Conc.: 3.33 gL⁻¹pH: 10Light source: photoreactor with six (8 W each) lamps, 316–400 nm	Photodegradation:k(0st) = 4.57×10⁻⁶ min⁻¹	[91]
Nano-sized TiO₂ supported on single wall carbon nanotubes (SWCNTs/TiO₂)	Composition: anatase	Medium: ultrapure waterAnt. Conc.: 200–500 µgL⁻¹Cat. Conc.: 0.1 gL⁻¹Light source: low pressure mercury lamp, 17 W, 254 nm, 0.1 Wcm⁻²	Photodegradation:k(1st) = 0.42 min⁻¹Stability: 5 cycles	[142]
Cu–TiO₂ P25	Particle size: 21 ± 4 nm (TEM)BET surface: 52–59 m²g⁻¹	Medium: oxygen-saturated waterAnt. Conc.: 0.004 gL⁻¹Cat. Conc.: 1 gL⁻¹pH: 5.2Light source: eight 8 W lamps, 77 mWcm⁻², 460 nm visible lightCat.: Cu–TiO₂ P25 (0.045 wt.% Cu)	Photodegradation: 100% (90 min)k(1st) = 0.0506 min⁻¹Stability: 4 cycles	[82]
TiO₂ P25	Particle size: 21 nm	Medium: distilled waterAnt. Conc.: 0.05 gL⁻¹ Cat. Conc.: 0.1 gL⁻¹Light source: UVA 15 W, 360 nm, 1.15 mWcm⁻²Temperature: 25 ± 2 °C	Photodegradation: 92.3% (510 min)t_1/2 = 132 ± 5 minMineralization: 34% (510 min) CODEcotoxicity: Vibrio fischeri	[105]
TiO₂ P25	-	Medium: purified waterAnt. Conc.: 400 ppb each (mixture of two antibiotics)Cat. Conc.: 0.05 gL⁻¹pH: 5.6Light source: thirty two 1W UVA/LED chips, peak emission at 365 nmTemperature: 25 °C	Photodegradation: 91% (20 min)Disinfection: Escherichia coliEcotoxicity: Vibrio fischeri	[75]
Trimethoprim
TiO₂ P25	-	Medium: purified waterAnt. Conc.: 400 ppb each (mixture of two antibiotics)Cat. Conc.: 0.05 gL⁻¹pH: 5.6Light source: thirty two 1W UVA/LED chips, peak emission at 365 nmTemperature: 25 °C	Photodegradation: 96% (20 min)Disinfection: Escherichia coliEcotoxicity: Vibrio fischeri	[75]
Biobased-PET-TiO₂ P25 composite films	TiO₂ P25:Particle size: 20–30 nmComposition: 80% anatase; 20% rutileBET surface: 56 m²g⁻¹ZPC: 6.3–6.8	Medium: ultrapure waterAnt. Conc.: 0.001 gL⁻¹ (for each antibiotic, mixture of eight)Cat. Conc.: 0.05 gL⁻¹Light source: solar simulator with xenon lamp (1.5 kW, 500 Wm⁻²)Cat.: PET-10%-TiO₂ (10 wt.% of TiO₂)	Photodegradation: 90% (6 h)k(1st) = 0.007 min⁻¹t_1/2 = 99.0 minStability: 5 cycles	[103]
Keratin-TiO₂ nanocomposite	Composition: 85% anatase; 15% rutile	Ant. Conc.: 0.172 mM Cat. Conc.: 1 gL⁻¹Light source: Xenon lamp, simulated solar light, 28 kLuxTemperature: 25 °CCat.: K-TiO₂ 10%	Photodegradation: 100% (4 h)Stability: 4 cycles	[143]
S-TiO₂	Particle size: 12–22 nm (TEM)Composition: anatase; brookite	Ant. Conc.: 0.01 gL⁻¹Cat. Conc.: 0.5 gL⁻¹pH: 7.0Light source: UV-Vis 400 W metal halide lamp	Photodegradation: 98.2% (4 h)Stability: 5 cycles (for immobilized catalyst)	[126]
TiO₂ P25	Composition: 80% anatase, 20% rutile	Medium: secondary urban wastewaterAnt. Conc.: 100 µgL⁻¹Cat. Conc.: 1 gL⁻¹Light source: 4 UVA-LEDs, 9 W, 381 nm	Photodegradation: 100% (60 min)k(1st) = 0.1171 min⁻¹Disinfection: total heterotrophs, E. coli and enterococci	[111]
Nitrogen-doped TiO₂ immobilized on polystyrene spheres (N-TiO₂)	Composition: anataseBand gap: 2.5 eV	Medium: distilled waterAnt. Conc.: 200 µgL⁻¹Cat. Conc.: 160.74 gL⁻¹pH: 6.13–6.38 (not adjusted)Light source: natural sunlightTemperature: 25.7–36.1 °C (not adjusted)	Photodegradation: 100% (150 min)k(1st) = 0.0167 min⁻¹t_1/2 = 42 min	[117]
Porous titanium-titanium dioxide (PTT) substrates	Composition: anataseBand gap: 3.0 eVIsoelectric point: 6.0	Medium: ultrapure waterAnt. Conc.: 300 mL of 2 µgL⁻¹ (eighteen pharmaceuticals)pH: ~5Light source: UV-LEDs, 1.7×10⁻³ W, 365 nm	Photodegradation: 31.78% (300 min)k(1st) = 0.00132 min⁻¹	[138]
Nano-sized TiO₂ supported on single wall carbon nanotubes (SWCNTs/TiO₂)	Composition: anatase	Medium: ultrapure waterAnt. Conc.: 200–500 µgL⁻¹Cat. Conc.: 0.1 gL⁻¹Light source: low pressure mercury lamp, 17 W, 254 nm, 0.1 Wcm⁻²	Photodegradation:k(1st) = 0.075 min⁻¹Stability: 5 cycles	[142]
Immobilized TiO₂ P25 on poly(vinylidene fluoride) (PVDF) dual layer hollow fibre membrane	Composition: anatase, rutileBET surface: 50 m²g⁻¹	Medium: ground water and secondary wastewater effluentAnt. Conc.: 200–400 μgL⁻¹Cat. Conc.: 0.057 gL⁻¹Light source: low-pressure mercuryUV lamp, 40 W, 254 nm	Photodegradation:k(1st) = 0.045 min⁻¹ (ground water)k(1st) = 0.095 min⁻¹ (secondary WW effluent)Stability: 5 cycles	[144]
TiO₂ film deposited on stainless steel mesh (nanoTiO₂-SS)	-	Medium: filtered ground waterAnt. Conc.: 200–400 µgL⁻¹Cat. Conc.: 0.1 gL⁻¹pH: 7.85Light source: 40 W Hg low pressure UV lamp, 254 nm, 50 mWcm⁻²Temperature: 24.8 °C	Photodegradation:k(1st) = 0.107 min⁻¹Ecotoxicity: AMES Fluctuation, Fish Embryo, Green alga Selenastrum capricornutum, Daphnia magna. Vibrio fischeriStability: 10 cycles	[131]
Ciprofloxacin
Black Ti³⁺/N-TiO₂ P25 (b-N-TiO₂)	Particle size: < 100 nm (FE-SEM)Composition: anataseBET surface: 100 m²g⁻¹PZC: 7.9Band gap: 2.0 eV	Medium: ultrapure waterAnt. Conc.: 0.5 mgL⁻¹ Cat. Conc.: 0.43 gL⁻¹pH: 6.7Light source: 5 W visible LED lamp, 550 nm	Photodegradation: 100% (70 min)k(1st) = 0.0778 min⁻¹Mineralization: 82% (140 min) TOCEcotoxicity: Daphnia magnaStability: 5 cycles	[24]
3D tripyramid TiO₂ architectures	Particle size: 10 nm (TEM)Composition: anataseBET surface: 84 m²g⁻¹Band gap: 3.2 eV	Ant. Conc.: 32.6 μMCat. Conc.: 0.1 gL⁻¹Light source: UV-vis light	Photodegradation: 90% (60 min)k(1st) = 0.0403 min⁻¹Stability: 5 cycles	[97]
TiO₂ nanorod/g-C₃N₄ nanosheet (TiO₂ nanorod-CN)	Composition: anataseZPC: 6.3Band gap: 2.95 eV	Ant. Conc.: 15 µmolL⁻¹Cat. Conc.: 0.2 gL⁻¹pH: 6.3Light source: simulated sunlight irradiation, 500 W Xenon lampCat.: 30 wt.% g-C₃N₄	Photodegradation: 93.4% (60 min)k(1st) = 0.0389 min⁻¹	[94]
TiO₂-Ag NPs	Particle size: 80–100 nm (SEM)Composition: >90% anatase, rutile (XRD)Grain size: 51.56 nm anatase, 21.87 nm rutile (XRD)Band gap: 3.26–3.30 eV	Ant. Conc.: 1.0 mMCat. Conc.: 0.001 gL⁻¹pH: 7.0Light source: 120 W UV Hg lamp; natural sunlight	Photodegradation:85.21% UV lightk = 1.53 mMs⁻¹75.58% visible lightk = 1.20 mMs⁻¹	[109]
TiO₂ P25 (Acros) and γ-Fe₂O₃ co-doped graphene oxide (GO) nanosheets(TiO₂/γ-Fe₂O₃/GO)	Composition: anatase, rutileBand gap: 2.43 eV	Ant. Conc.: 0.01 gL⁻¹Cat. Conc.: 0.4 gL⁻¹pH: 6.6Light source: 300 W xenon lamp, 420 nm cutoff filterCat.: 0.03TiO₂/γ-Fe₂O₃/GO	Photodegradation: 99% (140 min)k(1st) = 0.019 min⁻¹Stability: 4 cycles	[114]
TiO₂/kaolinite	Crystallite size: 16.631 nm (XRD)Composition: anataseBET surface: 60.21 m²g⁻¹Band gap: 3.18 eV	Medium: deionized waterAnt. Conc.: 20 ppmCat. Conc.: 2 gL⁻¹Light source: high-pressure mercury lampCat.: TK-6.0 (58.54% TiO₂)	Photodegradation: 93.14% (40 min)k(1st) = 0.04549 min⁻¹Stability: 4 cycles	[145]
Fe–N–TiO₂	Particle size: ~25 nm (HR-TEM)Composition: 61.1% anatase, 21.4% rutile, 17.5% brookiteCrystallite size: 28.1 nm anatase, 28.7 nm rutile, 34.3 nm brookite (XRD)BET surface: 90.2 m²g⁻¹Band gap: 2.7 eV	Ant. Conc.: 0.02 gL⁻¹Cat. Conc.: 1 gL⁻¹Light source: 12 W LED daylight lamps (visible)Temperature: ambientCat.: 2.5%N–1.5%Fe	Photodegradation: 67.72% (6 h)k(1st) = 0.00552 min⁻¹Mineralization: 49.07% (6 h) TOC	[146]
N-TiO₂	Particle size: 15 ± 0.56 nm (TEM)Composition: anatase, rutileCrystallite size: 12–18 nm (XRD)BET surface: 24.59 m²g⁻¹PZC: 5.76Band gap: 2.84 eV	Medium: distilled waterAnt. Conc.: 0.01 gL⁻¹ Cat. Conc.: 0.5 gL⁻¹pH: 7.0Light source: three 14 W blue LEDs, 457 nmCat.: Ti:N molar ratio 1:1	Photodegradation: 55% (180 min)Mineralization: 24% (180 min) TOC	[76]
B-TiO₂ P25	Particle size: 61.8 nm (nanoparticle size analyzer)25–34 nm (TEM)Composition: anatase, rutileCrystallite size: 19.82 nm anatase, 26.24 nm rutile (XRD)BET surface: 30.1 m²g⁻¹Band gap: 2.89 eV	Ant. Conc.: 0.01 gL⁻¹Cat. Conc.: 1 gL⁻¹pH: 7.0Light source: natural sunlightTemperature: 33 °CCat.: 1B-TiO₂ (1 at.% B)	Photodegradation: 93.16% (180 min)k(1st) = 0.0249 min⁻¹Mineralization: 93% (180 min) CODDisinfection: 95–99.99% efficiency against E.coliStability: 3 cycles	[118]
Ce-TiO₂ P25	Particle size:89.5 nm (nanoparticle size analyzer), 19–39 nm (TEM)Composition: anatase, rutileCrystallite size: 16.95 nm anatase, 23.21 nm rutile (XRD)BET surface: 41.5 m²g⁻¹Band gap: 2.50 eV	Ant. Conc.: 0.01gL⁻¹Cat. Conc.: 0.5 gL⁻¹pH: 7.0Light source: natural sunlightTemperature: 33 °CCat.: 1Ce-TiO₂ (1 at.% Ce)	Photodegradation: 93.22% (180 min)k(1st) = 0.0266 min⁻¹Mineralization: 92% (180 min) CODDisinfection: 95–99.99% efficiency against E.coliStability: 3 cycles	[118]
Graphene/TiO₂/g–C₃N₄ (GTOCN)	Particle size: 227.18 nmBET surface: 26.41 m²g⁻¹PZC: 4.16	Medium: ultrapure waterAnt. Conc.: 0.003 gL⁻¹ Cat. Conc.: 0.6 gL⁻¹Light source: 300 W Xenon lamp (>400 nm), 300 mWcm⁻²Cat.: GTOCN3 (40 mg Ti₃C₂)	Photodegradation: 61.7% (60 min)k(1st) = 0.01675 min⁻¹Mineralization: 41.8% (60 min) TOCStability: 3 cycles	[77]
TiO₂ nanotube arrays (TiO₂ NTAs) with Ag₃PO₄ nanoparticles	Composition: anataseBET surface: 4.7 m²g⁻¹Band gap: < 3.25 eV	Ant. Conc.: 0.01 gL⁻¹Cat. Conc.: 1 gL⁻¹Light source: 300 W Xenon lamp, visible light, 200 mWcm⁻²Cat.: 0.6Ag₃PO₄/TiO₂ (Ag₃PO₄:TiO₂ mass ratio 0.6:1)	Photodegradation: 85.3% (60 min)k(1st) = 0.02499 min⁻¹Disinfection: E. coli 100% (120 min)Stability: 3 cycles	[99]
Nitrogen and carbon co-doped TiO₂ nano-catalysts (NCD-TiO₂)	Particle size: 9 nm (HR-TEM)Composition: anataseCrystallite size: 8.8 nm (XRD)BET surface: 116.5 m²g⁻¹Band gap: 2.94 eV	Medium: ultrapure waterAnt. Conc.: 75 µMCat. Conc.: 1 gL⁻¹pH: 5.7Light source: four 8 W fluorescent lamps with UV light filter, 11.58 mWcm⁻²Cat.: NCD200-430 (N:Ti molar ratio 2:1 calcined at 430 °C)	Photodegradation: 68.7% (120 min)k(1st) = 0.0093 min⁻¹	[98]
TiO₂/Graphene oxide	Composition: anataseCrystallite size: 12.5 nm (XRD)BET surface: 91.25 m²g⁻¹Band gap: 2.47 eV	Medium: distilled waterAnt. Conc.: 0.005 gL⁻¹ Cat. Conc.: 0.5 gL⁻¹Light source: visible lightCat.: TiO₂/GO (8%)	Photodegradation: 96.73% (60 min)Stability: 6 cycles	[108]
TiOF₂/TiO₂ nanosheets	Composition: anataseCrystallite size: 26.2 nm (XRD)BET surface: 119 m²g⁻¹Band gap: 3.285 eV	Ant. Conc.: 0.02 gL⁻¹Cat. Conc.: 1 gL⁻¹Light source: 300 W Xenon lamp (UV+visible)Cat.: S-160 (hydrothermal treatment at 160 °C)	Photodegradation: 95.3% (90 min)k(1st) = 0.034 min⁻¹	[147]
Cu-TiO₂	Particle size:10 nm Cu (TEM)200–400 nm Cu (SEM)Composition: anataseBET surface: 170.15 m²g⁻¹Band gap: 3.0 eV	Medium: deionized waterAnt. Conc.: 0.08 gL⁻¹ Cat. Conc.: 0.25 gL⁻¹Light source: 500 W Xenon lamp (simulated sunlight)Cat.: 0.1-Cu-TiO₂ (weight ratio of Cu)	Photodegradation: 97% (4 h)k(1st) = 0.63 min⁻¹Stability: 6 cycles	[100]
Ag-SrTiO₃/TiO₂(SrTiO₃ nanocubes supported on TiO₂ nanosheets with Ag nanoparticles deposited on both)	Particle size: 30 nm Ag (SEM)Composition: anataseBET surface: 28.3 m²g⁻¹	Ant. Conc.: 0.02 gL⁻¹Cat. Conc.: 0.4 gL⁻¹Light source: 300 W Xenon lamp (simulated sunlight)	Photodegradation: 97.6% (60 min)k(1st) = 0.07 min⁻¹	[115]
TiO₂-modified Bi₂MoO₆ nanocrystals TiO₂/Bi₂MoO₆	Particle size: 20 nm Bi₂MoO₆, 12.4 nm rutile, 6.1 nm diameter and 19.1 nm length rod-like anatase (TEM)Composition: Bi₂MoO₆, rutile, anataseBET surface: 17.7 m²g⁻¹Band gap: 2.60–2.68 eV	Ant. Conc.: 0.01 gL⁻¹Cat. Conc.: 0.6 gL⁻¹Light source: 350 W Xenon lamp, λ ≥ 420 nmCat.: TiO₂(0.41 wt%)/Bi₂MoO₆	Photodegradation: 88% (150 min)k(1st) = 0.008 min⁻¹	[148]
N–TiO₂	Particle size: 180 nm length and 50 nm width (FIB/FESEM)Composition: anataseBET surface: 42.70 m²g⁻¹Band gap: 3.17 eV	Medium: deionized waterAnt. Conc.: 20 ppmCat. Conc.: 1 gL⁻¹pH: 5.5Light source: three 20 W UV-A lamps 365 nm, 0.493 mWcm⁻²Cat.: 12.5% N	Photodegradation: 94.29% (420 min)Mineralization: 66.31% (420 min) TOC	[149]
Graphitized mesoporous carbon (GMC)-TiO₂	Particle size: 15 nm (TEM)Composition: anataseCrystallite size: 12 nm (XRD)BET surface: 286 m²g⁻¹	Medium: deionized waterAnt. Conc.: 0.015 gL⁻¹ Cat. Conc.: 0.35 gL⁻¹Light source: 14 W UV lamp, 254 nm	Photodegradation: 100% (90 min)k(1st) = 0.102 min⁻¹Mineralization: 100% (120 min) TOCEcotoxicity: Vibrio fischeri	[129]
Mesoporous nano-TiO₂	Composition: anataseCrystallite size: 13.5 nm (XRD)BET surface: 191.4 m²g⁻¹Band gap: 2.95 eV	Medium: deionized waterAnt. Conc.: 0.16 gL⁻¹ Cat. Conc.: 0.25 gL⁻¹Light source: 500 W Xe lamp, 200–1000 nmCat.: TiO₂ (hydr)–hydrothermal post-treatment	Photodegradation: 96.05% (6 h)k(1st) = 0.45 min⁻¹Mineralization: 76.66% (6 h) TOCEcotoxicity: Staphylococcus aureus	[130]
P-doped TiO₂	Particle size: 12 nm (TEM)Composition: anataseBET surface: 88.54 cm²g⁻¹Band gap: 3.02 eV	Ant. Conc.: 5 ppmCat. Conc.: 0.5 gL⁻¹Light source: visible lightCat.: PT-50 (50 mg of NaH₂PO₂)	Photodegradation: 100% (60 min)k(1st) = 0.065 min⁻¹Mineralization: 72.7% (60 min) TOC	[150]
g-C₃N₄/TiO₂/kaolinite	Composition: anataseCrystallite size: 14.21 nm TiO₂ (XRD)BET surface: 51.596 m²g⁻¹Band gap: 2.72 eV	Medium: deionized waterAnt. Conc.: 10 ppmCat. Conc.: 2 gL⁻¹Light source: Xenon lamp with 400 nm cut-off filter, 90 mWcm⁻²	Photodegradation: 92% (240 min)k(1st) = 0.00813 min⁻¹Disinfection: Stapheylococcus aureusStability: 4 cycles	[151]
N-TiO₂ immobilized on glass spheres	Composition: anataseCrystallite size: 5.69 nmBET surface: 140.47 m²g⁻¹Band gap: 2.84 eV	Ant. Conc.: 0.02 gL⁻¹Cat. Conc.: 3 gL⁻¹Light source: 500 W Xe lamp, λ > 420 nmCat.: N/Ti weight ratio 0.34%	Photodegradation: 93.5% (90 min)k(1st) = 0.02859 min⁻¹Ecotoxicity: Toxicity Estimation Software Tool (T.E.S.T.)Stability: 5 cycles	[133]
TiO₂ P25	Particle size: 30 nmComposition: 80% anatase, 20% rutileBET surface: 50 ± 15 m²g⁻¹	Medium: deionized waterAnt. Conc.: 0.02 gL⁻¹ Cat. Conc.: 0.1 gL⁻¹pH: 6.00 ± 0.10Light source: twelve 3 W UVA/LED lamps, 365 nm, 10mWcm⁻²	Photodegradation:k(1st) = 0.2217 ± 0.0179 min⁻¹Mineralization:76% (240 min) TOC	[96]
TiO₂ P25/Fe⁰	-	Medium: double distilled waterAnt. Conc.: 0.03 gL⁻¹Cat. Conc.: 1 gL⁻¹pH: 3.0Light source: 10 W UV lamp, 254 nm,2.0 Wm⁻²	Photodegradation: 94.6% (60 min)Stability: 5 cycles	[107]
Fe₃O₄/SiO₂/TiO₂	Particle size: 293 ± 81 nm (SEM)Composition: anatase, rutileBET surface: 19 m²g⁻¹PZC: 3–5Band gap: 2.8 eV	Medium: Millipore waterAnt. Conc.: 0.005 gL⁻¹ Cat. Conc.: 1 gL⁻¹pH: 5.5Light source: six 8 W blacklight blue lamps, 365 nm, 1.6–1.7 mWcm⁻²	Photodegradation: 95% (90 min)k(1st) = 0.032 min⁻¹Stability: 5 cycles	[152]
TiO₂ P25 immobilized on glass plates	Composition: 80% anatase, 20% rutileBET surface: 48.3 m²g⁻¹	Medium: deionized waterAnt. Conc.: 60 µmolL⁻¹ Cat. Conc.: 0.75 gL⁻¹ (3.8 × 48 cm rectangle glass plate with 20.5 gm⁻² TiO₂)pH: 9Light source: 15 W UV-C lamp, 254 nm	Photodegradation: 100% (120 min)k(1st) = 4.0×10⁻⁴ s⁻¹	[112]
Mono- (Au, Ag and Cu) and bi-metallic Au–Ag and Au–Cu nanoparticles deposited on TiO₂	Particle size: 2.5–4 nmComposition: anatase, brookiteBET surface: 52–64 m²g⁻¹PZC: 6.3Band gap: 3.1–3.19 eV	Medium: tridistilled waterAnt. Conc.: 0.03 gL⁻¹Cat. Conc.: 0.5 gL⁻¹Light source: 15 W UV-C low pressure Hg lamp, 254 nm, 44 Wm⁻²; solar simulator with a 1500 W Xenon lamp, 500 Wm⁻²Temperature: 25 °C (<35 °C under simulated sunlight)Cat.: 1.5 wt.% for Au and Ag, 1.0 wt.% for Cu (mono-metallic);1.0 wt.% for Au and 0.5 wt.% for Ag and Cu (bi-metallic)	Photodegradation: 100% (90 min) UV-Ck(1st) = 0.06 min⁻¹ 1.5% Au/TiO₂k(1st) = 0.117 min⁻¹ 1.5% Ag/TiO₂k(1st) = 0.072 min⁻¹ 1.0% Cu/TiO₂k(1st) = 0.053 min⁻¹ Au-Ag/TiO₂k(1st) = 0.099 min⁻¹ Au-Cu/TiO₂100% (240 min) simulated sunlightk(1st) = 0.042 min⁻¹ 1.5% Au/TiO₂k(1st) = 0.04 min⁻¹ 1.5% Ag/TiO₂k(1st) = 0.023 min⁻¹ 1.0% Cu/TiO₂k(1st) = 0.021 min⁻¹ Au-Ag/TiO₂k(1st) = 0.022 min⁻¹ Au-Cu/TiO₂Mineralization: 100% (180 min) TOC (UV-C)100% (360 min, bi-metallic) TOC (simulated sunlight)Ecotoxicity: V. fischeri, E. coli Stability: 3 cycles	[121]
TiO₂ P25	Particle size: 30 nmBET surface: 35–65 m²g⁻¹	Medium: ultra-pure waterAnt. Conc.: 300 µgL⁻¹Cat. Conc.: 1 gL⁻¹Light source: UV-A, 365 nm, 1.6–1.7 mWcm⁻²	Photodegradation: 100% (6 min)Ecotoxicity: Vibrio fischeri	[132]
TiO₂/MMT (montmorillonite)	Particle size: 40–60 nmComposition: anataseBET surface: 53.058 m²g⁻¹PZC: 8.4	Ant. Conc.: 0.02 gL⁻¹Cat. Conc.: 0.1 gL⁻¹pH: 5.0Light source: 16 W UV-C lamp	Photodegradation: 61.7% (120 min)k(1st) = 0.0069 min⁻¹t_1/2 = 100.46 minStability: 5 cycles	[78]
TiO₂ P25	BET surface: 56 m²g⁻¹	Ant. Conc.: 12.5 µMCat. Conc.: 1 gL⁻¹Light source: six 8 W mercurial fluorescent tubes, UVA 365 nm, 18–19 Wm⁻²	Photodegradation: 100% (30 min)	[153]

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The extensive application of antibiotics in human and veterinary medicine has led to their widespread occurrence in a natural aquatic environment. Global health crisis is associated with the fast development of antimicrobial resistance, as more and more infectious diseases cannot be treated more than once. Sulfamethoxazole, trimethoprim and ciprofloxacin are the most commonly detected antibiotics in water systems worldwide. The persistent and toxic nature of these antibiotics makes their elimination by conventional treatment methods at wastewater treatment plants almost impossible. The application of advanced oxidation processes and heterogeneous photocatalysis over TiO₂-based materials is a promising solution. This highly efficient technology has the potential to be sustainable, cost-efficient and energy-efficient. A comprehensive review on the application of various TiO₂-based photocatalysts for the degradation of sulfamethoxazole, trimethoprim and ciprofloxacin is focused on highlighting their photocatalytic performance under various reaction conditions (different amounts of pollutant and photocatalyst, pH, light source, reaction media, presence of inorganic ions, natural organic matter, oxidants). Mineralization efficiency and ecotoxicity of final products have been also considered. Further research needs have been presented based on the literature findings. Among them, design and development of highly efficient under sunlight, stable, recyclable and cost-effective TiO₂-based materials; usage of real wastewaters for photocatalytic tests; and compulsory assessment of products ecotoxicity are the most important research tasks in order to meet requirements for industrial application.

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Application of TiO₂-Based Photocatalysts to Antibiotics Degradation: Cases of Sulfamethoxazole, Trimethoprim and Ciprofloxacin

Author

Kutuzova, Anastasiya¹

; Dontsova, Tetiana²; Kwapinski, Witold³

¹ Leibniz Institute for Catalysis (LIKAT), Albert-Einstein-Straße 29a, 18059 Rostock, Germany; Department of Technology of Inorganic Substances, Water Treatment and General Chemical Technology, Igor Sikorsky Kyiv Polytechnic Institute, Prosp. Peremohy 37, 03056 Kyiv, Ukraine; [email protected]
² Department of Technology of Inorganic Substances, Water Treatment and General Chemical Technology, Igor Sikorsky Kyiv Polytechnic Institute, Prosp. Peremohy 37, 03056 Kyiv, Ukraine; [email protected]
³ Department of Chemical Sciences, Faculty of Science and Engineering, Bernal Institute, University of Limerick, Limerick V94 T9PX, Ireland; [email protected]

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2021

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2021

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MDPI AG

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20734344

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English

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https://doi.org/10.3390/catal11060728

ProQuest document ID

2544622662

Application of TiO₂-Based Photocatalysts to Antibiotics Degradation: Cases of Sulfamethoxazole, Trimethoprim and Ciprofloxacin

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Application of TiO₂-Based Photocatalysts to Antibiotics Degradation: Cases of Sulfamethoxazole, Trimethoprim and Ciprofloxacin