Enhanced Photocatalytic H2-Production Activity of TiO2 by Ni(OH)2

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Enhanced Photocatalytic H2-Production Activity of TiO2 by Ni(OH)2 Cluster Modification Jiaguo Yu,* Yang Hai, and Bei Cheng Sate Key Laboratory of Advanced Technology for Material Synthesis and Processing, Wuhan University of Technology, Wuhan 430070, P. R. China. ABSTRACT: Ni(OH)2 cluster-modified TiO2 (Ni(OH)2/TiO2) nanocomposite photocatalysts were fabricated by a simple precipitation method using Degussa P25 TiO2 powder (P25) as support and Ni(NO3)2 as precursor. The effect of Ni(OH)2 cluster loading content on the photocatalytic hydrogen production rates of the as-prepared samples in methanol aqueous solution was investigated. The results showed that the photocatalytic H2-production activity of TiO2 was significantly enhanced by loading Ni(OH)2 clusters. The optimal Ni(OH)2 loading content was found to be 0.23 mol %, giving a H2-production rate of 3056 μmol h-1 g-1 with quantum efficiency (QE) of 12.4%, exceeding that on pure TiO2 by more than 223 times. This high photocatalytic H2production activity is due to the deposition of Ni(OH)2 clusters on the surface of TiO2. The enhanced mechanism is because the potential of Ni2þ/Ni (Ni2þ þ 2e- = Ni, Eo = 0.23 V) is slightly lower than conduction band (CB) (-0.26 V) of anatase TiO2, meanwhile higher than the reduction potential of Hþ/H2 (2Hþ þ 2e- = H2, Eo = -0.00 V), which favors the electron transfer from CB of TiO2 to Ni(OH)2 and the reduction of partial Ni2þ to Ni0. The function of Ni0 is to help the charge separation and to act as cocatalyst for water reduction, thus enhancing the photocatalytic H2-production activity.

1. INTRODUCTION Hydrogen energy is a storable, clean, and environmentally friendly fuel for the future. Photocatalytic water splitting into hydrogen and oxygen using semiconductor photocatalysts has been considered as a promising and attractive approach to produce hydrogen energy since Honda and Fujishima discovered the photocatalytic splitting of water on TiO2 electrodes in 1972.1-6 For effective H2 production, the CB edge of semiconductor should be more negative than the H2 production potential and the valence band (VB) edge should be more positive than the oxygen oxidation potential. Among various semiconductor photocatalysts, titania has proven to be the most suitable for photocatalytic water splitting for its biological and chemical inertness, strong oxidizing and reducing power, cost effectiveness, and long-term stability against photocorrosion and chemical corrosion.7-10 Unfortunately, the photocatalytic H2-production efficiency on a bare TiO2 is very low, mainly due to the following reasons: (1) rapid recombination of photogenerated electrons and holes, (2) fast backward reaction between hydrogen and oxygen, and (3) large H2 production overpotential.11,12 To overcome these shortcomings and enhance photocatalytic H2 production activity, more efforts have been made, including deposition of noble metals,13-16 addition of sacrificial reagents,15-17 dye sensitization,11 metal cation doping,18-20 carbon and nitrogen doping,21,22 etc. Both the deposition of noble metals and the addition of sacrificial reagents have been intensively investigated for photocatalytic H2 production and have been proven to be very effective methods.12 It is well-known that the deposition of r 2011 American Chemical Society

Pt on TiO2 surface significantly increases the H2-production efficiency when photocatalytic water splitting is carried out in the presence of sacrificial reagents.7,13 However, Pt is a rare and expensive noble metal. Therefore, numerous efforts have been undertaken to replace Pt with low-cost additives.11,23 The previous work has shown that NiO-modified TiO2 is a cost-effective and efficient photocatalyst for water splitting to produce H2. For example, Sreethawong et al.24 reported synthesis of mesoporous TiO2-supported NiO photocatalyst by singlestep sol-gel process for photocatalytic H2-production from methanol aqueous solution. Jing et al.25 investigated fabrication of Ni-doped mesoporous TiO2 and its photocatalytic activity for hydrogen evolution in methanol aqueous solution. In 2009, Jang et al.26 reported the enhanced photocatalytic H2-production efficiency of nickel-intercalated titanate nanotube from methanol aqueous solution. However, to our knowledge, there are few reports on the photocatalytic H2 production over Ni(OH)2 cluster-modified TiO2 up to now. Herein, we for the first time report a simple precipitation method for the fabrication of Ni(OH)2/TiO2 nanocomposite photocatalysts at room temperature and their enhanced photocatalytic H2-production activity. Also a new mechanism for photocatalytic H2 production in Ni(OH)2/TiO2 is proposed and totally different from the previous reported mechanism.

Received: December 5, 2010 Revised: January 20, 2011 Published: February 28, 2011 4953

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Table 1. Effects of R on Physicochemical Properties and QE of Ni(OH)2/TiO2 Samples

a

samples

R

Ni(OH)2 (mol %) (ICP-AES)

ACS (nm) a

SBET (m2/g)

PV (cm3/g)

porosity (%)

QE (%)

N0

0.06

0

0

24.0 (A), 46.2 (R)

44.6

0.17

38.6

N0.1

0.1

0.04

24.0 (A), 46.2 (R)

44.8

0.50

64.9

3.0

N0.5

0.5

0.23

24.0 (A), 46.0 (R)

44.8

0.50

64.9

12.4

N1.0

1.0

0.38

23.9 (A), 46.0 (R)

45.2

0.50

64.9

9.3

N1.6

1.6

0.60

23.9 (A), 45.6 (R)

45.6

0.49

64.5

7.3

7.45

23.8 (A), 45.5 (R)

45.8

0.47

63.5

0.9

0.42

63.6

0

N10

10

N100

100

100

4.0 (Ni(OH)2)

188

A and R denote anatase and rutile, respectively. ACS: average crystallite size. PV: pore volume.

2. EXPERIMENTAL SECTION 2.1. Sample Preparation. All the reagents were of analytical grade and were used without further purification. Distilled water was used in all experiments. Commercially available P25 was used as the source of TiO2. Ni(OH)2/TiO2 photocatalysts were prepared by the conventional precipitation method. In a typical synthesis, 1.0 g of P25 was dispersed in 50 mL of 1.0 M NaOH aqueous solution, and then a certain volume of 0.05 M Ni(NO3)2 aqueous solution was added. The molar ratios of Ni(OH)2 to (TiO2 þ Ni(OH)2), which hereafter were designated as R, were 0, 0.1, 0.5, 1.0, 1.6, and 10 nominal molar % (mol %) (see Table 1); the obtained samples were labeled as samples N0, N0.1, N0.5, N1.0, N1.6, and N10, respectively. The mixed solutions were stirred for 24 h at room temperature. After that, the precipitates were collected by centrifuge and washed with distilled water and alcohol 10 times, respectively. The washed precipitates were dried at 80 °C for 24 h. The actual chemical compositions of the prepared samples were measured by inductively coupled plasma atomic emission spectrometry (ICP-AES) using an Optima 4300 DV spectrometer (Perkin-Elmer) (see Table 1). Pure Ni(OH)2 sample was also prepared for the purpose of comparison under the same experimental conditions, and the resulting Ni(OH)2 sample was labeled as N100. 2.2. Characterization. X-ray diffraction (XRD) patterns, obtained on an X-ray diffractometer (Rigaku, Japan) using Cu KR irradiation at a scan rate of 0.05° 2θ s-1, were used to determine the phase structures of the obtained samples. The average crystallite size was calculated using the Scherrer formula after correcting the instrumental broadening. Transmission electron microscopy (TEM) observations were conducted by a F20 S-TWIN electron microscope (Tecnai G2, FEI Company), using a 200 kV accelerating voltage. The Brunauer-EmmettTeller (BET) specific surface area (SBET) of the powders was analyzed by nitrogen adsorption in a Micromeritics ASAP 2020 nitrogen adsorption apparatus (U.S.A.). All the as-prepared samples were degassed at 180 °C prior to nitrogen adsorption measurements. The BET surface area was determined by a multipoint BET method using the adsorption data in the relative pressure (P/P0) range of 0.05-0.3. A desorption isotherm was used to determine the pore size distribution via the BarretJoyner-Halender (BJH) method, assuming a cylindrical pore model. The nitrogen adsorption volume at the relative pressure (P/P0) of 0.994 was used to determine the pore volume and average pore size. UV-vis diffused reflectance spectra of the samples were obtained for the dry-pressed disk samples by a UV-vis spectrotometer (UV2550, Shimadzu, Japan). BaSO4 was used as a reflectance standard in a UV-vis diffuse reflectance experiment. The X-ray photoelectron spectroscopy (XPS)

measurement was done in an ultrahigh-vacuum VG ESCALAB 210 electron spectrometer equipped with a multichannel detector. The spectra were excited using Mg KR (1253.6 eV) radiation (operated at 200 W) of a twin anode in the constant analyzer energy mode with a pass energy of 30 eV. Photoluminescence (PL) spectra were measured at room temperature on an F-7000 fluorescence spectrophotometer (Hitachi, Japan). 2.3. Photocatalytic H2 Production. The photocatalytic H2production experiments were performed in a 100 mL Pyrex flask at ambient temperature and atmospheric pressure, and openings of the flask were sealed with a silicone rubber septum. Four lowpower UV-LEDs (3 W, 365 nm) (Shenzhen LAMPLIC Science Co. Ltd., China), which were positioned 1 cm away from the reactor in four different directions, were used as light sources to trigger the photocatalytic reaction. The focused intensity and areas on the flask for each UV-LED were ca. 80.0 mW/cm2 and 1 cm2, respectively. In a typical photocatalytic experiment, 50 mg of Ni(OH)2/TiO2 photocatalyst was suspended in 80 mL of a mixed solution of methanol (20 mL) and water (60 mL). Prior to irradiation, the suspension of the catalyst was dispersed in an ultrasonic bath for 5 min and then bubbled with nitrogen through the reactor for 30 min to completely remove the dissolved oxygen and ensure that the reactor was in an anaerobic condition. A continuous magnetic stirrer was applied at the bottom of the reactor in order to keep the photocatalyst particles in suspension status during the whole experiment. A 0.4 mL gas was intermittently sampled through the septum, and hydrogen was analyzed by gas chromatograph (GC-14C, Shimadzu, Japan, TCD, nitrogen as a carrier gas and 5 Å molecular sieve column). All glassware was carefully rinsed with distilled water prior to use. The QE was measured and calculated according to eq 1: QE ð%Þ ¼ ¼

number of reacted electrons  100 number of incident photons number of evolved H2 molecules  2  100 number of incident photons

ð1Þ

3. RESULTS AND DISCUSSION 3.1. Phase Structures and Morphology. XRD was used to investigate the changes of phase structures and crystallite size of the as-prepared samples. Figure 1 presents the comparison of XRD patterns of N0, N0.1, N0.5, N1.0, N1.6, N10, and N100. Only anatase and rutile phases of TiO2 are observed for pure TiO2 (N0) and Ni(OH)2/TiO2 composite photocatalysts 4954

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Figure 1. XRD patterns of samples N0, N0.1, N0.5, N1.0, N1.6, N10, and N100.

Figure 2. High-magnification TEM image of N0.5.

(N0.1, N0.5, N1.0 N1.6, and N10). No characteristic diffraction peaks of Ni(OH)2 are observed because of its lower loading content and weak crystallization, on the other hand, also implying the good dispersion of the very small Ni(OH)2 clusters on the TiO2 surface. Further observation from Figure 1 shows that there is no obvious change observed in the diffraction peak position of anatase and rutile, suggesting that the deposited Ni(OH)2 clusters do not incorporate into the lattice of TiO2. For N100, five main diffraction peaks near at 2θ = 16.7, 23.8, 34.1, 35.9, and 53.2° can be observed, respectively corresponding to (001), (100), (101), (110), and (111) plane diffraction of Ni(OH)2 [JCPDS No. 14-117, hexagonal, space group P3mL (164)], indicating the presence of Ni(OH)2. The average crystallite sizes of anatase, rutile, and Ni(OH)2, calculated using Scherrer’s equation, respectively, from the main diffraction peak of anatase (101), rutile (110), and Ni(OH)2 (001), are listed in Table 1. Pure TiO2 and Ni(OH)2/TiO2 samples have almost the same crystallite size (ca. 24 nm for anatase, 46 nm for rutile), indicating that the deposition of Ni(OH)2 clusters on the surface of TiO2 has no obvious influence on its crystallite size and morphology. This is easy to be understood because ambient temperature deposition of Ni(OH)2 does not have enough energy to stimulate the growth of

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Figure 3. UV-vis diffuse reflection spectra of samples N0, N0.1, N0.5, N1.0, N1.6, N10, and N100. Inset shows the UV-vis absorption spectrum of 0.01 M Ni(NO3)2 aqueous solution.

TiO2 crystal.27 TEM was further used to observe the morphology and microstructures of as-prepared samples. Figure 2 presents a typical TEM image of Ni(OH)2/TiO2 sample (N0.5), clearly indicating many small Ni(OH)2 clusters with size of ca. 1-3 nm uniformly deposited on the surface of TiO2 nanoparticles. The chemical composition of N0.5 was determined by energy-dispersive X-ray spectroscopy (EDS) (not shown here), indicating that the peaks of Ni, O, and Ti elements are observed. 3.2. UV-Vis Diffuse Reflection Spectra. Figure 3 presents the comparison of UV-vis diffuse reflectance spectra of pure TiO2, Ni(OH)2/TiO2, and Ni(OH)2. It can be seen that the absorption intensity of TiO2 starts to increase rapidly at ∼410 nm, which corresponds to the intrinsic band gap absorption of rutile TiO2, indicating a band gap of 3.0 eV. This coincides with the reported value of rutile in the literature.28 As for Ni(OH)2/TiO2, the spectra show absorption in the ∼600800 nm region and an small absorption shoulder at ∼450 nm, in addition to the onset of absorption at ∼410 nm. The 600800 nm absorption can be assigned to the Ni(II) d-d transition.28 In fact, the absorption spectra of the Ni(NO3)2 aqueous solution (inset of Figure 3) and C100 solid powder sample also show absorption centered at ∼650-750 nm. According to Irie et al.’s viewpoint,28 the absorption at ∼450 nm of Ni(OH)2/TiO2 can be ascribed to the direct interfacial charge transfer (IFCT) from the VB of TiO2 to Ni(II). Therefore, it is not surprising that the absorption at ∼450 and 600-800 nm increases with increasing amounts of the deposited Ni(OH)2. In comparison to pure TiO2 (N0), no great change in the absorption edge of the Ni(OH)2/TiO2 samples was observed, also implying that Ni(OH)2 was not incorporated into the lattice of TiO2, only deposited on its surface. 3.3. BET Surface Areas and Pore Size Distributions. Generally, a catalyst with high specific surface area and big pore volume is indispensable to the enhancement of catalytic performance.29 Therefore, the effect of Ni(OH)2 loading on the pore structure and BET surface areas of as-prepared samples was investigated by the adsorption-desorption measurement. Figure 4 shows the nitrogen adsorption/desorption isotherms and the corresponding pore-size distribution curves (inset) of N0 and N0.5. N0 and N0.5 samples have isotherms of type IV from Brunauer-Deming-Deming-Teller (BDDT) classification, indicating the presence of mesopores (2-50 nm).30,31 The shapes of hysteresis loops are of type H3 at a high relative pressure range of 0.8-1.0, indicating the presence of slitlike 4955

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Figure 4. Nitrogen adsorption-desorption isotherms and the corresponding pore-size distribution curves (inset) of samples N0 and N0.5.

Figure 5. High-resolution XPS spectra of Ni 2p of the samples (a) N0.5, (b) N0.5 heated at 400 °C for 2 h and (c) N0.5 after 2 h photocatalytic hydrogen production from methanol aqueous solution under UV-LED irradiation.

pores. The isotherms show high absorption at high relative pressure (P/P0) range (approaching 1.0), implying the formation of large mesopores.30,31 Further observation shows that at high P/P0 range (approaching 1.0), isotherms of N0.5 shift up comparing with N0, suggesting N0.5 with bigger pore volume. The pore size distribution curves (see inset of Figure 4), calculated from the desorption branch of the nitrogen isotherms by the BJH method, show a wide range of 10-60 nm with a peak pore diameter of about 30 nm for N0 and N0.5 samples, confirming the presence of mesopores.30,31 Table 1 shows quantitative details on BET surface area, pore volume, and porosity of TiO2 and Ni(OH)2/TiO2 samples. The Ni(OH)2/ TiO2 samples show an increase in pore volume and porosity compared with pure TiO2, which can be ascribed to the formation of more aggregates of TiO2 crystallites. Such aggregated porous structures and increased porosity might be extremely useful in the photocatalytic process since they might provide efficient transport pathways to reactant and product molecules.32 3.4. XPS Analysis. In order to analyze chemical composition of the prepared Ni(OH)2/TiO2 photocatalysts and to identify the chemical status of Ni element in the samples, the samples were characterized by XPS. The XPS survey spectrum (not shown here) of N0.5 indicates that Ti, Ni, O, and C elements are observed and their corresponding photoelectron peaks respectively appear at binding energies of 458.8 (Ti 2p), 856

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Figure 6. Comparison of PL spectra of samples N0 and N0.5.

(Ni 2p), 531 (O 1s), and 285 eV (C1s). The atomic ratio of Ti to O is about 1:2, which is in good agreement with the nominal atomic composition of TiO2. The carbon peak is attributed to the residual carbon from the sample and adventitious hydrocarbon from the XPS instrument itself. The nickle peak is from the presence of Ni(OH)2. Figure 5 presents the comparison of highresolution XPS spectra of Ni 2p of the samples N0.5, N0.5 heated at 400 °C and N0.5 after photocatalytic reaction for 2 h. For the as-prepared N0.5, the measured binding energies of Ni 2p3/2 and Ni 2p1/2 are equal to 855.7 and 873.2 eV, respectively. The Ni 2p3/2 peak at 855.7 eV indicates that the deposited nickel compound on the surface of TiO2 nanoparticles consists mainly of Ni(OH)2, which reportedly exhibits a peak at about 856 eV.33-35 When N0.5 was heat-treated at 400 °C for 2 h, its XPS spectrum exhibits a significant change (see Figure 5b). The binding energies of Ni 2p3/2 and Ni 2p1/2 shift to 854.6 and 872.5 eV, respectively, indicating the formation of NiO due to the dehydration of Ni(OH)2.34,35 The above XPS result can convincingly confirm Ni(OH)2 existing on the surface of TiO2 (for the as-prepared N0.5). This is also in agreement with the above XRD and UV-vis diffuse reflectance spectra results. However, after photocatalytic reaction for 2 h, the XPS spectrum of sample N0.5 also exhibits a significant change. The binding energies of Ni 2p3/2 further shift to 854.1 eV (see Figure 5c), indicating the formation of NiO. One possible explanation is that the photoindued electrons in the CB of TiO2 transfer to Ni(OH)2 clusters and reduce partial Ni2þ to Ni0 atoms, finally forming Ni clusters. These Ni clusters are unstable in air and are easily oxidated. Therefore, it is not surprising that NiO is observed for sample N0.5 after photocatalytic reaction. Further observation from Figure 5c indicates that the binding energies of Ni 2p1/2 remain unchanged (at 873.2 eV), also suggesting that the N0.5 sample after photocatalytic reaction for 2 h still contains Ni(OH)2. 3.5. PL Spectra. Figure 6 presents a comparison of PL spectra of N0 and N0.5 in the wavelength range of 350-550 nm. The shape of two PL spectra is similar. A fluorescence decrease (or quenching) is observed for Ni(OH)2/TiO2. Six main emission peaks appear at about 398, 410, 451, 468, 483, and 492 nm, being equivalent to 3.12, 3.03, 2.75, 2.65, 2.57, and 2.52 eV, respectively. Two of the major emission peaks at about 398 and 410 nm are respectively ascribed to the band-band PL phenomenon with the energy of light approximately equal to the band gap energy of anatase and rutile. The other four PL peaks at 451, 468, 483, and 492 nm are attributed to excitonic PL.36 The PL intensity of Ni(OH)2/TiO2 exhibits a decrease compared with pure TiO2, implying that Ni(OH)2/TiO2 has a lower 4956

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Figure 7. Comparison of the photocatalytic activity of samples N0, N0.1, N0.5, N1.0, N1.6, N10, and N100 for the photocatalytic H2 production from methanol aqueous solution under UV-LEDs irradiation.

recombination rate of electrons and holes under UV light irradiation. This is due to the fact that the electrons are excited from the valence band to the conduction band of TiO2 and then migrate to Ni(OH)2 clusters, which prevent the direct recombination of electrons and holes.36 3.6. Photocatalytic H2-Production Activity. Photocatalytic H2-production activity on various samples was evaluated under UV-LED irradiation using methanol as a scavenger. Figure 7 and Table 1 present the comparison of the photocatalytic H2production activity of TiO2, Ni(OH)2/TiO2, and Ni(OH)2 samples. Control experiments indicate that no appreciable H2 production is detected in the absence of either irradiation or photocatalyst. R exhibits a significant influence on the photocatalytic activity. For pure TiO2 (R = 0), it shows a very low photocatalytic activity because of the rapid recombination between CB electrons and VB holes, the fast backward reaction (recombination of hydrogen and oxygen into water), and the presence of a large H2-production overpotential on TiO2 surface.7 All of these factors make TiO2 alone inactive in photocatalytic water splitting and H2 production. After loading only 0.1 mol % of Ni(OH)2 on TiO2, the H2-production activity of N0.1 is markedly enhanced and increased by up to 53 times. With further increasing R from 0.1 to 0.5, the rate of H2 production on Ni(OH)2/TiO2 is increased and achieves a maximum at R = 0.5 (N0.5). The highest H2-production rate is 3056 μmol h-1 g-1 with 12.4% QE. This value exceeds that of pure TiO2 (N0) by a factor of 223. When R is higher than 0.5, a further increase in R leads to a reduction of the photocatalytic activity. Especially, at R = 10, the photocatalytic activity of N10 has a drastic decrease. This is probably due to the following causes: (i) deposition of excessive Ni(OH)2 clusters gives rise to the decrease (or shield) of the TiO2 surface active sites; (ii) deterioration of the catalytic properties of Ni(OH)2 cluster or disappearance of surface effect due to the increase of their particle size;6,37 (iii) increase in the opacity leading to a decrease of irradiation passing through the reaction suspension solution.38 No hydrogen can be detected when Ni(OH)2 alone was used as the catalyst, suggesting that pure Ni(OH)2 was not active for photocatalytic H2 production. From what has been observed and discussed above, we can obtain several important conclusions: (1) The TiO2 without Ni(OH)2 cluster modification is inactive for photocatalytic H2 production under UV light irradiation, although the conduction

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Figure 8. Schematic illustration for the charge transfer and separation in the Ni(OH)2 cluster-modified TiO2 system under UV-LEDs irradiation.

band level of anatase TiO2 is more negative than the reduction potential of Hþ/H2. (2) After Ni(OH)2 cluster modification, the photocatalytic activity of samples is greatly enhanced. (3) The content of Ni(OH)2 cluster significantly influences H2-production activity of TiO2. Based on the above results, the photocatalytic mechanisms are proposed and shown in Figure 8. Although the conduction band edge of anatase is higher (or more negative) than the reduction potential of Hþ/H2, the rate of H2 production was negligible over bare TiO2 in the absence of Ni(OH)2 cluster due to the rapid recombination rate of CB electrons and VB holes and the presence of a large H2-production overpotential. Once TiO2 was modified by Ni(OH)2, under UV light irradiation, the VB electrons of TiO2 are excited to CB. Because the potential (-0.23 V vs SHE, pH = 0) of Ni2þ/Ni is slightly lower than the CB level (about -0.26 V) of anatase TiO2, the photoindued electrons in the CB can transfer to Ni(OH)2 clusters and then effectively reduce partial Ni2þ to Ni0 atoms, finally forming Ni clusters.39 These Ni clusters can act as cocatalyst to promote the separation and transfer of photogenerated electrons from TiO2 CB to the Ni(OH)2/Ni cluster, where Hþ is reduced to hydrogen molecules. The major reaction steps in this photocatalytic mechanism under UV light irradiation are summarized by eqs 2 and 3. Ni2þ þ 2e- ¼ Ni

ð2Þ

Ni þ 2e- þ 2Hþ ¼ Ni þ H2

ð3Þ

Therefore, it is not surprising that high photocatalytic H2production activity is achieved over TiO2 samples as a result of surface Ni(OH)2 modification. When R is lower than 0.5, with increasing Ni(OH)2 content, more Ni(OH)2 clusters are deposited on the TiO2 surface, thus resulting in the increase of the activity. On the contrary, when R is higher than 0.5, with further increasing Ni(OH)2 content, the particle size of Ni(OH)2 clusters increases and the surface effect becomes weak or disappears, causing the reduction of the photocatalytic activity. The above fluorescence quenching experiments further confirm the transfer of photogenerated electrons from TiO2 to Ni(OH)2 clusters.

4. CONCLUSIONS In conclusion, a simple precipitation method is developed for the synthesis of highly active Ni(OH)2 cluster-modified TiO2 nanocomposite photocatalysts for photocatalytic H2 production. 4957

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The Journal of Physical Chemistry C Ni(OH)2 clusters can act as effective cocatalysts to enhance the photocatalytic H2-production activity of TiO2, and its content exhibits a significant influence on the activity. The optimal Ni(OH)2 loading content is determined to be about 0.23 mol %, and the corresponding H2-production rate is 3056 μmol h-1 g-1 with QE of 12.4%, which exceeds that of pure TiO2 by more than 223 times. The potential of Ni2þ/Ni is deemed to be lower than the conduction band of TiO2 and more negative than the Hþ/H2 potential, which favors the electron transfer from CB of TiO2 to Ni(OH)2 and the reduction of Hþ, thus enhancing photocatalytic H2-production activity. This work not only shows a possibility for the utilization of low-cost Ni(OH)2 cluster as a substitute for noble metals in photocatalytic hydrogen production but also exhibits a facile method for fabricating highly active H2-production photocatalysts by a simple precipitation reaction.

’ AUTHOR INFORMATION Corresponding Author

*Tel. 0086-27-87871029; fax 0086-27-87879468; e-mail jiaguoyu@ yahoo.com.

’ ACKNOWLEDGMENT This work was partially supported by the National Natural Science Foundation of China (50625208, 20773097, 20877061, and 51072154) and the Natural Science Foundation of Hubei Province (2010CDA078). This work was also financially supported by the National Basic Research Program of China (2007CB613302). ’ REFERENCES (1) Fujishima, A.; Honda, K. Nature 1972, 238, 37. (2) Park, J. H.; Kim, S.; Bard, A. J. Nano Lett. 2006, 6, 24. (3) Li, Y. X.; Xie, Y. Z.; Peng, S. Q.; Lu, G. X.; Li, S. B. Chemosphere 2006, 63, 1312. (4) Kudo, A.; Miseki, Y. Chem. Soc. Rev. 2009, 38, 253. (5) Bard, A. J.; Whitesides, G. M.; Zare, R. N.; McLafferty, F. W. Acc. Chem. Res. 1995, 28, 91. (6) Yu, J. G.; Zhang, J.; Jaroniec, M. Green Chem. 2010, 12, 1611. (7) Yu, J. G.; Qi, L. F.; Jaroniec, M. J. Phys. Chem. C 2010, 114, 13118. (8) Hoffmann, M. R.; Martin, S. T.; Choi, W.; Bahnemann, D. W. Chem. Rev. 2002, 95, 69. (9) (a) Ksibi, M.; Rossignol, S.; Tatibouet, J. M.; Trapalis, C. Mater. Lett. 2008, 62, 4204. (b) Yu, J. G..; Yu, J. C.; Leung, M. K. P.; Ho, W. K.; Cheng, B.; Zhao, X. J.; Zhao, J. C. J. Catal. 2003, 217, 69. (c) Yu, J. G.; Su, Y. R.; Cheng, B. Adv. Funct. Mater. 2007, 17, 1984. (d) Yu, J. G.; Wang, W. G.; Cheng, B. Chem. Asian J. 2010, 5, 2499. (10) Ni, M.; Leung, M. K. H.; Leung, D. Y. C.; Sumathy, K. Renew. Sustain. Energy Rev. 2007, 11, 401. (11) Jin, Z. L.; Zhang, X. J.; Li, Y. X.; Li, S. B.; Lu, G. X. Catal. Commun. 2007, 8, 1267. (12) Xu, S. P.; Sun, D. D. Int. J. Hydrogen Energy 2009, 34, 6096. (13) (a) Bamwenda, G. R.; Tsubota, S.; Nakamura, T.; Haruta, M. J. Photochem. Photobiol., A 1995, 89, 177. (b) Sakata, T.; Kawai, T. Chem. Phys. Lett. 1981, 80, 341. (c) Sakthivel, S.; Shankar, M. V.; Palanichamy, M.; Arabindoo, B.; Bahnemann, D. W.; Murugesan, V. Water Res. 2004, 38, 3001. (14) Sreethawong, T.; Yoshikawa, S. Int. J. Hydrogen Energy 2006, 31, 786. (15) Sayama, K.; Arakawa, H. J. Chem. Soc., Faraday Trans. 1997, 93, 1647. (16) Galinska, A.; Walendziewski, J. Energy Fuels 2005, 19, 1143.

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