Fluorine

Fluorine is the chemical element with atomic number 9 and the symbol F. It is a very pale yellow-green gas in group 17, the halogens, and it is the most reactive of all the elements. In nature it occurs only in compounds, and the fluoride ions in those compounds are familiar from toothpaste, drinking water and non-stick cookware.

History of fluorine

Early chemists knew that certain metal compounds contained an unidentified element similar to chlorine, but they could not isolate it. The French scientist André Ampère coined the name fluorine in 1812, taking it from the Latin fluere, meaning to flow, because fluorspar, a fluoride mineral, was used as a flux to make metals flow more easily when they were smelted. Even the great English chemist Humphry Davy failed to produce the element, and he became ill trying to isolate it from hydrofluoric acid.

In 1869 the British chemist George Gore passed an electric current through liquid hydrogen fluoride and saw a gas that reacted violently with his apparatus. He thought it was fluorine but could not collect it or prove it. The decisive step came in 1886, when the French chemist Henri Moissan obtained fluorine by the electrolysis of potassium bifluoride dissolved in liquid hydrogen fluoride. He is credited with the discovery.

Properties of fluorine

Fluorine has a standard atomic weight of about 18.998 and the electron configuration [He] 2s2 2p5, one electron short of a full outer shell. That gap drives its extreme reactivity. The element pulls electrons from other atoms so strongly that it attacks all metals quickly, and steel wool bursts into flames when exposed to fluorine gas. In compounds, it almost always has an oxidation state of minus one.

Fluorine is a gas at room temperature with a density of about 0.0016 grams per cubic centimeter. It condenses to a liquid at minus 188.1 degrees Celsius and freezes at minus 219.7 degrees Celsius. The atom is small, with a covalent radius of about 57 picometers, and its first ionization energy is 17.423 electronvolts. Because fluorine is so reactive, it must be handled in specially prepared equipment.

The reactivity of fluorine also explains its unusual place in chemistry. Of the halogens, it is the lightest, and it forms the strongest single bonds to carbon. That is why fluorinated plastics and coatings are so durable and why compounds such as CFCs lasted so long in the atmosphere. PubChem reports that fluorine can even oxidize oxygen, an unusual property that sets it apart from every other element.

Occurrence and production of fluorine

Fluorine is not found as the free element in nature. It is combined in minerals such as calcium fluoride, called fluorspar or fluorite, and the crust holds roughly 585 milligrams of fluorine per kilogram, according to PubChem. Fluorspar is the common starting material for fluorine chemicals.

The element itself is made by electrolysis, the method Moissan pioneered. There was no commercial production of fluorine until the Second World War, when the atomic bomb and other nuclear energy projects made it necessary to produce large quantities. Before then, fluorides were used for a long time in welding and for frosting glass.

Uses of fluorine

The nuclear power industry uses fluorine to make uranium hexafluoride, a compound needed to separate the isotopes of uranium. Fluorine is also used to make sulfur hexafluoride, an insulating gas for high-power electricity transformers. Many fluorochemicals, including solvents and high-temperature plastics, rely on the strength of the carbon-fluorine bond.

The best-known fluoroplastic is Teflon, the trade name for polytetrafluoroethene, or PTFE, which is prized for its non-stick surface in frying pans. It is also used for cable insulation and plumber’s tape, and it is the basis of the waterproof fabric Gore-Tex. Hydrofluoric acid etches glass, including most of the glass used in light bulbs. Calcium fluoride crystals are used in optical applications.

Chlorofluorocarbons, or CFCs, were once used as aerosol propellants, refrigerants and foaming agents for expanded polystyrene. Their inertness allowed them to drift into the stratosphere, where they destroyed part of the Earth’s ozone layer, and they are now banned.

Isotopes of fluorine

Fluorine has only one stable isotope, fluorine-19, so it is a monoisotopic element. Every natural fluorine atom has nine protons and ten neutrons. Because there is only one stable isotope, the atomic weight of fluorine is known to unusual precision, about 18.998403163 according to IUPAC data compiled by PubChem. Nuclear magnetic resonance studies take advantage of fluorine-19, and a radioactive isotope, fluorine-18, is used in medical imaging.

Uranium hexafluoride deserves special mention, because it is a gas at relatively low temperatures and so can be passed through equipment that separates uranium-235 from uranium-238. The nuclear fuel supply chain therefore depends on fluorine chemistry, as does the production of the insulating gas used in electrical transformers and switchgear.

Health and environmental effects of fluorine

Fluoride, the ion formed when fluorine gains an electron, is essential for animals, where it strengthens teeth and bones. It is added to drinking water in some areas at a level of about 1 part per million, and it is added to toothpaste in compounds such as sodium fluoride, stannous fluoride and sodium monofluorophosphate. The Royal Society of Chemistry says that fluoride levels in drinking water below 2 parts per million are believed to prevent tooth cavities, but that above this concentration they may cause mottling of children’s tooth enamel. The average human body contains about 3 milligrams of fluoride.

The element itself is dangerous. Fluorine gas is highly toxic and corrosive, and hydrofluoric acid requires careful handling. Recommendations about fluoridation and fluoride products are made by health authorities and differ among countries, so this entry offers background only and is not personal advice. The persistence of some fluorine-containing chemicals in the environment is the subject of continuing scientific and regulatory attention.

This entry explains a medical term. It is not medical advice. If you are concerned about your health, speak with a qualified health professional.

Editorially reviewed by Specialty Digest Editorial TeamLast reviewed November 4, 2026Researched and drafted with AI assistanceReport an issue