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. 2023 Jul 25;57(47):18898–18908. doi: 10.1021/acs.est.3c00670

Mechanistic Insight for Disinfection Byproduct Formation Potential of Peracetic Acid and Performic Acid in Halide-Containing Water

Junyue Wang 1, Jiale Xu 1, Juhee Kim 1, Ching-Hua Huang 1,*
PMCID: PMC10690735  PMID: 37489812

Abstract

graphic file with name es3c00670_0007.jpg

Peracetic acid (PAA) and performic acid (PFA) are two major peroxyacid (POA) oxidants of growing usage. This study reports the first systematic evaluation of PAA, PFA, and chlorine for their disinfection byproduct (DBP) formation potential in wastewater with or without high halide (i.e., bromide or iodide) concentrations. Compared with chlorine, DBP formation by PAA and PFA was minimal in regular wastewater. However, during 24 h disinfection of saline wastewater, PAA surprisingly produced more brominated and iodinated DBPs than chlorine, while PFA effectively kept all tested DBPs at bay. To understand these phenomena, a kinetic model was developed based on the literature and an additional kinetic investigation of POA decay and DBP (e.g., bromate, iodate, and iodophenol) generation in the POA/halide systems. The results show that PFA not only oxidizes halides 4–5 times faster than PAA to the corresponding HOBr or HOI but also efficiently oxidizes HOI/IO– to IO3–, thereby mitigating iodinated DBP formation. Additionally, PFA’s rapid self-decay and slow release of H2O2 limit the HOBr level over the long-term oxidation in bromide-containing water. For saline water, this paper reveals the DBP formation potential of PAA and identifies PFA as an alternative to minimize DBPs. The new kinetic model is useful to optimize oxidant selection and elucidate involved DBP chemistry.

Keywords: peroxyacids, peracetic acid, performic acid, halogenated disinfection byproducts, saline wastewater, kinetic modeling

Short abstract

This paper reveals the disinfection byproduct formation potential of peracetic acid in saline wastewater and identifies performic acid as an alternative, through evaluation of experimental data with a comprehensive kinetic model.

Introduction

Disinfection byproducts (DBPs) are organic compounds generated from the interaction between oxidants and organic matter or halides, during food, water, and wastewater disinfection.1−4 Thus far, more than 700 toxic, mutagenic, or carcinogenic DBPs have been identified in drinking water,2,5−7 wastewater effluents,8,9 swimming pools,10 and food processing facilities,3,11 leading to chronic public health problems and ecological risks in aquatic ecosystems.1,4,12,13

Free chlorine (HOCl/ClO–), monochloramine (NH2Cl), ozone (O3), and UV irradiation are among the most widely applied disinfectants, and their capacity for DBP formation is discussed briefly as follows. Free chlorine produces DBPs in three major pathways: (1) oxidizing bulk aromatic organic matter to smaller carbonyl14−16 or phenolic compounds,17 with further conversion to carbonaceous DBPs (C-DBPs), including trihalomethanes (THMs), haloacetic acids (HAAs), haloacetaldehydes (HALs), and chlorinated aromatic compounds,18−21 (2) oxidizing amines to nitrogenous DBPs (N-DBPs), such as halonitromethanes (HNMs) and haloacetonitriles (HANs) (could be further hydrolyzed to haloacetamides (HAMs)),22,23 and (3) when applied as breakpoint chlorination in ammoniacal water, producing dichloramine and reactive nitrogen species (e.g., HNO, ONOOH), which in turn generates nitro(so) DBPs.24,25 In contrast, monochloramine is mainly responsible for N-DBP formation, owing to its reactions with amines26,27 and aldehydes28 to produce N-nitrosamines and HANs, respectively. Ozone leads to the formation of nitromethanes29−31 and bromate.32 UV irradiation may produce reactive nitrogen species (e.g., •NO, •NO2) in the presence of nitrite, nitrate, or chloramines, hence enhancing the formation of nitro(so) DBPs.33−35

Recently, peroxyacids (POAs) (i.e., peracetic acid (CH3C(O)OOH, PAA) and performic acid (HC(O)OOH, PFA)) have been proposed as an alternative to the aforementioned traditional disinfectants.9,11,36−38 PAA has a pKa of 8.2, and its bacterial inactivation capacity is comparable to that of chlorine.39,40 PFA has a pKa of 7.3;41 its bacterial inactivation efficacy usually outperforms PAA, despite its faster self-decay.38,42 POAs do not contain halogen elements, and the coexistent hydrogen peroxide (H2O2) in the POA solution can reduce halogenating agents in the water matrix.43 Furthermore, PAA’s sluggish reactivity toward amine compounds limits the potential of N-DBP formation.40,44 Nevertheless, in halide-containing water, PAA could oxidize bromide (Br–) and iodide (I–) to free bromine (HOBr/BrO–) and free iodine (HOI/IO–), respectively, which in turn produces Br- and I-DBPs. In some previous studies, PAA indeed generated more dibromoacetic acid45 and iodinated THMs (I-THMs)36 than chlorine in halide-containing waters; hence, the capacity of PAA on DBP control deserves further investigation. Moreover, a literature search indicates that the DBP formation potential of PFA has never been studied.

As mentioned above, the major DBP formation pathway of PAA and PFA is their oxidation of halides, particularly Br–/I–, into halogenating agents (i.e., HOBr and HOI). Moreover, POA may further oxidize HOBr and HOI, which in turn shifts the byproducts from organic Br- and I-DBPs to bromate (BrO3–, toxic) and iodate (IO3–, nontoxic), respectively.46−48 Thus, to understand the DBP production by PAA and PFA, their reactivity with halides (particularly Br–, I–), HOBr/BrO–, and HOI/IO– (rate constants denoted as kBr–, kI–, kBrO–, and kIO– in Table 1) is of great importance. However, thus far, only the reaction rate constants between PAA and halides have been reported, while the reactions of PAA with HOBr/BrO– and HOI/IO– and the reactions involving PFA have not been studied (Table 1).43 These rate constants for other typical oxidants, including free chlorine,47−49 monochloramine,50,51 ferrate(VI),52,53 permanganate,54 and ozone32,45 have all been well-studied (Table 1).

Table 1. Apparent Second-Order Rate Constants between Oxidants and Br–, I–, BrO–, and IO– (pH 7.0–7.2).

  kBr– kBrO– ref kI– kIO– ref
HOCl/ClO– 1550 slow (32) 4.3 × 108 21 (67)
NH2Cl 1.4 × 10–2 a 2.7 × 105 b (50) 2.4 × 103 <2 × 10–3 (67)
HOBr/BrO–   slow (46) 5.0 × 109 752 (47)
O3 160–258 1.58 (32) 2.0 × 109 3.7 × 104 (67)
KMnO4 unknown unknown   7.0 6.9 (67)
ferrate(VI) 0.1–10.9 unknown (32) 2.0 × 104 2.0 × 103 (67)
ClO2 <0.05 unknown (46) 1900c unknown (67)
PMS 0.7 slow (32) 1.1 × 103 7.9 × 102 (67)
PFA 1.03 slow (43), this study 2.47 × 103 120 this study
PAA 0.22 slow (43), this study 3.89 × 102 16 this study
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a

This reaction produces NH2Br and is reported at pH 7.5;48 other reactions for kBr– all produce HOBr.

b

This reaction produces NHBrCl and is reported at pH 7.5;48 other reactions for kBrO– all produce BrO2–.

c

The reaction produces I•;49 other reactions for kI– all produce HOI.

Therefore, the objectives of this study were to (1) comprehensively compare the DBP formation by POAs [i.e., PAA and PFA in this study] and chlorine in regular and halide-containing wastewater; (2) determine the rate constants of POAs with halides, HOBr/BrO–, HOI/IO–, and other water matrix constituents [e.g., effluent organic matter (EfOM), NH4+]; and (3) establish a kinetic model that can predict the levels of HOBr and HOI during POA oxidation of halide-containing water.

Materials and Methods

Chemicals and Reagents

PAA solution, H2O2 solution, and free chlorine (NaOCl) solution were purchased from Sigma-Aldrich (St. Louis, MO). PFA (w/60% H2O2, molar ratio) was freshly synthesized following the method described in Text S1 in the Supporting Information.38 It should be noted that additional H2O2 was added to PAA working solutions in DBP formation experiments to reach a similar coexistent H2O2 concentration as PFA for a fair comparison. The additional H2O2 did not affect PAA concentration40,55,56 and was confirmed by experiments. Free bromine was freshly synthesized according to a method by Guo et al. by mixing sodium hypochlorite and bromide at a molar ratio of 1:1.05 and adjusted the pH to 11.0.57

Standard chemicals for DBPs, including THM4 (trichloromethane (TCM), dichlorobromomethane (DCBM), dibromochloromethane (DBCM), tribromomethane (TBM)), three HANs (dichloroacetonitrile (DCAN), bromochloroacetonitrile (BCAN), dibromoacetonitrile (DBAN), trichloroacetonitrile (TCAN)), one HNM (trichloronitromethane (TCNM)), two haloketones (1,1-dichloro-2-propanone (1,1-DCP) and 1,1,1-trichloro-2-propanone (1,1,1-TCP)), and three I-THMs (dichloroiodomethane (DCIM), diiodochloromethane (DICM), and triiodomethane (TIM)) were obtained from Sigma-Aldrich, Agilent, and Toronto Research Chemicals. Other chemicals used in this study are listed in Text S1.

DBP Formation Experiments

20 mL of regular or saline wastewater (Table S1, Text S1) was buffered with phosphate to the desired pH and spiked with 100 μM of PAA, PFA, or NaOCl in amber glass vials. It is worth noting that these three disinfectants provide comparable bacterial inactivation at the same dosage as demonstrated by our recent study.38 The reactors were protected from ambient light, stirred vigorously, and sealed with caps with a headspace less than 5 mL. The reaction was quenched by excess sodium thiosulfate (Na2S2O3) after a defined reaction time.

Batch Kinetic Experiments

The reactivities of PAA and PFA with halides, HOBr/BrO–, and HOI/IO– was investigated. The reactions between PAA and halides have been studied previously.43 Thus, in this study, we investigated (i) PFA decay with and without excess Cl–/Br–, (ii) iodophenol formation in the PFA/I–/phenol system, (iii) BrO3– formation in POA/HOBr systems, and (iv) IO3– formation in POA/I– systems. These experiments yielded results to simulate the rate constants of POAs with Cl–/Br–, I–, HOBr/BrO–, and HOI/IO–, respectively. The employed experimental conditions are summarized in Table S3.

The reaction solutions were phosphate-buffered, magnetically stirred, and measured for pH right after oxidant addition, as well as throughout the experiments. Periodically, 1 mL samples were taken. The samples in group (ii) was quenched by Na2S2O3, while samples for experiments (iii) and (iv) were quenched by excess phenol. It is worth noting that although phenol has been shown to be not highly effective for elimination of PAA and PFA, we found the other quenching agents, including Fe(II), cysteine, ascorbic acid, and methionine, all affected BrO3– and/or IO3– measurement by reducing them to Br–/I– (data not shown). Therefore, we chose to use phenol (20 mM) to quench HOBr and HOI. Even if there were residual PAA/PFA in the samples, the measurement of BrO3– and IO3– could not be affected due to the rapid scavenging of their precursors (i.e., HOBr, HOI) by phenol.

Kinetic Model Simulation

Kinetic modeling was performed to simulate the unknown rate constants of POAs with halides, HOBr/BrO–, and HOI/IO– (Table 2, see discussion later), using the “FIT:2:3:FITDATA.TXT” command in Kintecus 4.55.31. Then, HOBr and HOI levels in POA/halide systems were predicted by the kinetic model with the reactions listed in Table 2.

Table 2. Principal Reactions in the POA Oxidation of Halide-Containing Ammonia-Free Water.

    apparent rate constanta
no. reaction pH 5.5 pH 7.1 pH 7.8 ref
R1 PFA → products (3.3 ± 0.3) × 10–2 (4.7 ± 0.3) × 10–2 (7.4 ± 0.3) × 10–2 Figure 2, Figure S2
  (a) PFA + H2O → H2O2 + FAb (3.3 ± 0.3) × 10–2 (1.8 ± 0.4) × 10–2 (1.6 ± 0.2) × 10–2 Figure 2
  (b) PFA → CO2 + H2O (0.3 ± 0.0) × 10–2 (2.9 ± 0.6) × 10–2 (5.8 ± 0.8) × 10–2 Figure 2
  (c) 2 PFA → 2 FA + O2 negligible when [PFA]0 < 200 μM Figure S2
R2 H2O2 + FA → PFA + H2O negligible when [H2O2]0 = [HCOO–]0 < 1 mM not shown
R3 PFA + Cl– → HOCl/ClO– + FA   negligible   Figure S5
R4 PFA + Br– → HOBr/BrO– + FA 1.12 ± 0.24 1.03 ± 0.11 0.35 ± 0.09 Figure 3
R5 PFA + I– → HOI/IO– + FA (3.18 ± 0.19) × 103 (2.48 ± 0.35) × 103 (1.34 ± 0.09) × 103 Figure 4
R6 PFA + HOBr/BrO– → BrO2– + FA   negligible   Figure 5a
R7 PFA + HOI/IO– → IO2– + FA   1.20 × 102   Figure 5b
R8 PFA + IO2– → IO3– + FA   1.00 × 103   assumed
R9 PFA + H2O2 → H2O + FA + O2 negligible negligible (3.16 ± 0.28) × 10–1 Figure 2
R10 PFA + EfOM → products   slow   not shown
R11 PAA + Cl– → HOCl/ClO– + AA   negligible   (43)
R12 PAA + Br– → HOBr/BrO– + AAc 2.40 × 10–1 2.20 × 10–1 1.70 × 10–1 (43), Text S3
R13 PAA + I– → HOI/IO– + AA 4.19 × 102 3.89 × 102 3.00 × 102 (43), Text S3
R14 PAA + HOBr/BrO– → BrO2– + AA   negligible   Figure 5a
R15 PAA + HOI/IO– → IO2– + AA   1.60 × 101   Figure 5b
R16 PAA + IO2– → IO3– + AA   1.00 × 103   assumed
R17 PAA + EfOM → products   slow   not shown
R18 HOBr/BrO– + H2O2 → Br– 6.03 × 102 2.36 × 104 1.09 × 105 (43), Text S3
R19 HOI/IO– + H2O2 → I– 1.59 × 102 6.32 × 103 3.16 × 104 (52), Text S3
R20 HOBr/BrO– + I– → HOI/IO– + Br– 5.00 × 109 4.99 × 109 4.55 × 109 (47), Text S3
R21 HOI/IO– + HOBr/BrO– → IO2– + Br–   7.40 × 102   (47)
R22 HOBr/BrO– + IO2– → IO3– + Br–   5.00 × 103   assumed
R23 HOBr/BrO– + EfOM → products   3.00 × 101   (57)
R24 HOI/IO– + EfOM → products   3.00 × 101   assumed
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a

The units for apparent rate constants: R1 in min–1, R2–R24 in M–1 s–1.

b

FA = formic acid/formate, AA = acetic acid/acetate

c

The self-decay of PAA,51 HOBr/BrO–,46 and HOI49 are all negligible at the concentrations employed in this study ([PAA] < 200 μM, [HOBr/BrO–] < 400 μM, [HOI/IO–] < 0.3 μM).

Analytical Methods

The concentrations of PAA and PFA were measured following the KI-DPD method using a UV–visible spectrophotometer detailed in our previous study.38 The total concentrations of PFA and H2O2 were determined by a horseradish peroxidase-2,2′-azino-bis(3-ethylbenzothiazoline-6-sulfonic) acid (HRP-ABTS) method which was validated in our previous study,58 and the H2O2 concentration was calculated by subtracting the PFA concentration from the total peroxides.

Iodophenols were analyzed using high-performance liquid chromatography equipped with a diode-array detector (HPLC-DAD) (Text S2). Anions (i.e., Cl–, Br–, I–, BrO3–, and IO3–) were measured by ion chromatography (Text S2). The DBPs were measured by a gas chromatograph equipped with an electron capture detector (GC-ECD) after liquid–liquid extraction of DBPs from water by 2 mL of methyl tert-butyl ether (MtBE, >99.8% purity), with 1,2-dibromopropane as the internal standard, following the protocol reported by Xu et al.35

Results and Discussion

DBP Formation during Disinfection by PAA, PFA, and Chlorine

DBP formation by the addition of PAA, PFA, and free chlorine was, for the first time, compared in the same water matrix in this study. As expected, chlorine induced considerable DBP formation in regular wastewater ([Cl–] = 5 mM, [Br–] = 9 μM, [I–] = nondetectable), among which THM4 and DCAN accounted for the major part (Figure S1). Although HAAs and HAMs were not measured in our study, their formation could be expected due to their correlation with THMs59 and HANs,22 respectively. The DBP formation by dosing free chlorine (NaOCl) should be ascribed to both chlorination and chloramination due to the presence of NH4+ (3.49 mg/L as N) in the wastewater (Table S1). Herein, we will use the word “chlorine” to refer to the overall effect of free and combined chlorines in the matrices. In contrast to chlorine, PAA and PFA produced little halogenated DBPs (Figure S1), suggesting that DBP formation through POA oxidation of halides was not significant at regular halide levels in typical domestic wastewater.

Subsequently, we dosed additional halides to the wastewater to create a saline matrix ([Cl–] = 400 mM, [Br–] = 500 μM, [I–] = 0.3 μM),43,60,61 which demonstrated the worst-case scenarios for Br-/I-DBP formation. There are several possible scenarios of saline wastewaters that could be subject to oxidation and/or disinfection: (1) combination of shale gas wastewater, produced by hydraulic fracturing and horizontal drilling for natural gas production, into wastewater treatment plants (WWTPs),62 (2) chemical oxidation treatment of ballast wastewater in ocean-going vessels,43 (3) application of seawater to flush toilets in fresh-water-stressed coastal cities,60,61 (4) decentralized treatment of urine,63−65 and (5) reclamation of high-salinity streams along with wastewater.66 The evaluation of DBP formation in the saline water matrix in this study could assist in reasonable oxidant selection for the aforementioned scenarios. While the real saline wastewaters may differ to some degree from the saline sample in this study (e.g., in organic matter compositions), our main objective to compare the DBP formation trends of POAs and chlorine should not be affected.

As shown in Figure 1a,c, during 30 min of oxidation of the synthetic saline wastewater, chlorine resulted in the highest total DBP concentration, where THM4 were the major contributors. Although PAA generated much less DBPs, it produced more TIM than chlorine at both pHs (7.1 and 7.8). The shift from THM4 to I-THMs during PAA oxidation indicated a higher HOI exposure level, which is probably associated with other I-DBP formation. Although the overall toxicity was not quantitatively calculated due to the large unknown portion of halogenated compounds, considering the high toxicity of I-DBPs,67 PAA probably resulted in a higher overall toxicity than chlorine in the saline wastewater, which challenges the intention to replace chlorine by PAA for DBP control in saline matrix. On the contrary, PFA controlled most DBPs effectively, only with modest TBM formation.48,50

Figure 1.

Figure 1

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Production of 14 DBPs during disinfection of saline wastewater. Experimental conditions: [oxidant]0 = 100 μM, [phosphate buffer] = 10 mM, temperature 23 ± 2 °C, PAA and PFA both contained 60 μM H2O2. Saline wastewater parameters: [Cl–] = 400 mM, [Br–] = 500 μM, [I–] = 0.3 μM, [NH4+] = 3.49 mg/L as N, COD = 38.48 mg/L.

As for the long-term (24 h) disinfection, relevant to ballast wastewater oxidation in ocean-going vessels, we surprisingly found that PAA led to a higher total DBP concentration than that of chlorine (Figure 1b,d). Although PAA slightly reduced the formation of chlorinated DBPs (Cl-DBPs), it enhanced the formation of most Br- and I-DBPs, including TBM, BCAN, and I-THMs. In particular, PAA triggered TIM formation at 13.15 and 10.81 μg/L at pH 7.1 and 7.8, respectively, whereas chlorine only resulted in 1.69 and 2.65 μg/L. Notably, 24 h PAA oxidation gave rise to an ∼25% iodine incorporation into I-THMs, which is higher than that reported for chloramination (<20%) at similar iodide levels (24 h, [NH2Cl] = 15 μM).47 Interestingly, contrary to PAA, PFA controlled all tested DBPs successfully during long-term disinfection.

To sum up, these results suggest that replacing chlorine by either PAA or PFA could alleviate DBP problems in regular wastewater effluents; however, with elevated halide concentrations, PFA remarkably outperformed PAA and chlorine for DBP control. Particularly, PAA induced an unwanted shift from Cl-DBPs to Br-/I-DBPs in saline wastewater, which can potentially make the disinfected effluents even more toxic than the chlorinated ones. It should be noted that, due to the presence of ammonia (3.49 mg/L as N) in the saline wastewater of this study, the generated HOCl/HOBr could be rapidly converted to corresponding halamines that could be major species contributing to Cl- and Br-DBP formation, while HOI was always the major iodinating agent due to the lack of reactivity with ammonia.48,50 Overall, the high ammonia concentration used in this study may inhibit the formation of carbonaceous Cl- and Br-DBPs (e.g., THM4) for POAs as well as free chlorine. The DBP formation potential of free chlorine could be particularly mitigated by ammonia due to direct and rapid consumption of the free chlorine oxidant. In fully nitrified wastewater (low ammonia levels), the advantages of POAs for DBP control (over free chlorine) may be more significant. However, we still expect PFA to outperform PAA for DBP control regardless of ammonia levels (see later discussion).

Reactivity of POAs with Water Matrices

To understand DBP formation by PAA and PFA, we first investigated their reactivity with the major components in the saline wastewater matrix, i.e., Br–, I–, NH4+, and EfOM.

Reactivity of PAA with Water Matrices

The reactivity between PAA and halides has been reported by Shah et al.43 and hence was not repeated in this study (Table 2, R11–R13). Additionally, we found the 30 min PAA decay in the regular wastewater and in a NH4+ solution (10 mM) were both less than 5% at pH 7.1 (data not shown); thus, the PAA self-decay and PAA consumption by NH4+ and EfOM could be neglected under the experimental conditions.

PFA Self-Decay

Unlike PAA, PFA decomposition is relatively fast and has to be studied. First, we found the change of initial PFA concentration from 50 to 200 μM did not affect PFA self-decay (Figure S2), and hence, the PFA self-decay was fitted into a pseudo-first-order kinetic model (eqs 1 and 2)

graphic file with name es3c00670_m001.jpg 1
graphic file with name es3c00670_m002.jpg 2

where kdecay is the first-order rate constant for PFA self-decay (in min–1), [PFA] is the concentration of PFA (in M), and t is the reaction time (in min). kdecay increased from (3.3 ± 0.3) × 10–2 to (7.4 ± 0.3) × 10–2 min–1 as the pH increased from 5.5 to 7.8. As the bimolecular PFA decay has been ruled out (when [PFA] < 200 μM), PFA decay could be ascribed to R1a, R1b, and/or R9 (Table 2).41 First, the possible reaction between PFA and H2O2 (R9) was examined by adding additional H2O2 to the PFA solution, and additional H2O2 only affected PFA self-decay at pH 7.8 (Figure 2 and Table 2 R9). Then, relative contributions from R1a and R1b to PFA self-decay were differentiated by measuring the change in the H2O2 concentration during PFA self-decay. We found that the H2O2 pathway (Table 2, R1a) dominated PFA self-decay at pH 5.5 (94.7 ± 1.2%), while the fraction of the bicarbonate pathway (Table 2, R1b) rose remarkably with the increase in pH (i.e., from 5.3 ± 1.2% at pH 5.5 to 79.2 ± 3.8% at pH 7.8) (Figure 2).41 The corresponding rate constants for the two self-decay pathways (R1a,b) at three tested pHs were calculated as shown in Table 2. In contrast, the PAA self-decay was negligible within 30 min.

Figure 2.

Figure 2

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PFA decay and the percentage of conversion to H2O2 at pH 5.5 (a), 7.1 (b), and 7.8 (c). Experimental conditions: [PFA]0 = 100 μM, [phosphate buffer] = 10 mM, temperature 23 ± 2 °C. Error bars represent standard deviations between parallel experiments.

Reactivity of PFA with Water Matrices

The reactivity of inorganic constituents with PFA was examined by testing their effects on the PFA decay. First of all, the presence of Br– significantly accelerated PFA self-decay (Table 2, R4) (Figure 3); for those experiments an excess amount of Br– was applied to achieve a pseudo-first-order reaction. We also applied additional H2O2 to reduce the produced HOBr back to Br–, which kept a constant Br– concentration and avoided the interference of HOBr during the PFA measurement by DPD (Table 2, R18). Then, the apparent second-order rate constants were calculated by a pseudo-first-order model (eqs 3 and 4).

graphic file with name es3c00670_m003.jpg 3
graphic file with name es3c00670_m004.jpg 4

where kdecay is the PFA self-decay rate constant (Table 2, R1) (in s–1), kapp,PFA,Br– is the apparent second-order rate constant for the reaction between PFA and Br– (in M–1 s–1),[PFA] and [Br–] are the concentrations of PFA and Br–, respectively (in M), kapp,FA,H2O2 is the apparent second-order rate constant for the reaction between PFA and H2O2 (Table 2, R9) (in M–1 s–1) (ote that this reaction is negligible at pH 5.5 and 7.1, but non-negligible at pH 7.8), [H2O2] is the concentration of additionally spiked H2O2 (i.e., 1.0 mM), and t is the reaction time (in s). As a result, the apparent second-order rate constants between PFA and Br– (kapp,PFA,Br–, Table 2 R4) were determined to be 0.35–1.12 M–1 s–1 in the pH range 5.5–7.8 and are approximately linear to the fraction of protonated PFA (pKa = 7.341) at each pH (Figure S3a and Text S3), suggesting that only the protonated PFA was the reactive form for oxidation of Br–.

Figure 3.

Figure 3

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Impact of bromide (0–4.0 mM) on PFA decay. Experimental conditions: [phosphate buffer] = 10 mM, [PFA]0 = 100 μM, [H2O2 (additional)]0 = 1.0 mM, temperature 23 ± 2 °C. Error bars represent standard deviations between parallel experiments, and the solid lines represent the linear regression.

The above experimental approach for Br– could not be applied to study the reaction of PFA with I–, because excess I– consumed PFA extremely rapidly. Alternatively, we used iodophenol formation to quantify HOI formation and indicate I– loss,52 where PFA was dosed in excess to I– to enable the first-order model, and phenol was dosed in excess as well to react with all the HOI/IO– and minimize IO3– formation. The iodophenol formation and I– loss (Figure 4) can be described by eqs 5 and 6.

graphic file with name es3c00670_m005.jpg 5
graphic file with name es3c00670_m006.jpg 6

where [iodophenols] is the total concentration of 2-, and 4-iodophenol (in M), kapp,PFA I– is the apparent second-order rate constant for the reaction between PFA and I– (in M–1 s–1), [PFA] and [I–] are the concentrations of PFA and I–, respectively (in M), and t is the reaction time (in s). As a result, the rate constants between PFA and I– (kapp,PFA,I–, Table 2 R4) were determined to be (1.34–3.18) × 103 M–1 s–1 in the pH range 5.5–7.8 and are linearly related to PFA protonation (Figure S3b and Text S3).

Figure 4.

Figure 4

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Production of iodophenols from iodide oxidation by PFA (a–c) and the corresponding iodide loss (d). Experimental conditions: [phosphate buffer] = 10 mM, [iodide]0 = 10 μM, [PFA]0 = 50 μM, [phenol]0 = 4.0 mM, temperature 23 ± 2 °C. Error bars represent standard deviationdbetween parallel experiments, and the solid lines in (d) represent the linear regression modeling.

Finally, we found Cl–, NH4+, EfOM, and phenols all had very limited impacts on PFA self-decay (Figures S4–S7). Thus, R3 and R10 could be neglected in the kinetic model for the PFA oxidation of saline wastewater.

Interestingly, the reaction of PFA with halides was 4–5 times faster than that of PAA at each pH. As the initial coexistent H2O2 concentrations in PAA and PFA solutions were similar for the DBP formation experiments, PFA would be expected to result in higher HOBr/HOI levels and more Br- and I-DBP formation. However, PFA oxidation produced much fewer DBPs than PAA in halide-containing water, prompting us to further investigate and develop a comprehensive kinetic model to elucidate this phenomenon.

Understanding the Distinct DBP Formation Behaviors of PAA and PFA

To understand why PFA controlled DBP formation better in halide-containing water despite its higher reactivity with halides, we proposed three hypotheses and examined them one by one.

Reducing Effects of Formate and Acetate

PAA and PFA solutions contain considerable amounts of formate and acetate, respectively. Particularly, formate is a weak reducing agent due to the formaldehyde moiety. Therefore, it was possible that coexistent formate reduced HOBr/HOI and controlled DBP formation. However, this possibility is ruled out by the minimal impacts of formate and acetate on HOBr/HOI decay (Figure S8).

Reaction between HOBr and HOI

It has been well-documented that HOBr could oxidize HOI/IO– to produce BrO3– and IO3– (Table 2, R21 and R22).48 Thus, if large amounts of HOBr/BrO– and HOI/IO– are produced by PFA, they can be canceled out into nonreactive forms, which hence controls DBP formation. To test this hypothesis, we dosed the oxidants into synthetic wastewater containing only 0.3 μM I– [without Br–]. Nonetheless, PAA still produced more TIM than did PFA and chlorine (Figure S9).

Production of BrO3– and IO3–

Finally, as mentioned above, the byproducts will shift from organic Br- and I-DBPs to BrO3– and IO3–, respectively, if POAs further oxidize HOBr/BrO– and HOI/IO–. The oxidation of HOBr/BrO– and HOI/IO– typically proceeds through two-electron transfer and produces BrO2– (Table 2, R6 and R14) and IO2– (Table 2, R7 and R15), which will be rapidly oxidized to BrO3– and IO3–, respectively (Table 2, R8 and R16, k assumed to be ∼103 M–1 s–1).48,50,52 Thus, to further understand DBP formation by PAA and PFA, we simulated the rate constants for R6 and R14 in the POA/HOBr system and R7 and R15 in the POA/I– system at pH 7.1. HOBr is easy to synthesize and relatively stable.57 Therefore, we mixed preprepared HOBr and POAs at high concentrations to maximize the BrO3– formation. However, less than 2.5% conversion from HOBr to BrO3– was observed during the 2 h interaction (Figure 5a). It is noteworthy that in the above experiment, we already tried intentionally to maximize BrO3– formation through spiking of 200 μM POAs and 400 μM HOBr. In fact, the HOBr concentration did not exceed 0.2 μM during POA oxidation of saline water, which was verified by the DPD measurement. Therefore, the reactions between HOBr and POAs (Table 2, R6 and R14) should be negligible during POA oxidation of bromide-containing water.

Figure 5.

Figure 5

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Production of bromate in POA/HOBr systems (a) and production of iodate from POA/iodide systems (b). Experimental conditions: [phosphate buffer] = 10 mM, pH = 7.1, temperature 23 ± 2 °C. Error bars represent standard deviations between parallel experiments.

Contrary to HOBr, HOI could not be presynthesized due to its self-decay (note that this self-decay is negligible when [HOI] < 0.3 μM49 and was thus not incorporated into the kinetic model). Thus, we mixed excess POAs with I– to test IO3– formation (Figure 5b). Significant IO3– formation was observed for both POAs. By kinetic model simulation, we found that PFA and PAA oxidized HOI/IO– at 1.20 × 102 M–1 s–1 (R7, RMSD = 0.118) and 1.60 × 101 M–1 s–1 (R15, RMSD = 0.133) at pH 7.1, respectively.

The different reactivities of PAA and PFA toward HOI/IO– explain why PFA controls I-DBP formation better even if it oxidizes I– faster. PFA oxidizes HOI/IO– into the nonreactive IO3– rapidly and hence controls I-DBP formation, while PAA exhibits rather slow oxidation of HOI and thus induces accumulation of HOI and production of I-THM. Even chlorine produces fewer I-DBPs than PAA, which is probably due to the faster oxidation of HOI/IO– by chlorine (Table 1) and competition of susceptible precursors with HOI.

Modeling the Oxidant Concentrations

So far, the inhibition of I-DBP formation by PFA has been attributed to its fast oxidation of HOI/IO– into IO3–. However, as both PAA and PFA could hardly oxidize HOBr/BrO–, PFA is expected to produce more Br-DBPs than PAA due to its faster oxidation of Br– to HOBr, which is inconsistent with the DBP formation results during 24 h disinfection (Figure 1b,d). Therefore, we built up a comprehensive kinetic model with all of the relevant rate constants retrieved from previous studies or simulated in this study (Table 2).

To systematically compare PAA and PFA oxidation, we simulated oxidant levels with the kinetic model in a representative halide-containing water matrix ([EfOM] = 1.0 mM as C, [Br–]0 = 0.5 mM, [I–]0 = 0.2 μM, no NH4+) with two simplifications justified as below. First, the presence of NH4+ can significantly complicate the model by incorporating the reactions of NH2Cl, NH2Br, NHBrCl, and NHCl2. The kinetic model can be exploited with the addition of halamine reactions,48 if needed; however, we believe the involvement of NH4+ will not change the major conclusion on PAA/PFA comparison. Cl– is not included because it could not change the modeling results due to its negligible reactions with POAs. Moreover, we would clarify that HOBr and HOI are not the only halogenating agents in the system because their speciation results in exotic electrophile formation (i.e., BrCl, BrOCl, Br2O, H2OI+, and ICl).68 Despite their low concentrations (at least 4 orders of magnitude lower than those of HOBr or HOI at pH 7.1–7.8), these species have non-negligible contributions to DBP formation due to their high reactivity. However, as these species always reach rapid equilibrium with their precursors (HOBr or HOI),68 the HOBr/HOI levels would not be affected and the overall DBP formation should still be proportional to HOBr/HOI levels. Overall, we believe the kinetic simulation in this representative matrix could provide valuable insight into PAA/PFA comparison, despite these acceptable simplifications.

Our model shows that PFA produces more HOBr and HOI than PAA in the short term (Figure 6a,b). However, the HOBr and HOI levels in the PFA/halide system reach the peak points at ∼0.88 and 0.18 h, respectively; then their concentrations decrease and are always lower than those with PAA in the long term. These simulation results concur with the DBP formation experiments, where the overall higher HOBr/HOI concentrations during PAA oxidation gave rise to its higher Br-/I-DBP formation in 24 h, whereas the rapid HOBr generation by PFA yielded more Br-DBPs in the first 30 min (Figure 1).

Figure 6.

Figure 6

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Kinetic simulation for PAA and PFA oxidation of ammonia-free saline water. The cumulative exposure to HOBr and HOI was calculated by integrating [HOBr] or [HOI] over time. Simulation time = 20 h.

Overall, the gradually decreasing HOBr/HOI levels during PFA oxidation remarkably control the DBP formation and could be attributed to three mechanisms. First, as indicated by the model, PFA is quickly consumed by the self-decay and faster reactions with halides, while the PAA decay is much slower (Figure 6c). The elimination of PFA cuts off the source for HOBr and HOI, while the relatively stable PAA maintains the oxidation of halides and the formation of HOBr/HOI. Second, H2O2 is the major reductant for HOBr/HOI in the POA disinfection processes and benefits the control of Br- and I-DBPs. Although we intentionally controlled the initial H2O2 in PAA and PFA at the same concentration (for both DBP experiments and kinetic simulation), the model showed that the H2O2 in the PFA system could not be totally consumed (Figure 6d), while the H2O2 in PAA is depleted within 3 h. Briefly, in the PFA system, the depletion of the HOBr/HOI source (i.e., PFA) inhibits H2O2 consumption by suppressing R18 and R19 (Table 2), and additional H2O2 could be released during PFA self-decay (Table 2, R1a). In contrast, the relatively stable PAA maintains halide oxidation and the catalytic depletion of H2O2 (Table 2, R18 and R19). The higher H2O2 level during PFA oxidation definitely contributes to the HOBr/HOI reduction and DBP control. Third, for I-DBPs, PFA oxidizes HOI/IO– into IO3– faster than PAA and hence limits the iodination of EfOM.

To sum up, the kinetic model shows that PFA’s faster decay, slow release of H2O2, and faster oxidation of HOI/IO– all limit the HOBr/HOI levels and control Br- and I-DBPs formation. In contrast, the stability of PAA results in continuous halide oxidation and catalytic depletion of H2O2, leading to higher HOBr/HOI levels and intensive DBP formation.

Environmental Significance

This study, for the first time, systematically compares the DBP formation potential of PAA, PFA, and chlorine. Both PAA and PFA showed a satisfactory DBP control capacity in regular wastewater. However, only PFA controlled DBP formation in saline wastewater, while PAA produced more Br- and I-DBPs than chlorine after 24 h oxidation.

Additionally, we developed a kinetic model for POA/halide systems. The reactivity of PAA and PFA toward halides can be compared with those of other oxidants (Table 1). PAA joins monochloramine as a significant contributor to I-DBPs, while PFA exhibits several properties beneficial for DBP control, including (i) low reactivity with HOBr/BrO–, (ii) self-decay and slow release of H2O2, and (iii) efficient oxidation of HOI to IO3–. As BrO3– formation has been reported as a problem for both ferrate(VI)53 and ozone,32 PFA seems to be the only disinfectant that produces neither BrO3– or I-DBPs in saline effluents. So far, in-depth studies have shown PFA to be more effective than PAA for pathogen inactivation38 and DBP control (this study). Hence, PFA is a strong candidate for disinfection of halide-containing water; however, on-site generation and immediate application of PFA are required due to its fast self-decay.

Acknowledgments

This work was supported by the National Science Foundation Grants CHE-1609361 and CHE-2108701. Any opinions, findings, and conclusions or recommendations expressed in this material are those of the authors and do not necessarily reflect the views of the National Science Foundation. The authors are grateful for instrumental support from Dr. Chaoyang Huang.

Supporting Information Available

The Supporting Information is available free of charge at https://pubs.acs.org/doi/10.1021/acs.est.3c00670.

  • Chemicals, analytical methods, calculation of rate constants, water quality parameters, and PFA decay experiements (PDF)

Author Present Address

† Department of Civil, Construction and Environmental Engineering, North Dakota State University, Fargo, North Dakota 58105, United States

The authors declare no competing financial interest.

Supplementary Material

es3c00670_si_001.pdf (715.4KB, pdf)

References

  1. Li X. F.; Mitch W. A. Drinking Water Disinfection Byproducts (DBPs) and Human Health Effects: Multidisciplinary Challenges and Opportunities. Environ. Sci. Technol. 2018, 52 (4), 1681–1689. 10.1021/acs.est.7b05440. [DOI] [PubMed] [Google Scholar]
  2. Krasner S. W.; Mitch W. A.; Westerhoff P.; Dotson A. Formation and control of emerging C- and N-DBPs in drinking water. Journal - American Water Works Association 2012, 104 (11), E582–E595. 10.5942/jawwa.2012.104.0148. [DOI] [Google Scholar]
  3. Simpson A. M. A.; Mitch W. A. Chlorine and ozone disinfection and disinfection byproducts in postharvest food processing facilities: A review. Critical Reviews in Environmental Science and Technology 2022, 52 (11), 1825–1867. 10.1080/10643389.2020.1862562. [DOI] [Google Scholar]
  4. Allen J. M.; Plewa M. J.; Wagner E. D.; Wei X.; Bokenkamp K.; Hur K.; Jia A.; Liberatore H. K.; Lee C. T.; Shirkhani R.; Krasner S. W.; Richardson S. D. Drivers of Disinfection Byproduct Cytotoxicity in U.S. Drinking Water: Should Other DBPs Be Considered for Regulation?. Environ. Sci. Technol. 2022, 56 (1), 392–402. 10.1021/acs.est.1c07998. [DOI] [PubMed] [Google Scholar]
  5. Krasner S. W.; Mitch W. A.; McCurry D. L.; Hanigan D.; Westerhoff P. Formation, precursors, control, and occurrence of nitrosamines in drinking water: a review. Water Res. 2013, 47 (13), 4433–50. 10.1016/j.watres.2013.04.050. [DOI] [PubMed] [Google Scholar]
  6. Dong H.; Cuthbertson A. A.; Richardson S. D. Effect-Directed Analysis (EDA): A Promising Tool for Nontarget Identification of Unknown Disinfection Byproducts in Drinking Water. Environ. Sci. Technol. 2020, 54 (3), 1290–1292. 10.1021/acs.est.0c00014. [DOI] [PubMed] [Google Scholar]
  7. Richardson S. D. Tackling unknown disinfection by-products: Lessons learned. Journal of Hazardous Materials Letters 2021, 2, 100041. 10.1016/j.hazl.2021.100041. [DOI] [Google Scholar]
  8. McCurry D. L.; Bear S. E.; Bae J.; Sedlak D. L.; McCarty P. L.; Mitch W. A. Superior Removal of Disinfection Byproduct Precursors and Pharmaceuticals from Wastewater in a Staged Anaerobic Fluidized Membrane Bioreactor Compared to Activated Sludge. Environmental Science & Technology Letters 2014, 1 (11), 459–464. 10.1021/ez500279a. [DOI] [Google Scholar]
  9. Dominguez Henao L.; Turolla A.; Antonelli M. Disinfection by-products formation and ecotoxicological effects of effluents treated with peracetic acid: A review. Chemosphere 2018, 213, 25–40. 10.1016/j.chemosphere.2018.09.005. [DOI] [PubMed] [Google Scholar]
  10. Daiber E. J.; DeMarini D. M.; Ravuri S. A.; Liberatore H. K.; Cuthbertson A. A.; Thompson-Klemish A.; Byer J. D.; Schmid J. E.; Afifi M. Z.; Blatchley E. R. 3rd; Richardson S. D. Progressive Increase in Disinfection Byproducts and Mutagenicity from Source to Tap to Swimming Pool and Spa Water: Impact of Human Inputs. Environ. Sci. Technol. 2016, 50 (13), 6652–62. 10.1021/acs.est.6b00808. [DOI] [PubMed] [Google Scholar]
  11. Lee W.-N.; Huang C.-H. Formation of disinfection byproducts in wash water and lettuce by washing with sodium hypochlorite and peracetic acid sanitizers. Food Chemistry: X 2019, 1, 100003. 10.1016/j.fochx.2018.100003. [DOI] [Google Scholar]
  12. Pruss-Ustun A.; Wolf J.; Bartram J.; Clasen T.; Cumming O.; Freeman M. C.; Gordon B.; Hunter P. R.; Medlicott K.; Johnston R. Burden of disease from inadequate water, sanitation and hygiene for selected adverse health outcomes: An updated analysis with a focus on low- and middle-income countries. Int. J. Hyg Environ. Health 2019, 222 (5), 765–777. 10.1016/j.ijheh.2019.05.004. [DOI] [PMC free article] [PubMed] [Google Scholar]
  13. Mitch W. A. Tap water and bladder cancer in China. Nature Sustainability 2022, 5, 643–644. 10.1038/s41893-022-00900-0. [DOI] [Google Scholar]
  14. Marron E. L.; Van Buren J.; Cuthbertson A. A.; Darby E.; von Gunten U.; Sedlak D. L. Reactions of alpha, beta-Unsaturated Carbonyls with Free Chlorine, Free Bromine, and Combined Chlorine. Environ. Sci. Technol. 2021, 55 (5), 3305–3312. 10.1021/acs.est.0c07660. [DOI] [PMC free article] [PubMed] [Google Scholar]
  15. Marron E. L.; Mitch W. A.; Gunten U. V.; Sedlak D. L. A Tale of Two Treatments: The Multiple Barrier Approach to Removing Chemical Contaminants During Potable Water Reuse. Acc. Chem. Res. 2019, 52 (3), 615–622. 10.1021/acs.accounts.8b00612. [DOI] [PMC free article] [PubMed] [Google Scholar]
  16. Marron E. L.; Prasse C.; Buren J. V.; Sedlak D. L. Formation and Fate of Carbonyls in Potable Water Reuse Systems. Environ. Sci. Technol. 2020, 54 (17), 10895–10903. 10.1021/acs.est.0c02793. [DOI] [PMC free article] [PubMed] [Google Scholar]
  17. Jiang J.; Han J.; Zhang X. Nonhalogenated Aromatic DBPs in Drinking Water Chlorination: A Gap between NOM and Halogenated Aromatic DBPs. Environ. Sci. Technol. 2020, 54 (3), 1646–1656. 10.1021/acs.est.9b06403. [DOI] [PubMed] [Google Scholar]
  18. Han J.; Zhang X.; Jiang J.; Li W. How Much of the Total Organic Halogen and Developmental Toxicity of Chlorinated Drinking Water Might Be Attributed to Aromatic Halogenated DBPs?. Environ. Sci. Technol. 2021, 55 (9), 5906–5916. 10.1021/acs.est.0c08565. [DOI] [PubMed] [Google Scholar]
  19. Wu Y.; Wei W.; Luo J.; Pan Y.; Yang M.; Hua M.; Chu W.; Shuang C.; Li A. Comparative Toxicity Analyses from Different Endpoints: Are New Cyclic Disinfection Byproducts (DBPs) More Toxic than Common Aliphatic DBPs?. Environ. Sci. Technol. 2022, 56 (1), 194–207. 10.1021/acs.est.1c03292. [DOI] [PubMed] [Google Scholar]
  20. Deborde M.; von Gunten U. Reactions of chlorine with inorganic and organic compounds during water treatment-Kinetics and mechanisms: a critical review. Water Res. 2008, 42 (1–2), 13–51. 10.1016/j.watres.2007.07.025. [DOI] [PubMed] [Google Scholar]
  21. Gao J.; Proulx F.; Rodriguez M. J. Halogenated acetaldehydes in water: A review of their occurrence, formation, precursors and control strategies. Critical Reviews in Environmental Science and Technology 2019, 49 (15), 1331–1385. 10.1080/10643389.2019.1571353. [DOI] [Google Scholar]
  22. Shah A. D.; Mitch W. A. Halonitroalkanes, halonitriles, haloamides, and N-nitrosamines: a critical review of nitrogenous disinfection byproduct formation pathways. Environ. Sci. Technol. 2012, 46 (1), 119–31. 10.1021/es203312s. [DOI] [PubMed] [Google Scholar]
  23. Huang H.; Wu Q. Y.; Hu H. Y.; Mitch W. A. Dichloroacetonitrile and dichloroacetamide can form independently during chlorination and chloramination of drinking waters, model organic matters, and wastewater effluents. Environ. Sci. Technol. 2012, 46 (19), 10624–31. 10.1021/es3025808. [DOI] [PubMed] [Google Scholar]
  24. Zhang R.; Meng T.; Huang C. H.; Ben W.; Yao H.; Liu R.; Sun P. PPCP Degradation by Chlorine-UV Processes in Ammoniacal Water: New Reaction Insights, Kinetic Modeling, and DBP Formation. Environ. Sci. Technol. 2018, 52 (14), 7833–7841. 10.1021/acs.est.8b00094. [DOI] [PubMed] [Google Scholar]
  25. Szczuka A.; Huang N.; MacDonald J. A.; Nayak A.; Zhang Z.; Mitch W. A. N-Nitrosodimethylamine Formation during UV/Hydrogen Peroxide and UV/Chlorine Advanced Oxidation Process Treatment Following Reverse Osmosis for Potable Reuse. Environ. Sci. Technol. 2020, 54 (23), 15465–15475. 10.1021/acs.est.0c05704. [DOI] [PubMed] [Google Scholar]
  26. Park S.-H.; Wei S.; Mizaikoff B.; Taylor A. E.; Favero C.; Huang C.-H. Degradation of Amine-Based Water Treatment Polymers during Chloramination as N-Nitrosodimethylamine (NDMA) Precursors. Environ. Sci. Technol. 2009, 43 (5), 1360–1366. 10.1021/es802732z. [DOI] [PubMed] [Google Scholar]
  27. Spahr S.; Cirpka O. A.; von Gunten U.; Hofstetter T. B. Formation of N-Nitrosodimethylamine during Chloramination of Secondary and Tertiary Amines: Role of Molecular Oxygen and Radical Intermediates. Environ. Sci. Technol. 2017, 51 (1), 280–290. 10.1021/acs.est.6b04780. [DOI] [PubMed] [Google Scholar]
  28. Vu T. N.; Kimura S. Y.; Plewa M. J.; Richardson S. D.; Marinas B. J. Predominant N-Haloacetamide and Haloacetonitrile Formation in Drinking Water via the Aldehyde Reaction Pathway. Environ. Sci. Technol. 2019, 53 (2), 850–859. 10.1021/acs.est.8b02862. [DOI] [PubMed] [Google Scholar]
  29. Shi J. L.; Plata S. L.; Kleimans M.; Childress A. E.; McCurry D. L. Formation and Fate of Nitromethane in Ozone-Based Water Reuse Processes. Environ. Sci. Technol. 2021, 55 (9), 6281–6289. 10.1021/acs.est.0c07895. [DOI] [PubMed] [Google Scholar]
  30. Lim S.; Shi J. L.; von Gunten U.; McCurry D. L. Ozonation of organic compounds in water and wastewater: A critical review. Water Res. 2022, 213, 118053. 10.1016/j.watres.2022.118053. [DOI] [PubMed] [Google Scholar]
  31. McCurry D. L.; Quay A. N.; Mitch W. A. Ozone Promotes Chloropicrin Formation by Oxidizing Amines to Nitro Compounds. Environ. Sci. Technol. 2016, 50 (3), 1209–17. 10.1021/acs.est.5b04282. [DOI] [PubMed] [Google Scholar]
  32. Yang J.; Dong Z.; Jiang C.; Wang C.; Liu H. An overview of bromate formation in chemical oxidation processes: Occurrence, mechanism, influencing factors, risk assessment, and control strategies. Chemosphere 2019, 237, 124521. 10.1016/j.chemosphere.2019.124521. [DOI] [PubMed] [Google Scholar]
  33. Shah A. D.; Dotson A. D.; Linden K. G.; Mitch W. A. Impact of UV disinfection combined with chlorination/chloramination on the formation of halonitromethanes and haloacetonitriles in drinking water. Environ. Sci. Technol. 2011, 45 (8), 3657–64. 10.1021/es104240v. [DOI] [PubMed] [Google Scholar]
  34. Chen C.; Du Y.; Zhou Y.; Wu Q.; Zheng S.; Fang J. Formation of nitro(so) and chlorinated products and toxicity alteration during the UV/monochloramine treatment of phenol. Water Res. 2021, 194, 116914. 10.1016/j.watres.2021.116914. [DOI] [PubMed] [Google Scholar]
  35. Xu J.; Kralles Z. T.; Dai N. Effects of Sunlight on the Trichloronitromethane Formation Potential of Wastewater Effluents: Dependence on Nitrite Concentration. Environ. Sci. Technol. 2019, 53 (8), 4285–4294. 10.1021/acs.est.9b00447. [DOI] [PubMed] [Google Scholar]
  36. Xue R.; Shi H.; Ma Y.; Yang J.; Hua B.; Inniss E. C.; Adams C. D.; Eichholz T. Evaluation of thirteen haloacetic acids and ten trihalomethanes formation by peracetic acid and chlorine drinking water disinfection. Chemosphere 2017, 189, 349–356. 10.1016/j.chemosphere.2017.09.059. [DOI] [PubMed] [Google Scholar]
  37. Maffettone R.; Manoli K.; Santoro D.; Passalacqua K. D.; Wobus C. E.; Sarathy S. Performic Acid Disinfection of Municipal Secondary Effluent Wastewater: Inactivation of Murine Norovirus, Fecal Coliforms, and Enterococci. Environ. Sci. Technol. 2020, 54 (19), 12761–12770. 10.1021/acs.est.0c05144. [DOI] [PubMed] [Google Scholar]
  38. Wang J.; Chen W.; Wang T.; Reid E.; Krall C.; Kim J.; Zhang T.; Xie X.; Huang C. H. Bacteria and Virus Inactivation: Relative Efficacy and Mechanisms of Peroxyacids and Chlor(am) ine. Environ. Sci. Technol. 2023, 10.1021/acs.est.2c09824. [DOI] [PMC free article] [PubMed] [Google Scholar]
  39. Ragazzo P.; Chiucchini N.; Piccolo V.; Spadolini M.; Carrer S.; Zanon F.; Gehr R. Wastewater disinfection: long-term laboratory and full-scale studies on performic acid in comparison with peracetic acid and chlorine. Water Res. 2020, 184, 116169. 10.1016/j.watres.2020.116169. [DOI] [PubMed] [Google Scholar]
  40. Kim J.; Huang C.-H. Reactivity of Peracetic Acid with Organic Compounds: A Critical Review. ACS ES&T Water 2021, 1 (1), 15–33. 10.1021/acsestwater.0c00029. [DOI] [Google Scholar]
  41. Santacesaria E.; Russo V.; Tesser R.; Turco R.; Di Serio M. Kinetics of Performic Acid Synthesis and Decomposition. Ind. Eng. Chem. Res. 2017, 56 (45), 12940–12952. 10.1021/acs.iecr.7b00593. [DOI] [Google Scholar]
  42. Luukkonen T.; Heyninck T.; Ramo J.; Lassi U. Comparison of organic peracids in wastewater treatment: Disinfection, oxidation and corrosion. Water Res. 2015, 85, 275–85. 10.1016/j.watres.2015.08.037. [DOI] [PubMed] [Google Scholar]
  43. Shah A. D.; Liu Z. Q.; Salhi E.; Hofer T.; von Gunten U. Peracetic acid oxidation of saline waters in the absence and presence of H (2) O (2): secondary oxidant and disinfection byproduct formation. Environ. Sci. Technol. 2015, 49 (3), 1698–705. 10.1021/es503920n. [DOI] [PubMed] [Google Scholar]
  44. West D. M.; Wu Q.; Donovan A.; Shi H.; Ma Y.; Jiang H.; Wang J. N-nitrosamine formation by monochloramine, free chlorine, and peracetic acid disinfection with presence of amine precursors in drinking water system. Chemosphere 2016, 153, 521–7. 10.1016/j.chemosphere.2016.03.035. [DOI] [PubMed] [Google Scholar]
  45. Shah A. D.; Liu Z.-Q.; Salhi E.; Höfer T.; Werschkun B.; von Gunten U. Formation of disinfection by-products during ballast water treatment with ozone, chlorine, and peracetic acid: influence of water quality parameters. Environmental Science: Water Research & Technology 2015, 1 (4), 465–480. 10.1039/C5EW00061K. [DOI] [Google Scholar]
  46. Heeb M. B.; Criquet J.; Zimmermann-Steffens S. G.; von Gunten U. Oxidative treatment of bromide-containing waters: formation of bromine and its reactions with inorganic and organic compounds-a critical review. Water Res. 2014, 48, 15–42. 10.1016/j.watres.2013.08.030. [DOI] [PubMed] [Google Scholar]
  47. Criquet J.; Allard S.; Salhi E.; Joll C. A.; Heitz A.; von Gunten U. Iodate and iodo-trihalomethane formation during chlorination of iodide-containing waters: role of bromide. Environ. Sci. Technol. 2012, 46 (13), 7350–7. 10.1021/es301301g. [DOI] [PubMed] [Google Scholar]
  48. Zhu X.; Zhang X. Modeling the formation of TOCl, TOBr and TOI during chlor(am) ination of drinking water. Water Res. 2016, 96, 166–76. 10.1016/j.watres.2016.03.051. [DOI] [PubMed] [Google Scholar]
  49. Bichsel Y.; von Gunten U. Oxidation of Iodide and Hypoiodous Acid in the Disinfection of Natural Waters. Environ. Sci. Technol. 1999, 33 (22), 4040–4045. 10.1021/es990336c. [DOI] [Google Scholar]
  50. Zhai H.; Zhang X.; Zhu X.; Liu J.; Ji M. Formation of brominated disinfection byproducts during Chloramination of drinking water: new polar species and overall kinetics. Environ. Sci. Technol. 2014, 48 (5), 2579–88. 10.1021/es4034765. [DOI] [PubMed] [Google Scholar]
  51. Cai M.; Sun P.; Zhang L.; Huang C. H. UV/Peracetic Acid for Degradation of Pharmaceuticals and Reactive Species Evaluation. Environ. Sci. Technol. 2017, 51 (24), 14217–14224. 10.1021/acs.est.7b04694. [DOI] [PubMed] [Google Scholar]
  52. 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. 10.1021/acs.est.8b01565. [DOI] [PubMed] [Google Scholar]
  53. 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. 10.1016/j.chemosphere.2016.04.093. [DOI] [PubMed] [Google Scholar]
  54. Zhao X.; Salhi E.; Liu H.; Ma J.; von Gunten U. Kinetic and Mechanistic Aspects of the Reactions of Iodide and Hypoiodous Acid with Permanganate: Oxidation and Disproportionation. Environ. Sci. Technol. 2016, 50 (8), 4358–65. 10.1021/acs.est.6b00320. [DOI] [PubMed] [Google Scholar]
  55. Kim J.; Du P.; Liu W.; Luo C.; Zhao H.; Huang C. H. Cobalt/Peracetic Acid: Advanced Oxidation of Aromatic Organic Compounds by Acetylperoxyl Radicals. Environ. Sci. Technol. 2020, 54 (8), 5268–5278. 10.1021/acs.est.0c00356. [DOI] [PubMed] [Google Scholar]
  56. Kim J.; Zhang T.; Liu W.; Du P.; Dobson J. T.; Huang C. H. Advanced Oxidation Process with Peracetic Acid and Fe(II) for Contaminant Degradation. Environ. Sci. Technol. 2019, 53 (22), 13312–13322. 10.1021/acs.est.9b02991. [DOI] [PubMed] [Google Scholar]
  57. Guo K.; Zheng S.; Zhang X.; Zhao L.; Ji S.; Chen C.; Wu Z.; Wang D.; Fang J. Roles of Bromine Radicals and Hydroxyl Radicals in the Degradation of Micropollutants by the UV/Bromine Process. Environ. Sci. Technol. 2020, 54 (10), 6415–6426. 10.1021/acs.est.0c00723. [DOI] [PubMed] [Google Scholar]
  58. Wang J.; Kim J.; Ashley D. C.; Sharma V. K.; Huang C. H. Peracetic Acid Enhances Micropollutant Degradation by Ferrate(VI) through Promotion of Electron Transfer Efficiency. Environ. Sci. Technol. 2022, 56 (16), 11683–11693. 10.1021/acs.est.2c02381. [DOI] [PubMed] [Google Scholar]
  59. Hua G.; Reckhow D. A.; Abusallout I. Correlation between SUVA and DBP formation during chlorination and chloramination of NOM fractions from different sources. Chemosphere 2015, 130, 82–9. 10.1016/j.chemosphere.2015.03.039. [DOI] [PubMed] [Google Scholar]
  60. Yang M.; Liu J.; Zhang X.; Richardson S. D. Comparative Toxicity of Chlorinated Saline and Freshwater Wastewater Effluents to Marine Organisms. Environ. Sci. Technol. 2015, 49 (24), 14475–83. 10.1021/acs.est.5b03796. [DOI] [PubMed] [Google Scholar]
  61. Yang M.; Zhang X. Halopyrroles: a new group of highly toxic disinfection byproducts formed in chlorinated saline wastewater. Environ. Sci. Technol. 2014, 48 (20), 11846–52. 10.1021/es503312k. [DOI] [PubMed] [Google Scholar]
  62. Parker K. M.; Zeng T.; Harkness J.; Vengosh A.; Mitch W. A. Enhanced formation of disinfection byproducts in shale gas wastewater-impacted drinking water supplies. Environ. Sci. Technol. 2014, 48 (19), 11161–9. 10.1021/es5028184. [DOI] [PubMed] [Google Scholar]
  63. Zhang R.; Sun P.; Boyer T. H.; Zhao L.; Huang C. H. Degradation of pharmaceuticals and metabolite in synthetic human urine by UV, UV/H2O2, and UV/PDS. Environ. Sci. Technol. 2015, 49 (5), 3056–66. 10.1021/es504799n. [DOI] [PubMed] [Google Scholar]
  64. Luo C.; Feng M.; Sharma V. K.; Huang C. H. Oxidation of Pharmaceuticals by Ferrate(VI) in Hydrolyzed Urine: Effects of Major Inorganic Constituents. Environ. Sci. Technol. 2019, 53 (9), 5272–5281. 10.1021/acs.est.9b00006. [DOI] [PubMed] [Google Scholar]
  65. Zhang R.; Yang Y.; Huang C.-H.; Li N.; Liu H.; Zhao L.; Sun P. UV/H2O2 and UV/PDS Treatment of Trimethoprim and Sulfamethoxazole in Synthetic Human Urine: Transformation Products and Toxicity. Environ. Sci. Technol. 2016, 50 (5), 2573–2583. 10.1021/acs.est.5b05604. [DOI] [PubMed] [Google Scholar]
  66. Wei X.; Sanders K. T.; Childress A. E. Reclaiming wastewater with increasing salinity for potable water reuse: Water recovery and energy consumption during reverse osmosis desalination. Desalination 2021, 520, 115316. 10.1016/j.desal.2021.115316. [DOI] [Google Scholar]
  67. Dong H.; Qiang Z.; Richardson S. D. Formation of Iodinated Disinfection Byproducts (I-DBPs) in Drinking Water: Emerging Concerns and Current Issues. Acc. Chem. Res. 2019, 52 (4), 896–905. 10.1021/acs.accounts.8b00641. [DOI] [PubMed] [Google Scholar]
  68. Rose M. R.; Lau S. S.; Prasse C.; Sivey J. D. Exotic Electrophiles in Chlorinated and Chloraminated Water: When Conventional Kinetic Models and Reaction Pathways Fall Short. Environ. Sci. Technol. Lett. 2020, 7 (6), 360–370. 10.1021/acs.estlett.0c00259. [DOI] [Google Scholar]

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