Huilei
Zhao
a,
Lianjun
Liu
a,
Jean M.
Andino
bc and
Ying
Li
*a
aMechanical Engineering Department, University of Wisconsin-Milwaukee, 3200 N Cramer St, Milwaukee, Wisconsin, USA. E-mail: liying@uwm.edu; Fax: +1-414-229-6958; Tel: +1-414-229-3716
bChemical Engineering, Arizona State University, USA
cCivil, Environmental, and Sustainable Engineering, Arizona State University, USA
First published on 28th May 2013
Among the three naturally existing phases of TiO2, brookite is the least studied as a photocatalyst. In this study, single-phase anatase and brookite, and mixed-phase anatase–brookite TiO2 nanomaterials were synthesized through a hydrothermal method. The anatase–brookite phase content was controlled by adjusting the concentration of urea in the precursor solution. XRD, Raman spectroscopy, and high-resolution TEM were used to confirm the crystal structures. SEM and TEM analyses demonstrated that anatase TiO2 were nearly spherical nanoparticles while brookite TiO2 were rod-shaped nanoparticles. UV-vis diffuse reflectance spectroscopy showed a blue shift in absorption spectra with increasing brookite content. The photocatalytic activities of the prepared bicrystalline TiO2 were evaluated for CO2 photoreduction in the presence of water vapor for production of solar fuels (CO and CH4). The activities were compared with those of pure anatase, pure brookite, and a commercial anatase–rutile TiO2 (P25). The results showed that bicrystalline anatase–brookite was generally more active than single-phase anatase, brookite, and P25. The bicrystalline mixture with a composition of 75% anatase and 25% brookite showed the highest photocatalytic activity, likely due to the enhanced interfacial charge transfer between anatase and brookite nanocrystals. In situ DRIFTS analysis showed that CO2− and HCO3− species were active reaction intermediates for CO2 photoreduction while the accumulation of non-reactive CO32− species on the TiO2 surface may be detrimental.
Anatase, brookite, and rutile are naturally existing TiO2 polymorphs.6–9 Anatase based catalysts have been demonstrated to be highly active, while rutile is less active,10,11 mainly due to the fast e− and h+ recombination in rutile.12 Mixtures of anatase–rutile such as commercially available TiO2 nanopowder, Evonik P25 (approximately 75% anatase and 25% rutile) have demonstrated higher photocatalytic activity than single-phase anatase or rutile crystals in various photocatalytic applications.13,14 Brookite is rarely studied in photocatalysis likely due to the past difficulties in synthesizing high purity brookite because of its metastable property.15–18 A few studies demonstrated the high activity of brookite in photocatalytic oxidation applications.19,20 Our recent study4 was the first one to investigate the activity of pure brookite TiO2 for CO2 photoreduction with water; brookite showed higher activity than rutile, and surface-defective brookite was even more active than anatase.
Mixed-phase TiO2 nanomaterials other than commercial P25 TiO2 nanoparticles are scarcely studied in CO2 photoreduction, and the interactions of the different phases during the CO2 photoreduction reactions have not been explored. It is believed that the junctions between different phases of TiO2 play a significant role in improving its photo-oxidation activity.6,7 Several studies have explored the mechanism for the enhanced activity of anatase and rutile mixtures such as P25; however, the findings are not conclusive. Some studies suggested that photoinduced electrons tend to transfer from the higher-level anatase conduction band (CB) to the lower-level rutile CB,21 while the holes in the anatase valance band (VB) could migrate to the rutile VB.11 On the contrary, other studies suggested that electrons transfer from rutile to anatase.12,13,22,23 Among them, Hurum et al.12,13 used electron paramagnetic resonance (EPR) spectroscopy to monitor the direction of electron migration; they suggested that electrons transfer from rutile to anatase because the energy levels of the electron trapping sites in anatase are lower than that of the rutile CB. In addition, the interfacial sites between anatase and rutile facilitate the electron transport and prevent the recombination of e− and h+.12,13 However, some recent studies reported that there are no interactions between the two phases in P25, and anatase and rutile independently catalyze reactions.24
Analogous to anatase–rutile mixed-phase TiO2, anatase–brookite heterojunctions are expected to enhance charge separation as well, probably even superior to anatase–rutile. Possible reasons are: (1) our previous study4 has discovered that brookite itself is more active than rutile in CO2 photoreduction; (2) the CB edge of brookite is slightly above that of anatase,25 making it possible for electrons to transfer from brookite to anatase even in the absence of anatase trapping sites. Some experimental evidence exists in the literature indicating that anatase–brookite mixtures are more active than anatase–rutile mixtures (e.g. P25)26 and pure anatase7,27 for photo-oxidation of organic compounds. While bicrystalline anatase–brookite TiO2 has demonstrated certain promising photocatalytic oxidation abilities, the relationship between phase composition and catalytic activity is not clear. In the case of anatase and rutile mixtures, there is an optimum phase content (i.e., anatase–rutile = 77/23) that corresponds to the highest photoactivity.28,29 The optimum phase composition in an anatase–brookite has never been explored; moreover, no studies have been reported on its application in CO2 photoreduction and the correlation of material properties with photocatalytic activities.
In this work, we have synthesized pure phase anatase, brookite, and mixed-phase anatase–brookite catalysts with a controllable brookite fraction. For the first time in the literature, we have investigated and report the effect of TiO2 phase fractions and nanostructures on CO2 photoreduction, including the optimum phase composition for CO2 photoreduction. We have also conducted in situ spectroscopic studies to investigate the CO2 reduction mechanism.
Catalyst | Urea conc. (M) | Anatase phase (%) | Brookite phase (%) | BET specific surface area (m2 g−1) | Pore size (nm) | Pore volume (cm3 g−1) | Band gap (eV) |
---|---|---|---|---|---|---|---|
A100 | 0.1 | 100 | 0 | 145.6 | 6.2 | 0.23 | 3.08 |
A96B4 | 0.25 | 96 | 4 | 134.5 | 5.9 | 0.20 | 3.12 |
A75B25 | 0.5 | 75 | 25 | 140.3 | 5.5 | 0.20 | 3.15 |
A50B50 | 1.0 | 50 | 50 | 84.5 | 7.0 | 0.18 | 3.15 |
A37B63 | 1.5 | 37 | 63 | 70.2 | 10.7 | 0.19 | 3.20 |
B100 | 7.0 | 0 | 100 | 76.6 | 18.4 | 0.35 | 3.22 |
Raman spectroscopy was carried out with a Renishaw 1000B system in the range of 100–700 cm−1. The specific surface areas of all synthesized catalysts were obtained by nitrogen adsorption–desorption at 77 K using the Brunauer–Emmett–Teller (BET) method (Micrometrics, ASAP 2020). UV-vis spectra of the catalysts were obtained by a diffuse reflectance UV-vis spectrophotometer (Ocean Optics) using BaSO4 as the background. The morphologies of TiO2 polymorphs were characterized by scanning electron microscopy (SEM, Hitachi S4800), using a secondary electron detector (SE) at an accelerating voltage range of 5–10 kV. Transmission electron microscopy (TEM) and high-resolution transmission electron microscopy (HRTEM) were carried out with 300 keV electrons in a Hitachi H9000NAR instrument with 0.18 nm point and 0.11 nm lattice resolutions, to obtain crystal structure, morphology, and lattice information of TiO2 nanocrystals.
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Fig. 1 XRD patterns of the prepared TiO2 catalyst powders with different fractions of anatase and brookite phases (a); Rietveld refinement result for the A50B50 sample – solid line: fitting curve; blue dots: experimental data; bottom panel: goodness of fitting (GOF) (b). |
At the lowest concentration of urea (0.1 M), pure phase anatase (JCPDS no. 21-1272), i.e., A100, was obtained. With an increasing concentration of urea up to 1.5 M, bicrystalline anatase–brookite mixtures with an increasing brookite fraction, i.e., A96B4, A75B25, A50B50, and A37B63 were formed. At the highest urea concentration (7.0 M), pure brookite (JCPDS no. 29-1360) was obtained. The existence of brookite in the resultant powders is readily discernible from its unique (121), (221), and (032) diffraction peaks at 30.81, 42.34, 46.07° (2θ), respectively (Fig. 1a). Clearly, through adjusting the concentration of urea, anatase, brookite or mixtures with controllable anatase–brookite phase fractions were successfully prepared.
The morphologies of the prepared TiO2 nanocrystals were examined by SEM, as shown in Fig. 2. Pure anatase, A100, consisted of agglomerate sphere-shaped nanoparticles, while pure brookite, B100, was rod-shaped nanoparticles. For anatase–brookite mixtures, the brookite nanorods appeared to have a lower aspect ratio (brick-shaped) than those observed in pure brookite (rod-shaped). With an increasing brookite fraction in the anatase–brookite mixtures, the number of nanobricks/nanorods increased and the size became larger. Eventually, the brookite crystals evolved from nanobricks to nanorods as the brookite content approached 100%.
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Fig. 2 SEM images of A100 (a), A96B4 (b), A75B25 (c), A50B50 (d), A37B63 (e), and B100 (f). The circled areas in black in (b), (c), and (d) indicate brookite nanobricks. |
The particle size, morphology, and lattice structure of A75B25 were further evaluated by TEM and HRTEM, as shown in Fig. 3. Mixtures of spherical and rod-shaped nanoparticles were observed (Fig. 3a). The spherical particles are in the range of 5–10 nm and the nanorod has a width of 20 nm and length of 60 nm. HRTEM images in Fig. 3b and c clearly demonstrated that the spherical nanoparticle was anatase (lattice spacing = 0.35 nm, corresponding to the (101) plane) and the nanorod was brookite (lattice spacing = 0.91, 0.35 nm, corresponding to (100) and (210) plane, respectively).4
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Fig. 3 TEM images of A75B25 (a), and HRTEM images of A75B25 (b and c). |
Fig. 3b and c also demonstrated the overlapping anatase and brookite particles and the interface between them. In Fig. 3b the lattice fringe becomes fuzzy and even discontinued (Region 1) where the anatase and brookite crystals overlap. An anatase particle was also located at the edge of brookite (Region 2). Besides those simple attachments of anatase particles to brookite particles, a coherent interface between anatase and brookite could be formed. As shown in Fig. 3c, the lattice spacing of the anatase (101) plane matches that of the brookite (210) plane (0.35 nm) in the region where the two particles overlap (Region 3). This observation agrees with theoretical simulation results that the brookite (210) surface and the most stable anatase (101) plane have the same type of building block, and both phases may share a common boundary.33 On anatase (101) all the octahedron units are closely packed, while on brookite (210) the rotated units are closely packed only along the direction.33,34 This difference may cause a distortion at the interface between the two crystal phases. The HRTEM images in Fig. 3 demonstrated that when anatase and brookite nanocrystals grow in the same orientation, a unique interface between them could be possibly formed.
Raman spectra (Fig. 4) have further proved the existence of both anatase and brookite crystals in the A–B mixture (A75B25). Brookite has the most intense Raman bands at ∼245, ∼320, ∼399, ∼410, and ∼637 cm−1, while anatase has intense Raman bands at ∼399, ∼519, and ∼639 cm−1.35 As compared with the spectra of A100 and B100, the Raman spectrum of A75B25 included all the bands of the two crystal phases, thus indicating a mixed phase of anatase and brookite.
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Fig. 4 Raman spectra of A100, B100, and A75B25 samples. |
Fig. 6 compares the UV-vis spectra, displayed in absorbance units, of the prepared TiO2 catalysts. The absorption edge of A100 is around 402 nm, corresponding to a band gap at 3.08 eV, close to the value reported in the literature.4,36 An increase in the brookite fraction in the anatase–brookite mixture results in a blue shift in the absorption spectra and an increase in the band gap (listed in Table 1). B100 has the largest band gap at 3.22 eV. The observed larger band gap of brookite than anatase agrees with that reported in the literature,19 as the conduction band of brookite is positioned at 0.14 eV more negative than that of anatase.37 The catalysts after CO2 photoreduction experiments were also characterized by UV-vis, and no changes in the catalyst band gap were observed.
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Fig. 6 The UV-vis diffuse reflectance spectra (displayed in absorbance) of the prepared TiO2 catalysts. |
As consistent with our previous studies,4,5 CO was the major product and CH4 was the minor product formed during CO2 photoreduction with H2O vapor, and no other products (such as H2) were detected. The absence of H2 production was consistent with the literature that H2 production is prominent only in the presence of noble metal co-catalysts (e.g., Pt, Ag)1,38 or in the presence of sacrificial agents (e.g., CH3OH).1,39 CH4 production was much lower than CO production, likely because the formation of CH4 requires the complete reduction of CO2 to C, and a subsequent reaction with available H atoms from the decomposition of water vapor on the same catalyst surface. If the CO desorbs from the surface before it can be fully reduced, then the formation of CH4 is impacted.
Fig. 7 compares the production rates of CO on various TiO2 catalysts. For each test 100 mg of catalyst was used (see Fig. S5 in the ESI† regarding the additional experiments using different amount of catalysts and the justification of using 100 mg to study the intrinsic catalyst activity). The test for each sample was repeated twice (a fresh catalyst was used each time), and the errors were within reasonable experimental uncertainties. The CH4 production rates are not plotted in Fig. 7 because they were too low (less than 0.005 μmol h−1) and did not vary significantly for different catalysts. B100 had the lowest activity among all the catalysts, with a CO production rate at 0.07 μmol h−1. A100 had a CO production rate at 0.12 μmol h−1, higher than that of B100. The bicrystalline samples with dominating anatase phase (A96B4 and A75B25) or equal anatase–brookite content (A50B50) were the most active ones, having a CO production rate from 0.16 to 0.21 μmol h−1. The activity of A75B25 was nearly twice as high as that of A100 and three times as high as that of B100. Further increasing the brookite content (i.e., brookite-rich A37B63) led to a lower CO production rate at 0.12 μmol h−1. Commercial TiO2 P25 was also tested as comparison; it had a CO production rate of 0.13 μmol h−1, lower than the anatase-rich bicrystalline anatase–brookite samples. The above results demonstrate that the anatase-rich bicrystalline anatase–brookite mixtures are superior to single crystalline anatase or brookite and anatase–rutile mixtures (i.e., P25). Considering that pure anatase A100 had the largest specific surface area and the smallest band gap (Table 1), the higher activity of bicrystalline anatase–brookite is very likely ascribed to the interactions between the anatase and brookite nanocrystals. This interaction between anatase and brookite seems to be more effective than that between anatase and rutile (as in P25). In addition, it is reasonable to find that anatase-rich A–B mixtures are more active than brookite-rich A–B mixtures since pure anatase is more active than pure brookite.
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Fig. 7 The production rates of CO using various TiO2 catalysts (the reactants were a gas mixture of 97.7 vol% CO2 and 2.3 vol% H2O; 100 mg catalysts were used for each test). |
The experimental results in Fig. 7 indicate that anatase rich A–B mixture with 75% anatase and 25% brookite, i.e., A75B25 is the most active among all the prepared catalysts. There are three possible reasons for its superior activity: (1) the more negative CB band of brookite may induce electron transfer from brookite to anatase, thus enhancing charge separation; (2) the facilitated formation of the intrinsic defect sites in brookite may enhance electron trapping,40 and (3) the distorted interfaces between anatase and brookite crystals (as evidenced in the HRTEM images) may facilitate interfacial electron transport and prevent the recombination of electron–hole pairs.41 It is also noticed from Fig. 6 that anatase-rich A–B mixtures such as A75B25 showed less light absorption in the UV region compared to other samples; the highest photocatalytic activity of A75B25 implies the importance of interfacial charge transfer between the A and B nanocrystals.
Additional experiments were conducted to explore the potential O2 production from H2O oxidation using the sample of A75B25. Prior to the photocatalytic reaction, the photoreactor was purged with a CO2–H2O gas mixture to eliminate the air inside the reactor. However, even after purging for a few hours, there was still background O2 (in a few hundred ppm range) together with N2 detected in the reactor effluent gas. Hence, a better indicator of O2 production by the catalyst is the volumetric ratio of O2/N2 in the effluent gas, as also suggested in the literature and in our previous work.4,42Fig. 8 shows the O2/N2 ratio before, during and after the photoreaction, when the reactor effluent gas was sampled every 15 min. Before turning on the light (in the dark), the O2/N2 ratio was steady. Immediately upon photo-illumination, the O2/N2 ratio dramatically decreased in the first 60 min and then gradually increased. The sharp decrease is probably due to the consumption of residual O2 in the reactor through the reaction with photogenerated electrons O2 + e− → O2− and consecutive reactions.4 In the meantime upon photo-illumination, the production of CO and CH4 occurred, and the rates increased to their maximum in 30 to 60 min (see Fig. S6, ESI†). This indicates that although O2 competes with CO2 for photogenerated electrons, it cannot completely block the electron transfer to CO2, since the concentration of CO2 was several orders of magnitude higher than O2 in the reactor. The gradual increase in the O2/N2 ratio after 60 min photo-irradiation (Fig. 8) indicates the generation of O2 through oxidation of H2O with photogenerated holes, H2O + 2h+ → 2H+ + (1/2)O2. Because the increasing amount of generated O2 overweighed the concurrent O2 consumption process, the net result was that the O2/N2 ratio gradually increased with time and finally reached in a steady state after 200 min. After turning off the light (the second dark period), the O2/N2 ratio dropped slightly. This result again verified the generation of O2 during the photocatalytic CO2 reduction with H2O vapor.
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Fig. 8 Time dependence of the volumetric ratio of O2/N2 before, during and after the photoreduction of CO2 with H2O on the A75B25. |
As shown in Fig. 9, in the dark, CO2–H2O adsorption on A75B25 resulted in the formation of bicarbonate (HCO3−, 1220 and 1420 cm−1), bidentate carbonate (b-CO32−, 1360 and 1570 cm−1), monodentate carbonate (m-CO32−, 1300 and 1540 cm−1), and carboxylate (CO2−, 1250 and 1673 cm−1).4,43,44 The intensities of the peaks remained almost unchanged from 20 min to 30 min, indicating that the adsorption of CO2–H2O reached the saturated level on the surface of A75B25. Immediately upon photo-irradiation for 15 min, the peaks for HCO3− and CO2− species decreased or even disappeared, while the peak for surface H2O (1639 cm−1)4 and CO32− increased. Upon increasing the photo-illumination time from 15 min to 3 h, the intensities of the H2O and CO32− peaks gradually increased, while that of HCO3− was gradually decreased and CO2− completely disappeared.
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Fig. 9 In situ DRIFTS spectra of CO2 and H2O adsorption on A75B25 for 30 min in the dark (short-dash lines), and subsequent photoreduction process under the UV-vis light irradiation for 3 h (solid lines). |
In our previous studies,45 we found that HCO3− and CO2−, formed from CO2 interaction with OH groups and Ti3+ sites, are the possible intermediates for CO2 photoreduction to CO and C1 fuels (e.g., CH4) once the dissociative H atom is available. The DRIFTS results in Fig. 9 again demonstrated the importance of the HCO3− and CO2− as the CO2 reduction intermediates, evidenced by the weakened HCO3− peak and the disappearance of CO2− upon photo-illumination. The increase of CO32− species could partially result from HCO3− transformation. The increase in surface H2O may be explained as follows. It is known that H2O and CO2 competitively adsorb on TiO2. H2O and CO2 adsorption reached equilibrium in the dark. Upon photo-illumination, this equilibrium was broken when adsorbed CO2 species was reduced to CO and desorbed into the gas phase. Hence, some surface sites like OH groups became available for incoming gas molecule adsorption. Because H2O has a higher binding energy with surface OH groups than CO2, once the OH groups were re-exposed, most of them may be occupied by H2O, and thus preventing subsequent CO2 adsorption on the surface. The lack of continuous formation of active CO2 adsorption species (HCO3− and CO2−, not CO32−) species may cause the observed catalyst deactivation over a few hours observed in the photocatalytic activity measurements. The above described DRIFTS results have provided significant insights into the reaction mechanism and catalyst deactivation, which have been rarely studied in the literature.
Footnote |
† Electronic supplementary information (ESI) available: Detailed description of the Rietveld refinement of the XRD data, and the investigation of the effect of catalyst mass on the photocatalytic activity. See DOI: 10.1039/c3ta11226h |
This journal is © The Royal Society of Chemistry 2013 |