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An amorphous TiO2 sol sensitized with H2O2 with the enhancement of photocatalytic activity
Journal of Alloys and Compounds 497 (2010) 420–427
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An amorphous TiO2 sol sensitized with H2 O2 with the enhancement of photocatalytic activity Jian Zou a,b,∗ , Jiacheng Gao b , Fengyu Xie b a b
School of Chemistry and Chemical Engineering, Southwest University, Chongqing, 400715, PR China College of Materials Science and Engineering, Chongqing University, Chongqing, 400044, PR China
a r t i c l e
i n f o
Article history: Received 16 November 2009 Received in revised form 6 March 2010 Accepted 8 March 2010 Available online 15 March 2010 Keywords: TiO2 sol H2 O2 Sensitizing Visible light Amorphous
a b s t r a c t An amorphous free-impurity TiO2 sol was synthesized only by means of ultrasonic dispersing of Ti(OH)4 precipitation without any peptizing agents. Anatase sol was obtained by hydrothermally treating the amorphous TiO2 sol. Photocatalytic tests showed that the amorphous sol exhibited higher photocatalytic activity for the photodegradation of methylene blue (MB) under visible light and UV irradiation than anatase sol and P25 TiO2 . The photocatalysts were characterized by UV–visible spectroscopy, Raman spectrometer, photoluminescence (PL), FTIR, TEM, XRD, and XPS. The results showed the peroxide titanium complexes were formed on TiO2 sol particles sensitized with H2 O2 . More peroxide complexes and stronger vis absorption were observed for the amorphous TiO2 sol sensitized with H2 O2 .The sensitization of H2 O2 would completely quench PL of the amorphous sol, but little impact on the anatase sol. The surface structure with fewer physisorbed water molecules facilitated the sensitization of H2 O2 to the amorphous TiO2 sol particles, which resulted in the generation of more OH radicals and higher photocatalytic activity under both visible light and UV irradiation. Crown Copyright © 2010 Published by Elsevier B.V. All rights reserved.
1. Introduction It is generally accepted that nanosize titanium dioxide (TiO2 ) is a promising photocatalyst for photocatalytic reactions, own to its low cost, nontoxicity, unique photocatalytic efficiency and high stability under photoirradiation. For nanosize TiO2 , its high activity was primarily governed by the well-known quantum confinement [1]. Some photocatalytic studies have demonstrated that the smaller anatase nanoparticle exhibited higher activity [2,3], and the size of particles could be controlled only by regulating the pH of solution [4]. However, since the small-sized TiO2 nanoparticles facilitate the agglomeration in synthesis and application, its photocatalytic activity usually decreases dramatically or even negligible. In addition, nothing but UV light of the sunlight can be absorbed by TiO2 due to its wide intrinsic band gap. These seriously limit its application and development. It is well known that TiO2 nanoparticles can be dispersed well in TiO2 hydrosol, and the evident quantum size effect was observed for the transparent TiO2 sols [5]. The TiO2 hydrosols could be facilely prepared by using titanium alkoxides and inorganic salts as precursors [5–9]. However, these TiO2 hydrosols contained a signif-
∗ Corresponding author at: School of Chemistry and Chemical Engineering, Southwest University, Chongqing, 400715, PR China. Tel.: +86 023 68252361; fax: +86 023 68252360. E-mail address: [email protected] (J. Zou).
icant amount of impurities which came from the above-mentioned precursors and peptizing agents, such as hydrochloric acid and nitric acid. It was reported that the impurities or low pH values, which were produced by hydrolysis of titanium precursors and during peptization, would retard the oxidation of the substrate [7]. Recently, it was reported that a neutral impurity-free sol could be obtained by introducing some SiO2 sol as stabilizer [10] or using H2 O2 as peptizing agent [11]. Lv et al. [8] reported that P-doped TiO2 prepared by introducing phosphoric acid into TiO2 sol displayed excellent activity. In addition, the crystal structure of TiO2 catalyst can play a crucial role in determining the efficiency of photocatalytic activity. It was well known that anatase was usually more photoactive than rutile, and the higher crystallinty would exhibit higher photocatalytic activity [12], but higher activity was also observed for rutile [13]. The enhancement of photocatalytic activity due to a synergistic effect of anatase and rutile mixed phases was supported by other experimental evidences [14]. In fact, a controllable phase composition of TiO2 could be selectively prepared by changing nucleation mode [13], but an amorphous phase usually accompanied with the crystalline phase [9,15]. It is reported that amorphous TiO2 is nearly inactive due to the facilitated recombination of the photoformed electron and hole [9,16], but one with large surface area and special microstructures is excepted [17]. The hydrothermal treatment could enhance the crystallization [18], but this would result in the precipitating of sols [19]. Sayılkan et al. [20] synthesized a complete anatase TiO2 crystalline phase by hydrothermal process at 200 ◦ C
0925-8388/$ – see front matter. Crown Copyright © 2010 Published by Elsevier B.V. All rights reserved. doi:10.1016/j.jallcom.2010.03.093
J. Zou et al. / Journal of Alloys and Compounds 497 (2010) 420–427
and prepared transparent sols by dispersing these nanopowders ultrasonically in water, but this sol could not degrade methylene blue (MB) due to the reduction–oxidation behavior of MB. The recombination of photogenerated holes and electrons has been regarded as an unfavorable or limiting process in photocatalysis of TiO2 catalyst. As mentioned above, the impurities and amorphous phase in TiO2 sol would depress the photocatalytic activity, and it was almost impossible to avoid those deficiencies during preparation of sol by the methods above. Therefore, a new method for preparing or modifying TiO2 sol was necessary in order to further enhance its activity. Some researches have demonstrated that adding an external electron acceptor such as hydrogen peroxide could enhance the performance of photocatalysis due to reducing the recombination of charge carriers [21,22]. Moreover, hydrogen peroxide molecules can be adsorbed on the TiO2 surface to form surface peroxide complexes, resulting in visible photoresponse [23–25]. In addition, it was also reported that H2 O2 could increase the stability of TiO2 sol [11]. In this study, we prepared the free-impurity amorphous TiO2 and anatase sol by ultrasonically dispersing Ti(OH)4 , and sensitized them with H2 O2 . It was expected that H2 O2 could inhibit the recombination of the exited charge carriers to enhance the activity of TiO2 sol. In addition, it was expected that sols sensitized with H2 O2 could degrade the MB under visible light irradiation. 2. Experimental 2.1. Preparation of TiO2 sol Titanium tetrachloride (Sinopharm Chemical Reagent Co. Ltd., China) was used as precursor without any further purification. The precursor TiCl4 was added dropwise into mixture of deionized water and ice under vigorous stirring to make resulting solution 0.1 M in TiCl4 . The mixture was kept stirring and heating on the agitating heater until the pale blue colloidal dispersion was observed, and the temperature was about 90 ◦ C. The dispersion was cooled to the room temperature. In order to ensure the complete hydrolysis reaction, a 10 wt% NH3 ·H2 O aqueous solution was added dropwise into the transparent sol to form a white precipitate Ti(OH)4 with ultimate aqueous mixture of pH 7–8. Subsequently, the resulting precipitates were separated from the suspension by using filtration and washed repeatedly with deionized water to remove the residual chloride ions, and the final pH value of filtrate agreed with deionized water. The sol containing 0.1 M TiO2 was recovered by dispersing the Ti(OH)4 precipitate into deionized water under ultrasonic condition, and the sol was denoted as AM sol. The AM sol was transferred to a stainless steel Teflon-lined autoclave. The autoclave was sealed and transferred into a 160 ◦ C oven. The sol was aged for 6 h without any agitation and cooled to room temperature at atmosphere. The hydrothermal sol was denoted as HT sol. All sols were uniform, stable and transparent. The sensitizing TiO2 sols were acquired by directly adding H2 O2 into AM and HT sols to make resulting solution 0.3 wt% in H2 O2 . The sols sensitized with H2 O2 were denoted as AM-H2 O2 sol and HT-H2 O2 sol respectively. Interestingly, the sol which was prepared by peptized Ti(OH)4 with inorganic acid [9] would deposit immediately when adding H2 O2 . The xerogels were gained by directly drying the sols at room temperature. 2.2. Characterization The phases of the products were characterized through X-ray diffraction method using CuK␣ radiation ( = 0.15418 nm) in a XD-3 diffractometer (Beijing Pgeneral). Transmission electron microscopy (TEM, Hitachi H7500) was performed at an accelerating voltage of 80 kV for electrons. Fourier transform infrared (FT-IR, Bruker TENSOR 27 FT-IR) spectra of TiO2 powders were obtained using a spectrometer with KBr pellets technique. The Raman spectra were recorded on a Bruker RFS100 spectrometer with a resolution of 4 cm−1 , using 1054 nm light as an excitation source. The precision of the wavenumber was 1 cm−1 . The UV–vis spectra were recorded on a spectrophotometer with an integrating sphere (Shmadzu UV-2550); BaSO4 was used as a reference sample. XPS spectrometer (VG Scientific ESCALAB 250) equipped with two ultra-high vacuum (UHV) chambers measurements were performed and all binding energies were referenced to the C1s peak at 284.6 eV of the surface adventitious carbon. The measurements of the amount of • OH were conducted according to the literature [21,26]. For UV irradiation, a 40 W UV lamp was used. For visible light irradiation, a 60W tungsten halogen lamp (Osram, HALOSPOT 111IRC 48837 ECO FL) with a filter ( > 420 nm) was used. 100 ml sol contained 0.1 wt% TiO2 , 0.01 M NaOH and 3 mM terephthalic acid. Before exposure to light irradiation, the sol was stirred in the dark for 30 min. And then, 10 ml of sol was taken out after 30 min for fluorescence spectroscopic measurements (Hatachi F-4500). The fluores-
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Fig. 1. XRD patterns of xerogel powders.
cence spectra of sols were recorded on fluorescence spectrophotometer (Hatachi F-4500). 2.3. Measurement of photocatalytic activity The photocatalytic activities of the samples were measured by the degradation of MB (20 mg L−1 ). For photodegradation experiments, the required volume of dye solution was added into TiO2 sol to result in 200 ml diluted sol containing 0.1 wt% TiO2 . A 450 W high pressure mercury lamp (Shanghai YaMin) with a filter was used as a visible-light source ( > 420 nm). For UV irradiation, a 40 W UV lamp was used. Before the photocatalytic experiment, the sol was stirred in the dark for 30 min and its concentration was original concentration (Co ) of MB. After a given irradiation time, 10 ml sol was withdrawn, and the concentration (C) of the MB was monitored by measuring the maximum absorbance of MB using the UV–vis spectrum. It must be noted here that no filtration and centrifuging were needed for the TiO2 sol, since it was transparent and do not absorb visible light beyond 500 nm (it was mentioned later) which can be absorbed by MB. But Degussa P-25 TiO2 needed to be removed from its suspension by centrifuging at 5000 rpm after photocatalysis procedure.
3. Results and discussion 3.1. Crystal structure and morphology Fig. 1 showed the X-ray powder diffraction (XRD) patterns of the as-prepared xerogel powders. The XRD patterns indicated that nothing could be observed for AM sol except for a weak peak located at 2 values of 25.3◦ which corresponded to anatase phase. It could be concluded that AM sol remained to be predominantly amorphous titania besides trace amounts of anatase. After hydrothermal treatment, the peaks of HT sol became sharper and more intensive, indicating better crystallization. And all diffraction peaks could be assigned to the anatase phase without any indication of other crystalline byproducts such as rutile or brookite, indicating the formation of predominant anatase TiO2 nanocrystals for HT sol. Generally, the anatase phase is considered to be the most photoactive. Traditionally, it was reported that the pH value [27] and chloride ions [28] would influence the phase composition for hydrothermal synthesis of TiO2 . At basic pH the main or exclusive anatase phase could be prepared, while at acidic pH, rutile formation was favored [27]. And existence of chloride ions also promoted rutile formation [28]. In this study, TiO2 sols were neutral, so anatase TiO2 benefited more from the absence of the residual chloride ions. The crystallite size estimated using the Scherrer equation [20] was about 6 nm for HT sol. The crystal structure of xerogel powders can also be identified by Raman scattering. Typically, for anatase TiO2 , there are six Raman active fundamental modes at 144 cm−1 (Eg ), 197 cm−1 (Eg ), 399 cm−1 (B1g ), 513 cm−1 (A1g ),519 cm−1 (B1g ), and 639 cm−1 (Eg ), respectively [29]. Fig. 2 showed the Raman spectra of HT sol and AM sol. Their spectra corresponded to the typical anatase TiO2 ,
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Fig. 4. IR spectra of as-prepared xerosols. Fig. 2. Raman spectra of xerogel powders (Inset: the expanded spectra for the lowest-frequency Eg mode).
3.2. Composition and chemical structure which indicated that only anatase phase was formed for these asprepared sols. And it was clear that the modes of the HT sol are much stronger than that of the AM sol, which implied that the crystallinity of hydrothermal sol TiO2 was enhanced. These conclusions were consistent with the XRD results shown in Fig. 1.In addition, a blue shift of the lowest-frequency Eg modes for as-prepared sols was also observed from the inset of Fig. 2. For HT sol, the mode shifted to 147 cm−1 . And more blue shift was observed for AM sol. In the case of anatase TiO2 nanoparticles the quantum size confinement effects would result in the blue shift of the low-frequency Eg mode [30]. Kelly et al. [31] also thought that the blue shift of the Eg modes was associated with changes in the particle size, i.e., with a decrease in particle size the peak shifted to the blue, and decreased in intensity. Therefore, the Raman results suggested that the particle size of the AM sol was smaller that of the HT sol. A similar behavior of frequency shift was observed for TiO2 nanoparticles by others [29]. The TEM micrographs of the as-prepared sols were presented in Fig. 3. It was clear that particles for the hydrothermal sol TiO2 had diameters about 6 nm, which was accordant with the results of XRD. For the sol before hydrothermal treatment, the particles size was about 3 nm, which was nearly a moiety as large as that of hydrothermal sol.
Fig. 4 showed the IR spectra of the sols before and after hydrothermal treatment. The xerogels exhibited similar IR spectra. The transmission bands in the range 400–800 cm−1 are characteristic of the formation of an O–Ti–O lattice. The peak at ∼1630 cm−1 results from O–H bending of molecularly physisorbed water [32,33]. The broad adsorption band observed at 3000–3700 cm−1 can be assigned to the stretch region of the surface hydroxyl groups and molecularly chemisorbed water [32,33]. The stronger O–H peak which resulted from the physisorbed water molecules in HT sol was observed, but an opposite result was also observed for stretch strength of the surface hydroxyl groups and chemisorbed water, compared with the sol before hydrothermal treatment. These results indicated that a higher content of physisorbed water was on HT sol particles, but more surface hydroxyl groups and chemisorbed water were on AM sol TiO2 . The Raman spectra in the range of 700–1200 cm−1 were recorded in Fig. 5. The Raman peaks of hydrogen peroxide in the different phase were detected at 884 cm−1 and at 874 cm−1 , corresponding to the O–O symmetric stretching vibration of H2 O2 in the liquid phase at ambient conditions and in the supercritical phase, respectively [34]. In the present work, for the as-prepared sols sensitized with hydrogen peroxide, a new vibration band at 840–920 cm−1 was observed, which indicated that peroxo groups
Fig. 3. TEM maps of (a) AM sol and (b) HT sol.
J. Zou et al. / Journal of Alloys and Compounds 497 (2010) 420–427
Fig. 5. Raman spectra of HT sol TiO2 and AM sol TiO2 treated and untreated with H2 O2 .
were on the surface of sols particles. Comparing with the Raman spectra of the sols treated with H2 O2 , we also noticed that the O–O stretching vibration for AM-H2 O2 sol particles was stronger than that of HT-H2 O2 sol. This was likely to be result from the formation of fewer surface peroxide complexes on the surface of the hydrothermal sol particles, due to more physisorbed water (Fig. 4). The physisorbed water molecules were bound by weaker hydrogen bond with the hydroxyl groups of TiO2 surface to form a multilayer, but the chemisorbed water molecules directly interacted with the TiO2 surface to form a monolayer [33,35]. In presence of H2 O2 , the –OOH groups of H2 O2 would replace –OH groups on the surface
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of titania to form surface complexes [23–25]. And Henderson et al. [36] thought the 2nd layer of water on the surface of TiO2 would block access of O2 molecules to the surface OH groups. We deduced this blocking effect was the same with H2 O2 molecules. The XPS technique can monitor the electron binding energy of sites with a few nanometer of materials surface, from which some information on surface composition and chemical structures would be inferred. We had examined the XPS region of O1s and Ti2p for asprepared sols. Fig. 6 showed the XPS spectra of O1s and Ti2p for used xerogels treated with H2 O2 .For the sols before hydrothermal treatment (Fig. 6a), O1s band could be separated into three peaks, located at 529.9, 531.3 and 532.8 eV respectively, which could be attributed to Ti–OH at 531.9 eV, Ti–O at 530.1 eV and peroxo groups at 532.8 eV [24,25]. Compared with AM-H2 O2 sol, the band of HTH2 O2 sol only could be separated into two peaks (Fig. 6b), and the absence of the band corresponding to peroxo groups. In fact, both of sols were sensitized with H2 O2 , and peroxo groups were also detected for them in Fig. 5, but it simultaneously showed that fewer complexes were on the surface of HT-H2 O2 sol. So, peroxo group peak in O1s XPS spectra was likely to be muffled by the OH peak. Herein, the absence of peroxo group peak in XPS of O1s reconfirmed the conclusion that the formation of surface peroxide complexes was restrained due to the overmuch physisorbed water on the HTH2 O2 sol particles. The Ti2p band for hydrothermal sol particles consists of two peaks of Ti2p3/2 and Ti2p1/2 located at 458.3 and 463.9 eV (Fig. 6c) [25], respectively. However, the Ti2p3/2 peak for the sol particles before hydrothermal treatment shifted to higher binding energy (BE). This shift in binding energy may be attributed to the change in the chemical environment. The similar shift of the Ti2p3/2 BE for TiO2 /SiO2 composite oxide after adding H2 O2 was also observed due to the formation of Ti–O–O–Si peroxo moieties [37]. It was
Fig. 6. High resolution XPS spectra of O1s , Ti 2p in HT-H2 O2 sol TiO2 and AM-H2 O2 sol TiO2 .
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reported that the formation of more peroxide complexes would result in larger shift of the Ti2p3/2 BE for sulfated TiO2 sensitized with H2 O2 [25]. Therefore, the more shift of the Ti2p3/2 BE for the AM-H2 O2 sol could be concluded to result from the formation of more peroxide complexes, which was consistent with the results above. 3.3. UV–vis spectra and PL The UV–vis absorption spectra of the as-prepared sols were illustrated in Fig. 7. All sols were adjusted to the concentration of 1.0 g L−1 TiO2 in aqueous solution for spectrum measurement. It could be seen that the no absorption in the visible region (wavelength > 420 nm) was observed but the strong absorption in the ultraviolet range for sols untreated with H2 O2 . It meant both sols before treatment with H2 O2 were highly transparent and could not drive the direct photocatalysis under visible light irradiation. The optical absorption edge was estimated at 347 nm for the HTH2 O2 sol and 337 nm for the AM-H2 O2 sol, and the corresponding band-gap energy was identified as 3.57 and 3.68 eV, respectively. The optical absorption edge appeared an obvious blue shift for the above sols relative to 387 nm (about 3.2 eV) [1] of the block anatase TiO2 .This was often ascribed to the quantum size effect of sol nanoparticles [5] due to their high dispersion and small size. After sensitized with hydrogen peroxide, both sols could absorb visible light (wavelength > 420 nm), but AM-H2 O2 sol displayed much stronger visible absorption than HT-H2 O2 sol. Generally, the change of energy band structure can result in vis absorption [38,39]. The electronic structure could be investigated by value band (VB) XPS spectrum [38,39]. The value band XPS spectra of sols treated with H2 O2 were presented in the inset of Fig. 7. It could be seen that the value band edge of hydrothermal sol shifted slightly to Fermi level, which meant narrower energy gap for the hydrothermal sol. This was incompatible with optical absorption of sols treated with H2 O2 . Obviously, the shift of VB edge could not explain HT-H2 O2 sol had poorer visible absorption. It was reported that the peroxide complexes on the surface of TiO2 could absorb visible light [23–25]. So, it could be deduced that the stronger vis absorption came form more peroxide complexes on the AM-H2 O2 sol, which had been confirmed by the results of Raman and XPS. Moreover, the hydrothermal treatment would weaken the quantum size effect of sols due to increase in particle size, which may be the reason that the VB edge shifted to Fermi level. Photoluminescence (PL) spectroscopy has been widely conducted to study charge carrier trapping, the role of surface states in photocatalysis and surface property for TiO2 particles because PL
Fig. 8. Photoluminescence of used sols treated and untreated with H2 O2 .
emission originate from the recombination of electron–hole pairs in the surface and bulk of TiO2 [40]. The photoluminescence spectra for our samples were showed in Fig. 8. For the hydrothermal sols, the stronger emission band at 390 nm is assigned to band-to-band transition, and the weaker emission band at 465 nm is related to the surface trap state [41]. Although the AM sols did not display a distinct peak around 390 nm, a weaker broad emission band around 350–600 nm was observed. This indicated the tiny crystalline TiO2 must be present in AM sol TiO2 because the amorphous TiO2 do not show fluorescence [42], which was also supported by above results of XRD and Raman. Otherwise, it could be seen that the PL spectrum of hydrothermal sol was much more intense than that of the sol before hydrothermal treatment, which could be attributed to the inferior crystallization (Figs. 1 and 2) and the larger density of surface states or defects for smaller AM sol TiO2 particles (Fig. 3) where the nonradiative emissions predominated [43]. Additionally, it could also be seen that the treatment with H2 O2 could reduce the PL intensity of sols from Fig. 8. PL emission of TiO2 is usually affected by surface chemistry. Deposition of Pt on TiO2 [44], adsorbed H2 O and N2 [45], organic modification [46] and doping of nitrogen [47] would quenched or weaken the PL emission of TiO2 . For sol sensitized with H2 O2 , its surface chemistry would be changed due to the formation of peroxide complexes on the TiO2 . PL emissions, attributed to indirect band-gap emission and surface defects, were nearly quenched for AM-H2 O2 sol TiO2 , which indicated that the effect of peroxide complexes on sols was not limited merely to surface. And Shu et al. [37] also reported that PL emission could be quenched for TiO2 /SiO2 composites treated with H2 O2 . However, an obvious difference for hydrothermal sols could be observed. The PL emissions were much less affected, especially for the lower energy emission associated with defect sites [42], suggesting that effect of H2 O2 on surface chemistry of HT sols TiO2 was negligible. Certainly, the decrease of effective interface between TiO2 and H2 O2 , resulting from the increase of particle size, would reduce peroxide complexes on the hydrothermal sol particles to weaken the effect of H2 O2 .However, the fact that the negligible decrease of the emission at 450–500 nm, being directly affected by surface chemistry, reconfirmed that overmuch physisorbed water molecules would block access of H2 O2 molecules to TiO2 surface. This was consistent with the results of Raman and XPS. 3.4. Photocatalytic activity
Fig. 7. Diffusion reflection absorption of used sols treated and untreated with H2 O2 (Inset: The values band XPS of as-prepared sol treated with H2 O2 ).
Photocatalytic activity tests were conducted by the degradation of methylene blue (MB) in aqueous solution under ultraviolet and visible light irradiation. Fig. 9 presented the photocatalytic results of the as-prepared sols in contrast to P25 under ultraviolet irradi-
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Table 1 The maximum absorbance (Co value) of MB in different sol before irradiation.
Fig. 9. Degradation of methylene blue under ultraviolet radiation for used samples.
ation. As could be seen, the photocatalytic degradation of MB in the hydrothermal sol behaved better than in the AM sol. About 60% discoloration for the hydrothermal sol was observed, but only about 13% degradation for the AM sol in the same time, which even was lower than that of pure H2 O2 . Interestingly, an opposite result was observed after the sols were sensitized with hydrogen peroxide. About 90% discoloration was acquired for AM-H2 O2 sols, which was nearly 7 times as much discoloration as ones before treatment with H2 O2 . However, the discoloration of HT-H2 O2 sols decreased to 38%, only 60% of sols before treatment with H2 O2 . Anatase was usually deemed to be the most photocatalytically active TiO2 . And activity of amorphous TiO2 was nearly negligible due to the facilitated recombination of the photoformed electron and hole at the traps on the surface and in the bulk of the particles [9,16]. PL emission is the result of the recombination of excited electrons and holes, and generally the higher PL intensity indicates a higher recombination rate of these electrons and holes under light irradiation. However, for amorphous TiO2 , the nonradiative emission predominated [43]. Therefore, PL emission of AM sol was not the result of the recombination of all excited electrons and holes. It was possible that much more recombination occurred for AM sol TiO2 .Therefore, the higher activity of HT sol can be ascribed to better anatase phase, as discussed above based on Figs. 1 and 2. Hydrogen peroxide may be split photocatalytically to produce hydroxyl radicals directly under UV light irradiation [22,48]. Those hydroxyl radicals could degrade MB as showed in Fig. 9, but if the decomposition of methylene blue was only from H2 O2 itself, then the efficiency of the decomposition using a different sol photocatalyst might not vary so much as that shown in Fig. 9. On the other hand, hydrogen peroxide is an electron acceptor, which would lower the electron–hole recombination rate to increase the photocatalyzed rate [21,22]. However, the enhancement of H2 O2 for photacatalytic activity of TiO2 had been in some cases controversial, and it appeared strongly dependent on substrate type and on various experimental parameters. Some researchers [24,27] found the enhancement practically depended on crystalline phase of catalysts, and rutile was preferable to anatase. Hirakawa et al. [26] suggested that the adsorption structure of H2 O2 on the rutile TiO2 surface was preferable to produce OH• . Augugliaro et al. [49] and Ban et al. [50] found the presence of hydrogen peroxide even did not affect appreciably the photodegradation rate of organic dye, and Chu and Wong [48] also observed that H2 O2 -assisted photocatalysis of dicamba was not enhanced under 350 nm UV irradiation. In addition, the adsorption of MB on TiO2 also can influence the degradation of dye. Table 1 showed the Co values of MB in different sols. However, as could be seen, a less than 10% difference of Co values indicated the effect was extremely limited.
Sample
AM sol
HT sol
AM-H2 O2 sol
HT-H2 O2 sol
Absorbance (a.u.)
1.561
1.672
1.533
1.642
In our study, we thought that an opposite effect of H2 O2 for both used sols was predominantly dependent on the surface chemistry of sol particles. For the HT-H2 O2 sol, as discussed above, more physisorbed water molecules would baffle the direct interaction of H2 O2 to TiO2 , which would result in less capture of electron. As a result, the separate effect of electron and hole pairs was negligible for HT-H2 O2 sol, as indicated in the result of PL in Fig. 8. For the AM-H2 O2 sol, the excited electrons would be recombined in the impurity states without H2 O2 , resulting in low activity as ascribed in Fig. 9.With H2 O2 , the recombination (even including nonradiative one) of charge carriers was remarkably restrained by H2 O2 , with the exhibition of quenching of PL emission due to fewer physisorbed water molecules on sol particles. In addition, the excess H2 O2 molecules would scavenge the valuable hydroxyl radicals (HO• ) to form much weaker oxidant HO2 • [48]. Therefore, the balance of two photochemistry action of H2 O2 for TiO2 sol systems would result in the increase of hydroxyl radicals for AM-H2 O2 sol but the decrease for HT-H2 O2 sol. The concentration of HO• could be estimated by terephthalic acid fluorescence (FL) probe method [21,26]. Since the FL spectra were identical to that of TAOH generated from TA by the reaction with OH• [21], the relative content of OH• formed in TiO2 sol was estimated from that of TAOH. Fig. 10 showed FL spectra of the used sols with 3 Mm TA after irradiated for 30 min.It could be seen that H2 O2 decreased the FL intensity for the HT-H2 O2 sol, but a prominent increase for the AM-H2 O2 sol was observed, which indicated that H2 O2 promoted authentically the formation of hydroxyl radicals for AM-H2 O2 sol as anticipated above. This was consistent with the results of photocatalytic degradation of MB. In addition, Sayılkan et al. [20] and Doushita and Kawahara [51] found the faded MB in TiO2 suspensions under illumination could reverted to its original blue color while stored in the dark for some hours. This reversible color change was also observed by others [7,52]. This reversible chromic phenomenon could be eliminated by enhancing oxidation–reduction potential of dye to the negative direction [20] or conducting simultaneously air-bubbling into MB/TiO2 suspensions under illumination [7]. Fig. 11 indicated the absorption spectrum changes of MB after UV irradiation and during storage in the dark. An absorption peak in AM-H2 O2 sol appeared
Fig. 10. Fluorescence spectra obtained for used sols containing 3 mM terephthalic acid under ultraviolet irradiation for 30 min. (Exited wavelength: 312 nm; Inset: under visible irradiation for 30 min).
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Fig. 11. The UV–vis absorption spectra of AM-H2 O2 sol containing MB at different time under UV irradiation and during storage in the dark (The spectra of (a) 0 min, (b) 120 min, (c) suspension (b) in the dark for 12 h).
at around 580 nm which shifted 85 nm relative to the peak of MB at 665 nm. This shift was also observed for TiO2 /SiO2 composites [7,51]. This absorption peak vanished nearly under UV irradiation for 120 min. After stored in the dark for 12 h, the peak did not only revert to its original shape but disappeared completely, indicating that no reversible discoloration of MB dye occurred to AM-H2 O2 sol without air-bubbling. The reason of the reversible chromic phenomenon for MB was that the colorless MB molecule reduced by photogenerated electrons would be reoxidized by oxygen back to its original blue color [7,20,51]. For AM-H2 O2 sol, as described in Fig. 10, an amount of OH radicals were generated. As a strong oxidant, OH radicals could photodegraded directly MB by oxidation, which avoided the blue MB to be reduced to colorless MB. According to Fig. 7, the sols treated with hydrogen peroxide could absorb visible light, which meant that they could drive the direct photocatalysis under visible light irradiation. Fig. 12 showed the results of photocatalytic degradation of MB aqueous solution for the used samples under visible light irradiation ( > 420 nm). For the reference of P25 without H2 O2 , only 18% discoloration was detected in 120 min, which hinted that the effect of dye sensitization [53] from MB was very limited. Similarly, the highest activity was observed for the AM-H2 O2 sol under visible light irradiation, especially in the primary stage of photodegradation, 75% discoloration for the AM-H2 O2 sol, but only 23% discoloration for the HT-H2 O2 sol within 40 min. In addition, Fig. 12 also showed H2 O2 without TiO2 sol was unable to degrade MB under visible irradia-
tion, which indicated the visible degradation of MB was not from H2 O2 .Besides, no reversible chromic phenomenon was observed for AM-H2 O2 under visible light irradiation. For TiO2 sensitized with H2 O2 , the conduction band electrons, which would react with the adsorbed H2 O2 on the TiO2 surface to generate the OH radical, only came from the excited surface peroxide complexes, not valence band of TiO2 under visible light irradiation [23]. Hereby, amounts of hydroxyl radicals were practically dependent on the surface peroxide complexes. As discussed above based on the results of IR, Raman, XPS and UV–vis, fewer peroxide complexes in response to visible light were formed on HT-H2 O2 sol particles owing to more physisorbed water molecules on their surface, resulting in inferior vis absorption. This was a great disadvantage to the generation of hydroxyl radicals for HT-H2 O2 sol TiO2 under visible light irradiation. It could be seen from the inset of Fig. 10 that the FL intensity at around 426 nm was lower for HTH2 O2 sol TiO2 than that for AM-H2 O2 sol TiO2 under visible light irradiation for 30 min.This indicated that fewer hydroxyl radicals were generated in the HT-H2 O2 sol. 4. Conclusions The stable, transparent TiO2 sol was prepared only by means of ultrasonic dispersion of Ti(OH)4 precipitation without any peptizing agents. The sol TiO2 was predominantly amorphous titania besides trace amounts of anatase. The sol with better crystallization was acquired by hydrothermal treatment. After sensitized with H2 O2 , the sols could absorb visible light due to formation of peroxide complexes on TiO2 surface. The hydrothermal sol treated with H2 O2 exhibited the inferior vis absorption and photoactivity under both UV and visible light irradiation. The hydrothermal treatment would increase the physisorbed water on TiO2 , resulting in formation of fewer peroxide complexes and hydroxyl radicals, which decreased photoactivity of the hydrothermal sol. Herein, the results presented here demonstrated that the sensitization of H2 O2 could enhance efficiently photoactivity of the amorphous TiO2 sol and realize the vis photocatalytic activity as expected. Acknowledgments The work described in this paper was supported by Doctor Foundation (SWU109017) and some other Foundations of Southwest University (SWNU2005010, XDJK2009C097). References [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20]
Fig. 12. Degradation of methylene blue under visible irradiation ( > 420 nm) for used samples.
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An amorphous TiO2 sol sensitized with H2 O2 with the enhancement of photocatalytic activity Jian Zou a,b,∗ , Jiacheng Gao b , Fengyu Xie b a b
School of Chemistry and Chemical Engineering, Southwest University, Chongqing, 400715, PR China College of Materials Science and Engineering, Chongqing University, Chongqing, 400044, PR China
a r t i c l e
i n f o
Article history: Received 16 November 2009 Received in revised form 6 March 2010 Accepted 8 March 2010 Available online 15 March 2010 Keywords: TiO2 sol H2 O2 Sensitizing Visible light Amorphous
a b s t r a c t An amorphous free-impurity TiO2 sol was synthesized only by means of ultrasonic dispersing of Ti(OH)4 precipitation without any peptizing agents. Anatase sol was obtained by hydrothermally treating the amorphous TiO2 sol. Photocatalytic tests showed that the amorphous sol exhibited higher photocatalytic activity for the photodegradation of methylene blue (MB) under visible light and UV irradiation than anatase sol and P25 TiO2 . The photocatalysts were characterized by UV–visible spectroscopy, Raman spectrometer, photoluminescence (PL), FTIR, TEM, XRD, and XPS. The results showed the peroxide titanium complexes were formed on TiO2 sol particles sensitized with H2 O2 . More peroxide complexes and stronger vis absorption were observed for the amorphous TiO2 sol sensitized with H2 O2 .The sensitization of H2 O2 would completely quench PL of the amorphous sol, but little impact on the anatase sol. The surface structure with fewer physisorbed water molecules facilitated the sensitization of H2 O2 to the amorphous TiO2 sol particles, which resulted in the generation of more OH radicals and higher photocatalytic activity under both visible light and UV irradiation. Crown Copyright © 2010 Published by Elsevier B.V. All rights reserved.
1. Introduction It is generally accepted that nanosize titanium dioxide (TiO2 ) is a promising photocatalyst for photocatalytic reactions, own to its low cost, nontoxicity, unique photocatalytic efficiency and high stability under photoirradiation. For nanosize TiO2 , its high activity was primarily governed by the well-known quantum confinement [1]. Some photocatalytic studies have demonstrated that the smaller anatase nanoparticle exhibited higher activity [2,3], and the size of particles could be controlled only by regulating the pH of solution [4]. However, since the small-sized TiO2 nanoparticles facilitate the agglomeration in synthesis and application, its photocatalytic activity usually decreases dramatically or even negligible. In addition, nothing but UV light of the sunlight can be absorbed by TiO2 due to its wide intrinsic band gap. These seriously limit its application and development. It is well known that TiO2 nanoparticles can be dispersed well in TiO2 hydrosol, and the evident quantum size effect was observed for the transparent TiO2 sols [5]. The TiO2 hydrosols could be facilely prepared by using titanium alkoxides and inorganic salts as precursors [5–9]. However, these TiO2 hydrosols contained a signif-
∗ Corresponding author at: School of Chemistry and Chemical Engineering, Southwest University, Chongqing, 400715, PR China. Tel.: +86 023 68252361; fax: +86 023 68252360. E-mail address: [email protected] (J. Zou).
icant amount of impurities which came from the above-mentioned precursors and peptizing agents, such as hydrochloric acid and nitric acid. It was reported that the impurities or low pH values, which were produced by hydrolysis of titanium precursors and during peptization, would retard the oxidation of the substrate [7]. Recently, it was reported that a neutral impurity-free sol could be obtained by introducing some SiO2 sol as stabilizer [10] or using H2 O2 as peptizing agent [11]. Lv et al. [8] reported that P-doped TiO2 prepared by introducing phosphoric acid into TiO2 sol displayed excellent activity. In addition, the crystal structure of TiO2 catalyst can play a crucial role in determining the efficiency of photocatalytic activity. It was well known that anatase was usually more photoactive than rutile, and the higher crystallinty would exhibit higher photocatalytic activity [12], but higher activity was also observed for rutile [13]. The enhancement of photocatalytic activity due to a synergistic effect of anatase and rutile mixed phases was supported by other experimental evidences [14]. In fact, a controllable phase composition of TiO2 could be selectively prepared by changing nucleation mode [13], but an amorphous phase usually accompanied with the crystalline phase [9,15]. It is reported that amorphous TiO2 is nearly inactive due to the facilitated recombination of the photoformed electron and hole [9,16], but one with large surface area and special microstructures is excepted [17]. The hydrothermal treatment could enhance the crystallization [18], but this would result in the precipitating of sols [19]. Sayılkan et al. [20] synthesized a complete anatase TiO2 crystalline phase by hydrothermal process at 200 ◦ C
0925-8388/$ – see front matter. Crown Copyright © 2010 Published by Elsevier B.V. All rights reserved. doi:10.1016/j.jallcom.2010.03.093
J. Zou et al. / Journal of Alloys and Compounds 497 (2010) 420–427
and prepared transparent sols by dispersing these nanopowders ultrasonically in water, but this sol could not degrade methylene blue (MB) due to the reduction–oxidation behavior of MB. The recombination of photogenerated holes and electrons has been regarded as an unfavorable or limiting process in photocatalysis of TiO2 catalyst. As mentioned above, the impurities and amorphous phase in TiO2 sol would depress the photocatalytic activity, and it was almost impossible to avoid those deficiencies during preparation of sol by the methods above. Therefore, a new method for preparing or modifying TiO2 sol was necessary in order to further enhance its activity. Some researches have demonstrated that adding an external electron acceptor such as hydrogen peroxide could enhance the performance of photocatalysis due to reducing the recombination of charge carriers [21,22]. Moreover, hydrogen peroxide molecules can be adsorbed on the TiO2 surface to form surface peroxide complexes, resulting in visible photoresponse [23–25]. In addition, it was also reported that H2 O2 could increase the stability of TiO2 sol [11]. In this study, we prepared the free-impurity amorphous TiO2 and anatase sol by ultrasonically dispersing Ti(OH)4 , and sensitized them with H2 O2 . It was expected that H2 O2 could inhibit the recombination of the exited charge carriers to enhance the activity of TiO2 sol. In addition, it was expected that sols sensitized with H2 O2 could degrade the MB under visible light irradiation. 2. Experimental 2.1. Preparation of TiO2 sol Titanium tetrachloride (Sinopharm Chemical Reagent Co. Ltd., China) was used as precursor without any further purification. The precursor TiCl4 was added dropwise into mixture of deionized water and ice under vigorous stirring to make resulting solution 0.1 M in TiCl4 . The mixture was kept stirring and heating on the agitating heater until the pale blue colloidal dispersion was observed, and the temperature was about 90 ◦ C. The dispersion was cooled to the room temperature. In order to ensure the complete hydrolysis reaction, a 10 wt% NH3 ·H2 O aqueous solution was added dropwise into the transparent sol to form a white precipitate Ti(OH)4 with ultimate aqueous mixture of pH 7–8. Subsequently, the resulting precipitates were separated from the suspension by using filtration and washed repeatedly with deionized water to remove the residual chloride ions, and the final pH value of filtrate agreed with deionized water. The sol containing 0.1 M TiO2 was recovered by dispersing the Ti(OH)4 precipitate into deionized water under ultrasonic condition, and the sol was denoted as AM sol. The AM sol was transferred to a stainless steel Teflon-lined autoclave. The autoclave was sealed and transferred into a 160 ◦ C oven. The sol was aged for 6 h without any agitation and cooled to room temperature at atmosphere. The hydrothermal sol was denoted as HT sol. All sols were uniform, stable and transparent. The sensitizing TiO2 sols were acquired by directly adding H2 O2 into AM and HT sols to make resulting solution 0.3 wt% in H2 O2 . The sols sensitized with H2 O2 were denoted as AM-H2 O2 sol and HT-H2 O2 sol respectively. Interestingly, the sol which was prepared by peptized Ti(OH)4 with inorganic acid [9] would deposit immediately when adding H2 O2 . The xerogels were gained by directly drying the sols at room temperature. 2.2. Characterization The phases of the products were characterized through X-ray diffraction method using CuK␣ radiation ( = 0.15418 nm) in a XD-3 diffractometer (Beijing Pgeneral). Transmission electron microscopy (TEM, Hitachi H7500) was performed at an accelerating voltage of 80 kV for electrons. Fourier transform infrared (FT-IR, Bruker TENSOR 27 FT-IR) spectra of TiO2 powders were obtained using a spectrometer with KBr pellets technique. The Raman spectra were recorded on a Bruker RFS100 spectrometer with a resolution of 4 cm−1 , using 1054 nm light as an excitation source. The precision of the wavenumber was 1 cm−1 . The UV–vis spectra were recorded on a spectrophotometer with an integrating sphere (Shmadzu UV-2550); BaSO4 was used as a reference sample. XPS spectrometer (VG Scientific ESCALAB 250) equipped with two ultra-high vacuum (UHV) chambers measurements were performed and all binding energies were referenced to the C1s peak at 284.6 eV of the surface adventitious carbon. The measurements of the amount of • OH were conducted according to the literature [21,26]. For UV irradiation, a 40 W UV lamp was used. For visible light irradiation, a 60W tungsten halogen lamp (Osram, HALOSPOT 111IRC 48837 ECO FL) with a filter ( > 420 nm) was used. 100 ml sol contained 0.1 wt% TiO2 , 0.01 M NaOH and 3 mM terephthalic acid. Before exposure to light irradiation, the sol was stirred in the dark for 30 min. And then, 10 ml of sol was taken out after 30 min for fluorescence spectroscopic measurements (Hatachi F-4500). The fluores-
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Fig. 1. XRD patterns of xerogel powders.
cence spectra of sols were recorded on fluorescence spectrophotometer (Hatachi F-4500). 2.3. Measurement of photocatalytic activity The photocatalytic activities of the samples were measured by the degradation of MB (20 mg L−1 ). For photodegradation experiments, the required volume of dye solution was added into TiO2 sol to result in 200 ml diluted sol containing 0.1 wt% TiO2 . A 450 W high pressure mercury lamp (Shanghai YaMin) with a filter was used as a visible-light source ( > 420 nm). For UV irradiation, a 40 W UV lamp was used. Before the photocatalytic experiment, the sol was stirred in the dark for 30 min and its concentration was original concentration (Co ) of MB. After a given irradiation time, 10 ml sol was withdrawn, and the concentration (C) of the MB was monitored by measuring the maximum absorbance of MB using the UV–vis spectrum. It must be noted here that no filtration and centrifuging were needed for the TiO2 sol, since it was transparent and do not absorb visible light beyond 500 nm (it was mentioned later) which can be absorbed by MB. But Degussa P-25 TiO2 needed to be removed from its suspension by centrifuging at 5000 rpm after photocatalysis procedure.
3. Results and discussion 3.1. Crystal structure and morphology Fig. 1 showed the X-ray powder diffraction (XRD) patterns of the as-prepared xerogel powders. The XRD patterns indicated that nothing could be observed for AM sol except for a weak peak located at 2 values of 25.3◦ which corresponded to anatase phase. It could be concluded that AM sol remained to be predominantly amorphous titania besides trace amounts of anatase. After hydrothermal treatment, the peaks of HT sol became sharper and more intensive, indicating better crystallization. And all diffraction peaks could be assigned to the anatase phase without any indication of other crystalline byproducts such as rutile or brookite, indicating the formation of predominant anatase TiO2 nanocrystals for HT sol. Generally, the anatase phase is considered to be the most photoactive. Traditionally, it was reported that the pH value [27] and chloride ions [28] would influence the phase composition for hydrothermal synthesis of TiO2 . At basic pH the main or exclusive anatase phase could be prepared, while at acidic pH, rutile formation was favored [27]. And existence of chloride ions also promoted rutile formation [28]. In this study, TiO2 sols were neutral, so anatase TiO2 benefited more from the absence of the residual chloride ions. The crystallite size estimated using the Scherrer equation [20] was about 6 nm for HT sol. The crystal structure of xerogel powders can also be identified by Raman scattering. Typically, for anatase TiO2 , there are six Raman active fundamental modes at 144 cm−1 (Eg ), 197 cm−1 (Eg ), 399 cm−1 (B1g ), 513 cm−1 (A1g ),519 cm−1 (B1g ), and 639 cm−1 (Eg ), respectively [29]. Fig. 2 showed the Raman spectra of HT sol and AM sol. Their spectra corresponded to the typical anatase TiO2 ,
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Fig. 4. IR spectra of as-prepared xerosols. Fig. 2. Raman spectra of xerogel powders (Inset: the expanded spectra for the lowest-frequency Eg mode).
3.2. Composition and chemical structure which indicated that only anatase phase was formed for these asprepared sols. And it was clear that the modes of the HT sol are much stronger than that of the AM sol, which implied that the crystallinity of hydrothermal sol TiO2 was enhanced. These conclusions were consistent with the XRD results shown in Fig. 1.In addition, a blue shift of the lowest-frequency Eg modes for as-prepared sols was also observed from the inset of Fig. 2. For HT sol, the mode shifted to 147 cm−1 . And more blue shift was observed for AM sol. In the case of anatase TiO2 nanoparticles the quantum size confinement effects would result in the blue shift of the low-frequency Eg mode [30]. Kelly et al. [31] also thought that the blue shift of the Eg modes was associated with changes in the particle size, i.e., with a decrease in particle size the peak shifted to the blue, and decreased in intensity. Therefore, the Raman results suggested that the particle size of the AM sol was smaller that of the HT sol. A similar behavior of frequency shift was observed for TiO2 nanoparticles by others [29]. The TEM micrographs of the as-prepared sols were presented in Fig. 3. It was clear that particles for the hydrothermal sol TiO2 had diameters about 6 nm, which was accordant with the results of XRD. For the sol before hydrothermal treatment, the particles size was about 3 nm, which was nearly a moiety as large as that of hydrothermal sol.
Fig. 4 showed the IR spectra of the sols before and after hydrothermal treatment. The xerogels exhibited similar IR spectra. The transmission bands in the range 400–800 cm−1 are characteristic of the formation of an O–Ti–O lattice. The peak at ∼1630 cm−1 results from O–H bending of molecularly physisorbed water [32,33]. The broad adsorption band observed at 3000–3700 cm−1 can be assigned to the stretch region of the surface hydroxyl groups and molecularly chemisorbed water [32,33]. The stronger O–H peak which resulted from the physisorbed water molecules in HT sol was observed, but an opposite result was also observed for stretch strength of the surface hydroxyl groups and chemisorbed water, compared with the sol before hydrothermal treatment. These results indicated that a higher content of physisorbed water was on HT sol particles, but more surface hydroxyl groups and chemisorbed water were on AM sol TiO2 . The Raman spectra in the range of 700–1200 cm−1 were recorded in Fig. 5. The Raman peaks of hydrogen peroxide in the different phase were detected at 884 cm−1 and at 874 cm−1 , corresponding to the O–O symmetric stretching vibration of H2 O2 in the liquid phase at ambient conditions and in the supercritical phase, respectively [34]. In the present work, for the as-prepared sols sensitized with hydrogen peroxide, a new vibration band at 840–920 cm−1 was observed, which indicated that peroxo groups
Fig. 3. TEM maps of (a) AM sol and (b) HT sol.
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Fig. 5. Raman spectra of HT sol TiO2 and AM sol TiO2 treated and untreated with H2 O2 .
were on the surface of sols particles. Comparing with the Raman spectra of the sols treated with H2 O2 , we also noticed that the O–O stretching vibration for AM-H2 O2 sol particles was stronger than that of HT-H2 O2 sol. This was likely to be result from the formation of fewer surface peroxide complexes on the surface of the hydrothermal sol particles, due to more physisorbed water (Fig. 4). The physisorbed water molecules were bound by weaker hydrogen bond with the hydroxyl groups of TiO2 surface to form a multilayer, but the chemisorbed water molecules directly interacted with the TiO2 surface to form a monolayer [33,35]. In presence of H2 O2 , the –OOH groups of H2 O2 would replace –OH groups on the surface
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of titania to form surface complexes [23–25]. And Henderson et al. [36] thought the 2nd layer of water on the surface of TiO2 would block access of O2 molecules to the surface OH groups. We deduced this blocking effect was the same with H2 O2 molecules. The XPS technique can monitor the electron binding energy of sites with a few nanometer of materials surface, from which some information on surface composition and chemical structures would be inferred. We had examined the XPS region of O1s and Ti2p for asprepared sols. Fig. 6 showed the XPS spectra of O1s and Ti2p for used xerogels treated with H2 O2 .For the sols before hydrothermal treatment (Fig. 6a), O1s band could be separated into three peaks, located at 529.9, 531.3 and 532.8 eV respectively, which could be attributed to Ti–OH at 531.9 eV, Ti–O at 530.1 eV and peroxo groups at 532.8 eV [24,25]. Compared with AM-H2 O2 sol, the band of HTH2 O2 sol only could be separated into two peaks (Fig. 6b), and the absence of the band corresponding to peroxo groups. In fact, both of sols were sensitized with H2 O2 , and peroxo groups were also detected for them in Fig. 5, but it simultaneously showed that fewer complexes were on the surface of HT-H2 O2 sol. So, peroxo group peak in O1s XPS spectra was likely to be muffled by the OH peak. Herein, the absence of peroxo group peak in XPS of O1s reconfirmed the conclusion that the formation of surface peroxide complexes was restrained due to the overmuch physisorbed water on the HTH2 O2 sol particles. The Ti2p band for hydrothermal sol particles consists of two peaks of Ti2p3/2 and Ti2p1/2 located at 458.3 and 463.9 eV (Fig. 6c) [25], respectively. However, the Ti2p3/2 peak for the sol particles before hydrothermal treatment shifted to higher binding energy (BE). This shift in binding energy may be attributed to the change in the chemical environment. The similar shift of the Ti2p3/2 BE for TiO2 /SiO2 composite oxide after adding H2 O2 was also observed due to the formation of Ti–O–O–Si peroxo moieties [37]. It was
Fig. 6. High resolution XPS spectra of O1s , Ti 2p in HT-H2 O2 sol TiO2 and AM-H2 O2 sol TiO2 .
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reported that the formation of more peroxide complexes would result in larger shift of the Ti2p3/2 BE for sulfated TiO2 sensitized with H2 O2 [25]. Therefore, the more shift of the Ti2p3/2 BE for the AM-H2 O2 sol could be concluded to result from the formation of more peroxide complexes, which was consistent with the results above. 3.3. UV–vis spectra and PL The UV–vis absorption spectra of the as-prepared sols were illustrated in Fig. 7. All sols were adjusted to the concentration of 1.0 g L−1 TiO2 in aqueous solution for spectrum measurement. It could be seen that the no absorption in the visible region (wavelength > 420 nm) was observed but the strong absorption in the ultraviolet range for sols untreated with H2 O2 . It meant both sols before treatment with H2 O2 were highly transparent and could not drive the direct photocatalysis under visible light irradiation. The optical absorption edge was estimated at 347 nm for the HTH2 O2 sol and 337 nm for the AM-H2 O2 sol, and the corresponding band-gap energy was identified as 3.57 and 3.68 eV, respectively. The optical absorption edge appeared an obvious blue shift for the above sols relative to 387 nm (about 3.2 eV) [1] of the block anatase TiO2 .This was often ascribed to the quantum size effect of sol nanoparticles [5] due to their high dispersion and small size. After sensitized with hydrogen peroxide, both sols could absorb visible light (wavelength > 420 nm), but AM-H2 O2 sol displayed much stronger visible absorption than HT-H2 O2 sol. Generally, the change of energy band structure can result in vis absorption [38,39]. The electronic structure could be investigated by value band (VB) XPS spectrum [38,39]. The value band XPS spectra of sols treated with H2 O2 were presented in the inset of Fig. 7. It could be seen that the value band edge of hydrothermal sol shifted slightly to Fermi level, which meant narrower energy gap for the hydrothermal sol. This was incompatible with optical absorption of sols treated with H2 O2 . Obviously, the shift of VB edge could not explain HT-H2 O2 sol had poorer visible absorption. It was reported that the peroxide complexes on the surface of TiO2 could absorb visible light [23–25]. So, it could be deduced that the stronger vis absorption came form more peroxide complexes on the AM-H2 O2 sol, which had been confirmed by the results of Raman and XPS. Moreover, the hydrothermal treatment would weaken the quantum size effect of sols due to increase in particle size, which may be the reason that the VB edge shifted to Fermi level. Photoluminescence (PL) spectroscopy has been widely conducted to study charge carrier trapping, the role of surface states in photocatalysis and surface property for TiO2 particles because PL
Fig. 8. Photoluminescence of used sols treated and untreated with H2 O2 .
emission originate from the recombination of electron–hole pairs in the surface and bulk of TiO2 [40]. The photoluminescence spectra for our samples were showed in Fig. 8. For the hydrothermal sols, the stronger emission band at 390 nm is assigned to band-to-band transition, and the weaker emission band at 465 nm is related to the surface trap state [41]. Although the AM sols did not display a distinct peak around 390 nm, a weaker broad emission band around 350–600 nm was observed. This indicated the tiny crystalline TiO2 must be present in AM sol TiO2 because the amorphous TiO2 do not show fluorescence [42], which was also supported by above results of XRD and Raman. Otherwise, it could be seen that the PL spectrum of hydrothermal sol was much more intense than that of the sol before hydrothermal treatment, which could be attributed to the inferior crystallization (Figs. 1 and 2) and the larger density of surface states or defects for smaller AM sol TiO2 particles (Fig. 3) where the nonradiative emissions predominated [43]. Additionally, it could also be seen that the treatment with H2 O2 could reduce the PL intensity of sols from Fig. 8. PL emission of TiO2 is usually affected by surface chemistry. Deposition of Pt on TiO2 [44], adsorbed H2 O and N2 [45], organic modification [46] and doping of nitrogen [47] would quenched or weaken the PL emission of TiO2 . For sol sensitized with H2 O2 , its surface chemistry would be changed due to the formation of peroxide complexes on the TiO2 . PL emissions, attributed to indirect band-gap emission and surface defects, were nearly quenched for AM-H2 O2 sol TiO2 , which indicated that the effect of peroxide complexes on sols was not limited merely to surface. And Shu et al. [37] also reported that PL emission could be quenched for TiO2 /SiO2 composites treated with H2 O2 . However, an obvious difference for hydrothermal sols could be observed. The PL emissions were much less affected, especially for the lower energy emission associated with defect sites [42], suggesting that effect of H2 O2 on surface chemistry of HT sols TiO2 was negligible. Certainly, the decrease of effective interface between TiO2 and H2 O2 , resulting from the increase of particle size, would reduce peroxide complexes on the hydrothermal sol particles to weaken the effect of H2 O2 .However, the fact that the negligible decrease of the emission at 450–500 nm, being directly affected by surface chemistry, reconfirmed that overmuch physisorbed water molecules would block access of H2 O2 molecules to TiO2 surface. This was consistent with the results of Raman and XPS. 3.4. Photocatalytic activity
Fig. 7. Diffusion reflection absorption of used sols treated and untreated with H2 O2 (Inset: The values band XPS of as-prepared sol treated with H2 O2 ).
Photocatalytic activity tests were conducted by the degradation of methylene blue (MB) in aqueous solution under ultraviolet and visible light irradiation. Fig. 9 presented the photocatalytic results of the as-prepared sols in contrast to P25 under ultraviolet irradi-
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Table 1 The maximum absorbance (Co value) of MB in different sol before irradiation.
Fig. 9. Degradation of methylene blue under ultraviolet radiation for used samples.
ation. As could be seen, the photocatalytic degradation of MB in the hydrothermal sol behaved better than in the AM sol. About 60% discoloration for the hydrothermal sol was observed, but only about 13% degradation for the AM sol in the same time, which even was lower than that of pure H2 O2 . Interestingly, an opposite result was observed after the sols were sensitized with hydrogen peroxide. About 90% discoloration was acquired for AM-H2 O2 sols, which was nearly 7 times as much discoloration as ones before treatment with H2 O2 . However, the discoloration of HT-H2 O2 sols decreased to 38%, only 60% of sols before treatment with H2 O2 . Anatase was usually deemed to be the most photocatalytically active TiO2 . And activity of amorphous TiO2 was nearly negligible due to the facilitated recombination of the photoformed electron and hole at the traps on the surface and in the bulk of the particles [9,16]. PL emission is the result of the recombination of excited electrons and holes, and generally the higher PL intensity indicates a higher recombination rate of these electrons and holes under light irradiation. However, for amorphous TiO2 , the nonradiative emission predominated [43]. Therefore, PL emission of AM sol was not the result of the recombination of all excited electrons and holes. It was possible that much more recombination occurred for AM sol TiO2 .Therefore, the higher activity of HT sol can be ascribed to better anatase phase, as discussed above based on Figs. 1 and 2. Hydrogen peroxide may be split photocatalytically to produce hydroxyl radicals directly under UV light irradiation [22,48]. Those hydroxyl radicals could degrade MB as showed in Fig. 9, but if the decomposition of methylene blue was only from H2 O2 itself, then the efficiency of the decomposition using a different sol photocatalyst might not vary so much as that shown in Fig. 9. On the other hand, hydrogen peroxide is an electron acceptor, which would lower the electron–hole recombination rate to increase the photocatalyzed rate [21,22]. However, the enhancement of H2 O2 for photacatalytic activity of TiO2 had been in some cases controversial, and it appeared strongly dependent on substrate type and on various experimental parameters. Some researchers [24,27] found the enhancement practically depended on crystalline phase of catalysts, and rutile was preferable to anatase. Hirakawa et al. [26] suggested that the adsorption structure of H2 O2 on the rutile TiO2 surface was preferable to produce OH• . Augugliaro et al. [49] and Ban et al. [50] found the presence of hydrogen peroxide even did not affect appreciably the photodegradation rate of organic dye, and Chu and Wong [48] also observed that H2 O2 -assisted photocatalysis of dicamba was not enhanced under 350 nm UV irradiation. In addition, the adsorption of MB on TiO2 also can influence the degradation of dye. Table 1 showed the Co values of MB in different sols. However, as could be seen, a less than 10% difference of Co values indicated the effect was extremely limited.
Sample
AM sol
HT sol
AM-H2 O2 sol
HT-H2 O2 sol
Absorbance (a.u.)
1.561
1.672
1.533
1.642
In our study, we thought that an opposite effect of H2 O2 for both used sols was predominantly dependent on the surface chemistry of sol particles. For the HT-H2 O2 sol, as discussed above, more physisorbed water molecules would baffle the direct interaction of H2 O2 to TiO2 , which would result in less capture of electron. As a result, the separate effect of electron and hole pairs was negligible for HT-H2 O2 sol, as indicated in the result of PL in Fig. 8. For the AM-H2 O2 sol, the excited electrons would be recombined in the impurity states without H2 O2 , resulting in low activity as ascribed in Fig. 9.With H2 O2 , the recombination (even including nonradiative one) of charge carriers was remarkably restrained by H2 O2 , with the exhibition of quenching of PL emission due to fewer physisorbed water molecules on sol particles. In addition, the excess H2 O2 molecules would scavenge the valuable hydroxyl radicals (HO• ) to form much weaker oxidant HO2 • [48]. Therefore, the balance of two photochemistry action of H2 O2 for TiO2 sol systems would result in the increase of hydroxyl radicals for AM-H2 O2 sol but the decrease for HT-H2 O2 sol. The concentration of HO• could be estimated by terephthalic acid fluorescence (FL) probe method [21,26]. Since the FL spectra were identical to that of TAOH generated from TA by the reaction with OH• [21], the relative content of OH• formed in TiO2 sol was estimated from that of TAOH. Fig. 10 showed FL spectra of the used sols with 3 Mm TA after irradiated for 30 min.It could be seen that H2 O2 decreased the FL intensity for the HT-H2 O2 sol, but a prominent increase for the AM-H2 O2 sol was observed, which indicated that H2 O2 promoted authentically the formation of hydroxyl radicals for AM-H2 O2 sol as anticipated above. This was consistent with the results of photocatalytic degradation of MB. In addition, Sayılkan et al. [20] and Doushita and Kawahara [51] found the faded MB in TiO2 suspensions under illumination could reverted to its original blue color while stored in the dark for some hours. This reversible color change was also observed by others [7,52]. This reversible chromic phenomenon could be eliminated by enhancing oxidation–reduction potential of dye to the negative direction [20] or conducting simultaneously air-bubbling into MB/TiO2 suspensions under illumination [7]. Fig. 11 indicated the absorption spectrum changes of MB after UV irradiation and during storage in the dark. An absorption peak in AM-H2 O2 sol appeared
Fig. 10. Fluorescence spectra obtained for used sols containing 3 mM terephthalic acid under ultraviolet irradiation for 30 min. (Exited wavelength: 312 nm; Inset: under visible irradiation for 30 min).
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Fig. 11. The UV–vis absorption spectra of AM-H2 O2 sol containing MB at different time under UV irradiation and during storage in the dark (The spectra of (a) 0 min, (b) 120 min, (c) suspension (b) in the dark for 12 h).
at around 580 nm which shifted 85 nm relative to the peak of MB at 665 nm. This shift was also observed for TiO2 /SiO2 composites [7,51]. This absorption peak vanished nearly under UV irradiation for 120 min. After stored in the dark for 12 h, the peak did not only revert to its original shape but disappeared completely, indicating that no reversible discoloration of MB dye occurred to AM-H2 O2 sol without air-bubbling. The reason of the reversible chromic phenomenon for MB was that the colorless MB molecule reduced by photogenerated electrons would be reoxidized by oxygen back to its original blue color [7,20,51]. For AM-H2 O2 sol, as described in Fig. 10, an amount of OH radicals were generated. As a strong oxidant, OH radicals could photodegraded directly MB by oxidation, which avoided the blue MB to be reduced to colorless MB. According to Fig. 7, the sols treated with hydrogen peroxide could absorb visible light, which meant that they could drive the direct photocatalysis under visible light irradiation. Fig. 12 showed the results of photocatalytic degradation of MB aqueous solution for the used samples under visible light irradiation ( > 420 nm). For the reference of P25 without H2 O2 , only 18% discoloration was detected in 120 min, which hinted that the effect of dye sensitization [53] from MB was very limited. Similarly, the highest activity was observed for the AM-H2 O2 sol under visible light irradiation, especially in the primary stage of photodegradation, 75% discoloration for the AM-H2 O2 sol, but only 23% discoloration for the HT-H2 O2 sol within 40 min. In addition, Fig. 12 also showed H2 O2 without TiO2 sol was unable to degrade MB under visible irradia-
tion, which indicated the visible degradation of MB was not from H2 O2 .Besides, no reversible chromic phenomenon was observed for AM-H2 O2 under visible light irradiation. For TiO2 sensitized with H2 O2 , the conduction band electrons, which would react with the adsorbed H2 O2 on the TiO2 surface to generate the OH radical, only came from the excited surface peroxide complexes, not valence band of TiO2 under visible light irradiation [23]. Hereby, amounts of hydroxyl radicals were practically dependent on the surface peroxide complexes. As discussed above based on the results of IR, Raman, XPS and UV–vis, fewer peroxide complexes in response to visible light were formed on HT-H2 O2 sol particles owing to more physisorbed water molecules on their surface, resulting in inferior vis absorption. This was a great disadvantage to the generation of hydroxyl radicals for HT-H2 O2 sol TiO2 under visible light irradiation. It could be seen from the inset of Fig. 10 that the FL intensity at around 426 nm was lower for HTH2 O2 sol TiO2 than that for AM-H2 O2 sol TiO2 under visible light irradiation for 30 min.This indicated that fewer hydroxyl radicals were generated in the HT-H2 O2 sol. 4. Conclusions The stable, transparent TiO2 sol was prepared only by means of ultrasonic dispersion of Ti(OH)4 precipitation without any peptizing agents. The sol TiO2 was predominantly amorphous titania besides trace amounts of anatase. The sol with better crystallization was acquired by hydrothermal treatment. After sensitized with H2 O2 , the sols could absorb visible light due to formation of peroxide complexes on TiO2 surface. The hydrothermal sol treated with H2 O2 exhibited the inferior vis absorption and photoactivity under both UV and visible light irradiation. The hydrothermal treatment would increase the physisorbed water on TiO2 , resulting in formation of fewer peroxide complexes and hydroxyl radicals, which decreased photoactivity of the hydrothermal sol. Herein, the results presented here demonstrated that the sensitization of H2 O2 could enhance efficiently photoactivity of the amorphous TiO2 sol and realize the vis photocatalytic activity as expected. Acknowledgments The work described in this paper was supported by Doctor Foundation (SWU109017) and some other Foundations of Southwest University (SWNU2005010, XDJK2009C097). References [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20]
Fig. 12. Degradation of methylene blue under visible irradiation ( > 420 nm) for used samples.
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