Photocatalytic Performance of N-Doped TiO2 Adsorbed with Fe3+ Ions

Jun 15, 2009 - Probing the Optical Property and Electronic Structure of TiO2 Nanomaterials for Renewable Energy Applications. Mukes Kapilashrami , Yan...
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Photocatalytic Performance of N-Doped TiO2 Adsorbed with Fe3+ Ions under Visible Light by a Redox Treatment Mingyang Xing, Jinlong Zhang,* and Feng Chen Key Laboratory for AdVanced Materials and Institute of Fine Chemicals, East China UniVersity of Science and Technology, 130 Meilong Road, Shanghai 200237, P. R. China ReceiVed: April 14, 2009; ReVised Manuscript ReceiVed: May 24, 2009

A simple method to prepare the N-TiO2 adsorbed with Fe3+ ions only on the surface of catalysts and modify the catalysts by a redox treatment (NaBH4 reduction and air oxidation treatment) was proposed. The samples were characterized by X-ray diffraction (XRD), UV-vis diffuse reflectance spectroscopy, FTIR, X-ray photoelectron spectroscopy (XPS), and high-resolution transmission electron micrograph (HRTEM). The photocatalytic activities of the samples were evaluated for degradation of methylene blue (MB) in aqueous solutions under visible light (λ > 420 nm). The results of XRD, FTIR, XPS, and HRTEM analysis indicated that the structure of Fe compounds changed from Fe2O3 to γ-FeOOH after redox treatment. Compared to N-TiO2 with Fe3+ ions, the catalysts after redox treatment showed higher photoactivity under visible light, and the formation of γ-FeOOH was responsible for the improvement of photocatalytic activity. Furthermore, to the catalysts after redox treatment, the mechanism for degradation of MB under visible light is an opening of a central aromatic ring attacked by OH• combined with a process of γ-FeOOH photoreductive dissolution. 1. Introduction Nanosized TiO2 as one of the most promising photocatalysts has long been investigated for photocatalytic degradation of organic pollutants, photocatalytic dissociation of water, solar energy conversion, and disinfection.1-5 However, the wide band gap of TiO2 limits the absorption wavelength less than 387 nm. Recently, a number of approaches have been undertaken to extend activity of TiO2 into the visible range.6-8 There is a growing number of publications reporting beneficial influence of substitution of oxygen in TiO2 by nonmetals such as nitrogen, carbon, and sulfur on bringing activity of the photocatalysts into the visible range.9-13 Doping with transition metals allows a tuning of the absorption properties of TiO2.14-20 Choi et al. conducted a systematic study on the photocatalytic activity of TiO2 nanoparticles doped with 21 transition-metal elements and found that doping with Fe3+, Mo5+, Ru3+, Os3+, Re5+, V4+, and Rh3+ significantly increased photoactivity in the liquid-phase photodegradation of CHCl3, while Co3+ and Al3+ doping decreased the photoactivity.21 Recently, Kisch et al. reported a new type of visible-light response photocatalyst developed by modification of platinum(IV) chloride surface complex.22-25 Among these transition-metal elements, Fe3+ was considered as an interesting dopant of TiO2 and most of the investigations have been carried out by preparation of Fe-doped TiO2 using chemical and physical methods.26-35 Ohno and co-workers found that the N-TiO2 adsorbed with Fe3+ ions shows high photocatalytic oxidation of 2-propanol due to the improvement of charge separation between electrons and holes.9 They also found that the photocatalytic activity of catalyst was further improved by reduction and air oxidation, but the main factor for this improvement was not clearly. Here, we modified the TiO2 by doping nitrogen and adsorbing ferric ions and then treated the catalysts with NaBH4 reduction and air oxidation. After the modification, the photocatalyst * To whom correspondence should be addressed. E-mail: jlzhang@ ecust.edu.cn.

shows a high photodegradation rate of methylene blue (MB) under visible light, and we investigated the reason for the improvement of visible-light photoactivity. 2. Experimental Section 2.1. Preparation of N-TiO2. The N-TiO2 was prepared by deposition procedure involving the reaction of 120 mL of H2O with 24 g of Ti(SO4)2, while stirring the solution for 2 h, followed by using ammonia to adjust the pH of the solution to 7.0. After filtration, washing, and calcination at 400 °C for 2 h, the nitrogen-doped TiO2 powder was obtained.36 A widely known material, the commercial pure anatase TiO2 (produced by Shanghai Kangyi Co. Ltd.) with specific surface area of 120 m2/g and primary particle size of 10 nm was treated in a similar way and considered as a reference. 2.2. Preparation of N-TiO2 Powders Adsorbed with Fe3+ Ions. A quantity of 2.0 g of nitrogen-doped TiO2 powder was suspended in 200 mL of FeCl3 aqueous solution with different concentrations of 0.3, 0.7, 3.6, and 7.1 mmol/L (the theoretical content of Fe3+ ions adsorbed on N-TiO2 surface is, respectively, 0.15, 0.40, 2.0, and 4.0 wt %), and the solution was stirred vigorously for 2 h. After filtration of the solution, the residue was washed with deionized water three times (it is difficult for the washing process to wash away the adsorbed Fe3+ ions via detecting the concentration of filtrates). The powders were dried under reduced pressure at 60 °C for 12 h. TiO2 was treated in a similar way and considered as a reference. 2.3. Redox Treatment (NaBH4 Reduction and Air Oxidation Treatment) of N-TiO2 Powders Adsorbed with Fe3+ Ions. A quantity of 1.0 g of N-doped TiO2 powder adsorbed with Fe3+ ions was suspended in 30 mL of deionized water. NaBH4, whose molar ratio of B/Fe is 5, was added to the solution. The solution was stirred for 2 h under aerated conditions. After filtration, the residue was washed with deionized water and then dried under reduced pressure at 60 °C for 12 h. TiO2 adsorbed with Fe3+ was treated in a similar way and considered as a reference.

10.1021/jp9034166 CCC: $40.75  2009 American Chemical Society Published on Web 06/15/2009

Performance of N-Doped TiO2 Adsorbed with Fe3+ Ions

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2.4. Measurement of OH•. A quantity of 0.02 g of catalyst was mixed with 20 mL of benzoic acid (1 mmol/L) and then was irradiated under visible light for 2 h. After filtration of the solution, the filtrate was studied by the photoluminescence (PL) emission spectra to indirectly measure the amount of OH•. 2.5. Characterization. X-ray diffraction (XRD) patterns of all samples were collected in the range 20-80° (2θ) using a Rigaku D/MAX 2550 diffractometer (Cu KR radiation, λ ) 1.5406 Å), operated at 40 kV and 100 mA. The crystallite size was estimated by applying the Scherrer equation to the full width at half-maximum (fwhm) of the (101) peak of anatase and the (110) peak of rutile, with R-silicon (99.9999%) as a standard for the instrumental line broadening. Fourier transform infrared (FTIR) spectra were recorded with KBr disks containing the powder sample with the FTIR spectrometer (Nicolet Magna 550). The instrument employed for X-ray photoelectron spectroscopy (XPS) studies was a Perkin-Elmer PHI 5000C ESCA system with Al KR radiation operated at 250 W. The shift of the binding energy due to relative surface charging was corrected using the C 1s level at 284.6 eV as an internal standard. The surface morphologies and particle sizes were observed by highresolution transmission electron microscopy (HRTEM, JEM2011), using an accelerating voltage of 200 kV. The UV-vis absorbance spectra were obtained for the dry-pressed disk samples using a scan UV-vis spectrophotometer (Varian, Cary 500) equipped with an integrating sphere assembly, using BaSO4 as the reflectance sample. The spectra were recorded at room temperature in air within the range 200-800 nm. The amount of OH• was indirectly studied by the PL emission spectra of hydroxybenzoic acid, which was measured on luminescence spectrometry (Cary Eclipse) at room temperature under the excitation light at 280 nm. The conditions were fixed as far as possible to compare the photoluminescence intensity directly. 2.6. Measurements of Photocatalytic Activities. The photocatalytic activity of each sample was evaluated in terms of the degradation of MB. MB was chosen as a model pollutant to evaluate the photocatalytic activities. The photocatalyst (0.07 g) was added into 100 mL of quartz photoreactor containing 70 mL of a 20 mg/L of MB solution. The mixture was stirred for 30 min in the dark to reach the adsorption-desorption equilibrium. A 1000-W tungsten halogen lamp equipped with a UV cutoff filters (λ > 420 nm) was used as a visible light source. The lamp was cooled with flowing water in a quartz cylindrical jacket around the lamp, and ambient temperature was maintained during the photocatalytic reaction. At the given time intervals, the analytical samples were taken from the mixture and immediately centrifuged, then filtered through a 0.22-µm Millipore filter to remove the photocatalysts. The filtrates were analyzed by recording variations in the absorption in UV-vis spectra of MB using a Cary 100 ultraviolet visible spectrometer. 3. Results and Discussion 3.1. X-ray Diffraction. The XRD patterns of the samples are shown in Figure 1. Within the detection limits of this technique, all samples present a homogeneous catalyst for the anatase crystal structure. The crystal sizes of all the samples are estimated using the Scherrer equation:37

D)

Kλ B cos θ

(1)

where B is the fwhm of the diffraction peak of anatase, K ) 0.89 is a coefficient, θ is the diffraction angle, and λ is the X-ray

Figure 1. XRD patterns of N-TiO2 and N-TiO2 adsorbed with 0.15, 0.40, and 2.0 wt % Fe3+ before and after redox treatment.

Figure 2. UV-vis diffuse reflectance spectra of N-TiO2 and N-TiO2 adsorbed with 0.15, 0.40, and 2.0 wt % Fe3+ before and after redox treatment.

wavelength corresponding to the Cu KR irradiation. The particle sizes of all the catalysts are approximately 15 nm. Before redox treatment, there is no peak assigned to Fe compound appearance, no matter the content of Fe because the amount of Fe3+ ions on N-TiO2 particle surfaces is not enough to analyze its structure by XRD technology. However, as discussed in section 3.4, the presence of Fe2O3 is attested by the XPS analysis. It is noteworthy that a new peak appears at 17.98° after the redox treatment and is assigned to γ-FeOOH.38-40 A change in crystal structure of Fe compounds loaded on the surface of N-TiO2 particles was observed before and after redox treatment. As discussed in section 3.7, the photocatalytic activities of N-TiO2 adsorbed with Fe3+ ions are related to the crystal structure of Fe compounds on the doped TiO2. That is, the crystal structure of Fe compounds on the surfaces of the catalysts is an important factor for determining the enhancement of photocatalytic activity of N-TiO2 adsorbed with Fe3+ ions under visible light. 3.2. UV-Vis Diffuse Reflectance Spectra. Figure 2 shows the absorption spectra of N-TiO2 adsorbed with Fe3+ ions before and after redox treatment. The absorbance of N-TiO2 in the visible light region increases with the increase of the amount of Fe3+ ions adsorbed on the surface of catalysts. That is because the Fe compounds on N-TiO2 are thought to show absorbance in the visible region.9 After the redox treatment, absorbance of N-TiO2 with Fe ions in the visible light region decreases slightly (Figure 2), which may be due to the slight

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Fe2O3 + NaBH4 + 3H2O f 2Fe(OH)2 V +3H2 v + NaBO2

Figure 3. FTIR spectra of N-TiO2 and N-TiO2 adsorbed with 2.0 wt % Fe3+ before and after redox treatment.

decrease of Fe after redox treatment as shown in Table 1 (the percentage of Fe changing from 1.7 to 1.4 atom %). Contrarily, Ohno and co-workers found that the absorbance of N-TiO2 with Fe3+ ions increased slightly in the visible light region after redox treatment, but they still investigated the reasons.9 3.3. FTIR. Before redox treatment, the N-TiO2 adsorbed with Fe3+ ions shows a peak at 1042.2 cm-1, which is the characteristic peak of R-Fe2O3 (Figure 3). After redox treatment, a new peak is presented at 1150.8 cm-1, which is ascribed to γ-FeOOH as shown in Figure 3.41 That is, the structure of Fe compounds loaded on the surface of N-TiO2 particles changed from R-Fe2O3 to γ-FeOOH after the redox treatment. 3.4. XPS Spectra. The XPS spectra for the Fe 2p region of N-doped TiO2 adsorbed with Fe3+ before redox treatment show two different peaks (Figure 4). Compared with normal XPS database, the peak at 709.4 eV is attributed to Fe2O3(normal energy of 709.9 eV), and another peak at 712.1 eV is attributed to Fe2(SO4)3 (normal energy of 713.3 eV). After redox treatment, the peak at 709.4 eV changes to 711.2 eV, which is attributed to the formation of FeOOH (normal energy of 711.5 eV), and another peak at 712.4 eV is almost unchanged compared with that before redox treatment. According to the XPS spectra of different elemental regions, the percentage of different elements in N-TiO2 adsorbed with 2.0 wt % Fe3+ ions is calculated in relation to the titanium as shown in Table 1. The contents of all the elements are almost unchanged after the redox treatment, and the XPS spectra of other elements such as N and S are also unchanged after redox treatment except that of Fe (the XPS spectra for N 1s region is not shown here). All the results indicate that the redox treatment visibly changes the structure of Fe compounds rather than the amount of Fe3+ ions adsorbing on the surface of N-TiO2. 3.5. HRTEM. Figure 5 shows that the HRTEM of N-TiO2 adsorbed with 2.0 wt % Fe3+ ions before and after redox treatment, the particle size of catalyst, and the distribution of Fe compounds on the surface of N-TiO2 are all uniform. After redox treatment, the interdistance of lattice planes of Fe compounds changes from 0.36 to 0.63 nm (Figure 5b). The interdistance of lattice planes of 0.63 nm is indexed to the (200) plan of γ-FeOOH. Therefore, all the above results indicate that the structure of Fe compounds adsorbing on the surface of N-TiO2 changes from Fe2O3 to γ-FeOOH after redox treatment, and the cause of γ-FeOOH formation in thermodynamics is described below:

(2)

4Fe(OH)2 + O2 f 4γ-FeOOH + 2H2O

(3)

∆Gγ-FeOOH(S) ) -112.5 kcal/mol

(4)

∆GFe(OH)2(S) ) -118.2 kcal/mol

(5)

∆GH2O(l) ) -56.7 kcal/mol

(6)

∆G295 ) 4∆Gγ-FeOOH(S) + 2∆GH2O(l) 4∆GFe(OH)2(S) - ∆GO2(g) ) -90.6 kcal/mol < 0

(7)

∆G < 0 indicates the reaction of γ-FeOOH formation is spontaneous in the process of redox treatment. 3.6. Evaluation of Photocatalytic Activity under Visible Light. Figure 6 shows the photocatalytic activities of catalysts adsorbed with Fe3+ ions before and after the redox treatment. Before redox treatment, the photocatalytic degradation rate of TiO2 is not improved by adsorbing Fe3+ ions as shown in Figure 6A,B. In contrast to this, the degradation rate of N-TiO2 improves slightly with the increment of Fe3+ ions (Figure 6C-G), and the catalysts adsorbed with 2.0 wt % Fe3+ present the highest photoactivity. After redox treatment, the degradation rate of TiO2 adsorbed with 2.0 wt % Fe3+ is almost unchanged compared with that before redox treatment (Figure 6B,b). However, after redox treatment, the photodegradation rates of N-TiO2 adsorbed with Fe3+ ions increase distinctly with the increasing of Fe3+ ions (Figure 6c-g). The catalyst adsorbed with 2.0 wt % Fe3+ ions shows better photocatalytic activity than others. Furthermore, the degradation rates of catalysts after redox treatment are much higher than those before redox treatment, as shown in Figure 6. 3.7. Mechanism for the Photocatalytic Activity under Visible Light. As we all know, the wide band gap of TiO2 limits the absorption wavelength less than 387 nm, resulting in the lower photocatalytic activity of TiO2 and TiO2 adsorbed with Fe3+ ions under visible light. However, the N-TiO2 shows a

Figure 4. XPS spectra for the Fe 2p region of N-TiO2 adsorbed with 2.0 wt % Fe3+ before and after redox treatment.

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Figure 6. Photocatalytic degradation rate of MB (20 mg/L) under visible light illumination for 2 h on TiO2, N-TiO2, and the catalysts adsorbed with Fe3+ ions before (A-G) and after (a-g) redox treatment. TiO2 (A, a), TiO2 adsorbed with 2.0 wt % Fe3+ ions (B, b), N-TiO2 (C, c), N-TiO2 adsorbed with 0.15 wt % Fe3+ ions (D, d), N-TiO2 adsorbed with 0.40 wt % Fe3+ ions (E, e), N-TiO2 adsorbed with 2.0 wt % Fe3+ ions (F, f), and N-TiO2 adsorbed with 4.0 wt % Fe3+ ions (G, g).

Figure 5. HRTEM for the N-TiO2 adsorbed with 2.0 wt % Fe3+ before (a) and after (b) redox treatment.

TABLE 1: Different Elemental Percentages of N-TiO2 Adsorbed with 2.0 wt % Fe3+ before and after Redox Treatment

Fe B N S

2.0 wt % before redox treatment (atom %)

2.0 wt % after redox treatment (atom %)

1.7 0.0 0.7 0.2

1.4 0.0 0.6 0.2

degradation rate of MB higher than that of TiO2 due to the presence of photoexcited electrons under longer visible light irradiation. By adsorbing Fe3+ ions on the surface of N-TiO2, the photoactivity is improved and as a result of the Fe3+ ions can be efficiently trapped by photoexcited electrons, resulting in improvement of the charge separation between electrons and holes.9,42 Above the amount of 2.0 wt %, the activity of N-TiO2 with Fe3+ ions gradually decreases with increase in Fe3+ ions because the large excess amount of Fe3+ ions is thought to occupy active sites on the surface of N-TiO2 and to function as a recombination center between electrons and holes.43 Kojima and co-workers found that the MB degraded by modified TiO2 under UV light is due to a process of demethylation because the methyl group can facilitate attack on MB by OH•,44,45 and the absorption peak of MB shows a distinct blueshift in UV light irradiation. However, in our investigation, the absorption peaks of MB degraded by the catalysts are all less blue-shifted under visible light irradiation (Figure 7a,b), which means the MB degraded by modified TiO2 under visible light is not a demethylation process. In the UV light irradiation, the modified TiO2 can produce a lot of photoexcited electrons to form OH•, which attack the methyl group easily; this is likely to be a major step in the photocatalytic degradation of MB under UV light.45 In our investigation, under visible light irradiation, the modified TiO2 cannot produce enough photoexcited electrons as much as that under UV light, resulting in not enough OH• being formed. When there is not enough quantity of OH•, the OH• will attack the CsS+dC functional group rather than the methyl group in MB.46 Therefore, the degradation of MB under

Figure 7. Absorption spectral changes of MB aqueous solution (10 mg/L, pH ) 7.0) degraded by N-TiO2 (a) and MB aqueous solution (20 mg/L, pH ) 7.0) degraded by N-TiO2 absorbed with 2.0 wt % Fe3+ ions after redox treatment under visible light (b).

visible light is mainly a process of the opening of a central aromatic ring rather than the demethylation of MB. Although there is a relatively small quantity of OH• produced by the catalysts in the irradiation of visible light as discussed above, OH• also plays an important role in the degradation of MB. This can be confirmed as follows. As an OH• scavenger, benzoic acid is used to scavenge OH• in the degradation of MB. Figure 8 shows the photodegradation rates of MB on N-TiO2 with 2.0 wt % Fe3+ ions after redox treatment in the presence of benzoic acid. As shown in Figure 8, the degradation rate of the catalyst with benzoic acid is obviously reduced compared with that without benzoic acid due to the decreasing of OH•, which can attack the CsS+dC functional group in MB.46 Therefore, the result indicates that OH• plays an important role in the degradation of MB under visible light. It is noteworthy that the photoactivity of N-TiO2 with Fe3+ ions after redox treatment is much higher than that before redox treatment, as shown in Figure 6, now that the OH• plays an important role in the degradation of MB, whether or not the OH• also induces the improvement of photoactivity of N-TiO2 with Fe3+ ions after redox treatment. To investigate the reason for the improvement of photocatalytic activity on the catalyst after redox treatment, we use benzoic acid to indirectly measure the amount of OH• produced by the catalyst. As we all know, benzoic acid can scavenge OH• to form hydroxybenzoic acid with fluorescence. The PL spectra of hydroxybenzoic acid can measure the amount of OH• produced by catalysts before and

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Figure 10. Photoreductive dissolution model of γ-FeOOH in the system containing organic ligand L where >FeOOH represents activated radicals of hydroxylation, L represents organic ligands, and L′ represents oxidized L.

Figure 8. Photodegradation rate of MB (20 mg/L) in the visible light illumination for 2 h on N-TiO2 absorbed with 2.0 wt % Fe3+ ions after redox treatment with and without benzoic acid in the solution of MB.

Figure 11. Absorption spectral changes of MB aqueous solution (20 mg/L, pH ) 7.0) in the presence of 1,10-phenanthroline degraded by N-TiO2 absorbed with 2.0 wt % Fe3+ ions after redox treatment under visible light irradiation for 2 h.

Figure 9. Photoluminescence spectra of benzoic acid and benzoic acid mixed with N-TiO2 adsorbed with 2.0 wt % before and after redox treatment in visible light irradiation for 2 h.

after redox treatment in visible light irradiation.47 As shown in Figure 9, pure benzoic acid does not show PL emission, and while it is mixed with catalysts in visible light irradiation, there appears a distinct PL peak. The peak intensity of N-TiO2 with 2.0 wt % Fe3+ after redox treatment is similar to that before redox treatment (i.e., the amount of OH• cannot be improved by the process of redox treatment). Therefore, the amount of OH• is not the reason for the improvement of the degradation rate of MB on the catalysts after redox treatment. The structure of Fe compounds changing from Fe2O3 to γ-FeOOH after redox treatment may induce the enhancing photoactivity of catalyst. This will be discussed below in detail. It is well-known that γ-FeOOH can produce some hydroxylactivated radicals such as >FeOH, >Fe3OH, >Fe2OH, >Fe(OH)2, and >Fe(OH)3.48,49 These activated radicals are coordinated with the organic dyes MB to form coordination compounds that induce photoreduction under visible light irradiation. Generally, the photoinduced electron transfers from ligand to Fe3+. Then the ligands induce oxidative degradation in the presence of O2. Simultaneously, the Fe3+ ions accept electrons to form Fe2+ diffusing into the solution and then coordinate with the dyes again. Afterward, these coordination compounds are oxidized by O2 to realize the circulation of photoreduction.50 A schematic diagram of the mechanism is shown in Figure 10.

In addition, the light-induced dissolution of Fe oxides also plays an important role in the reaction of photoreduction.50 To Fe2O3, the hydroxyl-activated radicals on the surface of catalysts also can coordinate with the organic dyes and then induce the transfer of electron from ligand to Fe3+. But the produced Fe2+ ions are oxidized by O2 rather than diffusing into the solution and are against the circulation of photoreduction.51 Compared with that of Fe2O3, γ-FeOOH has lower Madelung energy, resulting in the Fe2+ dissolution of γ-FeOOH to be better than that of Fe2O3.52 To prove the presence of dissolution of Fe2+, a 1,10-phenanthroline method with a detection limit of 1.8 × 10-7 mol/L for Fe2+ was used to measure the presence of Fe2+ ions in the solution.53 The 1,10-phenanthroline was added in the reaction solution. After 2 h visible light irradiation, a new peak appeared at 511 nm in the absorption spectral changes of MB (as shown in Figure 11), which is the characteristic peak of the coordination compound produced by Fe2+ and 1,10-phenanthroline.53 Consequently, to the catalysts after redox treatment, there is a better Fe2+ ions dissolution presence in the visible light irradiation. Anyhow, the reason for N-TiO2 with Fe3+ ions after redox treatment showing higher photoactivity under visible light than that before redox treatment is not only the photoreductive ability but also the better light-induced dissolution of γ-FeOOH than Fe2O3. 4. Conclusions N-TiO2 adsorbed with Fe3+ ions was successfully prepared by a simple method. After redox treatment (NaBH4 reduction and air oxidation), the structure of Fe compounds adsorbed on the surface of N-TiO2 changes from Fe2O3 to γ-FeOOH, and

Performance of N-Doped TiO2 Adsorbed with Fe3+ Ions the catalysts show a much higher photodegradation rate of MB than that without redox treatment under visible light. Furthermore, to N-TiO2 with Fe3+ ions after redox treatment, the mechanism for degradation of MB under visible light is not only an opening of central aromatic ring attacked by OH•, but also a process of γ-FeOOH photoreductive dissolution. The latter induces the improvement of photocatalytic activity of catalysts after redox treatment. In addition, the catalyst of N-TiO2 with 2.0 wt % Fe3+ ions after redox treatment shows high photoactivity in visible light irradiation. Acknowledgment. This work was supported by the Science andTechnologyCommissionofShanghaiMunicipality(07JC14015), Shanghai Nanotechnology Promotion Centre (0752nm001), National Nature Science Foundation of China (20773039), National Basic Research Program of China (973 Program, 2007CB613301, 2004CB719500), and the Ministry of Science and Technology of China (2006AA06Z379, 2006DFA52710). References and Notes (1) Fujishima, A.; Rao, T. N.; Tryk, D. A. J. Photochem. Photobiol., C 2000, 1, 1. (2) Wolfrim, E. J.; Huang, J.; Blake, D. M.; Maness, P.; Huang, Z.; Fiest, J. EnViron. Sci. Technol. 2002, 36, 3412. (3) Hu, C.; Yu, J. C.; Hao, Z.; Wong, P. K. Appl. Catal., B 2003, 42, 47. (4) Hoffmann, M. R.; Martin, S. T.; Choi, W.; Bahnemann, D. W. Chem. ReV. 1995, 95, 69. (5) Yu, J. C.; Ho, W.; Yu, J.; Yip, H.; Wong, P. K.; Zhao, J. EnViron. Sci. Technol. 2005, 39, 1175. (6) In, S.; Orlov, A.; Berg, R.; Garcia, F.; Pedrosa-Jimenez, S.; Tikhov, M. S.; Wright, D. S.; Lambert, R. M. J. Am. Chem. Soc. 2007, 129, 13790. (7) Kisch, H.; Sakthivel, S.; Janczarek, M.; Mitoraj, D. J. Phys. Chem. C 2007, 111, 11445. (8) Chen, D. M.; Jiang, Z. Y.; Geng, J. Q.; Wang, Q.; Yang, D. Ind. Eng. Chem. Res. 2007, 46, 2741. (9) Ohno, T.; Miyamoto, Z.; Nishijima, K.; Kanemitsu, H.; Feng, X. Y. Appl. Catal., A 2006, 302, 62. (10) Ohno, T.; Mitsui, T.; Matsumura, M. Chem. Lett. 2003, 32, 364. (11) Li, D.; Haneda, H.; Hishita, H.; Ohashi, N. Chem. Mater. 2005, 17, 2588. (12) Umebayashi, T.; Yamaki, T. H. Appl. Phys. Lett. 2002, 81, 454. (13) Cong, Y.; Zhang, J. L.; Chen, F.; Anpo, M.; He, D. J. Phys. Chem. C 2007, 111, 10618. (14) Kiwi, J.; Graetzel, M. J. Phys. Chem. 1986, 90, 637. (15) Karakitsou, K. E.; Verykios, X. E. J. Phys. Chem. 1993, 97, 1184. (16) Zhu, J. F.; Chen, F.; Zhang, J. L.; Chen, H. J.; Anpo, M. J. Mol. Catal. A: Chem. 2004, 216, 35. (17) Zhu, J. F.; Deng, Z. G.; Chen, F.; Zhang, J. L.; Chen, H. J.; Anpo, M.; Huang, J. Z. Appl. Catal., B 2006, 62, 329. (18) Nagaveni, K.; Hegde, M. S.; Madras, G. J. Phys. Chem. B 2004, 108, 20204. (19) Klosek, S.; Raftery, D. J. Phys. Chem. B 2001, 105, 2815. (20) Kisch, H.; Zang, L.; Lange, C.; Maier, W. F.; Antonius, C.; Meissner, D. Angew. Chem., Int. Ed. 1998, 37, 3034.

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