Evidence for H2O2 Generation during the TiO2-Assisted

Department of Chemistry & Biochemistry, Concordia UniVersity, Montreal, Canada H3G 1M8. ReceiVed: December 8, 1998; In Final Form: March 30, 1999...
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J. Phys. Chem. B 1999, 103, 4862-4867

Evidence for H2O2 Generation during the TiO2-Assisted Photodegradation of Dyes in Aqueous Dispersions under Visible Light Illumination Taixing Wu, Guangming Liu, and Jincai Zhao* Institute of Photographic Chemistry, Chinese Academy of Sciences, Beijing 100101, China

Hisao Hidaka Frontier Research Center for the Earth EnVironment Protection, Meisei UniVersity, 2-1-1 Hodokubo, Hino, Tokyo 191, Japan

Nick Serpone Department of Chemistry & Biochemistry, Concordia UniVersity, Montreal, Canada H3G 1M8 ReceiVed: December 8, 1998; In Final Form: March 30, 1999

Photodegradation of a series of dyes (Rhodamine B, Orange II, Sulfo-rhodamine B, Fluorescein, Alizarin red, Squarylium cyanine, and Eosin) in the presence of TiO2 particles under air-equilibrated controlled conditions and visible light illumination led to the formation of hydrogen peroxide. Combined with chemical oxygen demand (COD) measurements and in situ adsorption of added H2O2, the results reveal that H2O2 is a perfidious intermediate species in the process of dye photodegradation. The formation rate of H2O2 depends on the rate of dye degradation. It can also be decomposed by TiO2 particles under visible light irradiation because of a TiO2/H2O2 surface complex. H2O2 was detected because its formation rate was greater than its decomposition rate; failing this, H2O2 would not be observed. Its decomposition rate depends on the amount of substrates or intermediates formed during dye degradation and adsorbed on the surface of the TiO2 particles. The greater the degradation rate of the dye or the greater the quantity of substrates or intermediates adsorbed on the TiO2 surface, the less is the depletion of H2O2 with the consequence that a large quantity of H2O2 accumulates in the solution bulk permitting its facile detection. Also, pH appears to play a beneficial role in the observation of H2O2 formed.

Introduction Much attention has been focused in the past two decades on the photocatalytic degradation of organic pollutants mediated by TiO2 particles in aqueous dispersions irradiated by UV light irradiation. The large body of evidence that has been collected suggests that this method is undoubtedly a potential and effective approach toward the degradation or mineralization of a wide variety of harmful/toxic organic pollutants in wastewaters and toward the purification of drinking water.1-6 Mechanistically, it is now commonly accepted that the photocatalyst TiO2 is first excited by UV light and subsequently initiates the photodegradation processes. However, artificial UV light sources tend to be somewhat expensive, and the UV light component in sunlight reaching the surface of the earth and available to excite TiO2 is a relatively small component (ca. 3-5%) of the AM1 solar spectrum. Therefore, recent efforts have been focused on exploring means to utilize the inexpensive visible light sources more effectively or to use the inexhaustible sunlight for treating polluted waters. Dyestuffs represent a class of organic pollutants that absorb visible light. Electron transfer processes occurring between dyes and semiconductors, especially TiO2, have been examined and found to have practical potential.7-9 Kamat and co-workers10 reported the photodegradation of Acid Orange 7 and Naphthol Blue Black dyes preadsorbed on the surface of TiO2 particles, * Correspondence author fax, +86-10-6487-9375; e-mail, [email protected].

whereas Ross et al.11 have examined the degradation of terbutylazine under visible irradiation on TiO2 particles sensitized by Rose Bengal. Recently we reported the photodegradation of several dyes under exposure to visible light in aqueous TiO2 dispersions.12-14 The visible irradiation mechanism described by eqs 1-6 has been shown to be different from the UV irradiation pathway as portrayed previously;14a the dye and not the semiconductor TiO2 is excited by Visible light:

dye + hV f dye* +•

dye* + TiO2 f dye

+ TiO2 (e)

(1) (2)

TiO2 (e) + O2 f TiO2 + O2-•

(3)

O2-• + TiO2 (e) + 2H+ f H2O2

(4)

H2O2 + TiO2 (e) f •OH + OH-

(5)

dye+• + O2 (or O2-• or •OH) f peroxylated or hydroxylated intermediates f f degraded or mineralized products (6) Subsequently, TiO2 plays the pivotal role of an electron carrier leading to separation of injected electrons and dye cation radicals. The processes described by eqs 1-5, e.g., electron transfer,7-9 generation of the superoxide radical anion,14c H2O2,14c

10.1021/jp9846678 CCC: $18.00 © 1999 American Chemical Society Published on Web 05/22/1999

H2O2 Generation during Photodegradation of Dyes

and •OH radical formation14b,c have all been demonstrated earlier. Equation 6 is supported by recent evidence.15 Therefore, TiO2-assisted photoprocesses provide an attractive route to treat or pretreat dye pollutants using either artificial visible light or sunlight. To the extent that studies involving visible light irradiation to mediate the degradation of pollutants by this advanced oxidation technology are less reported than others, more extensive and detailed investigations seem appropriate and necessary to gain a further understanding of these processes. On the basis of the visible light irradiation mechanisms (eqs 1-6), H2O2 should be generated. In fact, when H2O2 formation was probed, it was surprising that not in all cases was the photogenerated H2O2 observed. To determine the causes of H2O2 generation, we systematically examined what decisiVe factors might affect the formation and thus observation of H2O2 in TiO2 dispersions during the photodegradation of various dyes under visible light irradiation. UV-visible spectra and chemical oxygen demand (COD) of the degraded solution were also examined to disclose some of the details of the H2O2 formation process. The results provide a determining contribution to our understanding of the photoassisted and TiO2-mediated pathway for the degradation of dyes under visible irradiation. Experimental Section Materials. TiO2 nanoparticulates (P25, ca. 80% anatase, 20% rutile; BET area, ca. 50 m2 g-1) were kindly supplied by Degussa Co. Horseradish peroxidase (POD) was purchased from Huamei Biologic Engineering Co. (China), and the N,Ndimethyl-p-phenylenediamine (DPD) reagent from Merck (p.a.). The dyes (see structures above) Orange II (Baker Co.), Rhodamine B (Beijing Chemicals Co.), Alizarin red (Beijing Chemicals Co.), Eosin (Beijing Chemicals Co.), and other chemicals used in the experiments were all of analytical reagent grade quality and used without further purification. The Sulforhodamine B dye was laser grade quality (Across Co.). The Squarylium cyanine dye was synthesized according to a method reported earlier.16 Deionized and doubly distilled water was used throughout this study. Photoreactor and Light Source. A 500-Watt halogen lamp (Institute of Electric Light Source, Beijing) was positioned inside a cylindrical Pyrex vessel surrounded by a circulating water jacket (Pyrex) to cool the lamp. A cutoff filter was placed outside the Pyrex jacket to remove radiation below 420 nm and to completely ensure irradiation of the dispersion only by visible light wavelengths.

J. Phys. Chem. B, Vol. 103, No. 23, 1999 4863

Procedures and Analyses Aqueous TiO2 suspensions were prepared by addition of a given weight of TiO2 powder to a 50-mL aqueous dye solution. Prior to irradiation, the suspensions were magnetically stirred in the dark for ca. 30 min to ensure the establishment of an adsorption/desorption equilibrium. The dispersions were kept under constant air-equilibrated conditions before and during irradiation. At given irradiation time intervals, 3-mL aliquots were collected, centrifuged, and then filtered through a Millipore filter (pore size 0.22 µm) to remove the TiO2 particulate. The filtrates were analyzed by recording variations at the wavelength of absorption band maximum in the UV-vis spectra of the dyes using a Lambda Bio 20 spectrophotometer (Perkin-Elmer Co.). COD was assayed using the potassium dichromate titration method.17 COD values of both the dye suspensions (45 mL) containing the TiO2 particulates and those of the bulk solution after removal of TiO2 particles were determined. The concentration of total peroxides (including organoperoxides and H2O2) formed during irradiation of a 50-mL solution of dye containing TiO2 was determined immediately after irradiation at various time intervals and removal of the TiO2 particles by centrifugation and filtration. The spectrophotometric DPD method18 was employed (λ ) 551 nm,  ) 21,000 M-1 cm-1) wherein the DPD reagent is oxidized by either H2O2 and/or the organoperoxides on the basis of the peroxidase catalyzed reaction.18 It is known that the catalase enzyme can eliminate H2O2 from the mixture of H2O2 and organoperoxides.19 If catalase is added to the degraded solution that contained H2O2 and organoperoxides before the DPD method is used, then this DPD method assays only the organoperoxides. Thus, we can discriminate between H2O2 and organoperoxides and determine whether one or both are formed during the temporal course of dye photodegradation. Results and Discussion The photodegradation of a series of dyes was carried out in TiO2 dispersions under visible light irradiation during which the generation of peroxides was examined. The DPD method employed for peroxide measurements was used for the detection of both H2O2 and any hydroperoxy organic intermediate that formed during the photodegradation of the dyes. Through the discrimination of total peroxides formed (see Experimental Section), the only peroxide detected was H2O2; no organoperoxides formed. For example, in the photodegradation of Squarylium cyanine the maximal concentration of peroxide appeared after ca. 30 min of irradiation (curve a, Figure 1); after

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Figure 1. Detection of H2O2 and organoperoxides in the photodegradation of Squarylium cyanine (8.3 × 10-5 M)/TiO2 dispersions (TiO2 20 mg/50 mL). Curve a was obtained by addition of phosphate buffer (pH ) 6.8), DPD, and POD to the degraded solution after ca. 30 min of irradiation; curve b was obtained by the same procedure except that catalase (0.1 mL, 0.5%; activity greater than 2.0 U mg-1) was added prior to the other reagents (DPD, POD).

addition of catalase to remove H2O2, no other peroxide was detected (curve b, Figure 1). Measurements of peroxides in other dye/TiO2 dispersions also showed that only H2O2 formed. When necessary, control experiments were also carried out. In all cases examined, except for Eosin, no dye photodegraded and no H2O2 formed in the absence of TiO2 particles, in the presence of visible light illuminated Al2O3 or SiO2 particles, or in TiO2 dispersions in the dark. For Eosin, degradation (initial concentration 2 × 10-4 M) and H2O2 formation were observed both in visible light irradiated aqueous dye solutions and in visible light irradiated dye/TiO2 dispersions. However, the degradation of Eosin with TiO2 particles was faster (only ca. 100 min for complete discoloration) than in a homogeneous media (ca. 250 min for complete discoloration). Evidently, the two relevant degradation mechanisms must be different. The presence of TiO2 significantly accelerates the degradation of Eosin. In this case, formation of H2O2 in the presence of TiO2 particles should be ascribed to photoassisted processes.The experimental conditions and the corresponding findings from H2O2 measurements are summarized in Table 1. The results show that for the Rhodamine B (pH ∼ 4.2) and Orange II dyes (pH ∼ 6.5), which adsorb little on the surface of TiO2 particles, the complete discoloration of the dyes necessitated longer irradiation times, and any H2O2 that may have formed was not observed. However, for the Fluorescein, Alizarin red, and Squarylium cyanine dyes, which strongly adsorb on TiO2, their complete discoloration required shorter irradiation times and H2O2 was observed; maximal values detected were 3.0 × 10-5 M, 5.3 × 10-5 M, and 6.6 × 10-5 M, respectively. For Rhodamine B, Orange II, or Sulforhodamine B, regardless of the dye concentration and TiO2 particle loading used, there was no evidence that any H2O2 formed at the original pH of 4.2, 6.5, and 4.2, respectively. The mechanism of dye degradation mediated by semiconductor TiO2 nanoparticles under visible light irradiation (eqs 1-5) indicates that the photodegradation of the dyes should generate the intermediate H2O2. That no hydrogen peroxide formed during the degradation of Orange II (for Rhodamine B and Sulfo-rhodamine B see below) is likely due to the little quantity of dye adsorbed on the TiO2 surface. Consequently, the pH of the Orange II/TiO2 dispersions was adjusted from 6.5 to 2.5 to enhance adsorption of Orange II on the TiO2 surface (the sulfonate moiety in this molecular structure has a strong

Wu et al. adsorption ability; dye adsorption should increase with an increase in the acidity of the dispersion) and to accelerate the degradation rate. The results plotted in Figure 2 show that H2O2 was generated, as expected, owing to the strong adsorption (Γ, 17.0 × 10-9 mol mg-1 TiO2) and a fast degradation rate (tend, 350 min for 8 × 10-5 M) after adjustment of pH. Similar results were obtained under otherwise similar changes in pH (adjusted from 4.2 to 2.5) for the TiO2 dispersions containing Sulforhodamine B (Figure 3). More interestingly, addition of the anionic surfactant DBS (at its cmc of 1.2 mM) to the Rhodamine B (1.0 × 10-4 M, pH 4.2) dispersions with TiO2 particles (previous studies showed that adsorption and the rate of degradation of Rhodamine B are not significantly enhanced by adjustment of pH but are greatly improved in the presence of an anionic surfactant such as DBS14b) led to a striking increase in both the extent of adsorption (∼ 100%) and in the rate of degradation of the dye (tend ) 80 min for complete discoloration of 1 × 10-4 M of dye); concomitantly, a considerable quantity of H2O2 formed with maximal [H2O2] ) 6.0 × 10-5 M after ca. 80 min of irradiation. These observations demonstrate that strong adsorption and fast degradation rates of the dyes are prerequisite for generation and detection of H2O2. Failure to observe formation of H2O2 during the degradation of the Sulfo-rhodamine B dye at the original pH (4.2) is somewhat enigmatic since the quantity of dye adsorbed (Γ ) 7.0 × 10-9 mol mg-1) compares with that of Fluorescein (10.0 × 10-9 mol mg-1) and the extent of adsorption is greater than that of Fluorescein by a factor of 2. Compared to the degradation of Fluorescein (maximal [H2O2 ] ) 3.0 × 10-5 M, ∆COD changed from an initial 7.7 mg L-1 to 27.0 mg L-1), the degradation of Sulfo-rhodamine B necessitated a longer irradiation time for its complete discoloration at the original pH of 4.2. Furthermore, addition of H2O2 had no effect on the degradation rates of all the dyes tested in the present experiments, indicating that H2O2 did not interact with the various substrates. The above two considerations suggest that H2O2 formed must have been photodecomposed by the TiO2 particulates during the degradation process and that its rate of decomposition was greater than its rate of formation. To confirm these assertions, we examined the photodecomposition of hydrogen peroxide in TiO2 dispersions irradiated with visible light containing only H2O2. The results depicted in Figure 4 illustrate that H2O2 decomposes under these conditions (note that no change in [H2O2] occurred either in the presence of TiO2 particles in the dark or in the absence of TiO2 under visible light irradiation). As well, addition of the surfactant DBS to the H2O2/TiO2 dispersions decreased the rate of decomposition of H2O2, which demonstrates that DBS covered the surface of TiO2 particles to some extent, thereby retarding the interaction between H2O2 and the TiO2 particles. Consideration of the quantity of H2O2 that should form in the Sulfo-rhodamine B/TiO2 suspensions at pH 2.5 suggests that equivalent concentrations of Sulfo-rhodamine B in dispersions, albeit containing different loadings of TiO2 particles, should nonetheless produce equal concentrations of H2O2 after complete discoloration of the dye if the H2O2 generated is not decomposed by the TiO2 particles (however, see curves e and f in Figure 3). Sulfo-rhodamine B dispersions containing three different loadings of TiO2 (100, 200, and 300 mg) produced maximal concentrations of H2O2 of 7.6 × 10-5 M, 4.2 × 10-5 M, and 2.5 × 10-5 M, respectively. After confirming adsorption of H2O2 on TiO2 particles in different dispersions by an in situ assay (20% for 100 mg, 28% for 200 mg, and 38% for 300 mg of TiO2), the maximal concentrations of hydrogen peroxide formed

H2O2 Generation during Photodegradation of Dyes

J. Phys. Chem. B, Vol. 103, No. 23, 1999 4865

TABLE 1. Data on the Formation of H2O2 in the TiO2-Mediated Photodegradation of Dyes dye

conc. (mol L-1)

pHa

adsorp (%)

TiO2 mg/50 mL

Γ (×10-9)b (mol mg-1 TiO2)

tendc (min)

CH2O2max d (mol L-1)

Rhodamine B Rhodamine Be Orange II Orange II Sulfo-rhodamine B Sulfo-rhodamine B Fluorescein Alizarin red Squarylium cyanine Eosin

1.0 × 10-5 1.0 × 10-4 2.0 × 10-5 8.0 × 10-5 5.0 × 10-5 5.0 × 10-5 8.0 × 10-5 2.0 × 10-4 8.3 × 10-5 2.0 × 10-4

4.2 4.2 6.5 2.5f 4.2 2.5g 6.1 3.8 4.2 4.8

8 ∼100 10 41.0 23 36 12.5 38 39 11

100 100 25 25 83 83 50 50 20 83

0.4 50.0 4.0 66.0 7.0 11.0 10.0 76.0 80.0 13.0

480 80 600 350 600 100 160 120 60 100

0 6.0 × 10-5 0 3.4 × 10-5 0 7.8 × 10-5 3.0 × 10-5 5.3 × 10-5 6.6 × 10-5 11.0 × 10-5

a Original pH when the dye dissolved in doubly distilled water, adjusted by neither acid nor base. b Γ, Adsorption amount per milligram of TiO . 2 tend, Irradiation time for the complete decoloration of dyes. d CH2O2max, maximal concentration of H2O2 formed. e With addition of surfactant DBS f ,g at its critical micelle concentration. After adjusting from pH 6.5 and 4.2, respectively.

c

Figure 2. Concentration changes of Orange II (8 × 10-5 M; curve a), H2O2 formation (curve b) and COD changes (curve c, total COD of dispersions containing TiO2; curve d, COD changes of bulk solution after removal of TiO2 particles) as a function of irradiation time during degradation in the presence of TiO2 particles (25 mg/50 mL) under visible light irradiation after adjustment of the pH to 2.5.

Figure 3. Concentration changes of Sulfo-rhodamine B (5 × 10-5 M) (curve a), H2O2 formation (curve b), and COD changes (curve c, total COD of dispersions containing TiO2; curve d, COD changes of bulk solution after removal of TiO2 particles) as a function of irradiation time during degradation in the presence of TiO2 particles (100 mg/60 mL) under visible light irradiation after adjustment of the pH to 2.5. Curves e and f represent H2O2 concentrations formed in the degradation under otherwise identical conditions except that the loading of TiO2 particles was 200 mg and 300 mg, respectively.

were reestimated, respectively, at 9.1 × 10-5 M, 5.4 × 10-5 M, and 3.5 × 10-5 M. This shows that the greater the loading of TiO2 particles, the greater the quantity of H2O2 that is decomposed, thereby emphasizing the photoassisted interaction between H2O2 and TiO2 particles.

Figure 4. H2O2 decomposition (initial concentration 6.0 × 10-5 M) with irradiation time in TiO2 dispersions (25 mg/50 mL) under visible light: (a) pH ) 2.5; (b) pH ) 6.0 with added surfactant DBS (at cmc of 1.2 mM); (c) pH ) 6.0; (d) pH ) 9.5.

COD Measurements. Changes in COD with illumination times represent the degree of dye degradation or mineralization from which we can gain some insights into the mechanism of H2O2 generation during visible light irradiation. Differences in COD (i.e., ∆COD) between the total COD of the dye/TiO2 dispersion and the COD of the bulk solution after removal of TiO2 particles, at the same irradiation times reflect the total organic carbon (substrates plus intermediates) adsorbed on the TiO2 surface. The greater this ∆COD, the greater is the quantity of organic carbon adsorbed on the TiO2 surface, which results in less or no direct interaction between H2O2 formed and TiO2 particulates, thereby favoring accumulation of formed H2O2 in the bulk solution and its subsequent detection. Indeed, when H2O2 was easily observed in the cases of Sulfo-rhodamine B (pH 2.5) and Fluorescein media, ∆COD changes were relatively large or even increased during the degradation of the dyes (except for Orange II and Alizarin red, see discussion below). In particular, in the case of Eosin degradation (see Figure 5), a greater final ∆COD (52 mg L-1) was obtained despite the small initial ∆COD (6.0 mg L-1). This demonstrates that an increase in ∆COD is an important contribution to the observation of H2O2 because the greater quantity of organic carbon (intermediates or degraded end products) adsorbed on the TiO2 surface hinders the reaction of H2O2 formed and TiO2 particles, even though the extent of adsorption is initially small (ca. 10%). Thus, H2O2 accumulated in solution and decayed slowly (note the plateau in the plot of H2O2 formation versus irradiation time). It is noteworthy that coverage of the surface of TiO2 particulates by a substrate and its intermediates varies in the different cases

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Wu et al. TABLE 2. In Situ Adsorption of H2O2 Added to the Dye/TiO2 Dispersions before Irradiation and after Complete Decoloration of the Dyes dye Orange II Fluorescein Sulfo-rhodamine B Alizarin red Eosin

TiO2 pH (mg/50 mL)

before irradiation

after complete decoloration

2.5 6.1 2.5 3.8 4.8

5% (8.8)a 50% (7.7) 27% (10.0) 11% (20.0) 38% (6.0)

5% (9.0) 13% (27.0) 20% (31.2) 37% (3.6) 14% (52.0)

25 50 83 50 83

a Data in parenthese are the ∆COD values in the corresponding stage.

Figure 5. Concentration changes (curve a), H2O2 formation (curve b), and COD changes (curve c, total COD of dispersions containing TiO2; curve d, COD of bulk solution after removal of TiO2 particles) as a function of irradiation time during the degradation of Eosin (2 × 10-4 M, pH ) 4.8) in the presence of TiO2 particles (100 mg/60 mL) under visible light irradiation.

Figure 6. Concentration changes of Alizarin red (2 × 10-4 M) and H2O2 formation as a function of irradiation time during degradation in the presence of TiO2 particles (50 mg/50 mL) under visible light irradiation. Inset: total COD changes containing TiO2 particles (curve c) and the COD changes of the bulk solution (curve d) after removal of TiO2 particles with illumination time under otherwise identical conditions.

examined, with different degrees of interaction between H2O2 and TiO2. Strong interactions (i.e., small ∆COD) led to rapid decomposition of the photogenerated H2O2. As a demonstration of this notion, we note the case of Alizarin red. The initial adsorption of substrate and ∆COD were remarkably large; subsequently, ∆COD decreased gradually with increase in irradiation time (see curves c and d in Figure 6). Note the sharp peak in the plot of H2O2 concentration versus irradiation time (curve b, Figure 6) because a large quantity of substrates was adsorbed on the TiO2 surface to hinder H2O2 decomposition during the initial stages; however, to the extent that the intermediates adsorbed less in the later stages, the H2O2 present in the solution bulk decomposed rapidly. For comparison, changes in COD with irradiation time were also examined at the original pH of 4.2 for Sulfo-rhodamine B. The COD changes in these dispersions display the same situation (large initial ∆COD ∼ 20.0 and final ∆COD near 0 after 6 h of irradiation) as that of Alizarin red. In the case of Rhodamine B/TiO2 dispersions at the original pH 4.2, our previous study had shown that ∆COD increased with increase in irradiation

time during the first 15 h.14a Complete discoloration of Sulforhodamine B and Rhodamine B needs too long an irradiation time, with the consequence that any H2O2 formed is decomposed during the photodegradation and is not detected. These above results demonstrate that when ∆COD is small and/or the rate of degradation of the dye is slow, the H2O2 is formed slowly but is rapidly and extensively decomposed. Diagnosis for an In Situ Adsorption of H2O2. The above results suggest that a large ∆COD should correspondingly lead to less adsorption of H2O2 since the ∆COD reflects the coverage of the TiO2 surface by substrates or intermediates. The in situ adsorption of H2O2 on the TiO2 surface was verified by addition of H2O2 to the dye/TiO2 dispersions before irradiation and after complete discoloration. Note that the final in situ adsorption of added H2O2 was measured after the TiO2 dispersions had been completely discolored and after the photogenerated H2O2 had been completely decomposed for prolonged irradiation times. The results are collected in Table 2. It is evident that the in situ adsorption of H2O2 depends on the ∆COD value in the degradation of a dye. When the initial ∆COD values were smaller than the final ∆COD, the in situ adsorption was greater before irradiation than after complete discoloration of the dye; conversely, the in situ adsorption was less before irradiation than after complete discoloration of the dye. Adsorption of H2O2 on the TiO2 surface did not vary (always at 5%), whereas the ∆COD remained constant (10 mg L-1) before irradiation and after complete discoloration of Orange II, which indicates that adsorption of H2O2 is closely related to ∆COD or to adsorption of substrates or intermediates on the surface of TiO2 particles. That is, adsorption of substrates or intermediates on the surface of TiO2 particles (reflected by ∆COD) directly affects the interaction of H2O2 with the TiO2 particles and hence significantly affects the decomposition rate of the H2O2 produced. Implication of pH Effects. Another factor that influences the adsorption and decomposition of H2O2 is the prevailing pH of the dispersions. As depicted in Figure 4, the adsorption and decomposition rates of H2O2 are smaller at the lower pH of 2.5. By contrast, at the higher pH of 9.5, the opposite results are obtained. These findings indicate that H2O2 displays a slower decomposition rate in higher acidic media, even though it can be adsorbed on the TiO2 surface. Accordingly, pH makes a nonnegligible contribution to the formation and thus observation of H2O2. A good example of this is the result from the Orange II dye (at pH 2.5): a large quantity of H2O2 (near the theoretical value) was detected after a relatively long irradiation time (Figure 2), albeit the ∆COD values are not large. Thus, it is possible that the slow decomposition rate of H2O2 at pH 2.5 played a positive role in the detection of H2O2 formed. The large extent of adsorption and decomposition of H2O2 in the TiO2 dispersion at the higher pH of 9.5 (Figure 4) implies that there is an interaction between H2O2 and TiO2 particles, and

H2O2 Generation during Photodegradation of Dyes even a stronger interaction by the complexing of titanium ions and H2O2 in the alkaline solution than seen in acidic media, in keeping with the results reported by Kiwi and Graetzel.20 We ascribe the decomposition of H2O2 formed to the photodecomposition by visible light of the (yellow) titanium(IV) peroxy surface complex, TiO2/H2O2, formed on the particles.21 Earlier, Hoffmann et al. reported the formation of H2O2 in semiconductor ZnO and TiO2 aqueous dispersions during the photocatalytic degradation of acetate under UV light irradiation.19b,22 It is relevant to note that the mechanism of H2O2 formation under UV irradiation is different from the present one under visible irradiation. In the UV irradiation case, H2O2 may be formed through both H2O and OH- ions by surface oxidation by the photogenerated holes and, in part, by the disproportion of the superoxide radical anion, whereas decomposition of H2O2 is caused by reduction by the conduction band electron, oxidation by the valence band holes, or by reaction with TiIII. 22b However, in the visible light irradiation mechanism, H2O2 is formed through continuous reduction of dioxygen adsorbed on the surface of TiO2 particles;14c H2O2 decomposition occurs mainly through the photodecomposition of the Ti(IV)/ H2O2 surface complex (see above). To our knowledge, formation of H2O2 during the photodegradation of dyes in the presence of TiO2 particles under visible light irradiation has not received much attention. Conclusions Photogenerated H2O2 has been observed during the TiO2mediated degradation of dyes by direct detection or by changing experimental conditions such as pH or by addition of a surfactant to the dye/TiO2 dispersions. The concentration of H2O2 depends on both the coverage of substrates and/or intermediates on the TiO2 surface and the photodegradation rate of the dye. Mechanistically, two moles of dyes produce one mole of H2O2 by reducing O2 via O2-• or •OOH, and subsequently the •OH radical is formed by reduction of H2O2 generated. On the basis of the present findings that a large quantity of H2O2 did form and that both discoloration and degradation of dyes occurred, we can unambiguously infer that electrons injected onto the TiO2 particles from excited state(s) of the dyes are depleted principally for the formation of H2O2. The dye cation radicals formed after the electron injection stage are subject to oxidation by molecular oxygen. Acknowledgment. The authors appreciate the generous financial support of this work by the National Natural Science

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