Low-Temperature Preparation and Characterization of Nanocrystalline

J. Phys. Chem. C 2009, 113, 4031–4037

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Low-Temperature Preparation and Characterization of Nanocrystalline Anatase TiO2 Sifang Li,* Guoliang Ye, and Guoqin Chen Department of Chemical and Biochemical Engineering, College of Chemistry and Chemical Engineering, Xiamen UniVersity, Xiamen 361005, China ReceiVed: August 29, 2008; ReVised Manuscript ReceiVed: January 16, 2009

Nanoanatase TiO2 of high crystallinity was prepared by a novel simple route at a temperature of 100 °C under mild conditions. Tetrabutyl titanate was used as a titanium precursor, Acetic acid was used as an inhibitor, and diethyl ether anhydrous was used as a solvent. X-ray powder diffraction (XRD), laser Raman spectroscopy (Raman), BET surface area analysis, scanning electron microscopy (SEM), Fourier transform infrared spectroscopy (FT-IR), and UV-vis diffuse reflectance spectra (DRS) were applied to characterize the crystallinity, morphology, surface structure, and other physicochemical properties of the nanoanatase TiO2. The photocatalytic activity of the prepared photocatalyst was evaluated by the photodegradation of formaldehyde in aqueous solution and compared with the commercial photocatalyst, namely, Degussa P-25. The result showed that the sample prepared at low temperature showed photocatalytic activity. The activity of the sample prepared at 120 °C was high and close to the sample calcined at 300 °C for 2 h. The degradation of formaldehyde by the highest active sample (calcined at 400 °C for 1 h) could almost achieve 100% within 80 min, which exhibited much higher photocatalytic activity than Degussa P-25. Heterogeneous semiconductor photocatalysis has been widely studied and applied recently in the environmental remediation field of removing organic and inorganic pollutants in water and air.1-4 It has been regarded as the most promising and efficient environmental purification technology. Among the semiconductor photocatalysis, TiO2 was the most widely studied because of its high quantum efficiency, high photocatalytic activity, and economical excellence.5-7 It has been shown that the performance of a TiO2 photocatalyst is strongly dependent on crystal phase, particle size, and surface structure (like surface hydroxyl, oxygen vacancy, etc.) and that the photocatalysis effectiveness is governed by the lifetime, or recombination probability, of the electron-hole pair.8 Anatase has been generally accepted to have a higher photocatalytic activity than rutile owing to its anomalously large Born effective charge tensor as the result of the presence of enhanced Ti(3d states)-O(2p states) hybridization.9 The preparation method and process of the photocatalyst have a great influence on the physicochemical properties, which would affect the photocatalytic activity of the photocatalyst.10,11 Among the preparation methods, sol-gel was the most widely employed to prepare the titanium dioxide because of its simplicity and low equipment requirement. However, the as-prepared titanium dioxide made by this method was generally amorphous and did not possess photocatalytic activity; it must be calcined generally at 500 °C to form highly crystalline anatase.12 The high temperature would seriously affect the particle size and surface structure and would result in a collapse of the mesoporous structure.13 So, a low-temperature preparation method is of great interest for preparing mesoporous nanocrystalline anatase. Many research studies have been conducted in this field. Bosc et al. developed a simple sol-gel route using surfactant for the synthesis of mesoporous and nanocrystalline anatase thin films at a low temperature of 350 °C.14 Beyers et al. synthesized * Corresponding author. Phone: +86 592 2186195. Fax: +86 592 2186195. E-mail: [email protected].

mesoporous titania applying the evaporation-induced selfassembly method. Anatase could be formed before calcination under certain circumstances. After a further thermal treatment at 400 °C, the obtained materials showed high photocatalytic activity.15 Peiro´ et al. applied microwave heating to enhance the crystallization to form anatase at 60 °C.16 Zhu et al. prepared the anatase TiO2 photocatalysts by a nonhydrolytic sol-gel reaction of TiCl4 and benzyl alcohol at low temperature, followed by subsequent calcination (400 °C) to form the high nanocrystalline anatase TiO2.17 Liu et al. treated the TiO2 colloid with UV-irradiation to form anatase TiO2 before calcination.18 After thermal treatment at 400 °C, the sample showed high crystallinity as well as high photocatalytic activity. Yoo et al. synthesized anatase-phase TiO2 at low temperature (120 °C) by a modified sol-gel method using ionic liquid as an effective template.19-21 Besides these, the most general method for synthesizing nanocrystalline titania at low temperature (usually below 150 °C) was hydrothermal synthesis,22,23 in which the solution or suspension was usually aged in the autoclave Teflon vessel. It generally required high pressure and special equipment. In this paper, a novel simple method is proposed for the preparation of nanocrystalline anatase of good crystallinity at a temperature as low as 100 °C. Tetrabutyl titanate was used as a titanium precursor, acetic acid was used as an inhibitor, and diethyl ether anhydrous was used as a solvent. Diethyl ether anhydrous was rarely used to synthesize TiO2 in the previous studies. Neither any other additive nor template agent, nor special instrument, was required, which showed a huge advantage over the other methods mentioned above. To our knowledge, this is the first report on the low-temperature synthesis of crystalline nanoanatase TiO2 in this way. Experimental Section Preparation. In this study, tetrabutyl titanate Ti(OC4H9)4 was used as a titanium precursor. Acetic acid was employed to control the acidity and restrain the intensive hydrolysis of the tetrabutyl titanate, and diethyl ether anhydrous was used as a

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4032 J. Phys. Chem. C, Vol. 113, No. 10, 2009 solvent. All the chemicals and solvents were of analytical grade. The synthesis method was as follows: At room temperature, a volume of 10 mL of tetrabutyl titanate was dissolved in 40 mL of diethyl ether anhydrous; then, 12 mL of acetic acid was added. After being stirred for 30 min, 15 mL of a mixture of deionized water and acetic acid in the ratio of 10:5 (v/v) was added dropwise to the precursor solution under vigorous stirring, and the white precipitate was formed in the solution. The suspension was aged under airtight conditions at room temperature for 24 h. Then, the upper clear diethyl ether liquid was removed, and the remainder was dried at different temperatures (40, 60, 80, 100, and 120 °C) for 24 h. After this, the samples were washed with water to remove the residual acid and dried at the same temperature for 24 h as before. Finally, the samples dried at 100 °C were calcined at 300, 400, 500, and 600 °C for 1 to 3 h. Characterization. The X-ray powder diffraction (XRD) of the samples was carried out on an X’Pert PRO X-ray diffractometer using Cu KR radiation (λ ) 0.154 06 nm) with a scanning angle (2θ) of 20-80° and voltage and current of 40 kV and 30 mA. The average size of anatase TiO2 crystallites was estimated by means of the Scherrer equation from broadening of the (101) anatase reflection. Raman spectroscopy was taken on a Renishaw UV-1000× spectrophotometer (equipped with Ar+ laser excitation wavelength of 514.5 nm) to investigate the structure of the samples. Surface areas calculated by the BET method were determined from nitrogen adsorptiondesorption isotherms at liquid nitrogen temperature by using a Micromeritics TriStar 3000 instrument. The morphologies of the catalysts were characterized by the LEO-1530 scanning electron microscope. The FT-IR spectra of the samples were measured by the KBr pellet method on a Nicolet Nexus FT-IR spectrophotometer in the wavenumber range 4000-400 cm-1. Optical absorption spectra (UV-vis) of the samples were recorded in the absorption mode using a Varian Cary 5000 UV-vis spectrophotometer in the span between 200 and 600 nm. Photocatalytic Evaluation. The photoactivity of TiO2 was evaluated by the photocatalytic degradation of formaldehyde under UV irradiation. The photoreactor was a glass cylinder (diameter 48 mm, length 288 mm) with a glass jacket covered by aluminum foil. A cylindrical quartz tube (diameter 25 mm, length 260 mm) was placed in the middle of the glass cylinder. The UV lamp (Philips, 11 W, 253.7 nm) was inserted to the quartz tube to irradiate the solution. A sample of 0.3 g was dispersed in 150 mL of aqueous formaldehyde solution with a concentration of 49 mg/L. Air was insufflated into the photoreactor to stir the suspension and provide O2 for the reaction. Adsorption equilibrium was reached in 10 min, and then the UV lamp was turned on. Each 3 mL of suspension was taken out and centrifuged at regular intervals, and 1.0 mL of clear liquid from each suspension was taken out to measure the concentration of formaldehyde via spectrophotometer. A similar experiment was carried out for the commercial Degussa P-25 to compare the photocatalytic activity. Meanwhile, the blank experiments were also carried out. Results and Discussion Characterization. Figure 1a shows the XRD spectroscopy of the samples dried at various temperatures from 40 to 120 °C for 24 h. Nanocrystalline anatase was formed when the sample dried at 40 °C, though the crystallinity was very low. Further, the crystallinity was enhanced gradually with the elevation of the drying temperature. Nanoanatase of high crystallinity was

Li et al.

Figure 1. XRD patterns of the thermal treatment of samples at different temperatures: (a) dried from 40 to 120 °C for 24 h and (b) before and after calcination from 300 to 600 °C for 2 h.

obtained for the sample dried at 100 °C and above. According to Debye-Scherrer, the crystallite sizes dried or calcined at different temperatures are summarized in Table 1. From Table 1, it can be seen that the crystallites grew as the drying temperature increased. The elevation of the temperature up to 120 °C contributed some changes to the crystallinity but also increased the particle size slightly. The obvious change can be seen from Figure 1b and Table 1 that the crystallinity became higher with the elevation of the heat treatment temperature, but the crystallite size increased slowly at temperatures below 500 °C. However, a significant increase in crystallite size can be observed (Table 1) when calcination temperature increased from 500 to 600 °C, as is shown by the sharp peaks in Figure 1b. Raman spectroscopy is a powerful technique for characterizing the microstructural and surface stoichiometric information of inorganic oxide.24 It has been considered to be sensitive to the TiO2 phase.25,26 Figure 2 shows the Raman spectra of the samples dried or calcined at different temperatures. The characteristic bands detected at 147, 197, 396, 515, and 638 cm-1 corresponding to A1g + 2B1g + 3Eg demonstrated the presence of anatase phase.27 Even the sample dried at 40 °C exhibited the characteristic bands of anatase phase, which correspond with its XRD pattern. This indicated that anatase TiO2 could be obtained at the temperature as low as 40 °C,

Nanocrystalline Anatase TiO2

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TABLE 1: Crystallite Sizes and BET Surface Areas of the TiO2 Thermally Treated at Various Temperatures and Durations temperature (°C)

60

80

100

120

300

400

400

400

500

600

time (h) crystallite size (nm) SBET (m2/g)

24 6.8 290.4

24 6.9 285.5

24 6.9 235.8

24 7.6 184.6

2 9.6 145.8

1 10.2 108.4

2 11.4 105.9

3 12.3 82.0

2 16.6 85.3

2 30.4 23.7

though the intensity of the Raman band was very low. The Raman spectra of the samples dried at relatively higher temperatures, especially the sample dried at 100 °C, demonstrated that highly crystalline nanoanatase TiO2 could be produced at low temperature, which was consistent with the characterization of XRD. Further, with an increase in drying and calcination temperature, a significant blue shift (toward the low wavenumber region) and a decreased broadening can be observed from the insert pattern of Figure 2, predominantly for the band around 147 cm-1 from 161.5 to 153.3 cm-1 and 146.8 to 141.8 cm-1, respectively. It has been known that the shift of the peak positions and the changes of the width correspond to changes of surface oxygen deficiency.24 The blue shift and decrease in band broadening demonstrated that the concentration of oxygen deficiency decreased as the thermal treatment temperature increased, which might be attributed to the formation of high crystallinity at the high temperature, leading to fewer oxygen vacancies on the surface of the crystal lattice.28

Figure 2. Raman spectra for the powders thermally treated at various temperatures: (a) dried from 40 to 120 °C for 24 h and (b) before and after calcination at different temperatures from 300 to 600 °C for 2 h. Insert: Raman peak shift in the range from 120 to 180 cm-1.

XRD and Raman studies concluded that nanocrystalline anatase not only can be obtained but can also be of high crystallinity at low temperature. A possible explanation for the anatase formation at low temperature can be found in the initial synthesis conditions. The acid and water took a very important part in the anatase formation. To verify the conclusion, another experiment has been done as follows: the obtained undried precipitate (the preparation method was mentioned above) was filtered and washed with deionized water 5 times to remove acetic acid or with ethanol to remove water and acetic acid, then dried at 100 °C for 24 h. It can be seen from Figure 3 that, after drying, the sample washed with ethanol was amorphous titania (Figure 3b), the sample washed with water was anatase (Figure 3c), and the sample that was not washed showed higher crystallinity (Figure 3d). This indicated that the acid and water were very important factors in the formation of anatase at low temperature. The water was necessary in the formation of titania, and the presence of the acetic acid accelerated the process. It was previously reported that the anatase phase had an edgeshared TiO62- octahedron.29 The linkage between TiO62octahedra was formed by the dehydration of Ti(OH)4 that was formed in the hydrolysis of the titanium precursor. Yanagisawa et al. reported two mechanisms of anatase nucleation: the solidstate mechanism and the dissolution precipitation mechanism, respectively.30 It was suggested that water was necessary in the transformation from amorphous TiO2 to anatase at low temperature. They considered that the presence of water in the crystallization reaction catalyzes the rearrangement of the TiO62octahedra in the amorphous titania by adsorption to the titania surface to form bridges between surface OH groups of different octahedra. When the acid concentration was relatively high, it could make the precipitate dissolve in the water; the presence of acid accelerated the formation of anatase. In this experiment, it could be seen that the precipitate dissolved gradually, then became gel. It was quite suitable for the dissolution precipitation mechanism. The sol-gel formation made the rearrangement of

Figure 3. XRD patterns of the samples thermally treated at different conditions: (a) ethanol as the solvent; (b) washed with ethanol; (c) washed with deionized water; and (d) not washed.

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Figure 4. SEM images of the samples before and after calcination: (a) dried at 100 °C and (b) calcined at 500 °C.

arbitrary bonds in the precipitation process toward the defined structure of anatase. The acetic acid may chelate to the Ti ions to facilitate a titania gel network, resulting in a less strong interaction between neighboring titania particles, and leading to a lower condensation rate of the titania. So, the condensation can be more slow and orderly to make the rearrangement of arbitrary bonds to the defined structure of anatase more complete. However, the anatase formation at low temperature can not only be attributed to the presence of water and acetic acid in this research. The solvent seems to also have an important influence. In this experiment, anatase formation at the low temperature mentioned above happened due to the hydrolysis conditions by using diethyl ether anhydrous as a solvent. There is a comparison experiment to verify the result as follows: A sample was prepared by the method mentioned in the experimental section, but the diethyl ether anhydrous was replaced by ethanol, which was frequently used in the preparation of TiO2 in previous studies. The sample was dried at 100 °C for 24 h. The XRD pattern (Figure 3a) shows that the sample prepared by this method was amorphous TiO2. So, it can be concluded that the diethyl ether anhydrous plays a very important role in the transformation from amorphous TiO2 to anatase at low temperature. It has been discussed in the Raman spectra (Figure 2) that the oxygen deficiency has existed in the samples prepared at various thermal treatment conditions. Ahonen et al. suggested that the oxygen deficiency was caused by the stress and strain effect induced by rapid oxidation and compaction or uncontrolled oxidation.31-33 Rapid hydrolyzation was found by using diethyl ether anhydrous as the solvent, while slow hydrolyzation was found by using ethanol as the solvent. It is expected that the rapid hydrolyzation in this system can generate Ti3+ defect sites on the TiO2 surface. Once the oxygen vacancies are formed, they can act as nucleation sites to promote the transition from amorphous to anatase phase.18 So, it can be concluded that the diethyl ether anhydrous used as the solvent can promote the formation of oxygen deficiency. Surface area of TiO2 has a large effect on the photocatalytic activity. Table 1 showed the surface areas of the samples prepared at different conditions, which were calculated by the BET method. The samples prepared at low temperature showed high BET surface areas, while those dried at 60 °C were 290.4 m2/g, which was far superior to results with the commercial photocatalyst Degussa P-25 (50 m2/g). As the thermal treatment temperature and time increased, BET surface area of the sample decreased gradually, reaching the lowest value 23.7 m2/g when calcined at 600 °C for 2 h. Morphology and size of nanocrystalline titanium dioxide before and after calcination were measured by the scanning electron microscope (SEM). Figure 4a shows that the average sizes of the primary particles for the samples before and after

Figure 5. FT-IR spectra of the samples: (a) before and (b) after calcinations.

calcination are about 10 and 20 nm, respectively, which roughly corresponds to the results of the XRD data (6.9 and 16.6 nm). It can also be seen from Figure 4a that only a slight amount of clustering was achieved by thermal treatment at low temperature. However, the high-temperature treatment made the aggregate phenomenon more serious, as is shown in Figure 4b. Infrared (IR) spectroscopy is an effective method to study adsorbed species on solid surfaces. Figure 5 and Figure 10 show the Fourier transform infrared (FT-IR) spectra of TiO2 before and after calcination at different temperatures for 2 h. A broadband peak around 3420 cm-1 is attributed to the O-H

Nanocrystalline Anatase TiO2

Figure 6. UV-vis spectra of TiO2 powders before and after calcination at different temperatures.

Figure 7. Photocatalytic degradation of formaldehyde on various samples dried at different temperatures.

stretching of physisorbed water on the TiO2 surface, and a relatively sharp band at 1637 cm-1 corresponds to the O-H bending modes of water molecules. It is well-known that the TiO2 superficial hydroxyl group plays an important role in the photocatalytic activity.34 The FT-IR spectra (Figure 5a) strongly indicate the presence of hydroxyl groups especially for the samples dried at low temperatures. From Figure 5b, it is observed that there was a significant reduction in the adsorbance intensity of the superficial hydroxyl groups with the elevation of the calcination temperature, which indicates that the hightemperature treatment resulted in a loss of surface hydroxyl groups. The strong absorption observed below 850 cm-1 is due to lattice vibrations of TiO2.35 The absorption at 1384 cm-1 can be assigned to the stretching vibration of -CH3 groups.36 Two strong bands (Figure 10) near 2300 cm-1 appearing in the sample calcined at 400 °C for 1 h are derived from bidentate carbonate and bicarbonate species.37 Two weak bands (Figure 5a) at 1530 cm-1 and 1420 cm-1 appearing in the samples dried at 60 and 80 °C are derived from bidentate (chelating or bridging) acetate ligands.38 Figure 6 shows the UV-vis diffuse reflectance spectra of the samples before and after calcination. It indicates that there is almost no absorbance of visible light, but there is an excellent optical response to UV light. From the pattern, it can also be observed that there are obvious differences to the absorbance

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Figure 8. Photocatalytic degradation of formaldehyde on various samples calcined at different temperatures for 2 h.

Figure 9. Photocatalytic degradation of formaldehyde on different samples calcined at 400 °C for different durations. The commercial photocatalyst Degussa P-25 was used for comparison, and blank experiments were also conducted.

in the UV light range between the samples. The sample dried at 120 °C showed the highest absorbance to the UV light, followed by the sample dried at 100 °C for 24 h and the sample calcined at 400 °C for 2 h. The response of the sample calcined at 300 °C to the UV light seems almost the same as that of the sample dried at 80 °C. The relatively low absorbance appeared in the samples calcined at 400 °C, which corresponds to the increase of the particle size. This may be attributed to the combined effect of the particle size and crystallinity. When the particle size is very small, the crystallinity may play a major role. When the crystallinity is high enough, the smaller the particle, the higher the response, which is due to the larger surface in the limit areas. Photocatalytic Activity. The photocatalytic degradation of formaldehyde has been conducted to evaluate the photocatalytic activities of the samples. The decision to choose formaldehyde was based on the further aim toward photodegradation of formaldehyde gas. Figure 7 and Figure 8 show the evolution curves of photocatalytic degradation of formaldehyde on various samples dried at low temperatures and calcined at different temperatures for 2 h, respectively. It indicates that the samples prepared at low temperature had photocatalytic activity. The photocatalytic activity was enhanced gradually with the elevation

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Figure 10. FT-IR spectra of the samples calcined at 400 °C for different durations.

Li et al.

Figure 11. UV-vis spectra of TiO2 powders calcined at 400 °C for different durations.

Conclusions of the drying temperature. The activity of the sample dried at 120 °C was close to that of the sample calcined at 300 °C for 2 h. Figure 8 shows the evaluation curves of photocatalytic degradation of formaldehyde on various samples calcined at different temperatures for 2 h. It indicates that the sample calcined at 400 °C for 2 h showed the highest activity. The activity declined gradually with the elevation of the calcination temperature, which corresponded to the particle size, BET surface areas, surface hydroxyl groups, and response to the UV-vis light. The small particle size, high BET surface areas, high response to the UV-vis light, and abundance of surface hydroxyl groups are beneficial to the photocatalytic activity. It is noticeable that the samples calcined at 300 °C or dried at low temperatures showed small particle size and high BET surface areas and were rich in surface hydroxyl groups; some (like the samples dried at 100 °C, 120 °C) even showed high response to the UV-vis light. In spite of this, they showed lower activity than the sample calcined at 400 °C for 2 h. This may be due to their lower crystallinity, which resulted in the existence of some amorphous TiO2. The presence of amorphous TiO2 could accelerate the electron-hole recombination, which lowered the photocatalytic activity. It can also be concluded from Figure 9 that the calcination time has influence on the photocatalytic activity. It reveals that the sample calcined at 400 °C for 1 h showed the highest activity, which could degrade almost 100% of formaldehyde within 80 min, while the commercial photocatalyst Degussa P-25 took approximately 110 min. It was considered that the crystallites grew gradually with the prolonging of the calcination time, leading to a series of changes in the physicochemical properties of the TiO2, which can be seen from the characterization of XRD and BET (Table 1), FT-IR (Figure 10), and DRS (Figure 11). When the crystallinity was relatively high, the sample calcined for 1 h showed more advantages in particle size, BET surface areas, surface hydroxyl groups, and response to the UV-vis light; so, it showed higher photocatalytic activity than others. A blank experiment was also conducted in this study, as we can see from Figure 9 in which the pure sorption of the catalyst and the photodegradation of the UV irradiation contributed little to the degradation of formaldehyde; almost all the degradation of formaldehyde is attributed to the photocatalytic degradation of the TiO2.

Nanoanatase TiO2 of high crystallinity has been successfully synthesized at a temperature as low as 100 °C by a novel simple method using diethyl ether anhydrous as the solvent. The prepared photocatalyst, which is rich in surface hydroxyl groups, has been found to exhibit high activity for the degradation of formaldehyde. The degradation of formaldehyde solution (about 49 mg/L) by the prepared photocatalyst (calcined at 400 °C for 1 h) could almost achieve 100% within 80 min, which showed much higher photocatalytic activity than the commercial photocatalyst Degussa P-25. Furthermore, the blank experiments revealed that almost all of the degradation of formaldehyde is attributed to the photocatalytic degradation of the TiO2. Acknowledgment. The authors thank Jianyang Wu, Umereweneza Daniel, Mujyambere Jean Marie Vianney, and Yujun Xu for their help with the revision of the paper. References and Notes (1) Agustina, T. E.; Ang, H. M.; Vareek, V. K. J. Photochem. Photobiol., C 2005, 6, 264. (2) Akyol, A.; Yatmaz, H. C.; Bayramoglu, M. Appl. Catal., B 2004, 54, 19. (3) Kumar, A.; Jain, A. K. J. Mol. Catal., A 2001, 165, 265. (4) Ao, C. H.; Lee, S. C. Appl. Catal., B 2003, 44, 191. (5) Fujishima, A.; Rao, T. N.; Tryk, D. A. J. Photochem. Photobiol., C 2000, 1, 1. (6) Zhang, D.-Y.; Qi. L.-M. Chem. Commun. 2005, 2735. (7) Makarova, O. V.; Rajh, T.; Thurnauer, M. C. EnViron. Sci. Technol. 2000, 34, 4797. (8) Dalton, J. S.; Janes, P. A.; Jones, N. G.; Nicholson, J. A.; Hallam, K. R.; Allen, G. C. EnViron. Pollut. 2002, 120, 415. (9) Mikami, M.; Nakamura, S.; Kitao, O.; Arakawa, H. Phys. ReV., B 2002, 66, 155213. (10) Arnal, P. A.; Corriu, R. J.P.; Leclercq, D.; Mutin, P. H.; Vioux, A. J. Mater. Chem. 1996, 6, 1925. (11) Watson, S. S.; Beydoun, D.; Scott, J. A.; Amal, R. Chem. Eng. J. 2003, 95, 213. (12) Sheng, Q.-R.; Cong, Y.; Yuan, S.; Zhang, J.-L.; Anpo, M. Microporous Mesoporous Mater. 2006, 95, 220. (13) Chen, Y.-G.; Dionysiou, D. D. J. Mol. Catal., A 2006, 244, 73. (14) Bosc, F.; Ayral, A.; Albouy, P.-A.; Guizard, C. Chem. Mater. 2003, 15, 2463. (15) Beyers, E.; Cool, P.; Vansant, E. F. J. Phys. Chem. B 2005, 109, 10081. (16) Peiro´, A. M.; Peral, J.; Domingo, C.; Dome`nech, X.; Ayllo´n, J. A. Chem. Mater. 2001, 13, 2567. (17) Zhu, J.; Yang, J.; Bian, Z.-F.; Ren, J.; Liu, Y.-M.; Cao, Y.; Li, H.-X.; He, H.-Y.; Fan, K.-N. Appl. Catal., B 2007, 76, 82. (18) Liu, H.-M.; Yang, W.-S.; Ma, Y.; Cao, Y.-A.; Yao, J.-N.; Zhang, J.; Hu, T.-D. Langmuir 2003, 19, 3001.

Nanocrystalline Anatase TiO2 (19) Yoo, K.; Choi, H.; Dionysiou, D. D. Chem. Commun. 2004, 2000. (20) Yoo, K. S.; Choi, H.; Dionysiou, D. D. Catal. Commun. 2005, 6, 259. (21) Choi, H.; Kim, Y.-J.; Varma, R. S.; Dionysiou, D. D. Chem. Mater. 2006, 18, 5377. (22) Wang, G.-H. J. Mol. Catal., A 2007, 274, 185. (23) Puddu, V.; Mokaya, R.; Puma, G. L. Chem. Commun. 2007, 4749. (24) Parker, J. C.; Siegel, R. W. Appl. Phys. Lett. 1990, 57, 943. (25) Iida, Y.; Furukawa, M.; Kato, K.; Morikawa, H. Appl. Spectrosc. 1997, 51, 673. (26) Balachandran, U.; Eror, N. J. Solid State Chem. 1982, 42, 276. (27) Burlacov, I.; Jirkovsky´, J.; Mu¨ller, M.; Heimann, R. B. Surf. Coat. Technol. 2006, 201–255. (28) Nag, M.; Guin, D.; Basak, P.; Manorama, S. V. Mater. Res. Bull. 2008, 02, 021. (29) Gopal, M.; Chan, W. J. M.; Jonghe, L. C. D. J. Mater. Sci. 1997, 32, 6001.

J. Phys. Chem. C, Vol. 113, No. 10, 2009 4037 (30) Yanagisawa, K.; Ovenstone, J. J. Phys. Chem. B 1999, 103, 7781. (31) Ahonen, P. P.; Kauppinen, E. I. J. Mater. Res. 1999, 14, 3938. (32) Parker, J. C.; Siegel, R. W. J. Mater. Res. 1990, 5, 1246. (33) Poniatowski, E. H.; Talavera, R. R.; Heredia, M. C.; Corona, O. C.; Murillo, R. A. J. Mater. Res. 1994, 9, 2102. (34) Szczepankiewicz, S. H.; Colussi, A. J.; Hoffmann, M. R. J. Phys. Chem. B 2000, 104, 9842. (35) Nakamura, R.; Imanishi, A.; Murakoshi, K.; Nakato, Y. J. Am. Chem. Soc. 2003, 125, 7443. (36) Kozlov, D. V.; Vorontsov, A. V.; Smirniotis, P. G.; Savinov, E. N. Appl. Catal., B 2003, 42, 77. (37) Rethwischt, D. G.; Dumesic, J. A. Langmuir 1986, 2, 73. (38) Guo, G.-Q.; Whitesell, J. K.; Fox, M. A. J. Phys. Chem. B 2005, 109, 18781.

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