Catalyzed oxidative degradation of methylene blue by in situ generated cobalt (II)-bicarbonate complexes with hydrogen peroxide

Applied Catalysis B: Environmental 102 (2011) 37–43

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Applied Catalysis B: Environmental journal homepage: www.elsevier.com/locate/apcatb

Catalyzed oxidative degradation of methylene blue by in situ generated cobalt (II)-bicarbonate complexes with hydrogen peroxide Aihua Xu a,∗ , Xiaoxia Li b , Shuang Ye c , Guochuan Yin c , Qingfu Zeng a a b c

Engineering Research Center for Clean Production of Dyeing and Printing, Ministry of Education, Wuhan Textile University, 1 Fangzhi Road, Wuhan, Hubei 430073, China Department of Environmental Science, College of Resource and Environmental Science, Wuhan University, Wuhan 430079, China School of Chemistry and Chemical Engineering, Huazhong University of Science and Technology, Wuhan 430074, China

a r t i c l e

i n f o

Article history: Received 23 July 2010 Received in revised form 7 November 2010 Accepted 12 November 2010 Available online 19 November 2010 Keywords: Cobalt Bicarbonate Hydrogen peroxide Degradation Methylene blue Hydroxyl radicals

a b s t r a c t Oxidative degradation of methylene blue (MB) by Co2+ –HCO3 − system with H2 O2 in aqueous solution was studied. Nearly complete decolorization of the dye was obtained in less than 50 min in diluted NaHCO3 solution (25 mM) in the presence of only 20 ␮M Co2+ ions. Meanwhile, the conjugated structure and phenyl rings of the MB molecule were destroyed or even broken down into small organic acids and inorganic ions, as indicated by FT-IR spectra and ion-chromatography. Photoluminescence probing and radical scavenging technologies suggested that the reaction of MB degradation in this system mainly involved the generation and participation of hydroxyl radicals. Furthermore, by cyclovoltammetric measurements, the in situ formed different complexes between Co2+ and HCO3 − were observed at different HCO3 − concentrations, and the complex [Co(HCO3 )]+ formed at intermediate HCO3 − concentrations (5–10 mM) was suggested to be more active than the others. © 2010 Elsevier B.V. All rights reserved.

1. Introduction Industrial activities generate wastewater with a wide variety of contaminants including hydrocarbons, phenol and its derivatives, sulfur, nitrogen and halogen containing organic compounds and heavy metals [1]. Among them, textile industries produce a great deal of dye wastes, and have presented significant environmental problems, due to that the dyes are highly colored, designed to resist chemical, biochemical and photochemical degradation. Therefore the degradation of organic dyes has attracted much attention. H2 O2 is one of the most commonly used oxidants owing to its eco-friendly nature. In recent decades, activation of H2 O2 by transition-metal ions for organic dyes degradation has been explored actively [2–5]. The most widely used process is homogeneous Fenton reaction (Fe2+ /H2 O2 ), in which the organic pollutants are degraded by the reactive oxygen species such as hydroxyl radical and hydroperoxyl radical [2,3]. Collins and co-workers have described a novel and highly effective method based on tetraamido macrocylic ligand (TAML) iron to destroy Orange II into small biodegradable and non-toxic organic products by H2 O2 [6]. After coordinated with succinic acid, copper ions were highly efficient in the decolorization of azo and anthroquinone-based dyes with H2 O2 through hydroxyl radical similar to Fenton reagent [7]. Co2+ ions alone do not effi-

∗ Corresponding author. E-mail address: [email protected] (A. Xu). 0926-3373/$ – see front matter © 2010 Elsevier B.V. All rights reserved. doi:10.1016/j.apcatb.2010.11.022

ciently generate • OH radicals from H2 O2 [8], and there are very few papers about organic pollutants degradation by Co2+ –H2 O2 system. While after properly chelated it seems clear that the oxygen radicals are indeed generated [9,10]. Some cobalt complexes with different phthalocyanines [11–13] and ascorbic acid [14] have been reported to be efficient catalysts for the oxidation of organic pollutants. However, cost effectiveness and robustness are often the major obstacles associated with the application of these metal complexes in industrial processes. Bicarbonate (HCO3 − ), a relative nontoxic anion, is one of the most abundant anions in natural water. It is also an efficient activator for H2 O2 to generate active oxygen species such as peroxymonocarbonate (HCO4 − ) [15,16] and singlet oxygen (1 O2 ) [17]. This simple and green system has been successfully used for many oxidation reactions [18,19]. In addition, HCO3 − is an important complexing ligand for some metal ions, leading to significant changes in their redox potentials. For example, upon addition of NaHCO3 to an aqueous solution of MnII , the potential for oxidation of MnII to MnIII was shifted from 1.19 to 0.63 V at fixed pH 8.3 as a result of formation of (bi)carbonate complexes [20]. The Mn2+ -bicarbonate complexes have been hypothesized to play an important role in the evolutionary origin of oxygenic photosynthesis in the archean era [21]. The objective of the present study was to evaluate the efficiency of Co2+ –HCO3 − system for degradation of the selected organic dye, methylene blue (MB) in aqueous solution with H2 O2 . The system has been shown to be an effect catalyst for H2 O2 decomposition

A. Xu et al. / Applied Catalysis B: Environmental 102 (2011) 37–43

2. Experimental 2.1. Materials MB was purchased from Shanghai Matex Chemical Co. Ltd., and used without further purification. Cobalt acetate, hydrogen peroxide (30%, w/w), sodium bicarbonate, terephthalic acid and other reagents were of analytical grade. The sample solutions were prepared using deionized water throughout the experiments. 2.2. General procedure Oxidation reactions were performed in a laboratory scale bath reactor equipped with a magnetic stirrer. In a typical degradation experiment, MB and cobalt acetate were combined in a NaHCO3 solution and stirred at 25 ◦ C; then the reaction was initialized by adding H2 O2 . In the reaction mixture, the concentrations of the reagents were: NaHCO3 25 mM, Co2+ 20 ␮M, MB 0.134 mM and H2 O2 20 mM. To monitor the organic dye degradation process, 1 ml of aqueous solution were withdrawn from the reaction mixtures periodically, and then diluted to the total volume of 10 ml with deionized water immediately. A decrease in the MB absorption band at the maximum (664 nm) was monitored by a Varian Cary 500 Scan UV–vis spectrophotometer. The percentage of degradation was reported as Ct /C0 , where Ct is the absorption of MB at each reaction time and C0 is the absorption of the starting concentration. 2.3. Analysis The mineralization degree of MB after degradation was evaluated by measuring total organic carbon (TOC) concentration with an Apolllo 9000 TOC analyzer. The chemical oxygen demand (COD) value of the sample solution was measured by the standard determination procedure using potassium dichromate as oxidant. For IR analysis, the collected solution before or after degradation, was adjusted to a pH value of 2.0 with HCl to completely decompose HCO3 − , and then freeze-dried. The resulting samples were recorded as KBr pellets on an FTIR 170S spectrometer. The analysis of small organic acids and inorganic ions after the degradation was performed with an ionic chromatography (ICS-9000). For the above-mentioned analyses, the degradation was performed at higher concentrations of MB (1.07 mM) and H2 O2 (0.1 M). The detection of • OH radical formation was carried out on a Hitachi F-4500 fluorescence spectrophotometer at an excitation wavelength of 315 nm and an emission wavelength of 425 nm. Before analysis, the selected amount of terephthalic acid was dissolved in a NaOH aqueous solution with the final pH of 7.0 adjusted by HCl. CV experiments were recorded on a 797-VA computrace using a three-electrode cell with a hanging mercury drop as working electrode and a Pt plate as counter electrode. The working volume of 10 ml was deaerated by passing a stream of high purity N2 through

1.0

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min 0 2 4 9 14 19 24 29 39 49

0.8 Ct / C0

[22], luminescence of luminol [23] and peroxidation of NADH [24]. However, its application in elimination of organic pollutants has not been previously reported. Here, our study reveals that MB can be effectively degraded with trace Co2+ ions as the catalyst in diluted NaHCO3 solution. By cyclovoltammetric (CV) measurements, in situ formed different complexes between Co2+ and HCO3 − were observed at different HCO3 − concentrations. The relationship between chemical structure of the Co2+ –HCO3 − complex and its catalytic activity was also discussed. This provides a simple, inexpensive and relatively nontoxic ligand for activation of Co2+ ions compared to other cobalt complexes [9–14]. It also provides us an effective and economic system ideal for the treatment of toxic and nonbiodegradable organic pollutants in water.

0.6 0.4 0.2 0.0

0

Abs

38

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20 30 Time (min)

40

50

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0.0 200

300

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500

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Wavelength (nm) Fig. 1. UV–vis spectral changes for MB degradation with Co2+ –HCO3 − system, MB 0.134 mM, Co2+ 20 ␮M, HCO3 − 25 mM, H2 O2 20 mM, 25 ◦ C.

the solution for 5 min prior to the measurements. The scan rate was 50 mV/s. Potentials were measured in a 0.1 M NaCl electrolyte solution at room temperature and are reported vs. an Ag/AgCl electrode. 3. Results and discussion 3.1. Catalytic activity for MB degradation Representative UV–vis spectra changes observed during MB degradation by Co2+ –HCO3 − system in the presence of H2 O2 are depicted in Fig. 1. The intense absorbance of MB at 250, 296 and 664 nm dropped rapidly and disappeared almost completely within 50 min in the presence of 20 ␮M Co2+ ions. The inset in Fig. 1 shows lineal decreasing of absorbance at 664 nm during the first 25 min and then exponential. This behavior, also observed for other H2 O2 based degradation systems [25], can be explained by the fact, that the formation of the catalytic active species is the rate-determining step. Meanwhile, the blue-shifted of the absorbance at 664 nm observed as the reaction time was prolonged, implying the formation of demethylated dyes [26]. The degradation rate decreased with cobalt concentration decreasing; however, even at a low concentration of 5 ␮M, complete decolorization was also observed after 120 min. The value can meet the national emission standard for cobalt pollutants in wastewater (lower than 1 mg/L) in China (GB 25467-2010), and between the amounts of cobalt ions leached from Co3 O4 catalysts that used for activation of peroxymonosulfate in recent studies [27,28]. These results indicated efficient degradation of the dye molecules by the Co2+ –HCO3 − system. Control experiments with HCO3 − in the absence of Co2+ ions, or with Co2+ ions in the absence of HCO3 − , implied no degradation of MB under these experimental conditions (shown in Fig. 2). Furthermore, instead of NaHCO3 with other basic solutions at the same pH, such as NaOH, Na2 CO3 or K2 HPO4 did not show the observed degradation. The catalytic activity of the system decreased dramatically when Co2+ was replaced by other transition metal ions, as shown in Fig. 3. In the case of Ag+ , nearly no degradation was observed, as Ag+ ion is not a good catalyst for the activation of H2 O2 to active radicals [29]. While in the case of Fe2+ and Cu2+ , the concentration of MB slowly decreased, with the decolorization of 9% and 26% after 90 min respectively. Mn2+ ions has been reported as an efficient catalyst for organic dyes degradation with H2 O2 in the presence of HCO3 − , and the in situ formed high-valent manganese intermediate was suggested as the key reactive species [30,31]. However, under the experimental conditions, the decolorization was only 49% after

A. Xu et al. / Applied Catalysis B: Environmental 102 (2011) 37–43

39

1320 1240 1384 1224

1400

1.0 1600

90

1482

1182

0.6 (1)

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(2) (3) (4)

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(5)

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Ct / C0

0.8

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81 After

72

(6) 0.0 0

20

40

60

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120

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Fig. 2. Degradation of MB under different conditions: (1) NaHCO3 − + H2 O2 ; (2) Co2+ /H2 O2 ; (3) Co2+ /H2 O2 + NaOH (pH 8.2); (4) Co2+ /H2 O2 + Na2 CO3 (pH 8.2); (5) Co2+ /H2 O2 + K2 HPO4 (pH 8.2); (6) Co2+ /HCO3 + H2 O2 . MB 0.134 mM, Co2+ 20 ␮M, HCO3 − 25 mM, H2 O2 20 mM, 25 ◦ C.

90 min. It was revealed that the Co2+ –HCO3 − system performed the higher catalytic activity for MB degradation than other tested metals ions. 3.2. Degradation products analyses The products remained in solution after MB degradation by the Co2+ –HCO3 − system were monitored by FT-IR spectra and ionic chromatography, as it is useful for evaluating the efficiency of the catalytic system. In order to get more accurate results, the initial concentration of MB was promoted from 0.134 mM to 1.07 mM. Accordingly, the concentration of H2 O2 as well as the reaction time needed for complete degradation MB increased. Fig. 4 shows the FT-IR spectra of MB before and after degradation. Prior to degradation, MB exhibited the prominent bands of the characteristic C N central ring stretching at 1600 cm−1 , C C side ring stretching at 1482 cm−1 , multiple ring stretching at 1384 cm−1 , CAr –N (the bond between the side aromatic ring and the nitrogen atom) stretching at 1320 cm−1 , and N–CH3 stretching at 1240 and 1182 cm−1 [32]. After treatment by the Co2+ –HCO3 − system, the disappearance of the above-mentioned characteristic bands indicated complete destruction of MB molecule by breaking of central and side aromatic rings and demethylation. Meanwhile, two new intense peaks at 1400

1500

900

Fig. 4. FT-IR spectrum of MB before and after degradation, MB 1.07 mM, Co2+ 20 ␮M, HCO3 − 25 mM, H2 O2 100 mM, 25 ◦ C.

and 1224 cm−1 appeared, which might be due to the stretching vibration of COO− and SO4 2− respectively [32], suggesting that the carboxylic acids and SO4 2− were the main final products. The ion chromatography after MB degradation is depicted in Fig. 5. As it shown, some measurable reaction products were identified, such as CH3 COO− , SO4 2− , SO3 2− and NO3 − , consistent with the analysis of FT-IR spectra. The degree of MB mineralization by Co2+ /HCO3 − + H2 O2 system was evaluated by determining the TOC and COD values. After degradation under the same experimental conditions as described in FT-IR spectra analysis, about 17% and 41% of TOC and COD removal were obtained respectively, indicating that considerable amount of carbon in MB were converted to CO2 and other intermediates. As the production of organic acid was detected during degradation, the decrease of HCO3 − amount is expected. The value of total inorganic carbon (TIC) dropping 24% after reaction confirmed that part of HCO3 − ions were neutralized by the products.

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Co 2+ Mn 2+ Cu 2+ Fe + Ag

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Time (min) Fig. 3. Degradation of MB with different metal ions in the presence of HCO3 − and H2 O2 , MB 0.134 mM, metal ions 20 ␮M, HCO3 − 25 mM, H2 O2 20 mM, 25 ◦ C.

4

6

Time (min)

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Fig. 5. Products detected by ion chromatography after MB degradation, MB 1.07 mM, Co2+ 20 ␮M, HCO3 − 25 mM, H2 O2 100 mM, 25 ◦ C.

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A. Xu et al. / Applied Catalysis B: Environmental 102 (2011) 37–43

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Fig. 6. Photoluminescence spectral changes as function reaction time, (A) TA 0.134 mM, Co2+ 20 ␮M, HCO3 − 25 mM, 25 ◦ C; (B) MB 0.134 mM, TA 0.134 mM, Co2+ 20 ␮M, HCO3 − 25 mM, 25 ◦ C.

3.3. Hydroxyl radicals determination

Ct / C0

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0.4 2-propanol 0 g 2-propanol 0.4 g 2-propanol 4.0 g methanol 4.0 g

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Time (min) Fig. 7. Inhibited effects of 2-propanol and methanol on MB degradation by Co2+ –HCO3 − system, MB 0.134 mM, Co2+ 20 ␮M, HCO3 − 25 mM, H2 O2 100 mM, 25 ◦ C.

10000

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HCO3 (mM) 0 5 10 15 25 50 100 0 with MB

8000

Fluorescent intensity

H2 O2 is a common oxidizing agent and is often converted to free radicals along with UV irradiation or the redox reaction of transition-metal ions. In order to identify the reactive oxygen species formed in the Co2+ –HCO3 − system, disodium salt of terephthalic acid (NaTA) photoluminescence probing technology and radical scavenging measurements were carried out. NaTA could react with • OH to give 2-hydroxyterephthalic acid (HTA), which exhibits a bright stable fluorescence [33]. This reaction is unaffected by the presence of other reactive species such as H2 O2 , HO2 • and O2 •− , so it could be used as a sensitive probe in detecting • OH radicals [34]. Fig. 6 shows the fluorescence spectra of the solution containing Co2+ , HCO3 − , H2 O2 and NaTA. It can be seen that the fluorescence intensity increases sharply to 2500 within 10 min, implying that • OH radicals were indeed generated in the system. Moreover, when MB and NaTA were simultaneously added into the solution, the generated fluorescence significantly decreased, with the intensity only increasing to 150 within 50 min. One of the reasons may be that part of • OH radicals reacted rapidly with MB, as the decolorization was observed during fluorescence measurement. Another possible reason was that the rate of • OH radicals production decreased after the formation of complex between MB and Co2+ , for MB is one of the most frequently used counter ions for ionpair formation and has been postulated for removal of metal ions from waste water [35]. The inhibited effects of 2-propanol and methanol, which are known to be effective scavenging agents for hydroxyl radicals [36,37], were also measured. As shown in Fig. 7, the rate of MB degradation became slower at relatively low concentration of 2-propanol; while at high concentrations of 2-propanol and methanol, great inhibition was observed. These results further proved that MB degradation in Co2+ –HCO3 − system mainly involved the generation and participation of • OH radicals. As mentioned in the above section of 3.1, HCO3 − was absolutely necessary in the Co2+ –HCO3 − system; it was revealed here that the amount of • OH radicals measured by the fluorescent intensity, was also strongly influenced by its concentrations. The fluorescent intensity, which was proportional to the amount of • OH radicals [33], detected at different HCO3 − concentrations is shown in Fig. 8. It can be seen that the intensity in the absence of HCO3 − was considerably low, and nearly no fluorescence produced after MB was added. On increasing HCO3 − concentration, the intensity increased.

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Time (min) Fig. 8. Effect of NaHCO3 concentration on fluorescence intensity produced by Co2+ –HCO3 − system, TA 0.134 mM, Co2+ 20 ␮M, H2 O2 20 mM, 25 ◦ C.

A. Xu et al. / Applied Catalysis B: Environmental 102 (2011) 37–43

1.0

-3.50E-008 -3.00E-008

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Time (min) Fig. 9. Effect of NaHCO3 concentration on MB degradation by Co2+ –HCO3 − system, MB 0.134 mM, Co2+ 20 ␮M, HCO3 − 25 mM, H2 O2 20 mM, 25 ◦ C.

For example, when 5 mM HCO3 − was present in the solution the fluorescent intensity increased sharply to 6100 after 10 min. However, when further increasing HCO3 − concentration, the intensity at a given time dramatically decreased, but no significant change was observed when the concentration of HCO3 − was higher than 50 mM. The effect of HCO3 − concentration on MB degradation was also determined and is shown in Fig. 9. In Fig. 10, the initial rates of • OH radicals production (k• OH ) and MB degradation (kMB ), determined by the slop of the lines shown in Figs. 8 and 9 respectively, were plotted against HCO3 − concentration. A quite similar behavior between k• OH and kMB was observed, in which they both first increased and then decreased. But the highest k• OH value obtained at HCO3 − concentration of 10 mM, was slightly different from the highest kMB value obtained at HCO3 − concentration of 15 mM, probably due to that MB took part in the reaction of • OH radicals production. The results provided a much stronger evidence that MB was mainly degraded by • OH radicals. Besides, the high catalytic activity obtained at so low concentrations of NaHCO3 makes the system a promising method of organic dyes abatement. 3.4. In situ complex formation between Co2+ and HCO3 − As Co2+ ions were actually not able to degrade MB with H2 O2 in the absence of HCO3 − , the complex between HCO3 − and Co2+ may form and serve as the catalyst to active H2 O2 to • OH radi-

Fig. 11. Voltage–current curves of Co2+ reduction in aqueous solution of 0.1 M NaCl at different concentrations of NaHCO3 : 0 mM (curve 0); 2.5 mM (curve 1); 5.0 mM (curve 2); 7.5 mM (curve 3); 10.0 mM (curve 4); 15.0 mM (curve 5); 25.0 mM (curve 6); 37.5 mM (curve 7); 50.0 mM (curve 8); 100 mM (curve 9); Co2+ 20 ␮M, 25 ◦ C. The curves are shifted on the x axis for clarity. The peak potentials are indicated at the top of each curve.

cals in the Co2+ –HCO3 − system. This assumption was confirmed by CV measurements at different HCO3 − concentrations. The voltagecurrent curves corresponding to the electrochemical reduction of Co2+ ions in the voltammetric experiments are shown in Fig. 11, in which the peak potential was used for data analysis. It can be seen that the CV of Co2+ in the absence of HCO3 − exhibited the peak potential at −1.132 V, which corresponded to the two-electron reduction of Co2+ to Co0 , in agreement with the literature values [38]. As bicarbonate is a negatively charged ligand, the formed complex is expected to have a more negative reduction potential compared to the fully aquated Mn2+ . Accordingly, it was found that the peak potential in the Co2+ –HCO3 − solutions shifted to more negative potentials concomitantly with the peak current decreasing at increasing HCO3 − concentrations, indicating the complex formation between Co2+ and HCO3 − . According to the Lingane equation [39], the presence of different complex species at different HCO3 − concentrations was revealed by their different slops when plotting the peak potential as a logarithmic function of HCO3 − concentration (shown in Fig. 12). Similar to the Mn–HCO3 − complex [21], at low HCO3 − concentrations (0–5 mM), Co2+ ions is mainly presented as the 1.22

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NaHCO3 concentration (mM) Fig. 10. Plot of the initial rates of • OH radicals production (k• OH ) and MB degradation (kMB ) against HCO3 − concentration.

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Co

2+ aq

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[Co(HCO 3)] -2.0

+

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log[HCO3- ] Fig. 12. Plot of peak potentials from Fig. 11 as a logarithmic function of HCO3 − concentration.

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A. Xu et al. / Applied Catalysis B: Environmental 102 (2011) 37–43

CoII + HCO3-

[CoII(HCO3-)]+

(1)

[CoII(HCO 3-)]+ + MB

[CoII(HCO 3-)(MB)]+

(2)

[CoII(HCO 3-)(MB)]+ + H2O2

[Co II(HCO3-)(MB)(H2O2)]+

(3)

[CoII(HCO 3-)(MB)(H 2O2)]+

[CoIII(HCO3-)(MB)( OH)]2+ + OH-

(4)

[CoIII(HCO3-)(MB)( OH)]2+

[CoIII(HCO3-)]2+ + Oxidation products

(5)

[CoIII(HCO3-)]2+ + H2O2

[Co II(HCO 3-)]+ + HOO + H +

(6)

Scheme 1. . Proposed catalyzed mechanism for MB degradation by Co2+ ions with H2 O2 at intermediate HCO3 − concentrations.

free aquated form, for the peak potential was insensitive to added HCO3 − . At intermediate HCO3 − concentrations (5–10 mM), [Co(HCO3 )]+ is the dominant species in solution, indicated by a slop of 0.03321V/log[HCO3 − ]. At the highest HCO3 − concentrations (10–100 mM), a linear phase of the potential yielding a slope of 0.06759V/log[HCO3 − ] was observed. Thus, [Co(HCO3 )2 ] is the most abundant species at these high concentrations. It should be noted that under the experimental conditions, there was not any precipitate (insoluble CoCO3 or Co(OH)2−2x (CO3 )x ) observed during testing, even when the concentration of Co2+ up to 1 mM in 100 mM bicarbonate solution. Besides, the stability of the Co2+ –HCO3 − complex could be enhanced with the presence of organic dye as a coordinating ligand [30,31]. Comparing the coordination number of Co2+ –HCO3 − complex, the efficiency of • OH radicals production and the rate of MB degradation at different HCO3 − concentrations, provided valuable mechanistic information for MB degradation by the Co2+ –HCO3 − system. At low HCO3 − concentrations (0-5 mM), the aquated Co2+ ion did not efficiently generate free radicals from H2 O2 , as it can form dimers or oligomers in solution, where some coordination positions are unavailable [40]. The higher rates of • OH radicals production and MB degradation observed at intermediate HCO3 − concentrations (5–10 mM), suggest that the [Co(HCO3 )]+ complex is relative more stable and can fast produce the active • OH radicals. When the [Co(HCO3 )2 ] complex is formed at highest HCO3 − concentrations (10–100 mM), the rates of • OH radicals production and MB degradation were expected to decrease, as there are no strongly binding equatorial positions available to bind the H2 O2 . In addition, it has been reported that HCO3 − ions are a radical scavenger that can trap • OH radicals [41]. The decrease of the concentration of • OH radicals and the rate of MB degradation could be observed with the concentration of HCO3 − increasing.

3.5. A possible catalytic mechanism Based on the obtained experimental results, Scheme 1 was proposed as a tentative pathway for oxidative degradation of MB catalyzed by Co2+ ions with H2 O2 at intermediate HCO3 − concentrations. The first step involves the initial formation of the [CoII (HCO3 − )]+ complex between Co2+ and HCO3 − , and the next is coordination of MB to the complex to produce [CoII (HCO3 − )(MB)]+ . When H2 O2 is present in aqueous solution, the complex may react directly with it to form the stable [CoII (HCO3 − )(MB)(H2 O2 )]+ complex and then dissociate it to a transient caged • OH radical through transfer of an electron from Co2+ , similar to the mechanism proposed by Yim for H2 O2 dismutation by Mn2+ –HCO3 − system [42]. Once formed, the high active • OH radical reacts rapidly with the bonded substrate and degrade it effectively. As a result, the rate of MB degradation is determined by the rate of • OH radical formation, giving the zero-

order substrate dependence. The [CoII (HCO3 − )]+ complex can be finally regenerated through the interaction between [CoIII (HCO3 − )] 2+ and H O [43]. 2 2 4. Conclusions We have demonstrated an efficient system, Co2+ –HCO3 − for elimination of MB in aqueous solution with H2 O2 . Similar to Fenton reagent, the system can produce hydroxyl radicals, as confirmed by photoluminescence probing and radical scavenging measurements. It seems that the in situ formed different complexes between Co2+ and HCO3 − were responsible for the observed rates of • OH radicals production and MB degradation at different HCO3 − concentrations. When Co2+ ions were replaced by other transition metal ions or using other basic solutions instead of HCO3 − at the same pH, the high activity dramatically decreased. Compared to other metal–ligand systems reported so far, using HCO3 − as the ligand is relative simple and economic, making the applied treatment method a promising way of pollutant abatement, though part of HCO3 − can be neutralized by the produced organic acids. However, detailed mechanistic studies are still essential to provide more clear information when MB is present and forms complex with Co2+ ; and the catalytic activity of the system towards other organic pollutants should be tested in future work. Acknowledgments This work was supported by the National High Technology Research and Development Program of China (2009AA063904). References [1] P. Bautista, A.F. Mohedano, J.A. Casas, J.A. Zazo, J.J. Rodriguez, J. Chem. Technol. Biotechnol. 83 (2008) 1323–1338. [2] S. Caudo, G. Centi, C. Genovese, S. Perathoner, Top. Catal. 40 (2006) 207–219. [3] M.S. Lucas, J.A. Peres, J. Hazard. Mater. 168 (2009) 1253–1259. [4] J.M. Monteagudo, A. Durán, C. López-Almodóvar, Appl. Catal. B 83 (2008) 46–55. [5] A.H. Xu, H. Xiong, G.C. Yin, J. Phys. Chem. A 113 (2009) 12243–12248. [6] N. Chahbane, D.L. Popescu, D.A. Mitchell, A. Chanda, D. Lenoir, A.D. Ryabov, K.W. Schramm, T.J. Collins, Green Chem. 9 (2007) 49–57. [7] V. Shah, P. Verma, P. Stopka, J. Gabriel, P. Baldrian, F. Nerud, Appl. Catal. B 46 (2003) 287–292. [8] S. Leonard, P.M. Gannett, Y. Rojanasakul, D. Schwegler-Berry, V. Castranova, V. Vallyathan, X.L. Shi, J. Inorg. Biochem. 70 (1998) 239–244. [9] X.G. Shi, N.S. Dalal, K.S. Kasprzakt, Chem. Res. Toxicol. 6 (1993) 277–283. [10] P.M. Hanna, M.B. Kadiiska, R.P. Mason, Chem. Res. Toxicol. 5 (1992) 109–115. [11] B. Agboola, K.I. Ozoemena, T. Nyokong, J. Mol. Catal. A: Chem. 227 (2005) 209–216. [12] M.A. Zanjanchi, A. Ebrahimian, M. Arvand, J. Hazard. Mater. 175 (2010) 992–1000. [13] W.Y. Lu, W.X. Chen, N. Li, M.H. Xu, Y.Y. Yao, Appl. Catal. B 87 (2009) 146–151. [14] P. Verma, P. Baldrian, F. Nerud, Chemosphere 50 (2003) 975–979. [15] D.E. Richardson, H.R. Yao, K.M. Frank, D.A. Bennett, J. Am. Chem. Soc. 122 (2000) 1729–1739. [16] B. Balagam, D.E. Richardson, Inorg. Chem. 47 (2008) 1173–1178.

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