Oxidation of Per- and Polyfluoroalkyl Ether Acids and Other Per- and Polyfluoroalkyl Substances by Sulfate and Hydroxyl Radicals: Kinetic Insights from Experiments and Models

Abstract

Per- and polyfluoroalkyl substances (PFAS) are widely used anthropogenic chemicals. Because of the strength of the carbon─fluorine bond, PFAS are not destroyed in typical water treatment processes. Sulfate $({\text{SO}}_4^{•-})$ and hydroxyl $({}^{•}\text{OH})$ radicals can oxidize some PFAS, but the behavior of per- and polyfluoroalkyl ether acids (PFEAs) in processes involving ${\text{SO}}_4^{•-}$ and ${}^{•}\text{OH}$ is poorly understood. In this study, we determined second-order rate constants $(k)$ describing the oxidation of 18 PFAS, including 15 novel PFEAs, by ${\text{SO}}_4^{•-}$ and ${}^{•}\text{OH}$. Among the studied PFAS, 6:2 fluorotelomer sulfonate reacted most readily with ${}^{•}\text{OH}[k_{•\text{OH}}=(1.1-1.2)\times {10}^7{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}]$, while polyfluoroalkyl ether acids containing an -O-CFH- moiety reacted more slowly $[k_{•\text{OH}}=(0.5-1.0)\times {10}^6{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}]$. In the presence of ${\text{SO}}_4^{•-}$, polyfluoroalkyl ether acids with an -O-CFH- moiety reacted more rapidly $[{k_{\text{SO}4}}^{•-}=(0.89-4.6)\times {10}^6{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}]$ than perfluoroalkyl ether carboxylic acids (PFECAs) and a chloro-perfluoro-polyether carboxylic acid (ClPFPECA) $[{k_{\text{SO}4}}^{•-}=(0.85-9.5)\times {10}^4{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}]$. For homologous series of perfluoroalkyl carboxylic acids, linear and branched monoether PFECAs, and multiether PFECAs, PFAS chain length had little impact on second-order rate constants. ${\text{SO}}_4^{•-}$ reacted with the carboxylic acid headgroup of perfluoroalkyl carboxylic acids and PFECAs. In contrast, for polyfluoroalkyl ether carboxylic and sulfonic acids with an -O-CFH- moiety, the site of ${\text{SO}}_4^{•-}$ attack was the -O-CFH- moiety. Perfluoroalkyl ether sulfonic acids were not oxidized by ${\text{SO}}_4^{•-}$ and ${}^{•}\text{OH}$ under the conditions evaluated in this study.

Graphical Abstract

INTRODUCTION

Per- and polyfluoroalkyl substances (PFAS) are extensively used anthropogenic chemicals that contain at least one fully fluorinated methyl (-CF₃) or methylene (-CF₂₋) group.¹ Because of their unique stability and surface tension lowering properties, PFAS have been and continue to be widely used in industrial and manufacturing processes.^2-7 Two of the most common PFAS, perfluorooctanesulfonic acid (PFOS) and perfluorooctanoic acid (PFOA), have received extensive attention as a result of their environmental persistence, bioaccumulation potential, ubiquity, and adverse effects on human and environmental health.^3,8 Fluorochemical manufacturers have therefore shifted production to fluorinated alternatives, including short-chain PFAS and per- and polyfluoroalkyl ether acids (PFEAs).^2,9,10 For example, hexafluoropropylene oxide-dimer acid (HFPO–DA) and its ammonium salt (“GenX”) are PFOA substitutes, and HFPO–DA is also a byproduct of fluorochemical manufacturing.^11,12 However, GenX is also associated with adverse health outcomes. In June 2022, the U.S. EPA released new final lifetime health advisory levels (HALs) for GenX (10 ng/L) and perfluorobutanesulfonic acid (PFBS, 2000 ng/L) in drinking water along with interim lifetime HALs for PFOA (0.004 ng/L) and PFOS (0.02 ng/L).¹³ HFPO–DA has been detected in surface water downstream of fluorochemical manufacturing sites.^11,14-16 The production of building blocks for fluoropolymers has also led to environmental releases of other PFEAs, such as perfluoroalkyl mono- and multiether carboxylic acids (mono- and multiether PFECAs) and polyfluoroalkyl ether acids (Table 1, Table S1, Figure S1).¹⁷ In North Carolina, discharges to the environment have contaminated important drinking water sources, including the Cape Fear River and thousands of private wells.^2,17-21 Also, nontargeted analysis revealed the existence of a series of chloro-perfluoro-polyether carboxylates (ClPFPECA) and related PFAS species in the vicinity of a fluorochemical manufacturer in New Jersey.²² Given the increasing environmental occurrence of novel PFAS, such as PFEAs, it is important to understand the effectiveness of potential remediation strategies.

Table 1.

Examples of Per- and Polyfluoroalkyl Ether Acids (PFEAs)

Compound	Formula	CAS #	Molecular structure
Perfluoroalkyl mono-ether carboxylic acids (mono-ether PFECAs)
Perfluoro-2-methoxyacetic acid (PFMOAA)	C3HF5O3	674-13-5
Hexafluoropropylene oxide-dimer acid (HFPO-DA)	C6HF11O3	13252-13-6
Perfluoroalkyl multi-ether carboxylic acid (multi-ether PFECA)
Perfluoro(3,5-dioxahexanoic) acid (PFO2HxA)	C4HF7O4	39492-88-1
Chloro-perfluoro-polyether carboxylate (ClPFPECA)
Chloro-perfluoro-polyether carboxylate (ClPFPECA) 0,1	C8HClF14O4	-
Perfluoroalkyl ether sulfonic acid (PFESA)
Perfluoro(2-ethoxyethane)sulfonic acid	C4F9KO4S	117205-07-9
Polyfluoroalkyl ether acids
Ethanesulfonic acid, 2-[1-[difluoro(l,2,2,2-tetrafluoroethoxy)methyl]-1,2,2,2-tetrafluoroethoxy]-1,1,2,2-tetrafluoro- (Nafion by-product 2)	C7H2F14SO5	749836-20-2
4,8-dioxa-3H-perfluorononanoic acid (ADONA)	C7H2F12O4	919005-14-4

Because of the strength of the carbon─fluorine bond, PFAS are not destroyed by the majority of oxidants used in drinking water and wastewater treatment.^23-25 Generally, PFAS are removed from bulk water by concentration techniques, such as granular activated carbon adsorption, ion exchange, or high-pressure membranes.²⁶ These technologies generate concentrated residuals that can be further treated by processes capable of PFAS destruction. Sulfate radicals $({\text{SO}}_4^{•-})$ and hydroxyl radicals $({}^{•}\text{OH})$ are generated in some destructive technologies. For example, ${}^{•}\text{OH}$ is generated in processes involving nonthermal plasmas,²⁷ electrochemical oxidation,²⁸ and supercritical water oxidation²⁹ that are being explored for PFAS destruction.^30,31 Furthermore, ${\text{SO}}_4^{•-}$ can contribute to PFAS destruction and can be generated in electrochemical oxidation processes³² as well as in advanced oxidation processes (AOPs) relying on UV- or heat-activation of persulfate (PS).^33,34

Hydroxyl radicals preferably add to C═C bonds or abstract H from C─H bonds.³⁵ Because of their fully fluorinated alkyl chain, perfluoroalkyl acids (PFAAs), such as PFOA and PFOS, are recalcitrant toward ${}^{•}\text{OH}$ attack.^25,36,37 In contrast, ${}^{•}\text{OH}$ can oxidize polyfluorinated compounds, such as 6:2 fluorotelomer sulfonate (6:2 FtS).^38,39 Reactions between ${}^{•}\text{OH}$ and PFEAs have not been well characterized. For HFPO–DA, a recent study indicated that it was not oxidized in systems involving ${}^{•}\text{OH}$.⁴⁰ Another study evaluating the behavior of 15 PFEAs in the total oxidizable precursor (TOP) assay, in which ${}^{•}\text{OH}$ is the principal oxidant, showed that perfluoroalkyl ether acids are resistant to ${}^{•}\text{OH}$ attack, while polyfluoroalkyl ether acids with a -O-CFH- moiety were readily oxidized.²⁰ The TOP assay was originally developed to determine the presence of PFAS that are not targeted by standard analytical methods for PFAS and that may be transformed through biotic or abiotic environmental processes to commonly measured PFAAs. In the TOP assay, oxidizable PFAA precursors are converted to PFAAs via ${}^{•}\text{OH}$ formed during thermolysis of PS at pH >12, a condition at which ${\text{SO}}_4^{•-}$ preferentially reacts with ${\text{OH}}^{-}$ to form ${}^{•}\text{OH}$.⁴¹

Sulfate radicals react more selectively with contaminants via direct one-electron transfer to form sulfate, making them stronger direct electron transfer oxidants than ${}^{•}\text{OH}$.²⁵ Under acidic conditions (pH <3), perfluoroalkyl carboxylic acids (PFCAs) can be degraded by ${\text{SO}}_4^{•-}$ generated via UV photolysis or thermolysis of PS;^23,24,42 while under basic (i.e., pH 13) and circumneutral pH conditions (i.e., pH 8.2), PFCA transformation was not observed.^6,43 Perfluoroalkyl sulfonic acids (PFSAs) are recalcitrant to ${\text{SO}}_4^{•-}$ attack.^6,24 There are only a few peer-reviewed papers that have attempted to develop second-order rate constants $({k_{\text{SO}4}}^{•-})$ that describe reactions between ${\text{SO}}_4^{•-}$ and PFAS, and these papers were limited to a small number of PFCAs. For example, ${k_{\text{SO}4}}^{•-}$ values have been previously determined for PFCAs containing 3–8 carbon atoms from kinetic models (2.59 × 10⁵ to 9.31 × 10⁷ M⁻¹ s⁻¹),²³ for PFCAs containing 4–8 carbon atoms from competition kinetics [(1.7–4.4) × 10⁴ M⁻¹ s⁻¹],²⁴ and for perfluorobutanoic acid (PFBA) and trifluoroacetic acid (TFA) from laser flash photolysis [(0.9 ± 0.2) × 10⁴ M⁻¹ s⁻¹].⁴² Values for ${k_{\text{SO}4}}^{•-}$ differ by 1 order of magnitude or more among studies;^23,24 and there is disagreement on the chain length-dependence of ${k_{\text{SO}4}}^{•-}$; one paper argued that rate constants increase with decreasing PFCA chain length²³ while the other stated that rate constants are independent of PFCA chain length.²⁴ To the best of our knowledge, very few papers have reported the behavior of PFEAs in ${\text{SO}}_4^{•-}$-based AOPs. The fully halogenated fluoroether sulfonic acid 9-chlorohex-adecafluoro-3-oxanonane-1-sulfonate (trade name F-53B) is recalcitrant to ${\text{SO}}_4^{•-}$ attack.⁴⁴ Bao et al. (2018)⁴⁰ demonstrated that <5% HFPO–DA was oxidized in 3 h by UV photolysis of PS (UV/PS) while Ding et al. (2022)⁴⁵ found that HFPO–DA was oxidized in a thermally activated PS process and that the pseudo-first-order oxidation rate was impacted by coexisting anions such as chloride. However, second-order rate constants that describe the oxidation of PFEAs by ${\text{SO}}_4^{•-}$ and ${}^{•}\text{OH}$ are lacking.

The overarching goal of this research was to investigate the oxidation rates of 15 structurally diverse and poorly understood PFEAs as well as 2 PFCAs and 1 FtS in systems containing ${}^{•}\text{OH}$ and ${\text{SO}}_4^{•-}$. Specific objectives were to (1) experimentally determine PFAS oxidation rates by ${}^{•}\text{OH}$ and ${\text{SO}}_4^{•-}$, (2) derive second-order rate constants describing PFAS oxidation by ${}^{•}\text{OH}$ and ${\text{SO}}_4^{•-}$, and (3) assess the effect of the solution pH on PFAS oxidation rates using experimental data and kinetic models. To meet these objectives, we conducted batch experiments in which ${}^{•}\text{OH}$ and ${\text{SO}}_4^{•-}$ were generated by UV irradiation of H₂O₂ and PS. Resulting data were described with two independent modeling approaches to derive second-order rate constants describing PFAS oxidation by ${}^{•}\text{OH}$ and ${\text{SO}}_4^{•-}$.

MATERIALS AND METHODS

Materials

Eighteen PFAS in three classes (i.e., 2 PFCAs, 1 FtS, and 15 PFEAs) were targeted in this study (Table S1 and Figure S1). Analytical standards were obtained from Wellington Laboratories (Guelph, ON, Canada), The Chemours Company (Wilmington, DE), SynQuest Laboratories (Alachua, FL), Fluoryx Laboratories (Carson City, NV), and the U.S. EPA (Research Triangle Park, NC) (see Table S1 for details). Isotopically labeled internal standards were purchased from Wellington Laboratories (Guelph, ON, Canada). All other chemicals and solvents were purchased from Fisher Scientific (Hampton, NH) and Sigma-Aldrich (St. Louis, MO).

Experimental Approach

To determine the absorbance spectra and decadic molar absorption coefficients of target PFAS with sufficient availability, we performed experiments using a UV–vis spectrophotometer with a 1 cm standard quartz cuvette (PFAS = 1 g/L, λ = 200–400 nm, pH 3.0, 7.0, or 12.0). Batch UV photolysis, UV/H₂O₂, UV/PS, and control experiments were conducted in a 50 mL glass Petri dish in a bench scale quasi-collimated beam (QCB) apparatus.⁴⁶ Standardized methods were used to calculate UV irradiance at the water surface and UV fluence (dose) delivered to the sample.⁴⁶ The QCB was equipped with four low-pressure (LP) UV lamps (λ = 254 nm), and a UV radiometer (International Light Technologies, Peabody, MA) was used to measure the UV irradiance at the surface of the sample. Before each experiment, UV lamps were warmed up for 2 h to ensure stable lamp output. The average irradiance [incident irradiance corrected for the water absorbance and path length (1.3 cm)] in the aqueous solution was calculated as previously described.⁴⁶ In batch UV photolysis experiments, UV fluence (mJ/cm²) delivered to the sample was calculated as the average irradiance (1.12 mW/cm²) multiplied by exposure time.^46,47 However, in UV/H₂O₂ and UV/PS experiments, decreasing oxidant concentrations led to decreased solution absorbance and increased irradiance over time. The average irradiance during 8 h in UV/H₂O₂ and UV/PS experiments was 0.88 and 1.04 mW/cm², respectively (details provided in Text S1, Figures S2-S3).

Analytical standards of PFAS were either used as received in methanol or water or they were diluted with methanol or water, respectively, to create PFAS stock solutions (Table S1). An aliquot of a stock solution containing 50 ng/μL of an individual PFAS was added to a Petri dish and evaporated to dryness if the PFAS stock solution was prepared in methanol. PFAS were subsequently dissolved in UPW to yield a starting concentration of 1 or 0.1 mg/L depending on PFAS availability. Initial concentrations of H₂O₂ and potassium persulfate were 20 and 5 mM, respectively. Dark control experiments were conducted with H₂O₂ or PS. Except when mentioned, the pH was monitored but not adjusted or buffered to avoid radical scavenging from buffer species and possible buffer degradation products. PFAS degradation by UV/PS was also investigated under alkaline conditions similar to those employed in the TOP assay; here, 150 mM of sodium hydroxide was added to maintain pH >12 during the experiment.^20,41 All experiments were conducted at room temperature (20 ± 2 °C). Samples were taken at predetermined times, diluted with UPW, and stored at room temperature until analysis (details provided in Text S2).

Analytical Methods

Calibration standards were used to determine all of the PFAS concentrations. Quantification was achieved by an isotope dilution approach when possible, in which the analyte response was normalized to that of an isotopically labeled analog.^18,20 For other PFAS, the analyte response was normalized to that of an isotopically labeled PFAS with an LC retention time similar to that of the analyte (Table S1).⁴⁸ Details of the analytical methods, including approaches to identify reaction products by liquid chromatography–high resolution mass spectrometry (LC–HRMS), are provided in the Supporting Information (Texts S3-S4).

H₂O₂ concentrations were determined with the Ghormley method,⁴⁹ which is based on the spectrophotometric determination of ${\mathrm{I}}_3^{-}$ formed when H₂O₂ reacts with I⁻. PS concentrations were measured using the KI colorimetric method.⁵⁰ The concentration of fluoride ions (F⁻) released from PFAS was determined with an ion selective electrode (ISE; Thermo Scientific, Waltham, MA). Additional details of analytical methods and defluorination calculations are listed in Text S5.

Mathematical Modeling

PFAS oxidation kinetics were described mathematically using both pseudo-first-order and second-order models. Both time- and fluence-based pseudo-first-order reaction rate constants ($k^{\prime }$ and $k_{\mathrm{f}}^{\prime }$, respectively) were determined from linear regression analysis of log-transformed concentration data that were plotted as a function of time or fluence. Second-order reaction rate constants, $k_{•\text{OH}}$ and ${k_{\text{SO}4}}^{•-}$, that describe PFAS oxidation by ${}^{•}\text{OH}$ and ${\text{SO}}_4^{•-}$, respectively, were obtained with kinetic models based on (1) calculated steady-state ${}^{•}\text{OH}$ or ${\text{SO}}_4^{•-}$ concentrations and (2) systems of differential equations describing known reactions in UV/H₂O₂ and UV/PS systems²³ (Kintecus 6.8).⁵¹ Details of the two independent modeling approaches are provided in Texts S6-S7 and Table S2. Using Kintecus, we predicted PFAS degradation, product formation and degradation, PS and/or H₂O₂ decay, pH changes, and radical concentrations. In addition, to understand the importance of reactions relative to each species in the model, we performed sensitivity analyses using the built-in function in Kintecus. Using PFOA oxidation in the UV/PS system as an example, we identified major sources and sinks for each species. Details are provided in Text S8 and Figures S6-S7.

RESULTS AND DISCUSSION

Behavior of PFEAs and Other PFAS during UV Photolysis and Hydroxyl Radical Exposure

We determined absorbance spectra and decadic molar absorption coefficients (λ = 200–400 nm at pH 3.0, 7.0, and 12.0) for selected PFAS and evaluated the photolytic behavior of all 18 targeted PFAS at λ = 254 nm using a QCB apparatus. None of the studied PFAS, including the 15 PFEAs, showed measurable degradation by UV photolysis after ~8 h, corresponding to a UV dose of ~32 000 mJ/cm² (detailed discussion provided in Text S9). Furthermore, we studied the behavior of 15 PFEAs, 2 PFCAs, and 1 FtS in UV/H₂O₂ experiments. Except for 6:2 FtS, none of the studied PFAS exhibited measurable degradation over ~8 h (~25 000 mJ/cm²). The second-order reaction rate constant $(k_{•\text{OH}})$ describing 6:2 FtS oxidation by ${}^{•}\text{OH}$ was determined to be (1.1–1.2) × 10⁷ M⁻¹ s⁻¹, which exceeded a previously reported value (4.0 × 10⁶ M⁻¹ s⁻¹, see Text S10 for details) by a factor of about 3.

Behavior of PFEAs during Sulfate Radical Exposure

The behavior of PFEAs in UV/PS experiments was investigated in unbuffered UPW. Upon activation of PS by UV photolysis, a free radical pathway is initiated through the formation of ${\text{SO}}_4^{•-}$ (eq 2),⁵² followed by the reaction of ${\text{SO}}_4^{•-}$ with H₂O to form ${}^{•}\text{OH}$ and molecular oxygen (eqs 3-5).^52,53

$${\mathrm{S}}_2{\mathrm{O}}_8^{2-}+hv\to 2{\text{SO}}_4^{•-}[\Phi ({\text{SO}}_4^{•-})=0.70]$$ (2)

$${\text{SO}}_4^{•-}+{\mathrm{H}}_2\mathrm{O}\to {}^{•}\text{OH}+{\text{SO}}_4^{2-}+{\mathrm{H}}^{+}$$ (3)

$$2{}^{•}\text{OH}\to {\mathrm{H}}_2\mathrm{O}+0.5{\mathrm{O}}_2$$ (4)

$$\text{Overall:}{\mathrm{S}}_2{\mathrm{O}}_8^{2-}+{\mathrm{H}}_2\mathrm{O}\to {2\mathrm{H}}^{+}+2{\text{SO}}_4^{2-}+{0.5\mathrm{O}}_2$$ (5)

The production of ${\mathrm{H}}^{+}$ leads to a continuously decreasing solution pH during the reaction (starting pH ~4, final pH ~2–3).⁵² In UV/PS experiments, ${}^{•}\text{OH}$ was generated, but at levels that had no measurable impact on ${k_{\text{SO}4}}^{•-}$ values as described below. In the following sections, the behavior of PFEAs in UV/PS experiments will be discussed for each PFEA subclass targeted in this study.

Monoether PFECAs + $S{O}_4^{•-}$

Oxidation rates and associated rate constants of linear and branched monoether PFECAs were determined in UV/PS experiments (Figure 1, Figures S10-S11, and Tables S9-S10). For linear monoether PFECAs (PFMOAA, PFMOPrA, and PFMOBA), in which the carboxylic acid moiety was adjacent to a secondary carbon (−CF₂─COOH), parent compound transformation was 18–54% at ~30 000 mJ/cm² (Figure 1), and $k_{\mathrm{f}}^{\prime }$ values were (0.836–2.59) × 10⁻⁵ cm² mJ⁻¹ (Table S11). Using both modeling approaches (steady-state ${\text{SO}}_4^{•-}$ and Kintecus), second-order rate constants were determined to be (0.85–4.3) × 10⁴ M⁻¹ s⁻¹ (Table 2), and no correlations were found between ${k_{\text{SO}4}}^{•-}$ and the chain length of linear monoether PFECAs. During the degradation of PFMOPrA and PFMOBA in the UV/PS process, formation of PFMOAA and PFMOPrA, respectively, was observed using LC–HRMS (Figures S12-S13,Table S12), suggesting that the initial site of attack by $\text{SO}{4}^{•-}$ is the carboxylate headgroup of monoether PFECAs. We used kinetic models and ${k_{\text{SO}4}}^{•-}$ values determined in our study to simulate the degradation of linear monoether PFECAs and the formation of products in the UV/PS process. Overall, model results agreed well with experimental data (Figures S11-S13).

Figure 1.

Fluence-based degradation rates of (A) PFCAs, (B) linear monoether PFECAs, (C) branched monoether PFECAs, (D) multiether PFECAs, (E) ClPFPECA 0,1 and perfluoroalkyl ether sulfonic acids, and (F) polyfluoroalkyl ether acids in UV/PS experiments. The initial PFAS concentration was ~1 or 0.1 mg/L (Table S9), and the initial PS concentration was 5 mM. The pH was monitored but not adjusted. Curves describe pseudo-first-order kinetics, and the corresponding rate constants are given in Table S11. Duplicate experiments were conducted for PFOA, HFPO–DA, ClPFPECA 0,1, and Nafion byproduct 2 and triplicate experiments were conducted for ADONA (Table S9); data points for all replicate experiments are shown. Approximately 20% of samples were collected in duplicate, and error bars represent the range obtained for duplicate samples.

Table 2.

Second-Order Reaction Rate Constants $({k_{\text{SO}4}}^{•-})$ Describing PFAS Oxidation by ${\text{SO}}_4^{•-}$

Class	Compound	${k_{\text{SO}4}}^{•-}({\mathrm{M}}^{-1}{\mathrm{s}}^{-1})$ obtained from steady-state approach (this study)	${k_{\text{SO}4}}^{•-}({\mathrm{M}}^{-1}{\mathrm{s}}^{-1})$ obtained from Kintecus (this study)	${k_{\text{SO}4}}^{•-}({\mathrm{M}}^{-1}{\mathrm{s}}^{-1})$ (other studies)a
PFCAs	PFBA	(1.8 ± 0.5) × 104b	(3.0 ± 0.5) × 104	(1.1 ± 0.2) × 107 (KM)23
				1.3 × 104 (CK)55
				(1.7–4.4) × 104 (CK)24
				(0.9 ± 0.2) × 104 (LFP)42
	PFOA	(2.9 ± 0.8) × 104	(4.7 ± 0.8) × 104	(2.6 ± 0.2) × 105 (KM)23
				(1.7–4.4) × 104 (CK)24
Monoether PFECAs	PFMOAA	(0.85 ± 0.34) × 104	(1.4 ± 0.5) × 104
	PFMOPrA	(2.6 ± 0.6) × 104	(4.3 ± 0.4) × 104
	PMPA	(1.2 ± 0.3) × 104	(2.0 ± 0.2) × 104
	PFMOBA	(1.5 ± 0.5) × 104	(2.5 ± 0.6) × 104
	PEPA	(1.2 ± 0.4) × 104	(1.9 ± 0.5) × 104
	HFPO–DA (“GenX”)	(1.7 ± 0.5) × 104	(2.7 ± 0.5) × 104
Multiether PFECAs	PFO2HxA	(2.1 ± 0.6) × 104	(3.5 ± 0.5) × 104
	PFO4DA	(2.7 ± 0.7) × 104	(4.4 ± 0.6) × 104
ClPFPECA	ClPFPECA 0,1	(5.8 ± 1.6) × 104	(9.5 ± 1.6) × 104
Perfluoroalkyl ether sulfonic acids	Perfluoro(2-ethoxyethane)sulfonic acid		No measurable degradation in ~8 h
	Perfluoro(4-methyl-3,6-dioxaoctane) sulfonic acid		No measurable degradation in ~8 h
Polyfluoroalkyl ether acids	Nafion byproduct 2	(0.89 ± 0.03) × 106	(1.7 ± 0.2) × 106
	NVHOS	(1.1 ± 0.1) × 106	(2.1 ± 0.3) × 106
	ADONA	(3.0 ± 0.3) × 106	(4.6 ± 0.7) × 106
	HydroEVE	(0.93 ± 0.07) × 106	(1.8 ± 0.2) × 106

^aKM: kinetic modeling, CK: competition kinetics, LFP: laser flash photolysis.

^bMean ± one standard deviation (SD). The SD was determined via error propagation from SDs for pseudo-first-order rate constants and ${[{\text{SO}}_4^{•-}]}_{\text{ss}}$.

For branched monoether PFECAs (PMPA, PEPA, and HFPO–DA) containing a carboxylic acid functional group adjacent to a tertiary carbon (>CF-COOH), 24–37% of the parent compound was transformed at ~30000 mJ/cm² (Figure 1C). The $k_{\mathrm{f}}^{\prime }$ values for branched monoether PFECAs were (1.16–1.67) × 10⁻⁵ cm² mJ⁻¹ (Table S11), similar to the linear monoether PFECAs. For HFPO–DA, defluorination over ~8 h (30 000 mJ/cm²) was 6.09%, and fluorine mass balances were 80–100% over the course of the experiments (Figure S14). Mass balances <100% may be a result of the formation of products/intermediates that were not amenable to detection by our LC-HRMS method (e.g., TFA), potential losses of volatile intermediates, and/or variability commonly associated with LC-HRMS analysis. The ${k_{\text{SO}4}}^{•-}$ value for HFPO–DA [(1.7–2.8) × 10⁴ M⁻¹ s⁻¹, Table 2] was statistically lower than that for PFOA [(2.9–4.7) × 10⁴ M⁻¹ s⁻¹ as described in the next section, t test, p < 0.05], which agrees with a previous study demonstrating HFPO–DA degraded more slowly than PFOA in UV/PS experiments, possibly because of steric hindrance associated by the branched ${\text{CF}}_{3^{-}}$ moiety at the C2 carbon, which may shield the carboxylate group.⁴⁰ The proposed initial site of oxidation by ${\text{SO}}_4^{•-}$ is the carboxylate headgroup rather than the ether bond on HFPO–DA, which is supported by DFT calculations.⁴⁵ Formation of PFPrA and TFA from the degradation of HFPO–DA was simulated via kinetic modeling, assuming a stepwise degradation from HFPO–DA to PFPrA, and from PFPrA to TFA (Figure S14). Our model results were in good agreement with experimental data. For PMPA and PEPA, ${k_{\text{SO}4}}^{•-}$ values were (1.2–2.0) × 10⁴ M⁻¹ s⁻¹ (Table 2), and no correlations were found between ${k_{\text{SO}4}}^{•-}$ and the chain length of branched monoether PFECAs.

Multiether PFECAs + $S{O}_4^{•-}$

Multiether PFECAs (PFO2HxA and PFO4DA), in which the carboxylic acid was adjacent to repeating – CF₂O– groups, degraded at a similar rate $[k_{\mathrm{f}}^{\prime }=(2.07-2.67)\times {10}^{-5}{\text{cm}}^2{\text{mJ}}^{-1}]$ as monoether PFECAs in UV/PS experiments (Figure 1D, Table S11). Our observations of similar rate constants for mono- and multiether PFECAs regardless of the number of ether oxygen atoms further suggest ${\text{SO}}_4^{•-}$ preferably attacks the carboxylate headgroup rather than the ether bond(s) on PFECAs.⁴⁵ The ${k_{\text{SO}4}}^{•-}$ value for PFO4DA [(2.7–4.4) × 10⁴ M⁻¹ s⁻¹] was statistically larger (t test, p < 0.05) than that for PFO2HxA [(2.1–3.5) × 10⁴ M⁻¹ s⁻¹], but the difference was within a factor of 2 (Table 2).

ClPFPECA + $S{O}_4^{•-}$

The recently identified ClPFPECA 0,1²² has a fluorine-to-chlorine substitution on the C7 carbon (Figure S1) and the pseudo-first-order rate constant describing its oxidation (5.71 × 10⁻⁵ cm² mJ⁻¹) was more than twice that for PFECAs (Figure 1E, Table S11). Similarly, the ${k_{\text{SO}4}}^{•-}$ for ClPFPECA 0,1 [(5.8–9.5) × 10⁴ M⁻¹ s⁻¹ (Table 2)] was more than twice that for PFECAs. Hence, the result for ClPFPECA 0,1 suggests the fluorine-to-chlorine substitution enhanced PFAS reactivity toward ${\text{SO}}_4^{•-}$ in comparison with perfluorinated PFECAs. In contrast to ClPFPECA 0,1, F-53B, a chlorofluoroalkyl ether sulfonic acid, was recalcitrant in the UV/PS process.^44,54 This result suggests that the sulfonic and carboxylic acid headgroups strongly impact the reactivity of chlorofluoroalkyl ether acids that lack −C─H bonds.

PFESAs + $S{O}_4^{•-}$

Similar to PFSAs, the two perfluoroalkyl ether sulfonic acids (PFESAs) selected in this study did not degrade in UV/PS experiments (Figure 1E), suggesting the introduction of ether oxygen atoms into PFSAs has a negligible effect on reactivity toward ${\text{SO}}_4^{•-}$.

Polyfluoroalkyl Ether Acids + $S{O}_4^{•-}$

Polyfluoroalkyl ether acids containing the -O-CFH- moiety [Nafion byproduct 2, NVHOS, HydroEVE, and 4,8-dioxa-3H-perfluorononanoic acid (ADONA)] exhibited pseudo-first-oder rate constants that were ~2 orders of magnitude larger than those of PFECAs (Figure 1F, Table S11). For example, ADONA, a polyfluoroalkyl ether carboxylic acid, was no longer detectable at ~1000 mJ/cm² $(k_{\mathrm{f}}^{\prime }=3.93\times {10}^{-3}{\text{cm}}^2{\text{mJ}}^{-1}$, Table S11). Using LC–HRMS, we found the disappearance of ADONA was accompanied by the formation of PFMOPrA (Figure S15, Table S12). The molar yield of PFMOPrA from ADONA was ~100% from ~0.4–1.5 h before PFMOPrA was oxidized further by ${\text{SO}}_4^{•-}$ (Figure S15). Formation of PFMOPrA suggests the site of ${\text{SO}}_4^{•-}$ attack on ADONA is the -O-CFH- moiety, a moiety that is also susceptible to ${}^{•}\text{OH}$ attack.²⁰ Using ${k_{\text{SO}4}}^{•-}$ values obtained from kinetic models (4.6 × 10⁶ and 4.3 × 10⁴ M⁻¹ s⁻¹ for ADONA and its oxidation product PFMOPrA, respectively), model results describing ADONA and PFMOPrA oxidation agreed well with the experimental data (Figure S15). Defluorination of ADONA (initial concentration ~0.1 mg/L) was 41.2% at ~6 h (22 000 mJ/cm²), and the fluorine mass balance was 76% at ~6 h (Figure S15).

The pseudo-first-order rate constant for HydroEVE (1.21 × 10⁻³ cm² mJ⁻¹) was about one-third of that obtained for ADONA (Figure 1F, Table S11), and the ${k_{\text{SO}4}}^{•-}$ for HydroEVE was (0.93–1.8) × 10⁶ M⁻¹ s⁻¹ (Table 2). Possible explanations for the difference in ${k_{\text{SO}4}}^{•-}$ values for ADONA and HydroEVE are (1) the branched structure of HydroEVE introduces a steric hindrance and (2) the -O-CFH- moiety on HydroEVE is located further away from the carboxylate headgroup, thus changing the electron density at the -O-CFH- moiety (Figure S1). Rate constants for polyfluoroalkyl ether sulfonic acids (Nafion byproduct 2 and NVHOS) $[k_{\mathrm{f}}^{\prime }=(1.16-1.42)\times {10}^{-3}{\text{cm}}^2{\text{mJ}}^{-1}$, and ${k_{\text{SO}4}}^{•-}=[(0.89-2.1)\times {10}^6{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}]$ were similar to those for HydroEVE (Tables 1 and S11), indicating that sulfonic and carboxylic acid headgroups have a relatively small effect on ${k_{\text{SO}4}}^{•-}$ values describing the oxidation of polyfluoroalkyl ether acids containing the -O-CFH- group. Defluorination of Nafion byproduct 2 (initial concentration ~1 mg/L) was 18.1% at 3 h (11 000 mJ/cm²), while fluoride was below the method reporting limit (0.02 mg/L) for HydroEVE and NVHOS because of the relatively low initial concentration (~0.1 mg/L). No oxidation product(s) could be identified by LC–HRMS for HydroEVE, Nafion byproduct 2, and NVHOS, suggesting oxidation of these compounds produced low-molecular weight organofluorine species, volatile organofluorine species, and/or fluoride that could not be detected by the analytical methods we employed.

Behavior of PFCAs during Sulfate Radical Exposure

In addition to 15 PFEAs, we conducted UV/PS experiments with PFBA and PFOA to (1) link our results to those obtained in prior studies and (2) resolve discrepancies observed in prior studies in terms of both the magnitude of second-order rate constants and the dependence of second-order rate constants on PFCA chain length. For PFBA, the $k_{\mathrm{f}}^{\prime }$ value was 1.80 × 10⁻⁵ cm² mJ⁻¹ (Table S11), and formation of perfluoropropanoic acid (PFPrA) was observed from the degradation of PFBA (Figure S16). TFA was not targeted in this study, but the formation of TFA from the degradation of PFBA in ${\text{SO}}_4^{•-}$-based AOPs was reported previously.²⁴ In our study, defluorination from the degradation of PFBA after ~8 h was 3.5% (Figure S16), and fluorine mass balances for the oxidation of PFBA in the UV/PS process were 80–100% over the course of the experiment. Using both modeling approaches, the second-order rate constant describing PFBA oxidation by ${\text{SO}}_4^{•-}$ $({k_{\text{SO}4}}^{•-})$ was determined to be (1.8–3.0) × 10⁴ M⁻¹ s⁻¹ (Table 2), which agrees well with values reported from competition kinetics [(1.7–4.4) × 10⁴ M⁻¹ s⁻¹] relative to the rate of TFA oxidation (1.6 × 10⁴ M⁻¹ s⁻¹, obtained from laser flash photolysis),^24,55 and from laser flash photolysis [(0.9 ± 0.2) × 10⁴ M⁻¹ s⁻¹].⁴² Qian et al. (2016)²³ derived the ${k_{\text{SO}4}}^{•-}$ for PFBA (1.1 × 10⁷ M⁻¹ s⁻¹) by fitting a kinetic model to PFOA degradation and product formation data, assuming a stepwise degradation of PFCAs and a 1:1 conversion ratio (i.e., 1 mol of PFOA forms 1 mol of PFHpA, 1 mol of PFHpA forms 1 mol of PFHxA, etc.). Using this approach, progressively lower experimental yields of PFCAs with shorter chain lengths could only be described with second-order rate constants that increased with decreasing perfluoroalkyl chain length, and the ${k_{\text{SO}4}}^{•-}$ value for PFBA in Qian et al. (2016)²³ is most likely an overestimate.²⁴ An understanding of the reaction pathways is therefore needed to accurately determine second-order rate constants for reaction intermediates. To demonstrate this point, we used the kinetic model to simulate product formation (PFPrA and TFA) from the degradation of PFBA in UV/PS experiments (Figure S16) using two assumptions: (1) PFBA forms PFPrA, and PFPrA forms TFA (all at a 1:1 conversion ratio); and (2) PFBA forms PFPrA and TFA, and PFPrA forms TFA based on conversion ratios calculated by Kutsuna and Hori (2007).⁵⁵ As previously reported by Kutsuna and Hori (2007)⁵⁵ conversion ratios for PFPrA and TFA formation from PFBA are 0.75 ± 0.05 and 0.17 ± 0.02, respectively, and the conversion ratio for TFA formation from PFPrA is 0.88 ± 0.05. For scenario 1, the second-order rate constant for PFPrA degradation to TFA (6.0 × 10⁴ M⁻¹ s⁻¹) was determined from kinetic modeling and was about twice that for PFBA degradation to PFPrA (i.e., rate constants increased with decreasing carbon chain length of PFCAs, Figure S16), In contrast, for scenario 2, the second-order rate constants for PFBA degradation to PFPrA and for PFPrA degradation to TFA were the same (i.e., rate constants were not correlated with carbon chain length of PFCAs, Figure 2). In both scenarios, models described experimentally determined PFBA disappearance and PFPrA appearance well, demonstrating that accurate conversion ratios are important to include in kinetic modeling to determine the second-order rate constants associated with sequential PFCA degradation in the UV/PS process.

Figure 2.

Concentration profiles of PFBA, PFPrA, and TFA for PFBA degradation in the UV/PS process. Results were obtained experimentally and by kinetic modeling. PFBA initial concentration was ~1 mg/L and PS initial concentration was 5 mM. The pH was monitored but not adjusted. Approximately 20% of samples were collected in duplicate, and error bars represent the range of duplicate samples.

During the degradation of PFOA in UV/PS experiments, the formation of shorter PFCAs was observed, with PFHpA being the dominant PFCA, followed by PFHxA, PFPeA, and PFBA (Figure S17). PFPrA was targeted but not found from the degradation of PFOA, most likely due to concentrations being below our method reporting limit (10 ng/L). Defluorination from the degradation of PFOA after ~8 h was 3.97% (Figure S17). The ${k_{\text{SO}4}}^{•-}$ for PFOA was (2.9–4.7) × 10⁴ M⁻¹ s⁻¹ (Table 2), which agrees well with the value reported previously from competition kinetics [(1.7–4.4) × 10⁴ M⁻¹ s⁻¹].²⁴ Although ${k_{\text{SO}4}}^{•-}$ for PFOA was statistically higher than that for PFBA (t test, p < 0.05) in our study, the values were within a factor of 2. Lutze et al. (2018)²⁴ suggested that rate constants describing PFCA oxidation by ${\text{SO}}_4^{•-}$ exhibit a negligible correlation with PFCA chain length, and our results for PFBA and PFOA are consistent with this finding. In Qian et al. (2016),²³ ${k_{\text{SO}4}}^{•-}$ for PFOA was determined to be 2.6 × 10⁵ M⁻¹ s⁻¹ from kinetic modeling, which is about 1 order of magnitude higher than values determined from our study and in Lutze et al. (2018).²⁴ Uncertainty in the rate constants associated with the 37 reactions in the kinetic model (Table S2) may explain the disparity. Based on results of our sensitivity analysis, reactions 18, 20, and 27 in Table S2 are important sources or sinks for several modeled species, including ${}^{•}\text{OH}$ and ${\text{SO}}_4^{•-}$ (Figures S6-S7). Values for the rate constants associated with reactions 18, 20, and 27 (Table S2) used in our model were within the range or similar to previously reported values but differed from those employed by Qian et al. (2016).²³ Specifically, we used rate constant values of 6.0 × 10⁴ M⁻¹ s⁻¹, 150 s⁻¹, and 1.0 × 10⁶ M⁻¹ s⁻¹ for reactions 18, 20, and 27 in Table S2, respectively, while Qian et al. (2016)²³ used 1.8 × 10⁵ M⁻¹ s⁻¹, 1800 s⁻¹, and 1.2 × 10⁷ M⁻¹ s⁻¹, respectively. It is important to note that the rate constant values we selected for reactions 18, 20, and 27 in Table S2 implicitly assume that directly determined second-order rate constants describing the oxidation of TFA⁵⁵ and PFBA⁴² by ${\text{SO}}_4^{•-}$ are correct.

Effect of pH on PFEA Oxidation by Sulfate and Hydroxyl Radicals

To determine the effect of pH on PFAS oxidation by ${\text{SO}}_4^{•-}$ and ${}^{•}\text{OH}$, we conducted UV/PS experiments with HFPO–DA, Nafion byproduct 2, and HydroEVE under alkaline conditions similar to those employed in the TOP assay; i.e., 150 mM of sodium hydroxide was added to maintain pH >12 (no additional pH adjustment).⁴¹ For HFPO–DA in UPW (without pH adjustment), $k_{\mathrm{f}}^{\prime }$ was 1.67 × 10⁻⁵ cm² mJ⁻¹ (Table S13). In contrast, HFPO–DA did not exhibit measurable degradation under TOP assay conditions (pH >12) in UV/PS experiments (Figures 3-4 and Figure S18), which agrees with our previous work demonstrating the stability of HFPO–DA in the TOP assay using heat-activated persulfate.²⁰ Under highly alkaline conditions such as those employed in the TOP assay, ${\text{SO}}_4^{•-}$ reacts rapidly with ${\text{OH}}^{-}$ to form ${}^{•}\text{OH}$,^6,41 and ${}^{•}\text{OH}$ cannot degrade HFPO–DA at measurable rates.²⁰ Furthermore, our results suggest PS degradation was independent of pH at our tested pH conditions (Figure S18), which agrees with past work demonstrating pH had a relatively small impact on calculated pseudo-first-order degradation rate constants for persulfate over the pH range of ~3–13.⁵²

Figure 3.

Effect of solution pH on the degradation of HFPO–DA and Nafion byproduct 2 in the UV/PS process. Results were obtained both experimentally and by kinetic modeling. Ultrapure water (UPW): unbuffered, pH was monitored but not adjusted, starting pH ~4, final pH ~2—3; TOP assay: 150 mM of sodium hydroxide was added to maintain pH >12. PFAS initial concentration = 1 mg/L. PS initial concentration = 5 mM. Duplicate experiments were conducted for HFPO–DA in UPW and for Nafion byproduct 2 in UPW and under TOP assay conditions, and data points for all replicate experiments are shown. Approximately 20% of samples were collected in duplicate, and error bars represent the range of duplicate samples.

Figure 4.

Effect of solution pH on PFEA degradation: (A) steady-state concentrations of sulfate and hydroxyl radicals in UV/PS and UV/H₂O₂ experiments; (B) contributions of sulfate and hydroxyl radicals to HFPO–DA degradation in UV/PS experiments; and (C) contributions of sulfate and hydroxyl radicals to Nafion byproduct 2 degradation in UV/PS experiments. Ultrapure water (UPW): unbuffered, pH was monitored but not adjusted, starting pH ~4, final pH ~2—3; TOP assay: 150 mM of sodium hydroxide was added to maintain pH >12. PFEA initial concentration = 1 mg/L. PS initial concentration = 5 mM in UV/PS experiments and H₂O₂ initial concentration = 20 mM in UV/H₂O₂ experiment.

The $k_{\mathrm{f}}^{\prime }$ value for Nafion byproduct 2 in UPW without pH adjustment was 1.16 × 10⁻³ cm² mJ⁻¹ (Table S13). However, under TOP assay conditions, $k_{\mathrm{f}}^{\prime }$ for Nafion byproduct 2 (4.86 × 10⁻³ cm² mJ⁻¹) was 4–5 times that in UPW (Figure 3, Figure S19, Table S13). Based on model results, in UPW, 80% of ${\text{SO}}_4^{•-}$ reacted with PS and 19% of ${\text{SO}}_4^{•-}$ reacted with water; while at TOP assay conditions, > 99.9% of ${\text{SO}}_4^{•-}$ reacted with ${\text{OH}}^{-}$ (Figure S20). Steady-state ${\text{SO}}_4^{•-}$ and ${}^{•}\text{OH}$ concentrations (${[{\text{SO}}_4^{•-}]}_{\text{ss}}$ and ${[{}^{•}\text{OH}]}_{\text{ss}}$, respectively) were calculated via kinetic modeling: ${[{\text{SO}}_4^{•-}]}_{\text{ss}}=6.3\times {10}^{-10}\mathrm{M}$ in UPW and 1.1 × 10⁻¹² M at TOP assay conditions and ${[{}^{•}\text{OH}]}_{\text{ss}}=2.2\times {10}^{-11}\mathrm{M}$ in UPW and 3.9 × 10⁻⁹ M at TOP assay conditions (Figure 4). At TOP assay conditions, ${[{\text{SO}}_4^{•-}]}_{\text{ss}}$ was 2–3 orders of magnitude lower than in UPW, which cannot explain the higher $k_{\mathrm{f}}^{\prime }$ value for Nafion byproduct 2 oxidation observed under TOP assay conditions. Therefore, Nafion byproduct 2 likely reacted with ${}^{•}\text{OH}(k_{•\text{OH}}=1.0\times {10}^6{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}$, obtained from kinetic modeling). We calculated contributions of ${\text{SO}}_4^{•-}$ and ${}^{•}\text{OH}$ to the oxidation of Nafion byproduct 2. In UPW, ${\text{SO}}_4^{•-}$ primarily contributed (>98%) to Nafion byproduct 2 oxidation while at TOP assay conditions, Nafion byproduct 2 oxidation was mainly attributed to ${}^{•}\text{OH}$ (>99%, Figure 4). These calculations, based on the second-order rate constants determined here, explain results in our previous study, in which we found Nafion byproduct 2 was readily oxidized in the TOP assay.²⁰ Moreover, ${[{\text{OH}}^{•}]}_{\text{ss}}$ in our UV/H₂O₂ experiments (5.7 × 10⁻¹² M, Figure 4) was ~3 orders of magnitude lower than at TOP assay conditions, which explains why Nafion byproduct 2 did not measurably degrade in our UV/H₂O₂ experiments. For HydroEVE, the carboxylic acid analogue of Nafion byproduct 2, we observed similar results, i.e., $k_{\mathrm{f}}^{\prime }$ for HydroEVE under TOP assay conditions (1.69 × 10⁻³ cm² mJ⁻¹) was slightly higher than that in UPW (1.21 × 10⁻³ cm² mJ⁻¹) (Table S13). Model results suggest that the disappearance of HydroEVE at TOP assay conditions was primarily attributable to reactions with ${}^{•}\text{OH}$ $(k_{•\text{OH}}=5.0\times {10}^5{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}$, Figure S21) while ${\text{SO}}_4^{•-}$ (>99%) was primarily responsible for the oxidation of HydroEVE in UPW. Results obtained for Nafion byproduct 2 and HydroEVE in UPW without pH control highlight that the presence of ${}^{•}\text{OH}$ had a negligible impact on PFAS oxidation rates and therefore the ${k_{\text{SO}4}}^{•-}$ values we are reporting.

Implications

We assessed the degradation of 18 PFAS including 15 novel PFEAs in ${\text{SO}}_4^{•-}$ and ${}^{•}\text{OH}$-based AOPs. Among the studied PFAS, 6:2 FtS most readily reacted with ${}^{•}\text{OH}$ $[k_{•\text{OH}}=(1.1-1.2)\times {10}^7{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}]$, while Nafion byproduct 2 and HydroEVE, two recently identified polyfluoroalkyl ether acids containing the -O-CFH- moiety, reacted more slowly $[k_{•\text{OH}}=(0.5-1.0)\times {10}^6{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}]$. The ${k_{\text{SO}4}}^{•-}$ values for PFAS that reacted with ${\text{SO}}_4^{•-}$ ranged from (0.85 to 9.5) × 10⁴ M⁻¹ s⁻¹ for PFCAs, PFECAs, and a ClPFPECA to (0.89 to 4.6) × 10⁶ M⁻¹ s⁻¹ for polyfluoroalkyl ether acids containing the -O-CFH- moiety. The reaction rate constants describing PFAS oxidation by ${\text{SO}}_4^{•-}$ are several orders of magnitude smaller than those for ${\text{SO}}_4^{•-}$ reactions with other organic contaminants (e.g., 10⁷ to 10⁹ M⁻¹ s⁻¹).⁵⁶ Thus, high ${\text{SO}}_4^{•-}$ exposures are required for PFAS oxidation. For example, to achieve ~50% parent compound degradation, ${\text{SO}}_4^{•-}$ exposures of (0.7 to >1.8) × 10⁻⁵ M s are needed for PFCAs, PFECAs, and a ClPFPECA (Figure S22). These ${\text{SO}}_4^{•-}$ exposures are ~ 4–5 orders of magnitude larger than those required for 50% degradation of other micropollutants such as atrazine (~1.0 × 10⁻⁹ M s).²⁴ For polyfluoroalkyl ether acids containing the -O-CFH- moiety, ~50% parent compound degradation requires ${\text{SO}}_4^{•-}$ exposures of (1.0–4.0) × 10⁻⁷ M s.

Our study contributes to the understanding of PFAS oxidation in destructive processes in which ${}^{•}\text{OH}$ and/or ${\text{SO}}_4^{•-}$ are present (e.g., nonthermal plasmas, supercritical water oxidation, and electrochemical oxidation) and gives new insights into the behavior of PFEAs in the TOP assay. Further research should focus on identifying pathways and conversion ratios for PFAS oxidation by ${}^{•}\text{OH}$ and ${\text{SO}}_4^{•-}$ to support the development of models that can predict (1) reaction kinetics of a wide range of PFAS and (2) formation of PFAS degradation intermediates/products, including (semi)volatile compounds, as well as matrix-related oxidation byproducts.