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Fluorine is a chemical element with the symbol F and it has the atomic number of 9. It is extremely reactive. It can react with almost all elements, including some noble gases, forming fluorides most of the time. Fluorine ranks as the 24th in abundance. The mineral Fluorite , which is mostly calcium fluoride, is the primary source of fluorine, and it was described in 1529 that adding fluorite to metal ores can lower their melting points for use in smelting. Fluorine and its compounds have plenty of industrial uses.
Fluorine is extremely poisonous. Several early experimenters died or sustained injuries from their attempts. But in 1886, Henri Moissan, a French chemist, managed to isolate fluorine by electrolyzing the mixture of potassium bifluoride(KHF2) and hydrogen fluoride.
Reactivity[]
The bond energy of difluorine is much lower than that of either dichlorine or dibromine 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. Conversely, bonds to other atoms are very strong because of fluorine's high electronegativity. 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.
Reactions of elemental fluorine with metals require varying conditions. Alkali metals cause explosions and alkaline earth metalsdisplay 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 °C (572–842 °F).Some solid nonmetals (sulfur, phosphorus) react vigorously in liquid fluorine.Hydrogen sulfide and sulfur dioxide combine readily with fluorine, the latter sometimes explosively; sulfuric acid exhibits much less activity, requiring elevated temperatures.
Hydrogen, like some of the alkali metals, reacts explosively with fluorine.Carbon, as lamp black, reacts at room temperature to yield tetrafluoromethane. Graphite combines with fluorine above 400 °C (752 °F) to produce non-stoichiometric carbon monofluoride; higher temperatures generate gaseous fluorocarbons, sometimes with explosions. Carbon dioxide and carbon monoxide react at or just above room temperature, whereas paraffins and other organic chemicals generate strong reactions: even completely substituted haloalkanes such as carbon tetrachloride, normally incombustible, may explode.Although nitrogen trifluoride is stable, nitrogen requires an electric discharge at elevated temperatures for reaction with fluorine to occur, due to the very strong triple bond in elemental nitrogen; ammonia may react explosively. Oxygendoes not combine with fluorine under ambient conditions, but can be made to react using electric discharge at low temperatures and pressures; the products tend to disintegrate into their constituent 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; however, it does form a compound with fluorine, argon fluorohydride.
Applications[]
Nuclear[]
The largest application of fluorine gas, consuming up to 7,000 metric tons annually, is in the preparation of UF6 for the nuclear fuel cycle. Fluorine is used to fluorinate uranium tetrafluoride, itself formed from uranium dioxide and hydrofluoric acid. Fluorine is monoisotopic, so any mass differences between UF6 molecules are due to the presence of 235U or 238U, enabling uranium enrichment via gaseous diffusion or gas centrifuge. About 6,000 metric tons per year go into producing the inert dielectric SF6 for high-voltage transformers and circuit breakers, eliminating the need for hazardous polychlorinated biphenyls associated with oil-filled devices. Several fluorine compounds are used in electronics: rhenium and tungsten hexafluoride in chemical vapor deposition, tetrafluoromethane in plasma etching and nitrogen trifluoride in cleaning equipment. Fluorine is also used in the synthesis of organic fluorides, but its reactivity often necessitates conversion first to the gentler ClF3, BrF3, or IF5, which together allow calibrated fluorination. Fluorinated pharmaceuticals use sulfur tetrafluoride instead.
Agrichemicals[]
About 30% of agrichemicals contain fluorine, most of them herbicides and fungicides with a few crop regulators. Fluorine substitution, usually of a single atom or at most a trifluoromethyl group, is a robust modification with effects analogous to fluorinated pharmaceuticals: increased biological stay time, membrane crossing, and altering of molecular recognition. Trifluralin is a prominent example, with large-scale use in the U.S. as a weedkiller, but it is a suspected carcinogen and has been banned in many European countries. Sodium monofluoroacetate(1080) is a mammalian poison in which one sodium acetate hydrogen is replaced with fluorine; it disrupts cell metabolism by replacing acetate in the citric acid cycle. First synthesized in the late 19th century, it was recognized as an insecticide in the early 20th century, and was later deployed in its current use. New Zealand, the largest consumer of 1080, uses it to protect kiwis from the invasive Australian common brushtail possum.
Medical[]
There are medical usages of fluorine compounds, such as in dentistry, where it is used to protect the enamel from rotting due to plaque and tartar. Sodium fluoride and Tin(II) fluoride are used to properly clean and whiten teeth.
Twenty percent of modern pharmaceuticals contain fluorine. One of these, the cholesterol-reducer atorvastatin (Lipitor), made more revenue than any other drug until it became generic in 2011. The combination asthma prescription Seretide, a top-ten revenue drug in the mid-2000s, contains two active ingredients, one of which – fluticasone – is fluorinated.Many drugs are fluorinated to delay inactivation and lengthen dosage periods because the carbon–fluorine bond is very stable. Fluorination also increases lipophilicity because the bond is more hydrophobic than the carbon–hydrogen bond, and this often helps in cell membrane penetration and hence bioavailability.