Fluorine is a chemical element with the symbol F and atomic number 9. It is the lightest halogen and exists as a highly toxic pale yellow diatomic gas at standard conditions. The fluoride ion from dissolved fluoride salts inhibits dental cavities, and so finds use in toothpaste and water fluoridation.
About Fluorine in brief
Fluorine is a chemical element with the symbol F and atomic number 9. It is the lightest halogen and exists as a highly toxic pale yellow diatomic gas at standard conditions. As the most electronegative element, it is extremely reactive, as it reacts with all other elements, except for argon, neon, and helium. Among the elements, fluorine ranks 24th in universal abundance and 13th in terrestrial abundance. Industrial production of fluorine gas for uranium enrichment, its largest application, began during the Manhattan Project in World War II. The fluoride ion from dissolved fluoride salts inhibits dental cavities, and so finds use in toothpaste and water fluoridation. Global fluorochemical sales amount to more than US$15 billion a year. Fluorocarbon gases are generally greenhouse gases with global-warming potentials 100 to 23,500 times that of carbon dioxide, SF6 having the highest global warming potential of any known substance. The element has no known metabolic role in mammals; a few plants and sea sponges synthesize organofluarine poisons that help deter predation. It has a small covalent radius of around 60 picometers, similar to those of its period neighbors oxygen and neon. Some pure fluorine and metals require 300–450 °C of metal ivation to form. Some solid nonmetals such as aluminium and iron react vigorously in liquid fluorine. Hydrogen sulfide and sulfur react vigorously with fluorine, the latter sometimes explosively with sulfuric acid.
Carbon monofluoride and fluorine generate gaseous fluorocarbons, sometimes with explosions. Carbon dioxide and carbon monomethane combine readily, sometimes at higher temperatures to produce non-stoichiometric fluorine gaseosorbents. fluorine reacts at room temperature to yield black fluorine black, as lamphane, or fluorine monofluminescent. Graphite combines with fluorine to yield fluorine-based graphite, which is used as a lubricant in the manufacture of cars and other vehicles. The carbon–fluorine bond often has very high chemical and thermal stability; their major uses are as refrigerants, electrical insulation and cookware, the last as PTFE. Pharmaceuticals such as atorvastatin and fluoxetine contain C−F bonds. The outer electrons are ineffective at nuclear shielding, and experience a high effective nuclear charge of 9 − 2 = 7; this affects the atom’s physical properties. It also has a high electron affinity, second only to chlorine, and tends to capture an electron to become isoelectronic with the noble gas neon; it has the highest electr onegativity of any element. The bond energy of difluorines is much lower than that of either Cl2 or Br2 and similar to the easily cleaved peroxide bond; this accounts for fluorine’s easy dissociation, high reactivity, and strong bonds to non-fluorines.