Preparation and Band Energetics of Transparent Nanostructured

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J. Phys. Chem. C 2010, 114, 815–819

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Preparation and Band Energetics of Transparent Nanostructured SrTiO3 Film Electrodes Shuming Yang,*,† Huizhi Kou,† Hongjun Wang,† Ke Cheng,‡ and Jichao Wang† College of Chemistry and Chemical Engineering, Institute of Applied Chemistry, Xinyang Normal UniVersity, Henan, China, 464000, and Key Laboratory of Special Functional Materials, Henan UniVersity, Kaifeng, Henan, China, 475001 ReceiVed: September 21, 2009; ReVised Manuscript ReceiVed: NoVember 22, 2009

Transparent nanostructured SrTiO3 film electrodes were fabricated on conductive substrates from SrTiO3 nanocrystals. Band energetics were determined with spectroelectrochemical measurements, and its trap states were investigated with electrochemistry. The flat band edges (Efb) of nanostructured SrTiO3 electrodes were highly dependent on the pH of the electrolyte, and measured to be -0.71, -0.96, and -1.19 V versus saturated Ag/gCl at pH 3.0, 6.8, and 13.0, respectively. The time-resolved current at different applied potentials positive of the flat band edge clearly indicate a trap-filling process. The results showed that trap state densities are also highly pH dependent. Total trap state densities of 2.0 × 1015, 8.08 × 1015, and 1.21 × 1016 cm-2 were determined at pH 3.0, 6.8, and 13.0, respectively with maximum located at -0.25, -0.50, and -0.70 V. The results obtained from cyclic voltammograms are in good agreement with that obtained from the measurements of time-resolved currents. The size of the peak potentials related to the trap states (-0.25, -0.51, and -0.72 V for pH 3.0, 6.8, and 13.0, respectively) in the cyclic voltammograms increases dramatically with increase of pH, indicating that traps are mostly surface-related. Introduction There has been increasing interest in the research field of nanostructured semiconductor metal oxide films. Nanostructured films combine a small primary particle size (1-100 nm) with a high porosity, resulting in a large surface area. This gives nanostructured electrodes great advantages over conventional electrodes in many applications such as photoelectrochemical solar cells1-3 and electrochromic windows.4 The metal oxides that have been employed are typically wide band gap n-type semiconductors such as TiO2,1,5 SnO2, ZnO, 7 and so forth. A striking characteristic of nanostructured electrodes is the relatively large density of surface states. Surface states are electronic energy levels located at the semiconductor surface. They are most important when the energies of these states lie in the band gap. Generally, their natures and properties are measured with cyclic voltammetry,8 charge recombination kinetics,9,10 and spectroelectrochemistry.5,11 Spectroelectrochemical methods are valuable in the study of transparent nanostructured metal oxide electrodes. Liu and Bard measured the absorption spectra of CdS and observed a reversible bleach in the bandgap region at negative applied potentials.12 This bleach was explained by electron presence in the conduction-band states, and this electron presence resulted in a Burstein shift.13 Fitzmaurice et al. studied electron accumulation in nanostructured TiO2 electrodes with a similar approach.7,14-19 In addition to the expected Burstein shift, they observed a monotonically increasing absorption in near-infrared regions. The former can be accounted for by band filling, and the latter was assigned principally to the transitions in the conduction band. It was assumed that the applied potential at which the observed spectral changes are first detected is related to the potential of the flat band edge (Efb) at the semiconductor* Corresponding author. Tel: +86-0376-6393736; e-mail address: [email protected]. † Xinyang Normal University. ‡ Henan University.

liquid electrolyte interface. To obtain a value for Efb, a model was developed, which accounted for the potential dependence of the absorption spectrum of a nanostructured TiO2 electrode.15 Fitting experimental absorbance-potential curves using this model yields Efb for nanostructured TiO2 electrodes in aqueous electrolytes. Fitzmaurice and co-workers determined Efb for nanostructured TiO2 electrodes under a wide range of conditions with the model.7,15-19 They found that, in protic solvents, Efb displayed a Nernstian dependence on pH and was independent of the nature and concentration of electrolyte cation.15 Strontium titanate (SrTiO3) with a perovskite structure shares more structural similarities with anatase TiO2 and can be to some extent thought of as a highly doped TiO2 structure. This perovskite material contains titanium atoms in 6-fold octahedral coordination, similar to the titanium arrangement in anatase. The two materials have nearly identical band gaps (Eg ) 3.2 eV),20 but thorough investigation of the band energetics of nanostructured SrTiO3 films has rarely been reported so far. In this work, SrTiO3 nanoparticles with narrow size distribution were synthesized and used to fabricate transparent nanostructured SrTiO3 film electrodes. The band energetics and the trap state properties at different pHs were investigated with electrochemical and spectroelectrochemical techniques. Experimental Section 1. Materials and Solutions. Optically transparent electrodes (OTE) were fabricated on an F-doped SnO2-coated glass substrate. Water (R ) 18.3 MΩ) from an Easy Pure RF water purification system from Thermo Scientific was used in the preparation of all solutions. Ti(n-OC4H9)4, HNO3, NaOH and ethyl cellulose were purchased from Tianjin Chemical Company. Sr(NO3)2, LiClO4, HClO4, terpineol, and tetramethylammonium hydroxide purchased from Shanghai Nuotai Chemical Company. All the chemicals used were reagent grade. 2. Preparation of Nanostructured SrTiO3 Films. Prepara-

10.1021/jp910204q  2010 American Chemical Society Published on Web 12/22/2009

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tion of SrTiO3 Nanocrystals. The preparation of TiO2 precursor and SrTiO3 nanocrystals was similar to that reported.20 Typically 48 mL of Ti(n-OC4H9)4 was poured into 250 mL water under vigorously stirring. After addition of 2 mL of 70% nitric acid, the mixture was heated to 80 °C for 2 h. After cooling, 17.5 mL of the solution was adjusted to pH ) 13 using tetramethylammonium hydroxide and degassed with nitrogen gas for at least 30 min. Sr(NO3)2 (1.95 g) was dissolved in 10 mL boiling distilled water, and added to the titanium dioxide suspension. This suspension was stirred at room temperature for at least 12 h and then heated in an autoclave at 190 °C for 4.5 h. Preparation of Nanostructured SrTiO3 Films. The resulting precipitate from the autoclave was washed with ethanol. Ethanol was removed by a rotary evaporator, and 6.84 g of terpineol and 8.44 g of 10% ethyl cellulose ethanol solution were added to the SrTiO3 paste. This mixture was then dispersed under sonication, and ethanol was removed by a rotary evaporator. The final paste was grounded with an agate mortar for 15 min. SrTiO3 paste was spread on the substrates by a glass rod with adhesive tapes as spacers. The films were dried at 125 °C and sintered at 500 °C for 30 min in air and finally cooled to room temperature. 3. Instrumentation. X-ray diffraction (XRD) measurements were performed on a D8 diffractometer (Bruker Co.) with Cu KR (λ ) 1.5405 Å) to identify the phase structure of samples. Transmission electron microscope (TEM) analysis was carried out with a JEOL-2010 microscope (Japan) operating at an accelerating voltage of 200 kV. The absorption spectra were recorded on a UV-1240 spectrophotometer (Shimadzu, Japan). Electrochemical experiments were performed on a CH800 electrochemical analyzer (CH Instrument). All electrochemical and spectroelectrochemical experiments were carried out in a typical three-electrode system, in which a nanostructured SrTiO3 electrode, a platinum wire, and a saturated Ag/AgCl electrode acted as working, counter, and reference electrodes, respectively. Spectroelectrochemistry measurements were undertaken according to published literature.7 A quartz cell with three electrodes and electrolyte was incorporated into the sample compartment of a Shimadzu UV-vis spectrophotometer and connected to a CHI 800 potentiostat. All aqueous electrolyte solutions were prepared in ion-free water with LiClO4 as the supporting electrolyte, and the pH of the electrolyte solutions was adjusted by HClO4 (for pH 3.0) or NaOH (for pH 13.0). The electrolyte solutions used were thoroughly deaerated by bubbling with N2 prior to experiments. All potentials are hereafter given with reference to the saturated Ag/AgCl electrode. The working area of SrTiO3 electrodes were 3 cm2. Results and Discussions 1. Preparation of Nanostructured SrTiO3 Film. In the synthesis of SrTiO3 nanoparticles, various parameters such as the concentrations of Ti and Sr precursors, pH, and the time and temperature of autoclaving were varied in order to investigate their effects on the quality of SrTiO3 nanoparticles. The pH is crucial for the formation of SrTiO3 phase and should be over 13, in agreement with the literature.21 It is very important to control the time and temperature of autoclaving in order to obtain SrTiO3 samples of narrow size distribution. We synthesized SrTiO3 nanoparticles under different conditions and found that too high temperature or too long time at autoclaving would easily yield defocused and polydispersed SrTiO3 samples due to Ostwald ripening. XRD analysis (Figure 1) shows that the synthesized SrTiO3 nanoparticles have a high pure cubic perovskite structure with

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Figure 1. XRD pattern of SrTiO3 nanoparticles.

Figure 2. TEM image of SrTiO3 nanoparticles.

Figure 3. Absorption spectrum of a nanostructured SrTiO3 film.

a space group of Pm3m in good agreement with the JCPDS (Joint Committee on Powder Diffraction Standards), card 350734. The TEM image (Figure 2) shows that the size distribution is very narrow with roughly spherical particles of about 40 nm. Figure 3 is the absorption spectrum of a nanostructured SrTiO3 film deposited on a quartz substrate. The absorption tail at longer wavelength is due to the diffusion and reflection of nanostructured SrTiO3 film. The onset is at 375 nm, corresponding to a band gap of 3.31 eV.20

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Figure 5. Flat band potential of nanostructured SrTiO3 films as a function of pH.

wavelength can be used to monitor the filling of electrons in the conduction band and, as a result, the conduction band edge of a semiconductor can be calculated. In Figure 4, the absorbance at 800 nm was plotted against applied potentials and shown in the inset figures. It can be seen that the flat band edges are closely dependent on the pH of the electrolyte and calculated to be -0.71, -0.96, and -1.19 V at pH 3.0, 6.8, and 13.0, respectively. It is known that flat band potential is dependent on the pH of the electrolyte, as shown by the following equations:22

E0f + ∆EH q ∆EH ) 0.0592(pHPZC - pH) Efb ) -

Figure 4. Differential spectra of nanostructured SrTiO3 electrodes in (a) 0.2 mol L-1 LiClO4 at pH 3.0, (b) 0.2 mol L-1 LiClO4 at pH 6.8, and (c) 0.1 mol L-1 LiClO4 at pH 13.0. The insets show absorbance changes at 800 nm. Spectra are recorded after being polarized for 5 min at indicated potentials. The spectrum measured after stabilization for 15 min at +0.8 V has been subtracted.

2. Spectroelectrochemistry and Flat Band Determination. Shown in Figure 4 are the potential-dependent absorption spectra of nanostructured SrTiO3 electrodes measured in aqueous electrolytes of different pHs. Application of a sufficiently negative potential results in an absorbance increase at longer wavelengths. The spectra measured at different pHs are similar, but the absorption onset shifted to more negative potential as pH increased. Spectroelectrochemical measurement was normally applied for the monitoring of the electron filling in the conduction band.12-19 Because the excitation of electrons in the conduction band needs only small energy, the absorbance changes at long

(1) (2)

where pHPZC is the pH at the point of zero charge (PZC), Ef0 is the Fermi level at pHPZC, and ∆EH is the potential drop in the Helmholtz layer. At pHPZC, the net charge on the surface is zero, and there is no potential drop in the Helmholtz layer. At pH smaller than pHPZC, there are net positive charges on the surface, so ∆EH > 0. The smaller the pH, the larger the ∆EH, and the flat band potential shifts more positively. On the other hand, at pH larger than pHPZC, there are net negative charge on the surface, so ∆EH < 0. The larger the pH, the smaller the ∆EH, and the flat band potential shifts more negatively. It was reported that the flat band edge of TiO2 is correlated linearly with pH and has the relation Efb ) -0.40 - 0.06pH versus saturated calomel electrode (SCE).15 The linear correlation was also found for the nanostructured SrTiO3 films. Figure 5 shows the dependence of the flat band edge of a nanostructured SrTiO3 film on the pH of the electrolyte. A linear relation was obtained from fitting data linearly and is expressed as Efb) -0.59 - 0.04pH. 3. Time-Resolved Current and Trap State Distribution. Time ResolWed Current at pH 13.0. The current resulting from each applied potential was measured versus time for a nanostructured SrTiO3 electrode in 0.1 mol L-1 LiClO4 solution of pH 13.0 under different potentials, and the results are shown in Figure 6a. It is shown that the current is significantly influenced by the applied potential. At potential from 0 to -0.6 V, the currents decrease to almost zero within a few seconds. At -0.7 V, the decrease in current slows down, and this behavior is found at all potentials more negative than -0.7 V. The results can be understood in terms of trap filling in the band gap region. A nanostructured SrTiO3 electrode has a flat band edge of -1.19 V at pH 13.0 (seen from Figure 5). At potentials more positive

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Figure 6. (a) Current-time curves of a nanostructured SrTiO3 electrode in 0.1 mol L-1 LiClO4 of pH 13.0. The electrode was initially polarized at 0.8 V for 5 min and then measured at different applied potentials. (b) Cathodic charges at different potentials derived by integrating the current-time curves in Figure 6a. The inset shows dQ/dU distribution against potential.

than -0.6 V, the density of traps is low, and thus the trapfilling time is short, resulting in fast decay of the time-resolved current.23 On the other hand, at potentials more negative than -0.6 V, trap density increases, so a longer time is required to fill these traps. The longest trap-filling time was at -1.1 V just below the conduction band edge, and a further negative shift of the potential significantly shortens the trap-filling time. This should be related to the kinetics of trap filling. A faster trap filling is expected at a more negative potential since the driving force for the trap filling is larger. The accumulated charge Q under the current-time curves in Figure 6a is calculated, and some interesting features appear. Figure 6b shows the accumulated charge Q against potentials. At potentials more positive than -0.5 V, the accumulated charge is pretty small. It is striking that there is a sharp increase in accumulated charge up to -0.9 V. At potential more negative than -0.9 V, the accumulated charges increase slowly (Figure 6b). If the accumulated charge Q from trap-filling reflects the density of states, eq 3 can be obtained:11

Ntrap(U) )

1 dQ q dU

(3)

where Q is the accumulated charge, Ntrap(U) is the density of trap states at potential U, and q is electron charge. Equation 3 clearly indicates that trap density is directly proportional to dQ/ dU, which provides a direct measurement of trap distribution. By differentiating the accumulated charge to the applied potential, a plot of dQ/dU against U is obtained and shown in

Figure 7. Cathodic charges accumulated at different potentials as derived by integrating the current-time curves in 0.2 mol L-1 LiClO4 at (a) pH 3.0 and (b) pH 6.8. The inset shows dQ/dU distribution against potential.

the inset of Figure 6b. This plot reflects the distribution of traps. It is seen that most traps are located at potentials around -0.7 V. A value of totally trapped electrons was calculated to be 1.21 × 1016 cm-2. As we will see below, the trap density is strongly pH dependent. Time-ResolWed Currents at pH 3.0 and 6.8. The current-time curves in the pH 3.0 and 6.8 solutions are similar to that in the pH 13.0 solution (not shown), and the cathodic charges at different potentials by integrating current-time curves measured in both solutions are shown in Figure 7a,b, respectively. Figure 7a,b shows the charge required to fill the traps, which is only about 1/10 and 1/3 than that required for the trap filling in the pH 13.0 solution. Again, a trap-filling process is observed at potentials negative of the level where most traps are located (around -0.25 V for pH 3.0 and -0.50 for pH 6.8). It is noticed that the trap-filling process is faster in the pH 3.0 solution and at pH 6.8 than that in the pH 13.0 solution. By differentiating the accumulated charge to the applied potential at pH 3.0 and 6.8, plots of dQ/dU against U are obtained and are shown in the insets of Figure 7a,b, respectively. Similar to the inset of Figure 6b, the inset plots of Figure 7a,b reflect the trap distribution at pH 3.0 and 6.8. Therefore, the total amount of trapped electrons were calculated to be 2.0 × 1015 cm-2 and 8.08 × 1015 cm-2, respectively. Effect of pH of the Electrolytes. The above estimation of trapping density leads us to consider the origin of traps, which has been related to the nanostructured semiconductor/electrolyte interface.24 Comparing trap densities at different pHs, it is noticed that much lower values are obtained in neutral and acid electrolytes. To get detailed information on this point, the cyclic voltammograms of nanostructured SrTiO3 electrodes were recorded at different pHs. The results are shown in Figure 8.

Transparent Nanostructured SrTiO3 Film Electrodes

J. Phys. Chem. C, Vol. 114, No. 2, 2010 819 electrochemistry. The results showed that trap-state densities are also highly pH dependent. Total trap state densities of 2.0 × 1015, 8.08 × 1015 and 1.21 × 1016 cm-2 were determined at pH 3.0, 6.8, and 13.0, respectively, with maxima located at -0.25, -0.50, and -0.70 V. The results obtained from cyclic voltammograms are in good agreement with those obtained from the measurements of time-resolved currents. The size of the peak at the potentials related to the trap states shown at -0.25, -0.51, and -0.72 for pH 3.0, 6.8, and 13.0, respectively, in the cyclic voltammograms, increases dramatically with increase in pH, indicating that traps are mostly surface-related.

Figure 8. Cyclic voltammograms of nanostructured SrTiO3 electrodes measured at different pH values. The electrodes were initially polarized at 0.8 V for 15 min before scanning, and the scan rate was 5 mV/s.

Cyclic voltammetry is a choice for detecting and characterizing surface traps in nanocrystalline electrodes. A shoulder in the current-potential curves that appears at potentials more positive than the conduction band edge has been generally assigned to the presence of electron traps.8,11 The conductionband edge of the nanostructured SrTiO3 electrode at pH 3.0 is approximately at -0.71 V (see Figure 5), and a feature at -0.25 V in the cyclic voltammogram corresponds to trap-state filling below the conduction band edge, confirming the presence of surface traps.5,25,26 To explain the peaks in Figure 8, eq 4 can be rewritten as11

Ntrap(U) )

1 dQ/dt 1 i ) q dU/dt q dU/dt

(4)

where i is the current density, t is time, and dU/dt is the scanning rate in the cyclic voltammograms. Equation 4 illustrates that the trap distribution at constant scanning rates is directly proportional to current density. Alternatively, current density distribution is a direct measure of trap density. Since the scanning rate was the same for all the experiments (5 mV/s), the curves in Figure 8 make it possible to compare the trap distribution at different pHs. The most trap state distributions at each pH are located at -0.25, -0.51, and -0.72 V at pH 3.0, 6.8, and 13.0, respectively, in good agreement with the results obtained from Figures 7a,b and 6b. It should be noticed that the size of the peak increases dramatically with increasing pH, so this strongly indicates that the traps are mostly surfacerelated. Conclusions Optical transparent nanostructured SrTiO3 film electrodes have been prepared. The band energetics of nanostructured SrTiO3 electrodes have been investigated with electrochemical and spectroelectrochemical methods. The flat band edges (Efb) of nanostructured SrTiO3 electrodes have been determined by using the spectroelectrochemical technique. The flat band edges of the nanostructured SrTiO3 electrodes strongly depend on the pH of electrolytes and shifts toward more negative potentials with the increase of pHs. The trap state distribution was investigated by the measurements of time-resolved current and

Acknowledgment. This work was supported financially by the National Natural Science Foundation of China (Grant No. 20773103), the Scientific Research Foundation for the Returned Overseas Chinese Scholars, State Education Ministry (2008101), the selected programs for scholars back from overseas, Ministry of Personnel (2006164), and the Science & Technology Program of Education Department of Henan (2008A150022). References and Notes (1) O’Regan, B.; Gratzel, M. Nature (London) 1991, 353, 737. (2) Gratzel, M. Inorg. Chem. 2005, 44, 6841. (3) Lagemaat, J.; Park, N. G.; Frank, A. J.; Boschloo, G. K.; Goossens, A. J. Phys. Chem. B 2000, 104, 2044. (4) Cinnsealach, R.; Boschloo, G.; Rao, S. N.; Fitzmaurice, D. Sol. Energy Mater. Sol. Cells 1999, 57, 107. (5) Boschloo, G.; Fitzmaurice, D. J. Phys. Chem. B 1999, 103, 2228. (6) Ferrere, S.; Zaban, A.; Gregg, B. A. J. Phys. Chem. B 1997, 101, 4490. (7) Enright, B.; Fitzmaurice, D. J. Phys. Chem. 1996, 100, 1027. (8) de la Garza, L.; Saponjic, Z. V.; Dimitrijevic, N. M.; Thurnauer, M. C.; Rajh, T. J. Phys. Chem. B 2006, 110, 680. (9) Haque, S. A.; Tachibana, Y.; Willis, R. L.; Moser, J. E.; Gratzel, M.; Klug, D. R.; Durrant, J. R. J. Phys. Chem. B 2000, 104, 538. (10) Tachibana, Y.; Haque, S. A.; Mercer, I. P.; Durrant, J. R.; Klug, D. R. J. Phys. Chem. B 2000, 104, 1198. (11) Wang, H.; He, J.; Boschloo, G.; Lindstrom, H.; Hagfeldt, A.; Lindquist, S. E. J. Phys. Chem. B 2001, 105, 2529. (12) Liu, C.; Bard, A. J. Phys. Chem. 1989, 93, 7749. (13) Burstein, E. Phys. ReV. 1952, 93, 632. (14) O’Regan, B.; Fitzmaurice, D.; Gratzel, M. J. Phys. Chem. 1991, 95, 10525. (15) Rothenberger, G.; Fitzmaurice, D.; Gratzel, M. J. Phys. Chem. 1992, 96, 5983. (16) Redmond, G.; Fitzmaurice, D. J. Phys. Chem. 1993, 97, 1426. (17) Redmond, G.; Gratzel, M.; Fitzmaurice, D. J. Phys. Chem. 1993, 97, 6951. (18) Enright, B.; Redmond, G.; Fitzmaurice, D. J. Phys. Chem. 1994, 98, 6195. (19) Flood, R.; Enright, B.; Allen, M.; Barry, S.; Dalton, A.; Doyle, H.; Fitzmaurice, D. Sol. Energy Mater. Sol. Cells 1995, 39, 83. (20) Burnside, S.; Moser, J. E.; Brooks, K.; Gratzel, M.; Cahen, D. J. Phys. Chem. B 1999, 103, 9328. (21) Srdic, V. V.; Djenadic, R. R. J. Optoelectron. AdV. Mater. 2005, 7, 3005. (22) Morrison, S. R. Electrochemistry at Semiconductor and Oxidized Metal Electrodes; Plenum Press: New York, 1980. (23) O’Regan, B.; Moser, J.; Anderson, M.; Gratzel, M. J. Phys. Chem. 1990, 94, 8720. (24) De Jongh, P. E.; Vanmaekelbergh, D. J. Phys. Chem. B 1997, 101, 2716. (25) Kavan, L.; Kratochvilova, K.; Gratzel, M. J. Electroanal. Chem. 1995, 394, 93. (26) Kavan, L.; Gratzel, M.; Rathousky, J.; Zukal, A. J. Electrochem. Soc. 1996, 143, 394.

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