1 Photocatalytic Degradation of Estriol using 2 Iron-doped TiO 2 under High and Low UV-irradiation 3 4

Iron Doped TiO2 nanoparticles (Fe-TiO2) were synthesized and photocatalitically 12 investigated under high and low fluence values of UV-radiation. The Fe-TiO2 physical 13 characterization was performed using X-ray Powder Diffraction (XRD), Brunauer-Emmett-Teller 14 (BET) surface area analysis, Transmission Electron Microscope (TEM), Scanning Electron 15 Microscope (SEM), Diffuse Reflectance Spectroscopy (DRS), and X-Ray Photoelectron Spectroscopy 16 (XPS) technique. The XPS evidenced that ferric ion (Fe3+) was in the lattice of TiO2 and co-dopants 17 no intentionally added were also present due to the precursors of the synthetic method. The Fe3+ 18 concentration played a key role in the photocatalytic generation of hydroxyl radical (•OH) and 19 estriol (E3) degradation. Fe-TiO2 materials accomplished E3 degradation, and it was found that the 20 catalyst with 0.3 at. % content of Fe (0.3 Fe-TiO2) enhanced the photocatalytic activity under low 21 UV-irradiation compared with no intentionally Fe-added TiO2 (zero-iron TiO2) and Aeroxide® TiO2 22 P25. Furthermore, the enhanced photocatalytic activity of 0.3 Fe-TiO2 under low UV-irradiation 23 may have applications when radiation intensity must be controlled, as in medical applications, or 24 when strong UV absorbing species are present in water. 25


Introduction
In recent years, society and the scientific community have concerned of Emerging Contaminants (ECs, also called Contaminants of Emerging Concern), which are chemicals that threaten the environment, human health, and water safety and are not currently covered by existing local or international water quality regulations [1].ECs include chemical species such as algae toxins, illegal drugs, industrial compounds, flame retardants, food additives, nanoparticles, pharmaceuticals (human and veterinary), personal care products, pesticides, biocides, steroids, synthetic and natural hormones, and surfactants [2].
Natural hormones (e.g., estrone (E1), 17β-estradiol (E2), and estriol (E3)) as ECs are susceptible of persisting and bioaccumulating in the environment, and could induce endocrine disruption in humans and wildlife (vertebrates [3][4][5] and invertebrates [6,7]).Natural attenuation, drinking water purification, and conventional municipal wastewater treatment processes are either incapable or only partially capable of removing estrogens from water [8].As result, water treatment techniques are being developed to manage, reduce, degrade, and mineralize low-concentrated ECs (including natural estrogen) in drinking and wastewater [9].Advanced Oxidation Processes (AOPs) are promising techniques to treat ECs in aqueous phase, which include well-known processes such as Fenton and Fenton-like processes, UV/H 2 O 2 , ozonation, and photocatalysis using semiconductors, peroxone processes (H 2 O 2 /O 3 ), and cavitation [10,11].Although there are many known AOPs, since Coleman's work [12], photocatalysis using titanium dioxide (TiO 2 ) has been identified as one of the most effective methods to degrade estrogens in water [13].Several reports recognized that TiO 2 can degrade estrogens, which prevents increases in estrogenic activity in water [14,15] and partially or completely mineralizing estrogens [14,16].
Titanium dioxide is the most commonly used photocatalyst because of its reasonable optical and electronic properties, good photocatalytic activity, insolubility in water, chemical and photochemical stability, nontoxicity, low cost, and high efficiency in pollutant mineralization [17][18][19][20].However, the band gap energy (E g ) of TiO 2 , frequently reported as 3.2 eV [21], restrains the photocatalytic activation to energy sources with a portion of spectrum emission below 387.5 nm [22].
In general the photocatalytic mechanism is as shown in Figure 1.According to Density Functional Theory (DFT) computations, the valence band (VB) and conduction band (CB) of pure TiO 2 are mainly composed of O2p orbitals and Ti3d orbitals, respectively.Hence, the Fermi level (EF) is located in the middle of the band gap (BG), indicating that VB is full filled while CB is empty [23].When using photons with energy higher than 3.2 eV, photoexcitation of the semiconductor promotes electrons from VB to CB creating a charge vacancy or hole (h + ) in the VB.The h + in the VB can react with hydroxide ion to form hydroxyl radical ( • OH) or can also be filled by donor absorbed organic molecule (OM ads ).Photogenerated electrons in the CB can be transferred to acceptor of electrons and bring about • OH.
Catalysts 2018, 8, x FOR PEER REVIEW 2 of 24 peroxone processes (H2O2/O3), and cavitation [10,11].Although there are many known AOPs, since Coleman's work [12], photocatalysis using titanium dioxide (TiO2) has been identified as one of the most effective methods to degrade estrogens in water [13].Several reports recognized that TiO2 can degrade estrogens, which prevents increases in estrogenic activity in water [14,15] and partially or completely mineralizing estrogens [14,16].Titanium dioxide is the most commonly used photocatalyst because of its reasonable optical and electronic properties, good photocatalytic activity, insolubility in water, chemical and photochemical stability, nontoxicity, low cost, and high efficiency in pollutant mineralization [17][18][19][20].However, the band gap energy (Eg) of TiO2, frequently reported as 3.2 eV [21], restrains the photocatalytic activation to energy sources with a portion of spectrum emission below 387.5 nm [22].
In general the photocatalytic mechanism is as shown in Figure 1.According to Density Functional Theory (DFT) computations, the valence band (VB) and conduction band (CB) of pure TiO2 are mainly composed of O2p orbitals and Ti3d orbitals, respectively.Hence, the Fermi level (EF) is located in the middle of the band gap (BG), indicating that VB is full filled while CB is empty [23].When using photons with energy higher than 3.2 eV, photoexcitation of the semiconductor promotes electrons from VB to CB creating a charge vacancy or hole (h + ) in the VB.The h + in the VB can react with hydroxide ion to form hydroxyl radical ( • OH) or can also be filled by donor absorbed organic molecule (OMads).Photogenerated electrons in the CB can be transferred to acceptor of electrons and bring about • OH.Consequently, reducing the photon energy needed for TiO2 photoactivation has been the focus of the scientific community until now.Doping is one of the techniques that has been tested to control or modify the surface properties or internal structure of TiO2.Doping introduces a foreign element into TiO2 to cause an impurity state in the band gap.The most frequently used doping materials are transition-metal cations (e.g., Cr, V, Fe, and Ni) at Ti sites, and anions (e.g., N, S, and C) at O sites [24].Among anion-and cation-dopants, the ferric ion (Fe 3+ ) is one of the most often used because the ionic radius of Fe 3+ (0.69 A) is similar to Ti 4+ (0.745 A) [25].Therefore, Fe 3+ can be easily incorporated into the TiO2 crystal lattice.
The main reported effects of iron-doped TiO2 is a rapid increase in photocatalytic activity that increases with increased Fe doping, which then reaches a maximum value, and finally decreases Consequently, reducing the photon energy needed for TiO 2 photoactivation has been the focus of the scientific community until now.Doping is one of the techniques that has been tested to control or modify the surface properties or internal structure of TiO 2 .Doping introduces a foreign element into TiO 2 to cause an impurity state in the band gap.The most frequently used doping materials are transition-metal cations (e.g., Cr, V, Fe, and Ni) at Ti sites, and anions (e.g., N, S, and C) at O sites [24].Among anion-and cation-dopants, the ferric ion (Fe 3+ ) is one of the most often used because the ionic radius of Fe 3+ (0.69 A) is similar to Ti 4+ (0.745 A) [25].Therefore, Fe 3+ can be easily incorporated into the TiO 2 crystal lattice.
Although several theoretical and experimental Fe-TiO 2 studies have been developed, the trade-off between doping ratio and radiation intensity is scarcely mentioned.Furthermore, Fe-TiO 2 photocatalyst has rarely been considered to be a useful technique for the degradation of E3 [42].
In this work, Fe-TiO 2 nanoparticles were synthesized to increase the understanding of the relationship between doping ratio and radiation intensity for hydroxyl radical ( • OH) generation and E3 degradation.Therefore, we investigated the photocatalytic degradation of E3 using Fe-TiO 2 under high and low UV irradiation.We highlight the term low UV irradiation to avoid confusion with the term "photocatalytic processes under visible light" because we did not intentionally use UV cutoff filters for the experiments.

Characterization of Iron-Doped TiO 2
Figure 2 shows X-ray Photoelectron Spectroscopy (XPS) general spectra of TiO 2 without added Fe (zero-iron TiO 2 ) and Fe-TiO 2 materials (b, c, and d).For the experimental condition used, Fe did not affect the bonding structure between titanium and oxygen because the main peaks for all samples were Ti2p and O1s with the proportion 1:2.2, which is in agreement with the atomic formula of TiO 2 .
Although several theoretical and experimental Fe-TiO2 studies have been developed, the trade-off between doping ratio and radiation intensity is scarcely mentioned.Furthermore, Fe-TiO2 photocatalyst has rarely been considered to be a useful technique for the degradation of E3 [42].
In this work, Fe-TiO2 nanoparticles were synthesized to increase the understanding of the relationship between doping ratio and radiation intensity for hydroxyl radical ( • OH) generation and E3 degradation.Therefore, we investigated the photocatalytic degradation of E3 using Fe-TiO2 under high and low UV irradiation.We highlight the term low UV irradiation to avoid confusion with the term "photocatalytic processes under visible light" because we did not intentionally use UV cutoff filters for the experiments.

Characterization of Iron-Doped TiO2
Figure 2 shows X-ray Photoelectron Spectroscopy (XPS) general spectra of TiO2 without added Fe (zero-iron TiO2) and Fe-TiO2 materials (b, c, and d).For the experimental condition used, Fe did not affect the bonding structure between titanium and oxygen because the main peaks for all samples were Ti2p and O1s with the proportion 1:2.2, which is in agreement with the atomic formula of TiO2.XPS detected unintentionally added elements such as carbon, sulfur, and nitrogen (Table 1) as co-dopants of zero-iron TiO 2 and Fe-TiO 2 , which were introduced into TiO 2 via precursors of the synthesis.Carbon and sulfur could come from sodium dodecyl sulfate (SDS), and nitrogen could come from iron (III) nitrate (Fe(NO 3 ) 3 •9H 2 O) and HNO 3 , all of them used in the synthesis process.High-resolution XPS spectra for the iron region (Figure 3) was studied only for 1.0 Fe-TiO 2 because no Fe2p signals were detected for zero-iron TiO 2 , 0.3 Fe-TiO 2 , or 0.6 Fe-TiO 2 .The deconvolution of high-resolution XPS spectra (Figure 3) was developed for previously reported peaks of Fe 2+ and Fe 3+ [43].Shirley baseline was subtracted before peak fitting.The Gaussian-Lorentzian mix function was used with a 40% factor.Charge compensation was set by the O1s peak charge with −0.58 eV.As a result, the correlation between the experimental signal and the theoretic model (Σχ 2 ) was 8.43 × 10 −2 .
XPS detected unintentionally added elements such as carbon, sulfur, and nitrogen (Table 1) as co-dopants of zero-iron TiO2 and Fe-TiO2, which were introduced into TiO2 via precursors of the synthesis.Carbon and sulfur could come from sodium dodecyl sulfate (SDS), and nitrogen could come from iron (III) nitrate (Fe(NO3)3•9H2O) and HNO3, all of them used in the synthesis process.High-resolution XPS spectra for the iron region (Figure 3) was studied only for 1.0 Fe-TiO2 because no Fe2p signals were detected for zero-iron TiO2, 0.3 Fe-TiO2, or 0.6 Fe-TiO2.The deconvolution of high-resolution XPS spectra (Figure 3) was developed for previously reported peaks of Fe 2+ and Fe 3+ [43].Shirley baseline was subtracted before peak fitting.The Gaussian-Lorentzian mix function was used with a 40% factor.Charge compensation was set by the O1s peak charge with −0.58 eV.As a result, the correlation between the experimental signal and the theoretic model (Σχ 2 ) was 8.43 × 10 −2 .According to the theoretical model (sum of fitting peaks), both Fe 3+ and Fe 2+ were present in the lattice of 1.0 Fe-TiO2.We suggest that Fe 3+ was incorporated into the lattice of TiO2 to form Ti-O-Fe bonds, because the ionic radius of Fe 3+ (0.69 A) is similar to the ionic radius of Ti 4+ (0.745 A) [25].The XPS technique detected Fe 2+ because Fe 3+ underwent reduction to Fe 2+ during XPS measurement in vacuum [44].
The band gap energy (Eg) obtained with the Kubelka-Monk method (Figure 4) for Aeroxide ® TiO2 P25 was 3.2 eV, which is consistent with the value reported previously [45].For Aeroxide ® TiO2 P25 Eg, red-shifts were detected as 0.22, 0.24, 0.25, and 0.3 eV for zero-iron TiO2, 0.3 Fe-TiO2, 0.6 Fe-TiO2, and 1.0 Fe-TiO2, respectively, which is consistent with values reported by Shi et al. of 0.25 eV [46] and with density functional theory calculations that suggested the hybridized band of Ti3d and Fe3d reduces Eg approximately 0.3-0.5 eV [44], or 0.2-0.34eV [47].According to the theoretical model (sum of fitting peaks), both Fe 3+ and Fe 2+ were present in the lattice of 1.0 Fe-TiO 2 .We suggest that Fe 3+ was incorporated into the lattice of TiO 2 to form Ti-O-Fe bonds, because the ionic radius of Fe 3+ (0.69 A) is similar to the ionic radius of Ti 4+ (0.745 A) [25].The XPS technique detected Fe 2+ because Fe 3+ underwent reduction to Fe 2+ during XPS measurement in vacuum [44].
The band gap energy (E g ) obtained with the Kubelka-Monk method (Figure 4) for Aeroxide ® TiO 2 P25 was 3.2 eV, which is consistent with the value reported previously [45].For Aeroxide ® TiO 2 P25 E g , red-shifts were detected as 0.22, 0.24, 0.25, and 0.3 eV for zero-iron TiO 2 , 0.3 Fe-TiO 2 , 0.6 Fe-TiO 2 , and 1.0 Fe-TiO 2 , respectively, which is consistent with values reported by Shi et al. of 0.25 eV [46] and with density functional theory calculations that suggested the hybridized band of Ti3d and Fe3d reduces E g approximately 0.3-0.5 eV [44], or 0.2-0.34eV [47].For zero-iron TiO2, Eg for Fe-TiO2 materials (Table 2) decreased as long as the Fe content increased, so the Fe content generated red-shift.For Aeroxide ® TiO2 P25 Eg, the red-shift of Fe-TiO2 agreed with previously reported values, but it agreed less for zero-iron TiO2.Therefore, red-shift was not only related to Fe content, but also to the synthesis method and unintentionally co-doped TiO2.
XRD patterns in Figure 5 revealed zero-iron TiO2 and Fe-TiO2 materials had both anatase and rutile phases.No XRD Fe2O3 peaks (2θ equal to 33.0°, 35.4°, 40.7°, 43.4°, and 49.2°) were observed, concluding that Fe 3+ replaced Ti 4+ in the TiO2 crystal framework [48,49].The synthesis method allowed uniform distribution of Fe within TiO2.The anatase:rutile phase ratio calculated by Spurr and Myers' method showed that zero-iron TiO2 and Fe-TiO2 materials were a mixture of anatase and rutile phases (Table 2).The amount of anatase was less in Fe-TiO2 materials than in Aeroxide ® TiO2 P25.The smaller proportion of anatase could lead to a reduction of photocatalytic activity because the anatase phase has higher photocatalytic activity than rutile TiO2 [50,51].However, it is accepted that the optimal photocatalytic activity of TiO2 is reached with an optimal mixture of anatase and rutile phases [52].Moreover, the increased anatase proportion in 0.3 Fe-TiO2 and 0.6 Fe-TiO2 compared with zero-iron TiO2 could improve photocatalytic activity.The increased anatase proportion was attributable to Fe doping disturbing the arrangements of TiO2 phases [53].This trend has also been observed when Fe-doped TiO2 was synthesized using sol-gel [54] or co-precipitation methods [32].For zero-iron TiO 2 , E g for Fe-TiO 2 materials (Table 2) decreased as long as the Fe content increased, so the Fe content generated red-shift.For Aeroxide ® TiO 2 P25 E g , the red-shift of Fe-TiO 2 agreed with previously reported values, but it agreed less for zero-iron TiO 2 .Therefore, red-shift was not only related to Fe content, but also to the synthesis method and unintentionally co-doped TiO 2 .[48,49].The synthesis method allowed uniform distribution of Fe within TiO 2 .The anatase:rutile phase ratio calculated by Spurr and Myers' method showed that zero-iron TiO 2 and Fe-TiO 2 materials were a mixture of anatase and rutile phases (Table 2).The amount of anatase was less in Fe-TiO 2 materials than in Aeroxide ® TiO 2 P25.The smaller proportion of anatase could lead to a reduction of photocatalytic activity because the anatase phase has higher photocatalytic activity than rutile TiO 2 [50,51].However, it is accepted that the optimal photocatalytic activity of TiO 2 is reached with an optimal mixture of anatase and rutile phases [52].Moreover, the increased anatase proportion in 0.3 Fe-TiO 2 and 0.6 Fe-TiO 2 compared with zero-iron TiO 2 could improve photocatalytic activity.The increased anatase proportion was attributable to Fe doping disturbing the arrangements of TiO 2 phases [53].This trend has also been observed when Fe-doped TiO 2 was synthesized using sol-gel [54] or co-precipitation methods [32].The average particle size of Fe-TiO2 materials obtained by Scherrer's formula was 6.9 nm, which is less than the particle size of Aeroxide ® TiO2 P25 (Table 2).Fe-TiO2 materials should increase photocatalytic activity because of their higher surface area and the short migration distance of the photogenerated charge carriers (electron/hole (e − /h + )) from the bulk material to the surface.
Further BET analysis (Figure 6) confirmed that average surface area of Fe-TiO2 materials was 77.9 m 2 g −1 , higher than zero-iron TiO2 and Aeroxide ® TiO2 P25.BET isotherms followed a type IV shape according to the Langmuir classification, which is associated with the characteristics of mesoporous material [55].The observed hysteresis is probably due to gas cooperative adsorption or condensation inside the pores of material [56].BET analysis showed pore sizes (Table 2) were in the mesoporous range (2-50 nm, according to IUPAC classification) for zero-iron TiO2 and 1.0 Fe-TiO2, and the microporous range (0.2-2 nm, according to IUPAC classification) for 0.3 Fe-TiO2 and 0.6 Fe-TiO2.Mesoporous pore size should facilitate the mass transfer of reactants and products in the reaction system, so photocatalytic improvement based on this property could improve zero-iron TiO2 and Fe-TiO2 materials with respect to Aeroxide ® TiO2 P25 [31].The average particle size of Fe-TiO 2 materials obtained by Scherrer's formula was 6.9 nm, which is less than the particle size of Aeroxide ® TiO 2 P25 (Table 2).Fe-TiO 2 materials should increase photocatalytic activity because of their higher surface area and the short migration distance of the photogenerated charge carriers (electron/hole (e − /h + )) from the bulk material to the surface.
Further BET analysis (Figure 6) confirmed that average surface area of Fe-TiO 2 materials was 77.9 m 2 g −1 , higher than zero-iron TiO 2 and Aeroxide ® TiO 2 P25.BET isotherms followed a type IV shape according to the Langmuir classification, which is associated with the characteristics of mesoporous material [55].The observed hysteresis is probably due to gas cooperative adsorption or condensation inside the pores of material [56].BET analysis showed pore sizes (Table 2) were in the mesoporous range (2-50 nm, according to IUPAC classification) for zero-iron TiO 2 and 1.0 Fe-TiO 2 , and the microporous range (0.2-2 nm, according to IUPAC classification) for 0.3 Fe-TiO 2 and 0.6 Fe-TiO 2 .Mesoporous pore size should facilitate the mass transfer of reactants and products in the reaction system, so photocatalytic improvement based on this property could improve zero-iron TiO 2 and Fe-TiO 2 materials with respect to Aeroxide ® TiO 2 P25 [31].
Patra et al. [49] developed a similar nanoparticle synthesis procedure, which generated surface area values ranging from 126 to 385 m 2 g −1 and mesoporous size distribution values ranging from 3.1 to 3.4 nm.Particles obtained in our work were different, probably because of the application of a mild thermal treatment and the use of SDS at critical micelle concentration as a template.
Figure 7 shows SEM images of agglomerated and assembled nanoparticles of zero-iron TiO 2 .The different amounts of Fe in the TiO 2 lattice changed neither the particle size nor the morphology of the zero-iron TiO 2 .Although the average pore size allowed an increase of the superficial area, agglomeration could lead to lower photocatalytic activity.Patra et al. [49] developed a similar nanoparticle synthesis procedure, which generated surface area values ranging from 126 to 385 m 2 g −1 and mesoporous size distribution values ranging from 3.1 to 3.4 nm.Particles obtained in our work were different, probably because of the application of a mild thermal treatment and the use of SDS at critical micelle concentration as a template.
Figure 7 shows SEM images of agglomerated and assembled nanoparticles of zero-iron TiO2.The different amounts of Fe in the TiO2 lattice changed neither the particle size nor the morphology of the zero-iron TiO2.Although the average pore size allowed an increase of the superficial area, agglomeration could lead to lower photocatalytic activity.Patra et al. [49] developed a similar nanoparticle synthesis procedure, which generated surface area values ranging from 126 to 385 m 2 g −1 and mesoporous size distribution values ranging from 3.1 to 3.4 nm.Particles obtained in our work were different, probably because of the application of a mild thermal treatment and the use of SDS at critical micelle concentration as a template.
Figure 7 shows SEM images of agglomerated and assembled nanoparticles of zero-iron TiO2.The different amounts of Fe in the TiO2 lattice changed neither the particle size nor the morphology of the zero-iron TiO2.Although the average pore size allowed an increase of the superficial area, agglomeration could lead to lower photocatalytic activity.Transmission electron microscopy (TEM) images confirmed nanoparticle clusters and particle sizes of zero-iron TiO 2 (Figure 8b) and 0.3 Fe-TiO 2 (Figure 8a) between 5 and 10 nm (between 1.2 and 9.4 nm according to Scherrer's formula).The lattice fringe spacing was 0.35 nm, as shown in Figure 8b, which was consistent with the d-spacing (101) of anatase [25].The lattice fingers of the nanoparticles showed that Fe-TiO 2 materials were highly crystallized.
Transmission electron microscopy (TEM) images confirmed nanoparticle clusters and particle sizes of zero-iron TiO2 (Figure 8b) and 0.3 Fe-TiO2 (Figure 8a) between 5 and 10 nm (between 1.2 and 9.4 nm according to Scherrer's formula).The lattice fringe spacing was 0.35 nm, as shown in Figure 8b, which was consistent with the d-spacing (101) of anatase [25].The lattice fingers of the nanoparticles showed that Fe-TiO2 materials were highly crystallized.

Characterization of Irradiation Source
Figure 9 shows the emission spectra of irradiation sources used in this study.Using the main peaks reported for a fluorescent lamp (Figure 9a), the calibration of the spectrometer generated an R 2 value equal to 0.999.The emission spectrum of the GE F15T8 BLB lamp (Figure 9b) was in the 356-410 nm range.However, the emission spectrum of the GE F15T8 D lamp (Figure 9c) was continuous broadband between 380 and 750 nm.The light intensity of the GE F15T8 lamp was reported to be between 3440 µW cm −2 [57] and 4000 µW cm −2 [58], from which 6% was UV radiation [59].The intensity of the GE F15T8 lamp was 1500 µW cm −2 .This lamp has an internal coating that absorbs 78% of visible light (as specified by the manufacturer) in the spectrum below 400 nm, as shown in Figure 9b.Therefore, the GE F15T8 BLB and GE F15T8 D lamps were designated as high and low UV irradiation sources, respectively.
Because Eg of Aeroxide ® TiO2 P25 is 3.2 eV (387.5 nm), see Figure 9, both the GE F15T8 BLB and GE F15T8 D lamps emitted photons that could photoactivate Aeroxide ® TiO2 P25.However, the proportion of the emission spectrum that Aeroxide ® TiO2 P25 could use for photocatalytic activity was different.An approximation of the amount of radiative intensity used for photocatalytic activity was obtained with the area under the curve-spectrum below the Eg value.Consequently, Aeroxide ® TiO2 P25 could take advantage of 36.4% of the emission spectrum of the GE F15T8 BLB lamp and 0.8% of the emission spectrum of the GE F15T8 D lamp.Table 2 lists amount of radiative spectrum used by zero-iron TiO2 and Fe-TiO2 materials according to each Eg.
Based on morphological and crystalline structure analysis, the favorable characteristics to enhance photocatalytic activity of Fe-TiO2 material are effective insertion of the Fe 3+ ion into the TiO2 lattice, red-shift (2.90-2.96eV), nanoparticle size (6.9-7.1 nm), specific surface area (73.0-83.1 nm), pore size (1.2-9.4 nm), and radiation absorbance below the equivalent Eg wavelength (8.21-10.63% of daylight lamp spectrum).Its main disadvantageous characteristics are expected to be high particle agglomeration and lower anatase phase compared with zero-iron TiO2.Further, photocatalytic activity is very sensitive to crystalline array and particle size and shape; differences in the density of hydroxyl groups on the particle surface and the number of water molecules hydrating the surface; the surface area and surface charge; differences in the number and nature of trap sites; the dopant

Characterization of Irradiation Source
Figure 9 shows the emission spectra of irradiation sources used in this study.Using the main peaks reported for a fluorescent lamp (Figure 9a), the calibration of the spectrometer generated an R 2 value equal to 0.999.The emission spectrum of the GE F15T8 BLB lamp (Figure 9b) was in the 356-410 nm range.However, the emission spectrum of the GE F15T8 D lamp (Figure 9c) was continuous broadband between 380 and 750 nm.The light intensity of the GE F15T8 lamp was reported to be between 3440 µW cm −2 [57] and 4000 µW cm −2 [58], from which 6% was UV radiation [59].The intensity of the GE F15T8 lamp was 1500 µW cm −2 .This lamp has an internal coating that absorbs 78% of visible light (as specified by the manufacturer) in the spectrum below 400 nm, as shown in Figure 9b.Therefore, the GE F15T8 BLB and GE F15T8 D lamps were designated as high and low UV irradiation sources, respectively.
Catalysts 2018, 8, x FOR PEER REVIEW 9 of 24 concentration, localization, and chemical state of the dopant ions; radiation intensity; particle aggregation and superficial charge; and scavenger species in media [39,60].Consequently, material characterization alone could not predict photocatalytic activity [28].Therefore, in this research, we used the N,N-dimethyl-p-nitrosoaniline (pNDA) probe and E3 to evaluate the photocatalytic activity by following • OH production, which is one of the most significant reactive oxygen species (ROS), and E3, which is an EC.

Hydroxyl Radical Generation under High and Low UV Irradiation
The generation of • OH was measured using pNDA, which is a well-characterized • OH scavenger as mentioned in Section 3.5.In brief, pNDA undergoes bleaching when reacting with • OH according to Muff et al. mechanism of the oxidation of pNDA by • OH [61].
In this work, pNDA bleaching followed a pseudo-first-order equation, so the apparent rate Because E g of Aeroxide ® TiO 2 P25 is 3.2 eV (387.5 nm), see Figure 9, both the GE F15T8 BLB and GE F15T8 D lamps emitted photons that could photoactivate Aeroxide ® TiO 2 P25.However, the proportion of the emission spectrum that Aeroxide ® TiO 2 P25 could use for photocatalytic activity was different.An approximation of the amount of radiative intensity used for photocatalytic activity was obtained with the area under the curve-spectrum below the E g value.Consequently, Aeroxide ® TiO 2 P25 could take advantage of 36.4% of the emission spectrum of the GE F15T8 BLB lamp and 0.8% of the emission spectrum of the GE F15T8 D lamp.Table 2 lists amount of radiative spectrum used by zero-iron TiO 2 and Fe-TiO 2 materials according to each E g .
Based on morphological and crystalline structure analysis, the favorable characteristics to enhance photocatalytic activity of Fe-TiO 2 material are effective insertion of the Fe 3+ ion into the TiO 2 lattice, red-shift (2.90-2.96eV), nanoparticle size (6.9-7.1 nm), specific surface area (73.0-83.1 nm), pore size (1.2-9.4 nm), and radiation absorbance below the equivalent E g wavelength (8.21-10.63% of daylight lamp spectrum).Its main disadvantageous characteristics are expected to be high particle agglomeration and lower anatase phase compared with zero-iron TiO 2 .Further, photocatalytic activity is very sensitive to crystalline array and particle size and shape; differences in the density of hydroxyl groups on the particle surface and the number of water molecules hydrating the surface; the surface area and surface charge; differences in the number and nature of trap sites; the dopant concentration, localization, and chemical state of the dopant ions; radiation intensity; particle aggregation and superficial charge; and scavenger species in media [39,60].Consequently, material characterization alone could not predict photocatalytic activity [28].Therefore, in this research, we used the N,N-dimethyl-p-nitrosoaniline (pNDA) probe and E3 to evaluate the photocatalytic activity by following • OH production, which is one of the most significant reactive oxygen species (ROS), and E3, which is an EC.

Hydroxyl Radical Generation under High and Low UV Irradiation
The generation of • OH was measured using pNDA, which is a well-characterized • OH scavenger as mentioned in Section 3.5.In brief, pNDA undergoes bleaching when reacting with • OH according to Muff et al. mechanism of the oxidation of pNDA by • OH [61].
In this work, pNDA bleaching followed a pseudo-first-order equation, so the apparent rate constant was calculated by ln(C/C 0 ) = k 1 t, where C 0 is the initial concentration, C is the reaction concentration at a given time, and k 1 is the pseudo-first-order reaction rate constant.The slope of the plot after applying a linear fit represents the rate constant, k 1 .
Because the relationship between pNDA bleaching and • OH production follows a 1:1 stoichiometry [61], the steady-state of • OH generation ([ • OH] ss ) can be considered equal to the initial velocity (r 0 ) according to Equation (1) and reported in Table 3: Fe-TiO 2 materials showed a similar anatase:rutile phase ratio, particle size, and specific surface area, and therefore the variation in r 0 values was due to the difference of Fe content inside TiO 2 .The generation of • OH radicals (r 0 ) was feasible using zero-iron TiO 2 , Fe-TiO 2 materials, and Aeroxide ® TiO 2 P25 under both high (Figure 10a) and low UV irradiation (Figure 10b).When high UV irradiation was used, the maximum r0 was 0.58 µM•OH min −1 for 0.3 Fe-TiO2.The enhancement in photocatalytic activity of 0.3 at.%Fe-TiO2, compared with zero-iron TiO2 was by the extended lifetime values of the photogenerated charge carriers (e − and h + ) produced by Fe 3+ ions, which played a role as charge carriers trapped at or near the particle surface.The trapping mechanisms are shown in Equations ( 2)-( 5) [62].
The mechanism suggested for • OH generation is shown in Figure 11.When TiO2 contains a Fe 3+ ion, the Fe3d orbitals split into two bands, one is a hybrid band (A2g) and one is midgap band (T2g), which induce a new localized BG state [23].Therefore, when TiO2 absorbs photons with energy less than 3.2 eV, photoexcitation of the semiconductor promotes an electron from the VB to the midgap band (T2g), also called a shallow trap, creating an electron-hole pair.The hole in the valence band (VB) can react with hydroxide ions to form • OH, absorbed organic molecules, or trap Fe 3+ following Equations ( 4) and (5).Additionally, photogenerated electrons in the midgap band (T2g) can be transferred to Fe 3+ following a dark redox reaction at the interface, as suggested by Neubert et al. [63] and consequently bring about • OH.   Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degrada general, the photocatalytic activity first increased and then decreased as the Fe concen increased, which is similar to the behavior found with the • OH probe in Section 2.3 and h previously reported using other organic molecules [23,29,64].Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degradation.In general, the photocatalytic activity first increased and then decreased as the Fe concentration increased, which is similar to the behavior found with the • OH probe in Section 2.3 and has been previously reported using other organic molecules [23,29,64].Under high UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iron TiO2, 0.3 Fe-TiO2, and 1.0 Fe-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was related with   Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degradatio general, the photocatalytic activity first increased and then decreased as the Fe concentr increased, which is similar to the behavior found with the • OH probe in Section 2.3 and has previously reported using other organic molecules [23,29,64].
Fe 3+ +h vb + → Fe 4+ hole trap (4) The mechanism suggested for • OH generation is shown in Figure 11.When TiO 2 contains a Fe 3+ ion, the Fe3d orbitals split into two bands, one is a hybrid band (A2g) and one is midgap band (T2g), which induce a new localized BG state [23].Therefore, when TiO 2 absorbs photons with energy less than 3.2 eV, photoexcitation of the semiconductor promotes an electron from the VB to the midgap band (T2g), also called a shallow trap, creating an electron-hole pair.The hole in the valence band (VB) can react with hydroxide ions to form • OH, absorbed organic molecules, or trap Fe 3+ following Equations ( 4) and (5).Additionally, photogenerated electrons in the midgap band (T2g) can be transferred to Fe 3+ following a dark redox reaction at the interface, as suggested by Neubert et al. [63] and consequently bring about • OH.Increasing the Fe 3+ doping content of Fe-TiO2 to 0.6 and 1.0 at.%,Fe-TiO2 was unfavorable to the photocatalytic activity because the additional Fe 3+ doping in the TiO2 sample inhibited the extended lifetime of charge carriers, acted as recombination sites and consequently decreased the photocatalytic efficiency [29], as proposed in Equations ( 6)-( 9) [39].
Fe 4+ + ecd − → Fe 3+ recombination (7) Fe 4+ + Ti 3+ → Fe 3+ + Ti 4+ recombination ( When low UV irradiation conditions were used, the r0 values for zero-iron TiO2 and Fe-TiO2 materials were lower than the value estimated for Aeroxide ® TiO2 P25.Compared with the effects of high UV irradiation, the reduction in r0 value observed was related both to pNDA adsorption of UV-visible radiation (lowered the number of photons available to activate the photocatalyst), and the augmented Fe content, which increased the recombination rate.

Photocatalytic Degradation of Estriol under High and Low UV Irradiation
E3 photocatalytic degradation curves are shown in Figure 12a,b using both high and low UV irradiation, respectively.In both cases, E3 photocatalytic degradation followed a pseudo-first-order model and the rate constant, k1 (Table 4), was obtained by fitting experimental data to ln ([E3]/[E30]) = k1t.Fe content influenced k1 for both high and low UV irradiation.Increasing the Fe 3+ doping content of Fe-TiO 2 to 0.6 and 1.0 at.%,Fe-TiO 2 was unfavorable to the photocatalytic activity because the additional Fe 3+ doping in the TiO 2 sample inhibited the extended lifetime of charge carriers, acted as recombination sites and consequently decreased the photocatalytic efficiency [29], as proposed in Equations ( 6)-( 9) [39].
Fe 4+ + e cd − → Fe 3+ recombination (7) Fe 4+ + Fe 2+ → 2Fe 3+ recombination (8) When low UV irradiation conditions were used, the r 0 values for zero-iron TiO 2 and Fe-TiO 2 materials were lower than the value estimated for Aeroxide ® TiO 2 P25.Compared with the effects of high UV irradiation, the reduction in r 0 value observed was related both to pNDA adsorption of UV-visible radiation (lowered the number of photons available to activate the photocatalyst), and the augmented Fe content, which increased the recombination rate.

Photocatalytic Degradation of Estriol under High and Low UV Irradiation
E3 photocatalytic degradation curves are shown in Figure 12a,b using both high and low UV irradiation, respectively.In both cases, E3 photocatalytic degradation followed a pseudo-first-order model and the rate constant, k 1 (Table 4), was obtained by fitting experimental data to ln ([E3]/[E3 0 ]) = k 1 t.Fe content influenced k 1 for both high and low UV irradiation.Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degradation.In general, the photocatalytic activity first increased and then decreased as the Fe concentration increased, which is similar to the behavior found with the • OH probe in Section 2.3 and has been previously reported using other organic molecules [23,29,64].Under high UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iron TiO2, 0.3 Fe-TiO2, and 1.0 Fe-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was related with  Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degradation.In general, the photocatalytic activity first increased and then decreased as the Fe concentration increased, which is similar to the behavior found with the • OH probe in Section 2.3 and has been previously reported using other organic molecules [23,29,64].Under high UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iron TiO2, 0.3 Fe-TiO2, and 1.0 Fe-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was related with zero-iron TiO 2 , re 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degradation.In the photocatalytic activity first increased and then decreased as the Fe concentration , which is similar to the behavior found with the • OH probe in Section 2.3 and has been ly reported using other organic molecules [23,29,64].er high UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iron TiO2, 0.3 and 1.0 Fe-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was related with  Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degradation.In general, the photocatalytic activity first increased and then decreased as the Fe concentration increased, which is similar to the behavior found with the • OH probe in Section 2.3 and has been previously reported using other organic molecules [23,29,64].Under high UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iron TiO2, 0.3 Fe-TiO2, and 1.0 Fe-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was related with  Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degradation.In general, the photocatalytic activity first increased and then decreased as the Fe concentration increased, which is similar to the behavior found with the • OH probe in Section 2.3 and has been previously reported using other organic molecules [23,29,64].Under high UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iron TiO2, 0.3 Fe-TiO2, and 1.0 Fe-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was related with  Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degradation.In general, the photocatalytic activity first increased and then decreased as the Fe concentration increased, which is similar to the behavior found with the • OH probe in Section 2.3 and has been previously reported using other organic molecules [23,29,64].Under high UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iron TiO2, 0.3 Fe-TiO2, and 1.0 Fe-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was related with  Figure 13 shows the pseudo-first-order rate constant (k 1 ) of E3 photocatalytic degradation.In general, the photocatalytic activity first increased and then decreased as the Fe concentration increased, which is similar to the behavior found with the • OH probe in Section 2.3 and has been previously reported using other organic molecules [23,29,64].Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degradation.In general, the photocatalytic activity first increased and then decreased as the Fe concentration increased, which is similar to the behavior found with the • OH probe in Section 2.3 and has been previously reported using other organic molecules [23,29,64].Under high UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iron TiO2, 0.3 Fe-TiO2, and 1.0 Fe-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was related with    Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degradation.In general, the photocatalytic activity first increased and then decreased as the Fe concentration increased, which is similar to the behavior found with the • OH probe in Section 2.3 and has been previously reported using other organic molecules [23,29,64].Under high UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iron TiO2, 0.3 Fe-TiO2, and 1.0 Fe-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was related with    Under high UV irradiation (Figure 13a), 0.6 Fe-TiO 2 k 1 was higher than for zero-iron TiO 2 , 0.3 Fe-TiO 2 , and 1.0 Fe-TiO 2 .The increase in photocatalytic performance of 0.6 Fe-TiO 2 was related with the increase in the lifetime of electron-hole pairs because Fe created additional energy levels near the conduction band of TiO 2 , as the mechanism suggests in Figure 11.
Under low UV irradiation (Figure 13b), zero-iron TiO 2, 0.3 Fe-TiO 2, and 0.6 Fe-TiO 2 showed more photocatalytic activity than Aeroxide ® TiO 2 P25 because those materials had enhanced superficial properties, such as particle size, and superficial area, as mentioned in Section 2.1.Furthermore, 0.3 Fe-TiO 2 enhanced photocatalytic activities with k 1 values as high as 0.005 min −1 .The high photocatalytic activity of 0.3 Fe-TiO 2 was due to the synergistic effect of unintentionally added co-dopants, superficial properties, and Fe content that increased the lifetime of photogenerated charge carriers and the efficiency of electron transfer.
The photocatalytic degradation rate of E3 using Aeroxide ® TiO 2 P25 was reported to be 0.25 min −1 [65], 0.134 min −1 [66], and 0.12 min −1 [67].However, the experimental setups and catalyst loads were different.Besides these few studies, E3 degradation using Fe-TiO 2 nanoparticles is scarcely reported.Only comparing magnitudes of k 1 , the first-order rates to degrade pharmaceuticals using Fe-TiO 2 nanoparticles were 0.001 min −1 for ibuprofen, 0.0015 min −1 for carbamazepine, and 0.0014 min −1 for sulfamethoxazole [68], which are in the order of magnitude obtained in this work (see Table 4).
Regarding unintentionally added co-dopants, Fe-TiO 2 co-doping demonstrated a synergistic effect to increase photocatalytic activity under visible light for sulfur [69], nitrogen [44], and Fe x T i1-x O 2-y N y co-doping [70].Surface properties of the material, such as a particle size (6.9 nm) and surface area (77.6 m 2 g −1 ), also facilitated the mass transfer between interface, E3, and sub-products.
The relationship between the • OH radical system and E3 kinetic degradation was determined via linear fit between • OH initial rate generation (r 0,OH ) and initial E3 degradation (r 0,E3 ).In general, the procedure to correlate r 0,OH and r 0,E3 was first to sort pair values (r 0,OH , r 0,E3 ), and then fit the data to linear regression, as shown Figure 14a,b.
Catalysts 2018, 8, x FOR PEER REVIEW 13 of 24 the increase in the lifetime of electron-hole pairs because Fe created additional energy levels near the conduction band of TiO2, as the mechanism suggests in Figure 11.Under low UV irradiation (Figure 13b), zero-iron TiO2, 0.3 Fe-TiO2, and 0.6 Fe-TiO2 showed more photocatalytic activity than Aeroxide ® TiO2 P25 because those materials had enhanced superficial properties, such as particle size, and superficial area, as mentioned in Section 2.1.Furthermore, 0.3 Fe-TiO2 enhanced photocatalytic activities with k1 values as high as 0.005 min −1 .The high photocatalytic activity of 0.3 Fe-TiO2 was due to the synergistic effect of unintentionally added co-dopants, superficial properties, and Fe content that increased the lifetime of photogenerated charge carriers and the efficiency of electron transfer.
The photocatalytic degradation rate of E3 using Aeroxide ® TiO2 P25 was reported to be 0.25 min −1 [65], 0.134 min −1 [66], and 0.12 min −1 [67].However, the experimental setups and catalyst loads were different.Besides these few studies, E3 degradation using Fe-TiO2 nanoparticles is scarcely reported.Only comparing magnitudes of k1, the first-order rates to degrade pharmaceuticals using Fe-TiO2 nanoparticles were 0.001 min −1 for ibuprofen, 0.0015 min −1 for carbamazepine, and 0.0014 min −1 for sulfamethoxazole [68], which are in the order of magnitude obtained in this work (see Table 4).
The relationship between the • OH radical system and E3 kinetic degradation was determined via linear fit between • OH initial rate generation (r0,OH) and initial E3 degradation (r0,E3).In general, the procedure to correlate r0,OH and r0,E3 first to sort pair values (r0,OH, r0,E3), and then fit the data to linear regression, as shown Figure 14a,b.Under high UV irradiation, the linear fit correlation was r0,E3 = 0.091 r0,OH + 0.040 with R 2 = 0.197.Under low UV irradiation, the linear fit correlation was r0,E3 = 0.066 r0,OH + 0.012 with R 2 = 0.975.The correlation between the pair (r0,OH, r0,E3) under high UV irradiation was too low to be considered a linear relationship.We suggest the low correlation was because not only • OH caused E3 degradation, but holes (h + ) or other reactive oxygen species also caused E3 degradation.
However, a linear relationship under low UV irradiation was attributable to • OH being the main reactive oxygen species responsible for photocatalytic activity.Therefore, the contribution of h + to photocatalytic activity was lower because oxidation power was lower due to reduced Eg.This suggestion supports the mechanisms proposed in Figure 11, in which adding Fe into the lattice of TiO2 reduced the Eg with a consistent reduction of redox potential, as mentioned by others [28].

Catalyst
Load High UV Irradiation k1 R 2 r0,E3 mg L −1 min −1 μME3 min −1 TiO2 Aeroxide ® P25 20 0.021 0.996 0.21 Zero-iron TiO2 320 0.007 0.997 0.069 0.3 Fe-TiO2 320 0.009 0.994 0.090 0.6 Fe-TiO2 320 0.011 0.997 0.099 1.0 Fe-TiO2 320 0.003 0.979 0.027 Figure 13 shows the pseudo-first-order rate constant (k1) o general, the photocatalytic activity first increased and then increased, which is similar to the behavior found with the • OH previously reported using other organic molecules [23,29,64].Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degr general, the photocatalytic activity first increased and then decreased as the Fe con increased, which is similar to the behavior found with the • OH probe in Section 2.3 an previously reported using other organic molecules [23,29,64].Under high UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iro Fe-TiO2, and 1.0 Fe-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was re     s the pseudo-first-order rate constant (k1) of E3 photocatalytic degradation.In catalytic activity first increased and then decreased as the Fe concentration similar to the behavior found with the • OH probe in Section 2.3 and has been using other organic molecules [23,29,64].Under high UV irradiation, the linear fit correlation was r 0,E3 = 0.091 r 0,OH + 0.040 with R 2 = 0.197.Under low UV irradiation, the linear fit correlation was r 0,E3 = 0.066 r 0,OH + 0.012 with R 2 = 0.975.The correlation between the pair (r 0,OH , r 0,E3 ) under high UV irradiation was too low to be considered a linear relationship.We suggest the low correlation was because not only OH caused E3 degradation, but holes (h + ) or other reactive oxygen species also caused E3 degradation.
However, a linear relationship under low UV irradiation was attributable to • OH being the main reactive oxygen species responsible for photocatalytic activity.Therefore, the contribution of h + to photocatalytic activity was lower because oxidation power was lower due to reduced E g .This suggestion supports the mechanisms proposed in Figure 11, in which adding Fe into the lattice of TiO 2 reduced the E g with a consistent reduction of redox potential, as mentioned by others [28].
The main mechanism of E3 degradation under low UV irradiation was via electron (e − ) transfer to give rise • OH.Additionally, the enhanced photocatalytic activity of 0.3 Fe-TiO 2 under low UV irradiation provides evidence that the trapping-recombination mechanism of Fe-TiO 2 can be controlled by irradiation intensity.Therefore, we suggest that there is a trade-off between irradiation intensity, the trapping-recombination rate, and • OH production that is worthy of further research.
The efficiency resource of the Fe-TiO 2 /Low UV system was obtained through dimensional analysis of the slope of the linear fit of data shown in Figure 14b.The units of slope are E3 moles degraded per • OH mol generated at initial time, so 0.662 E3 molecules underwent degradation when one • OH was generated for the photocatalytic system independent of Fe doping content in TiO 2 .A sustainable process was also achieved, for which 0.3 Fe-TiO 2 since absorbed 8.21% of emission spectra of the lamp below the equivalent E g wavelength over 0.8% or 7.64% of Aeroxide ® TiO 2 P25 and zero-iron TiO 2 , respectively.

Relationship between Fe Content and Kinetic Constant
Photonic efficiency has been suggested to increase linearly with the doping ratio due to the formation of the charge carrier trapping centers, while it concurrently decreases quadratically with the doping ratio because to the creation of recombination centers [71].Alternatively, we suggest an empirical relationship between the E3 degradation pseudo-first-order rate constant (k 1 ) and Fe content (at.%) in TiO 2 , as described in Equation ( 10): where k 1 is the pseudo-first-order constant, k e is the electron trap constant, k a is the electron recombination constant, δ at.% is the Fe doping amount in TiO 2 , and c and α are system constants.To solve the model described in Equation a numerical approximation by root-mean-square error minimization method was used according to Equation ( 11): where [k 1.i ] is the theoretical k 1 value, [k 1.i ] is the experimental k 1 value, n is the number of data, and ε is the root-mean-square error.The solution of Equation ( 10) was performed by simultaneously solving k e , k a , c, and α using Excel Solver ® (Frontline Systems, NV, US).As an example, photocatalytic degradation of E3 under low UV irradiation was fitted to Equation (10), as shown in Figure 15.
The empirical model solved in Equation (12) shows that electron trap constant (k e ) overcome electron recombination (k a ) before optimal catalyst load.This model could lead to experimental work using iron-doped TiO 2 in which the optimal content of Fe gives rise to the maximum E3 degradation.k 1 (δ) = −1.99 e −2.81(δ+0.197)− e −2. 78(δ+0.197)(12) The empirical model solved in Equation (12) shows that electron trap constant (ke) overcome electron recombination (ka) before optimal catalyst load.This model could lead to experimental work using iron-doped TiO2 which the optimal content of Fe gives rise to the maximum E3 degradation.k (δ) = −1.99 e . (. ) − e . (. )  (12)

Photoreactor Setup
Figure 16 depicts the photoreactor, which was a cylindrical water-jacketed glass vessel (318 mL) with 102 mm and 63 mm of interior height and diameter, respectively.The horizontal and vertical position of the photoreactor was constant for all experiments.Lamps were set horizontally and centered above the photoreactor.Two 15 W GE F15T8 BLB lamps (also called black-light lamps, Boston, MA, USA) supplied high UV irradiation, and two 15 W GE F15T8 D lamps (also called daylight lamps) provided low UV irradiation.The overall system was in a closed box to avoid the effects of sunlight or any artificial radiation sources.Lamp emission spectra were measured using a lab-made spectrophotometer using a CMOS webcam with a diffraction grating of 1000 lines mm −1 [72,73].Emission spectra calibration of the spectrophotometer was developed using a 9 W fluorescent lamp (Tecnolite, Jalisco, Mexico).The temperature of all experiments was set at 20 °C  Figure 13 shows the pseudo-fi general, the photocatalytic activity increased, which is similar to the b previously reported using other org  UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iron TiO2, 0.3 e-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was related with  Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degradation.In general, the photocatalytic activity first increased and then decreased as the Fe concentration increased, which is similar to the behavior found with the • OH probe in Section 2.3 and has been previously reported using other organic molecules [23,29,64].Under high UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iron TiO2, 0.3 Fe-TiO2, and 1.0 Fe-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was related with  Figure 13 shows the pseudo-first-order rate constant (k1) of E3 photocatalytic degradation.In general, the photocatalytic activity first increased and then decreased as the Fe concentration increased, which is similar to the behavior found with the • OH probe in Section 2.3 and has been previously reported using other organic molecules [23,29,64].Under high UV irradiation (Figure 13a), 0.6 Fe-TiO2 k1 was higher than for zero-iron TiO2, 0.3 Fe-TiO2, and 1.0 Fe-TiO2.The increase in photocatalytic performance of 0.6 Fe-TiO2 was related with 1.0 Fe-TiO 2 ; at pH 6 ± 0.1; and 20 • C.

Photoreactor Setup
Figure 16 depicts the photoreactor, which was a cylindrical water-jacketed glass vessel (318 mL) with 102 mm and 63 mm of interior height and diameter, respectively.The horizontal and vertical position of the photoreactor was constant for all experiments.Lamps were set horizontally and centered above the photoreactor.Two 15 W GE F15T8 BLB lamps (also called black-light lamps, Boston, MA, USA) supplied high UV irradiation, and two 15 W GE F15T8 D lamps (also called daylight lamps) provided low UV irradiation.The overall system was in a closed box to avoid the effects of sunlight or any artificial radiation sources.Lamp emission spectra were measured using a lab-made spectrophotometer using a CMOS webcam with a diffraction grating of 1000 lines mm −1 [72,73].Emission spectra calibration of the spectrophotometer was developed using a 9 W fluorescent lamp (Tecnolite, Jalisco, Mexico).The temperature of all experiments was set at 20 • C using a thermostatic bath with recirculation (Polystat, Cole-Palmer, Vernon Hills, IL, USA).An optical filter was not used in the experiments, so visible light condition was not simulated.

Synthesis of Materials
The synthesis method of iron-doped TiO2 (Fe-TiO2) materials followed the hydrothermal sol-gel synthetic approach proposed by Patra et al. with some differences in precursor and thermal treatment [49].Our synthesis method used iron (III) nitrate instead of FeCl3 and absolute ethanol instead of isopropyl alcohol.The thermal treatment was a programmed cycle of 31 h (increasing ramp-drying-increasing ramp-calcination-decreasing ramp) instead of direct calcination for 6 h.First, solution A was prepared by dissolving 1.44 g of SDS in 10 mL of deionized water.Then, four different solutions B were prepared to dissolve iron (III) nitrate in 2 mL of absolute ethanol (≥99.8 %) and 3 mL of TTIP was added slowly.The amounts of iron (III) nitrate were 0, 0.4, 4.3, and 42.6 mg of Fe(NO3)3•9H2O identified as zero-iron TiO2, 0.3 Fe-TiO2, 0.6 Fe-TiO2, and 1.0 Fe-TiO2, respectively.Once ready, solution A was continuously stirred and solution B was slowly dropped into solution A. The pH the resulting mixture was adjusted to 1 using concentrated HNO3 and stirred for 3 h.The mixture was kept at 3 °C for 36 h.The precipitated solid was collected by filtration using Whatman Quantitative Filter Paper Grade 42.The materials were simultaneously dried and calcinated with a programmed thermal treatment (Isotemp ® Programmable Muffle Furnace, Fisher Scientific, Dubuque, IA, USA) following first the temperature increase from ambient temperature to 353 K, with a temperature ramp of 1 K min −1 that was held for 720 min.The temperature was then increased from 353 K to 773 K with a temperature ramp of 1 K min −1 that was held for 360 min.Finally, the temperature was decreased from 773 K to 353 K with a temperature ramp of −1 K min −1 , and then the furnace was turned off.The materials were washed with 50:50 methanol-water and dried to 377 K overnight.5), an optical filter (if needed) (6), stirring plate (7), cooling fan (8), horizontal position template (9), and lab jack lifting platform (10).

Synthesis of Materials
The synthesis method of iron-doped TiO 2 (Fe-TiO 2 ) materials followed the hydrothermal sol-gel synthetic approach proposed by Patra et al. with some differences in precursor and thermal treatment [49].Our synthesis method used iron (III) nitrate instead of FeCl 3 and absolute ethanol instead of isopropyl alcohol.The thermal treatment was a programmed cycle of 31 h (increasing ramp-drying-increasing ramp-calcination-decreasing ramp) instead of direct calcination for 6 h.First, solution A was prepared by dissolving 1.44 g of SDS in 10 mL of deionized water.Then, four different solutions B were prepared to dissolve iron (III) nitrate in 2 mL of absolute ethanol (≥99.8 %) and 3 mL of TTIP was added slowly.The amounts of iron (III) nitrate were 0, 0.4, 4.3, and 42.6 mg of Fe(NO 3 ) 3 •9H 2 O identified as zero-iron TiO 2 , 0.3 Fe-TiO 2 , 0.6 Fe-TiO 2 , and 1.0 Fe-TiO 2 , respectively.Once ready, solution A was continuously stirred and solution B was slowly dropped into solution A. The pH of the resulting mixture was adjusted to 1 using concentrated HNO 3 and stirred for 3 h.The mixture was kept at 3 • C for 36 h.The precipitated solid was collected by filtration using Whatman Quantitative Filter Paper Grade 42.The materials were simultaneously dried and calcinated with a programmed thermal treatment (Isotemp ® Programmable Muffle Furnace, Fisher Scientific, Dubuque, IA, USA) following first the temperature increase from ambient temperature to 353 K, with a temperature ramp of 1 K min −1 that was held for 720 min.The temperature was then increased from 353 K to 773 K with a temperature ramp of 1 K min −1 that was held for 360 min.Finally, the temperature was decreased from 773 K to 353 K with a temperature ramp of −1 K min −1 , and then the furnace was turned off.The materials were washed with 50:50 methanol-water and dried to 377 K overnight.

Materials Characterization
X-ray photoelectron spectroscopy (XPS) was performed using a Thermo Fisher Scientific K-Alpha X-ray photoelectron spectrometer (Waltham, MA, USA) with a monochromatized Al K α X-ray source (1487 V).The deconvolution of high-resolution XPS spectra was developed using the software XPSpeak 4.1.(Raymund W.M. Kwok, Shatin, Hong Kong).
UV-visible reflectance spectroscopy was obtained with Video-Barrelino integrating sphere coupled to Cary 50 spectrophotometer (Varian Inc, Palo Alto, CA, USA).Diffuse reflectance spectra were transformed using the Kubelka-Munk method to obtain E g of zero-iron TiO 2 and Fe-TiO 2 materials.Kubelka-Munk method plots (F(R)hv) 1/2 versus hv, draws a tangent at the inflection point on the curve and estimates E g with the hv value at the intersection with abscissa.In this case, F(R) is a reflectance function equal to (1 − R) 2 /2R, R the reflectance percentage, h is the Planck's constant, and v is frequency.
XRD patterns were recorded in a Siemens D-5000 diffractometer (Munich, Germany) using Cu K α radiation (λ = 1.54060Å) from 10 • to 85 • .The procedure for phase identification used the QualX2.0software with database developed by Altomare et al. [74].The cards used for identification were 00-901-5929, 00-900-1681, and 00-900-4140 for anatase, rutile, and brookite, respectively.The quantification phases followed the method proposed by Spurr and Myers according to Equation ( 13): where f is the anatase percentage, I A is intensity at a diffraction angle 2θ of 25.36The particle size was estimated by Scherrer's formula described in Equation ( 14), where β is the full width at half of the maximum of the diffraction peaks (radians), k is the shape constant, λ is the wavelength of the incident Cu K α radiation (λ = 1.54060Å), θ is the Bragg's angle (radians), and D is the particle size (Å).
Brunauer-Emmett-Teller (BET) isotherms were obtained in Nova Station A equipment (Quantachrome Instruments, Boynton Beach, FL, USA).The surface morphology was observed by SEM in a JEOL ultrahigh resolution field emission electron microscope JSM-7800 F (JEOL, Tokyo, Japan) with 20 kV accelerating voltage, and 3 mm WD.Transmission electron microscopy (TEM) images were obtained in a JEM-2100 LaB6 electron microscope (JEOL, Tokyo, Japan).

Hydroxyl Radical Generation
In this study, pNDA bleaching was selected as an • OH probe because pNDA was useful for measuring the photocatalytic performance of TiO 2 [51,76,77] because of the following advantages: (1) it is selective of the reaction of pNDA with • OH [78]; (2) its high reaction rate with • OH on the order of 10 10 M −1 s −1 [51,79]; (3) its easy application through observable bleaching at 440 nm following Beer's Law, in which pNDA bleaching a yellowish solution to transparent; and (4) its 1:1 stoichiometry, meaning that one • OH can bleach one pNDA molecule [51,[80][81][82].
The pNDA absorption (Figure 17) measurements were obtained using a UV-visible spectrophotometer (Hatch DR/4000U, Loveland, CO, USA) at 440 nm following Beer-Lambert law.The pNDA test solution was 10 µM initial concentration and pH 6.0 ± 0.1 adjusted using NaOH or HCl when needed.No buffer solutions were used because they can compete for • OH.Final pH was verified at the end of tests to discharge pH-pNDA bleaching.The photocatalytic standard was Aeroxide ® TiO2 P25, and the load was 20 mg L −1 .The choice of catalyst load was based on our previous work on • OH generation of Aeroxide ® TiO2 P25 [16].For zero-iron TiO2 and Fe-TiO2 materials, the catalyst load used was 320 mg L −1 , which produced a • OH generation rate under high UV irradiation to set a baseline.Catalyst load differences were attributable to the aggregation of lab-made TiO2, superficial properties, and optical properties of suspensions, as shown in Figure 18.
The photocatalytic experiments were conducted as follows.First, a pNDA test solution was set at 20 °C, the catalyst was added, and the suspension was mixed for 20 min without radiation.To evaluate the adsorption of pNDA on TiO2, an aliquot was withdrawn and centrifuged.Then, the system was fully illuminated, and aliquots were withdrawn after specific periods.Each sample was centrifuged at 6000 rpm for 15 min (Biofuge Primo, Sorvall, Hanau, Germany) and measured in the UV-visible spectrophotometer.Once the catalyst load was used and after the dark phase, no adsorption of pNDA was detected near the detection limit of UV-visible spectrophotometer.The photocatalytic standard was Aeroxide ® TiO 2 P25, and the load was 20 mg L −1 .The choice of catalyst load was based on our previous work on • OH generation of Aeroxide ® TiO 2 P25 [16].For zero-iron TiO 2 and Fe-TiO 2 materials, the catalyst load used was 320 mg L −1 , which produced a • OH generation rate under high UV irradiation to set a baseline.Catalyst load differences were attributable to the aggregation of lab-made TiO 2 , superficial properties, and optical properties of suspensions, as shown in Figure 18.The photocatalytic standard was Aeroxide ® TiO2 P25, and the load was 20 mg L −1 .The choice of catalyst load was based on our previous work on • OH generation of Aeroxide ® TiO2 P25 [16].For zero-iron TiO2 and Fe-TiO2 materials, the catalyst load used was 320 mg L −1 , which produced a • OH generation rate under high UV irradiation to set a baseline.Catalyst load differences were attributable to the aggregation of lab-made TiO2, superficial properties, and optical properties of suspensions, as shown in Figure 18.
The photocatalytic experiments were conducted as follows.First, a pNDA test solution was set at 20 °C, the catalyst was added, and the suspension was mixed for 20 min without radiation.To evaluate the adsorption of pNDA on TiO2, an aliquot was withdrawn and centrifuged.Then, the system was fully illuminated, and aliquots were withdrawn after specific periods.Each sample was centrifuged at 6000 rpm for 15 min (Biofuge Primo, Sorvall, Hanau, Germany) and measured in the UV-visible spectrophotometer.Once the catalyst load was used and after the dark phase, no adsorption of pNDA was detected near the detection limit of UV-visible spectrophotometer.The photocatalytic experiments were conducted as follows.First, a pNDA test solution was set at 20 • C, the catalyst was added, and the suspension was mixed for 20 min without radiation.To evaluate the adsorption of pNDA on TiO 2 , an aliquot was withdrawn and centrifuged.Then, the system was fully illuminated, and aliquots were withdrawn after specific periods.Each sample was centrifuged at 6000 rpm for 15 min (Biofuge Primo, Sorvall, Hanau, Germany) and measured in the UV-visible spectrophotometer.Once the catalyst load was used and after the dark phase, no adsorption of pNDA was detected near the detection limit of UV-visible spectrophotometer.

Photolysis and Photocatalytic Degradation of E3
The initial E3 concentration was 10 µM because (1) this research was part of a project focused on the removal of E3 in water using sequentially coupled membrane filtration; (2) the solubility limit of E3 in water was previously reported to be 11.1 µM [83], and 45.1 µM [8,84], and (3) the sensitivity of the analytical techniques used in this work.The E3 solution was prepared to dissolve 2.88 mg of E3 in 1 L of deionized water by stirring at room conditions in the dark for six hours.Working solutions were stored in an amber flask.
Each photocatalytic experiment used 100 mL of E3 working solution.Initial pH was adjusted to obtain a similar surface charge of TiO 2 [85].Depending on the initial water conditions, the initial pH value was adjusted to 6.0 ± 0.1 using NaOH or HCl when needed.A dark period (no radiation) was allowed for 20 min.Then, similar experimental conditions were carried out as described in Section 3.5.Additionally, the aliquots withdrawn from suspension were filtered using a 0.1 µm syringe filter (MillexVV, Millipore, Billerica, MA, USA).A blank experiment without irradiation and TiO 2 photocatalyst was conducted for comparison.The blank experiment showed that E3 cannot be degraded in absences of either TiO 2 or UV light.Once the catalyst was loaded and after the dark phase, no adsorption of E3 was detected near the detection limit of HPLC.

Analytical Methods
The E3 concentration was monitored using an HPLC system (Waters 1515; Milford, MA, USA) equipped with a UV detector (Waters 2787) that has an injection volume of 20 µL.The analytical method was performed in isocratic analytical mode using an Inertsil ® ODS-3 column (GL Science, Tokyo, Japan; 150 mm × 4.6 mm, 5 µm) thermostated at 25 • C. The wavelength was at 280 nm according to E3 maximum absorbance.The mobile phase was methanol (49%) and deionized water (51%) at a flow rate of 1 mL min −1 .The retention time of E3 was 10 min, and the limit of E3 detection was 0.1 µM (0.029 mg L −1 ).The detection limit was obtained by developing two calibration curves: the first between 10 and 0.1 and second between 1 and 0.01.Both calibration curves followed area = 2928[E3] with R 2 = 0.9899, but areas below 0.1 were not detected.

Conclusions
This study provided an understanding of the relationship between the Fe doping ratio and radiation intensity for • OH generation and estriol (E3) degradation.The main results were that:

•
E3 degradation using 0.3 Fe-TiO 2 was feasible and can be improved by controlling irradiation intensity which was found closely related with light absorption and the catalytic reaction rate; • the synthesis method and thermal treatment allowed nanoparticles with large superficial areas and the incorporation of iron ions into the TiO 2 lattice.; and • changes in trapping recombination centers could be controlled with irradiation intensity to enhance the photocatalytic activity.
Therefore, our findings provide the opportunity to reconsider studies in which iron-doped TiO 2 impaired photocatalytic activity and to improve an application in which irradiation should be controlled.For example, Fe-TiO 2 can potentially be applied to medical uses in which low irradiation intensity should be used to avoid adverse effects in humans or wildlife, which has also been suggested by others [86].In the field of water treatment, we propose that Fe-TiO 2 is an efficient material that could harvest low-energy photons to degrade and mineralize dyes [87], biocides [88], pharmaceuticals [89], industrial chemicals [90], and estrogens-as shown in this study-to create an energetically green water treatment process.

Figure 1 .
Figure 1.Photocatalytic mechanism of TiO 2 for • OH generation.Where E g : Band gap energy; E: photon energy; OM ads : adsorbed organic molecule; and OM oxi : oxidized organic molecule.

Figure 3 .
Figure 3. High-resolution XPS spectra for the iron region for 1.0 Fe-TiO2.

Figure 3 .
Figure 3. High-resolution XPS spectra for the iron region for 1.0 Fe-TiO 2 .

Figure 7 .
Figure 7. SEM image of zero-iron TiO2 after mechanical grinding and sonication.

Figure 7 .
Figure 7. SEM image of zero-iron TiO2 after mechanical grinding and sonication.Figure 7. SEM image of zero-iron TiO 2 after mechanical grinding and sonication.

Figure 7 .
Figure 7. SEM image of zero-iron TiO2 after mechanical grinding and sonication.Figure 7. SEM image of zero-iron TiO 2 after mechanical grinding and sonication.

Figure 9 .
Figure 9. Emission spectrum and intensity graph of the irradiation source of Tecnolite fluorescent lamp (a), GE F15T8 BLB lamp (b), and GE F15T8 D lamp (c).

Figure 9 .
Figure 9. Emission spectrum and intensity graph of the irradiation source of Tecnolite fluorescent lamp (a), GE F15T8 BLB lamp (b), and GE F15T8 D lamp (c).

Figure 11 .
Figure 11.Photocatalytic mechanism of Fe-TiO2 and • OH generation.Eg is band gap energy, E is photon energy, OMads is adsorbed organic molecule, OMoxi is oxidized organic molecule.

Figure 11 .
Figure 11.Photocatalytic mechanism of Fe-TiO 2 and • OH generation.E g is band gap energy, E is photon energy, OM ads is adsorbed organic molecule, OM oxi is oxidized organic molecule.

Figure 13 .
Figure 13.Photocatalytic reaction rate (k 1 ) for degradation of E3 under high UV irradiation (a), and low UV irradiation (b); where

Figure 15 .Figure 12 .
Figure 15.Experimental relationship between pseudo first order constant and at.% content; where

Table 1 .
Surface elemental composition determined by XPS.

Table 1 .
Surface elemental composition determined by XPS.

Table 2 .
Structural and optical properties of zero-iron TiO2, and Fe-TiO2.
[23,29,64]ytic activity first increased and then decreased as the Fe concentration is similar to the behavior found with the • OH probe in Section 2.3 and has been ed using other organic molecules[23,29,64].
• , and I R is intensity at a diffraction angle 2θ of 27.46 • [75].