Ensemble Effects in Nanostructured TiO - American Chemical

4608

J. Phys. Chem. B 2002, 106, 4608-4616

Ensemble Effects in Nanostructured TiO2 Used in the Gas-Phase Photooxidation of Trichloroethylene King Lun Yeung,*,† A. Javier Maira,‡ Jennifer Stolz,† Eppie Hung,† Nick Ka-Chun Ho,§ An Chyi Wei,| Javier Soria,‡ Kuei-Jung Chao,| and Po Lock Yue† Department of Chemical Engineering, The Hong Kong UniVersity of Science and Technology, Clear Water Bay, Kowloon, Hong Kong, P.R. China, Instituto de Cata´ lisis y Petroleoquı´mica, CSIC, Campus UAM, Cantoblanco, Madrid 28047, Spain, Material Preparation and Characterization Facility, The Hong Kong UniVersity of Science and Technology, Clear Water Bay, Kowloon, Hong Kong, P.R. China, and Department of Chemistry, National Tsinghua UniVersity, Hsinchu, Taiwan, R.O. China ReceiVed: August 10, 2001; In Final Form: January 18, 2002

Nanostructured TiO2 with different crystal sizes (i.e., 3, 5, and 6 nm) but identical aggregate size (i.e., 100 nm) and shape (sphere) were successfully made using a modified sol-gel method. TiO2 catalysts with crystal size of 5 nm and aggregate size of 100 nm were also prepared. This batch of catalysts has similar phase structure, surface area, band-gap energy, and degree of surface hydroxylation, and differs mainly in their crystallinity. Synchrotron XRD, XANES and EXAFS, Raman spectroscopy and transmission electron microscopy (TEM) were used to characterize the structural properties of the nanometer sized TiO2, whereas ultraviolet, X-ray photoelectron and electron paramagnetic resonance spectroscopy provided important information on the electronic properties of these nanoparticles. The catalytic performance of the nanostructured TiO2 for gas-phase photocatalytic oxidation of trichloroethylene exhibited strong dependency on both crystal size and crystallinity.

Introduction Heterogeneous photocatalytic oxidation (PCO) is a promising technology for degradation of volatile organic compounds (VOCs) in process air stream.1-3 Using a semiconductor catalyst (e.g., TiO2) and a UV light source, this technology can oxidize airborne VOCs into carbon dioxide, water, and mineral compounds (e.g., HX, where X ) Cl, Br) at room temperature. The UV irradiated catalyst absorbs photons having energy greater than its band-gap, exciting electrons (e-) from the valence to the conduction band and generating holes (h+). The resulting electron-hole pairs migrate toward the catalyst surface and initiate redox reactions that oxidize the adsorbed organic molecules. Titanium dioxide has been widely used as a photocatalyst for the oxidation of organic pollutants due to its photostability, commercial availability, low cost, and relatively good activity. Among commercial samples, Degussa TiO2 P25 has been the most commonly used photocatalyst for organic pollutant photooxidation. Most commercial catalysts have moderate quantum yield and total mineralization of certain substrates is difficult.4-7 Catalyst deactivation is also a problem for PCO of a number of pollutants.8-10 The efficiency of the TiO2 photocatalyst depends on many factors including the number and energy of absorbed photons during irradiation, electron-hole pair formation and recombination rates, reactants and products adsorption and desorption rates, and charge transfer rate to the adsorbed species. Most of these factors are strongly * Author to whom correspondence should be addressed. Tel: 852-23587123. Fax: 852-2358-0054. E-mail: [email protected]. † Department of Chemical Engineering, The Hong Kong University of Science and Technology. ‡ Instituto de Cata ´ lisis y Petroleoquı´mica, CSIC. § Material Preparation and Characterization Facility, The Hong Kong University of Science and Technology. | National Tsinghua University.

affected by the structural and electronic properties of the catalyst that depend on particle size11-17 and preparation method.18-20 Recent reports have shown that nanostructured anatase TiO2 particles exhibit higher activity11,12,18,21 and selectivity12 compared with the commercial TiO2 (P25, Degussa) for gas-phase photocatalytic oxidation of volatile organic compounds. Although there is a significant interest in the use of nanostructured TiO2 in photocatalysis, the exact role of structure, electronic and surface chemistry on the photocatalytic properties of the nanometer sized TiO2 has not yet been fully resolved. Smaller TiO2 crystals possess larger surface area for adsorption and degradation of organic pollutants. As the size of TiO2 catalyst is decreased to nanometer range, its electronic and structural properties display a significant deviation from that of the bulk material. Discretization of the energy band structure occurs,22 accompanied by a significant change in the optical and chemical properties of TiO2.23 It also produces a net shift in the semiconductor band-gap to higher energies,11 which means that for the same light source, there are fewer photons with the required energy to generate the e-/h+ pairs needed for the reaction resulting in slower reaction rate. The altered band gap structure also modifies the redox potential of the e-/h+ pairs participating in the reaction process.24 Ensemble effects also become important as the particle size decreases. Small crystals have larger surface curvature and contain greater proportions of edge and corner sites, while the large crystals consist mainly of planar sites. This directly affects the types of adsorption and catalytic sites available for the reaction. In addition, the local structural environment of the atoms in nanometer sized crystal is distinctly different from the bulk material.25 This work attempts to elucidate the effects of crystal size, structure, and crystallinity on the photocatalytic activity of nanostructured TiO2 for gas-phase PCO of TCE.

10.1021/jp0131121 CCC: $22.00 © 2002 American Chemical Society Published on Web 04/17/2002

Nanostructured TiO2 Used in the Photooxidation of TCE

J. Phys. Chem. B, Vol. 106, No. 18, 2002 4609

TABLE 1: Preparation of Nanostructured TiO2 Catalysts titania catalysts

hydrothermal treatment composition (mole ratios)

crystallization temperature (K)

crystallization time (h)

P3 P5a P5b P5c P6

18 mL H2O + 7 mL 2-propanol 15 mL H2O + 10 mL 2-propanol 17 mL H2O + 8 mL 2-propanol 18 mL H2O + 7 mL 2-propanol 10 mL H2O + 15 mL 2-propanol

373 423 423 423 423

9 8 8 8 7

Experimental Section Synthesis and Characterization of Nanostructured TiO2. In this study, a modified sol-gel method was employed in the synthesis of nanostructured TiO2 photocatalysts of controlled crystal and aggregate sizes.11 The procedure involves the precipitation of amorphous titanium oxide gel spheres followed by the controlled crystallization of the anatase TiO2 nanocrystals. The concentrations of the precursor (i.e., titanium isopropoxide) and water have direct effects on the hydrolysis, condensation, and polymerization reactions that dictate the size of precipitated gel spheres. In this study, the synthesis conditions were adjusted to give 100 nm gel spheres of uniform size. Table 1 summarizes the treatment conditions used to crystallize the anatase TiO2. During the hydrothermal treatment, the amorphous gel spheres were transformed into aggregates of TiO2 nanocrystals with no significant change in size and shape. Thus, TiO2 photocatalysts with different crystal sizes but fixed aggregate size (i.e., 100 nm) and shape (i.e., spheres) were obtained. This preparation procedure was also adopted for the preparation of TiO2 with similar crystal size but different crystallinity. The crystal structure, particle size, morphology, and surface area of TiO2 photocatalysts were characterized using X-ray diffraction (XRD), X-ray absorption spectroscopy (XAS), microRaman, transmission electron microscopy (TEM), and N2 physiadsorption. X-ray diffraction and X-ray absorption analyses of the TiO2 powder were conducted on beamline BL17A at the Synchrotron Radiation Research Center (SRRC) in Taiwan. The X-ray radiation (λ ) 1.3271 Å) with a beam current of 120200 mA was supplied from a 1.5 GeV storage ring. The XRD patterns were recorded for 15° < 2θ < 55° by step-scanning at 0.05° increments. The phase structure, crystallinity, and grain size of the TiO2 were determined from the X-ray diffraction patterns. The X-ray absorption spectra of the TiO2 samples were measured in transmission mode at room temperature. To eliminate the effects of sample thickness, the TiO2 powder was rubbed uniformly onto a Scotch tape and folded to get the desired thickness that satisfies ∆µx e 1, where ∆µx is the edge step. The spectra were collected using a gas ionization detector. The ion chambers used for measuring the incident (I0) and transmitted (I) synchrotron beam intensities were filled with a He/N2 gas mixture and N2 gas, respectively. Data were collected from 200 eV below the Ti K-edge (4966 eV) to 1100 eV above the edge. Standard titanium metal foil and commercial anatase TiO2 were used as reference standards. Information on the bond structure (i.e., bond lengths, numbers, and types of neighboring atoms) could be obtained from the X-ray absorption data. The crystal size was also measured from Raman line broadening and compared with the results of the X-ray diffraction experiments. The Raman spectra were measured using a Renishaw 3000 micro-Raman system with an Olympus BH-2 microscope. The objective lenses with 20× and 50× magnifications were selected. The excitation source used was an argon laser operating at 514.5 nm with an output power of 25 mW. Raman spectra were taken of the powders on a glass microscope slide. The spectral resolution was set at approximately 1.0 cm-1 and the spot size was about 2 µm in diameter. For the TEM

imaging, the sample powder was dispersed in 2-propanol and placed onto a carbon-coated copper grid. The excess liquid was removed using a paper wick and the deposit was dried in air prior to imaging. The deposited TiO2 particles were examined with a Philips CM20 transmission electron microscope at an accelerating voltage of 200 kV. Routine chemical analysis using an in-situ energy-dispersive X-ray spectrometer (EDXS, Philips XL30) detected only the energy signatures corresponding to the Ti and O elements. The average TiO2 crystal and aggregate sizes, as well as their morphologies were obtained from the TEM micrographs. The BET surface area of the TiO2 photocatalyst was measured using the N2 physi-adsorption technique (Micromeritics ASAP 2010). The dried powder was outgassed under vacuum at 523 K for 4 h prior to the analysis. The surface composition and chemical state of the catalyst samples were determined by X-ray photoelectron spectroscopy (XPS, Physical Electronics PHI 5600) using a monochromatic aluminum X-ray source. The catalyst powder was pressed onto a double-sided carbon tape to make a flat surface and then introduced into the instrument for analysis. The change in the band-gap energy of the TiO2 as a result of the material preparation was also measured using a Philips/Unicam Pu8700 UV/VIS spectrometer with diffuse and specular reflection accessories to provide information on the electronic band gap. A Ba2SO4 standard white plate was used as background. The spectra were recorded at a scan speed of 125 nm/min and a bandwidth of 2 nm. Evaluation of Photocatalytic Activity of TiO2. Gas-phase, photocatalytic oxidation of trichloroethylene (TCE) was used as a probe reaction to study the catalytic activity of nanostructured TiO2. The reaction was conducted in a flat-plate photoreactor11 using TiO2-coated stainless steel plates. Catalyst powder (30 mg) was uniformly coated onto the 25 mm × 25 mm plate. A metered amount of TCE vapor (0.5 µL/min) was mixed with flowing oxygen (200 mL/min) before entering the photoreactor to give a TCE partial pressure of 70 Pa (4.5 ppmv). The reactant mixture then enters the reactor and flows over the catalyst plate through a narrow rectangular channel (2 mm deep × 112 mm wide). After the feed concentration had stabilized (t ∼ 1 h), the photocatalyst was illuminated by five fluorescent black lamps (6 W, BLB Sankio Denki) located 7 mm above the reactor window. The outlet gases were separated using a GS-GASPRO capillary column (0.32 mm × 30 m) and analyzed using a gas chromatograph (HP 6890) equipped with thermal conductivity and flame ionization detectors. The reactant and product concentrations were monitored continuously with time. The photocatalytic reaction was also investigated using an in-situ infrared photoreaction cell. The cell consists of two KBr windows and a sample holder for the catalyst wafer. The reactants and products enter and exit the cell through a set of inlet and outlet ports. The compressed air used in the reaction was metered (800 mL/min) and pretreated through a series of moisture and hydrocarbon traps that included a liquid N2 trap, a drying column packed with desiccant, and molecular sieves. Trichloroethylene was fed using a syringe pump (kdScientific 1000) to a mixing tee where it was vaporized and mixed with

4610 J. Phys. Chem. B, Vol. 106, No. 18, 2002

Yeung et al.

Figure 1. (a) Tapping-mode atomic force microscopy image of P5a dispersed onto mica substrate (inset: individual 100 nm aggregates of 5 nm crystals), (b) synchrotron X-ray diffraction pattern, and (c) XANES spectra of the TiO2 samples (inset: magnified preedge region of commercial anatase TiO2).

the dry air. The reactant mixture then flowed through the insitu infrared photoreaction cell and allowed to equilibrate at room temperature (298 K). Once the reactant concentration had stabilized, the inlet and outlet ports were shut-off and the UV lamp was turned on to allow the photoreaction to proceed under batch conditions. The infrared spectra were continuously collected by a Fourier transform infrared spectrometer (FTS 6000, Bio-Rad Laboratories) during the reaction until the intensity of the characteristic peaks of adsorbed TCE molecule was decreased to less than 0.2 of the original. The UV lamp was then turned off and the cell purged with flowing dry air. The catalyst wafers used in the reaction experiments were prepared by mixing 0.08 g of TiO2 with 0.70 g of KBr. Three wafers were prepared in each batch, each weighing 0.20 g. The wafer thickness as measured by scanning electron microscopy (JEOL 6300) was 280 ( 22 µm, thus giving the wafers an average density of 2.28 g/cm3. The wafers were pretreated in flowing dry air for 2 h at room temperature prior to the reaction to remove adsorbed CO2. Materials. The chemicals used in the synthesis of nanostructured TiO2 photocatalysts included titanium isopropoxide (97%, Aldrich), 2-propanol (99.5%, Aldrich), and distilled water. Anatase TiO2 (99.9+ %) from Aldrich Chemicals was used as

reference standard for the synchrotron X-ray experiments. Standard FTIR-grade potassium bromide (Aldrich) was used as catalyst diluent in the in-situ IR experiments. The reactants used for catalyst testing were trichloroethylene (99.5%, Aldrich) and oxygen (HP, HKO). Gases used in the GC were helium (UHP, CW), hydrogen (UHP, HKO), and synthetic air (HP, HKSP). Results and Discussion Characterization of Nanostructured TiO2. Figure 1a displays typical TiO2 aggregate clusters (P5a) prepared by the modified sol-gel process. The atomic force microscope image reveals that the clusters are made up of individual nanocrystals with 5 nm diameter (Figure 1a inset). It is clear from the presence of crystal facets that the TiO2 are crystalline. Indeed, this is confirmed by the X-ray diffraction patterns (Figure 1b) which show that the primary crystalline phase of the nanostructured TiO2 is anatase. Brookite is also present in small amount in all the samples. The average crystal size calculated from the broadening of the anatase (101) peak is displayed in Table 2. It is also clear from the diffraction patterns that although the P5a, P5b, and P5c have comparable crystal size, they exhibit different crystallinity. The Raman spectra of 5 and 6 nm TiO2 samples



TABLE 2: Structural Properties of Nanostructured TiO2 Prepared by Modified Sol-Gel Method Ti K-edge EXAFSd

titania catalysts

primarily particle size (nm)a

secondary particle size (nm)b

P3

2.9

100

--

P5a

4.6

100

1.0e,f

P5b

5.0

100

0.8e

P5c

5.1

100

0.5e

P6

5.7

100

1.0e

crystallinityc

shell

R(Å)

C. N.

σ2 ×10-3 (Å2)

Ti-O Ti-Ti Ti-O Ti-O Ti-Ti Ti-Ti Ti-O Ti-O Ti-Ti Ti-Ti Ti-O Ti-O Ti-Ti Ti-Ti Ti-O Ti-O Ti-Ti Ti-Ti

1.93 3.05 1.94 1.99 3.05 3.79 1.93 1.99 3.05 3.79 1.94 1.99 3.05 3.79 1.95 2.00 3.07 3.81

4.5 1.5 3.0 2.8 1.9 0.7 3.0 3.0 2.1 1.2 2.8 3.2 2.2 1.0 4.2 2.1 2.4 2.9

15 9.9 10.3 10.5 6.9 7.1 9.4 9.6 6.3 6.4 8.9 9.1 5.9 6.1 9.8 10.1 6.6 6.8

BET surface area (m2/g) 330 275

240

220

175

a The crystal size was calculated from XRD line broadening and confirmed by micro-Raman and TEM analyses. b The average aggregate size was obtained from 50 direct measurements of TiO2 clusters imaged by TEM. c The crystallinity was based on the intensity of the anatase (101) XRD peak, i.e., crystallinity ) I/I0e, using P5a as the reference f. d EXAFS values were obtained from model fitting using a standard software package (UWXAFS 3.0).

display a distinct shift in their spectral lines when compared with coarser TiO2 crystals (i.e., 11 nm P11). The peak intensity decreases and the spectral line broadens with decreasing particle size. A strong correlation between Raman line broadening and TiO2 particle size has been reported.26 Iida and co-workers26 claim that this method is more sensitive for nanometer sized crystals than X-ray line broadening. The average crystal size calculated from the Raman peak broadening is in good agreement with the synchrotron XRD results. Although TEM was used mainly to measure the size and shape of the TiO2 aggregates, the crystal size could also be determined from the high-resolution images. Their values usually agree well with that calculated from X-ray and Raman peak broadening. The BET surface area of the TiO2 measured by N2 physi-adsorption displays a monotonic increase with decreasing crystal size (Table 2). Figure 1c displays the normalized Ti K-edge X-ray absorption near-edge spectra (XANES) of the TiO2 samples listed in Table 2, including that of a commercial crystalline anatase TiO2 being used as a standard. The XANES spectrum of the commercial TiO2 (Figure 1c) indicates a somewhat distorted TiO6 octahedron atomic arrangement with four short Ti-O (1.93 Å) and two long Ti-O bonds (1.98 Å). The preedge region of the anatase Ti K-edge XANES usually consist of four preedge features labeled A1, A2, A3, and B (cf. Figure 1c inset). The feature labeled A2 is sometimes observed as a weak shoulder on the low energy side of the A3 peak. The intensity and position of the preedge transitions are sensitive to the symmetry of the surrounding atoms, being dipole forbidden but quadrupole allowed, are weak in symmetrical environments and display an increase in intensity as the environment is distorted. The postedge region of the Ti K-edge XANES from 4980 to 5020 eV contains a number of resolved and understood intense 3s f np dipole-allowed transitions. The normalized Ti K-edge XANES spectra of TiO2 samples with different crystal size (P3, P5a, and P6) and crystallinity (P5a, P5b, and P5c) are also shown in the figure. It is clear from the spectra that as the crystal size of TiO2 samples decreases, the features in the postedge region become poorly resolved. A pronounced difference in the feature of the preedge region is evident in the spectra of TiO2

samples with different crystal size. The 4-, 5-, 6-fold coordinated Ti can be identified according to the position and intensity of the preedge peaks, which are located at 4969.5 (A1), 4970.5 (A2), and 4971.5 (A3) eV, respectively. As the TiO2 crystal size decreases, the interfacial region comprising titanium atoms with low coordination number increases. Thus, the titanium atoms in smaller crystals will exhibit lower coordination numbers. Indeed, the 3 nm TiO2 (P3) has the most intense A2 peak (4970.75 eV) that is consistent with 5-fold coordinated Ti. Similar structural geometry has been reported for 1.9 nm TiO2.27 Larger crystals, which have smaller interfacial area, exhibit a weaker A2 peak but a more intense A3 peak (4971.75 eV) that corresponds to 6-fold coordinated Ti. The preedge features of the 5 nm TiO2 with different crystallinity are identical. Extended X-ray absorption fine structure (EXAFS) analysis of the TiO2 samples provides detailed local structural information not available from the X-ray diffraction data. The analysis was done using UWXAFS 3.0 software package and FEFF8 program and the results are tabulated in Table 2. The magnitude of Fourier transform of Ti K-edge k3-weighted EXAFS data (uncorrected for phase shift) for the nanostructured TiO2 exhibit a general increase in amplitude of the FT peaks with increasing particle size, in particular the relative intensity of the peaks corresponding to the higher shells. The fitting results for the 3 nm TiO2, P3 (Table 2) show that the Ti-O bond length is 1.93 Å and the coordination number is 4.5, which is consistent with the XANES result. This catalyst, which is comprised of 3800 TiO2 units, has a surface region (Ap) that is comparable to the bulk (Vp) with an Ap/Vp ratio of 1. In the other TiO2 samples, the first peak must be fitted with two shells of Ti-O. The P5a, P5b, and P5c samples (Ap/Vp ) 0.6) have the same results showing the presence of three short bonds (1.93 Å) and three long Ti-O bonds (1.98 Å) in the first shell despite the difference in their crystallinity. Although the Ti in these samples are 6-fold coordinated, they are not like the octahedron of the anatase TiO2 that is characterized by four short and two long Ti-O bonds. This could be due to the contribution from the interfacial atoms or the possible presence of structural defects. The P6 (Ap/Vp ) 0.5) has the same structure as the commercial anatase TiO2 (four


Figure 2. UV spectra of the nanometer sized TiO2.

short Ti-O bonds of 1.93 Å and two long Ti-O bonds of 1.98 Å). The high concentrations of low coordinated Ti4+ cations, as well as the distorted anatase crystal structure observed by XANES and EXAFS, are strongly suggestive of the presence of amorphous materials in the nanometer sized TiO2. However, a recent EPR study conducted on the same set of nanostructured TiO217 was unable to detect the broad Ti3+ signal that is characteristic of UV irradiated amorphous TiO2 in a vacuum, although it did confirm the presence of low coordinated Ti4+ cations. The EPR spectra for the nanostructured TiO2 contained mainly the signals corresponding to Ti3+ cations formed by the trapping of photogenerated electrons at Ti4+ sites on the crystal surface (i.e., corners, steps, edges, kinks and defects). These

Yeung et al. signals disappeared when annealing the TiO2 at high temperatures eliminates the low coordinated surface sites. Subsurface defects containing low coordinated Ti4+ were detected indirectly from the formation of subsurface O- species (O-b) when the nanocrystals were irradiated in the presence of adsorbed oxygen. The trapping of photogenerated h+ at these subsurface O2anions was shown to be mediated by the surface hydroxyl groups (-O2--Ti4+-OH + h+ f -O--Ti4+-OH) that were abundant on the surface of nanostructured TiO2.17 Only titanium and oxygen were detected by in-situ EDXS analysis of the TiO2 particles during TEM imaging. The XPS measurements also gave similar results for the surface composition of the crystalline samples with adsorbed atmospheric carbons being the main surface contaminant. Due to the larger mean free path of photoelectrons from the valence band, the valence band spectra will contain signals from both the surface and bulk, and can be used to probe the electronic structure of TiO2. The XPS results indicated that the valence bands of P6 and P5a TiO2 were identical, but the energy distribution of the valence band of P3 displayed a distinct difference from the larger TiO2. The crystallinity has little effect on the structure of the valence band. Figure 2 shows that there is a clear shift in the semiconductor band-gap toward higher energy when the TiO2 crystal size is decreased from 11 to 6 and 3 nm, but there is only a slight difference between the 5 and 6 nm TiO2. The occurrence of the blue-shift means that for the same light source, there are fewer photons with the required energy to generate the e-/h+ pairs needed for a reaction resulting in a lower utilization of photons. Photocatalytic Activity of Nanostructured TiO2. Crystal Size and Structure. Figure 3a plots the areal conversion rate of

Figure 3. (a) Areal TCE conversion rates as a function of time for different crystal sizes, (b) areal CO2 production rates as a function of time for different crystal sizes, (c) areal TCE conversion rates as a function of time for different crystallinity, (d) areal CO2 production rates as a function of time for different crystallinity.



Figure 4. (a) Transient in situ FTIR transmittance spectra of TCE photooxidation on P5a and (b) FTIR transmittance spectra collected immediately after turning the UV off [1] and after 32 h of purging in dry air [2].

TCE as a function of UV irradiation time for 5 and 6 nm TiO2 catalysts. The areal rate was used for convenience since the exact nature and population of active sites were not known. The experiments were conducted in the flat-plate photoreactor11 at room temperature and under steady-state flow conditions. The photocatalytic oxidation reaction was run under dry condition. Both catalysts exhibit activation behavior, reaching their maximum conversion rates after 20 min of irradiation. Similar reaction behavior has been reported for photocatalytic oxidation of TCE over commercial anatase TiO2 powder.28 After activation, P6 has a stable activity but the 5 nm TiO2 (i.e., P5a) displays a slow deactivation that was accompanied by a color change from white to light yellow. It is clear from Figure 3a that the 6 nm TiO2 is more active than the 5 nm catalyst. It also exhibits a higher mineralization rate as shown by the plot of CO2 production rate versus irradiation time (Figure 3b). The TCE conversion over commercial Degussa P25 catalyst is six times slower than P6 with an areal rate of 3 × 10-10 mol1 s-1 m-2. The corresponding CO2 production rate over P25 is 1 × 10-10 mol1 s-1 m-2. In-situ infrared reaction study provides real-time monitoring of transient events that are occurring on the catalyst during the reaction. The information on surface adsorbed species (i.e., reactants, intermediates, and products) obtained from the insitu study give important insight on the reaction mechanisms and kinetics.29 Infrared spectroscopy is also an invaluable technique for the characterization of surface hydroxyl groups that populate the oxide surfaces (e.g., TiO2). Figure 4a displays a set of infrared transmittance spectra obtained during the

photocatalytic oxidation of TCE over 5 nm TiO2 catalyst (P5a). The spectra were corrected using the clean TiO2 as the background. Prior to UV illumination (t ) 0), the spectrum displays the characteristic TCE bands at 850 (δC-H), 950 (νC-Cl), 1255 (δCH-Cl), 1550 (νCdC), 1580 (νCdC), and 3070 cm-1 (νC-H).30 Upon irradiation, the intensity of TCE bands (e.g., 850, 950, 1255, and 3070 cm-1) display a monotonic decrease with time, while the signals corresponding to CO2 (i.e., 2340 and 2360 cm-1) and dichloroacetaldehyde (i.e., 1230 (δC-H), 1400 (δCHdO), and 1740 cm-1 (νCdO)29) increase as the reaction progresses. It is evident from the time-resolved spectra that both associated and isolated hydroxyl groups native to the catalyst (i.e., the broad band (3200-3500 cm-1) centered at 3400 cm-1 and the band at 3687 cm-1, respectively19) were either being consumed or desorbed during the reaction. In contrast, water associated with the bands at 1630 and 3100 cm-1 was produced, as is expected for the photooxidation of TCE. Figure 4b displays the postreaction infrared spectrum of TiO2 catalyst immediately after turning the UV illumination off (Figure 4b-1), and after 32 h of purging using dry air (800 cm3/ min) at room temperature (Figure 4b-2). Both spectra used the clean TiO2 wafer as the background reference. Spectrum 1 shows the presence of adsorbed TCE reactant as well as carbon dioxide, dichloroacetaldehyde, and water. TCE and carbon dioxide were removed within 10 minutes of purging, whereas the water was desorbed after 80 min. An additional peak at 1600 cm-1 was revealed after desorption of water (Figure 4b) which had been assigned to νc)c of dichloroethylene by Phillips and Raupp,29 however the accompanying signal for δCH-Cl (1302


Yeung et al.

Figure 5. (a) Plots of the intensity of νC-Cl (950 cm-1) for TCE molecule and (b) Ln (I950) versus irradiation time for different TiO2 crystal sizes (] P3, O P5a, 0 P6), (c) plots of the intensity of I950, and (d) Ln (I950) versus irradiation time for different TiO2 crystallinity (O P5a, 0 P5b, ] P5c).

cm-1) was missing. The signals at 1600 and 1360 cm-1 had also been assigned to adsorbed formic acid on TiO2. However, gas chromatographic analyses indicate that dichloroethylene is the more likely candidate. Continuous purgings with dry air were unable to remove the adsorbed dichloroacetaldehyde (i.e., 1230, 1400, and 1740 cm-1) and dichloroethylene. These signals remained unchanged even after UV irradiation in flowing dry air (i.e., room temperature and 6 h) and were only removed when the sample was irradiated under humid condition (80 % R. H.). It is plausible that these strongly adsorbed molecules play a role in the observed catalyst deactivation. Figure 5a plots the intensity of the characteristic νC-Cl peak (950 cm-1) for TCE as a function of irradiation time for the 3, 5, and 6 nm TiO2 catalysts. The intensity of the νC-Cl peak displays a monotonic decrease with time as TCE is consumed by the photocatalytic reaction. It is clear from the plots that the three catalysts differ in their activity. TCE conversion is highest for P6 followed by the P5a with the 3 nm catalyst (P3) being the least active. Several studies31-33 have reported that the photocatalytic oxidation of TCE over anatase TiO2 fit to a Langmuir-Hinshelwood expression where the reaction is first order with respect to trichloroethylene at low TCE concentrations. Figure 5b shows the results of fitting the intensity of the νC-Cl peak to a first-order kinetic model. The agreement between the experimental data and model is good, and the slopes (m) of the lines are 0.006, 0.010, and 0.011 for 3, 5, and 6 nm TiO2, respectively. Since the surface area of P6 is half that of the 5 nm TiO2 (cf. Table 2), it means that it has a higher areal rate than the latter catalyst. Figure 6a-c plots the intensity of the signals corresponding to δCHdO (1400 cm-1) and νCdO (1740 cm-1) belonging to dichloroacetaldehyde, νCdO (2340 cm-1) of carbon dioxide and δH-O-H (1630 cm-1) of strongly bounded water for 3, 5, and 6 nm TiO2 catalysts, respectively. The plots show that the amount of dichloroacetaldehyde, carbon dioxide,

and water increases steadily as the reaction progresses. The presence of dichloroethylene cannot be ascertained due to the signal interference from adsorbed water molecules. It is clear from the figures that the product distributions are different for the three catalysts. The 3 and 5 nm TiO2 catalysts produce more dichloroacetaldehyde as compared to the 6 nm TiO2, which generates more carbon dioxide during the photoreaction. Greater quantities of water are also produced by the 3 and 5 nm TiO2 at the expense of the weakly bound water (i.e., ∼3400 cm-1) native to the catalyst. The signal for δH-O-H for P6 at 1630 cm-1 is weak and there is no noticeable change in the intensity of the signal centered at 3400 cm-1. Unlike the other two catalysts, P6 displays little deactivation during TCE photocatalytic oxidation (up to 5 h) and the catalyst does not show any discoloration after the reaction. A closer inspection of Figure 6a indicates that dichloroacetaldehyde is produced first, before carbon dioxide. The latter only becomes significant after 40 min of UV irradiation. This suggests a stepwise oxidation of trichloroethylene toward the final mineral products (i.e., CO2, H2O, Cl2, and HCl). This is not inconsistent with the previous mechanism proposed in the literature.29,34 Water (δH-O-H) is detected immediately upon UV irradiation and must have originated from the catalyst rather than the product of photomineralization of TCE. Indeed, water production is usually accompanied by a decrease in the signal intensity of 3400 cm-1 assigned to hydroxyl groups associated by hydrogen bonding on the TiO2 catalyst surface. The most noticeable difference between the 5 and 3 nm TiO2 is the production of more carbon dioxide (Figure 6b). Figure 6c shows a significant increase in carbon dioxide production for the 6 nm TiO2 at the expense of dichloroacetaldehyde, indicating the high mineralization rate that could be obtained from this catalyst. The intensity of the water peak at 1640 cm-1 is significantly lower when compared to the other two catalysts. Qualitatively,



Figure 6. Product distribution from TCE photocatalytic oxidation over (a) P3 (3 nm TiO2), (b) P5a (5 nm TiO2), (c) P6 (6 nm TiO2), (d) P5b (crystallinity ) 0.8), and (e) P5c (crystallinity ) 0.5). (O I1640, b I2343, 0 I1400, 9 I1740).

the results of the flow reactor (i.e., flat-plate photoreactor) and batch reactor (infrared cell) are in agreement that the 6 nm TiO2 is the better photocatalyst for TCE mineralization. The nanostructured TiO2 catalyst displayed a progressive decrease in their areal conversion rates despite the increase in their surface area as the particle size diminishes. This phenomenon could be attributed to electronic and ensemble effects in nanometer sized particles. It is clear from the UV spectra and X-ray absorption spectroscopy (i.e., XANES and EXAFS) that the decrease in TiO2 crystal size is accompanied by a blue shift (Figure 2), and a significant change in the local structural environment of the catalyst (Table 2). The number of defect sites that stabilize the low coordinated Ti4+ cations was also observed by EPR spectroscopy to increase with decreasing crystal size.17 The presence of subsurface defects directly affects the photocatalytic properties of the TiO2 nanoparticles because they act as trapping sites for photogenerated holes, preventing them from reaching the surface region to participate in the reaction. OH• radicals on TiO2 surface are considered to be important active species in photocatalytic oxidation reactions. They are formed from the interactions between photogenerated holes and surface hydroxyl groups (i.e., basic OH). The trapping of photogenerated holes in subsurface defect sites not only hinders the formation of OH• radicals, but also lead to electronhole recombination. This explains the significant difference in the performance of the 5 and 6 nm TiO2 (Figure 3a,b). It is safe to attribute most of the observed activity and selectivity behavior of 5 and 6 nm TiO2 to the catalyst structure (i.e., ensemble effects) since they differ in crystal size by less than a nanometer (Table 2) and have comparable band-gap energies (Figure 2). Crystallinity. A set of TiO2 catalysts (i.e., P5a, P5b, and P5c) with similar crystal size, structure, and band-gap energy, but

of different crystallinity (cf. Table 2 and Figure 1b) were prepared using the modified sol-gel method. The TiO2 samples also have comparable BET surface area (Table 2). Using P5a as the reference, the crystallinity of P5b and P5c were assigned the values of 0.8 and 0.5, respectively. Figure 3c,d displays the plots of areal TCE conversion rate and CO2 production rate as a function of UV irradiation time for TCE photocatalytic oxidation in a flat-plate photoreactor. The plots show that the three 5 nm photocatalysts have comparable activity for TCE conversion (Figure 3c), but the least crystalline TiO2 (i.e., P5cS100) exhibits a significantly lower mineralization rate (Figure 3d). All three catalysts experienced slow deactivation and displayed a color change from white to light yellow at the end of reaction. The results of the in-situ infrared reaction study (Figure 5c,d) also show that the TCE conversions for the three catalysts are comparable and unaffected by the difference in TiO2 crystallinity. Figures 6b,d,e plot the intensity of the infrared signals corresponding to δCHdO (1400 cm-1), δH-O-H (1630 cm-1), νCdO (1740 cm-1) and νCdO (2343 cm-1) for P5a, P5b, and P5c, respectively. The plots show that the amount of dichloroacetaldehyde, carbon dioxide, and water increases steadily with time, however the product distribution is different for the three catalysts. The P5a and P5b produce the most carbon dioxide, while the P5c produces the least. This is in general agreement with the results of the flat-plate photoreactor experiments (Figure 3d). It is interesting to note that the signal for δH-O-H at 1630 cm-1 is lower for less crystalline TiO2 (Figure 6e), although infrared study indicates that the surface of all three catalysts is comparably hydroxylated. XRD, XANES, and EXAFS analyses indicated that the three TiO2 catalysts have comparable size, crystal phase, and atomic structure (Table 2). The catalysts also have similar BET surface area and surface hydration as measured by N2 physi-adsorption

4616 J. Phys. Chem. B, Vol. 106, No. 18, 2002 experiment and infrared spectroscopy. UV spectroscopy and XPS analyses of the catalysts also yield identical band gap energy and valence band structure. The TiO2 differs mainly in their crystallinity as measured by XRD. Crystallinity gave an indirect qualitatiVe measure of the structural defects present in the catalyst materials. However, the exact nature and number of these defects are difficult if not impossible to characterize and quantify. Their presence is known to affect the migration of photogenerated charges and their recombination rate. This is expected to directly impact the photocatalytic activity of the nanostructured TiO2. It is clearly manifested by the increased trapping of photogenerated holes at subsurface region of the catalyst and concomitant decrease in OH• radicals as detected by EPR.17 This suggests that the lower mineralization rate of the less crystalline 5 nm TiO2 is directly related to the greater number of defects in its structure. Concluding Remarks It is important to elucidate the role of structure, chemistry and electronic on catalytic activity of TiO2 nanoparticles in order to further enhance their performance as an environmental photocatalyst. The ability to engineer the phase structure, size and crystallinity of the nanometer sized TiO2 enabled us to separately probe the effects of structure (i.e., crystal and aggregate size)11,12 and electronics,12,17 as well as surface hydroxyl group19 on the photocatalytic activity of the nanostructured TiO2. This work demonstrated that the crystal size and crystallinity of the nanometer sized TiO2 can be precisely engineered using the modified sol-gel method. Reaction study conducted using gas-phase photooxidation of TCE as a probe reaction show that both crystal size and crystallinity affect the catalyst activity. Although the influence of crystallinity on TCE conversion rate is weak, it has a significant effect on the TCE mineralization. The results of this study suggest that an effective TiO2 photocatalyst must have large surface area populated by low coordinated surface sites (i.e., corners, edges, kinks, and defects) that are abundantly hydroxylated. This can be achieved by decreasing the crystal size and by using hydrothermal crystallization method. However, it is also necessary that the nanometer sized TiO2 should be well crystallized with minimal bulk defects that may trap migrating photogenerated charges. This latter requirement is far more difficult to satisfy as size of the TiO2 crystal diminishes. Further study will be conducted to formulate new synthesis strategy for reducing the number of bulk defect sites in nanostructured TiO2 (dp < 6 nm). Acknowledgment. The authors gratefully acknowledge the following funding from the Hong Kong Research Grant Councils (RGC-HKUST 6247/00P) and the Institute of Nano Science and Technology (INST). We also thank the Synchrotron Radiation Research Center (SRRC) of Taiwan for providing the beam time

Yeung et al. and facility for conducting the XRD, XANES, and EXAFS studies. We also acknowledge the help of Dr. Miguel Angel Ban˜ares of Instituto de Cata´lisis y Petroleoquı´mica for conducting the MicroRaman study, and Ms. Willma Lau Wai Ngar for conducting some of the photoreaction experiments. References and Notes (1) Hoffmann, M. R.; Martin, S. T.; Choi, W.; Bahnemann, D. W. Chem. ReV. 1995, 95, 69. (2) Mills, A.; Davies, R. H.; Worsley, D. Chem. Soc. ReV. 1993, 417. (3) Fox, M. A.; Dulay, T. Chem. ReV. 1993, 93, 341. (4) Nimlos, M. R.; Jacoby, W. A.; Blake, D. M.; Milne, T. A. EnViron. Sci. Technol. 1993, 27, 732. (5) Muggli, D. S.; Larson, S. A.; Falconer, J. L. J. Phys. Chem. 1996, 100, 15886. (6) Maira, A. J.; Soria, J.; Augugliaro, V.; Loddo, V. Chem. Biochem. Eng. Q. 1997, 11 (2), 89. (7) Djeghri, N.; Teichner, S. J. J. Catal. 1980, 6, 99. (8) Blake, N. R.; Griffin, G. L. J. Phys. Chem. 1988, 92, 5697. (9) Luo, Y.; Ollis, D. F. J. Catal. 1996, 163, 1. (10) Ameen, M. M.; Raupp, G. B. J. Catal. 1999, 184, 112. (11) Maira, A. J.; Yeung, K. L.; Yan, C. Y.; Yue, P. L.; Chan, C. K. J. Catal. 2000, 192, 185. (12) Maira, A. J.; Yeung, K. L.; Soria, J.; Coronado, J. M.; Belver, C.; Lee, C. Y.; Augugliaro, V. Appl. Catal. B: EnViron. 2001, 29, 327. (13) Anpo, M.; Shima, T.; Kodama, S.; Kubokama, Y. J. Phys. Chem. 1987, 91, 4305. (14) Xu, N.; Shi, Z.; Fan, Y.; Dong, J.; Shi, J.; Hu, M. Z.-C. Ind. Eng. Chem. Res. 1999, 38, 373. (15) Wang, C.-C.; Zhang, Z.; Ying, Y. J. Nanostr. Mater. 1997, 9, 583. (16) Zhang, Z.; Wang, C.-C.; Zakaria, R.; Ying, Y. J. Phys. Chem. B 1998, 102, 10871. (17) Coronado, J. M.; Maira, A. J.; Conesa, J. C.; Augugliaro, V.; Soria, J. Langmuir 2001, 17, 5368. (18) Cao, L.; Huang, A.; Spiess, F. J.; Suib, S. L. J. Catal. 1999, 188, 48. (19) Maira, A. J.; Coronado, J. M.; Augugliaro, V.; Yeung, K. L.; Conesa, J. C.; Soria, J. J. Catal. 2001, 202, 413. (20) Zaban, A.; Aruna, S. T.; Tirosh, S.; Gregg, B. A.; Mastai, Y. J. Phys. Chem. B. 2000, 104, 4130. (21) Cao, L.; Gao, Z.; Suib, S. L.; Obee, T. N.; Hay, S. O.; Freihaut, J. D. J. Catal. 2000, 196, 253. (22) Brus, L. J. Phys. Chem. 1986, 90, 2555. (23) Kariyone, T.; Anpo, M.; Chiba, K.; Tomonari, M. Hyomen (surface) 1991, 29, 156. (24) Dagn, G.; Sampath, S.; Lev, O. Chem. Mater. 1995, 7, 446. (25) Henglein, A. Chem. ReV. 1989, 89, 1861. (26) Iida, Y.; Furukawa, M.; Aoki, T.; Sakai, T. Appl. Spectrosc. 1998, 52, 673. (27) Chen L. X.; Rajh, T.; Ja¨ger, W.; Nedeljkovic, J.; Thurnauer, M. C. J. Synchrotron Radiat. 1999, 6, 445. (28) D’Hennezel, O.; Ollis, D. F. J. Catal. 1997, 167, 118. (29) Phillips, L. A.; Raupp, G. B. J. Mol. Catal. 1992, 77, 297. (30) Dreiessen, M. D.; Goodman, A. L.; Miller, T. M.; Zaharoies, G. A.; Grassian, V. H. J. Phys. Chem. B 1998, 102, 549. (31) Larson, S. A.; Falconer, J. L. Appl. Catal. B: EnViron. 1994, 4, 325. (32) Wang, K.-H.; Tsai, H.-H.; Hsieh, Y.-H. Appl. Catal. B: EnViron. 1998, 17, 313. (33) Yamazaki-Nishida, S.; Nagano, K. J.; Phillips, L. A.; CerveraMarch, S.; Anderson, M. A. J. Photochem. Photobiol. A: Chem. 1993, 70, 95. (34) Pruden, A.; Ollis, D. J. Catal. 1983, 60, 404.

Ensemble Effects in Nanostructured TiO - American Chemical

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