Degradation of Perfluorooctanoic Acid by Reactive Species Generated

Letter pubs.acs.org/journal/estlcu

Degradation of Perfluorooctanoic Acid by Reactive Species Generated through Catalyzed H2O2 Propagation Reactions Shannon M. Mitchell, Mushtaque Ahmad,† Amy L. Teel, and Richard J. Watts* Department of Civil and Environmental Engineering, Washington State University, Pullman, Washington 99163-2910, United States S Supporting Information *

ABSTRACT: Perfluorinated compounds, which are environmentally persistent and bioaccumulative contaminants, cannot currently be treated in the subsurface by in situ technologies. Catalyzed H2O2 propagation (CHP) reactions, which generate hydroxyl radical, hydroperoxide anion, and superoxide anion, were investigated for treating perfluorooctanoic acid (PFOA) as a basis for in situ chemical oxidation remediation of groundwater. Using 1 M H2O2 and 0.5 mM iron(III), PFOA was degraded by 89% within 150 min. Hydroxyl radical does not react with PFOA, but systems producing only superoxide promoted 68% PFOA degradation within 150 min. In systems producing only hydroperoxide, the level of PFOA degradation was 80% over 150 min. The generation of near-stoichiometric equivalents of fluoride during PFOA degradation and the lack of detectable degradation products suggest PFOA may be mineralized by CHP. CHP process conditions can be adjusted during treatability studies to increase the flux of superoxide and hydroperoxide to treat PFOA, providing an easily implemented technology.

■

date15−18 as well as these ex situ technologies cannot be applied to in situ subsurface treatment. Because PFOA is highly oxidized, reductive processes such as microbial reductive dehalogenation and zero valent iron have been investigated for in situ PFOA remediation but have been found to be largely ineffective.22,23 Because current PFOA treatment methods are not amenable to in situ application, an alternative process is needed. The effectiveness of catalyzed H2O2 propagations (CHP; modified Fenton’s reagent) for in situ chemical oxidation (ISCO) of contaminated soil and groundwater is well-established.24 CHP is based on the standard Fenton reaction, in which dilute hydrogen peroxide is decomposed by iron(II), generating hydroxyl radical (OH•):

INTRODUCTION Perfluorinated compounds (PFCs) are industrial chemicals that are becoming an environmental and public health concern because of their persistence and potential for bioaccumulation.1 PFCs exhibit immunotoxicity,2 neurotoxicity,3 endocrine disrupting effects,4 and developmental effects.5 Perfluorooctanoic acid (PFOA) is a typical PFC used in surfactants, fire retardants, high-temperature lubricants, and Teflon production.6,7 PFOA differs from many other persistent pollutants in that it is water-soluble; it readily migrates to groundwater8 and bioaccumulates in blood serum and liver rather than fatty tissue.9 PFOA has been detected in air, water, and sediment, human blood, liver, and breast milk, and numerous species of wildlife around the world.10 High concentrations of PFOA are common in point sources and industrial wastewaters; for example, groundwater collected from military bases that used PFCs for fire training contained 6.6 mg/L PFOA.11 PFCs are highly recalcitrant, and their treatment is challenging, especially in soils and subsurface environments. Advanced oxidation processes (AOPs) are ineffective because PFOA has negligible reactivity with hydroxyl radical, the reactant characteristic of AOPs.12−14 Activated persulfate systems have also been investigated for the treatment of PFCs, with effective destruction of PFOA reported in persulfate systems activated by UV light,15 heat,16,17 and microwaves.18 Ex situ processes that have successfully treated PFOA-contaminated groundwater include granular activated carbon,19 reverse osmosis,20 photolysis and photocatalysis,21 and sonification.14 Unfortunately, the activated persulfate processes studied to © 2013 American Chemical Society

Fe 2 + + H 2O2 → Fe3 + + OH• + OH−

(1)

Hydroxyl radical oxidizes many organic pollutants but is unreactive with PFOA (kOH• ≤ 105 M−1 s−1).12−14 Fenton’s reagent is modified for CHP by using high concentrations of hydrogen peroxide and initiators such as soluble iron(III), iron chelates, or minerals, which initiate propagation reactions that generate a variety of reactive oxygen species in addition to hydroxyl radical, including perhydroxyl radical (HO2•), superoxide radical anion (O2•−), and hydroperoxide anion (HO2−): Received: Revised: Accepted: Published: 117

August 2, 2013 October 15, 2013 October 18, 2013 October 18, 2013 dx.doi.org/10.1021/ez4000862 | Environ. Sci. Technol. Lett. 2014, 1, 117−121

Environmental Science & Technology Letters

Letter

Fe3 + + H 2O2 → Fe2 + + HO2• + H+

(2)

OH• + H 2O2 → HO2• + H 2O

(3)

HO2• ↔ O2•− + H+ (pK a = 4.8)

(4)

HO2• + Fe 2 + → Fe3 + + HO2−

(5)

peroxide, was generated in isolation by increasing the pH of hydrogen peroxide (0.5−2 M) to 12.8:34 H 2O2 ↔ HO2− + H+ (pK a = 11.8)

(6)

Two methods were used for generating superoxide alone: (1) a birnessite (γ-MnO2)−hydrogen peroxide system,30 consisting of 5 mg of birnessite and 1 or 2 M hydrogen peroxide at pH 6.8, and (2) a hydroxyl radical-scavenged Fenton-like system at basic pH,34 consisting of 1 or 2 M hydrogen peroxide, 3 M ethanol to scavenge hydroxyl radical, and 5 mM iron(III)EDTA at pH 9.5. Analysis. Samples were extracted using SPE cartridges and eluted in methanol35 for analysis by liquid chromatography and mass spectrometry (LC−MS). Twenty microliters of the eluent was injected in an Agilent 1200 series high-performance liquid chromatograph fit with an Agilent XDB-C18 column (50 mm × 4.6 mm, 1.8 μm particle size). The flow rate was 0.4 mL/min; the mobile phase was 10 mM formic acid in deionized water and 10 mM formic acid in methanol. The eluent isocratic gradient started and ended with 70% methanol for 6.5 min. The column temperature was held at 40 °C. Samples separated chromatographically were analyzed by electrospray ionization using an Agilent 6460 Triple Quad instrument in negative monitoring mode for the PFOA molecular ion (m/z 413) with a nebulizer pressure of 35 psi, a gas flow rate of 10 L/min, a gas temperature of 300 °C, a capillary voltage of 3500 V, and a nozzle voltage of 500 V. Mass spectrometer conditions were changed to detect PFOA degradation products: multiplereaction monitoring mode was changed to positive ion scan mode, and the LC run times were increased to scan for products. Hydrogen peroxide concentrations were determined by iodometric titration.36 Fluoride was quantified in triplicate samples using a Metrohm 761 compact ion chromatograph. A Fisher Accumet AB15 pH meter was used to measure pH. The SAS 9.1.3 package was used to calculate the differences between the experimental data sets. Results are reported as means of triplicate reactions at each time point; error bars represent the standard error of the mean.

Superoxide has high nucleophilic reactivity in organic solvents,25,26 but the conventional view has been that in deionized water superoxide has negligible reactivity with halogenated organic compounds.27,28 However, recent results have documented that the reactivity of superoxide in water is significantly increased in the presence of hydrogen peroxide concentrations characteristic of CHP reactions (0.3−4.0 M)29 and in the presence of solids, including birnessite.30 Hydroperoxide is a strong nucleophile in aqueous systems.31 Superoxide and hydroperoxide generated in CHP reactions have the potential to degrade PFCs; therefore, CHP may be effective in degrading PFOA for in situ applications. The purpose of this research was to investigate the effectiveness of CHP for degrading PFOA and the reactive oxygen species responsible for its degradation.

■

MATERIALS AND METHODS Materials. Perfluorooctanoic acid (96%), hydrogen peroxide (50%), iron(II) perchlorate, and iron(III) perchlorate were purchased from Sigma-Aldrich (St. Louis, MO). Iron perchlorate salts were used to minimize scavenging of hydroxyl radical; however, perchlorate is not typically used for CHP ISCO.24 Sodium hydroxide, iron(III)-EDTA, 2-propanol, and methanol were obtained from J. T. Baker (Phillipsburg, NJ). Purified titanium sulfate was purchased from GFS Chemicals (Columbus, OH). Ethanol was obtained from Fisher Scientific (Fair Lawn, NJ). Birnessite (γ-MnO2) was prepared by dropwise addition of 2 M hydrochloric acid to boiling 1 M potassium permanganate that was being vigorously stirred;32 mineralogical structure was confirmed by X-ray diffraction. SepPak Vac tC18 (6 cm3, 1 g) solid phase extraction (SPE) cartridges were purchased from Waters (Milford, MA). Doubledeionized water (>18 MΩ cm) was prepared using a Barnstead NANO pure II Ultrapure system. Reaction Systems. PFOA degradation reactions were conducted at 20 ± 2 °C in 45 mL capped polypropylene tubes. All reaction mixtures contained 100 μg/L PFOA in a volume of 40 mL. The reaction pH was adjusted using 2 M sulfuric acid or 2 M sodium hydroxide. Reactions were conducted in triplicate; control experiments were conducted in parallel for each system using deionized water in place of hydrogen peroxide. A 5 mL sample was drawn at selected time points, adjusted to pH 7, and applied to SPE cartridges. CHP reactions were conducted with 0.25, 0.5, and 1 M hydrogen peroxide and 0.5 mM iron(III) at pH 3.5. These hydrogen peroxide concentrations are significantly lower than the concentration of 12% (4 M) typically applied in the field.24 Parallel reaction mixtures contained 1 M 2-propanol to scavenge hydroxyl radical. Several well-documented systems were used to generate single reactive oxygen species (hydroxyl radical, superoxide, and hydroperoxide anion). Hydroxyl radical alone was generated by a standard Fenton’s reaction, with 5 mM iron(II) slowly added to 5 mM hydrogen peroxide (pH 2).33 Hydroperoxide anion, the conjugate base of hydrogen

■

RESULTS AND DISCUSSION

PFOA Destruction by CHP Reactions. PFOA loss in CHP reactions over a range of hydrogen peroxide concentrations is shown in Figure 1. PFOA degraded by 68, 85, and 89% over 150 min at initial hydrogen peroxide concentrations of 0.25, 0.5, and 1 M, respectively. PFOA loss was 15% in parallel control systems, likely because of volatilization. These results demonstrate that CHP reactions can rapidly degrade PFOA, although hydroxyl radical has been shown to be unreactive with PFOA (kOH• ≤ 105 M−1 s−1).12−14 To confirm that hydroxyl radical was not involved in PFOA degradation, the hydroxyl radical scavenger 2-propanol was added (Figure 1). In the presence of 1 M 2-propanol, only 24% of the PFOA was lost over 150 min, compared to 89% in the system without 2-propanol, suggesting hydroxyl radical was necessary for PFOA degradation in CHP reactions. Because these results were unexpected, the role of hydroxyl radical in PFOA degradation was further investigated in a Fenton system generating only hydroxyl radical (Figure S1 of the Supporting Information), which confirmed that hydroxyl radical does not react with PFOA. The results of Figure 1 and Figure S1 of the Supporting Information suggest that non-hydroxyl radical 118

dx.doi.org/10.1021/ez4000862 | Environ. Sci. Technol. Lett. 2014, 1, 117−121

Environmental Science & Technology Letters

Letter

loss over 150 min were 22 and 59%, respectively; loss in the control system was 8%. In iron(III)-EDTA systems with 1 and 2 M hydrogen peroxide, the levels of PFOA degradation were 41 and 68%, respectively. These results of superoxide degrading PFOA in aqueous systems are similar to those of Smith et al.29 and Furman et al.,30 who documented superoxide-mediated destruction of the perhalogenated compounds carbon tetrachloride and hexachloroethane in CHP systems. PFOA Degradation by Hydroperoxide. To investigate the role of hydroperoxide in PFOA degradation by CHP reactions, hydroperoxide was generated as the sole reactant in hydrogen peroxide solutions at pH 12.8. The losses of PFOA over 150 min in systems containing 0.5, 1, and 2 M hydrogen peroxide were 27, 41, and 80%, respectively, while the loss in the control system was 14% (Figure 3). As a strong

Figure 1. Degradation of PFOA in CHP systems [100 μg/L PFOA, 0.5 mM iron(III), 0 M (control), 0.25 M, 0.5 M, or 1 M hydrogen peroxide at pH 3.5, 0 or 1 M 2-propanol, total volume of 40 mL, 20 ± 2 °C].

reactive oxygen species may be responsible for PFOA degradation. The lower rate of PFOA degradation in the 2propanol-scavenged system may be due to inhibition of hydroxyl radical activity in reaction 3, resulting in a lower flux of hydroperoxide and superoxide generation in subsequent reactions. Some PFOA degradation still occurred in the 2propanol-scavenged reaction, likely because of the superoxidedriven CHP initiation reaction (reaction 2). This pathway is relatively slow33 but generates a low flux of superoxide, which in turn generates hydroperoxide, either of which may potentially degrade PFOA. Therefore, superoxide and hydroperoxide were investigated as reactants for PFOA destruction. PFOA Degradation by Superoxide. To investigate the role of superoxide in PFOA degradation by CHP reactions, superoxide was generated in hydrogen peroxide systems catalyzed by birnessite at neutral pH30 and by iron(III)EDTA with excess ethanol at pH 9.534 (Figure 2). In birnessite systems with 1 and 2 M hydrogen peroxide, the levels of PFOA

Figure 3. Degradation of PFOA in hydroperoxide anion systems [100 μg/L PFOA, 0 M (control), 0.5 M, 1 M, or 2 M hydrogen peroxide at pH 12.8, total volume of 40 mL, 20 ± 2 °C].

nucleophile, hydroperoxide anion readily attacks electrondeficient carbon atoms, such as those of PFOA.31 These results demonstrate that hydroperoxide rapidly degrades PFOA and are consistent with previous findings of hydroperoxide reactivity in CHP reactions in acidic pH regimes.37 Together, the data of Figures 2 and 3 suggest that CHP reactions degrade PFOA through the activity of both superoxide and hydroperoxide, although reactivity in the different systems cannot be directly compared because each of these systems has different cosolvents, ionic strengths, and solids, which affect the reactivity of superoxide and hydroperoxide. Fluoride Release during PFOA Degradation. Free fluoride concentrations were quantified in hydroperoxide systems containing 0.5, 1, and 2 M hydrogen peroxide at pH 12.8. Free fluoride, as a percentage of total fluoride initially present in the system, is shown in Figure 4a in relation to the percent of initial PFOA degraded. The percent fluoride generated was the same as the percent PFOA degraded in each system (p < 0.05), which suggests the release of essentially all the fluorine atoms of the degraded PFOA. The stoichiometry of fluoride release for the same systems shows approximately 15 mol of fluoride released per mole of PFOA degraded (Figure 4b), confirming release of all or nearly all 15 fluorine atoms from each PFOA molecule degraded, likely because of nucleophilic attack by hydroperoxide.

Figure 2. Degradation of PFOA in superoxide systems [100 μg/L PFOA, 0 M (control), 1 M, or 2 M hydrogen peroxide, 5 mg of birnessite at pH 6.8 or 5 mM of Fe(III)-EDTA at pH 9.5, total volume of 40 mL, 20 ± 2 °C]. 119

dx.doi.org/10.1021/ez4000862 | Environ. Sci. Technol. Lett. 2014, 1, 117−121

Environmental Science & Technology Letters

Letter

Figure 4. Generation of fluoride in the hydroperoxide anion systems of Figure 3 compared to PFOA degradation. (a) Fluoride released as a percent of the total fluoride added to the system as PFOA, compared to PFOA degradation as a percent of the total PFOA. (b) Ratio of moles of fluoride released to moles of PFOA degraded (1 mol of PFOA contains 15 mol of fluorine).

Present Address

Such effective decomposition of PFOA to products containing no fluorine is likely the result of excess hydroperoxide, a strong nucleophile, in relation to the concentration of PFOA. Furthermore, product studies conducted with CHP, superoxide, and hydroperoxide systems revealed no detectable concentrations of PFOA degradation products at any time point. Using 14 C-labeled contaminants, Watts et al.38 documented mineralization of benzo[a]pyrene by CHP; they postulated that mineralization of many of the benzo[a]pyrene products (e.g., keto acids), which are unreactive with hydroxyl radical, was possible through reactions with superoxide and hydroperoxide. The results of this research suggest that PFOA may be similarly mineralized by CHP systems. CHP ISCO technology can be successfully implemented to treat groundwater at contaminated sites. CHP reactions generate superoxide and hydroperoxide, which degrade PFOA with no detectable degradation products and release nearstoichiometric equivalents of fluoride; CHP ISCO therefore has the potential to mineralize PFOA. Practitioners of ISCO can adjust CHP conditions to effectively generate superoxide and hydroperoxide, including conducting reactions at neutral pH with hydrogen peroxide stabilization to increase the radius of influence at injection wells,39 using manganese dioxide as a CHP catalyst,30 or increasing the hydrogen peroxide concentration to drive the propagation reactions that generate superoxide and hydroperoxide.24

■

†

M.A.: ARCADIS U.S., Inc., Ten Friends Lane, Newtown, PA 18940. Notes

The authors declare no competing financial interest.

■

ASSOCIATED CONTENT

S Supporting Information *

Treatment of PFOA in a traditional Fenton system generating only hydroxyl radical (Figure S1) and pseudo-first-order rate constants for PFOA destruction in CHP, superoxide, and hydroperoxide systems (Table S1). This material is available free of charge via the Internet at http://pubs.acs.org.

■

REFERENCES

(1) Post, G. B.; Cohn, P. D.; Cooper, K. R. Perfluorooctanoic acid (PFOA), an emerging drinking water contaminant: A critical review of recent literature. Environ. Res. 2012, 116, 93−117. (2) Dewitt, J. C.; Peden-Adams, M. M.; Keller, J. M.; Germolec, D. R. Immunotoxicity of perfluorinated compounds: Recent developments. Toxicol. Pathol. 2012, 40, 300−311. (3) Onischenko, N.; Fischer, C.; Wan Ibrahim, W. N.; Negri, S.; Spulber, S.; Cottica, D.; Ceccatelli, S. Prenatal exposure to PFOS or PFOA alters motor function in mice in a sex-related manner. Neurotoxic. Res. 2011, 19 (4), 452−461. (4) White, S. S.; Fenton, S. E.; Hines, E. P. Endocrine disrupting properties of perfluorooctanoic acid. J. Steroid Biochem. 2011, 127, 16− 26. (5) Lee, Y. J.; Kim, M. K.; Bae, J.; Yank, J. H. Concentrations of perfluoroalkyl compounds in maternal and umbilical cord sera and birth outcomes in Korea. Chemosphere 2013, 90 (5), 1603−1609. (6) Moody, C. A.; Field, J. A. Perfluorinated surfactants and the environmental implications of their use in fire-fighting foams. Environ. Sci. Technol. 2000, 34, 3864−3870. (7) Giesy, J. P.; Kannan, K. Global distribution of perfluorooctane sulfonate in wildlife. Environ. Sci. Technol. 2001, 35, 1339−1342. (8) Davis, K. L.; Aucoin, M. D.; Larsen, B. S.; Kaiser, M. A.; Hartten, A. S. Transport of ammonium perfluorooctanoate in environmental media near a fluoropolymer manufacturing facility. Chemosphere 2007, 67, 2011−2019. (9) Lau, C.; Anitole, K.; Hodes, C.; Lai, D.; Pfahles-Hutchens, A.; Seed, J. Perfluoroalkyl acids: A review of monitoring and toxicological findings. Toxicol. Sci. 2007, 99, 366−394. (10) Mak, Y. L.; Taniyasu, S.; Yeung, L. W. Y.; Lu, G.; Jin, L.; Yang, Y.; Lam, P. K. S.; Kannan, K.; Yamashita, N. Perfluorinated compounds in tap water from China and several other countries. Environ. Sci. Technol. 2009, 43 (13), 4824−4829. (11) Moody, C. A.; Field, J. A. Determination of perfluorocarboxylates in groundwater impacted by fire-fighting activity. Environ. Sci. Technol. 1999, 33 (16), 2800−2806.

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone: (509) 335-3761. Fax: (509) 335-7632. 120

dx.doi.org/10.1021/ez4000862 | Environ. Sci. Technol. Lett. 2014, 1, 117−121

Environmental Science & Technology Letters

Letter

tion in modified Fenton’s systems. Environ. Sci. Technol. 2004, 38 (20), 5465−5469. (30) Furman, O.; Laine, D. F.; Blumenfeld, A.; Teel, A. L.; Shimizu, K.; Cheng, I. F.; Watts, R. J. Enhanced reactivity of superoxide in water-solid matrices. Environ. Sci. Technol. 2009, 43 (5), 1528−1533. (31) David, M. D.; Seiber, J. N. Accelerated hydrolysis of industrial organophosphates in water and soil using sodium percarbonate. Environ. Pollut. 1999, 105 (1), 121−128. (32) McKenzie, R. M. The synthesis of birnessite, cryptomelane, and some other oxides and hydroxides of manganese. Mineral. Mag. 1971, 38, 493−502. (33) Walling, C. Fenton’s reagent revisited. Acc. Chem. Res. 1975, 8 (1), 125−131. (34) Smith, B. A.; Teel, A. L.; Watts, R. J. Mechanism for the destruction of carbon tetrachloride and chloroform DNAPLs by modified Fenton’s reagent. J. Contam. Hydrol. 2006, 85 (3−4), 229− 246. (35) Cheng, J.; Vecitis, C. D.; Park, H.; Mader, B. T.; Hoffmann, M. R. Sonochemical degradation of perfluorooctane sulfonate (PFOS) and perfluorooctanoate (PFOA) in groundwater: Kinetic effects of matrix inorganics. Environ. Sci. Technol. 2010, 44 (1), 445−450. (36) Schumb, W. C.; Stratterfield, C. N.; Wentworth, R. L. Hydrogen Peroxide; American Chemical Society: Washington, DC, 1955. (37) Hui, Y. Mechanisms of Fenton-like reactions for reduction of refractory compounds. Ph.D. Dissertation, Washington State University, Pullman, WA, 2001. (38) Watts, R. J.; Stanton, P. C.; Howsawkeng, J.; Teel, A. L. Mineralization of a sorbed polycyclic aromatic hydrocarbon in two soils using catalyzed hydrogen peroxide. Water Res. 2002, 36 (12), 4283−4292. (39) Watts, R. J.; Finn, D. D.; Cutler, L. M.; Schmidt, J. T.; Teel, A. L. Enhanced stability of hydrogen peroxide in the presence of subsurface solids. J. Contam. Hydrol. 2007, 91, 312−326.

(12) Wardman, P. Reduction potentials of one-electron couples involving free radicals in aqueous solution. J. Phys. Chem. Ref. Data 1989, 18, 1637−1755. (13) Schroder, H. F.; Meesters, R. J. W. Stability of fluorinated surfactants in advanced oxidation processes: A follow up of degradation products using flow injection-mass spectrometry, liquid chromatography-mass spectrometry and liquid chromatography-multiple stage mass spectrometry. J. Chromatogr., A 2005, 1082 (1), 110− 119. (14) Vecitis, C. D.; Park, H.; Cheng, J.; Mader, B. T.; Hoffmann, M. R. Treatment technologies for aqueous perfluorooctanesulfonate (PFOS) and perfluorooctanoate (PFOA). Front. Environ. Sci. Eng. China 2009, 3 (2), 129−151. (15) Hori, H.; Yamamoto, A.; Hayakawa, E.; Taniyasu, S.; Yamashita, N.; Kutsuna, S. Efficient decomposition of environmentally persistent perfluorocarboxylic acids by use of persulfate as a photochemical oxidant. Environ. Sci. Technol. 2005, 39, 2383−2388. (16) Hori, H.; Nagaoka, Y.; Murayama, M.; Kutsuna, S. Efficient decomposition of perfluorocarboxylic acids and alternative fluorochemical surfactants in hot water. Environ. Sci. Technol. 2008, 42, 7438−7443. (17) Liu, C. S.; Higgings, C. P.; Wang, F.; Shih, K. Effect of temperature on oxidative transformation of perfluorooctanoic acid (PFOA) by persulfate activation in water. Sep. Purif. Technol. 2012, 91, 46−51. (18) Lee, Y.; Lo, S.; Kuo, J.; Hsieh, C. Decomposition of perfluorooctanoic acid by microwave-activated persulfate: Effects of temperature, pH, and chloride ions. Front. Environ. Sci. Eng. China 2012, 6 (1), 17−25. (19) Yu, Q.; Zhang, R. Q.; Deng, S. B.; Huang, J.; Yu, G. Sorption of perfluorooctane sulfonate and perfluorooctanoate on activated carbons and resin: Kinetic and isotherm study. Water Res. 2009, 43 (4), 1150− 1158. (20) Thompson, J.; Eaglesham, G.; Reungoat, J.; Poussade, Y.; Bartkow, M.; Lawrence, M.; Mueler, J. F. Removal of PFOS, PFOA, and other perfluoroalkyl acids at water reclamation plants in South East Queensland, Australia. Chemosphere 2010, 82 (1), 9−17. (21) Cho, I. H. Degradation and reduction of acute toxicity of environmentally persistent perfluorooctanoic acid (PFOA) using VUV photolysis and TiO2 photocatalysis in acidic and basic aqueous solutions. Toxicol. Environ. Chem. 2011, 93 (5), 925−940. (22) Liou, J. S.; Szostek, B.; DeRito, C. M.; Madsen, E. L. Investigating the biodegradability of perfluorooctanoic acid. Chemosphere 2010, 80 (2), 176−183. (23) Hori, H.; Nagaoka, Y.; Mamoto, A.; Sano, T.; Yamashita, N.; Taniyasu, S.; Kutsuna, S. Efficient decomposition of environmentally persistent perfluorooctanesulfonate and related fluorochemicals using zerovalent iron in subcritical water. Environ. Sci. Technol. 2006, 40, 1049−1054. (24) Watts, R. J.; Teel, A. L. Chemistry of modified Fenton’s reagent (catalyzed H2O2 propagationsCHP) for in situ soil and groundwater remediation. J. Environ. Eng. 2005, 131 (4), 612−622. (25) Roberts, J. L., Jr.; Sawyer, D. T. Facile degradation by superoxide ion of carbon tetrachloride, chloroform, methylene chloride, and p,p′-DDT in aprotic media. J. Am. Chem. Soc. 1981, 103 (3), 712−714. (26) Roberts, J. L., Jr.; Calderwood, T. S.; Sawyer, D. T. Oxygenation by superoxide ion of CCl4, FCCl3, HCCl3, p,p′-DDT and related trichloromethyl substrates (RCCl3) in aprotic solvents. J. Am. Chem. Soc. 1983, 105 (26), 7691−7696. (27) Asmus, K. D. In Reactive Oxygen Species in Chemistry, Biology, and Medicine; Quintanilha, A., Ed.; NATO ASI Series; Plenum Press: New York, 1988. (28) Monig, J.; Bahnemann, D.; Asmus, K. D. One electron reduction of CCl4 in oxygenated aqueous solutions: A CCl3O2-free radical mediated formation of Cl− and CO2. Chem.-Biol. Interact. 1983, 47 (1), 15−27. (29) Smith, B. A.; Teel, A. L.; Watts, R. J. Identification of the reactive oxygen species responsible for carbon tetrachloride degrada121

dx.doi.org/10.1021/ez4000862 | Environ. Sci. Technol. Lett. 2014, 1, 117−121