Reductive defluorination of perfluorooctanoic acid by zero-valent iron and zinc: A DFT-based kinetic model

Chemical Engineering Journal 335 (2018) 248–254

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Reductive defluorination of perfluorooctanoic acid by zero-valent iron and zinc: A DFT-based kinetic model

MARK

⁎

Jens Blotevogela, , Robert J. Giraudb, Thomas Borcha,c,d a

Department of Civil and Environmental Engineering, Colorado State University, Fort Collins, CO 80523, USA The Chemours Company, Wilmington, DE 19899, USA c Department of Chemistry, Colorado State University, Fort Collins, CO 80523, USA d Department of Soil and Crop Sciences, Colorado State University, Fort Collins, CO 80523, USA b

G RA P H I C A L AB S T R A C T

A R T I C L E I N F O

A B S T R A C T

Keywords: PFOA PFAS Reductive defluorination Zero-valent iron Permeable reactive barrier

Over the past two decades, groundwater contaminated with chlorinated organic compounds has been successfully remediated via reductive dehalogenation by zero-valent iron. While reductive defluorination of the environmentally persistent perfluorooctanoic acid (PFOA) by Fe0 and Zn0 is thermodynamic favorable, no successful zero-valent metal applications have been reported yet. Consequently, we developed a combined experimental-theoretical approach based on density functional theory to predict the kinetics of reductive PFOA defluorination as a function of reduction potential. The theoretical model was calibrated with experimental results for the reductive dehalogenation of the structurally similar compound tetrachloroperfluorooctanoic acid to account for the typical non-standard conditions in remedial systems, such as increased pH and metal surface passivation. Our model estimate reveals that the half-lives for the first reductive PFOA defluorination step are ∼8 years for Zn0 and ∼500,000 years for Fe0 at metal-to-water ratios typical for permeable reactive barriers. Therefore, we conclusively document that – in contrast to chlorinated solvents – reductive dehalogenation by zero-valent metals is not a viable remedial approach for PFOA unless suitable catalysts are identified.

1. Introduction Poly- and perfluoroalkyl substances (PFASs) are ubiquitous

⁎

anthropogenic contaminants of extraordinary environmental persistence due to the unique strength of the CeF bond [1-4]. Among them are perfluoroalkyl acids (PFAAs) such as perfluorooctanoic acid (PFOA,

Corresponding author at: Department of Civil and Environmental Engineering, 1320 Campus Delivery, Colorado State University, Fort Collins, CO 80523-1320, USA. E-mail address: [email protected] (J. Blotevogel).

http://dx.doi.org/10.1016/j.cej.2017.10.131 Received 29 June 2017; Received in revised form 18 October 2017; Accepted 20 October 2017 Available online 21 October 2017 1385-8947/ © 2017 Elsevier B.V. All rights reserved.

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C− X+ e− ↔ C% + X−

C8HF15O2), which has been used as a processing aid in the production of polytetrafluoroethylene (PTFE) and other fluoropolymers [3,4], but is also formed during biotransformation of fluorotelomer precursors [5]. U.S. EPA has established a drinking water health advisory at 0.07 μg/L and included it in its updated Contaminant Candidate List 4 (CCL4) on November 17, 2016 [6,7]. Since listing on CCL4 may require future regulation under the Safe Drinking Water Act (SDWA), technologies for remediation of groundwater sources of drinking water warrant evaluation. Several advanced oxidation processes such as electrochemical treatment [8-10], photo(cata)lysis [11], and heat-activated persulfate [12] have shown promise in degrading PFOA. Reports on successful reductive defluorination, however, are scarce and limited to applications involving the solvated electron as reductant (EH0 = −2.9 V) [1315]. Hori and co-workers [16] observed decomposition of perfluorooctanesulfonate (PFOS) by Fe0, but only in subcritical water (350 °C, > 20 MPa). Yet, reductive dehalogenation by zero-valent metals (ZVMs) has been successfully applied to remediate groundwater contaminated with chlorinated volatile organic compounds (CVOCs) for over two decades [17-19]. Zero-valent iron is the most commonly used reductant, transforming chlorinated aliphatics either by dihaloelimination, where two vicinal halogen atoms are released under formation of an additional carbonecarbon bond, or by hydrogenolysis, where a halogen is replaced by hydrogen. While the reductive defluorination of PFOA by Fe0 (EH0 = −0.44 V) [19] and Zn0 (EH0 = −0.76 V) [16] is thermodynamically favorable (Fig. S1), it is expected to be slower than reductive dechlorination reactions because fluorine is missing a lowlying vacant d orbital to accept an electron. Given the paramount success of ZVMs in driving reductive dehalogenation reactions of CVOCs, it appears reasonable to assume that many researchers have tried to apply this process to PFOA and other PFASs. We therefore interpret the absence of affirmative reports as a failure of ZVMs to achieve reductive defluorination within practical time frames. Nevertheless, documentation of well-executed, but failed treatment approaches is critical to save further resources and guide better solutions. Consequently, the overarching objective of our study is to estimate the kinetics for the reductive defluorination of PFOA by the widely used reductants Fe0 and Zn0. Here, we have developed a theoretical approach based on density functional theory (DFT) to predict the kinetics of reductive dehalogenation as a function of the reduction potential. This free energy relationship is then calibrated using experimental results for the partially chlorinated, structurally related compound 3,5,7,8-tetrachloroperfluorooctanoic acid (TCPFO, C8HCl4F11O2) to account for non-standard conditions in remedial systems (e.g., permeable reactive barriers, PRBs), such as increased pH and metal surface passivation [18,20]. By developing this modeling approach, we also address another critical environmental research challenge: the quantification of (degradation) reactions that are too slow to practically measure [21].

and a stepwise mechanism (where halogen release occurs after electron transfer, and a linearized reaction coordinate ξ is used [24]):

C− X+ e− ↔ C−X% − ↔ C% + X− need to be evaluated to determine the more favorable one. The appeal of this model is that it renders cumbersome and costly explicit theoretical modeling of a complex metal surface unnecessary. All calculations were performed in Gaussian09 at the M06-2X/6311++G(2d,2p) level of theory, with potential energy minima verified by frequency calculations. Unrestricted methods were used for all openshell systems. Geometries were optimized without constraints, except for bond lengths and angles along the respective reaction coordinates. The M06-2X functional was chosen due to its excellent performance regarding main-group thermochemistry and kinetics involving radical species [25,26], and as best compromise between accuracy and computational time. This functional, along with the 6-311++G(2d,2p) basis set, has been applied successfully in comparable studies [24,27,28]. The implicit SMD model was used to account for aqueous solvation energies [29]. Bond dissociation energies indicated that the carbon-halogen bonds at C2 in (linear) PFOA and C5 in TCPFO would be the most favorable to be broken (Tables S1 and S2), and only C-X dissociations at these positions were further considered to minimize the number of potential energy surfaces to be investigated. 3. Materials and methods 3.1. Batch experiments To investigate the reactivity of fluorochemicals towards zero-valent metals, triplicate batch experiments in 60-mL wide-mouth screw-cap amber HDPE bottles (VWR) were set up, containing either 10 mg/L PFOA (96% purity, > 98% linear, Sigma Aldrich) or 15 mg/L TCPFO (95%, Synquest Laboratories) in 30 mL degassed and N2-purged DI water. Fe0 (50–70 mesh filings, Fisher Scientific) and Zn0 (30 mesh granular, certified ACS, Fisher Scientific) were prepared according to Fennelly & Roberts [30] (acid-washing of Fe0 with 1 M HCl, Zn0 with 0.4% H2SO4, followed by triple DI water and final acetone rinsing) except for drying under 98:2 N2:H2 instead of argon atmosphere. Both ZVMs were added to reach a final concentration of 125 g/L, except for TCPFO/zinc batches, where a lower concentration of 25 g/L Zn0 was used due to fast reduction kinetics. BET analyses revealed specific surface areas (as) of 3.71 m2/g for acid-washed Fe0 and 0.280 m2/g for acid-washed Zn0. Long-term PFOA batches were vented weekly to release pressure build-up from evolving H2 gas. Metal preparation, batch setup, venting and storage (23 ± 2 °C) occurred in an anoxic chamber holding a 98:2 N2:H2 atmosphere. TCPFO batches were sampled within the anoxic chamber. PFOA batches were removed from the chamber and sacrificed at the time of sampling. For organofluorine analysis, 2-mL aqueous samples were filled into 2-mL borosilicate glass LC headspace vials. For ion chromatography analysis, 300-μL aqueous samples were filled into 0.5-mL PolyVials (Dionex). For ion selective electrode analysis, 15-mL aqueous samples were filled into 40-mL borosilicate glass vials and adjusted with total ionic strength adjustment buffer (TISAB, Cole-Parmer) to pH 5.8 according to manufacturer’s specifications. For PFOA, the solid phases (i.e., interior walls of the batch containers, plus zero-valent metals when present) were extracted by adding 10 mL methanol followed by one minute of manual agitation and ten minutes on a vortexer.

2. Theoretical model To predict dehalogenation kinetics as a function of reductant strength, we used an approach suggested by Savéant [22]. In this approach, the crossing points between potential energy surfaces (PES) for the neutral parent compound and the first radical reduction intermediate along a reaction coordinate are determined. The difference in free energy between PES crossing point and ground state of the neutral parent equals the free energy of activation (Δ‡G, Fig. 3). This approach assumes that the first one-electron transfer from the reductant to the halogenated species is rate-limiting and thus the only step to be kinetically considered, independent of whether overall (two-electron) dehalogenation occurs via hydrogenolysis or dihaloelimination [23]. However, both a concerted mechanism (where electron transfer and halogen release occur concurrently, and the carbon-halogen bond is the reaction coordinate):

3.2. Chemical analyses Organofluorine compounds were analyzed with an Agilent 1100 Series liquid chromatograph equipped with a 4.6 × 150 mm, 5 μm ZORBAX Eclipse XDB-C18 column (Agilent) in combination with an 249

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(partial) replacement by hydrogen. Generation of aqueous chloride and fluoride (Figs. S3 and S4) provided further evidence that the aqueous removal of TCPFO in the batch reactors was due to transformation reactions. The detected intermediates suggest that TCPFO was reduced both via hydrogenolysis (e.g. INT1; loss of chlorine, gain of hydrogen) and dichloroelimination (e.g. INT2; loss of two chlorines). Hydrogenolysis intermediates of reductive defluorination were not detected. Consequently, the release of fluoride is likely due to HF elimination after hydrogenolysis at a chlorinated carbon (e.g. INT4; loss of chlorine and fluorine). While dihaloelimination of vicinal fluorine and chlorine cannot be excluded in the formation of INT4, ClF elimination is typically substantially slower than HF elimination [31]. Furthermore, several intermediates revealed loss of up to four fluorine substituents, confirmed by quantification of about the fourfold molar fluoride concentration in relation to the initial TCPFO formation after 5 h of reaction time (Fig. S4), indicating final alkyne or allene formation. No further notable increase in fluoride concentration was detected over the course of 48 h. However, the generated alkyne or allene species may be prone to subsequent oxidative degradation reactions. These findings reveal that substitution of isolated fluorine atoms by chlorine enables rapid defluorination of the carbon backbone by both zero-valent metals. Thus, it can be expected that the environmental persistence of partially chlorinated polyfluorochemicals, which have been used for instance as mist suppressant by the chrome plating industry [32], is lower compared to their perfluorinated counterparts. In the PFOA batches, decreasing concentrations of the organofluorine parent were also observed, but over a substantially longer time period (Fig. 1b). In the control batches not containing any zero-valent metals, a loss of 39% PFOA occurred over the course of the 152-day experiment. Since PFOA is practically completely deprotonated at pH > 3 [33], these losses are likely due to sorption to the HDPE surface of the batch containers rather than volatilization, although methanol extraction was unable to recover any sorbed PFOA as indicated by equal aqueous and total PFOA concentrations. In the Fe0 and Zn0 batches, decreases in PFOA concentration of 50% over 152 days and 100% over 61 days, respectively, were observed. Neither organofluorine intermediates nor aqueous fluoride were detected, implying that PFOA removal was not due to reductive defluorination by ZVMs. However, a reduction of the terminal carboxylate by ZVMs cannot be excluded. Decarboxylated organofluorine compounds with high volatility have previously been reported during reduction of PFAAs by solvated electrons [34]. Nevertheless, partial recovery of PFOA by

Agilent G3250AA MSD TOF system in negative electrospray ionization mode (LC/ESI−-TOF-MS). The injection volume was 50 µL and the flow rate was 0.9 mL/min. Separation was carried out isocratically with 0.1% formic acid in water/acetonitrile (60:40). The capillary and fragmentation voltages were 2.5 kV and 150 V, respectively. Chloride and fluoride were analyzed on a Dionex DX-500 ion chromatograph equipped with a Dionex IonPac AS14A anion-exchange column (4 × 250 mm), using an injection volume of 50 μL. The limits of detection (LOD) were 55.9 μg/L for fluoride and 26.0 μg/L for chloride. For PFOA batch experiments, fluoride was additionally analyzed with a fluoride ion selective electrode (Cole-Parmer) connected to an Orion 4 Star Meter (Thermo Electron) with an LOD of 10 μg/L. For accurate calibration, sacrificial batches had been set up containing 125 g/L of the respective zero-valent metal without fluorochemicals. These batches were sacrificed at the same time of fluorochemical sampling and used to prepare fluoride standards, simulating the background matrix. 4. Results and discussion 4.1. Experimental investigations Batch experiments with zero-valent iron and zinc were conducted to investigate the potential for reductive dehalogenation of PFOA and its partially chlorinated surrogate TCPFO. Fig. 1a shows that the parent TCPFO was readily removed from the aqueous phase. No losses were observed in the control batches not containing any reductants, indicating that sorption or volatilization losses of TCPFO over the 5-h experiment were negligible. In case of both metals, statistical testing of the experimental data revealed a better fit for pseudo-first than pseudo-zero order kinetics (see Supplementary data, Fig. S2), with half-lives t1/2 of 9.5 h for Fe0 and 2.1 min for the stronger reductant Zn0 (Fig. 2). Based on observed pseudo-first order rate constants kobs, surface-area normalized rate constants kSA of 1.58·10−4 L m−2 h−1 for Fe0 and 2.81 L m−2 h−1 for Zn0 were determined using

kSA =

kobs as ϱm

(1)

where ρm is the mass concentration (g L−1) or bulk density of the ZVM. A total of 14 organofluorine intermediates in the Fe0 batches and 20 organofluorine intermediates in the Zn0 batches were detected by accurate mass spectrometry (Table S3), indicating losses of multiple chlorine and fluorine atoms from the carbon backbone as well as their

b) 1.0

1.0

0.8

0.8

0.6

0.6

C/C0

C/C0

a)

Control

0.4

Control AQ

Fe(0) AQ

Zn(0) AQ

Control TOT

Fe(0) TOT

Zn(0) TOT

0.4

Fe(0) Zn(0)

0.2

0.2

0.0

0.0 0

1

2

3

4

0

5

Time (h)

30

60

90

120

150

Time (d)

Fig. 1. a) Normalized aqueous TCPFO concentrations over time in the presence of zero-valent metals. b) Normalized PFOA concentrations over time in the presence of zero-valent metals. Shown are both aqueous concentrations and total concentrations, which are the sum of aqueous and solid phase-extracted species.

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Zn0

Time (min) 0.0

0.5

1.0

1.5

2.0

0.00

1.0

2.0

3.0

4.0

5.0

0.00

y = -0.328x - 0.002 R² = 0.985

-0.20

y = -0.0732x - 0.043 R² = 0.954

-0.10

-0.40

-0.20

ln C/C 0

ln C/C 0

Time (h)

Fe0 0.0

2.5

-0.60 -0.80

-0.40

kobs = 0.328 min-1 t1/2 = 2.1 min kSA = 2.81 L m-2 h-1

-1.00

-1.20

-0.30

-0.50 -0.60

kobs = 0.0732 h-1 t1/2 = 9.5 h kSA = 0.000158 L m-2 h-1

Fig. 2. Pseudo-first order kinetic plots and descriptors for the reduction of TCPFO by zero-valent zinc (left) and zero-valent iron (right). Note that the Zn0 concentration was 25 g/L, while the Fe0 concentration was 125 g/L.

methanol extraction (up to 3% with Fe0 and up to 18% with Zn0, Fig. 1) suggested that sorption does contribute to PFOA removal from the aqueous phase, likely due to electrostatic interactions between the negatively ionized PFOA and the positively charged oxidized metal surfaces [35,36] or ligand exchange of carboxylate groups at metal surface hydroxyl groups [37].

ΔG(aq) = −nFEH

where ΔG (aq) is the aqueous Gibbs free energy, n is the number of electrons transferred (here 1), F is the Faraday constant, and EH is the reduction potential. Consequently, reductive dehalogenation rates will increase with increasing reductant strength. These free energy relationships are plotted in Fig. 4. The figure illustrates that the activation energies required for reductive defluorination of PFOA are higher than for TCPFO at potentials > −1.5 V vs. SHE. It also shows that the difference in free energy of activation ΔΔ‡G(aq) decreases with decreasing reductant potential due to the steeper PES of the TCPFO radical intermediate at shorter C-X bond distance (Fig. 3). The two free energy relationships also reveal a crossing point at strongly reducing conditions at −1.6 V vs. SHE, below which reductive dehalogenation of PFOA would be faster than of TCPFO. However, this may be an artifact as it is well documented that the radical intermediate PESs systematically underestimate activation barriers at short C-X distances (i.e., for extremely strong reducing conditions) [24]. From the free energy relationships shown in Fig. 4, reductive dehalogenation rates can be directly predicted by superimposing the known (standard) potential of a specific reductant [24]. However, two limitations to this direct estimation exist. Firstly, since the PESs calculated in Fig. 3 are for the first one-electron transfer step, the reductant’s one-electron redox potentials (i.e., ionization energies) in the

4.2. Kinetic model development and estimation To estimate the reductive defluorination kinetics of PFOA by ZVMs, we calculated aqueous free energies of activation as a function of reductant potential. The potential energy surfaces for both concerted and stepwise mechanisms revealed that concerted electron transfer was more favorable for both defluorination of PFOA and dechlorination of TCPFO (Tables S4–S7, Figs. S5 and S6). Since reduction potentials in aqueous solution are reported relative to the standard hydrogen electrode (SHE), an ionization potential of 98.6 kcal/mol is added to the PES of the radical intermediate [24]. Its crossing point with the parent’s PES then yields the aqueous free energy of activation Δ‡G (aq) at 0.0 V (Fig. 3). With decreasing reduction potential, the PES of the radical intermediate is lowered relative to the reactant’s PES, leading to decreasing free energies of activation, via:

100

භ

60 50 40 30

ȴ‡G(aq)

70

60 50 40 30 20

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0

0 1.30

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1.70

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TCPFO

Cභ + Cl- ± 0.0 V vs. SHE

80

ȴG(aq) (kcal/mol)

PFOA

70

20

C-Cl + e-

90

-

C + F ± 0.0 V vs. SHE

80

ȴG(aq) (kcal/mol)

100

C-F + e-

90

(2)

2.30

2.50

2.70

ȴ‡G(aq) 1.70

1.90

2.10

2.30

2.50

2.70

2.90

3.10

C-Cl Bond Distance (Å)

C-F Bond Distance (Å)

Fig. 3. Aqueous-phase potential energy surfaces for concerted one-electron transfer (i.e., the more favorable mechanism) to PFOA (left) and TCPFO (right) calculated at the SMD/M062X/6-311++G(2d,2p) level of theory. Shown are the neutral parent compounds (green) and the radical intermediate products at standard hydrogen electrode potential (red). (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)

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40

Table 1 Theoretical and experimental kinetic data used to calculate final aqueous-phase free energies of activation and surface area-normalized reaction rate constants for the reductive defluorination of PFOA by zero-valent iron and zinc. Also shown are predicted half-lives for reductive defluorination of PFOA in our experimental system and in a permeable reactive barrier (PRB) with maximum solid-to-liquid ratio of the respective reductant (i.e., best case scenario).

ȴ‡G(aq) (kcal/mol)

35 30 25 20 15

TCPFO kSA exp. (L m−2 h−1) TCPFO kobs PRB* (h−1) TCPFO Δ‡G (aq) (kcal/mol) PFOA Δ‡G (aq) (kcal/mol) PFOA kobs PRB* (yr−1) PFOA t1/2 PRB* (yr)

10 5 0 -1.50

-1.20

-0.90

-0.60

-0.30

0.00

ȴ‡G(PFOA,aq) (kcal/mol)

Fig. 4. Free energy relationships for reductive defluorination of PFOA (red) and reductive dechlorination of TCPFO (green). (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)

1.58 · 10−4 1.70 22.0 35.7 1.3 · 10−6 520,000

2.81 2600 17.6 29.1 9.1 · 10−2 7.6

therefore not yield any information regarding which stable products are formed. Potential transformation pathways include dihaloelimination, hydrogenolysis, decarboxylation, and several others depending on the respective environmental setting [40].

45 40 35 30 25 20 15 10 5 0

4.3. Model accuracy

0

5

10

15

20

25

The accuracy of our kinetic estimates depends on both the theoretical model and variations in the experimental or remedial environment. In general, kinetic predictions are highly sensitive to even small errors in Δ‡G as 1.36 kcal/mol (at standard temperature of 298.15 K) corresponds to one order of magnitude in reaction rate. While modern density functionals such as M06-2X in combination with implicit solvation models may still produce errors in absolute Δ‡G calculations of 5 kcal/mol [41], the advantage of the quantum mechanical approach outlined here is that effectively only the difference in free energy of activation (ΔΔ‡G) is used for the kinetic estimate due to inclusion of experimental data (Fig. 5). Much like an isodesmic reaction scheme, this will typically lower estimation errors to around 1–2 kcal/mol [42], but nevertheless still a factor of around ten in either direction for reaction rates and half-lives. More accurate ab initio quantum mechanical models are available, but their use for molecules the size of PFOA is still impractical and even challenging for supercomputers due to their excessive computational cost. While the inclusion of experimental or environmental conditions increases the accuracy of the theoretical contribution to the kinetic estimate, it will vice versa impact the estimate’s transferability to other environmental settings. For instance, various water constituents such as (bi)carbonate, chloride, phosphate and nitrate are known to effect longterm ZVM reactivity [39,43]. Furthermore, different product types such as nanoscale ZVMs exhibit varying reactivities, specific surface areas, and bulk densities [44]. We also note that while our model is generally applicable to both other reductive dehalogenation reactions and different reductants, the M06-2X functional used here for PFOA has been reported to lack accuracy for redox reactions involving sulfur [45]. Thus, other levels of theory should be considered in the investigation of, for instance, perfluoroalkyl sulfonic acids.

30

ȴ‡G(TCPFO,aq) (kcal/mol) Fig. 5. Correlation of reductive dechlorination kinetics for TCPFO with reductive defluorination kinetics of PFOA according to Eq. (3).

aqueous phase must be known [38]. For Fe0 and Zn0, these are still unreported. Secondly, ZVM surfaces are passivated due to oxidation over time, and their actual reactivity and pH diverge substantially from aqueous standard conditions [39]. To overcome these limitations, the two free energy relationships in Fig. 4 are equated and the reduction potential cancels out. This leads to a direct correlation of the reductive dechlorination kinetics for TCPFO with the reductive defluorination kinetics of PFOA (Fig. 5):

Δ‡G(PFOA,aq) = −0.0252·(Δ‡G(TCPFO,aq) )2 + 2.49·Δ‡G(TCPFO,aq)−6.93

(3)

Eq. (3) now enables estimation of PFOA defluorination kinetics based on experimental TCPFO kinetics, through which actual redox and pH conditions are accounted for. As an example for a relevant remediation setting, we estimated PFOA kinetics in a PRB (Table 1), using the bulk densities of the ZVMs used in this study (Fe0 2900 g/L; Zn0 3300 g/L). From the kSA values for TCPFO measured above and Eq. (1), kobs within a PRB is first determined. The free energy of activation can then be calculated from the Eyring equation [21]:

kobs h kB T

Zn0

* Bulk densities: Fe0 2.9 kg/L; Zn0 3.3 kg/L.

EH (V)

Δ‡G(aq) = −RTln

Fe0

5. Conclusions (4)

The half-life estimate of 520,000 years for reductive PFOA defluorination even at maximum metal-to-water ratios applied in iron PRBs is a clear statement that this remedial approach will not be effective. Here, we used micron-size zero-valent metals, and increased mass-normalized kinetics may be achieved using nanoscale materials [46]. However, use of nano-size particles for groundwater remediation is largely confined to injection-based applications at much lower massto-water ratios than in PRBs, and their efficiency is typically limited by non-uniform distribution in heterogeneous aquifer settings [44,46].

where R is the universal gas constant, T is absolute temperature, h is Planck’s constant, and kB is the Boltzmann constant. Use of Eq. (3) yields Δ‡G (aq) for the reductive defluorination of PFOA by Fe0 of 35.7 kcal/mol and Zn0 of 29.1 kcal/mol. This equates to PFOA half-lives of 520,000 years in iron PRBs and 7.6 years in zinc PRBs for reductive defluorination as the sole removal mechanism. As stated above in Section 2, the kinetic model only considers the formation of the rate-limiting first one-electron transfer and does 252

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Considering that a t1/2 estimation error on the order of a factor of 10 is possible, it appears possible that reductive PFOA defluorination by Zn0 may be observed over reasonable time frames. However, the halflife estimates reported here refer to the first defluorination reaction only, and several more are necessary for complete defluorination of the molecule. The initial reductive defluorination step may make the molecule susceptible to subsequent oxidation, but documented oxidation reactions of fluorotelomer compounds suggest that this would lead to shorter-chain per- and polyfluorocarboxylates [5,47]. Overall, as demonstrated by the batch experiments, removal of PFOA from the aqueous phase inside a PRB is more likely to occur via sorption to the metal surface and possibly even decarboxylation [34]. While substitution of isolated fluorine atoms by chlorine enables rapid defluorination of the carbon backbone by both Zn0 and Fe0, reductive defluorination of perfluorinated compounds remains kinetically challenged. Remedial treatment of PFOA by ZVM application is therefore not a viable option unless suitable catalysts are identified. Acknowledgements Funding for this work was provided by E. I. du Pont de Nemours and Company and The Chemours Company. We thank Yury Desyaterik for guidance in the use of LC/ESI-TOF-MS and ion chromatography, Jeramy Jasmann for BET specific surface area analyses, and Paul Tratnyek for insightful discussions on electron transfer at zero-valent metal surfaces. Conflict of interest statement: R.J.G. is an employee of The Chemours Company, a maker of fluoroproducts. Appendix A. Supplementary data Supplementary data associated with this article can be found, in the online version, at http://dx.doi.org/10.1016/j.cej.2017.10.131. References [1] B.D. Key, R.D. Howell, C.S. Criddle, Fluorinated organics in the biosphere, Environ. Sci. Technol. 31 (1997) 2445–2454. [2] E.F. Houtz, C.P. Higgins, J.A. Field, D.L. Sedlak, Persistence of perfluoroalkyl acid precursors in AFFF-impacted groundwater and soil, Environ. Sci. Technol. 47 (2013) 8187–8195. [3] G.B. Post, P.D. Cohn, K.R. Cooper, Perfluorooctanoic acid (PFOA), an emerging drinking water contaminant: a critical review of recent literature, Environ. Res. 116 (2012) 93–117. [4] R.C. Buck, J. Franklin, U. Berger, J.M. Conder, I.T. Cousins, P. de Voogt, A.A. Jensen, K. Kannan, S.A. Mabury, S.P.J. van Leeuwen, Perfluoroalkyl and polyfluoroalkyl substances in the environment: terminology, classification, and origins, Integr. Environ. Assess. Manage. 7 (2011) 513–541. [5] M.J.A. Dinglasan, Y. Ye, E.A. Edwards, S.A. Mabury, Fluorotelomer alcohol biodegradation yields poly- and perfluorinated acids, Environ. Sci. Technol. 38 (2004) 2857–2864. [6] S.D. Richardson, S.Y. Kimura, Water analysis: emerging contaminants and current issues, Anal. Chem. 88 (2016) 546–582. [7] United States Environmental Protection Agency, Drinking water contaminant candidate list 4 – Final, 81 Federal Register 81099, November 17, 2016. [8] C.E. Schaefer, C. Andaya, A. Urtiaga, E.R. McKenzie, C.P. Higgins, Electrochemical treatment of perfluorooctanoic acid (PFOA) and perfluorooctane sulfonic acid (PFOS) in groundwater impacted by aqueous film forming foams (AFFFs), J. Hazard. Mater. 295 (2015) 170–175. [9] C.E. Schaefer, C. Andaya, A. Burant, C.W. Condee, A. Urtiaga, T.J. Strathmann, C.P. Higgins, Electrochemical treatment of perfluorooctanoic acid and perfluorooctane sulfonate: Insights into mechanisms and application to groundwater treatment, Chem. Eng. J. 317 (2017) 424–432. [10] J. Niu, Y. Li, E. Shang, Z. Xu, J. Liu, Electrochemical oxidation of perfluorinated compounds in water, Chemosphere 146 (2016) 526–538. [11] 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. China 3 (2009) 129–151. [12] C.S. Liu, C.P. Higgins, F. Wang, K. Shih, Effect of temperature on oxidative transformation of perfluorooctanoic acid (PFOA) by persulfate activation in water, Sep. Purif. Technol. 91 (2012) 46–51. [13] L. Huang, W. Dong, H. Hou, Investigation of the reactivity of hydrated electron toward perfluorinated carboxylates by laser flash photolysis, Chem. Phys. Lett. 436 (2007) 124–128. [14] Y. Qu, C. Zhang, F. Li, J. Chen, Q. Zhou, Photo-reductive defluorination of perfluorooctanoic acid in water, Water Res. 44 (2010) 2939–2947.

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