Visible-Light-Assisted Degradation of Dye Pollutants over Fe(III

cyanine (FePcS) (9) or to amidomacrocylic ligand (FeTAML). (10) have been .... on an Apollo 9000 TOC analyzer; for that the sample was prepared by sim...
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Environ. Sci. Technol. 2004, 38, 1569-1575

Visible-Light-Assisted Degradation of Dye Pollutants over Fe(III)-Loaded Resin in the Presence of H2O2 at Neutral pH Values MINGMING CHENG, WANHONG MA, JING LI, YINGPING HUANG, AND JINCAI ZHAO* Laboratory of Photochemistry, Center for Molecular Sciences, Institute of Chemistry, The Chinese Academy of Sciences, Beijing 100080, China YU XIANG WEN Institute of Physics, The Chinese Academy of Sciences, Beijing 100080, China YIMING XU Chemistry Department, Zhejiang University, Hangzhou, Zhejiang 310027, China

A novel catalyst was synthesized by direct exchange of ferric ions onto a cationic resin (Amberlite IRA200). Upon visible light irradiation (λ > 420 nm) in the presence of H2O2, this catalyst was found to be highly effective for the degradation of nonbiodegradable cationic dyes, Malachite green, Rhodamine B, and Methylene blue, even at neutral pH values. It was also easy to separate from the degraded solution. By total organic carbon, FT-IR, and GC-MS analysis, the degradation process of Malachite green was shown to proceed with demethylation and phenyl ring openings into CO2 and small molecular compounds. EPR studies revealed that •OH radicals, other than •OOH/O2•-, were involved as the active species. A possible reaction mechanism is proposed on the basis of all the information obtained under various experimental conditions.

Introduction The Fenton reaction (Fe2+/H2O2) has been developing for more than a century since it was first reported in 1894 (1). Attention focusing on its possible application in wastewater detoxification (2), destruction of organic pollutants (3), and biodegradability improvement of wastewater (4) in recent decades has proved this method an effective as well as environmentally benign choice. A Fenton-like system (Fe3+/ H2O2) also exhibits high efficiency for the degradation of organic compounds (5). The organic pollutant is degraded by some reactive oxygen species (ROS) such as hydroxyl radical (•OH) and hydroperoxyl radical (•OOH/O2•-) that are generated from H2O2-Fe3+/Fe2+ reactions (eqs 1 and 2) (6). While •OH is generally accepted as the most active in this process (6, 7), additional pathways involving high-valence iron species, such as FeO3+ and ferryl complexes (L)FeIVdO and (L•+)FeIVdO formed from iron peroxide complex by * Corresponding author phone: +86-10-8261-0080; fax: +86-108261-6495; e-mail: [email protected]. 10.1021/es034442x CCC: $27.50 Published on Web 01/17/2004

 2004 American Chemical Society

internal two-electron transfer, have also been proposed (8).

Fe2+ + H2O2 f

Fe3+ + OH- + •OH (k1 ) 58 mol-1 dm3 s-1) (1)

Fe3+ + H2O2 f

Fe2+ + •OOH + H+ (k2 ) 0.02 mol-1 dm3 s-1) (2)

However, the Fenton reaction suffers from the drawback of a narrow pH range (pH < 3) and Fe-sludge disposal and/ or regeneration. The Fe ions bound covalently to phthalocyanine (FePcS) (9) or to amidomacrocylic ligand (FeTAML) (10) have been reported to be highly active for catalytic oxidation of chlorophenols by H2O2 in the dark. However, the organic ligand moiety of the complex was found to decompose gradually during the reactions, which might bring about secondary contamination in environmental application. Heterogeneous immobilization of ferric ions on derivatized silica fabrics, Nafion, and polyethylene membranes, on the other hand, has been studied by Kiwi and co-workers for the use as an efficient photocatalyst for Orange II degradation (11-13). Also, Fe(III)-loaded TiO2 shows high efficiency for the photocatalytic oxidation of organic compounds (14) and photodegradation of dyes (15) under UV irradiation. Our group has examined Fe complexes (FePcS) in solution or supported on an ion exchange resin (Amberlite IRA200) that are active for degradation of organics upon visible light irradiation (16). We have explored recently a novel catalyst (FeR) of ferric ions supported directly on cationic exchange resin (Amberlite IRA200), which shows a highly catalytic activity for the degradation of small organic molecules, dichlorophenol, and benzyltrimethylammonium chloride (BTAC), in the presence of H2O2 upon UV light irradiation (17). Related to this work, we report in the present study that this catalyst is also highly effective as a photocatalyst for the photosensitized degradation of dyes under visible light irradiation. Dyes represent some of the principal pollutants in the textile and photographic industry (18). Recently we have reported that dyes could be degraded effectively in the presence of TiO2 particles under visible irradiation (19). Dyes absorb visible light and could be excited under visible irradiation, which could lead to electron transfer between the dye molecule and the ferric ion on the resin and the degradation of dyes. The degradation pathway under visible irradiation is different from that under UV light irradiation. It was found that, in the presence of H2O2 under visible irradiation, the photodegradation of Malachite green occurred efficiently even at neural pH, with a total organic carbon (TOC) removal yield of 59% and a very high efficiency of H2O2 usage. The catalyst was stable and could be repeatedly used. The experiments were also carried out using other cationic dyes, Rhodamine B and Methylene blue, anionic dyes Sulforhodamine B and Orange II, and the small molecule BTAC as the model pollutants. The degradation process was characterized by TOC, H2O2 analysis, FT-IR, and GC-MS techniques. EPR was used to detect the reactive species involved in the reactions. Our finding provides an approach for the treatment or pretreatment of organic dyes under visible irradiation.

Experimental Section Materials and Reagents. The cationic exchange resin (Amberlite IRA200) was obtained from Alfa Aesar Co. Rhodamine B (RhB), Malachite green (MG), Methylene blue VOL. 38, NO. 5, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 1. Schematic illustration of the photoreactor and light source: 1, halogen lamp; 2, cylindrical Pyrex flask surrounded by a circulating water jacket; 3, light filter embedded in the chipboard; 4, stirrer; 5, magnetic stirrer.

(MB), Orange II, Sulforhodamine B (SRB), BTAC, hydrogen peroxide, and FeCl3‚7H2O were of laboratory reagent grade and used without further purification. The reagent 5,5dimethyl-1-pyrroline N-oxide (DMPO), used as the spintrapping agent in EPR studies, was purchased from the Sigma Chemical Co. Deionized and doubly distilled water was used throughout this study. The solution pH was adjusted by diluted aqueous solutions of NaOH or HClO4. The ion exchange resin before use was pretreated with 1 M HCl and water in turn three times, followed by washing with 1 M NaOH and water. After being dried at room temperature, the resin was triturated into powder. The resin was then loaded with ferric ions in a vial (ca. 100 mL) containing 1 g of resin powder and ferric chloride. After the mixture was stirred for 24 h, the resin was washed out with water till no ferric ion was detected spectrophotometrically in the solution (20). The Fe-exchanged resin was finally dried in air. Following this procedure, the catalysts with different Fe content from 20% to 100% of the exchange capacity were prepared (note that the exchange capacity of the resin was about 1 mmol of Fe3+/g) and applied in the FeR/H2O2/MG system. Since the resin loaded with 50% ferric ion exchange capacity exhibited the highest activity, this catalyst was used for all of the experiments. Procedures and Analysis. The irradiation source was a 500 W halogen lamp fixed inside a cylindrical Pyrex flask, which was surrounded by a circulating water jacket to cool the lamp (Figure 1). The exterior of the cylindrical Pyrex flask was wrapped by black paper, just leaving a small window (φ ) 4 cm) at the side face. A light filter (to cut off the light below 420 nm wavelength) was put at the small window to ensure the irradiation only by visible light. All irradiation experiments were carried out in a cylindrical Pyrex vessel (60 mL) which was 10 cm to the light center. The samples (3 mL) after reaction were analyzed immediately on a Hitachi U-3010 spectrometer. The content of TOC was determined on an Apollo 9000 TOC analyzer; for that the sample was prepared by simple filtration to remove the catalyst particles. The EPR study was performed on a Bruker model EPR 300E spectrometer equipped with an irradiation source of QuantaRay ND:YAG laser system (λ ) 532 nm), and the same quartz capillary tube was used to minimize the errors during the measurements. GC-MS data were obtained on a Trio-2000 apparatus (Micromass UK Ltd.) equipped with a BPX70 column (size 30 m × 0.25 mm). The sample after a certain period of reaction time was filtered, and the filtrate was then concentrated under reduced pressure at a temperature below 323 K. The residue was finally dissolved in 0.5 mL of methanol for GC-MS analysis. The infrared spectrum was recorded on a TENSOR 27 FTIR spectrophotometer (Bruker). For this measurement, the dry residue was obtained similarly to the above, and supported on a KBr pellet. Hydrogen peroxide was analyzed photometrically by the POD (horseradish peroxidase)-catalyzed oxidation product of DPD (N,Ndiethyl-p-phenylenediamine) at λ 551 nm ( ) 21000 M-1 cm-1) (21, 22). 1570

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FIGURE 2. MG degradation under various conditions: (a) in the presence of FeR under visible irradiation, (b) in H2O2 solution under visible light irradiation, (c) in the presence of FeR and H2O2 in the dark, (d) in the presence of FeR and H2O2 under visible irradiation. All the reactions were performed at pH 6.0, with 2.5 mg of FeR, [H2O2] ) 2.5 mM, and [MG] ) 0.1 mM. The inset is the UV/vis spectral changes recorded for (d) as a function of irradiation time.

FIGURE 3. Cyclic degradation of MG as a function of irradiation time: MG, 0.1 mM/run; H2O2, 2.5 mM/run (amount of FeR 2.5 mg, pH 6.0).

Results and Discussion Kinetic Study of MG Photodegradation. The MG photodegradation was carried out after the adsorption/desorption equilibrium had been reached between MG and the catalyst FeR. MG did not react considerably with H2O2 over FeR in the dark (Figure 2, curve c), or under visible light irradiation in the absence of the catalyst (curve b). However, MG could be photodegraded completely in the presence of H2O2 over the catalyst (curve d), whereas without H2O2 the photodegradation hardly occurred (curve a). The temporal UV/vis spectra showed that the MG characteristic band centered at 616 nm was decreased promptly upon light irradiation (inset plot in Figure 2), together with a blue shift at the maximum absorption. As will be shown by GC-MS analysis, this is attributed to the N-demethylation of MG during the degradation process. The solution color, after the reaction, turned to colorless from blue, but the solution pH changed only slightly from 6.0 to 5.8. No free Fe ions were found in the solution phase by using the ferrozine spectrophotometric method, indicating that the immobilized Fe species was stable and did not drop off from the resin during the reaction. This was further supported by six repetitive experiments (Figure 3), during which the catalyst activity decreased 3-5% in each run. The results above demonstrate that the Fe3+-exchanged

FIGURE 4. Decomposition of H2O2 under various conditions: (a) in the presence of MG under visible irradiation, (b) in the presence of MG and FeR in the dark, (c) in the presence of FeR under visible irradiation, (d) in the presence of MG and FeR under visible irradiation, (e) in the presence of Fe3+ under visible irradiation. All the reactions were at pH 6.0 (in (e) pH 3.0), with 2.5 mg of FeR, [Fe3+] ) 5 × 10-6 M, [H2O2] ) 2.5 mM, and [MG] ) 0.1 mM. resin is an effective and stable catalyst for MG degradation under visible light irradiation. An interesting behavior in H2O2 decomposition was observed during the photodegradation of MG (Figure 4). In the presence of MG (curve a) or FeR (curve c) alone, no substantial decomposition of H2O2 was observed upon visible light irradiation. In the presence of both MG and FeR, however, the irradiation caused H2O2 decomposition obviously (curve d), while in the dark still no decomposition was observed (curve b). It is known that the ferric ion can catalyze H2O2 decomposition even in the dark independent of the presence of organic substrates (eqs 1 and 2), which was confirmed by H2O2 decomposition upon light irradiation in homogeneous solution containing the same concentration of free ferric ions as that on the Fe-loaded resin (curve e). Moreover, H2O2 decomposition was found to proceed synchronously with MG degradation. On curve d, point 1 corresponds to the time when MG was completely discolored. If MG was not added at this time till point 2, the concentration of H2O2 remained unchanged even under the irradiation (a quasi horizontal line was seen in this period of irradiation time for 180 min). After MG was added artificially into the system at point 2, the H2O2 decomposition occurred again, accompanied by MG degradation under irradiation. Thus, for the H2O2 decomposition, all the components of organic substrate, catalyst, and visible light irradiation are necessary, and the catalyst provides a high efficiency in H2O2 utilization. Degradation of Other Dyes. The experiments were then expanded to use other organic compounds of different types as model pollutants. As shown in Figure 5, the cationic dyes RhB and MB were degraded efficiently in the FeR/H2O2 system upon visible light irradiation as MG, whereas the anionic dyes SRB and Orange II and the small cationic molecule BTAC as well were not degradable under the same conditions. Accordingly, H2O2 decomposition was observed only with RhB and MB degradations. The resin had 50% of its exchange capacity and still had free exchange sites; the catalyst was negatively charged. It was determined that SRB and Orange II had no adsorption on the negatively charged catalyst, implying that the preadsorption of the substrate is a requirement for efficient reaction on the catalyst surface. On the other hand, the small molecule BTAC with a lack of light absorption in the visible region was not degradable, despite the fact that the cationic BTAC molecules could adsorb easily onto the catalyst. This suggests that the reaction occurs from

FIGURE 5. Photodegradation of various organic substrates (0.1 mM) accompanied with simultaneous decomposition of H2O2 (2.5 mM). the excited states of the organic molecules because the catalyst cannot absorb the visible light. In fact, BTAC on this catalyst could be degraded rapidly under UV irradiation (17). TOC Analysis. A decrease in TOC as measured by TOC with the irradiation time during MG and RhB photodegradations in the FeR/H2O2 system is shown in Table 1. However, the TOC decrease was stopped when the dye was completely discolored (500 and 1100 min for MG and RhB degradation, respectively), where the yield of TOC removal was determined to be 59% and 65% respectively. This was realized by the spectra of the irradiated solutions that exhibited no absorption in the visible light region. The result offers further evidence that the dye degradation is initiated from its excited state on the catalyst surface, which is different from that under UV irradiation. It is the catalyst other than the substrate that is excited and leads to the substrate degradation under UV irradiation (17). Effect of pH. The photodegradations of MG, MB, and RhB were performed at different pH values from 2 to 10 in the FeR/H2O2 system (Figure 6). The rates of conversion are listed in Table 2. Clearly, the reaction occurred at a wide pH range from acidic to alkaline, which was distinct from the Fenton system (pH ≈ 3). However, the degradation rate was decreased with the increase in initial pH for all the cases. No free Fe ions were detected in all cases. This tendency was checked to be consistent with the adsorption decreasing also with pH. Little degradation of RhB observed at pH higher than 6 was due to its negligible adsorption over the catalyst. EPR Measurements during MG Reaction. The electron paramagnetic trapping (EPR) technique is an effective method in identification of active radicals. For this study, all the EPR spectra were recorded in situ by laser irradiation (λ ) 532 nm) using DMPO as the radical scavenger. During MG photodegradation, the EPR spectra exhibited a 4-fold charVOL. 38, NO. 5, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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TABLE 1. TOC Degradation of MG and RhB (0.1 mM, 50 mL) in the Presence of H2O2 and FeR MG RhB

irradiation time (min) TOC (ppm) irradiation time (min) TOC (ppm)

0 23.67 0 24.51

60 18.44 150 20.4

120 16.32 300 17.25

210 14.96 400 14.49

300 11.85 500 12.54

390 10.41 700 8.48

500 9.88 960 7.60

600 9.56 1100 6.54

700 9.66 1300 6.21

TABLE 2. Rates of Conversion of Dyes in 60 min MG MB RhB

pH conversion %t)60min pH conversion %t)60min pH conversion %t)60min

2.0 75.40 2.0 79.08 1.75 15.99

2.8 56.44 2.93 70.12 2.47 9.09

4.0 46.21 4.17 42.30 3.7 5.64

5.0 34.34 5.2 31.12 3.9 4.91

6.0 27.20 6.0 23.1 4.6 4.06

7.2 16.32 7.1 13.32 5.0 3.45

8.2 6.12 8.1 5.62 5.9 0.91

10 3.41 9.8 4.20 7.0 0.24

FIGURE 7. DMPO spin-trapping EPR spectra under visible irradiation in aqueous solutions: (a) H2O2/MG system under visible irradiation, (b) H2O2/FeR/MG system under visible irradiation ([MG] ) 0.1 mM, [H2O2] ) 2.5 mM).

FIGURE 6. Effect of pH on the degradation of MG, MB, and RhB ([dye] ) 0.1 mM, [H2O2] ) 2.5 mM, 2.5 mg of FeR). acteristic peak of DMPO-•OH adducts with an intensity ratio of 1:2:2:1 (Figure 7b) consistent with similar spectra reported before (23). The EPR signal was not displayed in the H2O2/ FeR/MG system in the dark, or under visible irradiation in either the H2O2/MG (Figure 7a) or the H2O2/FeR system. This indicates that •OH is formed only in the irradiated H2O2/ FeR/MG system, consistent with the results obtained in Figures 2 and 4. Furthermore, the signal intensity was enhanced gradually with irradiation time (Figure 7b). On the other hand, a very weak signal was obtained in the H2O2/ FeR/Orange II system, explaining that Orange II photodegradation was quite slow under the same conditions (Figure 5). To verify whether •OH was the dominant ROS, the formation of •OOH/O2•- radicals was also detected in methanol, because it could not be detected in water owing to its instability in water at room temperature (24). The sextet peaks of DMPO-•OOH/O2•- adducts (25) were observed as expected in the H2O2/Fe3+/MG system under visible irradiation (Figure 8a). However, no such EPR signal appeared in the H2O2/FeR/MG system under the same conditions (Figure 1572

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FIGURE 8. Formation of •OOH/O2•- radicals determined by DMPO spin-trapping EPR in methanol solutions under visible irradiation: (a) H2O2/Fe3+/MG system under visible irradiation, (b) H2O2/FeR/MG system under visible irradiation ([MG] ) 0.1 mM, [H2O2] ) 2.5 mM, and [Fe3+] ) 2.5 × 10-5 mol L-1). 8b). This suggests that the cycling pathway of iron species is significantly altered over the resin, and therefore, only •OH, other than •OOH/O •-, radicals are the active species 2 involved in the MG photodegradation over FeR in the presence of H2O2. FT-IR Data. The process of MG degradation was further examined with FT-IR. After photoreaction, the characteristic

TABLE 3. Retention Times and the Identified Products Corresponding to Figure 10

peaks at 1586 and 1373 cm-1, due to the stretching vibrations of the molecular frame and the stretching vibrations of the Ar-N bond, respectively (26), most clearly disappeared (Figure 9). Meanwhile, several new peaks were generated such as at 1679, 1401, and 1082 cm-1, which were attributed to the stretching vibration of CdO groups, O-H bending vibration of carboxylic acids, and C-N stretching vibration of aromatic monamines (Ar-NH2), respectively. The broad band from 2350 to 3000 cm-1 was due to the N-C stretching vibration of the N(Me)2Cl-1 group; it disappeared finally. The absorptions due to methyl groups at 1476 and 1446 cm-1 were shifted to 1466 cm-1. The peak at 798 cm-1 was attributed to the out-of-plane bending vibration of the CH bonds of 1,4-disubstituted benzenes. All these changes combined with the GC-MS results below gave the evidence that MG degradation underwent demethylation and the cleavage of the central carbon. The increased absorption in the region from 3100 to 3500 cm-1 suggests intermediates containing hydroxyl groups (such as hydroxyl adducts, carboxylic acids, etc). GC-MS Data. The intermediates formed during the MG degradation were further determined by GC-MS. For this analysis, two samples were prepared after MG was degraded by 20% and 90%, as evaluated from UV/vis spectra measurement. The GC chromatograms and the identified intermediates are shown, respectively, in Figure 10 and Table 3, together

FIGURE 9. IR spectral changes recorded for the photodegradation of MG: (a) MG before reaction, (b) 50% of MG degraded, (c) 90% of MG degraded ([MG] ) 0.1 mM). with the MS data listed in Table 4 for each identified intermediate. The most abundant intermediate, as seen from VOL. 38, NO. 5, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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TABLE 4. MS Data of the Identified Intermediates Corresponding to Figure 10 peak

m/z (rel intens of predominant ions in fragmentation pattern/%)

A B C D E F G H I J K X1 X2

179 (90); 148 (100); 136 (4); 132 (12); 118 (10); 104 (15); 77 (26) 211 (96); 210 (89); 194 (15); 134 (67); 105 (10); 91 (100); 77 (57); 51 (51) 233 (36); 231 (100); 197 (8); 181 (9); 156 (23); 154 (79); 128 (7); 105 (5); 91 (32); 77 (28) 136 (91); 120 (3); 105 (100); 77 (40); 51 (10) 260 (35); 259 (100); 183 (10); 182(89); 166 (12); 105 (44); 77 (88); 51 (23) 225 (62); 181 (11); 148 (100); 132 (8); 105 (65); 77 (90); 51 (65) 121 (47); 120 (100); 105 (10); 91 (11); 77 (38); 51 (35); 39 (19) 107 (72); 106 (100); 91 (8); 77 (32); 51 (29); 39 (28) 94 (100); 66 (47); 51 (13); 39 (56) 149 (12); 148 (41); 105 (100); 77 (95); 51 (44); 39 (11) 122 (73); 105 (93); 77 (94); 51 (59); 45 (100); 39 (17) 149 (4); 136 (8); 121 (8); 109 (7); 81 (44); 69 (100); 41 (23); 28 (32) 225 (42); 148 (72); 105 (38); 77 (100); 51 (43)

SCHEME 1. Proposed Mechanism for the Photodegradation of Organic Pollutants over FeR in the Presence of H2O2 under Visible Light Irradiation

might be the actual intermediates formed in the reaction. There were also some chlorinated compounds detected, such as chloro-4-aminobenzophenone (peak C), an indication of chlorine radicals involved in the degradation. Chlorine radicals may be initiated by the reaction between chloride ions and •OH. Small molecular compounds such as HCOOH, HCONHCH3, and HCONH(CH3)2 were not observed due to the evaporation during the sample preparation. The results from FT-IR and GC-MS analysis suggest that MG photodegradation proceeds by cleavage of the central carbon atom to generate 4-dimethylaminobenzophenone, followed by demethylation and opening of phenyl rings to form small molecular compounds.

FIGURE 10. GC chromatograms of the sample obtained from the MG photodegradation in FeR/H2O2 under visible irradiation after (a) 20% of MG and (b) 90% of MG was degraded ([MG] ) 0.2 mM, [H2O2] ) 2.5 mM, 2.5 mg of FeR). Figure 10, was 4-dimethylaminobenzophenone (peak F at 28.92 min), which was degradable further in the system. Together with IR information, this intermediate is produced by the cleavage of the bond between the central carbon and dimethylaminophenyl group. Chloro-4-aminobenzophenone (peak C) and 4-(N-methylformamido)benzophenone (peak B) are analogous degradation products. Methyl 4-(N,Ndimethyl)aminobenzolate (peak A) and methyl benzolate (peak D) are two identified ester compounds. Since methanol was used as a solvent in the sample preparation, which has a tendency to react with a carboxyl group to produce an ester, 4-(N,N-dimethyl)aminobenzolic acid and benzolic acid 1574

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Mechanism Discussion. A possible reaction mechanism, based on all the information obtained above, is proposed in Scheme 1. Since BTAC and the colorless intermediates formed from the dye degradation are not degradable upon visible light irradiation, it is the excited dye molecule that interacts with ferric species on the catalyst surface, generating Fe2+ via one-electron transfer. The ferrous species formed then react with H2O2, producing •OH radicals that initiate subsequent degradation and mineralization of the dye. Different from the photodegradation of dyes under visible irradiation in the free Fe3+ ions/H2O2 solution (27), however, the regenerated ferric species on the resin cannot react directly with H2O2, which was evidenced by EPR results that there were no •OOH/O2•- radicals involved in the reaction. Instead, it reacts only with the excited dye if there is any on the surface. This mechanism can also explain the experimental result that H2O2 does not decompose when the dye is absent or when the dye is discolored completely. The proposed visiblelight-induced mechanism here is distinct from that followed under UV light illumination. Therefore, the resin not only acts as a support for ferric ions and an adsorbent toward the pollutants in solution, but also provides a special microenvironment for active iron centers, enhancing thereafter the catalytic decomposition of H2O2 even at neutral pH values.

Acknowledgments This work was supported financially by the Ministry of Science and Technology of China (Grant 2003CG415006), NSFC (Grants 20371048, 50221201, 20133010, and 20373074), and CAS.

Literature Cited (1) Fenton, H. J. H. J. Chem. Soc. 1894, 6, 899. (2) (a) Kang, S. F.; Liao, C. H.; Chen, M. C. Chemosphere 2002, 46, 923. (b) Rivas, F. J.; Beltra’n, F. J.; Gimeno, O.; Frades, J. J. Agric. Food Chem. 2001, 49, 1873. (3) (a) Joseph, J. M.; Destaillats, H.; Hung, H. M.; Hoffmann, M. R. J. Phys. Chem. A 2000, 104, 301. (b) Theruvathu, J. A.; Aravindakumar, C. T.; Flyunt, R.; Sonntag, J. V.; Sonntag, C. V. J. Am. Chem. Soc. 2001, 123, 9007. (4) (a) Kao, C. M.; Wu, M. J. J. Hazard. Mater. 2000, 74, 197. (b) Lee, B. D.; Hosomi, M. Water Res. 2001, 35, 2314. (5) (a) Murphy, A. P.; Boegli, W. J.; Price, M. K.; Moody, C. D. Environ. Sci. Technol. 1989, 23, 166. (b) Pignatello, J. J. Environ. Sci. Technol. 1992, 26, 944. (6) (a) Walling, C. Acc. Chem. Res. 1975, 8, 25. (b) Walling, C. Acc. Chem. Res. 1998, 31, 155. (7) Zepp, R. G.; Faust, B. C.; Hoigne´, J. Environ. Sci. Technol. 1992, 26, 313. (8) Dong, Y.; Fujii, H.; Hendrich, M. P.; Leising, R. A.; Pan, G.; Randall, C. R.; Wilkinson, E. C.; Zang, Y.; Que, L., Jr.; Fox, B. G.; Kaufmann, K.; Mu ¨ nck, E. J. Am. Chem. Soc. 1995, 117, 2778. (9) Sorokin, A.; De Suzzoni-Dezard, S.; Meunier, B. Science 1995, 268, 1163. (10) Collins, T. J. Acc. Chem. Res. 2002, 35, 782. (11) Fernandez, J.; Dhananjeyan, M. R.; Kiwi, J.; Senuma, Y.; Hiborn, J. J. Phys. Chem. B 2000, 104, 5298. (12) Fernandez, J.; Bandara, J.; Lopez, A.; Buffat, Ph.; Kiwi, J. Langmuir 1999, 15, 185. (13) Pulgarin, C.; Kiwi, J. Langmuir 1995, 11, 519. (14) Ohono, T.; Izumi, S.; Fujihara, K.; Masaki, Y.; Matsumura, M. J. Phys. Chem. B 2000, 104, 6801.

(15) Morwetz, M.; Selli, E. J. Photochem. Photobiol., A, in press. (16) (a) Tao, X.; Ma, W.; Zhang, T.; Zhao, J. Angew. Chem. 2001, 113, 3103. (b) Tao, X.; Ma, W.; Zhang, T.; Zhao, J. Chem.sEur. J. 2002, 8, 1321. (c) Tao, X.; Ma, W.; Li, J.; Huang, Y.; Zhao, J. Chem. Commun. 2003, 80. (17) Ma, W.; Huang, Y.; Li, J.; Cheng, M.; Song, W.; Zhao, J. Chem. Commun. 2003, 1582. (18) Tincher, W. C. Text. Chem. Color. 1989, 21, 33. (19) (a) Liu, G.; Li, X.; Zhao, J.; Hidaka, H.; Serpone, N. Environ. Sci. Technol. 2000, 34, 3982. (b) Wu, T.; Liu, G.; Zhao, J.; Hidaka, H.; Serpone, N. J. Phys. Chem. 1999, 103, 4862. (20) Moss, M. L.; Mellon, M. G. Ind. Eng. Chem. 1942, 14, 862. (21) Bander, H.; Sturzenegger, V.; Hoigne´, J. Water Res. 1988, 22, 1109. (22) Zuo, Y.; Hoigne, J. Science 1993, 260, 71. (23) (a) Jiang, J. J.; Bank, J. F.; Scholes, C. P. J. Am. Chem. Soc. 1993, 115, 4742. (b) Yamazaki, I.; Piette, L. H. J. Am. Chem. Soc. 1991, 113, 7588. (24) Sawyer, D. T.; Valentine, J. S. Acc. Chem. Res. 1981, 14, 393. (25) Chen, F.; Xie, Y.; He, J.; Zhao, J. J. Photochem. Photobiol., A 2001, 138, 139. (26) Thomas, J. J. Chem. Soc. 1970, 979. (27) Wu, K.; Xie, Y.; Zhao, J.; Hisao, H. J. Mol. Catal. A 1999, 144, 77.

Received for review May 6, 2003. Revised manuscript received October 31, 2003. Accepted December 8, 2003. ES034442X

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