oxygen ← fluorine → neon - ↑ F ↓ Cl 9F Periodic table Appearance Tan or yellow gas General properties Name, symbol, number fluorine, F, 9 Pronunciation /ˈflʊəriːn/, /ˈflʊərɪn/, /ˈflɔːriːn/ Element category halogen Group, period, block 17, 2, p Standard atomic weight 18.9984032g·mol−1 Electron configuration 1s2 2s2 2p5 Electrons per shell 2, 7 (Image) Physical properties Phase gas Density (0 °C, 101.325 kPa) 1.7 g/L Liquid density at b.p. 1.5051 g·cm−3 Melting point 53.53 K, −219.62 °C, −363.32 °F Boiling point 85.03 K, −188.12 °C, −306.62 °F Critical point 144.13 K, 5.172 MPa Heat of fusion (F2) 0.510 kJ·mol−1 Heat of vaporization (F2) 6.62 kJ·mol−1 Specific heat capacity (25 °C) (F2) 31.304 J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 38 44 50 58 69 85 Atomic properties Oxidation states −1 (oxidizes oxygen) Electronegativity 3.98 (Pauling scale) Ionization energies (more) 1st: 1681.0 kJ·mol−1 2nd: 3374.2 kJ·mol−1 3rd: 6050.4 kJ·mol−1 Covalent radius 57±32 pm Van der Waals radius 147 pm Miscellanea Crystal structure cubic Magnetic ordering nonmagnetic Thermal conductivity (300 K) 27.7 m W·m−1·K−1 CAS registry number 7782-41-4 Most stable isotopes Main article: Isotopes of fluorine iso NA half-life DM DE (MeV) DP 18F syn 109.77 min β+ (97%) 0.64 18O ε (3%) 1.656 18O 19F 100% 19F is stable with 10 neutrons v · d · e Fluorine ( /ˈflʊəriːn/, /ˈflʊərɪn/, or /ˈflɔrɪn/) is the chemical element with atomic number 9, represented by the symbol F. It is the lightest halogen. The free element, never found in nature, is a yellow-green gas. Under normal conditions elemental fluorine forms a diatomic molecule with chemical formula F2. Chemically, fluorine is one of the strongest oxidants known, and is similar to, but even more reactive and dangerous than elemental chlorine. Affinity for electrons of fluorine leads it to direct reactions with all other elements in which the reaction has been attempted, except for helium, neon and argon. Fluorine's most common mineral fluorite has been known since 1530, but elemental fluorine was not prepared until 1886 by Henri Moissan, whose method of electrolysis remains the only industrial production method. Fluorine's high electronegativity and small atomic radius gives distinctive properties to many of its compounds. Ionic compounds of flourine are water-soluble halide salts. In contrast, covalent fluorine compounds are often low-melting molecular compounds that show physical and chemical properties more similar to hydrocarbons, without being as flammable. The unusual strength of the carbon–fluorine bond confers particularly high chemical and thermal stability to these compounds. Fluorine was named after fluorite (fluorospar), in turn given its name after the Latin noun fluo, meaning "stream or current", because fluorite had a metallurical use as a flux to facilitate liquification of ores in smelting. Thus the word fluorine is related to the English word "flow," and many compounds of fluorine are characterized by their tendency to facillitate flow by decreasing frictional forces between substances. For example, perfluorocarbons and perfluorinated compounds, organic molecules in which a molecule has reacted with a maximal amount of fluorine so that the outer parts of the molecule are fluorinated, are substances in which only covalently-bonded fluorine atoms are exposed to other molecules. This creates compounds of exceptionally low adhesiveness, due to the low van der Waals attraction between covalent fluorine atoms and most other atoms (including other fluorines). Such molecular properties cause the slipperiness of Teflon, which is used as a non-stick coating. The low interaction forces between fluorinated molecules also causes the low viscosity of perfluorocarbon liquids, and the unusual volatility of many high molecular weight fluorine compounds, such as Freons, fluorinated anesthetics, sulfur hexafluoride, tungsten hexafluoride and uranium hexafluoride. Despite its well-established role in the prevention of dental caries,3 fluoride is not considered an essential mineral element for mammals, including humans, in the sense of being necessary for life. Small amounts of fluoride provided by contamination may be beneficial for bone strength, but this is an issue only for formulation of artifical diets.4 Several fluorine compounds and elemental fluorine itself are dangerously toxic, whereas a number of other artificially-produced fluorinated organic compounds are useful pharmaceuticals.5 Contents 1 Characteristics 1.1 Physical 1.2 Chemical 1.3 Isotopes 1.4 Occurrence 2 Etymology and history 3 Production 3.1 Industrial 3.2 Chemical routes 4 Compounds 4.1 Inorganic compounds 4.1.1 Major noble gas compounds 4.2 Organofluorine compounds 5 Applications 5.1 Uses of isotopes of fluorine 5.2 Other industrial uses of fluorine and fluorine-containing compounds 5.2.1 Elemental fluorine 5.2.2 Fluorine containing compounds 6 Fluorine in natural biology 6.1 Inorganic fluorine (fluoride) 6.2 Bioorganic fluorine 6.3 Dental uses 6.4 Pharmaceuticals, agricultural chemicals, and poisons 7 Environmental and health issues 8 Precautions 8.1 Elemental fluorine, fluoride ion, and fluoroacetate 8.2 Hydrogen fluoride and hydrofluoric acid 9 See also 10 References 11 External links // Characteristics Physical

A fluorine atom has nine electrons, which falls just before the extremely stable ten-electron configuration of neon, so a fluorine atom is much more likely to receive the "missing" electron rather than to lose one; its first ionization energy is 1681.0 kJ·mol−1, which makes a fluorine atom extremely difficult to oxidize into a monopositive cation, F+. Such oxidation of fluorine by extraction of an electron requires such energy that no oxidant is known that can to oxidize fluorine to any positive oxidation state.6 Flourine's second and third ionization energies are even higher, 3374.2 and 6050.4 kJ·mol−1, respectively. Such high electron affinities are also reflected in the relatively small radius of a fluorine atom, about 60 picometres, that is roughly comparable to that of hydrogen. Fluorine therefore usually replaces hydrogen in hydrocarbons without significant changes in molecular size. When a fluorine atom becomes a fluoride anion (F-), it gains one electron and therefore has a larger electron cloud radius.2 The covalent radius of fluorine in difluorine molecules, about 71 picometres, is significantly larger than that of a free fluorine atom. This reflects that fact that the bonding between the two atoms is relatively weak—a fact that contributes to the high reactivity of fluorine.2 Fluorine atoms form diatomic molecules with the chemical formula F2. At standard temperature and pressure, this is a corrosive yellow-brown gas. Fluorine gas liquifies at −188.12 °C (−306.62 °F), comparable to oxygen and nitrogen, and solidifies at −219.62 °C (−363.32 °F). Fluorine gas has a density of 1.695 grams per liter and mass of 37.997 u per molecule, which makes fluorine gas about 1.3 times as dense as air.7 Chemical See also: Category:Fluorine compounds Fluorine is the most reactive and most electronegative of the elements, making elemental flourine a dangerously powerful oxidant. For example, fluorine reacts explosively with hydrogen to form hydrogen fluoride,8 notable for bonding very close to ionic in solid, unlike any nonmetals binary compound, including other hydrogen halides. It is so reactive that metals, water, as well as most other substances, burn with a bright flame in a jet of fluorine gas.9 All metals react with fluorine, forming ionic fluorides, but depending on metal, different conditions are required. Alkali metals react with it violently, forming fluorides with formulae MF, while alkaline earth metals do not react violently, but nevertheless react at room temperature. Ruthenium, rhodium, palladium, platinum and gold, the most nonreactive metals to fluorine, react in an atmosphere of pure fluorine only at 300—450 °C.10 Metal tetra- and lower fluorides are highly ionic and soluble, with exception of calcium fluoride, which is hard and insoluble because of its large lattice energy. Metal higher fluorides are known for their volatility; this unique property of fluoride is caused by its small radius.11 Most elements reacts with fluorine readily and directly, including noble gases krypton, xenon, and radon.10 It is known to form compounds with all elements in which the reaction has been attempted, up to einsteinium, element 99.12 No attempt has been made to oxidize astatine, francium, four latter actinides, and all the transactinides with fluorine, due to the radioactive instability of these elements, though such oxidations are thought to be possible in theory. 13 Computational studies have suggested that helium could form a bond with fluorine,14 and excited states containing neon-fluorine bonds have been observed in a mixture of neon and fluorine irradiated with electrons.15 Argon forms argon fluorohydride at low temperature.16 Isotopes Main article: Isotopes of fluorine Fluorine has only one naturally-occurring isotope, fluorine-19;17 making fluorine a mononuclidic element. At least 17 radioisotopes have also been synthesized, ranging in mass number from 14 to 31.18 Fluorine-18 is the lightest unstable nuclide with equal odd numbers of protons and neutrons; it is the most stable radioisotope of fluorine, with a half-life of 109.77 minutes, and commercially an important source of positrons, finding its major use in positron emission tomography scanning.1920 To date, only one nuclear isomer has been characterized: fluorine-18m.21 Its half-life is approximately 160 nanoseconds. All ground states of isotopes from fluorine-17 to fluorine-30, except for fluorine-28, have longer half lives.21 Occurrence From the perspective of cosmology, fluorine is relatively rare because the solar temperatures needed to make it also enable it to quickly fuse with hydrogen to form oxygen and helium, or with helium to become neon and hydrogen. Most fluorine is created either during a supernova when a neutrino hits an atom of neon, or when a blue Wolf-Rayet star massing over 40 solar masses has a stellar wind blowing the fluorine out of the star before hydrogen or helium can destroy it.2223 Fluorite crystals In total abundance fluorine is the thirteenth most common element in Earth crust, making up 0.08% of the crust by mass. There exist three minerals which were mined and contain enough fluorine to be used as industrial resources: one source for fluoride is fluorite, which is widespread. It is used in smelting, construction, and the manufacture of hydrogen fluoride. 24 25 Cryolite is the least abundant of three and it is directly used for the production of aluminium. These two minerals are mined from meteoric water; cryolite is also found in magmatic waters.26 The last important mineral is fluorapatite, which is mined along with other apatites because of its phosphate content, mostly for production of phosphate fertilizers. The hexafluorosilicates produced as by-product of the phosphoric acid production are mostly disposed of as waste. Fluorocarbon-containing chlorofluorocarbons and tetrafluoromethane have been reported in rocks, presumably having formed without action of living organisms. However, they are not a commerically or environmentally important source of fluorine.23 Etymology and history

Fluorine is /ˈflʊəriːn/, /ˈflʊərɪn/, or commonly /ˈflɔrɪn/, a word that ultimately derives from from the Latin noun fluo, meaning a stream or flow of water. In verb form this was fluor or fluere, meaning to flow. The mineral fluorospar (now fluorite, a natural form of calcium fluoride) was first discussed in print in a 1530 work by Georgius Agricola and named for its usefulness as a flux.2728 A flux (from the Latin noun fluxus, a wash or current of water) is used in metallurgy to lower the melting point and promote the fusion of metals and minerals in slag, during smelting. Agricola, a German scientist with expertise in philology, mining, and metallurgy, naming fluospar as a Latinization of the German Flussespar from Flusse (stream, river) and "Spar" (meaning a nonmetallic mineral akin to gypsum, spærstān, spear stone, referring to its crystalline projections). The first recorded preparation of hydrofluoric acid occurred in 1720 by an unknown English glassworker. In 1771 the Swedish chemist Carl Wilhelm Scheele obtained impure hydrofluoric acid by heating fluorite with sulfuric acid in a glass retort, which was greatly corroded by the product; as a result, vessels made of metal were used in subsequent experiments with the substance. The nearly anhydrous acid was reported in 1809, and in 1811 the French physicist André-Marie Ampère suggested that it was a compound of hydrogen with an unknown element, analogous to chlorine, for which he suggested the name fluorine.29 Fluorite was then recognized to be mostly composed of calcium fluoride.630 Henri Moissan Owing to its extreme reactivity, elemental fluorine was not isolated until many years after the characterization of fluorite. Progress in isolating elemental fluorine was slowed because it could be prepared only electrolytically and even then under stringent conditions, since the gas attacks all but certain exotic materials. In 1886, the isolation of elemental fluorine was reported by Henri Moissan after almost 74 years of effort by other chemists.31 The generation of elemental fluorine from hydrofluoric acid proved to be exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. These individuals came to be referred to as "fluorine martyrs".30 For Moissan, the feat earned the 1906 Nobel Prize in chemistry.32 With the introduction of the industrial steel production fluorite was used to improve the slag properties produced during the process.33 Besides the production of hydrofluoric acid this was for some time the two largest applications. With the need to produce artificial cryolite for the aluminium production in the Hall–Héroult process the demand for fluorite increased. The two most prominent uses of organofluorine compounds, Teflon (invented 1938), and hydrofluorocarbons and chlorofluorocarbons (for example Freon-12, introduced as a refrigerant in the late 1920s), are still major applications for fluorine. Bromo and chlorofluorocarbons are being phased out in favor of hydrofluorocarbons, due to environmental protection reasons (see below). The first large-scale production of fluorine was undertaken in support of the Manhattan project, where the compound uranium hexafluoride, a chemical with formula UF6, had been selected as the form of uranium for the separation of uranium-235 from uranium-238.6 Today both the gaseous diffusion process and the gas centrifuge process use gaseous uranium hexafluoride to enrich uranium-235 for nuclear power applications. During the Manhattan Project, it was found that uranium hexafluoride had the corrosive properties of elemental fluorine because it existed in equilibrium with small amounts of uranium tetrafluoride (UF4) and elemental fluorine gas; the latter attacked all chemical compounds which did not already contain a maximal amount of fluorine. The corrosion problem was eventually solved by electrolytically coating all uranium hexafluoride-carrying piping with nickel metal, which forms nickel difluoride that prevents the metal underneath from further attack. Joints and flexible parts were fabricated from teflon, then a new fluorocarbon plastic that (because of its fluorine content) is also not attacked by fluorine.34 Production Industrially, fluorine (F) is used either directly as the mined mineral fluorite (calcium fluoride) or as hydrogen fluoride. Only a very small fraction of industrial F is ever electrolyzed to elemental fluorine (F2). Most fluorine contained in fine organofluorine chemicals is ultimately derived from hydrogen flouride, produced from fluorite (or secondarily as a byproduct of phosphoric acid manufacture from fluoroapatite).353637 Industrial Fluorine cell room at F2 Chemicals Ltd, Preston, UK Industrial production of fluorine in the middle temperature process entails the electrolysis of a solution of potassium bifluoride in hydrogen fluoride. Hydrogen fluoride required for the electrolysis is mainly obtained by treatment of fluorite with sulfuric acid. Several thousand tonnes of fluorine are produced annually. Potassium bifluoride forms spontaneously from potassium fluoride and the hydrogen fluoride: HF + KF → KHF2 The mixture with the approximate composition KF•2HF melts at 70°C and is electrolysed in the temperature range between 70 and 130°C.37 KHF2 increases the electrical conductivity of the solution and provides the anion that is oxidized at the anode, while hydrogen forms at the cathode. These electrolytes and method are based on the pioneering studies by Moissan. Improvements have since been made in electrodes and containment: while Moissan used platinum group metal electrodes and carved fluorite containers, the modern process uses the steel cell itself as cathode, while blocks of carbon are used as anode (the Söderberg carbon electrodes are similar to those used in the electrolysis of aluminium). The voltage for the electrolysis varies between 8 and 12 Volts. If the fluorine gas is cleaned of hydrogen fluoride and oxygen impurities, it may be stored in steel cylinders, where the inside surface becomes coated (passivated) with a metal fluoride layer that resists further attack.3537 Chemical routes

In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl Christe discovered a purely chemical preparation involving the reaction of potassium hexafluoromanganate and antimony pentafluoride at 150 °C, conducted in hydrogen fluoride: 2 K2MnF6 + 4 SbF5 → 4 KSbF6 + 2 MnF3 + F2↑ Although of no practical value, this synthetic route is is conceptually interesting as a rare chemical preparation of elemental fluorine— a reaction not previously thought possible.38 Compounds Fluorine is found in two oxidation states: 0 when only bounded to another fluorine atom in difluorine molecule, and -1 in all other derivatives, as shown in a table below: Oxidation state Name Formula Illustrative compounds 0 fluorine F2 elemental fluorine -1 fluoride F− any other purpose, such as calcium fluoride, ammonium hexafluorosilicate and tetrafluoromethane Fluorine as a freely reacting oxidant gives the strongest oxidants known, which all contain fluoride ions. Chlorine trifluoride, for example, can burn water and sand, both compounds of a weaker oxidant, oxygen. The "superacid" fluoroantimonic acid is known for its oxidizing properties and is the strongest acid presently known. It has an extremely low pKa of −31.3 and is twenty quintillion (20,000,000,000,000,000,000) times stronger than pure sulfuric acid, which has pKa of −12.39 The fluoride ion is basic, therefore hydrofluoric acid is a weak acid in water solution, with pKa equal to 3.18.40 Despite that, it is still very corrosive and attacks glass. Consequently, fluorides of alkali metals produce basic solutions.41 However, water is not an inert solvent in case of hydrogen fluoride: when less basic solvents such as anhydrous acetic acid are used, hydrofluoric acid is the strongest of the hydrohalic acids. Also, owing to the basicity of the fluoride ion, soluble fluorides give basic water solutions. The fluoride ion is a Lewis base, and has a high affinity to certain elements such as calcium and silicon. For example, deprotection of silicon protecting groups is achieved with a fluoride. The fluoride ion is poisonous.42 Inorganic compounds Since the fluorides of virtually all elements are known, a large variety of inorganic fluorides are known. The alkali metal fluoride often resemble the chlorides in terms of structure and solubilities. Reflecting the high basicity of fluoride anion, many alkali metal fluorides form bifluorides with the formula MHF2. Of such metal fluorides, calcium fluoride from the mineral fluorite is preeminant in occurrence and usefulness. Fluorides of the transition metals often resemble oxides in that they are generally polymeric and refractory. Metal hexafluorides are volatile, with important examples being tungsten hexafluoride, which is widely used to deposit coatings of tungsten, and uranium hexafluoride, volatility of which provides its first major technique of uranium enrichment. Major noble gas compounds Main article: Noble gas compound The reactivity of fluorine-containing platinum hexafluoride toward the noble gas xenon was first reported by Neil Bartlett in 1962. The compound he prepared he called xenon hexafluoroplatinate, but since then that has been revealed to be mixture of different chemicals.43 Later that year, xenon was oxidized directly with fluorine, to form xenon difluoride. Since then, xenon tetrafluoride and xenon hexafluoride also been prepared.4445 Additionally, number of oxyfluorides and oxyfluoroxenates is known, including xenon oxytetrafluoride, XeOF4.4647 Krypton has been oxidized by fluorine to krypton difluoride, from which some positive ions are derived, including monofluorokryptyl(II), KrF+, and trifluorodikryptyl(II),48 Kr2F3+. Radon readily reacts with fluorine to form a solid compound, which is generally thought to be radon difluoride. However, it decomposes on attempted vaporization and its exact composition is uncertain.49 Calculations have shown that radon difluoride can be ionic, unlike all other binary noble gas fluorides.50 Both helium and argon can react with hydrogen fluoride, to form helium fluoride hydride and argon fluoride hydride. Argon fluoride hydride was observed at cryogenic temperatures for first time in 2000, being the first chemical compound with argon;51 helium fluoride hydride is metastable.14 The lifetime of HHeF is estimated to be more than 120 ps. The deuterium analogue DHeF is predicted to be more stable with lifetime having 14 ns.52 The compounds of neon bonded to fluorine remains unknown. Organofluorine compounds Main article: Organofluorine chemistry Fluorine can replace hydrogen in organic compounds. The carbon–fluorine bond is the strongest covalent bond in organic chemistry and is very stable. The range of organofluorine compounds is thus diverse, in part because the area is driven by commercial value of such compounds in materials science and pharmaceutical chemistry. Methods for introducing fluorine into organic compounds include both direct fluorination, which can be dangerous, and less direct methods relying on fluorine-containing reagents such as sulfur tetrafluoride. The large inductive effect (electron-withdrawing effect) of the trifluoromethyl group results in the high strength of many fluorinated organic acids, which may be comparable to mineral acids. In these compounds, the affinity of the acid cation for the acid proton is decreased by the cation's fluorine content, which increase its affinity for the extra electron left when the acidic proton leaves. For example, acetic acid is a weak acid, with pKa equal to 4.76, while its fluorinated derative, trifluoroacetic acid has pKa of -0.25, giving it 100,000 times greater formal acidic potential.53 Applications Uses of isotopes of fluorine

Natural fluorine is monoisotopic, consisting of fluorine-19. This isotope has a nuclear spin of 1/2 and a high nuclear magnetic moment, thus fluorine compounds are highly amenable to nuclear magnetic resonance spectroscopy.54 Compounds containing fluorine-18, a radioactive isotope that emits positrons, are often used in positron emission tomography, because its half-life of 110 minutes is long by the standards of positron-emitters. One such species is fluorodeoxyglucose.55 Other industrial uses of fluorine and fluorine-containing compounds Elemental fluorine Elemental fluorine is too dangerous to be used in most laboratory chemistry, although it is occasionally used as a fluorination agent in industrial processes. Elemental fluorine is a precursor to uranium hexafluoride, which is used in the production of nuclear fuels. Sulfur hexafluoride is used as an inert dielectric medium in high voltage switch gear.56 Elemental fluorine is used for plasma etching in semiconductor manufacturing, flat panel display production, and microelectromechanical systems fabrication.57 Together with xenon difluoride, it is also used for fabricating microelectromechanical systems Some researchers including United States and Soviet space scientists in the early 1960s have studied elemental fluorine gas as a possible rocket propellant due to its exceptionally high specific impulse. The experiments failed because fluorine proved difficult to handle, and its combustion products proved extremely toxic and corrosive.58596061 Fluorine containing compounds Inorganic fluorides and organofluorides, a fraction of which are prepared from elemental fluorine, find use in a variety of materials and chemicals, including important pharmaceuticals, agrochemicals, lubricants, and textiles. Polytetrafluoroethylene is often used to coat non-stick frying pans as it is hydrophobic and possesses fairly high heat resistance. The first recorded use of fluorides have been used in the past to help molten metal flow by helping to break the oxide coating of metals, and offering a lower-friction surface, hence the name of fluorite, which derives from Latin verb fluere, meaning to flow.28 Hydrofluoric acid and certain fluoride-containing salts are useful etchants for glass, e.g. for light bulbs. Compounds of fluorine such as potassium fluoride, and sodium hexafluoroaluminate are utilized in applications such as anti-reflective coatings and dichroic mirrors, due to their unusually low refractive index. Sodium hexafluoroaluminate (which is widespread as mineral cryolite), is used in the electrolysis of aluminium because it acts to lower the melting point of aluminum oxide. Nafion, a strongly acidic fluorinated polymer, is a component of fuel cells. Perfluorooctanoic acid and tetrafluoroethylene are directly used in water resistant coatings, and also in the production of low friction plastics such as Teflon, or PTFE. The low van der Waals forces in solid Teflon give it odd antiadhesive properties: for example, it is the only known material surface to which a gecko cannot stick.6263 Other fluorine-based compounds are used in the production of haloalkanes such as chlorofluorocarbons, which are used extensively in air conditioning and in refrigeration. They have been banned for these applications because they contribute to ozone destruction (see below). Fluorine in natural biology Inorganic fluorine (fluoride) Despite its well-established role in the natural and pharmacological prevention of dental caries,3 fluoride is not considered an essential mineral element for mammals and humans, in the sense of being necessary for life. Small amounts of fluoride provided by contamination may be beneficial for bone strength, but this is an issue only for formulation of artifical diets.4 Bioorganic fluorine Fluoroacetic acid compounds are the most common example of organofluorines in plants Biologically synthesized organofluorines have been found in microorganisms and plants,64 but not animals.23 The most common example is fluoroacetate, which occurs as a plant defence against herbivores in at least 40 plants in Australia, Brazil and Africa.65 Other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate which are all believed to be biosynthesized in biochemical pathways from the intermediate fluoroacetaldehyde.23 Adenosyl-fluoride synthase is an enzyme capable of biologically synthesizing the carbon–fluorine bond.66 Dental uses Inorganic compounds of fluoride, including sodium fluoride, tin(II) fluoride, and, most commonly, sodium monofluorophosphate, are in toothpaste to prevent dental cavities. These or related compounds are also added to some municipal water supplies, a process called water fluoridation, although the practice has remained controversial since its beginnings in 1945. This technology was implemented to prevent tooth disease, which affects the majority of people. Despite some controversies, its importance in preventing tooth decay is well-recognized.367 Pharmaceuticals, agricultural chemicals, and poisons A number of pharmaceuticals and organic pesticides contain fluorine. Because of the considerable stability of the carbon-fluorine bond, many drugs are fluorinated to prevent their metabolism and prolong their half-lives, allowing for longer times between dosing or application. For example, an aromatic ring may be a useful drug addition to prevent metabolism, but presents a safety problem: enzymes in the body metabolize some of them into poisonous epoxides. When the para position is substituted with fluorine, however, the aromatic ring is protected and epoxide is no longer produced.68 Another reason for fluorination of biologically active organics is that many organic compounds show have increased lipophilicity due to addition of fluorine (the C-F bond is even more hydrophobic than the C-H bond), and this effect often increases a drugs' bioavailability due to increased cell membrane penetration.69 Of the drugs that were commercialized in the past 50 years, 5-15% contain fluorine, and this percentage is increasing.

For example, fludrocortisone is one of the most common mineralocorticoids, a class of drugs that mimics the actions of aldosterone. The antiinflammatories dexamethasone and triamcinolone, which are among the most potent of the synthetic corticosteroids class of drugs, contain fluorine.70 Several inhaled general anesthetic agents, including the most commonly used inhaled agents, contain fluorine. Examples are sevoflurane, desflurane, and isoflurane, which are hydrofluorocarbon derivatives. Most SSRI antidepressants are fluorinated organics, such as citalopram, escitalopram, fluoxetine, fluvoxamine, and paroxetine. Fluoroquinolones are a commonly used family of broad-spectrum antibiotics. Because of the difficulty of biological systems in dealing with metabolism of fluorinated molecules, fluorinated pharmaceuticals (often antibiotics and antidepressants) are among the major fluorinated organics found in treated city sewage and wastewater. In addition to pharmaceuticals, an estimated thirty percent of agrochemical compounds contain fluorine. Although this water is usually not treated along with sewage, it does contaminate rivers with runoff organofluorines.71 Sodium fluoroacetate (compound 1080) has been used as an insecticide, especially against cockroaches. This compound is also effective as a bait-poison against mammalian pests. As noted above, its anion form, fluoroacetate, is synthesized naturally by some poisonous plants. Today a number of insecticides still contain sodium fluoride, which is much less toxic than fluoroacetate.72 Environmental and health issues Clorofluorocarbons ("CFC"s) and bromofluorocarbons have recently come under heavy enviromental regulation due to their long residence times in the atmosphere, and their contribution to ozone depletion in the upper atmosphere. Since it is specifically chlorine and bromine radicals that harm the ozone layer, not fluorine, compounds that do not contain chlorine or bromine but contain only fluorine, carbon, and hydrogen (called hydrofluorocarbons) are not on the United States Environmental Protection Agency list of ozone-depleting substances,73 and have been widely used as replacements for the chlorine- and bromine-containing halocarbons. Hydrofluorocarbons and perfluorocarbons do, however, function as another type of pollutant: they are greenhouse gases about 4,000 to 10,000 times that of carbon dioxide.74 Abiotic processes can also result in organofluorines considered as "problem molecules." Fluorocarbon based chlorofluorcarbons and tetrafluoromethane have been reported in igneous and metamorphic rock.23 However, environmental and health issues still face many organofluorines. Because of the strength of the carbon–fluorine bond, many synthetic fluorocarbons and fluorcarbon-based compounds are persistent in the environment. The fluorosurfactants PFOS and PFOA (used in waterproofing sprays), and other related chemicals, are persistent global contaminants. PFOS is a persistent organic pollutant and may be harming the health of wildlife; the potential health effects of PFOA to humans are under investigation by the C8 Science Panel. Precautions Elemental fluorine, fluoride ion, and fluoroacetate See also: Fluoride poisoning and sodium fluoroacetate Elemental fluorine (fluorine gas) is a highly toxic, corrosive oxidant, which can cause ignition of organic material (not perfluorinated materials). 56 Fluorine gas is so reactive to the sense of smell that concentrations as low as one part in 50 billions are detectable by odor.75 Soluble flourides are moderately toxic. In the case of the simple salt sodium fluoride, about 28 milligrams per kilogram of body mass (5–10 g for a 70 kg human) are toxic. Fluoride ion is readily absorbed by the stomach, intestines and excreted through the urine. Urine tests have been used to ascertain rates of excretion in order to set upper limits in exposure to fluoride compounds and associated detrimental health effects.76 Ingested fluoride initially acts locally on the intestinal mucosa, where it forms hydrofluoric acid in the stomach.77 Thereafter it binds calcium and interferes with various enzymes.77 Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluoride,78 or (more rarely) rodenticides containing sodium fluoroacetate ("Compound 1080")79) containing organoflourine. Currently, most flouride poisonings are due to the ingestion of fluoride-containing toothpaste.77 Malfunction of water fluoridation equipment has occurred several times, including a notable incident in Alaska.80 Hydrogen fluoride and hydrofluoric acid Main article: Hydrofluoric acid Hydrofluoric acid is a highly corrosive liquid and is a contact poison. It must be handled with extreme care far beyond that accorded to other mineral acids, even the analogous halogen acid HCl. Owing to its lesser chemical dissociation in water, hydrogen fluoride (HF) penetrates tissue more quickly than typical acids, because non-ionized molecular HF has a significant presense even in in water solution, and the HF molecule quickly dissolves in lipids and penetrates them. Poisoning can occur readily through exposure of skin or eyes, or when inhaled or swallowed. Symptoms of exposure to hydrofluoric acid may not be immediately evident. Hydrogen fluoride interferes with nerve function, meaning that burns may not initially be painful. Accidental exposures can go unnoticed, delaying treatment and increasing the extent and seriousness of the injury.81

Once absorbed into blood through the skin, HF reacts with blood calcium and may cause cardiac arrest. Formation of insoluble calcium fluoride is proposed as the etiology for both precipitous fall in serum calcium and the severe pain associated with tissue toxicity.82 In some cases, exposures can lead to hypocalcemia. Burns with areas larger than 160 square centimeters (25 in2) have the potential to cause serious systemic toxicity from interference with blood and tissue calcium levels.83 Thus, hydrofluoric acid exposure is often treated with calcium gluconate, a source of Ca2+ that sequesters the fluoride ions. Hydrogen fluoride chemical burns can be treated with a water wash and 2.5% calcium gluconate gel848586 or special rinsing solutions.8788 However, because it is absorbed, medical treatment is necessary; in some cases, amputation may be required.83 See also Chemistry portal Halogen — chemical series fluorine belongs to Chlorine — a heavier halogen with atomic number 17, like fluorine, very reactive and dangerous, but not as much as fluorine is Teflon — a fluorine-containing material used in everyday life Water fluoridation — process and technology of adding fluorine in water to prevent people's dental problems References ^ Jarry, Roger L.; Miller, Henry C. (1956). Journal of the American Chemical Society 78: 1552. doi:10.1021/ja01589a012.  ^ a b c see covalent radius of fluorine ^ a b c Olivares M and Uauy R (2004). "Essential nutrients in drinking-water (Draft)". WHO. http://www.who.int/water_sanitation_health/dwq/en/nutoverview.pdf. Retrieved 2008-12-30.  ^ a b Micronutrients in parenteral nutrition: boron, silicon, and fluoride. Nielsen FH. Gastroenterology. 2009 Nov;137(5 Suppl):S55-60. PMID 19874950 ^ Obikoya, George. "Fluoride Benefits". The Benefits of Fluoride. http://www.vitamins-nutrition.org/vitamins/fluoride.html.  ^ a b c Britannica (2011). "Fluorine". Britannica.com. http://www.britannica.com/EBchecked/topic/211394/fluorine. Retrieved 2011.  ^ "Fluorine - Compound Summary". PubChem. http://pubchem.ncbi.nlm.nih.gov/summary/summary.cgi?cid=24524.  ^ Greenwood, Norman N.; Earnshaw, Alan. (1997), Chemistry of the Elements (2nd ed.), Oxford: Butterworth-Heinemann, ISBN 0080379419  ^ "Fluorine Facts". About.com. http://chemistry.about.com/od/elementfacts/a/fluorine.htm. Retrieved 2011.  ^ a b (Russian)Lidin P.A., Molochko V.A., Andreeva L.L.. Химические свойства неорганических веществ. p. 442-455.  ^ For example, larger oxide, which is a weaker oxidant and is more likely to form covalent bonds, forms those only in four compounds (manganese heptoxide, technetium heptoxide, ruthenium tetroxide and osmium tetroxide), unlike fluorine, which forms covalent bonds to twelve metals; see fluoride volatility ^ Young, J. P.; Haire, R. G.; Peterson, J. R.; Ensor, D. D.; Fellow, R. L. (1981). "Chemical consequences of radioactive decay. 2. Spectrophotometric study of the ingrowth of berkelium-249 and californium-249 into halides of einsteinium-253". Inorganic Chemistry 20: 3979. doi:10.1021/ic50225a076.  ^ N.B. Francium, as an alkali metal, would surely oxidize, and astatine, as a halogen, should form compounds analogous to the various iodine fluorides. ^ a b Lewars, Errol G. (2008-11-17). Modelling Marvels. Springer. pp. 70–71. 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Retrieved 2008-01-03. dead link ^ Hult�n, Peter; H�jer, J.; Ludwigs, U.; Janson, A. (2004). "Hexafluorine vs. standard decontamination to reduce systemic toxicity after dermal exposure to hydrofluoric acid". J. Toxicol. Clin. Toxicol. 42 (4): 355–361. doi:10.1081/CLT-120039541. PMID 15461243.  ^ "News & Views". Chemical Health and Safety 12 (5): 35–37. September–October 2005. doi:10.1016/j.chs.2005.07.007.  External links Wikimedia Commons has media related to: Fluorine Look up fluorine in Wiktionary, the free dictionary. The Periodic Table of Videos video of Fluorine at YouTube WebElements.com – Fluorine It's Elemental – Fluorine Picture of liquid fluorine – chemie-master.de Chemsoc.org Fluorine: An Element-ary Mystery by Ken Croswell, published in Sky and Telescope September 2003 v · d · eDiatomic chemical elements

Hydrogen H2 | Nitrogen N2 | Oxygen O2 | Fluorine F2 | Chlorine Cl2 | Bromine Br2 | Iodine I2 | Astatine At2 | v · d · e Periodic table H   He Li Be   B C N O F Ne Na Mg   Al Si P S Cl Ar K Ca   Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr   Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals Other metals Metalloids Other nonmetals Halogens Noble gases Large version 

Hydrogen H2 | Nitrogen N2 | Oxygen O2 | Fluorine F2 | Chlorine Cl2 | Bromine Br2 | Iodine I2 | Astatine At2 | v · d · e Periodic table H   He Li Be   B C N O F Ne Na Mg   Al Si P S Cl Ar K Ca   Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr   Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals Other metals Metalloids Other nonmetals Halogens Noble gases Large version 

Hydrogen H2 | Nitrogen N2 | Oxygen O2 | Fluorine F2 | Chlorine Cl2 | Bromine Br2 | Iodine I2 | Astatine At2 | v · d · e Periodic table H   He Li Be   B C N O F Ne Na Mg   Al Si P S Cl Ar K Ca   Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr   Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals Other metals Metalloids Other nonmetals Halogens Noble gases Large version