Oxidative Removal of Bisphenol A by Manganese Dioxide: Efficacy

Apr 14, 2009 - Bisphenol A (BPA) is a ubiquitous environmental contaminant with endocrine disruption potential. In this study, exploiting the outstand...
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Environ. Sci. Technol. 2009, 43, 3860–3864

Oxidative Removal of Bisphenol A by Manganese Dioxide: Efficacy, Products, and Pathways K U N D E L I N , * ,†,‡ W E I P I N G L I U , ‡ A N D JAY GAN† Department of Environmental Sciences, University of California, Riverside, California 92521, and Research Center of Green Chirality, Zhejiang University of Technology, Hangzhou 310032, China

Received January 26, 2009. Revised manuscript received March 14, 2009. Accepted March 17, 2009.

Bisphenol A (BPA) is a ubiquitous environmental contaminant with endocrine disruption potential. In this study, exploiting the outstanding oxidative capacity of manganese dioxide (δMnO2), we explored for the first time the efficacy and mechanisms of BPA removal by MnO2. In aqueous solutions, MnO2 demonstrated an extremely efficient capacity to remove BPA. Nearly all BPA (>99%) was eliminated in 6 min in a pH 4.5 solution initially containing 800 µM MnO2 and 4.4 µM BPA. While humic acid showed negligible inhibition on BPA removal, coexisting metal ions such as Mn2+, Ca2+, Mg2+, and Fe3+ displayed suppressive effects and the inhibitive capacity followed the order Mn2+ > Ca2+ > Mg2+ ≈ Fe3+. A total of 11 products or intermediates were indentified and a detailed reaction scheme was suggested. The products could be ascribed to a suite of reactions of radical coupling, fragmentation, substitution, and elimination, triggered by the BPA radical formed through electron transfers to MnO2. The exceptional efficiency of MnO2 in removing BPA represents a potential use of MnO2 to treat waters containing phenolic compounds and also suggests a potentially important role of oxide-facilitated abiotic transformations in BPA attenuation in natural soil and sediment environments.

Introduction Bisphenol A [2,2-bis(4-hydroxyphenyl)propane or BPA] is an important industrial chemical that is primarily used as an intermediate in the production of polycarbonate plastic and epoxy resins (1). Products of BPA range from polycarbonate bottles to the linings of metallic cans and from dental sealants to cell phones (1). Due to the continuous growing demand, BPA production has rapidly increased in recent years. The total global production capacity of BPA was about 3.2 × 109 kg in 2003, about one-third of which was in the United States (2). Bisphenol A can be released into the environment from bottles, packaging, landfill leachates, paper, and plastic plants (3-5). According to the U.S. Environmental Protection Agency’s Toxics Release Inventory (6), about 2.0 × 106 kg of BPA was released into the environment in 2006 across the United States due to on-site and off-site disposal by various industrial sectors, municipal wastewater treatment plants, and landfill sites. * Corresponding author phone: 951-827-3860; fax: 951-827-3993; e-mail: [email protected]. † University of California. ‡ Zhejiang University of Technology. 3860

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Controversy surrounding BPA mainly involves its potential human health risk due to its estrogenic activities (7-10). Such estrogenic potencies can mimic or interfere with hormonal activities and thus have adverse effects on ecosystems and human health by disrupting growth, development, and reproduction (11, 12). For example, Lee et al. (13) found that women had low levels of BPA (2.8 ( 3.7 µg/L residue in urine samples from 125 pregnant women) during early pregnancy may contribute to the risks of retarded fetal growth. Another recent study (14) found that mice treated with 400 ng of BPA per day from 11.5 days of gestation onward resulted in severe genetic abnormality and reproductive problems. Similarly, Hunt et al. (15) demonstrated that laboratory mice exposed to about 20 µg/L of BPA developed chromosomal abnormalities in eggs. The increasing evidence of adverse effects caused by low levels of BPA raises challenges to the environmental regulation of BPA. Bisphenol A is ubiquitous in the aquatic environment. According to studies from the United States, Germany, Japan, Spain, China, and The Netherlands, BPA levels in river water were 8 µg/L or less but reached as high as 21 µg/L in one sample (9). In addition, BPA is commonly found at ppb-level in effluents from paper or plastic production plants and domestic sewage treatment plants, because BPA is not completely removed during treatment (4, 16-18). Leachates from hazardous waste landfills usually contain high concentrations of BPA. For example, up to 17 200 µg/L of BPA was found in leachates from a landfill in Japan (3). Several studies reported that elimination rates of BPA in wastewater treatment plants ranged from 37 to 94% (4, 16, 19). The poor and inconsistent removal of phenolic compounds is generally unsatisfactory using current water and wastewater treatment techniques, because phenols are relatively hydrophilic and therefore less vulnerable to traditional physical/ chemical treatment processes (20). To this end, several advanced oxidation processes have been evaluated for the removal of aqueous BPA (21-23). However, efficient techniques for removing BPA from water are still urgently needed. Manganese oxide (δ-MnO2) has shown powerful oxidative capacity for phenolic and aniline compounds (24-32), because phenolic or aniline compounds are susceptible to oxidation by MnO2, forming an intermediate radical that may subsequently trigger a series of radical reactions (24-29). MnO2 was recently examined for removing steroid estrogens from water (33). However, the study did not consider oxidative products, and the reaction mechanism and estrogen transformation intermediates were not investigated. The objective of this study was to systematically explore oxidative removal of BPA with δ-MnO2 by investigating removal efficacy and influencing factors such as MnO2 loading, pH, and coexisting solutes (e.g., CaCl2, MgCl2, MnCl2, FeCl3, and humic acid). Furthermore, the principal reaction intermediates and products were identified, and a detailed reaction scheme was proposed.

Materials and Methods Chemicals. BPA standard (>99%) was purchased from SigmaAldrich (St. Louis, MO). Other chemicals and reagents used are given in the Supporting Information. Manganese oxide (δ-MnO2) was synthesized according to Murray’s method (34). Detailed information on the synthetic procedure and its characterization are also available in the Supporting Information. Reaction Setup. All glassware was soaked with 5 N HNO3, thoroughly rinsed with reagent-grade water, and baked at 400 °C for 4 h prior to use. All experiments were conducted 10.1021/es900235f CCC: $40.75

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in 100-mL amber borosilicate glass bottles under ambient O2 conditions in the absence of light and at room temperature (22 ( 1 °C). Reaction mixtures were constantly stirred with Teflon-coated magnetic stir bars at 480 rpm. Reaction solutions were maintained at constant pH with 10 mM buffers: acetic acid/sodium acetate for pH 4.5 and 5.5, 4-morpholinepropanesulfonic acid and its sodium salt for pH 6.5 and 7.5, and 2-(cyclohexylamino)ethanesulfonic acid and its sodium salt for pH 8.6 and 9.6. An appropriate amount of NaCl was also added to the solution to maintain the reaction solution at a constant ionic strength of 0.01 M. Prior to initiation of reactions, MnO2 suspensions in buffer solutions (40 mL) with or without cosolutes (e.g., CaCl2, MnCl2, MgCl2, FeCl3, humic acid) were mixed on a shaker at a low speed for 20 h. Reactions were initiated by adding 40 µL of 4.40 mM BPA solution into the pre-equilibrated and continuously stirred MnO2 suspensions. Aliquots of 1.0 mL of reaction mixture was periodically withdrawn, transferred to 2-mL HPLC vials containing 5 µL of L-ascorbic acid solution (50 mg/mL) and immediately vortexed for 10 s. The typical quench times for 800, 400, 200, and 100 µM MnO2 were about 3, 3, 2, and 2 s, respectively. Dissolution of the MnO2 released the adsorbed BPA and reaction products. All samples were analyzed within 24 h. Samples quenched by L-ascorbic were used to determine residue concentrations of BPA. The effect of initial MnO2 loading on BPA removal was studied by separately employing various MnO2 concentrations (0, 100, 200, 400, and 800 µM) under the fixed conditions of pH 5.5, 4.4 µM BPA, and 0.01 M ionic strength. pH effect was evaluated in solutions with various pH values (4.5, 5.5, 6.5, 8.6, and 9.6), under the fixed conditions of 800 µM MnO2, 4.4 µM BPA, and 0.01 M ionic strength. The cosolute effects of metal ions were investigated by separately fortifying two different concentrations (e.g., 0.01 and 0.05 M) of cosolutes including CaCl2, MgCl2, FeCl3, and MnCl2 under the fixed conditions of 800 µM MnO2, 4.4 µM BPA, and pH 5.5. The effect of humic acid was studied at pH 5.5 and 8.6 by treating the solution with humic acid at 0.1 or 10 mg/L under the fixed conditions of 800 µM MnO2, 4.4 µM BPA, and 0.01 M ionic strength. Chemical Analysis. Analysis of BPA disappearance in the reaction solutions and identification of reaction products are given in the Supporting Information. Molecular Modeling. Molecular modeling was performed using the Gaussian 03 program (Gaussian, Inc., Wallingford, CT). The molecular structure of a BPA free radical was first optimized by the density functional theory method at the level of B3LYP with basis sets of 6-31g(d) and was then used to calculate the charge and spin densities.

Results and Discussion Efficacy of BPA Removal by MnO2. In the absence of MnO2, BPA remained stable in water and no degradation was observed within 24 h under the employed experimental conditions. However, when the solutions contained a certain amount of MnO2, BPA rapidly disappeared from the solution. A typical time course of BPA removal from the aqueous phase by MnO2 is shown in Figure 1. For example, 90% of BPA was removed in 5 min and >99% was removed in 10 min in a pH 5.5 reaction system originally containing 800 µM MnO2 and 4.4 µM BPA. As will be described below, the removal efficiency was further improved by optimizing experimental conditions, such as the initial loading amount of MnO2 and pH of the reaction solution. These results clearly suggested that MnO2 is capable of efficiently removing BPA from the aqueous phase. Effect of Initial MnO2 Concentration. It is generally accepted that the oxidation of organic compounds by MnO2 is initially triggered by surface reactions on the oxide, where

FIGURE 1. Effect of MnO2 initial loading on bisphenol A (BPA) removal at 22 ( 1 °C in pH 5.5 solutions initially containing 4.40 µM BPA. Data for MnO2-free treatment are from single measurements, whereas the other data points are given as means ( standard deviations (n ) 3). the reductant complexes with the oxide and subsequently loses electrons to the oxide to form a phenoxy radical (24-29). The precursor complex formation and electron transfer are believed to be the rate-limiting processes (24-28), suggesting that the accessible sites on the oxide can greatly impact the reaction rate. The removal of BPA strongly depended on the loading of MnO2 in the reaction system (Figure 1). High concentrations of MnO2 facilitated the removal of BPA. For example, the removal of BPA in 60 min was increased from 49% for a system amended with 100 µM MnO2 to 73% with 200 µM MnO2 (Figure 1). The removal efficiency was further enhanced with higher concentrations of MnO2, and near complete removal was attained in 10 and 20 min for systems amended with 800 and 400 µM MnO2, respectively. On the basis of the empirical formula of synthetic MnO2 and its measured specific area, the initial ratios of surface area to moles of BPA were estimated to be 5.5 × 105, 1.1 × 106, 2.2 × 106, and 4.4 × 106 m2/mol for 100, 200, 400, and 800 µM MnO2, respectively. The steady increasing of BPA removal efficiency with the increasing ratio of surface area to BPA moles suggested that at the low MnO2 levels, a saturation with BPA was approached, and a further increase of MnO2 concentration led to accelerated BPA removal. Effect of pH. Bisphenol A removal rates were apparently sensitive to the change of solution pH (Figure 2). Overall, BPA removal rates in different pH solutions followed the order of pH 4.5 > pH 5.5 > pH 8.6 > pH 9.6 > pH 7.5 > pH 6.5, suggesting that acidic or basic conditions facilitated BPA removal, whereas neutral conditions disfavored the reaction. Similar to our finding, the oxidation power of MnO2 toward other organic compounds has been demonstrated to be pHdependent in previous studies (24-32). The variation of BPA removal by MnO2 with pH may be attributed to two major factors. First, the oxidation of BPA by MnO2 strongly depended on its oxidation potential and thereby solution pH, because reduction of MnO2 to Mn(II) requires the participation of protons (e.g., 1/2MnO2(s) + 2H+ + e- f 1/2Mn2+(aq) + H2O). For example, when the pH was decreased from 8.0 to 4.0, MnO2 reduction potential was increased from 0.76 to 0.99 V (35). Likely for this reason, BPA removal increased as the pH was increased from 4.5 to 6.5. On the other hand, pH governs the speciation of BPA, which affects the electron transfer process between BPA and MnO2 during surface reactions. Bisphenol A has the first and second acid dissociations of pKa1 ) 9.6 and pKa2 ) 10.2 (36), indicating that the distribution of BPA species, i.e., undissociated BPA, monobasic BPA, and dibasic BPA, varies with the solution pH (Figure S1, Supporting Information). For example, almost VOL. 43, NO. 10, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 2. Effect of pH on bisphenol A (BPA) removal at 22 ( 1 °C in solutions initially containing 4.40 µM BPA and 800 µM MnO2. Data points are given as means ( standard deviations (n ) 3).

FIGURE 3. Effect of metal ions on bisphenol A (BPA) removal at 22 ( 1 °C in pH 5.5 solutions initially containing 4.40 µM BPA and 800 µM MnO2. Data points are given as means ( standard deviations (n ) 3). all BPA (>99.9%) remains undissociated in solutions at pH e 6.5. The percentages of monobasic BPA are 0.8%, 9.07%, and 44.4% at pH 7.5, 8.6, and 9.6, respectively, while the corresponding percentages of undissociated BPA were 99.2%, 99.7%, and 44.4%. In this study, BPA removal did not decrease when pH was increased from 6.5 to 8.6, suggesting that the deprotonated BPA was more susceptible to losing electrons to the oxide during surface reactions and, therefore, more capable of triggering the oxidation process. Similar to our finding, Stone (28) found that the protonated form of phenol (ArOH) was considerably less reactive than the phenolate anion (ArO-). It is likely that the combined effects decreased BPA removal in the pH 9.6 solution. Effect of Cosolutes. Bisphenol A removal was generally inhibited by coexisting metal ions, i.e., Ca2+, Mg2+, Fe3+, and Mn2+ (Figure 3). For example, after 10 min of reaction, BPA removal rates were 99.5%, 88.1%, 86.2%, 61.9%, and 24.7% for the control and solutions fortified with 0.01 M MgCl2, 0.01 M FeCl3, 0.01 M CaCl2, and 100 µM MnCl2, respectively. 3862

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FIGURE 4. Effect of humic acid (HA) on bisphenol A (BPA) removal at 22 ( 1 °C in solutions initially containing 4.40 µM BPA and 800 µM MnO2. Data points are given as means ( standard errors (n ) 3). It was reported that metal ions suppressed the oxidative power of MnO2 to organic compounds by complexing with both dissolved and sorbed reactants or by occupying reactive surface sites on the oxide (37). In this study, the cosolute effect was studied at pH 5.5, where more than 99.99% of BPA remained undissociated and thereby BPA removal resulting from the complexation of BPA with metal ions should be negligible. Under the used experimental conditions, the metal ions were expected to bind to the reactive sites via electrostatic interactions between metal ions and the negatively charged MnO2 surface and, therefore, decrease the accessible sites for BPA. The competitive binding effect was also considered to be responsible for the inhibition of coexisting metal ions (Cu2+, Zn2+, and Fe3+) to 17β-estadiol removal by MnO2 (33). The inhibitive effect of metal ions varied with species and concentrations (Figure 3). Higher concentrations of metal ions resulted in a higher inhibitory effect. For example, after 20 min of reaction, removal of BPA by treatments with 1.0 × 104 µM MgCl2 and CaCl2 were 98.7% and 85.1%, respectively, and the corresponding values decreased to 96.8% and 72.5% when the metal ion concentrations were increased to 5.0 × 104 µM. The concentration dependence may be attributed to the fact that higher concentrations of metal ions led to increased occupation of reactive sites on MnO2. The inhibitive capacity of the metal ions generally followed the order Mn2+ > Ca2+ > Mg2+ ≈ Fe3+. Therefore, Mn2+ exhibited the strongest suppressive effect compared to the other metal ions studied. For example, no significant disappearance of BPA was observed in 30 min for solutions pre-equilibrated with 1.0 × 103 µM MnCl2. Even for solutions fortified with MnCl2 at 100 µM, BPA removal was reduced by 75% after 10 min of reaction compared to the Mn2+-free treatment (Figure 3). The much lower affinity of MnO2 for Ca2+ than for Mn2+ (38) may partially explain the strong inhibition of Mn2+. Besides, Mn2+ is the final reduced product of MnO2, and therefore, an increasing concentration of Mn2+ may reduce the redox potential of MnO2/Mn(II). Due to the same reason, the autoinhibitive effect of Mn2+ inevitably accompanying the oxidation of organic compounds by MnO2 would result in a progressively slower reaction rate (27, 30). Humic acid (HA), a natural organic matter mostly found in water, is another factor that may affect the oxidative efficiency of MnO2 (29, 33). In this study, only the treatment of 10 mg/L HA in the pH 8.6 system showed weak inhibition on the BPA removal, and the other treatments were generally not affected by HA (Figure 4). However, HA was found to facilitate the removal of 17β-estadiol by MnO2, due possibly to the complexation of HA with Mn2+ as the oxidative reaction

progressed (33). On the other hand, HA suppressed the oxidation of substituted anilines by MnO2, likely due to the dissolution of MnO2 by HA during the pre-equilibrating period (29). Our results suggested that these interactions did not occur in the reaction of BPA with MnO2 under the conditions tested, or the interactions may have counteracted with each other, leading to the generally negligible effect. Reaction Intermediates and Scheme. To identify reaction products, the reaction solution was extracted and possible polar products were silylated with N,O-bis(trimethylsilyl)trifluoroacetamide, followed by GC-MS analysis. Overall, 10 products were identified in their silylated derivatives and one in its original form. The mass spectra and the possible fragments during GC-MS analysis for these products are given in Figure S2 (Supporting Information). The indentified compounds correspond to the formation products 1-11, as identified in Figure 5. Products 6 and 7 were further confirmed using authentic standards of 4,4′-biphenol and hydroquinone, and their mass spectra of the silylated standards are given in the Supporting Information (Figure S2M,L, respectively). However, authentic standards of the other products were not available. To assist the elucidation of mass spectra of other products, a mass spectrum of silylated BPA is also given in the Supporting Information (Figure S2N). It should be noted that some reaction products may not have been efficiently extracted into methylene chloride and may not be detected in this study. However, on the basis of the detected products, several reaction schemes are possible (Figure 5). In the suggested scheme, oxidation of BPA is initiated by transferring an electron to the oxide, forming a BPA radical R1-0 (pathway I). The BPA radical is susceptible to interchange to other transition forms via resonance. The calculation of charge and spin densities via molecular modeling provides support for the resonance of the radical. It is known that the charge density denotes net charges associated with each atom, and the spin density indicates the possibility of the presence of a spin-unpaired electron. The calculation results (Figure S3, Supporting Information) suggest that the unpaired electron in the BPA radical R1-0 is likely to delocalize to positions 8, 10, and 12 of the BPA radical, forming transition forms R1-1, R1-2, and R1-3 (pathway II). The radicals thereafter trigger a suite of reactions. For example, R1-1 could undergo β scission by releasing an olefin 1 (Figure S2A, Supporting Information) and a new radical R2 which is involved in the other reactions to be discussed later (pathway III). Radical couplings are common reactions for the oxidation of phenolic compounds by MnO2 (25). Generally, the coupling processes involve carbon-carbon (C-C) and carbon-oxygen (C-O) coupling of the mesomeric radicals followed by enolization to yield the resultant dimers as shown in pathways i-iv. Taking both charge and spin densities into account, coupling of R1-0 with R1-1 is the one most likely to happen. Such coupling forms an unstable intermediate which, to decrease the steric instability, subsequently eliminates cationic isopropylphenol R3 and releases 2 (Figure S2H, Supporting Information; pathway i). Coupling reactions between other radicals are also possible. For example, coupling of R1-0 with R1-2 generates 3 (Figure S2K, Supporting Information; pathway ii) and coupling R1-2 with R1-3 gives 4 (Figure S2J, Supporting Information; pathway iii). The radical R2 formed in pathway III maybe subject to a coupling reaction releasing 6 (Figure S2E, Supporting Information; pathway v) or further oxidation by MnO2 giving 7 (Figure S2B, Supporting Information; pathway V). The cationic carbon R3 formed in pathway i could also trigger a suite of substitution or elimination reactions. For example, R3 may be subject to deprotonation, yielding 8 (Figure S2C, Supporting Information; pathway vi), or substitution of a proton of water, forming 9 (Figure S2, Supporting Information; pathway vii). Such

FIGURE 5. Proposed reaction scheme for removal of bisphenol A by MnO2 in pH 5.5 solutions. R1-0, R1-1, R1-2, R1-3 are four transition forms of BPA radical. R2 and R3 are new radical and cationic carbon formed during further reactions, respectively. Each number 1-11 represents an intermediate or product. substitution reactions may also involve the primary products such as 7 and 8 correspondingly forming 10 and 11 (Figure S2F,G, Supporting Information; pathways viii and ix, respectively). Since products or intermediates of 2-11 all are phenolic compounds, they may be further oxidized in similar pathways as mentioned above. For example, the phenoxy radical of 7 couples with R1-3, forming 5 (Figure S2I, Supporting Information; pathway iv). As the reactions proceed, it will become increasingly complicated because of the further oxidation of primary products or intermediates. VOL. 43, NO. 10, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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Environmental Implications. This study demonstrated that MnO2 efficiently removes BPA from water, presenting a promising use of MnO2 to remove phenolic compounds in a range of applications. The rapid disappearance of BPA likely resulted from a series of reactions, such as substitution, elimination, and radical fragmentation and coupling, which was triggered after the formation of BPA radical by transferring electrons to MnO2. Most of the formation intermediates were identified as phenolic compounds, and thus may still have estrogenic potencies. Because of the lack of standard references for these intermediates, the possibility of their further oxidation by MnO2 and their toxicity were not evaluated in this study and therefore merit future research.

Acknowledgments This work was supported by California State Water Resources Control Board. The authors thank Dr. Haiming Zhang for his assistance in molecular modeling.

Supporting Information Available Procedures, analysis, and Figures S1-S3 described in this paper. This material is available free of charge via the Internet at http://pubs.acs.org.

Literature Cited (1) Bisphenol A information & resources. Accessible at http:// www.bisphenol-a.org/?gclid)CJuhpJ6Y5ZcCFQ8Qagodlw8UEA. (2) Tsai, W. T. Human health risk on environmental exposure to bisphenol-A: A review. J. Environ. Sci. Health C 2006, 24, 225– 255. (3) Yamamoto, T.; Yasuhara, A.; Shiraishi, H.; Nakasugi, O. Bisphenol A in hazardous waste landfill leachates. Chemosphere 2001, 42, 415–418. (4) Fu ¨ rhacker, M.; Scharf, S.; Weber, H. Bisphenol A: Emissions from point sources. Chemosphere 2000, 41, 751–756. (5) Staples, C. A.; Dorn, P. B.; Klecka, G. M.; O’Block, S. T.; Harris, L. R. A review of the environmental fate, effects, and exposures of bisphenol A. Chemosphere 1998, 36, 2149–2173. (6) U.S. Environmental Protection Agency (EPA) Toxics Release Inventory (TRI) Explore Report. Accessible at http://www. epa.gov/tri/. (7) Ben-Jonathan, N.; Steinmetz, R. Xenoestrogens: The emerging story of bisphenol A. Trends Endocrinol. Metab. 1998, 9, 124– 128. (8) Levy, G.; Lutz, I.; Kru ¨ ger, A.; Kloas, W. Bisphenol A induces feminization in Xenopus laevis tadpoles. Environ. Res. 2004, 94, 102–111. (9) Kang, J. H.; Aasi, D.; Katayama, Y. Bisphenol A in the aquatic environment and its endocrine-disruptive effects on aquatic organisms. Crit. Rev.Toxicol. 2007, 37, 607–625. (10) Robaire, B. Is it my grandparents’ fault? Nat. Med. 2008, 14, 1186–1187. (11) Jobling, S.; Nolan, M.; Tyler, C. R.; Brighty, G.; Sumpter, J. P. Widespread sexual disruption in wild fish. Environ. Sci. Technol. 1998, 32, 2498–2506. (12) Welshons, W. V.; Thayer, K. A.; Judy, B. M.; Taylor, J. A.; Curran, E. M.; vom Saal, F. S. Large effects from small exposures. I. Mechanisms for endocrine-disrupting chemicals with estrogenic activity. Environ. Health Perspect. 2003, 111, 994–1006. (13) Lee, B.; Ha, E.; Park, H.; Kim, B.; Seo, J.; Chang, M.; Ha, M.; Kim, Y.; Roh, Y.; Hong, Y. Exposure to bisphenol A in pregnant women and early fetal growth. Epidemiology 2008, 19, S365-S365. (14) Susiarjo, M.; Hassold, T. J.; Freeman, E.; Hunt, P. A. Bisphenol A exposure in utero disrupts early oogenesis in the mouse. Plos. Genet. 2007, 3, 0063-0070. (15) Hunt, P. A.; Koehler, K. E.; Susiarjo, M.; Hodges, C. A.; Ilagan, A.; Voigt, R. C.; Thomas, S.; Thomas, B. F.; Hassold, T. J. Bisphenol A exposure causes meiotic aneuploidy in the female mouse. Curr. Biol. 2003, 13, 546–553.

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(16) Lee, H. B.; Peart, T. E. Determination of bisphenol A in sewage effluent and sludge by solid-phase and supercritical fluid extraction and gas chromatography/mass spectrometry. J. AOAC Int. 2000, 83, 290–297. (17) Rigol, A.; Latorre, A.; Lacorte, S.; Barcelo, D. Determination of toxic compounds in paper-recycling process waters by gas chromatography-mass spectrometry and liquid chromatographymass spectrometry. J. Chromatgr. A 2002, 963, 265–275. (18) Quinn, B. P.; Booth, M. M.; Delfino, J. J.; Holm, S. E.; Gross, T. S. Selected resin acid in effluent and receiving waters derived from a bleached and unbleached kraft pulp and paper mill. Environ. Toxicol. Chem. 2003, 22, 214–218. (19) Fuerhacker, M. Bisphenol A emission factors from industrial sources and elimination rates in a sewage treatment plant. Water Sci. Technol. 2003, 47, 117–122. (20) Snyder, S. A.; Westerhoff, P.; Yoon, Y.; Sedlak, D. L. Pharmaceuticals, personal care products, and endocrine disruptors in water: Implications for the water industry. Environ. Eng. Sci. 2003, 20, 449–469. (21) Huang, Q. G.; Weber, W. J. Transformation and removal of bisphenol A from aqueous phase via peroxidase-mediated oxidative coupling reactions: Efficacy, products and pathways. Environ. Sci. Technol. 2005, 39, 6029–6036. (22) Deborde, M.; Rabouan, S.; Mazellier, P.; Duguet, J. P.; Legube, B. Oxidation of bisphenol A by ozone in aqueous solution. Water Res. 2008, 42, 4299–4308. (23) Rosenfeldt, E. J.; Linden, K. G. Degradation of endocrine disrupting chemicals bisphenol A, ethinyl estradiol, and estradiol during UV photolysis and advanced oxidation processes. Environ. Sci. Technol. 2004, 38, 5476–5483. (24) Zhang, H. C.; Chen, W. R.; Huang, C. H. Kinetic modeling of oxidation of antibacterial agents by manganese oxide. Environ. Sci. Technol. 2008, 42, 5548–5554. (25) Zhang, H. C.; Huang, C. H. Reactivity and transformation of antibacterial N-oxides in the presence of manganese oxide. Environ. Sci. Technol. 2005, 39, 593–601. (26) Zhang, H. C.; Huang, C. H. Oxidative transformation of fluroquinolone antibacterial agents and structurally related amines by manganese oxide. Environ. Sci. Technol. 2005, 39, 4474–4483. (27) Zhang, H. C.; Huang, C. H. Oxidative transformation of triclosan and chlorophene by manganese oxides. Environ. Sci. Technol. 2003, 37, 2421–2430. (28) Stone, A. T. Reductive dissolution of manganese(III/IV) oxides by substituted phenols. Environ. Sci. Technol. 1987, 21, 979– 988. (29) Klausen, J.; Haderlern, S. B.; Schwarzenbach, R. P. Oxidation of substituted anilines by aqueous MnO2: Effect of co-solutes on initial and quasi-steady-state kinetics. Environ. Sci. Technol. 1997, 31, 2642–2649. (30) Rubert, K. F.; Pedersen, J. A. Kinetics of oxyteracycline reaction with a hydrous manganese oxide. Environ. Sci. Technol. 2006, 40, 7216–7221. (31) Laha, S.; Luthy, R. G. Oxidation of aniline and other primary aromatic amines by manganese dioxide. Environ. Sci. Technol. 1990, 24, 363–373. (32) Li, H.; Lee, L. S.; Schulze, D. G.; Guest, C. A. Role of soil manganese in the oxidation of aromatic amines. Environ. Sci. Technol. 2003, 37, 2686–2693. (33) Xu, L.; Xu, C.; Zhao, M. R.; Qiu, Y. P.; Sheng, G. D. Oxidative removal of aqueous steroid estrogens by manganese oxides. Water Res. 2008, 42, 5038–5044. (34) Murray, J. W. Surface chemistry of hydrous manganese dioxide. J. Colloid Interface Sci. 1974, 46, 357–371. (35) Stumm, W.; Morgan, J. J. Aquatic Chemistry, 3rd ed.; Wiley: New York, 1996. (36) Kosky, P. G.; Silva, J. M.; Guggenheim, E. A. The aqueous phase in the interfacial synthesis of polycarbonates. 1. Ionic equilibria and experimental solubilities in the BPA-NaOH-H2O system. Ind. Eng. Chem. Res. 1991, 30, 462–467. (37) Barrett, K. A.; McBride, M. B. Oxidative degradation of glyphosate and aminomethylphosphonate by manganese oxide. Environ. Sci. Technol. 2005, 39, 9223–9228. (38) Morgan, J. J.; Stumm, W. J. Colloid-chemical properties of manganese dioxide. J. Colloid Interface Sci. 1964, 19, 347–359.

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