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Role of Ferrate(IV) and Ferrate(V) in Activating Ferrate(VI) by Calcium Sulfite for Enhanced Oxidation of Organic Contaminants Binbin Shao,†,‡ Hongyu Dong,†,‡ Bo Sun,§ and Xiaohong Guan*,†,‡,∥ †

Environ. Sci. Technol. Downloaded from pubs.acs.org by UNIV OF KANSAS on 01/04/19. For personal use only.

State Key Laboratory of Pollution Control and Resources Reuse, College of Environmental Science and Engineering, Tongji University, Shanghai 200092, P. R. China ‡ Shanghai Institute of Pollution Control and Ecological Security, Shanghai 200092, P. R. China § Department of Civil and Environmental Engineering, The Hong Kong University of Science and Technology, Clear Water Bay, Kowloon, Hong Kong ∥ International Joint Research Center for Sustainable Urban Water System, Tongji University, Shanghai 200092, P. R. China S Supporting Information *

ABSTRACT: Although the Fe(VI)−sulfite process has shown great potential for the rapid removal of organic contaminants, the major active oxidants (Fe(IV)/ Fe(V) versus SO4•−/•OH) involved in this process are still under debate. By employing sparingly soluble CaSO3 as a slow-releasing source of SO32−, this study evaluated the oxidation performance of the Fe(VI)−CaSO3 process and identified the active oxidants involved in this process. The process exhibited efficient oxidation of a variety of compounds, including antibiotics, pharmaceuticals, and pesticides, at rates that were 6.1−173.7-fold faster than those measured for Fe(VI) alone, depending on pH, CaSO3 dosage, and the properties of organic contaminants. Many lines of evidence verified that neither SO4•− nor •OH was the active species in the Fe(VI)−CaSO3 process. The accelerating effect of CaSO3 was ascribed to the direct generation of Fe(IV)/Fe(V) species from the reaction of Fe(VI) with soluble SO32− via one-electron steps as well as the indirect generation of Fe(IV)/Fe(V) species from the self-decay of Fe(VI) and Fe(VI) reaction with H2O2, which could be catalyzed by uncomplexed Fe(III). Besides, the Fe(VI)−CaSO3 process exhibited satisfactory removal of organic contaminants in real water, and inorganic anions showed negligible effects on organic contaminant decomposition in this process. Thus, the Fe(VI)−CaSO3 process with Fe(IV)/Fe(V) as reactive oxidants may be a promising method for abating various micropollutants in water treatment.

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substrates (e.g., NOM and Cl−) because SO4•− is very reactive toward these substrates.2,21 Consequently, developing a more efficient and cost-effective activation process is necessary for sulfite application. Recently, there has been increasing interest in iron-mediated oxidation of refractory organic contaminants involving highvalent iron-based oxidants (Fe(IV) and Fe(V)) due to their high reactivity and selectivity.22−25 Fe(V) and Fe(IV) are known to be several orders of magnitude more reactive than ferrate (Fe(VI)).26,27 Although Fe(VI) is widely regarded as an effective green oxidant to treat a wide range of organic contaminants, the effectiveness of Fe(VI) treatment is highly variable and is even inefficient for some refractory contaminants.28 Therefore, the induction of Fe(V) and Fe(IV) generation may contribute to more efficient oxidation of organic contaminants than Fe(VI). The discovery of appropriate one-electron or two-electron reductants for promoting the transient generation of competent Fe(V) and

INTRODUCTION Sulfate radical (SO4•−)-based advanced oxidation processes (SR-AOPs) have been extensively studied for the remediation of organic pollutants in groundwater and wastewaters.1,2 The most frequently employed SO4•− precursor is either peroxydisulfate or peroxymonosulfate. Recently, sulfite has attracted considerable attention for the production of SO4•− as an active oxidant and has been viewed as a supplement to persulfateassociated AOPs.3 Various transition metals (TMs) (e.g., iron(II/III),4−8 cobalt(II),9 copper(II),10 chromium(VI),11−13 manganese(II)14) and TM-based heterogeneous catalysts (e.g., zerovalent iron,15,16 cobalt ferrite,17 zinc−copper ferrite,18 copper ferrite19) have been reported to activate sulfite and then produce SO4•− through chain-propagation steps (eqs 1 and 2 in Table 1).20 Compared to persulfates, sulfite has significant advantages of low toxicity, competitive price, and easy preparation, thus making it both environmentally friendly and economic.3 However, the low efficiency of these activation processes at near-neutral pH, intensive energy input (e.g., UV6−8), and secondary contamination (e.g., chromium(VI)) are still major limitations for the pilot-scale application. Moreover, one major disadvantage associated with SR-AOPs is that the performance is greatly affected by some coexisting © XXXX American Chemical Society

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September 5, 2018 December 18, 2018 December 20, 2018 December 20, 2018 DOI: 10.1021/acs.est.8b04990 Environ. Sci. Technol. XXXX, XXX, XXX−XXX

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Environmental Science & Technology

Table 1. Major Reactions Occurred in the Fe(VI)−CaSO3 Process and the Corresponding Second-Order Rate Constants no.

reactions •−

•−

ref

109 108 (alkaline) 108 (alkaline) 107(alkaline) 108 (alkaline) 109 (alkaline)

20 20 20 47 20 20

3 4 5

SO3 + O2 → SO5 SO5•− + SO32− → (a) SO52− + SO3•− SO5•− + SO32− → (b) SO42− + SO4•− SO4•− + OH− → SO42− + •OH SO4•− + SO32− → SO42− + SO3•− • OH + SO32− → SO3•− + OH−

1.1 1.0 5.6 7.3 7.5 5.4

6

2HFe VIO4− + 4H 2O → 2H3Fe IV O4− + 2H 2O2

26 (pH = 7.0)a

40

7 8 9

HFeVIO4− + H2O2 → H3FeIVO4− + O2 H3FeIVO4− + H2O2 + H+ → FeII(OH)2 + O2 + 2H2O HFeVIO4− + FeII(OH)2 + H2O → H2FeVO4− + FeIII(OH)3

10 (pH = 7.0)a ∼104 (pH = 7.0)a ∼107 (pH = 7.0)a

40 40 40

10

2H 2Fe V O4− + 2H 2O + 2H+ → 2Fe III(OH)3 + 2H 2O2

5.8 × 107 (pH = 7.0)a

40

5.6 × 10 (pH = 7.0)

40

1 2

V

2−

+ H2O2 + 2H → Fe (OH)3 + O2 + H2O +

III

11

HFe O4

12

2Fe(IV) → 1/3H 2O2 + 2Fe(III) + 1/3O2

13 14 15 16 17

× × × × × ×

k (M−1 s−1)

5

a

1.0 × 106 (pH = 7.5, 9.0)a

43

H2O2 + SO32− → SO42− + H2O HFeVIO4− + SO32− → HFeVO42− + SO3•− FeVIO42− + SO3•− → FeVO43− + SO3 HFeVO42− + SO3•− → HFeIVO43− + SO3

103−104 (pH = 9.0)a 1.9 × 108 (pH = 11.4)

3 57 58 35

2HFe V O4 2 − + 2SO32 − + 4H 2O → 2Fe III(OH)3 + 2SO4 2 − + 4OH−

3.9 × 104 (pH = 11.4)

58

a

In phosphate-buffered solution.

involvement of Fe(V) and Fe(IV) in the Fe(VI)−sulfite system, is far from clear and therefore warrants further study. On one hand, a high dosage of Na2SO3 may consume the generated active oxidants in the Fe(VI)−Na2SO3 system (eqs 4 and 5 in Table 1), while Na2SO3 at low concentration may not be enough to induce sufficient active oxidants to achieve the efficient oxidation of organic contaminants. On the other hand, the oxidation in the Fe(VI)−Na2SO3 system is too rapid to obtain solid evidence for active oxidants easily. Thus, calcium sulfite (CaSO3), which is sparingly soluble and can release SO32− slowly and continuously, was proposed to activate Fe(VI) in this study. The objectives of this study were to (1) evaluate the capability and selectivity of Fe(VI)−CaSO3 system for degrading various organic contaminants; (2) identify the contribution of Fe(V)/Fe(IV) species; (3) propose the mechanisms of active oxidants generation based on the process of iron and sulfur transformation; and (4) show the applicability of Fe(VI)−CaSO3 process.

Fe(IV) in the Fe(VI) oxidation process is therefore of prime importance. Several recent investigations have shown that sulfite could activate permanganate (Mn(VII)) to achieve rapid decomposition of organic pollutants via a nonradical pathway.29−32 The authors concluded that the active oxidant was Mn(III)aq and that this process could not only degrade the test organic contaminants 5 to 7 orders of magnitude faster than conventional AOPs but also destruct some organic contaminants that are refractory to Mn(VII).30−32 Since Fe(VI) is analogous to Mn(VII) with several high-valent reactive intermediates, sodium sulfite (Na2SO3) has been used to activate Fe(VI) for the enhanced removal of organic contaminants.33−35 All three studies reported that the selected organic contaminants could be decomposed within seconds in the Fe(VI)−Na2SO3 system, but different active oxidants have been proposed for the rapid oxidation of contaminants. Zhang et al.33 showed that both tert-butanol and methanol greatly inhibited sulfamethoxazole degradation and concluded that both SO4•− and •OH, generated via eq 3 in Table 1, contributed to the rapid oxidation of organic contaminants. However, Sun et al.34 demonstrated that SO4•− was the primary active species, mainly based on the negative effect of excess ethanol and the negligible influence of tert-butanol on target organic contaminant degradation. Nevertheless, Feng et al.35 reported that Fe(V) and Fe(IV) species, besides SO3•−, SO4•−, and •OH, were responsible for antibiotic oxidation in the Fe(VI)−Na2SO3 system under oxic conditions, but only Fe(V), Fe(IV), and SO3•− were involved under anoxic conditions, where the generation of oxysulfur species via radical chain reactions was cut off. The different conclusions reported by different researchers may be ascribed to the different reaction conditions (e.g., Na2SO3/Fe(VI) molar ratio) employed in their studies. Furthermore, little direct evidence for Fe(V) and Fe(IV) generation and their contribution to organic contaminants was provided.35 Therefore, the mechanisms of organics decomposition, especially the

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EXPERIMENTAL SECTION Chemicals and Reagents. A complete listing of reagents is provided in Text S1 of the Supporting Information (SI). Experimental Procedure. Batch degradation experiments were performed in 250 mL glass bottles equipped with a magnetic stirrer in a temperature-controlled room at 22 ± 2 °C. The experiments were conducted in air atmosphere unless otherwise noted. The initial concentrations of dissolved oxygen (DO) in working solution in air atmosphere and N 2 atmosphere were ∼8.9 and ∼0 mg/L, respectively. The experiments investigating the oxidation performance of Fe(VI)−Fe(III), Fe(VI)−H2O2, or Fe(VI)−CaSO3 processes were initiated by dosing filtered Fe(VI) solution immediately after addition of Fe(III), H2O2, or CaSO3 into the working solution containing the target contaminant. The pH of the working solution was buffered with either 10 mM borate or 10 mM borate/10 mM phosphate and was adjusted to the predetermined value by dropwise addition of sodium hydroxide or sulfuric acid, when necessary. During all of the B

DOI: 10.1021/acs.est.8b04990 Environ. Sci. Technol. XXXX, XXX, XXX−XXX

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Environmental Science & Technology tests in this study, pH change of the reaction solutions was within ±0.05. At defined time intervals, ∼ 1.5 mL of the sample was collected and immediately mixed with 20 μL of 200 mM Na2S2O3 to quench the reaction and then filtered through a 0.22 μm filter. All degradation experiments were performed in duplicate or triplicate. The obtained data were averaged, and the corresponding standard deviations were determined and presented. Analytical Methods. The concentrations of pollutants in the samples taken from the batch experiments were quantified with UPLC (Waters Co.). The details of this method are summarized in Text S2 and Table S1. The analytical methods for Fe(VI), H2O2, sulfite, formaldehyde, Fe(II), and Fe(III) as well as the details for collecting the electron spin resonance (ESR) spectra by using 5,5-dimethyl-1-pyrrolidine N-oxide (DMPO) as spin-trapping agent are presented in Text S2. Ultraperformance liquid chromatography-quadrupoletime-of flight premier mass spectrometer (UPLC-Q-TOF MS/MS, Waters Co.) was used to identify the degradation products of contaminants (detailed information is provided in the Supporting Information). DO was monitored with an oxygen microsensor (OX-25, Unisense A/S) equipped with a one-channel oxygen sensor amplifier (OXY-Meter, Unisense A/S). The competition kinetics method was employed to examine the contribution of SO4•− (Text S3).

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RESULTS AND DISCUSSION Degradation of Organic Contaminants in Fe(VI)− CaSO3 Process. The degradation kinetics of a variety of organics, including sulfamethoxazole (SMX), enrofloxacin (ENR), carbamazepine (CBZ), diclofenac sodium (DCF), atrazine (ATZ), ibuprofen (IBU),and benzoic acid (BA), by Fe(VI) alone (Figure 1a, b), and Fe(VI)−CaSO3 process (Figure 1c,d) were determined in either borate- or borate/ phosphate-buffered solutions at pH 8.0. All of the tested organic contaminants, except BA, were more rapidly degraded by Fe(VI) alone in borate buffer than their counterparts obtained in borate/phosphate buffer (Figure 1a,b). The application of CaSO3 greatly accelerated the degradation of all of the tested contaminants, except BA, by Fe(VI) in both buffers, yet the CaSO3-induced enhancement was dependent on the structure of the organic contaminants and the background buffers (Figure 1c,d). The main portion of each data set, except the tailing part, was well simulated with the pseudo-first-order rate law, as illustrated by the lines in Figure 1a−d, and the obtained rate constants (kobs) are summarized in Table S2 and Figure 1e,f. The Fe(VI)−CaSO3 process with borate/phosphate as buffer and that with borate as buffer oxidized organic contaminants at pH 8.0 with rates that were 6.1−12.6- and 16.5−173.7-fold faster than those measured for Fe(VI) alone, respectively. Among the test compounds, SMX was selected as the target pollutant to investigate the performance and mechanisms of the Fe(VI)−CaSO3 process in the following parts because its oxidation rate in the Fe(VI)− CaSO3 process was modest (Figure 1c) and its oxidation rate constants and degradation products by Fe(VI) and different reactive radical species (e.g., SO4•−, •OH) were welldocumented.16,36−38 As the initial CaSO3 dosage varied from 50 to 500 μM, the release rate of SO32− increased dramatically (Figure S1a), accompanied by the increase of initial degradation rate of SMX in Fe(VI)−CaSO3 process (Figure S1b). The crucial role of soluble SO32− in the oxidation process was demonstrated by a

Figure 1. Degradation kinetics of selected pollutants by Fe(VI) (a and b) and Fe(VI)−CaSO3 process (c and d) in borate or borate/ phosphate buffer and the corresponding kobs in Fe(VI) (e) and Fe(VI)−CaSO3 (f) processes. Reaction conditions: [Fe(VI)]0 = 50 μM, [SMX]0 = [ENR]0 = [CBZ]0 = [DCF]0 = [BA]0 = 5 μM, [ATZ]0 = [IBU]0 = 2.5 μΜ, [CaSO3]0 = 150 μM, pH = 8.0.

Ca2+ addition experiment (Figure S2a) in which excess Ca2+ inhibited SMX oxidation in the Fe(VI)−CaSO3 process due to its inhibition on SO32− release (Figure S2b). The maximum removal of SMX was achieved at a CaSO3 dosage of 150 μM (Figure S1b). The deteriorating effect of CaSO3 at a higher dosage on SMX removal may be associated with the competition between SO32− and SMX for the active oxidants in this process. Furthermore, the application of CaSO3 could considerably enhance the removal of SMX by Fe(VI) at pH 7.0−9.0 (Figure S3a−c), while the degradation of SMX in the Fe(VI)−CaSO3 process was slower at higher pH. This could be ascribed to a decrease of Fe(VI) oxidizing ability with increasing pH39 and the drop in the release rate of SO32− from CaSO3 particles as pH increased from 7.0 to 9.0 (Figure S3d). The kinetics of SMX degradation by Fe(VI) alone and via the Fe(VI)−CaSO3 process as functions of pH and CaSO3 dosages also followed the pseudo-first-order rate law, and the Fe(VI)− CaSO3 process oxidized SMX with rates that were 8.7−32.1fold greater than those measured for Fe(VI) alone (Table S2). Role of Fe(III) and H2O2 in the Absence of Phosphate. Lee et al. studied Fe(VI) decomposition in phosphate-buffered water at pH 7.0 and determined that Fe(III), H2O2, and O2 were the major end products (eqs 6−11 in Table 1).40 Uncomplexed Fe(III) could catalyze the self-decomposition of Fe(VI)41 and the reaction of H2O2 with Fe(VI).42 Because phosphate is a well-known complexing agent for Fe(III), the influence of phosphate on the decay of Fe(VI) as well as the generation and consumption of H2O2 in the Fe(VI)−CaSO3 process were determined (Figure 2a). Obviously, Fe(VI) C

DOI: 10.1021/acs.est.8b04990 Environ. Sci. Technol. XXXX, XXX, XXX−XXX

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were quantified in both buffers. Figure S5 showed that Fe(III) precipitate was the dominant form in borate buffer. In contrast, almost all of Fe(III) existed in soluble form in the presence of 10 mM phosphate (Figure S5). Therefore, the decomposition of H2O2 was likely associated with the catalyzing effect of Fe(VI) decomposition products (assumed to be Fe(III) precipitate),42 and the persisitence of H2O2 in phosphate buffer should be ascribed to the negligible generation of Fe(III) precipitate. The influence of phosphate on the decay rate of Fe(VI) and the decomposition of H2O2 in stage II (Figure 2a) indicates that Fe(VI) decomposition products in the Fe(VI)−CaSO3 process (Fe(III) precipitate and H2O2) might play an important role in accelerating organic contaminants oxidation (Figure 1). To verify this hypothesis, the influence of freshly prepared soluble Fe(III) and H2O2 on SMX degradation kinetics by Fe(VI) alone in both buffers was investigated (Figure S6). The removal of SMX by Fe(III) alone or Fe(III)/ H2O2 was negligible, while the presence of Fe(III) or H2O2 accelerated the degradation of SMX by Fe(VI) only in borate buffer. The rate constants (kobs) for SMX oxidation increased progressively from 0.14 to 0.33 min−1 as the added Fe(III) concentration increased from 0 to 50 μM and were enhanced from 0.14 to 0.32 min−1 as the H2O2 dosage was increased from 0 to 15 μM in borate buffer (Figure 2b). However, Fe(III) of 50 μM and H2O2 of 15 μM had little effect on SMX oxidation by Fe(VI) in borate/phosphate buffer. The selfdecay of Fe(VI) to Fe(IV), Fe(V), and H2O2 was catalyzed by Fe(III) precipitate,41 and H2O2 would further react with Fe(VI) to generate Fe(IV) (eq 7 in Table 1), which was also catalyzed by Fe(III) precipitate.42 When soluble Fe(III) was added to the Fe(VI)−CaSO3 process in borate buffer, Fe(III) precipitate was the dominant form and the amount of Fe(III) precipitate increased with the increase of initial soluble Fe(III) concentration. However, the dominant Fe(III) species was soluble phosphate−Fe(III) complex in phosphate buffer even when Fe(III) was spiked (Figure S5). Thus, the accelerating effects of both Fe(III) and H2O2 on SMX oxidation by Fe(VI) may be ascribed to the enhanced formation of Fe(IV) and/or Fe(V)these phenomena were only obvious without the presence of phosphate. Consequently, the enhancing effect of CaSO3 should be partially ascribed to the promoting effect of the products of its reaction with Fe(VI) in the absence of phosphate. Identification of Active Oxidants. In the process of SO32− oxidation by Fe(VI) open to air, the possible active oxidants in Fe(VI)−CaSO3 process include SO3•−, SO5•−, SO4•−, •OH, Fe(V), and Fe(IV) according to eqs 1−3 and eqs 14−16 (Table 1). Due to the poor oxidizing ability of SO3•− and SO5•−,44 neither of these radicals can be the oxidants contributing to the oxidation of the tested organic contaminants. Even in optimized AOPs, the maximum concentrations of SO4•− and •OH are low (10−14−10−12 M).45 Since the second-order rate constants of organic contaminants by SO4•− and •OH generally fall in the range of 109−1010 M−1 s−1,44,46 the concentrations of SO4•− and •OH should be as high as 10−12−10−10 M, calculated from kobs summarized in Figure 1c, if they were the major active oxidants in the Fe(VI)−CaSO3 process. Such high concentrations are unrealistic for highly reactive SO4•− and •OH in traditional AOPs. The negligible degradation of BA suggested that the concentrations of SO4•− and •OH generated in the Fe(VI)−CaSO3 process (if any) should be negligible.

Figure 2. Fe(VI) decay and H2O2 production in Fe(VI)−CaSO3 process in different buffered solutions (a) and the kobs of SMX oxidation by Fe(VI) alone, Fe(VI) with the presence of Fe(III) at various concentrations, and Fe(VI) with the presence of H2O2 at various concentrations (b). Reaction conditions: [Fe(VI)]0 = 50 μM, [CaSO3]0 = 150 μM, [SMX]0 = 5 μM, [Fe(III)]0 = 10, 30, 50 μM, [H2O2]0 = 3, 6, 10, 15 μM, pH = 8.0.

decomposed slower in the Fe(VI)−CaSO3 process with phosphate/borate as buffer than its counterpart with borate as buffer. Since less than 2 μM of Fe(VI) disappeared due to self-decay in 5 min (Figure S4), the decay of Fe(VI) with the presence of CaSO3 is likely associated with the reduction of Fe(VI) by CaSO3. Furthermore, rapid accumulation of H2O2 at the beginning of the reaction of Fe(VI) with CaSO3 was observed (Figure 2a). Because the rate of H2O2 generation resulting from the reaction of Fe(VI) with H2O (eq 6 in Table 1) was much slower than that from the self-decay of Fe(V) and Fe(IV) (eqs 10 and 12 in Table 1),40,43 it was suspected that Fe(VI) was reduced by CaSO3 to generate Fe(V) and Fe(IV), which quickly yielded H2O2 and Fe(III) by self-decay. Although the presence of phosphate had little influence during the H2O2 accumulation stage (stage I), it greatly affected the concentration profile of H2O2 in the following stage (stage II). The concentration of H2O2 remained fairly stable in borate/ phosphate buffer, which might be ascribed to the balance between H2O2 production from Fe(IV)/Fe(V) decay and H2O2 consumption by Fe(IV)/Fe(V) (eqs 8 and 11 in Table 1) and SO32− (eq 13 in Table 1). In contrast, the concentration of H2O2 decreased progressively in borate buffer in stage II (Figure 2a). A similar influence of phosphate buffer on the concentration profile of H2O2 accumulation and consumption during the reaction of Fe(VI) with bromide ion was reported by Jiang et al.42 In the Fe(VI)−CaSO3 process, the amounts of Fe(III) precipitate and soluble Fe(III) D

DOI: 10.1021/acs.est.8b04990 Environ. Sci. Technol. XXXX, XXX, XXX−XXX

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Environmental Science & Technology To further examine the role of •OH and SO4•−, alcoholscavenging experiments were performed by using TBA and MeOH as quenching agents. The second-order rate constants for the reaction of •OH with TBA (k•OH−TBA) and SMX (k•OH−SMX) are about (3.8−7.6) × 108 M−1 s−1 and 8.0 × 109 M−1 s−1, respectively.38,46 The oxidation of •OH toward 5 μM SMX should be completely inhibited by 10 mM TBA ((k•OH−SMX × [SMX])/(k•OH−TBA × [TBA]) < 1.05%). However, TBA exhibited slight effects on SMX degradation by Fe(VI), Fe(VI)−Fe(III), Fe(VI)-H2O2, and Fe(VI)-CaSO3 processes (Figure 3), implying that •OH was not the active

organics oxidation in the Fe(VI)−CaSO3 process. However, there is more evidence against the contribution of SO4•− in this process. First, the results of competition kinetics experiments with ATZ and SMX showed that the ratio of the second-order rate constant of ATZ oxidation by SO4•− to that of SMX oxidation by SO4•− at pH 8.0 (i.e., k SO4•− − ATZ /k SO4•− − SMX = 0.449) was much higher than that obtained in the Fe(VI)− CaSO3 process (0.059) (Figure S7). Second, formaldehyde, the oxidation product of MeOH by SO4•−,48 was not detectable in the Fe(VI)−CaSO3 process in the presence of 10 mM MeOH. Third, methyl phenyl sulfoxide (PMSO) at a much lower concentration (10 μM) had a much more significant adverse effect on SMX degradation than MeOH did (Figure 3), but it was resistant to oxidation in the Co2+− peroxymonosulfate system (Figure S8), which is a well-known SO4•−-dominated process.49 Moreover, to clarify the active oxidants contributing to SMX oxidation in Fe(VI)-CaSO3 process, UPLC-Q-TOF MS/MS was employed to analyze the transformation products (TPs) of SMX. Based on the molecular weights and degradation products reported in previous studies,16,36−38 six TPs of SMX were identified. The MS spectra and possible structures of the fragments of SMX and its TPs are presented in Figure S9. The accurate m/z values are provided in Table S3. The oxidation of SMX by SO4•− and •OH was generally initiated by the electrophilic attack at the olefinic double bond in the isoxazole ring to generate dihydroxylated product (m/z = 288 ([M + H]+)).16,38 This product was not detected during SMX oxidation in the Fe(VI)−CaSO3 process, further verifying the negligible contribution of SO4•− and •OH in this process. The observed suppressing effect of MeOH and anaerobic conditions on SMX oxidation in Figure 3 may be due to other mechanisms. ESR spectroscopy was also employed to monitor the evolution of reactive species using DMPO as a spin trap. ESR spectrum with distinctive signals of DMPO−OH adduct (αN = 14.9, αβ‑H = 14.9)50 was detected in the Fe(VI)−CaSO3 process but not in the Fe(VI) oxidation process within 2 min (Figure S10a,b). In view of the negligible role of SO4•− and • OH in the Fe(VI)−CaSO3 process, the formation of false positive signals of •OH might be ascribed to inverted spin trapping or the Forrester−Hepburn mechanism.50 Inverted spin trapping involves a direct electron-transfer oxidation reaction of DMPO (DMPO → DMPO•+) followed by nucleophilic attack by water (DMPO •+ → DMPO− OH),51,52 while the Forrester−Hepburn mechanism is the

Figure 3. Influence of various conditions on kobs of SMX degradation by Fe(VI), Fe(VI)−Fe(III), Fe(VI)-H2O2, and Fe(VI)-CaSO3 processes. Reaction conditions: [Fe(VI)]0 = 50 μM, [Fe(III)]0 = 50 μM, [H2O2]0 = 15 μM, [CaSO3]0 = 150 μM, [TBA]0 = [MeOH]0 = 10 mM, [DMSO]0 = [PMSO]0 = 10 μM, pH = 8.0 (10 mM borate).

oxidant for SMX degradation in these processes. MeOH is relatively reactive toward SO4•− (k SO4•− − MeOH = 2.0 × (106− 107) M−1 s−1).47 An obvious inhibition of SMX degradation by the Fe(VI)−CaSO3 process was observed when 10 mM MeOH was present (Figure 3). Meanwhile, SMX degradation was inhibited in the Fe(VI)−CaSO3 process in the absence of oxygen (Figure 3). The ineffective decomposition of organic contaminants in the Fe(VI)−Na2SO3 process under anoxic conditions has also been reported in the literature,33−35 which was regarded as a line of evidence for the role of SO4•− as an active oxidant. Without the presence of O2, the transformation of SO3•− to SO4•− would be completely stopped due to the chain reactions (eqs 1 and 2 in Table 1).35 This evidence seemingly supports the conclusion that SO4•− contributes to

Figure 4. Oxidation of PMSO and production of PMSO2 in various oxidation processes. Reaction conditions: [Fe(VI)]0 = 50 μM, [PMSO]0 = 10 μM, [Fe(III)]0 = 50 μM, [H2O2]0 = 15 μM, [CaSO3]0 = 150 μM, [MeOH]0 = 10 mM, pH = 8.0 (10 mM borate). E

DOI: 10.1021/acs.est.8b04990 Environ. Sci. Technol. XXXX, XXX, XXX−XXX

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Environmental Science & Technology reverse of this process (DMPO → DMPO−OH− → DMPO− OH).53,54 A previous report showed that the DMPO−OH− product of nucleophilic addition could be oxidized to the DMPO−OH adduct by compounds with low redox potential (−0.26 to 0.74 V/SHE),52 which is lower than that of highvalent iron−oxo species under alkaline conditions (e.g., E°(FeVI/FeIII) = 0.70 V/SHE, E°(FeIV/FeIII) = 1.0−1.4 V/ SHE).27 Therefore, the formation of the DMPO−OH adduct in the Fe(VI)−CaSO3 process was possibly due to the oxidation of DMPO by Fe(V) and/or Fe(IV). Sulfoxides (e.g., dimethyl sulfoxide (DMSO), PMSO) are known to be oxidized by high-valent iron to produce the corresponding sulfones (e.g., dimethyl sulfone (DMSO2) and methyl phenyl sulfone (PMSO2)) through an oxygen-atomtransfer step, which is markedly different from the radicalbased oxidation pathway.24,55,56 The oxidation of SMX by all of the studied processes was significantly inhibited by 10 μM DMSO or PMSO (Figure 3). Since DMSO or PMSO had little effect on Fe(VI) decay, especially during the first 5 min of the reaction (Figure S4), the reaction of Fe(VI) with DMSO or PMSO was very slow. The inhibiting effect of DMSO and PMSO on SMX oxidation for the Fe(VI)−CaSO3 process could likely be ascribed to the competition between DMSO/ PMSO and SMX for Fe(IV) and/or Fe(V) species, which were the major reactive species for SMX degradation. An obvious drop in the intensity of DMPO−OH adduct peaks was observed when 100 μM PMSO was present in the Fe(VI)− CaSO3 system (Figure S10b,c), indicating that PMSO was readily oxidized by Fe(V) and/or Fe(IV) species. During the oxidation of PMSO by Fe(VI), Fe(VI)−Fe(III), Fe(VI)− H2O2, and Fe(VI)−CaSO3 processes, almost all of the oxidized PMSO was transformed to PMSO2 (Figure 4), confirming that PMSO was oxidized via the oxygen-atom-transfer reaction. The presence of Fe(III) or H2O2 accelerated PMSO oxidation by Fe(VI) (Figure 4a−c), and the application of CaSO3 significantly improved the oxidation rate of PMSO by Fe(VI) (Figure 4d). Complete transformation of PMSO to PMSO2 was achieved within 0.5 min in the Fe(VI)−CaSO3 process. The oxidation of PMSO by Fe(VI) alone within such a short period was 10 μM within 1 min, Figure 2a) is over 5 orders of magnitude higher than that of SO5•− under the experimental conditions in this study. Therefore, the reaction of SO5•− with SO32− and the subsequent yield of SO4•− was inhibited by the presence of Fe(VI) (kFe(VI) − SO32−[Fe(VI)]≫k SO5•− − SO32−[SO5•−]). To further verify the proposed mechanisms, the oxidation of PMSO by the Fe(VI)−Na2SO3 process with various initial concentrations of Na2SO3 (10−500 μM) was investigated (Figure S13a). The contribution of Fe(V)/Fe(IV) was quantified by the ratio of formed PMSO2 to degraded PMSO (i.e., Δ[PMSO2]/Δ[PMSO]). Figure S13a indicates that Δ[PMSO2]/Δ[PMSO] decreased gradually from ∼96% to ∼54% with increasing doses of Na2SO3 from 10 to 500 μM. Since PMSO was resistant to SO4•− oxidation (Figure S8) but readily reacted with •OH (Figure S14), and •OH could be generated from the reaction of SO4•− with OH− (eq 3 in Table 1), both •OH and Fe(V)/Fe(IV) were the active oxidants for PMSO oxidation in the Fe(VI)−Na2SO3 process when SO32− was applied at higher concentrations (50−500 μM). The oxidation of BA by the Fe(VI)−Na2SO3 process with various initial concentrations of Na2SO3 (10−500 μM) was also investigated (Figure S13b). The oxidation of BA was negligible at 10 μM Na2SO3, indicating that BA was resistant to Fe(IV)/ Fe(V) oxidation. However, the oxidation of BA was promoted when SO32− was applied at higher concentrations (50−500 μM), which was ascribed to the generation of reactive radicals (SO4•−/•OH). This was in line with conclusions drawn in previous studies,33,34 where a high [SO32−]/[Fe(VI)] ratio (≥4:1) was employed. When Na2SO3 was applied at high concentrations, the evolution of oxysulfur species via radical chain reaction (eqs 1−3 in Table 1) was facilitated because Fe(VI) was depleted rapidly and completely by excess SO32−. The release of SO32− became faster as CaSO3 dosage increased from 50 to 500 μM (Figure S1). To examine the potential role of SO4•−/•OH in the Fe(VI)−CaSO3 process at a high CaSO3 concentration of 500 μM, the oxidation of PMSO was investigated (Figure S15). A balance between PMSO disappearance and PMSO2 generation was observed, indicating that Fe(IV) and Fe(V) were still the predominant oxidants. Because the release of soluble SO32− was limited by the solubility of CaSO3 and the released SO32− was immediately oxidized by Fe(VI), the [SO32−]/[Fe(VI)] ratio was always very low, and thus the formation SO4•−/•OH was suppressed.

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ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.est.8b04990.

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Three texts, three tables, and 19 figures with experimental details and additional results (PDF)

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel: +86-21-65980956. Fax: 86-21-65986313. ORCID

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Xiaohong Guan: 0000-0001-5296-423X

ENVIRONMENTAL IMPLICATIONS Fe(IV) and Fe(V) have been of great interest because of their much higher reactivity than Fe(VI), yet the application of Fe(IV)/Fe(V) in solid form has far been limited, partially due to the difficulty in synthesis and storage.27 This study showed that Fe(IV) and Fe(V) could be generated effectively in situ by using CaSO3 to activate Fe(VI). Moreover, the generated Fe(IV) and Fe(V) could efficiently and rapidly oxidize various

Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the National Natural Science Foundation of China (Grant 21522704) and State Key Laboratory of Pollution Control and Resource Reuse Foundation (No. PCRRK16001). The ESR spectra were G

DOI: 10.1021/acs.est.8b04990 Environ. Sci. Technol. XXXX, XXX, XXX−XXX

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Environmental Science & Technology

vis light irradiation: Green strategy to generate SO4•−. Appl. Catal., B 2018, 221, 380−392. (19) Chen, L.; Luo, T.; Yang, S.; Xu, J.; Liu, Z.; Wu, F. Efficient metoprolol degradation by heterogeneous copper ferrite/sulfite reaction. Environ. Chem. Lett. 2018, 16 (2), 599−603. (20) Das, T. N. Reactivity and role of SO5•− radical in aqueous medium chain oxidation of sulfite to sulfate and atmospheric sulfuric acid generation. J. Phys. Chem. A 2001, 105, 9142−9155. (21) Zhou, D.; Zhang, H.; Chen, L. Sulfur-replaced Fenton systems: Can sulfate radical substitute hydroxyl radical for advanced oxidation technologies? J. Chem. Technol. Biotechnol. 2015, 90 (5), 775−779. (22) de Oliveira, F. T.; Chanda, A.; Banerjee, D.; Shan, X.; Mondal, S.; Que, L., Jr; Bominaar, E. L.; Münck, E.; Collins, T. J. Chemical and spectroscopic evidence for an FeV-oxo complex. Science 2007, 315, 835−838. (23) de Sousa, D. P.; Miller, C. J.; Chang, Y.; Waite, T. D.; McKenzie, C. J. Electrochemically generated cis-carboxylato-coordinated iron(IV) oxo acid-base congeners as promiscuous oxidants of water pollutants. Inorg. Chem. 2017, 56 (24), 14936−14947. (24) Li, H.; Shan, C.; Pan, B. Fe(III)-doped g-C3N4 mediated peroxymonosulfate activation for selective degradation of phenolic compounds via high-valent iron-oxo species. Environ. Sci. Technol. 2018, 52 (4), 2197−2205. (25) Feng, M.; Jinadatha, C.; McDonald, T. J.; Sharma, V. K. Accelerated oxidation of organic contaminants by ferrate(VI): The overlooked role of reducing additives. Environ. Sci. Technol. 2018, 52 (19), 11319−11327. (26) Rush, J. D.; Cyr, J. E.; Zhao, Z.; Bielski, B. H. J. The oxidation of phenol by ferrate(VI) and ferrate(V). A pulse radiolysis and stopped-flow study. Free Radical Res. 1995, 22 (4), 349−360. (27) Sharma, V. K. Oxidation of inorganic contaminants by ferrates (VI, V, and IV)-kinetics and mechanisms: A review. J. Environ. Manage. 2011, 92 (4), 1051−1073. (28) Yang, B.; Ying, G. G.; Zhao, J. L.; Liu, S.; Zhou, L. J.; Chen, F. Removal of selected endocrine disrupting chemicals (EDCs) and pharmaceuticals and personal care products (PPCPs) during ferrate(VI) treatment of secondary wastewater effluents. Water Res. 2012, 46 (7), 2194−2204. (29) Sun, B.; Bao, Q.; Guan, X. Critical role of oxygen for rapid degradation of organic contaminants in permanganate/bisulfite process. J. Hazard. Mater. 2018, 352, 157−164. (30) Sun, B.; Dong, H.; He, D.; Rao, D.; Guan, X. Modeling the kinetics of contaminants oxidation and the generation of manganese(III) in the permanganate/bisulfite process. Environ. Sci. Technol. 2016, 50 (3), 1473−1482. (31) Sun, B.; Guan, X.; Fang, J.; Tratnyek, P. G. Activation of manganese oxidants with bisulfite for enhanced oxidation of organic contaminants: The involvement of Mn(III). Environ. Sci. Technol. 2015, 49 (20), 12414−12421. (32) Sun, B.; Li, D.; Linghu, W.; Guan, X. Degradation of ciprofloxacin by manganese(III) intermediate: Insight into the potential application of permanganate/bisulfite process. Chem. Eng. J. 2018, 339, 144−152. (33) Zhang, J.; Zhu, L.; Shi, Z.; Gao, Y. Rapid removal of organic pollutants by activation sulfite with ferrate. Chemosphere 2017, 186, 576−579. (34) Sun, S.; Pang, Y.; Jiang, J.; Ma, J.; Huang, Z.; Zhang, J.; Liu, Y.; Xu, C.; Liu, Q.; Yuan, Y. The combination of ferrate(VI) and sulfite as a novel advanced oxidation process for enhanced degradation of organic contaminants. Chem. Eng. J. 2018, 333, 11−19. (35) Feng, M.; Sharma, V. K. Enhanced oxidation of antibiotics by ferrate(VI)-sulfur(IV) system: Elucidating multi-oxidant mechanism. Chem. Eng. J. 2018, 341, 137−145. (36) Kim, C.; Panditi, V. R.; Gardinali, P. R.; Varma, R. S.; Kim, H.; Sharma, V. K. Ferrate promoted oxidative cleavage of sulfonamides: Kinetics and product formation under acidic conditions. Chem. Eng. J. 2015, 279, 307−316.

collected with the Steady High Magnetic Field Facilities, High Magnetic Field Laboratory, CAS.

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REFERENCES

(1) Anipsitakis, G. P.; Dionysiou, D. D. Radical generation by the interaction of transition metals with common oxidants. Environ. Sci. Technol. 2004, 38 (13), 3705−3712. (2) Oh, W. D.; Dong, Z.; Lim, T. T. Generation of sulfate radical through heterogeneous catalysis for organic contaminants removal: Current development, challenges and prospects. Appl. Catal., B 2016, 194, 169−201. (3) Zhou, D.; Chen, L.; Li, J.; Wu, F. Transition metal catalyzed sulfite auto-oxidation systems for oxidative decontamination in waters: A state-of-the-art minireview. Chem. Eng. J. 2018, 346, 726− 738. (4) Chen, L.; Peng, X.; Liu, J.; Li, J.; Wu, F. Decolorization of orange II in aqueous solution by an Fe(II)/sulfite system: Replacement of persulfate. Ind. Eng. Chem. Res. 2012, 51 (42), 13632−13638. (5) Zhou, D.; Yuan, Y.; Yang, S.; Gao, H.; Chen, L. Roles of oxysulfur radicals in the oxidation of acid orange 7 in the Fe(III)− sulfite system. J. Sulfur Chem. 2015, 36 (4), 373−384. (6) Zhou, D.; Chen, L.; Zhang, C.; Yu, Y.; Zhang, L.; Wu, F. A novel photochemical system of ferrous sulfite complex: kinetics and mechanisms of rapid decolorization of Acid Orange 7 in aqueous solutions. Water Res. 2014, 57, 87−95. (7) Guo, Y.; Lou, X.; Fang, C.; Xiao, D.; Wang, Z.; Liu, J. Novel photo-sulfite system: toward simultaneous transformations of inorganic and organic pollutants. Environ. Sci. Technol. 2013, 47 (19), 11174−11181. (8) Zhang, L.; Chen, L.; Xiao, M.; Zhang, L.; Wu, F.; Ge, L. Enhanced decolorization of orange II solutions by the Fe(II)−sulfite system under xenon lamp irradiation. Ind. Eng. Chem. Res. 2013, 52 (30), 10089−10094. (9) Yuan, Y.; Zhao, D.; Li, J.; Wu, F.; Brigante, M.; Mailhot, G. Rapid oxidation of paracetamol by Cobalt(II) catalyzed sulfite at alkaline pH. Catal. Today 2018, 313, 155−160. (10) Chen, L.; Tang, M.; Chen, C.; Chen, M.; Luo, K.; Xu, J.; Zhou, D.; Wu, F. Efficient bacterial inactivation by transition metal catalyzed auto-oxidation of sulfite. Environ. Sci. Technol. 2017, 51 (21), 12663− 12671. (11) Jiang, B.; Liu, Y.; Zheng, J.; Tan, M.; Wang, Z.; Wu, M. Synergetic transformations of multiple pollutants driven by Cr(VI)sulfite reactions. Environ. Sci. Technol. 2015, 49 (20), 12363−12371. (12) Yuan, Y.; Yang, S.; Zhou, D.; Wu, F. A simple Cr(VI)-S(IV)-O2 system for rapid and simultaneous reduction of Cr(VI) and oxidative degradation of organic pollutants. J. Hazard. Mater. 2016, 307, 294− 301. (13) Dong, H.; Wei, G.; Fan, W.; Ma, S.; Zhao, H.; Zhang, W.; Guan, X.; Strathmann, T. J. Reinvestigating the role of reactive species in the oxidation of organic co-contaminants during Cr(VI) reactions with sulfite. Chemosphere 2018, 196, 593−597. (14) Connick, R. E.; Zhang, Y. X. Kinetics and mechanism of the oxidation of HSO3− by O2. 2. The manganese(II)-catalyzed reaction. Inorg. Chem. 1996, 35 (16), 4613−4621. (15) Xiong, X.; Gan, J.; Zhan, W.; Sun, B. Effects of oxygen and weak magnetic field on Fe0/bisulfite system: performance and mechanisms. Environ. Sci. Pollut. Res. 2016, 23 (16), 16761−16770. (16) Du, J.; Guo, W.; Wang, H.; Yin, R.; Zheng, H.; Feng, X.; Che, D.; Ren, N. Hydroxyl radical dominated degradation of aquatic sulfamethoxazole by Fe0/bisulfite/O2: Kinetics, mechanisms, and pathways. Water Res. 2018, 138, 323−332. (17) Liu, Z.; Yang, S.; Yuan, Y.; Xu, J.; Zhu, Y.; Li, J.; Wu, F. A novel heterogeneous system for sulfate radical generation through sulfite activation on a CoFe2O4 nanocatalyst surface. J. Hazard. Mater. 2017, 324, 583−592. (18) Huang, Y.; Han, C.; Liu, Y.; Nadagouda, M. N.; Machala, L.; O’Shea, K. E.; Sharma, V. K.; Dionysiou, D. D. Degradation of atrazine by ZnxCu1−xFe2O4 nanomaterial-catalyzed sulfite under UV− H

DOI: 10.1021/acs.est.8b04990 Environ. Sci. Technol. XXXX, XXX, XXX−XXX

Article

Environmental Science & Technology (37) Sharma, V. K.; Mishra, S. K.; Nesnas, N. Oxidation of sulfonamide antimicrobials by ferrate(VI) [FeVIO42‑]. Environ. Sci. Technol. 2006, 40 (23), 7222−7227. (38) Yang, Y.; Lu, X.; Jiang, J.; Ma, J.; Liu, G.; Cao, Y.; Liu, W.; Li, J.; Pang, S.; Kong, X.; Luo, C. Degradation of sulfamethoxazole by UV, UV/H2O2 and UV/persulfate (PDS): Formation of oxidation products and effect of bicarbonate. Water Res. 2017, 118, 196−207. (39) 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 (22), 8978− 8984. (40) Lee, Y.; Kissner, R.; von Gunten, U. Reaction of ferrate(VI) with ABTS and self-decay of ferrate(VI): Kinetics and mechanisms. Environ. Sci. Technol. 2014, 48 (9), 5154−5162. (41) Jiang, Y.; Goodwill, J. E.; Tobiason, J. E.; Reckhow, D. A. Effect of different solutes, natural organic matter, and particulate Fe(III) on ferrate(VI) decomposition in aqueous solutions. Environ. Sci. Technol. 2015, 49 (5), 2841−2848. (42) Jiang, Y.; Goodwill, J. E.; Tobiason, J. E.; Reckhow, D. A. Bromide oxidation by ferrate(VI): The formation of active bromine and bromate. Water Res. 2016, 96, 188−197. (43) Shin, J.; von Gunten, U.; Reckhow, D. A.; Allard, S.; Lee, Y. Reactions of ferrate(VI) with iodide and hypoiodous acid: kinetics, pathways, and implications for the fate of iodine during water treatment. Environ. Sci. Technol. 2018, 52 (13), 7458−7467. (44) Neta, P.; Huie, R. E. Free-radical chemistry of sulfite. Environ. Health Perspect. 1985, 64, 209−217. (45) Schwarzenbach, R. P.; Gschwend, P. M.; Imboden, D. M. Environmental Organic Chemistry, 2nd ed.; John Wiley & Sons: New York, 2005. (46) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. 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 1988, 17 (2), 513−886. (47) Neta, P.; Huie, R. E.; Ross, A. B. Rate constants for reactions of inorganic radicals in aqueous solution. J. Phys. Chem. Ref. Data 1988, 17 (3), 1027−1284. (48) Bartlett, P. D.; Cotman, J. D. The kinetics of the decomposition of potassium persulfate in aqueous solutions of methanol. J. Am. Chem. Soc. 1949, 71 (4), 1419−1422. (49) Anipsitakis, G. P.; Dionysiou, D. D. Degradation of organic contaminants in water with sulfate radicals generated by the conjunction of peroxymonosulfate with cobalt. Environ. Sci. Technol. 2003, 37 (20), 4790−4797. (50) Jing, Y.; Chaplin, B. P. Mechanistic study of the validity of using hydroxyl radical probes to characterize electrochemical advanced oxidation processes. Environ. Sci. Technol. 2017, 51 (4), 2355−2365. (51) Chandra, H.; Symons, M. C. R. Hydration of spin-trap cations as a source of hydroxyl adducts. J. Chem. Soc., Chem. Commun. 1986, 16, 1301−1302. (52) Eberson, L. Formation of hydroxyl spin adducts via nucleophilic addition-oxidation to 5,5-Dimethyl-1-pyrroline N-Oxide (DMPO). Acta Chem. Scand. 1999, 53, 584−593. (53) Forrester, A.; Hepburn, S. Spin traps. A cautionary note. J. Chem. Soc. C 1971, 701−703. (54) Leinisch, F.; Ranguelova, K.; DeRose, E. F.; Jiang, J.; Mason, R. P. Evaluation of the Forrester-Hepburn mechanism as an artifact source in ESR spin-trapping. Chem. Res. Toxicol. 2011, 24 (12), 2217−2226. (55) Pang, S. Y.; Jiang, J.; Ma, J. Oxidation of sulfoxides and arsenic(III) in corrosion of nanoscale zero valent iron by oxygen: Evidence against ferryl ions (Fe(IV)) as active intermediates in Fenton reaction. Environ. Sci. Technol. 2011, 45, 307−312. (56) Sharma, V. K.; Luther, G. W., III; Millero, F. J. Mechanisms of oxidation of organosulfur compounds by ferrate(VI). Chemosphere 2011, 82 (8), 1083−1089. (57) Sharma, V. K. Oxidation of inorganic compounds by ferrate(VI) and ferrate(V): One-electron and two-electron transfer steps. Environ. Sci. Technol. 2010, 44 (13), 5148−5152.

(58) Sharma, V. K.; Cabelli, D. Reduction of oxyiron(V) by sulfite and thiosulfate in aqueous solution. J. Phys. Chem. A 2009, 113, 8901−8906. (59) Huang, X.; Deng, Y.; Liu, S.; Song, Y.; Li, N.; Zhou, J. Formation of bromate during ferrate(VI) oxidation of bromide in water. Chemosphere 2016, 155, 528−533.

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DOI: 10.1021/acs.est.8b04990 Environ. Sci. Technol. XXXX, XXX, XXX−XXX