PFAS have become a regular news story: found in drinking water, rivers, even remote wildlife. It is easy to treat them as a mysterious villain. The chemistry is actually straightforward, and it starts with one bond.
PFOA (perfluorooctanoic acid, C₈HF₁₅O₂), one of the best-known PFAS: a water-loving acid head and a seven-carbon tail wrapped in fluorine. Simplified 2D sketch.
The chemistry chain
1. Structure
A carbon chain wrapped in strong, polar C–F bonds.
2. Property
Chemically stable; repels water and oil.
3. Behaviour
Resists breakdown, so it persists and spreads.
4. Impact
Useful products, but widespread contamination.
5. Limitation
Capturing PFAS is easier than destroying it.
The carbon–fluorine bond
Fluorine is the most electronegative element, and it is small, so it sits very close to the carbon it bonds with. The result is a short, highly polar and very strong bond. Average bond enthalpy is an approximate measure of the energy associated with breaking that type of bond in the gas phase. On that measure C–F is roughly 485 kJ/mol, compared with about 413 for C–H and 347 for C–C, so a molecule carrying many C–F bonds is resistant to being taken apart.
Average bond enthalpy (kJ/mol) — higher means harder to break
- C–Cl346
- C–C347
- C–O358
- C–H413
- C–F485
Typical average bond enthalpies from standard data books. Exact C–F values differ slightly between data books (roughly 450–490 kJ/mol) but it is always the strongest here.
Quick check
Which of these bonds is hardest to break?
There is a second reason the tail is so durable: shielding. Each fluorine atom is small, very electronegative and holds its electrons tightly, and the carbon atoms in the chain are wrapped in a dense sheath of them. That sheath physically blocks other molecules from reaching the carbon backbone, so even reagents that would normally attack a carbon chain find little to grab.
Why that made PFAS so useful
A chain of carbon atoms wrapped in fluorine is chemically inert and repels both water and oil. That is exactly what you want in these everyday products:
- non-stick cookware, where the coating is a fluoropolymer such as PTFE (Teflon);
- firefighting foams (AFFF) that smother a fuel fire without reacting with it;
- stain- and water-resistant fabrics and carpets;
- some food packaging that must resist grease.
The same stability that makes them valuable is the source of the trouble.
Why they became a problem
Bacteria, sunlight and everyday chemical reactions break down most pollutants over time. Many PFAS resist all three, so they persist and spread through water and soil. Some also build up in living organisms (bioaccumulation). In large US health surveys, at least one PFAS was found in more than 99% of blood samples[4]. For most people, drinking water and food (including food packaging) are the main routes of exposure, with some consumer products also contributing[4].
US exampleUS regulation, and it is still moving
In April 2024 the US Environmental Protection Agency finalised legally enforceable drinking-water limits for several PFAS, including 4 parts per trillion each for PFOA and PFOS[2]. In May 2025 it announced it would keep the PFOA and PFOS limits, intends to rescind and reconsider the limits for some others, and plans to extend the compliance deadline from 2029 to 2031[1].
UK contextWhat about the UK?
In England and Wales the Drinking Water Inspectorate updated its guidance so that a combined 100 nanograms per litre applies across 48 PFAS, from January 2025. Scotland applies 100 ng/l to a smaller group of 20 PFAS[5]. Two cautions: this is regulator guidance rather than the same kind of legal limit as the US rules, and the numbers are not directly comparable. 100 ng/l is 100 parts per trillion, but it is a combined figure, whereas the US limits of 4 ppt apply to PFOA and PFOS individually. Rules change, so check the current position with the Drinking Water Inspectorate.
Bond strength is not the whole story
The strong C–F bond is a major reason PFAS resist breaking down, but scientists also point to biology. Microbes have had no evolutionary reason to develop enzymes that attack highly fluorinated compounds, there is little energy to gain from doing so, and breaking the bond releases fluoride, which is toxic to many microbes. Persistence is best explained by the strong bonds, the fluorine shielding and these biological obstacles together[3].
The distinction students often miss
Persistent is not the same as toxic, and capture is not destruction.
Three different ideas, often blurred together
| Term | Meaning | PFAS example |
|---|---|---|
| Persistence | Resists breaking down in the environment | The strong, shielded C–F bond means they last for a very long time |
| Bioaccumulation | Builds up in living organisms over time | Some PFAS accumulate in blood and tissue |
| Exposure pathway | How people actually come into contact with it | Mainly contaminated drinking water and food (including via packaging), plus some consumer products |
Clean-up is really a bond-breaking problem
| Approach | What it does | Limitation |
|---|---|---|
| Activated carbon | PFAS stick to the carbon's surface | Moves PFAS into waste carbon that still needs handling |
| Ion-exchange resins | Charged PFAS swap onto the resin | Same: capture, not destruction |
| High-pressure membranes | Filter PFAS from water | Leaves a concentrated waste stream |
| Advanced oxidation and other emerging chemistry | Aims to break the C–F bonds | Mostly research or pilot scale; can be energy-hungry |
| High-temperature thermal destruction | Supplies enough energy to break C–F bonds | Costly, and must be done carefully so fluorinated by-products are not released |
On its own, capturing PFAS just moves them somewhere else; destroying them means supplying enough energy, or the right chemistry, to break the C–F bond. Finding cheaper, lower-energy ways to do that is a live research challenge, and a fine example of how bond enthalpy is not just an exam topic. Materials researchers are also investigating porous materials such as metal-organic frameworks as adsorbents for PFAS — promising, but still an area of research rather than a proven fix.
Exam-style question
Explain why PFAS can remain in the environment for long periods. [3 marks]
Show model answer and mark breakdown
- C–F bonds are very strong (high bond enthalpy), so a lot of energy is needed to break them (1 mark)
- the fluorine atoms shield the carbon chain, making the molecule chemically unreactive (1 mark)
- so PFAS are not broken down by bacteria, sunlight or normal water treatment, and persist (1 mark)
Examiner focus: cause (strong, shielded C–F bonds) → chemical principle (energy needed to break them; unreactive) → consequence (persistence).
Common mistake: saying PFAS last because they are toxic. Toxicity is a separate idea from persistence.
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Sources and further reading
- US EPA: EPA announces it will keep maximum contaminant levels for PFOA and PFOS (14 May 2025) — US position after 2025
- Federal Register: PFAS National Primary Drinking Water Regulation (26 April 2024) — the 2024 US rule
- Wackett (2024), Microbial Biotechnology: Evolutionary obstacles and not C–F bond strength make PFAS persistent — the biological argument
- US CDC/ATSDR: PFAS and your health — blood levels and exposure routes (US)
- Chemistry World: Water regulator applies stricter PFAS limits in England and Wales — UK Drinking Water Inspectorate guidance
- Nature: 'Forever' chemicals can be destroyed with clever chemistry — now test these techniques outside the lab — destruction methods
Last reviewed: 26 September 2026. Figures are typical published values and rules can change. Exam wording varies by specification, so check yours.
Frequently asked questions
What are PFAS?
PFAS stands for per- and polyfluoroalkyl substances: a large family of man-made chemicals whose carbon chains are heavily bonded to fluorine atoms. They have been used in non-stick coatings, waterproofing, some food packaging and firefighting foams.
Why are they called 'forever chemicals'?
Because the carbon–fluorine bond is very strong and the fluorine atoms shield the carbon chain, many PFAS resist the processes that break down most pollutants. Researchers also point out that microbes have not evolved suitable enzymes to break the bond, so they persist in the environment for a very long time.
How strong is the carbon–fluorine bond?
A typical C–F bond enthalpy is roughly 485 kJ/mol, compared with about 347 kJ/mol for C–C and about 413 kJ/mol for C–H, making it one of the strongest single bonds in organic chemistry.
Can PFAS be removed from water?
They can be captured using methods such as activated carbon, ion-exchange resins and high-pressure membranes, but capture only moves the chemicals somewhere else. Actually destroying them requires breaking those very strong bonds, which is an active area of research.
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Sudershan Soni
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