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Oxidation of Phenolic Endocrine Disrupting Chemicals by Potassium Permanganate in Synthetic and Real Waters Jin Jiang,† Su-Yan Pang,†,‡ Jun Ma,*,† and Huiling Liu† †

State Key Laboratory of Urban Water Resource and Environment, School of Municipal and Environmental Engineering, Harbin Institute of Technology, Harbin 150090, China ‡ College of Chemical and Environmental Engineering, Harbin University of Science and Technology, Harbin 150080, China S Supporting Information *

ABSTRACT: In this study, five selected environmentally relevant phenolic endocrine disrupting chemicals (EDCs), estrone, 17β-estradiol, estriol, 17αethinylestradiol, and 4-n-nonylphenol, were shown to exhibit similarly appreciable reactivity toward potassium permanganate [Mn(VII)] with a second-order rate constant at near neutral pH comparable to those of ferrate(VI) and chlorine but much lower than that of ozone. In comparison with these oxidants, however, Mn(VII) was much more effective for the oxidative removal of these EDCs in real waters, mainly due to the relatively high stability of Mn(VII) therein. Mn(VII) concentrations at low micromolar range were determined by an ABTS [2,2-azino-bis(3-ethylbenzothiazoline)-6sulfonic acid diammonium] spectrophotometric method based on the stoichiometric reaction of Mn(VII) with ABTS [MnVII + 5ABTS → MnII + 5ABTS•+] forming a stable green radical cation (ABTS•+). Identification of oxidation products suggested the initial attack of Mn(VII) at the hydroxyl group in the aromatic ring of EDCs, leading to a series of quinone-like and ring-opening products. The background matrices of real waters as well as selected model ligands including phosphate, pyrophosphate, NTA, and humic acid were found to accelerate the oxidation dynamics of these EDCs by Mn(VII). This was explained by the effect of in situ formed dissolved Mn(III), which could readily oxidize these EDCs but would disproportionate spontaneously without stabilizing agents.

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INTRODUCTION The widespread occurrence of endocrine disrupting chemicals (EDCs) in the aquatic environments has raised serious concerns about their adverse effect on aquatic ecology and public health.1 Recent studies suggest that trace levels of EDCs can change the endocrine functions of wildlife.2,3 The effluents of municipal wastewater treatment plants (WWTPs) have been identified as one major source of EDCs released into the aquatic environments.4 Although significant amounts of EDCs can be removed from wastewater via biological adsorption/ oxidation processes in WWTPs, considerable quantities still remain in the effluents.5 As a result, many EDCs have been frequently detected in effluent-affected surface waters.6−8 Chemical oxidation processes can be applied to control many EDCs during water and wastewater treatment. Many studies have demonstrated that selective oxidants such as ozone, chlorine dioxide, chlorine, and ferrate [Fe(VI)] are effective in treating emerging micropollutants such as steroid estrogens, pharmaceuticals, and personal care products containing electron-rich moieties, including phenol, olefin, aniline, and thiol groups.9−15 In particular, Lee et al.16 reported that these oxidants could readily oxidize the phenolic moiety of 17αethinylestradiol (EE2), a synthetic steroid estrogen used in the oral contraceptive pill, resulting in a consequent rapid decrease of its estrogenic activity. The pH-dependent second-order rate © 2011 American Chemical Society

constants for the reactions of selective oxidants with several environmentally relevant phenolic EDCs [i.e., four natural and synthetic estrogens including estrone (E1), 17β-estradiol (E2), estriol (E3), and EE2, as well as the two industrial chemical products 4-n-nonylphenol (4-n-NP) and bisphenol A (BPA)] available in literature are summarized in the Supporting Information (SI) Figure S1. However, to date little kinetic information is known on the common oxidant permanganate [Mn(VII); KMnO4], which can address some of the concerns of currently used oxidants;17,18 for instance, the formation of chlorinated and brominated byproduct during chlorination and ozonation. Mn(VII) has been already widely used in water utilities over the past decades for control of dissolved Mn(II), taste and odor compounds, and cyanotoxins, due to its comparative stability, ease of handling, and relatively low cost.19−21 In addition, Mn(VII) is attractive because its reduction product MnIVO2 can serve as a coagulant aid to enhance the coagulation performance or as an adsorbent to remove heavy metals simultaneously.22,23 Mn(VII) can selectively react with electron-rich Received: Revised: Accepted: Published: 1774

October 7, 2011 December 21, 2011 December 27, 2011 December 27, 2011 dx.doi.org/10.1021/es2035587 | Environ. Sci. Technol. 2012, 46, 1774−1781

Environmental Science & Technology

Article

Mn(VII) concentrations at low levels. Finally, the influence of selected model ligands on Mn(VII) oxidation of phenolic EDCs and the reactivity of identified dissolved Mn(III) species with them were examined to further confirm and better understand the ligand-enhanced Mn(VII) reactions reported in our previous studies when dealing with TCS and BPA,27,28 which are structurally similar to these EDCs investigated in this study.

moieties via several reaction pathways such as hydrogen abstraction, electron exchange, and direct oxygen transfer.24 In recent years, the potential application of Mn(VII) for the oxidative removal of emerging micropollutants containing electron-rich moieties during water and wastewater treatment has received great attention. The very recent work conducted by Hu et al.25 demonstrated that Mn(VII) was fairly effective in treating three antibiotics with apparent second-order rate constants kMn(VII) at pH 8 and 25 °C of 0.67 M−1 s−1 (ciprofloxacin), 0.92 M−1 s−1 (trimethoprim), and 26.6 M−1 s−1 (lincomycin). Identification of reaction intermediates and products suggested that Mn(VII) initially attacked the tertiary aromatic and secondary aliphatic amine, thioether, and olefin groups of these micropollutants.26 Our recent studies showed that Mn(VII) could readily oxidize triclosan (TCS; a widely used phenolic biocide) and BPA with kMn(VII) at pH 8 and 25 °C of 350 and 240 M−1 s−1, respectively.27,28 Interestingly, we found that phosphate buffer as well as pyrophosphate and humic acid significantly enhanced the oxidation of TCS and BPA by Mn(VII). In contrast, these ligands had negligible influence on the oxidation of carbamazepine (CBZ; an anticonvulsant drug containing an olefin group).28,29 The discrepancy regarding the role of ligands in Mn(VII) oxidation of TCS (or BPA) and CBZ was mainly attributed to the effect of identified dissolved Mn(III), which would otherwise disproportionate spontaneously in the absence of stabilizing ligands to yield Mn(II) and solid MnO2 (reaction 1).30

2Mn III → Mn IV O2 + Mn II

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EXPERIMENTAL SECTION Materials. Selected EDCs and model compounds (see SI Figure S2 for their chemical structures) were purchased from Sigma-Aldrich with a purity >97%. A soil-humic acid which has been purified and well characterized in our previous study32 was used. The preparation of stock solutions of EDCs and oxidants is described in the SI (Text S1). Oxidation in Synthetic Waters. Kinetic reactions were initiated by addition of Mn(VII) in excess into solutions containing one target substrate and/or ligand of interest. The pH of reaction solutions was controlled using acetate buffer (10 mM) for pH 5 and borate buffer (10 mM) for pH 8−10. However, no buffer was used for pH 6−7, which was maintained almost constant (±0.1) by addition of HClO4 or NaOH when necessary, because phosphate buffer could significantly affect the reaction kinetics while organic alternatives (e.g., MOPS and HEPES) would appreciably react with Mn(VII).27−29 No buffer was used for solution pH 11 either; however, it changed negligibly during the kinetic runs. For kinetic experiments involving Mn(III), reactions were started by addition of individual EDCs into Mn(III) solutions, which were freshly prepared by the stoichiometric reaction of Mn(VII) with Mn(II) in excess ligands (reaction 2) and adjusted to the desired pH (in most cases excess ligands served as pH buffers).28,30

(1)

One-electron oxidant Mn(III) species were shown to be able to rapidly oxidize TCS and BPA but exhibited negligible reactivity toward the olefinic moiety for which the two-electron oxygen donation may be the primary oxidation mechanism.28 In these regards, it is not difficult to understand that the oxidation kinetics of olefin-containing organics such as chlorinated ethylenes and microcystins by Mn(VII) were widely examined in previous studies, but no impact of phosphate buffer was reported.21,31 Also, the different reactivity of Mn(III) species with phenolics versus olefins could reasonably explain the contrasting finding that the background matrice of river water significantly accelerated the oxidation of TCS and BPA by Mn(VII),27,28 while the rate constants obtained from pure water precisely predicted the oxidation kinetics of micropollutants containing the olefinic moiety by Mn(VII) in real waters (e.g., Zurich lake water).21 This study was conducted to assess the ability of Mn(VII) to oxidize phenolic EDCs (i.e., E1, E2, E3, EE2, and 4-n-NP) during water and wastewater treatment. First, the reaction kinetics of Mn(VII) with these EDCs was determined in synthetic buffer solutions over a wide pH range from 5 to 11. Then, reaction products of E2 and EE2 with Mn(VII) were identified with liquid chromatography tandem mass spectrometry (LC−MS/MS) and the help of the model compound 5,6,7,8-tetrahydro-2-naphthol (THN) that represented the reactive phenolic moiety. Further, the effectiveness and kinetics of the oxidation of phenolic EDCs by Mn(VII) in real waters under the conditions relevant to water utility practice were examined and compared with other selective oxidants including ozone, Fe(VI), and chlorine. For this purpose, an ABTS [2,2azino-bis(3-ethylbenzothiazoline)-6-sulfonic acid diammonium] spectrophotometric method was developed to measure

ligand

Mn VII + 4Mn II ⎯⎯⎯⎯⎯⎯→ 5Mn III

(2)

Aliquots were periodically collected and quenched with ascorbic acid and analyzed for residual EDCs with highperformance liquid chromatography with fluorescence detection (HPLC-FD). Identification of Oxidation Products. To identify reaction products, series of E2, EE2, and the model compound THN solutions (80 μM; pH 7) containing 10% acetonitrile (v/ v) to aid dissolution were treated with Mn(VII) of 40−160 μM. Sufficient time (15−20 min) was allowed for the reactions to reach completion [i.e., Mn(VII) was consumed completely]. Test demonstrated that Mn(VII) showed negligible reactivity with acetonitrile within this time. The resulting solutions were then filtered through 0.45 μm glass fiber membranes before analyzing by LC−MS/MS with negative mode electrospray ionization (ESI). Details of analytical procedures are provided in the SI (Text S2). Oxidation in Real Waters. Surface water (DOC = 5.0 mg C/L, alkalinity = 150 mg/L as CaCO3, and pH = 7.7) was taken from Songhua river, Harbin, China. The secondary wastewater effluent (DOC = 7.1 mg C/L, alkalinity = 230 mg/L as CaCO3, [NH4+/NH3] = 1.4 μM, [NO2−] = 5.4 μM, and pH = 7.8) was obtained from Wencang WWTP in Harbin. After vacummfiltering through 0.45 μM glass fiber membranes, these waters were buffered to pH 8 with borate buffer (10 mM) and stored at 4 °C. 1775

dx.doi.org/10.1021/es2035587 | Environ. Sci. Technol. 2012, 46, 1774−1781


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The transformation efficiency of EDCs in real waters treated by Mn(VII) at various doses was tested and compared with ozone, Fe(VI), and chlorine. Each phenolic EDC was spiked at the concentration of 0.15 μM into a series of vials containing 20 mL of real waters. The concentrations of phenolic EDCs originally present in real waters were negligible relative to that artificially spiked. Different doses of one oxidant were then injected into vials under vigorous stirring. After 30 min, the reaction solutions were quenched with ascorbic acid and analyzed for residual EDC concentrations with HPLC-FD. Further, the oxidation kinetics of phenolic EDCs by Mn(VII) in real waters was examined following the same procedure as that in synthetic buffer solutions except that the oxidant concentration during kinetic runs was measured simultaneously in a second sample. Oxidant concentrations were determined by several colorimetric methods: chlorine by the ABTS method developed by Pinkernell et al.33 with trace level of iodide as catalyst; Fe(VI) and Mn(VII) by the ABTS method following the suggestion of Lee et al.,34 except that in the case of Mn(VII), samples were filtered prior to analysis to eliminate the interference from MnO2 precipitates; and ozone by the indigo method.35 Detailed procedure for the ABTS method is provided in the SI (Text S3). All experiments were conducted at 25 °C in duplicates or triplicates and the averaged data were presented. The standard deviations were always 0.99), but their slopes varied significantly, indicating the different stoichiometries of the reactions of these oxidants with ABTS. The stoichiometric factors (i.e., Δ[ABTS•+]/Δ[oxidant]) can be obtained by dividing the slopes in Figure 2 by the molar absorption coefficient of ABTS•+.34 The determined stoichiometric factors were about 1, 2, and 5 for Fe(VI), chlorine, and Mn(VII), respectively; the former two values were in agreement with the literature.33,34 The stoichiometric factor for Mn(VII) was determined at different pH conditions and also found to be 5 in the pH range of 2−6. A stoichiometric factor of 5 for Mn(VII) suggests that 1 mol of Mn(VII) produces 5 mol of ABTS•+ after its reaction with excess ABTS (reaction 3).

Mn VII + 5ABTS → Mn II + 5ABTS•+ 1777

(3)



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Figure 3. The relative residual concentration of E2 (0.15 μM) during oxidative treatment as a function of oxidant dose in surface water (a) and wastewater effluent (b) at pH 8. Consumption kinetics of selective oxidants (∼12 μM) in (c) surface water and (d) wastewater effluent.

Fe(VI) are slowly consumed along the reaction time. The fast consumption of chlorine is expected to be caused by ammonia, nitrite, and primary and secondary amine moieties in real waters, while in the latter phase the contribution of phenolic and tertiary amine moieties becomes important.13,41 The smooth decrease of Mn(VII) and Fe(VI) is consistent with their relatively low and narrowly distributed reactivity to various electron-rich moieties.24,41,42 No experimental data for ozone are presented in Figure 3c,d because of its complete consumption within 10 s of the first data point time. The extremely low stability of ozone in real waters is consistent with its high reactivity to electron-rich moieties such as phenols, anilines, olefins, and amines.43 Oxidation Dynamics of Phenolic EDCs in Real Waters by Mn(VII). Kinetic experiments were conducted to examine whether the rate constants determined in synthetic buffer solutions could predict the oxidation kinetics of phenolic EDCs by Mn(VII) in real waters. The following equation can be used to predict the extent of oxidation under the condition of [EDCs] ≪ [Mn(VII)]:

This can be rationalized by considering that manganese intermediates generated via sequential one-electron reduction of Mn(VII) rapidly oxidize additional ABTS to produce ABTS•+. However, the stoichiometric factor of 1 for Fe(VI) indicates that Fe(V) in situ formed by one-electron reduction of Fe(VI) is not involved in the production of ABTS•+. The discrepancy can be explained by the following tentative mechanisms: (i) Fe(V) is much more powerful than manganese intermediates, so that it attacks nonselectively at various active sites of the ABTS molecule that do not produce ABTS•+,34 and/or (ii) Fe(V) is much less stable than manganese intermediates, so that it decomposes rapidly rather than be involved in the oxidation of additional ABTS.27,28 Oxidant Consumption Kinetics and Transformation Efficiency of Phenolic EDCs in Real Waters by Mn(VII): Comparison with Ozone, Chlorine, and Fe(VI). A comparison was made for the transformation efficiency of phenolic EDCs in real waters by selective oxidants including Mn(VII), Fe(VI), chlorine, and ozone. Figure 3a,b shows the relative residual concentration of E2 (as a representative of selected phenolic EDCs) during treatment of surface water and wastewater effluent at pH 8 as a function of oxidant doses. As can be seen, when the oxidant doses were relatively low, E2 decreased slightly, especially in the cases of chlorination and ozonation. This could be ascribed to the strong competition between E2 and various electron-rich moieties in real waters.41 A further increase of oxidant doses led to the depletion of electron-rich moieties, and thus, the residual concentration of E2 started to decrease significantly. Generally, Mn(VII) was more efficient than chlorine, Fe(VI), and ozone in terms of the oxidant doses required to achieve more than 90% transformation in both waters. This result, inconsistent with the trend of the relative reactivity of these selective oxidants with E2 as shown in Figure 1b, was mainly attributed to their different stability in real waters. Figure 3c,d shows the consumption kinetics of selective oxidants in surface water and wastewater effluent for a dose of ∼12 μM. As can be seen, chlorine shows a fast consumption followed by a slow decrease over 30 min, while Mn(VII) and

t ⎧ ⎫ C = exp⎨ − kMn(VII) [Mn(VII)]dt ⎬ ⎩ ⎭ C0 0

∫

(4)

where ∫ 0t [Mn(VII)]dt is the Mn(VII) exposure [i.e., Mn(VII) concentration integrated over time].11 As shown in Figure 4a, the predictions significantly underestimated the level of E2 removal, although the background matrices of real waters slightly consumed Mn(VII) of 2−3 μM within the reaction time of 4 min (i.e., slightly decreased the Mn(VII) exposure) (Figure 3c,d). Comparable predictions were also observed in the cases of BPA and phenol, where Mn(VII) was relatively greatly consumed along the much longer reaction time of 15 and 60 min, respectively (SI Figure S4). Further, we found that when real waters were pretreated by ozonation, leading to negligible Mn(VII) consumption (and also negligible change of DOC), the oxidation kinetics of E2 were slowed down (but still slightly faster than those in synthetic water). It should be mentioned 1778



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Effect of Selected Model Ligands on Mn(VII) Oxidation of Phenolic EDCs. To better understand the proposed mechanism involving ligand-stabilized manganese intermediates, the influence of selected model ligands including phosphate, pyrophosphate, EDTA, NTA, and humic acid on the oxidation of phenolic EDCs by Mn(VII) was examined. These ligands except for EDTA exerted significant oxidation enhancement at acidic pH 5 and 6, while a negligible effect was observed at pH 7−9 (Figure 4b and SI Figure S7). As expected, the addition of strongly reduced species further accelerated the oxidation kinetics at acidic pH 5 and 6 (see SI Figure S8 for example). By combining the online UV−vis scanning, the capillary electrophoresis technique, and an ozonation method as done previously,28 the occurrence of Mn(III) species in the reaction solutions involving pyrophosphate, EDTA, and NTA was confirmed (see SI Figures S9 and S10, for example). However, in the cases of phosphate and humic acid, as well as real waters, colloidal MnO2 rather than dissolved Mn(III) was identified. This result was explained by the relatively low stability of Mn(III) complexes with phosphate, humic acid, and unknown ligand(s) in real waters against disproportionation as compared to Mn(III)−pyrophosphate, Mn(III)−EDTA, and Mn(III)− NTA complexes.46 The reactions of Mn(III) species with these EDCs were found to be first-order with respect to both oxidant and reductant (SI Figure S11). Measured second-order rate constants kMn(III) for E2 are shown in Table 2, while the kinetic data for the other four phenolic EDCs exhibiting the comparable reactivity are not presented.

Figure 4. Oxidation kinetics of E2 by Mn(VII) in real waters and the effect of pretreatment by ozonation (a), and the influence of selected model ligands on the oxidation kinetics of E2 by Mn(VII) (b). Symbols represent measured data and solid lines show the trend, while dashed lines represent the model predictions. Experimental conditions for panel a, pH = 8, [E2]0 = 0.15 μM, and [Mn(VII)]0 = 12 μM; for panel b, pH = 5, [E2]0 = 0.15 μM, [Mn(VII)]0 = 60 μM, [phosphate] = 5 mM, [pyrophosphate] = [EDTA] = [NTA] = 5 μM, and [humic acid] = 0.25 mg C/L.

Table 2. Measured Second-Order Rate Constants (kMn(III), M−1 s−1) for Mn(III) Reactions with E2a

that the biological degradation of E2 in real waters (which had been filtered when transported to the lab within several hours of sampling) in the control experiments was found to be negligible within the time investigated of several minutes. So, the possibility that the preozonation reduced the biological activity and thus slowed down the kinetics of E2 degradation is ruled out. It seems likely that MnO2 in situ formed upon Mn(VII) reduction plays a significant role as a mild oxidant, since several studies have reported that manganese oxides can readily oxidize phenolic EDCs in water.44,45 However, the addition of reduced species such as Mn(II), sulfite, and p-hydroquinone in low micromolar range, which could competitively consume Mn(VII) to yield MnO2, was noted to have negligible effect on the oxidation kinetics of E2 in synthetic buffer solutions at near neutral pH (SI Figure S5). The addition of preformed MnO2 colloids at low levels had negligible influence in both synthetic and real waters (SI Figure S6). By considering the different experimental conditions of this study versus previous reports for the reactions between EDCs and MnO2 (i.e., near neutral pH and reaction time of several minutes versus acidic pH and reaction time of several hours), it is not difficult to understand the negligible contribution of MnO2 at applied pH 8. According to our previous reports,27,28 the oxidation enhancement exerted by the background matrices of real waters can be rationalized by considering the contribution of dissolved Mn(III), of which the stability was enhanced in the presence of ligands in real waters, and the formation upon Mn(VII) reduction was accelerated by the reduced species therein.

pHb

kMn(III)−NTA

kMn(III)−pyrophosphate

kMn(III)−EDTA

5 6 7 8

19735 12181 5288 783

15199 1947 501 97

59 373 217 40

a

The molar ratio of [Mn(III)]:[ligand] was set at 1:10. bThe kinetic data at higher pH were not obtained due to the fast disproportionation and/or negligible reactivity of Mn(III).

The kMn(III) values were a function of pH and the complexing ligand, as well as the molar [Mn(III)]:[ligand] ratio,28 generally consistent with the trend of ligand-enhanced Mn(VII) oxidation (Figure 4b and SI Figure S7). This also means that the relative contribution of Mn(III) can reasonably explain the oxidation enhancement (i.e., kMn(III)∫ 0t [Mn(III)]dt versus kMn(VII)∫ 0t [Mn(VII)]dt). Since ∫ 0t [Mn(III)]dt would be much lower than ∫ 0t [Mn(VII)]dt under the condition of [EDCs] ≪ [Mn(VII)] and kMn(III)−EDTA ≪ kMn(VII) (Table 2 versus SI Table S1), it is not difficult to understand the negligible enhancement of EDTA in Mn(VII) reactions over the wide p H r a n g e ( i . e . , k M n ( I I I ) − E D T A ∫ 0t [ M n ( I I I ) ] d t ≪ kMn(VII)∫ 0t [Mn(VII)]dt). In contrast, in the cases of NTA and pyrophosphate at acidic pH 5 and 6, where kMn(III) ≫ kMn(VII), the oxidation enhancement may occur. The much lower reactivity of Mn(III)−EDTA complex relative to Mn(III) complexes with NTA and pyrophosphate (Table 1) may be related to their redox potentials [e.g., the standard reduction potential of 1.15 V for Mn(III)−pyrophosphate versus 0.824 V for Mn(III)−EDTA in acidic medium].30 Although no stable Mn(III) species are identified in the presence of phosphate and 1779



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(8) Koplin, D. W.; Furlong, E. T.; Meyer, M. T.; Thurman, E. M.; Zaugg, S. D.; Barber, L .B.; Buxton, H. T. Pharmaceuticals, hormones, and other organic wastewater contaminants in US streams, 1999− 2000: A national reconnaissance. Environ. Sci. Technol. 2002, 36, 1202−1211. (9) Huber, M. M.; Ternes, T. A.; von Gunten, U. Removal of estrogenic activity and formation of oxidation products during ozonation of 17α-ethinylestradiol. Environ. Sci. Technol. 2004, 38, 5177−5186. (10) Deborde, M.; Rabouan, S.; Duguet, J.-P.; Legube, B. Kinetics of aqueous ozone-induced oxidation of some endocrine disruptors. Environ. Sci. Technol. 2005, 39, 6086−6092. (11) Lee, Y.; Yoon, J.; von Gunten, U. Kinetics of the oxidation of phenols and phenolic endocrine disruptors during water treatment with ferrate (Fe(VI)). Environ. Sci. Technol. 2005, 39, 8978−8984. (12) Sharma, V. K.; Li, X. Z.; Graham, N.; Doong, R. A. Ferrate(VI) oxidation of endocrine disruptors and antimicrobials in water. J. Water Supply: Res. Technol.AQUA 2008, 57, 419−426. (13) Deborde, M.; von Gunten, U. Reactions of chlorine with inorganic and organic compounds during water treatment-Kinetics and mechanisms: A critical review. Water Res. 2008, 42, 13−51. (14) Deborde, M.; Rabouan, S.; Gallard, H.; Legube, B. Aqueous chlorination kinetics of some endocrine disruptors. Environ. Sci. Technol. 2004, 38, 5577−5583. (15) Huber, M. M.; Korhonen, S.; Ternes, T. A.; von Gunten, U. Oxidation of pharmaceuticals during water treatment with chlorine dioxide. Water Res. 2005, 39, 3607−3617. (16) Lee, Y.; Escher, B.; von Gunten, U. Efficient removal of estrogenic activity during oxidative treatment of waters containing steroid estrogens. Environ. Sci. Technol. 2008, 42, 6333−6339. (17) Sharma, V. K.; Mishra, S. K.; Nesnas, N. Oxidation of sulfonamide antimicrobials by ferrate(VI) [FeVIO42‑]. Environ. Sci. Technol. 2006, 40, 7222−7227. (18) Zhang, H.; Yamada, H.; Tsuno, H. Removal of endocrinedisrupting chemicals during ozonation of municipal sewage with brominated byproducts control. Environ. Sci. Technol. 2008, 42, 3375− 3380. (19) Benschoten, J. E.; Lin, W.; Knocke, W. R. Kinetic modeling of manganese(II) oxidation by chlorine dioxide and potassium permanganate. Environ. Sci. Technol. 1992, 26, 1327−1333. (20) Dietrich, A. M.; Hoehn, R. C.; Dufresne, L. C.; Buffin, L. W.; Rashash, D. M. C.; Parker, B. C. Oxidation of odorous and nonodorous algal metabolites by permanganate, chlorine, and chlorine dioxide. Water Sci. Technol. 1995, 31, 223−228. (21) Rodriguez, E.; Majado, M. E.; Meriluoto, J.; Acero, J. L. Oxidation of microcystins by permanganate: Reaction kinetics and implications for water treatment. Water Res. 2007, 41, 102−110. (22) Ma, J.; Graham, N.; Li, G. Effect of permanganate preoxidation in enhancing the coagulation of surface waterLaboratory case studies. J. Water Supply Res. Technol.AQUA 1997, 46, 1−10. (23) Korshin, G. V.; Chang, H.-S.; Frenkel, A. I.; Ferguson, J. F. Structural study of the incorporation of heavy metals into solid phase formed during the oxidation of EDTA by permanganate at high pH. Environ. Sci. Technol. 2007, 41, 2560−2565. (24) Waldemer, R. H.; Tratnyek, P. G. Kinetics of contaminant degradation by permanganate. Environ. Sci. Technol. 2006, 40, 1055− 1061. (25) Hu, L.; Martin, H.; Strathmann, T. J. Oxidation kinetics of antibiotics during water treatment with potassium permanganate. Environ. Sci. Technol. 2010, 44, 6416−6422. (26) Hu, L.; Stemig, A. M.; Wammer, K. H.; Strathmann, T. J. Oxidation of antibiotics during water treatment with potassium permanganate: Reaction pathways and deactivation. Environ. Sci. Technol. 2011, 45, 3635−3642. (27) Jiang, J.; Pang, S.-Y.; Ma, J. Oxidation of triclosan by permanganate (Mn(VII)): Importance of ligands and in situ formed manganese oxides. Environ. Sci. Technol. 2009, 43, 8326−8331. (28) Jiang, J.; Pang, S.-Y.; Ma, J. Role of ligands in permanganate oxidation of organics. Environ. Sci. Technol. 2010, 44, 4270−4275.

humic acid, it seems likely that the proposed mechanism can also account for ligand-enhanced oxidation in these cases in terms of the relatively low stability but much higher reactivity of Mn(III) (i.e., kMn(III)∫ 0t [Mn(III) dt] ≫ 0). Nevertheless, no oxidation enhancement by specific ligands was isolated to the condition of pH 8, where the real water experiments were conducted. One possible reason is that some other unknown Mn(III)-stabilizing ligands may be present in real waters, while other explanations including the contribution of oxidizing species such as Mn(V) and Mn(VI) intermediates,47 and phenoxyl radicals formed in nature organic matter48,49 may also be plausible. Further studies will be needed to test these hypotheses.

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ASSOCIATED CONTENT

S Supporting Information *

Additional text, figures, and tables addressing supporting data as mentioned in the text. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Phone: 86-451-86283010; fax: 86-451-86283010; e-mail: [email protected].

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ACKNOWLEDGMENTS This work was financially supported by the National Natural Science Foundation of China (51008104), China and Heilongjiang Province Postdoctoral Science Foundation (20110490106 and LBH-Z10150), the Fundamental Research Funds for the Central Universities (HIT.NSRIF.201188), the Open Project of State Key Laboratory of Urban Water Resource and Environment, HIT (QA201012), and the Funds for Creative Research Groups of China (51121062). We gratefully acknowledge Dr. Feng Qin and Lijun Li (Greater China Application Support Center of AB SCIEX) for the kind help in LC−MS/MS analysis.

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REFERENCES

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