J. Phys. Chem. C 2009, 113, 16293–16298
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Temperature-Dependent Growth of Self-Assembled Hematite (r-Fe2O3) Nanotube Arrays: Rapid Electrochemical Synthesis and Photoelectrochemical Properties Thomas J. LaTempa, Xinjian Feng, Maggie Paulose, and Craig A. Grimes* Department of Electrical Engineering, and The Materials Research Institute, The PennsylVania State UniVersity, UniVersity Park, PennsylVania 16802 ReceiVed: May 15, 2009; ReVised Manuscript ReceiVed: July 30, 2009
We report on the self-assembled fabrication and photoelectrochemical properties of R-Fe2O3 (hematite) nanotube arrays, prepared by potentiostatic anodization of iron foil in an ethylene glycol electrolyte containing NH4F and deionized water. Vertically oriented nanotube arrays provide a highly ordered material architecture of high surface area that is nearly ideal for efficient transport and separation of photogenerated charges. We have achieved iron oxide nanotube arrays over a potential range of 30-60 V, using an electrolyte comprised of 0.2-0.5 wt % NH4F, and 2-4% DI water. The resulting nanotube arrays have a pore diameter ranging between 30 and 80 nm, with a minimum wall thickness of ∼10 nm. Nanotube formation is strongly dependent on the anodization bath temperature and fluoride concentration, with higher temperatures leading to increased rates of nanotube array growth, and with lower temperatures required in order to obtain identifiable nanotubes arrays for increased F- ion concentration. Crystallization of the as-synthesized, amorphous, iron oxide nanotube arrays to hematite is achieved through annealing in an oxygen-deficient ambient. Introduction Our interest is in materials suitable for scale production of hydrogen by water photoelectrolysis.1,2 While Fujishima and Honda first reported water photoelectrolysis using n-type TiO2 in 1972,3 more than 35 years later it remains an unsolved challenge to produce a photocorrosion-stable material system that captures a significant portion of solar spectrum energy and that can efficiently use that collected solar energy to split water. While not critical for laboratory demonstrations, should we desire to achieve hydrogen generation on a significant scale a further requirement of this material system is that it be abundant in nature. Hematite, R-Fe2O3 has a bandgap of ≈2.2 eV (indirect), well suited for capturing a significant portion of the solar spectrum energy with a maximum theoretical photoconversion efficiency of 12.9%.4 Unfortunately, iron oxide suffers from rapid electron-hole pair recombination, poor conductivity, and a low absorption coefficient.5 Furthermore the flat band potential of hematite is more positive then the hydrogen reduction potential, necessitating application of an external bias for water splitting,6,7 although it has been suggested that 1D quantum confinement effects may be sufficient for direct water splitting if an upward shift of the conduction band edge is realized.8 Recent efforts toward developing iron oxide for water splitting have focused on fabrication of leaflet,9 wire,10 and porous11 nanoarchitectures with desired feature sizes commensurate with the minority carrier diffusion length of R-Fe2O3, ≈4 nm,5 which should significantly reduce unwanted recombination of the photogenerated charge carriers. Synthesis of randomly oriented hematite nanotubes via a hydrothermal technique have been developed,12 while more recently the transformation of nanoporous hematite into nanotubes has been demonstrated through sonoelectrochemical anodization of iron foil.13 * To whom correspondence should be addressed. E-mail: cgrimes@ engr.psu.edu.
Our interest is in the photoelectrochemical application of selfassembled vertically oriented hematite nanotube arrays. The nanotube array geometry offers large surface areas,14 vectorial charge transport of majority carriers, short radial diffusion distances for minority carriers, and lengths sufficient to absorb incident sunlight.15,16 Nanotube arrays have been achieved through electrochemical anodization of valve metals such as titanium17-19 and tantalum.20 Titania nanotube arrays have been successfully applied in water photoelectrolysis achieving incident photon to electron (IPCE) conversion efficiencies of over 80%,21,22 heterojunction solar cells,23 room temperature hydrogen gas sensing,24-26 solar conversion of carbon dioxide and water vapor to hydrocarbon fuels,27 and drug delivery.28,29 Ti-Fe-O nanotube arrays have been reported, fabricated by anodization of cosputtered titanium and iron thin films.30 Drawing from our experience in anodization techniques, we report on the fabrication of highly ordered, vertically oriented iron oxide nanotube arrays that are subsequently crystallized to form hematite nanotubes under an oxygen deficient ambient. Experimental Section Iron foil (99.99% pure, 0.25 mm thick), ethylene glycol (EG; 99.8%, anhydrous), and ammonium fluoride (98%) were purchased from Sigma Aldrich. Anodization was conducted in a two-electrode electrochemical cell with a platinum foil cathode under constant DC potential, with a water circulating heater (MGW Lauda, RM6) used to control the bath temperature. The nanotube array films were subsequently annealed in ultra high purity nitrogen or argon at a rate of 5 °C/min with a dwell time of 30 min. The surface, as well as lateral topography, of the nanotubes was verified using field-emission scanning electron microscopy (FESEM; JSM6300, JEOL, Tokyo, Japan). Images of a single nanotube were obtained using transmission electron microscopy (TEM; Phillips, EM 420) and used to confirm the wall thickness of the as-anodized samples. Diffuse reflectance UV-vis spectroscopy was performed with a Perkin-Elmer Lambda 950 spectrophotometer equipped with an integrating
10.1021/jp904560n CCC: $40.75 2009 American Chemical Society Published on Web 08/13/2009
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sphere. Sample crystallinity was verified using glancing angle X-ray diffraction (GAXRD, Scintag X2, California, USA). IPCE was measured using a 300 W xenon lamp (Spectra physics) integrated with parabolic reflector, passing through a monochromator, using 0.5 V DC bias. Samples were characterized with the nanotube arrays as the working photoelectrode and platinum foil as the counter electrode, 1 M NaOH solution was used as the electrolyte. Results and Discussion Film Morphology. Panels A-B of Figure 1 show, respectively, FESEM top-view and cross-sectional images of iron oxide nanotubes prepared by anodization of iron foil at a potential of 50 V in EG + 0.3 wt % NH4F + 3.0% deionized (DI) water. Anodization was carried out for 180 s at a starting temperature of 60 °C. The surface morphology demonstrates a nanotube structure with an average pore diameter of 55 nm and length exceeding 3 µm, while the TEM image of a single amorphous iron oxide tube, Figure 1C, indicates a wall thickness of ∼10 nm. The nanotube pore diameters vary with anodization voltage; in otherwise identical conditions pore diameters are found to vary from 35 to 65 nm over a potential range of 30-60 V. Rapid growth rates are achieved using elevated electrolyte temperatures: at an anodization temperature of 75 °C nanotube lengths of 1.67 µm are obtained after 30 s anodization, and 2.75 µm after 60 s. The minimum temperature required to form nanotubes was found to vary with the fluoride and water concentration of the electrolyte. Panels A-B of Figure 2, respectively, show FESEM top-view and cross-sectional images of iron oxide nanotubes prepared by anodization at a potential of 50 V in EG + 0.2 wt % NH4F + 2.0% DI water for 180 s with a starting anodization temperature of 75 °C. The resulting nanotubes with a maximum length of ≈4 µm had an average pore diameter of ≈42 nm and wall thickness in excess of 20 nm. Panels C-D of Figure 2, respectively, show FESEM top-view and cross-sectional images of iron oxide nanotubes prepared in an electrolyte containing EG + 0.5 wt % NH4F + 3.0% DI water for 180 s with a starting temperature of 45 °C. The resulting ≈2.6 µm long nanotubes had an average pore diameter of ≈60 nm and wall thickness in excess of 8 nm. By varying the electrolyte bath temperature, we have achieved iron oxide nanotube arrays over a potential range of 30-60 V, using an electrolyte comprised of 0.2-0.5 wt % NH4F and 2-4% DI water. The onset of nanotube formation was found to be strongly dependent on the electrolyte bath temperature and fluoride content, with factors such as anodization potential and water content playing lesser roles. Figure 3A-C shows FESEM top-view images of samples prepared using a starting temperature of (A) 35, (B) 45, and (C) 55 °C in EG + 0.3 wt % NH4F + 3.0% DI water at 50 V for a duration of 180 s. A nanoporous surface is obtained at the lower electrolyte temperature, with transformation to a nanotube structure at higher anodization temperatures. It is important to note that the iron sample self-heats during anodization. Figure 3D,E shows a sample prepared at room temperature (≈22 °C) using EG + 0.35 wt % NH4F + 3.0% DI water at 50 V for a duration of 1200 s. A nanoporous surface is obtained for shorter anodization times; however, the electrolyte undergoes self-heating during anodization, hence nanotube formation becomes favorable with increased anodization duration. For the 1200 s anodization, nanotubes are found upon a nanoporous base with a total film thickness of ≈4.5 µm. Figure 4A shows the variation of iron oxide nanopore/nanotube array
Figure 1. (A, B) FESEM top view and cross-sectional images of iron oxide nanotube arrays prepared by anodization of iron foil at a potential of 50 V in ethylene glycol + 0.3 wt % NH4F + 3.0% DI water, for a duration of 180 s at 60 °C. (C) TEM image of a single, amorphous iron oxide nanotube prepared under identical conditions.
length with anodization temperature for anodizations conducted at 50 V for 180 s in an ethylene glycol bath containing 0.3 wt % NH4F and 3% H2O. The nanopore/nanotube length increases linearly with temperature, as does the pore diameter until saturating at ∼50 °C when tube formation occurs. Figure 4B shows the variation in inner pore diameter as a function of applied potential for otherwise identical anodization conditions but at 55 °C. Nanotube formation was found to occur between 30 and 60 V with a near linear relationship observed within
Growth of Self-Assembled Hematite (R-Fe2O3) Nanotube Arrays
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Figure 2. (A, B) FESEM top view and cross-sectional images of iron oxide nanotube arrays prepared by anodization of iron foil at a potential of 50 V in ethylene glycol + 0.2 wt % NH4F + 2.0% DI water, for a duration of 180 s at 75 °C. (C, D) FESEM top view and cross-sectional images of iron oxide nanotube arrays prepared by anodization of iron foil at a potential of 50 V in ethylene glycol + 0.5 wt % NH4F + 3.0% DI water, for a duration of 180 s at 45 °C.
Figure 3. FESEM top view images of iron oxide nanotube arrays prepared by anodization of iron foil at a potential of 50 V in ethylene glycol + 0.3 wt % NH4F + 3.0% DI water, for a duration of 180 s at (A) 35, (B) 45, and (C) 55 °C. (D, E) FESEM top view and cross-sectional images of iron oxide nanotube arrays prepared by anodization of iron foil, initially at room temperature (≈22 °C) at a potential of 50 V in ethylene glycol + 0.35 wt % NH4F + 3.0% DI water, for a duration of 1200 s.
this potential range. The wall thickness was not found to be directly controllable. Although it appears to be dependent on fluoride concentration, with higher fluoride concentration yielding nanotubes having thinner walls, the anodization temperature was necessarily scaled at each fluoride concentration (ranging from 45 to 75 °C) to achieve tubes. For example, 0.2 wt % flouride, 75 °C results in long nanotubes having relatively thick walls, while 0.5 wt % flouride, 45 °C results in shorter tubes having thin walls. If a constant temperature is used, nanotubes cannot be formed over a large range of fluoride concentration, and vice versa. Elucidating the interrelated effects of temperature and fluoride concentration is a focus of our ongoing research.
Nanotube Array Formation. Figure 5A illustrates the current-time behavior during anodization of iron foil at three (starting) anodization bath temperatures, using the same anodization parameters as Figure 1 (50 V, 0.3 wt % NH4F, 3% H2O in EG). Increased current densities are observed with a rise in bath temperature. Figure 5B is an expanded view of the current behavior within the first 90 s of anodization at 60 °C; the current-time behavior closely resembles that observed in the anodization of titanium foil,31 with an initial drop in current due to the formation of an insulating oxide layer and the subsequent current increase due to oxide pitting by fluoride ions. While for formation of titania nanotubes arrays by Ti anodization
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LaTempa et al. density continued to rise slowly due to additional heating of the electrolyte bath temperature, until ∼20 min beyond which the film dissolves into the solution. Nanotube formation in fluoride-bearing electrolytes can be achieved through the delicate balance of three simultaneous processes, namely (1) the field assisted oxidation of Fe metal to form iron oxide, (2) the field assisted dissolution of Fe metal ions in the electrolyte, and (3) chemical dissolution of Fe and Fe2O3 due to etching by fluoride ions, which is substantially enhanced by the presence of H+ ions. For aqueous, acidic electrolytes, an initial oxide layer is formed according to the reaction:
2Fe + 3H2O f Fe2O3 + 6H+
(1)
Conversely, localized chemical dissolution of the oxide occurs due to fluoride ions:
Fe2O3 + 12F- + 6H+ f 2[FeF6]3- + 3H2O
Figure 4. (A) Variation of iron oxide nanopore/nanotube array length and average inner pore diameter with anodization temperature. Anodization was conducted at 50 V for 180 s in an ethylene glycol bath containing 0.3 wt % NH4F and 3% H2O. (B) Variation of nanotube inner pore diameter with anodization potential for samples prepared under conditions otherwise identical to (A), at a constant temperature of 55 °C.
Figure 5. (A) Constant-voltage current density of iron foil anodized at 50 V in an electrolyte containing ethylene glycol + 0.3 wt % NH4F + 3.0% DI water at 60 °C for 180 s. (B) Magnified view of part a showing the first 90 s of anodization.
a steady-state current is obtained when maximum nanotube array length is achieved,32,33 for the iron oxide nanotubes the current
(2)
An increase in chemical dissolution will lead to thinner barrier layers at the pore bottoms keeping the electrochemical processes active. Electrolytes having lower diffusion coefficients are known to prevent local acidification at the semiconductorelectrolyte interface limiting the barrier layer thickness.11 The effect of electrolyte bath temperature has been investigated for polar organic electrolytes where it has been found to have a significant impact on nanotube array formation.34 It was hypothesized that at lower electrolyte temperatures fluoride ion mobility was suppressed in viscous electrolytes, thereby limiting the rate of chemical dissolution; longer lengths and larger pore diameters have been obtained with increased electrolyte temperature.35 Although chemical dissolution is necessary for nanotube formation to occur, the field-assisted oxidation and dissolution rates will also increase with temperature and water content.30 Therefore, an optimal temperature range exists for nanotubes formation that is dependent on a number of factors including fluoride ion concentration, water content, and applied bias; reaction kinetics were previously described for anodization of titanium foil in ethylene glycol solvents.36 Film Structure. The as-anodized samples are amorphous, with partially crystallized samples achieved through sample annealing. GAXRD data were obtained for samples prepared using EG + 0.3 wt % NH4F + 3.0% DI water bath at a temperature of 60 °C, 50 V, for 180 s, with different annealing ambients used for the otherwise identical samples. Figure 6 shows the GAXRD pattern of samples annealed at 400 °C for 1800 s in air, UHP argon, UHP nitrogen or a mixture containing 95% UHP argon and 5% hydrogen. Oxygen deficient ambients were investigated to reduce unwanted filling of the nanotubes pores due to oxidation of the underlying metal foil and to minimize the barrier layer thickness between the nanotubes and the underlying substrate. Hematite peaks are obtained with predominately (104) and (110) diffraction peaks along with several other orientations and smaller peaks corresponding to magnetite (Fe3O4) for the samples annealed in argon and nitrogen. The use of hydrogen was investigated to introduce additional vacancies with the intent of enhancing the material conductivity and to form predominately hematite (110).13 The hematite lattice can be modeled as alternating bilayers of iron and oxygen parallel to the (001) basal plane; an anisotropic conductivity is observed with preferred orientation along the
Growth of Self-Assembled Hematite (R-Fe2O3) Nanotube Arrays
J. Phys. Chem. C, Vol. 113, No. 36, 2009 16297 also be more favorable for the 3.0 eV transition.10 A further reduction in the wall thickness in addition to enhanced crystallinity can be expected to improve the observed photoconversion efficiencies through the reduction of recombination and/or trapping centers across grain boundaries, various defects and remaining amorphous regions. In addition, n-type doping can be used to substitute for iron sites, i.e., Si4+ for Fe3+, resulting in higher charge carrier concentrations.9 Conclusions
Figure 6. GAXRD pattern of iron oxide nanotubes annealed in air, UHP argon, UHP nitrogen, and a mixture containing 95% UHP argon and 5% hydrogen for 1800 s at 400 °C, with a ramp rate of 5 °C/min.
Figure 7. UV-vis absorbance spectra and IPCE of iron oxide nanotubes prepared using EG + 0.3 wt % NH4F + 3.0% DI water at a potential of 50 V for 180 s, annealed in pure argon at 400 °C for duration of 1800 s with a ramp rate of 5 °C/min. IPCE was measured in 1 M NaOH with 0.4 V DC bias applied.
[110] direction.37,38 However, the hydrogen-annealed samples were heavily reduced and showed, primarily, an iron diffraction peak along with magnetite. The samples annealed in air predominately showed hematite peaks. Further investigations are underway to determine optimal annealing conditions. Photoelectrochemical Properties. IPCE measurements were used to confirm the visible light activity of iron oxide nanotubes. Figure 7 plots IPCE vs UV-vis absorbance data for hematite nanotubes prepared using EG + 0.3 wt % NH4F + 3.0% DI water at a potential of 50 V for 180 s, annealed in pure argon at 400 °C for duration of 1800 s. IPCE measurements are calculated using the following equation: IPCE )
(1240 eV · nm)(photocurrent density µA/cm2) × 100% (λ nm)(irradiance µW/cm2)
(3) In 1 M NaOH solution, Fe2O3 nanotubes exhibit visible light activity, although the IPCE quickly drops after ∼400 nm. This may be attributed to differences observed between the two charge-transfer processes; a direct transition (3.0 eV) corresponding to the ligand-metal transfer O2- f Fe3+, and the indirect transition (2.2 eV) from the metal-metal transfer of 2Fe3+ f Fe2+ + Fe4+. It has been suggested that electrons trapped at Fe2+ sites are not free to conduct and holes from the Fe4+ may not contribute to oxygen evolution because of low transfer rates to the electrolyte.5 Charge carrier generation may
Hematite nanotube arrays have been prepared using temperature-controlled anodization of iron foil in an ethylene glycol electrolyte containing NH4F and DI water. We have achieved iron oxide nanotube arrays over a potential range of 30-60 V, using an electrolyte comprised of 0.2-0.5 wt % NH4F and 2-4% DI water. Nanotube formation has a strong dependence on both the electrolyte bath temperature and fluoride concentration, with lower temperature required for increased flouride ion concentration. Crystallization of the as-synthesized, amorphous, iron oxide nanotube arrays is achieved through annealing in an oxygen-deficient ambient. Future efforts will focus on modifying the electrolyte bath chemistry to obtain thinner walled nanotube arrays of shorter length. In addition, alternative routes for crystallization and doping are currently being investigated to optimize the photoelectrochemical properties. Acknowledgment. Support of this work by DE-FG0206ER15772 is gratefully acknowledged. The authors thank Sorachon Yoriya for her assistance with the FESEM images, and Trevor Clark of the Materials Research Institute for collecting TEM images. Thanks to various members of Grimes research group for valuable discussions. References and Notes (1) Grimes, C. A.; Varghese, O. K.; Ranjan, S. Light, Water, Hydrogen: The Solar Generation of Hydrogen by Water Photoelectrolysis; Springer: Norwell, MA, 2008. (2) Lewis, N. S. Nature 2001, 414, 589. (3) Fujishima, A.; Honda, K. Nature 1972, 238, 37. (4) Murphy, A. B.; Barnes, P. R. F.; Randeniya, L. K.; Plumb, I. C.; Grey, I. E.; Horne, M. D.; Glasscock, J. A. Int. J. Hydrogen Energy 2006, 31, 1999. (5) Kennedy, J. H.; Frese, K. W. J. Electrochem. Soc. 1978, 125, 709. (6) Gratzel, M. Nature 2001, 414, 338. (7) Ingler, W. B.; Khan, S. U. M. Electrochem. Solid State Lett. 2006, 9, G144. (8) Vayssieres, L.; Sathe, C.; Butorin, S. M.; Shuh, D. K.; Nordgren, J.; Guo, J. H. AdV. Mater. 2005, 17, 2320. (9) Cesar, I.; Kay, A.; Martinez, J. A. G.; Gratzel, M. J. Am. Chem. Soc. 2006, 128, 4582. (10) Beermann, N.; Vayssieres, L.; Lindquist, S. E.; Hagfeldt, A. J. Electrochem. Soc. 2000, 147, 2456. (11) Prakasam, H. E.; Varghese, O. K.; Paulose, M.; Mor, G. K.; Grimes, C. A. Nanotechnology 2006, 17, 4285. (12) Jia, C. J.; Sun, L. D.; Yan, Z. G.; You, L. P.; Luo, F.; Han, X. D.; Pang, Y. C.; Zhang, Z.; Yan, C. H. Angew. Chem., Int. Ed. 2005, 44, 4328. (13) Mohapatra, S. K.; John, S. E.; Banerjee, S.; Misra, M. Chem. Mater. 2009, 21, 3048. (14) Shankar, K.; Mor, G. K.; Prakasam, H. E.; Yoriya, S.; Paulose, M.; Varghese, O. K.; Grimes, C. A. Nanotechnology 2007, 18, 065707. (15) Grimes, C. A. J. Mater. Chem. 2007, 17, 1451. (16) Shankar, K.; Basham, J. I.; Allam, N. K.; Varghese, O. K.; Mor, G. K.; Feng, X. J.; Paulose, M.; Seabold, J. A.; Choi, K. S.; Grimes, C. A. J. Phys. Chem. C 2009, 113, 6327. (17) Gong, D.; Grimes, C. A.; Varghese, O. K.; Hu, W. C.; Singh, R. S.; Chen, Z.; Dickey, E. C. J. Mat. Res. 2001, 16, 3331. (18) Paulose, M.; Shankar, K.; Yoriya, S.; Prakasam, H. E.; Varghese, O. K.; Mor, G. K.; Latempa, T. A.; Fitzgerald, A.; Grimes, C. A. J.Phys. Chem. B 2006, 110, 16179. (19) Grimes, C. A.; Mor, G. K. TiO2 Nanotube Arrays: Synthesis, Properties and Applications; Springer: Norwell, MA, 2009. (20) Allam, N. K.; Grimes, C. A. J. Phys. Chem. C 2007, 111, 13028.
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(21) Mor, G. K.; Varghese, O. K.; Paulose, M.; Shankar, K.; Grimes, C. A. Sol. Energy Mater. Sol. Cells 2006, 90, 2011. (22) Varghese, O. K.; Paulose, M.; Shankar, K.; Mor, G. K.; Grimes, C. A. J. Nanosci. Nanotechnol. 2005, 5, 1158. (23) Shankar, K.; Mor, G. K.; Prakasam, H. E.; Varghese, O. K.; Grimes, C. A. Langmuir 2008, 24, 14321. (24) Varghese, O. K.; Mor, G. K.; Grimes, C. A.; Paulose, M.; Mukherjee, N. J. Nanosci. Nanotechnol. 2004, 4, 733. (25) Paulose, M.; Varghese, O. K.; Mor, G. K.; Grimes, C. A.; Ong, K. G. Nanotechnology 2006, 17, 398. (26) Varghese, O. K.; Yang, X.; Kendig, J.; Paulose, M.; Zeng, K.; Palmer, C.; Ong, K. G.; Grimes, C. A. Sens. Lett. 2006, 4, 120. (27) Varghese, O. K.; Paulose, M.; LaTempa, T. J.; Grimes, C. A. Nano Lett. 2009, 9, 731. (28) Popat, K. C.; Eltgroth, M.; LaTempa, T. J.; Grimes, C. A.; Desai, T. A. Biomaterials 2007, 28, 4880. (29) Popat, K. C.; Eltgroth, M.; La Tempa, T. J.; Grimes, C. A.; Desai, T. A. Small 2007, 3, 1878.
LaTempa et al. (30) Mor, G. K.; Prakasam, H. E.; Varghese, O. K.; Shankar, K.; Grimes, C. A. Nano Lett. 2007, 7, 2356. (31) Mor, G. K.; Varghese, O. K.; Paulose, M.; Grimes, C. A. AdV. Funct. Mater. 2005, 15, 1291. (32) Paulose, M.; Mor, G. K.; Varghese, O. K.; Shankar, K.; Grimes, C. A. J. Photochem. Photobiol. 2006, 178. (33) Ruan, C.; Paulose, M.; Varghese, O. K.; Grimes, C. A. Sol. Energy Mater. Sol. Cells 2006, 90, 1283. (34) Wang, J.; Lin, Z. J. Phys. Chem. C 2009, 113, 4026. (35) Macak, J. M.; Schmuki, P. Electrochim. Acta 2006, 52, 1258. (36) Prakasam, H. E.; Shankar, K.; Paulose, M.; Varghese, O. K.; Grimes, C. A. J. Phys. Chem. C 2007, 111, 7235. (37) Iordanova, N.; Dupuis, M.; Rosso, K. M. J. Chem. Phys. 2005, 122. (38) Kay, A.; Cesar, I.; Gratzel, M. J. Am. Chem. Soc. 2006, 128, 15714.
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