Fluorine

Fluorine is a chemical element, represented by the symbol F and carrying atomic number 9. As the lightest member of the halogen group, it appears under standard conditions as a pale yellow gas made of diatomic molecules. This element is extraordinarily reactive, capable of combining with every other element except the lighter noble gases, and in its pure form it is highly poisonous.

In the universe, fluorine ranks 24th in abundance, while in Earth’s crust it is the 13th most common element. Its name comes from the mineral fluorite, first described in 1529, which was added to metal ores to help them melt more easily; the Latin verb *fluo*, meaning “to flow,” inspired the mineral’s name. Fluorine was proposed as an element in 1810, but isolating it from its compounds proved difficult and dangerous—several early researchers died or were injured trying. It was not until 1886 that French chemist Henri Moissan succeeded in isolating elemental fluorine using low-temperature electrolysis, a method still used today. Large-scale industrial production of fluorine gas began during World War II as part of the Manhattan Project, primarily for uranium enrichment.

Because purifying fluorine is expensive, most commercial uses rely on fluorine compounds. Roughly half of all mined fluorite goes into steelmaking. The remainder is turned into hydrogen fluoride, which is then used to make various organic fluorides, or into cryolite, essential for refining aluminium. The carbon–fluorine bond is generally very stable. Organofluorine compounds are common in refrigerants, electrical insulation, and PTFE (Teflon). Pharmaceuticals such as atorvastatin and fluoxetine also contain carbon–fluorine bonds. Dissolved fluoride salts help prevent dental cavities, so they are added to toothpaste and drinking water. Global sales of fluorochemicals exceed US$15 billion each year.

Fluorocarbon gases are potent greenhouse gases, with global-warming potentials ranging from 100 to 23,500 times that of carbon dioxide; sulfur hexafluoride (SF₆) has the highest global-warming potential of any known substance. Organofluorine compounds often persist in the environment because the carbon–fluorine bond is so strong. Fluorine has no known metabolic role in mammals, though a few plants and marine sponges produce organofluorine poisons—most often monofluoroacetates—to deter predators.

**Characteristics**

**Electron configuration** A fluorine atom has nine electrons, one fewer than neon, arranged as 1s²2s²2p⁵: two electrons fill the inner shell, and seven occupy the outer shell, which needs one more electron to be complete. The outer electrons do not shield the nucleus well, so they experience a high effective nuclear charge of 9 − 2 = 7, influencing the atom’s physical properties.

Fluorine’s first ionization energy is the third highest among all elements, after helium and neon, making it hard to remove an electron from a neutral fluorine atom. It also has a high electron affinity, second only to chlorine, and readily captures an electron to become isoelectronic with neon. Fluorine has the highest electronegativity of any reactive element. Its covalent radius is small—about 60 picometers—similar to those of its period neighbors oxygen and neon.

**Reactivity** The bond energy of difluorine (F₂) is much lower than that of Cl₂ or Br₂ and is comparable to the easily broken peroxide bond. This, combined with fluorine’s high electronegativity, explains why it dissociates easily, reacts so readily, and forms strong bonds with non-fluorine atoms. Conversely, bonds to other atoms are very strong because of fluorine’s high electronegativity. Even normally unreactive materials like powdered steel, glass fragments, and asbestos fibers react quickly with cold fluorine gas; wood and water burst into flame under a jet of fluorine.

Reactions between elemental fluorine and metals depend on the conditions. Alkali metals cause explosions, and alkaline earth metals react vigorously in bulk. To prevent a protective layer of metal fluoride from forming, most other metals—such as aluminium and iron—must be powdered, while noble metals require pure fluorine gas at 300–450 °C (572–842 °F). Some solid nonmetals, like sulfur and phosphorus, react violently in liquid fluorine. Hydrogen sulfide and sulfur dioxide combine readily with fluorine, sometimes explosively; sulfuric acid reacts much less and needs elevated temperatures.

Hydrogen, like some alkali metals, reacts explosively with fluorine. Carbon in the form of lamp black reacts at room temperature to produce tetrafluoromethane. Graphite combines with fluorine above 400 °C (752 °F) to form non-stoichiometric carbon monofluoride; higher temperatures yield gaseous fluorocarbons, sometimes with explosions. Carbon dioxide and carbon monoxide react at or just above room temperature, while paraffins and other organic chemicals react strongly—even completely substituted haloalkanes like carbon tetrachloride, normally nonflammable, may explode. Nitrogen trifluoride is stable, but elemental nitrogen requires an electric discharge at high temperatures to react with fluorine because of its very strong triple bond; ammonia may react explosively. Oxygen does not combine with fluorine under ordinary conditions but can be made to react using an electric discharge at low temperatures and pressures; the products tend to break apart into their elements when heated. Heavier halogens react readily with fluorine, as does the noble gas radon. Of the other noble gases, only xenon and krypton react, and only under special conditions. Argon does not react with fluorine gas, though it does form a compound, argon fluorohydride.

**Phases** At room temperature, fluorine is a gas of diatomic molecules. When pure, it is pale yellow (sometimes described as yellow-green) and has a sharp, biting, halogen-like odor detectable at 20 ppb. Fluorine condenses into a liquid at low temperatures.