Mechanistic Studies of Photocatalytic Reaction of Methanol for

Photoassisted production of hydrogen from water or organic compounds has been of ..... the presence of CH3OH, the photogenerated holes in TiO2 film de...
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J. Phys. Chem. C 2007, 111, 8005-8014

8005

Mechanistic Studies of Photocatalytic Reaction of Methanol for Hydrogen Production on Pt/TiO2 by in situ Fourier Transform IR and Time-Resolved IR Spectroscopy Tao Chen, Zhaochi Feng, Guopeng Wu, Jianying Shi, Guijun Ma, Pinliang Ying, and Can Li* State Key Laboratory of Catalysis, Dalian Institute of Chemical Physics, Chinese Academy of Sciences, P.O. Box 110, Dalian 116023, China ReceiVed: February 6, 2007; In Final Form: March 30, 2007

The anaerobic photocatalytic reaction of methanol on Pt/TiO2 catalyst was studied by in situ Fourier transform IR and time-resolved IR spectroscopy. For the Pt/TiO2 catalysts reduced at high temperature, the capacity of methanol adsorption decreases with the increase of the Pt loading, indicating that Pt particles or atoms occupy some of the active sites on TiO2 for methanol adsorption. Surface species CH2O(a), CH2OO(a), and HCOO(a) are derived from the photocatalytic reaction of methanol on Pt/TiO2. The increase of gas-phase methanol or water accelerates the photoreaction and improves the activity of H2 production. When the catalysts are exposed to methanol, the strong electronic absorption on nanosecond to second time scale is observed, indicating that a great amount of long-lived electrons are produced in TiO2-based photocatalysts after band gap excitation. The decay rate of the long-lived electrons correlates well with the activity of H2 production. These results show that the long-lived electrons contribute to the H2 generation and the decays of the long-lived electrons on millisecond to second time scale in Pt/TiO2 are ascribed to the reaction for H2 evolution: etr-[Pt] + H+fH‚f1/2H2. The function of the molecularly adsorbed methanol or water is found to mediate the proton transfer on the TiO2 surface. The activities of H2 production under steady-state irradiation conditions were also measured, and it is deduced that the yield of the long-lived electrons could be responsible for the activity of H2 production.

1. Introduction Photoassisted production of hydrogen from water or organic compounds has been of considerable interest due to the potential utilization of solar energy.1-9 One of the most extensively studied photocatalysts is the TiO2-based catalyst, which has been widely used in many photocatalytic reaction systems due to its high activity and high stability.10 The activity of H2 production is dependent on many factors such as sacrificial reagents, surface properties of TiO2, cocatalysts, catalyst preparation, and so on.2-6,11-14 Because the photocatalytic processes are very complicated, the detailed mechanisms of the photoreactions are still not well understood. In-situ IR is a powerful tool for the investigation of the surface properties of catalysts and the nature of the adsorbed species.15 Reaction mechanisms of photocatalytic degradation of organic compounds have been intensively studied by Fourier transform infrared (FTIR) method16-20 but few FTIR studies on anaerobic photooxidation of organic compounds for H2 production were reported.21,22 The kinetics of the photogenerated charge carriers reflects the intrinsic properties of the photocatalysts and directly influences the efficiency of the photocatalytic reactions. The studies on the kinetics of the photogenerated charge carriers are crucial for understanding the photocatalytic processes. It was reported that the excited electrons in conduction band (CB) and the positive holes in valence band (VB) can be trapped by respective shallow trap states in femtosecond time scale.23-25 The trapped holes and electrons show broad absorption in the * To whom correspondence should be addressed. Email: [email protected]. Tel: +86 411 84379070. Fax: +86 411 84694447.

UV-vis region, centered around 430-520 nm24,26,27 and 600800 nm,24,26-30 respectively. The absorption in mid-IR region is due to the conduction electrons and/or shallowly trapped electrons.31-34 Because there is no interference of the holes in mid-IR region and the shape of the mid-IR absorption spectra is not sensitive to surrounding conditions, time-resolved midIR absorption spectroscopy is potentially a powerful tool to accurately trace the kinetics of the electrons. The ultrafast electron transfer from adsorbed sensitizers to the nanocrystalline TiO2 semiconductor has been investigated by femtosecond timeresolved mid-IR absorption spectroscopy.35-37 Recently, Yamakata et al. utilized a homemade mid-IR absorption spectroscopy with nanosecond to second time resolution to study the decay kinetics of photogenerated electrons in bare TiO2,38 platinized TiO2,38-40 dye-sensitized TiO2,41,42 and NaTaO3 catalysts.43 Their results demonstrated the powerful method of nanosecond to second time-resolved IR spectroscopy being able for kinetic analysis of the photocatalytic reactions. When methanol is present, they observed that the amount of longlived electrons in platinized TiO2 increases significantly after 50 ns.44,45 They proposed that adsorbed methoxy groups capture the holes within 50 ns. This has been demonstrated by a recent femtosecond time-resolved UV-vis spectroscopy study in which the authors directly observed that most of the holes are captured by adsorbed methanol within 1 ns.46 Yamakata et al. also found that the slow decay of the long-lived electrons on millisecond to second time scale can be accelerated by increasing methanol or water in the gas phase, and the slow electron decay was ascribed to electron capture by molecularly adsorbed methanol or water.44,45 Because the long-lived electrons successfully escape from the fast recombination with the holes and they have more opportunities to react with the surface adsorbed molecules,

10.1021/jp071022b CCC: $37.00 © 2007 American Chemical Society Published on Web 05/16/2007

8006 J. Phys. Chem. C, Vol. 111, No. 22, 2007 the long-lived electrons may play an important role in the photocatalytic reactions. In this work, the in-situ FTIR and nanosecond to second timeresolved IR measurements were combined to investigate the mechanisms of photocatalytic reaction of methanol for H2 production on Pt/TiO2 catalyst. Our studies were focused on the slow kinetics of the long-lived electrons and the relationship between the decays of the long-lived electrons and the activities of H2 production. We found that the long-lived electrons contribute to the H2 generation and the electron decays on millisecond to second scale in Pt/TiO2 correspond to the electron consumption for H2 generation. The function of the molecularly adsorbed methanol or water is found to mediate the proton transfer on TiO2 surface. A possible reaction model for anaerobic photocatalytic reaction of methanol is proposed. 2. Experimental Section 2.1. Sample Preparation. Pt-loaded TiO2 catalysts, Pt/TiO2, were prepared by impregnation method. The commercially available TiO2, generally known as degussa P25 (20% rutile, 80% anatase), was used as the support. Briefly, 1 g of P25 was dispersed into 30 mL of H2PtCl6 aqueous solution at a given concentration and stirred for 6 h. After evaporation to dryness at 373 K, the obtained dry powder was then reduced with H2 at 673 K for 2 h. The final products are designated as Pt/TiO2 (n wt %), where n represents the value of weight percentage of Pt. The TiO2 sample used in the measurements was prepared and pretreated in an identical way to other Pt/TiO2 (n wt %) samples, except without coadding Pt. 2.2. In situ FTIR Measurement. In-situ IR spectra were recorded on a Nicolet 870 FTIR spectrometer with a MCT detector. The TiO2 or Pt/TiO2 sample was pressed into a selfsupporting wafer (about 20 mg) and put into a quartz IR cell with BaF2 windows. The IR cell was connected to the vacuum system. Before each run of the experiment, the sample in the cell was heated to 573 K under vacuum (10-3 Torr) for 1 h. After the heating, when the sample was cooled to room temperature under the protection of N2 atmosphere, the IR cell with the sample was evacuated and the methanol vapor was then introduced. In the photocatalysis study, the laser at 355 nm with 10 ns width from third harmonic generation of a Q-switched Nd: YAG laser (Labeit Beijing) was used as the light source. The frequency of the pulse laser was set to 5 Hz and the power was 70 mW at 5 Hz. The diameter of the light spot at the sample was 8 mm. The speed of the photocatalytic reaction increases with increasing of the laser power from 30-100 mw. To prevent the scattered light from reaching interferometer and detector optics, two AR-coated Ge plates were placed in the openings of the sample chamber.47,48 After the sample was irradiated for an indicated time, the difference IR spectra were collected using the background spectrum taken before the irradiation. All the infrared spectra were collected with a resolution of 4 cm-1 and 64 scans in the region of 4000-1000 cm-1. 2.3. Time-Resolved IR Measurement. The transient IR absorption signals were recorded on the same Nicolet 870 FTIR spectrometer with the MCT detector. This instrument provides a step scan time-resolved measurement mode. In the step scan time-resolved FTIR experiments, the AC output (20 MHz to 110 Hz) of the photovoltaic MCT was used to measure the transient signal changes. The AC output, further amplified by a factor of 50 with a SR560 preamplifier (Stanford Research Systems, 1 MHz-0.03 Hz), was recorded with a 50 MHz, 20bit A/D converter supplied with the spectrometer. The static

Chen et al. interferogram was recorded from the DC output of the detector preamplifier. The absorbance difference spectrum was calculated by the relation: ∆A(ν˜ , t) ) -log{[I(ν˜ ) + ∆I(ν˜ , t)/g]/I(ν˜ )} (I(ν˜ ), the static single beam spectrum is the Fourier transform of DCcoupled interferogram; ∆I(ν˜ , t), laser-induced spectrum is the Fourier transform of AC-coupled interferogram; and g is the gain of the AC-coupled preamplifier).47,48 The same pulse laser at 355 nm (3 Hz, 8 mJ/pulse) was used to photoexcite the samples. The synchronization between laser excitation and data acquisition was achieved with a Stanford Research model DG 535 pulse generator. The instrument operation and data acquisition were all controlled by Nicolet Omnic 6.2 software package. Because of the limited low cutoff frequency of the AC amplifier in the MCT and some other technical difficulties, very slow signals ranging from millisecond to second are not suitable to be measured by step scan measurement mode. To observe the slow signals, we slightly changed the measurement method. The stepping mirror was stopped to one fixed position. The DC output (20 MHz to 0 Hz) of the MCT amplified by AC-coupled SR560 preamplifier (1 MHz to 0.03 Hz) was used to measure the transient signals (∆I′(t)). The frequency of the 355 nm laser pulse was set to 0.1 or 0.01 Hz and the energy was 2-3 mJ/ pulse. The laser-induced transient signals were accumulated and averaged in the external oscilloscope (Tektronix, TDS5104). From the knowledge of circuit analysis, the recorded transient signals (∆I′(t)) is the convolution of ∆I(t) and g(t)

∆I′(t) )

∫-∞∞ ∆I(t′)g(t - t′) dt′

where g(t) represents the unit-impulse response of the whole electronic systems and ∆I(t) represents the undistorted curve. g(t) was determined experimentally. Through the deconvolution, the true transient curve ∆I(t) was obtained. The deconvolution process was accomplished by a Fourier Transform method49,50 without using any filter functions. Before measuring the transient signals, the DC output of MCT modulated by a handmade chopper was used to measure the intensity of the whole IR light in 1000-4000 cm-1 (I) and the average transient absorbance changes were calculated as ∆A(t) ) -log{[I + ∆I(t)/g]/I}. The transient average absorbance changes, ∆A(t), reflect the transient absorption of the probe light from 1000 to 4000 cm-1 by the excited samples. All the samples measured in time-resolved IR experiments were pretreated in the IR cell under the same condition as that for in-situ FTIR experiments. 2.4. Photocatalytic Activity. The photocatalytic H2 production under steady-state irradiation conditions was investigated in a Pyrex reaction cell connected to a closed gas circulation and evacuation system.51 The light source was a 300 W Xe lamp with a water-cooled quartz jacket. The catalyst (0.3 g) was suspended in an aqueous solution containing 160 mL of H2O and 40 mL of CH3OH. Before the experiment, the reaction mixture was deaerated thoroughly. The evolved H2 was measured by online gas chromatography. The H2 produced in the in-situ IR cell was sampled and analyzed by the same evacuation system and online gas chromatography. The apparent quantum yields were estimated from the ratio of twice the number of hydrogen molecules to the number of incident photons. 3. Results and Discussion 3.1. FTIR Studies of Methanol Adsorption. Figure 1 shows the FTIR spectra taken after the adsorption of CH3OH on Pt/

Photocatalytic Reaction of Methanol on Pt/TiO2

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Figure 1. FTIR spectra of adsorbed CH3OH on Pt/TiO2 samples at room temperature. The subtracted spectra of the catalysts exposed to about 2 Torr CH3OH vapor followed by evacuation for 15 min are shown.

TABLE 1: BET Surface Area of Various Pt/TiO2 Samples sample surface area (m2 g-1)

Figure 2. The difference IR spectra of TiO2 adsorbed with CH3OH(a) and CH3O(a) before and after being irradiated for 28, 38, 52, and 65 min by 5 Hz, 70 mW, 355 nm pulse laser. The CH3O(a)/TiO2 surface was prepared by exposing a pretreated sample surface to adequate CH3OH vapor followed by evacuation at room temperature for 15 min.

Pt/TiO2 Pt/TiO2 Pt/TiO2 Pt/TiO2 TiO2 (0.02 wt %) (0.1 wt %) (0.2 wt %) (1.0 wt %) 56

55

53

54

52

TiO2 and then after the evacuation at room temperature. The spectra of the catalysts recorded in vacuum are subtracted. It is observed that undissociated adsorbate CH3OH(a) and dissociated adsorbate CH3O(a) are formed on the catalyst surface. They are distinguishable by the characteristic frequencies of symmetric and antisymmetric CH3 stretching at 2843 and 2944 cm-1 for CH3OH(a) and 2818 and 2923 cm-1 for CH3O(a).52-54 The C-O stretching vibrations of CH3O(a) and CH3OH(a) in 10001200 cm-1 region can also be clearly observed (not shown). In the O-H stretching region for surface hydroxyl groups, two negative peaks appear at 3726 and 3675 cm-1. There are mainly two reasons for the decrease of the surface HO(a) groups. First, the HO(a) groups have become hydrogen bonded to oxygen atoms of CH3OH(a) or CH3O(a) species, accompanied by developing a broad band in the hydrogen-bonded O-H stretching region (3100-3500 cm-1), as can be seen in Figure 1.52,53 Second, HO(a) groups react with methanol as53,55

OH-(a) + CH3OH(g) f CH3O-(a) +H2O(g)

(1)

where (a) and (g) represent the adsorbate and gas-phase species, respectively. As shown in Figure 1, the IR absorbance of CH3OH(a) and CH3O(a) dramatically decreases with the Pt loading increasing. Though it is difficult to precisely control the thickness of the sample wafer in every IR adsorption experiment, the drastic difference and well-reproducible results can still reflect some useful information. The Brunauer-Emmett-Teller (BET) surface area data given in Table 1 show that the loading of Pt on TiO2 particles does not induce appreciable changes in the surface area, indicating that the slight decrease of surface area is not responsible for the much lower adsorption capacity of CH3OH on Pt/TiO2 than on TiO2. Transmission electron microscopy (TEM) images (not shown) indicate that the Pt particles on TiO2 are smaller than 1 nm in diameter. The possible reason for the decrease of CH3OH adsorption capacity on Pt/TiO2 is that the Pt particles occupy some of the

surface active sites for methanol adsorption. The active sites are mainly oxygen vacancies detected under the ultrahigh vacuum conditions as H2O or CH3OH preferably dissociates on the oxygen vacancy sites.56-58 Because the concentration of oxygen vacancies on the TiO2 surface is relatively small, the small amount of Pt loadings may effectively affect these active sites. The oxygen vacancies are considered to be the nucleation centers for small metal clusters.59 As for the Pt-loaded TiO2 catalysts reduced in H2 at 673 K, the strong metal-support interaction effect occurs60 and many experimental results indicated that the bonding interaction between metal and reduced titanium ions (e.g., Ti3+) exists.61 The strong interaction of PtTi3+ leads to the decrease of Ti3+ sites (oxygen vacancy sites) for methanol adsorption. 3.2. Photocatalytic Reaction of Methanol on Pt/TiO2 Samples. Figure 2 shows the difference spectra taken before and after the UV irradiation with a different time for the anaerobic photocatalytic reaction of methanol adsorbed on TiO2 sample. During the UV irradiation, the temperature of the sample was only ∼50 °C, so the possible contribution from thermal reaction can be ruled out. The UV light-induced broad IR absorption from 4000 to 1000 cm-1 is originated from the photogenerated electrons in the CB and/or in shallow traps below the CB, which will be discussed further in the next section. The negative IR bands in 2800-2950 cm-1 and a weak and broad positive band in 1400-1100 cm-1 appear after the UV irradiation. The negative bands can be attributed to the desorption and reaction of CH3OH(a) and CH3O(a). The assignment of the broad positive band is unclear yet. All these features show that the photoreaction of CH3OH adsorbed on TiO2 is very slow. When the CH3O(a)/TiO2 surface was exposed to 1 Torr H2O, new bands at 1563, 1381, and 1360 cm-1 appear and grow with the UV irradiation as shown in Figure 3. The bands at 1560 and 1360 cm-1 are ascribed to -COOantisymmetric stretching mode and -COO-symmetric stretching mode of adsorbed formate species respectively, and the band at 1381 cm-1 is due to C-H deformation or -COO- rocking of adsorbed formate.18,62 Compared to the photoreaction without H2O, the presence of H2O slightly enhances the formate formation.

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Figure 3. The difference IR spectra taken before and after indicated UV irradiation time during the photocatalytic reaction of methanol on TiO2. The 5 Hz, 70 mW, 355 nm pulse laser was used as the UV photolysis light. Before UV irradiation, a pretreated sample surface was exposed to adequate CH3OH vapor followed by evacuation. Then ∼1 Torr H2O vapor was introduced in the cell.

Figure 4. The difference IR spectra of Pt/TiO2 (0.5 wt %) adsorbed with CH3OH(a) and CH3O(a) before and after being irradiated for 20, 60, 100, and 156 min by 5 Hz, 70 mW, 355 nm pulse laser. The CH3O(a)/Pt/TiO2 (0.5 wt %) surface was prepared by exposing a pretreated sample surface to adequate CH3OH vapor followed by evacuation at room temperature for 15 min.

Figure 4 shows the difference spectra taken before and after the UV irradiation for the photocatalytic reaction of methanol adsorbed on Pt/TiO2 (0.5 wt %). The characteristic bands of CH3OH(a) and CH3O(a) greatly decrease in intensity and new bands due to formate species at 2960, 2859, 2733, 1579, 1561, 1380, and 1361 cm-1 grow remarkably with the UV irradiation time. In addition to these bands, some weak bands at 1206, 1174, 1146, 1114, and 1065 cm-1 also appear. The band at 1174 cm-1 is ascribed to molecularly adsorbed formaldehyde [CH2O(a)] and the bands at 1206, 1146, 1114, and 1065 cm-1 are ascribed to dioxymethylene species [CH2OO(a), DOM(a)].62,63 In the photoreaction process, both monodentate formate and bidentate formate species are formed on the surface of the catalyst. The bands at 1579 and 1561 cm-1 correspond to the -COOantisymmetric stretch mode of monodentate formate and bidentate formate species, respectively.64 From Figure 4, it can

Chen et al.

Figure 5. The difference IR spectra taken before and after indicated UV irradiation time during the photocatalytic reaction of methanol on Pt/TiO2 (0.5 wt %). The 5 Hz, 70 mW, 355 nm pulse laser was used as the UV photolysis light. Before UV irradiation, a pretreated sample surface was exposed to adequate CH3OH vapor followed by evacuation. Then ∼1 Torr H2O vapor was introduced in the cell.

Figure 6. Relative amounts of DOM(a) as a function of the UV irradiation time. The relative amounts are represented by the peak intensity of DOM(a) at 1117 cm-1 divided by the initial absorbance value of CH3O(a) at 2818 cm-1.

be seen that the bidentate formate species appear first, then the monodentate formate species appear. When the CH3O(a)/Pt/ TiO2 (0.5 wt %) surface was exposed to 1 Torr H2O, the same species are formed during the phtoreaction, as shown in Figure 5. The new band at 1315 cm-1 is assigned to molecularly adsorbed HCOOH(a).18,65 Figures 6 and 7 show the relative amount of DOM(a) and formate species as a function of irradiation time irrespectively. The amount of DOM(a) increases at first, then decreases with the UV irradiation time. The amount of formate continuously increases during the photoreaction. These phenomena manifest that DOM(a) species are a kind of surface intermediates and the adsorbed formate species are relatively stable on the surface of TiO2. During the photoreaction, DOM(a) species are formed and then transformed to other species quickly but the formate species transform to carbonate or CO2 slowly. The hydrogen generated in the IR cell was measured online. The apparent quantum yields of H2 evolution are also presented in Figure 7. When the TiO2 sample was used as the photocatalyst, no H2 was detected. When the CH3O(a)/Pt/TiO2 (0.5 wt %) surface was irradiated by 355 nm pulse laser, some H2 was generated.

Photocatalytic Reaction of Methanol on Pt/TiO2

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Figure 7. Relative amounts of HCOO(a) as a function of the UV irradiation time. The CH3O(a)/sample surface was prepared by (a) CH3OH adsorbed on Pt/TiO2 (0.5 wt %) followed by evacuation, then the CH3O(a)/Pt/TiO2 (0.5 wt %) surface was exposed to ∼1 Torr H2O vapor; (b) CH3OH adsorbed on Pt/TiO2 (0.5 wt %) followed by evacuation; (c) CH3OH adsorbed on TiO2 followed by evacuation, then the CH3O(a)/TiO2 surface was exposed to ∼1 Torr H2O vapor; (d) CH3OH adsorbed on TiO2 followed by evacuation. The relative amounts are represented by the peak intensity of HCOO(a) at 1360 cm-1 divided by the initial absorbance value of CH3O(a) at 2818 cm-1. The apparent quantum yields of H2 evolution are also shown.

When the CH3O(a)/Pt/TiO2 (0.5 wt %) surface was exposed to ∼1 Torr H2O and was irradiated by the same laser for the same time, more H2 was produced. From Figure 7 we can see that the photocatalytic reaction of methanol on Pt/TiO2 is much faster than on TiO2 and increasing H2O(g) can promote the photoreaction significantly. It was found that increasing pressure of gas-phase CH3OH can also promote the photoreaction and enhance the H2 generation. On the basis of the above results, possible photoreactions on the surface are proposed as

CH3OH f CH3O-(a) + H+ CH3O-(a) + h+ f CH3O‚(a) CH3O‚(a) + h+ f CH2O(a) + H+

(2)

‚‚‚‚‚ +

H + e- f 1/2 H2 At first, methanol adsorbs on TiO2 to form methoxy groups. Methoxy groups can capture the holes to form methoxy radicals.44,45,66 CH3O‚(a) can release protons and go on capturing the holes to form CH2O(a). CH2O(a) then transform to CH2OO(a) [DOM(a)] on titania surface.62 The generated CH2OO(a) species can further consume the holes and can release protons to form formate species that is transformed into carbonate or CO2 finally. In these photooxidation reactions, methanol-derived species are oxidized by holes directly and release H+ stepwise. For electron-consuming reactions, it is accepted that Pt can trap the photogenerated electrons within a few picoseconds25,67,68 and can catalyze the H2 formation reaction2,4,11

e- + H+ f H‚(a) f 1/2H2

(3)

Therefore, it is reasonable that Pt loaded on TiO2 strongly

Figure 8. Time-resolved IR absorption spectra of TiO2 in vacuum recorded by step scan measurement mode. 3 Hz, 8 mJ/pulse, 355 nm laser was used to excite the sample. The insert diagram shows the normalized decay curves at 1100, 1300, and 1500 cm-1.

promotes the whole photoreaction. But it is very interesting that increasing the pressure of CH3OH or H2O vapor accelerates the photoreaction effectively. Increasing the pressure of CH3OH or H2O in gas phase mainly increases the amount of molecularly adsorbed CH3OH(a) or H2O(a) species on the surface of titania.44,45 How does the molecularly adsorbed CH3OH(a) or H2O(a) promote the photoreaction? Can they accelerate the electron-consuming reaction or hole-consuming reaction? To further investigate the possible mechanisms, time-resolved IR experiments were carried out under the same conditions as in in-situ FTIR experiments. 3.3. Decay Kinetics of Photogenerated Electrons. Figure 8 shows the transient spectra of TiO2 in vacuum recorded by step scan measurement mode. A broad UV-induced IR absorption appears, and the absorption monotonously increases in intensity with lower wavenumbers. The broad and structureless IR absorption is assigned to intraband transitions of CB electrons or excitation of shallowly trapped electrons to the CB31,33,39,69 It has been reported31-34 that the broad IR absorption obeys the equation ∆Absorbance ) Aν˜ -p, where ν˜ is the wavenumber, A is a constant, and p is the scattering constant. According to the equation, the transient absorption decay profiles should be wavenumber independent. The inset of Figure 8 shows the decay curves at 1100, 1300, and 1500 cm-1 and they are identical indeed after normalization. Pt-loaded TiO2 sample gives the similar UV-induced IR absorption spectra (not shown). Figure 9 shows the transient average absorbance changes of Pt/TiO2 (0.2 wt %) measured in vacuum, in the presence of ∼1 Torr H2O and in the presence of ∼20 Torr CH3OH. The temporal profiles can reflect the decay tendency of the photogenerated electrons. The electron decay curve of the catalyst in vacuum shows a fast decay before 10 µs and the subsequent decay is very slow. The fast decay within 10 µs is mainly due to the recombination of the electrons with the holes. The minor part of the absorbance that decays slowly corresponds to the long-lived electrons. It has been reported that the photogenerated holes can be captured by surface hydroxyl groups.28,39,70 Once the holes are captured by surface hydroxyl groups, an equal number of the electrons become the long-lived electrons because the subsequent recombination of electrons with the holes captured by hydroxyls is very slow. In the presence of H2O, the decay curve is similar to that obtained in vacuum.

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Chen et al.

Figure 9. Transient profiles of average IR absorption of Pt/TiO2 (0.2 wt %) excited by 355 nm laser pulse of 10 ns duration. The pretreated sample in vacuum, exposed to ∼1 Torr H2O and exposed to ∼20 Torr CH3OH were measured, respectively. The curves were obtained at room temperature by accumulating 30 traces.

Figure 10. Normalized transient profiles of average IR absorption of TiO2 excited by 355 nm laser pulse of 10 ns duration. The pretreated sample exposed to (a) ∼20 Torr CH3OH and (b) followed by evacuation for 15 min were measured, respectively. The curves were obtained at room temperature by accumulating 20 traces repeated at 0.01 Hz.

In the presence of CH3OH, the initial absorbance is enhanced about 17 times as much as that observed in vacuum and the absorbance does not decay on the time scale shown in Figure 9. Comparing the absorbance at 50 µs, the amount of longlived electrons is enhanced about 50 times when CH3OH is present. Yamakata et al. reported that adsorbed methoxy groups capture the holes within 50 ns.44,45 Once the holes are captured by methoxy groups, the complementary electrons become the long-lifetime electrons because the subsequent recombination of electrons with the holes captured by methoxy groups is extremely slow. Above results also indicate that methoxy groups are much more effective than surface hydroxyls to capture the photogenerated holes, which is consistent with the results observed by Tamaki et al.,46 who found in the presence of CH3OH, the photogenerated holes in TiO2 film decay clearly within 1 ns but in the case of TiO2 film in air or water, the holes do not decay on that time scale. Rothenberger et al.28 reported that for the TiO2 semiconductor, the capturing of holes presumably by OH- groups required an average time of 250 ns. According to literature,25,28,30,46,67,68,71 the main ultrafast processes after band gap excitation (in the presence of CH3OH) are listed as follows and the characteristic time domains are given to the right of each step:

Figure 11. Normalized transient profiles of average IR absorption of Pt/TiO2 (0.2 wt %) excited by 355 nm laser pulse of 10 ns duration. The pretreated sample exposed to (a) ∼ 20 Torr CH3OH, followed by evacuated for (b) 0.5 min, (c) 2.5 min, and (d) 15 min were measured respectively. The curves were obtained at room temperature by accumulating 50 traces repeated at 0.1 Hz.

e- + h+ f hν or ∆ (light or heat) e- + Pt f etr-[Pt]

ps-10 µs

fs-10ps

h+ + CH3O(a)- f CH3O‚(a) + hν or ∆

(4) (5)

ps-1 ns

(6)

where e- represents CB electrons and shallowly trapped electrons, h+ represents VB holes and shallowly trapped holes, etr-[Pt] and CH3O‚(a) refer to electrons trapped by Pt and holes captured by CH3O(a)-, respectively. Equation 6 represents the process of methoxy groups capturing the photogenerated holes. In this process, methoxy groups provide the electrons to recombine with the photogenerated holes, and this ultrafast reaction is crucial for increasing the yield of long-lived CB electrons and shallowly trapped electrons. Here, the efficient factor of long-lived electrons, f, is defined as the ratio of the number of long-lived electrons to the total number of the

photogenerated electrons. Reaction 5 may also compete with reaction 4 and suppress the ultrafast recombination process.25,67,68 Figures 10 and 11 show the normalized IR absorption decay curves of TiO2 and Pt/TiO2 (0.2 wt %) measured under different methanol pressures. The electron decay in TiO2 is tentatively assigned to the recombination reaction

e- + CH3O‚(a) f CH3O(a)- + hν or ∆

(7)

The recombination process is extremely slow as shown in Figure 10. The decay rate of the electrons shows less change with increasing methanol pressures. Pt loading brings two distinct features: the electrons decay much faster, and the decays of the electrons are dramatically accelerated with increasing of methanol pressures. When CH3O(a)/Pt/TiO2 (0.2 wt %) surface is exposed to water vapor, the electron decays are also promoted

Photocatalytic Reaction of Methanol on Pt/TiO2 effectively (not shown), which is very different from that observed for TiO2 without Pt loading. Above results indicate that Pt loaded on TiO2 and the presence of CH3OH or H2O can significantly promote the decay of the long-lived electrons. Because electrons are already trapped by Pt particles on a picosecond time scale, reaction 5 is not responsible for the accelerating effect on the electron-consuming process in millisecond to second time region. Another explanation that Pt particles can promote the recombination of the longlived electrons with the holes trapped by methoxy groups (reaction 7) is also ruled out because the electrons are separated with the trapped holes by Pt trapping, and Pt loaded on TiO2 should restrain reaction 7 but not promote it. It can be seen from Figure 7 that the yields of H2 generated under the similar conditions correlate well with the rate of the electron decay. These results strongly support the explanation that Pt particles can catalyze reaction 3. Hence, the electron-consuming reaction on millisecond to second time scale is mainly via reaction 3. Some similar results that the reactant-induced acceleration of the electron-consuming reaction qualitatively correlates with the H2 production observed in steady-state UV irradiation were reported in studies of electron decay in NiO-loaded NaTaO343 and K3Ta3B2O1272 catalysts, but we first report that the decays of the electrons on millisecond to second time scale in Pt/TiO2 are mainly due to the direct reaction of electrons with protons to produce H2. Because the bottom of the conduction band of TiO2 lies slightly above the redox potential of the H+/H2 couple, only the conduction electrons and the shallowly trapped electrons are thermodynamically suitable to reduce H+ but the potential level of those deeply trapped electrons is too positive to reduce H+. Hence, above results further confirm that the electrons observed by mid-IR absorption spectroscopy are mainly conduction electrons and/or the shallowly trapped electrons rather than deeply trapped electrons. Now the question is why the photoreaction can be accelerated effectively by increasing the pressure of CH3OH or H2O vapor. It was reported that H2O dissociates on oxygen vacancy sites to form pairs of bridge-bonded hydroxyls57,58,73 and CH3OH dissociates on oxygen vacancy sites to form methoxy groups and bridge-bonded hydroxyls on the surface of TiO2.55,74 From the images of TiO2 (110) surface with high-resolution scanning tunneling microscopy, Wendt and co-workers58 observed that the protons in the pairs of bridge-bonded hydroxyls are stable without molecularly adsorbed H2O. When interacting with molecularly adsorbed H2O that can easily diffuse on the surface, the proton transfers can be observed with the result of net OHbr diffusion. Zhang et al.74 observed that the methoxy and hydroxyls occupy primarily neighboring bridge-bonded oxygen sites and the protons in the hydroxyls are very stable with methanol coverages below the bridge-bonded oxygen vacancy coverage. At higher methanol coverages, the molecularly adsorbed methanol appears and the facile hydroxyl migration through proton transfers assisted by mobile CH3OH(a) is observed. According to above results and discussions, it is suggested that the molecularly adsorbed CH3OH and H2O equilibrated with CH3OH(g) and H2O(g) can mediate H+ transfers on the surface of the catalysts. So, reaction 3 is accelerated significantly with increasing CH3OH(g) or H2O(g). In our experiments, it was observed that the effect of reactantinduced acceleration of electron decay is more sensitive to the pressure of H2O vapor than CH3OH vapor, which indicates H2O(a) mediate H+ transfers more effectively than CH3OH(a). Figure 12 shows the normalized IR absorption decay curves of Pt/TiO2 with a different loading amount of platinum. The

J. Phys. Chem. C, Vol. 111, No. 22, 2007 8011

Figure 12. Normalized transient profiles of average IR absorption of Pt/TiO2 samples excited by 355 nm laser pulse of 10 ns duration. The decay curves were recorded at room temperature by accumulating 50 traces repeated at 0.1 Hz when the pretreated samples were exposed to ∼20 Torr CH3OH.

TABLE 2: The Lifetime and Relative Amplitudes from the Two Exponential Fits of the Decay Curves sample

A1(%)

τ1 (s)

A2(%)

τ2 (s)

Pt/TiO2 (0.02 wt %) Pt/TiO2 (0.1 wt %) Pt/TiO2 (0.2 wt %) Pt/TiO2 (0.5 wt %) Pt/TiO2 (1.0 wt %)

51 56 63 54 65

0.38 0.081 0.13 0.31 0.42

49 44 37 46 35

2.77 0.59 1.11 2.51 7.82

pressure of methanol was fixed to 20 Torr. All the curves decay much faster than that observed for TiO2. So, reaction 3 is the dominant process for the electron consumption, and the recombination process (reaction 7) can be neglected. The shape of the decay curves is insensitive to the UV pulse energy in the range of 2-6 mJ, which suggests the first-order kinetics of electron decay. If we suppose that the long-lived electrons are composed by two first-order decaying species with the lifetime of τ1 and τ2, then the decay curves can be fitted well by the corresponding two-exponential formula y ) A1* exp(-x/τ1) + A2* exp(-x/τ2). Because not all the electrons are consumed in the interval of laser pulses (10 s) and a small amount of electrons might be accumulated in the catalysts, the curves are fitted by the adjusted formula

y ) A1* exp(-x/τ1) + A2* exp(-x/τ2) A1* exp(-10/τ1) - A2* exp(-10/τ2) (8) The lifetime and relative amplitudes from the nonlinear leastsquare fits are listed in Table 2. The lifetimes of both the fast component and the slow one vary with Pt loading and shows similar trends. Figure 13 shows the dependence of the simulated rate constants of the fast component on the Pt-loading amount. It can be seen that there exists an optimal amount of Pt loading for reaction 3 and the maximum rate was obtained with 0.1 wt % Pt-loaded TiO2. With further increase of the Pt-loading amount, the rate declines dramatically. Because the amount of Pt loaded on our catalysts is very low, Pt particles only block off a small part of the incident light. Furthermore, the effect of blocking off the incident light by loaded Pt may only decrease the initial absorption intensity of the decay curves but not change the shape of the curves with first-order kinetics. To explain the

8012 J. Phys. Chem. C, Vol. 111, No. 22, 2007

Figure 13. Dependence of the simulated rate constants of the fast component on Pt-loading amount.

above observations, the reaction mechanisms for anaerobic photocatalytic reaction of methanol are proposed and shown in Scheme 1. As shown in Scheme 1, the H2 evolution reaction mainly takes place on Pt sites and the hole-consuming reactions mainly take place on methanol-derived adsorbates. Methanol-derived species are oxidized by holes directly and release H+ stepwise. Molecularly-adsorbed water and methanol mediate the proton transfer. It is assumed that H2O(a) species receive the protons released from methanol (Hm+) to form H2HmO+(a), which diffuses on the surface and quickly exchanges the protons with other H2O(a) to form H3O+(a). When H3O+(a) arrives at Pt sites, the protons are liberated (Hw+) to take part in reaction 3. So, most of the protons that react with the long-lived electrons to generate H2 come from water but not methanol.2 CH3OH(a) can also mediate proton transfer through a similar mechanism, but the speed of the proton-transfer mediated by CH3OH(a) is slower than that mediated by H2O(a) under the same conditions. Increasing Pt loading amount increases the reactive sites for reaction 3, but both of the concentration of the electrons trapped by each Pt particle and the probability of the protons that can arrive at each Pt particle decrease, which may decelerate reaction 3. In addition, Pt particles or atoms occupy the active sites on the surface of TiO2 (Figure 1), the number of the surface active sites on TiO2 and the concentration of the released protons

Chen et al.

Figure 14. Dependence of hydrogen evolution rate upon Pt-loading amount. 0.3 g catalyst suspended in aqueous solution containing 160 mL H2O and 40 mL CH3OH was irradiated by 300 W Xe lamp covered with a water-cooled quartz jacket.

decrease drastically with increasing the Pt-loading amount. It is also a possible explanation for the results that the apparent rate constant for reaction 3 decreases with increasing the Ptloading amount. 3.4. Correlation Between Kinetics of the Long-Lived Electrons and the Activity of H2 Production under SteadyState UV Irradiation. In practice, the H2 is generated from a water-methanol solution, and the liquid-solid reaction systems are usually irradiated by continuous wave light. It is necessary to simulate or analyze the H2 evolution rate of the practical photoreaction by using the kinetic results obtained from timeresolved IR experiments. So, the activities of H2 production under steady-state UV irradiation were also investigated. From the estimation, the charge carrier density in the steady-state photocatalytic reactions is comparable or lower than that in timeresolved IR experiments. Figure 14 shows the effect of Pt-loading amount on the rate of the H2 evolution measured under steady-state irradiation conditions. It can be seen that Pt/TiO2 with Pt loading of 0.10.5 wt % show a high activity of H2 production, and the maximum rate was obtained for Pt/TiO2 with 0.2 wt % Pt loading. Pt/TiO2 (0.02 wt %) and Pt/TiO2 (1.0 wt %) show relatively poor activity of H2 production. Comparing the results with the simulated rate constant of reaction 3 shown in Figure

SCHEME 1: A Possible Reaction Mechanism for Photocatalytic Production of H2 from Methanol-Water Solution

Photocatalytic Reaction of Methanol on Pt/TiO2

J. Phys. Chem. C, Vol. 111, No. 22, 2007 8013

13, we can see that there are some relationships between the apparent rate constant of reaction 3 evaluated by time-resolved IR method and the rate of the H2 evolution estimated under the steady-state irradiation, but the photocatalytic activity does not completely accord with the apparent rate constant. The fact that the long-lived electrons accumulate in the photoreaction systems has been observed by many researchers.32,34,75,76 On the basis of steady state approximation, the rate of H2 generation from long-lived electrons under steady-state irradiation conditions can be deduced as

d[e-]/dt ) kPf - K1[e-] - K2[e-] ) 0

(9)

where [e-] represents the concentration of the long-lived electrons, k is a constant that correlates with the light utilization efficiency of the systems, P is the power of the irradiation light, f is the yield of the long-lived electrons, K1 is the apparent rate constant of the pseudo first-order electron-consuming reaction for H2 generation (reaction 3), K2 is the total apparent rate constant of the pseudo first-order recombination reaction (reaction 7) and other unwanted reactions. K1 and K2 can be evaluated based on the time-resolved IR measurement. From above equation, the equilibrium concentration of the long-lived electrons is obtained

[e-]eq ) kPf /(K1 + K2)

(10)

Then, the rate equation of H2 evolution under steady-state UV irradiation conditions is obtained

1 1 γH2 ) K1[e-]eq ) K1kPf /(K1 + K2) 2 2

(11)

For Pt-loaded TiO2 catalysts, K1 . K2. So, the rate of H2 generation can be represented as a simple equation

1 γH2 ) kPf 2

(12)

From the equation, the yield of the long-lived electrons could be responsible for the activity of H2 evolution. Because K1 is reduced in the equation, it seems that the rate of reaction 3 evaluated by time-resolved IR method does not affect the H2 evolution under steady-state irradiation conditions. But we believe K1 is an important factor that affects the activity of H2 evolution. If the K1 value is too small, from eq 10 it can be seen that the equilibrium concentration of the long-lived electrons will be very high. The accumulated electrons will occupy the Pt sites so as to decrease reaction 5. For the same reason, the counterpart holes that are not consumed will occupy CH3O-(a) sites and reaction 6 is declined. So, the yield of longlived electrons, f, will decrease. If electrons are consumed soon through reaction 3, the accumulated electrons will decrease, which is beneficial for promoting the activity of H2 evolution. Increasing the Pt-loading amount can improve reaction 5 and tolerate more accumulated electrons, but the active sites for methanol adsorption will decrease, which is not good for improving f. The contradiction between the Pt-loading amount and the number of the active sites for methanol adsorption is an important reason to explain why there exists an optimal Ptloading amount for H2 evolution. 4. Conclusions The capacity of methanol adsorption on Pt/TiO2 decreases with increasing Pt-loading amount, which indicates Pt particles

or atoms selectively occupy the active sites for methanol adsorption. Surface species, CH2O(a), CH2OO(a), and HCOO(a) are formed during the photocatalytic reaction of CH3OH on Pt/TiO2. Increasing the pressure of methanol or water can accelerate the photocatalytic reaction and increase the activity of H2 production. Pt particles on TiO2 trap the excited electrons, and the adsorbate species from methanol capture the holes, which are important for improving the yield of long-lived electrons. These long-lived electrons decay on the millisecond to second scale in Pt/TiO2, and the decays correspond to the electron consumption process for H2 generation. The function of the molecularly adsorbed methanol or water is found to mediate the proton transfer on the TiO2 surface. There exists an optimal amount of Pt loading for the decays of the long-lived electrons, and the maximum decay rate was obtained with 0.1 wt % Pt-loaded TiO2. The yield of the long-lived electrons could be responsible for the activity of H2 production under steady-state irradiation conditions. Acknowledgment. T.C. is grateful to Professor Taka-aki Ishibashi and Mr. Kentaro Tanabe for their helpful discussions on the time-resolved IR instrument. This work was financially supported by National Natural Science Foundation of China (NSFC, Grant No. 20373069), Knowledge Innovation Program of Chinese Academy of Sciences, DICP (Grant No. K2006E2), and National Basic Research Program of China (Grant No. 2003CB615806). References and Notes (1) Fujishima, A.; Honda, K. Nature 1972, 238, 37-38. (2) Sakata, T.; Kawai, T. Chem. Phys. Lett. 1981, 80, 341-344. (3) Kiwi, J.; Gratzel, M. J. Phys. Chem. 1984, 88, 1302-1307. (4) Baba, R.; Nakabayashi, S.; Fujishima, A.; Honda, K. J. Phys. Chem. 1985, 89, 1902-1905. (5) Bamwenda, G. R.; Tsubota, S.; Kobayashi, T.; Haruta, M. J. Photochem. Photobiol., A 1994, 77, 59-67. (6) Dickinson, A.; James, D.; Perkins, N.; Cassidy, T.; Bowker, M. J. Mol. Catal. A: Chem. 1999, 146, 211-221. (7) Kato, H.; Asakura, K.; Kudo, A. J. Am. Chem. Soc. 2003, 125, 3082-3089. (8) Liu, M.; You, W.; Lei, Z.; Zhou, G.; Yang, J.; Wu, G.; Ma, G.; Luan, G.; Takata, T.; Hara, M.; Domen, K.; Li, C. Chem. Commun. 2004, 2192-2193. (9) Lu, D.; Takata, T.; Saito, N.; Inoue, Y.; Domen, K. Nature 2006, 440, 295. (10) Linsebigler, A.; Lu, G.; Yates, J. T., Jr. Chem. ReV. 1995, 95, 735758. (11) Kaway, T.; Sakata, T. J. C. S. Chem. Comm. 1980, 694-695. (12) Bamwenda, G.; Tsubota, S.; Nakamura, T.; Haruta, M. J. Photochem. Photobiol., A 1995, 89, 177-189. (13) Ohtani, B.; Iwai, K.; Nishimoto, S.; Sato, S. J. Phys. Chem. B 1997, 101, 3349-3359. (14) Galinska, A.; Walendziewski, J. Energy Fuels 2005, 19, 11431147. (15) Ryczkowski, J. Catal. Today 2001, 68, 263-381. (16) Ekstrom, G. N.; McQuillan, A. J. J. Phys. Chem. B 1999, 103, 10562-10565. (17) Rusu, C. N.; Yates, J. T., Jr. J. Phys. Chem. B 2000, 104, 1229912305. (18) Liao, L. F.; Wu, W. C.; Chen, C. Y.; Lin, J. L. J. Phys. Chem. B 2001, 105, 7678-7685. (19) Coronado, J. M.; Kataoka, S.; Tejedor-Tejedor, I.; Anderson, M. A. J. Catal. 2003, 219, 219-230. (20) Kataoka, S.; Tejedor-Tejedor, M. I.; Coronado, J. M.; Anderson, M. A. J. Photochem. Photobiol., A 2004, 163, 323-329. (21) Kawai, M.; Naito, S.; Tamaru, K. Chem. Phys. Lett. 1983, 98, 377380. (22) Sata, S.; Ueda, K.; Kawasaki, Y.; Nakamura, R. J. Phys. Chem. B 2002, 106, 9054-9058. (23) Skinner, D. E.; Colombo, D. P., Jr.; Cavaleri, J. J.; Bowman, R. M. J. Phys. Chem. 1995, 99, 7853-7856. (24) Yang, X.; Tamai, N. Phys. Chem. Chem. Phys. 2001, 3, 33933398.

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