A class of chemicals known as per- and polyfluoroalkyl compounds (PFAS) are particularly worrisome in the environment because they do not biodegrade. For this reason, scientists often refer to PFAS as “forever chemicals.” PFAS’s staying power results from the carbon–fluorine (C–F) bond being one of the strongest and most stable bonds in organic chemistry. More than 5000 types of PFAS are used in a wide range of products, including nonstick pans, stain-resistant carpet, waterproof cosmetics, fast-food wrappers, dental floss, and fire-fighting foam. As a result of product use and waste released from manufacturing sites, PFAS occur widely in the environment, and, in particular, in drinking water—the most common exposure route. Because PFAS accumulate in the human body, exposure has been linked (in limited studies) to low infant birth weight, endocrine disruption, kidney cancer, and thyroid hormone disruption. Scientists have acknowledged that reducing the likelihood of PFAS exposure requires an effective water-treatment strategy. See also: Biodegradation; Chemical bonding; Environmental toxicology; Fluorocarbon
Nanofiltration, which is a process that removes molecules in the nanometer range during water treatment, could remove PFAS. However, the collected PFAS would still need to be destroyed by some means. As reported in the journal Physical Chemistry Chemical Physics (January 2020), researchers have used modeling experiments to understand how to break C–F bonds, and have shown that excess electrons can break both the carbon–carbon (C–C) and C–F bonds of PFAS in water almost instantly—that is, in less than 100 femtoseconds (10-15 seconds). In the 2020 study, the researchers ran simulations for two of the most common PFAS pollutants: perfluorooctanoic acid (PFOA) and perfluorooctanesulfonic acid (PFOS). See also: Electron; Water treatment
A successful treatment process would require breaking all the C–F bonds, because short-chain PFAS are known to be toxic. Results showed that for both PFOA and PFOS molecules, C–F bonds broke almost instantly in the presence of excess electrons and that the reaction formed hydrogen fluoride (HF) molecules. To be clear, there is still much work to be done, such as determining whether any carcinogens are formed in actual practice. Still, the research indicates that breaking down PFAS quickly and economically is potentially feasible. A process for producing excess electrons could be as simple as shining an ultraviolet light on metal-containing compounds suspended in water, after which the metallic compounds could be recovered and reused. So perhaps even forever chemicals need not last forever. See also: Fluorine; Hydrogen fluoride; Mutagens and carcinogens; Simulation; Ultraviolet radiation
