Synthesis and photoactivity of nanostructured CdS–TiO2 composite catalysts

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Synthesis and photoactivity of nanostructured CdS–TiO2 composite catalysts Xin Li a,b , Ting Xia a , Changhui Xu a , James Murowchick c , Xiaobo Chen a,∗ a b c

Department of Chemistry, University of Missouri – Kansas City, Kansas City, MO 64110, USA Institute of Biomaterial, College of Science, South China Agricultural University, Guangzhou 510642, China Department of Geosciences, University of Missouri – Kansas City, Kansas City, MO 64110, USA

a r t i c l e

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Article history: Received 31 May 2013 Received in revised form 10 October 2013 Accepted 11 October 2013 Available online xxx Keywords: Titanium dioxide CdS Amorphous Composite Photocatalyst

a b s t r a c t In this paper, a series of nanostructured TiO2 –CdS composite photocatalysts have been synthesized using an aqueous solution route at room temperature, and characterized with X-ray diffraction (XRD), Raman spectroscopy, transmission electron microscopy (TEM), UV–VIS absorption and Fourier transform infrared spectroscopy (FTIR). Their photocatalytic activities have been investigated on the degradation of methylene blue under simulated solar light irradiation. CdS shows the highest activity when the ratio of the starting materials Na2 S to Cd(NO3 )2 is close to 1.0, TiO2 shows the best activity when prepared under pH = 13.0, and the CdS/TiO2 composite with the molar ratio of TiO2 to CdS of 50 displays the highest activity. © 2013 Elsevier B.V. All rights reserved.

1. Introduction The depletion of fuel energy and the deterioration of natural environment are two most urgent challenges facing modern society [1]. Both problems could be eased by using photocatalysts to generate hydrogen from water [2–4], to photodegrade toxic pollutants [5–7] and to convert CO2 into useable fuels [8–10]. Among all the semiconductor photocatalysts, TiO2 and CdS have attracted great attention and been studied extensively in the past decade years. TiO2 is generally accepted as the most widely used photocatalyst because of its high activity, good stability, low cost and environmental friendliness [11]. However, its use as photocatalytic materials is limited [12], because it only absorbs about 5% of sunlight in the ultraviolet region due to its wide bandgap (3.0–3.2 eV). Therefore, it is important to improve its optical properties in the visible light range. The efforts include reducing its bandgaps by hydrogenation [13,14], metal doping [2,15] and nonmetal doping [16–19], and sensitizing with a low bandgap semiconductor material [2,11,20] or dye molecule [21,22]. CdS has a high activity and quantum efficiency in the visible light region due to its smaller band-gap energy (Eg ≈ 2.30 eV) [23], and proper positions of the conduction band minimum (CB: −0.75 V vs NHE) and valence band maximum (VB: 1.75 V vs NHE) [24–26]. However, because

∗ Corresponding author. Tel.: +1 816 235 6420; fax: +1 816 235 2290. E-mail address: [email protected] (X. Chen).

of photocorrosion, CdS is not stable in aqueous media during the photocatalytic reactions where CdS is itself oxidized by the photogenerated holes [27–29]. Therefore, it is important to improve the stability and suppress the photocorrosion of CdS nanoparticles, such as by the addition of sacrificial reagents of S2− and SO3 2− in aqueous solutions, incorporation of CdS into these layered oxides [30,31], growth of CdS in varieties of microporous and mesoporous materials [32–35], coupling with other wide bandgap semiconductors with suitable bandgap structure [36–38] and formation of core/shell structures with protective layers on the surface of CdS nanoparticles [27,39], and so on. The CdS/TiO2 composite configuration has attracted great attention in various solar energy applications because of their matched band structures and complementary optical and photocatalytic properties [40–47]. Various methods have been developed to prepare CdS/TiO2 composites, such as chemical vapor deposition [48], chemical bath deposition [47,49,50], photodeposition [51,52], electrodeposition [53], successive ionic layer adsorption and reaction method [54], co-precipitation [55,56], sonochemical method [57,58], and hydrothermal [59,60]. Most of these methods require long reaction times (at least 20–24 h), high temperature calcinations (200–400 ◦ C), and the presence of different surfactants [57,61]. The calcinations can lead to the increase of the crystallite size and the decrease of the specific surface area, and thus hampers the enhancement of photocatalytic activity. Furthermore, CdS in CdS/TiO2 composite can be oxidized during thermal treatment [61]. The presence of different surfactants is also harmful to their

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photocatalytic activities due to the possible block of the active sites on the surface and the limit of efficient charge transfer across the interface [55]. Therefore, from the application point of view, surfactant-free, room-temperature preparation of CdS/TiO2 composites with highly exposed surfaces is attractive for their photocatalytic applications. Here, we report a facile approach to fabricate CdS/TiO2 composite photocatalysts. The effects of the synthetic parameters on the crystal structure, optical properties, photocatalytic activities of CdS/TiO2 composites and their relationships are investigated in details.

(FTIR) spectra are collected using a Thermo-Nicolet iS10 FT-IR spectrometer with an attenuated total reflectance (ATR) unit. The photocatalytic activities of the samples are determined by measuring the photocatalytic decomposition process of methylene blue (1.46 × 10−5 M) under simulated solar light irradiation. The solar simulator (81094, Newport) has a 150 W Xe lamp with an AM 1.5 air mass filter. 1.0 mg of catalyst is added into 3.0 ml methylene blue solution (optical density of 1.0). The UV–vis absorption spectrum of methylene blue is monitored over time after the photocatalytic reaction starts. The UV–vis spectrum of methylene blue is measured with an Agilent Cary 60 UV-Vis spectrometer with a spectrum range from 400 nm to 800 nm.

2. Experimental 2.1. Synthesis A series of nanostructured CdS photocatalysts are synthesized by reacting Na2 S to Cd(NO3 )2 under basic conditions [28]. The molar ratio of starting materials (Na2 S:Cd(NO3 )2 ) is changed from 1:25 to 2:1. In a typical synthesis (molar ratio of Na2 S:Cd(NO3 )2 = 1.0), 10 mL of 0.1 M Cd(NO3 )2 aqueous solution is added into 100 mL of 0.1 M NaOH solution at room temperature, then 10 mL of 0.1 M Na2 S solution is quickly poured in under stirring. The obtained CdS precipitate is then filtered, washed with D.I. water, and dried at 70 ◦ C overnight. Second, a series of nanostructured TiO2 photocatalysts are prepared in water under pH values from 9.0 to 13.0. The pH of the solution is controlled with 1.0 M NaOH solution. In a typical synthesis (pH = 12.0), 30 ␮L titanium(IV) isopropoxide is added dropwise into 120 mL of 0.010 M NaOH solution under stirring at room temperature and stirred for 1 h. The obtained white TiO2 precipitate is then filtered and dried at 100 ◦ C overnight. Finally, the molar ratio of CdS to TiO2 is varied to prepare a series of nanostructured CdS/TiO2 composites. When preparing the series of CdS–TiO2 samples, the molar ratio of Na2 S·9H2 O to Cd(NO3 )2 ·4H2 O is set to constantly 1:1. In a typical experiment, 30 ␮L titanium(IV) isopropoxide is added dropwise into the solution of CdS under stirring for 1 h at room temperature. The obtained precipitate is filtered and collected after dried in an oven at 70 ◦ C. 2.2. Characterization The properties of the samples are investigated with powder Xray diffraction (PXRD), Raman spectroscopy, Fourier-transformed infrared spectroscopy (FTIR), UV–vis diffusive reflectance and transmission electron microscopy (TEM). The PXRD is performed using a Rigaku Miniflex PXRD machine with Cu K␣ as the X-ray ˚ and the 2-theta range is from 15◦ sources (wavelength = 1.5418 A) to 85◦ with a step width of 0.08 and count time of 3 s/step. The Raman spectra are collected on an EZRaman-N benchtop Raman spectrometer (Enwave Optronics, Inc.). The Raman spectrometer is equipped with a 300 mW diode laser and the excitation wavelength is 785 nm. The spectrum range was from 100 cm−1 to 3100 cm−1 . The spectrum collection time is 4 s and is averaged over three measurements to improve the signal-to-noise ratio. The TEM study is performed on a FEI Tecnai F20 TEM machine. The electron accelerating voltage is at 200 kV. Small amount of sample is first dispersed in water by sonication. One drop of the aqueous suspension is then dropped onto a thin holey carbon film. The girds are then dried at 60 ◦ C overnight before TEM measurement. The optical properties of the samples are characterized by diffusive reflectance measurement. The reflectance spectra are collected with an Agilent Cary 60 UV-Vis spectrometer with an optical reflectance fiber unit. Fourier-transformed infrared spectroscopy

3. Results and discussions The XRD patterns of the CdS samples obtained with different molar ratios (from 0.04 to 2.00) of Na2 S to Cd(NO3 )2 are shown in Fig. 1A. When the molar ratio is less than 0.15 (curve a–b), the diffraction peaks are mainly from Cd(OH)2 (JCPDS card No. 310228), with tiny peaks from CdS (curve b). This suggests small amount of CdS is formed and majority of the product is the unreacted Cd(OH)2 formed from the precipitation of Cd(NO3 )2 by NaOH. The crystalline grain size can be calculated using the Scherrer equation:  = (K)/(ˇ cos ), where  is the mean size of the ordered (crystalline) domains, which may be smaller or equal to the grain size, K is the shape factor with a typical value of 0.9,  is the X-ray wavelength, ˇ is the line broadening full width at half maximum (FWHM) peak height in radians, and  is the Bragg angle [62]. Calculated with the main peak with 2 at 35.43◦ ((2 2 0) plane of Cd(OH)2 ), the primary grain size of nanostructured Cd(OH)2 is around 26.9 nm (from curve a) to 32.1 nm (from curve b). When the molar ratio of Na2 S to Cd(NO3 )2 equals to 0.48 (curves c), small amount of nanostructured CdS also evolved out, and their sizes are around 2.4 nm when estimated from the (1 1 0) line with 2 of 43.0 (from curve d). Apparently, smaller grain sizes (2.4 nm) of CdS evolved from the larger Cd(OH)2 (29.5 nm) grains in the intermediate steps. When the molar ratio of Na2 S to Cd(NO3 )2 equals to or larger than 0.70, the diffraction peaks are found to be in good agreement with pure hexagonal CdS (JCPDS card No. 41-1049). The three very broadening diffraction peaks at 2 values of 26.9◦ , 43.8◦ and 51.8◦ , correspond to the (0 0 2), (1 1 0) and (1 1 2) planes of a hexagonal phase of CdS, respectively [63–65]. The shape of the XRD pattern keeps unchanged even with increasing molar ratio of Na2 S to Cd(NO3 )2 from 0.70 to 2.00 (curves d–i). The average grain sizes of the CdS is around 3.5 nm regardless of the molar ratio of Na2 S to Cd(NO3 )2 . Therefore, it is concluded from the XRD analysis that the molar ratio of Na2 S to Cd(NO3 )2 does not have an important effect on the average grain size of the nanostructured CdS. No other impurities are detected in the product. The excess Cd(OH)2 or Na2 S may be washed away during the cleaning steps. The XRD patterns of the as-synthesized TiO2 samples are shown in Fig. 1B. All the samples prepared in solutions at room temperature under different pH exhibit broad and low intensity X-ray diffraction bands, which are the characteristics for long-range highly disordered inorganic structure from the hydrolysis and aggregation reactions in the sol–gel process. These featureless bands suggest that all the TiO2 samples have amorphous structures, as reported previously [66,67]. The positions and the widths of these diffraction bands display the extent of the hydrolysis and aggregation reactions before forming the ordered inorganic lattice structures which show well-defined crystal structures of either anatase or rutile, or brookite. The XRD patterns of the CdS/TiO2 nano-heterostructures are shown in Fig. 1C. As the molar ratio of CdS to TiO2 changes from 2:1 to 1:20, the intensity of the CdS gets weaker, and the widths of

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2θ / degree Fig. 1. (A) XRD patterns of CdS samples prepared from a series of various molar ratios of Na2 S to Cd(NO3 )2 : (a) 0.04, (b) 0.15, (c) 0.48, (d) 0.70, (e) 0.84, (f) 1.00, (g) 1.26, (h) 1.68, (i) 2.00. (*): Cd(OH)2 , (o): CdS, (B) XRD patterns of as-prepared TiO2 under different pH: (a) 13.0, (b) 12.0, (c) 11.0, (d) 10.0, and (e) 9.0. (C) XRD patterns of CdS/TiO2 compositions with different molar ratios of CdS to TiO2 : (a) 2:1, (b) 1:1, (c) 1:2, (d) 1:5, (e) 1:8, (f) 1:10, (g) 1:20.

the diffraction peaks keep almost unchanged. As expected, no characteristic peaks of well-crystallized TiO2 phases are found, as the TiO2 formed under such condition has mainly amorphous structure. The weakened intensity of nanostructured CdS is due to the increased thickness of the amorphous TiO2 layer on the surface. The unchanged peak widths of the CdS phase suggest that the addition of the amorphous TiO2 layer on the surface of CdS does not cause the growth of CdS primary grains. The Raman spectra of CdS samples with different molar ratios of Na2 S to Cd(NO3 )2 are shown in Fig. 2A.The Raman peaks near 298.9 cm−l and 596.3 cm−l corresponds to the first (A1 1LO) and second-order (A1 2LO) longitudinal optical (LO) phonon modes of CdS, respectively [68]. The A1 1LO phonon at 298.9 cm−1 is related to vibrations along the c axis [69,70]. The Raman peaks near 216.9 cm−1 , and 465.9 cm−1 , and 979.1 cm−1 are due to the multiphonon response [68]. The peak near 382 cm−1 is attributed to the A1g (T) translational lattice mode of the Cd(OH)2 phase [71]. The vibrational modes of CdS are clearly identified even when the molar ratio of Na2 S to Cd(NO3 )2 is 0.04. Compared to the XRD results where the diffraction peaks from CdS is hardly seen, Raman seems more sensitive to the small amount of CdS formed. The Raman intensity near 382 cm−1 of Cd(OH)2 decreases as the molar ratio of Na2 S to Cd(NO3 )2 increases, and disappear when the ratio increase to above 0.48 (curve e). This suggests the gradual consumption of the Cd(OH)2 . Meanwhile, the intensity of the vibrational modes at 216.9 cm−1 and 298.9 cm−1 (A1 1LO) of CdS increases as the molar ratio of Na2 S to Cd(NO3 )2 increases. However, the vibrational mode of CdS due to the A1 transverse optical (TO) mode at 235 cm−l , E1 TO mode, or their combination, due to the E2 phonon mode at 255 cm−l [68] do not appear for all the samples. The intensity of the multiphonon scattering peak at 216.9 cm−l is large compared to those of the LO and TO modes, suggesting the enhanced multiphonon responses in all the nanostructured CdS samples [68]. Fig. 2B shows the Raman spectra of TiO2 prepared under pH from 9.0 to 13.0. The TiO2 sample prepared at pH = 9.0 shows broad peaks centered at 187.7 cm−1 , 419.5 cm−1 , 458.0 cm−1 , and 612.1 cm−1 , and one weak peak centered at 871.1 cm−1 , accompanied

with large luminescence background. The TiO2 sample prepared at pH = 10.0 shows broad peaks centered at 187.0 cm−1 , 276.7 cm−1 , 443.4 cm−1 , and 609.5 cm−1 , and one weak peak centered at 802.7 cm−1 , accompanied with large luminescence background. The TiO2 samples prepared at pH = 11.0–13.0 show broad peaks centered at 157.0 cm−1 , 280.0 cm−1 , 388.4 cm−1 , 449.4 cm−1 , 660.8 cm−1 , and 903.5 cm−1 , and one weak peaks centered at 824.6 cm−1 . Typically, well-crystalline anatase TiO2 has six Raman-active fundamentals peaks at Eg (144 cm−1 , 197 cm−1 ), B1g (400 cm−1 ), A1g (507 cm−1 ), B1g (519 cm−1 ) and (Eg ) 640 cm−1 and two-phonon scattering bands at 320 cm−1 and 695 cm−1 with a first overtone at 796 cm−1 ; rutile TiO2 crystals show four Ramanactive modes: B1g (143 cm−1 ), Eg 447 (cm−1 ), A1g (612 cm−1 ) and B2g (826 cm−1 ), two-phonon scattering band at 247 cm−1 ; and brookite TiO2 crystals have A1g (127, 154, 194, 247, 412, 497, 640 cm−1 ), B1g (133, 159, 215, 320, 415, 502 cm−1 ), B2g (254, 329, 366, 395, 463, 476, 584 cm−1 ), and B3g (172, 287, 452, 545, 618 cm−1 ) [72]. Some of the vibrational peaks of the TiO2 samples match well with anatase or rutile, or brookite, and other peaks do not match either phase of TiO2 . This suggests that the TiO2 samples formed are full of long-range disorder as seen from the large bandwidths of each peak and the large luminescence background. Their structures are somewhat partially like anatase, or rutile or brookite in some local environment, but different from each of the common TiO2 crystals as a whole. Similar results were reported previously that no well-defined peak can be seen in the spectra of amorphous TiO2 films deposited at substrate temperatures below 350 ◦ C [73]. The Raman spectra of CdS/TiO2 nanocomposites are shown in Fig. 2C. For all the CdS/TiO2 composite samples with different molecular ratios of CdS to TiO2 from 2:1 to 1:20, the Raman spectra of the composites are similar to that of amorphous TiO2 nanostructures, and the Raman vibrational modes of the nanostructured CdS disappear. This, combined with the crystalline nature of the CdS in the composites from the XRD results in Fig. 1C, suggests that most likely that the TiO2 is most likely amorphous and the amorphous TiO2 layer completely covers the surface of the nanostructured CdS in forming CdS/TiO2 crystalline/amorphous nanostructures.

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Fig. 2. (A) Raman spectra of CdS samples prepared with different molar ratios of Na2 S to Cd(NO3 )2 : (a) 0.04, (b) 0.15, (c) 0.48, (d) 0.70, (e) 0.84, (f) 1.00, (g) 1.26, (h) 1.68, (i) 2.00. (B) Raman spectra of as-prepared TiO2 under different pH: (a) 13.0, (b) 12.0, (c) 11.0, (d) 10.0, and (e) 9.0. (C) Raman spectra of CdS/TiO2 compositions with different molar ratios of CdS to TiO2 : (a) 2:1, (b) 1:1, (c) 1:2, (d) 1:5, (e) 1:8, (f) 1:10, (g) 1:20.

Large luminescence background seen for the sample with molecular ratios of CdS to TiO2 from 2:1 and 1:1 also suggests the defective feature of the TiO2 amorphous structures. The FTIR spectra of pure CdS are shown in Fig. 3A.The broad adsorption band center at 3400 cm−1 is related to the stretch region of the surface hydroxyl groups with hydrogen bonds and chemisorbed water [74–76] and the peak centered at 1630 cm−1 from O H bending of physisorbed water[75–78]. The weak band observed around 2979.94 cm−1 is possibly due to the Cd O vibrations [79]. The two peaks at 1130 cm−1 and 996 cm−1 can be attributed to the Cd S bond [80]. It is interesting for the sample obtained at the molar ratio of Na2 S to Cd(NO3 )2 equal 0.25 (curve a), no OH bands from water adsorbed are observed, and strong absorption at 1455 cm−1 , 1382 cm−1 , and 858 cm−1 . As seen from the XRD and Raman results, we know that only small amount of CdS is formed, and the majority of the product is Cd(OH)2 , and the intensities of these absorption bands decreases as the ratio of molar ratio of Na2 S to Cd(NO3 )2 increases. So we can assign these peaks to Cd(OH)2 compounds, either from Cd O or OH vibrations [81,82]. The FTIR spectra of TiO2 are shown in Fig. 3B. The broad adsorption band centered at 3219 cm−1 is related to the stretch region of the surface hydroxyl groups and molecularly chemisorbed water [74,75]. The peak at 1644 cm−1 results from O H bending of molecularly physisorbed water [75,77,78]. The large absorption in the range 600–800 cm−1 are characteristic of the formation of an O Ti O [74]. As the rate of hydrolysis of titanium isopropoxide is faster in higher pH basic solution, the extent of hydrolysis and condensation of the Ti OH is larger. In the first step of hydrolysis, Ti(OH)4 is most likely formed, then followed by condensation reactions in forming polymeric O Ti O network toward the TiO2 structure. Thus more Ti OH is seen in the early stage of condensation or lower pH solution than in the later stage or higher pH solution. As the condensation reaction is more complete in higher pH (curve a) than in lower pH (curve d), the intensities of the peaks of 1107 cm−1 , 1053 cm−1 and 858 cm−1 decrease and the intensities of the peaks of 1540 cm−1 and 1333 cm−1 increases, we can deduce

that 1107 cm−1 , 1053 cm−1 and 858 cm−1 are more related to the free or oligomeric Ti(OH)4 structures [83,84], while 1540 cm−1 and 1333 cm−1 are more related to the O Ti O or O Ti OH vibrations in the polymeric network [85,86]. The FTIR spectra of CdS/TiO2 powders are shown in Fig. 3C. As the molar ratio of CdS to TiO2 decreases from 2:1 to 1:2 and 1:8, the position of the OH vibration bands from the chemisorbed water moves from 3378 to 3338 and 3231 cm−1 . As the vibration of the OH groups in the CdS is around 3400 cm−1 , and the vibration of the OH in the TiO2 is around 3219 cm−1 , this suggests the increasing TiO2 influence on the OH vibrations as the molar ratio of TiO2 to CdS in the composite increases. Meanwhile, for the OH in the physisorbed water, it position keeps relatively unchanged around 1639 cm−1 even the molar ratio of TiO2 to CdS increases, as the physisorbed water shows similar OH vibrational positions either in CdS (1630 cm−1 ) or in TiO2 (1644 cm−1 ). The 1359 cm−1 peak seems to evolve from the 1333 cm−1 from TiO2 , since its intensity increases as the molar ratio of TiO2 to CdS in the composite increases, and the peaks at 1113, 1051 and 1007 cm−1 seem to evolve from CdS since their intensities decrease as the molar ratio of TiO2 to CdS in the composite increases. Figs. 4A and B shows the TEM images of the nanostructured CdS samples along with the selected area electron diffraction (SAED) patterns. Fig. 4A shows the CdS sample is made of hollow nanocuboids of 10–40 nm in width and 40–100 nm in length. The thickness of the wall is around 5–10 nm. The high-resolution TEM (HRTEM) image in Fig. 4B shows that these nanocuboids are well crystallized and are made of aggregates of small crystalline grains of 2.0–5.0 nm in diameter. These results are consistent with the XRD and Raman measurements and are in good agreement with those previously reported by Domen et al. [28]. The SAED pattern in the inset of Fig. 4A testifies the polycrystalline nature of these CdS nanocuboids. Fig. 4C and D shows the TEM images of nanostructured TiO2 samples along with the selected area electron diffraction (SAED) pattern. Large aggregates with the size of 100–200 nm in diameter

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Fig. 3. (A) FTIR spectra of the CdS samples prepared with different molar ratios of Na2 S to Cd(NO3 )2 : (a) 0.15, (b) 0.48, (d) 1.00, (d) 1.26 (e) 2.00. (B) FTIR spectra of as-prepared TiO2 under different pH: (a) 13.0, (b) 12.0, (c) 11.0, and (d) 10.0. (C) FTIR spectra of CdS/TiO2 compositions with different molar ratios of CdS to TiO2 : (a) 2:1, (b) 1:2, (c) 1:8.

are observed and are made of smaller grains of 20–50 nm in diameter as shown in Fig. 4C. The SAED pattern in the inset of Fig. 4C shows diffraction cloud instead of well-defined rings or dots, suggesting the amorphous nature of the TiO2 structure. The HRTEM of the TiO2 shown in Fig. 4C further confirm its amorphous nature with no lattice fringe or ordered atomic alignment seen. The amorphous structure of TiO2 consistent with the XRD and Raman measurements. Fig. 4E and F shows the TEM images of nanostructured CdS/TiO2 composite samples along with the selected area electron diffraction (SAED) pattern. The TEM image in Fig. 4E shows that the CdS/TiO2 composite is still of hollow structures with 20–30 in width and 40–60 nm in length. The SAED pattern in the inset shows satellite and rings of diffraction dots, displaying the polycrystalline nature of the sample, and the center diffraction cloud, also testifying the amorphous nature of the sample on the other hand. Careful comparison of the SAED patterns in Fig. 4A, C and E clearly tells the difference. The crystalline nature of the CdS/TiO2 composite is clearly seen in the HRTEM image in Fig. 4F with well resolved lattice fringes and crystalline grains. Compared to the TEM images of pure CdS nanocuboids in Fig. 4A, that the hollow structures of the CdS/TiO2 composite in Fig. 4E are less clearly seen as the surface of the hollow CdS nanocuboids are covered TiO2 layers. Meanwhile, compared to the HRTEM images of pure CdS nanocuboids in Fig. 4B, some areas of the crystalline lattice fringe in the CdS/TiO2 composite in the HRTEM images in Fig. 4F are less clearly seen, as the surface of the crystalline CdS nanocuboids is covered with amorphous TiO2 layers. These comparisons show that the surface of the hollow crystalline CdS nanocuboids is covered with amorphous TiO2 structure. The UV–vis diffuse reflectance spectra of the as-prepared CdS samples with different molar ratio of Na2 S to Cd(NO3 )2 are shown in Fig. 5A. All the CdS samples absorbs visible and ultraviolet (UV) light and the absorption edge of the nanostructured CdS is related to the ratio of Na2 S to Cd(NO3 )2 during the synthesis. The nanostructured CdS prepared with molar ratio of Na2 S to Cd(NO3 )2 equal to 0.04 starts to absorb light around 506.5 nm. The absorption edge of the nanostructured CdS shifts to the red as the molar ratio of Na2 S to Cd(NO3 )2 increases. When the molar ratio of Na2 S to Cd(NO3 )2

equal to 2.0, the prepared nanostructured CdS shows absorption edge of around 563.8 nm and onset above 600 nm. The optical bandgaps can be quantitatively determined using the Tauc plot by the following equation [87,88]: (˛h)

2/n

∝ h − Eg

(1)

where ˛, , A, and Eg are the absorption coefficient, light frequency, proportionality constant, and bandgap, respectively. In the equation, n decides the characteristics of the transition in a semiconductor, i.e., direct transition (n = 1) or indirect transition (n = 4) [24,89,90]. The intersection of the tangent of the function (˛h)2/n on the photon energy axis (h) gives the value of the bandgap. The Tauc plots of the nanostructured CdS are shown in the inset of Fig. 5B and the change of the bandgap of CdS as a function of the molar ratio of Na2 S to Cd(NO3 )2 in the reaction is shown in Fig. 5B. Clearly, as the molar ratio of Na2 S to Cd(NO3 )2 in the starting reagents increases from 0.04 to 2.0, the bandgap of the resulting nanostructured CdS decreases to 2.21 eV from 2.46 eV. The CdS sample prepared from the molar ratio of Na2 S to Cd(NO3 )2 equals to 1.0 has the smallest bandgap, about 2.19 eV. As the molar ratio further increases, the bandgap slightly increases. The UV–vis diffuse reflectance spectra of the amorphous TiO2 prepared under various pH are shown in Fig. 5C. All the amorphous TiO2 absorbs light in the UV region except the one prepared at pH = 9.0 which shows some absorption in the visible-light range. The Tauc plots of the TiO2 samples are shown in the inset of Fig. 5D, and the derived bandgap of the amorphous TiO2 shows an apparent dependence on the pH value in the hydrolysis reaction. As the pH value decreases from 13.0 to 9.0, the bandgap of the formed TiO2 decreases almost linearly from around 3.30 to 3.14 eV. This bandgaps of the TiO2 samples are close to single crystal anatase (3.20 eV) and the previously reported values for amorphous TiO2 (3.3–3.5 eV) [91–94]. As the hydrolysis and condensation rates in the sol–gel process increases proportionally to the concentration of the [OH− ] (which acts as catalyst in the hydrolysis reaction) in the basic solution [95,96], the extent of the completion of the hydrolysis and condensation reaction in forming the polymeric O Ti O network in the amorphous TiO2 powder increases exponentially with

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Fig. 4. TEM images of nanostructured CdS (A, B, Cd:S = 1), TiO2 (C, D, pH = 13), CdS/TiO2 (E, F, CdS:TiO2 = 1:8). Insets show the SAED patterns.

the value of pH. It is thus expected the amorphous TiO2 powder prepared at lower pH values has more short-range and long-range disorders than the amorphous TiO2 powder prepared at higher pH values. Previous studies on amorphous Si have shown smaller bandgaps than crystalline Si [97], and recent hydrogenated disordered TiO2 has also demonstrated the long-wavelength absorption from the amorphous phase [13]. It is thus reasonable to observe the smaller bandgap of the amorphous nanostructured TiO2 in this study. The UV–vis diffuse reflectance spectrums of the CdS/TiO2 composites are shown in Fig. 5E. All the CdS/TiO2 composites start to absorb light from around 571.5 nm and the absorption bandedge shifts lightly to the blue as the content of TiO2 in the composites increases. As TiO2 only absorbs in the UV region, the absorption in the visible-light region of the composites mainly comes from the CdS component. As the content of the TiO2 on the composite increases, the visible-light absorption from the CdS slightly decreases (from curve a to curve g in Fig. 5E). On the other hand, 9.1% of CdS in the composite is sufficient to bring large amount of absorption in the visible-light region (curve f). This indicates that small amount of CdS can be used to greatly improve the absorbing capability of the amorphous TiO2 for the visible light in the sunlight. The Tauc plots of the CdS/TiO2 composites are shown in the

inset of Fig. 5F, as the visible-light absorption is mainly from the CdS component, n here should equal 1. The change of the bandgap of the CdS/TiO2 nanocomposites as a function of the molar ratio of TiO2 to CdS in the reaction is shown in Fig. 5F. The bandgap of the CdS/TiO2 nanocomposites increases slightly as the molar ratio of TiO2 to CdS in the reaction increases, which causes the increase of the content in the composite and the thickness of the TiO2 layer on the outside of the nanoporous CdS. The increased bandgap of the embedded nanoporous CdS is probably due to the quantum well confinement effects it is facing as the wrapping TiO2 has a large bandgap [98], and possible with some contribution of the dielectric effect of the surrounding TiO2 matrix as well [99]. The decomposition of methylene blue (MB) is used as a probe to evaluate and compare the photocatalytic activities of the nanostructured CdS, TiO2 and CdS/TiO2 composites. When the initial concentration of dye is very small, the degradation of dyes can be described by an apparent first-order equation with a simplified Langmuir–Hinshelwood model [100,101]:

 ln

C0 C

 = ka t

(2)

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40 20

50

(

60

75

2.4

)

(a) (b) (c) (d) (e) (f) (j) (h) (i)

100

25

2.3

0

2.2

0 300 400 500 600 700 800 900 1000

0.0

0.5

D 3.29

80

TiO2 (a) (b) (c)

20

(d)

Bandgap / eV

Reflectance / %

C

40

3.25

1.0

2.0

2

0

3.17

2.6 2.8 3.0 3.2 3.4 3.6 3.8 Photon Energy / eV

3.13

13

12

11

10

9

pH value

Wavelength / nm

F

E

2.22 15

CdS/TiO2 (a) (b) (c) (d) (e) (f) (g)

40 20 0

300 400 500 600 700 800 900 1000

Wavelength / nm

(αhv )2

80

Bandgap / eV

Reflectance / %

1.5

3

TiO2

300 400 500 600 700 800 900 1000

60

2.6

1

3.21

(e)

0

2.4 hv/ ev

Molar ratio of Na 2S to Cd(NO3)2

Wavelength / nm

60

2.2

CdS

(αhv )1/2

80

B

CdS

Bandgap / eV

Reflectance / %

A

7

2.20

(a) (b) (c) (d) (e) (f) (g)

10 5

2.0 2.1 2.2 2.3 2.4 2.5 2.6

2.18

Photon Energy / eV

CdS/TiO2 0

5

10

15

20

Molar ratio of TiO2 to CdS

Fig. 5. (A) UV–vis diffusive reflectance spectra of the CdS samples prepared with different molar ratio of Na2 S to Cd(NO3 )2 : (a) 0.04, (b) 0.15, (c) 0.48, (d) 0.70, (e) 0.84, (f) 1.00, (g) 1.26, (h) 1.68, (i) 2.00. (B) The change of the bandgap of CdS as a function of the molar ratio of Na2 S to Cd(NO3 )2 in the reaction. (C) UV–vis diffusive reflectance spectra of as-prepared TiO2 under different pH: (a) 13.0, (b) 12.0, (c) 11.0, (d) 10.0, and (e) 9.0. (D) The change of the bandgap of TiO2 as a function of the pH of the reaction. (E) UV–vis diffusive reflectance spectra of CdS/TiO2 composite prepared in different molar ratio of CdS to TiO2 : (a) 2:1, (b) 1:1, (c) 1:2, (d) 1:5, (e) 1:8, (f) 1:10, (g) 1:20. (F) The change of the bandgap of CdS/TiO2 as a function of the molar ratio of TiO2 to CdS in the reaction. The insets show the Tauc plots of the UV–vis spectra.

where C0 is the initial concentration of dye, ka is the apparent firstorder rate constant, C is the concentration of the dye and t is the illumination time. Here, we use this equation to analysis the decomposition kinetics of methylene blue on the photocatalysts. Fig. 6A and B gives the time-profile of photodegradation of MB molecules and ln(C0 /C) versus irradiation time for CdS prepared in the solutions with different molar ratio of Na2 S to Cd(NO3 )2 , respectively. It can be seen from Fig. 6B that the rate constant of CdS sample is the largest when the molar ratio of Na2 S to Cd(NO3 )2 equals to 1. This suggests the photocatalytic activities of CdS samples are obviously influenced by the molar ratio of Na2 S to Cd(NO3 )2 . When the molar ratio of Na2 S to Cd(NO3 )2 is below 0.84, the photocatalytic activity of CdS samples increases with the increase of the molar ratio. When the molar ratio of Na2 S to Cd(NO3 )2 is above 1.26, the photocatalytic activity of CdS decreases with the increase of the molar ratio. When the molar ratio of Na2 S to Cd(NO3 )2 is between 0.84 and 1.26, the CdS sample shows the best photocatalytic performance. As these three samples also show the smallest bandgaps, the possible reason could be that these CdS samples absorb the most light in decomposing the methylene blue. The CdS sample prepared with the molar ratio of Na2 S to Cd(NO3 )2 of 0.70 also shows similar bandgap, but its photocatalytic activity is lower. The CdS sample prepared near the molar ratio of Na2 S to Cd(NO3 )2 of 1.0 has closer charge balance on the surface of the CdS photocatalyst, the other CdS samples far from the unit molar ratio of S2− to Cd2+ may suffer large charge imbalance on the surface, which may prevent the

efficient charge transfer across the interface. On the other hand, the CdS samples prepared from the molar ratio of Na2 S to Cd(NO3 )2 of 1.68 and 2.00 show large surface adsorption capability for methylene blue. This can be explained with the effective ionic interaction between their surfaces and the methylene blue molecules as their surfaces are likely negatively charged due to the excess amount of S2− in the solution, and the chromophore moiety of methylene blue is positively charged. So the surface chemistry and physical properties may also contribute to the photocatalytic activity. Fig. 6C and D shows the time-profile of photodegradation of MB molecules and ln(C0 /C) versus irradiation time for amorphous TiO2 prepared under different pH values. It is clear that the photocatalytic performance of the amorphous TiO2 increases monotonically with the pH under which it is prepared. The concentration of methylene blue decreases most with the TiO2 prepared under pH 13.0, followed by pH 12.0, pH 11.0, pH 10.0 and pH 9.0. As the TiO2 sample prepared under pH 9.0 shows the largest photon absorption and the bandgap of the amorphous TiO2 increases with the pH value under which it is prepared, the optical absorption properties would be able to explain the photocatalytic trend of these amorphous TiO2 samples. As seen from XRD and Raman, the amorphous TiO2 samples prepared under higher pH values have better longrange ordering as the hydrolysis and condensation reactions are more complete, thus the degree of amorphous phase may be less. The amorphous phase is capable of trapping excited electrons and holes insides and preventing them in migrating to the surface and

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Fig. 6. (A) Time-profile of photodegradation of MB molecules and (B) ln(C0 /C) versus irradiation time for CdS prepared with different molar ratios of Na2 S to Cd(NO3 )2 : (a) 0.04, (b) 0.15, (c) 0.48, (d) 0.70, (e) 0.84, (f) 1.00, (g) 1.26, (h) 1.68, (i) 2.00; (C) Time-profile of photodegradation of MB molecules and (D) ln(C0 /C) versus irradiation time for TiO2 prepared in different pH condition: (a) 13.0, (b) 12.0, (c) 11.0, (d) 10.0, (e) 9.0; (E) Time-profile of photodegradation of MB molecules and (F) ln(C0 /C) versus irradiation time for CdS/TiO2 composite prepared in different molar ratios of CdS to TiO2 :(a) 2:1, (b) 1:1, (c) 1:2, (d) 1:5, (e) 1:8, (f) 1:10, (g) 1:20, (h) 1:50.

decomposing the methylene blue molecules. So the photocatalytic activity trend of the amorphous TiO2 can mainly be related to the crystallinity of these samples. As the CdS prepared with the molar ratio of Na2 S to Cd(NO3 )2 equal 1.0 and the TiO2 prepared under pH 13.0 show most photocatalytic activities among their series, we prepared the CdS/TiO2 composites under these conditions and only vary their relative molecular ratio. The photocatalytic activities of the CdS/TiO2 composites over the photodegradation of MB molecules are shown in Fig. 6E and F. The photocatalytic activity of the CdS/TiO2 composites increases with the ratio of TiO2 over CdS increases with two exceptions at the ratio of 10 (curve f) and 20 (curve g). Compared the rate constant, it is easy to find that the photocatalytic activity decreases in the order of CdS/TiO2 > TiO2 > CdS. As noted, the sample with around 2% of CdS ingredient (the molar ratio of CdS/TiO2 = 1:50) has the largest rate constant, which demonstrates that the photocatalytic activities of the CdS/TiO2 increase with the thickness of amorphous TiO2 protective layer in the compositions due to the better charge separation across the junction of CdS/TiO2 and the good charge trapping capability of the TiO2 protective layer, which allows longer times for the excited charges to decompose the methylene blue on the surface. The mechanism of photocatalytic photodegradation of MB molecules in CdS–TiO2 system have been discussed in the previous reports [40,49,102], which was also shown in Fig. 7. The carrier of CdS can be excited from the valence band to the conduction band by

visible light due to its low direct bandgap (Eg < 2.4). It is well known that the conduction band (CB) of TiO2 is more positive than that of CdS [40,49]. The photogenerated electrons can be transferred effectively from the CB of CdS to that of TiO2 , while the hole remains in the CdS particles. The photogenerated electrons can react with the adsorbed O2 to form radicals such as • O2− , • OH and then these radicals further oxidize the MB molecules [103]. Meanwhile, the holes on the surface of CdS can react with OH− or H2 O to generate HO• radicals, which can also take part in the degradation reactions of MB in the photocatalytic process. The proposed reactions for the

Fig. 7. The mechanism of photocatalytic photodegradation of MB molecules over CdS/TiO2 composite.

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photocatalytic decomposition of MB by CdS/TiO2 heterostructures can be expressed as follows [49,102–105]: TiO2 /CdS + h␯ → eCB − + hVB +

(3)

eCB − + O2 → O2 •−

(4)

2O2 •− + 2H+ → O2 + H2 O2

(5)

O2 •− + 2H+ + 2eCB − → H2 O2

(6)

H2 O2 + eCB − → • OH + OH−

(7)

H2 O + hVB + → • OH + H+

(8)

In a word, the HO• generated from hole oxidation and electron reduction can be used to oxidize the MB to CO2 and H2 O. In this way, the efficient separation of photogenerated electron–hole pairs can be achieved, which contributes to the improvement of the photodegradation capacity. 4. Conclusions In summary, we have synthesized a series of nanostructured CdS/TiO2 photocatalysts. The CdS photocatalysts have hollow and polycrystalline nanostructures made of small crystalline grains, the TiO2 is mainly amorphous, and the CdS/TiO2 photocatalysts have the crystalline CdS/amorphous TiO2 structures. The optical properties, bandgaps, and photocatalytic activities of the samples depend on their preparation conditions. The bandgap of the pure CdS photocatalysts decreases as the ratio of the starting materials Na2 S to Cd(NO3 )2 increases from 0.04 to 2.00, the bandgap of the amorphous TiO2 photocatalysts decreases almost linearly with the pH of the synthetic solution, and the bandgap of the CdS/TiO2 composite photocatalysts increases slightly with the ratio of the TiO2 over CdS. The photocatalytic activity of CdS, TiO2 and CdS/TiO2 also depend on the ratio of the starting materials Na2 S to Cd(NO3 )2 , the pH, and the ratio of the TiO2 over CdS, respectively. CdS shows the highest activity when the ratio of the starting materials Na2 S to Cd(NO3 )2 is close to 1.0, TiO2 shows the best activity when prepared under pH 13.0, and the CdS/TiO2 composite with the molar ratio of TiO2 to CdS of 50 displays the highest activity. Among the samples studied, CdS/TiO2 composite shows best activity, followed by TiO2 and then CdS. Acknowledgements X. Chen thanks the support from College of Arts and Sciences, University of Missouri – Kansas City and University of Missouri Research Board. X. Li thanks the National Science Foundation of China (No. 20906034) and the National Scholarship Fund of China Scholarship Council (No. 2011844194) for support. References [1] X. Chen, C. Li, M. Gratzel, R. Kostecki, S.S. Mao, Chemical Society Reviews 41 (2012) 7909–7937. [2] X. Chen, S. Shen, L. Guo, S.S. Mao, Chemical Reviews 110 (2010) 6503–6570. [3] A. Kudo, Y. Miseki, Chemical Society Reviews 38 (2009) 253–278. [4] K. Maeda, K. Domen, Journal of Physical Chemistry Letters 1 (2010) 2655–2661. [5] M.R. Hoffmann, S.T. Martin, W. Choi, D.W. Bahnemann, Chemical Reviews 95 (1995) 69–96. [6] A. Fujishima, T. Rao, D. Tryk, Journal of Photochemistry and Photobiology C: Photochemistry Reviews 1 (2000) 1–21. [7] A. Mills, S. Le Hunte, Journal of Photochemistry and Photobiology A: Chemistry 108 (1997) 1–35. [8] T. Inoue, A. Fujishima, S. Konishi, K. Honda, Nature 277 (1979) 637–638.

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