Fluorine

oxygen ← fluorine → neon - ↑ F ↓ Cl 9F Periodic table
Appearance
Liquid fluorine at cryogenic temperatures
General properties
Name, symbol, number	fluorine, F, 9
Pronunciation	/ˈflʊəriːn/, /ˈflʊərɪn/, /ˈflɔːriːn/
Element category	halogen
Group, period, block	17, 2, p
Standard atomic weight	18.9984032(5)[1]
Electron configuration	1s2 2s2 2p5[2]
Electrons per shell	2, 7[2] (Image)
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa) 1.696[3] g/L
Liquid density at b.p.	1.505[4] g·cm−3
Melting point	53.53 K, −219.62 °C, −363.32[5] °F
Boiling point	85.03 K, −188.12 °C, −306.62[5] °F
Critical point	144.00 K, 5.220[6] MPa
Heat of fusion	0.51[7] kJ·mol−1
Heat of vaporization	3.27[7] kJ·mol−1
Molar heat capacity	(Cp) (21.1 °C) 825[8] J·mol−1·K−1 (Cv) (21.1 °C) 610[8] 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[9] (Pauling scale)
Ionization energies (more)	1st: 1681.0[10] kJ·mol−1
	2nd: 3374.2[10] kJ·mol−1
	3rd: 6050.4[10] kJ·mol−1
Covalent radius	60[11] pm
Van der Waals radius	147[12] pm
Miscellanea
Crystal structure	cubic
Crystal structure note	the structure refers to solid fluorine, at boiling point, 1 atm[13]
Magnetic ordering	diamagnetic[14]
Thermal conductivity	0.02591[15] W·m−1·K−1
CAS registry number	7782-41-4[2]
Most stable isotopes
Main article: Isotopes of fluorine
iso NA half-life DM DE (MeV) DP 18F trace 109.77 min β+ (97%) 0.633 18O ε (3%) 1.656 18O 19F 100% 19F is stable with 10 neutrons reference[16]
· r

Fluorine is the chemical element with atomic number 9, represented by the symbol F. It is the lightest element of the halogen column of the periodic table and has a single stable isotope, fluorine-19. At standard pressure and temperature, fluorine is a pale yellow gas composed of diatomic molecules, F₂. In stars, fluorine is rare compared to other light elements. On Earth, fluorine is more common; it is the 13th most abundant element in the crust.

Fluorine's most important mineral, fluorite, was first formally described in 1530, in the context of metal smelting. The mineral's name derives from the Latin verb fluo, which means "stream" or "flow," because fluorite was added to metal ores to lower their melting points. Suggested to be a chemical element in 1811, "fluorine" was named after the source mineral. Several chemists (called the "fluorine martyrs") died in accidents while trying to isolate the element. In 1886, French chemist Henri Moissan succeeded in generating fluorine. His method of electrolysis remains the only industrial production method of fluorine gas. The main use of elemental fluorine, uranium enrichment, was developed during the Manhattan Project.

Because of the difficulty in making elemental fluorine, the vast majority of commercial fluorine is never reduced to the element. Fluorine has the highest electron affinity of any element but chlorine and is an extremely strong oxidizing agent. Fluorine forms stable compounds, fluorides, with all elements except helium and neon. Hydrofluoric acid, in contrast to hydrochloric acid, is only a weak acid, but it is nonetheless extremely corrosive. Fluorides of lighter metal elements are ionic compounds (salts). Heavier metal elements such as uranium form volatile coordination compounds (molecules with several fluorine atoms surrounding a metal atom).

Organic fluorine compounds tend to have high chemical and thermal stability and water-repellent properties. Several have large-scale commercial application, such as the fluorinated plastic polytetrafluoroethylene ("Teflon") used in cookware. Another major application is fluorinated refrigerants. Here, traditional chlorofluorocarbons ("Freons"), which cause ozone depletion, have been largely replaced by hydrofluorocarbons. A few organofluorine compounds are synthesized in microorganisms and plants. But, although it helps prevent tooth decay, fluorine is not an essential mineral for mammals. Several fluorine compounds, as well as elemental fluorine itself, are dangerously toxic. Nevertheless, an increasing number of pharmaceuticals (about 10% of new drugs) contain fluorine.

Characteristics

Electronic structure

A fluorine atom has nine protons and nine electrons, arranged in electronic configuration [He]2s²2p⁵, one fewer than neon.^[17] 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 relatively high effective nuclear charge. Because of this, fluorine is reluctant to ionize and has an attraction for one more electron to achieve the extremely stable neon-like arrangement.^[17]

Fluorine's first ionization energy (energy required to remove an electron to form F⁺) is 1,681 kilojoules per mole, which is higher than for any other element except neon and helium.^[10] The second and third ionization energies of fluorine are 3,374 and 6,147 kilojoules per mole, respectively.^[10] Fluorine's electron affinity (energy released by adding an electron to form F^–) is 328 kilojoules per mole, which is higher than that of any other element except chlorine.^[18] Fluorine has a relatively small covalent radius, about 60 picometers, which is more than neon and less than oxygen.^[19]

Isotopes

Fluorine occurs naturally on Earth exclusively in the form of its only stable isotope, fluorine-19,^[20] which makes the element monoisotopic and mononuclidic. Seventeen radioisotopes have been synthesized: mass numbers 14–18 and 20–31.^[21] 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.^[22]

The lightest fluorine isotopes, 14–16, decay by electron capture. F-17 and F-18 undergo beta plus decay (positron emission). All isotopes heavier than the stable fluorine-19 decay by beta minus mode (electron emission). Some of them also decay by neutron emission.^[21]

Only one nuclear isomer (long-lived excited nuclear state), fluorine-18m, has been characterized.^[23] 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 numbers 14–16, 28, and 31.^[23]

Physical properties

Fluorine form diatomic molecules that are gaseous at room temperature. The density is about 1.3 times that of air.^{[24][note 1]} Though sometimes cited as yellow-green, fluorine gas is actually a very pale yellow. Its color can only be observed in concentrated fluorine gas when looking down the axis of long tubes. It appears transparent when observed from the side in normal glass tubes or if allowed to escape into the atmosphere.^[26] The element has a "pungent" characteristic odor that is noticeable in concentrations as low as 20 ppb.^[27]

Fluorine condenses to a bright yellow liquid at −188.1 °C (−306.6 °F),^[5] a comparable temperature to the boiling points of oxygen and nitrogen. Fluorine solidifies at −219.6 °C (−363.3 °F)^[5] into a cubic structure, called beta-fluorine. This phase is transparent and soft, with significant disorder of the molecules. At −227.5 °C (−377.5 °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 than to the other halogens.^[28][29]

Elemental fluorine

An 1892 observation of fluorine gas (middle), compared to air (left) and chlorine (right): the observation was done looking down 5 meter long tubes.  Solid alpha-fluorine's crystal structure: fluorine molecules lie in shingled layers.

Chemical reactivity

Halogen bond energies (kJ/mol)^[30]

X	XX	HX	BX3	AlX3	CX4
F	159	574	645	582	456
Cl	243	428	444	427	327
Br	193	363	368	360	272
I	151	294	272	285	239

Fluorine's chemistry is dominated by its tendency to gain an electron. It is the most electronegative element and a strong oxidant.^[9] The removal of an electron from a fluorine atom requires so much energy that no known oxidant can oxidize fluorine to any positive oxidation state.^[31]

Fluorine gas is highly reactive with other substances because of its oxidizing power—which leads to strong bonds with other atoms—and because of the relative weakness of the fluorine–fluorine bond. That bond energy is similar to the easily cleaved oxygen–oxygen bonds of peroxides or nitrogen–nitrogen bonds of hydrazines and significantly weaker than those of dichlorine or dibromine molecules.^[32] The covalent radius of fluorine in difluorine molecules, about 71 picometers, is significantly larger than that in other compounds because of the weak bonding between fluorine atoms.^[33]

External videos
	Fluorine video from the University of Nottingham: Cold gas impinging on several substances causes bright flames.

Reactions with fluorine are often sudden or explosive.^[24] Many generally non-reactive substances such as bricks, water, and frozen meat will burn with a bright flame in a jet of fluorine gas.^[24]

Fluorine forms compounds, fluorides, with all elements except neon and helium. All of the elements up to einsteinium, element 99, have been checked except for astatine and francium.^[34] Fluorine is also known to form compounds with rutherfordium, element 104,^[35] and seaborgium, element 106.^[36] Several heavy radioactive elements have not been fluoridated because of their extreme rarity, but such reactions are theoretically possible.^[37]

All metals react with fluorine, but conditions vary with the metal. Often, the metal must be powdered because many metals passivate (form protective layers of the metal fluoride that resist further fluoridation). Alkali metals react with fluorine violently. Alkaline earth metals react at room temperature as well but do not release as much heat. The noble metals ruthenium, rhodium, palladium, platinum, and gold react least readily, requiring pure fluorine gas at 300–450 °C (575–850 °F).^[38]

Fluorine reacts explosively with hydrogen in a manner similar to that of alkali metals.^[39] The halogens react readily with fluorine gas^[40] as does the heavy noble gas radon.^[41] The lighter noble gases xenon and krypton can be made to react with fluorine under special conditions and argon will combine with hydrogen fluoride.^[42] Nitrogen, with its very stable triple bonds, requires electric discharge and high temperatures to combine directly with fluorine.^[43]

Origin and occurrence

In the universe

Abundance in the Solar System^[44]

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.^[45][46]

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.^[45][46]

Even though fluorine, due to its chemical activity, does not exist in its elementary state on Earth, it can be found in the interstellar medium.^[47] Fluorine cations have been seen in planetary nebulae and in stars, including our Sun.^[48]

On Earth

Fluorine is the thirteenth most common element in Earth's crust, comprising between 600 and 700 ppm of the crust by mass. Due to its reactivity, it is found as fluoride ion rather than as the element. Three minerals exist that are industrial sources of fluorine.^[49][50]

• Fluorite (also called fluorspar) is the most significant source of commercial fluorine. It is used directly in metal smelting and is the source mineral for hydrogen fluoride, which is further converted to other fluorine chemicals. Fluorite is a colorful mineral associated with hydrothermal deposits. It is common and found worldwide. China currently supplies more than half of the world's demand; Mexico is the second-largest producer. The last United States mine, in Illinois, shut down in 1995.^[50][51]

• Fluorapatite 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 mostly discarded as waste. Only in the United States is there significant recovery. There the hexafluorosilicates produced as byproducts are used to supply municipal water fluoridation.^[50][51]

• Cryolite 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.^[50][51]

Notable 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.^[50]

Organofluorines have been observed in volcanic eruptions and in geothermal springs. Their ultimate origin (physical formation under geological conditions or initial biological production and deposition in sediments) is unclear. They are not a commercially important source of fluorine, but are trace environmental contaminants whose amount is being studied.^[52]

History

"Fluorine" is a word that ultimately derives from the Latin verb fluo, meaning "flow." The mineral fluorite, a natural form of calcium fluoride, was first mentioned in 1529 by Georgius Agricola, who named it after its use as a "flux"—an additive that helps melt ores and slags during smelting.^[53][54] Agricola first named the mineral "fluorspar" as a latinization of the German Flußspat.^[55] Since then, the mineral has been renamed "fluorite," although "fluorspar" is still sometimes used.^[56]

In 1670, Heinrich Schwanhard, a German glass cutter, found that a vapor was produced when he treated "fluorspar" with strong acid. He used this property to develop a new art form, covering portions of glassware with protective varnish then exposing it to vapor, ending with clear figures on a cloudy background.^[57][58]

Andreas Sigismund Marggraf made the first recorded preparation of "fluoric acid" (hydrofluoric acid in modern nomenclature) in 1764 when he heated fluorite with sulfuric acid in glass, which was greatly corroded by the product.^[59] In 1771, Swedish chemist Carl Wilhelm Scheele repeated this reaction.^[59] In 1810, French physicist André-Marie Ampère suggested that the acid was a compound of hydrogen with an unknown element, analogous to chlorine.^[60] Fluorite was then shown to be mostly composed of calcium fluoride.^[61][62]

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. (However, Greek, Russian, and several other languages use the ftor or deratives, which comes from Greek φθόριος (phthorios), meaning "destructive.") The new Latin name (fluorum) gave the element its current symbol, F, although the symbol Fl is seen in early papers.^[63]

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 it could only be prepared electrolytically and because the gas reacted with most materials. The generation of elemental fluorine from fluorspar proved to be exceptionally dangerous, killing or blinding several early experimenters. Jean Dussaud referred to these men as "fluorine martyrs," a term still used.^[62] Edmond Frémy decided that it might be preferable to pass the electric current through hydrofluoric acid instead. Unfortunately, hydrogen fluoride was only available in an aqueous solution. Frémy therefore devised a method for producing anhydrous hydrogen fluoride by acidifying potassium hydrogen fluoride (now known as potassium bifluoride or KHF₂). Unfortunately, KHF₂ will not pass an electric current.^[57]

French chemist Henri Moissan, formerly one of Frémy's students, then took up the battle. After trying many different elements, he took Frémy's potassium bifluoride (created from hydrogen fluoride) and mixed it into a solution of hydrogen fluoride (so that the insulating potassium bifluoride could have an electric charge passed through the fluorides which formed the liquid of the solution) in equipment crafted from a mixture of platinum and iridium (even more non-reactive with fluorine than platinum), with flurospar stoppers (which wouldn't carry an electrical current and wouldn't react with fluorine). After 74 years of effort by other chemists, on 26 June 1886, Moissan reported the isolation of elemental fluorine.^[57][64] He later discovered a less expensive method of producing fluorine, by using copper equipment overlain with copper fluoride. He received the 1906 Nobel Prize in chemistry for his fluorine isolation and invention of the electric arc furnace.^[65]

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 started a joint venture in 1930 to market the new product, but soon DuPont took over the business. The substance proved to be a marketplace hit, rapidly replacing earlier, more toxic, refrigerants and growing the overall market for kitchen refrigerators.^[59][67][68]

In 1938, polytetrafluoroethylene (DuPont brand name Teflon) was discovered by accident by a recently-hired DuPont Ph.D., Roy J. Plunkett. While working with tetrafluoroethylene gas, he noticed missing weight. Scraping down his container, he found white flakes of a new-to-the-world polymer. 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.^[59][67][68]

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.^[69] 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.^[61] Because uranium hexafluoride releases small quantities of corrosive fluorine, the separation plants were built with special materials. All pipes were coated with nickel, which forms a protective fluoride layer on its surface after exposure to fluorine. Joints and flexible parts were fabricated from Teflon.^[68][70]

In 1958, a DuPont research manager in the Teflon business, Bill Gore, left the company because of its unwillingess 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 expanded PTFE membrane which led to the large Gore-tex business in breathable rainwear. The company developed many other uses of PTFE and grew the market for the substance.

In the 1970s and 80s, 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.^[71][72]

Industry and applications

Fluorite mining was estimated in 2003 to be a $550 million industry, extracting 4.5 million tons per year. Two grades of mined fluorite exist: metspar (for iron smelting) is less than 97% CaF₂; acidspar (for HF production) is more than 97% CaF₂. Approximately half of mined fluorite is used to help molten metal flow, especially in iron smelting. The rest is converted to hydrofluoric acid by reaction with sulfuric acid. HF is primarily used to produce organofluorides and synthetic cryolite. Only about 1% of mined fluorine is ever converted to the element.^[51][59]

Uses of inorganic fluorides

About 3 kg of metspar grade fluorite, added directly to the batch, are used for every tonne of steel made. The fluoride ions from CaF₂ lower the melt's temperature and viscosity (make the liquid runnier). (The calcium helps remove sulfur and phosphorus, but other additives such as lime do as well.) 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.^[51]

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.^[51][73]

One third of HF (one sixth of mined fluorine) is used to make synthetic cryolite (sodium hexafluoroaluminate) and aluminum trifluoride. These compounds are used in the electrolysis of aluminium. About 23 kg are required for every tonne of aluminum. These compounds are also used as a flux for glass.^[51][74]

Fluorosilicate is the next most significant inorganic fluoride formed from HF. The material is used for water fluoridation and is also an intermediate for synthetic cryolite. Other inorganic fluorides made in large quantities include cobalt difluoride (for organofluorine synthesis), nickel difluoride (electronics), lithium fluoride (a flux), sodium fluoride (water fluoridation), magnesium difluoride (flux and optical use), potassium fluoride (flux), and ammonium fluoride (various).^[51]

Uses of organofluorines

Refrigerant gases

Traditionally chloroflorocarbons (CFCs) have been the predominant fluorinated organic chemical. CFCs (and HFCs and HCFCs) are identified by a system of numbering that explains the amount of fluorine, cholorine, carbon and hydrogen in the molecules. The term Freon has been colloquially used for CFCs and similar halogenated molecules. However, strictly speaking this is just a DuPont brand name; 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).^[51]

Production of CFCs grew strongly through the 1980s, primarily for refrigeration and air conditioning but also for propellants, and solvents. However, 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-80s peak, with remaining use primarily as an intermediate for other chemicals. The banning of CFCs initially depressed the overall demand for fluorite. However, 21st century production of the source mineral has recovered to 1980s levels.^[51]

Hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) now serve as replacements for CFC refrigerants, even though 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).^[75][51]

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.^[76]

Polymers and coatings

Fluoropolymers were, as of 2008, a $1.4 billion market in the United States alone.^[77] PTFE (polytetrafluoroethylene) is 85% of the world's fluoropolymer production. The term Teflon is sometimes used genarically for the substance, but is a DuPont brand name—other PTFE producers exist and also DuPont sometimes uses the Teflon brand for non-PTFE 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 ETFE (ethylene tetrafluoride, the starting point for PTFE).^[78]

The largest application for PTFE is in electrical insulation. The material has excellent dielectric properties and is not flammable. It is also used extensively in the chemical process industry where corrosion resistance is needed: to coat pipes and in tubing and gaskets. PTFE withstands attack from all chemicals other than alkali metals. Another major use in architectural fabric (PTFE-coated fiberglass cloth used for stadium roofs and such). The major consumer application is in non-stick cookware.^[78]

Major PTFE applications

PTFE dielectric separating core and outer metal in a specialty coaxial cable	First Teflon frying pan, 1961	The interior of the Tokyo Dome. The roof is PTFE-coated fiberglass, air-supported.[79]

When stretched with a jerk, PTFE films makes a material with fine pores: ePTFE (expanded PTFE) membrane. The term "Gore-tex" is sometimes used generically for this membrane, 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 drawn into fibers which are used in pump packing (seals) and industrial bag house filters for industries with corrosive exhausts.^[78]

Other fluoroplymers 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.^[80][81][78]

Elastomers (rubber-like substances) exist that are composed of mixtures of several fluoropolymers. Chemical-resistant O-rings are the primary application. Fluoroelastomers tend to be more stiff and less durable than conventional elastomers. Viton is a prominent example.^[78][82]

Fluorinated ionomers 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 ton 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.^{[83][84][85][86]}

Fluorinated surfactants are small organofluorine molecules, principally used in DWR (durable water repellent). DWR is a finish (thin coating) put on fabrics that makes them lightly rain resistant, that makes water bead. First developed in then 1950s, by 1990 fluorosurfacts were 90% of the DWR industry. In that same year, the market for fluorosurfactants was $100 million in the U.S. textile industry alone. 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). Scotchgard is a prominent brand.^[87][88]

Uses of fluorine gas

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.^[89] 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.^[51] About 7,000 tons per year of fluorine gas are used for this application.

The second largest application for fluorine gas is sulfur hexafluoride, which is used as an 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 switchear, has no PCB (a hazardous chemical). 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 Moissan, remains the commercial practice. About 6,000 tons per year of fluorine gas are consumed.^[90][51]

Several compounds made from elemental fluorine serve the electronics industry. Tetrafluoromethane,^[91] is used for plasma etching in semiconductor manufacturing,^[92] flat panel display production, and microelectromechanical systems fabrication.^[93][94][89] Nitrogen trifluoride is increasingly used for cleaning equipment at display manufacturing plants. The hexafluorides of rhenium and tungsten are used for chemical vapor deposition (CVD) of thin metal films onto semiconductors. Elemental fluorine, itself, is used sometimes for cleaning equipment.^[51]

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 pharmacueticals. BF₃ and SbF₃ are not fluorinating agents, but they are formed from fluorine gas and used as catalysts for some organic reactions.^[51]

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.^[95][96][97]

Production of fluorine gas

Commercial producers of fluorine gas continue to use the method of electrolysis pioneerd by Moissan, with some modifications in the cell design. Chemical routes to the elemental form are curiosities that rely on mechanisms to release fluorine gas from compounds previously made using fluorine gas.

Electrolytic synthesis

Several thousand tons of elemental fluorine are produced annually by electrolysis of potassium bifluoride in hydrogen fluoride.^[51]

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).^[59] 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.^[98]

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.^[99]

Pure fluorine gas may be stored in steel cylinders where the inside surface is passivated by a metal fluoride layer that resists further attack.^[98][59] In the G7 countries, about 17,000 tonnes (17,000,000 kg) of fluorine are produced per year by 11 companies.^[89] Fluorine is relatively inexpensive, costing about $5–8 per kilogram when sold as uranium hexafluoride or sulfur hexafluoride. Because of difficulties in storage and handling, the price of pure fluorine gas is much higher.^[89]

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. It involved the reaction of potassium hexafluoromanganate and antimony pentafluoride at 150 °C, in an atmosphere of hydrogen fluoride:

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.^[100] The manganese(IV) fluoride has to be prepared by reaction with fluorine gas itself,^[101] or with krypton difluoride,^[102] which is synthesized by reaction with elemental fluorine. This reaction is therefore not an industrially viable way to produce fluorine. Fluorine can by also be chemically synthesized by the reaction of hexafluoronickelate ion NiF2−
6^[49] with the fluorides of the heavier noble gases (krypton^[103] and xenon^[104]), which can only be produced by a reaction of krypton or xenon with fluorine gas.

Biological aspects

Living organisms

Fluoride is not considered an essential mineral element for mammals and humans, although its role in prevention of tooth decay is well-established.^[105] Sodium fluoride, tin(II) fluoride, and, most commonly, sodium monofluorophosphate, are used in toothpaste. These or related compounds, such as fluorosilicates, are added to many municipal water supplies, a process called water fluoridation, which has been controversial since its inception in 1945.^[105][106] Small amounts of fluoride may be beneficial for bone strength, but this is an issue only in the formulation of artificial diets.^[107]

Biologically synthesized organofluorines have been found in microorganisms and plants,^[52] but not in animals.^[108] The most common example is fluoroacetate, which is used as a defense against herbivores by at least 40 plants in Australia, Brazil and Africa.^[109] Other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate, all of which are believed to be biosynthesized from the intermediate fluoroacetaldehyde.^[108] The enzyme adenosyl-fluoride synthase is capable of biologically synthesizing the carbon–fluorine bond.^[110]

Pharmaceuticals, agricultural chemicals, and poisons

Several important pharmaceuticals contain fluorine.^[112] 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 and activation. For example, an aromatic ring may add to prevent the metabolism of a drug, but this presents a safety problem, because enzymes in the body metabolize some aromatic compounds into poisonous epoxides. Substituting a fluorine into a para position, however, protects the aromatic ring and prevents the epoxide from being produced.^[113] Adding fluorine to biologically active organics increases their lipophilicity, because the carbon–fluorine bond is even more hydrophobic than the carbon–hydrogen bond. This effect often increases a drug's bioavailability due to increased cell membrane penetration.^[114] Since the carbon–fluorine bond is strong, organofluorides are generally very stable, although the potential of the fluorine to be released as a fluoride leaving group is heavily dependent on the its position in the molecule.

Of drugs that have been commercialized in the past 50 years, 5–15% contain fluorine, and the percentage of currently available fluorine-containing drugs is increasing.^[115] For example, fludrocortisone is one of the most common mineralocorticoids, a class of drugs that mimics the actions of aldosterone. The anti-inflammatories dexamethasone and triamcinolone, which are among the most potent of the synthetic corticosteroids class of drugs, contain fluorine.^[116] Several inhaled general anesthetic agents, including the most commonly used inhaled agents, also contain fluorine. Examples include sevoflurane, desflurane, and isoflurane, which are hydrofluorocarbon derivatives.^[115]

Many SSRI antidepressants are fluorinated organics,^[117] including citalopram, escitalopram, fluoxetine, fluvoxamine, and paroxetine. Fluoroquinolones are a commonly used family of broad-spectrum antibiotics.^[118] 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.^[119]

In addition to pharmaceuticals, an estimated 30% of agrochemical compounds contain fluorine.^[120] Because of these, water from agricultural sites contaminates rivers with runoff organofluorines. Synthetic sodium fluoroacetate has been used as an insecticide, especially against cockroaches, and is effective as a bait-poison against mammalian pests.^[121] Several other insecticides contain sodium fluoride, which is much less toxic than fluoroacetate.^[122]

Isotope use in scanning

Compounds containing fluorine-18, a radioactive isotope that emits positrons, are often used in PET scanning, because its half-life of about 110 minutes is long by positron-emitter standards. One such species is 2-deoxy-2-(¹⁸F)fluoro-D-glucose, commonly abbreviated as ¹⁸F-FDG.^[123][124] In PET imaging, ¹⁸F-FDG can be used for the assessment of glucose metabolism in the brain and for imaging tumors in oncology. This radiopharmaceutical is retained by cells and is taken up by tissues with a high need for glucose, such as the brain and most types of malignant tumors.^[125] Tomography can thus be used for diagnosis, staging, and monitoring treatment of many types of cancers, particularly in Hodgkin's disease, 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, which allows it to make measurements quickly, comparable with a similar effect based on hydrogen-1.^[126] Fluorine-19 is commonly used in NMR study of protein structures and conformational changes.^[127]

Chronology

Because groundwater contains fluorine ions, organic items such as bone that are buried in soil will absorb those ions over time. As such, it is possible to determine the relative age of an object by comparing the amount of fluoride with another object found in the same area. It is important as a separation technique in intra-site chronological analysis and inter-site comparisons.^[128]

However, if no actual age of any object is known, the ages can only be expressed in terms of one of the objects being older or younger than the other. The fluctuating amount of fluoride found in groundwater means the objects being compared must be in the same local area in order for the comparisons to be accurate. This technique is not always reliable, given that not all objects absorb fluorine at the same rates.^[129]

Hazards

Elementary state, fluoride ion, and fluoroacetate

Elemental fluorine is a highly toxic, corrosive oxidant, and is extremely reactive to organic material (except for perfluorinated substances) even at very low concentrations and can cause ignition at higher ones.^[131] Significant irritation to humans can be caused by concentration of fluorine of 25 ppm; at this and higher concentrations fluorine attacks the eyes, respiratory tract, lungs, liver and kidneys. At a concentration of 100 ppm, human eyes and noses are irritated and seriously damaged.^[132]

Soluble fluorides are moderately toxic. In the case of the simple salt sodium fluoride, the lethal dose for most adult humans is estimated at 5 to 10 g, which is equivalent to 32 to 64 mg/kg elemental fluoride/kg body weight.^{[133][134][135]} A toxic dose that may lead to adverse health effects is estimated at 3 to 5 mg/kg of fluoride.^[136] The fluoride ion is somewhat toxic, in part because of its ability to form, by equilibration, small amounts of hydrogen fluoride in water. This mobile uncharged species diffuses across cell membranes to attack intracellular calcium. The fluoride ion is readily absorbed by the stomach, intestines and excreted through 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.^[137] Ingested fluoride initially acts locally on the intestinal mucosa, where it forms hydrofluoric acid in the stomach.^[138] Thereafter it binds with calcium and interferes with various enzymes.^[138] Excess fluoride consumption can lead to skeletal fluorosis, which currently affects millions of people.^[139]

Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluoride,^[140] or (more rarely) rodenticides containing sodium fluoroacetate ("Compound 1080").^[141] Currently, most fluoride poisonings are due to the ingestion of fluoride-containing toothpaste.^[138] Malfunction of water fluoridation equipment has occurred several times, including an Alaskan incident, which affected nearly 300 people and killed one person.^[142]

Hydrofluoric acid

Hydrofluoric acid is a contact poison, and must be handled with extreme care far beyond that accorded to other mineral acids, even the analogous hydrochloric acid, HCl. 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. 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.^[143]

Once absorbed into blood through the skin, hydrogen fluoride reacts with blood calcium and may cause cardiac arrest.^[144] Formation of insoluble calcium fluoride possibly causes both a fall in calcium serum and the strong pain associated with tissue toxicity.^[145] 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.^[146]

Hydrofluoric acid exposure is often treated with calcium gluconate, a source of Ca²⁺ that binds with the fluoride ions. Hydrogen fluoride chemical burns to the skin can be treated with a water wash and 2.5% calcium gluconate gel^[147][148] or special rinsing solutions.^[149][150] However, because it is absorbed, medical treatment is necessary; in some cases, amputation may be required.^[146]

Environmental concerns

Chlorofluorocarbons and bromofluorocarbons have come under strict environmental regulation due to their long residence times in the atmosphere, and their contribution to ozone depletion. 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,^[151] and have been widely used as replacements for halocarbons containing chlorine and bromine. Hydrofluorocarbons and perfluorocarbons are greenhouse gases about 4,000 to 10,000 times as potent as carbon dioxide.^[152] Sulfur hexafluoride exhibits an even stronger effect, having 20,000 times the global warming potential of carbon dioxide.^[152]

Because of the strength of the carbon–fluorine bond, many synthetic fluorocarbons and fluorocarbon-based compounds are persistent in the environment. The fluorosurfactants perfluorooctanesulfonic acid (PFOS) and perfluorooctanoic acid (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.^[153] The potential health effects of PFOA to humans are unclear; its tissue distribution in humans is unknown, but studies in rats suggest it is likely to be present primarily in the liver, kidney, and blood, as it is absorbed easily via the gastrointestinal tract in rats. PFOA has been shown not to metabolize in the body, and, unlike chlorinated hydrocarbons, it is neither genotoxic nor lipophilic. It binds to serum albumin and is excreted primarily from the kidney.^[154]

Compounds

Fluorine exists in the −1 oxidation state in all compounds except for elemental fluorine, where the atoms are bonded to each other and thus at oxidation state 0. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds. Higher bonding can occur; for example boron monofluoride features a triple bond.^[155] 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, non-metals, and even noble gases; as well as a diverse set of organic compounds.^{[note 2]}

Inorganic compounds

Inorganic acids

Unlike other hydrohalic acids, such as hydrochloric acid, hydrofluoric acid is only a weak acid in water solution, with acid dissociation constant (pK_a) equal to 3.18.^[156] Despite its weakness as an acid in water, hydrogen fluoride is very corrosive, even attacking glass. Because of the basicity of the fluoride ion, soluble fluorides give basic water solutions.

HF's weakness as an acid is seemingly paradoxical considering the polar of the HF bond. Fluoride ions and hydrogen fluoride forms strong hydrogen bonds, unlike the other halides and hydrogen halides. Fluoride ion tends to participate in homoconjugation, forming the highly stable bifluoride ion (HF₂^-) With less basic solvents such as dry acetic acid, these interactions become less important and hydrofluoric acid then becomes the strongest of the hydrohalic acids.^[157][158]

Fluorine forms a family of inorganic acids that contain only a hydrogen, fluorines, and some other element as the central atom. They are generally very strong because of the stable fluorinated anion. One such acid, fluoroantimonic acid (HSbF₆), is a "superacid" and the strongest acid known.^[159] It has an extremely low pK_a of −31.3 and is 20 quintillion (2×10¹⁹) times stronger than pure sulfuric acid.^[159] Fluoroantimonic acid is so strong that it protonates otherwise inert compounds like hydrocarbons. Hungarian-American chemist George Olah received the 1994 Nobel Prize in chemistry for investigating such reactions.^[160]

Metal fluorides

Metal fluorides have similarities with other metal halides and with metal oxides but are more ionic. In general, the other halides (chlorides, bromides, iodides) are more similar to each other than to the fluorides, which often adopt bonding and structures more like oxides.^[161] Metal fluorides are solids,^[162] liquids,^[163] or gases^[164] at room temperature.

The fluorides of transition metal elements 25-29

Manganese difluoride   Iron trifluoride   Cobalt difluoride   Nickel difluoride   Copper difluoride

Ionic metal fluorides

The solubility of ionic fluorides varies greatly but tends to decrease as the charge on the metal ion increases.

The alkali metals form monofluorides. All have the sodium chloride (rock salt) structure and are soluble.^[165] Because the fluoride anion is highly basic, many alkali metal fluorides form bifluorides with the formula MHF₂. Sodium and potassium bifluorides are significant to the chemical industry.^[166] Among other monofluorides, only silver(I)^[167] and thallium(I)^[168] fluorides are well-characterized. Both are very soluble, unlike the other halides of those metals.

The alkaline earth metals form difluorides that, except for beryllium's, are insoluble.^[169] In contrast, the alkaline earth chlorides are readily soluble.^[63] Several transition metal difluorides, such as those of copper(II) and nickel(II), are soluble.^[63]

Many metals form trifluorides, such as iron, the rare earth elements, and the metals in the aluminium and scandium columns of the periodic table. No trifluoride is soluble in water, but several are soluble in other solvents.^[170]

While metal tri- and lower fluorides are all ionic solids, some tetrafluorides are ionic and some molecular. For example, zirconium tetrafluoride is an ionic solid,^[171] but germanium tetrafluoride is a molecular gas.^[172]

Molecular metal fluorides

Metal penta- and higher fluorides are all molecular and volatile. This behavior contrasts with the corresponding oxides. Oxygen is a weaker oxidant and inherently more likely to form covalent bonds, but it only forms molecules with four metals (manganese heptoxide, technetium heptoxide, ruthenium tetroxide and osmium tetroxide). Fluorine forms molecules with twelve metals because its small size and single charge as an ion allows surrounding metal atoms with more fluorines than oxygen can. (See fluoride volatility.)

Vanadium, niobium, and tantalum lie in a periodic table column that typically reaches +5 as the highest oxidation state. They form pentafluorides as their highest fluoride. Many metals that form hexafluorides also can form pentafluorides.

The metals that make hexafluorides include nine metals in the center of the periodic table (molybdenum, technetium, ruthenium, rhodium, tungsten, rhenium, osmium, iridium, and platinum) along with elements 92–94: uranium, neptunium, and plutonium. At room temperature, tungsten hexafluoride is a gas. Molybdenum hexafluoride and rhenium hexafluoride are liquids. The rest are volatile solids. Metal hexafluorides are oxidants because of their tendency to release fluorines: for example, platinum hexafluoride was the first compound to oxidize molecular oxygen^[173] and xenon.^[174]

Rhenium is the only metal known to form bonds with seven fluorides, which is the record for number of charged ligands for a charge-neutral metal compound.^[175] Rhenium heptafluoride adopts a pentagonal bipyramid molecular geometry. Calculations shows that the currently unknown but perhaps possible compounds osmium heptafluoride and iridium heptafluoride will also have this structure.^[176]

Nonmetal fluorides

The nonmetal fluorides are volatile molecules that do not always follow the octet rule of bonding. For instance boron trifluoride has only six electrons around the central boron atom (and thus an incomplete octet). However, the later Period 2 elements form fluorides that follow the octet rule: carbon tetrafluoride,^[177] nitrogen trifluoride,^[178] and oxygen difluoride.^[179][180] Lower periods, however, may form fluorides that are hypervalent molecules, such as phosphorus pentafluoride.^[181] The reactivity of such species varies greatly: sulfur hexafluoride is inert, while chlorine fluorides are oxidants, but there are some tendencies and trends based on periodic table columns.

Boron trifluoride is a planar molecule. It is a notable Lewis acid: because of its incomplete octet, it attracts Lewis bases such as ammonia or another fluoride ion which can donate two more electrons.^[182][183]

Carbon tetrafluoride and silicon tetrafluoride have tetrahedral structures. CF₄ is relatively inert and stable, but SiF₄ is a weak acid that is not thermally stable.^[177]

Among pnictogens (nitrogen's periodic table column), reactivity and acidity of fluorides increases down the group. Bismuth is an exception—its fluorides are less acidic then those of antimony because bismuth pentafluoride is polymeric,^[184] and bismuth trifluoride is ionic (bismuth is a borderline metal).^[185] Nitrogen forms no pentafluoride, although tetrafluoroammonium ion, NF+
4, with nitrogen in the formal oxidation state of +5, is known.^[186]

Chalcogens (oxygen's periodic table column) show analagous characteristics: hexafluorides increase in acidity and reactivity down the group.

The halogens form hypervalent compounds. Chlorine^[187] and bromine^[188] make pentafluorides, both strong fluorinators. Iodine may be fluorinated up to iodine heptafluoride.^[189] Astatine is not well-studied, and astatine fluoride has not been produced, even though this should be possible.^[190]

The polymeric structure of bismuth pentafluoride

Structure of a (BiF5)n chain	Packing of chains

Noble gas compounds

The noble gases are generally non-reactive because they have fully filled electronic shells, which are extremely stable. Until the 1960s, no chemical bond with a noble gas was known. In 1962, Neil Bartlett used fluorine-containing platinum hexafluoride to react with xenon. He called the compound he prepared xenon hexafluoroplatinate, but since then the product has been revealed to be mixture of different chemicals. Bartlett probably synthesized a mixture of monofluoroxenyl(II) hexafluoroplatinate, [XeF]⁺[PtF₆]^–, monofluoroxenyl(II) undecafluorodiplatinate, [XeF]⁺[Pt₂F₁₁]^–, and trifluorodixenyl(II) hexafluoroplatinate, [Xe₂F₃]⁺[PtF₆]^–.^[191] Bartlett's fluorination of xenon has been called one of the ten most beautiful experiments in the history of chemistry.^[192] Later in 1962, xenon was reacted directly with fluorine to form the di- and tetra- fluorides. Since then, chemists have made extensive efforts to form other noble gas fluorides.

Helium and neon form no stable chemical compounds. Helium, forms a temporary helium fluoride hydride that is stable for a few nanoseconds.^[193][194] Neon, the least reactive element,^{[note 3]}forms a metastable chemical compound, neon monofluoride, NeF^•.^[196]

Argon forms no binary fluoride, but reacts in extreme conditions with hydrogen fluoride, to form argon fluoride hydride.^[197] Argon also forms a short-lived, metastable binary argon monofluoride, ArF^•, which is used in the argon fluoride laser.^[198]

Krypton forms a well-characterized binary noble gas–fluorine compound: krypton difluoride.^{[199][200][201]} Krypton tetrafluoride was reported in 1963,^[202] but was subsequently shown to be a mistaken identification.^[203][204] A krypton monofluoride radical and cation have been observed at low temperature.^[204]

Xenon has the most known noble gas–fluorine compounds. Binary compounds include xenon difluoride, xenon tetrafluoride, and xenon hexafluoride, as well as their derivatives.^{[199][200][201]} Xenon forms several oxyfluorides, such as xenon oxytetrafluoride, XeOF₄, by reaction of xenon hexafluoride with water.^[205][206]

Radon readily reacts with fluorine to form a solid compound, which is generally thought to be radon difluoride. The exact composition is uncertain—the compound is prone to decompose.^[207] Calculations indicate that radon difluoride may be ionic, unlike all other binary noble gas fluorides.^[208]

Ununoctium, the last currently known group 18 element, is predicted to form ununoctium difluoride, UuoF₂, and ununoctium tetrafluoride, UuoF₄, which is likely to have the tetrahedral molecular geometry.^[209] However, only a few atoms of ununoctium have been synthesized,^[210] and the chemical properties have not been examined.

Highest oxidation states

Elements frequently have their highest oxidation state in the form of a binary fluoride. Several elements show their highest oxidation state only in a few compounds, one of which is the fluoride; and some elements' highest known oxidation state is seen exclusively in a fluoride.^[176]

Fluorine was the first element able to oxidize a group 12 element to an oxidation state above +2; making the element's d-electrons participate in bonding.^[211] Mercury(IV) fluoride was produced by this reaction, the first ever mercury(IV) compound; its discovery has heated the debate over whether mercury, cadmium, and zinc are transition metals.^[212] Another unique oxidation state available for fluorine only is gold(V).^[213] It is only known in the hexafluoroaurate(V) ion, which can be synthesized indirectly under extreme conditions, and gold(V) fluoride, which is obtained during hexafluoroaurate(V) decomposition. Because of fluorine's high oxidizing potential it was suggested that gold heptafluoride contained gold(VII),^[214] but current calculations show that the claimed AuF₇ molecule was AuF₅·F₂.^[215]

Aside from these examples, fluorine is the only element that is known to oxidize the respective elements to palladium(VI),^[216] platinum(VI),^[217] copper(IV),^[218] silver(IV),^[219] nickel(IV),^[220] iridium(VI),^[221] and krypton(II).^[201] It is possible that element 113, ununtrium, will be the first boron group element to form a species in the +5 oxidation state, the fluorine-based hexafluoroununtrate(V), UutF−
6;^[222] no other ununtrium(V) species is expected.

For groups 1—6, 13, 14, and 16, the highest oxidation states of oxides and fluorides are always equal. Differences are only seen in groups 7—11, mercury, nitrogen, halogens, and the noble gases. Fluorination allows elements to achieve relatively low^{[note 4]} oxidation states that are, however, hard to achieve. For example, no binary oxide is known for krypton, but krypton difluoride is well-studied.^[201] At the same time, very high oxidation states are known for oxygen-based species only. For example, ruthenium and xenon octafluorides have not been synthesized, but ruthenium tetroxide and xenon tetroxide are both well-studied.^[223] Also, the halogens tend to achieve higher oxidation states as oxides than as fluorides. The main problem that prevents fluorine from forming the highest states in covalent hepta- and octafluorides is that it is hard to attach such a large number of ligands around a single atom; the number of ligands is halved in analogous oxides.^{[176][note 5]}

Organic compounds

Organofluorine compounds are defined by having a carbon–fluorine chemical bond. This bond is the strongest bond in organic chemistry and is very stable.^[224] Fluorine replaces hydrogen in hydrocarbons even at room temperature. After the reaction, the molecular size is not changed significantly. Organofluorine compounds are synthesized via both direct reaction with fluorine gas, which can be dangerously reactive, or reaction with fluorinating reagents such as sulfur tetrafluoride.^[68] 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.^[68]

Small molecules

Organic compounds with all C-H bonds changed to C-F are called fluorocarbons or perfluorocarbons. They are extremely insoluble in water. The derivatives of alkanes (hydrocarbons with single bonds) have higher density and melting and boiling points and are more thermally and chemically stable. In contrast, fluorinating hydrocarbons which contain double bonds (alkenes) or triple bonds (alkynes) gives rise to very easily attacked molecules. Difluoroacetylene, which is explosive even at low temperatures, is a notable example. Partially fluorinated alkanes exist as well, the hydrofluorocarbons (HFCs).

The Fowler process produces commercial perfluoroalkanes. Hydrocarbon feedstocks are passed over a bed of cobalt trifluoride, which loses fluorine and is reduced to cobalt difluoride. Fluorine reacts with the hydrocarbon producing perfluorocarbon and HF. The cobalt bed is periodically refluorinated to the trifluoride, in a separate step, by reaction with concentrated HF.

Substituting other halogens in combination with fluorine gives rise to chlorofluorocarbons (CFCs) or bromofluorocarbons (BFCs) and the like. (Or 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.^[225] As with fluorocarbons, chlorofluorocarbons and bromofluorocarbons are not flammable: they don't have carbon–hydrogen bonds to react and released halides quench flames.^[225]

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. The perfluoro parts of the molecule tend to be hydrophobic and slippery. The functional group allows reactions, attachment to solids, solubility, etc. Electrochemical fluorination (ECF), essentially electrolysis in HF, is the commercial production method.

The fluorosurfactants are notable perfluorinated compounds. They have a medium length straight chain of perfluoroalkane terminated by an acid group. This type of arrangement, like a fatty acid, gives rise to a combination of properties. For instance perfluorooctanesulfonic acid (PFOS, formerly the active component in brand "Scotchgard") has eight fluorinated carbons in a row ending with a sulfonic acid group.

For fluorinated organic acids, the large inductive effect of the trifluoromethyl group results in high acid strengths, which may be comparable to mineral acids. In these compounds, the anion's affinity for the acid proton is decreased by the acid's fluorine content, which increases its affinity for the extra electron left when the acidic proton leaves. For example, acetic acid is a weak acid, with pK_a equal to 4.76, while its fluorinated derivative, trifluoroacetic acid has pK_a of −0.23, giving it 33,000 times greater acid strength.^[226]

Polymers

Fluoropolymers are organic polymers ("plastics") containing fluorine. Polytetrafluoroethylene (PTFE, DuPont brand Teflon) 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. The non-stick nature of PTFE results from the repulsion of highly charged fluorine atoms in polymeric chains. Its resistance to van der Waals forces makes PTFE the only known surface to which a gecko cannot stick.^[227][228]

Several other fluoropolymers exist that are more complicated structurally than PTFE. In general, they have similar properties to Teflon, but are less thermally and chemically stable, while also more processable (lower melting). Three fluoropolymers that contain hydrogen are polyvinylidene fluoride (PVDF, structural unit: –(CF₂CH₂)–), polyvinyl fluoride (PVF, structural unit: –(CF₂CHF)–), and ethylene tetrafluoroethylene (ETFE, structural unit: –(CF₂CF₂CH₂CH₂)–).

FEP (fluorinated ethylene propylene) and PFA (perfluoroalkoxy) are similar to PTFE in not containing any hydrogens and contain a PTFE backbone. FEP has branching chains of fluoroalkane. PFA also has branched fluoralkane but with an ether (oxygen) link.

Nafion is a complicated polymer, structurally. It has a PTFE-like backbone, but also contains side chains of perfluoro ether that end in sulfonic acid (SO₃H) groups. It has similar chemical stability as PTFE, but because of the acid side chains, is also an ionic conductor, the first ionomer.

See also

• Halogens—the periodic group to which fluorine belongs
• Caesium—the least electronegative element, fluorine's "opposite"

Books View or order collections of articles	Fluorine Period 2 elements Halogens Chemical elements (sorted alphabetically) Chemical elements (sorted by number)
Portals Access related topics	Chemistry portal
Find out more on Wikipedia's Sister projects	[//commons.wikimedia.org/wiki/Special:Search/Fluorine Images and media] from Commons [//en.wiktionary.org/wiki/Special:Search/fluorine Definitions] from Wiktionary [//en.wikibooks.org/wiki/Special:Search/Wikijunior:The_Elements/Fluorine Textbooks] from Wikibooks [//en.wikiversity.org/wiki/Special:Search/Fluorine_atom Learning resources] from Wikiversity

Notes

1. ^ Density of air at 100 kilopascal and 0 °C is 1.2724 grams per liter.^[25]
2. ^ In this article, metalloids are not treated separately from metals and nonmetals, but among elements they are closer to. For example, germanium is treated as a metal, and silicon as a nonmetal. Antimony, bismuth and polonium are included for comparison among nonmetals, even though they are closer to metals chemically than to nonmetals. The noble gases are treated separately from nonmetals and hydrogen is discussed in the Inorganic acids section. P-block period 7 elements have not been studied and thus are not included. This is illustrated by the image to the right: the dark gray elements are metals, the green ones are nonmetals, the light blue ones are the noble gases, the purple one is hydrogen, and the light gray elements have unknown properties.
3. ^ Helium and neon are the only elements that have no known "stable" compounds; i. e., the compounds that do not decay on specific conditions. Calculations show that helium, unlike neon, may form some compounds that will not decay over time in specific conditions; additionally, clathrates are known for every noble gas but neon.^[195]
4. ^ There is no general line where oxidation states are "relatively low" or "relatively high," they rely on specific elements (and defined only for elements that have highest oxides and fluorides are in different oxidation states); in general, +7 and +8 are high, while +4 and below are low. States +5 and +6 rely on element properties, like atomic radius; for a small nitrogen atom, +5 is what is called "high" here, but for larger palladium and platinum +6 is still "low."
5. ^ Note that aside from the molecular one, other forms of the highest fluorides are not considered to be possible. Such compounds are extremely unlikely to be ionic due to large sixth, seventh, eight oxidation energies of all elements, which make the reactions that may produce such compounds highly unfavorable. An even higher fluoride of any element seems to be impossible for similar reasons.^[176]

References

Bibliography

Diatomic chemical elements Hydrogen H2 Nitrogen N2 Oxygen O2 Fluorine F2 Chlorine Cl2 Bromine Br2 Iodine I2 Astatine At2

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 Post-transition metals Metalloids Other nonmetals Halogens Noble gases Unknown chem. properties
Large version