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Fluorine (symbol F, pron. Fluu-reen) is the chemical element with atomic number 9. It is the lightest halogen. At standard pressure and temperature, fluorine is a pale yellow gas composed of diatomic molecules, F2. Fluorine is the most electronegative element and is extremely reactive, requiring great care in handling. It has a single stable isotope, fluorine-19. In stars, fluorine is rare compared to other light elements. In Earth's crust, fluorine is the thirteenth-most abundant element. Fluorine's most-important mineral, fluorite, was first formally described in 1530, in the context of smelting. The mineral's name derives from the Latin verb "fluo" meaning "flow" because fluorite was added to metal ores to lower their melting points. Suggested as a chemical element in 1811, fluorine was named after the source mineral, but resisted many attempts to isolate it. In 1886, French chemist Henri Moissan succeeded. His method of electrolysis remains the industrial production method for fluorine gas. The largest use of elemental fluorine, uranium enrichment, was developed during the Manhattan Project. Because of the difficulty in making elemental fluorine, most fluorine used in commerce is never converted to free fluorine. Instead, hydrofluoric acid is the key intermediate for the US$16 billion-per-year, global fluorochemical industry. The fluorides of low-charged metals are ionic compounds (salts); those of high-charged metals are volatile molecular compounds. The largest uses of inorganic fluoride compounds are steel making and aluminium refining. Organic fluorine compounds tend to have high chemical and thermal stability. The largest commercial use is in refrigerant gases (the many types of "Freons"). Although traditional chlorofluorocarbons are widely banned, the replacement gases still contain fluorine. Polytetrafluoroethylene (Teflon) is the most-important fluoropolymer and is used in electrical insulation, chemical-resistant parts, stadium roofs, and cookware. A growing fraction of modern pharmaceuticals contain fluorine; Lipitor and Prozac are prominent examples. While a few plants and bacteria synthesize organofluorine poisons, fluorine has no metabolic role in mammals. The fluoride ion, when directly applied to teeth, reduces decay and for this reason is used in toothpaste and water fluoridation. Characteristics Physical properties Fluorine forms diatomic molecules (F2) that are gaseous at room temperature with a density about 1.3 times that of air. Though sometimes cited as yellow-green, pure fluorine gas is actually a very pale yellow. The color can only be observed in concentrated fluorine gas when looking down the axis of long tubes, as it appears transparent when observed from the side in normal tubes or if allowed to escape into the atmosphere. The element has a "pungent" characteristic odor that is noticeable in concentrations as low as 20 ppb. Fluorine condenses to a bright yellow liquid at −188 °C (−307 °F), which is near the condensation temperatures of oxygen and nitrogen. It solidifies at −220 °C (−363 °F) into a cubic structure, called beta-fluorine. This phase is transparent and soft, with significant disorder of the molecules. At −228 °C (−378 °F) fluorine undergoes a solid–solid phase transition into a monoclinic structure called alpha-fluorine. This phase is opaque and hard, with close-packed layers of molecules. The solid state phase change requires more energy than the melting point transition and can be violent, shattering samples and blowing out sample holder windows. In general, fluorine's solid state is more similar to oxygen's than to the other halogens' Atomic structure A fluorine atom has nine protons and nine electrons, one fewer than neon, arranged in the electronic configuration [He]2s22p5. Fluorine's outer electrons are relatively separate from each other, and thus they do not shield each other from the nucleus. Therefore, they experience a high effective nuclear charge. Fluorine has a relatively small covalent radius, around 60 picometers. This makes fluorine atoms similar in size to those of oxygen and neon. Structure of the fluorine atom Fluorine is reluctant to ionize. Instead, it exhibits a very strong prefference for one more electron and thus achieve the extremely stable neon-like electron configuration. Fluorine's first ionization energy (energy required to remove an electron to form F+) is 1,681 kJ/mol, which is higher than for any other element except neon and helium. Fluorine's electron affinity (energy released by adding an electron to form F–) is 328 kJ/mol, which is higher than that of any other element except chlorine. Molecular structure While an individual fluorine atom has one unpaired electron, molecular fluorine has all the electrons paired. This makes it diamagnetic (slightly repelled by magnets) with the magnetic susceptibility of −1.2×10−4 (SI), which is close to theoretical predictions. In contrast, the diatomic molecules of the neighboring element oxygen, with two unpaired electrons per molecule, are paramagnetic (attracted to magnets). The fluorine–fluorine bond of the difluorine molecule is relatively weak when compared to the bonds of heavier dihalogen molecules. The bond energy is significantly weaker than those of Cl2 or Br2 molecules and similar to the easily cleaved oxygen–oxygen bonds of peroxides or nitrogen–nitrogen bonds of hydrazines. The covalent radius of fluorine of about 71 picometers found in F2 molecules is significantly larger than that in other compounds because of this weak bonding between the two fluorine atoms. This is a result of the relatively large electron and internuclear repulsions, combined with a relatively small overlap of bonding orbitals arising due to the small size of the atoms. The F2 molecule is commonly described as having exactly one bond (in other words, a bond order of 1) provided by one p electron per atom, as are other halogen X2 molecules. However, the heavier halogens' p electron orbitals partly mix with those of d orbitals, which results in an increase effective bond order; for example, chlorine has a bond order of 1.12. Fluorine's electrons cannot exhibit this d character since there are no such d orbitals close in energy to fluorine's valence orbitals. This also helps explain why bonding in F2 is weaker than in Cl2. Chemical reactivity Fluorine's chemistry is dominated by its strong tendency to gain an electron. It is the most electronegative element and elemental fluorine is a strong oxidant. The removal of an electron from a fluorine atom requires so much energy that no known reagents are known to oxidize fluorine to any positive oxidation state. Reactions with elemental fluorine are often sudden or explosive. Many substances that are generally regarded as unreactive, such as powdered steel, glass fragments, and asbestos fibers, are readily consumed by cold fluorine gas. Wood and even water burn with flames when subjected to a jet of fluorine, without the need for a spark. Reactions of elemental fluorine with metals require different conditions that depend on the metal. Often, the metal (such as aluminium, iron, or copper) must be powdered because many metals passivate by forming protective layers of the metal fluoride that resist further fluoridation. The alkali metals react with fluorine with a bang (small explosion), while the alkaline earth metals react not quite as aggressively. The noble metals ruthenium, rhodium, palladium, platinum, and gold react least readily, requiring pure fluorine gas at 300–450 °C (575–850 °F). Fluorine reacts explosively with hydrogen in a manner similar to that of alkali metals. The halogens react readily with fluorine gas as does the heavy noble gas radon. The lighter noble gases xenon and krypton can be made to react with fluorine under special conditions, while argon will undergo chemical trasformations only with hydrogen fluoride. Nitrogen, with its very stable triple bonds, requires electric discharge and high temperatures to combine with fluorine directly. Isotopes Fluorine occurs naturally on Earth exclusively in the form of its only stable isotope, fluorine-19, which makes the element monoisotopic and mononuclidic. Seventeen radioisotopes have been synthesized: mass numbers 14–18 and 20–31. Fluorine-18 is the most stable radioisotope of fluorine, with a half-life of 109.77 minutes. It is also the lightest unstable nuclide with equal odd numbers of protons and neutrons. The lightest fluorine isotopes, those with mass numbers of 14–16, decay via electron capture. 17F and 18F undergo beta plus decay (positron emission). All isotopes heavier than the stable fluorine-19 decay by beta minus mode (electron emission), while some also decay by neutron emission. Only one nuclear isomer (long-lived excited nuclear state), fluorine-18m, has been characterized. Its half-life before gamma ray emission is 160 nanoseconds. This is less than the decay half-life of any of the fluorine radioisotope nuclear ground states except for mass numbers 14–16, 28, and 31. Origin and occurrence

In the universe Abundance in the Solar System Atomic number	Element	Relative amount 6	Carbon	4,800 7	Nitrogen	1,500 8	Oxygen	8,800 9	Fluorine	1 10	Neon	1,400 11	Sodium	24 12	Magnesium	430 From the perspective of cosmology, fluorine is relatively rare with 400 ppb in the universe. Within stars, any fluorine that is created is rapidly eliminated through nuclear fusion: either with hydrogen to form oxygen and helium, or with helium to make neon and hydrogen. The presence of fluorine at all—outside of temporary existence in stars—is somewhat of a mystery because of the need to escape these fluorine-destroying reactions. Three theoretical solutions to the mystery exist. In type II supernovae, atoms of neon are hit by neutrinos during the explosion and converted to fluorine. In Wolf-Rayet stars (blue stars over 40 times heavier than the Sun), a strong solar wind blows the fluorine out of the star before hydrogen or helium can destroy it. In asymptotic giant branch (a type of red giant) stars, fusion reactions occur in pulses and convection lifts fluorine out of the inner star. Only the red giant hypothesis has supporting evidence from observations. In space, fluorine commonly combines with hydrogen to form hydrogen fluoride. (This compound has been suggested as a tracer to enable tracking reservoirs of hydrogen in the universe.) In addition to HF, monatomic fluorine has been observed in the interstellar medium. Fluorine cations have been seen in planetary nebulae and in stars, including our Sun. On Earth Fluorine is the thirteenth most common element in Earth's crust, comprising between 600 and 700 ppm of the crust by mass. Because of its reactivity, it is essentially only found in compounds. Three minerals exist that are industrially relevant sources: fluorite, fluorapatite, and cryolite. Fluorite (CaF2), also called fluorspar or Blue John, is the main source of commercial fluorine. Fluorite is a colorful mineral associated with hydrothermal deposits. It is common and found worldwide. China supplies more than half of the world's demand; Mexico is the second-largest producer. The United States produced most of the world's fluorite in the early 20th century, but the last mine, in Illinois, shut down in 1995. Fluorapatite (Ca5(PO4)3F) is mined along with other apatites for its phosphate content and is used mostly for production of fertilizers. Most of the Earth's fluorine is bound in this mineral, but because the percentage within the mineral is low (3.5%), the fluorine is discarded as waste. Only in the United States is there significant recovery. There the hexafluorosilicates produced as byproducts are used to supply water fluoridation. Cryolite (Na3AlF6) is the least abundant of the three, but is a concentrated source of fluorine. It was formerly used directly in aluminium production. However, the main commercial mine, on the west coast of Greenland, closed in 1987. Major fluorine-containing minerals Fluorite	Fluorapatite	Cryolite Several other minerals, such as the gemstone topaz, contain fluoride. Fluoride is not significant in seawater or brines, unlike the other halides, because the alkaline earth fluorides precipitate out of water. Organofluorines have been observed in volcanic eruptions and in geothermal springs. Their ultimate origin varies from physical formation under geological conditions to initial biological production and deposition in sediments. However, the provenance is still being studied, as is the natural organofluorine distribution. They are not found in large quantities (compare also the number of known natural organofluorines, 30, to that of organochlorines, 2150), so they are not commercially important source of fluorine. The possibility of small amounts of gaseous fluorine within crystals has been debated for many years. One form of fluorite, antozonite, has a smell suggestive of fluorine when crushed. The mineral also has a dark black color, perhaps from free calcium (not bonded to fluoride). In 2012, a study reported detection of trace quantities (0.04% by weight) of diatomic fluorine in antozonite. It was suggested that radiation from small amounts of uranium within the crystals had caused the free fluorine defects. History

Smelting illustration from Agricola's De re metallica, where fluorite was first described The word "fluorine" derives from the Latin stem of the main source mineral, fluorite, which was first mentioned in 1529 by Georgius Agricola, who described it as a flux—an additive that helps melt ores and slags during smelting. Fluorite stones were called schone flusse in the German of the time. Agricola, writing in Latin but describing 16th century industry, invented several hundred new Latin terms. For the schone flusse stones, he used the Latin noun fluores, "fluxes", because they made metal ores flow when in a fire. After Agricola, the name for the mineral evolved to fluorspar (still commonly used) and then to fluorite. Some sources claim that the first production of hydrofluoric acid was by Heinrich Schwanhard, a German glass cutter, in 1670. A peer-reviewed study of Schwanhard's writings, though, showed no specific mention of fluorite and only discussion of an extremely strong acid. It was hypothesized that this was probably nitric acid or aqua regia, both capable of etching soft glass. Andreas Sigismund Marggraf made the first recorded preparation of hydrofluoric acid in 1764 when he heated fluorite with sulfuric acid in glass, which was greatly corroded by the product. In 1771, Swedish chemist Carl Wilhelm Scheele repeated this reaction. Scheele recognized the product of the reaction as an acid, which he called "fluss-spats-syran" (fluor-spar-acid); in English, it was known as "fluoric acid". In 1810, French physicist André-Marie Ampère suggested that the acid was a compound of hydrogen with an unknown element, analogous to chlorine. Fluorite was then shown to be mostly composed of calcium fluoride. Sir Humphry Davy originally suggested the name fluorine, taking the root from the name of "fluoric acid" and the -ine suffix, similarly to other halogens. This name, with modifications, came to most European languages. (Greek, Russian, and several other languages use the name ftor or derivatives, which was suggested by Ampère and comes from the Greek φθόριος (phthorios), meaning "destructive".) The New Latin name (fluorum) gave the element its current symbol, F, although the symbol Fl has been used in early papers. The symbol Fl is now used for the super-heavy element flerovium. Owing to its extreme reactivity, elemental fluorine was not isolated until many years after the characterization of fluorite. Progress in isolating elemental fluorine was slow because its electrolytic preparation was hard to perform and because the gas reacted with most materials. The generation of elemental fluorine proved to be exceptionally dangerous, killing or blinding several early experimenters. Edmond Frémy thought that passing electric current through pure hydrofluoric acid might work. Previously, hydrogen fluoride was only available in a water solution. Frémy therefore devised a method for producing dry hydrogen fluoride by acidifying potassium bifluoride (KHF2). Unfortunately, pure hydrogen fluoride did not pass an electric current.

Henri Moissan, fluorine discover French chemist Henri Moissan, formerly one of Frémy's students, continued the search. After trying many different approaches, he built on Frémy's earlier attempt by combining potassium bifluoride and hydrogen fluoride. The resultant solution conducted electricity. Moissan also constructed especially corrosion-resistant equipment: containers crafted from a mixture of platinum and iridium (more chemically resistant than pure platinum) with fluorite stoppers. After 74 years of effort by many chemists, on 26 June 1886, Moissan reported the isolation of elemental fluorine. Moissan's report to the French Academy of making fluorine showed appreciation for the feat: One can indeed make various hypotheses on the nature of the liberated gas; the simplest would be that we are in the presence of fluorine Moissan later devised a less expensive apparatus for making fluorine: copper equipment coated with copper fluoride. In 1906, two months before his death, Moissan received the Nobel Prize in chemistry for his fluorine isolation as well as the invention of the electric arc furnace. During the 1930s and 1940s, the DuPont company commercialized organofluorine compounds at large scales. Following trials of chlorofluorcarbons as refrigerants by researchers at General Motors, DuPont developed large-scale production of Freon-12. DuPont and GM formed a joint venture in 1930 to market the new product; in 1949 DuPont took over the business. Freon proved to be a marketplace hit, rapidly replacing earlier, more toxic, refrigerants and growing the overall market for kitchen refrigerators. In 1938, polytetrafluoroethylene (Teflon) was discovered by accident by a recently-hired DuPont PhD, Roy J. Plunkett. While working with a cylinder of tetrafluoroethylene, he was unable to release the gas, although the weight had not changed. Scraping down the container, he found white flakes of a polymer new to the world. Tests showed the substance was resistant to corrosion from most substances and had better high temperature stability than any other plastic. By early 1941, a crash program was making commercial quantities.

The Manhattan Project's K-25 gaseous diffusion plant in Oak Ridge, Tennessee Large-scale productions of elemental fluorine began during World War II. Germany used high-temperature electrolysis to produce tons of chlorine trifluoride, a compound planned to be used as an incendiary. The Manhattan project in the United States produced even more fluorine for use in uranium separation. Gaseous uranium hexafluoride was used to separate uranium-235, an important nuclear explosive, from the heavier uranium-238 in centrifuges and diffusion plants. Because uranium hexafluoride releases small quantities of corrosive fluorine, the separation plants were built with special materials. All pipes were coated with nickel; joints and flexible parts were fabricated from Teflon. In 1958, a DuPont research manager in the Teflon business, Bill Gore, left the company because of its unwillingness to develop Teflon as wire-coating insulation. Gore's son Robert found a method for solving the wire-coating problem and the company W. L. Gore and Associates was born. In 1969, Robert Gore developed an expanded polytetrafluoroethylene (ePTFE) membrane which led to the large Gore-tex business in breathable rainwear. The company developed many other uses of ePTFE. In the 1970s and 1980s, concerns developed over the role chlorofluorocarbons play in damaging the ozone layer. By 1996, almost all nations had banned chlorofluorocarbon refrigerants and commercial production ceased. Fluorine continued to play a role in refrigeration though: hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) were developed as replacement refrigerants. Industry and applications

The global market for fluorochemicals was about US$16 billion per year as of 2006. Historically, the industry has grown a few percent per year and is predicted to do so in the future. Although fluorochemical demand contracted during the 2008–2009 global recession, the industry was predicted to reach 2.6 million metric tons per year by 2015. The largest market is the United States. Western Europe is the second largest. Asia Pacific is the fastest growing region of production. China in particular has experienced significant growth as a fluorochemical market and is becoming a producer of them as well. Fluorite mining (the main source of fluorine) was estimated in 2003 to be a $550 million industry, extracting 4.5 million tons per year. Most ores must be processed to concentrate the fluorite from other minerals by various methods of flotation separation. However, only about 1% of mined fluorite is ever converted to elemental fluorine.

Fluorine industry supply chain: major sources, intermediates and applications. Click for links to related articles. Inorganic fluorides Mined fluorite is separated into two main grades, with about equal production of each. Acidspar is at least 97% CaF2; metspar is much lower purity, 60–85%. (A small amount of the intermediate, ceramic, grade is also made.)  Metspar is used almost exclusively for iron smelting. Acidspar is primarily converted to hydrofluoric acid (by reaction with sulfuric acid). The resultant HF is mostly used to produce organofluorides and synthetic cryolite. About 3 kg (6.5 lb) of metspar grade fluorite, added directly to the batch, are used for every metric ton of steel made. The fluoride ions from CaF2 lower the melt's temperature and viscosity (make the liquid runnier). The calcium content has a tangential benefit in removing sulfur and phosphorus, but other additives such as lime are still needed. Metspar is similarly used in cast iron production and for other iron-containing alloys. Fluorite of the acidspar grade is used directly as an additive to ceramics and enamels, glass fibers and clouded glass, and cement, as well as in the outer coating of welding rods. Acidspar is primarily used for making hydrofluoric acid, which is a chemical intermediate for most fluorine-containing compounds. Significant direct uses of HF include pickling (cleaning) of steel, cracking of alkanes in the petrochemical industry, and etching of glass.

Aluminium smelting process: cryolite (a fluoride) is required to dissolve alumnium oxide. One third of HF (one sixth of mined fluorine) is used to make synthetic cryolite (sodium hexafluoroaluminate) and aluminium trifluoride. These compounds are used in the electrolysis of aluminium by the Hall–Héroult process. About 23 kg (51 lb) are required for every metric ton of aluminium. These compounds are also used as a flux for glass. Fluorosilicates are the next most significant inorganic fluorides formed from HF. The most common one, that of sodium, is used for water fluoridation, as an intermediate for synthetic cryolite and silicon tetrafluoride, and for treatment of effluents in laundries. MgF2 and, to a lesser extent, other alkaline earth difluorides are specialty optical materials. Magnesium difluoride is widely used as an antireflection coating for spectacles and optical equipment. The compound is also a component in newly devised constructions (negative index metamaterials) which are the subject of "invisibility" research. The layered structures can curve light around objects. Other inorganic fluorides made in large quantities include cobalt difluoride (for organofluorine synthesis), nickel difluoride (electronics), lithium fluoride (a flux), sodium fluoride (water fluoridation), potassium fluoride (flux), and ammonium fluoride (various). Sodium and potassium bifluorides are significant to the chemical industry. Fluorocarbons Making organic fluorides is the main use for hydrofluoric acid, consuming over 40% of it (over 20% of all mined fluorite). Within organofluorides, refrigerant gases are still the dominant segment, consuming about 80% of HF. Even though chlorofluorocarbons are widely banned, the replacement refrigerants are often other fluorinated molecules. Fluoropolymers are less than one quarter the size of refrigerant gases in terms of fluorine usage, but are growing faster. Fluorosurfactants are a small segment in mass but are significant economically because of very high prices. Freons and Halon Traditionally chlorofluorocarbons (CFCs) were the predominant fluorinated organic chemical. CFCs are identified by a system of numbering that explains the amount of fluorine, chlorine, carbon and hydrogen in the molecules. The term Freon has been colloquially used for CFCs and similar halogenated molecules, though strictly speaking this is just a DuPont brand name, and many other producers exist. Brand neutral terminology is to use "R" as the prefix. Prominent CFCs included R-11 (trichlorofluoromethane), R-12 (dichlorodifluoromethane), and R-114 (1,2-dichlorotetrafluoroethane).

A Halon fire suppression system in a ship's machinery room Production of CFCs grew strongly through the 1980s, primarily for refrigeration and air conditioning but also for propellants and solvents. Since the end use of these materials is banned in most countries, this industry has shrunk dramatically. By the early 21st century, production of CFCs was less than 10% of the mid-1980s peak, with remaining use primarily as an intermediate for other chemicals. The banning of CFCs initially depressed the overall demand for fluorite but 21st century production of the source mineral has recovered to 1980s levels. Hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) now serve as replacements for CFC refrigerants; few were commercially manufactured before 1990. Currently more than 90% of fluorine used for organics goes into these two classes (in about equal amounts). Prominent HCFCs include R-22 (chlorodifluoromethane) and R-141b (1,1-dichloro-1-fluoroethane). The main HFC is R-134a (1,1,1,2-tetrafluoroethane). A bromofluoroalkane, "Halon" (bromotrifluoromethane) is still widely used in ship and aircraft gaseous fire suppression systems. Because Halon production has been banned since 1994, systems are dependent on the pre-ban stores and on recycling. Fluoropolymers Fluoropolymers are less than 0.1% of all polymers produced in terms of weight. Compared to other polymers, they are more expensive and their consumption is growing at a higher rate. As of about 2006–2007, estimates of the global fluoropolymer production varied from over 100,000 to 180,000 metric tons per year. Yearly revenue estimates ranged from over $2.5 billion to over $3.5 billion. Polytetrafluoroethylene (PTFE) is 60–80% of the world's fluoropolymer production on a weight basis. The term Teflon is sometimes used generically for the substance, but is a DuPont brand name—other PTFE producers exist and DuPont sometimes uses the Teflon brand for other materials. PTFE gets its fluorine without the need for fluorine gas: chloroform (trichloromethane) is treated with HF to make chlorodifluoromethane (R-22, an HFC); this chemical when heated makes tetrafluoroethylene (abbreviated TFE), the starting point for PTFE. The largest application for PTFE is in electrical insulation. It is an excellent dielectric and very chemically stable. It is also used extensively in the chemical process industry where corrosion resistance is needed: in coating pipes, in tubing, and gaskets. Another major use is architectural fabric (PTFE-coated fiberglass cloth used for stadium roofs and such). The major consumer application is non-stick cookware. Major PTFE applications PTFE dielectric separating core and outer metal in a specialty coaxial cable	First Teflon branded frying pan, 1961	The interior of the Tokyo Dome. The roof is PTFE-coated fiberglass and air-supported.

When stretched with a jerk, a PTFE film makes a fine-pored membrane: expanded PTFE (ePTFE). The term "Gore-tex" is sometimes used generically for this material, but that is a specific brand name. W.L. Gore is not the only producer of ePTFE and furthermore "Gore-tex" often refers to more complicated multi-layer membranes or laminated fabrics. ePTFE is used in rainwear, protective apparel and liquids and gas filters. PTFE can also be formed into fibers which are used in pump packing seals and bag house filters for industries with corrosive exhausts. Other fluoropolymers tend to have similar properties to PTFE—high chemical resistance and good dielectric properties—which leads to use in the chemical process industry and electrical insulation. They are easier to work with (to form into complex shapes), but are more expensive than PTFE and have lower thermal stability. Fluorinated ethylene propylene (FEP) is the second most produced fluoropolymer. Films from two different fluoropolymers serve as glass-replacements in solar cells.

Microscopic structure of ePTFE ("Gore-Tex") Fluorinated ionomers (polymers that include charged fragments) are expensive, chemically resistant materials used as membranes in certain electrochemical cells. Nafion, developed in the 1960s, was the first example and remains the most prominent material in the class. The initial Nafion application was as a fuel cell material in spacecraft. Since then, the material has been transforming the 55 million tons per year chloralkali industry; it is replacing hazardous mercury-based cells with membrane cells that are also more energy efficient. While older technology plants continue to run, new plants typically use membrane cells. By 2002, more than a third of the global capacity for the industry was membrane-cell based. Recently, the fuel cell application has reemerged; significant research is being conducted and investments made related to getting proton exchange membrane (PEM) fuel cells into vehicles. Fluoroelastomers are rubber-like substances that are composed of crosslinked mixtures of fluoropolymers. Viton is a prominent example. Chemical-resistant O-rings are the primary application. Fluoroelastomers tend to be more stiff than conventional elastomers, but with superior chemical and heat resistance. Surfactants

Drop of water on a fabric treated with fluorinated surfactant Fluorinated surfactants are small organofluorine molecules, principally used in durable water repellent (DWR). Fluorosurfactants form a large market, over $1 billion per year as of 2006. Scotchgard is a prominent brand, with over $300 million revenue in 2000. Fluorosurfactants are expensive chemicals, comparable to pharmaceutical chemicals: $200–2000 per kilogram ($90–900 per pound). Fluorosurfactants make a very small part of the overall surfactant market, most of which is hydrocarbon based and much cheaper. Some potential applications (e.g. low cost paints) are unable to use fluorosurfactants because of the price impact of compounding in even small amounts of fluorosurfactant. Use in paints was only about $100 million as of 2006. DWR is a finish (very thin coating) put on fabrics that makes them lightly rain resistant, that makes water bead. First developed in the 1950s, fluorosurfactants were 90% of the DWR industry by 1990. DWR is used with garment fabrics, carpeting, and food packaging. DWR is applied to fabrics by "dip-squeeze-dry" (immersion in DWR-water bath, pressing water out, and then drying). Fluorine gas For countries with available data (free-market countries), about 17,000 metric tons of fluorine are produced per year by 11 companies, all G7-resident. Fluorine is relatively inexpensive, costing about $5–8 per kilogram ($2–4 per pound) when sold as uranium hexafluoride or sulfur hexafluoride. Because of difficulties in storage and handling, the price of pure fluorine gas is much higher. Processes demanding large amounts of fluorine gas generally vertically integrate and produce the gas onsite for direct use. The largest application for elemental fluorine is the preparation of uranium hexafluoride, which is used in the production of nuclear fuels. To obtain the compound, uranium dioxide is first treated with hydrofluoric acid, to produce uranium tetrafluoride. This compound is then further fluorinated by direct exposure to fluorine gas to make the hexafluoride. Fluorine's monoisotopic natural occurrence makes it useful in uranium enrichment, because uranium hexafluoride molecules will differ in mass only because of mass differences between uranium-235 and uranium-238. These mass differences are used to separate uranium-235 and uranium-238 via diffusion and centrifugation. Up to 7,000 metric tons per year of fluorine gas are used for this application.

SF6 transformers at a Russian railway The second largest application for fluorine gas is sulfur hexafluoride, which is used as a dielectric medium in high voltage switching stations. SF6 gas has a much higher dielectric strength than air. It is extremely inert and, compared to oil-filled switchgear, has no hazardous polychlorinated biphenyls (PCBs). Sulfur hexafluoride is also used in soundproof windows, in the electronics industry, as well as niche medical and military applications. The compound can be made without using fluorine gas, but the reaction between pure sulfur and pure fluorine gas, first developed by Henri Moissan, remains the commercial practice. About 6,000 metric tons per year of fluorine gas are consumed. Several compounds made from elemental fluorine serve the electronics industry. Rhenium and tungsten hexafluorides are used for chemical vapor deposition of thin metal films onto semiconductors. Tetrafluoromethane, is used for plasma etching in semiconductor manufacturing, flat panel display production, and microelectromechanical systems fabrication. Nitrogen trifluoride is increasingly used for cleaning equipment at display manufacturing plants. Elemental fluorine, itself, is used sometimes for cleaning equipment. For making niche organofluorines and fluorine-containing pharmaceuticals, direct fluorination is usually too hard to control. Preparation of intermediate strength fluorinators from fluorine gas solves this problem. The halogen fluorides ClF3, BrF3, and IF5 provide gentler fluorination, with a series of strengths. They are also easier to handle. Sulfur tetrafluoride is used particularly for making fluorinated pharmaceuticals. United States and Soviet space scientists in the early 1960s studied elemental fluorine as a possible rocket propellant because of the higher specific impulse generated when fluorine replaced oxygen in combustion. The experiments failed because fluorine proved difficult to handle, and its combustion product (typically hydrogen fluoride) was extremely toxic and corrosive. Production of fluorine gas

Fluorine cell room at F2 Chemicals, Preston, England. Commercial producers of fluorine gas continue to use the method of electrolysis pioneered by Moissan, with some modifications in the cell design. Owing to the gas's corrosiveness, special containment materials and handling precautions are required. Chemical routes to the elemental form were published in 1986. Electrolytic synthesis Several thousand metric tons of elemental fluorine are produced annually by electrolysis of potassium bifluoride in hydrogen fluoride. Potassium bifluoride forms spontaneously from potassium fluoride and the hydrogen fluoride: HF + KF → KHF2 A mixture with the approximate composition KF•2HF melts at 70 °C (158 °F) and is electrolyzed between 70 °C and 130 °C (160–265 °F). Potassium bifluoride increases the electrical conductivity of the solution and provides the bifluoride anion, which releases fluorine at the anode (negative part of the cell). If HF alone is electrolyzed, hydrogen forms at the cathode (positive part of the cell) and the fluoride ions remain in solution. After electrolysis, potassium fluoride remains in solution. 2 HF2– → H2↑ + F2↑ + 2 F– The modern version of the process uses steel containers as cathodes, while blocks of carbon are used as anodes. The carbon electrodes are similar to those used in the electrolysis of aluminium. An earlier version of fluorine production process, by Moissan, uses platinum group metal electrodes and carved fluorite containers. The voltage for the electrolysis is between 8 and 12 volts. Handling Pure fluorine gas may be stored in steel cylinders where the inside surface is passivated by a metal fluoride layer that resists further attack. Passivated steel will withstand fluorine provided the temperature is kept below 200 °C (400 °F). Above that temperature, nickel is required. Regulator valves are made of nickel. Fluorine piping is generally made of nickel or Monel (nickel-copper alloy). Care must be taken to passivate all surfaces frequently and to exclude any water or greases. In the laboratory, fluorine gas can be used in glass tubing provided the pressure is low and moisture is excluded,  although some sources recommend systems made of nickel, Monel, and PTFE. 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 of fluorine gas; however, he stated in his work that the basics were known 50 years before the actual reaction. The main idea is that some metal fluoride anions do not have a neutral counterpart (or those are very unstable) and their acidifying would result in chemical oxidation, rather than formation of the expected molecules. Christe lists the following reactions as a possible way: 2 KMnO4 + 2 KF + 10 HF + 3 H2O2 → 2 K2MnF6 + 8 H2O + 3 O2↑ 2 K2MnF6 + 4 SbF5 → 4 KSbF6 + 2 MnF3 + F2↑ This synthetic route is a rare chemical preparation of elemental fluorine, a reaction not previously thought possible. Biological aspects Fluoride is not considered an essential mineral element for mammals and humans. Small amounts of fluoride may be beneficial for bone strength, but this is an issue only in the formulation of artificial diets. Food and drinking water typically contain at least small amounts of fluorides, which are naturally present. Several other biological aspects of fluorine exist. Fluoride is widely used for prevention of dental cavities. In pharmaceuticals and agrichemicals, fluorine sees increasing use in new molecules. Poisons containing fluorine are well known for killing insects and rodents and the very few organisms that incorporate fluorine in their biochemistry do so to make natural poisons. Both radioactive and natural fluorine isotopes are important in respective scanning applications. Oxygen-carrying perfluorocarbons present a possibility for human liquid breathing. Dental care Fluoride ions in contact with teeth have been long thought to limit cavities by turning the forming the hydroxyapatite of teeth into less soluble fluorapatite. The more up-to-date studies show no difference in caries levels between teeth with enamel fluoridated to different degrees, while low levels of fluoride in plaque fluid and saliva definitely do help to fight the early caries. The process of fluoride absorption works only by direct contact (topical treatment). Fluoride ions that are swallowed do not benefit the teeth. Water fluoridation is the controlled addition of fluoride to a public water supply to reduce tooth decay. Its use began in the 1940s, following studies of children in a region where water is naturally fluoridated. It is now used for about two-thirds of the U.S. population on public water systems  and for about 5.7% of people worldwide. Although the best available evidence shows no association with adverse effects other than fluorosis (dental and, in worse cases, skeletal), most of which is mild,  water fluoridation has been contentious for ethical, safety, and efficacy reasons,   and opposition to water fluoridation exists despite its support by public health organizations. The benefits of water fluoridation have lessened recently, presumably because of the availability of fluoride in other forms, but are still measurable, particularly for low income groups. Systematic reviews in 2000 and 2007 showed significant reduction of cavities in children associated with water fluoridation. Sodium fluoride, tin difluoride, and, most commonly, sodium monofluorophosphate, are used in toothpaste. In 1955, the first fluoride toothpaste was introduced, in the United States. Now, almost all toothpaste in developed countries is fluoridated. For example, 95% of European toothpaste contains fluoride. Gels and foams are often advised for special patient groups, particularly those undergoing radiation therapy to the head (cancer patients). The patient receives a four-minute application of a high amount of fluoride. Varnishes exist that perform a similar function, but are more quickly applied. Fluoride is also contained in prescription and non-prescription mouthwashes and is a trace component of foods manufactured from fluoridated water supplies. Pharmaceuticals

Prozac: one of several notable fluorine-containing drugs Of all commercialized pharmaceutical drugs, 20% contain fluorine, including important drugs in many different pharmaceutical classes. Fluorine is added to drug molecules as even a single atom of it can greatly change the chemical properties of the molecule in ways that are desirable. Because of the considerable stability of the carbon-fluorine bond, many drugs are fluorinated to delay their metabolism, which is the chemical process in which the drugs are turned into compounds that allows them to be eliminated. This prolongs their half-lives and allows for longer times between dosing and activation. For example, an aromatic ring may prevent the metabolism of a drug, but this presents a safety problem, because some aromatic compounds are metabolized in the body into poisonous epoxides by the organism's native enzymes. Substituting a fluorine into a para position, however, protects the aromatic ring and prevents the epoxide from being produced. Adding fluorine to biologically active organics increases their lipophilicity (ability to dissolve in fats), because the carbon–fluorine bond is even more hydrophobic than the carbon–hydrogen bond. This effect often increases a drug's bioavailability because of increased cell membrane penetration. Although the potential of fluorine being released in a fluoride leaving group depends on its position in the molecule,  organofluorides are generally very stable, since the carbon–fluorine bond is strong. Fluorines also find their uses in common mineralocorticoids, a class of drugs that increase the blood pressure. Adding a fluorine increases both its medical power and anti-inflammatory effects. Fluorine-containing fludrocortisone is one of the most common of these drugs. Dexamethasone and triamcinolone, which are among the most potent of the related synthetic corticosteroids class of drugs, contain fluorine as well. Several inhaled general anesthetic agents, including the most commonly used inhaled agents, also contain fluorine. The first fluorinated anesthetic agent, halothane, proved to be much safer (neither explosive nor flammable) and longer-lasting than those previously used. Modern fluorinated anesthetics are longer-lasting still and almost insoluble in blood, which accelerates the awakening. Examples include sevoflurane, desflurane, enflurane, and isoflurane, all hydrofluorocarbon derivatives. Prior to 1980s, antidepressants altered not only the serotonin uptake (lack of serotonin is a reason for a depression), but also the altered norepinephrine uptake; this caused most antidepressants' side effects. The first drug to only alter the serotonin uptake was Prozac; it gave birth to the extensive selective serotonin reuptake inhibitor (SSRI) antidepressants class and is the best-selling antidepressant. Many other SSRI antidepressants are fluorinated organics, including Celexa, Luvox, and Lexapro. Fluoroquinolones are a commonly used family of broad-spectrum antibiotics. Molecular structures of several fluorine-containing pharmaceuticals Lipitor (atorvastatin)	5-FU (fluorouracil)	Florinef (fludrocortisone)	Isoflurane Agrichemicals and natural poisons

South Africa's gifblaar is one of the few organisms that makes fluorine compounds. An estimated 30% of agrichemical compounds contain fluorine. Most of them are poisons, but a few stimulate the growth instead. It is expected that how often the fluorine agrichemicals will be used depends on two factors: if the synthesis reaction will be improved (to reduce the prices) and if green chemistry will be taken in account to a larger scale (fluorochemicals are more environment-friendly). Synthetic sodium fluoroacetate has been used as an insecticide but is especially effective against mammalian pests. The name "1080" refers to the catalogue number of the poison, which became its brand name. Fluoroacetate is similar to acetate, which has a pivotal role in the Krebs cycle (a key part of cell metabolism). Fluoroacetate halts the cycle and causes cells to be deprived of energy. Several other insecticides contain sodium fluoride, which is much less toxic than fluoroacetate. Currently the compound is banned. Another important agrichemcial is Trifluralin. It was once very important (for example, in 1998 over a half of U.S. cotton field area was coated with the chemical ); however, its suspected carcinogenic properties caused some Northern European countries to ban it in 1993. Currently, the whole European Union has it banned, although there was a case intended to cancel the decision. The currently used agrichemicals utilize another tactic: instead of being poisonous themselves, e.g., by directly affecting metabolism, they transform the metabolism so the organism produces poisonous compounds. For example, insects fed 29-fluorostigmasterol produce the fluoroacetates from it. If a fluorine is transferred to a body cell, it blocks metabolism at the position occupied. Biologically synthesized organofluorines have been found in microorganisms and plants, but not in animals. The most common example is fluoroacetate, with an active poison molecule identical to commercial "1080". It is used as a defense against herbivores by at least 40 green plants in Australia, Brazil, and Africa;  other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate. In bacteria, the enzyme adenosyl-fluoride synthase, which makes the carbon–fluorine bond, was isolated. The discovery was touted as possibly leading to biological routes for organofluorine synthesis. Scanning Compounds containing fluorine-18, a radioactive isotope that emits positrons, are often used in positron emission tomography (PET) scanning, because the isotope's half-life of about 110 minutes is long by positron-emitter standards. One such radiopharmaceutical is 2-deoxy-2-(18F)fluoro-D-glucose (generically referred to as fludeoxyglucose), commonly abbreviated as 18F-FDG, or simply FDG. In PET imaging, FDG can be used for assessing glucose metabolism in the brain and for imaging cancer tumors. After injection into the blood, FDG is taken up by "FDG-avid" tissues with a high need for glucose, such as the brain and most types of malignant tumors. Tomography, often assisted by a computer to form a PET/CT (CT stands for "computer tomography") machine, can then be used to diagnose or monitor treatment of cancers; especially Hodgkin's lymphoma, lung cancer, and breast cancer. Natural fluorine is monoisotopic, consisting solely of fluorine-19. Fluorine compounds are highly amenable to nuclear magnetic resonance (NMR), because fluorine-19 has a nuclear spin of ½, a high nuclear magnetic moment, and a high magnetogyric ratio. Fluorine compounds typically have a fast NMR relaxation, which enables the use of fast averaging to obtain a signal-to-noise ratio similar to hydrogen-1 NMR spectra. Fluorine-19 is commonly used in NMR study of metabolism, protein structures and conformational changes. In addition, inert fluorinated gases have the potential to be a cheap and efficient tool for imaging lung ventilation. Oxygen transport research Liquid fluorocarbons have a very high capacity for holding gas in solution. They can hold more oxygen or carbon dioxide than blood does. For that reason, they have attracted ongoing interest related to the possibility of artificial blood or of liquid breathing. Blood substitutes are the subject of research because the demand for blood transfusions grows faster than donations. In some scenarios, artificial blood may be more convenient or safe. Because fluorocarbons do not normally mix with water, they must be mixed into emulsions (small droplets of perfluorocarbon suspended in water) to be used as blood. One such product, Oxycyte, has been through initial clinical trials. Possible medical uses of liquid breathing (which uses pure perfluorocarbon liquid, not a water emulsion) involve assistance for premature babies or for burn victims (because the normal lung function is compromised). Both partial filling of the lungs and complete filling of the lungs have been considered, although only the former has any significant tests in humans. Several animal tests have been performed and some human partial liquid ventilation trials. One effort, by Alliance Pharmaceuticals, reached clinical trials but was abandoned because of insufficient advantage compared to other therapies. Nanocrystals represent a possible method of delivering water or fat soluble drugs within a perfluorochemical fluid. The particles usage is being developed to help treat babies with damaged lungs. Other posited applications include deep sea diving and space travel, applications that would both require total liquid ventilation, not partial ventilation. The 1989 film The Abyss showed a fictional use of perfluorocarbon for human diving but also filmed a real rat surviving while cooled and immersed in perfluorocarbon. (See also list of fictional treatments of perfluorocarbon breathing.) Hazards Fluorine gas and hydrogen fluoride Elemental fluorine is highly toxic. Above a concentration of 25 ppm, fluorine causes significant irritation while attacking the eyes, respiratory tract, lungs, liver and kidneys. At a concentration of 100 ppm, human eyes and noses are seriously damaged. Hydrogen fluoride is a gas, but upon contact with moisture, including tissue, hydrogen fluoride immediately converts to hydrofluoric acid, which is a contact poison and must be handled with extreme care, far beyond that accorded to other mineral acids. Owing to its lesser chemical dissociation in water (remaining a neutral molecule), hydrogen fluoride penetrates tissue more quickly than typical acids. 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, with 8-hour delay for 50% HF and up to 24-hour if the concentration is smaller. 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. If the burn has been initially noticed, then HF should be washed off with a forceful stream of water for ten to fifteen minutes, to prevent its further penetration into the body. Clothing used by the person burned may also exhibit danger. Once in the blood, hydrogen fluoride reacts with calcium and magnesium, resulting in electrolyte unbalance, cardiac arrhythmia, and potentially, death. Formation of insoluble calcium fluoridepossibly causes both a fall in calcium serum and the strong pain associated with tissue toxicity. In some cases, exposures can lead to hypocalcemia. Burns with areas larger than 160 cm2 (25 in2) can cause serious systemic toxicity from interference with blood and tissue calcium levels. Hydrofluoric acid exposure is often treated with calcium gluconate, a source of Ca2+ that binds with the fluoride ions. Skin burns can be treated with a water wash and 2.5% calcium gluconate gel    or special rinsing solutions. However, because HF is absorbed, medical treatment is necessary; sometimes amputation may be required. Fluoride ion Soluble fluorides are moderately toxic. For sodium fluoride, the lethal dose for adults is 5–10 g, which is equivalent to 32–64 mg of elemental fluoride per kilogram of body weight. The dose that may lead to adverse health effects is about one fifth the lethal dose. Chronic excess fluoride consumption can lead to skeletal fluorosis, a disease of the bones that affects millions in Asia and Africa. The fluoride ion is readily absorbed by the stomach and intestines. Ingested fluoride forms hydrofluoric acid in the stomach. In this form, fluoride crosses cell membranes and then binds with calcium and interferes with various enzymes. Fluoride is excreted through urine. Fluoride exposure limits are based on urine testing which has determined the human body's capacity for ridding itself of fluoride. Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluoride,  Currently, most calls to poison control centers for possible fluoride poisoning come from the ingestion of fluoride-containing toothpaste. Malfunction of water fluoridation equipment has occurred several times, including an Alaskan incident, which affected nearly 300 people and killed one. Environmental concerns Atmosphere Chlorofluorocarbons (CFCs) and bromofluorocarbons (BFCs) have been strictly regulated via a series of international agreements, the Montreal Protocol, because they deplete the ozone layer. It is the chlorine and bromine from these molecules that cause harm, not fluorine. Because of the inherent stability of these fully halogenated molecules (which makes them so nonflammable and useful), they are able to reach the upper reaches of the atmosphere, before decomposing, and then release chlorine and bromine to attack the ozone at those altitudes. Predictions are that generations will be required, even after the CFC ban, for these molecules to leave the atmosphere and for the ozone layer to recover. Early indications are that the CFC ban is working—ozone depletion has stopped and recovery is underway. Hydrochlorofluorocarbons (HCFCs) are current replacements for CFCs; HCFCs have about one tenth the ozone damaging potential (ODP) of CFCs. They were originally scheduled for elimination by 2030 in developed nations (2040 in undeveloped). In 2007, a new treaty was signed by almost all nations to move that phaseout up by ten years because HFCs, which have no chlorine and thus zero ODP, are available. Meanwhile, individual HCFCs with the highest anti-ozone depleting potential are being phased out first. For example, in 2003, HCFC-141b was phased out in the U.S. by Environmental Protection Agency regulation. Many of the other HCFCs are now being produced at a fraction of their previous production rates. Fluorocarbon gases of all sorts (CFCs, HFCs, etc.) are greenhouse gases about 4,000 to 10,000 times as potent as carbon dioxide. Sulfur hexafluoride exhibits an even stronger effect, about 20,000 times the global warming potential of carbon dioxide. Biopersistance

Bottlenose dolphins have high PFOS. Because of the strength of the carbon–fluorine bond, organofluorines endure in the environment. Perfluorooctanoic acid (PFOA) and perfluorooctanesulfonic acid (PFOS), used in waterproofing sprays, are persistent global contaminants. Trace quantities of these substances have been detected worldwide, from polar bears in the Arctic to the global human population. One study indicates that PFOS levels in wildlife are starting to go down because of the recent reduced production of that chemical. PFOA's tissue distribution in humans is unknown, but studies in rats suggest it is likely to be present primarily in the liver, kidney, and blood. In the body, PFOA binds to a protein, serum albumin; it has been detected in breast milk and the blood of newborns. PFOA is not metabolized by the body, but is excreted by the kidneys. The potential health effects of PFOA are unclear. Unlike chlorinated hydrocarbons, PFOA is not lipophilic (stored in fat), nor is it genotoxic (damaging genes). While both PFOA and PFOS cause cancer in high quantities in animals, studies on exposed humans have not been able to prove an impact at current exposures. Bottlenose dolphins have some of the highest PFOS concentrations of any wildlife studied; one study suggests an impact on their immune systems. Because biological systems do not metabolize fluorinated molecules easily, fluorinated pharmaceuticals (often antibiotics and antidepressants) are among the major fluorinated organics found in treated city sewage and wastewater. Fluorine-containing agrichemicals are measurable in farmland runoff and nearby rivers. Compounds Fluorine's only common oxidation state is −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine has a rich chemistry including inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state is only achieved in as their fluorides. Inorganic Hydrogen fluoride Boiling points of the hydrogen halides (blue) and hydrogen chalcogenides (red): HF and H2O break trends. Freezing point of HF/ H2O mixtures: arrows indicate compounds in the solid state. Fluorine combines with hydrogen to make a compound called hydrogen fluoride (HF) or, especially in the context of water solutions, hydrofluoric acid. The HF molecules interact weakly through hydrogen bonds, thus creating extra clustering associations with other HF molecules. Because of this, hydrogen flouride behaves more like water than like other hydrogen halides, such as HCl. This hydrogen bonding between HF molecules gives rise to high viscosity in the liquid phase and lower than expected pressure in the gas phase. Hydrogen fluoride does not boil until 20 °C in contrast to the heavier hydrogen halides which boil between −85 °C (−120 °F) and −35 °C (−30 °F). Hydrogen fluoride is fully miscible with water (dissolves in any proportion), while the other hydrogen halides have large solubility gaps with water. Hydrogen fluoride and water also form several compounds in the solid state, most notably a 1:1 compound that does not melt until −40 °C (−40 °F), which is 44 °C (79 °F) above the melting point of pure HF. Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in water solution, with acid dissociation constant (pKa) of 3.19. This is in part a result of the strength of the hydrogen-fluorine bond, but other factors such as the tendency of HF, H2O, and F− anions to form clusters. At high concentrations, HF molecules undergo homoassociation to form polyatomic ions (such as bifluoride, HF− 2) and protons, thus greatly increasing the acidity. This leads to protonation of very strong acids like hydrochloric, sulfuric, or nitric when using concentrated hydrofluoric acid solutions. Although hydrofluoric acid is regarded as a weak acid, it is very corrosive, even attacking glass when hydrated. Dry hydrogen fluoride readily dissolves low-valent metal fluorides as well as several molecular fluorides. Many proteins and carbohydrates can be dissolved in dry HF and recovered from it. In contrast, most non-fluoride inorganic chemicals react with HF rather than dissolving. Metal fluorides Metal fluorides share some similarities with other metal halides but are more ionic. In some aspects though, metal fluorides differ from the other metal halides (chlorides, bromides, iodides), which are very similar to each other. In such cases, the fluorides are more similar to the oxides, often having similar bonding and crystal structures. The solubility of fluorides varies greatly but tends to decrease as the charge on the metal ion increases. Dissolved fluorides produce basic solutions. (F- is a weak base because HF is a weak acid.) Alkali metal fluorides are very ionic and soluble and behave similar to oxides such as CaO. Other ionic metals fluorides show clear difference than their chlorides stemming from their increased ionic character. For example, thallium and silver monofluorides are soluble, while the chlorides are not; aluminium and gold form ionic trifluorides while the trichlorides are covalent. The increased ionic character leads to alkaline earth metal fluorides (such as CaF2) to have very low solubilities. At larger oxidation numbers, the metal halides do exhibit covalent character: titanium, niobium, and vanadium tetrafluorides are polymeric. They normally have low melting points or decompose easily. Beryllium fluoride (BeF2) also exhibits significant covalent character, and forms similar structures to SiO2 (quartz), but unlike the other alkaline earth salts, it is very soluble in water. Higher fluorides are normally discrete molecules, which contrasts the behavior of corresponding oxides. While oxygen forms discrete molecules with only five metals (manganese heptoxide, technetium heptoxide, ruthenium tetroxide, osmium tetroxide, and iridium tetroxide), fluorine forms molecules with fifteen metals. This is because its small size and single charge as an ion allows surrounding metal atoms with more fluorines than oxygen can. These compounds are highly reactive, acting like acids. For example, platinum hexafluoride was the first compound to oxidize molecular oxygen and xenon. Nonmetal fluorides The nonmetal binary fluorides are volatile compounds. Nonmetals from period 3 and below form fluorides which are hypervalent, such as phosphorus pentafluoride or sulfur hexafluoride. Their reactivity varies greatly: sulfur hexafluoride is relatively inert, while chlorine trifluoride is extremely reactive. Boron trifluoride is a planar molecule, where the central boron atom has only six electrons (and thus an incomplete octet). It readily accepts a Lewis base, forming adducts with molecules that have lone-pairs such as ammonia. With another fluoride ion it completes its octet to form the relatively unreactive BF−4 anion. Boron monofluoride is an unstable molecule that is isoelectronic with N2 and has an unusual (higher than single) bond to fluorine, and an unusual, positive partial charge on fluorine (see the image). Silicon tetrafluoride adopts a molecular tetrahedral structure and demonstrates a weak acidic character, while fluorosilicic acid (H2SiF6) is a strong acid. Pnictogens (nitrogen's periodic table column) fluorides become more reactive as the pnictogen becomes heavier and are weak Lewis bases. The pentafluorides are much more reactive than the trifluorides, with antimony pentafluoride being the strongest Lewis acid of all charge-neutral compounds. Nitrogen is different than other pnictogens, as it forms a triflouride that stable against hydrolysis and is not a Lewis base, and does not form pentafluoride. The metastable nitrogen monofluoride has been observed in laser studies and is isoelectronic with O2, and like BF, it contains a multiple bond. The chalcogen (oxygen's periodic table column) fluorides exhibit similar trends to the pnictogens. The tetafluorides are thermally unstable and hydrolyze, form adducts with other (acidic) fluorides through their lone pair. Sulfur and selenium tetrafluorides are molecular while TeF4 is a polymer. The hexafluorides are the result of direct fluorination of sulfur, selenium, and tellurium, while other hexahalides of the elements do not even exist. SF6 is extremely inert, while SeF6 and TeF6 show increasingly higher reactivity. The only binary fluoride for oxygen is oxygen difluoride, but the OF+3 cation (with a +4 oxidation state at oxygen) was predicted to be theoretically stable. The well-characterized heavier halogens (chlorine, bromine, and iodine) all form mono-, tri-, and pentafluorides: XF, XF3, and XF5. Of the neutral +7 species, only iodine heptafluoride is known. The corresponding cations ClF+6 and BrF+6, are known and are extremely strong oxidizers. For the radioactive astatine only the non-volatile astatine monofluoride has been studied, but its existence is debated. Many of the halogen fluorides are powerful fluorinators (sources of fluorine atoms). ClF3 readily fluorinates asbestos and refractory oxides, and industrial uses require special precautions similar to those for fluorine gas due to its corrosiveness and hazards to humans. Superacids Several important very strong inorganic acids contain fluorine. One such acid, fluoroantimonic acid (HSbF6), is the strongest charge-neutral acid known. The dispersion of the charge on the anion affects the acidity of the solvated proton (in form of H2F+), and result in an extremely low pKa of −28, making it 10 quadrillion (1016) times stronger than pure sulfuric acid. This acid, as well as several related ones, are so strong that they protonate otherwise inert compounds like hydrocarbons; an early superacid was given the nickname (and eventual tradename) of "magic acid" after a 1966 demonstration in which this fluorinated acid dissolved a paraffin wax candle. Hungarian-American chemist George Olah received the 1994 Nobel Prize in chemistry for investigating such reactions. Noble gas compounds The noble gases are generally non-reactive because they have complete electronic shells, and until the 1960s, no chemical bond with a noble gas was known. In 1962, Neil Bartlett reported the first chemical compound of xenon, xenon hexafluoroplatinate. Later in 1962, xenon was reported to react directly with fluorine to form the di- and tetrafluorides. Since then, xenon hexafluoride, various oxyfluorides, and their derivatives have been prepared. Krypton, xenon's lighter homolog, also forms difluoride and a few more complicated fluorine-containing compounds; the possibility of the existence of tetrafluoride and hexafluoride has been debated. Radon, xenon's heavier homolog has been shown to readily react with fluorine to form a solid compound, generally thought to be radon difluoride, but its exact structure has not been clearly established; if radon were not as radioactive and difficult to collect, its chemistry could be at least as extensive as xenon's. The lightest noble gases do not form stable binary fluorides. Argon, however, reacts in extreme conditions with hydrogen fluoride to form argon fluorohydride. Helium and neon do not form any stable chemical compounds at all, but helium fluorohydride has been observed and it is unstable in gas phase, but it may be stable under enormous pressure. Neon is considered to be even less reactive than helium, and is not expected to form a stable compound capable of synthesis. Highest oxidation states: fluorine versus oxygen Elements generally do exhibit their highest oxidation state in binary compounds with fluorine. For several elements, the highest oxidation state has been reported only in a few compounds, including the fluoride, but for some elements, the highest known oxidation state has been reported exclusively for fluorides. For groups 1–5, 10, 13–16, the highest oxidation states of oxides and fluorides are the same. Differences are only seen for chromium, copper, mercury, and for the groups 7–9 and the noble gases. In fluorides, elements can achieve relatively low oxidation states that are, however, hard to achieve. For example, no binary oxide is known for krypton, while krypton difluoride is well-studied. In addition, mercury tetrafluoride is the only compound for a group 12 element with an oxidation state above +2, while the CoF+4 cation, is the only observed species with a cobalt atom at the +5 oxidation state. For oxidation states above +6 however, binary fluorides are not accessible largely due to the limited space around the central atom available for fitting more than six fluorine atoms. For oxides, achieving an oxidation state of +8 only requires four oxygen atoms around the central atom. As such, ruthenium octafluoride is unlikely to be ever synthesized, while ruthenium tetroxide has even found an industrial use. Organic The carbon–fluorine chemical bond in organofluorine compounds is the strongest bond in organic chemistry. This C–F bond stability and the low polarizability of the molecules containing it are the most important factors contributing to the great stability of the organofluorines. The carbon–fluorine bond of the smaller molecules is formed in three principal ways: fluorine replaces a halogen or hydrogen, or adds across a multiple bond. The direct reaction of hydrocarbons with fluorine gas can be dangerously reactive, so the temperature may need to be lowered even to −150 °C (−240 °F). Hydrogen fluoride or "solid fluorine carriers", compounds that can release fluorine upon heating, notably cobalt trifluoride,  may be used instead. After the reaction, the molecular size is not changed significantly, as the elements have very similar van der Waals radii. Direct fluorination becomes even less important when it comes to organohalogens or unsaturated compounds reactions, or when a prefluorocarbon is desired (then HF-based electrolysis is typically used). In contrast, the fluoropolymers are formed by polymerizing free radicals; other techniques used for hydrocarbon polymers do not work in that way with fluorine. The range of organofluorine compounds is diverse, reflecting the inherent complexity of organic chemistry. A vast number of small molecules exist with varying amounts of fluorine substitution, as well as many polymers—research into particular areas is driven by the commercial value of applications. Small molecules Monofluoroalkanes (alkanes with one hydrogen replaced with fluorine) may be chemically and thermally unstable, yet are soluble in many solvents; but as more fluorines are in instead of hydrogens, the stability increases, while melting and boiling points, and solubility decrease. While the densities and viscosities are increased, the dielectric constants, surface tensions, and refractive indices fall. Partially fluorinated alkanes are the hydrofluorocarbons (HFCs). Substituting other halogens in combination with fluorine gives rise to chlorofluorocarbons (CFCs) or bromofluorocarbons (BFCs) and the like (if some hydrogen is retained, HCFCs and the like). Properties depend on the number and identity of the halogen atoms. In general, the boiling points are even more elevated by combination of halogen atoms because the varying size and charge of different halogens allows more intermolecular attractions. As with fluorocarbons, chlorofluorocarbons and bromofluorocarbons are not flammable: they do not have carbon–hydrogen bonds to react and released halides quench flames. When all hydrogens are replaced with fluorine to achieve perfluoroalkanes, a great difference is revealed. Such compounds are extremely stable, and only sodium in liquid ammonia attacks them at standard conditions. They are also very insoluble, with few organic solvents capable of dissolving them. However, if a perfluorocarbon contains double or triple bonds (perfluoroalkenes or -alkynes), it becomes very reactive towards ligand accepting, even less stable than corresponding hydrocarbons. Difluoroacetylene, which decomposes even under liquid nitrogen temperatures,  is a notable example. If such a molecule is asymmetric, then the more fluorinated carbon is attacked, as it holds positive charge caused by the C–F bonds and is shielded weakly. Perfluorinated compounds, as opposed to perfluorocarbons, is the term used for molecules that would be perfluorocarbons—only carbon and fluorine atoms—except for having an extra functional group (even though another definition exists ). They share most of perfluorocarbon properties (inertness, stability, non-wettingness and insolubility in water and oils, slipperiness, etc.),  but may differ because of the functional group properties, although the perfluorocarbon tail differ the group-specific properties as compared to those of hydrocarbon-tailed compounds. Perfluoroalkanic acids organic acids may also be seen as perfluorinated compounds, with perfluoroalkyl joined to a carboxyl group. As the fluorines are added to the acid, its strength grows: consider acetic acid and its mono-, di-, and trifluoroacetic derivatives and their pKa values (4.74, 2.66, 1.24, and 0.23). This happens because with fluorines, the anion formed after giving the proton off becomes stable, an effect caused by fluorine's great inductive effect. Because of this, the trifluoro derivative is 33,800 times stronger an acid than acetic. Similarly, the acidity is greatly increased for other perfluorocarboxyl acids, as well as the amines (which are not acids but become less basic if fluorinated). The perfluoroalkanesulfonic acids are also very notable for their acidity. The sulfonic acid derivative, trifluoromethanesulfonic acid, is comparable in strength to perchloric acid. These compounds lower surface energy; for this reason, they, especially perfluorooctanesulfonic acid (PFOS, formerly the active component in brand "Scotchgard") have found industrial use as surfactants (see above). If a perfluorinated compound has a fluorinated tail, but also a few non-fluorinated carbons (typically two) near the functional group, it is called a fluorotelomer. Industrially, such compounds are treated as perfluorinated. The chain end may similarly be attached to different functional groups (via the hydrogenized terminal carbon), such as hydroxyl resulting in fluorotelomer alcohols, sulfonate resulting in fluorotelomer sulfonates, etc. Polymers Fluoropolymers are similar in many regards with smaller molecules; adding fluorine to a polymer affects the properties in the same manner as in small molecules (increasing chemical stability, melting point, reducing flammability, solubility, etc.). Each fluoropolymer has own characteristic properties. The simplest fluoroplastic is polytetrafluoroethylene (PTFE, DuPont brand Teflon), which is a simple linear chain polymer with the repeating structural unit: –CF2–. PTFE has a backbone of carbons single bonded in a long chain, with all side bonds to fluorines. It contains no hydrogens and can be thought of as the perfluoro analog of polyethylene (structural unit: –CH2–). PTFE has high chemical and thermal stability, as expected for a perfluorocarbon, much stronger than polyethylene. Its resistance to van der Waals forces makes PTFE the only known surface to which a gecko cannot stick. The compound, however, lacks an ability to transform upon melting, which is not a problem for various PTFE derivatives, namely FEP (fluorinated ethylene propylene, structurally similar to PTFE but has some fluorines replaced with the –CF3 groups) or PFA (perfluoroalkoxy, some fluorines replaced with –OCF3). They share most properties with PTFE, but there are still differences, namely maximum usage temperature (highest for the non-flexible PTFE). There are other fluoroplastics other than perfluorinated. Polyvinylidene fluoride (PVDF, structural unit: –CF2CH2–), is an analog of PTFE with half the fluorines. PVF (polyvinyl fluoride, structural unit: –CH2CHF–) contains one one-fourth the fluorines of PTFE. Despite this, it still has many properties of more fluorinated compounds. PCTFE (polychlorotrifluoroethylene, structural unit: –CF2CFCl–) is another important compound. It differs from PTFE by having a quarter of fluorines replaced with chlorines, yet this difference brings even greater hardness, creep resistance, and moisture persistence. Mild fluorination of polyethylene gives does not make all of the plastic lose its hydrogens for fluorine; only a thin layer (0.01 mm at maximum) is then affected. This is somewhat similar to metal passivation: the bulk properties are not affected, but the surface properties are, most notably, a greater vapor barrier. Therefore, they are a cheaper alternative to the perfluoro plastics if only surface is important. Nafion is a structurally complicated polymer. It has a PTFE-like backbone, but also contains side chains of perfluoro ether that end in sulfonic acid (–SO2OH) groups. It also possesses great chemical stability, while exact properties vary with morphology. However, because of the difficult chemical structure, it is also relatively easily converted to an ionomer (shows conductivity) by adding cations like Na+ or by converting into the sulfonic acid rather than the given sulfonyl fluoride. The conductivity is due to that the main carbon chain separates from the side chains, thus forming polar and non-polar regions. This form is also very hydroscopic.