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Under ordinary conditions fluorine is a gas a little heavier than air, with a pale yellow colour; inhalation except in very low concentrations is dangerous. Upon cooling, fluorine becomes a yellow liquid. Fluorine occurs combined in the widely distributed mineral fluorite ( calcium fluoride, fluorspar), its chief source, in the minerals cryolite and fluorapatite, and in small amounts in seawater, bones, and teeth. Not a rare element, it
makes up about 0.065 percent of the Earth's crust. Only one isotope occurs in nature, stable fluorine-19.

Fluorine is difficult to isolate from its compounds, and in fact it is impossible to free it by chemical means. No other element is powerful enough, as an oxidizing agent, to replace it. The French chemist Henri Moissan first isolated fluorine in 1886 by electrolysis of anhydrous hydrogen fluoride (HF), in which potassium hydrogen fluoride (KHF2) had been dissolved to make it conduct a current. Elemental fluorine of high purity is prepared commercially by Moissan's procedure. The elemental gas is used as an oxidizer in rocket fuels and to prepare
fluorides.

Fluorine, composed of two-atom molecules (F2), combines with all other elements except helium, neon, and argon to form ionic or covalent fluorides. Its chemical activity can be attributed to its extreme ability to attract electrons (it is the most electronegative element) and to the small size of its atoms. The oxidation state of -1 is the only one observed in fluorine compounds. Because of the small size of the fluoride ion (F-), it forms many stable complexes with positive ions; for example, hexafluorosilicate(IV) (SiF62-) and hexafluoroaluminate(III)
(AlF63-).

One of the principal industrial compounds of fluorine is hydrogen fluoride, obtained by treating fluorite with sulfuric acid. It is employed in the preparation of numerous inorganic and organic fluorine compounds of commercial importance, e.g., sodium aluminum fluoride (Na3AlF6), used as an electrolyte in the electrolytic smelting of aluminum metal; and uranium hexafluoride (UF6), utilized in the gaseous diffusion process of
separating uranium-235 from uranium-238 for reactor fuel. A solution of hydrogen fluoride gas in water is called hydrofluoric acid, large quantities of which are consumed in industry for cleaning metals and for polishing, frosting, and etching glass.

Boron trifluoride (BF3) and antimony trifluoride (SbF3), like hydrogen fluoride, are important catalysts for organic reactions; cobalt trifluoride (CoF3) and chlorine trifluoride (ClF3) are useful fluorinating agents; and sulfur hexafluoride (SF6) is used as a gaseous electrical insulator. Sodium fluoride (NaF) is used to treat dental caries and is often added in small amounts to fluoride-deficient water supplies (fluoridation) to reduce tooth
decay.

Elemental fluorine, often diluted with nitrogen, reacts with hydrocarbons to form corresponding fluorocarbons in which some or all hydrogen has been replaced by fluorine. The resulting compounds are usually char-acterized by great stability, chemical inertness, high electrical resistance, and other valuable physical and chemical properties. This fluorination may be accomplished also by treating organic compounds with cobaltic fluoride or by electrolyzing their solutions in anhydrous hydrogen fluoride. Useful plastics with non-sticking qualities, such as polytetrafluoroethylene ([CF2CF2)x]; known by the commercial name Teflon), are readily made from unsaturated fluorocarbons. Organic compounds containing chlorine, bromine, or iodine are fluorinated to produce compounds such as dichlorodifluoromethane (Cl2CF2), the coolant used in most
household refrigerators and air conditioners.

Atomic number 9

Atomic weight 18.9984

Melting point -219.62° C (-363.32° F)

Boiling point -188° C (-306° F)

Density (1 atm, 0 C) 1.696 g/litre

Oxidation states -1

Electronic config. 2-7 or 1s22s22p5