![]() Carbon, as lamp black, reacts at room temperature to yield tetrafluoromethane. ![]() Hydrogen, like some of the alkali metals, reacts explosively with fluorine. Hydrogen sulfide and sulfur dioxide combine readily with fluorine, the latter sometimes explosively sulfuric acid exhibits much less activity, requiring elevated temperatures. Some solid nonmetals (sulfur, phosphorus) react vigorously in liquid fluorine. Alkali metals cause explosions and alkaline earth metals display vigorous activity in bulk to prevent passivation from the formation of metal fluoride layers, most other metals such as aluminium and iron must be powdered, and noble metals require pure fluorine gas at 300–450 ☌ (575–850 ☏). Reactions of elemental fluorine with metals require varying conditions. Unreactive substances like powdered steel, glass fragments, and asbestos fibers react quickly with cold fluorine gas wood and water spontaneously combust under a fluorine jet. Conversely, bonds to other atoms are very strong because of fluorine's high electronegativity. The bond energy of difluorine is much lower than that of either ClĢ and similar to the easily cleaved peroxide bond this, along with high electronegativity, accounts for fluorine's easy dissociation, high reactivity, and strong bonds to non-fluorine atoms. Main article: Chemical characteristics of fluorine External videos Fluorine has no known metabolic role in mammals a few plants and sea sponges synthesize organofluorine poisons (most often monofluoroacetates) that help deter predation. Organofluorine compounds often persist in the environment due to the strength of the carbon–fluorine bond. Global fluorochemical sales amount to more than US$69 billion a year.įluorocarbon gases are generally greenhouse gases with global-warming potentials 100 to 23,500 times that of carbon dioxide, and SF 6 has the highest global warming potential of any known substance. The fluoride ion from dissolved fluoride salts inhibits dental cavities, and so finds use in toothpaste and water fluoridation. Pharmaceuticals such as atorvastatin and fluoxetine contain C−F bonds. Molecules containing a carbon–fluorine bond often have very high chemical and thermal stability their major uses are as refrigerants, electrical insulation and cookware, and PTFE (Teflon). The rest of the fluorite is converted into corrosive hydrogen fluoride en route to various organic fluorides, or into cryolite, which plays a key role in aluminium refining. Owing to the expense of refining pure fluorine, most commercial applications use fluorine compounds, with about half of mined fluorite used in steelmaking. Industrial production of fluorine gas for uranium enrichment, its largest application, began during the Manhattan Project in World War II. Only in 1886 did French chemist Henri Moissan isolate elemental fluorine using low-temperature electrolysis, a process still employed for modern production. Proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, and several early experimenters died or sustained injuries from their attempts. Fluorite, the primary mineral source of fluorine which gave the element its name, was first described in 1529 as it was added to metal ores to lower their melting points for smelting, the Latin verb fluo meaning 'flow' gave the mineral its name. As the most electronegative reactive element, it is extremely reactive, as it reacts with all other elements except for the light inert gases.Īmong the elements, fluorine ranks 24th in universal abundance and 13th in terrestrial abundance. It is the lightest halogen and exists at standard conditions as a highly toxic, pale yellow diatomic gas.
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