Journal of Hazardous Materials 271 (2014) 9–15

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Ferric ion mediated photodecomposition of aqueous perfluorooctane sulfonate (PFOS) under UV irradiation and its mechanism Ling Jin, Pengyi Zhang ∗ , Tian Shao, Shiliang Zhao State Key Joint Laboratory of Environment Simulation and Pollution Control, School of Environment, Tsinghua University, Beijing 100084, China

h i g h l i g h t s

g r a p h i c a l

a b s t r a c t

• Photodecomposition of PFOS under

PFOS-

UV increased 50 times by ferric ions.

• Addition of ferric ion results in reduc-

1.0

tion of PFOS signal in UPLC–MS/MS.

0.8

UV leads to PFOS decomposition.

• The main intermediates, i.e. per-

C /C 0

• Excitation of PFOS–Fe3+ complex by

fluorocarboxylic acids were further decomposed.

Fe3+

[PFOS- Fe3+] hv O2 OH C8F17

UV

0.6

0.2

PFOA-

UV+Fe3+

0.0 0

a r t i c l e

i n f o

Article history: Received 9 September 2013 Received in revised form 21 January 2014 Accepted 28 January 2014 Available online 14 February 2014 Keywords: Perfluorooctane sulfonate Photodecomposition Ferric ion Mechanism

12

24

36 48 Time (h)

60

72

Fe3+

hv

C7F15

... ...

OH

O2

a b s t r a c t Perfluorooctane sulfonate (PFOS) recently has received much attention due to its global distribution, environmental persistence and bioaccumulation. The methods for PFOS decomposition are very limited due to its inertness. In this report we first found the photodecomposition of PFOS under UV was greatly accelerated by addition of ferric ions. In the presence of ferric ion (100 ␮M), PFOS (20 ␮M) decreased to below the detection limit within 48 h, with the rate constant of 1.67 d−1 , which was 50 times higher than that by direct photolysis (0.033 d−1 ). Besides fluoride and sulfate ions, C2–C8 perfluorocarboxylic acids (PFCAs) were identified as the main intermediates. It was found that addition of PFOS into the FeCl3 aqueous solution led to reduction of UV absorption, and the presence of ferric ion reduced the response of PFOS as analyzed by UPLC–MS/MS, which indicated that PFOS formed a complex with ferric ion. The ESR detection indicated that the electronic state of Fe3+ –PFOS complex changed during reaction. And the role of oxygen and hydroxyl radical on the defluorination of PFOS was investigated. Accordingly the mechanism for PFOS photodecomposition in the presence of ferric ion was proposed. © 2014 Elsevier B.V. All rights reserved.

1. Introduction Perfluorinated compounds (PFCs) have been widely used as raw materials and products, such as emulsifying agents in polymer synthesis, waterproof textiles, protective coatings on metals, firefighting foams, and semiconductor etching [1,2], because of their

∗ Corresponding author. Tel.: +86 10 62796840x601; fax: +86 10 62797760/+86 10 62796840. E-mail address: [email protected] (P. Zhang). http://dx.doi.org/10.1016/j.jhazmat.2014.01.061 0304-3894/© 2014 Elsevier B.V. All rights reserved.

C8F17O2

C8F17OH

0.4

specific characteristics including high surface-active effect, high thermal and chemical stability. In recent years, perfluorinated sulfonic and carboxylic acids have received much attention because they have been found to be globally distributed, persistent and bioaccumulative [3]. Perfluorooctane sulfonate (PFOS), one of main PFCs, has been widely detected in humans, wildlife, and the environment [4–7]. Furthermore, its toxicity and likely carcinogenic effect on human and wildlife have been proved [8]. In 2009, PFOS and its salts together with its precursor, perfluorooctane sulfonyl fluoride (PFOSF), were added to Annex B of the Stockholm Convention, calling for restricted production and use in the world.

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L. Jin et al. / Journal of Hazardous Materials 271 (2014) 9–15

PFOS is resistant to conventional oxidation processes which utilize the hydroxyl radical and biological degradation [9–11] due to its high strength of C–F bond. The techniques for PFOS decomposition are very limited now. Moriwaki et al. [12] first showed PFOS can be efficiently decomposed by using high-power sonochemical action, in which PFOS was transformed into PFOA and shorter-chain perfluorocarboxylic acids (PFCAs). Hori et al. [13] reported that PFOS can be reductively decomposed by zerovalent iron in subcritical water. Photolysis is another possible method. However, the decomposition of PFOS (0.37 mM) was not observed under irradiation of a 200 W Hg–Xe lamp (main wavelength 220–460 nm) within 12 h. Furthermore, Yuan et al. reported that PFOS was not photodecomposed when it was used for titania surface modification [14]. Nonetheless, Yamamoto found that PFOS (40 ␮M) in alkaline 2-propanol solution decreased by 76% after 1 d and 92% after 10 d under irradiation of a low-pressure mercury lamp (254 nm, 32 W) [15]. Park et al. [16] reported that PFOS can be reductively photolyzed by aquated electrons generated from the photolysis of iodide by UV light. Transition metal ions are able to coordinate with many natural or anthropogenic organic matters. The environmental photochemistry of these complexes is of crucial importance for pollution abatement because it may cause photodecomposition of persistent pollutants [17,18]. Hori et al. [19] reported that photodecomposition of short-chain PFCAs were induced by Fe3+ together with 220–460 nm light. Wang et al. also reported perfluorooctanoic acid (PFOA) can be decomposed by UV light in the presence of Fe3+ [20]. Perfluoroalkyl sulfonates such as PFOS were quite different from PFCAs in their photochemical characteristics, and they are more difficult to be photolyzed as previously reported. In the present study, we first found the photochemical decomposition of PFOS in the presence of ferric ion under UV irradiation, and its mechanism was discussed. 2. Experimental 2.1. Chemical reagents Heptadecafluorooctanesulfonic acid potassium salt (CF3 (CF2 )7 SO3 K, PFOS) was purchased from Tokyo Kasei Kogyo (Tokyo, Japan). Iron chloride hexahydrate (FeCl3 ·6H2 O) was purchased from Zhongguancun Chemical Co. (Beijing, China). Perfluorooctanoic acid (C7 F15 COOH, PFOA, 96%), perfluoroheptanoic acid (C6 F13 COOH, PFHpA, 96%), perfluorohexanoic acid (C5 F11 COOH, PFHxA, 97%), perfluoropentanoic acid (C4 F9 COOH, PFPeA, 97%), perfluorobutanoic acid (C3 F7 COOH, PFBA, 99%) and perfluoropropionic acid (C2 F5 COOH, PFPA, 97%) were purchased from Aldrich Chemical Co. (New Jersey, USA). Oxygen gas (99.9%) for the reaction was supplied by Beijing Longhui Jingcheng Gas Company (Beijing, China). All aqueous solutions were prepared with high purity water (18.2 M) produced by a Thermo Barnstead Nanopure Diamond water purification system. 2.2. Photocatalytic procedures The photochemical decomposition of PFOS was investigated by using a tubular glass reactor with the inner diameter of 45 mm. A low-pressure mercury lamp (23 W, Cnlight Co. Ltd, China) emits 254 nm UV light (hereafter referred as UV) was placed in the center of the reactor with a quartz envelope. In a typical photochemical reaction experiment, 400 mL mixture of PFOS (20 ␮M) and FeCl3 with various ferric ion concentration was filled into the reactor. The reaction temperature was kept at 25 ◦ C with a cooling water jacket around the reactor, and gas (oxygen or nitrogen) was continuously bubbled into the reactor all the time. Aliquots of sample

were taken at a regular time interval. The outlet gas containing the reaction intermediates flowed through a adsorption tube, in which an adsorbent called Carbopack B was used to adsorb the intermediates. After the end of the reaction, the purge gas was turned off and the adsorption tube was sealed for analysis with ATD–GC/MS.

2.3. Analyses Electrospray ionization mass spectrometry (ESI-MS) was used to identify the intermediate products of PFOS after 24 h of photochemical decomposition in the presence of ferric ion. The full scan (m/z 100–550) mass spectrum was obtained using a triple quadrupole mass spectrometer (Quattro Premier XE, Waters Corp., USA) with an electrospray negative ionization mode. The mobile was 100% methanol, and the flow rate was 0.1 mL/min. The capillary potential was 2.1 kV, and the cone voltage was 35 V. The source temperature was 120 ◦ C, and the desolvation temperature was 400 ◦ C. Concentrations of PFOS and its intermediates were measured using ultra performance liquid chromatography coupled with a triple quadrupole mass spectrometer (UPLC–MS/MS, Quattro Premier XE, Waters Corp., USA) equipped with an Acquity UPLC BEH C18 column (2.1 mm i.d. × 50 mm, 1.7 ␮m particles). To eliminate the interference of ferric ion on the detection of PFOS, prior to analysis, each sample was adjusted to pH 10.5 by NaOH to precipitate ferric ion and filtered through 0.22 ␮m membrane. Aqueous standard solutions were prepared to make external calibration curves for PFOS in the range of 20–800 ␮g/L and C2–C8 PFCAs in the range of 1–200 ␮g/L. And the limits of PFCAs are below 1 ␮g/L. The column temperature was set at 50 ◦ C. The flow rate was maintained at 0.3 mL/min with a mobile phase of eluent A (2 mmol/L ammonium acetate in 100% methanol) and B (2 mmol/L ammonium acetate in 100% water). The eluent gradient started with 30% A for 0.5 min, then was linearly increased to 90% A in 4.5 min, and further increased to 100% A in 1 min and finally back to 30% A in 3 min. The sample volume was 10 ␮L injected with an automatic sampler. Ionization of the analyte was achieved by electrospray in the negative-ion mode. Quantification was performed using multiple reaction monitoring (MRM) of the transitions m/z 499 → 80 (PFOS), m/z 413 → 369 (PFOA), m/z 363 → 319 (PFHpA), m/z 313 → 269 (PFHxA), m/z 263 → 219 (PFPeA), m/z 213 → 169 (PFBA), m/z 163 → 119 (PFPrA) and m/z 113 → 69 (TFA).The optimum mass parameters obtained were as follows: capillary voltage, 2.1 kV; source temperature, 120 ◦ C; desolvation temperature, 400 ◦ C; and desolvation gas flow rate, 650 L/h. The concentrations of F− and SO4 2− was measured with an ion chromatography (Dionex ICS-2000, USA) consisting of a degasser, an autosampler (250 ␮L injection volume), a guard column (IonPac AG11-HC, 4 × 50 mm), a separation column (IonPac AS11-HC, 4 × 250 mm), a column heater (30 ◦ C), and a conductivity detector with a suppressor. The mobile phase was an aqueous solution of KOH (25 mmol/L), and the flow rate was 1.2 mL/min. The suppressor current was set at 75 mA. The qualitative analysis of gas-phase intermediates was conducted on a GC/MS (GCMS-QP2010 Plus, Shimadzu, Japan). The intermediates adsorbed in the adsorption tube were first desorbed by an automatic thermal desorber (Turbo Matrix ATD 650, PerkinElmer, USA) and then detected by GC/MS. The electron spin resonance (ESR) spectra of the reaction solutions were taken at 77 K on an ESR spectrometer (A300, Bruker, German) equipped with a coaxial quartz Dewar under photoirradiation of a mercury lamp. A few milliliters of the sample solutions were poured into quartz tubes, frozen in liquid nitrogen, and subjected to measurement. Typical spectrometer parameters were as follows: center field 2500 G, sweep width 3000 G, microwave

L. Jin et al. / Journal of Hazardous Materials 271 (2014) 9–15

C /C 0

0.6 0.4

UV

F-

UV+Fe3+

SO42-

2 1

0.2 0.0 0

12

24

36

48

60

0 72

Time (h)

(a)

499

PFOS

Relative intensity (%)

3

0.8

Amount of ions (mg/L)

4 1.0

100

200

300

(b)

400

500

PFHxA 313

TFA 113

PFPeA PFPrA

Fig. 1. Time-dependence of PFOS, fluoride ion, and sulfate ion under UV irradiation. [PFOS]0 = 20 ␮M, [Fe(III)]0 = 100 ␮M, pH0 = 3.6.

163

PFBA 213

263

PFHpA 363

PFOS 499

PFOA 413

frequency 9.445 GHz, modulation frequency 100 kHz, and power 10 mW. 3. Results and discussion

11

100

200

300

400

500

m/z Fig. 2. Mass spectra with negative ESI of (a) PFOS standard and (b) its decomposition products after 24 h of reaction. Reaction conditions were the same as in Fig. 1.

3.1. Decomposition of PFOS by UV in the presence of Fe(III)

3.2. Intermediates formed during PFOS decomposition In addition to fluoride and sulfate detected by IC, aqueous solution of PFOS after 24 h of photochemical decomposition in the presence of ferric ion was analyzed by ESI-MS. As shown in Fig. 2, the PFOS degradation solution gave the ion peaks at m/z 413, 363, 313, 263, 213, 163, and 113. These peaks which were not obtained in the mass spectra of the PFOS standard solution were assigned to PFOA and shorter-chain PFCAs, namely, CF3 (CF2 )n COO− (n = 0–5), resulting from dissociation of CF2 from [C7 F15 COO]− . No obvious peaks were observed that could be assigned to compounds containing sulfur, suggesting that C–S bonds in the PFOS molecule completely dissociated.

The time-dependent change of PFOS intermediates, i.e. shorterchain PFCAs are shown in Fig. 3. Though PFOA was identified as one of intermediates, its concentration was too low to be quantified, which can be partly attributed to its fast photodecomposition by Fe3+ /UV as reported in our previous research [20]. And the other reason may relate with the decomposition pathway, which will be discussed in the following section. PFHpA concentration increased rapidly within the first 8 h and reached a maximum value after 32 h of reaction, and then it gradually decreased. PFHxA reached maximum in 24 h and thereafter it almost keep constant, while concentrations of other PFCAs with shorter carbon-chain increased continuously over time. It is worth noting that concentrations of all intermediates (≤ 0.4 mg/L) were very low compared with the initial concentration of PFOS (10 mg/L). These results indicate that PFOS intermediates underwent a further and fast decomposition in a stepwise manner losing a CF2 unit one by one. The mass balance of fluorine element in aqueous phase was calculated, and the results are summarized in Table 1. The total amount of fluorine element in the aqueous phase consists of three parts, i.e. remaining PFOS, shorter-chain PFCAs and F− . After 12 h of reaction, the total amount of fluorine was about 80.3% of the original, while it decreased to 73.7% after 72 h. The imbalance of fluorine mass was largely due to unknown intermediates and purge-off of fluorine-containing products from the aqueous solution. The pH value of PFOS solution was ∼3.6 after addition of iron salt. Since

TFA PFPeA

500 PFCAs (ugL -1)

Fig. 1 shows the photochemical decomposition of PFOS under irradiation of 254 nm UV light. A slight direct photolysis of PFOS (12% after 72 h) by 254 nm UV was observed, which is in agreement with result in the literature [15]. After 100 ␮M ferric ions were added in the reaction solution, the PFOS decomposition was greatly accelerated and it decreased to below the detection limit of UPLC–MS/MS within 48 h. The decomposition of PFOS under UV irradiation followed first order kinetics, with the rate constants of 1.67 d−1 in the presence of 100 ␮M ferric ion and 0.033 d−1 for the direct photolysis, respectively. The decomposition rate constant of PFOS in the presence of ferric ion is nearly 50 times higher than that of direct photolysis, and it is also significantly higher that by UV photolysis in alkaline 2-propanol (0.93 d−1 ) [15]. The energy consumption required to decompose PFOS is 82.3 kJ/␮mol, which is smaller than most values summarized in the literature [21]. These results indicate that the ferric ion mediated photodecomposition is an efficient and mild method to decompose PFOS. In addition, ion chromatography analysis confirmed the simultaneous formation of fluoride and sulfate ions during PFOS photodecomposition. After 72 h of reaction, PFOS with initial concentration of 20 ␮M completely disappeared and the corresponding concentration of fluoride ions increased to 197.9 ␮M. The F index, i.e. −nF− produced /nPFOS decomposed [16], which is used to reflect the mineralization extent of decomposed PFOS, was 9.9 and the defluorination ratio reached 58.2%. Incomplete defluorination implied that there were some fluorinated intermediates formed.

PFPrA PFHxA

PFBA PFHpA

400 300 200 100 0

0

12

24

36 48 Time (h)

60

72

Fig. 3. Time-dependent change of intermediates of PFOS. Reaction conditions were the same as in Fig. 1.

12

L. Jin et al. / Journal of Hazardous Materials 271 (2014) 9–15

Table 1 Mass balance of fluorine element in aqueous phase during PFOS decomposition. F contents from different parts (%)

Reaction time (h)

0 12 24 48 72

Remaining PFOS

PFCAs

Fluoride ion

100 47.2 24.8 3.6 0

0 10.4 12.7 16.8 17.2

0 22.7 37.4 51.6 57.5

Total F contents in aqueous phase (%) 100 80.3 74.9 72.0 73.7

The reaction conditions were the same as in Fig. 1.

3.3. Photodecomposition mechanism

Absorbance (a.u.)

0.3

2

0 20 50 3.0

3.5 4.0 4.5 Retention time (min)

100

c on

of

c rri

( µM

fe

C

1 - PFOS 2 - FeCl 3

3

3 - FeCl 3 + 20 uM PFOS

0.2

4 - FeCl 3 + 100 uM PFOS 5 - FeCl 3 + 200 uM PFOS

4

)

Abundance (a.u.)

Fig. 4 shows the UV–Vis absorption spectra of aqueous solution of PFOS, FeCl3 , and their mixture, respectively. PFOS has week absorption from the deep UV-region to 220 nm and no appreciable absorption from 220 nm to 600 nm. Aqueous solution of FeCl3 has very strong absorption from deep UV region to 400 nm. The mixture of FeCl3 (50 ␮M) and PFOS (20 ␮M) had absorption spectra similar to that of FeCl3 , however the absorption intensity in the UV region were somewhat lower than that of FeCl3 . Moreover, the UV absorption of the mixture further decreased with increase of PFOS concentration. When the PFOS concentration increased to 100 ␮M or 200 ␮M, the absorption from 350 nm to 600 nm increased slightly, while it decreased obviously from deep UV-region to 350 nm. These observations implied that PFOS coordinated with ferric ion to form a complex, which led to the reduction of absorption in the UV region. It is common that ferric ion forms a complex with an organic acid, and the complex may be photolyzed to form Fe2+ and an organic radical through the ligand-to-metal charge transfer. Hori et al. proposed that the complex between Fe3+

and short-chain PFCAs was a prerequisite species for their photodecomposition based on the UV–Vis and ESR spectra of the mixture of iron and PFCAs [19]. Wang et al. [20] also proposed that the complex between Fe3+ and PFOA formed and it accelerated the photodecomposition of PFOA. In addition, it is well known that ferric ion may form an [Fe(SO4)]+ complex with sulfate ions, which can be photolyzed forming ferrous and sulfate radical [24]. The formation of complex between ferric ion and PFOS was further confirmed by the UPLC–MS/MS measurement. As shown in Fig. 5(a), when measured with UPLC–MS/MS, the relative abundance of a PFOS sample with known concentration was significantly lowered after addition of ferric ion, and it decreased with the increase of ferric ion concentration. This observation indicated that ferric ion formed a complex with PFOS, which resulted in the reduction of PFOS concentration as reflected by the UPLC–MS/MS measurement. Moreover, as shown in Fig. 5(b), the relative abundance of a PFOS solution containing ferric ion can be recovered by adjusting the pH from 3.6 to 10, at which ferric ion was completely precipitated and removed from the aqueous phase, indicating that the complexation of ferric ion with PFOS was reversible and the reduction of PFOS response was not due to the adsorption on ferric precipitates.

Abundance (a.u.)

the pKa value of HF was 3.2, a large part of fluoride was present in the form of HF [22], which might be purged off with bubbling. Besides, gaseous products may be formed [23], which was trapped by the Carbopack B adsorbent and analyzed by ATD–GC/MS. Several trace fluorine alkane and perfluorocarboxylic acid such as Cn F2n+2 , Cn HF2n+1 and Cn F2n+1 COOH (n = 2–7) were detected in the gas phase. The imbalance of fluoride mass is commonly reported in literature [22,23], and the recovery of fluorine mass in the present work is close to those reported in literatures. No sulfur-containing organic intermediates were detected either in the aqueous solution or in the gas phase, which confirmed that C–S bonds in the PFOS molecule completely dissociated during photodecomposition. The mass imbalance of sulfur element was possibly due to the specific adsorption of sulfate ion on ferriccolloids. In addition, the presence of ferric ion also interferes the detection of sulfate, leading to underestimation of sulfate ions.

5

0.1

10 6

1

3.6

0.0

3.0

200

300

400

500

600

3.5 4.0 4.5 Retention time (min)

pH

Wavelength (nm) Fig. 4. UV–Vis absorption spectra of aqueous solution of PFOS, FeCl3 and their mixture.

Fig. 5. (a) Effect of ferric ion on PFOS response, [PFOS] = 20 ␮M, pH 3.6; (b) effect of pH on PFOS response, [PFOS] = 20 ␮M, [Fe(III)] = 100 ␮M. All samples were filtrated with 0.22 ␮M filter before UPLC–MS/MS measurement.

L. Jin et al. / Journal of Hazardous Materials 271 (2014) 9–15

g = 4.2

16

Fe 3+

Intensity (a.u.)

13

(µmol/L)

Fe 3+ + PFOS UV irradiation

g = 2.0

Fe

2+

after irradiation

12 8 4 0

1000

1500

2000

2500

3000

3500

4000

Magnetic Field (G) Fig. 6. ESR spectra of an aqueous solution of FeCl3 and aqueous solutions containing PFOS and FeCl3 under various conditions. [Fe(III)]0 = 100 ␮M, [PFOS] = 100 ␮M.

The ESR spectra of the aqueous solution of FeCl3 and the reaction solution containing PFOS and FeCl3 under different conditions were measured at 77 K, and the results are shown in Fig. 6. The ESR spectrum of the aqueous solution of FeCl3 showed a resonance signal at g = 4.2, which can be ascribed to the high-spin electronic state of Fe3+ [25]. The ESR spectrum of the aqueous solution containing FeCl3 and PFOS before irradiation was similar to that of FeCl3 solution, which indicates that the resonance signal of Fe3+ –aqua complex was similar to that of Fe3+ –PFOS complex. However, when the UV lamp was turned on, the ESR spectrum of the reaction solution containing FeCl3 and PFOS showed dramatic changes, a new distinct sharp peak at g = 2.0 appeared. This observation suggests that a dramatic change in the electronic state of the Fe3+ ions in PFOS solution occurred during UV irradiation. The resonance signal at g = 2.0 is ascribable to a low spin ferric complex with PFOS [25,26]. When the lamp was turned off for 30 min, the resonance signal at g = 2.0 was almost unchanged, which indicated that the low spin ferric complex with PFOS was stable at low temperature (77 K). These results mentioned above not only shows that Fe3+ ion can coordinate with PFOS, but also suggests that the electronic state of Fe3+ –PFOS complex changes under UV irradiation. To elucidate the photodecomposition pathway of PFOS induced by Fe3+ , we investigated the influences of reaction atmosphere and the sacrificial agent on the PFOS decomposition at the early reaction stage. As shown in Fig. 7, when the photoreaction was carried out under nitrogen atmosphere, the decomposition of PFOS was almost the same as that under oxygen atmosphere. This result is different from the earlier report [19], which is most likely due to the fact that the concentration of Fe3+ ion was much higher than that of PFOS in our study, the reoxidation step by oxygen, i.e. Fe2+ to Fe3+ , was not essential for PFOS decomposition in the presence of excessive ferric ion.

25

degradation ratio defluorination ratio

Ratio (%)

20 15 10 5 0 O2

N2

O2+H2O2 O2+meth N2+meth

Fig. 7. Effect of atmosphere and HO· on decomposition and defluorination of PFOS in the presence of 100 ␮M Fe(III) at reaction time of 2 h. [PFOS] = 20 ␮M and meth means methanol.

0

12

24

36 48 Time (h)

60

72

Fig. 8. Time-dependent change of concentration of Fe2+ . Reaction conditions were the same as in Fig. 1.

As shown in Fig. 7, when hydrogen peroxide was added in the reaction solution under oxygen atmosphere, both the decomposition and defluorination of PFOS especially the latter was greatly enhanced. In the presence of H2 O2 , a large amount of hydroxyl radicals can be formed during the reaction. Thus, the result above suggests that hydroxyl radical may play an important role on the defluorination although hydroxyl radical itself is inert to PFOS decomposition [21]. Furthermore, when methanol was added to trap hydroxyl radicals under different atmospheres, both the decomposition and defluorination of PFOS decreased much more remarkably under nitrogen atmosphere than that under oxygen atmosphere, indicating the role of oxygen on the PFOS decomposition in the absence of hydroxyl radicals. The amount of Fe2+ during the reaction under oxygen atmosphere was quantified. As shown in Fig. 8, Fe2+ concentration increased rapidly in the first 2 h and reached a maximum value at 6 h of reaction, and then it gradually decreased. The rapid increase of Fe2+ can be attributed to the photolysis of PFOS–Fe3+ complex through the ligand-to-metal charge transfer [17] and the photolysis of Fe(III) in aqueous solution under UV light [24]. The maximum concentration Fe2+ was 15.5 ␮M, less than the PFOS initial concentration (20 ␮M), which indicates that Fe2+ can be re-oxidized to Fe3+ under oxygen atmosphere. And the subsequent gradual decrease is due to the gradual reduction of the PFOS and its intermediates. The balance between Fe3+ and Fe2+ seems to dynamically change with the concentration of PFOS and its perfluorocarboxylic acids, which also form complex with ferric ions and decompose via the excitation of the complex under UV light [20]. Based on the results mentioned above, a mechanism for photodecomposition of PFOS in the presence of ferric ion is proposed in Scheme 1. First, PFOS forms a complex with ferric ion, which is excited by UV light and photolyzed to Fe2+ and an organic radical through the ligand-to-metal charge transfer [17]. The unstable PFOS radical is subsequently desulfonated to form a perfluoroalkyl radical. And the dissociated SO3 • group is further transformed to sulfate (SO4 2− ), which is confirmed by the detection of sulfate as shown in Fig. 1. In the meanwhile, Fe2+ can be recovered to Fe3+ in the presence of oxygen. Under oxygen atmosphere, the perfluoroalkyl radical (C8 F17 •) then may follow two reaction pathways [22,27]. One way is that, it may fast react with oxygen to form a perfluoroalkylperoxy radical, which undergoes a bimolecular radical–radical reaction, yielding two perfluoroalkoxy radicals [27]. As a tertiary oxyl radical, perfluoroalkoxy radical may undergo fast ␤-cleavage to form C7 F15 • and CF2 O. While CF2 O will further hydrolyze to produce fluoride ions, and C7 F15 • may repeat the same reactions as C8 F17 • to be mineralized finally. C8 F17 • + O2 → C8 F17 O2 •

(1)

14

L. Jin et al. / Journal of Hazardous Materials 271 (2014) 9–15

C8F17SO3-

4. Conclusions

Fe3+

The addition of ferric ion greatly enhanced the decomposition and defluorination of PFOS under UV irradiation, which fitted the first-order kinetics. Besides fluoride and sulfate ions, C2–C8 perfluorocarboxylic acids (PFCAs) were identified and quantified as the main intermediates. It is proposed that PFOS forms a complex with ferric ion, which is excited by UV light and photolyzed to Fe2+ and an organic radical through the ligand-to-metal charge transfer. The PFOS radical is unstable and subsequently desulfonated to form a perfluoroalkyl radical. The perfluoroalkyl radical (C8 F17 •) then may react with oxygen or HO•, which leads to PFOS defluorination. Compared with other existing technologies for PFOS decomposition, this method works efficiently under milder conditions, i.e. room temperature and atmospheric pressure. And this well-known method can be easily scaled up and conveniently operated.

[C8F17SO3- Fe3+] hv OH

C8F17

C8F17OH

RFOO

C7F15COF

C8F17O

H2O

C7F15COOFe3+ -

O2

C8F17O2

HF

HF

SO42-

SO3

Fe2+

Fe3+

COF2 3+

[C7F15COO Fe ] hv 2+ Fe

C7F15

...

O2

H2O

CO2+2HF

C7F15COO

Acknowledgments

CO2

This work was supported by National Natural Science Foundation of China (No. 21221004, 21177071), National Basic Research Program of China (2013CB632403) and the Collaborative Innovation Center for Regional Environmental Quality.

C7F15

...

OH

Scheme 1. Proposed mechanism for decomposition of PFOS by UV in the presence of ferric ion.

2C8 F17 O2 • → 2 C8 F17 O • + O2

(2)

C8 F17 O• → C7 F15 • + COF2

(3)

COF2 + H2 O → CO2 + 2 HF

(4)

Alternatively, the perfluoroalkyl radical (C8 F17 •) may react with HO• to form perfluoroalkyl alcohol, i.e. C8 F17 OH. Hydroxyl radicals (HO•) can be formed via the photolysis of Fe(III) [24]. The thermally unstable perfluoroalkyl alcohol undergoes HF elimination to form C7 F15 COF, which then hydrolyzes to yield PFOA. Fe3+ + OH– → FeOH2+

(5)

FeOH2+ + hv → Fe2+ + OH•

(6)

C8 F17 • + OH• → C8 F17 OH

(7)

C8 F17 OH → C7 F15 COF + HF

(8)

C7 F15 COF + H2 O → C7 F15 COOH + HF

(9)

Under oxygen atmosphere, perfluoroalkyl radical (C8 F17 •) may undergo subsequent transformation via the above two pathways. It is hard to say which is more important. However, in the absence of oxygen, i.e. under nitrogen atmosphere, the perfluoroalkyl radical (C8 F17 •) can only continue transformation via the hydroxyl radical pathway. Thus, though hydroxyl radical itself cannot induce the decomposition of PFOS, it can play an important role via involving the subsequent reaction of perfluoroalkyl radical (C8 F17 •), which is resulted from the excitation of the complex of PFOS with ferric ion under UV irradiation. As for the subsequent decomposition of PFOA, it has been reported that PFOA can be decarboxylated to form perfluoroheptyl radicals (C7 F15 •) by UV light in the presence of ferric ion [20,28]. The perfluoroheptyl radical then undergoes the same reactions as C8 F17 · radicals. Accordingly various shorter-chain PFCAs and more fluoride ions are formed. As a result, PFOS is finally decomposed into fluoride ions, sulfate and various shorter-chain PFCAs, which can be further decomposed and completely mineralized via the reactions similar to those of PFOA.

References [1] J.P. Giesy, K. Kannan, Perfluorochemical surfactants in the environment, Environ. Sci. Technol. 36 (2002) 146A–152A. [2] A.B. Lindstrom, M.J. Strynar, E.L. Libelo, Polyfluorinated compounds: past, present, and future, Environ. Sci. Technol. 45 (2011) 7954–7961. [3] K. Kannan, S. Corsolini, J. Falandysz, G. Fillmann, K.S. Kumar, B.G. Loganathan, M.A. Mohd, J. Olivero, N. Van Wouwe, J.H. Yang, K.M. Aldous, Perfluorooctanesulfonate and related fluorochemicals in human blood from several countries, Environ. Sci. Technol. 38 (2004) 4489–4495. [4] G.W. Olsen, T.R. Church, E.B. Larson, G. van Belle, J.K. Lundberg, K.J. Hansen, J.M. Burris, J.H. Mandel, L.R. Zobel, Serum concentrations of perfluorooctanesulfonate and other fluorochemicals in an elderly population from Seattle, Washington, Chemosphere 54 (2004) 1599–1611. [5] J.P. Giesy, K. Kannan, Global distribution of perfluorooctane sulfonate in wildlife, Environ. Sci. Technol. 35 (2001) 1339–1342. [6] J.P. Benskin, L.W.Y. Yeung, N. Yamashita, S. Taniyasu, P.K.S. Lam, J.W. Martin, Perfluorinated acid isomer profiling in water and quantitative assessment of manufacturing source, Environ. Sci. Technol. 44 (2010) 9049–9054. [7] T. Zhang, Q. Wu, H.W. Sun, X.Z. Zhang, S.H. Yun, K. Kannan, Perfluorinated compounds in whole blood samples from infants, children, and adults in China, Environ. Sci. Technol. 44 (2010) 4341–4347. [8] J.R. Thibodeaux, R.G. Hanson, J.M. Rogers, B.E. Grey, B.D. Barbee, J.H. Richards, J.L. Butenhoff, L.A. Stevenson, C. Lau, Exposure to perfluorooctane sulfonate during pregnancy in rat and mouse. I: Maternal and prenatal evaluations, Toxicol. Sci. 74 (2003) 369–381. [9] H.F. Schroder, R.J.W. Meesters, Stability of fluorinated surfactants in advanced oxidation processes—a follow up of degradation products using flow injectionmass spectrometry, liquid chromatography-mass spectrometry and liquid chromatography-multiple stage mass spectrometry, J. Chromatogr. A 1082 (2005) 110–119. [10] R. Dillert, D. Bahnemann, H. Hidaka, Light-induced degradation of perfluorocarboxylic acids in the presence of titanium dioxide, Chemosphere 67 (2007) 785–792. [11] M.M. Schultz, C.P. Higgins, C.A. Huset, R.G. Luthy, D.F. Barofsky, J.A. Field, Fluorochemical mass flows in a municipal wastewater treatment facility, Environ. Sci. Technol. 40 (2006) 7350–7357. [12] H. Moriwaki, Y. Takagi, M. Tanaka, K. Tsuruho, K. Okitsu, Y. Maeda, Sonochemical decomposition of perfluorooctane sulfonate and perfluorooctanoic acid, Environ. Sci. Technol. 39 (2005) 3388–3392. [13] H. Hori, Y. Nagaoka, A. Yamamoto, T. Sano, N. Yamashita, S. Taniyasu, S. Kutsuna, I. Osaka, R. Arakawa, Efficient decomposition of environmentally persistent perfluorooctanesulfonate and related fluorochemicals using zerovalent iron in subcritical water, Environ. Sci. Technol. 40 (2006) 1049–1054. [14] Q.Z. Yuan, R. Ravikrishna, K.T. Valsaraj, Reusable adsorbents for dilute solution separation. 5. Photodegradation of organic compounds on surfactant-modified titania, Sep. Purif. Technol. 24 (2001) 309–318. [15] T. Yamamoto, Y. Noma, S.-I. Sakai, Y. Shibata, Photodegradation of perfluorooctane sulfonate by UV irradiation in water and alkaline 2-propanol, Environ. Sci. Technol. 41 (2007) 5660–5665. [16] H. Park, C.D. Vecitis, J. Cheng, W. Choi, B.T. Mader, M.R. Hoffmann, Reductive defluorination of aqueous perfluorinated alkyl surfactants: effects

L. Jin et al. / Journal of Hazardous Materials 271 (2014) 9–15

[17]

[18] [19]

[20]

[21]

[22]

of ionic headgroup and chain length, J. Phys. Chem. A 113 (2009) 690–696. P. Ciesla, P. Kocot, P. Mytych, Z. Stasicka, Homogeneous photocatalysis by transition metal complexes in the environment, J. Mol. Catal. A: Chem. 224 (2004) 17–33. J. Sima, J. Makanova, Photochemistry of iron(III) complexes, Coord. Chem. Rev. 160 (1997) 161–189. H. Hori, A. Yamamoto, K. Koike, S. Kutsuna, I. Osaka, R. Arakawa, Photochemical decomposition of environmentally persistent short-chain perfluorocarboxylic acids in water mediated by iron(II)/(III) redox reactions, Chemosphere 68 (2007) 572–578. Y. Wang, P. Zhang, G. Pan, H. Chen, Ferric ion mediated photochemical decomposition of perfluorooctanoic acid (PFOA) by 254 nm UV light, J. Hazard. Mater. 160 (2008) 181–186. C.D. Vecitis, H. Park, J. Cheng, B.T. Mader, M.R. Hoffmann, Treatment technologies for aqueous perfluorooctanesulfonate (PFOS) and perfluorooctanoate (PFOA), Front. Environ. Sci. Eng. 3 (2009) 129–151. Q. Zhuo, S. Deng, B. Yang, J. Huang, G. Yu, Efficient electrochemical oxidation of perfluorooctanoate using a Ti/SnO2 –Sb–Bi anode, Environ. Sci. Technol. 45 (2011) 2973–2979.

15

[23] X. Li, P. Zhang, L. Jin, T. Shao, Z. Li, J. Cao, Efficient photocatalytic decomposition of perfluorooctanoic acid by indium oxide and its mechanism, Environ. Sci. Technol. 46 (2012) 5528–5534. [24] L. Deguillaume, M. Leriche, K. Desboeufs, G. Mailhot, C. George, N. Chaumerliac, Transition metals in atmospheric liquid phases: sources, reactivity, and sensitive parameters, Chem. Rev. 105 (2005) 3388–3431. [25] W.S. Szulbinski, ESR and resonance Raman spectroscopic studies of peroxide intermediates derived from iron diazaoxacyclononane, Spectrochim. Acta, Part A 56 (2000) 2117–2124. [26] V.N. Vasyukov, V.P. D’Yakonov, V.A. Shapovalov, E.I. Aksimentyeva, H. Szymczak, S. Piechota, Temperature-induced change in the ESR spectrum of the Fe3+ ion in polyaniline, Low Temp. Phys. 26 (2000) 265–269. [27] Y. Mao, C. Schoneich, K.D. Asmus, Identification of organic acids and other intermediates in oxidative degradation of chlorinated ethanes on TiO2 surfaces en route to mineralization—a combined photocatalytic and radiation chemical study, J. Phys. Chem. 95 (1991) 10080–10089. [28] Y. Wang, P. Zhang, Effects of inorganic anion, isopropanol and TiO2 on photochemical degradation of perfluorooctanoic acid (PFOA) mediated by ferric ion, Acta Chim. Sin. 68 (2010) 345–350.