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Reductive Defluorination of Branched Per- and Polyfluoroalkyl Substances with Cobalt Complex Catalysts Jinyong Liu, Daniel J Van Hoomissen, Tianchi Liu, Andrew Maizel, Xiangchen Huo, Seth R. Fernández, Changxu Ren, Xin Xiao, Yida Fang, Charles Schaefer, Christopher P. Higgins, Shubham Vyas, and Timothy J. Strathmann Environ. Sci. Technol. Lett., Just Accepted Manuscript • DOI: 10.1021/acs.estlett.8b00122 • Publication Date (Web): 28 Mar 2018 Downloaded from http://pubs.acs.org on March 28, 2018

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Environmental Science & Technology Letters

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Reductive Defluorination of Branched Per- and Polyfluoroalkyl Substances

2

with Cobalt Complex Catalysts

3 4 5

Jinyong Liu,*,†,‡ Daniel J. Van Hoomissen,§ Tianchi Liu,† Andrew Maizel,‡ Xiangchen Huo,‡ Seth R. Fernández,† Changxu Ren,† Xin Xiao,#,‡ Yida Fang,‡ Charles Schaefer,∆ Christopher P. Higgins,‡ Shubham Vyas,*,§ and Timothy J. Strathmann*,‡

6 7 8 9 10 11

†

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Abstract

13

This

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polyfluoroalkyl substances (PFASs) undergoing cobalt-catalyzed reductive defluorination

15

reactions. Experimental results and theoretical calculations reveal correlations between the extent

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of PFAS defluorination, the local C−F bonding environment, and calculated bond dissociation

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energies (BDEs). In general, BDEs for tertiary C−F bonds < secondary C−F bonds < primary

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C−F bonds. A tertiary C−F bond adjacent to three fluorinated carbons (or two fluorinated

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carbons and one carboxyl group) has a relatively low BDE that permits an initial defluorination

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to occur. Both a biogenic cobalt-corrin complex (B12) and an artificial cobalt-porphyrin complex

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(Co-PP) are found to catalytically defluorinate multiple C−F bonds in selected PFASs. In general,

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Co-PP exhibits higher initial rate of defluorination than B12. Neither complex induced significant

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defluorination in linear perfluorooctanoic acid (PFOA; no tertiary C−F bond) or a perfluoroalkyl

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ether carboxylic acid (tertiary C−F BDEs too high). These results open new lines of research,

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including (1) designing branched PFASs and cobalt complexes that promote complete

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defluorination of PFASs in natural and engineered systems, and (2) evaluating potential impacts

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of branched PFASs in biological systems where B12 is present.

Department of Chemical and Environmental Engineering, University of California, Riverside, California 92521, United States ‡ Department of Civil and Environmental Engineering and §Department of Chemistry, Colorado School of Mines, Golden, Colorado 80401, United States # Department of Environmental Science, Zhejiang University, Hangzhou 310058, China ∆ CDM Smith, 110 Fieldcrest Avenue, No. 8, Sixth Floor, Edison, New Jersey 08837, United States

study

investigates

structure-reactivity

relationships

1 ACS Paragon Plus Environment

within

branched

per-

and


28 29

Introduction

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Since the original development in the 1940s, per- and polyfluoroalkyl substances (PFASs) have

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been widely used in industrial and consumer products.1,

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environment has been extensively documented.2-6 Substantial research efforts on the two C8

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legacy compounds, perfluorooctanoic acid (PFOA) and perfluorooctanesulfonic acid (PFOS),

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have confirmed a variety of adverse health effects,7 leading to the phase-out of ≥C8 PFASs in

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North America and Europe8 and the USEPA’s recent issuance of drinking water health advisory

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levels for PFOA and PFOS.9,

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sulfonic acids, perfluoroalkyl ether carboxylic acids, and other novel structures) have already

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been detected in aquatic environments and are considered recalcitrant.11-15 Recent studies have

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also indicated that the emerging PFASs exhibit variable toxicities and environmental

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mobilities.11, 16-20 Still, knowledge on emerging PFASs remains limited.11, 21

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Despite having received much less attention than linear PFASs, branched PFASs (Figure 1) have

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also been extensively applied and detected in the environment. For example, perfluoro-3,7-

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dimethyloctanoic acid (PFMe2OA) serves, along with linear PFASs, as an ingredient of well

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treatment fluids.22 This compound is also on the “STANDARD 100 by OEKO-TEX” list,

10

2

Detection of PFASs in the global

New alternative PFASs (e.g., shorter chain carboxylic and


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indicating its wide application in textile production,23 and it has been detected in European water

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bodies.24 Perfluoroethylcyclohexane sulfonate (PFECHS)25 also contains two branched carbons

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on the cyclic structure, and its detection in Canadian Arctic lakes has been attributed to its use in

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aircraft anti-erosion fluid.26 Industrial PFOS products often contain variable fractions of

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branched isomers, which have been detected in both environmental waters and human tissues.27,

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28

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backbone.29, 30 Recently reported perfluoroalkyl ether carboxylic acid pollutants such as GenX12

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and its longer analog15 also contain branched structures.

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The recalcitrance of PFASs to biological and chemical degradation is attributed to the high

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stability of C−F bonds.31 However, Ochoa-Herrera et al.32 observed >70% fluoride ion (F−)

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release from a mixture of branched PFOS isomers by reaction with B12 (a corrin-CoIII complex,

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catalyst precursor) and TiIII citrate (reductant). More recently, Park et al.33 reported cleavage of

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multiple sp3 C−F bonds from analytical standards of mono-branched PFOS with a –CF3 at the 3-,

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4-, 5-, or 6-position (see the 6-brPFOS structure in Figure 1) using B12 and nanosized Zn0 as an

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alternative reductant. Since B12 is an essential component for microorganisms and animals, these

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findings suggest that reductive defluorination of branched PFASs could occur in natural

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environments or biological systems. Furthermore, if the initial defluorination replaces one or

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more F atoms with H atoms, additional defluorination mechanisms could be triggered. For

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example, HF elimination from fluorotelomers (i.e., −CH2−CF2− into –CH=CF−) has been

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observed both in vivo and abiotically in the environment,34, 35 and such a transformation could

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significantly alter the toxicity of the PFASs.34 Hence, it is both scientifically intriguing and

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practically imperative to further investigate critical structural factors determining Co-mediated

Typical branched PFOS isomers contain one or two −CF3 branching from the perfluorinated



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defluorination of branched PFASs. The findings of this work will help evaluate the fate of

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emerging PFASs and aid in the design of less persistent and toxic PFASs.

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Materials and Methods

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Chemicals and solution preparation. PFASs (SynQuest Laboratories), B12 (Alfa Aesar) and

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other cobalt species (Sigma-Aldrich), 12% TiCl3 solution (Acros Organics), and other chemicals

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(Fisher Chemical) were used as received. Detailed chemical information (Table S1) and

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preparation of stock solutions of PFASs, cobalt species, and TiIII citrate in carbonate buffer36 are

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described in the Supporting Information (SI).

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Defluorination reaction and sample analysis. The procedure was modified from Ochoa-

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Herrera et al.32 In an anaerobic glove bag, a series of 9-mL serum bottles were loaded with 4 mL

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of solution containing specific PFASs (0.1 mM), TiIII citrate (~36 mM) with carbonate buffer

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(~40 mM), and cobalt catalyst (0.25 mM). More operational details are provided in the SI. Each

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serum bottle was sealed and transferred to a 70°C oven. At designated reaction times, individual

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reactors were sacrificed for analysis. Each reactor was used for a single measurement, such that a

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typical reaction series of five time points (0, 1, 3, 7, and 15 days) began with five replicates of

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each reaction mixture. Each reaction series was repeated at least twice. Due to the interference of

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TiIII citrate matrix to the fluoride-selective electrode, F− release was analyzed by ion

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chromatography.

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chromatography−quadrupole time-of-flight mass spectrometer (LC−QToF-MS). Details of

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instrumental analyses are provided in SI.

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C‒F bond dissociation energy (BDE) calculation. BDEs of C‒F bonds in PFAS structures

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(anion for carboxylic acids) were calculated using Grimme’s GD3-BJ empirical dispersion

Degradation

of

parent

PFASs


was

analyzed

by

liquid

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corrected37 hybrid density functional theory (DFT) at the B3LYP/6-311+G(2d,2p) level of

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theory.38-41 Truhlar’s SMD solvent model was chosen to implicitly model the aqueous

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environment.42 The BDE for each bond was calculated through Eq. (1):

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∗ ∗ ∗ = [ ] + −

(1)

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where H* represents the enthalpy of formation.

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Results and Discussion

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Initial experiments tested the defluorination of PFMe2OA (1), a well-defined branched structure,

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at rate-optimized reaction conditions reported previously (pH 9.0 and 70°C).32 Defluorination

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was evaluated by the concentration ratio of released F− ions to the F initially present in the PFAS

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substrate. Significant defluorination occurred only in the presence of both TiIII and B12 (Figure

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S1). Figure 2a shows that a maximum of 85% defluorination from 1 was achieved in 7 d.

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Because LC−QToF-MS analysis indicated complete degradation of 1, this high defluorination

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ratio corresponds to an average of sixteen of the nineteen F atoms within the PFMe2OA structure

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being released as F−. According to the mechanisms proposed for Co-catalyzed dehalogenation

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reactions (X = Cl and Br),43-45 TiIII reduces the CoIII in B12 to CoI, which then interacts with

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either the carbon or halide (X) to cleave the C−X bonds. Assuming a similar mechanism for the

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defluorination reactions, the turnover number (TON) for each Co center is estimated to be 6.5 for

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the reaction with 1, demonstrating the catalytic nature of reaction. The TON could be further

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increased for at least 10 times (i.e., TON=65) because elevating the concentration of 1 from 0.1

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mM to 1.0 mM still achieved the same defluorination ratio. The reaction slowed at room

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temperature (21 ± 2°C), but still resulted in at least 44% defluorination within 8 mo (Figure S2),



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suggesting the potential for slow defluorination of branched PFASs in low redox potential

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natural environments, where microorganisms employ B12 for dechlorination.46

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To elucidate the relationship between PFAS structure and susceptibility to B12-catalyzed

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defluorination, commercially available PFASs with different “branched structures” were reacted

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with B12 and TiIII (Figure 2). It was soon realized that a simple classification as “branched” was

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insufficient as a predictor for susceptibility to defluorination. For example, a branched

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fluorotelomer acid 4 (Figure 2d) released negligible F− under the same reaction condition as 1.

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In stark contrast, replacing the –CH2–CH2– moiety within the fluorotelomer structure with –

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CF=CF– (2) led to the rapid and complete degradation of the parent compound within 1 day and

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a maximum of 91% defluorination (an average of eight of the nine C−F bonds cleaved in each

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molecule, Figure 2b). Cyclic 3 is a carboxylic acid analog of PFECHS. It contains one branched

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carbon with two −CF2− and a –COOH neighbors. The parent compound was partially degraded

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(70%), and the F− release from 3 corresponded to an average of three out of eleven F atoms in

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each molecule (Figure 2c). In comparison, negligible F− release was observed from the cyclic

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amine 6, in which the N atom can be considered as a “branched” point. Negligible F− release was

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observed for the perfluoroalkyl ether compound 5 possessing two branched carbons. Hence, it

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can be inferred that the local chemical environment surrounding the tertiary C−F is critical to

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their reactivity with B12. No appreciable defluorination was observed in experiments with the

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linear PFOA (7) and several shorter chain linear acids lacking branched carbon atoms (CF3SO3H,

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CF3COOH, and CF3CF2COOH). LC–QToF-MS analysis showed no significant degradation of 7

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for up to 30 days.

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In comparison to the carboxylic acid 1, B12-catalyzed F− release was much slower for the

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analogous telomer alcohol structure 8 (Figure 2f versus 2a), and the F− release from alcohol 9, 6 ACS Paragon Plus Environment

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with only one branched carbon, was even slower (Figure 2g). The comparison between cyclic 3

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and 10 also shows much slower defluorination from the telomer alcohol than from the carboxylic

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acid (Figure 2h versus 2c). Degradation of alcohols in aqueous solution was not readily

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observed by LC−QToF-MS. Reactions with these alcohols were also conducted with variable

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headspace volumes (e.g., 9 mL liquid + 0 mL headspace versus 2 mL liquid + 7 mL headspace in

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the sealed 9-mL bottles). All conditions yielded similar defluorination results, excluding the

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possibility that the slower defluorination resulted from volatilization of alcohol substrates into

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the headspace.

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Calculated C−F bond dissociation energies (BDEs) provide further insights into mechanisms for

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the initial step of defluorination. As shown in Figure 3, the general order of C−F BDEs is

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tertiary < secondary < primary. The lowest secondary C−F bond BDE (451.9 kJ mol−1) in the 2-

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position of 7 can be attributed to its proximity to –COOH. In 1, 8 and 9, secondary C−F bonds

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adjacent to branched carbons have even lower BDEs (414.2 to 431.0 kJ mol−1). Importantly,

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BDEs for tertiary C−F bonds within the structures for which defluorination was observed

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(compound 1−3 and 8−10) range from 364.4 to 431.0 kJ mol−1, whereas higher BDEs ranging

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from 431.4 to 443.9 kJ mol−1 were found in the two branched structures with no defluorination (4

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and 5). The higher BDEs for those tertiary C−F bonds are due to the presence of relatively weak

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electron-withdrawing hydrocarbon and oxo moieties nearby. We propose that the initial

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defluorination steps occur at tertiary C−F bonds with low BDEs. Such weak bonds can dissociate

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upon interaction with CoI. A Comparison of the three non-cyclic structures 1, 8, and 9 suggests a

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rough correlation between the BDEs of the tertiary C−F bonds (1 < 8 < 9) and the reaction rates

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of defluorination (1 > 8 > 9). Comparison of the cyclic 3 and 10 reveals a similar trend.



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Compound 2 has two sp2 C−F bonds with high BDEs (478.2 and 477.8 kJ mol−1) and exhibited

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rapid and extensive defluorination. In comparison to 1, the presence of sp2 C−F bonds seems to

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promote defluorination. The bonding with the C=C double bond significantly weaken the tertiary

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C−F bond (364.4 kJ mol−1, the lowest BDE among all structures). Similar results have been

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reported by Im et al.,47 who reported B12-catalyzed defluorination of the only sp2 C−F in the

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refrigerant HFO-1234yf (H2C=CF−CF3) yielding H2C=CH−CF3, and further defluorination of

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one sp3 C−F in H2C=CH−CF3 yielding H2C=CH−CF2H. Additionally, sp2 C−F bonds can be

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cleaved with H2 gas and a Rh/Al2O3 catalyst48, 49 while sp3 C−F bonds cannot. Defluorination

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involving unsaturated bonds probably follows reaction mechanisms similar to the Co-catalyzed

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dechlorination of chlorinated ethenes.43 Future studies are necessary for mechanistic elucidation.

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In addition, the close BDE values for the two tertiary C−F bonds in 9 and 4 (431.0 versus 431.4

166

kJ mol−1) suggest that the neighboring atoms are also critical to initiate defluorination reactions.

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The experimental results described above indicate that, B12-catalyzed defluorination requires a

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tertiary C−F branch surrounded by either three fluorinated carbons or two fluorinated carbons

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plus one carboxyl group. If one surrounding atom is changed to hydrocarbon or oxygen, the

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defluorination reaction cannot be initiated. Since typical ether compounds synthesized from C3

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building blocks (e.g., the two shown in Figure 1) do not contain the susceptible C4 branched

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structure, they are expected to be recalcitrant toward B12-catalyzed defluorination.

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It is important to emphasize that BDEs of the parent PFASs can only be used to interpret the

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initiation of defluorination reactions. For example, the experimental data for 1 suggests that,

175

among the sixteen F− released per 1, at least six derived from primary C−F bonds with high

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BDEs ranging from 481.2 to 492.5 kJ mol−1. In contrast, although the linear 7 contains twelve

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secondary C−F bonds with BDEs ranging from 451.9 to 458.1 kJ mol−1, no defluorination was 8 ACS Paragon Plus Environment

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observed. These results collectively indicate that B12-catalyzed defluorination reactions are

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initiated at tertiary C−F bonds with suitable local chemical environments.

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Defluorination intermediates or end products were not observed by LC−QToF-MS (like those

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reported by Park et al. on branched PFOS isomers33). One probable reason for the lack of

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observed products could be that, the fragmentation of PFASs during reaction led to the loss of

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ionizable groups (e.g., −COO−) that enable MS detection. We emphasize that the actual reaction

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mechanisms for the multiple step reactions are complicated because the interpretation goes

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beyond prediction with C−F bond BDEs of the parent structures. For example, the cyclic 3 does

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not contain any primary C−F bonds, but the defluorination ratio observed was much lower than

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that for 1, which contains nine primary C−F bonds. Furthermore, if F is replaced by H during

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defluorination, BDEs of the adjacent C−F bonds in the resulting structures are mostly elevated or

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unchanged (Figure S3), and this would appear to inhibit subsequent defluorination by the same

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mechanism. However, based on the unsaturated intermediate structures proposed by Park et al.,33

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other mechanisms such as HF elimination might be involved in further defluorination. As

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discussed earlier for compound 2 and HFO-1234yf,47 the presence of C=C bonds could promote

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defluorination. Clear mechanistic elucidation requires further experiments with more model

194

compounds and theoretical calculations.

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A protoporphyrin-coordinated complex, Co-PP (Figure 2e),50 exhibited higher initial rate of

196

defluorination than B12 for most PFASs. This trend is apparent for the three examined alcohols

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(Figure 2f−h). For 1, although the maximum defluorination by Co-PP (an average of thirteen of

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nineteen C−F bonds cleaved) was lower than by B12, defluorination by Co-PP in the 1st day was

199

higher than by B12 (Figure 2a). For 3, Co-PP was superior to B12 in both initial reaction rate and

200

the maximum defluorination ratio (Figure 2c). Thus, the two Co complexes demonstrate 9 ACS Paragon Plus Environment


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selectivity toward specific PFAS structures. As with B12, structures 4 to 7 (Figure 2d) were not

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reactive with Co-PP.

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The difference between Co-PP and B12 might be attributed to two factors. First, the lack of axial

204

benzimidazole lower ligand as in B12 may allow faster electron transfer from TiIII to the CoIII

205

precursor to form the reactive CoI.51 Second, the porphyrin ligand has one more bridging carbon

206

than corrin, and the π-electron resonance is circular. Thus, the N4 ligand cavity size, ring

207

flexibility, and electronic effects in Co-PP are all different from B12,52 thus influencing the

208

defluorination

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bis(salicylidene)ethylenediamine ligand (Co-salen, Figure 2e), inorganic CoCl2, and Co3O4

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nanopowder did not show any defluorination activity, suggesting again the critical role of the

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Co-coordinating ligand in defluorination activity. Systematic investigations are required to probe

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the effect of ligands. Nevertheless, the above results have clearly shown that even the highly

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recalcitrant primary sp3 C−F bonds in PFASs could be cleaved by natural and artificial N4-

214

coordinated cobalt species.

215

Although branched structures containing tertiary C−F bonds with low BDEs are subject to Co-

216

catalyzed defluorination, it is worth mentioning that at ambient temperature the branched PFASs

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would still be relatively recalcitrant in the environment. Results in this letter will be valuable in

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two aspects of environmental research. First, rational molecular design may be applied to

219

develop more readily degradable PFASs and Co catalysts that are active for the rapid and

220

complete defluorination in both natural and engineered systems. On the other hand, caution

221

should also be taken as branched PFASs may be reactive in cells, tissues, and organs where B12

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or other catalytic metal species are present.

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ASSOCIATED CONTENT

activity.

An

N2O2-coordinated


Co

complex

with

a

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Supporting Information

225 226

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.estlett.xxxxxxx.

227

Additional experimental procedures, tables, and figures

228

Coordinates of optimized geometries of PFASs and their corresponding radicals

229

AUTHOR INFORMATION

230 231 232 233

Corresponding Authors *(J.L.) E-mail: [email protected]; [email protected]. *(S.V.) E-mail: [email protected]. *(T.J.S.) E-mail: [email protected].

234 235

Notes The authors declare no competing financial interest.

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ACKNOWLEDGEMENTS

237 238 239 240 241

Financial support was provided by the Strategic Environmental Research and Development Program (ER-2424) and the National Science Foundation (CHE-1709719, CHE-1710079). All of the computations were performed using allocated resources at high performance computing facility at Colorado School of Mines. T. Liu and X. Xiao received scholarship support from the China Scholarship Council.



F F

F3 C F F F F CF3 COOH

F3 C

PFMe2OA

F3 C F F CF3 F F SO3 H

COOH

F F F CF3

F F F CF3 F3 C O COOH O F F F F F CF3

HFPO dimer acid (GenX) 242 243 244

SO 3H

di-branched PFOS (3,5-dibrPFOS)

mono-branched PFOS (6-brPFOS) F F

F3 C

F F F F

F F F F F F

F3 C

F FF F

F F

PFECHS

F3 C F F F F F

O

SO3 H F

F F3 C

F F F F F F

F3 C

F F

Page 12 of 17

HFPO trimer acid

*HFPO=Hexafluoropropylene oxide

Figure 1. Examples of branched PFASs detected in the environment. Branched carbons are highlighted in grey.


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245 246 247 248 249


Figure 2. Degradation and defluorination for each PFAS with cobalt catalysts shown in (e). Branches that are effective and ineffective in promoting defluorination are highlighted in green and red, respectively. Reaction conditions: PFAS (0.1 mM), Co catalyst (0.25 mM), TiIII citrate (~36 mM), and carbonate buffer (~40 mM) in water; pH 9.0; 70°C.



250 251 252 253

Figure 3. Calculated bond dissociation energies (in kJ mol−1) at B3LYP/6-311+G(2d,2p)/SMD level of theory (BDEs) of C−F bonds in the PFASs shown in Figure 2. The displayed terminal group with two C=O bonds represent charge-delocalized −COO− anion.


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