Evaluating Intrinsic Photocatalytic Activities of Anatase and Rutile TiO

Oct 18, 2010 - liquid-solid interface are the conduction band electron (ecb. -) and valence band hole (hvb. +), generated on TiO2 upon UV light excita...
1 downloads 0 Views 1MB Size
Download PDF
J. Phys. Chem. C 2010, 114, 18911–18918

18911

Evaluating Intrinsic Photocatalytic Activities of Anatase and Rutile TiO2 for Organic Degradation in Water Qiong Sun and Yiming Xu* State Key Laboratory of Silicon Materials and Department of Chemistry, Zhejiang UniVersity, Hangzhou, Zhejiang 310027, China ReceiVed: May 25, 2010; ReVised Manuscript ReceiVed: September 27, 2010

Various factors that influence the photocatalytic activity of TiO2 for organic degradation in aerated aqueous solution have been reported. For instance, anatase is considered to be much more active than rutile. However, no attention has been paid to the difference in their sorption capacities toward O2 in water, which might be critical to the activity determination. In this work, Ag+ has been used as an electron scavenger, for the photocatalytic degradation of 4-chlorophenol in the N2-purged aqueous suspension of TiO2 under UV light. Three different TiO2 samples in the crystal forms of anatase and rutile, containing micro- and mesopores, were prepared, followed by sintering at different temperatures (Ts). The initial rate of 4-chlorophenol photodegradation, per the initial amount of Ag+ adsorbed, increased exponentially with Ts. Such Ts-dependent normalized rates were observed with three differently prepared TiO2 samples, and three curves were almost overlapped. Comparatively, in the aerated aqueous suspension of TiO2, the initial rate of 4-chlorophenol photodegradation, per surface area of the catalyst, first increased and then decreased with Ts, the trend similar to those widely reported. Moreover, from the literature data of water photosplitting over TiO2, the initial rate of O2 evolution, per the initial amount of Ag+ adsorbed, increased exponentially with Ts. Evidence clearly shows that with the same amount of electron scavenger on the catalyst surfaces, anatase and rutile actually have a similar photocatalytic activity at a given Ts, for organic degradation or water oxidation. It is recommended that to evaluate the photocatalytic activities of different TiO2 samples in an aerated aqueous solution, the difference in O2 adsorption needs to be taken into account. Introduction Titanium dioxide is a powerful and robust photocatalyst for the degradation of various organic pollutants in aerated aqueous solution1,2 and has received considerable attention for development as a potential technology for water treatment. However, the overall efficiency achieved so far with TiO2-based systems is not sufficiently high to enable practical applications. This low efficiency is mainly due to fast recombination of charge carriers on irradiated TiO2 and thus low quantum yield in the generation of reactive species for organic degradation. For example, phenol oxidation over a benchmark TiO2 (Degussa P25), in an acidic aerated aqueous suspension, only has a quantum yield of 0.14 at 365 nm.3 Many factors that influence the efficiency of organic photodegradation over TiO2 have been reported during the past 30 years. The main focus has been the physical properties of TiO2, including crystal phases, crystal facets, crystallinity, particle size, surface area, porosity, and morphology.5 In general, anatase is considered to be much more photoactive than rutile,6-8 while a mixture of anatase and rutile with a sintered interface, as like Degussa P25, is claimed to be more active than anatase.9-12 Brookite TiO2 and TiO2(B) as a photocatalyst have been rarely studied, probably because of the difficulty in synthesis.13-15 However, it is not easy to make a reliable correlation between the structure and photoactivity of TiO2, because the solid physical properties often affect one another. More recently there has been controversy between several research groups about the model substrate used for assessment of the photocatalyst * To whom correspondence should be addressed. Phone: +86-57187952410. Fax: +86-571-87951895. E-mail: [email protected].

activity.16-18 Different conclusions about the trend in photoactivity among various TiO2 samples may be reached with different model substrates or reactions.8,18 On the other hand, for water splitting into O2 over irradiated TiO2 in the presence of Ag(I) or Fe(III), rutile appears to be more photoactive than anatase,7,8 contrary to that observed for organic degradation. In principle, the primary species that initiate redox reactions at a liquid-solid interface are the conduction band electron (ecb-) and valence band hole (hvb+), generated on TiO2 upon UV light excitation. Water oxidation is a four-hole process,19 while organic degradation occurs mainly through hvb+, or through hydroxyl radical (•OH), originated from the hvb+ reaction with surface OH-/H2O on TiO2.1-4 In this regard, two separate reactions of organic degradation and water oxidation should result in the same trend in photoactivity among different TiO2 samples, because each catalyst has a definite intrinsic photoactivity. However, this is not the case in practice, as shown above. Something must have masked the determination of the real reaction rate. It has been widely observed that the photocatalytic degradation of organic pollutants over TiO2 only proceeds in the presence of O2, and that the reaction rate increases toward a plateau with the partial pressure of O2 in a gas phase.20,21 This is consistent with the recognized mechanism that ecb- and hvb+ are formed in a pair, followed either by recombination to release heat, or migration into the surface to react with a suitable electron acceptor and donor.1-4 Then, an increase in the surface concentration of the adsorbed O2 (and organic substrate) on TiO2 would result in enhancement in the rate of interfacial charge transfer and thus in the rate of organic degradation. However,

10.1021/jp104762h  2010 American Chemical Society Published on Web 10/18/2010

18912

J. Phys. Chem. C, Vol. 114, No. 44, 2010

a routine assessment of the relative photoactivity of TiO2 among various samples is usually made in an aqueous solution exposed to air or purged with O2, by comparing the rate of model substrate loss or its product formation. Although O2 entering into the reactor is constant in concentration, the amount of O2 adsorbed on each catalyst might be significantly different. This might result in a misjudgement of the relative photoactivity of TiO2 among different samples. A catalyst that has a high intrinsic photoactivity, but a low capacity for uptake of O2 from water, may display a low apparent photoactivity for organic degradation in an aerated aqueous medium. For a whole photocatalytic reaction, the rates of ecb- and hvb+ reactions should be equivalent and balanced.22-24 Therefore, assessment of the intrinsic photoactivity of TiO2 among different samples should be made under a unique condition, such as the same surface concentration of the adsorbed O2 on each sample. However, to our knowledge, there have been no studies of this issue reported in literature, probably because of the difficulty in measuring the O2 adsorption onto TiO2 from aqueous medium. In a previous study, we have shown that for phenol photodegradation in an aerated aqueous suspension, rutile is indeed much less photoactive than anatase, but its activity becomes similar to or higher than that of anatase on the addition of AgNO3.5g This implies that removal of ecbfrom irradiated TiO2 is very important to organic oxidation, and the reaction rates have to be compared under an appropriate condition. In this work, Ag+ has been used as an electron scavenger, and the photocatalytic degradation of 4-chlorophenol (4-CP) as a model reaction. The amount of Ag+ adsorbed on TiO2 in the aqueous phase can be quantitatively measured, so that the photocatalytic activities of different TiO2 samples can be compared at the same level of electron acceptor adsorbed on the catalysts. It has been reported that the initial rate of Ag+ photoreduction, in the Ar-purged aqueous suspension of TiO2, is linearly proportional to the initial amount of its adsorption.5c,25As a model substrate, 4-CP has been extensively studied.26-30 In a neutral aqueous solution, 4-CP does not have absorption toward UV light at λ > 300 nm, and it also hardly adsorbs on TiO2. This would make it convenient to perform the TiO2-photocatalyzed reaction at a fixed concentration of 4-CP, without the need to examine the effect of 4-CP photolysis and its initial concentration on the activity determination. Experiments were carried out in a N2-purged aqueous suspension of TiO2 and AgNO3 under UV light (λ g 320 nm). For a comparison, parallel experiments in the aerated aqueous suspension of TiO2 without AgNO3 were also performed. In order to examine the effect of the TiO2 physical properties, three series of samples with different crystal phases (anatase and rutile) and porosities (micro- and mesopores) were prepared, followed by sintering in air at various temperatures. As expected, the reactions in the presence of AgNO3 under N2 are remarkably different from those under air in the absence of AgNO3. However, after the reaction rate is normalized by the amount of Ag+ adsorbed on TiO2, the result shows that anatase and rutile actually have a similar photocatalytic activity for 4-CP degradation at a given sintering temperature. It is therefore proposed that for evaluating the relative photocatalytic activity of TiO2 among different samples, variation in the catalyst sorption capacity toward O2 in aerated aqueous solution needs to be taken into account. Experimental Section Materials. Three series of TiO2 samples were prepared by following literature procedures.5g,31,32 In the text, SAT and SRT represent the as-synthesized microporous TiO2 in the crystal

Sun and Xu forms of anatase and rutile, respectively, while MAT represents the as-prepared mesoporous TiO2 in the form of anatase. Briefly, SAT was synthesized in an iced bath by hydrolysis of TiCl4 in aqueous solution containing (NH4)2SO4 and HCl. The suspension was refluxed for 2 h, cooled to room temperature, and mixed with aqueous ammonia until pH 6. SRT was prepared similarly in an iced bath, but TiCl4 was added into the aqueous solution of HCl. After the suspension was heated at 60 °C for 2 h, it was stored overnight. The white particles were then separated and washed thoroughly with distilled water until the filtrate was free of chloride ions. MAT was synthesized through a hydrothermal route, using titanium butoxide as a precursor, and cetyltrimethylammonium bromide (CTMA) as a template.32 After the suspension was autoclaved at 120 °C for 36 h, the particles were separated and washed with water. The dried powder was further treated with an aqueous solution of ethylenediamine (EN) at 95 °C for 36 h, followed by a thorough washing with distilled water. All the powders were first dried at 110 °C, ground, and sieved with a 200 mesh sieve. Finally, the powder was sintered in air at various temperatures for 3 h. Characterization. Powder X-ray diffraction (XRD) patterns were recorded on a D/max-2550/PC diffractometer (Rigaku) using a Cu KR as X-radiation source (40 kV, 100 mA). The content of anatase phase was estimated with the integrated intensities of anatase (101) and rutile (110). The crystallite size was calculated using the Scherrer equation. The adsorptiondesorption isotherm of N2 was measured at 77 K on a Micromeritics ASAP2020 apparatus. The Brunauer-Emmett-Teller (BET) surface area was calculated by the 12-point method, and total pore volume was estimated at a relative pressure of 0.99. The BJH pore size distribution was derived from the desorption branch. Thermogravimetric analysis (TGA) was performed on a PE PYRIS I apparatus at a rate of 10.00 °C/min. Diffuse reflectance spectra were recorded on a Shimadzu UV-2550 spectrophotometer using BaSO4 as a reference. The band gap energy of TiO2 for allowed direct transition was estimated by following literature procedures.33 Photocatalysis. Photocatalytic reactions were carried out in a thermostatted Pyrex-glass reactor under constant magnetic stirring.5g The aqueous suspensions containing all necessary components were first stirred in the dark for 1 h and then irradiated with a prewarmed high pressure mercury lamp (375 W, Shanghai Yamin). Except when stated otherwise, the experimental condition was set at 1.0 g/L TiO2, 0.23 mM 4-CP, and 1.0 mM AgNO3. For the reaction in the absence of O2, the suspension was continuously purged with N2 (99.99%). At given intervals, small aliquots were withdrawn and filtered through a membrane (pore size 0.22 µm). Organic compounds in the filtrate were immediately analyzed by high performance liquid chromatograph (HPLC) on a Dionex P680 (Apollo C18 reverse column, and 3:2 CH3OH-H2O eluent at 1.0 mL/min). Silver ion concentration was measured on an Agilent 8451 spectrometer, by recording the absorbance at 467 nm of its complex with p-dimethylaminobenzal rhodanine.34 Results and Discussion Characterization. The physical properties of TiO2 often vary with its history of preparation and post-treatment. Therefore, it is necessary first to characterize the samples synthesized in the present study. Figure 1 shows the XRD patterns of representative samples. All the as-prepared powders exhibit a weak diffraction pattern, due to either fine crystallites or some amorphorous solids present in the sample. The crystalline phases in SAT and MAT are identified as anatase (PDF no. 12-1272), while the phase in

Anatase and Rutile TiO2 for Organic Degradation

Figure 1. XRD patterns of SAT (a, b), MAT (c, d), and SRT (e, f), sintered at 110 °C (a, c, e), and 600 °C (b, d, f), respectively.

SRT is rutile (PDF no. 12-1276). After calcination in air for 3 h, all diffraction peaks become more intensive and acute, indicative of crystallization and particle growth. Meanwhile, SAT and MAT samples begin to undergo phase transformation from anatase to rutile, at a sintering temperature (Ts) higher than 400 °C (Figure S1, Supporting Information). At the temperature between 500 and 800 °C, SAT and MAT are a mixture of anatase and rutile, and the content of anatase in the samples decreases with Ts (Figure S2, Supporting Information). Moreover, all the samples (SAT, MAT, and SRT) show a progressive increase in the primary particle size with Ts. These changes of phase composition and crystallite size as a function of Ts are commonly observed with synthetic TiO2 in literature. Figure 2 shows the isotherms of N2 adsorption-desorption on SAT, SRT, and MAT at 77 K. All the isotherms are Type VI, characteristic of mesoporous structure. However, as Ts increases, the hysteresis loop gradually shifts toward the high pressure side, indicative of an increase in pore size. With each sample, the distribution of the BJH pore size is wide, but both the average pore size and the peak width at half height increase with Ts (Figure S3, Supporting Information). For example, the average pore sizes in diameter for SAT, MAT, and SRT are increased from 4.74, 5.18, and 8.98 nm at 110 °C, to 30.79, 33.75, and 39.94 nm at 800 °C, respectively. Moreover, a t-plot analysis shows that SAT and SRT also contains micropores with a volume decreasing with Ts, while such a micropore is nearly absent in MAT (Figure S4, Supporting Information). Results indicate that mesopores are formed by the sticking of primary particles. Upon thermal treatment, small micropores collapse, while densification among intra- and interagglomerates leads to a constitution of mesopores with different shapes and pore sizes. The corresponding BET surface area (Asp) and total pore volume (Vp) are summarized in Figure 3. Because the surface area measured by N2 is largely external to the agglomerates, a

J. Phys. Chem. C, Vol. 114, No. 44, 2010 18913

Figure 3. BET surface area (solid symbols) and total pore volume (open symbols) of (a) SAT, (b) SRT, and (c) MAT, sintered at different temperatures. Data were from N2 adsorption isotherms at 77 K.

progressive decrease in Asp with Ts indicates a thermal-induced agglomeration. However, the values of Vp change differently with Ts for the three series of the samples. SRT and MAT show a gradual decrease in Vp with Ts, but SAT displays a volcano curve. Since the pore volume is mainly located within the agglomerates, this might be due to the fact that SAT has a fine primary particle, in relative to SRT and MAT (Figure S2, Supporting Information), so that the SAT particles would easily join together to constitute a mesoporous network upon thermal treatment. However, as crystallization occurs, some mesopores with a narrow size may collapse. As a result, the capillary condensation effect is reduced, and the total pore volume is decreased. The balance between mesopore reconstruction and damage renders SAT a maximal pore volume at 600 °C. The as-prepared MAT has a relatively well-defined mesopore, as demonstrated by its pore size distribution (Figure S3, Supporting Information). Upon thermal treatment at 400 °C, the Vp value of MAT is only decreased by 7%, probably due to the inhibition effect of CTMA and EN used during its synthesis and posttreatment.32 Thermal analysis shows that residual organics in the as-prepared MAT are less than 5%, and they burn off at about 400 °C (Figures S5, Supporting Information). Figure 4 shows the UV diffuse reflectance spectrum of TiO2, expressed in Kubelka-Munk units (FR). The intensitive absorption band in the UV region is ascribed to charge transfer from O2- to Ti4+. As Ts increases from 400 to 800 °C, the spectra of SAT and MAT gradually shift to a longer wavelength, while the spectral onset of SRT remains nearly unchanged. This might be due to the fact that rutile has a lower band gap energy (3.0 eV) than that of anatase (3.2 eV), and that the phase content of rutile in SAT and MAT increases with Ts. To confirm this, the band gap energy (Eg) was estimated by following literature methods.33,35 For each series of the samples, the estimated value of Eg decreases gradually with Ts. For instance, as Ts is increased

Figure 2. Adsorption-desorption isotherms of N2 at 77 K on (A) SAT, (B) SRT, and (C) MAT, sintered at 110 °C (O), 300 °C (b), 400 °C (0), 500 °C (9), and 700 °C (2), respectively.

18914

J. Phys. Chem. C, Vol. 114, No. 44, 2010

Sun and Xu

Figure 4. Diffuse reflectance spectra of (A) SAT, (B) SRT, and (C) MAT, sintered at the temperature as indicated in the legend. The Kubelka-Munk unit is used, FR ) (1 - R)2/2R, where R is the reflectance.

Figure 5. Time profiles of 4-CP (solid bars) and Ag+ (open bars) under different conditions: (a) in the aerated aqueous suspension of TiO2, (b) in the N2-purged aqueous suspension of TiO2 and AgNO3, (c) in the N2-purged aqueous suspension of TiO2, and (d) in the N2-purged aqueous solution of AgNO3.

from 110 to 800 °C, the value of Eg is decreased from 3.18 to 3.02 eV for SAT, from 3.29 to 3.02 eV for MAT, and from 3.10 to 3.02 eV for SRT. Moreover, the spectra in Figure 4 also show a gradual decrease in absorbance at λ < 330 nm with Ts. The reason for that is not known, but it might be due to the particle growth that leads to enhancement in light scattering. Effect of AgNO3 on the Reaction Process. Photocatalytic reaction was carried out using 4-chlorophenol (4-CP) as a model substrate, and SAT sintered at 600 °C as a catalyst. Figure 5 shows the time profiles of 4-CP loss under different conditions. We see that the rate of 4-CP photodegradation, in the N2-purged aqueous suspension of TiO2 and AgNO3, is much faster than that in the aerated aqueous suspension of TiO2. This is in agreement with the fact that as an electron scavenger, Ag+ has a higher driving force than O2. The standard redox potentials for the Ag+/Ag and O2/HO2• couples in aqueous solution are 0.799 and -0.05 V versus NHE,36 respectively, while the potential of TiO2 electrons is about -0.12 V.4 An improvement in charge carrier separation would result in enhancement in the rate of 4-CP oxidation by hvb+ or •OH. Photoreduction of Ag+ on TiO2 was demonstrated by the decrease in Ag+ concentration with irradiation time, by a color change from white to gray, and by silver crystals formed in the suspension (Figure S6, Supporting Information). Control experiments in the solution of AgNO3, or in the N2-purged suspension of TiO2, showed that there was negligible (photo)degradation of 4-CP (Figure 5). Thus, the observed degradation of 4-CP is due to the reaction initiated by TiO2 photocatalysis, for which O2 or Ag+ functions as an electron scavenger of irradiated TiO2. However, the photocatalytic degradation of 4-CP in the N2purged aqueous suspension of TiO2 and AgNO3 appears to stop at 60 min, while the parallel reaction in the aerated suspension

Figure 6. Benzoquinone (cycle bars) and hydroquinone (square bars) produced from 4-CP degradation in the suspension as indicated in Figure 5. Only the intermediates in the aqueous phase were detected.

of TiO2 still proceeds efficiently. At the same time, Ag+ concentration also drops nearly to zero at 60 min. This is consistent with the model where ecb- and hvb+ reactions are equivalent in rates.22-24 Once Ag+ is depleted, hvb+ will recombine with ecb-, consequently resulting in a null cycle for organic degradation. In an air-open reactor, O2 is immediately supplied once it is consumed by ecb-. Thus, 4-CP can degrade continuously upon further irradiation. On the other hand, the result also indicates that removal of O2 from the aqueous suspension by N2 bubbling is quite efficient. Under N2, no photodegradation of 4-CP was observed either with bare TiO2 or with Ag-deposited TiO2. The mole ratio of Ag+ consumed to 4-CP that disappeared is about 6 at 60 min (Figure 5). If Ag+ is used only to capture ecb- on TiO2, and chloride ion is released from 4-CP, the mole ratio would not be more than 2. This implies that Ag+ might also react with the intermediates of 4-CP photodegradation. Simultaneous detection of the intermediates by HPLC showed that benzoquinone (BQ) was the main intermediate, while hydroquinone (HQ) was formed when BQ began to disappear (curve b, Figure 6). It has been reported that BQ can be reduced into HQ over the irradiated TiO2.28b However, accumulation of BQ in the early stage is a hint that BQ is not able to compete with Ag+ for ecb- on TiO2. Comparatively, in the aerated aqueous suspension of TiO2, 4-CP photodegradation gave HQ as the main intermediate (curve a, Figure 6). A photocatalytic balance exists between HQ and BQ.29 It is highly possible that HQ is the primary intermediate of 4-CP photodegradation, followed by hvb+ oxidation into BQ. A control experiment in the dark showed that HQ could be oxidized into BQ by Ag+ or O2, but the reaction with Ag+ was much faster than that with O2. Thus, Ag+ can react with HQ and other unidentified intermediates, and its presence in the system does not lead to significant change in the primary pathway of 4-CP photodegradation, as compared with O2.

Anatase and Rutile TiO2 for Organic Degradation

J. Phys. Chem. C, Vol. 114, No. 44, 2010 18915

Figure 7. (A) Initial rate of 4-CP photodegradation in the aerated aqueous suspension of SAT (circle bars), SRT (square bars), and MAT (triangle bars). (B) The normalized initial rate with the BET surface area of TiO2, measured at 77 K with N2.

Figure 8. (A) Initial rate of 4-CP photodegradation under N2 in the presence of AgNO3, over SAT (circle bars), SRT (square bars), and MAT (triangle bars). (B) The normalized initial rate with the initial amount of Ag+ adsorbed. The insert is the plot of ln (the normalized R0) vs Ts.

The decay of 4-CP concentration with time in the aerated suspension of TiO2 is well fitted to the first-order rate equation. This is commonly observed in TiO2 photocatalysis for organic photodegradation at low concentration.37,38 However, in the presence of AgNO3 under N2, the reaction kinetics does not follow the first-order. This is because Ag+ concentration is not constant but rather decreases with time. Therefore, in the following study, only the initial rate of 4-CP photodegradation, determined at the first 10 min, will be used for evaluating the photocatalytic activities of different TiO2 catalysts. For simplicity, the initial rates of 4-CP photodegraded in the aerated aqueous suspension of TiO2 is denoted as R0(O2), while the rate measured in the N2-purged aqueous suspension of TiO2 and AgNO3 is denoted as R0(Ag+). Photoactivity Assessment under Air without AgNO3. Figure 7A shows the result of R0(O2) obtained with three differently prepared catalysts. The initial rate measured with SAT or MAT first increases, and then decreases with Ts, while the rate obtained with SRT decreases with Ts. A maximum rate of 4-CP degradation appears at 700, 500, and 110 °C for SAT, MAT, and SRT, respectively. At a given Ts, nearly all SRT samples have a much lower photoactivity than that of SAT and MAT. Mesoporous MAT has a higher photoactivity than that of SAT only at Ts e 600 °C. This kind of R0(O2) often appears in the literature for the activity determination. However, in some papers, the initial rate is normalized with the surface area of TiO2, considering that different catalysts have different surface areas, and that all surface areas have taken part in the photocatalytic reaction. The initial rates in Figure 7A are then normalized with the corresponding BET surface area of the catalyst in Figure 3, and the result is shown in Figure 7B. This specific rate, R0(O2)/Asp as a function of Ts, shows a very

different trend from that in Figure 7A. The photocatalytic activity of SRT not only increases with Ts but also is slightly higher than or similar to those of SAT and MAT at low Ts. The optimum Ts now appears at 700, 600, and 750 °C for SAT, MAT, and SRT, respectively. At this Ts, both SAT and MAT have a similar phase composition to P25 TiO2 (the content of anatase is about 72, 76, and 80%, respectively). The present result obtained in an aerated aqueous suspension is essentially similar to those reported in the literature. That is, in a certain region of Ts, as a photocatalyst for organic degradation in water, rutile is less active than anatase, while mesoporous TiO2 is better than nonporous and/or microporous TiO2. Photoactivity Assessment under N2 with AgNO3. Figure 8A shows the result of R0(Ag+) as a function of Ts, measured with the same catalysts used in Figure 7. At first glance, the plot results are remarkably different from those obtained in the absence of AgNO3 (Figure 7A). The initial rates of 4-CP degradation in the presence of 1.0 mM AgNO3 are about 1 order of magnitude larger than those obtained in an aerated aqueous suspension, similar to that shown in Figure 5. Interestingly, the photocatalytic activity of SRT now becomes very close to those of SAT and MAT at a given Ts. Moreover, the optimum Ts at which R0(Ag+) is maximal is the same for three differently prepared TiO2 samples. Since 4-CP degradation is the outcome of Ag+ reduction (Figure 5), the initial rate of 4-CP degradation was then divided by the initial amount of Ag+ adsorbed on TiO2, q0(Ag+), measured before light illumination. Surprisingly, this specific rate, R0(Ag+)/q0(Ag+), increases exponentially with Ts, and three curves obtained with SAT, SRT, and MAT are almost overlapped (Figure 7B). Moreover, a linear relationship between the logarithmic specific rate and Ts is satisfactory (insert in Figure

18916

J. Phys. Chem. C, Vol. 114, No. 44, 2010

Figure 9. Initial amount of Ag+ adsorbed on (a) SAT, (b) SRT, and (c) MAT, sintered at different temperatures. Data were measured just before light irradiation.

7B), which gives an average slope of 9.54 × 10-3. Data scattering at low and high Ts might be due to errors in measuring low R0(Ag+) and q0(Ag+), respectively (see below). The result indicates that at the same amount of Ag+ adsorbed on TiO2, anatase and rutile actually have a similar photocatalytic activity at a given Ts for 4-CP degradation in aqueous suspension. This specific photoactivity of TiO2 appears only to be determined by its Ts or crystallinity. As Ts increases, the solid becomes better crystallized, and the number of defect sites is reduced. As a result, the rate of charge carrier recombination would be decreased,39 and thus the rate of 4-CP degradation is increased. The SAT and MAT samples sintered at Ts ) 400-800 °C are a mixture of anatase and rutile, but they display a similar photoactivity to pure rutile (SRT). Moreover, well-defined mesoporous MAT also has a similar photoactivity to that of SAT or SRT at a given Ts. The corresponding q0(Ag+) used above as a function of Ts is shown in Figure 9. As Ts is increased, q0(Ag+) is decreased with each series of the samples. This might be ascribed to the decrease in the number of OH- groups on TiO2.5g,40 Ohtani and co-workers have reported that the adsorption isotherm of Ag+ on TiO2 (anatase and rutile) in aqueous suspension is a Langmuir-type, and the amount of Ag+ adsorbed increases with the suspension’s pH, ascribed to the increase in surface OHdeprotonation.5c,25 Comparatively, SRT has a lower capacity for Ag+ adsorption than SAT and MAT, probably due to its lower number of strongly acidic OH- groups, as compared to that on anatase.5g,25 In this regard, SAT seems to possess OH- groups slightly stronger in acidity than those of MAT; thus, SAT shows a q0(Ag+) value higher than that of MAT (Figure 9). Therefore, upon thermal treatment, SAT particles may easily join together to constitute a mesoprous network, as compared to MAT (Figure 3). The q0(Ag+)-normalized initial rate of 4-CP photodegradation might be taken as a measure of TiO2 intrinsic photoactivity. In principle, the rate of hvb+ reaction should be equal to the rate of ecb- reaction.22 That is, for the reactions using O2 and Ag+ as an electron acceptor, eqs 1 and 2 should hold, respectively. Then, the ratio, R0(O2)/q0(O2) or R0(Ag+)/q0(Ag+), would represent the activities of different catalysts in the production of ecb-/hvb+, provided that the catalysts have the same k(O2) or k(Ag+). Otherwise, the ratio could only be regarded as a specific activity of TiO2 per electron acceptor adsorbed on its surface. Moreover, although the light intensity entering into the reactor is controlled at a constant flux, the exact number of photons adsorbed by each catalyst is unknown, which might affect the activity determination among different catalysts.

Sun and Xu

R0(O2) ) k(O2)[ecb-]q0(O2)

(1)

R0(Ag+) ) k(Ag+)[ecb-]q0(Ag+)

(2)

where q0(O2) ) the initial amount of O2 adsorbed on TiO2; [ecb-] ) the surface concentration of ecb- on TiO2; k(O2) and k(Ag+) are the interfacial electron-transfer rate constants of O2 and Ag+, respectively. Slow reduction of O2 by TiO2 electron has been proposed to the rate-determing step in heterogeneous photoreaction.1,22-24 This could result from small k(O2), low q(O2), and/or both (refer to eq 1). The samples that display different photoactivies in an aerated aqueous suspension (Figure 7A) may relate to their difference in sorption capacity toward O2 in aqueous solution.41,42 Studies by Smith and Ford have shown that anatase provides heat of O2 adsorption higher than that of rutile in a gas-solid phase.43 Assuming that this trend also exists in an aqueous phase, then anatase would have a larger q(O2) and thus a higher rate of 4-CP degradation compared to rutile at a given Ts (Figure 7A). As Ts is increased, the “intrinsic” or specific photoactivity of TiO2 is increased (Figure 8B), and q(O2) on TiO2 would be decreased. If the catalyst has a high q(O2), the rate of 4-CP degradation would increase with Ts, as observed with SAT at Ts < 700 °C, and with MAT at Ts < 500 °C as well (curves a and c, Figure 7A). On the contrary, if the catalyst has a low q(O2), the rate of 4-CP degradation would decrease with Ts, as observed with SRT (curve b, Figure 7A). The balance between q(O2) and the real photoactivity of TiO2 would result in an optimum Ts for 4-CP degradation in aerated solution, as observed with SA and MAT at 700 and 500 °C, respectively. The above proposal can be also used to interpret the literature data about O2 evolution from water in the presence of Ag+ or Fe3+. First, rutile is more photoactive than anatase.7,8 Because rutile has a weaker affinity toward O2 in water than anatase, the rate of O2 reduction by ecb- would be slower, thus resulting in a higher rate of O2 evolution, as compared to anatase. Second, according to the reaction stoichiometry, the molar ratio of reduced Ag+ or Fe3+ to evolved O2 would be 4. This has been observed with the samples sintered at high Ts, but the ratio obtained with the samples sintered at low Ts is about 5.4-8.8.5c,44-46 The catalyst sintered at low Ts would have a large q(O2), a high rate of O2 reduction, and thus a low rate of O2 evolution. Third, it has been reported that with pure rutile or anatase, the rate of O2 evolution from water first increases, and then decreases with Ts.44,45 However, when the initial rate of O2 evolved is normalized by the initial amount of Ag+ adsorbed,45 we find that the normalized rate also increases exponentially with Ts. A linear relationship between the normalized rate in logarithmic form and Ts gives a slope of 9.73 × 10-3. Surprisingly, this slope is almost the same as that of 9.54 × 10-3 determined for organic degradation in Figure 8B. Because both water oxidation and organic degradation occur through hvb+, evidence provides strong support to our finding that anatase and rutile have a similar “intrinsic” photoactivity at a given Ts. Since the “intrinsic” photoactivity of TiO2 increases, and q(O2) on TiO2 decreases with Ts, respectively, the rate of O2 evolved in an aerated solution is expected to continuously increase with Ts. This has been observed at low Ts, but the rate of O2 reduction decreases with Ts after a certain Ts.44,45 The latter could be ascribed to the decrease of Ag+ adsorption on TiO2, as shown in Figure 9. In the absence of Ag+, the evolution of O2 from an irradiated aqueous suspension of TiO2 is very slow or negligible.44-46

Anatase and Rutile TiO2 for Organic Degradation

J. Phys. Chem. C, Vol. 114, No. 44, 2010 18917 Acknowledgment. Financial support from the National Science Foundation of China (Nos. 20525724 and 20873124) and the National Basic Research Program of China (No. 2009CB825300) is gratefully acknowledged. Supporting Information Available: XRD patterns, pore size distribution, and TGA curves. This material is available free of charge via the Internet at http://pubs.acs.org. References and Notes

Figure 10. Estimated amount of O2 adsorbed on (a) SAT, (b) SRT, and (c) MAT in aqueous suspension, on the basis of eq 3 and the data in Figures 7-9.

The trend in q0(O2) among different TiO2 samples might be estimated through eq 3, a combination of eqs 1 and 2, provided that the catalysts have the same ratios of k(Ag+) to k(O2). The result is shown in Figure 10. For each series of TiO2, the calculated q0(O2) decreases with Ts. Among the catalysts at given Ts, SRT has the lowest q0(O2). This is in agreement with the above proposal. Because of a well-defined mesopore, the MAT samples sintered at 300 and 400 °C have a q0(O2) higher than that of SAT. Although the trend in q0(O2) as a function of Ts is similar to that in surface area (Figure 3), the difference in calculated q0(O2) among samples is notably larger than the difference in surface area. The powders dispersed in solution would have a much different surface area from those dried in a vacuum. Therefore, the BET surface area-normalized rate of 4-CP degradation in Figure 7B, as a measure of TiO2 photoactivity, is not reliable. The present estimation of q0(O2) clearly has a significant error but may be taken as a rough indicator of q0(O2) for different catalysts in water.

q0(O2) )

k(Ag+)R0(O2) k(O2)R0(Ag+)

q0(Ag+)

(3)

Conclusions In this work, we have shown that the initial rate of 4-CP photodegradation, per the initial amount of Ag+ adsorbed on TiO2, increases exponentially with Ts, ascribed to the increase in crystallinity (Figure 1). This observation is made with three differently prepared catalysts and also found accidentally in the literature data of water splitting. Evidence strongly suggests that at the same amount of electron acceptor adsorbed on the catalyst surfaces, anatase and rutile actually have a similar photocatalytic activity at a given Ts, for both reactions of organic and water oxidation. The observed variation in photoactivity among the samples in aerated suspension is ascribed to the catalyst difference in sorption capacity toward O2. Because the real photoactivity and O2 sorption capacity of TiO2 increases and decreases with Ts, respectively, there appears an optimum Ts for organic degradation or water oxidation in an aerated aqueous suspension. It is recommended that to evaluate the photocatalytic activities of different TiO2 in an aerated aqueous solution, the difference in q(O2) needs to be taken into account. For future development of an advanced photocatalyst, anatase with porous structure and high crystallinity would be preferred for organic degradation, while rutile with a nonporous network and high crystallinity would be preferred for water splitting into molecular oxygen.

(1) Hoffmann, M. R.; Martin, S. T.; Choi, W.; Bahnemann, D. W. Chem. ReV. 1995, 95, 69–96. (2) Carp, O.; Huisman, C. L.; Reller, A. Prog. Solid State Chem. 2004, 32, 33–177. (3) Emeline, A. V.; Zhang, X.; Jin, M.; Murakami, T.; Fujishima, A. J. Phys. Chem. B 2006, 110, 7409–7413. (4) Tachikawa, T.; Fujitsuka, M.; Majima, T. J. Phys. Chem. C 2007, 111, 5259–5275. (5) For examples, (a) Sclafani, A.; Palmisano, L.; Schiavello, M. J. Phys. Chem. 1990, 94, 829–834. (b) Serpone, N.; Lawless, D.; Khairutdinov, R.; Pelizzetti, E. J. Phys. Chem. 1995, 99, 16655–16661. (c) Kominami, H.; Murakami, S.; Kato, J.; Kera, Y.; Ohtani, B. J. Phys. Chem. B 2002, 106, 10501–10507. (d) Ohno, T.; Sarukawa, K.; Matsumura, M. New J. Chem. 2002, 26, 1167–1170. (e) Inagaki, M.; Nonaka, R.; Tryba, B.; Morawski, A. W. Chemosphere 2006, 64, 437–445. (f) Dong, W.; Sun, Y.; Lee, C. W.; Hua, W.; Lu, X.; Shi, Y.; Zhang, S.; Chen, J.; Zhao, D. J. Am. Chem. Soc. 2007, 129, 13894–13904. (g) Xu, Y.; Lv, K.; Xiong, Z.; Leng, W.; Du, W.; Liu, D.; Xue, X. J. Phys. Chem. C 2007, 111, 19024–19032. (6) Tanaka, K.; Capule, M. F. V.; Hisanaga, T. Chem. Phys. Lett. 1991, 187, 73–76. (7) (a) Ohno, T.; Sarukawa, K.; Matsumura, M. J. Phys. Chem. B 2001, 105, 2417–2420. (b) Ohno, T.; Sarukawa, K.; Matsumura, M. New J. Chem. 2002, 26, 1167–1170. (8) Prieto-Mahaney, O. O.; Murakami, N.; Ohtani, B. Chem. Lett. 2009, 38, 238–239. (9) Kawahara, T.; Konishi, Y.; Tada, H.; Tohge, N.; Nishii, J.; Ito, S. Angew. Chem., Int. Ed. 2002, 41, 2811–2813. (10) (a) Hurum, D. C.; Agrios, A. G.; Gray, K. A.; Rajh, T.; Thurnauer, M. C. J. Phys. Chem. B 2003, 107, 4545–4549. (b) Hurum, D. C.; Gray, K. A.; Rajh, T.; Thurnauer, M. C. J. Phys. Chem. B 2005, 109, 977–980. (c) Li, G.; Gray, K. A. Chem. Phys. 2007, 339, 173–187. (11) Zhang, J.; Xu, Q.; Feng, Z.; Li, M.; Li, C. Angew. Chem., Int. Ed. 2008, 47, 1766–1769. (12) (a) Zachariah, A.; Baiju, K. V.; Shukla, S.; Deepa, K. S.; James, J.; Warrier, K. G. K. J. Phys. Chem. C 2008, 112, 11345–11356. (b) Baiju, K. V.; Zachriah, A.; Shukla, S.; Biju, S.; Reddy, M. L. P.; Warrier, K. G. K. Catal. Lett. 2009, 130, 130–136. (13) Kominami, H.; Ishii, Y.; Kohno, M.; Konishi, S.; Kera, Y.; Ohtani, B. Catal. Lett. 2003, 91, 41–47. (14) Kandiel, T. K.; Feldhoff, A.; Robben, L.; Dillert, R.; Bahnemann, D. W. Chem. Mater. 2010, 22, 2050–2060. (15) Yang, D.; Liu, H.; Zheng, Z.; Yuan, Y.; Zhao, J.; Waclawik, E. R.; Ke, X.; Zhu, H. J. Am. Chem. Soc. 2009, 131, 17885–17893. (16) Agrios, A. G.; Pichat, P. H. J. Photochem. Photobiol., A 2006, 180, 130–135. (17) Hathway, T.; Jenks, W. S. J. Photochem. Photobiol., A 2008, 200, 216–224. (18) Ryu, J.; Choi, W. EnViron. Sci. Technol. 2008, 42, 294–300. (19) Tang, J. W.; Durrant, J. R.; Klug, D. R. J. Am. Chem. Soc. 2008, 130, 13885–13891. (20) Mills, A.; Hunte, S. L. J. Photochem. Photobiol., A 1997, 108, 1– 35. (21) Mills, A.; Wang, J.; Ollis, D. F. J. Catal. 2006, 243, 1–6. (22) Wang, C. M.; Heller, A.; Gerischer, H. J. Am. Chem. Soc. 1992, 114, 5230–5234. (23) Kesselman, J. M.; Shreve, G. A.; Hoffmann, M. R.; Lewis, N. S. J. Phys. Chem. 1994, 98, 13385–13395. (24) Ahmed, S.; Fonseca, S. M.; Kemp, T. J.; Unwin, P. R. J. Phys. Chem. B 2003, 107, 5892–5900. (25) Ohtani, B.; Okugawa, Y.; Nishimoto, S.-I.; Kagiya, T. J. Phys. Chem. 1987, 91, 3550–3555. (26) Mills, A.; Morris, S.; Davies, R. J. Photochem. Photobiol., A 1993, 70, 183–191. (27) Stafford, U.; Gray, K. A.; Kamat, P. V. J. Phys. Chem. 1994, 88, 6343–6351. (28) (a) Richard, C.; Boule, P. New J. Chem. 1994, 18, 547–552. (b) Richard, C. N. J. Chem. 1994, 18, 443–445. (29) Theurich, J.; Lindner, M.; Bahnemann, D. W. Langmuir 1996, 12, 6368–6376.

18918

J. Phys. Chem. C, Vol. 114, No. 44, 2010

(30) Li, X.; Cubbage, J. W.; Tetzlaff, T. A.; Jenks, W. S. J. Org. Chem. 1999, 64, 8509–8524. (31) Anpo, M.; Shima, T.; Kodama, S.; Kubokawa, Y. J. Phys. Chem. 1987, 91, 4305–4310. (32) Tian, G.; Fu, H.; Jing, L.; Xin, B.; Pan, K. J. Phys. Chem. C 2008, 112, 3083–3085. (33) Chakrabarti, S.; Ganguli, D.; Chaudhuri, S. Phys. E 2004, 24, 333– 342. (34) Cave, G. C.; Hume, D. N. Anal. Chem. 1952, 24, 1503–1505. (35) Hagfeldt, A.; Gratzel, M. Chem. ReV. 1995, 95, 49–68. (36) Sawyer, D. T.; Valentine, J. S. Acc. Chem. Res. 1981, 14, 393. (37) (a) Turchi, C. S.; Ollis, D. F. J. Catal. 1990, 122, 178–192. (b) Ollis, D. F. J. Phys. Chem. B 2005, 109, 2439–2444. (38) (a) Emeline, A. V.; Ryabchuk, V. K.; Serpone, N. J. Photochem. Photobiol., A 2000, 133, 89–97. (b) Emeline, A. V.; Ryabchuk, V. K.; Serpone, N. J. Phys. Chem. B 2005, 109, 18515–18521.

Sun and Xu (39) Ohtani, B.; Bowman, R. M.; Colombo, D. P.; Kominami, H.; Noguchi, H.; Uosaki, K. Chem. Lett. 1998, 27, 579–580. (40) Carneiro, J. T.; Savenije, T. J.; Moulijn, J. A.; Mul, G. J. Phys. Chem. C 2010, 114, 327–332. (41) Yu, J. C.; Lin, J.; Lo, D.; Lam, S. K. Langmuir 2000, 16, 7304– 7308. (42) Peiro, A. M.; Colombo, C.; Doyle, G.; Nelson, J.; Mills, A.; Durrant, J. R. J. Phys. Chem. B 2006, 110, 23255–23263. (43) Smith, W. R.; Ford, D. G. J. Phys. Chem. 1965, 69, 3587–3592. (44) Oosawa, Y.; Gratzel, M. J. Chem. Soc. Faraday Trans. 1 1988, 84, 197–205. (45) Nishimoto, S.; Ohtani, B.; Kajiwara, H.; Kagiya, T. J. Chem. Soc., Faraday Trans. 1 1985, 81, 61–68. (46) Kominami, H.; Murakami, S.; Kera, Y.; Ohtani, B. Catal. Lett. 1998, 56, 125–129.

JP104762H