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Properties of O2•- and OH• Formed in TiO2 Aqueous Suspensions by Photocatalytic Reaction and the Influence of H2O2 and Some Ions Tsutomu Hirakawa and Yoshio Nosaka* Department of Chemistry, Nagaoka University of Technology, Kamitomioka, Nagaoka, 940-2188, Japan Received November 18, 2001. In Final Form: January 23, 2002 We have investigated the effect of H2O2 on the behavior of O2•- and OH• produced in photocatalysis of aqueous TiO2 suspensions by means of luminol chemiluminescence probing and terephthalic acid fluorescence probing, respectively. The reduction of O2 by photoinduced conduction band electrons (e-) was increased by the addition of H2O2, since the consumption of photoinduced valence band holes (h+) in the oxidation of H2O2 caused the repression of e--h+ recombination. After the end of the light irradiation, the amount of O2•- decreased based on the fractal-like kinetics at the heterogeneous surface of the TiO2 particle. The decay process might be caused by trapped h+, which cannot react with water and then remains on the TiO2 particle after the irradiation. The energy level of the trapped h+ was estimated to be above the redox potential of SCN-, since it could react with the adsorbed H2O2 and I- ions but not with other ions such as SCN-, Br-, and Cl-. The formation rate of OH• was increased by the addition of H2O2, indicating the direct reduction of adsorbed H2O2 on the TiO2 surface. The quantum efficiencies of the formation of O2•and OH• were increased by 3.0 and 3.6 times by the addition of 0.2 mM H2O2.
Introduction The TiO2 photocatalytic reaction has been studied with much attention in recent years because it can decompose and mineralize pollutant and/or undesirable compounds in air and wastewater.1 In general, it is known that the TiO2 photocatalytic reactions proceed mainly by the contributions of active oxygen species, such as OH•, O2•-, and H2O2.2 These species are formed by the following schemes, where e- represents photoinduced conduction band electrons. The photoinduced valence band holes are trapped at the surface of TiO2 forming trapped holes that could not oxidize water. In the present study, the hole in the valence band state, h+vb, is distinguished from the trapped hole, h+tr.
TiO2 + hν f e- + h+vb h
+
+
vb
fh
(1) (1′)
tr
O2 + e- f O2•-
(2)
O2•- + O2•- + 2H+ f H2O2 + O2
(3)
•-
O2
+
+h
vb
f O2
(4)
O2•- + h+tr f O2 -
+
OH + h •
(4′) •
vb
f OH
•
(5)
OH + OH f H2O2
(6)
e- + h+tr f recombination
(7)
Since the reaction ability of OH• is high enough to attack any organic molecules, it has been assigned as a key species in the mineralization mechanism of many hazardous chemical compounds.3 Thus, to explore the TiO2 photocatalytic reactions, OH• has been investigated with several * To whom correspondence should be addressed. E-mail:
[email protected]. Fax: +81-258-47-9315.
methods: analysis of hydroxylated products from methanol4 and UV absorption analysis with 2,3-dihydroxybenzoic acid5 as well as the radical detection methods by electron spin resonance spectroscopy.6 In recent years, Ishibashi et al. have developed a fluorescence technique with terephthalic acid to study the photocatalysis of TiO2 thin films in aerated aqueous solution and demonstrated that this method can detect OH• selectively.7 Then, it is interesting to apply this method to quantify OH• which was formed photocatalytically in a TiO2 aqueous suspension. The behaviors of O2•- in a TiO2 aqueous suspension9-12 and on TiO2 thin film13-15 have been recently studied by (1) (a) Fujishima, A.; Rao, T. N.; Tryk, D. A. J. Photochem. Photobiol. C 2000, 1, 1-21. (b) Fujishima, A.; Hashimoto, K.; Watanabe, T. Photocatalysis; BKC Inc.: Tokyo, 1999. (c) Photocatalysis: Science and Technology; Kaneko, M., Okura, I., Eds.; Kodansha-Scientific: Tokyo, in press. (2) (a) Hoffmann, M. R.; Martin, S. T.; Choi, W.; Bahnemann, D. W. Chem. Rev. 1995, 95, 69-96. (b) Serpone, N.; Pelizzetti, E.; Hidaka, H. Photocatalytic Purification and Treatment of Water and Air; Ollis, D. F., Al-Ekabi, H., Eds.; Elsevier: London, 1993; pp 225-250. (3) (a) O’Shea, K. E.; Pernas, E.; Saiers, J. Langmuir 1999, 15, 20712076. (b) Lindner, M.; Theurich, J.; Bahnemann, D. W. Water Sci. Technol. 1997, 35, 79-86. (c) Bekbolet, M.; Lindner, M.; Weichgrebe, D.; Bahnemann, D. W. Sol. Energy 1996, 56, 455-469. (d) Mills, A.; Davies, R. H.; Worsley, D. Chem. Soc. Rev. 1993, 417-425. (4) Sun, L.; Bolton, R. J. J. Phys. Chem. 1996, 100, 4127-4134. (5) Dai, Q.; Wang, D.; Yuan, C. Supramol. Sci. 1998, 5, 469-473. (6) (a) Nosaka, Y.; Kishimoto, M.; Nishino, J. J. Phys. Chem. B 1998, 102, 10279-10283. (b) Grela, M. A.; Coronel, M. E.; Colussi, A. J. J. Phys. Chem. 1996, 100, 16940-16946. (c) Schwarz, P. F.; Turro, N. J.; Bossmann, S. H.; Braun, A. M.; Abdel Wahab, A. A; Durr, H. J. Phys. Chem. B 1997, 101, 7127-7134. (d) Brezova, V.; Stasko, A. J. Catal. 1994, 147, 156-162. (7) Ishibashi, K.; Fujishima, A.; Watanabe, T.; Hashimoto, K. J. Photochem. Photobiol., A 2000, 134, 139-142. (8) Ishibashi, K.; Fujishima, A.; Watanabe, T.; Hashimoto, K. Electrochem. Commun. 2000, 2, 207-210. (9) Nosaka, Y.; Yamashita, Y.; Fukuyama, H. J. Phys. Chem. B 1997, 101, 5822-5827. (10) Hirakawa, T.; Nakaoka, Y.; Nishino, J.; Nosaka, Y. J. Phys. Chem. B 1999, 101, 4399-4403. (11) Hirakawa, T.; Kominami, H.; Ohtani, B.; Nosaka, Y. J. Phys. Chem. B 2001, 105, 6993-6999. (12) Nosaka, Y. In Photocatalysis: Science and Technology; Kaneko, M., Okura, I., Eds.; Kodansha-Scientific: Tokyo, in press. (13) Ishibashi, K.; Nosaka, Y.; Hashimoto, K.; Fujishima, A. J. Phys. Chem. B 1998, 102, 2117-2120.
10.1021/la015685a CCC: $22.00 © 2002 American Chemical Society Published on Web 03/21/2002
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means of the luminol chemiluminescence (CL) probe method developed by the present authors. Since the reduction of O2 by e- (eq 2) and oxidation of O2•- by h+ (eqs 4 and/or 4′) might determine the quantum efficiency of TiO2 photocatalysts under irradiation,16 the reduction of O2 is important process to determine the efficiency of TiO2 photocatalytic reactions. To investigate the role of O2•-, in the present study we measured the absolute amount of O2•- and compared it with that of OH•. The decay profile of O2•- was also measured to explore the effect of additives on the reaction of trapped h+. In the photocatalytic reactions, H2O2 would be formed in both processes of reduction and oxidation, that is, dimerization of OH• (eq 6), disproportionation of O2•- (eq 3), and reduction of O2•-. Then, H2O2 might be taken into account to understand the details in TiO2 photocatalysis. In the literature, it has been suggested that H2O2 is reduced photocatalytically by O2•- or photoinduced e- to give OH•17-19 while it is also oxidized by OH• or photoinduced h+ to give O2•-.17-21 However, the contribution of these reactions to photocatalytic reactions at the TiO2 surface remains unclear.22,23 Therefore, to know how the active oxygen species influence the mechanism of TiO2 photocatalysis, it is very important to observe directly the effect of H2O2 on TiO2 photocatalysis and to investigate the behavior of H2O2 in the photocatalytic reaction. Thus, in the present paper we measured the formation and decay of O2•- in the presence of H2O2 and other ions by means of the luminol CL probe method. In addition, terephthalic acid (TA) was applied, for the first time, to TiO2 aqueous suspensions for measuring the amount of OH• radicals. These experimental observations clarified the effect of H2O2 and the role of trapped holes in the photocatalytic reactions. Experimental Procedure Materials. Degussa P25 (Japan Aerosil) of TiO2 powder used in this study was a generous gift from the manufacturer. This photocatalyst has a Brunauer-Emmett-Teller (BET) surface area of 50 m2 g-1, a primary particle size of about 32 nm, and a crystal structure of 80% anatase and 20% rutile.10 Luminol and terephthalic acid (Nacalai Tesque, Ltd.), superoxide dismutase (SOD) (4400 units/mg, Tokyo Kasei, Ltd.), hemoglobin (Hb) from bovine (Wako Pure Chemical Industries, Ltd.), 2-bromoterephthalic acid (Aldrich, Ltd.), NaI, KSCN, KCl (Nacalai Tesqu, Ltd.), and KBr (Kantou Kagaku. Ltd.) were used (14) Ishibashi, K.; Fujishima, A.; Watanabe, T.; Hashimoto, K. J. Phys. Chem. B 2000, 104, 4934-4938. (15) Ishibashi, K.; Fujishima, A.; Watanabe, T.; Hashimoto, K. Electrochemistry 2001, 69, 160-164. (16) Courne, J. C.; Colussi, J. A.; Hoffmann, R. M. J. Phys. Chem. B 2001, 105, 1351-1354. (17) (a) Pichat, P.; Guillard, C.; Amalric, L.; Renard, A.-C.; Plaidy, O. Sol. Energy Mater. Sol. Cells 1995, 38, 391-399. (b) Harvey, P.; Rudham, R.; Ward, S. J. Chem. Soc., Faraday Trans. 1 1983, 79, 13811390. (c) Al-Ekabi, H.; Serpone, N. J. Phys. Chem. 1988, 92, 57265731. (18) (a) Doong, A.-R.; Chang, W.-H. J. Photochem. Photobiol., A 1997, 107, 239-244. (b) Wei, T.-Y.; Wan, C.-C. J. Photochem. Photobiol., A 1992, 69, 241-249. (c) Wei, T.-Y.; Wang, Y.-Y.; Wan, C.-C. J. Photochem. Photobiol., A 1990, 55, 115-126. (19) (a) Fujihira, M.; Satoh, Y.; Osa, T. Bull. Chem. Soc. Jpn. 1982, 55, 666-671. Okamoto, K.; Yamamoto, Y.; Tanaka, H.; Tanaka, H. Bull. Chem. Soc. Jpn. 1985, 58, 2015-2022. (c) Izumi, I.; Fan, F.-R. F.; Bard, A. J. J. Phys. Chem. 1981, 85, 218-223. (20) Rabani, J.; Yamashita, K.; Ushida, K.; Stark, J.; Kira, A. J. Phys. Chem. B 1998, 102, 1689-1695. (21) Sun, L.; Bolton, J. R. J. Phys. Chem. 1996, 100, 4127-4134. (22) (a) Matthew, R. W. J. Chem. Soc., Faraday Trans. 1 1984, 80, 457-471. (b) Chen, T.-F.; Doong, R.-A.; Lei, W.-G. Water Sci. Technol. 1998, 37, 189-194. (23) (a) Ohno, T.; Kigoshi, T.; Nakabeya, K.; Matsumura, M. Chem. Lett. 1998, 877-878. (b) Jia, J.; Ohno, T.; Masaki, Y.; Matsumura, M. Chem. Lett. 1999, 963-964. (c) Jia, J.; Ohno, T.; Masaki, Y.; Matsumura, M. Chem. Lett. 2000, 908-909.
Hirakawa and Nosaka without further purification. The water used was distilled followed by purification with a Milli-Q system. 2-Hydroxyterephthalic acid (TAOH) was synthesized according to the report.24 The purity of the synthesized TAOH was evaluated from the absorbance at 312 nm using a UV-vis spectrophotometer (Shimadzu UV-2500PC) with the absorption coefficient of 4000 M-1 cm-1.25 Luminol Chemiluminescence Probe. Fifteen milligrams of TiO2 powder was suspended in a 1 cm × 1 cm Pyrex glass cell containing 3.5 mL of 0.01 M NaOH aqueous solution of pH 11.5. The sample cell was placed on a magnetic stirrer in a dark box, and the suspension was stirred for 10 min before the measurement. In some cases, H2O2 or another reagent such as NaI, KSCN, KCl, or KBr was added into the TiO2 suspension before starting the stir, which was continued during the measurement. The light source for the excitation of TiO2 was a 150 W Xe lamp (Hamamatsu Photonics, C2499). The excitation wavelength was confined to 387 ( 11 nm by using a set of glass filters of U330 and L39 (HOYA). The intensity of incident light was measured with a photometer (Advantest, TQ8210) to be 40 mW cm-2. Luminol was dissolved into a 0.01 M NaOH aqueous solution to prepare a 7 mM aqueous solution. The essential points in the procedure of the luminol CL probe method are as follows. Fifty microliters of luminol solution was injected with a microsyringe into the TiO2 suspension after the end of the irradiation. At the same time as the injection, the luminol CL intensity was monitored with a photon-counting photomultiplier tube (PMT) (Hamamatsu, R2949) for 20 s. To measure the decay profile of active species, the injection of luminol was delayed after stopping the irradiation. In the other experiments, 50 µL of an aqueous solution of 3 mg of SOD was added into the TiO2 suspension before the injection of luminol. In this case, SOD was added into the TiO2 suspension 5 s before the injection of luminol. Other details in the experimental procedure have been described previously.9,10 To convert the observed CL intensity into the absolute concentration of the O2•- radicals in the TiO2 suspension, the apparatus factor was measured using a standard CL reaction. In this measurement, the sample cell containing 3.4 mL of Kolthoff buffer solution at pH 11.5 with 1.0 µM luminol was placed on the magnetic stirrer as described above. Similar to the procedure described above, 15 mg of TiO2 powder was suspended and stirred in the solution, but it was not irradiated by light. After opening the shutter for the PMT detector, 50 µL of 5 mM H2O2 solution was injected into the TiO2 suspension. When 50 µL of 6.2 µM Hb solution was injected into the sample at 30 s after the addition of H2O2, a strong CL of luminol was observed. The CL intensity was recorded for 15 min. Following the 15 min observation, the injection of Hb solution and CL observation were repeated three times to ensure the complete utilization of luminol molecules. In a separate experiment, it was confirmed that the CL intensity is proportional to the amount of luminol used. Terephthalic Acid Fluorescence Probe. It is known that OH• reacts with TA and generates TAOH which emits fluorescence at around 426 nm on the excitation of its own 312 nm absorption band.24,25
OH• + TA f TAOH
(8)
The measurements of the amount of OH• were performed for the TiO2 photocatalytic reaction by means of this TA fluorescence (FL) probe method as follows. An aqueous solution containing 0.01 M NaOH and 0.5-4 mM TA was prepared, and then 15 mg of TiO2 powder was suspended in 3.5 mL of this solution in a 1 cm × 1 cm Pyrex glass cell. The sample cell was placed on the magnetic stirrer in the dark box, and the suspension was stirred for 10 min. In some cases, H2O2 was added into the TiO2 suspension before the stirring. The excitation light source was the same as that in the luminol CL measurement. These experiments were performed for each sample irradiated at various periods, that is, 3, 5, 7, and 10 min. One-half gram of KCl was (24) Mason, J. T.; Lorimer, P. J.; Bates, M. D.; Zhao, Y. Ultrason. Sonochem. 1994, S91, 1. (25) Armstrong, A. W.; Facey, A. R.; Grant, W. D.; Humphreys, G. W. Can. J. Chem. 1963, 41, 1575-1577.
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Figure 1. Concentration of O2•- formed with (O) and without ()) 0.2 mM H2O2 in the P25 TiO2 aqueous suspension (15 mg/ 3.5 mL) as a function of irradiation time.
Figure 2. Concentration of O2•- formed in the irradiated TiO2 suspension (15 mg/3.5 mL) with various concentrations of H2O2. The irradiation duration was 100 s.
added into the TiO2 suspension after the irradiation, and the cell was kept in a refrigerator under dark for 24 h to sediment TiO2 powder from the suspension. Fluorescence spectra of the supernatant liquid were measured with a fluorescence spectrophotometer (Shimadzu RF-5300PC). By the addition of KCl, most TiO2 particles precipitated and then the light scattering became negligible to measure the fluorescence spectra. Since TAOH molecules were not adsorbed on the TiO2 particle in the presence of KCl at pH 11.5, the supernatant solution is supposed to contain all the TAOH molecules produced. By comparison of the fluorescence intensity with that of the known concentration of TAOH, the amount of produced TAOH was determined. The amount of OH• formed in TiO2 photocatalysis was estimated from that of TAOH by adopting the trapping factor of 80%.7
luminol solution was injected 10 s thereafter. If the observed CL actually originates from the reaction of O2•with luminol, CL would not be observed since SOD can scavenge O2•-. In this experiment, the CL was not observed. In addition, the lifetime of the active oxygen species is too short to be that of H2O2. Therefore, we concluded that the CL observed in the present experiment with 0.2 mM H2O2 is attributable to the reaction with O2•-. Consequently, it was proved that the ordinate in Figure 2 is convincing for the concentration of O2•-. The quantum efficiency (φsuperoxide) of O2•- formation can be estimated from the initial slope in Figure 1 divided by the amount of incident photon flux. The values are calculated to be about 0.13% and 0.4% without and with 0.2 mM H2O2, respectively. That is, φsuperoxide was increased by about 3 times on the addition of 0.2 mM H2O2. Effects of Additional Ions. O2•- is formed during the light irradiation on TiO2 photocatalysts, while simultaneously it decays by the reaction of h+vb and/or h+tr (eq 4 and/or 4′). Then, the existence of additional ions that can scavenge the holes would influence the amount of produced O2•-. By changing the redox potentials of the additives, therefore, the energy level of effective holes might be expected. Figure 3a shows a change of the amount of O2•as a function of the concentration of four additional ions. The amount of produced O2•- was decreased by the addition of these ions and reached a steady value at about 0.3 mM. Figure 3b shows the ratio of decreased O2•- concentration as a function of the redox potential of four additional ions, which are listed in Table 1. This figure clearly shows that the decrease in the O2•- concentration became drastically large by the addition of SCN- and I- ions, while it was small for Br- and Cl- ions. Decay Process of O2•-. To investigate the decay process of the photocatalytically formed O2•-, the addition of the luminol solution was delayed from 0.5 to 600 s after the end of the 100 s irradiation. Figure 4a represents the logarithmic plot for the concentration of O2•- observed in the absence of H2O2 as a function of the delay time. This semilogarithmic plot shows multiple linear correlations, and the reciprocal plot of the concentration for checking second-order kinetics was not linear. Thus, O2•- decays
Results Luminol Chemiluminescence Probe. O2•- Formation. The concentration of O2•- was calculated from the CL intensity observed on the addition of luminol after the irradiation since it is proportional to the amount of O2•-.9 In Figure 1, the concentrations of O2•- are plotted as a function of the duration of the irradiation. As the period of the irradiation increased, the amount of O2•- increased and reached a steady value at about 100 s. Therefore, hereafter the irradiation period of 100 s was used in the present luminol CL probe experiment. Effects of H2O2. Figure 2 shows the plot for the concentration of O2•- produced with 100 s irradiation as a function of the concentration of H2O2 in the reaction system. The amount of O2•- increased with the increasing concentration of H2O2 at first, and then it reached a maximum value at 0.2 mM H2O2. On the further addition of H2O2, the amount of O2•- decreased and reached a steady value at about 0.4 mM H2O2. On the basis of this observation, hereafter the H2O2 concentration of 0.2 mM was adopted in the luminol CL probe experiments. Since two-electron-oxidized luminol can emit CL by reacting with H2O2,9,26 contribution of this reaction to the observed CL intensity was checked as follows: Fifty microliters of SOD solution was added into the TiO2 suspension after the end of the irradiation, and then (26) Merenyi, G.; Lind, J.; Eriksen, E. T. J. Phys. Chem. 1990, 94, 748-752.
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Figure 3. (a) The concentration of produced O2•- is plotted as a function of the concentration of the three additional ions, (4) Br-, ()) SCN-, and (O) I-. (b) The ratio (%) of O2•- concentration decreased by the addition of 0.5 mM ions is plotted as a function of the redox potential of the ions. The arrows A and B indicate the redox potential of trapped holes, h+tr (1.5 V) (ref 27b), and that of OH•/OH- (1.9 V) (ref 28), respectively, and arrow C is the valence band potential of TiO2 (2.39 V) calculated from Ecb ) -0.13-0.059 pH (ref 29) with a pH of 11.5. Table 1. Redox Potential of Additional Ions and Parameters in Equation 12 for O2•- Decay redox potential/ V (vs NHE)a
additives non H2O2 NaI KSCN KBr KCl a
O2•-, H+/HO2-
1.00
I•/I-
1.35 1.64 2.00 2.47
SCN•/SCNBr•/BrCl•/Cl-
concn/ mM
[O2•-]0/ µM
β
0.03 0.2 0.5 0.5 0.5 0.5
0.92 2.00 3.25 0.52 0.66 1.10 b
0.69 0.74 0.82 0.80 0.64 0.69 b
Reference 28. b Not measured.
likely obeying a two-step pseudo-first-order rate kinetics. A similar decay profile for O2•- has been observed in the study for TiO2 thin films,13 where the h+tr remaining on the TiO2 surface is reported to eliminate O2•-. Then, similar
Figure 4. Time dependence of the O2•- concentration after stopping the irradiation on the TiO2 suspension without (a) and with (b) 0.2 mM H2O2. The solid line shows the curve fitted by eq 12.
to the case of the thin films, O2•- in the TiO2 suspension may decay by reaction 4′ even after the irradiation has been stopped. In the presence of 0.2 mM H2O2, the concentration of O2•- changed with time as shown in Figure 4b. The decay profile of O2•- seems to be of the two-step pseudo-firstorder rate kinetics even in the presence of H2O2, but the two decay rates became slow as compared with those in Figure 4a. The decay process of O2•- was examined in the presence of the additional ions. The decay profile showed the multistep processes even when the additional ions are present in the TiO2 suspension. Although SCN- and Brdid not have much influence on the decay processes of O2•-, significant elongation was observed for I-. This observation is consistent with the above explanation that the energy level of the h+tr locates between the redox potentials of SCN- and I- ions and it is consumed to oxidize I- ions during the light irradiation. Since the fraction of the two decay components did not change much in the presence of the additives, it may be
Properties of O2•- and OH• in TiO2 Suspensions
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Figure 6. The effect of the concentration of TA on the formation rate of TAOH calculated from the growth of the fluorescence intensity for the supernatant liquid of the TiO2 suspension.
Figure 5. (a) Fluorescence spectra obtained for the supernatant liquid of the irradiated TiO2 suspension containing 3 mM terephthalic acid at various irradiation periods and (b) the time dependence of the fluorescence intensity at 426 nm.
difficult to assume the existence of two kinds of h+tr which cause different decay rates. In a later section, the decay will be analyzed taking into account the heterogeneous situation. Terephthalic Acid Fluorescence Probe. OH • Probing. Figure 5a shows FL spectra observed for the supernatant solution of the TiO2 suspension containing 3 mM TA irradiated for various duration times. Since the observed fluorescence spectra are identical to that of TAOH, it is concluded that TAOH was generated from TA by the reaction with OH• (eq 8), where OH• was formed in the TiO2 photocatalysis. Figure 5b represents the fluorescence intensity as a function of the duration of irradiation. Since the fluorescence intensity increased linearly with the irradiation time, the formation rate of TAOH was calculated from this slope. Figure 6 shows the production rate of TAOH as a function of the concentration of TA. If the amount of TA is enough to react rapidly with all OH• radicals, the rate of TAOH formation is determined by that of OH• formation. Then, Figure 6 indicates that OH• can be quantitatively detected by more than 2 mM TA. Hereafter, 3 mM was adopted for TA in the TA-FL probe method. Thus, from Figure 5b, the growth rate of OH• was calculated to be 0.06 µM min-1.
Figure 7. The rate of OH• formation for the irradiated P25 TiO2 aqueous suspension is plotted as a function of the concentration of H2O2.
Effect of H2O2. When H2O2 was added to the TiO2 suspension containing the TA probe, the peak position and spectral shape in the obtained fluorescence spectra were not changed. This observation suggests that TA can probe OH• radicals even when H2O2 is present in the suspension. Figure 7 shows that the formation rate of OH• radicals increased linearly with the amount of H2O2 below 0.3 mM and then it reached a constant value of about 0.25 µM min-1. Provided that TiO2 absorbs the whole incident light, the quantum efficiencies of OH• without and with 0.2 mM H2O2 are calculated to be 2.8 × 10-5 and 1.0 × 10-4, respectively. The actual quantum efficiencies, however, may be larger than these values since the reflection of light by TiO2 in water is about 11%, and the excitation wavelength was so broad that all the light may not be absorbed by TiO2 particles. In any case, it can be concluded that the production rate of OH• increased by 3.6 times with the addition of 0.2 mM H2O2 in the TiO2 suspension.
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Discussion Luminol Chemiluminescence Probe. O2•- Formation and the Effect of H2O2. It was confirmed, by the experiment of the addition of SOD, that the amount of O2•- could be determined from the luminol CL intensity even in the presence of H2O2. Then, the increase of O2•with a small amount of H2O2 in Figure 2 is attributable to the oxidation of H2O2 as described by the following equations.
H2O2 + OH• + OH- f O2•- + 2H2O
(9)
HO2- + OH• f O2•- + H2O
(10)
H2O2 + h+ + 2OH- f O2•- + H2O
(11)
Although the molecular ratio HO2-/H2O2 is 0.6 in the alkaline solution of pH 11.5, HO2- is likely to be negligible in the reaction scheme. The negatively charged HO2- ions may not locate around the negatively charged TiO2 surface, where the ζ potential has been measured to be about -110 mV at this pH.10 However, since the formation rate of OH• was increased by the addition of H2O2 as shown in Figure 7, both of the reactions 9 and 10 consuming OH• are unlikely processes for the formation of O2•-. If O2•- is formed by reaction 11, the amount of O2•- would be changed with the concentration of H2O2 obeying the absorption isotherm since H2O2 has been reported to adsorb on the TiO2 surface in the alkaline solution of pH 11.30 However, as shown in Figure 2, the amount of O2•- was decreased by the addition of a higher amount of H2O2. Then, reaction 11 may not be the dominant formation process of O2•since the formed O2•- likely suffers further oxidation (eqs 4 and 4′). The remained possibility to explain the increase of O2•- is that the reduction of O2 (eq 2) was accelerated by the presence of H2O2 as an oxidizable species to diminish the e--h+ recombination (eq 7). The decrease of the amount of O2•- in Figure 2 at more than 0.2 mM H2O2 may be explained by blocking O2 by the adsorbed H2O2 as will be discussed later. Effect of Additional Ions on the O2•- Formation. In the present experiment, the effect of the addition of Cl-, Br-, SCN-, and I- ions was examined, and then the amount of O2•- was decreased with the addition of these ions. It is known that Cl- ions accelerate the charge carrier recombination under irradiation33 and I- and Br- ions scavenge the photoinduced h+.36 In the present study, however, Cl- ions did not significantly influence the formation process of O2•-. The explanation for this result is that the valence band potential of TiO2 (arrow C in Figure 3b) at pH 11.5 is more negative than the oneelectron oxidation potential of Cl- ions. Similar to the Cl(27) (a) Serpone, N.; Lawless, D.; Terziom, R.; Meisel, D. In Electrochemistry in Colloids and Dispersions; Mackay, R. A., Texter, J., Eds.; VCH: New York, 1992; Chapter 30, pp 399-416. (b) Lawless, D.; Serpone, N.; Meisel. D, J. Phys. Chem. 1991, 95, 5166. (28) Wardman, P. J. Phys. Chem. Ref. Data 1989, 18, 1637-1755. (29) Duonghong, D.; Ramsden, J.; Graetzel, M. J. Am. Chem. Soc. 1982, 104, 2977-2985. (30) Fujihira, M.; Muraki, H. Chem. Lett. 1986, 2001-2002. (31) Ferradini, C.; Foos, J.; Houee, C.; Pucheault, J. Photochem. Photobiol. 1978, 28, 697-700. (32) Melhuish, W. H.; Sutton, H. C. J. Chem. Soc., Chem. Commun. 1978, 970-971. (33) Herrmann, H. M.; Hoffmann, R. M. J. Chem. Soc., Faraday Trans. 1994, 90, 3323-3330. (34) Bahnemann, D. W.; Memming, R.; Higendroff, M. J. Phys. Chem. B 1997, 101, 4265-4275. (35) Kamat, P. V. Langmuir 1985, 1, 608-611. (36) Herrmann, J. M.; Pichat, P. J. Chem. Soc., Faraday Trans. 1 1980, 76, 1138-1146.
ions, Br- ions also did not influence the formation and decay process of O2•-. Since the redox potential of Br- ions locates close to that of OH- (arrow B) and more positive than that of h+tr (arrow A), Br- could not reduce h+tr and then the recombination reaction (eq 7) is still dominant. SCN- ions are oxidized by photoinduced h+, and then the product, SCN•, reacts with the SCN- ion to form (SCN)2•-. Then, since the lifetime of (SCN)2•- is reported to be 0.03-40 µs from laser photolysis,34,35 no influence is caused by (SCN)2•- on the production of O2•- after the end of the irradiation. In addition, no reaction of (SCN)2•with O2•- was reported so far. Therefore, SCN- ions might serve as a recombination center since the steady amount of O2•- was decreased by the addition of SCN- ions. Similar to the SCN- ions, I- ions are oxidized to produce I•, which reacts with I- in a diffusion-controlled process forming I2•-.35,38 However, I2•- cannot react with O2•although I2, which is formed by the dimerization of I•, can react with O2•-.37 Then, in separate experiments the presence of I2 was checked by the iodine-starch method. Since I2 was not detected by the test, it was not produced by photocatalytic oxidation of I- ions in the present experimental conditions. Consequently, the additional ions used in this study act only as the charge carrier recombination center under irradiation but not as the reaction intermediates. The redox potential of h+tr is reported to be 1.5 V as that of adsorbed OH• in an acidic condition (Ti-O•H/Ti-OH).27b This redox potential is likely to be applicable in the present study at pH 11.5 where the redox couple of Ti-O•/Ti-Omay be adopted, since the reduction potential of O•- is close to that of OH• in solution.28 Then, in the present experiment, the reported value of the redox potential for h+tr was verified as observed in Figure 3b. Decay Process of O2•-. Two-step pseudo-first-order rate kinetics in the O2•- decay has been suggested by Figure 4 in which the logarithm of O2•- concentration was plotted. Although in the previous study we reported that the lifetime of O2•- is about 17 s,9 the present detailed study revealed the existence of the slower decay. In general, it is known that O2•- in homogeneous bulk solution decays by disproportionation (eq 3). Since the concentration has been measured to be 1.2 µM in the absence of H2O2 at the irradiation time of 100 s, the halflifetime of O2•- at pH 11.5 is calculated to be about 11.3 h from the reported rate constants. On the other hand, the observed decay rate is not so slow as shown in Figure 4a. Even in the slow decay region, the lifetime was 182 s. Thus, the disproportionation, at least in bulk solution, may not be responsible for the observed decay of O2•-. In the presence of 0.2 mM H2O2, the lifetime of O2•- of the slow decay was about 990 s. This value of the lifetime is too small to be attributed to the reaction with HO2-, which is calculated to be 1.8 h in the bulk solution. Thus, neither the disproportionation nor the reaction with H2O2 is responsible for the decay of O2•-. Therefore, O2•- in an aqueous TiO2 suspension likely decays only by h+tr as suggested by Ishibashi et al. for TiO2 films13-15 or some other adsorbed species. Then, how can we explain the two-decay processes for O2•-? One possibility is that there are two kinds of h+tr having different oxidation abilities and one of them disappears for 3 s as the initial fast decay of O2•-. Another possibility is that the slow decay is caused by the reaction with O2•(37) Schwarz, H. A.; Bielski, B. H. J. J. Phys. Chem. 1986, 90, 14451448. (38) Kormann, C.; Bahnemann, D. W.; Hoffmann, M. R. J. Phys. Chem. B 1988, 92, 5196-5201.
Properties of O2•- and OH• in TiO2 Suspensions
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adsorbed on the surface of the TiO2 particle. The experimental results for the effect of additives could be used to identify the counterpart in the reaction of O2•- after the irradiation. By the addition of ions that can scavenge h+tr, the O2•- decay showed similar multiexponential properties as described above in the experimental result. SCN- and Br- ions did not influence the decay process of O2•-, though the concentration of produced O2•- has been decreased to a certain extent as shown in Figure 3b. This experimental result can be explained by the increase in the adsorption of formed O2•- on the surface of the TiO2 particle. On the other hand, I- ions decreased the decay rate of O2•- in the fast region. This observation is explained by the decreases in the amount of h+tr because the added I- ions can react with the h+tr. The decay rate in the hundreds-of-seconds region also became slower with I- ions, where the O2•concentration was not altered much. If the slow decay occurs by the reaction with surface adsorbed O2•-, the decay rate depends on the O2•- concentration because of the probable adsorption equilibrium. Therefore, the slow decay process is not attributable to the reaction with adsorbed O2•- on the TiO2 surface. Consequently, it is reasonable to assume that the decay process of the whole time range examined in the present study might be attributed to a certain single process, that is, the reaction between O2•- and the h+tr. In the presence of 0.2 mM H2O2, the decay process became significantly slower as shown in Figure 4b similar to the case of the addition of I- ions. This is due to the reaction of h+tr with the adsorbed H2O2, since the redox potential of H2O2 at pH 11.5 is about 1 V (vs NHE) and negative against 1.5 V of h+tr. In Figure 4a,b, the decay profiles did not fit to a linear line. Since the reciprocal plot of the O2•- concentration did not show a linear correlation either, the decay process of O2•- is not obeying completely both the first- and the second-order kinetics. Since the above discussion led to the conclusion that h+tr is responsible for the decay process of O2•-, the observed decay indicates that the rate constant should be time dependent. This may be caused by the migration of O2•- from the formation site to the h+tr site over the heterogeneous surface of the TiO2 particle. To account for the heterogeneous surface reaction, we adopted a fractal distribution theory,39 which is characterized by the nonexponential decay and random work model. This theory has been widely used to analyze e--h+ recombination for photoluminescence40 and transition photocurrent41 observed in TiO2. According to the fractal-like kinetics, the following equation can be assumed.
[O2•-] ) [O2•-]0 t β-1
(12)
where [O2•-] 0 and [O2•-] are the concentrations of O2•- at the end of the irradiation and at the delay time t, respectively, and β (0 e β e 1) is the extent of rate-constant distribution. When β is 1, classical rate kinetics can be applied. Namely, [O2•-] becomes equal to [O2•-]0, which means that the rate constant is independent of time. In this case, the decay rate should be expressed by k[O2•-][h+tr] with a rate constant k, and eq 12 is not needed to fit the experimental data. On the other hand, when β comes near 0, the rate constant has a distribution, since the migration of O2•- molecules on the surface controlled the reaction. Note that the β does not have a physicochemical (39) Kopelman, R. Science 1988, 241, 1620-1626. (40) Fujihara, K.; Izumi, S.; Ohno, T.; Matsumura, M. J. Photochem. Photobiol., A 2000, 132, 99-104. (41) Nelson, J. Phys. Rev. B 1999, 59, 15374-15380.
implication but merely means the degree of distribution in the rate constant. Equation 12 was fitted to the experimental data as shown in Figure 4a,b. The two parameters, thus obtained [O2•-]0 and β, are shown in Table 1 besides those for other additives where the data were also fitted in a fairly good precision. The value of β was increased from 0.69 to about 0.8 by some additives such as H2O2 and I- ions. This increase indicates that H2O2 and I- ions prevent the migration of O2•- and disturb the reaction between O2•and h+tr. Thus, the above observation suggests that the O2•molecules are adsorbed on the TiO2 surface and the decay of O2•- occurs by the reaction with h+tr (eq 4′). This reaction with h+tr is likely to be dominant relative to that with h+vb (eq 4) even under irradiation, since the trapping hole (eq 1′) occurs faster than the formation of O2•- (eq 2).2 Terephthalic Acid Fluorescence Probe. The amount of OH• produced in the photoirradiated TiO2 suspension was estimated by measuring the amount of TAOH, which is generated by the reaction of OH• with TA. In the alkaline solution of pH 11.5, OH• is partly dissociated since pKa is 11.8. Then, the molar ratio of O•-/OH• is calculated to be 0.45. TAOH would be produced from O•- as well as OH• because O•- reacts rapidly with water to form OH•. Then, the ionic dissociation may not affect the amount of TAOH, and therefore O•- is not distinguished from OH• in the present report. Two kinds of OH• are suggested in the literature; the adsorbed OH• and the free OH•. In the previous section, it was suggested that h+tr is equivalent to the adsorbed OH• and the redox potential is 1.5 V.27b Since the redox potential of free OH• is 1.9V,28 it could be distinguished from h+tr. Then, we suggest that TA could probe only the free OH• and it cannot react with h+tr or the adsorbed OH•. The photo-Fenton reaction that is initiated by the excitation of H2O2 cannot occur since the excitation wavelength is confirmed to 387 nm. Therefore, in the present experimental conditions, OH• radicals are formed by photocatalytic reduction of the H2O2 adsorbed on the TiO2 particles as follows.
H2O2 + O2•- f OH• + OH- + O2
(13)
H2O2 + e- f OH• + OH-
(14)
As shown in Figure 7, when H2O2 is present in the TiO2 suspension, the formation rate of OH• is increased. If OH• was formed by the Harber-Weiss reaction (eq 13), the formation rate of OH• is calculated to be 0.006 µM min-1 by using the reaction rate32 of 0.23 M-1 s-1 and the O2•concentration of about 3.5 µM in Figure 1. Since the observed formation rate for OH• was 0.22 µM min-1, reaction 13 is not responsible for the increase in the rate by the addition of H2O2. The reduction of H2O2 to form OH• by photoinduced e(eq 14) was reported in photoelectrochemical studies. In general, the irradiated TiO2 has an ability to reduce H2O2 into OH• since the reduction potential of H2O2/OH• is about 0.87 V (vs NHE).28 Therefore, our experimental observation in Figure 7 leads to the conclusion that in the aqueous TiO2 suspension OH• was formed by the photocatalytic reduction of H2O2 (eq 14) as well as the oxidation of OH(eq 5). The formation rate of OH• was increased by the addition of H2O2, and the data profile in Figure 7 seems to represent the shape of the adsorption isotherm. Therefore, the
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Hirakawa and Nosaka
Figure 8. Energy level diagram showing the role of H2O2 and trapped holes, h+tr, in the photocatalytic reactions at photoirradiated TiO2 particles (R, recombination; R-D, recombination was decreased; A, acceleration).
decomposition rate of H2O2 by TiO2 photocatalysis was influenced by the amount of adsorbed H2O2. Iliz et al. reported that the kinetics of H2O2 decomposition in TiO2 photocatalysis extended from Langmuir-Hinshelwood to Freundrich kinetics.42 Although they obtained the best fitting curve for the Freundrich kinetics, neither the Langmuir kinetics nor Freundrich kinetics fits to the data in Figure 7. Thus, in the present experimental conditions, many reaction processes such as oxidations (eqs 9 and 11) and reductions (eqs 13 and 14) may be involved to a certain extent. Figure 2 shows that the amount of O2•- was decreased with the addition of H2O2 at the concentration above 0.2 mM, while the formation rate of OH• was increased with the addition of the same amount of H2O2 as shown in Figure 7. Further addition of H2O2 up to 0.35 mM caused the decrease of the amount of O2•- in contrast to the increase in the formation rate of OH•. Thus, the decrease of O2•- is explained by the decrease in the amount of adsorbed O2 which is caused by the adsorption of H2O2. Figure 8 summarizes the proposed reaction mechanism in the present study with an energy diagram. In the presence of 0.2 mM H2O2, the amounts of O2•- and OH• were increased to about 3.6 and 0.37 µM, respectively. The increment of OH• by the addition of H2O2 was smaller than that of O2•-, indicating that the reduction of O2 by e- (eq 2) was accelerated by the direct oxidation of the adsorbed H2O2 (eq 11) where both reactions produce O2•-. Parallel to the direct oxidation, the direct reduction of adsorbed H2O2 (eq 14) to produce OH• also occurs to a certain extent, since the production of OH• was increased despite the possible decrease in the OH- oxidation by competing with the H2O2 oxidation. The adsorbed H2O2 can release electrons to fill the h+tr, and thus the decay of O2•- by h+tr after the irradiation became slow in the presence of H2O2. The P25 TiO2 used in the present study is composed of two types of crystals, namely, 80% anatase and 20% rutile. (42) Iliz, I.; Foglein, K.; Dombi, A. J. Mol. Catal. A: Chem. 1998, 135, 55-61.
The different crystalline types may cause different behavior in the production and decay processes of O2•and OH• as well as the energy level of h+tr. Investigation of the difference in the anatase-rutile structure is important and is now in progress. In the present study, however, the obtained result might be that of rutile crystals since the excitation wavelength was 387 nm which is absorbed mainly by rutile crystals. Conclusions In the present work, the following behaviors of O2•- and OH• formed in TiO2 photocatalysis were revealed by means of luminol chemiluminescence and terephthalic acid fluorescence probe methods, respectively. The reduction of O2 by e- was accelerated by the oxidation of H2O2 based on the decrease of e--h+ recombination. When the amount of adsorbed H2O2 was so large that the adsorption of O2 was restrained, the production of O2•- was decreased. The O2•- formed in the irradiated TiO2 suspension decays by the reaction with h+tr remaining on TiO2 particles after stopping the irradiation. Then, the observed decay could be fitted by the equation based on the fractal-distribution theory. The OH• formed by photocatalysis in the TiO2 suspension could be probed by the terephthalic acid. The production rate of OH• was significantly increased with the addition of H2O2. We confirmed that the increase is responsible for the direct reduction of the adsorbed H2O2 on the TiO2 surface. That is, the Fenton reaction by e- occurs, and the Harber-Weiss reaction was calculated to be negligible in the production rate of OH•. Acknowledgment. We thank Dr. K. Takenaka and Dr. T. Kobayashi for their helpful technical suggestions and Mr. M. Abe for his technical assistance in TAOH synthesis. The present work is partly defrayed by the Grant-in-Aid for Scientific Research on Priority-AreaResearch of “Photo-Functional Interfaces” from the Japanese Ministry of Science, Education and Culture. LA015685A
