Defluorination of Aqueous Perfluorooctanesulfonate by Activated Persulfate Oxidation

Abstract

Activated persulfate oxidation technologies based on sulfate radicals were first evaluated for defluorination of aqueous perfluorooctanesulfonate (PFOS). The influences of catalytic method, time, pH and K₂S₂O₈ amounts on PFOS defluorination were investigated. The intermediate products during PFOS defluorination were detected by using LC/MS/MS. The results showed that the S₂O₈ ²⁻ had weak effect on the defluorination of PFOS, while the PFOS was oxidatively defluorinated by sulfate radicals in water. The defluorination efficiency of PFOS under various treatment was followed the order: HT (hydrothermal)/K₂S₂O₈ > UV (ultraviolet)/K₂S₂O₈ > Fe²⁺/K₂S₂O₈ > US (ultrasound)/K₂S₂O₈. Low pH was favorable for the PFOS defluorination with sulfate radicals. Increase in the amount of S₂O₈ ²⁻ had positive effect on PFOS defluorination. However, further increase in amounts of S₂O₈ ²⁻ caused insignificant improvement in PFOS defluorination due to elimination of sulfate radicals under high concentration of S₂O₈ ²⁻. CF₃(CF₂)_nCOOH (n = 0–6) were detected as intermediates during PFOS defluorination. Sulfate radicals oxidation and hydrolysis were the main mechanisms involved in defluorination process of PFOS.

Introduction

Perfluorooctane sulfonate (C₈F₁₇SO₃ ⁻, PFOS) have been widely used in industry as emulsifying agents, surface treatment agents, food containers, fabrics, paper coatings, waxes, fire-fighting foams and polishes [1]. Because of specific characteristics such as unique high surface activity, thermal and acid resistance, and hydro- and lipophobic properties [2], PFOS are considered almost un-degradable in nature. Most conventional degradation technologies are ineffective for degradation of aqueous PFOS since they are inherently recalcitrant to chemical and microbiological treatment [3]–[6]. Even Advanced oxidation processes (AOPs), which utilize the hydroxyl radical, such as UV/O₃, O₃/H₂O₂, UV/H₂O₂ or Fenton's reagent, are also relatively ineffective in PFOS destruction [7]. In fact, the decomposition resistance of PFCs to conventional AOTs is evidenced by the use of PFOS as a surfactant to increase the adsorption of organic pollutants on TiO₂ to obtain the accelerated AOT effects [2]. Direct photolysis [8], VB₁₂/Ti³⁺reduction [9] and permanganate oxidation [2] also show low defluorination efficiencies and incompletely mineralization of PFOS. In contrast, photo reduction [10], [11], sonolysis [12]–[17], sub-critical elemental iron reduction [18], electrochemical oxidation [19] and alkaline ozonation [20] have been proved effective.

Persulfate anion (S₂O₈ ²⁻) is a strong oxidizing agent with a redox potential of 2.0 V, and can be reduced to sulfate anions as shown below (Eq.1) [21]:

(1)

S₂O₈ ²⁻ can be activated to produce sulfate radicals (SO₄ ^•−) with higher redox potential of 2.6–3.1 V, which are very reactive with a wide range of contaminants. There are two general ways of activating S₂O₈ ²⁻: homolysis of the peroxide bond using heat, ultrasound or light (Eq.2) and an oxidation reduction process (analogous to the Fenton reaction) with electron donors, including e⁻ (aq) from radiolysis of water or low-valent metals such as Fe²⁺ and Ag⁺ (Eq.3).

(2)

(3)

Previous studies show that SO₄ ^•− is more prone to redox reactions than hydroxyl radicals (•OH) under neutral conditions, as summarized in the following reactions (Eq.4–6):

(4)

(5)

(6)

SO₄ ^•− is a strong oxidizing radicals and has been known to react with electron-rich moieties through several reaction pathways, including electron exchange, hydrogen abstraction, and direct oxygen transfer. Because of comparative stability of persulfate, strong oxidation characteristics, and pH-independent effectiveness, persulfate and sulfate radical oxidation have been utilized for the oxidative degradation of a number of organics [22]–[26]. It also had been reported that perfluorooctanoic acid (C₇F₁₅COOH, PFOA) and other short-chain perfluorocarboxylic acid (C_nF_2n+1COOH, PFCAs) could be decomposed by sulfate radicals [27]–[30]. The results show that PFOA and other PFCAs were effectively decomposed to F⁻ and CO₂ by using SO₄ ^•−, and almost all of the initial S₂O₈ ²⁻ was transformed to SO₄ ²⁻ during the reaction. However, no studies have been reported on persulfate oxidation of PFOS and perfluoroalkyl sulfonate (C_nF_2n+1SO₃ ⁻).

The present study made first attempt to evaluate the PFOS defluorination effect of S₂O₈ ²⁻ or SO₄ ^•− using the activated K₂S₂O₈ oxidation systems under air atmosphere. The influence of catalytic method (ultraviolet, ultrasound, hydrothermal, ferrous ion), time, pH and K₂S₂O₈ concentration on PFOS defluorination were investigated. The concentrations of intermediate products were detected using LC/MS/MS. Finally, the possible defluorinaiton mechanism of PFOS with sulfate radicals was proposed.

Materials and Methods

Materials

PerfluorooctaneSulfonate (PFOS, C₈F₁₇SO₃K, 98%), Perfluorooctanoic acid (PFOA, C₇F₁₅COOH, 97%), Perfluoroheptanoic acid (PFHpA, C₆F₁₃COOH, 98%), Tridecafluoroheptanoic acid (PFHxA, C₅F₁₁COOH, 98%), perfluoropentanoic acid (PFPA, C₄F₉COOH, 97%) and perfluorobutyric acid (PFBA, C₃F₇COOH, 98%) and pentafluoropropionic acid (PFPrA, C₂F₅COOH, 97%) was purchased from Sigma (USA). Trifluoroacetic acid (TFA, CF₃COOH, 99%) was purchased from Sigma (USA). Perfluorohexanesulfonate (PFHxS, C₆F₁₃SO₃K, 98%), Nonafluorobutanesulfonate (PFBS, C₄F₉SO₃K, 98%) were purchased form TCI(Japan). And ammonium acetate (>99%) and LC-MS grade methanol (HR-GC, >99.99%) were obtained from Merck Chemicals (Germany). Sodium carbonate (Na₂CO₃), Sodium bicarbonate (NaHCO₃), Potassium persulfate (K₂S₂O₈, 98%), Ferrous chloride (FeCl₂, 98%), Hydrogen chloride (HCl) and sodium hydroxide (NaOH) were purchased from Guangzhou Chemical Reagent Factory(Guangzhou, China). Milli-Q water prepared by Millipore with a conductivity of 18.2 mΩ-cm at 25°C was used in all experiments. All aqueous solutions were prepared with ultrapure water prepared using a Thermo Barnstead Nanopure Diamond water purification system.

Methods

PFOS stock solution (100 mg/L) was prepared with ultrapure water and then stored in a refrigerator (5°C) before used. The reactions were conducted in batch experiments in polytetrafluoroetylene (PTFE) reactor, due to their superior non-sorption property. Ultraviolet (UV)-activated K₂S₂O₈ oxidation system was shown in Fig. 1A: Two 15W low-pressure mercury lamps (Direct-immersion, H model, SUNSHINE, China) emitting light at 254 nm were placed in the reactor. The photoelectric conversion efficiency was 35%∼38%. The effective radiating length of the UV tube was 9.5 cm (with a radiating surface area of 113.04 cm²) and UV fluence rate was around 40 mW/cm². During the irradiation, the UV device was switched on for the first 15 min to warm up the UV lamps. The reactor was placed on constant temperature magnetic stirrer at desired temperature during the reaction time. Ultrasound (US)-activated K₂S₂O₈ oxidation system was shown in Fig. 1B: The ultrasonic apparatus consisted of an ultrasonic generator and an oscillator operated at an applied (calorimetric) power of 100 (86)W (TOSO model, China, 40 kHz). The average power density delivered to the reactor was 2W/cm². The solution temperature was maintained by a thermostatic container. A sink was placed in the a thermostatic container to keep the temperature of both ultrasound transducer and reactor. During the US process, the thermostatic container device was switched on for the first about 20 to 30 min to keep desired temperature (20°C). According to results of control experiments, the temperature of reaction solution remained about 20°C, was no more than 21°C during the reaction (measured by using mercury's thermometer). Hydrothermal (HT)-activated K₂S₂O₈ oxidation system: The reactor was placed on constant temperature magnetic stirrer at desired temperature during the reaction time. And the reactor was closed during the reaction process to prevent loss of H₂O and gas product by volatilization. Ferrous ion (Fe²⁺)-activated K₂S₂O₈ oxidation system: the addition quantity of Fe²⁺ catalyst was 3 mM. The reactor was placed on constant temperature magnetic stirrer at desired temperature during the reaction time.

Figure 1. Schematic representation of reactor. (A) UV/K₂S₂O₈. (B) US/K₂S₂O₈.

In all K₂S₂O₈ oxidation systems, an aqueous (Milli-Q) solution (100 mL) of PFOS (100 mg/L; 0.186 mM) were introduced into the reactor. The samples were withdrawn from the reactor at fixed time intervals and quickly quenched in iced water to end reaction [28], [29]. Control experiments also had been conducted to make sure the reactions could be quenched in iced water (Fig. S1). The PFOS and F⁻ concentrations were determined by LC/MS/MS and IC, respectively. All the results discussed in this article were average values form at least two experiments. The initial pH values were adjusted at about 7.0 by using standard NaOH and HCl solutions except for “solution pH effect”.

LC/MS/MS. The concentrations of PFOS and the intermediates were measured using liquid chromatography tandem mass spectrometry. Liquid chromatography was performed on an HPLC apparatus equipped with an Agilent model 1200 series (RRLC/6410B Triple Quard MS, USA) and was used for the LC separation of PFOS. The HPLC separation was carried out at 30°C using a gradient composed of solution A (10 mM ammonium acetate solution adjusted to pH 4 with the addition of acetic acid) and solvent B (acetonitrile, GR). The gradient expressed as changes in solvent B was as follows: 0 to 2.0 min, a linear increase from 10% to 20% B; 2.0 to 4.0 min, 20% to 45% B; 4.0 to 5.0 min, 45% to 60% B; 5.0 to 6.0 min, 60% to 95% B; 6.0 to 6.1 min, 95% B to 10% B, hold at 4 min. The flow rate was 0.3 mL/min. Ionization was achieved by electro spray in the negative-ion mode. The electro spray conditions were as follows: nitrogen curtain gas flow: 10.0 L/min; gas temperature: 350°C, nebulizer pressure: 275.8 kPa (40.0psi), capillary voltage: 1000 V. The LC/MS/MS acquisition was performed in the multiple reaction monitor (MRM) mode by following the reactions m/z 499.0–79.9, which are characteristic of PFOS. The sonochemical products of PFOS were also measured by LC/MS/MS. The LC and electrospray conditions were the same as those used in the PFOS analysis. The LC/MS/MS acquisition was performed in MRM mode by following the reaction m/z 413.0–368.8 (PFOA), m/z 363.0–318.9 (PFHpA), m/z 313.0–268.8 (PFHxA), m/z 263.0–218.9 (PFPA), m/z 213.0–168.9 (PFBA), m/z 163.0–118.9 (PFPrA), m/z 113.0–68.9 (TFA), m/z 398.9–98.9 (PFHxS) and m/z 298.9–79.9 (PFBS). Mass spectrometry acquisition parameters of the intermediates were shown in Table 1.

Table 1. Mass spectrometry acquisition parameters of the intermediates. * Quantitative product ion.

No.	Compound	tR(min)	Prec Ion(m/z)	Prod Ion (m/z)	Frag(V)	Cllision energy(V)	Dwell (ms)
1	PFOA	6.45	413.0	368.8*, 168.8	58	0, 12	80
2	PFHxS	6.61	398.9	98.8*, 79.9	129	40, 52	80
3	PFHpA	6.07	363.0	318.9*, 168.8	56	0, 12	80
4	PFHxA	5.59	313.0	268.8*,168.8	62	0, 12	80
5	PFBS	5.66	298.9	98.8, 79.9*	114	28, 36	80
6	PFPA	4.85	263.0	218.9*	50	0	80
7	PFBA	2.92	213.0	168.9*	55	4	80
8	PFPrA	1.47	163.0	118.9*	60	4	80
9	TFA	1.39	113.0	68.9*	60	4	80

The sample solution and standard solution were injected into the liquid chromatography tandem mass spectrometry, respectively. MRM spectra of standard solutions of nine kinds of PFCs were showed in Fig. 2. The linear regression equations between the concentration of each PFCs and peak area of characteristic peaks were built with the correlation coefficient R² more than 0.99. The linear ranges of each PFCs were 0.001–1 mg/L.

Figure 2. MRM spectra of standard solutions of nine kinds of PFCs.

Ion Chromatography. The concentrations of F⁻ and SO₄ ²⁻ were determined by an anion-chromatography system (Dionex, ICS-1000, USA) consisting of a degasser, a sampler (1 mL injection volume), a guard column (AS4A-SC 4×50 mm, Dionex), a separation column (AS4A-SC 4×250 mm, Dionex), a column heater (30°C), and a conductivity detector with a suppressor. A mixture solution containing 3.5 mM Na₂CO₃ and 1.0 mM NaHCO₃ as the mobile phase was delivered at a flow rate of 1.0 mL·min⁻¹.The lowest detection limit of F⁻ and SO₄ ²⁻ was 0.02 mg/L and 0.1 mg/L, respectively.

Results and Discussion

Defluorinationof PFOS by K₂S₂O₈ oxidation

The defluorination of PFOS (100 mg/L, 0.186 mM) in each system with the initial K₂S₂O₈ amount of 5 g/L (18.5 mM) were shown in Fig. 3.

Figure 3. Time course of PFOS defluorination in activated K₂S₂O₈ oxidation systems.

The S₂O₈ ²⁻ had weak effect on the defluorination of PFOS. The defluorination efficiency (moles of F⁻ formed)/(moles of fluorine content in initial PFOS) of PFOS were only 1.45% in solo K₂S₂O₈ system after 12 h, indicating that S₂O₈ ²⁻ was not an efficient oxidant for degrading PFOS at the room temperature of 20°C. In contrast, the K₂S₂O₈ showed different defluorination efficiencies of PFOS in the presence of catalyst, such as UV, US, HT and Fe²⁺. It should be mentioned that these treatments had poor effects on PFOS defluorination when they are used alone. Thus, it can be concluded that the sulfate radical oxidation was mainly responsible for the defluorination of PFOS in activated K₂S₂O₈ oxidation. On the basis of data (Fig. 3) of this research, the varying tendencies of F⁻ concentration were fitted using first-order kinetics equation. According to the fitting results, the apparent rate constants of F⁻ in UV/K₂S₂O₈, US/K₂S₂O₈, HT/K₂S₂O₈ and Fe²⁺/K₂S₂O₈ were 0.016, 0.004, 0.023 and 0.010 h⁻¹, respectively. Thus, PFOS was defluorinated faster in HT/K₂S₂O₈ system than others. The highest PFOS defluorination efficiency reached 22.52% in HT/K₂S₂O₈ system after 12 h. Temperature has an important influence on free radical and its inducing oxidization reactions. The PFOS move more quickly at high temperature, which increases their probability to react with SO₄ ^•−. Thus, high temperature could not only activate S₂O₈ ²⁻ but also promote the reaction process, and HT/K₂S₂O₈ system had showed better defluorination effect than other activated K₂S₂O₈ oxidation systems.

The formation of SO₄ ²⁻ in each system with the initial K₂S₂O₈ amount of 5 g/L (18.5 mM) was also shown in Fig. 4.

Figure 4. Time course of SO₄ ²⁻ amounts in activated K₂S₂O₈ oxidation systems.

In activated K₂S₂O₈ oxidation, SO₄ ^•− radical anions were formed from S₂O₈ ²⁻, which was then react with PFOS. The SO₄ ²⁻ could be formed through two main ways. One way is that one-electron transfer from reducible agent such as PFOS, S₂O₈ ²⁻ and H₂O to SO₄ ^•− (total concentration = 37 mM) and the other way is sulfonic acid group deprivation to form PFOS (total concentration = 0.186 mM). By contrast, the SO₄ ²⁻ generation form sulfonic acid group deprivation could be negligible, therefore we could assume that all the SO₄ ²⁻ was product of S₂O₈ ²⁻.

Low formation of SO₄ ²⁻ under sole K₂S₂O₈ oxidization treatment, implied that S₂O₈ ²⁻ was the major oxidants during the reaction. The activated K₂S₂O₈ oxidation treatment showed different production of SO₄ ²⁻. Based on these data (Fig. 2), the concentrations of SO₄ ²⁻ were fitted using first-order kinetics equation. According to the fitting results, the apparent rate constants of SO₄ ²⁻ in UV/K₂S₂O₈, US/K₂S₂O₈, HT/K₂S₂O₈ and Fe²⁺/K₂S₂O₈ were 0.162, 0.026, 0.233 and 0.131 h⁻¹, respectively. Li et al. found that K₂S₂O₈ could be activated by ultrasonic irradiation and degradation rates of TCA was increased from 0.0014 to 0.396 min⁻¹ with an increase in the ultrasound frequency from 50 to 400 kHz [31]. However, in present study, low frequency of US (40 KHz) could not effectively decomposes S₂O₈ ²− to sulfate radicals. In contrast, S₂O₈ ²⁻ was fastly converted to SO₄ ²− under Fe2+/K₂S₂O₈ treatment. However, the PFOS defluorination efficiency was only 23.50% after 12 h. One reason for this phenomenon could be that there is lower production of SO₄ ^•− under Fe²⁺/K₂S₂O₈ treatment than UV/K₂S₂O₈ and HT/K₂S₂O₈ treatment. Homolysis of 1 mol S₂O₈ ²⁻ produce two moles SO₄ ^•− using heat or light but only one mole SO₄ ^•− production using Fe²⁺ catalyst (Eq.7):

(7)

Effect of initial solution pH

Effects of initial pH (3∼11) on defluorination of PFOS 0.186 mM (100 mg/L) under each activated K₂S₂O₈ oxidation treatment was shown in Table 2.

Table 2. The calculated pseudo-first-order constants and defluorination efficiency (%) of PFOS with 18.5 mM persulfate in activated K₂S₂O₈ oxidation systems at different initial pH values.

system	Initial pH	Rate constant (per hour)	Reaction time (h)
			1	4	8	12
UV/K2S2O8	3.13	0.018	0.94	5.57	15.91	19.32
	6.91	0.016	0.70	4.27	13.63	17.21
	11.11	0.009	-	3.23	6.89	10.56
US/K2S2O8	3.11	0.005	-	2.47	4.90	5.82
	7.04	0.004	0.25	1.93	3.95	5.22
	10.92	0.003	-	0.98	3.10	3.51
HT/K2S2O8	3.11	0.025	2.18	10.56	19.96	25.71
	6.94	0.023	2.47	9.35	20.33	22.52
	10.95	0.013	0.16	4.87	10.15	15.24
Fe2+/K2S2O8	3.11	0.013	0.44	4.93	12.57	13.88
	7.12	0.010	-	4.36	10.31	10.68
	11.24	0.005	-	1.77	4.02	6.07

During the experimental runs, pH of the PFOS solution dropped as the reaction progressed due to formation of lager amount of SO₄ ²⁻. Change of pH in K₂S₂O₈ solution during reaction process in activated oxidation systems was shown in Fig. 5. With the reaction going on, pH of the solution dropped, and showed different changing tendency with different initial pH in each activated persulfate oxidation system. The pH of US/K₂S₂O₈ system slowly decreased as compared with other three systems. It inferred that US had a weaker activating capability than UV, HT and Fe²⁺. Moreover, there was no obvious change in pH in Fe²⁺/K₂S₂O₈ system in 120 min at initial pH of 11.0. Thus, a fixed initial pH was just maintained at the initial stage of PFOS defluorination in UV/K₂S₂O₈ and HT/K₂S₂O₈ systems, and later reaction was conducted under the acidic condition. Even so, PFOS defluorination was varied with initial pH under each activated K₂S₂O₈ oxidation treatment. As shown in Table 2, the PFOS defluorination efficiency decreased with the increase of initial pH under all treatment, especially for Fe²⁺/K₂S₂O₈.

Figure 5. Change of pH of K₂S₂O₈ solution during reaction process in activated oxidation systems.

(A) UV/K₂S₂O₈; (B) HT/K₂S₂O₈; (C) Fe²⁺/K₂S₂O₈; (D) US/K₂S₂O₈.

The PFOS defluorination efficiency increased with deceasing of solution pH, because additional sulfate radicals were formed due to acid catalyzation [32]. In addition, sulfate radicals may react with OH⁻ to form more hydroxyl radical under alkaline conditions [28], as shown by Eq. (8) below:

(8)

The generated hydroxyl radicals generally attack organic molecules through the H-atom abstraction to form water. However, PFOS contain no hydrogen to be abstracted, therefore hydroxyl radicals have a very poor reactivity with PFOS in aqueous solution and slow down the PFOS decomposition rate. Thus, alkaline conditions were unfavorable for the defluorination of PFOS by sulfate radicals. The pH of PFOS solution decrease during the defluorination process, but high initial concentration of OH⁻ need to consume SO₄ ^•−, which could weaken the oxidant effect of activated K₂S₂O₈ oxidization. Thus, the PFOS defluorination efficiency decreased with the increase of initial pH under all treatment. In addition, compared with UV/K₂S₂O₈ and HT/K₂S₂O₈ system, the more decrease in PFOS defluorination under Fe²⁺/K₂S₂O₈ treatment under alkaline conditions could be explained by the formation of Fe(OH)₂ which cause the deactivation of catalyst.

Effect of S₂O₈ ²⁻ amounts

The effect of initial S₂O₈ ²⁻ amount (0–12.5 g/L) on defluorination of PFOS (100 mg/L) under each activated K₂S₂O₈ oxidation treatment was shown in Fig. 6.

Figure 6. Effect of K₂S₂O₈ amounts on defluorination of PFOS in activated K₂S₂O₈ oxidation systems.

The defluorination efficiency increased under all treatments when the initial amount of S₂O₈ ²⁻ was increased. However, further increase in the initial amount of S₂O₈ ²⁻ resulted in saturation, there is no further increase in the PFOS defluorination efficiency. The saturated concentrations of S₂O₈ ²⁻ were different with the treatment used. Hori et al. reported that sulfate radicals react with S₂O₈ ²⁻ or themselves (Eq.9, 10), In addition, excess SO₄ ^•− can also react with H₂O (Eq.11) [27]:

(9)

(10)

(11)

The SO₄ ^•−/S₂O₈ ²⁻, SO₄ ^•−/SO₄ ^•− and SO₄ ^•−/H₂O reaction could compete with the SO₄ ^•−/PFOS for electrons [28]. Thus, a large amounts of SO₄ ^•− were produced along with high concentration of K₂S₂O₈ under activated K₂S₂O₈ oxidation treatment. The scavenging reactions by SO₄ ^•− themselves and with the remaining S₂O₈ ²⁻ may become significant to decrease PFOS defluorination efficiency. In this case, the systems might have generated much S₂O₈ ^•− and •OH when sulfate radicals too excessive to react with PFOS. Oxidation by S₂O₈ ^•− and •OH are not effective for PFOS and other PFCs because of low rates of reaction at reasonable S₂O₈ ^•− and •OH and their lower oxidation capacity compared with SO₄ ^•−. Therefore, further increase in S₂O₈ ²⁻ concentration resulted in production of larger amounts of SO₄ ^•−. The scavenging reactions by SO₄ ^•− themselves and with the excess S₂O₈ ²⁻ or H₂O might become significant to decrease in PFOS defluorination efficiency.

PFOS defluorination by-products

We measured the concentrations of PFOS and its intermediates under UV/K₂S₂O₈ treatment during UV irradiation. Fig. 7 shows the MRM spectra of assigned ionic compounds in the degraded solutions after 4 h by UV/K₂S₂O₈ oxidation.

Figure 7. MRM spectra of assigned ionic compounds in the degraded solutions after 4h by UV/K₂S₂O₈ oxidation.

The concentrations of PFOS and the intermediates at different irradiation time were determined using external standard method. The results were shown in Fig. 8.

Figure 8. Changes in the concentration of PFOS and intermediates through the PFOS defluorination in UV/K₂S₂O₈ system.

The concentrations of PFOA, PFHpA and PFPA first increased and then decreased during the reaction process, and the concentrations of PFHxA, PFBA, PFPrA and TFA gradually increased with the increasing of time. As intermediates, PFCAs were not only the hydrolysis product but also the reactant of following reactions. The formed perfluorinated carboxylic acid undergo a further degradation and shortening of the perfluorocarbon chain. Based on the reaction dynamics of continuous reaction, the concentration of PFCAs should first increase and then decreased during the reaction process. From the results, it can infer that PFHxA, PFBA, PFPrA and TFA could not reached the maximum at 12 h. On the other hand, there is low concentration of PFHxS and PFBS were maintained at the constant concentration over time. The formation of trace levels of PFHxS and PFBS would be due to the recombination of CF₃(CF₂)₅·and CF₃(CF₂)₃·with SO₃ ⁻ group [8].

Because of the lack of standards of some compounds we could not measure their actual concentration. The unknown fluorines are those not to be confirmed by above analysis. Thus, the irradiation-time dependence of the mass balance of fluorine was also studied in this study. The mass concentration of the inorganic fluorine in the solutions was analyzed using IC at different reaction time. The organic fluorine referred mainly to the fluorine in PFOS and intermediates which had been detected by LC/MS/MS and their mass concentration was calculated from mass ratio between fluorine and parent PFCs (Eq.12).

(12)

The result of analysis of mass balance of fluorine was given in Table 3. The sum of the fluorine in PFOS, PFCAs and inorganic fluorine accounted for greater than 95% of the fluorine from the degraded PFOS at any point in time during UV/K₂S₂O₈ system. According to the results of analyses of the intermediates and the mass balance of fluorine, it can be inferred that PFCAs were the main intermediates.

Table 3. Irradiation-time dependence of fluorine element mass balance during decomposition of PFOS in UV/K₂S₂O₈ systems.

Time (h)	Organic fluorine (mg/L)		Inorganic fluorine (mg/L)	Unknown fluorine (mg/L)
	PFOS	Intermediates	F−
0	60	0	0	0
1	58.34	0.37	0.42	0.87
2	55.42	0.99	1.6	1.99
4	54.46	2.15	2.57	0.82
8	48.96	1.84	8.19	1.01
12	45.72	2.53	10.34	1.41

PFOS defluorination mechanisms

Previous investigations have provided that PFOS could lose electrons to anode or strong oxidizers, but the mechanism was still in controversy. Kimberly et al. found that an alternative oxidation mechanism is the direct transfer of electrons from PFOS to the BDD (Boron-Doped Diamond) anode as shown in Eq.13 [19].

(13)

Lin et al. reported that hydroxyl radicals attack the perfluoro-anion and become quenched into hydroxide ions, leaving perfluorinated radicals for continuing chain reactions that lead to further decomposition [20]. Liu et al. found that the first oxidative attack of PFOS by permanganate may have occurred at the C-S and C-C bonds, which released the -CF₂ units and oxidation product of SO₄ ²⁻ to form the shorter chain perfluoroalkyl sulfonates and perfluorochemical products subject to further transformation reactions such as hydrolysis [2].

As another most extensive and typical of PFCs, PFOA oxidation mechanism with SO₄ ^•− have previously reported. Both Hori et al. and Lee et al. reported that the PFOA oxidation mechanism with SO₄ ^•− was described as follows (Eq.14–18) [27], [28]:

(14)

(15)

(16)

(17)

(18)

Based on the results discussed above, the mechanism of oxidative defluorination of PFOS with sulfate radicals could be illuminated. The defluorination of PFOS was as Fig. 9: The oxidation potential of SO₄ ^•− is 2.5–3.1eV, while the C-F bond in PFOS is the most oxidation resistant bond (E0 = 3.6 eV) and is difficult to be dissociated by SO₄ ^•−. Therefore, the first oxidative attack of PFOS by SO₄ ^•− may have occurred at the C-S bonds, therefore, SO₄ ^•– oxidize PFOS to form C₈F₁₇•, the unstable C₈F₁₇• may react with H₂O to form unstable C₈F₁₇OH, which undergo HF elimination to form C₇F₁₅COF. Moreover, C₇F₁₅COF further undergo hydrolysis, resulting in the formation of short chain PFOA(C₇F₁₅COOH). PFOA is first formed by dissociation of two fluorine of PFOS, and long-chain PFCAs are decomposed stepwise to form short-chain perfluorocarboxylic acid such as PFHpA, PFHxA, PFPA and PFBA are formed successively by similar reaction. Moreover, previous investigations have provided that PFOA and other short-chain perfluorocarboxylic acid could be effectively degraded by sulfate radicals [28]. Hori et al. found that the defluorination of PFOA (1.35 mM) reached 73.8% in UV/K₂S₂O₈ system after 12 h of irradiation. And they also reported that the defluorination of PFOA (0.374 mM) was 77.5% in HT/K₂S₂O₈ system at 80°C after 12 h [33]. In addition, Lee et al. found that short-chain PFCAs are easier to degrade and mineralize than long-chain PFCAs [28]. Thus, the SO₄ ^•−/PFCAs reaction could compete with the SO₄ ^•−/PFOS reaction. Our results show that PFOS was much more stable and more difficult to be defluorinated by sulfate radicals. It can be inferred that SO₄ ^•−/PFCAs could be more likely to occurred than the SO₄ ^•−/PFOS.

Figure 9. Possible defluorination pathways of PFOS in the activated K₂S₂O₈ oxidation systems.

Conclusion

The S₂O₈ ²⁻ showed weak effect on defluorination of PFOS while sulfate radicals could oxidatively decompose PFOS. The defluorination efficiencies were varied with the K₂S₂O₈ oxidation activated by different treatment. The defluorination performance of treatments follow the orders: HT/K₂S₂O₈ > UV/K₂S₂O₈ > Fe²⁺/K₂S₂O₈ > US/K₂S₂O₈. Acidic conditions was favorable for the defluorination of PFOS by sulfate radicals. The defluorination efficiency increased with increases of the initial amount of S₂O₈ ²⁻ under all activated K₂S₂O₈ oxidation treatment. However, further increase in the initial amounts of S₂O₈ ²⁻ did not resulted in further increase in PFOS defluorination efficiency. The main intermediates during the degradation of PFOS were short-chain PFCs. Oxidization and hydrolysiswere the main defluorination mechanism of PFOS. PFOS was first lost electron to sulfate radicals and then desulfonated to C₈F₁₇ radical. During the hydrolytic reaction, long-chain PFCAs are decomposed stepwise to form short-chain PFCAs.

Supporting Information

Figure S1

The effects of low temperature (0°C) on the defluorination of PFOS (0.186 mM) with S₂O₈ ²⁻ (18.5 mM) and Fe²⁺ (3 mM) in ice water were investigated. IC spectra of F⁻, S₂O₈ ²⁻ and SO₄ ²⁻ before and after reaction were shown in Fig. S1. The results showed that no F⁻ and SO₄ ²⁻ was observed after 20 min, and there was less change in concentration of S₂O₈ ²⁻. Thus, it could be considered that S₂O₈ ²⁻ was stable in iced water, and the ice water could quench the formation SO₄ ^•−. Due to the time for sampling was controlled in 10 min, we can considered that this method for quenching reaction could ensure the accuracy of the F⁻ detection.

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Funding Statement

The authors wish to gratefully acknowledge the Project of Science and Technology Excellent Scholar of Zhu Jiang (No. 0501-64) and Fundamental research funds for the National Central Universities (No.2116520) for financial support. The funders had no role in study design, data collection and analysis, decision to publish, or preparation of the manuscript.