Titanium Dioxide Nanomaterials: Synthesis ... - ACS Publications

Chem. Rev. 2007, 107, 2891−2959

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Titanium Dioxide Nanomaterials: Synthesis, Properties, Modifications, and Applications Xiaobo Chen* and Samuel S. Mao† Lawrence Berkeley National Laboratory, and University of California, Berkeley, California 94720 Received March 27, 2006

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Contents 1. Introduction 2. Synthetic Methods for TiO2 Nanostructures 2.1. Sol−Gel Method 2.2. Micelle and Inverse Micelle Methods 2.3. Sol Method 2.4. Hydrothermal Method 2.5. Solvothermal Method 2.6. Direct Oxidation Method 2.7. Chemical Vapor Deposition 2.8. Physical Vapor Deposition 2.9. Electrodeposition 2.10. Sonochemical Method 2.11. Microwave Method 2.12. TiO2 Mesoporous/Nanoporous Materials 2.13. TiO2 Aerogels 2.14. TiO2 Opal and Photonic Materials 2.15. Preparation of TiO2 Nanosheets 3. Properties of TiO2 Nanomaterials 3.1. Structural Properties of TiO2 Nanomaterials 3.2. Thermodynamic Properties of TiO2 Nanomaterials 3.3. X-ray Diffraction Properties of TiO2 Nanomaterials 3.4. Raman Vibration Properties of TiO2 Nanomaterials 3.5. Electronic Properties of TiO2 Nanomaterials 3.6. Optical Properties of TiO2 Nanomaterials 3.7. Photon-Induced Electron and Hole Properties of TiO2 Nanomaterials 4. Modifications of TiO2 Nanomaterials 4.1. Bulk Chemical Modification: Doping 4.1.1. Synthesis of Doped TiO2 Nanomaterials 4.1.2. Properties of Doped TiO2 Nanomaterials 4.2. Surface Chemical Modifications 4.2.1. Inorganic Sensitization 5. Applications of TiO2 Nanomaterials 5.1. Photocatalytic Applications 5.1.1. Pure TiO2 Nanomaterials: First Generation 5.1.2. Metal-Doped TiO2 Nanomaterials: Second Generation 5.1.3. Nonmetal-Doped TiO2 Nanomaterials: Third Generation * Corresponding author. E-mail: [email protected]. † E-mail: [email protected].

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5.2. Photovoltaic Applications 5.2.1. The TiO2 Nanocrystalline Electrode in DSSCs 5.2.2. Metal/Semiconductor Junction Schottky Diode Solar Cell 5.2.3. Doped TiO2 Nanomaterials-Based Solar Cell 5.3. Photocatalytic Water Splitting 5.3.1. Fundamentals of Photocatalytic Water Splitting 5.3.2. Use of Reversible Redox Mediators 5.3.3. Use of TiO2 Nanotubes 5.3.4. Water Splitting under Visible Light 5.3.5. Coupled/Composite Water-Splitting System 5.4. Electrochromic Devices 5.4.1. Fundamentals of Electrochromic Devices 5.4.2. Electrochromophore for an Electrochromic Device 5.4.3. Counterelectrode for an Electrochromic Device 5.4.4. Photoelectrochromic Devices 5.5. Hydrogen Storage 5.6. Sensing Applications 6. Summary 7. Acknowledgment 8. References

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1. Introduction Since its commercial production in the early twentieth century, titanium dioxide (TiO2) has been widely used as a pigment1 and in sunscreens,2,3 paints,4 ointments, toothpaste,5 etc. In 1972, Fujishima and Honda discovered the phenomenon of photocatalytic splitting of water on a TiO2 electrode under ultraviolet (UV) light.6-8 Since then, enormous efforts have been devoted to the research of TiO2 material, which has led to many promising applications in areas ranging from photovoltaics and photocatalysis to photo-/electrochromics and sensors.9-12 These applications can be roughly divided into “energy” and “environmental” categories, many of which depend not only on the properties of the TiO2 material itself but also on the modifications of the TiO2 material host (e.g., with inorganic and organic dyes) and on the interactions of TiO2 materials with the environment. An exponential growth of research activities has been seen in nanoscience and nanotechnology in the past decades.13-17 New physical and chemical properties emerge when the size of the material becomes smaller and smaller, and down to

10.1021/cr0500535 CCC: $65.00 © 2007 American Chemical Society Published on Web 06/23/2007

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Dr. Xiaobo Chen is a research engineer at The University of California at Berkeley and a Lawrence Berkeley National Laboratory scientist. He obtained his Ph.D. Degree in Chemistry from Case Western Reserve University. His research interests include photocatalysis, photovoltaics, hydrogen storage, fuel cells, environmental pollution control, and the related materials and devices development.

Dr. Samuel S. Mao is a career staff scientist at Lawrence Berkeley National Laboratory and an adjunct faculty at The University of California at Berkeley. He obtained his Ph.D. degree in Engineering from The University of California at Berkeley in 2000. His current research involves the development of nanostructured materials and devices, as well as ultrafast laser technologies. Dr. Mao is the team leader of a high throughput materials processing program supported by the U.S. Department of Energy.

the nanometer scale. Properties also vary as the shapes of the shrinking nanomaterials change. Many excellent reviews and reports on the preparation and properties of nanomaterials have been published recently.6-44 Among the unique properties of nanomaterials, the movement of electrons and holes in semiconductor nanomaterials is primarily governed by the well-known quantum confinement, and the transport properties related to phonons and photons are largely affected by the size and geometry of the materials.13-16 The specific surface area and surface-to-volume ratio increase dramatically as the size of a material decreases.13,21 The high surface area brought about by small particle size is beneficial to many TiO2-based devices, as it facilitates reaction/interaction between the devices and the interacting media, which mainly occurs on the surface or at the interface and strongly depends on the surface area of the material. Thus, the performance of TiO2-based devices is largely influenced by the sizes of the TiO2 building units, apparently at the nanometer scale. As the most promising photocatalyst,7,11,12,33 TiO2 materials are expected to play an important role in helping solve

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many serious environmental and pollution challenges. TiO2 also bears tremendous hope in helping ease the energy crisis through effective utilization of solar energy based on photovoltaic and water-splitting devices.9,31,32 As continued breakthroughs have been made in the preparation, modification, and applications of TiO2 nanomaterials in recent years, especially after a series of great reviews of the subject in the 1990s.7,8,10-12,33,45 we believe that a new and comprehensive review of TiO2 nanomaterials would further promote TiO2-based research and development efforts to tackle the environmental and energy challenges we are currently facing. Here, we focus on recent progress in the synthesis, properties, modifications, and applications of TiO2 nanomaterials. The syntheses of TiO2 nanomaterials, including nanoparticles, nanorods, nanowires, and nanotubes are primarily categorized with the preparation method. The preparations of mesoporous/nanoporous TiO2, TiO2 aerogels, opals, and photonic materials are summarized separately. In reviewing nanomaterial synthesis, we present a typical procedure and representative transmission or scanning electron microscopy images to give a direct impression of how these nanomaterials are obtained and how they normally appear. For detailed instructions on each synthesis, the readers are referred to the corresponding literature. The structural, thermal, electronic, and optical properties of TiO2 nanomaterials are reviewed in the second section. As the size, shape, and crystal structure of TiO2 nanomaterials vary, not only does surface stability change but also the transitions between different phases of TiO2 under pressure or heat become size dependent. The dependence of X-ray diffraction patterns and Raman vibrational spectra on the size of TiO2 nanomaterials is also summarized, as they could help to determine the size to some extent, although correlation of the spectra with the size of TiO2 nanomaterials is not straightforward. The review of modifications of TiO2 nanomaterials is mainly limited to the research related to the modifications of the optical properties of TiO2 nanomaterials, since many applications of TiO2 nanomaterials are closely related to their optical properties. TiO2 nanomaterials normally are transparent in the visible light region. By doping or sensitization, it is possible to improve the optical sensitivity and activity of TiO2 nanomaterials in the visible light region. Environmental (photocatalysis and sensing) and energy (photovoltaics, water splitting, photo-/electrochromics, and hydrogen storage) applications are reviewed with an emphasis on clean and sustainable energy, since the increasing energy demand and environmental pollution create a pressing need for clean and sustainable energy solutions. The fundamentals and working principles of the TiO2 nanomaterials-based devices are discussed to facilitate the understanding and further improvement of current and practical TiO2 nanotechnology.

2. Synthetic Methods for TiO2 Nanostructures 2.1. Sol−Gel Method The sol-gel method is a versatile process used in making various ceramic materials.46-50 In a typical sol-gel process, a colloidal suspension, or a sol, is formed from the hydrolysis and polymerization reactions of the precursors, which are usually inorganic metal salts or metal organic compounds such as metal alkoxides. Complete polymerization and loss of solvent leads to the transition from the liquid sol into a solid gel phase. Thin films can be produced on a piece of

Titanium Dioxide Nanomaterials

substrate by spin-coating or dip-coating. A wet gel will form when the sol is cast into a mold, and the wet gel is converted into a dense ceramic with further drying and heat treatment. A highly porous and extremely low-density material called an aerogel is obtained if the solvent in a wet gel is removed under a supercritical condition. Ceramic fibers can be drawn from the sol when the viscosity of a sol is adjusted into a proper viscosity range. Ultrafine and uniform ceramic powders are formed by precipitation, spray pyrolysis, or emulsion techniques. Under proper conditions, nanomaterials can be obtained. TiO2 nanomaterials have been synthesized with the solgel method from hydrolysis of a titanium precusor.51-78 This process normally proceeds via an acid-catalyzed hydrolysis step of titanium(IV) alkoxide followed by condensation.51,63,66,79-91 The development of Ti-O-Ti chains is favored with low content of water, low hydrolysis rates, and excess titanium alkoxide in the reaction mixture. Threedimensional polymeric skeletons with close packing result from the development of Ti-O-Ti chains. The formation of Ti(OH)4 is favored with high hydrolysis rates for a medium amount of water. The presence of a large quantity of Ti-OH and insufficient development of three-dimensional polymeric skeletons lead to loosely packed first-order particles. Polymeric Ti-O-Ti chains are developed in the presence of a large excess of water. Closely packed firstorder particles are yielded via a three-dimensionally developed gel skeleton.51,63,66,79-91 From the study on the growth kinetics of TiO2 nanoparticles in aqueous solution using titanium tetraisopropoxide (TTIP) as precursor, it is found that the rate constant for coarsening increases with temperature due to the temperature dependence of the viscosity of the solution and the equilibrium solubility of TiO2.63 Secondary particles are formed by epitaxial self-assembly of primary particles at longer times and higher temperatures, and the number of primary particles per secondary particle increases with time. The average TiO2 nanoparticle radius increases linearly with time, in agreement with the Lifshitz-SlyozovWagner model for coarsening.63 Highly crystalline anatase TiO2 nanoparticles with different sizes and shapes could be obtained with the polycondensation of titanium alkoxide in the presence of tetramethylammonium hydroxide.52,62 In a typical procedure, titanium alkoxide is added to the base at 2 °C in alcoholic solvents in a threeneck flask and is heated at 50-60 °C for 13 days or at 90100 °C for 6 h. A secondary treatment involving autoclave heating at 175 and 200 °C is performed to improve the crystallinity of the TiO2 nanoparticles. Representative TEM images are shown in Figure 1 from the study of Chemseddine et al.52 A series of thorough studies have been conducted by Sugimoto et al. using the sol-gel method on the formation of TiO2 nanoparticles of different sizes and shapes by tuning the reaction parameters.67-71 Typically, a stock solution of a 0.50 M Ti source is prepared by mixing TTIP with triethanolamine (TEOA) ([TTIP]/[TEOA] ) 1:2), followed by addition of water. The stock solution is diluted with a shape controller solution and then aged at 100 °C for 1 day and at 140 °C for 3 days. The pH of the solution can be tuned by adding HClO4 or NaOH solution. Amines are used as the shape controllers of the TiO2 nanomaterials and act as surfactants. These amines include TEOA, diethylenetriamine, ethylenediamine, trimethylenediamine, and triethylenetetramine. The morphology of the TiO2 nanoparticles

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changes from cuboidal to ellipsoidal at pH above 11 with TEOA. The TiO2 nanoparticle shape evolves into ellipsoidal above pH 9.5 with diethylenetriamine with a higher aspect ratio than that with TEOA. Figure 2 shows representative TEM images of the TiO2 nanoparticles under different initial pH conditions with the shape control of TEOA at [TEOA]/ [TIPO] ) 2.0. Secondary amines, such as diethylamine, and tertiary amines, such as trimethylamine and triethylamine, act as complexing agents of Ti(IV) ions to promote the growth of ellipsoidal particles with lower aspect ratios. The shape of the TiO2 nanoparticle can also be tuned from roundcornered cubes to sharp-edged cubes with sodium oleate and sodium stearate.70 The shape control is attributed to the tuning of the growth rate of the different crystal planes of TiO2 nanoparticles by the specific adsorption of shape controllers to these planes under different pH conditions.70 A prolonged heating time below 100 °C for the as-prepared gel can be used to avoid the agglomeration of the TiO2 nanoparticles during the crystallization process.58,72 By heating amorphous TiO2 in air, large quantities of single-phase anatase TiO2 nanoparticles with average particle sizes between 7 and 50 nm can be obtained, as reported by Zhang and Banfield.73-77 Much effort has been exerted to achieve highly crystallized and narrowly dispersed TiO2 nanoparticles using the sol-gel method with other modifications, such as a semicontinuous reaction method by Znaidi et al.78 and a twostage mixed method and a continuous reaction method by Kim et al.53,54 By a combination of the sol-gel method and an anodic alumina membrane (AAM) template, TiO2 nanorods have been successfully synthesized by dipping porous AAMs into a boiled TiO2 sol followed by drying and heating processes.92,93 In a typical experiment, a TiO2 sol solution is prepared by mixing TTIP dissolved in ethanol with a solution containing water, acetyl acetone, and ethanol. An AAM is immersed into the sol solution for 10 min after being boiled in ethanol; then it is dried in air and calcined at 400 °C for 10 h. The AAM template is removed in a 10 wt % H3PO4 aqueous solution. The calcination temperature can be used to control the crystal phase of the TiO2 nanorods. At low temperature, anatase nanorods can be obtained, while at high temperature rutile nanorods can be obtained. The pore size of the AAM template can be used to control the size of these TiO2 nanorods, which typically range from 100 to 300 nm in diameter and several micrometers in length. Apparently, the size distribution of the final TiO2 nanorods is largely controlled by the size distribution of the pores of the AAM template. In order to obtain smaller and monosized TiO2 nanorods, it is necessary to fabricate high-quality AAM templates. Figure 3 shows a typical TEM for TiO2 nanorods fabricated with this method. Normally, the TiO2 nanorods are composed of small TiO2 nanoparticles or nanograins. By electrophoretic deposition of TiO2 colloidal suspensions into the pores of an AAM, ordered TiO2 nanowire arrays can be obtained.94 In a typical procedure, TTIP is dissolved in ethanol at room temperature, and glacial acetic acid mixed with deionized water and ethanol is added under pH ) 2-3 with nitric acid. Platinum is used as the anode, and an AAM with an Au substrate attached to Cu foil is used as the cathode. A TiO2 sol is deposited into the pores of the AMM under a voltage of 2-5 V and annealed at 500 °C for 24 h. After dissolving the AAM template in a 5 wt % NaOH solution, isolated TiO2 nanowires are obtained. In order to


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Figure 1. TEM images of TiO2 nanoparticles prepared by hydrolysis of Ti(OR)4 in the presence of tetramethylammonium hydroxide. Reprinted with permission from Chemseddine, A.; Moritz, T. Eur. J. Inorg. Chem. 1999, 235. Copyright 1999 Wiley-VCH.

Figure 2. TEM images of uniform anatase TiO2 nanoparticles. Reprinted from Sugimoto, T.; Zhou, X.; Muramatsu, A. J. Colloid Interface Sci. 2003, 259, 53, Copyright 2003, with permission from Elsevier.

fabricate TiO2 nanowires instead of nanorods, an AAM with long pores is a must. TiO2 nanotubes can also be obtained using the sol-gel method by templating with an AAM95-98 and other organic compounds.99,100 For example, when an AAM is used as the template, a thin layer of TiO2 sol on the wall of the pores of

the AAM is first prepared by sucking TiO2 sol into the pores of the AAM and removing it under vacuum; TiO2 nanowires are obtained after the sol is fully developed and the AAM is removed. In the procedure by Lee and co-workers,96 a TTIP solution was prepared by mixing TTIP with 2-propanol and 2,4-pentanedione. After the AAM was dipped into this



Figure 5. SEM of a TiO2 nanotube array; the inset shows the ZnO nanorod array template. Reprinted with permission from Qiu, J. J.; Yu, W. D.; Gao, X. D.; Li, X. M. Nanotechnology 2006, 17, 4695. Copyright 2006 IOP Publishing Ltd. Figure 3. TEM image of anatase nanorods and a single nanorod composed of small TiO2 nanoparticles or nanograins (inset). Reprinted from Miao, L.; Tanemura, S.; Toh, S.; Kaneko, K.; Tanemura, M. J. Cryst. Growth 2004, 264, 246, Copyright 2004, with permission from Elsevier.

Figure 4. SEM image of TiO2 nanotubes prepared from the AAO template. Reprinted with permission from Liu, S. M.; Gan, L. M.; Liu, L. H.; Zhang, W. D.; Zeng, H. C. Chem. Mater. 2002, 14, 1391. Copyright 2002 American Chemical Society.

solution, it was removed from the solution and placed under vacuum until the entire volume of the solution was pulled through the AAM. The AAM was hydrolyzed by water vapor over a HCl solution for 24 h, air-dried at room temperature, and then calcined in a furnace at 673 K for 2 h and cooled to room temperature with a temperature ramp of 2 °C/h. Pure TiO2 nanotubes were obtained after the AAM was dissolved in a 6 M NaOH solution for several minutes.96 Alternatively, TiO2 nanotubes could be obtained by coating the AAM membranes at 60 °C for a certain period of time (12-48 h) with dilute TiF4 under pH ) 2.1 and removing the AAM after TiO2 nanotubes were fully developed.97 Figure 4 shows a typical SEM image of the TiO2 nanotube array from the AAM template.97 In another scheme, a ZnO nanorod array on a glass substrate can be used as a template to fabricate TiO2 nanotubes with the sol-gel method.101 Briefly, TiO2 sol is

deposited on a ZnO nanorod template by dip-coating with a slow withdrawing speed, then dried at 100 °C for 10 min, and heated at 550 °C for 1 h in air to obtain ZnO/TiO2 nanorod arrays. The ZnO nanorod template is etched-up by immersing the ZnO/TiO2 nanorod arrays in a dilute hydrochloric acid aqueous solution to obtain TiO2 nanotube arrays. Figure 5 shows a typical SEM image of the TiO2 nanotube array with the ZnO nanorod array template. The TiO2 nanotubes inherit the uniform hexagonal cross-sectional shape and the length of 1.5 µm and inner diameter of 100120 nm of the ZnO nanorod template. As the concentration of the TiO2 sol is constant, well-aligned TiO2 nanotube arrays can only be obtained from an optimal dip-coating cycle number in the range of 2-3 cycles. A dense porous TiO2 thick film with holes is obtained instead if the dip-coating number further increases. The heating rate is critical to the formation of TiO2 nanotube arrays. When the heating rate is extra rapid, e.g., above 6 °C min-1, the TiO2 coat will easily crack and flake off from the ZnO nanorods due to great tensile stress between the TiO2 coat and the ZnO template, and a TiO2 film with loose, porous nanostructure is obtained.

2.2. Micelle and Inverse Micelle Methods Aggregates of surfactant molecules dispersed in a liquid colloid are called micelles when the surfactant concentration exceeds the critical micelle concentration (CMC). The CMC is the concentration of surfactants in free solution in equilibrium with surfactants in aggregated form. In micelles, the hydrophobic hydrocarbon chains of the surfactants are oriented toward the interior of the micelle, and the hydrophilic groups of the surfactants are oriented toward the surrounding aqueous medium. The concentration of the lipid present in solution determines the self-organization of the molecules of surfactants and lipids. The lipids form a single layer on the liquid surface and are dispersed in solution below the CMC. The lipids organize in spherical micelles at the first CMC (CMC-I), into elongated pipes at the second CMC (CMC-II), and into stacked lamellae of pipes at the lamellar point (LM or CMC-III). The CMC depends on the chemical composition, mainly on the ratio of the head area and the tail length. Reverse micelles are formed in nonaqueous media, and the hydrophilic headgroups are directed toward the core of the micelles while the hydrophobic groups are


directed outward toward the nonaqueous media. There is no obvious CMC for reverse micelles, because the number of aggregates is usually small and they are not sensitive to the surfactant concentration. Micelles are often globular and roughly spherical in shape, but ellipsoids, cylinders, and bilayers are also possible. The shape of a micelle is a function of the molecular geometry of its surfactant molecules and solution conditions such as surfactant concentration, temperature, pH, and ionic strength. Micelles and inverse micelles are commonly employed to synthesize TiO2 nanomaterials.102-110 A statistical experimental design method was conducted by Kim et al. to optimize experimental conditions for the preparation of TiO2 nanoparticles.103 The values of H2O/surfactant, H2O/titanium precursor, ammonia concentration, feed rate, and reaction temperature were significant parameters in controlling TiO2 nanoparticle size and size distribution. Amorphous TiO2 nanoparticles with diameters of 10-20 nm were synthesized and converted to the anatase phase at 600 °C and to the more thermodynamically stable rutile phase at 900 °C. Li et al. developed TiO2 nanoparticles with the chemical reactions between TiCl4 solution and ammonia in a reversed microemulsion system consisting of cyclohexane, poly(oxyethylene)5 nonyle phenol ether, and poly(oxyethylene)9 nonyle phenol ether.104 The produced amorphous TiO2 nanoparticles transformed into anatase when heated at temperatures from 200 to 750 °C and into rutile at temperatures higher than 750 °C. Agglomeration and growth also occurred at elevated temperatures. Shuttle-like crystalline TiO2 nanoparticles were synthesized by Zhang et al. with hydrolysis of titanium tetrabutoxide in the presence of acids (hydrochloric acid, nitric acid, sulfuric acid, and phosphoric acid) in NP-5 (Igepal CO-520)cyclohexane reverse micelles at room temperature.110 The crystal structure, morphology, and particle size of the TiO2 nanoparticles were largely controlled by the reaction conditions, and the key factors affecting the formation of rutile at room temperature included the acidity, the type of acid used, and the microenvironment of the reverse micelles. Agglomeration of the particles occurred with prolonged reaction times and increasing the [H2O]/[NP-5] and [H2O]/[Ti(OC4H9)4] ratios. When suitable acid was applied, round TiO2 nanoparticles could also be obtained. Representative TEM images of the shuttle-like and round-shaped TiO2 nanoparticles are shown in Figure 6. In the study carried out by Lim et al., TiO2 nanoparticles were prepared by the controlled hydrolysis of TTIP in reverse micelles formed in CO2 with the surfactants ammonium carboxylate perfluoropolyether (PFPECOO-NH4+) (MW 587) and poly(dimethyl amino ethyl methacrylate-block-1H,1H,2H,2H-perfluorooctyl methacrylate) (PDMAEMA-b-PFOMA).106 It was found that the crystallite size prepared in the presence of reverse micelles increased as either the molar ratio of water to surfactant or the precursor to surfactant ratio increased. The TiO2 nanomaterials prepared with the above micelle and reverse micelle methods normally have amorphous structure, and calcination is usually necessary in order to induce high crystallinity. However, this process usually leads to the growth and agglomeration of TiO2 nanoparticles. The crystallinity of TiO2 nanoparticles initially (synthesized by controlled hydrolysis of titanium alkoxide in reverse micelles in a hydrocarbon solvent) could be improved by annealing in the presence of the micelles at temperatures considerably lower than those required for the traditional calcination

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Figure 6. TEM images of the shuttle-like and round-shaped (inset) TiO2 nanoparticles. From: Zhang, D., Qi, L., Ma, J., Cheng, H. J. Mater. Chem. 2002, 12, 3677 (http://dx.doi.org/10.1039/b206996b). s Reproduced by permission of The Royal Society of Chemistry.

Figure 7. HRTEM images of a TiO2 nanoparticle after annealing. Reprinted with permission from Lin, J.; Lin, Y.; Liu, P.; Meziani, M. J.; Allard, L. F.; Sun, Y. P. J. Am. Chem. Soc. 2002, 124, 11514. Copyright 2002 American Chemical Society.

treatment in the solid state.108 This procedure could produce crystalline TiO2 nanoparticles with unchanged physical dimensions and minimal agglomeration and allows the preparation of highly crystalline TiO2 nanoparticles, as shown in Figure 7, from the study of Lin et al.108

2.3. Sol Method The sol method here refers to the nonhydrolytic sol-gel processes and usually involves the reaction of titanium chloride with a variety of different oxygen donor molecules, e.g., a metal alkoxide or an organic ether.111-119

TiX4 + Ti(OR)4 f 2TiO2 + 4RX

(1)

TiX4 + 2ROR f TiO2 + 4RX

(2)



Figure 8. TEM image of TiO2 nanoparticles derived from reaction of TiCl4 and TTIP in TOPO/heptadecane at 300 °C. The inset shows a HRTEM image of a single particle. Reprinted with permission from Trentler, T. J.; Denler, T. E.; Bertone, J. F.; Agrawal, A.; Colvin, V. L. J. Am. Chem. Soc. 1999, 121, 1613. Copyright 1999 American Chemical Society.

The condensation between Ti-Cl and Ti-OR leads to the formation of Ti-O-Ti bridges. The alkoxide groups can be provided by titanium alkoxides or can be formed in situ by reaction of the titanium chloride with alcohols or ethers. In the method by Trentler and Colvin,119 a metal alkoxide was rapidly injected into the hot solution of titanium halide mixed with trioctylphosphine oxide (TOPO) in heptadecane at 300 °C under dry inert gas protection, and reactions were completed within 5 min. For a series of alkyl substituents including methyl, ethyl, isopropyl, and tert-butyl, the reaction rate dramatically increased with greater branching of R, while average particle sizes were relatively unaffected. Variation of X yielded a clear trend in average particle size, but without a discernible trend in reaction rate. Increased nucleophilicity (or size) of the halide resulted in smaller anatase nanocrystals. Average sizes ranged from 9.2 nm for TiF4 to 3.8 nm for TiI4. The amount of passivating agent (TOPO) influenced the chemistry. Reaction in pure TOPO was slower and resulted in smaller particles, while reactions without TOPO were much quicker and yielded mixtures of brookite, rutile, and anatase with average particle sizes greater than 10 nm. Figure 8 shows typical TEM images of TiO2 nanocrystals developed by Trentler et al.119 In the method used by Niederberger and Stucky,111 TiCl4 was slowly added to anhydrous benzyl alcohol under vigorous stirring at room temperature and was kept at 40150 °C for 1-21 days in the reaction vessel. The precipitate was calcinated at 450 °C for 5 h after thoroughly washing. The reaction between TiCl4 and benzyl alcohol was found suitable for the synthesis of highly crystalline anatase phase TiO2 nanoparticles with nearly uniform size and shape at very low temperatures, such as 40 °C. The particle size could be selectively adjusted in the range of 4-8 nm with the appropriate thermal conditions and a proper choice of the relative amounts of benzyl alcohol and titanium tetrachloride. The particle growth depended strongly on temperature, and lowering the titanium tetrachloride concentration led to a considerable decrease of particle size.111 Surfactants have been widely used in the preparation of a variety of nanoparticles with good size distribution and dispersity.15,16 Adding different surfactants as capping agents, such as acetic acid and acetylacetone, into the reaction matrix

Figure 9. TEM of TiO2 nanorods. The inset shows a HRTEM of a TiO2 nanorod. Reprinted with permission from Cozzoli, P. D.; Kornowski, A.; Weller, H. J. Am. Chem. Soc. 2003, 125, 14539. Copyright 2003 American Chemical Society.

can help synthesize monodispersed TiO2 nanoparticles.120,121 For example, Scolan and Sanchez found that monodisperse nonaggregated TiO2 nanoparticles in the 1-5 nm range were obtained through hydrolysis of titanium butoxide in the presence of acetylacetone and p-toluenesulfonic acid at 60 °C.120 The resulting nanoparticle xerosols could be dispersed in water-alcohol or alcohol solutions at concentrations higher than 1 M without aggregation, which is attributed to the complexation of the surface by acetylacetonato ligands and through an adsorbed hybrid organic-inorganic layer made with acetylacetone, p-toluenesulfonic acid, and water.120 With the aid of surfactants, different sized and shaped TiO2 nanorods can be synthesized.122-130 For example, the growth of high-aspect-ratio anatase TiO2 nanorods has been reported by Cozzoli and co-workers by controlling the hydrolysis process of TTIP in oleic acid (OA).122-126,130 Typically, TTIP was added into dried OA at 80-100 °C under inert gas protection (nitrogen flow) and stirred for 5 min. A 0.1-2 M aqueous base solution was then rapidly injected and kept at 80-100 °C for 6-12 h with stirring. The bases employed included organic amines, such as trimethylamino-N-oxide, trimethylamine, tetramethylammonium hydroxide, tetrabutylammonium hydroxyde, triethylamine, and tributylamine. In this reaction, by chemical modification of the titanium precursor with the carboxylic acid, the hydrolysis rate of titanium alkoxide was controlled. Fast (in 4-6 h) crystallization in mild conditions was promoted with the use of suitable catalysts (tertiary amines or quaternary ammonium hydroxides). A kinetically overdriven growth mechanism led to the growth of TiO2 nanorods instead of nanoparticles.123 Typical TEM images of the TiO2 nanorods are shown in Figure 9.123 Recently, Joo et al.127 and Zhang et al.129 reported similar procedures in obtaining TiO2 nanorods without the use of catalyst. Briefly, a mixture of TTIP and OA was used to generate OA complexes of titanium at 80 °C in 1-octadecene.


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Figure 10. TEM images of TiO2 nanorods with lengths of (A) 12 nm, (B) 30 nm, and (C) 16 nm. (D) 2.3 nm TiO2 nanoparticles. Inset in parts C and D: HR-TEM image of a single TiO2 nanorod and nanoparticle. Reprinted with permission from Zhang, Z.; Zhong, X.; Liu, S.; Li, D.; Han, M. Angew. Chem., Int. Ed. 2005, 44, 3466. Copyright 2005 Wiley-VCH.

The injection of a predetermined amount of oleylamine at 260 °C led to various sized TiO2 nanorods.129 Figure 10 shows TEM images of TiO2 nanorods with various lengths, and 2.3 nm TiO2 nanoparticles prepared with this method.129 In the surfactant-mediated shape evolution of TiO2 nanocrystals in nonaqueous media conducted by Jun et al.,128 it was found that the shape of TiO2 nanocrystals could be modified by changing the surfactant concentration. The synthesis was accomplished by an alkyl halide elimination reaction between titanium chloride and titanium isopropoxide. Briefly, a dioctyl ether solution containing TOPO and lauric acid was heated to 300 °C followed by addition of titanium chloride under vigorous stirring. The reaction was initiated by the rapid injection of TTIP and quenched with cold toluene. At low lauric acid concentrations, bulletand diamond-shaped nanocrystals were obtained; at higher concentrations, rod-shaped nanocrystals or a mixture of nanorods and branched nanorods was observed. The bulletand diamond-shaped nanocrystals and nanorods were elongated along the [001] directions. The TiO2 nanorods were found to simultaneously convert to small nanoparticles as a function of the growth time, as shown in Figure 11, due to the minimization of the overall surface energy via dissolution and regrowth of monomers during an Ostwald ripening.

2.4. Hydrothermal Method Hydrothermal synthesis is normally conducted in steel pressure vessels called autoclaves with or without Teflon

liners under controlled temperature and/or pressure with the reaction in aqueous solutions. The temperature can be elevated above the boiling point of water, reaching the pressure of vapor saturation. The temperature and the amount of solution added to the autoclave largely determine the internal pressure produced. It is a method that is widely used for the production of small particles in the ceramics industry. Many groups have used the hydrothermal method to prepare TiO2 nanoparticles.131-140 For example, TiO2 nanoparticles can be obtained by hydrothermal treatment of peptized precipitates of a titanium precursor with water.134 The precipitates were prepared by adding a 0.5 M isopropanol solution of titanium butoxide into deionized water ([H2O]/ [Ti] ) 150), and then they were peptized at 70 °C for 1 h in the presence of tetraalkylammonium hydroxides (peptizer). After filtration and treatment at 240 °C for 2 h, the as-obtained powders were washed with deionized water and absolute ethanol and then dried at 60 °C. Under the same concentration of peptizer, the particle size decreased with increasing alkyl chain length. The peptizers and their concentrations influenced the morphology of the particles. Typical TEM images of TiO2 nanoparticles made with the hydrothermal method are shown in Figure 12.134 In another example, TiO2 nanoparticles were prepared by hydrothermal reaction of titanium alkoxide in an acidic ethanol-water solution.132 Briefly, TTIP was added dropwise to a mixed ethanol and water solution at pH 0.7 with nitric acid, and reacted at 240 °C for 4 h. The TiO2 nanoparticles



Figure 12. TEM images of TiO2 nanoparticles prepared by the hydrothermal method. Reprinted from Yang, J.; Mei, S.; Ferreira, J. M. F. Mater. Sci. Eng. C 2001, 15, 183, Copyright 2001, with permission from Elsevier.

Figure 13. TEM image of TiO2 nanorods prepared with the hydrothermal method. Reprinted with permission from Zhang, Q.; Gao, L. Langmuir 2003, 19, 967. Copyright 2003 American Chemical Society.

Figure 11. Time dependent shape evolution of TiO2 nanorods: (a) 0.25 h; (b) 24 h; (c) 48 h. Scale bar ) 50 nm. Reprinted with permission from Jun, Y. W.; Casula, M. F.; Sim, J. H.; Kim, S. Y.; Cheon, J.; Alivisatos, A. P. J. Am. Chem. Soc. 2003, 125, 15981. Copyright 2003 American Chemical Society.

synthesized under this acidic ethanol-water environment were mainly primary structure in the anatase phase without secondary structure. The sizes of the particles were controlled to the range of 7-25 nm by adjusting the concentration of Ti precursor and the composition of the solvent system.

Besides TiO2 nanoparticles, TiO2 nanorods have also been synthesized with the hydrothermal method.141-146 Zhang et al. obtained TiO2 nanorods by treating a dilute TiCl4 solution at 333-423 K for 12 h in the presence of acid or inorganic salts.141,143-146 Figure 13 shows a typical TEM image of the TiO2 nanorods prepared with the hydrothermal method.141 The morphology of the resulting nanorods can be tuned with different surfactants146 or by changing the solvent compositions.145 A film of assembled TiO2 nanorods deposited on a glass wafer was reported by Feng et al.142 These TiO2 nanorods were prepared at 160 °C for 2 h by hydrothermal treatment of a titanium trichloride aqueous solution supersaturated with NaCl. TiO2 nanowires have also been successfully obtained with the hydrothermal method by various groups.147-151 Typically, TiO2 nanowires are obtained by treating TiO2 white powders in a 10-15 M NaOH aqueous solution at 150-200 °C for 24-72 h without stirring within an autoclave. Figure 14 shows the SEM images of TiO2 nanowires and a TEM image of a single nanowire prepared by Zhang and co-workers.150 TiO2 nanowires can also be prepared from layered titanate particles using the hydrothermal method as reported by Wei


Figure 14. SEM images of TiO2 nanowires with the inset showing a TEM image of a single TiO2 nanowire with a [010] selected area electron diffraction (SAED) recorded perpendicular to the long axis of the wire. Reprinted from Zhang, Y. X.; Li, G. H.; Jin, Y. X.; Zhang, Y.; Zhang, J.; Zhang, L. D. Chem. Phys. Lett. 2002, 365, 300, Copyright 2002, with permission from Elsevier.

et al.152 In their experiment, layer-structured Na2Ti3O7 was dispersed into a 0.05-0.1 M HCl solution and kept at 140170 °C for 3-7 days in an autoclave. TiO2 nanowires were obtained after the product was washed with H2O and finally dried. In the formation of a TiO2 nanowire from layered H2Ti3O7, there are three steps: (i) the exfoliation of layered Na2Ti3O7; (ii) the nanosheets formation; and (iii) the nanowires formation.152 In Na2Ti3O7, [TiO6] octahedral layers are held by the strong static interaction between the Na+ cations between the [TiO6] octahedral layers and the [TiO6] unit. When the larger H3+O cations replace the Na+ cations in the interlayer space of [TiO6] sheets, this static interaction is weakened because the interlayer distance is enlarged. As a result, the layered compounds Na2Ti3O7 are gradually exfoliated. When Na+ is exchanged by H+ in the dilute HCl solution, numerous H2Ti3O7 sheet-shaped products are formed. Since the nanosheet does not have inversion symmetry, an intrinsic tension exists. The nanosheets split to form nanowires in order to release the strong stress and lower the total energy.152 A representative TEM image of TiO2 nanowires from Na2Ti3O7 is shown in Figure 15.152 The hydrothermal method has been widely used to prepare TiO2 nanotubes since it was introduced by Kasuga et al. in 1998.153-175 Briefly, TiO2 powders are put into a 2.5-20 M NaOH aqueous solution and held at 20-110 °C for 20 h in an autoclave. TiO2 nanotubes are obtained after the products are washed with a dilute HCl aqueous solution and distilled water. They proposed the following formation process of TiO2 nanotubes.154 When the raw TiO2 material was treated with NaOH aqueous solution, some of the Ti-O-Ti bonds were broken and Ti-O-Na and Ti-OH bonds were formed. New Ti-O-Ti bonds were formed after the Ti-O-Na and Ti-OH bonds reacted with acid and water when the material was treated with an aqueous HCl solution and distilled water. The Ti-OH bond could form a sheet. Through the dehydration of Ti-OH bonds by HCl aqueous solution, Ti-O-Ti bonds or Ti-O-H-O-Ti hydrogen bonds were generated. The bond distance from one Ti to the next Ti on the surface decreased. This resulted in the folding of the sheets and the

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Figure 15. TEM images of TiO2 nanowires made from the layered Na2Ti3O7 particles, with the HRTEM image shown in the inset. Reprinted from Wei, M.; Konishi, Y.; Zhou, H.; Sugihara, H.; Arakawa, H. Chem. Phys. Lett. 2004, 400, 231, Copyright 2004, with permission from Elsevier.

Figure 16. TEM image of TiO2 nanotubes. Reprinted with permission from Kasuga, T.; Hiramatsu, M.; Hoson, A.; Sekino, T.; Niihara, K. Langmuir 1998, 14, 3160. Copyright 1998 American Chemical Society.

connection between the ends of the sheets, resulting in the formation of a tube structure. In this mechanism, the TiO2 nanotubes were formed in the stage of the acid treatment following the alkali treatment. Figure 16 shows typical TEM images of TiO2 nanotubes made by Kasuga et al.153 However, Du and co-workers found that the nanotubes were formed during the treatment of TiO2 in NaOH aqueous solution.161 A 3D f 2D f 1D formation mechanism of the TiO2 nanotubes was proposed by Wang and co-workers.171 It stated that the raw TiO2 was first transformed into lamellar structures and then bent and rolled to form the nanotubes. For the formation of the TiO2 nanotubes, the two-dimensional lamellar TiO2 was essential. Yao and co-workers further suggested, based on their HRTEM study as shown in Figure



Figure 18. TEM micrographs of TiO2 nanoparticles prepared with the solvothermal method. Reprinted with permission from Li, X. L.; Peng, Q.; Yi, J. X.; Wang, X.; Li, Y. D. Chem.sEur. J. 2006, 12, 2383. Copyright 2006 Wiley-VCH.

Figure 17. (a) HRTEM images of TiO2 nanotubes. (b) Crosssectional view of TiO2 nanotubes. Reused with permission from B. D. Yao, Y. F. Chan, X. Y. Zhang, W. F. Zhang, Z. Y. Yang, N. Wang, Applied Physics Letters 82, 281 (2003). Copyright 2003, American Institute of Physics.

17, that TiO2 nanotubes were formed by rolling up the singlelayer TiO2 sheets with a rolling-up vector of [001] and attracting other sheets to surround the tubes.172 Bavykin and co-workers suggested that the mechanism of nanotube formation involved the wrapping of multilayered nanosheets rather than scrolling or wrapping of single layer nanosheets followed by crystallization of successive layers.156 In the mechanism proposed by Wang et al., the formation of TiO2 nanotubes involved several steps.176 During the reaction with NaOH, the Ti-O-Ti bonding between the basic building blocks of the anatase phase, the octahedra, was broken and a zigzag structure was formed when the free octahedras shared edges between the Ti ions with the formation of hydroxy bridges, leading to the growth along the [100] direction of the anatase phase. Two-dimensional crystalline sheets formed from the lateral growth of the formation of oxo bridges between the Ti centers (Ti-O-Ti bonds) in the [001] direction and rolled up in order to saturate these dangling bonds from the surface and lower the total energy, resulting in the formation of TiO2 nanotubes.176

2.5. Solvothermal Method The solvothermal method is almost identical to the hydrothermal method except that the solvent used here is nonaqueous. However, the temperature can be elevated much higher than that in hydrothermal method, since a variety of organic solvents with high boiling points can be chosen. The solvothermal method normally has better control than hydrothermal methods of the size and shape distributions and the crystallinity of the TiO2 nanoparticles. The solvothermal method has been found to be a versatile method for the

synthesis of a variety of nanoparticles with narrow size distribution and dispersity.177-179 The solvothermal method has been employed to synthesize TiO2 nanoparticles and nanorods with/without the aid of surfactants.177-185 For example, in a typical procedure by Kim and co-workers,184 TTIP was mixed with toluene at the weight ratio of 1-3:10 and kept at 250 °C for 3 h. The average particle size of TiO2 powders tended to increase as the composition of TTIP in the solution increased in the range of weight ratio of 1-3: 10, while the pale crystalline phase of TiO2 was not produced at 1:20 and 2:5 weight ratios.184 By controlling the hydrolyzation reaction of Ti(OC4H9)4 and linoleic acid, redispersible TiO2 nanoparticles and nanorods could be synthesized, as found by Li et al. recently.177 The decomposition of NH4HCO3 could provide H2O for the hydrolyzation reaction, and linoleic acid could act as the solvent/reagent and coordination surfactant in the synthesis of nanoparticles. Triethylamine could act as a catalyst for the polycondensation of the TiO-Ti inorganic network to achieve a crystalline product and had little influence on the products’ morphology. The chain lengths of the carboxylic acids had a great influence on the formation of TiO2, and long-chain organic acids were important and necessary in the formation of TiO2.177 Figure 18 shows a representative TEM image of TiO2 nanoparticles from their study.177 TiO2 nanorods with narrow size distributions can also be developed with the solvothermal method.177,183 For example, in a typical synthesis from Kim et al., TTIP was dissolved in anhydrous toluene with OA as a surfactant and kept at 250 °C for 20 h in an autoclave without stirring.183 Long dumbbell-shaped nanorods were formed when a sufficient amount of TTIP or surfactant was added to the solution, due to the oriented growth of particles along the [001] axis. At a fixed precursor to surfactant weight ratio of 1:3, the concentration of rods in the nanoparticle assembly increased as the concentration of the titanium precursor in the solution increased. The average particle size was smaller and the size distribution was narrower than is the case for particles synthesized without surfactant. The crystalline phase, diameter, and length of these nanorods are largely influenced by the precursor/surfactant/solvent weight ratio. Anatase nano-


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Figure 20. TEM images of TiO2 nanowires synthesized by the solvothermal method. From: Wen, B.; Liu, C.; Liu, Y. New J. Chem. 2005, 29, 969 (http://dx.doi.org/10.1039/b502604k) s Reproduced by permission of The Royal Society of Chemistry (RSC) on behalf of the Centre National de la Recherche Scientifique (CNRS).

Figure 19. TEM micrographs and electron diffraction patterns of products prepared from solutions at the weight ratio of precursor/ solvent/surfactant ) 1:5:3. Reprinted from Kim, C. S.; Moon, B. K.; Park, J. H.; Choi, B. C.; Seo, H. J. J. Cryst. Growth 2003, 257, 309, Copyright 2003, with permission from Elsevier.

rods were obtained from the solution with a precursor/ surfactant weight ratio of more than 1:3 for a precursor/ solvent weight ratio of 1:10 or from the solution with a precursor/solvent weight ratio of more than 1:5 for a precursor/surfactant weight ratio of 1:3. The diameter and length of these nanorods were in the ranges of 3-5 nm and 18-25 nm, respectively. Figure 19 shows a typical TEM image of TiO2 nanorods prepared from the solutions with the weight ratio of precursor/solvent/surfactant ) 1:5:3.183 Similar to the hydrothermal method, the solvothermal method has also been used for the preparation of TiO2 nanowires.180-182 Typically, a TiO2 powder suspension in an 5 M NaOH water-ethanol solution is kept in an autoclave at 170-200 °C for 24 h and then cooled to room temperature naturally. TiO2 nanowires are obtained after the obtained sample is washed with a dilute HCl aqueous solution and dried at 60 °C for 12 h in air.181 The solvent plays an important role in determining the crystal morphology. Solvents with different physical and chemical properties can influence the solubility, reactivity, and diffusion behavior of the reactants; in particular, the polarity and coordinating ability of the solvent can influence the morphology and the crystallization behavior of the final products. The presence of ethanol at a high concentration not only can cause the polarity of the solvent to change but also strongly affects the ζ potential values of the reactant particles and the increases solution viscosity. For example, in the absence of ethanol, short and wide flakelike structures of TiO2 were obtained instead of nanowires. When chloroform is used, TiO2 nanorods were obtained.181 Figure 20 shows representative TEM images of the TiO2 nanowires prepared from the solvothermal method.181 Alternatively, bamboo-shaped Agdoped TiO2 nanowires were developed with titanium butoxide as precursor and AgNO3 as catalyst.180 Through the electron diffraction (ED) pattern and HRTEM study, the Ag

Figure 21. SEM morphology of TiO2 nanorods by directly oxidizing a Ti plate with a H2O2 solution. Reprinted from Wu, J. M. J. Cryst. Growth 2004, 269, 347, Copyright 2004, with permission from Elesevier.

phase only existed in heterojunctions between single-crystal TiO2 nanowires.180

2.6. Direct Oxidation Method TiO2 nanomaterials can be obtained by oxidation of titanium metal using oxidants or under anodization. Crystalline TiO2 nanorods have been obtained by direct oxidation of a titanium metal plate with hydrogen peroxide.186-191 Typically, TiO2 nanorods on a Ti plate are obtained when a cleaned Ti plate is put in 50 mL of a 30 wt % H2O2 solution at 353 K for 72 h. The formation of crystalline TiO2 occurs through a dissolution precipitation mechanism. By the addition of inorganic salts of NaX (X ) F-, Cl-, and SO42-), the crystalline phase of TiO2 nanorods can be controlled. The addition of F- and SO42- helps the formation of pure anatase, while the addition of Cl- favors the formation of rutile.189 Figure 21 shows a typical SEM image of TiO2 nanorods prepared with this method.186 At high temperature, acetone can be used as a good oxygen source and for the preparation of TiO2 nanorods by oxidizing


Figure 22. SEM images of large-scale nanorod arrays prepared by oxidizing a titanium with acetone at 850 °C for 90 min. From: Peng, X.; Chen, A. J. Mater. Chem. 2004, 14, 2542 (http:// dx.doi.org/10.1039/b404750h) s Reproduced by permission of The Royal Society of Chemistry.

a Ti plate with acetone as reported by Peng and Chen.192 The oxygen source was found to play an important role. Highly dense and well-aligned TiO2 nanorod arrays were formed when acetone was used as the oxygen source, and only crystal grain films or grains with random nanofibers growing from the edges were obtained with pure oxygen or argon mixed with oxygen. The competition of the oxygen and titanium diffusion involved in the titanium oxidation process largely controlled the morphology of the TiO2. With pure oxygen, the oxidation occurred at the Ti metal and the TiO2 interface, since oxygen diffusion predominated because of the high oxygen concentration. When acetone was used as the oxygen source, Ti cations diffused to the oxide surface and reacted with the adsorbed acetone species. Figure 22 shows aligned TiO2 nanorod arrays obtained by oxidizing a titanium substrate with acetone at 850 °C for 90 min.192 As extensively studied, TiO2 nanotubes can be obtained by anodic oxidation of titanium foil.193-228 In a typical experiment, a clean Ti plate is anodized in a 0.5% HF solution under 10-20 V for 10-30 min. Platinum is used as counterelectrode. Crystallized TiO2 nanotubes are obtained after the anodized Ti plate is annealed at 500 °C for 6 h in oxygen.210 The length and diameter of the TiO2 nanotubes could be controlled over a wide range (diameter, 15-120 nm; length, 20 nm to 10 µm) with the applied potential between 1 and 25 V in optimized phosphate/HF electrolytes.229 Figure 23 shows SEM images of TiO2 nanotubes created with this method.208

2.7. Chemical Vapor Deposition Vapor deposition refers to any process in which materials in a vapor state are condensed to form a solid-phase material. These processes are normally used to form coatings to alter the mechanical, electrical, thermal, optical, corrosion resistance, and wear resistance properties of various substrates. They are also used to form free-standing bodies, films, and fibers and to infiltrate fabric to form composite materials. Recently, they have been widely explored to fabricate various nanomaterials. Vapor deposition processes usually take place within a vacuum chamber. If no chemical reaction occurs, this process is called physical vapor deposition (PVD);


Figure 23. SEM images of TiO2 nanotubes prepared with anodic oxidation. Reprinted with permission from Varghese, O. K.; Gong, D.; Paulose, M.; Ong, K. G.; Dickey, E. C.; Grimes, C. A. AdV. Mater. 2003, 15, 624. Copyright 2003 Wiley-VCH.

Figure 24. SEM images of TiO2 nanorods grown at 560 °C. Reprinted with permission from Wu, J. J.; Yu, C. C. J. Phys. Chem. B 2004, 108, 3377. Copyright 2004 American Chemical Society.

otherwise, it is called chemical vapor deposition (CVD). In CVD processes, thermal energy heats the gases in the coating chamber and drives the deposition reaction. Thick crystalline TiO2 films with grain sizes below 30 nm as well as TiO2 nanoparticles with sizes below 10 nm can be prepared by pyrolysis of TTIP in a mixed helium/oxygen atmosphere, using liquid precursor delivery.230 When deposited on the cold areas of the reactor at temperatures below 90 °C with plasma enhanced CVD, amorphous TiO2 nanoparticles can be obtained and crystallize with a relatively high surface area after being annealed at high temperatures.231 TiO2 nanorod arrays with a diameter of about 50-100 nm and a length of 0.5-2 µm can be synthesized by metal organic CVD (MOCVD) on a WC-Co substrate using TTIP as the precursor.232 Figure 24 shows the TiO2 nanorods grown on fused silica substrates with a template- and catalyst-free MOCVD method.233 In a typical procedure, titanium acetylacetonate (Ti(C10H14O5)) vaporizing in the low-temperature zone of a furnace at 200-230 °C is carried by a N2/O2 flow into the high-temperature zone of 500-700 °C, and TiO2 nanostructures are grown directly on the substrates. The phase and


Figure 25. SEM images of the TiO2 nanowire arrays prepared by the PVD method. Reprinted from Wu, J. M.; Shih, H. C.; Wu, W. T. Chem. Phys. Lett. 2005, 413, 490, Copyright 2005, with permission from Elsevier.

morphology of the TiO2 nanostructures can be tuned with the reaction conditions. For example, at 630 and 560 °C under a pressure of 5 Torr, single-crystalline rutile and anatase TiO2 nanorods were formed respectively, while, at 535 °C under 3.6 Torr, anatase TiO2 nanowalls composed of well-aligned nanorods were formed.233 In addition to the above CVD approaches in preparing TiO2 nanomaterials, other CVD approaches are also used, such as electrostatic spray hydrolysis,234 diffusion flame pyrolysis,235-239 thermal plasma pyrolysis,240-246 ultrasonic spray pyrolysis,247 laser-induced pyrolysis,248,249 and ultronsicassisted hydrolysis,250,251 among others.

2.8. Physical Vapor Deposition In PVD, materials are first evaporated and then condensed to form a solid material. The primary PVD methods include thermal deposition, ion plating, ion implantation, sputtering, laser vaporization, and laser surface alloying. TiO2 nanowire arrays have been fabricated by a simple PVD method or thermal deposition.252-254 Typically, pure Ti metal powder is on a quartz boat in a tube furnace about 0.5 mm away from the substrate. Then the furnace chamber is pumped down to ∼300 Torr and the temperature is increased to 850 °C under an argon gas flow with a rate of 100 sccm and held for 3 h. After the reaction, a layer of TiO2 nanowires can be obtained.254 A layer of Ti nanopowders can be deposited on the substrate before the growth of TiO2 nanowires,252,253 and Au can be employed as catalyst.252 A typical SEM image of TiO2 nanowires made with the PVD method is shown in Figure 25.252

2.9. Electrodeposition Electrodeposition is commonly employed to produce a coating, usually metallic, on a surface by the action of reduction at the cathode. The substrate to be coated is used as cathode and immersed into a solution which contains a salt of the metal to be deposited. The metallic ions are attracted to the cathode and reduced to metallic form. With the use of the template of an AAM, TiO2 nanowires can be obtained by electrodeposition.255,256 In a typical process, the electrodeposition is carried out in 0.2 M TiCl3 solution with

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Figure 26. Cross-sectional SEM image of TiO2 nanowires electrodeposited in AAM pores. Reprinted from Liu, S.; Huang, K. Sol. Energy Mater. Sol. Cells 2004, 85, 125, Copyright 2004, with permission from Elsevier.

pH ) 2 with a pulsed electrodeposition approach, and titanium and/or its compound are deposited into the pores of the AAM. By heating the above deposited template at 500 °C for 4 h and removing the template, pure anatase TiO2 nanowires can be obtained. Figure 26 shows a representative SEM image of TiO2 nanowires.256

2.10. Sonochemical Method Ultrasound has been very useful in the synthesis of a wide range of nanostructured materials, including high-surfacearea transition metals, alloys, carbides, oxides, and colloids. The chemical effects of ultrasound do not come from a direct interaction with molecular species. Instead, sonochemistry arises from acoustic cavitation: the formation, growth, and implosive collapse of bubbles in a liquid. Cavitational collapse produces intense local heating (∼5000 K), high pressures (∼1000 atm), and enormous heating and cooling rates (>109 K/s). The sonochemical method has been applied to prepare various TiO2 nanomaterials by different groups.257-269 Yu et al. applied the sonochemical method in preparing highly photoactive TiO2 nanoparticle photocatalysts with anatase and brookite phases using the hydrolysis of titanium tetraisoproproxide in pure water or in a 1:1 EtOH-H2O solution under ultrasonic radiation.109 Huang et al. found that anatase and rutile TiO2 nanoparticles as well as their mixtures could be selectively synthesized with various precursors using ultrasound irradiation, depending on the reaction temperature and the precursor used.259 Zhu et al. developed titania whiskers and nanotubes with the assistance of sonication as shown in Figure 27.269 They found that arrays of TiO2 nanowhiskers with a diameter of 5 nm and nanotubes with a diameter of ∼5 nm and a length of 200-300 nm could be obtained by sonicating TiO2 particles in NaOH aqueous solution followed by washing with deionized water and a dilute HNO3 aqueous solution.

2.11. Microwave Method A dielectric material can be processed with energy in the form of high-frequency electromagnetic waves. The principal


Figure 27. TEM images of TiO2 nanotubes (A) and nanowhiskers (B) prepared with the sonochemical method. From: Zhu, Y.; Li, H.; Koltypin, Y.; Hacohen, Y. R.; Gedanken, A. Chem. Commun. 2001, 2616 (http://dx.doi.org/10.1039/b108968b) s Reproduced by permission of The Royal Society of Chemistry.

frequencies of microwave heating are between 900 and 2450 MHz. At lower microwave frequencies, conductive currents flowing within the material due to the movement of ionic constituents can transfer energy from the microwave field to the material. At higher frequencies, the energy absorption is primarily due to molecules with a permanent dipole which tend to reorientate under the influence of a microwave electric field. This reorientation loss mechanism originates from the inability of the polarization to follow extremely rapid reversals of the electric field, so the polarization phasor lags the applied electric field. This ensures that the resulting current density has a component in phase with the field, and therefore power is dissipated in the dielectric material. The major advantages of using microwaves for industrial processing are rapid heat transfer, and volumetric and selective heating. Microwave radiation is applied to prepare various TiO2 nanomaterials.270-276 Corradi et al. found that colloidal titania nanoparticle suspensions could be prepared within 5 min to 1 h with microwave radiation, while 1 to 32 h was needed for the conventional synthesis method of forced hydrolysis at 195 °C.270 Ma et al. developed high-quality rutile TiO2 nanorods with a microwave hydrothermal method and found that they aggregated radially into spherical secondary nanoparticles.272 Wu et al. synthesized TiO2 nanotubes by microwave radiation via the reaction of TiO2 crystals of anatase, rutile, or mixed phase and NaOH aqueous solution under a certain microwave power.275 Normally, the TiO2 nanotubes had the central hollow, open-ended, and multiwall structure with diameters of 8-12 nm and lengths up to 200-1000 nm.275

2.12. TiO2 Mesoporous/Nanoporous Materials In the past decade, mesoporous/nanoporous TiO2 materials have been well studied with or without the use of organic


Figure 28. SEM image of the mesoporous TiO2 film synthesized from the acetic acid-modified precursor and autoclaved at 230 °C. Reprinted with permission from Barbe, C. J.; Arendse, F.; Comte, P.; Jirousek, M.; Lenzmann, F.; Shklover, V.; Gra¨tzel, M. J. Am. Ceram. Soc. 1997, 80, 3157. Copyright 1997 Blackwell Publishing.

surfactant templates.28,80,264,265,277-312 Barbe et al. reported the preparation of a mesoporous TiO2 film by the hydrothermal method as shown Figure 28.80 In a typical experiment, TTIP was added dropwise to a 0.1 M nitric acid solution under vigorous stirring and at room temperature. A white precipitate formed instantaneously. Immediately after the hydrolysis, the solution was heated to 80 °C and stirred vigorously for 8 h for peptization. The solution was then filtered on a glass frit to remove agglomerates. Water was added to the filtrate to adjust the final solids concentration to ∼5 wt %. The solution was put in a titanium autoclave for 12 h at 200-250 °C. After sonication, the colloidal suspension was put in a rotary evaporator and evaporated to a final TiO2 concentration of 11 wt %. The precipitation pH, hydrolysis rate, autoclaving pH, and precursor chemistry were found to influence the morphology of the final TiO2 nanoparticles. Alternative procedures without the use of hydrothermal processes have been reported by Liu et al.292 and Zhang et al.311 In the report by Liu et al., 24.0 g of titanium(IV) n-butoxide ethanol solution (weight ratio of 1:7) was prehydrolyzed in the presence of 0.32 mL of a 0.28 M HNO3 aqueous solution (TBT/HNO3 ∼ 100:1) at room temperature for 3 h. 0.32 mL of deionized water was added to the prehydrolyzed solution under vigorous stirring and stirred for an additional 2 h. The sol solution in a closed vessel was kept at room temperature without stirring to gel and age. After aging for 14 days, the gel was dried at room temperature, ground into a fine powder, washed thoroughly with water and ethanol, and dried to produce porous TiO2. Upon calcination at 450 °C for 4 h under air, crystallized mesoporous TiO2 material was obtained.292 Yu et al. prepared three-dimensional and thermally stable mesoporous TiO2 without the use of any surfactants.265 Briefly, monodispersed TiO2 nanoparticles were formed initially by ultrasound-assisted hydrolysis of acetic acidmodified titanium isopropoxide. Mesoporous spherical or globular particles were then produced by controlled conden-


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sation and agglomeration of these sol nanoparticles under high-intensity ultrasound radiation. The mesoporous TiO2 had a wormhole-like structure consisting of TiO2 nanoparticles and a lack of long-range order.265 In the template method used by the Stucky group278-280,287,295,302,306-307,313 and other groups,264,293,297,303,309 structure-directing agents were used for organizing networkforming metal oxide species in nonaqueous solutions. These structure-directing agents were also called organic templates. The most commonly used organic templates were amphiphilic poly(alkylene oxide) block copolymers, such as HO(CH2CH2O)20(CH2CH(CH3)O)70(CH2CH2O)20H (designated EO20PO70EO20, called Pluronic P-123) and HO(CH2CH2O)106(CH2CH(CH3)O)70(CH2CH2O)106H (designated EO106PO70EO106, called Pluronic F-127). In a typical synthesis, poly(alkylene oxide) block copolymer was dissolved in ethanol. Then TiCl4 precursor was added with vigorous stirring. The resulting sol solution was gelled in an open Petri dish at 40 °C in air for 1-7 days. Mesoporous TiO2 was obtained after removing the surfactant species by calcining the as-made sample at 400 °C for 5 h in air.306 Figure 29 shows typical TEM images of the mesoporous TiO2. Besides triblock copolymers as structure-directing agents, diblock polymers were also used such as [CnH2n-1(OCH2CH2)yOH, Brij 56 (B56, n/y ) 16/10) or Brij 58 (B58, n/y ) 16/20)] by Sanchez et al.285 Other surfactants employed to direct the formation of mesoporous TiO2 include tetradecyl phosphate (a 14-carbon chain) by Antonelli and Ying277 and commercially available dodecyl phosphate by Putnam and co-workers,298 cetyltrimethylammonium bromide (CTAB) (a cationic surfactant),281,283,296 the recent Gemini surfactant,294 and dodecylamine (a neutral surfactant).304 Carbon nanotubes310 and mesoporous SBA-15286 have also been used as the skeleton for mesoporous TiO2.

2.13. TiO2 Aerogels The study of TiO2 aerogels is worthy of special mention.314-326 The combination of sol-gel processing with supercritical drying offers the synthesis of TiO2 aerogels with morphological and chemical properties that are not easily achieved by other preparation methods, i.e., with high surface area. Campbell et al. prepared TiO2 aerogels by sol-gel synthesis from titanium n-butoxide in methanol with the subsequent removal of solvent by supercritical CO2.315 For a typical synthesis process, titanium n-butoxide was added to 40 mL of methanol in a dry glovebox. This solution was combined with another solution containing 10 mL of methanol, nitric acid, and deionized water. The concentration of the titanium n-butoxide was kept at 0.625 M, and the molar ratio of water/HNO3/titanium n-butoxide was 4:0.1: 1. The gel was allowed to age for 2 h and then extracted in a standard autoclave with supercritical CO2 at a flow rate of 24.6 L/h, at 343 K under 2.07 × 107 Pa for 2-3 h, resulting in complete removal of solvent. After extraction, the sample was heated in a vacuum oven at 3.4 kPa and 383 K for 3 h to remove the residual solvent and at 3.4 kPa and 483 K for 3 h to remove any residual organics. The pretreated sample had a brown color and turned white after calcination at 773 K or above. The resulting TiO2 aerogel, after calcination at 773 K for 2 h, had a BET surface area of >200 m2/g, contained mesopores in the range 2-10 nm, and was of the pure anatase form. Dagan et al. found the TiO2 aerogels obtanied by using a Ti/ethanol/H2O/nitric acid ratio of 1:20: 3:0.08 could have a porosity of 90% and surface areas of

Figure 29. TEM micrographs of two-dimensional hexagonal mesoporous TiO2 recorded along the (a) [110] and (b) [001] zone axes, respectively. The inset in part a is selected-area electron diffraction patterns obtained on the image area. (c) TEM image of cubic mesoporous TiO2 accompanied by the corresponding (inset) EDX spectrum. Reprinted with permission from Yang, P.; Zhao, D.; Margolese, D. I.; Chmelka, B. F.; Stucky, G. D. Chem. Mater. 1999, 11, 2813. Copyright 1999 American Chemical Society.

600 m2/g, as compared to a surface area of 50 m2/g for TiO2 P25.316,317 Figure 30 shows a typical SEM image of a TiO2 aerogel with a surface area of 447 m2/g and an interpore structure constructed by near uniform grains of elliptical shapes with 30 nm × 50 nm axes.326



Figure 30. SEM image of a TiO2 aerogel. Reprinted with permission from Zhu, Z.; Tsung, L. Y.; Tomkiewicz, M. J. Phys. Chem. 1995, 99, 15945. Copyright 1995 American Chemical Society.

2.14. TiO2 Opal and Photonic Materials The syntheses of TiO2 opal and photonic materials have been well studied by various groups.327-358 Holland et al. reported the preparation of TiO2 inverse opal from the corresponding metal alkoxides, using latex spheres as templates.334,335 Millimeter-thick layers of latex spheres were deposited on filter paper in a Buchner funnel under vacuum and soaked with ethanol. Titanium ethoxide was added dropwise to cover the latex spheres completely while suction was applied. Typical mass ratios of alkoxide to latex were between 1.4 and 3. After drying the composite in a vacuum desiccator for 3 to 24 h, the latex spheres were removed by calcination in flowing air at 575 °C for 7 to 12 h, leaving hard and brittle powder particles with 320- to 360-nm voids. The carbon content of the calcined samples varied from 0.4 to 1.0 wt %, indicating that most of the latex templates had been removed from the 3D host. Figure 31 shows an illustration of the simple synthesis of TiO2 inverse opal and an SEM image of TiO2 inverse opals. Similar studies have also been carried out by other researchers.327,356 Dong and Marlow prepared TiO2 inversed opals with a skeleton-like structure of TiO2 rods by a template-directed method using monodispersed polystyrene particles of size 270 nm.328-330,345 Infiltration of a titania precursor (Ti(i-OPr)4 in EtOH) was followed by a drying and calcination procedure. The precursor concentration was varied from 30% to 100%, and the calcination temperature was tuned from 300 to 700 °C. A SEM picture of the TiO2 inversed opal is shown in Figure 32.329 The skeleton structure consists of rhombohedral windows and TiO2 cylinders forming a highly regular network. The cylinders connect the centers of the former octahedral and tetrahedral voids of the opal. These voids form a CaF2 lattice which is filled with cylindrical bonds connecting the Ca and F sites. Wang et al. reported their study on the large-scale fabrication of ordered TiO2 nanobowl arrays.354 The process starts with a self-assembled monolayer of polystyrene (PS) spheres, which is used as a template for atomic layer deposition of a TiO2 layer. After ion-milling, toluene-etching, and annealing of the TiO2-coated spheres, ordered arrays of nanostructured TiO2 nanobowls can be fabricated as shown in Figure 33. Wang et al. fabricated a 2D photonic crystal by coating patterned and aligned ZnO nanorod arrays with TiO2.355 PS spheres were self-assembled to make a monolayer mask on

Figure 31. (A) Schematic illustration of the synthesis of a TiO2 inversed opal. (B) SEM image of the TiO2 inversed opal. Reprinted with permission from Holland, B. T.; Blanford, C.; Stein, A. Science 1998, 281, 538 (http://www.sciencemag.org). Copyright 1998 AAAS.

a sapphire substrate, which was then covered with a layer of gold. After removing the PS spheres with toluene, ZnO nanorods were grown using a vapor-liquid-solid process. Finally, a TiO2 layer was deposited on the ZnO nanorods by introducing TiCl4 and water vapors into the atomic layer deposition chamber at 100 °C. Figure 34 shows SEM images of a ZnO nanorod array and the TiO2-coated ZnO nanorod array. Li et al. reported the preparation of ordered arrays of TiO2 opals using opal gel templates under uniaxial compression at ambient temperature during the TiO2 sol/gel process.337 The aspect ratio was controllable by the compression degree, R. Polystyrene inverse opal was template synthesized using silica opals as template. The silica was removed with 40 wt % aqueous hydrofluoric acid. Monomer solutions consisting of dimethylacrylamide, acrylic acid, and methylenebisacrylamide in 1:1:0.02 weight ratios were dissolved in a water/


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Figure 32. SEM picture of a TiO2 skeleton with a cylinder radius of about 0.06a. a is the lattice constant of the cubic unit cell. Reprinted from Dong, W.; Marlow, F. Physica E 2003, 17, 431, Copyright 2003, with permission from Elsevier. Figure 34. (A) SEM images of short and densely aligned ZnO nanorod array on a sapphire substrate. Inset: An optical image of the aligned ZnO nanorods over a large area. (B) SEM image of the TiO2-coated ZnO nanorod array. Reprinted with permission from Wang, X.; Neff, C.; Graugnard, E.; Ding, Y.; King, J. S.; Pranger, L. A.; Tannenbaum, R.; Wang, Z. L.; Summers, C. J. AdV. Mater. 2005, 17, 2103. Copyright 2005 Wiley-VCH.

gels with correspondingly different properties can be produced. Water was completely removed from the opal hydrogel by repeatedly rinsing it with a large amount of ethanol. Afterward, the opal gel was put into a large amount of tetrabutyl titanate (TBT) at ambient temperature for 24 h. The TBT-swollen opal gel was then immersed in a water/ ethanol (1:1 wt/wt) mixture for 5 h to let the TiO2 sol/gel process proceed. Figure 35A shows the opal structure of the gel/titania composite spheres formed. After calcination, TiO2 opal with distinctive spherical contours could be found. The compression degree, R, was adjusted by the spacer height when the substrates were compressed. When the substrates were slightly compressed against each other to the extent of producing a 20% reduction in the thickness of the composition opal, the deformation of the template-synthesized titania spheres was not substantial (Figure 35B). When the compression degree was increased to the point of reaching 35% deformation in the opal gel, noticeably deformed titania opals could be obtained (Figure 35C and D). Figure 33. (A) Experimental procedure for fabricating TiO2 nanobowl arrays. (B) Low- and high- (inset) magnification SEM image of TiO2 nanobowl arrays. Reprinted with permission from Wang, X. D.; Graugnard, E.; King, J. S.; Wang, Z. L.; Summers, C. J. Nano Lett. 2004, 4, 2223. Copyright 2004 American Chemical Society.

ethanol mixture (4:7 wt/wt) with total monomer content 30 wt %. Ethanol was used to facilitate diffusion of the monomer solution into the inverse opal polystyrene. After the inverse opal was infiltrated by the monomer solution containing 1 wt % of the initiator AIBN and a subsequent free radical polymerization at 60 °C for 3 h, a solid composite resulted. The initial inverse opal polystyrene template was then removed with chloroform in a Soxhlet extractor for 12 h, whereupon the opal gel was formed. By using different compositions of the monomer solution, hole sizes, and stacking structures of the starting inverse opal templates, opal

2.15. Preparation of TiO2 Nanosheets The preparation of TiO2 nanosheets has also been explored recently.359-368 Typically, TiO2 nanosheets were synthesized by delaminating layered protonic titanate into colloidal single layers. A stoichiometric mixture of Cs2CO3 and TiO2 was calcined at 800 °C for 20 h to produce a precursor, cesium titanate, Cs0.7Ti1.82500.175O4 (0: vacancy), about 70 g of which was treated with 2 L of a 1 M HCl solution at room temperature. This acid leaching was repeated three times by renewing the acid solution every 24 h. The resulting acidexchanged product was filtered, washed with water, and airdried. The obtained protonic titanate, H0.7Ti1.82500.175O4‚H2O, was shaken vigorously with a 0.017 M tetrabutylammonium hydroxide solution at ambient temperature for 10 days. The solution-to-solid ratio was adjusted to 250 cm3 g-1. This procedure yielded a stable colloidal suspension with an



Figure 35. SEM of the TiO2 opals. (A) A gel/titania composite opal fabricated without compressing the opal gel template during the sol/gel process. (Inset) Image of the sample after calcination at 450 °C for 3 h. (B-D) (Main panel) Oblate titania opal materials after calcination at 450 °C for 3 h, subject to compression degree R of (B) 20%, (C) 35%, and (D) 50%. The images were taken for the fractured surfaces containing the direction of applied compression. (Inset) Image of the same sample, but with the fracture surface perpendicular to the direction of applied compression. From: Ji, L.; Rong, J.; Yang, Z. Chem. Commun. 2003, 1080 (http://dx.doi.org/10.1039/b300825h) s Reproduced by permission of The Royal Society of Chemistry.

opalescent appearance. Figure 36 shows TEM and AFM images of TiO2 nanosheets with thicknesses of 1.2-1.3 nm, which is the height of the TiO2 nanosheet with a monolayer of water molecules on both sides (0.70 + 0.25 × 2) thick.366

3. Properties of TiO2 Nanomaterials 3.1. Structural Properties of TiO2 Nanomaterials Figure 37 shows the unit cell structures of the rutile and anatase TiO2.11 These two structures can be described in terms of chains of TiO6 octahedra, where each Ti4+ ion is surrounded by an octahedron of six O2- ions. The two crystal structures differ in the distortion of each octahedron and by the assembly pattern of the octahedra chains. In rutile, the

octahedron shows a slight orthorhombic distortion; in anatase, the octahedron is significantly distorted so that its symmetry is lower than orthorhombic. The Ti-Ti distances in anatase are larger, whereas the Ti-O distances are shorter than those in rutile. In the rutile structure, each octahedron is in contact with 10 neighbor octahedrons (two sharing edge oxygen pairs and eight sharing corner oxygen atoms), while, in the anatase structure, each octahedron is in contact with eight neighbors (four sharing an edge and four sharing a corner). These differences in lattice structures cause different mass densities and electronic band structures between the two forms of TiO2. Hamad et al. performed a theoretical calculation on TinO2n clusters (n ) 1-15) with a combination of simulated


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Figure 36. (A) TEM of Ti1-δO24δ- nanosheets. (B and C) AFM image and height scan of the TiO2 nanosheets deposited on a Si wafer. (D) Structural model for a hydrated TiO2 nanosheet. Closed, open, and shaded circles represent Ti atom, O atom, and H2O molecules, respectively. All the water sites are assumed to be half occupied. Reprinted with permission from Sasaki, T.; Ebina, Y.; Kitami, Y.; Watanabe, M.; Oikawa, T. J. Phys. Chem. B 2001, 105, 6116. Copyright 2001 American Chemical Society.

Figure 37. Lattice structure of rutile and anatase TiO2. Reprinted with permission from Linsebigler, A. L.; Lu, G.; Yates, J. T., Jr. Chem. ReV. 1995, 95, 735. Copyright 1995 American Chemical Society.

annealing, Monte Carlo basin hopping simulation, and genetic algorithms methods.369 They found that the calculated global minima consisted of compact structures, with titanium atoms reaching high coordination rapidly as n increased. For n g 11, the particles had at least a central octahedron surrounded by a shell of surface tetrahedra, trigonal bipyramids, and square base pyramids. Swamy et al. found the metastability of anatase as a function of pressure was size dependent, with smaller crystallites preserving the structure to higher pressures.370 Three size regimes were recognized for the pressure-induced phase transition of anatase at room temperature: an anatase-

amorphous transition regime at the smallest crystallite sizes, an anatase-baddeleyite transition regime at intermediate crystallite sizes, and an anatase-R-PbO2 transition regime comprising large nanocrystals to macroscopic single crystals. Barnard et al. performed a series of theoretical studies on the phase stability of TiO2 nanoparticles in different environments by a thermodynamic model.371-375 They found that surface passivation had an important impact on nanocrystal morphology and phase stability. The results showed that surface hydrogenation induced significant changes in the shape of rutile nanocrystals, but not in anatase, and that the size at which the phase transition might be expected increased dramatically when the undercoordinated surface titanium atoms were H-terminated. For spherical particles, the crossover point was about 2.6 nm. For a clean and faceted surface, at low temperatures (a phase transition pointed at an average diameter of approximately 9.3-9.4 nm for anatase nanocrystals), the transition size decreased slightly to 8.9 nm when the surface bridging oxygens were H-terminated, and the size increased significantly to 23.1 nm when both the bridging oxygens and the undercoordinated titanium atoms of the surface trilayer were H-terminated. Below the cross point, the anatase phase was more stable than the rutile phase.371 In their study on TiO2 nanoparticles in vacuum or water environments, they found that the phase transition size in water (15.1 nm) was larger than that under vacuum (9.6 nm).373 In their predictions on the transition enthalpy of nanocrystalline anatase and rutile, they found that thermochemical results could differ for various faceted or spherical


Figure 38. Morphology predicted for anatase (top), with (a) hydrogenated surfaces, (b) hydrogen-rich surface adsorbates, (c) hydrated surfaces, (d) hydrogen-poor adsorbates, and (e) oxygenated surfaces, and for rutile (bottom), with (f) hydrogenated surfaces, (g) hydrogen-rich surface adsorbates, (h) hydrated surfaces, (i) hydrogen-poor adsorbates, and (j) oxygenated surfaces. Reprinted with permission from Barnard, A. S.; Curtiss, L. A. Nano Lett. 2005, 5, 1261. Copyright 2005 American Chemical Society.

nanoparticles as a function of shape, size, and degree of surface passivation.372 Their study on anatase and rutile titanium dioxide polymorphs passivated with complete monolayers of adsorbates by varying the hydrogen to oxygen ratio with respect to a neutral, water-terminated surface showed that termination with water consistently resulted in the lowest values of surface free energy when hydrated or with a higher fraction of H on the surface on both anatase and rutile surfaces, but conversely, the surfaces generally had a higher surface free energy when they had an equal ratio of H and O in the adsorbates or were O-terminated.375 They demonstrated that, under different pH conditions from acid to basic, the phase transition size of a TiO2 nanoparticle varied from 6.9 to 22.7 nm, accompanied with shape changes of the TiO2 nanoparticles as shown in Figure 38.374 Enyashin and Seifert conducted a theoretical study on the structural stability of TiO2 layer modifications (anatase and lepidocrocite) using the density-functional-based tight binding method (DFTB).376 They found that anatase nanotubes were the most stable modifications in a comparison of singlewalled nanotubes, nanostrips, and nanorolls. Their stability increased as their radii grew. The energies for all TiO2 nanostructures relative to the infinite monolayer followed a 1/R2 curve. Chen et al. found that severe distortions existed in Ti site environments in the structures of 1.9 nm TiO2 nanoparticles compared to those octahedral Ti sites in bulk anatase Ti using K-edge XANES.377 The distorted Ti sites were likely to adopt a pentacoordinate square pyramidal geometry due to the truncation of the lattice. The distortions in the TiO2 lattice were mainly located on the surface of the nanoparticles and were responsible for binding with other small molecules. Qian et al. found that the density of the surface states on TiO2 nanoparticles was likely dependent upon the details of the preparation methods.378 The TiO2 nanoparticles prepared from basic sol were found to have more surface states than those prepared from acidic sol based on a surface photovoltage spectroscopy study.

3.2. Thermodynamic Properties of TiO2 Nanomaterials Rutile is the stable phase at high temperatures, but anatase and brookite are common in fine grained (nanoscale) natural


Figure 39. Enthalpy of nanocrystalline TiO2. Reprinted with permission from Ranade, M. R.; Navrotsky, A.; Zhang, H. Z.; Banfield, J. F.; Elder, S. H.; Zaban, A.; Borse, P. H.; Kulkarni, S. K.; Doran, G. S.; Whitfield, H. J. Proc. Natl. Acad. Sci. U.S.A. 2002, 99, 6476. Copyright 2002 National Academy of Sciences, U.S.A.

and synthetic samples. On heating concomitant with coarsening, the following transformations are all seen: anatase to brookite to rutile, brookite to anatase to rutile, anatase to rutile, and brookite to rutile. These transformation sequences imply very closely balanced energetics as a function of particle size. The surface enthalpies of the three polymorphs are sufficiently different that crossover in thermodynamic stability can occur under conditions that preclude coarsening, with anatase and/or brookite stable at small particle size.73,74 However, abnormal behaviors and inconsistent results are occasionally observed. Hwu et al. found the crystal structure of TiO2 nanoparticles depended largely on the preparation method.379 For small TiO2 nanoparticles (973 K. Banfield et al. found that the prepared TiO2 nanoparticles had anatase and/or brookite structures, which transformed to rutile after reaching a certain particle size.73,380 Once rutile was formed, it grew much faster than anatase. They found that rutile became more stable than anatase for particle size > 14 nm. Ye et al. observed a slow brookite to anatase phase transition below 1053 K along with grain growth, rapid brookite to anatase and anatase to rutile transformations between 1053 K and 1123 K, and rapid grain growth of rutile above 1123 K as the dominant phase.381 They concluded that brookite could not transform directly to rutile but had to transform to anatase first. However, direct transformation of brookite nanocrystals to rutile was observed above 973 K by Kominami et al.382 In a later study, Zhang and Banfield found that the transformation sequence and thermodynamic phase stability depended on the initial particle sizes of anatase and brookite in their study on the phase transformation behavior of nanocrystalline aggregates during their growth for isothermal and isochronal reactions.74 They concluded that, for equally sized nanoparticles, anatase was thermodynamically stable for sizes < 11 nm, brookite was stable for sizes between 11 and 35 nm, and rutile was stable for sizes > 35 nm. Ranade et al. investigated the energetics of the TiO2 polymorphs (rutile, anatase, and brookite) by high-temperature oxide melt drop solution calorimetry, and they found the energetic stability crossed over between the three phases as shown in Figure 39.383 The dark solid line represents the phases of lowest enthalpy as a function of surface area. Rutile was energetically stable for surface area < 592 m2/mol (7 m2/g or >200 nm), brookite was energetically stable from


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592 to 3174 m2/mol (7-40 m2/g or 200-40 nm), and anatase was energetically stable for greater surface areas or smaller sizes ( 40 m2/g, it was metastable with respect to both anatase and rutile, and the sequence brookite to anatase to rutile during coarsening was energetically downhill. If anatase formed initially, it could coarsen and transform first to brookite (at 40 m2/g) and then to rutile. The energetic driving force for the latter reaction (brookite to rutile) was very small, explaining the natural persistence of coarse brookite. In contrast, the absence of coarse-grained anatase was consistent with the much larger driving force for its transformation to rutile.383 Li et al. found that only anatase to rutile phase transformation occurred in the temperature range of 973-1073 K.384 Both anatase and rutile particle sizes increased with the increase of temperature, but the growth rate was different, as shown in Figure 40. Rutile had a much higher growth rate than anatase. The growth rate of anatase leveled off at 800 °C. Rutile particles, after nucleation, grew rapidly, whereas anatase particle size remained practically unchanged. With the decrease of initial particle size, the onset transition temperature was decreased. An increased lattice compression of anatase with increasing temperature was observed. Larger distortions existed in samples with smaller particle size. The values for the activation energies obtained were 299, 236, and 180 kJ/mol for 23, 17, and 12 nm TiO2 nanoparticles, respectively. The decreased thermal stability in finer nanoparticles was primarily due to the reduced activation energy as the size-related surface enthalpy and stress energy increased.

3.3. X-ray Diffraction Properties of TiO2 Nanomaterials XRD is essential in the determination of the crystal structure and the crystallinity, and in the estimate of the crystal grain size according to the Scherrer equation

D)

Kλ β cos θ

(3)

where K is a dimensionless constant, 2θ is the diffraction angle, λ is the wavelength of the X-ray radiation, and β is the full width at half-maximum (fwhm) of the diffraction peak.385 Crystallite size is determined by measuring the broadening of a particular peak in a diffraction pattern associated with a particular planar reflection from within the crystal unit cell. It is inversely related to the fwhm of an individual peaksthe narrower the peak, the larger the crystallite size. The periodicity of the individual crystallite

Figure 40. (A) Changes in particle sizes of anatase and rutile phases as a function of the annealing temperatures. (B) Arrenhius plot of ln(AR/A0) vs 1/T for activation energy calculations as a function of the size of the TiO2 nanoparticles. AR and A0 are the integrated diffraction peak intensity from rutile (110), and the total integrated anatase (101) and rutile (110) peak intensity, respectively. Reused with permission from W. Li, C. Ni, H. Lin, C. P. Huang, and S. Ismat Shah, Journal of Applied Physics, 96, 6663 (2004). Copyright 2004, American Institute of Physics.

domains reinforces the diffraction of the X-ray beam, resulting in a tall narrow peak. If the crystals are randomly arranged or have low degrees of periodicity, the result is a broader peak. This is normally the case for nanomaterial assemblies. Thus, it is apparent that the fwhm of the diffraction peak is related to the size of the nanomaterials. Figure 41 shows the XRD patterns for TiO2 nanoparticles of different sizes111 and for TiO2 nanorods of different lengths.129 As the nanoparticle size increased, the diffraction peaks became narrower. In the anatase nanoparticle and nanorods developed by Zhang et al., the diameters of the TiO2 nanoparticles and nanorods were both around 2.3 nm. The nanorods were elongated along the [001] direction with preferred anisotropic growth along the c-axis of the anatase lattice, which was indicated by the strong peak intensity and narrow width of the (004) reflection and relatively lower intensity and broader width for the other reflections. With an increase in length of the nanorods, the (004) diffraction peak became much stronger and sharper, whereas other peaks remained similar in shape and intensity.129 Similar results have been observed by other groups.123,127,177,183

3.4. Raman Vibration Properties of TiO2 Nanomaterials As the size of TiO2 nanomaterials decreases, the featured Raman scattering peaks become broader.255,318,370,386-395 The size effect on the Raman scattering in nanocrystalline TiO2 is interpreted as originating from phonon confine-



selection rule for the excitation of Raman active optical phonons with long-range order and crystallite size.318,370 In a perfect “infinite” crystal, conservation of phonon momentum requires that only optic phonons near the Brillouin zone (BZ) center (q ≈ 0) are involved in first-order Raman scattering. In an amorphous material lacking long-range order, the q vector selection rule breaks down and the Raman spectrum resembles the phonon density of states. For nanocrystals, the strict “infinite” crystal selection rule is replaced by a relaxed version. This results in a range of accessible q vectors (as large as ∆q ≈ 1/L (L diameter)) due to the uncertainty principle. The anatase TiO2 has six Raman-active fundamentals in the vibrational spectrum: three Eg modes centered around 144, 197, and 639 cm-1 (designated here Eg(1), Eg(2), and Eg(3), respectively), two B1g modes at 399 and 519 cm-1 (designated B1g(1) and B1g(2d)), and an A1g mode at 513 cm-1.370 As the particle size decreases, the Raman peaks show increased broadening and systematic frequency shifts (Figure 42).370 The most intense Eg(1) mode shows the maximum blue shift and significant broadening with decreasing crystallite size. A small blue shift is seen for the Eg(2) mode, while the B1g(1) mode and the B1g(2)+A1g modes show very small blue shifts and red shifts (the latter peak represents a combined effect of two individual modes), respectively. Whereas the frequency shifts for the A1g and B1g modes are not pronounced, increased broadening with decreasing crystallite size is clearly seen for these modes. The Eg(3) mode shows significant broadening and a red shift with decreasing crystallite size. Choi et al. found a volume contraction effect in anatase TiO2 nanoparticles due to increasing radial pressure as particle size decreases, and they suggested that the effects of decreasing particle size on the force constants and vibrational amplitudes of the nearest neighbor bonds contributed to both broadening and shifts of the Raman bands with decreasing particle diameter.388

3.5. Electronic Properties of TiO2 Nanomaterials

Figure 41. (A) Powder XRD patterns of TiO2 samples of different diameters: (a) 5 nm; (b) 7 nm; (c) 13 nm. Reprinted with permission from Niederberger, M.; Bartl, M. H.; Stucky, G. D. Chem. Mater. 2002, 14, 4364. Copyright 2002 American Chemical Society. (B) Powder XRD patterns of TiO2 samples of diameter 2.3 nm: (a) spherical particles; (b) 16-nm nanorods; (c) 30-nm nanorods. Reprinted with permission from Zhang, Z.; Zhong, X.; Liu, S.; Li, D.; Han, M. Angew. Chem., Int. Ed. 2005, 44, 3466. Copyright 2005 Wiley-VCH.

ment,255,318,370,386,387,395 nonstoichiometry,391,392 or internal stress/surface tension effects.390 Among these theories, the most convincing is the three-dimensional confinement of phonons in nanocrystals.255,318,370,386,387,394,395 The phonon confinement model is also referred to as the spatial correlation model or q vector relaxation model. It links the q vector

The DOS of TiO2 is composed of Ti eg, Ti t2g (dyz, dzx, and dxy), O pσ (in the Ti3O cluster plane), and O pπ (out of the Ti3O cluster plane), as shown in Figure 43A.396 The upper valence bands can be decomposed into three main regions: the σ bonding in the lower energy region mainly due to O pσ bonding; the π bonding in the middle energy region; and O pπ states in the higher energy region due to O pπ nonbonding states at the top of the valence bands where the hybridization with d states is almost negligible. The contribution of the π bonding is much weaker than that of the σ bonding. The conduction bands are decomposed into Ti eg (>5 eV) and t2g bands (3.0 eV), electrons are excited from the valence band into the unoccupied conduction band, leading to excited electrons in the conduction band and positive holes in the valence band. These charge carriers can recombine, nonradiatively or radiatively (dissipating the input energy as heat), or get trapped and react with electron donors or acceptors adsorbed on the surface of the photocatalyst. The competition between these processes determines the overall efficiency for various

applications of TiO2 nanoparticles. These fundamental processes can be expressed as follows:415

TiO2 + hυ 98e- + h+

(9)

e- + Ti(IV)O-H f Ti(III)O-H-(X)

(10)

h+ + Ti(IV)O-H f Ti(IV)O•-H+(Y)

(11)

1 1 h+ + O2-lattice T O2(g) + vacancy 2 4

(12)

e-| + O2,s f O2,s-

(13)



O2,s- + H+ T HO2,s

(14)

h+ + Ti(III)O-H- f Ti(IV)O-H

(15)

e- + Ti(IV)O•-H+ f Ti(IV)O-H

(16)

O2,s + Ti(IV)O•-H+ f Ti(IV)O-H + O2,s

(17)

Reaction 8 is the photon absorption process. Reactions 1014 are photocatalytic redox pathways, whereas reactions 1517 represent the recombination channels. Reactions 11 and 12 are the competition pathways for holes, leading to bound OH radicals and O vacancies, respectively. The reverse of reaction 12 generates O adatom intermediates upon exposing defective surfaces to O2-(g).415 Electrons and holes generated in TiO2 nanoparticles are localized at different defect sites on the surface and in the bulk. Electron paramagnetic resonance (EPR) results showed that electrons were trapped as two Ti(III) centers, while the holes were trapped as oxygen-centered radicals covalently linked to surface titanium atoms.416-419 Howe and Gra¨tzel found that irradiation at 4.2 K in vacuo produced electrons trapped at Ti4+ sites within the bulk and holes trapped at lattice oxide ions immediately below the surface, which decayed rapidly in the dark at 4.2 K. In the presence of O2, trapped electrons were removed and the trapped holes were stable to 77 K. Warming to room-temperature caused loss of trapped holes and formation of O2- at the surface.416,417 Hurum et al. found that, upon band gap illumination, holes appeared at the surface and preferentially recombined with electrons in surface trapping sites for mixed-phase TiO2, such as Degussa P25, and recombination reactions were dominated by surface reactions that followed charge migration.419 Colombo and Bowman studied the charge carrier dynamics of TiO2 nanoparticles with femtosecond time-resolved diffuse reflectance spectroscopy and found a dramatic increase in the population of trapped charge carriers within the first few picoseconds.420,421 Skinner et al. found that the trapping time for photogenerated electrons on 2 nm TiO2 nanoparticles in acetonitrile by ultrafast transient absorption was about 180 fs.422 Serpone et al. found that localization (trapping) of the electron as a Ti3+ species occurred with a time scale of about 30 ps and about 90% or more of the photogenerated electron/ hole pairs recombined within 10 ns.409 They suggested that photoredox chemistry occurring at the particle surface emanated from trapped electrons and trapped holes rather than from free valence band holes and conduction band electrons. Bahnemann et al. found that, in 2.4 nm TiO2 nanoparticles, electrons were instantaneously trapped within the duration of the laser flash (20 ns). Deeply trapped holes were rather long-lived and unreactive, and shallowly trapped holes were in a thermally activated equilibrium with free holes which exhibited a very high oxidation potential.423 Szczepankiewicz and Hoffmann et al. found that O2 was an efficient scavenger of conduction band electrons at the gas/solid interface and the buildup of trapped carriers eventually resulted in extended surface reconstruction involving Ti-OH functionalities.415 They found that photogenerated free conduction band electrons were coupled with acoustic phonons in the lattice and their lifetimes were lengthened when dehydrated.424 The photoexcited charge carriers in TiO2 nanoparticles produced Stark effect intensity and wavelength shifts for surface TiO-H stretching vibrations. Although deep electron-trapping states affected certain

Figure 47. (A) (a) Absorption spectrum and (b) luminescence excitation spectrum (wavelength of emission light is 400 nm) of colloidal TiO2 nanotubes of different mean diameters: (1) 2.5 nm; (2) 3.1 nm; (3) 3.5 nm; (4) 5 nm. The curves are shifted vertically for clarity. (B) Photoluminescence spectra of colloidal TiO2 nanotubes of different mean diameters: (1) 2.5 nm; (2) 3.1 nm; (3) 3.5 nm; (4) 5 nm. Room temperature, excitation wavelength 237 nm, slits width 5 nm. The range of wavelengths, 455-490 nm, in the spectra is omitted due to the high signal of the second harmonic from scattered excitation light. The curves are shifted vertically for clarity. Vertical lines (5) show the positions of the peaks in the PL spectrum of the nanosheets. Reprinted with permission from Bavykin, D. V.; Gordeev, S. N.; Moskalenko, A. V.; Lapkin, A. A.; Walsh, F. C. J. Phys. Chem. B 2005, 109, 8565. Copyright 2005 American Chemical Society.

types of TiO-H stretch, shallow electron-trapping states produced a homogeneous electric field and were suggested not to be associated with localized structures, but rather delocalized across the TiO2 surface.424 Berger et al. studied UV light-induced electron-hole pair excitations in anatase TiO2 nanoparticles by electron paramagnetic resonance (EPR) and IR spectroscopy.425 The localized states such as holes trapped at oxygen anions (O-)


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Figure 48. Schematic presentation of the transformation of the electron band structure of the nanosheet semiconductor accompanying the formation of nanotubes: (a) band diagram of a 2-dimensional nanosheet; (b) band diagram of quasi-1-D nanotubes; (c) energy density of states for nanosheets (G2D) and nanotubes (G1D). EG1D and EG2D are the band gaps of the 1D and 2D structures, respectively. kx and ky are the wave vectors. Reprinted with permission from Bavykin, D. V.; Gordeev, S. N.; Moskalenko, A. V.; Lapkin, A. A.; Walsh, F. C. J. Phys. Chem. B 2005, 109, 8565. Copyright 2005 American Chemical Society.

trapped at localized sites, giving paramagnetic Ti3+ centers, or remained in the conduction band as EPR silent species which may be observed by their IR absorption and that the EPR-detected holes produced by photoexcitation were Ospecies, produced from lattice O2- ions. It was also found that, under high-vacuum conditions, the majority of photoexcited electrons remained in the conduction band. At 298 K, all stable hole and electron states were lost.

4. Modifications of TiO2 Nanomaterials

Figure 49. Scheme of UV-induced charge separation in TiO2. Electrons from the valence band can either be trapped (a) by defect states, which are located close to the conduction band (shallow traps), or (b) in the conduction band, where they produce absorption in the IR region. Electron paramagnetic resonance spectroscopy detects both electrons in shallow traps, Ti3+, and hole centers, O-. Reprinted with permission from Berger, T.; Sterrer, M.; Diwald, O.; Knoezinger, E.; Panayotov, D.; Thompson, T. L.; Yates, J. T., Jr. J. Phys. Chem. B 2005, 109, 6061. Copyright 2005 American Chemical Society.

and electrons trapped at coordinatively unsaturated cations (Ti3+ formation) were accessible to EPR spectroscopy. Delocalized and EPR silent electrons in the conduction band may be traced by their IR absorption, which results from their electronic excitation within the conduction band in the infrared region (Figure 49). They found that, during continuous UV irradiation, photogenerated electrons were either

Many applications of TiO2 nanomaterials are closely related to its optical properties. However, the highly efficient use of TiO2 nanomaterials is sometimes prevented by its wide band gap. The band gap of bulk TiO2 lies in the UV regime (3.0 eV for the rutile phase and 3.2 eV for the anatase phase), which is only a small fraction of the sun’s energy ( Cr > Mn > Fe > Ni.466-471 Anpo et al. found that the absorption band of Cr-ion-implanted TiO2 shifted smoothly toward the visible light region, with the extent of the red shift depending on the amount of metal ions implanted as shown in Figure 54A.470 Impregnated or chemically Cr-ion-doped TiO2 showed no shift in the absorption edge of TiO2; however, a new absorption band appeared at around 420 nm as a shoulder peak due to the formation of an impurity energy level within the band gap, with its intensity increasing with the number of Cr ions (Figure 54B).470 In the study by Umebayashi et al., visible light absorption of V-doped TiO2 was due to the transition between the VB and the V t2g level.509 The holes in the VB produced an anodic photocurrent. The photoexcitation processes under visible light of V-, Cr-, and Mn-doped TiO2 are illustrated in Figure 55. Photoexcitation for V-, Cr-, Mn-, and Fe-doped TiO2 occurred via the t2g level of the dopant. The visible light absorption for Mn- and Fe-doped TiO2 was due to the optical transitions from the impurity band tail into the CB. The Mn (Fe) t2g level was close to the VB and easily

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Figure 54. (A) The UV-vis absorption spectra of TiO2 (a) and Cr ion-implanted TiO2 photocatalysts (b-d). The amount of implanted Cr ions (µmol/g) was (a) 0, (b) 0.22, (c) 0.66, or (d) 1.3. (B) The UV-vis absorption spectra of TiO2 (a) and Cr ion-doped TiO2 (b′-d′) photocatalysts prepared by an impregnation method. The amount of doped Cr ions (wt%) was (a) 0, (b′) 0.01, (c′) 0.1, (d′) 0.5, or (e′) 1. Reprinted from Anpo, M.; Takeuchi, M. J. Catal. 2003, 216, 505, Copyright 2003, with permission from Elsevier.

overlapped in highly impure media. The visible light absorption for the Cr-doped TiO2 can be attributed to a donor transition from the Cr t2g level into the CB and the acceptor transition from the VB to the Cr t2g level. Stucky et al. found that up to 8 mol % Eu3+ ions could be doped into mesoporous anatase TiO2, and excitation of the TiO2 electrons within their band gap led to nonradiative energy transfer to the Eu3+ ions with a bright red luminescence.287 The mesoporous TiO2 acted as a sensitizer. 4.1.2.2.2. Optical Properties of Nonmetal-Doped TiO2 Nanomaterials. Nonmetal doped TiO2 normally has a color from white to yellow or even light gray, and the onset of the absorption spectra red shifted to longer wavelengths (refs 385, 426, 478, 483, 489, 494, 495, 497, 498, 505, 506, 512, 516, 518, 519, 521, and 529). In N-doped TiO2 nanomaterials, the band gap absorption onset shifted 600 nm from 380 nm for the undoped TiO2, extending the absorption up to 600 nm, as shown in Figure 56.426 The optical absorption of N-doped TiO2 in the visible light region was primarily located between 400 and 500 nm, while that of oxygendeficient TiO2 was mainly above 500 nm from their densityfunctional theory study.520 N-F-co-doped TiO2 prepared by spray pyrolysis absorbs light up to 550 nm in the visible light spectrum.518 The S-doped TiO2 also displayed strong absorption in the region from 400 to 600 nm.494 The red shifts in the absorption spectra of doped TiO2 are generally attributed to the narrowing of the band gap in the electronic structure after doping.489 C-doped TiO2 showed long-tail absorption spectra in the visible light region.472,543 Cl-, Br-, and Cl-Br-doped TiO2 had increased optical response compared to the case of pure TiO2 in the visible region.508 Livraghi et al. recently found that N-doped TiO2 contained single atom nitrogen impurity centers localized in the band gap of the oxide which were responsible for visible light



Figure 56. Reflectance spectra of N-doped TiO2 nanoparticles and pure TiO2 nanoparticles. Reprinted with permission from Burda, C.; Lou, Y.; Chen, X.; Samia, A. C. S.; Stout, J.; Gole, J. L. Nano Lett. 2003, 3, 1049. Copyright 2003 American Chemical Society.

Figure 57. IPCEλ and APCEλ curves for N-doped TiO2 and TiO2. SE stands for the substrate/electrode (SE) interface. The action spectra are recorded with light incident onto the SE interface. Reprinted from Lindgren, T.; Lu, J.; Hoel, A.; Granqvist, C. G.; Torres, G. R.; Lindquist, S. E. Sol. Energy Mater. Sol. Cells 2004, 84, 145, Copyright 2004, with permission from Elsevier.

4.1.2.3. Photoelectrical Properties of Doped TiO2 Nanomaterials. The photoelectrical properties of a material can be measured with an “action spectrum” curve using a phototo-current conversion setup.385,486,497,521 In this setup, light from a xenon lamp passing through a monochromator is radiated onto the electrode, and the photocurrents from the electrodes are measured as a function of wavelength.385,486,497,521 The incident photo-to-current efficiency as a function of wavelength, IPCEλ, is called an “action spectrum”. IPCEλ can be calculated by

IPCEλ ) Figure 55. Schematic diagram to illustrate the photoexcitation process under visible light of metal-doped TiO2: (a) Ti1-xVxO2; (b) Ti1-xFexO2; (c) Ti1-xCrxO2. Reprinted from Umebayashi, T.; Yamaki, T.; Itoh, H.; Asai, K. J. Phys. Chem. Solids 2002, 63, 1909, Copyright 2002, with permission from Elsevier.

absorption with promotion of electrons from the band gap localized states to the conduction band.547 Nick Serpone “proposed that the commonality in all...doped titanias rests with formation of oxygen vacancies and the advent of color centers...that absorb the visible light radiation, and he argued that the red shift of the absorption edge is in fact due to formation of the color centers.546

hc Iph,λ e Pλλ

(18)

where Iph,λ is the photocurrent, Pλ is the power intensity of the light at wavelength λ, and h, c, and e are Planck’s constant, the speed of light, and the elementary charge, respectively.385 The IPCEλ curve normally has a similar shape and trend as the absorption spectrum. When the IPCEλ is divided by the absorption, the absorbed photon-to-current efficiency (APCEλ; also called the quantum yield) is obtained.521 Figure 57 shows IPCEλ and APCEλ curves for N-doped TiO2 and TiO2.521 The photoelectrochemical onset for TiO2-xNx is shifted to around 550 nm into the visible region of the spectrum, and some ultraviolet (UV) efficiency for TiO2-xNx is lost compared to that of TiO2, suggesting


the TiO2-xNx has a typical photoelectrochemical behavior of a material with states in the band gap which act as recombination centers for light-induced charge carriers.521 In another study, the action spectrum of N-doped TiO2 also displayed a higher response in the visible region than that of pure TiO2.486 The photocurrent spectra for the pure and S-doped crystals showed that the photocurrent spectrum edge shifted to the low-energy region below 2.9 eV for the S-doped crystal, compared to 3.0 eV for pure TiO2, due to the transition of electrons across the narrowed band gap between the VB and the CB.497

4.2. Surface Chemical Modifications When a photocurrent is generated with light energy less than that of the semiconductor band gap, the process is known as sensitization and the light-absorbing dyes are referred to as sensitizers.9,10 TiO2 is a semiconductor with a wide band gap, with optical absorption in the UV region ( untreated TiO2 nanotubes > TiO2 nanoparticles, since TiO2 nanotubes treated with H2SO4 were composed of smaller particles and had higher specific surface areas.818 TiO2 aerogels were also suggested as promising candidates for photocatalysts.316,317,319 Degan et al. prepared TiO2 aerogels with a porosity of 90% and surface areas of 600 m2/g, and they found that the photodegradation of salicylic acid on TiO2 aerogels, after 1 h of near-UV illumination, was about 10 times faster than that on the Degussa TiO2.316,317 Figure 61 shows photodegradation profiles for the aerogel before and after annealing, as compared to the commercial Degussa P25 powder.

5.1.2. Metal-Doped TiO2 Nanomaterials: Second Generation Over the past decades, metal-doped TiO2 nanomaterials have been widely studied for improved photocatalytic performance on the degradation of various organic pollutants, i.e., under visible light irradiation (refs 21, 430, 433-435, 444, 446, 450-452, 455-457, 490, 515, 548, 810, 827836). Choi et al. conducted a systematic study on the photocatalytic activity of TiO2 nanoparticles doped with 21 transition metal elements on the oxidation of CHCl3 and the reduction of CCl4 and found that the photocatalytic activity was related to the electron configuration of the dopant ion in that dopant ions with closed electron shells had little or no effect on the activity.434,435 Doping with Fe3+, Mo5+, Ru3+, Os3+, Re5+, V4+, and Rh3+ at 0.1-0.5 at % significantly increased the photoreactivity, while Co3+ and Al3+ doping decreased the photoreactivity. The presence of metal ion dopants in the TiO2 matrix significantly influenced the charge carrier recombination rates and interfacial electron-transfer



Figure 61. Photodegradation profiles of salicylic acid on annealed (Ela) and nonannealed (El) TiO2 aerogels as compared to a commercial Degussa P25. Reprinted with permission from Dagan, G.; Tomkiewicz, M. J. Phys. Chem. 1993, 97, 12651. Copyright 1993 American Chemical Society.

rates. The photoreactivity of doped TiO2 appeared to be a complex function of the dopant concentration, the energy level of dopants within the TiO2 lattice, their d electronic configurations, the distribution of dopants, the electron donor concentrations, and the light intensity. Sn4+ ion-doped TiO2 nanoparticle films prepared by the plasma-enhanced CVD method displayed a higher photocatalytic activity for photodegradation of phenol than pure TiO2 under both UV and visible light, and the Sn4+ dopant was found profitable to the separation of photogenerated carriers under both UV and visible light excitation.433 Figure 62 shows the photocatalytic decomposition of phenol with reaction time under UV and visible light using Sn4+-doped TiO2 nanoparticles as photocatalyst.433 Fe-doped nanocrystalline TiO2 was shown to display higher photocatalytic activity with lower Fe content (optimal 0.05% mass fraction) than TiO2 in the treatment of papermaking wastewater,837 and it was shown to be more efficient in the photoelectrocatalytic disinfection of E. coli than pure TiO2.827 V-doped TiO2 photocatalyst photooxidized ethanol under visible radiation and had comparable activity under UV radiation to that of pure TiO2.548 Pt4+ ion-doped TiO2 nanoparticles exhibited higher visible light photocatalytic activities on the degradations of dichloroacetate and 4-chlorophenol,830 and Ag-TiO2 nanocatalysts displayed enhanced photocatalytic activity in the degradation of 2,4,6-trichlorophenol due to a better separation of photogenerated charge carriers and improved oxygen reduction inducing a higher extent of degradation of atoms.809 Wei et al. synthesized La- and N-co-doped TiO2 nanoparticles with superior catalytic activity under visible light, where N doping was responsible for the band gap narrowing of TiO2 and La3+ doping prevented the aggregation of nanoparticles.833 Chang et al. reported Cr- and N-co-doped TiO2 nanomaterials with visible light absorbance generally led to a reduction in photocatalytic efficacy in the decolorization of methylene blue, except at the low nitrogen doping concentration.490 Bessekhouad et al. found that low concentration alkaline (Li, Na, K)-doped TiO2 nanoparticles were promising materials for organic pollutants degradation.430 Peng et al. found that in Be2+-doped TiO2 nanomaterials,

Figure 62. Variation of phenol concentration with reaction time under (A) UV and (B) visible light: (a) pure TiO2 catalyst; (b) Sn4+-doped TiO2. From: Cao, Y.; Yang, W.; Zhang, W.; Liu, G.; Yue, P. New J. Chem. 2004, 28, 218 (http://dx.doi.org/10.1039/ b306845e) s Reproduced by permission of The Royal Society of Chemistry (RSC) on behalf of the Centre National de la Recherche Scientifique (CNRS).

when the doping ions were in the shallow surface, the doping was beneficial, while, in the deep bulk, the doping was detrimental.451 However, not all the metal-doped TiO2 nanomaterials showed higher photocatalytic activities than pure TiO2 nanomaterials. Martin found V-doped TiO2 nanoparticles had reduced photocatalytic activity on the photooxidation of 4-chlorophenol compared to pure TiO2 nanoparticles. Vanadium appeared to reduce the photoreactivity of TiO2 by promoting charge-carrier recombination with electron trapping at VO2+ centers or with hole trapping at V4+ impurity centers, which shunted charge carriers away from the solid/ solution interface.446 Hermann et al. found that although Crdoped (0.85 atomic %) TiO2 absorbed in the visible region, its activity for oxidation of oxalic acid, propene, and 2-propanol and for O isotope exchange was null under visible illumination and was smaller under UV light than that of pure TiO2, due to an increase in electron-hole recombination at the Cr3+ ion sites.440 Luo et al. reported that the photoactivity of TiO2 doped with 1.5 mol % Mo, 1 mol % V, 0.1 mol % V plus 1 mol % Al, or 0.1 mol % V plus 1 mol % Pb decreased, since the d electrons of Mo(4d) and V(3d), as majority carriers in TiO2, could effectively quench the high-energy photogenerated holes at the impurity levels introduced by doping within the band gap of TiO2.445

5.1.3. Nonmetal-Doped TiO2 Nanomaterials: Third Generation Nonmetal-doped TiO2 nanomaterials have been regarded as the third generation photocatalyst. Various nonmetal-


Figure 63. Photocatalytic properties of TiO2-xNx and TiO2 based on decomposition rates [measuring the change in absorption of the reference light (∆abs)] of methylene blue as a function of the cutoff wavelength of the optical high-path filters under fluorescent light. The inset shows the decomposition rates of methylene blue in the aqueous solution under visible light as a function of the ratio of the decomposed area in the XPS spectra with the peak at 396 eV to the total area of N 1s. The total N concentrations were 1.0 atom % (a), 1.1 atom % (b), 1.4 atom % (c), 1.1 atom % (d), and 1.0 atom % (e). Reprinted with permission from Asahi, R.; Morikawa, T.; Ohwaki, T.; Aoki, K.; Taga, Y. Science 2001, 293, 269 (http:// www.sciencemag.org). Copyright 2001 AAAS.

doped TiO2 nanomaterials have been widely studied for their visible light photocatalytic activities (refs 21, 385, 426, 428, 452, 472-474, 481-487, 490, 492-494, 496, 505, 518, 520, 524, 525, 527, 532, 533, 802, 838-847). Nonmetal-doped TiO2 nanomaterials have been demonstrated to have improved photocatalytic activities compared to those for pure TiO2 nanomaterials, especially in the visible light region.426,428,485,489,833,848 Figure 63 shows the decomposition of methylene blue using N-doped TiO2 as measured by Asahi and co-wokers.489 It was found that N-doped TiO2 had much higher photocatalytic activity than pure TiO2 in the visible light region, while displaying lower activity in the UV-light region. A nitrogen concentration dependent performance of the photocatalytic activity of the nitrogen-doped TiO2 was found in the visible region, and the active sites of N for photocatalysis under visible light were identified with the atomic β-N states peaking at 396 eV in the XPS spectra.489 In the study of Irie and co-workers, the concentration dependent photocatalytic activity of the N-doped TiO2 was attributed to the fact that the band structure of the N-doped TiO2 with lower nitrogen concentration (35%.670,860 Figure 70 shows a bilayer nanoporous electrode which consists of a nanoporous TiO2 matrix covered with a thin layer of Nb2O5 and the performance of three TiO2 electrodes coated with Nb2O5. For the best coating condition, the photocurrent increased from 10.2 to 11.4 mA/cm2, the photovoltage from 661 to 730 mV, and the fill factor from 51.0 to 56.5%. As a result, the conversion efficiency of the solar cell increased by 35% from 3.62 to 4.97%.860 They also found that sometimes the shell material shifted the conduction band potential of the core rather than forming an energy barrier. For example, coating of TiO2 with a SrTiO3 shell resulted in a shift of the TiO2 conduction band in the negative direction.861,862 Consequently, introduction of a SrTiO3-coated TiO2 electrode to a DSSC increased the open circuit photovoltage while reducing the short circuit photocurrent compared to that of the noncoated TiO2 electrode.861,862 Diamant et al. found that the mechanism by which the shell affected the electrode properties depended on the coating material. Coating materials included Nb2O5, ZnO, SrTiO3, ZrO2, Al2O3, and SnO2.862 The coating Nb2O5 formed a surface energy barrier, which slowed the recombination reactions, while the other shell materials each formed a surface dipole layer that shifted the conduction band potential of the core TiO2. The shift direction and magnitude depended on the dipole parameters

which were induced by the properties of the two materials at the core/shell interface.862 Palomares et al. found that the conformal growth of an overlayer of Al2O3 on a nanocrystalline TiO2 film resulted in a 4-fold retardation of interfacial charge recombination and a 30% improvement in photovoltaic device efficiency.870 Fabregat-Santiago et al. found that the alumina barrier reduced the recombination of photoinjected electrons to both the dye cations and the oxidized redox couple, due to two effects: (a) almost complete passivation of surface trap states in TiO2 that were able to inject electrons to acceptor species and (b) slowing down by a factor of 3-4 of the rate of interfacial charge transfer from conduction band states.868 O’Regan found that the Al2O3 layer acted as a tunnel barrier, thus increasing Voc and the fill factor.869 Palomares et al. prepared SiO2, Al2O3, and ZrO2 overlayers by dipping mesoporous nanocrystalline TiO2 films in organic solutions of their respective alkoxides, followed by sintering at 435 °C.865 The metal oxide overlayers acted as barrier layers for interfacial electron-transfer processes. The most basic overlayer coating, Al2O3 (pzc ) 9.2), was optimal for retarding interfacial recombination losses under negative applied bias, with an increase in open-circuit voltage of up to 50 mV and a 35% improvement in overall device efficiency. Diamant et al. found that SrTiO3-coated nanoporous TiO2 electrodes increased the open circuit photovoltage while reducing the short circuit photocurrent and resulting in a 15% improvement of the overall conversion efficiency of the solar cell.861 The SrTiO3 layer shifted the conduction band of the TiO2 in the negative direction due to a surface dipole rather than forming an energy barrier at the TiO2/electrolyte interface.861,862 The shell having a more negative conduction band potential acted as an energy barrier that slowed recombination reactions. Photoexcitation of dye molecules anchored to ultrathin (e1 nm) outer shells of insulators or semiconductors on n-type semiconductor crystallites resulted in electron transfer to the inner core material. However, there is still considerable recombination that increases with the distance between the electron injection point and the current collector. In other words, the limited lifetime of the injected electron and the slow diffusion rate inside the porous structure limit the effective thickness of the nanoporous electrode. Chappel et al. proposed a electrode design, shown in Figure 71, with a core shell configuration based on a conductive ITO or Sb-doped SnO2 matrix coated with TiO2.872 In principle, the conducting core extended the current collector into the nanoporous network and was



Figure 72. (a) SEM of a cross section of the bilayer photonic crystal-nano-TiO2 photoelectrode. The conductive glass is at the top of the image in part a. The photonic crystal layer and the nanocrystalline TiO2 layer are enlarged in parts b and c, respectively. Reprinted with permission from Nishimura, S.; Abrams, N.; Lewis, B. A.; Halaoui, L. I.; Mallouk, T. E.; Benkstein, K. D.; van de Lagemaat, J.; Frank, A. J. J. Am. Chem. Soc. 2003, 125, 6306. Copyright 2003 American Chemical Society.

denoted the nanoporous “collector shell electrode”. Consequently, the distance between the injection spot and the current collector should decrease to several nanometers throughout the nanoporous electrode, in contrast to several micrometers with the standard electrode. All electrons injected into the electrode, including those generated several micrometers away from the substrate, had to travel a very short distance before reaching the current collector. As shown by several studies, transport shorter than 1 µm provides 100% collection efficiency. In addition, the new collector-shell electrode contained inherent screening capability due to the high doping level of the conducting matrix. Theoretically, the new design should enable efficient charge separation and collection for thick nanoporous layers and solid electrochemical mediators. They found that, unless the TiO2 coating was thicker than 6 nm, the electrode performance was very low due to fast recombination.872 5.2.1.4.4. Electrode Coupled with Photonic Crystals. Development of photosensitizers with improved spectral response at the low-energy end of the solar spectrum has not proven so successful because dye molecules with high red absorbance have lower excited-state excess free energy, thus lowering the quantum yield for charge injection. Increasing the thickness of the film beyond 10-12 µm in order to increase the absorbance in the red results in an increase in the electron transport length and the recombination rate, and a decrease in the photocurrent. An alternative approach to improving efficiency was to increase the path length of light by enhancing light scattering in the TiO2 films.873-878 While the small size of TiO2 nanoparticles (1030 nm) employed to ensure a high surface area makes conventional nanocrystalline TiO2 films poor light scatterers, mixing the nanoparticles with larger particles or applying a scattering layer to the nanocrystalline film has been shown to increase light harvesting by enhancing the scattering of light.873-878

Nishimura347 and Halaoui333 reported an enhancement in the light conversion efficiency of dye-sensitized TiO2 solar cells by coupling a conventional nanocrystalline TiO2 film to a TiO2 inverse opal, with a 26% increase in the IPCE relative to that of a nanocrystalline film of the same overall thickness in the 550-800 nm spectral range. They found that the bilayer architecture, rather than enhanced light harvesting within the inverse opal structures, was responsible for the bulk of the gain in the IPCE.333 Figure 72 shows an SEM image of a cross section of the bilayer photonic crystal-nano-TiO2 photoelectrode.347 Figure 73 shows the sketch for the mechanism of the photonic crystal in enhancing absorption in certain regimes.347 The fact that light waves were localized in different parts of the structure, depending on their energy, implied that an absorber in the high dielectric medium should interact more strongly with light at wavelengths to the red of the stop band, and less strongly to the blue. Effectively, the red part of the spectrum of this absorber would “borrow” intensity from the blue part. Figure 74A shows the effect of the TiO2 photonic crystal as compared to a film of nanocrystalline TiO2 on the absorption spectra when dye is adsorbed to the surface.347 In a comparison of the spectrum of dye molecules adsorbed to the TiO2 photonic crystal film with that of a conventional nanocrystalline TiO2 film, there was a substantial enhancement absorbance on the red side of the stop band, as well as a slight attenuation of absorbance on the blue side of the stop band. The enhanced absorbance was most pronounced between 500 and 550 nm, but it persisted to a lesser degree at longer wavelengths. Figure 74B shows the enhancement of the performance of a bilayer electrode compared to a conventional nanocrystalline TiO2 photoelectrode.347 Between 400 and 530 nm, there was little difference between the two kinds of electrodes. The close similarity in the maximum photocurrent from the two electrodes was consistent with


Figure 73. (A) Simplified optical band structure of a photonic crystal. Near the Brillouin zone center, light travels with velocity c0/n, where c0 is the speed of light in a vacuum and n is the average refractive index. At photon energies approaching a full band gap or a stop band from the red side, the group velocity of light decreases and light can be increasingly described as a sinusoidal standing wave that has its highest amplitude in the high-refractiveindex part of the structure. At energies above the band gap or stop band, the standing wave is predominantly localized in the low index part of the photonic crystal, i.e., in the air voids. (B) Illustration of the effect of standing wave localization on dye absorbance. In an isotropic medium, the dye absorbs strongly in the blue but weakly in the red (heavy line). If the stop band is tuned to the position shown by the arrow, the blue absorbance is diminished and the red absorbance is increased when the dye is confined to the highrefractive-index part of the photonic crystal (dotted line). Reprinted with permission from Nishimura, S.; Abrams, N.; Lewis, B. A.; Halaoui, L. I.; Mallouk, T. E.; Benkstein, K. D.; van de Lagemaat, J.; Frank, A. J. J. Am. Chem. Soc. 2003, 125, 6306. Copyright 2003 American Chemical Society.

the fact that both contain the same amount of dye. Between 540 and 750 nm, the short circuit photocurrent was substantially increased in the bilayer electrode. The overall gain, integrated over the visible spectrum (400-750 nm), was about 30%. Localization of heavy photons at the edges of the photonic stop band347,879,880 from Bragg diffraction in the periodic lattice and multiple scattering events at disordered regions in the photonic crystal or at disordered films led ultimately to enhanced backscattering.333 This largely accounted for the enhanced light conversion efficiency in the red spectral range (600-750 nm), where the sensitizer was a poor absorber.333

5.2.2. Metal/Semiconductor Junction Schottky Diode Solar Cell McFarland and Tang reported a multilayer photovoltaic device structure in which photon absorption occurred in photoreceptors deposited on the surface of an ultrathin metal/ semiconductor junction Schottky diode.881 The device structure was a solid-state multilayer with a photoreceptor layer deposited on a 10-50 nm Au film, which capped 200 nm of TiO2 on an ohmic metal back contact (Figure 75). The photon-to-electron conversion in this device occurred in four

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Figure 74. (A) Absorption spectra of the TiO2 photonic crystal (a), the N719 dye adsorbed on the photonic crystal (b), and the dye adsorbed on a film of nanocrystalline TiO2 (c). The position of the stop band at 486 nm is indicated by the arrow. (B) Wavelength dependence of the short-circuit photocurrent in the bilayer electrode (a) and the conventional nanocrystalline TiO2 photoelectrode (b). The position of the stop band maximum in the bilayer electrode was 610 nm. Reprinted with permission from Nishimura, S.; Abrams, N.; Lewis, B. A.; Halaoui, L. I.; Mallouk, T. E.; Benkstein, K. D.; van de Lagemaat, J.; Frank, A. J. J. Am. Chem. Soc. 2003, 125, 6306. Copyright 2003 American Chemical Society.

steps. First, light absorption occurred in the surface-absorbed photoreceptors, giving rise to energetic electrons. Second, electrons from the photoreceptor excited state were injected into the conduction levels of the adjacent conductor, where they travelled ballistically through the metal at an energy, 1e, above the Fermi energy, Ef. Third, provided that 1e was greater than the Schottky barrier height, f, and the carrier mean-free path was long compared to the metal thickness, the electrons traversed the metal and entered the conduction levels of the semiconductor (internal electron emission). The absorbed photon energy was preserved in the remaining excess electron free energy when it was collected at the back ohmic contact, giving rise to the photovoltage, V. The photooxidized dye was reduced by transfer of thermalized electrons from states near Ef in the adjacent metal. Devices fabricated by using a fluorescein photoreceptor on an Au/ TiO2/Ti multilayer structure had typical open-circuit photovoltages of 600-800 mV and short-circuit photocurrents of 10-18 mA cm-2 under 100 mW cm-2 visible light illumination: the internal quantum efficiency (electrons measured per photon absorbed) was 10%. This alternative approach to photovoltaic energy conversion might provide the basis for durable low-cost solar cells using a variety of materials.

5.2.3. Doped TiO2 Nanomaterials-Based Solar Cell Lindgren et al. found that N-doped TiO2 nanocrystalline porous thin films showed visible light absorption in the wavelength range from 400 to 535 nm and generated an incident photon-to-current efficiency response in good agree-


Figure 75. Electron transfer in the operating photovoltaic device: (process A) photon absorption and electron excitation from the chromophore ground state, S, to the excited state, S*; (process B) energetic electron transfer from S* into and (ballistically) through the conducting surface layer and over the potential energy barrier into the semiconductor; (process C) conduction of electrons as majority carriers within the semiconductor to the ohmic back-contact and through the load; (process D) reduction of the oxidized chromophore, S, by a thermal electron from the conductor surface. Shown schematically are the relative energies of the electron levels within the device structures, the Schottky barrier, f, the Fermi energy, Ef, and the semiconductor band gap, Eg. Reprinted with permission from McFarland, E. W.; Tang, J. Nature 2003, 421, 616. Copyright Nature Publishing Group.

Figure 76. Reaction schemes for semiconductor photocatalysts. Reprinted Figure 2 from Kudo, A. Catal. SurV. Asia 2003, 7, 31, Copyright 2003, with kind permission of Springer Science and Business Media.

ment with the optical spectra.385 For the best nitrogen-doped TiO2 electrodes, the photoinduced current due to visible light and at moderate bias increased around 200 times compared to the behavior of pure TiO2 electrodes.

5.3. Photocatalytic Water Splitting 5.3.1. Fundamentals of Photocatalytic Water Splitting An enormous research effort has been dedicated to the study of the properties and applications of TiO2 under light illumination since the discovery of photocatalytic splitting of water on a TiO2 electrode in 1972 (Fujishima and Honda).6-8 Photocatalytic splitting of water into H2 and O2 using TiO2 nanomaterials continues to be a dream for clean and sustainable energy sources.882 Figure 76 shows the principle of water splitting using a TiO2 photocatalyst.761 When TiO2 absorbs light with energy larger than the band gap, electrons and holes are generated in the conduction and valence bands, respectively. The photogenerated electrons and holes cause redox reactions. Water molecules are reduced by the electrons to form H2


and oxidized by the holes to form O2, leading to overall water splitting.883-885 The width of the band gap and the potentials of the conduction and valence bands are important. The bottom level of the conduction band has to be more negative than the reduction potential of H+/H2 (0 V vs NHE), while the top level of the valence band has to be more positive than the oxidation potential of O2/H2O (1.23 V). The potential of the band structure of TiO2 is just the thermodynamical requirement. Other factors such as charge separation, mobility, and lifetime of photogenerated electrons and holes also affect the photocatalytic properties of TiO2. These factors are strongly affected by the bulk properties of the material such as crystallinity. Surface properties such as surface states, surface chemical groups, surface area, and active reaction sites are also important.768 The water-splitting process in return affects the local pH environment and surface structures of the TiO2 electrode.769 Salvador conducted a thermodynamic and kinetic consideration of water-splitting and competitive reactions in the photoelectrochemical cell, and they found that the overvoltage for evolution of O must be minimized, which was on the order of 0.6 eV for n-TiO2 electrodes loaded with RuO2.767 Cocatalysts such as Pt and NiO are often loaded on the surface in order to introduce active sites for H2 evolution. Thus, suitable bulk and surface properties and energy structure are demanded for photocatalysts. Laser-induced photocatalytic oxidation/splitting of water over TiO2 catalysts was studied.883,886,887 Sayama and Arakawa found that addition of carbonate salts to Pt-loaded TiO2 suspensions led to highly efficient water splitting.888 The carbonate ions affected both the Pt particles and the TiO2 surface. The Pt was covered with some titanium hydroxide compounds and the rate of the back reaction on the Pt was suppressed effectively in the presence of carbonate ions. The carbonate species aided desorption of O2 from the TiO2 surface.888 Khan and Akikusa found that bare n-TiO2 nanocrystalline film electrodes were unstable during watersplitting reactions under illumination of light and their stability could be significant improved when covered with Mn2O3.759

5.3.2. Use of Reversible Redox Mediators It has been reported that pure TiO2 could not easily split water into H2 and O2 in the simple aqueous suspension system.413,754,889 The main problem is the fast, undesired electron-hole recombination reaction.762 Therefore, it is important to prevent the electron-hole recombination process. The Pt-TiO2 system could be illustrated as a “shortcircuited” photoelectrochemical cell, where a TiO2 semiconductor electrode and a platinum-metal counterelectrode are brought into contact. Well-dispersed metal particles act as miniphotocathodes, trapping electrons, which reduces water to hydrogen. The role of sacrificial reagents is shown in Figure 77.761 When the photocatalytic reaction is carried out in aqueous solutions including easily oxidizable reducing reagents, photogenerated holes irreversibly oxidize the reducing reagents instead of water. This makes the photocatalyst electron-rich, and a H2 evolution reaction is enhanced as shown in Figure 77a. On the other hand, in the presence of electron acceptors such as Ag+ and Fe3+, the photogenerated electrons in the conduction band are consumed by them and an O2 evolution reaction is enhanced as shown in Figure 77b. These reactions using sacrificial reagents are regarded


Figure 77. Photocatalytic H2 (a) or O2 (b) evolution in the presence of sacrificial reagents. Reprinted Figure 5 from Kudo, A. Catal. SurV. Asia 2003, 7, 31, Copyright 2003, with kind permission of Springer Science and Business Media.

Figure 78. Proposed reaction mechanism for overall photocatalytic water splitting using a IO3-/I- redox mediator and a mixture of Pt-TiO2-antase and TiO2-rutile photocatalysts. Reprinted with permission from Abe, R.; Sayama, K.; Domen, K.; Arakawa, H. Chem. Phys. Lett. 2001, 344, 339. Copyright 2001 Elsevier.

as half reactions and are often employed for test reactions of photocatalytic H2 or O2 evolution. However, one should realize that the results do not guarantee a photocatalyst to be active for overall water splitting into H2 and O2 in the absence of sacrificial reagents. A sacrificial reagent helps to control the electron-hole recombination process. The photoefficiency of the process can be improved by the addition of sacrificial reagents.754,889,890 The sacrificial reagents help separation of the photoexcited electrons and holes. Various compounds such as methanol, ethanol, EDTA (an ethylenediaminetetraacetic derivative), Na2S, and Na2SO4 or ions such as I-, IO3-, CN-, and Fe3+ have been used as sacrificial reagents.753-755,757,890,891 Abe et al. conducted a series of experiments on water splitting under sunlight.753-755 They designed a new photocatalytic reaction that split water into H2 and O2 by a twostep photoexcitation system composed of an IO3-/I- shuttle redox mediator and two different TiO2 photocatalysts: Ptloaded TiO2-anatase for H2 evolution and TiO2-rutile for O2 evolution (Figure 78).753 Simultaneous gas evolution of H2 (180 mmol/h) and O2 (90 mmol/h) was observed from a basic (pH ) 11) NaI aqueous suspension of two different TiO2 photocatalysts under UV radiation. The overall water splitting proceeded by the redox cycle between IO3- and I- under basic conditions as follows: (a) water reduction to H2 and

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I- oxidation to IO3- over Pt-TiO2-anatase, and (b) IO3reduction to I- and water oxidation to O2 over TiO2-rutile. IO3- reduction to I- over Pt-TiO2-anatase is an undesirable reaction. If this reaction is suppressed, the total water-splitting reaction will take place more efficiently. The advantage of this system is that H2 gas is evolved over the Pt-TiO2anatase photocatalyst only and that O2 gas is evolved over the TiO2-rutile photocatalyst only, even from a mixture of IO3- and I- in a basic aqueous solution. Therefore, another undesirable backward reaction, H2O formation from H2 and O2 on Pt particles, was suppressed.753 They found that addition of a small amount of iodide anion, I-, into the aqueous suspension of Pt-TiO2-anatase photocatalyst significantly improved the splitting into H2 and O2 with a stoichiometric ratio. The iodide anion was adsorbed preferentially onto the Pt cocatalyst as iodine atom. This iodine layer effectively suppressed the backward reaction of water formation from H2 and O2 to H2O over the Pt surface.754 Fujihara et al. studied the photochemical splitting of water by combining the reduction of water to hydrogen using bromide ions and the oxidation of water to oxygen using FeIII ions.892 The bromide ions were oxidized to bromine on Pt-loaded TiO2 nanoparticles, and the FeIII ions were reduced to FeII ions on TiO2 nanoparticles. These two reactions were carried out in separated compartments and combined via platinum electrodes and cation-exchange membranes as shown in Figure 79. At the electrodes, FeII ions were oxidized by bromine, and protons were transported through the membranes to maintain the electrical neutrality and pH of the solutions in the two compartments. As a result, water was continuously split into hydrogen and oxygen under radiation. The reversible reactions on photocatalysts which often suffered from the effects of back reactions were largely prevented due to the low concentration of the products in solution. Lee et al. found that a considerable amount of photocatalytic H2 was produced from water over NiO/TiO2 in proportion to the hole scavenger CN-.890 Galinska and Walendziewski studied water splitting over a Pt-TiO2 catalyst with various sacrificial reagents, such as methanol, Na2S, EDTA, and I- and IO3- ions, and they found that the sacrificial reagents had a key role in hydrogen production via the photocatalyzed water-splitting reaction.757 Photocatalytic water splitting was obtained when EDTA and Na2S were used. They acted as effective hole scavengers, preventing oxygen formation and the recombination reaction of oxygen with hydrogen.

5.3.3. Use of TiO2 Nanotubes Mor et al. found that highly ordered TiO2 nanotube arrays efficiently decomposed water under UV radiation.198 The authors found that the nanotube wall thickness was a key parameter influencing the magnitude of the photoanodic response and the overall efficiency of the water-splitting reaction. For TiO2 nanotubes with 22-nm pore diameter and 34-nm wall thickness (Figure 80A), upon 320-400 nm illumination at an intensity of 100 mW/cm2, hydrogen gas was generated at the power-time normalized rate of 960 mmol/h W (24 mL/h W) at an overall conversion efficiency of 6.8% as shown in Figure 80B.198,199 They also claimed that, for illumination at 320-400 nm (98 mW/cm2), the TiO2 nanotube-array photoanodes could generate H2 by H2O photoelectrolysis with a photoconversion efficiency of 12.25%.212 Park et al. further found that, when doped with



Figure 80. (A) SEM images, top view, of 20 V TiO2 nanotube arrays anodized at 5 °C. (B) Photoconversion efficiency as a function of measured potential [vs Ag/AgCl] for 10 V samples anodized at four temperatures [i.e., 5, 25, 35, and 50 °C]. Reprinted with permission from Mor, G. K.; Shankar, K.; Paulose, M.; Varghese, O. K.; Grimes, C. A. Nano Lett. 2005, 5, 191. Copyright 2005 American Chemical Society. Figure 79. (A) Schematic of the photocatalytic reaction cell for splitting water. (B) Energy diagram of splitting of water by combined photocatalytic reactions. From: Fujihara, K.; Ohno, T.; Matsumura, M. Faraday Trans. 1998, 94, 3705 (http://dx.doi.org/ 10.1039/a806398b) s Reproduced by permission of The Royal Society of Chemistry.

carbon, TiO2-xCx nanotube arrays showed more efficient water splitting under UV and visible light illumination (>420 nm) than pure TiO2 nanotube arrays.475

5.3.4. Water Splitting under Visible Light 5.3.4.1. Water Splitting over Doped TiO2 Nanomaterials. In general, the conduction bands of stable oxide semiconductor photocatalysts consisting of metal cations with a d0 and d10 configuration consist of empty orbitals (LUMO) of the metal cations. On the other hand, the valence bands consist of O2p orbitals. The potential of this valence band (about +3 eV) is considerably more positive than the oxidation potential of H2O to O2 (E0 ) 1.23 V). Therefore, the band gaps of oxide semiconductor photocatalysts with the potential for H2 evolution inevitably become wide. Accordingly, a valence band or an electron donor level consisting of orbitals of some element, except for O2p, has to be formed to make the band gaps or the energy gaps

Figure 81. Strategy of the development of photocatalysts with a visible light response. Reprinted Figure 6 from Kudo, A. Catal. SurV. Asia 2003, 7, 31, Copyright 2003, with kind permission of Springer Science and Business Media.

narrow. New photocatalysts having the band structure shown in Figure 81 are necessary in order to develop materials for splitting water into H2 and O2 under visible light.761 The created levels have to possess not only the thermodynamical potential for oxidation of H2O but also the catalytic properties for the four-electron oxidation reaction. The following strategies can be considered for the development of visible light-driven photocatalysts: (i) forming a donor level above a valence band by doping some element into conventional photocatalysts with wide band gaps such as TiO2; (ii) creating


a new valence band employing some element; and (iii) controling the band structure by making a solid solution.761 Borgarello et al. found that water cleavage could be induced with visible light in colloidal solutions of Cr-doped TiO2 nanoparticles deposited with ultrafine Pt or RuO2.431 A pronounced synergistic effect in catalytic activity was noted when both RuO2 and Pt were co-deposited onto the particle. Jin and Lu found that Pt/B-doped TiO2 was a good system for water splitting under a B4O72- environment without sacrificial electron donor reagents.893 Luo et al. found that Br-- and Cl--co-doped nanocrystalline TiO2 with the absorption edge shifted to a lower energy region displayed higher efficiency for water splitting than pure TiO2.508 Jing et al. found that a Ni-doped mesoporous TiO2 photocatalyst with 0.2 wt % Pt accomplished hydrogen evolution at nearly 125.6 lmol/h compared to 81.2 lmol/h for TiO2 P25.894 N-, B-doped TiO2 nanomaterials have displayed higher activity than pure TiO2 in water splitting, i.e., under visible light illumination.529,889 Khan et al. found that a C-doped TiO2 nanocrystalline film with visible light response obtained by controlled combustion of Ti metal in a natural gas flame had a high water-splitting performance with a total conversion efficiency of 11% and a maximum photoconversion efficiency of 8.35% when illuminated at 40 mW/cm2,476 although there were questions about its solar-to-hydrogen conversion efficiency by other researchers.895-897 Matsuoka et al. developed visible light responsive TiO2 nanocrystalline thin films by the radio frequency magnetron sputtering method, which decomposed water when Pt-loaded and in the presence of a sacrificial reagent such as methanol or silver nitrate under visible light.763,764 5.3.4.2. Water Splitting over Dye-Sensitized TiO2. Duonghong et al. found that TiO2 loaded simultaneously with ultrafine Pt and RuO2 displayed extremely high activity as an H2O decomposition catalyst under band gap excitation of the TiO2 and that, when Ru(bipy)32+ or rhodamine B was used as a sensitizer, H2O was decomposed under visible light.898 Abe et al. investigated H2 production over merocyanine or coumarin dye C343 or Ru complex dye N3 dye-sensitized Pt/TiO2 photocatalysts under visible light in a wateracetonitrile solution containing iodide as an electron donor.756 They found that the rates of H2 evolution decreased with increasing proportion of water in the solutions because of a decrease in the energy gap between the redox potential of I3-/I- and the HOMO levels of the dyes, which decreases the efficiency of electron transfer from I- to dye. The energy diagram and the mechanism for the H2 production from water over the dye-sensitized Pt/TiO2 photocatalyst system are shown in Figure 82. The two key electron-transfer steps, electron injection from an excited state of the dye to the TiO2 conduction band and oxidation of I- to I3- (steps 2 and 5), occurred efficiently in acetonitrile solvent. The increased ratio of water hindered electron transfer from I- to the HOMO level of the oxidized dye (step 5). In addition, Park and Bard designed two different kinds of cells with bipolar dye-sensitized TiO2/Pt panels connected so that their photovoltages added to provide vectorial electron transfer for unassisted water splitting to yield the separated products H2 and O2.765

5.3.5. Coupled/Composite Water-Splitting System Akikusa et al. found that a self-driven system for water splitting under illumination could be achieved with the

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Figure 82. Energy diagram of H2 production from water over dyesensitized Pt/TiO2 photocatalysts in the presence of I- or EDTA as an electron donor. Reprinted with permission from Abe, R.; Sayama, K.; Sugihara, H. J. Sol. Energy Eng. 2005, 127, 413. Copyright 2005 by ASME.

combination of single-crystal p-SiC and nanocrystalline n-TiO2 photoelectrodes.899 Both photoelectrodes (p-SiC and n-TiO2) were placed side by side facing the light source and in contact with an electrolyte of 0.5 M H2SO4. The open circuit potential was found to be 1.24 V between the n-TiO2 and p-SiC photoelectrodes, with a maximum photocurrent density of 0.05 mA cm-2 under a closed circuit potential of 0.23 V, corresponding to an efficiency of 0.06%. The low cell photocurrent density and the photoconversion efficiency for the p-SiC/n-TiO2 self-driven system for the water-splitting reaction were due to the high band gap energies of both semiconductors and high recombination of photogenerated carriers mainly in the covalently bonded p-SiC. Takabayashi et al. proposed a solar water-splitting system based on a composite polycrystalline-Si/doped TiO2 thinfilm electrode for high-efficiency and low-cost by combining the advantages of Si and doped TiO2: (1) an n-Si electrode with surface alkylation and a metal nanodot coating gave an efficient and stable photovoltaic characteristic, and (2) TiO2 doped with other elements, such as nitrogen and sulfur, could cause water photooxidation (oxygen photoevolution) under visible light illumination.770 The structure and working mechanism of solar water splitting with this system is shown in Figure 83. Although a high solar-to-chemical conversion efficiency of more than 10% was calculated for this system, several major problems needed to be solved before the real device could show promising performance.770

5.4. Electrochromic Devices TiO2 nanomaterials have been widely explored as electrochromic devices, such as electrochromic windows and displays.611,634,772,900-916 Electrochromism can be defined as the ability of a material to undergo color change upon oxidation or reduction. Electrochromic devices are able to vary their throughput of visible light and solar radiation upon electrical charging and discharging using a low voltage. A small voltage applied to the windows will cause them to darken; reversing the voltage causes them to lighten. Thus, one can regulate the amount of energy entering through a “smart window” so that the need for air conditioning in a cooled building decreases. The energy efficiency inherent in this technology can be large, provided that the control strategy is adequate. Additionally, the transmittance regulation can impart glare control as well as user control of the indoor environment. The absorbance, rather than the reflec-



The second type is the electrochromism of nanocrystalline TiO2 electrodes modified with viologens and/or anthrachinons equipped with a surface anchoring group.902-904,906-908,910,917-919 This category usually has fast switching times and considerable optical dynamics, due to the combination of good conductivity between the TiO2 nanoparticles and the fast electron exchange between TiO2 and the monolayer of the electrochromic compound covering each particle.904 Bach et al. demonstrated high-quality paperlike electrochromic displays based on nanostructured TiO2 films modified with electrochromophores with excellent inkon-paper optical qualities, fast response times, and low power consumption.901 Moeller et al. demonstrated electrochromic pictures with unprecedented resolution (360 dpi) in transparent and reflective electrochromic displays (ECD) based on ink-jet printing technology and cascade-type crosslinking reactions of viologens in the mesopores of a TiO2 electrode, with a completely transparent counterelectrode based on mesoporous antimony tin oxide coated with CeO2.913

5.4.1. Fundamentals of Electrochromic Devices

Figure 83. (A) Schematic illustration and (B) the operation principle of solar water splitting with a composite polycrystallineSi/doped TiO2 semiconductor electrode. Reprinted from Takabayashi, S.; Nakamura, R.; Nakato, Y. J. Photochem. Photobiol., A: Chem. 2004, 166, 107, Copyright 2004, with permission from Elsevier.

tance, is modulated so that the electrochromic devices tend to heat up in their low-transparent state.752 Two types of electrochromism of nanocrystalline thin film TiO2 electrodes have been reported. The first type is the electrochromism of nanocrystalline TiO2 electrodes in Licontaining electrolytes related to the reversible insertion of Li+ into the anatase lattice of the nanoparticles.912 Hagfeldt et al. found that forward biasing of transparent nanocrystalline TiO2 films in lithium ion-containing organic electrolytes led to rapid and reversible coloration due to electron accumulation and Li+ intercalation in the anatase lattice.912 Absorption of >90% light throughout the visible and near IR could be switched on and off within a few seconds. The nanocrystalline morphology of the film played a role in enhancing the electrochromic process. Ottaviani et al. found that the rate of the electrochromic process was controlled by the diffusion of the Li+ ions throughout the TiO2 lattice.914 It was convenient to drive the electrochromic process with potentiostatic pulses, and under these conditions, many cycles with initially good color contrast and efficiencies which approached 100% were obtained with TiO2 thin film electrodes.

Figure 84A shows the principle of the electrochromism of a molecular monolayer adsorbed on TiO2.902 A molecule, which functions as the electrochromophore and exhibits different colors in different oxidation states, must be chosen such that its redox potential lies above the conduction band edge of the TiO2 nanocrystalline electrode at the liquid/solid interface. In this way, electrons can be transferred reversibly from the conduction band to the molecule. The TiO2 electrode in fact behaves like a conductor for the adsorbed electroactive species. If the redox potential is situated below the conduction band edge, the reduction process is irreversible. Figure 84B shows the TiO2 nanocrystalline electrochromic devices based on viologen (solvent: glutarodinitrile) with a counterelectrode made of Prussian blue.902 The device could be switched back and forth between the colorless and the colored states within 1 s. The nanocrystalline structure of the TiO2 film makes possible 100- to 1000-fold amplification compared to a flat surface as shown in Figure 85.902 The combination of high conductivity of the nanocrystalline TiO2 particles, fast electron exchange with the molecular monolayer, optical amplification by the porous structure, and fast charge compensation by ions in the contacting liquid makes the nanocrystalline electrodes highly attractive electrochromic elements. The principle of efficiency relies on fast interfacial electron transfer between the nanocrystalline TiO2 and the adsorbed modifier as well as on the high surface area of the TiO2 support that amplifies optical phenomena by 2 or 3 orders of magnitude.902 The investigated TiO2 nanocrystalline electrodes include ordered905 and disordered902-904 mesoporous films. Ordered mesoporous nanocrystalline TiO2 electrodes were found to display enhanced color contrast yet have similar conduction band edge energy levels and electron percolation ability as electrodes made from nanocrystalline TiO2, attributed to the uniform and ordered mesopore architecture and the large accessible surface area for tethering viologen molecules.905

5.4.2. Electrochromophore for an Electrochromic Device The viologen group (N,N′-disubstituted-4,4′-bipyridinium) has been commonly chosen as an electrochromophore, for its remarkable stability in both the oxidized and the reduced (radical cation) states (Figure 86).902,904,906,907 Oxidized vi-


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Figure 84. (A) Principle of the electrochromism of a molecular monolayer adsorbed on a semiconductor surface. Electrons are injected from the conducting substrate into the conduction band of the semiconductor and from there reduce the adsorbed electroactive molecule. Provided the redox potential of that molecule lies above the conduction band edge, the process is reversible by application of a positive potential to the conductive substrate. (B) Nanocrystalline electrochromic devices based on viologen (solvent: glutarodinitrile) with a counterelectrode made of Prussian blue, in the colorless and in the colored state. Reprinted from Bonhote, P.; Gogniat, E.; Campus, F.; Walder, L.; Gra¨tzel, M. Displays 1999, 20, 137, Copyright 1999, with permission from Elsevier.

ologen is colorless, while the radical cation can be blue, violet, purple, or green, depending on the substituents. The associated first reduction potential is between 0.2 and -0.6 V (vs NHE). The typical absorption spectrum of reduced N,N′-dialkylviologen in an organic solvent has a maximum around 600 nm. With N,N′-diarylviologens, the absorption band is shifted by about 50 nm to the red. In concentrated solution or in the solid state, viologen radical cations form dimers, with their blue-shifted absorption maximum in the 550 nm region. A second reduced state can be reached at potentials which are more negative by 0.2-0.4 V. This state is neutral and almost colorless (yellowish). This second reduction is reversible in organic solvents like acetonitrile but not in water. The anchoring groups with strong affinity toward TiIV include carboxylates, salicylates, or phosphonates.902 Bonhote et al. examined phosphonated triarylamine as an electrochromophore due to its oxidation by the stable

triarylamminum radical cation, which is accompanied by a blue coloration with the absorption band at 730 nm.903 Vayssieres et al. studied bis(phthalocyaninato)lutetium(III) complexes (Pc2Lu) as electrochromophores, and they found that the typical neutral green state of Pc2Lu was reduced to a brown state at potentials < -0.3 V vs Ag/AgCl at neutral pH when Pc2Lu was adsorbed onto a nanostructured TiO2 electrode.916 Ag-TiO2 films, prepared by loading nanoporous films with Ag nanoparticles by photocatalytic means, exhibited multicolor photochromism, which was related to the oxidation and reduction of Ag nanoparticles under UV and visible radiation.773 Please also see section 4.2.1.2 on Sensitization by Metal Nanoparticles.

5.4.3. Counterelectrode for an Electrochromic Device Closed cells are built by combining a transparent nanocrystalline electrode with a counterelectrode able to provide



Figure 85. Principle of signal amplification by a TiO2 nanocrystalline film. Sintered 20-nm particles of TiO2 form a several millimeter thick film characterized by a very high surface area. Once derivatized with a molecular adsorbate, the structure contains the equivalent of hundreds of superposed monolayers. Reprinted from Bonhote, P.; Gogniat, E.; Campus, F.; Walder, L.; Gra¨tzel, M. Displays 1999, 20, 137, Copyright 1999, with permission from Elsevier.

ing glass and modified by the electrochromophore [β-(10phenothiazyl)propoxy]phosphonic acid, which displayed cycles-switching times of

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