Chlorine vs Fluorine

Fluorine is the most reactive element there is and the most electronegative, and it beats chlorine at every measure of chemical aggression. It is also the one you will almost never encounter, buy, or see specified, because that same reactivity makes it nearly impossible to handle as an element.

Chlorine is the workhorse. It is made by the tens of millions of tonnes, shipped in tank cars, and involved at some stage in a large fraction of all industrial chemistry. Fluorine is made in comparatively tiny quantity and consumed within metres of where it is generated.

The most reactive element is the one nobody ships

Chlorine liquefies under modest pressure and travels. Global production runs to something like eighty million tonnes a year, essentially all of it from the electrolysis of brine, with sodium hydroxide as the co-product — which is why the two are made together in what the industry calls chlor-alkali plants and why their markets are permanently coupled.

Fluorine is produced on the order of a few tens of thousands of tonnes, by electrolysing a molten mixture of potassium fluoride and hydrogen fluoride — essentially the method Henri Moissan used to isolate the element in 1886, a feat that had defeated and in several cases badly injured the chemists who tried before him. Practically all of it is used immediately, principally to make uranium hexafluoride for enrichment, sulfur hexafluoride for high-voltage switchgear, and nitrogen trifluoride for cleaning semiconductor deposition chambers.

The commercial fluorine trade is therefore a trade in compounds. Fluorspar comes out of the ground, hydrogen fluoride comes out of fluorspar, and everything else comes out of hydrogen fluoride.

A weak bond in the molecule, strong bonds in the products

Fluorine's reactivity is often assumed to come from a strong F–F bond breaking. It is the opposite. The bond in F₂ is unusually weak — around 159 kilojoules per mole, well below chlorine's 243. The two atoms are so small that the lone pairs on each of them crowd the other, and that repulsion weakens the bond holding them together.

So fluorine starts from a low barrier and finishes at a very deep well. The bonds it forms are among the strongest in chemistry: the carbon-fluorine bond at roughly 485 kilojoules per mole is the strongest single bond carbon makes with anything, against about 327 for carbon-chlorine. Easy to start, enormously favourable to finish — that is the whole thermodynamic profile of a violent reagent.

One consequence is worth knowing because it trips people up. In water, hydrochloric acid dissociates completely and hydrofluoric acid does not — the fluorine one is the weak acid of the pair. The H–F bond is too strong to give up its proton readily, and the fluoride ion that does form is held tightly by hydrogen bonding to the solvent. Strength of acid and strength of bond are not the same property, and this pair is the classic illustration of the point.

Chlorine has an oxidation-state ladder; fluorine has one rung

Nothing is more electronegative than fluorine, so fluorine can never be the positive partner in a bond. It has exactly one oxidation state in compounds, and that single fact removes an entire branch of chemistry.

Chlorine ranges from −1 all the way to +7, and the oxyanions along that range are industrially enormous: hypochlorite is household and municipal bleach, chlorate is a herbicide and a match-head oxidiser, perchlorate is the oxidiser in solid rocket propellant. There are no fluorine analogues of any of them. When someone asks which halogen bleaches, the answer is chlorine, and the reason is not that fluorine is too weak but that the required compound cannot exist.

What the ozone hole revealed about the difference

Chlorofluorocarbons contain both elements, and when ultraviolet light breaks them apart in the stratosphere only one of the two causes trouble.

A chlorine atom released up there attacks ozone, is regenerated, and attacks again — a catalytic cycle in which a single atom can destroy on the order of a hundred thousand ozone molecules before it is finally removed. A fluorine atom does not get the chance. It reacts almost immediately with methane or water to give hydrogen fluoride, which is stable, unreactive towards ozone and eventually rained out. Fluorine's greater reactivity makes it less damaging, because it takes itself out of the game.

That asymmetry is precisely what the replacement refrigerants exploited. Hydrofluorocarbons keep the fluorine and drop the chlorine, and their ozone depletion potential is zero. It then turned out they are potent greenhouse gases, which is what the 2016 Kigali Amendment was written to deal with — a reminder that swapping one halogen for another solves one problem at a time.

Fluorine went into medicine; chlorine went into water

Nature makes thousands of organochlorine compounds and almost no organofluorines — fluoroacetate, produced by a handful of plants including the southern African gifblaar, is the famous exception. Chloride is meanwhile essential to animals, as the counter-ion of extracellular fluid and as the acid in the stomach. Fluoride has no established essential role in humans at all.

Synthetic chemistry has reversed that balance completely. Something like a fifth to a quarter of all marketed pharmaceuticals now contain fluorine, because a carbon-fluorine bond placed where the liver would otherwise attack a molecule blocks that metabolism, and because fluorine's electron withdrawal is a precise tool for tuning acidity and membrane permeability. Fluoxetine, atorvastatin and the fluoroquinolone antibiotics all owe their existence to it.

Chlorine's contribution to human health came from a different direction entirely. Continuous chlorination of municipal drinking water, first practised in the United States at Jersey City in 1908, is among the most consequential public health measures ever adopted, and waterborne typhoid and cholera collapsed wherever it followed. Fluoride's public health role is narrower and specific: it converts tooth enamel's hydroxyapatite into the less soluble fluorapatite, which is what water fluoridation and fluoride toothpaste are for.

Chlorine's darkest chapter belongs to the same period. Its release as a weapon at Ypres in 1915 opened chemical warfare on the Western Front.

The bond that will not break

Fluorine's strongest asset has become its most serious liability. Per- and polyfluoroalkyl substances persist in the environment for the same reason polytetrafluoroethylene makes a good non-stick surface and a good gasket: nothing biological can break a carbon-fluorine bond at any useful rate. The stability that chemists spent decades designing in is now the property regulators are trying to design out, and the phrase "forever chemicals" is a direct statement about a bond energy.

Chlorinated organics have their own persistence problems — PCBs and DDT among them — but they are degradable in a way that heavily fluorinated chains are not.

Which halogen the job needs

• Disinfecting water at municipal scale — chlorine, with no realistic competitor.
• Bleaching, or any oxidising anion — chlorine, because the fluorine equivalent does not exist.
• A polymer that must resist every chemical and every temperature — fluorine, accepting the end-of-life problem.
• Blocking metabolism at one position in a drug molecule — fluorine.
• Making PVC, solvents, or titanium dioxide by the chloride route — chlorine, which is the bulk of its market.
• Enriching uranium or insulating a substation — fluorine, in compound form, made and used on the same site.