Formation of H2O2 in TiO2 Photocatalysis of Oxygenated and

Publication Date (Web): April 16, 2014. Copyright © 2014 American ... Photocatalytic Oxidation of Glucose into Formate on Nano TiO2 Catalyst. Binbin ...
5 downloads 0 Views 1MB Size
Download PDF
Article pubs.acs.org/JPCC

Formation of H2O2 in TiO2 Photocatalysis of Oxygenated and Deoxygenated Aqueous Systems: A Probe for Photocatalytically Produced Hydroxyl Radicals Veronica Diesen and Mats Jonsson* School of Chemical Science and Engineering, Applied Physical Chemistry, KTH Royal Institute of Technology, SE-100 44 Stockholm, Sweden ABSTRACT: The formation of H2O2 in oxygenated and deoxygenated aqueous solutions using immobilized TiO2 illuminated by black light (365 nm) was studied to verify the presence of hydroxyl radicals in TiO2 photocatalysis. In oxygen containing systems, formation of H2O2 proceeds through reduction of molecular oxygen by conduction band electrons or by recombination of hydroxyl radicals. In oxygen free solutions recombination of hydroxyl radicals constitutes the only pathway to H2O2 formation. Detection of H2O2 in absence of oxygen therefore serves as an indicator for hydroxyl radical formation. The H2O2 concentration was determined using the Ghormley triiodide method. It was found that a significant amount of H2O2 was produced in the deoxygenated aqueous solutions supporting the hypothesis of hydroxyl radical production in photocatalysis. To further elucidate the origin of the H2O2, experiments using the radical scavenger tris(hydroxymethyl)aminomethane (Tris) were conducted. The results showed that the H2O2 concentration increased in the oxygenated system as Tris protects the H2O2 from decomposition by hydroxyl radicals. In the deoxygenated system, no H2O2 could be detected due to hydroxyl radical scavenging by Tris, which prevents H2O2 formation. The results presented support the hypothesis that the hydroxyl radical is the primary oxidant in aqueous TiO2 photocatalysis.

1. INTRODUCTION Heterogeneous photocatalysis on TiO2 in aqueous systems has long been assumed to proceed via the formation of hydroxyl radicals formed upon oxidation of adsorbed hydroxide or water by the initially formed positive hole.1−5 Despite the numerous experimental observations claimed to confirm the production of hydroxyl radicals,6−10 there is still some controversy regarding the nature of the primary product responsible for oxidation of solute molecules.11−14 A number of chemical probes have been used to quantify the photocatalytic efficiency and also to indirectly identify the primary reactive species.10,15−17 In principle, two approaches can be used to quantify the photocatalytic efficiency. Either the probe concentration or the product concentration can be monitored as a function of illumination time. Monitoring the product concentration is a prerequisite for a sensitive quantitative method.15 The sensitivity is attributed to the detection of a unique product and the fact that fairly high probe concentrations can be used. The latter implies quantitative scavenging of the primary oxidant. Heterogeneous photocatalysis of aerated aqueous solutions results in the formation of hydrogen peroxide via reduction of molecular oxygen by the electrons excited to the conduction band. This process is illustrated in reactions 1−3. O2 +

− eCB

→

O•− 2

(2)

HO•2 + HO•2 → H 2O2 + O2

(3)

Another potential route for formation of hydrogen peroxide is by recombination of hydroxyl radicals at the surface of the photocatalyst. This process is the only possible pathway for production of hydrogen peroxide in oxygen free systems. Hence, formation of hydrogen peroxide under oxygen free conditions could serve as an indicator for the formation of hydroxyl radicals in photocatalysis. The reactivity of hydrogen peroxide toward metal oxide surfaces has received increasing attention during the past decade.18−21 The main reason for this is the importance of such processes in nuclear technological installations.22 These studies have shown that hydrogen peroxide undergoes catalytic decomposition, primarily to form adsorbed hydroxyl radicals, on oxide surfaces of the type MxOy.23 The overall process is described by reactions 4 and 5. H 2O2 (aq) ⇌ H 2O2 (ads)

(4)

H 2O2 (ads) → 2HO•(ads)

(5)

Received: January 10, 2014 Revised: April 16, 2014 Published: April 16, 2014

(1) © 2014 American Chemical Society

+ • O•− 2 + H → HO2

10083

dx.doi.org/10.1021/jp500315u | J. Phys. Chem. C 2014, 118, 10083−10087

The Journal of Physical Chemistry C

Article

film was prepared in the same way using a transparent glass support to allow an absorption spectrum to be recorded. The area and thickness of the deposited TiO2 layer was 31.2 cm2 and 100 μm, respectively. The crystallographic phase of the film was studied by X-ray diffraction (PANanalytical X’pert instrument using Cu Kα irradiation). Data was collected over the range 5° < 2θ < 65°. A scanning electron microscopy (SEM) image of the film was taken using a JEOL JSM-6301F field emission SEM at an accelerating voltage of 5 kV. The TiO2 absorption spectrum was recorded using a JASCO V-630 spectrophotometer. Experiments with H2O2 were performed in a closed reaction beaker. The lid was made of glass and transparent to the wavelengths used in this study. The immobilized TiO2 film was placed at the bottom of the reaction beaker facing the solution and the entering light. All solutions used for the experiments in this work were prepared with Millipore Milli-Q water. The Ghormley triiodide method was used to measure the H2O2 concentration. In this method I− is oxidized to I3− in the presence of H2O2. The absorbance of the product, I3− is subsequently measured spectrophotometrically at 360 nm. A linear correlation between the absorbance and the H2O2 concentration was obtained by a calibration curve, where the I3− absorbance was plotted as a function of the H2O2 concentration in the range 0.1−500 μM. The H2O2 solutions for the calibration curve were prepared from a 30% Merck standard solution. N2 and O2 gas were provided by AGA Gas AB. A gas flow rate of 0.2 L min−1 was used for the experiments. The solutions were purged with N2 for 30 min prior to the experiment to reduce the oxygen content to a minimum. The solutions were stirred during experiments to maintain a homogeneous solution and to facilitate mass transfer to the surface. To determine the total amount H2O2 in solution, a sample volume of 0.2 mL was extracted and diluted in 1.6 mL of H2O and mixed with 0.1 mL of 1 M KI and 0.1 mL of 1 M HAc/NaAc (including a few drops of ammonium dimolybdate (ADM) acting as a catalyst for the oxidation of I− by H2O2). The HAc/NaAc is added to adjust the pH. The solution was left to react for 2 min before measuring the absorbance at 360 nm. To determine the total amount of H2O2 in the system, that is, in the solution and adsorbed on the TiO2, the reagents were added to the reaction beaker instead (with the same ratio as previously described) and left to react before measuring the absorbance. The light source consisted of two 20 W black light blue lamp tubes (Philips Blacklight) with a wavelength range measured to 345−385 nm with a peak at 365 nm. The total incident photonic flux was continuously monitored at the irradiation distance with a photometer (International light, Inc.) during the experiments and was measured to 0.10 ± 0.01 mW cm−2.

The efficiency in decomposing hydrogen peroxide depends on the nature of the metal oxide.23,24 Experimental and theoretical studies have shown that the adsorption energies for hydrogen peroxide as well as the hydroxyl radical are of key importance.23,24 Interestingly, TiO2 has been shown to be a very poor catalyst for decomposition of H2O2, and the affinity for hydroxyl radicals is much lower than for other metal oxides.23 This has been confirmed in experimental studies as well as in theoretical studies using density functional theory (DFT).23,24 Low surface affinity of the hydroxyl radical implies higher reactivity toward solutes, this being a prerequisite for an efficient photocatalyst. The overall chemistry of an oxygen-free aqueous photocatalytic system is often described by the following reactions, 6−11. + H 2O + hVB → H+ + OH•

(6)

+ OH− + hVB → OH•

(7)

OH• + OH• → H 2O2

(8)

OH• + H 2O2 → H 2O + HO•2

(9)

− eCB + H 2O2 → OH• + OH−

(10)

HO•2 + HO•2 → H 2O2 + O2

(11)

Under continuous illumination, the above mechanism (reactions 1−10) leads to steady-state where the rate of hydrogen peroxide production is equal to the rate of consumption. The photooxidation of TiO2 adsorbed H2O molecules and OH− ions as a source of adsorbed H2O2 has been the subject of debate during the past decade.11,25,26 It has recently been claimed that the photooxidation of TiO2 adsorbed hydroxide ions or water molecules (reactions 6 and 7) is thermodynamically and kinetically hindered, proposing as alternative mechanisms of water photooxidation the nucleophilic attack mechanism25,27 and the redox photooxidation mechanism.26 Although both mechanisms have important differences, both consider doubly coordinated terminal oxygen ions, known as bridging oxygen ions, as primary trapping sites of photogenerated holes, proposing that primary intermediate species of water photooxidation to H2O2 and O2 are structural radicals instead of adsorbed radicals.25−27 It has also been claimed that bridging oxygen atoms are incorporated into H2O2 and O2 photogenerated species, which means that bridging oxygen vacancies are induced during water photooxidation reaction, although they are rapidly healed via dissociative adsorption of water molecules at sites.26 The oxidizing role of ions in TiO2 photocatalysis has recently been confirmed via isotopic tracing experiments involving the photooxidation of H216O dissolved C6H6 by using Ti18O2 as catalyst.28 In this work we have measured the hydrogen peroxide concentration as a function of illumination time for oxygen-free and air- and oxygen-saturated aqueous solutions. Experiments using radical scavengers were carried out to elucidate the origin of hydrogen peroxide.

3. RESULTS AND DISCUSSION The photocatalyst used for this study was a screen-printed immobilized TiO2 film. A scanning electron microscope image of the surface is presented in Figure 1 and shows a very porous film consisting of particles of between 20 and 80 nm. Similar TiO2 films have previously been shown to be highly photocatalytically active.15 The absorbance spectrum of the film is shown in Figure 2. The film displays typical TiO2 absorption characteristics. The X-ray diffraction pattern of the TiO2 film shows reflections at 25.3, 38.6, 48.1, 53.9, 55.1, and 62.7° 2θ (Figure 3) corresponding to the anatase crystallographic phase. Anatase

2. EXPERIMENTAL SECTION Materials, Reagents, and Methods. Immobilized TiO2 was used for the experiments in this work. A TiO2 paste consisting of 13 nm particles was screen-printed onto a titanium support and heat-treated at 500 °C for 1 h. Another 10084

dx.doi.org/10.1021/jp500315u | J. Phys. Chem. C 2014, 118, 10083−10087

The Journal of Physical Chemistry C

Article

Figure 4. Formation of H2O2 from the TiO2 film in 20 mL of deoxygenated (N2), oxygenated (O2), and air-saturated aqueous solutions using black light (0.10 ± 0.01 mW cm−2).

Figure 1. Scanning electron micrograph of the TiO2 film prepared by screen-printing using a Solaronix commercial paste.

As can clearly be seen, the H2O2 concentration increases with time for about 20 min when it reaches a constant level. This level corresponds to the steady-state where the rate of H2O2 production is equal to the rate of consumption. The main route for hydrogen peroxide production in the oxygen containing systems is the formation of superoxide via trapping of the electron by molecular oxygen, according to reactions 1−3. The higher the oxygen content, the more efficient is the trapping of the electron. Consequently, it is hardly surprising that the steady-state concentration of H2O2 in the photocatalytic system depends on the oxygen concentration in aqueous solution. What is perhaps more surprising is the fact that there is a significant production of hydrogen peroxide also in the deoxygenated system. However, the deoxygenation by purging with N2 is not believed to render completely oxygen free solutions. Trace amounts of molecular oxygen will still be present and act as electron scavenger. It is also quite clear that there is no direct proportionality between the oxygen concentration and the steady-state concentration, that is, the H2O2 steady-state concentration is not five times higher for the system purged with 1 bar O2 compared to the system exposed to 1 bar air. The steady-state concentrations of H2O2 in the oxygen free and in the oxygen containing systems are given by eqs 12 and 13 based on the mechanism above.

Figure 2. Absorption spectrum for the screen-printed TiO2 film.

Figure 3. XRD pattern of the TiO2 film annealed at 500 °C. The anatase peak positions are displayed in brackets.

[H 2O2 ]SS =

k 8[OH•]2 k 9[OH•] + k10[e−]

(12)

[H 2O2 ]SS =

k 8[OH•]2 + k 3[HO•2 ]2 k 9[OH•] + k10[e−]

(13)

As the concentrations of all the radical species in the steadystate expressions above depend on the rate of primary photolysis as well as on the concentrations of O2 and H2O2, a direct quantitative comparison between the two equations is not straightforward. It should be noted that the effective rate of oxidant production increases with increasing electron scavenging capacity of the system, that is, formation of species consuming hydrogen peroxide is enhanced if oxygen is present or if the hydrogen peroxide concentration is increased. To determine whether the H2O2/TiO2 system is in equilibrium during continuous illumination, the hydrogen peroxide concentration in solution and in the whole system (solution and TiO2 surface) was measured under steady state conditions for the three systems. The results are presented in Figure 5.

is usually identified as the predominant phase of TiO2 when annealed at 500 °C and is also preferred for photocatalysis.2,29 The adsorption of H2O2 to the TiO2 surface used in these studies has been quantified and reported recently.30 The equilibrium constant for adsorption was determined to KH2O2−TiO2 = (6 ± 1) × 104 M−1 on the basis of the Langmuir isotherm.30 Given the fact that H2O2 adsorbs to the TiO2 surface, the fraction of adsorbed H2O2 must be accounted for when examining the total production of H2O2. In Figure 4, the solution concentration of H2O2 is plotted as a function of illumination time for three different systems: O2-saturated, airsaturated, and N2-purged. 10085

dx.doi.org/10.1021/jp500315u | J. Phys. Chem. C 2014, 118, 10083−10087

The Journal of Physical Chemistry C

Article

hydrogen peroxide concentration is reduced to below the detection limit. For the air saturated system, addition of Tris has the opposite effect. These observations are perfectly in line with the hypothesis that the hydroxyl radical is the primary oxidant and that formation of hydrogen peroxide in deoxygenated systems can primarily be attributed to recombination of hydroxyl radicals.

4. CONCLUSIONS Formation of H2O2 was used as a probe to evaluate the presence of hydroxyl radicals as the primary oxidizing species in aqueous TiO2 photocatalysis. The concentration of H2O2 was measured in oxygenated and deoxygenated aqueous systems using a porous, immobilized TiO2 film and black light (365 nm) illumination. It was found that the formation of H2O2 increases in all systems until a steady-state is reached where the rate of H2O2 formation is equal to the rate of consumption. The steady-state level depends on the O2 concentration in the solution and increases with increasing O2 content, however no direct proportionality between the O2 concentration and the steady-state concentration was found. The formation rate of species consuming H2O2 was enhanced as the rate of oxidant formation was increased. At the steady-state level, the amount of adsorbed H2O2 onto the TiO2 surface increases with a decreased total amount of H2O2 in the system, which agrees with the adsorption isotherm. As a significant amount of H2O2 was detected in the deoxygenated aqueous solutions, experiments with the hydroxyl radical scavenger Tris were performed to elucidate the origin of the H2O2. The results showed that the formation of H2O2 was significantly suppressed in the presence of Tris in deoxygenated solutions due to the hydroxyl radical scavenging effect of Tris. In oxygen-containing systems, enhanced H2O2 concentrations were observed as Tris protects the H2O2 from attack by hydroxyl radicals, which reduces the rate of H2O2 consumption. These results support the hypothesis that the hydroxyl radical is the primary oxidant in TiO2 photocatalysis.

Figure 5. Concentration H2O2 in solution and adsorbed on the TiO2 surface in solutions saturated with air, N2, and O2 after 60 min black light illumination.

As can be seen, the fraction of H2O2 adsorbed to the TiO2 surface is significant for all three systems. The fraction of adsorbed H2O2 increases with decreasing total amount of H2O2, as expected from the adsorption isotherm. The relative fractions are in good agreement with the equilibrium constant and the obvious conclusion is that the system is equilibrated throughout the illumination process. To further test the hypothesis of the two processes behind formation of H2O2, we performed experiments with and without tris(hydroxymethyl)aminomethane (Tris). Tris is an efficient scavenger for hydroxyl radicals that has been used to probe the efficiency of photocatalysis in recent papers15,30−32 and it has also been used to verify the existence of the hydroxyl radical in catalytic decomposition of H2O2 on metal oxide surfaces.23,24,33 The use of Tris as a probe for surface bound hydroxyl radicals was recently evaluated in detail.34 The hydroxyl radical abstracts hydrogen atoms from Tris. There are several abstractable hydrogen atoms and abstraction from some sites yield formaldehyde as a stable product. Production of formaldehyde has been used to probe hydroxyl radicals formed upon catalytic decomposition of H2O2 at oxide surfaces as well as the photocatalytic efficiency. The purpose of performing these experiments here is to be able to confirm the existence of the hydroxyl radical as the primary oxidizing species in aqueous TiO2 photocatalysis. The anticipated effect of adding Tris to the system is a reduction in the H2O2 steadystate concentration in the deoxygenated system where the major part of the H2O2 production can be attributed to hydroxyl radical recombination. For the oxygen containing systems, the opposite effect is anticipated since Tris will protect the hydrogen peroxide from attack by the hydroxyl radicals and thereby reduce the rate of H2O2 consumption. This can also be seen from the steady-state expressions for the two cases (eqs 11 and 12). The results are summarized in Table 1. As can be seen in the table, the effect of adding the hydroxyl radical scavenger is very clear. In the deoxygenated system where Tris will scavenge hydroxyl radicals, the steady-state

■

*Tel.: +46 8790 9123. Fax: +46 8790 8772. E-mail: matsj@kth. se. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS Wallenius Water AB is gratefully acknowledged for the financial support of the research described here.

■

amt of H2O2 detected (μM) atmosphere

H2O

200 mM Tris

0.97 ± 0.06 0.31 ± 0.01

1.99 ± 0.02 0 ± 0.001

REFERENCES

(1) Daneshvar, N.; Rabbani, M.; Modirshahla, N.; Behnajady, M. A. Photooxidative Degradation of Acid Red 27 in a Tubular ContinuousFlow Photoreactor: Influence of Operational Parameters and Mineralization Products. J. Hazard. Mater. 2005, 118, 155−160. (2) Hoffmann, M. R.; Martin, S. T.; Choi, W.; Bahnemann, D. W. Environmental Applications of Semiconductor Photocatalysis. Chem. Rev. 1995, 95, 69−96. (3) Fox, M. A.; Dulay, M. T. Heterogeneous Photocatalysis. Chem. Rev. 1993, 93, 341−357. (4) Sauer, T.; Cesconeto Neto, G.; José, H. J.; Moreira, R. F. P. M. Kinetics of Photocatalytic Degradation of Reactive Dyes in a TiO2 Slurry Reactor. J. Photochem. Photobiol., A 2002, 149, 147−154.

Table 1. Amount of H2O2 Formed in H2O and in Aqueous Tris Solutions under Air and N2 Saturation

air N2-saturated

AUTHOR INFORMATION

Corresponding Author

10086

dx.doi.org/10.1021/jp500315u | J. Phys. Chem. C 2014, 118, 10083−10087

The Journal of Physical Chemistry C

Article

(5) Saien, J.; Soleymani, A. R. Degradation and Mineralization of Direct Blue 71 in a Circulating Upflow Reactor by UV/TiO2 Process and Employing a New Method in Kinetic Study. J. Hazard. Mater. 2007, 144, 506−512. (6) Ceresa, E. M.; Burlamacchi, L.; Visca, M. An ESR Study on the Photoreactivity of TiO2 Pigments. J. Mater. Sci. 1983, 18, 289−294. (7) Jaeger, C. D.; Bard, A. J. Spin Trapping and Electron Spin Resonance Detection of Radical Intermediates in the Photodecomposition of Water at Titanium Dioxide Particulate Systems. J. Phys. Chem. 1979, 83, 3146−3152. (8) Xiang, Q.; Yu, J.; Wong, P. K. Quantitative Characterization of Hydroxyl Radicals Produced by Various Photocatalysts. J. Colloid Interface Sci. 2011, 357, 163−167. (9) Ishibashi, K.-i.; Fujishima, A.; Watanabe, T.; Hashimoto, K. Quantum Yields of Active Oxidative Species Formed on TiO2 Photocatalyst. J. Photochem. Photobiol., A 2000, 134, 139−142. (10) Ishibashi, K.-i.; Fujishima, A.; Watanabe, T.; Hashimoto, K. Detection of Active Oxidative Species in TiO2 Photocatalysis Using the Fluorescence Technique. Electrochem. Commun. 2000, 2, 207−210. (11) Salvador, P. On the Nature of Photogenerated Radical Species Active in the Oxidative Degradation of Dissolved Pollutants with TiO2 Aqueous Suspensions: A Revision in the Light of the Electronic Structure of Adsorbed Water. J. Phys. Chem. C 2007, 111, 17038− 17043. (12) Valencia, S.; Catano, F.; Rios, L.; Restrepo, G.; Marín, J. A New Kinetic Model for Heterogeneous Photocatalysis with Titanium Dioxide: Case of Non-Specific Adsorption Considering Back Reaction. Appl. Catal., B 2011, 104, 300−304. (13) Stafford, U.; Gray, K. A.; Kamat, P. V. Photocatalytic Degradation of Organic Contaminants: Halophenols and Related Model Compounds. Heterog. Chem. Rev. 1996, 3, 77−104. (14) Linsebigler, A. L.; Lu, G.; Yates, J. T., Jr Photocatalysis on TiO2 Surfaces: Principles, Mechanisms, and Selected Results. Chem. Rev. 1995, 95, 735−758. (15) Diesen, V.; Jonsson, M. Tris(hydroxymethyl)aminomethane as a Probe in Heterogeneous TiO2 Photocatalysis. J. Adv. Oxid. Technol. 2012, 15, 392−398. (16) Mills, A.; Wang, J.; McGrady, M. Method of Rapid Assessment of Photocatalytic Activities of Self-Cleaning Films. J. Phys. Chem. B 2006, 110, 18324−18331. (17) Houas, A.; Lachheb, H.; Ksibi, M.; Elaloui, E.; Guillard, C.; Herrmann, J.-M. Photocatalytic Degradation Pathway of Methylene Blue in Water. Appl. Catal., B 2001, 31, 145−157. (18) Rasti, N.; Toyserkani, E.; Ismail, F. Chemical Modification of Titanium Immersed in Hydrogen Peroxide using Nanosecond Pulsed Fiber Laser Irradiation. Mater. Lett. 2011, 65, 951−954. (19) Soleymani, M.; Moheb, A.; Babakhani, D. Hydrogen Peroxide Decomposition over Nanosized La1−xCaxMnO3 (0 ≤ x ≤ 0.6) Perovskite Oxides. Chem. Eng. Technol. 2011, 34, 49−55. (20) Pehrman, R.; Amme, M.; Roth, O.; Ekeroth, E.; Jonsson, M. Oxidative Dissolution of Actinide Oxides in H2O2 Containing Aqueous Solution−A Preliminary Study. J. Nucl. Mater. 2010, 397, 128−131. (21) Park, J.-N.; Shon, J. K.; Jin, M.; Hwang, S. H.; Park, G. O.; Boo, J.-H.; Han, T. H.; Kim, J. M. Highly Ordered Mesoporous α-Mn2O3 for Catalytic Decomposition of H2O2 at Low Temperatures. Chem. Lett. 2010, 39, 493−495. (22) Wada, Y.; Uchida, S.; Nakamura, M.; Akamine, K. Empirical Understanding of the Dependency of Hydrogen Water Chemistry Effectiveness on BWR Designs. J. Nucl. Sci. Technol. 1999, 36, 169− 178. (23) Lousada, C. M.; Johansson, A. J.; Brinck, T.; Jonsson, M. Mechanism of H2O2 Decomposition on Transition Metal Oxide Surfaces. J. Phys. Chem. C 2012, 116, 9533−9543. (24) Lousada, C. M.; Johansson, A. J.; Brinck, T.; Jonsson, M. Reactivity of Metal Oxide Clusters with Hydrogen Peroxide and Water−A DFT Study Evaluating the Performance of Different Exchange−Correlation Functionals. Phys. Chem. Chem. Phys. 2013, 15, 5539−5552.

(25) Imanishi, A.; Okamura, T.; Ohashi, N.; Nakamura, R.; Nakato, Y. Mechanism of Water Photooxidation Reaction at Atomically Flat TiO2 (Rutile) (110) and (100) Surfaces: Dependence on Solution pH. J. Am. Chem. Soc. 2007, 129, 11569−11578. (26) Salvador, P. Mechanisms of Water Photooxidation at n-TiO2 Rutile Single Crystal Oriented Electrodes under UV Illumination in Competition with Photocorrosion. Prog. Surf. Sci. 2011, 86, 41−58. (27) Nakamura, R.; Nakato, Y. Primary Intermediates of Oxygen Photoevolution Reaction on TiO2 (Rutile) Particles, Revealed by in Situ FTIR Absorption, and Photoluminescence Measurements. J. Am. Chem. Soc. 2004, 126, 1290−1298. (28) Montoya, J. F.; Ivanova, I.; Dillert, R.; Bahnemann, D. W.; Salvador, P.; Peral, J. Catalytic Role of Surface Oxygens in TiO2 Photooxidation Reactions: Aqueous Benzene Photooxidation with Ti18O2 under Anaerobic Conditions. J. Phys. Chem. Lett. 2013, 4, 1415−1422. (29) O’Neill, S. A.; Clark, R. J.; Parkin, I. P.; Elliott, N.; Mills, A. Anatase Thin Films on Glass from the Chemical Vapor Deposition of Titanium (IV) Chloride and Ethyl Acetate. Chem. Mater. 2003, 15, 46−50. (30) Diesen, V.; Jonsson, M. Effects of O2 and H2O2 on TiO2 Photocatalytic Efficiency Quantified by Formaldehyde Formation from Tris(hydroxymethyl)aminomethane. J. Adv. Oxid. Technol. 2013, 16, 16−22. (31) Diesen, V.; Dunnill, C. W.; Ö sterberg, E.; Parkin, I. P.; Jonsson, M. Silver Enhanced TiO2 Thin Films: Photocatalytic Characterization using Aqueous Solutions of Tris(hydroxymethyl)aminomethane. Dalton Trans. 2014, 43, 344−351. (32) Diesen, V.; Jonsson, M.; Parkin, I. P. Improved Texturing and Photocatalytic Efficiency in TiO2 Films Grown Using Aerosol-Assisted CVD and Atmospheric Pressure CVD. Chem. Vap. Deposition 2013, 19, 355−362. (33) Lousada, C. M.; Jonsson, M. Kinetics, Mechanism, and Activation Energy of H2O2 Decomposition on the Surface of ZrO2. J. Phys. Chem. C 2010, 114, 11202−11208. (34) Yang, M.; Jonsson, M. Evaluation of the O2 and pH Effects on Probes for Surface Bound Hydroxyl Radicals. J. Phys. Chem. C 2014, 118, 7971−7979.

10087

dx.doi.org/10.1021/jp500315u | J. Phys. Chem. C 2014, 118, 10083−10087