Activation of oxygen with sulfite for enhanced Removal of Mn(II): The involvement of SO4•-

Water Research 157 (2019) 435e444

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Activation of oxygen with sulfite for enhanced Removal of Mn(II): The involvement of SO
4 Dandan Rao a, Yuankui Sun a, Binbin Shao a, Junlian Qiao a, b, c, Xiaohong Guan a, b, c, * a State Key Laboratory of Pollution Control and Resources Reuse, College of Environmental Science and Engineering, Tongji University, Shanghai, 200092, China b Shanghai Institute of Pollution Control and Ecological Security, Shanghai, 200092, China c International Joint Research Center for Sustainable Urban Water System, Tongji University, Shanghai, 200092, China

a r t i c l e i n f o

a b s t r a c t

Article history: Received 7 September 2018 Received in revised form 20 March 2019 Accepted 30 March 2019 Available online 3 April 2019

Taking advantage of the active oxidants generated in the process of Mn(II)-catalyzed sulfite oxidation by oxygen, this study sought to enhance Mn(II) removal from water by activating oxygen with sulfite. The results revealed that Mn(II) can be effectively oxidized by oxygen to MnO2 with the addition of sulfite under environmentally relevant conditions, and the performance of this process is dependent on the dosage of sulfite and the initial pH. Mn K-edge XANES analysis indicates that Mn(II) removal is primarily due to the transformation of Mn(II) to MnO2 and, secondarily, to the adsorption of Mn(II) on generated MnO2. Co-existing NaCl and CaCl2 negatively affect Mn(II) removal, while the presence of Fe(II) considerably enhances Mn(II) removal by improving both Mn(II) oxidation and Mn(II) adsorption on the generated solids. Consequently, Mn(II) removal is as high as 98% in the presence of 1.0 mg/L of Fe(II) and both the residual Mn (<0.1 mg/L Mn) and Fe (<0.3 mg/L Fe) can meet China's drinking water standard. The experiments with real water samples also demonstrate the effectiveness of the sulfite-promoted Mn(II) removal process, especially in the presence of Fe(II). The enhancing effect of sulfite on Mn(II) oxidation by oxygen is mainly associated with the generation of HSO 5 , and the critical step for generating

HSO 5 is the rapid oxidation of SO3 by oxygen. EPR and radical scavenging studies demonstrate that SO4 radical is the key reactive oxygen species responsible for Mn(II) oxidation by HSO . 5 © 2019 Elsevier Ltd. All rights reserved.

Keywords: Mn(II) Oxidation Peroxymonosulfate Sulfite Sulfate radicals (SO
4)

1. Introduction Groundwater is a preferred drinking water source when the potable surface water is inaccessible, particularly in arid and semiarid regions worldwide (Lyu et al., 2019). Over 400 major cities in China utilize groundwater as drinking water source (Jia et al., 2018). Manganese is a common mineral in groundwater (Piazza et al., 2019) and exists mainly as reduced soluble form of Mn(II). Mn concentration of groundwater in China can range up to 2.0 mg/L, attributed to factitious pollution and dissolution from soil and rock (Jia et al., 2018). Although Mn is one of the essential trace elements for human health (Lyu et al., 2019), excess Mn in water supply system could result in aesthetic, operational, and health-related

* Corresponding author. State Key Laboratory of Pollution Control and Resources Reuse, College of Environmental Science and Engineering, Tongji University, Shanghai, 200092, China. E-mail address: [email protected] (X. Guan). https://doi.org/10.1016/j.watres.2019.03.095 0043-1354/© 2019 Elsevier Ltd. All rights reserved.

problems (Scholz, 2016). The Chinese government established a drinking water standard of 0.1 mg/L for Mn in 2006 (GB5749-2006) and thus the removal of excess Mn(II) from groundwater is a necessity. Oxidation is a conventional method used to separate manganese from water by transforming soluble Mn(II) into sparingly soluble manganese oxides, which can be removed by subsequent coagulation-sedimentation treatment (Patil et al., 2016). However, the oxidation of Mn(II) in homogeneous solutions by atmospheric oxygen is kinetically slow, with a half-life of about 400 days, which is likely due to the d5 electron configuration of Mn(II) (Diem and Stumm, 1984). Therefore, a number of methods have been proposed to enhance the transformation of Mn(II) into Mn oxides in the process of drinking water treatment, including chemical oxidation with strong oxidants (such as chlorine, chlorine dioxide, ozone, ferrate and permanganate) (Van Benschoten et al., 1992; Choo et al., 2005; Goodwill et al., 2016), biological oxidation (Hoyland et al., 2014), and surface-catalyzed oxidation with O2
(Jones and Knocke, 2017). Recent studies suggest that
O 2 , HO , and

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Fe(VI) produced from photochemical reactions, electrochemical processes, and extracellular microbiological processes may present an efficient pathway for the oxidation of Mn(II) (Baral et al., 1986; Van Genuchten and Pena, 2017; Learman et al., 2011; Nico et al., 2002; Wuttig et al., 2013). These highly reactive species can react with aqueous Mn(II) with rate constants of 107e108 M1s1 (Table S1). Thus, even at quite low concentrations, these species can appreciably affect the geochemical cycles of several metals, including Mn(II), in natural water (Wuttig et al., 2013). It is well known that Mn(II) is an active catalyst for S(IV) autooxidation in aqueous solutions (Connick and Zhang, 1996; Grgic and Ber ci c, 2001). The S(IV) oxidation process has been extensively studied to understand its kinetics and mechanisms, largely due to its involvement in the formation of acid rain, the desulfurization of flue gas, and biological systems (Siskos et al., 1984; Connick and Zhang, 1996; Brandt and Van Eldik, 1995). It was found that various oxidants, including SO
(E(SO
5 5/

HSO 5 ) ¼ 1.1 V (vs. NHE) at pH 7.0), strongly oxidizing SO4 (E(SO4 / 2

SO4 ) ¼ 2.5e3.1 V) (Neta et al., 1988), and HO (E(HO /H2O) ¼ 2.7 V in acidic conditions, and E(HO
/HO) ¼ 1.8 V in neutral conditions) (Buxton et al., 1988), could generate in the process of Mn(II)catalyzed oxidation of SO2 by O2 (Connick and Zhang, 1996) (Eqs. (3)e(9) in Table 1). Previous studies have proposed that the role of Mn(II) in catalyzing the oxidation of SO2 by O2 was as a chain propagator, i.e., Mn(II) can be oxidized by SO
5 into Mn(III) firstly,
and then the generated Mn(III) oxidize SO2 3 to give SO3 , which will
combine with O2 to form SO5 again (Eqs. (3)e(5) in Table 1) (Berglund et al., 1993; Brandt and Van Eldik, 1995; Connick and Zhang, 1996). In addition, it is well known that Mn(III) is prone to disproportionate to MnO2 and Mn(II), with an equilibrium constant log K z 7e9 (Reaction 1) (Klewicki and Morgan, 1998). Taking together, the oxidation of Mn(II) to MnO2 by O2 may be accelerated 2 with the addition of HSO 3 /SO3 , the aqueous species of SO2. Zhang et al. (2010) reported that oxidative precipitation with a SO2/air mixture was much more effective than that with air alone for Mn(II) removal from a typical synthetic laterite waste solution containing 2.0 g/L Mn(II). However, in their study, the oxidation experiments were carried out at 60  C, and the initial concentration of Mn(II) was three orders of magnitude higher than those involved in drinking water treatment. These conditions are quite different from the scenarios involved in drinking water treatment. 2Mn(III) þ H2O / MnO2 þ Mn2þ þ 4Hþ log K z 7e9

(1)

Considering that O2 is ubiquitous in nature and sodium sulfite (Na2SO3) is inexpensive and environmentally-friendly, here we propose using sulfite to enhance the oxidative removal of Mn(II) from drinking water by O2. Specifically, the objectives of this study were to 1) evaluate the feasibility of adding sulfite to enhance Mn(II) removal from both simulated and real groundwater; 2) determine the influence of operating conditions (such as sulfite concentration, initial pH (pHini), and co-existing ions) on the performance of sulfite-enhanced Mn(II) oxidation by O2; and 3) clarify the mechanisms of Mn(II) oxidation in the system by characterizing the precipitated solids, the transformed sulfur species and the 2 generated radicals. It should be specified that both HSO 3 and SO3 are present in solution over the typical pH range (6.0e9.0) of groundwater (Jia et al., 2018) because the second pKa of H2SO3 is 7.2 (Neta and Huie, 1985). Thus, the term sulfite or SO2 3 used in the following section refers to the equilibrium mixture of HSO 3 and SO2 3 , unless otherwise specified. The results of this study may provide an alternative method for Mn(II) removal in drinking water treatment. 2. Materials and methods 2.1. Chemicals A complete listing of reagents is provided in Supporting Information (SI) Text S1. 2.2. Experimental procedures The process of carrying out batch experiments for investigating sulfite-promoted Mn(II) removal is summarized in Fig. 1. Specifically, experiments were performed in open air in two transparent wide-mouth glass bottles (V ¼ 250 mL). Unless otherwise noted, the 200 mL-working solution only contained 36 mM of MnCl2 (2.0 mg Mn(II)/L), and its initial pH (pHini) was adjusted to the predetermined value using HCl and NaOH. After >5 min of stirring (to ensure the working solution was saturated with O2), the reaction was initiated by spiking the solution with a set number of aliquots of 150 mM Na2SO3 stock solution to reach the targeted sulfite concentration. The stock solution of Na2SO3 was freshly prepared and used within 2 h for each set of experiment to avoid its oxidation before use and its pH was adjusted to that of the working solution before dosing to the working solution. To investigate the influence of co-existing ions, the constituent of interest (NaCl, Na2SO4,

Table 1 Major reactions involved in the Mn(II)/sulfite system. Remarks

Equation

Rate constant (k) or equilibrium constant (K)

Ref.

Chain Initiation

þ (1) FeOH2þ þ HSO 3 4 [FeOHSO3H] 2þ (2) [FeOHSO3H]þ / SO
þ H2O 3 þ Fe
(3) O2 þ SO
3 / SO5

Log K ¼ 2.78 k2 ¼ 0.065 s1 k3 ¼ 1.5e2.5  109 M1s1

þ  (4) SO
5 þ Mn(II) þ H / HSO5 þ Mn(III) þ 
(5) Mn(III) þ SO2 3 /HSO3 / Mn(II) þ SO3 (þH )  

2
2(6) SO
þ HSO / HSO þ SO (7) SO þ SO 5 3 5 3 5 3 / SO4 þ SO4

k4 ¼ 2  108e2  1010 M1s1 (pH 3.0) k5 > 2.4  104 M1s1 k6þk7 ¼ 2.5  104 M1s1(pH 4.9) k6þk7 ¼ 3.0  106 M1s1(pH 6.8) k6þk7 ¼ 1.3  107 M1s1(pH 8.7) k8 ¼ (2.6e8)  108 M1s1 (pH 4.0e8.0) k9 ¼ 3.5  102 M1s1 (pH 8.0) k9 ¼ 9.1  103 M1s1 (pH 2.9) k10 ¼ 1.4  108 M1s1 k11 ¼ (1.8e2.5)  108 M1s1

Kraft and Rudi (1989) Warneck and Ziajka (1995) (Brandt and Eldik, 1995; Connick and Zhang, 1996) Brandt and Eldik (1995) Berglund et al. (1994) Huie and Neta (1987)

Chain Propagation

2  2
(8) SO
4 þ HSO3-/SO3 / HSO4 /SO4 þ SO3 þ 2  2 (9) HSO 5 þ SO3 /HSO3 /2SO4 þ H

Chain Termination

Mn(II) Oxidation

(10) (11) (12) (13) (14) (15)


2SO
5 þ SO5 / O2 þ S2O8
2SO
þ SO / S O 3 3 2 6
2 SO
5 þ SO3 / S2 presO6 þ O2  HSO5 þ Mn(II) / Mn(III) þ OH þ SO
4 2SO
4 þ Mn(II) / Mn(III) þ SO4 2þ 2Mn(III) þ H2O / MnO2 þ Mn þ 4Hþ

k14 ¼ (1.6e8.3)  107 M1s1 (pH 3.0) log K z 7e9

Brandt and Eldik (1995) Brandt and Eldik (1995) Huie and Neta (1987) Berglund et al. (1994) Berglund et al. (1994) Brandt and Eldik (1995) Klewicki and Morgan (1998)

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Fig. 1. Overview of the typical process for batch experiments investigating sulfite-promoted Mn(II) removal. In stage I, sulfite decay was monitored. In stage II, MnO2 production and the residual HSO 5 concentration in solution were monitored with time. After a pre-determined reaction time, samples were extracted and filtered to determine Mn removal. The working solution in the flask was well-mixed with magnetic stirring throughout the reaction.

NaNO3, NaClO4, or CaCl2, FeCl2, or FeCl3) was spiked into the working solution before pH adjustment. Since Fe is a common cocontaminant with Mn(II) in groundwater with concentration generally below 10.0 mg/L (Jia et al., 2018), we also investigated the influence of both FeCl2 and FeCl3 on the Mn(II)/sulfite process. The initial concentration of Mn(II) in groundwater samples, collected from cities of Xi'an and Langfang, was increased to 36 mM (2.0 mg/L) manually by adding MnCl2 stock solution to examine the efficiency of Mn(II) removal in real water with the sulfite-promoted oxygenation method. Detailed information about two natural groundwater samples used in this study was provided in Text S2. For all experiments, the solution in the flask was well mixed via magnetic stirring and its temperature was maintained at 25 ± 1  C with a circulating water bath. No buffer was used in this study to avoid its potential influence, and no attempt was made to keep pH constant during the reaction. All experiments were performed in at least duplicate unless otherwise noted, the obtained data were averaged, and the corresponding standard deviations were determined and presented. To quantify the kinetics of sulfite consumption, Mn(II) oxidation, and HSO 5 reduction, we collected samples at different time intervals to examine the residual sulfite as well as the concentrations of MnO2 and HSO 5 in the solution. The samples for sulfite determination did not need to be filtered because no MnO2 generated until the depletion of sulfite. The samples collected for analyzing the amount of produced MnO2 and HSO 5 were not filtered. It should be specified that Mn(II) can be removed by transforming to non-soluble MnO2 and adsorbing to the generated MnO2. Both the generated MnO2 and the absorbed Mn(II) can be easily separated from water by filtration. Consequently, this study not only quantified the amount of in situ produced Mn oxides but also determined the amount of filterable Mn (including both MnO2 and Mn(II) adsorbed on MnO2) to better understand the mechanisms involved in sulfite-enhanced Mn(II) removal process. 2.3. Solid phase characterization After specific tests, the precipitates were collected on a 0.22-

mm-pore cellulose acetate membrane filter, washed with deionized water, freeze-dried under vacuum, and then placed into zippered bags for solid phase characterization. The oxidation states of manganese species in the collected precipitates were analyzed with X-ray absorption near-edge spectroscopy (XANES, Shanghai Synchrotron Radiation Facility, China). The detailed information for XANES analysis is presented in Text S3. The morphology of the precipitates was characterized with a Hitachi 4800 Field Emission Scanning Electron Microscope (SEM) at 3 kV. 2.4. Chemical analysis A Shanghai Leici pH meter with a saturated KCl solution as the electrolyte was used to measure the solution's pH. The concentration of sulfite was quantified with a modified colorimetric procedure with DTNB as the chromogenic agent (Humphrey et al., 1970). The detailed procedure for sulfite detection is presented in Text S4 (for the standard curve see Fig. S1). At the end of experiment, a 10-mL sample was collected and filtered through a 0.22-mm filter, quenched with one drop of Na2S2O3 stock solution, acidified with one drop of 65% HNO3, then subjected to determine the residual Mn concentration using a Perkin Elmer Optima 5300 DV ICPOES. The filterable Mn was determined by subtracting the concentration of residual Mn from the initial concentration of Mn(II). Qualitative analysis of the oxidized sulfur species generated in this process with ion chromatography (IC) and liquid chromatography-tandem mass spectrometry under multiple reaction monitoring modes (LC-MS-MS/MRM) revealed that HSO 5 and SO2 4 were the oxidation products of sulfite in sulfite-enhanced Mn(II) removal process. Details of the qualitative analysis of the oxidized sulfur species are presented in Text S5. In this study, a modified ABTS method was adopted to determine the concentrations of generated MnO2 and HSO 5 . Colorless 2,20 -azino-bis(3-ethylbenzothiazoline-6-sulfonate) (ABTS) can react with oxidants via single-electron transfer process to give ABTS
þ, a stable green colored radical ion, which can be quantified spectrophotometrically at 645 nm (Fan et al., 2017b). In our previous study, we found that HSO 5 interferes with the ABTS analysis of

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MnO2 at pH 2.0 (Sun et al., 2018a). Here, we found that the ability of HSO 5 to oxidize ABTS was efficiently suppressed at pH 6.0, whereas the oxidation by MnO2 was unaffected. Therefore, we could first obtain the MnO2 concentration at pH 6.0, then the total concentration of MnO2 and HSO 5 was derived at pH 2.0 where both of them reacted with ABTS at a stoichiometry of 2:1. The concentration of HSO 5 was calculated by the subtraction method (Relevant standard curves of the ABTS method see Fig. S2). Moreover, to avoid changes in MnO2 and HSO 5 concentrations during measurement, EDTA, a chelating agent, was added to the samples immediately after they were collected in order to halt the Mn(II) oxidation by complexing Mn(II). The influence of co-existing Mn(II) (100 mM) and HSO 5 (40 mM) on the detection of MnO2 with the ABTS method at pH 6.0 was found negligible (Fig. S3), verifying the validity of this modified ABTS method for determining the concentrations of MnO2 and HSO 5 . Details for analyzing the concentrations of MnO2 and HSO 5 are presented in Text S6. For the experiments investigating the effects of Fe(II) and Fe(III) on sulfite-promoted Mn(II) removal, Fe(III) may affect the analysis of MnO2 and HSO 5 . Dissolved Fe(III) was found capable of oxidizing ABTS to ABTS
þ at pH 2.0 (for the standard curve, see Fig. S2) but not at pH 6.0, therefore, the contribution of Fe(III) to the ABTS oxidation needs to be deducted from the determined total concentration of MnO2 and HSO 5 . The concentration of Fe(III) was obtained by subtracting the concentration of Fe(II) from that of total iron. The concentration of Fe(II) was examined with a modified ferrozine method using a Persee TU-1901 UV/visible spectrophotometer at a wavelength of 560 nm (Fan et al., 2017a), and the concentration of total iron was analyzed using ICP-OES. Different from Fe(III), the presence of other metal ions, such as Ca(II), Na(I), and Fe(II), had negligible influences on the determination of MnO2 and HSO 5 concentrations with the ABTS method. DMPO was used as the spin-trapping agent in the electron paramagnetic resonance (EPR) experiments to determine the presence of radicals in the system (Oh et al., 2016). Further details of EPR analysis are given in Text S7. 3. Results and discussion 3.1. Feasibility of sulfite-enhanced Mn(II) removal and products characterization The oxidative removal of 36 mM Mn(II) by aeration in pure water was evaluated at pHini 7.0. As expected, no filterable manganese was detected even after 120-min aeration (Fig. S4A), which is in line with results reported in other studies (Kouzbour et al., 2017; Morgan, 2005). However, upon the addition of sulfite, the Mn(II) solution turns yellow quickly, implying the occurrence of Mn(II) oxidation. Since the filterable manganese may be a mixture of Mn(II), Mn(III), and Mn(IV), we collected the XANES spectrum of the precipitates produced from the Mn(II)/sulfite process and made linear combination fit analysis to identify the Mn species. As illustrated in Fig. 2A, the major Mn species in the collected sample are Mn(IV) (85.7%) and Mn(II) (14.3%). Mn(II) contained in the filterable precipitate can be mainly ascribed to its adsorption on freshly generated MnO2. Therefore, Mn(II) can be removed by both transformation to MnO2 and adsorption on the generated MnO2. In the following section, the generated MnO2 (Reo, defined as the amount of removed Mn(II) by oxidation), the filterable Mn (Ref, defined as the amount of removed Mn(II) by filtration), as well as the amount of Mn removed by adsorption (Rea ¼ Ref - Reo) were quantified to better elucidate the process of sulfite-promoted Mn(II) removal. With regard to the Mn(II)-catalyzed oxidation of S(IV), a series of

 2 2 2 sulfur oxides, including SO2 4 , HSO5 , S2O6 , S2O8 , and S2O7 , are likely to be involved (Brandt and Van Eldik, 1995). IC was employed to clarify the final sulfite oxidation products but only sulfate (SO2 4 ) was detected (Fig. 2B). In view of the possible effects of strong alkaline IC mobile phase on the detection of other possible sulfur oxidation products, we subsequently analyzed the reaction solution with LC-MS-MS. Interestingly, under MRM mode, signals assigned  to SO2 4 (m/z 97.1) and HSO5 (m/z 113.0) were detected in the spectrum (Fig. 2C), and their major abundant product ions are the  same as that of standard SO2 4 and HSO5 solutions, respectively 2  (Figs. S5AeB). Therefore, SO4 and HSO5 are believed to be the major oxidized products of sulfite. This result is in accordance with that reported in previous literature (Fronaeus et al., 1998), where HSO 5 is proposed to be an important intermediate participating in the catalytic cycle for the manganese-catalyzed auto-oxidation of sulfite under acidic conditions, and SO2 4 is suggested to be the stable oxidation product of sulfite. Since SO2 4 is redox-inert yet HSO 5 is an unsymmetrical oxidant with a standard redox potential of 1.82 V (Ghanbari and Moradi, 2017), HSO 5 may be an important active oxidant for the enhanced oxidative removal of Mn(II) in the presence of sulfite.

3.2. Factors affecting the process of sulfite-promoted Mn removal 3.2.1. Effect of sulfite dosage The kinetics of sulfite consumption and MnO2 generation were determined as a function of sulfite dosage at pHini 7.0 in Mn(II)/ sulfite system and the results are shown in Fig. 3A and B. Control experiments (Fig. S4B) revealed that the oxidation of sulfite by oxygen is very slow without Mn(II) and only 1.8 mM sulfite is oxidized within 4 min. However, complete disappearance of sulfite was observed within 4 min with the presence of Mn(II) and at the same initial sulfite concentration (36 mM) and pHini (7.0), as depicted in Fig. 3A, which should be ascribed to the catalyzing effect of Mn(II) (Connick and Zhang, 1996). Fig. 3A demonstrates that the concentration of sulfite drops linearly with time after a lag phase in Mn(II)/sulfite system, regardless of the initial sulfite concentration, clearly indicating a zero-order dependence of the reaction rate on sulfite concentration. This result corresponds well with the observed constant disappearance rate of O2 in the process of Mn(II)-catalyzed sulfite oxidation by O2 (Connick and Zhang, 1996). The disappearance kinetics of sulfite at various initial sulfite concentrations in the presence of 36 mM Mn(II) are simulated with the zero-order rate law, as illustrated by the solid lines in Fig. 3A, and the obtained sulfite consumption rates are summarized in Table S2. As the initial concentration of sulfite increased from 36 to 108 mM, the sulfite consumption rate increases from 10.2 to 42.0 mM/min. However, a further increase in the sulfite concentration results in a lower but almost steady sulfite consumption rate. Some researchers reported that the first-order dependence of sulfite appearance rate at low sulfur concentrations changes to a zeroorder dependence at high sulfite concentrations (Ibusuki and Barnes, 1984; Martin and Hill, 1987), similar to the phenomenon observed in this study. Fig. 3B shows that the generation of MnO2 is negligible without sulfite dosing, but ~7.0 mM MnO2 is generated after 90 min once 36 mM of sulfite is spiked into the solution at the beginning of the reaction. The maximum amount of MnO2 is produced when the initial sulfite concentration is 108 mM. Interestingly, MnO2 does not accumulate until the depletion of sulfite, which may be ascribed to the rapid reduction of generated Mn(III) by sulfite before the depletion of sulfite (Brandt and Eldik, 1995; Podkrajsek et al., 2004). The initial generation kinetics of MnO2 as a function of initial sulfite concentration is simulated with the zero-order rate law, as

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Fig. 2. (A) Mn K-edge XANES spectrum of the precipitates collected in the SO23 -promoted Mn(II) removal system after reaction time of 120 min; (B) The IC chromatogram and (C) LC-MS-MS chromatograms of sample collected from Mn(II)/SO2 3 system after 5-min reaction time. For LC-MS-MS analysis, m/z values are set according to a series of sulfur oxides 2 that are expected to form resulting from SO2 3 oxidation. Experimental conditions: [Mn(II)]0 ¼ 36 mM, [SO3 ]0 ¼ 108 mM, pHini ¼ 7.0.

 Fig. 3. (A) Depletion kinetics of SO2 3 , (B) Generation kinetics of MnO2, (C) Removal of Mn(II) after reaction for 90 min and the quantity of generated HSO5 at different sulfite dosages in Mn(II)/sulfite system. The solid lines are the results of simulating the data with zero-order rate law. Experimental conditions: [Mn(II)]0 ¼ 36 mM, [SO2 3 ]0 ¼ 0e180 mM, pHini ¼ 7.0.

illustrated by the solid lines in Fig. 3B. The obtained MnO2 generation rate constants are listed in Table S2. Fig. 3C illustrates that 27.6e39.5% Mn(II) can be removed by filtration after reaction with 36e180 mM sulfite for 90 min. The fraction of adsorbed Mn(II) constitutes 11.5e29.6% of the filterable Mn, consistent with the linear combination fits of the Mn K-edge

XANES spectrum (Fig. 2A). The amounts of HSO 5 accumulated in the Mn(II)/sulfite system at the time point, where sulfite is depleted, were determined for each experimental set at pHini 7.0 and the results are presented in Fig. 3C. Generally, a positive correlation is observed between the amount of generated MnO2 and  that of generated HSO 5 , indicating that HSO5 is responsible for

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(3)

constant at pHini 7.5e9.0. For sets investigating the effects of initial sulfite concentration and pHini, we found there is a linear regression between the quantity of HSO 5 generated and the sulfite consumption rates (Fig. S8), but the reason for this remains unknown at this stage. It is revealed in the section “the effect of sulfite dosage”, the transformation rate of Mn(II) into MnO2 is strongly dependent on the quantity of HSO 5 generated when the pHini is 7.0, however, no such dependency is found in trials investigating the effects of pHini. Over the pHini range of 6.0e9.0, although the rate of sulfite oxidation along with the quantity of HSO 5 generated reach a maximum at pHini 6.0, the formation of MnO2 is negligible at this pH level. As pHini increases from 6.0 to 7.0, the percentage of Mn(II) transformed into MnO2 ascends considerably from 3.4% to 39.4% then drops to ~20% at pHini 7.5e9.0, with the highest percentage of Mn(II) oxidation achieved at pHini 7.0, where the quantity of HSO 5 generated is ~35 mM. The minor formation of MnO2 at pHini 6.0e6.5 in spite of the presence of a large amount of HSO 5 is likely associated with the large energy cost of oxidizing Mn2þ to its corresponding Mn(III) species (Morgan, 2005). Above pH 6.5, the shift of Mn2þ to the hydroxide form can reduce the energy barrier for oxidation (Morgan, 2005). Accordingly, it is reasonable that many previous studies reported the oxidation of Mn(II) by O2, ClO2, KMnO4, and by the electrocoagulation-H2O2 system increases progressively with increasing pH (Van Benschoten et al., 1992; Morgan, 2005; Van Genuchten and Pena, 2017). In this study, similar amounts of MnO2 are generated at pHini 7.5e9.0, which differs from the pH dependence of Mn(II) oxidation reported in the literature. This is because the solution pH drops to ~7.0 after the depletion of sulfite in Mn(II)/sulfite process when pHini is in the range of 7.5e9.0 (Table S2). In summary, whereas the generation of HSO 5 , the oxidant responsible for Mn(II) oxidation in Mn(II)/sulfite process, is favored under acidic conditions, Mn(II) is more easily oxidized by HSO 5 under neutral and alkaline conditions. As a result, the optimum removal of Mn(II) in the presence of sulfite is achieved at pHini 7.0 in pure water.

Fig. 4B exhibits the MnO2 generation kinetics whilst Fig. 4C summarizes both the amount of Mn removed by oxidation and/or adsorption after 120-min reaction and the amount of HSO 5 generated at different pHini levels. About ~60% of consumed sulfite is transformed into HSO 5 at pHini 6.0e6.5, with a further increase in pH, however, the amount of generated HSO 5 decreases sharply from ~60 mM at pHini 6.5, to ~15 mM at pHini 7.5, and remaine almost

3.2.3. Effect of various metal ions The influences of co-existing metal ions (Na(I), Ca(II), Fe(II), and Fe(III)) on sulfite-promoted Mn(II) removal were investigated at pHini 7.0. It was found that Mn(II) removal is significantly inhibited with the concentration of NaCl increasing from 0 to 200 mg of Na/L (Fig. 5A). To unravel the influence of anions, the kinetics of sulfitepromoted Mn(II) removal in the presence of NaCl, Na2SO4, NaClO4,

Mn(II) oxidation. Moreover, it was found that the generation of MnO2 is accompanied by the consumption of HSO 5 , and the molar ratio of consumed HSO 5 to that of generated MnO2 is ~1:1 (Fig. S6), indicating the overall reaction of Mn(II) with HSO 5 follows Reaction 2. 2þ þ HSO þ H2O / SO2 5 þ Mn 4 þ MnO2 þ 3H

(2)

3.2.2. Effect of pHini level After determining the optimum sulfite dosage (108 mM) for Mn(II) removal at pHini 7.0, the influence of pHini on the performance of sulfite-promoted Mn(II) removal was examined and the results are summarized in Figs. 4AeC. Fig. 4A demonstrates that the oxidation of sulfite is strongly dependent on pHini. No obvious lag phase is observed at pHini 6.0e6.5 but the duration of the lag phase is prolonged with increasing pHini from 7.0 to 9.0, which can be ascribed to the more difficult chain initiation process at pHini  7.5 (Lim and Hamrick, 1984). Excluding the lag phase and tailing data, the main portion of each dataset for sulfite disappearance can be simulated with the zero-order rate law. The sulfite disappearance rates are found to decrease gradually from 94.2 to ~10 mM/min with elevating pHini from 6.0 to 9.0, which is likely associated with changes in the species distribution of sulfite and Mn(II) at different pH values. The distributions of S(IV) oxide species (Neta and Huie, 1985) and Mn(II) species (Zhang et al., 1994) as a function of pH are shown in Fig. S7. Specifically, at pH > 6.0, the fraction of SO2 3 increases with pH, and HSO 3 is commonly documented to be more readily oxidized to S(VI) species than SO2 3 in the presence of Mn(II) (Hoffmann and Edwards, 1975). In addition, the shift of the dominant form of Mn(II) from Mn2þ (reactive) to its hydroxide form (MnOHþ) at pHini > 7.0 via Reaction 3 (Zhang et al., 1994) may also contribute to the decrease in sulfite consumption rate. Mn2þ þ H2O ⇔ MnOHþ þ Hþ K ¼ ca. 2  1011

 Fig. 4. Influence of pHini on (A) Depletion kinetics of SO2 3 , (B) Generation kinetics of MnO2, and (C) Removal of Mn(II) after reaction for 120 min and the quantity of generated HSO5 . Experimental conditions: [Mn(II)]0 ¼ 36 mM, [SO2 3 ]0 ¼ 108 mM, pHini ¼ 6.0e9.0.

D. Rao et al. / Water Research 157 (2019) 435e444

Fig. 5. Influence of (A) Na(I), (B) Ca(II), (C) Fe(III), and (D) Fe(II) on Mn(II) removal in Mn(II)/sulfite process and the corresponding quantity of generated HSO 5 . Experimental conditions: [Mn(II)]0 ¼ 36 mM, [SO2 3 ]0 ¼ 108 mM, t ¼ 120 min, pHini ¼ 7.0.

or NaNO3 at two concentrations (100 mg Na/L and 200 mg Na/L) were determined and the results are shown in Fig. S9. Obviously, the counter-anions have no obvious differential effect on the kinetics of MnO2 generation in the presence of excess Na(I). Thus, the negative effect of NaCl on sulfite-promoted Mn(II) removal should be associated with either Na(I) or the ionic strength. Fig. 5B demonstrates that CaCl2 at the same mass concentration as NaCl has much greater inhibiting effect on Mn(II) removal as well as HSO 5 generation than NaCl. The influences of NaCl and CaCl2 on MnO2 generation and HSO 5 generation were further compared by converting their concentrations to ionic strength, as demonstrated in Fig. S10. Obviously, the inhibited Mn(II) removal at elevated NaCl or CaCl2 concentration should be primarily ascribed to the increased ionic strength and secondly due to the cation species. The result is consistent with the negative effect of ionic strength on the kinetics of sulfite oxidation by oxygen (Martin and Hill, 1987). The slightly stronger inhibiting effect of Ca(II) compared to Na(I) might be due to the stronger interaction of Ca(II) with sulfite (Millero et al., 1989). The presence of Fe(II) or Fe(III) has very different effects on Mn(II) removal than that of Na(I) and Ca(II), as shown in Fig. 5C and D. The quantity of generated HSO 5 decreases from 35.1 mM to 25.1e26.6 mM due to the presence of 0.5e3.0 mg/L of Fe(III). The presence of 0.5 mg/L of Fe(III) slightly depresses MnO2 generation while the amount of generated MnO2 in the presence of 1.0e3.0 mg/L of Fe(III) equals to or is lightly greater than that without Fe(III). Fig. S11 summarizes the data on generated MnO2 and adsorbed Mn(II) obtained under various conditions. There is a good correlation between them, verifying that the freshly formed MnO2 offers adsorption sites for Mn(II). As Fe(III) concentration elevates from 1.0 to 3.0 mg/L, the amount of adsorbed Mn(II) increases progressively and the removal efficiency of total Mn is increased by 24% in the presence of 3.0 mg/L of Fe(III) although the amount of generated MnO2 varies slightly, which may be ascribed to the coagulating function of Fe(III) (Wang et al., 2019). The coagulating function of Fe(III) is also revealed by the larger flocs formed in Mn(II)/sulfite process with the presence of Fe(III) compared to its counterpart without Fe(III), as illustrated in Figs. S12AeB. Compared to Fe(III), Fe(II) has a much greater enhancing effect on Mn(II) removal. The dosing of Fe(II) not only considerably enhances the oxidation of Mn(II) but also promotes the adsorption of Mn(II) on generated solids. With the presence of 1.0 mg/L Fe(II), the percentage of Mn(II) transformed into MnO2 is enhanced to 71.8%,

441

which is 32% higher than that obtained without Fe(II). This may be due to the following reasons. Firstly, sulfite is completely consumed within 20 s in the presence of 1.0 mg/L Fe(II) (data not shown) and thus the duration for Mn(II) oxidation by HSO 5 is prolonged. Secondly, the enhanced Mn(II) oxidation may be ascribed to the activation of HSO 5 by residual Fe(II) after sulfite depletion or the solids generated in the process. The amount of adsorbed Mn(II) increases from 6.3% to 26.7% by increasing the Fe(II) concentration from 0 to 1.0 mg/L. The greater enhancing effect of Fe(II) on Mn(II) removal via adsorption than Fe(III) should be due to the better coagulating effect of Fe(III) formed in-situ than commercial Fe(III) (Guan et al., 2009) and the larger amount of generated MnO2, which provides adsorption sites for Mn(II). Consequently, Mn(II) removal is as high as 98% with the addition of 1.0 mg/L of Fe(II) and both the residual Mn (<0.1 mg/L Mn) and Fe (<0.3 mg/L Fe) can meet China's drinking water standard. Moreover, it was observed that the presence of Fe(II) benefits the aggregation and precipitation of the generated MnO2, which is consistent with the phenomenon that larger, tighter, and denser flocs are generated in the presence of Fe(II), as revealed by the SEM images in Fig. S12C. 3.3. Identification of the active oxidants contributing to Mn(II) oxidation The above evidence indicates that HSO 5 is the key species contributing to the sulfite-enhanced Mn(II) oxidation process. In addition to direct oxidation by HSO 5 , Mn(II) may also be oxidized by various reactive species since it has been documented that

Mn(II) and MnO2 can activate HSO 5 to generate SO4 and SO5 radicals following Reactions 4,5 (Anipsitakis and Dionysiou, 2004; Ghanbari and Moradi, 2017). The generated SO
4 radicals can further react with H2O/OH to produce HO
(Oh et al., 2016). Moreover, some researchers have proposed that HSO 5 selfdecomposition without explicit activation could generate singletoxygen (1O2) (Yang et al., 2018). All these reactive oxygen species

1 (ROSs), including SO
4 , SO5 , HO , and O2, are potential active oxidants contributing Mn(II) oxidation in Mn(II)/sulfite system. 2þ 2þ / [MnIII(SO
þ OH HSO 5 þ Mn 4 )]

(4)


 HSO 5 þ2MnO2 /Mn2O3 þ SO5 þ OH

(5)

To identify the major oxidants involved in the process of sulfiteenhanced Mn(II) oxidation, EPR spectra were collected using DMPO
as a spin trap for SO
4 and HO (Oh et al., 2016). As shown in Fig. 6A(1), only the characteristic signal of DMPO-SO
3 (aN ¼ 14.7 G, aH ¼ 15.9 G) (Sun et al., 2018b) was observed for the Mn(II)/sulfite system at pHini 7.0 when DMPO was spiked into the sulfite solution before the addition of Mn(II). As proposed in the literature, the initial generation of SO
3 in the absence of Mn(III) can be ascribed to the oxidation of Mn(II) into Mn(III) by trace quantities of Fe(III) (z108 mol/L) (Eqs. (1) and (2) in Table 1) or by an interaction between HSO c and 3 and traces of Fe(III) (Reactions 6,7) (Grgi Bercic, 2001; Warneck and Ziajka, 1995). Such low concentration of Fe(III) are usually present as impurities in chemicals and purified water (Grgi c and Ber ci c, 2001). The generated SO
3 can be rapidly oxidized by O2 into SO
5 with a rate constant as large as (1.5e2.5)  109 M1 s1 (Eq. (3) in Table 1) (Brandt and Van Eldik, 1995; Connick and Zhang, 1996). SO
5 can be further transformed 2 into HSO 5 and SO4 via Eqs. (4) and (6e9) in Table 1 (Berglund et al., 1994; Brandt and Van Eldik, 1995). þ Fe(III) þ HSO 3 % [FeOHSO3H] K ¼ 600 L/mol

(6)

1 [FeOHSO3H]þ / Fe2þ þ SO
3 þ H2O k ¼ 0.065 s

(7)

442

D. Rao et al. / Water Research 157 (2019) 435e444

Fig. 6. (A) EPR spectra obtained for Mn(II)/sulfite at pHini ¼ 7.0 after reaction for (1) 0 s, (2) 13 min, and for (3) Mn(II)/HSO 5 system at pHini ¼ 7.0 after reacting for 0 s (: stands for $ DMPO-SO$3 adduct and + stands for DMPO-HO adduct); (B) Effect of radical scavengers on Mn(II) oxidation in Mn(II)/sulfite system. Scavengers were added after the depletion of  2  SO2 3 ; (C) H2O2-induced decomposition of HSO5 . Experimental conditions: [Mn(II)]0 ¼ 36 mM, [SO3 ]0 ¼ 108 mM, [HSO5 ]0 ¼ 108 mM, [DMPO]0 z 100 mM, pHini ¼ 7.0.

Since Mn(II) was not oxidized until sulfite was depleted, the EPR spectrum of the Mn(II)/sulfite system was also collected during the Mn(II) oxidation phase, i.e., DMPO was spiked 13 min after the initiation of the reaction between Mn(II) and sulfite. As shown in Fig. 6 A(2), the spectrum characteristic of the DMPO- HO
adduct (quartet lines with peak a height ratio of 1:2:2:1) was observed (Xiong et al., 2015). In addition, the typical spectrum of the DMPOHO
adduct was also observed by mixing Mn(II), HSO 5 , and DMPO, as demonstrated in Fig. 6A(3). These results further confirm that Mn(II) is oxidized by the radicals generated via HSO 5 activation. Although only DMPO-HO
was detected in the Mn(II) oxidation process, the contribution of SO
4 radicals for Mn(II) oxidation cannot be excluded because the DMPO-SO
4 adduct is unstable and readily transforms into the more stable DMPO-HO
adduct via nucleophilic substitution (Oh et al., 2016). Moreover, no EPR signal for SO
5 has been reported in the literature to the best of our knowledge. To further clarify the contribution of various radicals to Mn(II) oxidation, the influence of quenching agents on Mn(II) transformation to MnO2 was investigated. It is well known that tertbutyl alcohol (TBA) can effectively quench HO
(k ¼ 5 1 1 (3.8e7.6)  108 M1s1) but not SO
4 (k ¼ (4.0e9.1)  10 M s ) while ethanol (EtOH) can effectively quench both HO
6 1 1 (k ¼ 9.7  108 M1s1) and SO
4 (k ¼ 3.2  10 M s ) (Buxton
et al., 1988; Neta et al., 1988). In addition, SO5 is relatively nonreactive toward both TBA and EtOH (k < 103 M1s1) (Neta et al., 1988). Both TBA and EtOH significantly slow sulfite consumption in the Mn(II)/sulfite system (as shown in Fig. S13), which may be due to the inhibition effect of alcohols on the radical chain reactions involved in sulfite oxidation (Wang et al., 2009). Thus, EtOH and TBA were dosed into the Mn(II)/sulfite system after the depletion of sulfite to assess the contributions of different species to Mn(II)

oxidation. As shown in Fig. 6B, the presence of 100 mM of TBA has a negligible influence on Mn(II) oxidation but the process of Mn(II) transformation to MnO2 is completely suppressed by 100 mM of
EtOH. These results suggest that neither SO
5 nor HO is the major active oxidant in the Mn(II)/sulfite system. 1 H2O2 is known to react with HSO 5 to generate O2 (Reaction 8) (Yang et al., 2018). To exclude the dominant role of 1O2 in the Mn(II) oxidation process, H2O2 was added after the depletion of sulfite. Despite the enhanced rate of 1O2 production, Mn(II) oxidation is greatly depressed (Fig. 6B), which can be explained by the H2O2induced partial decomposition of HSO 5 (Fig. 6C). Consequently, SO
4 is identified as the major reactive species responsible for Mn(II) oxidation and the reactions involved in the Mn oxidation phase are illustrated by Eqs. 13e15 in Table 1. The mechanisms for the sulfite-promoted removal of Mn(II) are schematically illustrated in Fig. 7A and the influences of various factors on sulfiteenhanced Mn(II) removal are summarized in Fig. 7B. Since Fe(II) is much more efficient than Fe(III) and Mn(II) for activating HSO 5 to generate SO
4 (Reaction 9) (Anipsitakis and Dionysiou, 2004), the dosing of Fe(II) greatly promoted the oxidation of Mn(II) in Mn(II)/ sulfite system. 2 þ 1 HSO 5 þ H2O2 / H þ SO4 þ H2O þ O2

(8)


 Fe(II) þ HSO 5 / Fe(III) þ SO4 þ OH

(9)

3.4. Effectiveness in real water The efficiency of Mn(II) removal from three real water samples

D. Rao et al. / Water Research 157 (2019) 435e444

443

Fig. 7. (A) Illustration of the proposed mechanisms for sulfite-promoted Mn(II) removal; (B) Influence of various factors on the process of sulfite-promoted Mn(II) removal.

was also examined to demonstrate the effectiveness of sulfiteenhanced Mn(II) oxidation process. The properties of the three water samples, i.e., a tap water sample from our laboratory and two groundwater samples collected from Langfang city and Xi'an, respectively, are summarized in Table S3. As presented in Fig. 8, at pHini 7.0 and upon the application of 108 mM sulfite, quite similar Mn removal (~30%) is observed for two groundwater samples, and ~54% Mn removal is achieved in tap water. The ineffective removal of Mn(II) in these three water samples can be ascribed to the Ca(II) and Na(I) present in real water samples (Table S3). The higher removal efficiency of Mn(II) in tap water compared to that in groundwater samples may be due to the much lower Na(I) concentration (53.4 mg/L) compared to that in the two groundwater samples (113.3e126.6 mg/L). The performance of sulfite-enhanced Mn(II) oxidation in real water samples was determined with the presence of 1.0 mg/L of Fe(II). Similar to the results in pure water, the presence of 1.0 mg/L

Fig. 8. The removal of Mn(II) from real waters enhanced by sulfite addition (A) or spiking both sulfite and Fe(II). Group A: [Mn(II)]0 ¼ 36 mM, [SO2 3 ]0 ¼ 108 mM, pHini ¼ 7.0, Group B: [Mn(II)]0 ¼ 36 mM, [Fe(II)]0 ¼ 1.0 mg/L, [SO2 3 ]0 ¼ 108 mM, pHini ¼ 7.0.

of Fe(II) greatly improves Mn(II) removal in the three real water samples and the Mn(II) removal efficiency is as high as 86% in the tap water. Furthermore, the residual Fe(II) concentration is below 0.10 mg/L for all three treated real water samples. Therefore, sulfiteenhanced Mn(II) removal can be an effective method for removing Mn(II) from groundwater, especially for groundwater containing both Mn(II) and Fe(II). 4. Conclusions This study offers an innovative approach for Mn(II) removal from water via the oxidants generated in the process of Mn(II)catalyzed sulfite oxidation under aerobic conditions. The major conclusions include: 1) The enhancing effect of sulfite on Mn(II) oxidation by O2 is associated with the generation of SO
3 , which can rapidly  combine with O2 to form SO
5 and further transforms into HSO5 . HSO 5 is the major oxidant contributing to the oxidation of Mn(II) to MnO2, which initiates after the depletion of sulfite. 2) EPR and radical scavenging studies demonstrate that the SO
4 radical is the crucial active oxidant responsible for Mn(II) oxidation by HSO 5. 3) NaCl and CaCl2 thwart Mn(II) removal in the presence of sulfite, which is primarily due to the inhibiting effect of high ionic strength on HSO 5 generation. Fe(III) has a slight influence while Fe(II) considerably enhances Mn(II) removal since Fe(II) is much
more efficient than Fe(III) for activating HSO 5 to generate SO4 and thus the oxidation of Mn(II) and the adsorption of Mn(II) to the generated MnO2 are promoted. 4) The experiments with real water samples also demonstrate the effectiveness of the sulfite-promoted Mn(II) removal process, especially in the presence of Fe(II). In sum, the use of sulfite to enhance Mn(II) removal from water may represent a promising alternative to conventional methods since it is cost-effective, environmentally friendly, easily applied, and the final stable product of sulfite is sulfate, a common anion in groundwater. Moreover, the concentration of generated sulfate in Mn(II)/sulfite process is much below the maximum contaminant

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level of the China's drinking water standard for sulfate (250 mg/L). Further studies with natural groundwater samples with elevated Mn(II) concentration under weak alkaline conditions are needed to verify the effectiveness of sulfite-enhanced Mn(II) removal approach. Acknowledgment This work was supported by the National Natural Science Foundation of China (Grant No. 21522704). The authors thank the beamline BL14W1 (Shanghai Synchrotron Radiation Facility) for providing the beam time. Appendix A. Supplementary data Supplementary data to this article can be found online at https://doi.org/10.1016/j.watres.2019.03.095. References Anipsitakis, G.P., Dionysiou, D.D., 2004. Radical generation by the interaction of transition metals with common oxidants. Environ. Sci. Technol. 38 (13), 3705e3712. Baral, S., Lume-Pereira, C., Janata, E., Henglein, A., 1986. Chemistry of colloidal manganese oxides. 3. formation in the reaction of hydroxyl radicals with Mn2þ ions. J. Phys. Chem. 90 (22), 6025e6028. Berglund, J., Elding, L.I., Buxton, G.V., McGowan, S., Salmon, G.A., 1994. Reaction of peroxomonosulfate radical with manganese(II) in acidic aqueous-solutionda pulse-radiolysis study. J. Chem. Soc. Faraday. T. 90 (21), 3309e3313. Berglund, J., Fronaeus, S., Elding, L.I., 1993. Kinetics and mechanism for manganesecatalyzed oxidation of sulfur(IV) by oxygen in aqueous solution. Inorg. Chem. 32 (21), 4527e4538. Brandt, C., Van Eldik, R., 1995. Transition metal-catalyzed oxidation of sulfur(IV) oxides. Atmospheric-relevant processes and mechanisms. Chem. Rev. 95 (1), 119e190. Buxton, G.V., Greenstock, C.L., Helman, W.P., Ross, A.B., 1988. Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals (,OH/,O) in Aqueous Solution. J. Phys. Chem. Ref. Data 17 (2), 513e886. Choo, K., Lee, H., Choi, S., 2005. Iron and manganese removal and membrane fouling during UF in conjunction with prechlorination for drinking water treatment. J. Membr. Sci. 267 (1e2), 18e26. Connick, R.E., Zhang, Y., 1996. Kinetics and mechanism of the oxidation of HSO-3 by O2. 2. the manganese(II)-catalyzed reaction. Inorg. Chem. 35 (16), 4613e4621. Diem, D., Stumm, W., 1984. Is dissolved Mn2þ being oxidized by O2 in absence of Mn-bacteria or surface catalysts? Geochim. Cosmochim. Ac. 48 (7), 1571e1573. Fan, P., Sun, Y., Qiao, J., Lo, I., Guan, X., 2017a. Influence of weak magnetic field and tartrate on the oxidation and sequestration of Sb(III) by zerovalent iron: batch and semi-continuous flow study. J. Hazard. Mater. 343 (5), 266e275. Fan, W., Qiao, J., Guan, X., 2017b. Multi-wavelength spectrophotometric determination of Cr(VI) in water with ABTS. Chemosphere 171, 460e467. Fronaeus, S., Berglund, J., Elding, L.I., 1998. Iron-manganese redox processes and synergism in the mechanism for manganese-catalyzed autoxidation of hydrogen sulfite. Inorg. Chem. 37 (19), 4939e4944. Ghanbari, F., Moradi, M., 2017. Application of peroxymonosulfate and its activation methods for degradation of environmental organic pollutants: review. Chem. Eng. J. 310, 41e62. Goodwill, J.E., Mai, X., Jiang, Y., Reckhow, D.A., Tobiason, J.E., 2016. Oxidation of manganese(II) with ferrate: stoichiometry, kinetics, products and impact of organic carbon. Chemosphere 159, 457e464. Grgi c, I., Ber ci c, G., 2001. A simple kinetic model for autoxidation of S(IV) oxides catalyzed by iron and/or manganese ions. J. Atmos. Chem. 39 (2), 155e170. Guan, X., Ma, J., Dong, H., Jiang, L., 2009. Removal of arsenic from water: effect of calcium ions on As(III) removal in the KMnO4-Fe(II) process. Water Res. 43 (20), 5119e5128. Hoffmann, M.R., Edwards, J.O., 1975. Kinetics of the oxidation of sulfite by hydrogen peroxide in acidic solution. J. Phys. Chem. 79 (20), 2096e2098. Hoyland, V.W., Knocke, W.R., Falkinham, J.O., 3rd Pruden, A., Singh, G., 2014. Effect of drinking water treatment process parameters on biological removal of manganese from surface water. Water Res. 66, 31e39. Huie, R.E., Neta, P., 1987. Rate constants for some oxidations of S(IV) by radicals in aqueous-solutions. Atmos. Environ. 21 (8), 1743e1747. Humphrey, R.E., Ward, M.H., Hinze, W., 1970. Spectrophotometric determination of sulfite with 4,4'-dithio-dipyridine and 5,5'-dithiobis(2-nitrobenzoic acid). Anal. Chem. 42 (7), 698e702. Ibusuki, T., Barnes, H.M., 1984. Manganese(II) catalyzed sulfur dioxide oxidation in aqueous solution at environmental concentrations. Atmos. Environ. 18 (1), 145e151. Jia, Y., Xi, B., Jiang, Y., Guo, H., Yang, Y., Lian, X., Han, S., 2018. Distribution, formation

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