Efficient H2O2 Oxidation of Organic Dyes Catalyzed by Simple Copper

Feb 16, 2014 - *Phone: +86 27-59367334. Fax: +86 ... Bicarbonate (HCO3–), one of the most abundant anions in fresh water, is relatively nontoxic and...
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Efficient H2O2 Oxidation of Organic Dyes Catalyzed by Simple Copper(II) Ions in Bicarbonate Aqueous Solution Long Cheng,† Mingyu Wei,† Lianghua Huang,† Fei Pan,† Dongsheng Xia,† Xiaoxia Li,*,‡ and Aihua Xu*,† †

School of Environmental Engineering, Wuhan Textile University, Wuhan, 430073 Hubei, China School of Chemistry and Chemical Engineering, Wuhan Textile University, Wuhan, 430073 Hubei, China

‡

ABSTRACT: Bicarbonate (HCO3−), one of the most abundant anions in fresh water, is relatively nontoxic and cheap. In this work, the degradation of organic dyes with simple copper(II) ions as the catalyst in the HCO3− solution using H2O2 as the oxidant was investigated. It was found that the dyes such as Orange II (AOII), Methyl Orange, Methyl Red, and Toluidine Blue could be efficiently decolorized by the system. The rate of H2O2 decomposition was much slower in the presence of the dyes than that without the pollutants. The formed copper(II) species at different HCO3− concentrations were calculated, and CuCO3 was suggested to be more reactive. The radical scavenging measurements further implied that the produced higher oxidation state of copper, Cu(III), was to be responsible for the dye decolorization. A possible pathway of AOII degradation was also proposed based on the detected intermediate products by electrospray ionization mass spectrometry. This study can provide us a simple, effective, and economical system ideal for the treatment of toxic and nonbiodegradable organic dyes. reactivity of other metal ions in the HCO3− solution for wastewater treatment. Copper compounds, an environmentally relatively acceptable and cheap metal, are known to catalyze oxygenations of hydrocarbons with H2O2 as oxidant.18−20 Meanwhile, the copper complexes such as copper(II)−organic acid and copper(II)−ethylenediamine have also been demonstrated as efficient Fenton-like catalysts for the degradation of organic pollutants in wastewater.21,22 However, cost effectiveness and robustness are often the major obstacles associated with the application of these complex homogeneous catalysts in industrial processes. In order to provide an economical and effective copper based catalyst for the treatment of organic pollutants, there is an attempt made in this study to present the performance of copper(II) ions in the bicarbonate solution on H2O2 activation and organic dyes degradation. The possible catalytic mechanism was also discussed.

1. INTRODUCTION Due to the high water solubility and low volatility, sodium bicarbonate is found predominantly in the aquatic environment. It is also present at high concentration in the range of 14.7−25 mM in biological fluids.1 A major function of the HCO3−-CO2 couple in biological systems is to regulate pH, but recently the role of HCO3− in biological oxidations was recognized.2−5 For example, 1,4-dihydroxy-2-naphthoyl-CoA (DHNA-CoA) synthase is an essential enzyme in vitamin K biosynthesis, and there are two types of DHNA-CoA synthases with a distinctive bicarbonate dependence of their catalytic activity.5 In addition, HCO3− has been found to be an effective H2O2 activator for organic compounds oxidation. Richardson et al. investigated the activation mechanism of H2O2 by HCO3− or CO2, and such a method has been applied in oxidation of aliphatic amines, aryl sulfide, cysteine and related thiols, methionine, epoxidation of alkene, and even decontamination of chemical warfare agents.6−10 It was found that the formed peroxymonocarbonate between the two reagents oxidizes some substrates more rapidly than hydrogen peroxide and was suggested to be the reactive oxygen species in the system.7,8 Textile industries produce a great deal of wastewater polluted with dyes, which are known to be largely nonbiodegradable in aerobic conditions and to be reduced to more hazardous intermediates in anaerobic conditions. The simple and green system, HCO3−/H2O2, has been also successfully used for oxidative degradation of many organic dyes with O2•‑ as the suggested reactivity oxygen species from HCO4− decomposition,11 but high concentration of HCO3− and H2O2 is needed in the system. Later, it was further demonstrated that the incorporation of metal ions such as Mn2+ or Co2+ ions into the HCO3− solution can dramatically enhance the elimination rate of dyes by production of other reactive species such as the MnIVO intermediate and •OH rather than HCO4−.12−17 However, until now there have been no reports for the © 2014 American Chemical Society

2. MATERIALS AND METHODS 2.1. Materials. Copper sulfate, sodium bicarbonate, isopropyl alcohol, sodium azide, and other chemicals were of analytical grade if not noted otherwise. Superoxide dismutase (SOD) and 5,5-dimethyl-1-pyrroline N-oxide (DMPO) were obtained from Sigma Chemical Co. Hydrogen peroxide (30% w/w) was obtained from Sinapharm Chemical reagent Co. Ltd. Orange II (AOII) from Sigma Chemical Co. was used without further purification (>85%) for all experiments. Toluidine Blue (TB), Methyl Orange (MO), and Methyl Red (MR) were also used without further purification. The sample solutions were prepared using deionized water throughout the experiments. Received: Revised: Accepted: Published: 3478

November 11, 2013 February 15, 2014 February 15, 2014 February 16, 2014 dx.doi.org/10.1021/ie403801f | Ind. Eng. Chem. Res. 2014, 53, 3478−3485

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Figure 1. (A) Time evolution of AOII concentration under different conditions and (B) UV−vis spectral changes for AOll decolorization with the Cu2+-HCO3− system. Conditions: Cu2+ 0.03 mM, HCO3− 10 mM, H2O2 4 mM, AOll 0.05 mM, 29 °C.

Figure 2. (A) Decolorization efficiency at different AOII concentrations and (B) decolorization of other organic dyes with the Cu2+-HCO3− system. Conditions: Cu2+ 0.03 mM, HCO3− 10 mM, H2O2 4 mM, AOll 0.05 M, MR 0.074 mM, MO 0.061 mM, TB 0.065 mM, 29 °C.

adding of 10 μL of diluted H2O2 into 100 μL of a mixture solution containing Cu2+ ions, HCO3−, AOII, and DMPO. For identification of degradation products, the samples were analyzed by mass spectrometry. The experiments were performed on an Esquire LC−ion trap mass spectrometer (Bruker Daltonics, Bremen, Germany) equipped with an orthogonal geometry ESI source. Nitrogen was used as the drying (3 L/min) and nebulizing (6 psi) gas at 300 °C. The spray shield was set to 4.0 kV, and the capillary cap was set to 4.5 kV. Scanning was performed from m/z 70 to 800 in the standard resolution mode at a scan rate of 13 kDa/s. Before analysis, each sample was diluted five times.

The solution pH was adjusted by H2SO4 (0.05 mol/L) or NaOH (0.05 mol/L). 2.2. Experimental Procedure. Oxidation reactions were performed in a laboratory scale bath reactor equipped with a magnetic stirrer. The desired amounts of CuSO4, NaHCO3, and dye in 25 mL of the aqueous solution were added in a 50 mL flask and kept at 29 °C in a water bath; then the reaction was initialized by adding 0.25 mL of a diluted H2O2 solution (40 mM). In a typical degradation experiment, the concentrations of the reagents in the reaction mixture were as follows: Cu2+ 0.030 mM, HCO3− 10 mM, AOII 0.05 mM, and H2O2 4 mM. For the UV−vis analysis, about 3 mL of the reaction solution was withdrawn and then placed back in the reactor after the measurement. 2.3. Analysis. To monitor the pollutants degradation process, solution samples taken at different time intervals were measured on a UV−vis spectrophotometer at the maximum characteristic absorption wavelength (Beijing Rayleigh Analytical Instrument Co. Ltd., China). For measurement of UV−vis spectral changes for AOII degradation with the Cu2+-HCO3− system, the scan was carried out with a rate of 20000 nm/min from 600 to 200 nm. Spectrophotometry was used for determination of hydrogen peroxide concentration with titanium oxalate.23 Electron spin resonance (ESR) spectra were recorded at room temperature using a Bruker ESR A-300 spectrometer with the following parameters: center field 3516 G, sweep width 100 G, microwave frequency 9.86 G, modulation frequency 100 kHz, microwave power 1 mW; the measurement was carried out after a reaction of 5 min by

3. RESULTS AND DISCUSSION 3.1. Activity of the Cu2+-HCO3− System. The removal of AOII under different reaction conditions is reported in Figure 1(A). Due to the presence of aromatic groups, AOII is very stable, and almost no decrease of its concentration was observed in a blank experiment. For the experiments with HCO3− in the absence of Cu2+ ions, or with Cu2+ ions in the absence of HCO3− at natural solution pH, the results also indicated very slow degradation rates. For the latter reaction at a higher pH of 8.4 adjusted by NaOH, a fast removal of AOII after the beginning of the reaction but with no further decolorization after 0.5 min was observed. These results suggest that the formed complexes between Cu2+ and OH− ions exhibit reactivity for H2O2 activation, but it can be easily deactivated. While for the Cu2+-HCO3− system, the reaction was very fast and the solution became colorless after 10 min. When using 3479

dx.doi.org/10.1021/ie403801f | Ind. Eng. Chem. Res. 2014, 53, 3478−3485

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Figure 6. Effect of radical scavengers on AOII decolorization with the Cu2+-HCO3− system (A) and DMPO spin-trapping ESR spectra of HO• radicals (B). Conditions: Cu2+ 0.03 mM, HCO3− 10 mM, H2O2 4 mM, AOII 0.05 mM, DMPO 10 mM, isopropyl alcohol 2.5 M, SOD 120 u/mL. NaN3 25 mM, 29 °C.

Figure 3. H2O2 decomposition by the Cu2+-HCO3− system. Conditions: Cu2+ 0.03 mM, HCO3− 10 mM, H2O2 4 mM, AOll 0.05 mM, TB 0.065 mM, 29 °C.

Figure 4. Distribution of copper species under different HCO3− concentrations. Conditions: Cu2+ 0.03 mM, 29 °C. Figure 7. ESR spectra of wet-grinding mixture of CuSO4, NaHCO3, AOII, and H2O2. Conditions: CuSO4 0.042 mM, NaHCO3 12.5 mM, AOII 0.062 mM, H2O2 1 mM, 29 °C.

obtained after a 24 h incubation in the presence of 10 mM Cu(II), 200 mM succinic acid, and 100 mM H2O2.21 Thus, these results clearly evidence that the system is highly reactive for AOII decolorization with H2O2 as the oxidant. Representative UV−vis spectra changes during AOII decolorization by the Cu2+-HCO3− system are depicted in Figure 1(B). The main absorption bands at 484 nm, corresponding to the n-p* transition of the azo form, and two other bands at 310 and 280 nm, attributing to the p-p* transition of the naphthalene and benzene rings, respectively,25 diminished simultaneously, indicating complete destruction of the molecular structure of AOII. In the figure the red shift of the absorbance at 484 nm increased with reaction time increasing can also be observed. This behavior might be due to the formation and collapse of intermediates under the catalytic conditions. The variation of decolorization efficiency of the Cu2+-HCO3− system under different initial AOII concentrations, ranging from 0.05 to 0.30 mM, was investigated with the results shown in Figure 2(A). Similar to other oxidation systems,26,27 the rate decreased when the initial AOII concentration increased, which might be due to the formation of complexes between Cu2+ ions and AOII or its intermediate products with more coordination numbers, and thus there are no binding equatorial positions available to bind H2O2. Even at a high concentration of 0.30 mM, near complete dcolorization was observed after 300 min.

Figure 5. Comparison of initial rates of AOII decolorization and the content of CuCO3 under different HCO3− concentrations. Conditions: Cu2+ 0.03 mM, H2O2 4 mM, AOll 0.05 mM, 29 °C.

other copper salts such as copper nitrate and copper dichloride, no significant difference in catalytic activity was observed. The degradation rate decreased with Cu2+ ions concentration decreasing; however, even at a low concentration of 0.01 mM, nearly complete decolorization was observed after 120 min. While for other copper-based Fenton-like systems, the dosage of Cu2+ ions was higher than 0.05 mM. For example, the Cu(II) (0.1 mM)/H2O2 system exhibited high Reactive Black 5 removal under neutral and alkaline conditions.24 85−95% decolorization of Remazol Brilliant Blue R, Reactive Blue, Poly B-411, Chicago Sky Blue, Evans Blue, and Methyl Orange was 3480

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Figure 10. ESI (-) mass spectra of AOII solution during degradation with the Cu2+-HCO3− system. Conditions: Cu2+ 0.03 mM, HCO3− 10 mM, H2O2 10 mM, AOII 0.2 mM, 29 °C.

strong complexing groups −NH2 and −N− and −NN− and −OH, respectively, while MR contains only one −NN− group. Similar phenomenon has been observed for the Co2+HCO3− system.14 3.2. H2O2 Decomposition. The rate of H2O2 decomposition by the Cu2+-HCO3− system in the absence and presence of the dye was investigated and is shown in Figure 3. In the presence of TB and AOII, approximately 5% and 30% of the initial H2O2 was decomposed after 30 min, respectively; whereas the rate was much faster in the absence of the dye with 82% of H2O2 decomposed after 30 min reaction. The result is different from other advanced oxidation processes such as Fenton, which suffer from rapid H2O 2 decomposition regardless of the presence of substrates.28,29 This is probably due to the formation of stable complexes between the dye and

Figure 8. UV−vis spectra changes during AOII oxidation with the Cu2+-HCO3− system under different Cu2+ concentrations. Conditions: HCO3− 5 mM, H2O2 20 mM, AOII 0.05 mM, 30 °C.

The activity of the Cu2+-HCO3− system for degradation of other organic dyes such as MO, MR, and TB was examined under similar conditions. About 90% of the pollutants was decolorized as observed in Figure 2 (B). Among the four test dyes TB, AOII, MO, and MR, TB showed the highest decolorization rate and MR showed the slowest. If we assume that the more stable the formed complex between Cu2+ ions and the dye molecular, the higher the degradation rate, this trend can be further explained in terms of the electronic properties of these pollutants, as TB and AOII exhibit two

Figure 9. (A) The initial rate of Cu(I) formation and decreasing under different HCO3− concentrations; (B) the product of the two rates. Conditions: Cu2+ 0.1 mM, H2O2 2 mM, AOII 0.05 mM, 10 °C. 3481

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Scheme 1. Proposed Pathway of AOII Degradation by the Cu2+-HCO3− System

The effect of the HCO 3 − concentration on AOII decolorization was also determined, and the results are shown in Figure 5, in which the initial rates of AOII decolorization (kAOII) were determined by the slope of the lines. From the figure it was observed that the trend of kAOII versus HCO3− was very similar with that of CuCO3 shown in Figure 4. They both first increased and then decreased with the maximum value obtained at 5.0 mM HCO3−. Under other Cu2+ concentration such as 0.06 mM, or for other dyes such as MO, very similar trends were also obtained. These results suggested that the CuCO3 complex is the main species in the Cu2+-HCO3− system attributed for the formation of reactive species and for AOII decolorization, while other species including CuHCO3+ and Cu(CO3)22_ appear to have some minor contribution. The test of Cu2+ ions at pH 8.4 adjusted by NaOH also implied that the complexes between Cu2+ and OH− at this pH were not as reactive as CuCO3. It is then suspected that HCO3− and OH− under the conditions would be weak complexes compared with CO32‑ and likely change their oxidation states at a relatively slow rate. When the Cu(CO3)22‑ complex is formed at higher HCO3− concentrations, the decolorization rate decreases as predicted, probably due to that there are no strongly binding equatorial positions available to bind H2O2 and AOII. In the literature, Mn(II)-HCO3− is a very efficient Fentonlike system for organic pollutants degradation.12,30 In the system, peroxymonocarbonate (HCO4−) can be readily generated via a pre-equilibrium between HCO3− and H2O2 and can form a complex with the metal ions. Then the high activity Mn(IV)O species can be produced from the complex responsible for the pollutant degradation. However, the contribution of HCO4− or Cu2+-HCO4− might be not so important in the Cu2+-HCO3− system, due to that their concentrations will increase with HCO3− concentration increasing,30 and are not consistent with the trend of AOII decolorization.

Figure 11. ESI (+) mass spectra of AOII solution during degradation with the Cu2+-HCO3− system. Conditions: Cu2+ 0.03 mM, HCO3− 10 mM, H2O2 10 mM, AOII 0.2 mM, 29 °C.

Cu2+ ions in the HCO3− solution, leading to the number of coordinated H2O2 and its disproportionation rate decreased. As the two dyes could be nearly completely decolorized within 10 min, the utilization of H2O2 was high in the Cu2+-HCO3− system. For example, when 0.06 mM of TB was decolorized after 10 min, only 0.094 mM of H2O2 was decomposed. If one equivalent of H2O2 gives one equivalent of reactive species that can degrade one equivalent of TB, then nearly 65% of H2O2 was utilized. 3.3. Effect of HCO3− Concentration. As mentioned in the above section of 3.1, HCO3− was necessary in the Cu2+-HCO3− system for H2O2 activation to produce the reactive species, while all the ligands including OH−, HCO3−, and CO32‑ can coordinate with Cu2+ ions to form different complexes in the HCO3− solution. In order to determine the reactivity of these species, the distribution Cu(II) species was calculated with the Visual MINTEQ 3.0 computer program. Figure 4 shows the molar fraction of copper(II) complexes at various HCO3− concentrations. It can be seen that when the HCO3 − concentration was higher than 0.5 mM, CuCO3 and Cu(CO3)22‑ were the major species, and when it increased to 5 mM, the fraction of all other species was lower than 4%. However, the trends of the two major species were quiet different, the fraction for Cu(CO3)22‑ increased on increasing the HCO3− concentration, while for CuCO3, its fraction first increased and then decreased with the highest value obtained at 5.0 mM HCO3−. 3482

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3.4. Mechanism Discussion. In the literature for oxidative degradation of pollutants using H2O2 as the oxidant, there are two major alternate mechanisms: the classical free radical and the high-valent metal oxo species; but in some case the two pathways might be compatible.31 For the copper(II) complex in the presence of H2O2, it can occur in different oxidation states such as Cu(I) and Cu(III), and the mechanism is usually complex. For example, at least two pathways were suggested to be responsible for the oxidation of quinaldine blue in the copper(II)/amino acid/H2O2 systems.32 In order to verify the mechanism for the reaction of Cu2+ ions with H2O2 in the HCO3− solution, the inhibited effects of three selective scavengers including isopropyl alcohol, SOD, and azide were measured with the results shown in Figure 6 (A). Isopropyl alcohol is a strong HO• scavenger; SOD catalyzes the dismutation of O2•‑ to H2O2 and O2, and azide reacts with both 1O2 and HO•.33,34 There was no decrease of the decolorization rate observed when the scavengers added into the reaction mixture, implying that the contribution of HO•, O2•‑, and 1O2 was negligible. Then it can be suggested that the decolorization may involve a higher oxidation state of copper, Cu(III), formed from the reaction of Cu(I) and H2O2.31 The high reactive species HO• can also be produced by the same reaction and can even be detected by ESR spin-trapping technique. As shown in Figure 6(B), in the presence of both Cu2+ and HCO3−, the ESR signal displayed a weak four-peak spectrum with an intensity ratio of 1:2:2:1, which is the typical and characteristic ESR signal of the DMPO-•OH adducts.35 As its formation rate might be very slow, the concentration was low compared with Cu(III) species and then did not contribute to the dye decolorization. These results are consistent with that in the literature: at neutral pH, Cu(III) is the main reactive species, while in acid solutions, the mechanism involves the formation of HO•.36 In order to have a better understanding of the reaction mechanism, the state change of copper species during AOII oxidation was further studied. ESR analysis of the wet-grinding mixture of CuSO4, NaHCO3, and AOII at room temperature showed a Cu(II) signal as expected,37 while after the addition of H2O2, the signal nearly disappeared (shown in Figure 7), indicating the change of Cu(II) species. The UV−vis spectra were also used to follow these reactions. From Figure 8, one can see that besides the decrease of absorbance at 484 nm, the formation and decreasing of the Abs at 520 nm can be observed, with the maximum value increasing with the increase of Cu2+ ions concentration. According to the literature, the increase of the Abs at about 520 nm can be explained by the formation of Cu(I) species from Cu(II) reduction,37 and the decrease can be attributed to the oxidation of Cu(I) species, probably with the formation of Cu(III) as the main species under the current reaction conditions. The initial rate of Cu(I) formation and decreasing under different HCO3− concentrations is shown in Figure 9. These reactions were performed under a low temperature and a low concentration of H2O2 to get a more accurate measurement. From Figure 9 (A), it can be observed that, increasing with HCO3− concentration, the rate of Cu(I) formation decreased, indicating that after being complexed with HCO3−, the rate of Cu(II) reduction decreased, and the higher the coordination number, the more difficult the Cu(II) species reduction. While for the decreasing of Cu(I) species to Cu(III), the rate first increased and then decreased, with the maximum value obtained at 10 mM HCO3−. From the analysis of the change,

it can be suggested that the formed Cu(I) complexes with HCO3− are more easily oxidized by H2O2 than Cu(I), but at a high HCO3− concentration, the number of HCO3− or CO32‑ coordinated with Cu(I) ions increased and the formed complexes would be more stable. Interestingly, if we plot the product of the two rates against HCO3− concentration (Figure 9(B)), which can represent the amount of formed Cu(III) species, a similar trend with that shown in Figure 5 can be observed, all with the highest value appearing at 5 mM of HCO3−. These results further prove that CuCO3 is more active than other copper species by the formation of Cu(III) as the main reactive species. 3.5. Intermediate Products Analysis. We have examined the ESI-MS behavior of AOII and its degradation products at different times of incubation with the Cu2+-HCO3− system. In order to get more accurate results, the initial concentration of AOII was promoted from 0.05 mM to 0.2 mM. The results at negative ion mode are displayed in Figure 10. At the beginning of the reaction, an intense ion of m/z 327 corresponding to AOII [M−H] − was observed as expected. Here the signal of AOII was so strong that the signal of HCO3− was hardly seen. During oxidative treatment, the intensity of AOII at m/z 327 decreased significantly, indicating that the dye was degraded into some intermediate products. Meanwhile, the signal peaks of HCO3− at m/z 145, 163, 189, 286, and 347 showed up, and the peak at m/z 173 emerged with regularly increasing of the relative intensity. According to the values of m/z and the suggested structures, this product can be regarded as pphenolsulfonic acid (p-PSA),17,38 which can be found as a common product that comes from the destruction of AOII in the reported literature.39 In addition, no other products such as 4-aminobenzenesulfonate and 4-nitrosobenzenesulfonate were found. All these implied that p-PSA was the only product from the broken C−N bond between the benzene ring and the azo bond. Positive ionization mode was also tested to detect positively charged species, and the results are shown in Figure 11. Due to the typical fragments interference of HCO3− at m/z 129, 324, 340, and 380, the analysis of intermediates is relatively difficult. In spite of this, we could still find some useful information. At the beginning of the reaction, an intense ion of m/z 373 was corresponded to AOII [M+Na]+, which was identical to the result of [M−H]− at m/z 327. After 60 min of incubation, new peaks at m/z 181, 197, 213, and 253 were observed. They could be attributed to 1,2-naphthaquinone and its further oxidized compounds. The toxicity of naphthaquinone from AOII degradation is even higher than that of the azo dye;38 fortunately, the compound can be easily degraded by the Cu2+-HCO3− system. Maybe there were some intermediates not recognized at present, and other technologies such as HPLC-MS, GC-MS, and ion chromatography are needed to further confirm the products in future work. According to the above analysis, a schematic degradation pathway of AOII by the Cu2+-HCO3− system was proposed and is given in Scheme 1. During oxidation, AOII was attacked by the active species, especially Cu(III), and then decomposed to p-PSA and 1,2-naphthaquinone. Afterward, the 1,2-naphthaquinone was further attracted by radicals to hydroxylated or polyhydroxylated derivatives and finally to carboxylic acids by ring-opening. 3483

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(11) Xu, A.; Li, X.; Xiong, H.; Yin, G. Efficient degradation of organic pollutants in aqueous solution with bicarbonate-activated hydrogen peroxide. Chemosphere 2011, 82, 1190. (12) Xu, A.; Shao, K.; Wu, W.; Fan, J.; Cui, J.; Yin, G. Oxidative degradation of organic pollutants catalyzed by trace manganese(II) ion in sodium bicarbonate solution. Chin. J. Catal. 2010, 31, 1031. (13) Xu, A.; Li, X.; Ye, S.; Yin, G.; Zeng, Q. Catalyzed oxidative degradation of methylene blue by in situ generated cobalt(II)bicarbonate complexes with hydrogen peroxide. Appl. Catal., B 2011, 102, 37. (14) Li, X.; Xiong, Z.; Ruan, X.; Xia, D.; Zeng, Q.; Xu, A. Kinetics and mechanism of organic pollutants degradation with cobalt− bicarbonate-hydrogen peroxide system: Investigation of the role of substrates. Appl. Catal., A 2012, 411−412, 24. (15) Yang, Z.; Wang, H.; Chen, M.; Luo, M.; Xia, D.; Xu, A.; Zeng, Q. Fast degradation and biodegradability improvement of reactive brilliant red X-3B by cobalt(II)/bicarbonate/hydrogen peroxide system. Ind. Eng. Chem. Res. 2012, 51, 11104. (16) Long, X.; Yang, Z.; Wang, H.; Chen, M.; Peng, K.; Zeng, Q.; Xu, A. Selective degradation of Orange II with the cobalt (II)− bicarbonate−hydrogen peroxide system. Ind. Eng. Chem. Res. 2012, 51, 11998. (17) Luo, M.; Lv, L.; Deng, G.; Yao, W.; Ruan, Y.; Li, X.; Xu, A. The mechanism of bound hydroxyl radical formation and degradation pathway of Acid Orange II in Fenton-like Co2+-HCO3−system. Appl. Catal., A 2014, 469, 198. (18) Mirkhani, V.; Moghadam, M.; Tangestaninejad, S.; Mohammadpoor-Baltork, I.; Rasouli, N. A comparative study of oxidation of alkanes and alkenes by hydrogen peroxide catalyzed by Cu(salen) complex covalently bound to a Keggin type polyoxometalate and its neat counterpart. Catal. Commn. 2008, 9, 2411. (19) Kirillova, M. V.; Kozlov, Y. N.; Shul’pina, L. S.; Lyakin, O. Y.; Kirillov, A. M.; Talsi, E. P.; Pombeiro, A. J. L.; Shul’pin, G. B. Remarkably fast oxidation of alkanes by hydrogen peroxide catalyzed by a tetracopper(II) triethanolaminate complex: Promoting effects of acid co-catalysts and water, kinetic and mechanistic features. J. Catal. 2009, 268, 26. (20) Conde, A.; Vilella, L.; Balcells, D.; Díaz-Requejo, M. M.; Lledós, A.; Pérez, P. J. Introducing copper as catalyst for oxidative alkane dehydrogenation. J. Am. Chem. Soc. 2013, 135, 3887. (21) Shah, V.; Verma, P.; Stopka, P.; Gabriel, J.; Baldrian, P.; Nerud, F. Decolorization of dyes with copper (II)/organic acid/hydrogen peroxide systems. Appl. Catal., B 2003, 46, 287. (22) Li, Y.; Yi, Z.; Zhang, J.; Wu, M.; Liu, W.; Duan, P. Efficient degradation of Rhodamine B by using ethylenediamine-CuCl2 complex under alkaline conditions. J. Hazard. Mater. 2009, 171, 1172. (23) Sellers, R. M. Spectrophotometric determination of hydrogen peroxide using potassium titanium(IV) oxalate. Analyst 1980, 105, 950. (24) Lee, H.; Lee, H.-J.; Sedlak, D. L.; Lee, C. pH-Dependent reactivity of oxidants formed by iron and copper-catalyzed decomposition of hydrogen peroxide. Chemosphere 2013, 92, 652. (25) Xu, A.; Li, X.; Xiong, Z.; Wang, Q.; Cai, Y.; Zeng, Q. High catalytic activity of cobalt (II)-triethylamine complex towards orange II degradation with H2O2 as an oxidant under ambient conditions. Catal. Commun. 2012, 26, 44. (26) Badawy, M. I.; Gohary, F. E.; Ghaly, M. Y.; Ali, M. E. M. Enhancement of olive mill wastewater biodegradation by homogeneous and heterogeneous photocatalytic oxidation. J. Hazard. Mater. 2009, 169, 673. (27) Natarajan, K.; Natarajan, T. S.; Bajaj, H. C.; Tayade, R. J. Photocatalytic reactor based on UV-LED/TiO2 coated quartz tube for degradation of dyes. Chem. Eng. J. 2011, 178, 40. (28) Huang, Y.; Ma, W.; Li, J.; Cheng, M.; Zhao, J. Efficient H2O2 oxidation of organic pollutants catalyzed by supported iron sulfophenylporphyrin under visible light irradiation. J. Phys. Chem. B 2003, 107, 9409. (29) Chen, F.; Ma, W.; He, J.; Zhao, J. Fenton degradation of malachite green catalyzed by aromatic additives. J. Phys. Chem. A 2002, 106, 9485.

4. CONCLUSION In summary, the Cu2+-HCO3− system has been established as a potential candidate for elimination of organic dyes in aqueous solution. During the reaction with H2O2, the complex between Cu2+ ions and HCO3− may produce a higher oxidation state of copper species, Cu(III) as the main reactive species, attributing for AOII decolorization. Among the copper species, the CuCO3 complex is suggested to have more activity than other copper species for the high-valent copper oxo species formation. A possible pathway for AOII degradation was proposed based on the analysis ESI-MS data. However, the detailed characterization of the reactive species and its ability for organic dyes mineralization should be studied in future work.

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AUTHOR INFORMATION

Corresponding Authors

*Phone: +86 27-59367334. Fax: +86 27-59367334. E-mail: [email protected] (A.X.). *E-mail: [email protected] (X.L.). Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the National Nature Science Foundation of China (21207105).

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REFERENCES

(1) Arai, H.; Berlett, B. S.; Chock, P. B.; Stadtman, E. R. Effect of bicarbonate on iron-mediated oxidation of low-density lipoprotein. Proc. Natl. Acad. Sci. U.S.A. 2005, 102, 10472. (2) Sankarapandi, S.; Zweier, J. L. Evidence against the generation of free hydroxyl radicals from the interaction of copper, zinc-superoxide dismutase and hydrogen peroxide. J. Biol. Chem. 1999, 274, 34576. (3) Carvalho do Lago, L. C.; Matias, A. C.; Nomura, C. S.; Cerchiaro, G. Radical production by hydrogen peroxide/bicarbonate and copper uptake in mammalian cells: modulation by Cu(II) complexes. J. Inorg. Biochem. 2011, 105, 189. (4) Zhou, H.; Singh, H.; Parsons, Z. D.; Lewis, S. M.; Bhattacharya, S.; Seiner, D. R.; LaButti, J. N.; Reilly, T. J.; Tanner, J. J.; Gates, K. S. The biological buffer bicarbonate/CO2 potentiates H2O2-mediated inactivation of protein tyrosine phosphatases. J. Am. Chem. Soc. 2011, 133, 15803. (5) Sun, Y.; Song, H.; Li, J.; Jiang, M.; Li, Y.; Zhou, J.; Guo, Z. Active site binding and catalytic role of bicarbonate in 1,4-dihydroxy-2naphthoyl coenzyme a synthases from vitamin K biosynthetic pathways. Biochemistry 2012, 51, 4580. (6) Yao, H.; Richardson, D. E. Epoxidation of alkenes with bicarbonate-activated hydrogen peroxide. J. Am. Chem. Soc. 2000, 122, 3220. (7) Richardson, D. E.; Yao, H.; Frank, K. M.; Bennett, D. A. Equilibria, kinetics, and mechanism in the bicarbonate activation of hydrogen peroxide: oxidation of sulfides by peroxymonocarbonate. J. Am. Chem. Soc. 2000, 122, 1729. (8) Richardson, D. E.; Reqino, C. A.; Johnson, J. V. Methionine oxidation by peroxymonocarbonate, a reactive oxygen species formed from CO2/bicarbonate and hydrogen peroxide. Free Radical Biol. Med. 2003, 35, 1538. (9) Balagam, B.; Richardson, D. E. The mechanism of carbon dioxide catalysis in the hydrogen peroxide N-oxidation of amines. Inorg. Chem. 2008, 47, 1173. (10) Bakhmutova-Albert, E. V.; Yao, H.; Denevan, D. E.; Richardson, D. E. The kinetics and mechanism of peroxymonocarbonate formation. Inorg. Chem. 2010, 49, 11287. 3484

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Industrial & Engineering Chemistry Research

Article

(30) Ember, E.; Gazzaz, H. A.; Rothbart, S.; Puchta, R.; van Eldik, R. MnII-A fascinating oxidation catalyst: Mechanistic insight into the catalyzed oxidative degradation of organic dyes by H2O2. Appl. Catal., B 2010, 95, 179. (31) Pham, A. N.; Xing, G.; Miller, C. J.; Waite, T. D. Fenton-like copper redox chemistry revisited: Hydrogen peroxide and superoxide mediation of copper-catalyzed oxidant production. J. Catal. 2013, 301, 54. (32) Lin, T.-Y.; Wu, C.-H. Activation of hydrogen peroxide in copper(II)/amino acid/H2O2 systems: effects of pH and copper speciation. J. Catal. 2005, 232, 117. (33) Xu, Z.; Jing, C.; Li, F.; Meng, X. Mechanisms of photocatalytical degradation of monomethylarsonic and dimethylarsinic acids using nanocrystalline titanium dioxide. Environ, Sci. Technol 2008, 42, 2349. (34) Duan, J.; Liu, Y.; Pan, X.; Zhang, Y.; Yu, J.; Nakajim, K.; Taniguchi, H. High photodegradation efficiency of Rhodamine B catalyzed by bismuth silicate nanoparticles. Catal. Commun. 2013, 39, 65. (35) Chen, M.; Peng, K.; Wang, H.; Yang, Z.; Zeng, Q.; Xu, A. High performance of a simple cobalt(II)−monoethanolamine complex for orange II degradation with H2O2 as an oxidant at ambient conditions. Chem. Eng. J. 2012, 197, 110. (36) Meyerstein, D. Trivalent copper. I. Pulse radiolytic study of the chemical properties of the aquo complex. Inorg. Chem. 1971, 10, 638. (37) Hoover, J. M.; Ryland, B. L.; Stahl, S. S. Mechanism of copper(I)/TEMPO-catalyzed aerobic alcohol oxidation. J. Am. Chem. Soc. 2013, 135, 2357. (38) Rismayani, S.; Fukushima, M.; Ichikawa, H.; Tatsumi, K. Decolorization of orange II by catalytic oxidation using iron (III) phthalocyanine-tetrasulfonic acid. J. Hazard. Mater. 2004, 114, 175. (39) Chen, X.; Chen, J.; Qiao, X.; Wang, D.; Cai, X. Performance of nano-Co3O4/peroxymonosulfate system: Kinetics and mechanism study using Acid Orange 7 as a model compound. Appl. Catal., B 2008, 80, 116.

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