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Synthesis, Characterization, and Activity of Visible-Light-Driven Nitrogen-Doped TiO2-SiO2 Mixed Oxide Photocatalysts Thiam Peng Ang,* Choon Sian Toh, and Yi-Fan Han Institute of Chemical and Engineering Sciences, 1 Pesek Road, Jurong Island, Singapore 627833 ReceiVed: January 4, 2009; ReVised Manuscript ReceiVed: April 2, 2009

A series of nitrogen-doped TiO2-SiO2 mixed oxide catalysts were prepared by a sol-gel method through varying the TiO2/SiO2 molar ratio. The catalyst structure has been extensively characterized by using UV-visible spectroscopy, X-ray photoelectron spectroscopy (XPS), X-ray diffraction (XRD), transmission electron microscopy (TEM), thermogravimetry analysis (TGA), and Fourier transform infrared spectroscopy (FT-IR). UV-visible spectroscopy revealed that the absorption edge of the catalysts composed of only the anatase phase was red-shifted to the visible light region. In addition, the presence of SiO2 stabilized the anatase phase even at high temperatures. The results of XRD and TEM revealed a decrease in the size of TiO2 particles with increasing the SiO2 content, but the rise of SiO2/TiO2 ratio in the mixed oxides led to a decrease in the amount of surface sulfate species as identified by the FT-IR, XPS, and TGA. Moreover, XPS results also suggested that the TiO2 lattice was interstitially doped by nitrogen species. The catalysts as prepared exhibited better photocatalytic activity in degradation of methylene blue under visible light compared to the doped TiO2 catalysts. The addition of SiO2 results in (i) the improvement of thermal stability of the catalyst, (ii) the decrease of TiO2 particle size, and (iii) the formation of Brønsted acid sites. All these factors are responsible for an enhancement in photogenerated electron-hole separation and thus better photocatalytic activity. Introduction TiO2 is a widely studied photocatalyst because of its low cost, ultrastability, and nontoxicity.1-8 However, TiO2 can only be activated under UV light since it possesses a large band gap (about 3.0-3.2 eV). This is not favorable as UV only makes up to about 3% of the solar energy.9 To maximize utilization of the solar energy spectrum, activation of TiO2 under visible light is essential for application.10 Among all methods for preparing visible-light-driven TiO2, doping TiO2 with transition metal cations has been widely investigated.11-13 However, these cations often serve as recombination centers for photogenerated electrons and holes, thus resulting in poorer photocatalytic activity.14-17 Furthermore, these cation-doped TiO2 also suffered from thermal instability.16,18 Recently, there has been strong interest on doping TiO2 with anions such as N,14,16,19-27 S,9,15,28-31 B,32 P,33 C,34-36 and halogens.37-40 N was claimed to be the best dopant because (i) the band gap of TiO2 is narrowed from the mixing of N 2p state with O 2p state and (ii) the ionic sizes between N and O are similar thus making easy incorporation of N into TiO2 lattice.14 As such, the study of N as the main dopant for TiO2 activated under the visible light has attracted more attention. On the other hand, TiO2-SiO2 mixed oxides have been found to be active photocatalysts.41-49 Yu et al. suggested that the high activity is due to smaller TiO2 particle size and high content of surface hydroxyl.49 In addition, TiO2 phase transformation from anatase to rutile that shows relative lower photoactivity can also be retarded due to the improved thermal stability of the mixed oxides. Despite these advantages, it seems that most of the TiO2-SiO2 mixed oxides can only be activated under UV light thus limiting their usage in the solar energy. * Corresponding author. Tel./Fax: +65-67963832/+65-63166182. E-mail address: [email protected].

In this paper, we present a study on preparing TiO2-SiO2 mixed oxides that can be activated under the visible light after anionic doping using the sol-gel method. The TiO2/SiO2 ratio was varied to evaluate its impacts on the photocatalytic degradation of methylene blue under visible light. While we have found evidence that showed the TiO2 lattice was doped with nitrogen instead of sulfur (from thiourea), IR and XPS analysis pointed out that the latter exists as sulfate species. In addition, our prepared samples are thermally stable and have shown great photocatalytic activities when compared to TiO2 sample synthesized without SiO2. Experimental Section Sample Preparation. All chemicals were used as received and without any purification. TiO2-SiO2 mixed oxides were first prepared by adding titanium isopropoxide (Sigma Aldrich, 97%) and tetraethyl orthosilicate (Sigma Aldrich, 98%) into 20 mL of isopropanol. Amounts of 20 mL of 0.05 M aqueous ammonium carbonate and ca. 2.1 g of thiourea (Sigma Aldrich, g99%) were then added. White precipitate was formed instantaneously and was stirred for 2 h before aging overnight at room temperature. The precipitate was then rotovaped to form a slurry that will be dried in a 60 °C oven overnight. The sample was finally calcined in a muffle furnace at 500 °C (heating rate ) 1.5 °C/min) for 5 h. A series of mixed oxides was prepared in this study using different titanium isopropoxide to tetraethyl orthosilicate feed ratios. On the other hand, a similar method was used to prepare doped TiO2 samples by omitting tetraethyl orthosilicate. To facilitate our discussion, all prepared TiO2-SiO2 mixed oxides were abbreviated as Ti/Si (X) where X represents the tetraethyl orthosilicate molar feed ratio with respect to the fixed amount of titanium isopropoxide added. For consistency, doped TiO2 sample was quoted as Ti/Si(0).

10.1021/jp9000658 CCC: $40.75  2009 American Chemical Society Published on Web 05/19/2009

Mixed Oxide Photocatalysts

Figure 1. Representative XPS N 1s and S 2p peaks of the prepared mixed oxides as represented by Ti/Si (0.075).

Characterization. The catalyst surface was analyzed using X-ray photoelectron spectroscopy (XPS), a VG Scientific ESCALAB 250 where a monochromatic Al KR X-ray source at 1486.7 eV was used. The binding energies (BE) of all elements were calibrated using a C 1s peak of adventitious C at 284.8 eV. XPS spectra were recorded at θ ) 90° for X-ray source. The crystal phase of the catalysts was analyzed using an X-ray diffractometer (Bruker D8 Advance) operated at a voltage of 35 kV and a current of 40 mA. The diffraction pattern was taken in the Bragg’s angle (2θ) range from 1.5 to 90° at room temperature. FT-IR spectra were determined on KBr disks using a FT-IR spectrophotometer (Biorad Excalibur Series FTS 3000MX) while solid-state UV-vis spectra of the catalysts were recorded in BaSO4 using a Shimadzu UV-2550 with slit width of 1 nm. The specific surface area (BET) was measured from the adsorption and desorption N2 isotherms collected from a Quantachrome Autosorb-6B surface area & pore size analyzer at 77 K. Before the measurement, the samples were degassed at 200 °C overnight. The size and morphology of the catalysts were determined by transmission electron microscopy (TEM). Small amounts of samples were first sonicated in ethanol, and a drop of suspension was then dropped and dried over a copper grid. Electron micrographs were taken with a Tecnai, TF20 Super Twin transmission electron microscope at an accelerating voltage of 200 kV. The percentage weight loss of our photocatalysts was determined using a Q500 TGA analyzer (TA Instruments) under air (flow rate ) 60 mL/min) at a heating rate of 20 °C/min. Photocatalytic Activity Measurement. The photocatalytic activities of the samples were measured by the degradation of aqueous methylene blue under visible light. About 50 mg of the sample was first added to a reactor that contained 2 × 10-5 M of aqueous methylene blue. The mixture was then stirred

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Figure 2. (a) XPS S 2p peak of Ti/Si (0.075) after washing with deionized water. (b) XPS N 1s peak of Ti/Si (0.075) before and after washing with deionized water.

TABLE 1: Si 2p and Ti 2p Binding Energies of the Prepared Mixed Oxides with Si/Ti and S/N Ratios Measured from XPSa

sample ID Ti/Si (0) Ti/Si (0.01) Ti/Si (0.025) Ti/Si (0.05) Ti/Si (0.075) Ti/Si (0.10) SiO2 Ti/Si (0.075) washed a

binding binding specific energy energy of Si 2p of Ti 2p Si/Ti S/N surface area (m2/g) (eV) (eV) ratio ratio N.A. 102.3 102.5 102.6 102.5 102.5 103.2 102.3

458.8 458.6 459.4 458.9 458.5 458.6 N.A. 458.5

N.A. 0.02 0.06 0.10 0.13 0.16 N.A. 0.10

2.06 2.07 1.33 1.28 1.17 1.21 N.A. 0.37

105 118 141 133 155 163 N.A. N.A.

Specific surface areas were also included for comparison.

for 1 h in the dark. This is followed by irradiation of visible light obtained from a 300 W Xe lamp with a cutoff filter (λ g 420 nm). The concentration of methylene blue was monitored on the basis of its UV-visible absorption peak at 664 nm. The results were compared with the results obtained from the doped TiO2 sample prepared using the same method. To test the activity of the reused sample, the samples were first retrieved by centrifugation after the first run of activity. The retrieved samples were then dried in the oven overnight before running the methylene blue degradation in the visible light as described previously. Results and Discussion Figure 1 presents the typical XPS N 1s and S 2p peak of the sample Ti/Si (0.075). It can be observed that two N 1s peaks

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Figure 3. (a) FTIR spectra of (A) Ti/Si (0), (B) Ti/Si (0.01), (C) Ti/Si (0.025), (D) Ti/Si (0.05), (E) Ti/Si (0.075), and (F) Ti/Si (0.10). (b) Zoom-in spectra of (A) Ti/Si (0), (B) Ti/Si (0.01), (C) Ti/Si (0.025), (D) Ti/Si (0.05), (E) Ti/Si (0.075), and (F) Ti/Si (0.10) from 750 to 1500 cm-1. (c) FT-IR spectra of Ti/Si (0.10) and Ti/Si (0.40) where a Ti-O-Si peak at ca. 935 cm-1 was observed in the latter. (d) Zoom-in view of the unknown peak for (A) Ti/Si (0), (B) Ti/Si (0.01), (C) Ti/Si (0.025), (D) Ti/Si (0.05), (E) Ti/Si (0.075), and (F) Ti/Si (0.10). (e) FT-IR spectra of the unknown peaks for samples prepared under different conditions.

were present at ca. 400 and 402 eV, respectively, while a sulfur peak with binding energy of ca. 169 eV was detected. The position of the latter peak pointed to the presence of S6+ that might imply the presence of sulfate species on the surface.15,30 In order to verify the presence of sulfate species, Ti/Si (0.075) was washed thoroughly with deionized water before drying overnight in an oven. Indeed, a very small S 2p peak was detected for the washed sample (Figure 2a). On the other hand, assignment of a N 1s peak was shown to be more contradictory in the previous studies.50 Nonetheless, a recent study reported by Asahi et al. have conclusively pointed out that the N 1s peak at ca. 400 eV can be assigned to NO species occupying in the interstitial TiO2 site while the peak at ca. 402 eV is attributed to weakly physisorbed nitrogen species.51 It is noteworthy that the intensity of N 1s peak increases slightly after washing with

deionized water as shown in Figure 2b thus verifying the presence of surface sulfate species on the surface. Ti 2p and Si 2p peaks were also detected, and their positions are listed in Table 1. The corresponding Ti 2p position for Ti/Si (0) was included for comparison. There is almost no change in Ti 2p peak position between the mixed oxides and Ti/Si (0) since the Ti-O-Ti bond should be predominant in TiO2-SiO2 mixed oxides. XPS results indicate there is a layer of sulfate species on the mixed oxide surfaces that can be removed by washing with deionized water. There is no substitution of O by S in the TiO2 lattice during preparation because the Ti-S bond was not detected after washing. We have also noticed from Table 1 that the sample surfaces were enriched by SiO2, since TEOS is



Figure 4. (a) XRD diffractograms of the (A) Ti/Si (0), (B) Ti/Si (0.01), (C) Ti/Si (0.025), (D) Ti/Si (0.05), (E) Ti/Si (0.075), and (F) Ti/Si (0.10). (b) Zoom-in view of anatase peak at 2θ ∼ 25° for (A) Ti/Si (0), (B) Ti/Si (0.01), (C) Ti/Si (0.025), (D) Ti/Si (0.05), (E) Ti/Si (0.075), and (F) Ti/Si (0.10) with the average particle size calculated by Scherrer equation indicated. (c) Comparison of anatase peak at 2θ ∼ 25° of Ti/Si (0.075), Ti/Si (0.075) without thiourea and Ti/Si (0) with the average particle size calculated by the Scherrer equation indicated. (d) XRD diffractograms of Ti/Si (0.075) calcined at (A) 500 °C, (B) 600 °C, (C) 700 °C, (D) 800 °C, and (E) 1000 °C and (F) Ti/Si (0) calcined at 800 °C. (e) Comparison of XRD diffractograms of Ti/Si (0.075) and Ti/Si (0) after calcination at 800 °C.

known to be hydrolyzed and condensed much slower than titanium isopropoxide. The FT-IR spectra (Figure 3a) of the mixed oxides and Ti/Si (0) indicate that TiO2 structure is not altered by SiO2. In addition, no peak attributed to the Ti-O-Si bond (ca. 910-960 cm-1) was detected for all samples (Figure 3b), suggesting that either the TiO2 and SiO2 are physically mixed in oxides or the amount of Ti-O-Si bond is too small to register a peak. The FT-IR spectrum recorded from a sample with a higher SiO2 content [Ti/Si (0.40)] (Figure 3c) implies the formation of a Ti-O-Si bond in all mixed oxides. On the other hand, the peaks at ca. 1050 and 1125 cm-1 were observed (Figure 3b) for all samples as well as Ti/Si (0). In line with the study on a thiourea-modified

SnO2 powder, the presence of these two peaks after sintering at 600 °C was attributed to a sulfate group band.52 Thus, the FT-IR analysis agrees well with the XPS results. In addition, it has also been pointed out that the peak at ca. 1050 cm-1 might result from the active sulfate species while the peak at ca. 1125 cm-1 results from the inactive sulfate species.53 Since the peak intensity at ca. 1125 cm-1 decreases with increasing the SiO2 content, this means the amount of inactive sulfate species that can reduce the photocatalytic activity decreases in the same order. This observation is further confirmed by the decreasing XPS sulfur/nitrogen (S/N) ratio (Table 1) with increasing SiO2/ TiO2 molar ratio.

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Figure 5. (a) UV-visible spectra of (A) Ti/Si (0), (B) Ti/Si (0.01), (C) Ti/Si (0.025), (D) Ti/Si (0.05), (E) Ti/Si (0.075), and (F) Ti/Si (0.10). (b) Zoom-in spectra of (A) Ti/Si (0), (B) Ti/Si (0.01), (C) Ti/Si (0.025), (D) Ti/Si (0.05), (E) Ti/Si (0.075), and (F) Ti/Si (0.10) from 450 to 800 nm.

Noticeably, two weak peaks at ca. 2050 and 2065 cm-1 (Figure 3d) were detected for all samples, and their intensities seem to decrease with increasing the SiO2 content. In addition, the peak at ca. 2065 cm-1 almost disappeared for samples with the addition of over 5 mol % of Si precursor. In order to understand the origin of these peaks, mixed oxides were synthesized (i) without the use of thiourea, (ii) without the use of ammonium carbonate, and (iii) replacing thiourea with urea. The FT-IR spectra in Figure 3e showed the appearance of the unknown peaks for all samples except preparation without thiourea, suggesting that these peaks originated from thiourea. On the other hand, the prepared sample with urea only generates a peak at ca. 2050 cm-1. By comparison of the molecular structure of thiourea and urea, it seems that nitrogen from both molecules might link to the peak at 2050 cm-1. In a recent study, Belver et al. reported the presence of a peak at ca. 2070 cm-1 attributed to NO+ species that was suggested to occupy the interstitial position.54 Therefore, the unknown peak at 2050 cm-1 could be attributed to interstitial NO+ species that was previously identified by an XPS N 1s peak. In addition, we assigned the peak at ca. 2065 cm-1 to the NC stretch of HSCN species that might be produced in the following reactions55

It is noted that no HSCN species are observed when SiO2 is greater than 5 mol % in the mixed oxides.

Ang et al. The XRD diffractograms in Figure 4a showed that the mixed oxides were composed of only anatase phase while a peak broadening was observed with decreasing the TiO2/SiO2 molar ratio (Figure 4b). It is evident that the particle size of TiO2 anatase decreases with increasing the SiO2 content (from 9.1 nm in Ti/Si (0) to 6.8 nm in Ti/Si (0.1)), and this agrees well with the increase in surface area as observed in Table 1. On the other hand, the crystallization of TiO2 was reported to be inhibited by sulfate.44 To investigate the effects of sulfate species and SiO2 on the growth of TiO2 particle, we have prepared one control (Ti/Si (0.075)) sample without thiourea and the second one without TEOS. XRD patterns in Figure 4c showed the peak width changed little for the Ti/Si (0.075) sample prepared with or without thiourea where the particle size for both samples is ∼8.0 nm. In comparison, a significant decrease in the peak width was observed for the second control (particle size was measured as ∼9.1 nm). We can, therefore, conclude that SiO2 played an important role in inhibiting the crystallinity of TiO2 in the mixed oxides. Furthermore, the XRD patterns for the Ti/Si (0.075) sample calcined at different temperatures (Figure 4d) demonstrated that SiO2, indeed, is capable of preventing anatase phase from transforming to rutile phase, in comparison with the Ti/Si (0) calcined at 800 °C (Figure 4e). The UV-visible spectra in Figure 5a show an absorption edge (onset) situated at ca. 530-600 nm (estimated band gap energy ) 2.07-2.34 eV). The presence of interstitial N-doping has been previously deduced on the basis of XPS and FT-IR spectra. This N-doping could form discrete energy levels above the valence band of TiO2, and transition between these energy levels and the conduction band of TiO2 might result in absorption in visible light.14 Nonetheless, the number of these discrete energy levels is limited, resulting in the low absorption intensity in the visible light range. Comparing the spectra in Figure 5b, we noticed that the onset that located in the visible light seems to blue shift with increasing the SiO2 content. In a recent study, Kim et al. have found that an increase in the band gap of TiO2 supported on SiO2 is due to quantum size effect.45 Likewise, the particle size of TiO2 was reduced with a decrease in TiO2/SiO2 ratio. The corresponding TiO2 band gap in our samples is expected to increase which, in turn, will widen the gap between the discrete energy levels and the conduction band. Eventually, this led to a blue shift of the onset as observed in the visible light range (Figure 5b). TEM images in Figure 6 show the nanometer-sized TiO2 lattice fringes for all samples. In particular, a decrease in the nanocrystal size with rise in the SiO2 content agrees well with the XRD results, owing to the retardation of TiO2 particles growth inhibited by SiO2. Anderson and Bard have reported that the Ti-O species are locked at the interface by the SiO2 lattice, and Ti atoms are possibly substituted into the tetrahedral SiO2 lattice to form tetrahedral Ti sites at the interface of TiOSi.41 The latter interacts with the existing octahedral Ti sites in anatase, thus preventing the formation of a rutile phase. The photocatalytic activity of all catalysts was evaluated by measuring the absorbance (At) of methylene blue UV-visible spectrum at 664 nm during every hourly interval. With the assumption that Beer’s law was obeyed, the graph of At/Ao against t (where Ao is the intensity of methylene blue peak at 664 nm after stirring in the dark) is equivalent to the graph of Ct/Co against t and the latter was plotted as shown in Figure 7a. It can be seen that the mixed oxide catalysts outperformed the reference Ti/Si (0). On the other hand, it can also be clearly observed in Figure 7b that the activity of Ti/Si (0.05) is much



Figure 6. Typical TEM micrographs of (A) Ti/Si (0.01), (B) Ti/Si (0.025), (C) Ti/Si (0.05), (D) Ti/Si (0.075), and (E) Ti/Si (0.10).

better than photolysis and activity of Ti/Si (0.05) operated without light. As such, we can effectively rule out the contribution of photolysis and adsorption to the overall photocatalytic activity. In addition, it is noteworthy that Ti/Si (0.05) (Figure 7c) suffers a 20-30% drop in photocatalytic activity after the first run of reaction. Clearly, the photocatalytic property of TiO2-SiO2 mixed oxides is strongly affected by the addition of SiO2 on, at least, three aspects. First, as evidenced by XRD analysis, the thermal stability of TiO2-SiO2 was significantly improved by the addition of SiO2, and the anatase phase remained unchanged, even calcined at 800 °C. In this way, the photocatalytic capability of the mixed oxides was not compromised even if the samples were calcined at high temperatures since anatase is generally known to be more photocatalytically active than rutile.56 Moreover, calcination at higher temperatures promotes TiO2 crystallinity and help to reduce bulk defects,44,46 which often serve as recombination centers for photogenerated electrons and holes and thus lower the activity of the catalyst. Second, the addition of SiO2 leads to a decrease in the TiO2 particle size. Generally, it is advantageous for TiO2 to be present in very small sizes since it is easier for the photogenerated electrons and holes to arrive at the reaction site.46 Third, it is

well-known that Brønsted acidity is generated with the addition of SiO2.46 The generation of Brønsted acidity often induces an increase in the amount of surface hydroxyl groups, which could trap the photogenerated holes and form OH radicals to carry out oxidation. As a result, the photogenerated electron-hole recombination is retarded and the photocatalytic activity of the mixed oxides is enhanced. In addition, it is believed that the presence of surface hydroxyl groups could improve the adsorption of methylene blue to the surface of the mixed oxide near to TiO2.41 In this manner, methylene blue was more effectively degraded by TiO2, and that might explain the difference in photocatalytic methylene blue rate constant between the mixed oxides and Ti/Si (0). It was also found that the methylene blue degradation constant increases with increasing the SiO2 content. It reaches a maximum value for Ti/Si (0.075) and decreases thereafter. In principle, we would expect oxide with higher SiO2 content to result in a higher degradation rate constant because of smaller TiO2 particle size and higher surface area. In addition, Table 1 shows that the amount of SiO2 in the surface layers also increased with decreasing TiO2/SiO2 ratio. As such, it seemed that the mixed oxides should have an increasing methylene blue adsorption and better photogenerated electron-hole separation

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Figure 8. Thermogravimetric graphs of (A) Ti/Si (0.01), (B) Ti/Si (0.025), (C) Ti/Si (0.05), (D) Ti/Si (0.075), and (E) Ti/Si (0.10) in the temperature range from 550 to 900 °C with the weight change indicated.

Figure 7. (a) Plot of Ct/Co vs time for the mixed oxides at different SiO2/TiO2 ratio and washed Ti/Si (0.075). (b) Plot of Ct/Co vs time for Ti/Si (0.05), photolysis of methylene blue and Ti/Si (0.05) operated without light. (c) Plot of normalized In(Co/Ct) vs time for Ti/Si (0.05) in the first and second run of activity.

at the lower TiO2/SiO2 ratios. This would result in more efficient photocatalytic methylene blue degradation. Nonetheless, a drop in methylene blue degradation activity was observed with increasing the SiO2 content from 7.5 to 10 mol %. In fact, among all samples, TiO2 is the least crystallized in Ti/Si (0.10) despite having the smallest size. In addition, the amount of SiO2 present on the surface of the photocatalyst is expected to increase with SiO2 content. Since SiO2 is photocatalytically inactive, the number of surface-active sites will decrease correspondingly and thus compromise on the activity of Ti/Si (0.10) with respect to Ti/Si (0.075). The sulfate species may also involve in the improvement of the catalyst property. Thermogravimetry analysis was performed

for determining the amount of sulfate in the prepared samples. As shown in Figure 8, one transition was observed in the temperature range 500-900 °C, owing to the desorption of sulfate species.57,58 It can be seen that the amount of sulfate increases with the TiO2/SiO2 ratio, because sulfate species could be significantly decomposed on SiO2 sites above 250 °C.59 Our FT-IR and XPS results seem to agree with this observation. We demonstrated the effects of sulfate species on the photocatalytic activity of Ti/Si (0.075) through the photodegradation of the methylene blue. Figure 7a clearly showed that the methylene blue photodegraded more slowly on a sulfatefree Ti/Si (0.075) (removed by washing) relative to Ti/Si (0.075). Xie et al. has suggested that the surface sulfate species will retain the photoinduced electrons leaving the photogenerated holes on TiO2;59 thus the photogenerated electron-hole recombination will be retarded. This explains well the improvement of photocatalytic performance with the presence of sulfate species. On the contrary, the photocatalytic performance was enhanced from Ti/Si (0) to Ti/Si (0.05) even though the amount of sulfate species decreases. Thus, it seems that other factors such as surface SiO2 content, surface area, and TiO2 particle size outweigh this factor in determining the catalyst performance in methylene blue photodegradation. Only between Ti/Si (0.075) and Ti/Si (0.10), the difference in the amount of sulfate species could then bring about a noticeable difference in methylene blue photodegradation rate constant. On the other hand, it was speculatedthatsomesulfateleachingandstrongcatalyst-methylene blue interaction might have taken place that leads to the 20-30% photocatalytic activity drop after the first run of reaction. Conclusions A series of TiO2-SiO2 mixed oxides have been prepared by using a sol-gel method. It was observed that the mixed oxides were interstitially doped by N atoms that could probably lead to absorption in visible light. The methylene blue degradation rate constant initially increases with decreasing TiO2/SiO2 molar ratio and reaches the maximum at Ti/Si (0.075) before it decreases. The photoactivity was initially affected by factors such as the amount of surface SiO2, TiO2 particle size, and surface area. Subsequently, the decreasing amount of surface sulfate species with increase in SiO2/TiO2 ratio as well as the increasing amount of photocatalytically inactive SiO2 on the surface outweighed all other factors that led to the decrease in

Mixed Oxide Photocatalysts rate constant. In general, the addition of SiO2 into TiO2 can bring about (i) the improvement of thermal stability of the catalyst, (ii) the decrease of TiO2 particle size, and (iii) the formation of Brønsted acid sites. All these factors lead to an enhancement in photogenerated electron-hole separation and thus better photocatalytic activity. Acknowledgment. The research was supported by ICES and Agency for Science, Technology and Research in Singapore. The authors thank Professor J. Y. Lin, Dr. P. K. Wong, and Dr. Armando Borgna for their great support and advice. The authors would also like to thank Ms. Wang Zhan for the XPS analysis. References and Notes (1) Yamashita, H.; Takeuchi, M.; Anpo, M. In Encyclopedia of Nanoscience and Nanotechnology; Nalwa, H. S., Ed.; American Scientific Publishers: Stevenson Ranch, CA, 2004; Vol. 10, p 639. (2) Hoffmann, M. R.; Martin, S. T.; Choi, W. Y.; Bahnemann, D. W. Chem. ReV. 1995, 95, 69. (3) Herrmann, J.-M. Catal. Today 1999, 53, 115. (4) Anpo, M. Pure Appl. Chem. 2000, 7, 1265. (5) Fujishima, A.; Rao, T. N.; Tryk, D. A. J. Photochem. Phototbiol., C 2000, 1, 1. (6) Mills, A.; Lee, S. K. J. Photochem. Phototbiol., A 2002, 152, 233. (7) Anpo, M.; Takeuchi, M. J. Catal. 2003, 216, 505. (8) Anpo, M. Bull. Chem. Soc. Jpn. 2004, 77, 1427. (9) Yu, J. C.; Ho, W. K.; Yu, J. G.; Yip, H. Y.; Wong, P. K.; Zhao, J. C. EnViron. Sci. Technol. 2005, 39, 1175. (10) Sun, H. Q.; Bai, Y.; Cheng, Y. P.; Jin, W. Q.; Xu, N. P. Ind. Eng. Chem. Res. 2006, 45, 4971. (11) Anpo, M. Catal. SurV. Jpn. 1997, 1, 169. (12) Ohno, T.; Tanigawa, F.; Fujihara, K.; Izumi, S.; Matsumura, M. J. Photochem. Phototbiol., A 1999, 127, 107. (13) Dana, D.; Vlasta, B. Appl. Catal. A: Gen. 2001, 208, 335. (14) Asahi, R.; Morikawa, T.; Ohwaki, T.; Aoki, K.; Taga, Y. Science 2001, 293, 269. (15) Ohno, T.; Mitsui, T.; Matsumura, M. Chem. Lett. 2003, 32, 364. (16) Chen, S. F.; Chen, L.; Gao, S.; Cao, G. Y. Chem. Phys. Lett. 2005, 413, 404. (17) Sathish, M.; Viswanathan, B.; Viswanath, R. P.; Gopinath, C. S. Chem. Mater. 2005, 17, 6349. (18) Kudo, A. Catal. SurV. Jpn. 2003, 7, 31. (19) Yin, S.; Yamaki, H.; Komatsu, M.; Zhang, Q. W.; Wang, J. S.; Tang, Q.; Saito, F.; Sato, T. J. Mater. Chem. 2003, 13, 2996. (20) Yin, S.; Yamaki, H.; Komatsu, M.; Zhang, Q. W.; Wang, J. S.; Tang, Q.; Saito, F.; Sato, T. Solid State Sci. 2005, 7, 1479. (21) B Chen, X.; Burda, C. J. Phys. Chem. B 2004, 108, 15446. (22) Rhee, C. H.; Bae, S. W.; Lee, J. S. Chem. Lett. 2005, 34, 660. (23) Burda, C.; Lou, Y. B.; Chen, X. B.; Samia, A. C. S.; Stout, J.; Gole, J. L. Nano Lett. 2003, 3, 1049. (24) Li, D.; Haneda, H.; Hishita, S.; Naoki, Ohashi Mater. Sci. Eng., B 2005, 117, 67. (25) Livraghi, S.; Votta, A.; Paganini, M. C.; Giamello, E. Chem. Commun. 2005, 498.

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