oxygen ← fluorine → neon - ↑ F ↓ Cl 9F Periodic table
Appearance
gas: very pale yellow liquid: bright yellow 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 (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	144.4 K, 5.215[4] MPa
Heat of vaporization	6.51[3] kJ·mol−1
Molar heat capacity	(Cp) (21.1 °C) 825[4] J·mol−1·K−1 (Cv) (21.1 °C) 610[4] 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[2] (Pauling scale)
Ionization energies (more)	1st: 1,681[6] kJ·mol−1
	2nd: 3,374[6] kJ·mol−1
	3rd: 6,147[6] kJ·mol−1
Covalent radius	64[7] pm
Van der Waals radius	135[8] pm
Miscellanea
Crystal structure	cubic
Crystal structure note	the structure refers to solid fluorine, at boiling point, 1 atm[9]
Magnetic ordering	diamagnetic[10]
Thermal conductivity	0.02591[11] 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+ (96.9%) 0.634 18O ε (3.1%) 1.656 18O 19F 100% 19F is stable with 10 neutrons reference[12]
v t e · 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. In Earth's crust, fluorine is more common, being 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, which means "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, the "fluorine martyrs," died in accidents while trying to isolate the element. In 1886, French chemist Henri Moissan succeeded. His method of electrolysis remains the industrial production method. The main use of elemental fluorine, uranium enrichment, was developed during the Manhattan Project.

Fluorine is the most electronegative element and forms stable compounds, fluorides, with all elements except helium and neon. Because of the difficulty in making elemental fluorine, the vast majority of commercial fluorine is never oxidized to the element. Hydrofluoric acid is the key intermediate for the $13 billion fluorochemical industry. Although only a weak acid, HF eats through glass and is a worse burn danger than conventional strong acids. The fluorides of lighter metal elements are ionic compounds (salts); those of heavier metal elements are volatile molecular compounds. The largest uses of inorganic fluorides are steel making and aluminium refining.

Organic fluorine compounds tend to have high chemical and thermal stability and repel water. The largest use is in refrigerant gases ("Freon"). Although traditional chlorofluorocarbons—which cause ozone depletion—are widely banned, the replacements, hydrochlorofluorocarbons and hydrofluorocarbons, still contain fluorine. Polytetrafluoroethylene ("Teflon") is the most important fluoropolymer and is used in electrical insulation, chemical-resistant parts, stadium roofs, raincoats and cookware. Fluoride is not an essential mineral for mammals, but does prevent tooth decay. A growing fraction of drugs contain fluorine. Lipitor and Prozac are notable examples.

Characteristics[link]

Physical properties[link]

Fluorine forms diatomic molecules that are gaseous at room temperature. The density is about 1.3 times that of air.^{[13][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.^[15] The element has a "pungent" characteristic odor that is noticeable in concentrations as low as 20 ppb.^[16]

Fluorine condenses to a bright yellow liquid at −188 °C (−307 °F),^[17] close to the condensation temperatures of oxygen and nitrogen. Fluorine solidifies at −220 °C (−363 °F)^[17] 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 than to the other halogens.^[18][19]

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

Electronic structure[link]

A fluorine atom has nine protons and nine electrons, arranged in electronic configuration [He]2s²2p⁵, one fewer than neon.^[20] 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.^[20]

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.^[21] The second and third ionization energies of fluorine are 3,374 and 6,147 kilojoules per mole, respectively.^[21] 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.^[22] Fluorine has a revised relatively small covalent radius, slightly less than 60 picometers, which is less than oxygen and also less than neon.^[23]

Although an individual fluorine atom has one unpaired electron, in molecular fluorine all the electrons are paired. Because of this, elemental fluorine is diamagnetic (slightly repelled by magnets). In contrast, the molecules of neighbor element oxygen are paramagnetic (attracted to magnets).^[10] The measured value of fluorine's magnetic susceptibility is −9.6×10⁻⁶ (cgs), which is close to theoretical predictions. The experimental result was not accomplished until 1999 because of the difficulties in handling fluorine gas as well as the need to specially purify the fluorine of any trace of paramagnetic oxygen.^[24]

Electron arrangement in atomic and molecular fluorine
Structure of the fluorine atom: seven electrons are not self-shielding and result in a large atomic size.	Molecular orbitals of the fluorine molecule: all electrons are paired, resulting in diamagnetism ("nonmagnetic").

The fluorine molecule is unique among the halogens in having a bond order of exactly 1 (exactly single bonded diatomic molecules). For comparison, chlorine has a bond order of 1.12. In the heavier halogens, the presence of a d-subshell (first occurring at the third energy sublevel) allows for the slightly more than single bonding. Fluorine's outermost electrons are located at the second energy sublevel where there is no d subshell.^[25]

Isotopes[link]

Fluorine occurs naturally on Earth exclusively in the form of its only stable isotope, fluorine-19,^[26] which makes the element monoisotopic and mononuclidic. Seventeen radioisotopes have been synthesized: mass numbers 14–18 and 20–31.^[27] 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.^[28]

The lightest fluorine isotopes, 14–16, decay by 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). Some of them also decay by neutron emission.^[27]

Only one nuclear isomer (long-lived excited nuclear state), fluorine-18m, has been characterized.^[29] 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.^[29]

Chemical reactivity[link]

Fluorine's chemistry is dominated by its tendency to gain an electron. It is the most electronegative element and a strong oxidant. 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]

Reactions with fluorine are often sudden or explosive. Many generally non-reactive substances 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.^[13][34]

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.^[35] Fluorine is also known to form compounds with rutherfordium, element 104,^[36] and seaborgium, element 106.^[37] Several heavy radioactive elements have not been fluoridated because of their extreme rarity, but such reactions are theoretically possible.^[38]

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

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). The alkali metals react with fluorine violently, while the 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).^[39]

Fluorine reacts explosively with hydrogen in a manner similar to that of alkali metals.^[40] The halogens react readily with fluorine gas^[41] as does the heavy noble gas radon.^[42] The lighter noble gases xenon and krypton can be made to react with fluorine under special conditions and argon will combine with hydrogen fluoride.^[43] Nitrogen, with its very stable triple bonds, requires electric discharge and high temperatures to combine directly with fluorine.^[44]

Origin and occurrence[link]

In the universe[link]

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.^[46][47]

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.^[46][47]

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.^[48] Fluorine cations have been seen in planetary nebulae and in stars, including our Sun.^[49]

On Earth[link]

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 industrially relevant sources: fluorite, fluorapatite, and cryolite.^[50][51]

• 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; many countries currently produce fluorite. 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. Canada also exited production in the 1990s. The United Kingdom has declining fluorite mining and has been a net importer since the 1980s.^{[51][52][53][54][55]}
• 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 municipal water fluoridation.^[51][52]
• 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.^[51][52]

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.^[51]

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.^[56]

History[link]

Smelting illustration from Agricola's De re metallica, where fluorite was first described

"Fluorine" is a word that derives from an invented Latin term for the main source mineral. Fluorite, was first mentioned in 1529 by Georgius Agricola, who described it as a flux—an additive that helps melt ores and slags during smelting.^[57][58] 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 stones, he devised the noun fluores (from the Latin verb for flow fluo) because they made metal ores flow when in a fire. After Agricola, the name for the mineral evolved to fluorspar (still common) and then to fluorite.^[55][59][60]

Some sources claim that the first production of hydrofluoric acid was by Heinrich Schwanhard, a German glass cutter, in 1670.^[61] A peer-reviewed study of Schwanhard's writings, though, showed no specific mention of fluorite and just discussion of an extremely strong acid. It was hypothesized that this was probably nitric acid or aqua regia, capable of etching soft glass.^[62][63] 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.^[64] In 1771, Swedish chemist Carl Wilhelm Scheele repeated this reaction.^[64] In 1810, French physicist André-Marie Ampère suggested that the acid was a compound of hydrogen with an unknown element, analogous to chlorine.^[65] Fluorite was then shown to be mostly composed of calcium fluoride.^[66]

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 name ftor or derivatives, which was suggested by André-Marie Ampère and comes from the Greek φθόριος (phthorios), meaning "destructive.")^[67] The new Latin name (fluorum) gave the element its current symbol, F, although the symbol Fl is seen in early papers.^[68] The symbol Fl is now being used for the element flerovium.^[69]

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 proved to be exceptionally dangerous, killing or blinding several early experimenters. Jean Dussaud referred to these men as "fluorine martyrs," a term still used.^[66]

Edmond Frémy thought that passing the 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.^[61]

Henri Moissan, the first to isolate elemental fluorine.

French chemist Henri Moissan, formerly one of Frémy's students, then took up the battle. 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 resistant to fluorine than pure platinum) with fluorite stoppers. After 74 years of effort by other chemists, on 26 June 1886, Moissan reported the isolation of elemental fluorine.^[61][70] 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^[71]

Moissan later devised 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.^[72][73][74]

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.^{[64][75][76][77]}

In 1938, polytetrafluoroethylene (DuPont brand name Teflon) was discovered by accident by a recently-hired DuPont PhD, 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.^[64][76][78]

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.^[79] 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.^[62] 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.^[76][80]

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.^[81] 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.^[82]

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.^[83][84]

Industry and applications[link]

The global market for fluorochemicals was about $13 billion per year as of 2006–2008.^[85][86] 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 tons per year by 2015.^[87]

The largest market is the United States. Western Europe is the second largest. Asia Pacific is the fastest growing region of production.^[87] China in particular has experienced significant growth as a fluorochemical market and is becoming a producer of them as well.^[85][88][89]

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. Two main grades of mined fluorite exist, with about equal production of each. Acidspar is at least 97% CaF₂; metspar is much lower purity.^[52][64][89]

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. Only about 1% of mined fluorite is ever converted to elemental fluorine.^[52][64]

Fluorine industry supply chain: major sources, intermediates and applications. Click for links to related articles.

Uses of inorganic fluorides[link]

Aluminium smelting process: cryolite (a fluoride) is required to make it work.

About 3 kg (6.5 lbs) 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 from fluorite helps remove 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. Metspar varies in purity and the trend has been to higher CaF₂ content in recent years, especially in developed countries. In the United States, commercial metspar is generally at least 80% pure and may be as high as 93%. Worldwide, metspar is generally at least 60% pure.^[52]

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.^[52] Two ceramic grades of fluorite also exist that are at least 90% or 95% CaF₂.

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.^[52]

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. About 23 kg (51 lbs) are required for every tonne of aluminium. These compounds are also used as a flux for glass.^[52]

Fluorosilicate is the next most significant inorganic fluoride formed from HF. The material is used for water fluoridation, but the more significant use is as 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), potassium fluoride (flux), and ammonium fluoride (various).^[52] Sodium and potassium bifluorides are significant to the chemical industry.^[90]

Anti-reflection coated glass: 45° and 0°

MgF₂ and CaF₂ are important optical materials. Synthetic calcium difluoride is used for lenses and windows for spectrometers because of its ability to pass a very wide spectrum of light. It was used widely in semiconductor manufacturing equipment until about 2005, but is no longer suitable for the deeper UV wavelengths currently used. Magnesium difluoride is widely used as an antireflection coating for spectacles and optical equipment. That 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.^[91][92][93]

On the galaxy scale, hydrogen fluoride is a very promising way to track the reservoirs of gas, especially hydrogen. The compound, present in all of the Universe, emits radiation that is not emitted by the hydrogen itself. The compound commonly used today, carbon monoxide, even contains no hydrogen. The new knowledge may become so useful that it can help to find all (not only the largest) gas clouds and even become a standard one day.^[94]

Uses of fluorocarbons[link]

Making organic fluorides is the number one 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.^[52][87] Fluorosurfactants are a small segment in mass but are significant economically because of very high prices.

Refrigerant gases[link]

Traditionally chloroflorocarbons (CFCs) have been the predominant fluorinated organic chemical. CFCs 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).^[52]

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. 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-1980s 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.^[52]

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).^[52]

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.^[95]

Fluoropolymers[link]

Fluoropolymers are less than 0.1% of all polymers produced in terms of weight. However they are more expensive and have higher growth rates than average polymers. 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.^[96][97][98]

PTFE (polytetrafluoroethylene) is 60–80% of the world's fluoropolymer production on a weight basis.^[97][98] 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 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).^[99]

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.^[99]

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.[100]

Microscopic structure of ePTFE

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.^[99]

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.^{[99][101][102]}

Schematic of a PEM fuel cell. Nafion membrane is the key component.

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.^{[103][104][105][106]}

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 and less durable than conventional elastomers.^[99][107]

Fluorinated surfactants[link]

Fluorinated surfactants are small organofluorine molecules, principally used in stain repellents and DWR (durable water repellent). Fluorosurfactants are a large market economically, 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 (2.2 lbs).^[108]

A fabric treated with fluorinated surfactant to make it water-resistant

Fluorosurfactants are 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.^[108]

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).^[109][110]

Uses of fluorine gas[link]

In the G7 countries, about 17,000 tonnes (17,000,000 kg) of fluorine are produced per year by 11 companies.^[111] Fluorine is relatively inexpensive, costing about US$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.^[111] 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.^[111] 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.^[52] About 7,000 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 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.^[112][52]

Several compounds made from elemental fluorine serve the electronics industry. The hexafluorides of rhenium and tungsten are used for chemical vapor deposition (CVD) of thin metal films onto semiconductors. Tetrafluoromethane, is used for plasma etching in semiconductor manufacturing, flat panel display production, and microelectromechanical systems fabrication.^{[113][114][115]} Nitrogen trifluoride is increasingly used for cleaning equipment at display manufacturing plants. Elemental fluorine, itself, is used sometimes for cleaning equipment.^[52]

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.^[52]

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.^[116][117]

Production of fluorine gas[link]

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. A chemical routes to the elemental form was published in 1987.

Electrolytic synthesis[link]

Several thousand tons of elemental fluorine are produced annually by electrolysis of potassium bifluoride in hydrogen fluoride.^[52] 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).^[64] 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.^[118]

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.^[119]

Handling[link]

Pure fluorine gas may be stored in steel cylinders where the inside surface is passivated by a metal fluoride layer that resists further attack.^[118][64] 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). ^[120] 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^[120] although some sources recommend systems made of nickel, monel, and PTFE.^[121]

Chemical routes[link]

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.^[122] To make it, the scientist first needed a perfluoromanganese(IV) complex. A strong oxidizer, it released fluorine gas upon heating to 150 °C when treated with antimony pentafluoride in an atmosphere of hydrogen fluoride:

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.^[123]

Biological aspects[link]

South Africa's gifblaar is one of the few organisms that makes fluorine compounds.

Fluoride is not considered an essential mineral element for mammals and humans.^[124] Small amounts of fluoride may be beneficial for bone strength, but this is an issue only in the formulation of artificial diets.^[125]

Several other biological aspects of fluorine exist. Fluoride is widely used for prevention of cavitities. In pharmaceuticals and agrichemicals, fluorine sees increasing use in new molecules. Both radioactive and natural fluorine isotopes are important in respective scanning applications. A method of archeological dating of bones uses fluoride's tendency to accumulate in calcium-containing organic matter. Oxygen-carrying perfluorocarbons present a possibility for human liquid breathing. 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.

Dental care[link]

Topical fluoride treatment in Panama

Fluoride ions in contact with teeth limit cavities by forming fluorapatite. The enamel of the teeth is partially mineralized, which makes it more resistant to decay. The process works only by direct contact (topical treatment). Fluoride ions that are swallowed do not benefit the teeth.^[126][127]

Water fluoridation is the controlled addition of fluoride to a public water supply in order to reduce tooth decay.^[128] 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^[129] and for about 5.7% of people worldwide.^[130] Although the best available evidence shows no association with adverse effects other than fluorosis, most of which is mild,^[131] water fluoridation has been contentious for ethical, safety, and efficacy reasons,^[130] and opposition to water fluoridation exists despite its support by public health organizations.^[132] 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.^[133] Systematic reviews in 2000 and 2007 showed significant reduction of cavities in children associated with water fluoridation.^[131]

Sodium fluoride, tin(II) fluoride, 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.^[133] 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.^[134]

Pharmaceuticals and agrichemicals[link]

Diflunisal has a half-life of 13 hours, which is much longer than most other non-steroidal anti-inflammatory drugs.^[135]

Several important pharmaceuticals contain fluorine.^[136] 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. 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.^[137] 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.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. 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.^[138]

Many SSRI antidepressants are fluorinated organics,^[139] including citalopram, escitalopram, fluoxetine, fluvoxamine, and paroxetine. Fluoroquinolones are a commonly used family of broad-spectrum antibiotics.^[140] 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.^[141]

In addition to pharmaceuticals, an estimated 30% of agrochemical compounds contain fluorine.^[142] Because of these, water from agricultural sites contaminates rivers with runoff organofluorines.

Scanning[link]

Whole-body PET scan using fluorine-18

Compounds containing fluorine-18, a radioactive isotope that emits positrons, are often used in PET scanning, because F-18's 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.^[143] In PET imaging, ¹⁸F-FDG can be used for asessing glucose metabolism in the brain and for imaging cancer tumors. After injection into the blood, this radiopharmaceutical is taken up by tissues with a high need for glucose, such as the brain and most types of malignant tumors.^[144] Tomography (similar to a CAT scan) can then be used to diagnose (or monitor treatment of) cancers; especially 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. 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.^[145] Fluorine-19 is commonly used in NMR study of metabolism, protein structures and conformational changes.^[146] In addition, inert fluorinated gases have the potential to be a cheap and efficient tool for imaging lung ventilation.^[147]

Chronology[link]

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.^[148]

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 fluoride at the same rates.^[149]

Oxygen transport research[link]

Computer-generated model of nanocrystal of perflubron (red) and gentamicin (white, an antibiotic) in liquid suspension: nanocrystals represent a method of delivering water or fat soluble drugs within a perfluorochemical fluid. The particles may help treat babies with damaged lungs.^[150]

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.^[151]

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) in order to be used as blood.^{[152][153][154]} One such product, Oxycyte, has been through initial clinical trials.^[155][156]

Possible medical uses of liquid breathing 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. Several animal tests have been performed and some human trials.^[157] One effort, by Alliance Pharmaceuticals, reached clinical trials but was abandoned because of insufficient advantage compared to other therapies.^[158]

Other posited applications include deep sea diving and space travel.^{[159][160][161]} The 1989 film The Abyss showed a fictional use of perfluorocarbon for human diving but also filmed a real rat surviving while immersed in perfluorocarbon.^[162] Several other fictional treatments of perfluorocarbon breathing exist.

Poisons, artificial and natural[link]

Synthetic sodium fluoroacetate has been used as an insecticide but is especially effective against mammalian pests.^[163] The name "1080" refers to the catalogue number of the poison, which became its brand name.^[164] 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.^[164] Several other insecticides contain sodium fluoride, which is much less toxic than fluoroacetate.^[165]

Biologically synthesized organofluorines have been found in microorganisms and plants,^[56] but not in animals.^[166] The most common example is fluoroacetate, which is used as a defense against herbivores by at least 40 plants in Australia, Brazil, and Africa;^[164] other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate.^[166] 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 commercial organofluorine synthesis.^[167]

Hazards[link]

Fluorine gas and fluoride ion[link]

U.S. hazard signs for commercially transported fluorine^[168]

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.^[169]

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.^[170] The dose that may lead to adverse health effects is about one fifth the lethal dose.^[171] Chronic excess fluoride consumption can lead to skeletal fluorosis, a disease of the bones that affects millions in Asia and Africa.^[171][172]

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 interfere 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.^[171][173]

Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluoride,^[174] Currently, most calls to poison control centers for possible fluoride poisoning come from the ingestion of fluoride-containing toothpaste.^[171] Malfunction of water fluoridation equipment has occurred several times, including an Alaskan incident, which affected nearly 300 people and killed one.^[175]

Hydrofluoric acid[link]

HF burns, not evident until a day after

Hydrofluoric acid 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. 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.^[176]

Once in the blood, hydrogen fluoride reacts with calcium and may cause cardiac arrest.^[177] Formation of insoluble calcium fluoride possibly causes both a fall in calcium serum and the strong pain associated with tissue toxicity.^[178] 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.^[179]

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^[180][181] or special rinsing solutions.^[182][183] However, because HF is absorbed, medical treatment is necessary; sometimes amputation may be required.^[179]

Environmental concerns[link]

Atmosphere[link]

NASA projection of stratospheric ozone over North America if CFCs had not been banned

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.^[184] 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—that ozone depletion has stopped and recovery is underway.^[185][186]

Hydrochlorofluorocarbons (HCFCs) are current replacements for CFCs; HCFCs have about one tenth the ozone damaging potential (ODP) of CFCs.^[187] 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.^[188]

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.^[189][190]

Biopersistance[link]

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.^[191][192]

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.^[191][192]

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.^[191][192]

Compounds[link]

Fluorine as a straight (a) or bent (b) bridging ligand. Other patterns exist.^[193]

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, although at least one example of a higher order bond exists.^[194] 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[link]

Hydrogen fluoride[link]

HF and H2O similarities
Boiling points of the hydrogen halides (blue) and hydrogen chalcogens (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 (HF) called hydrogen fluoride or, especially in the context of water solutions, hydrofluoric acid. The H-F bond type is one of the few capable of hydrogen bonding (creating extra clustering associations with similar molecules). This influences various peculiar aspects of hydrogen fluoride's properties. In some ways the substance behaves more like water, also very prone to hydrogen bonding, than one of the other hydrogen halides, such as HCl.^{[195][196][197]}

Dashed lines show the hydrogen bonds of HF: The molecules align in short chains in the liquid state and as long chains in the solid state.

Hydrogen bonding amongst 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 and −35 °C (−120–30 °F). HF is fully miscible with water (will dissolve in any proportion), while the other hyrogen 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 degrees Celsius (79 degrees Fahrenheit) above the melting point of pure HF.^[198]

Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in water solution, with acid dissociation constant (pK_a) equal to 3.18.^[199] Because of the basicity of the fluoride ion, soluble fluorides give basic water solutions. HF's weakness as an aqueous acid is paradoxical considering how polar the HF bond is, much more so than the bond in HCl, HBr, or HI. The explanation for the behavior is complicated, having to do with various cluster-forming tendencies of HF, water, and fluoride ion, as well as thermodynamic issues.^{[note 3]} When less basic solvents such as dry acetic acid are used instead of water, these interactions become less important and hydrofluoric acid then becomes the strongest of the hydrohalic acids.^[201][202] Despite its weakness, hydrofluoric acid is very corrosive, even attacking glass.^[203]

Hydrofluoric acid also has a property, called homoconjugation, of becoming stronger (than its dissociation constant would predict) when highly concentrated. This occurs because HF begins to accept fluoride ion, forming bifluoride ion: HF₂^-.^[198]

Dry hydrogen fluoride dissolves low charged metal fluorides readily. Several molecular fluorides also dissolve in HF. Many proteins and carbohydrates can be dissolved in dry HF and can be recovered from it. Most non-fluoride inorganic chemicals react with HF rather than dissolving.^[204]

Metal fluorides[link]

Metal fluorides have similarities with other metal halides but are more ionic. In many respects, metal fluorides differ from other metal halides (chlorides, bromides, iodides), very similar to each other. Instead, fluorides are more similar to oxides, often having similar bonding and structures.^[205] Fluoride and oxide crystal structures often share names.

The metal fluorides show broad trends based on the charge of the metal. Metals in an oxidation state of +3 or lower tend to form ionic fluorides. Metals charged +5 or higher tend to form covalently bonded fluorides: polymers or discrete molecules. The bonding variations have physical implications: metal fluorides can be solids,^[206] liquids,^[207] or gases^[208] at room temperature. The solubility of fluorides varies greatly but tends to decrease as the charge on the metal ion increases.

The fluorides of transition metal elements 25–29
Manganese difluoride	Iron trifluoride	Cobalt difluoride	Nickel difluoride	Copper difluoride

Low oxidation state metal fluorides[link]

Rock salt structure of sodium fluoride: yellow is fluoride, purple is sodium. They are isoelectronic, but fluoride is bigger because its nuclear charge is lower.

The alkali metals form monofluorides. All are soluble and have the sodium chloride (rock salt) structure which is also adopted by some alkaline earth oxides such as CaO.^[209] Because the fluoride anion is highly basic, many alkali metal fluorides form bifluorides with the formula MHF₂. Among other monofluorides, only silver(I)^[68] and thallium(I)^[210] fluorides are well-characterized. Both are very soluble, unlike the other halides of those metals.

Unlike the monofluorides, the difluorides may be either soluble or insoluble. The alkaline earth metals form difluorides that are insoluble.^[68] In contrast, the alkaline earth chlorides are readily soluble.^[68] Several transition metal difluorides, such as those of copper(II) and nickel(II), are soluble.^[68]

Many of the difluorides adopt the fluorite structure, named after calcium fluoride (and also adopted by several metal dioxides such as CeO₂, UO₂, ThO₂, etc.), which surrounds each metal cation with 8 fluorides. Some difluorides adopt the tetragonal rutile structure, named after a form of titanium dioxide and adopted by several other metal dioxides also. The structure is tetragonal and puts metal atoms in octahedral coordination.

Beryllium difluoride is different than the other difluorides. In general, beryllium has a tendency to bond covalently, much more so than the other alkaline earths and its fluoride is partially covalent (although still more ionic than its other halides). BeF₂ has many similarities to SiO₂ (quartz) a mostly covalently bonded network solid. BeF₂ has tetrahedrally coordinated metal and forms glasses (is difficult to crystallize). When crystalline, beryllium fluoride has the same room temperature crystal structure as quartz and shares many higher temperatures structures also.

Beryllium difluoride is very soluble in water,^[68] unlike the other alkaline earths. (Although they are strongly ionic, they do not dissolve because of the especially strong lattice energy of the fluorite structure.) However, BeF₂ has much lower electrical conductivity when in solution or when molten than would be expected if it were fully ionic.^{[211][212][213][214]}

Order and disorder in difluorides
The strong and stable ionic fluorite structure adopted by calcium difluoride and many other difluorides	Disordered structure of beryllium glass (sketch, two dimensions)

The spiral chains in gold trifluoride

Many metals form trifluorides, such as iron, bismuth, the rare earth elements, and the metals in the aluminium and scandium columns of the periodic table. The trifluorides of many rare earths as well as bismuth have the Yttrium(III) fluoride|YF₃]] structure. Trifluorides of plutonium, samarium (at high temperature), and lanthanum adopt the LaF₃ structure. Iron and gallium trifluorides have the FeF₃ structure which is similar to rhenium trioxide. Gold trifluoride adopts a structure of linked AuF₄ squares that align in a helix (spiral chain).^[215]

In contrast to gold's distinctly ionic trifluoride, its trichloride and tribromide are volatile dimeric molecules.^[216] Aluminium trifluoride is a high melting point solid which is a monomer in the gas phase, while its other trihalides are low-melting, volatile molecules or linear polymeric chains that form dimers as gases phase.^[217] No trifluoride is soluble in water, but several are soluble in other solvents.^[218]

ZrF₄, common tetrafluoride structure

The tetrafluorides show a mixture of ionic and covalent bonding. Zirconium, hafnium, plus many of the actinides form tetrafluorides with an ionic structure that puts the metal cation in an 8-coordinate square antiprism. Melting points are around 1000 °C.^{[219][220][221]} Titanium and tin tetrafluorides are polymeric, with melting points below 400 °C. (In contrast, their tetrachlorides are molecular and liquids at room temperature.) Vanadium tetrafluoride has a similar structure to tin's^[222] and decomposes at 100–120 °C to form the trifluoride and the pentafluoride.^[223]

The tetrafluorides of iridium, platinum, palladium, and rhodium all share the same structure which was not known until 1975. They have octahedrally coordinated metal atoms with four of the fluorines shared and two unshared, thus somewhat molecular bonding.^[224] The melting points, where known, are below 300 °C. Manganese tetrafluoride is a solid that decomposes at 70 °C and is not well understood structurally. Germanium tetrafluoride forms tetrahedral molecules and is a gas at room temperature.^[225]

High oxidation state metal fluorides[link]

Metal penta- and higher fluorides are all covalently bonded 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 seven metals (manganese heptoxide, technetium heptoxide, ruthenium tetroxide, osmium tetroxide, plutonium tetroxide, iron tetroxide,^[226] and iridium tetroxide^[227]). Fluorine forms molecules with thirteen metals because its small size and single charge as an ion allows surrounding metal atoms with more fluorines than oxygen can.

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. Bismuth's highest fluoride is a volatile penta species that is a powerful fluorinating agent. In the solid state, it is polymeric, consisting of linear chains of octahedra, sharing axial fluorides. In combination with alkali metals, pentavalent bismuth can form hexafluorobismuthates: M[BiF₆].^{[224][228][229]} Many metals that form hexafluorides also can form pentafluorides. For instance, uranium, well known as a hexafluoride, also forms two different pentafluoride structures. The room-temperature (alpha) form has the same linear chain structure as bismuth pentafluoride. As a molecular (gas) species, UF₅ has square pyramidal structure.

The polymeric structure of bismuth pentafluoride
Structure of a (BiF5)n chain	Packing of chains

Uranium hexafluoride in sealed glass

The metals that make well-characterized 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^[230] and xenon.^[231] Polonium also forms a hexafluoride, but it is understudied.^[229] Chromium also forms an unstable hexafluoride which decomposes at temperatures higher than −100 °C.^[232][233]

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.^[234] Rhenium heptafluoride adopts a pentagonal bipyramid molecular geometry. Calculations shows that the currently unknown but perhaps possible iridium heptafluoride^[235] (report of synthesis is being prepared^[227]) and osmium heptafluoride^[236] will also have this structure.

Osmium octafluoride was first reported in 1913, but in 1958 that compound was shown to be actually osmium hexafluoride.^[237] A 1993 theoretical study predicted very weak bonds in osmium octafluoride and said that it would be difficult to ever detect experimentally. The study predicted, if made, OsF₈ would have Os-F bonds of two different lengths.^[238]

Nonmetal fluorides[link]

The nonmetal binary fluorides are volatile compounds that usually do not 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. Lower periods elements, however, may form fluorides that are hypervalent molecules, such as phosphorus pentafluoride.^[239] The reactivity of such species varies greatly: sulfur hexafluoride is inert, while chlorine trifluoride is far more reactive than even fluorine itself, but there are some tendencies and trends based on periodic table columns:

Lewis dot diagram structures show three formal alternatives for describing bonding in boron monofluoride

• Boron trifluoride is a planar molecule. It is a classic Lewis acid, forming adducts with Lewis bases such as ammonia or another fluoride ion which can donate two more electrons.^[240] Boron monofluoride is an unstable gas that has been isolated at low temperatures. It is a "subfluoride" (compound containing less than normal saturation with fluorine) and is prone to polymerizing. It is stable as a ligand in some metal complexes. The bond order has been described alternately as triple bond^[194] and as 1.4 (intermediate between a single and double bond).^[241]

• Silicon tetrafluoride adopts molecular tetrahedral structure. It is stable on its own against heating or electric spark, but yet reacts with water (even in most air), metals and alkalies, demonstrating weak acidic character.^[242] Reactions with organomagnesium compounds, alcohols, amines, and ammonia yield addiction compounds.^[242] Fluorosilicic acid, its derative, is a strong acid in aqueous solution (anhydrous does not exist).^[243]

• Among the pnictogens (nitrogen's periodic table column) trifluorides, reactivity and acidity of fluorides increases down the group: NF₃ is stable against hydrolysis,^[244] PF₃ hydrolyzes very slowly in moist air,^[245] while AsF₃ completely hydrolyzes.^[244] SbF₃ hydrolyzes only partially due to the increasing ionic character of the bond to fluorine. The compounds are weak Lewis bases, with NF₃ again being an exception.^[244] The pentafluorides of phosphorus^[245] and arsenic^[246] are much more reactive than their trifluorides; antimony pentafluoride is such a strong acid that it holds the title of the strongest Lewis acid,^[246] and its complex with HF gives the strongest superacid known (see below). Nitrogen does not form a pentafluoride, although the tetrafluoroammonium cation (NF+
  4) features nitrogen in the formal oxidation state of +5,^[247] not achieved in any other compound.

• The chalcogens (oxygen's periodic table column) form fluorides similar to pnictogens' in many respects: the two main kinds of fluorides are the tetrafluorides and hexafluorides (although others exist). The tertafluorides are thermally unstable and hydrolyze, and are alos ready to use their lone pair to form adducts to other (acidic) fluorides. Sulfur and selenium tetrafluorides molecules are trigonal pyramids, while TeF₄ is a polymer.^[248] The hexafluorides are the result of direct fluorination of the elements (compare: other hexahalides of the elements do even not exist). They, like the previous group's highest fluorides, increase in reactivity with atomic number: SF₆ is extremely inert, SeF₆ is less noble (for example, reacts with ammonia at 200 °C (400 °F), and TeF₆ easily hydrolyzes to give an oxoacid.^[248] Oxygen is (similarly to nitrogen) out of line: Its highest fluoride is oxygen difluoride,^[248] and fluorine again can (however, for this element only theoretically as of 2012) oxidize it to an uniquely high oxidation state of +4 in a fluorocation, OF+
  3.^[249]

Iodine (purple) can form a compound with seven surrounding fluorines (yellow).

• The halogens (chlorine, bromine, and iodine) all form mono-, tri-, and pentafluorides: XF, XF₃, and XF₅. Of the +7 species, iodine heptafluoride is known;^[250] while chlorine and bromine heptafluorides are not known, the corresponding cations ClF+
  6 and BrF+
  6, extremely strong oxidizers, are.^[251] Many of the halogen fluorides are powerful fluorinators. Chlorine trifluoride is the most powerful oxidizer among interhalogens (possibly even ahead of the pentafluoride).^[252] It is extremely reactive and explosive: it readily fluorinates asbestos and refractory oxides.^[253] Astatine is not well-studied, and although there is a report of a non-volatile astatine monofluoride,^[254] its existence is debated.^[255]

Superacids[link]

Several important inorganic acids contain fluorine. They are generally very strong because of the high electronegativity of fluorine. One such acid, fluoroantimonic acid (HSbF₆), is the strongest acid known.^[256] The acid is not an actual compound of its own, but a mixture of HF and SbF₅, which greatly increases the dissociation of HF by binding fluoride.^[257] It has an extremely low pK_a of −31.3 and is 20 quintillion (2×10¹⁹) times stronger than pure sulfuric acid.^[256] 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.^[258]

Noble gas compounds[link]

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]^–.^[259] Bartlett's fluorination of xenon has been called one of the ten most beautiful experiments in the history of chemistry.^[260] 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.

Xenon tetrafluoride crystals, 1962

Having started the noble gases–fluorine compounds class, xenon is also the noble gas to have the most well-known such compounds. Its binary compounds include xenon difluoride, xenon tetrafluoride, and xenon hexafluoride, as well as their derivatives.^[261] Xenon forms several oxyfluorides, such as xenon oxydifluoride, XeOF₂, by reaction of xenon tetrafluoride with water.^[262] Its upper neighbor, krypton, is the only other one to form a well-established compound: krypton difluoride.^[261][263] Krypton tetrafluoride was reported in 1963,^[264] but was subsequently shown to be a mistaken identification;^[265] the compound seems to be very hard to synthesize now (although even the hexafluoride may exist).^[265] A krypton monofluoride radical^[266] and cation^[267] have been observed at low temperature.

In strong accordance with the periodic trends, radon is notably more reactive to oxidizers, including fluorine. It has been shown to readily react with fluorine to form a solid compound, which is generally thought to be radon difluoride. The exact composition is uncertain as the compound is prone to decompose. Calculations indicate that radon difluoride may be ionic, unlike all other binary noble gas fluorides.^[254]

The lighter noble gases (helium through argon) do not form stable binary fluorides. Argon forms no binary fluoride, but reacts in extreme conditions with hydrogen fluoride, to form argon fluorohydride under extreme conditions; it is the only "stable" (i.e. not decaying chemically over time under certain conditions) argon compound. Argon also forms a short-lived, metastable binary argon monofluoride, ArF^•, which is used in the argon fluoride laser.^[268] Helium and neon do not form any stable chemical compounds at all. Helium forms a temporary helium fluoride hydride that is stable for a few nanoseconds.^[269][270] Neon, the least reactive element,^{[note 4]} forms a metastable chemical compound, neon monofluoride, NeF^•.^[272]

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.^[273] However, only a few atoms of ununoctium have been synthesized,^[274] and its chemical properties have not been examined.

Highest oxidation states[link]

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.

Mercury(IV) fluoride: its formation impacted debate about mercury's status as a transition metal.

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.^[275] 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.^[276] Another unique oxidation state available for fluorine only is gold(V).^[277] 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), but current calculations show that the claimed AuF₇ molecule was AuF₅·F₂.^[278]

Aside from these examples, fluorine is the only element that is known to oxidize the respective elements to platinum(VI),^[279] copper(IV),^[280] silver(IV)^[281] (and even silver(V), again in a perfluorinated species, is possible^[282]), nickel(IV),^[283] and krypton(II).^[263] 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;^[284] 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, and the noble gases. Fluorination allows elements to achieve relatively low^{[note 5]} oxidation states that are, however, hard to achieve. For example, no binary oxide is known for krypton, but krypton difluoride is well-studied.^[263] At the same time, very high oxidation states are known for oxygen-based species only. For the previously mentioned volatile oxides, there are no corresponding hepta- or octafluorides. (For example, ruthenium octafluoride has not been synthesized, while ruthenium tetroxide has even found an industrial use.^[285]) 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.^[286]

Organic compounds[link]

This unreferenced section requires citations to ensure verifiability.

Organofluorine compounds are defined by having a carbon–fluorine chemical bond. This bond is the strongest bond in organic chemistry and is very stable.^[287] 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 cobalt trifluoride.^[76] 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.^[76]

Some organofluorine compounds: (A) fluoromethane (methyl fluoride), used in semiconductor processing; (B) tetrafluoroethane (R-134a), an HFC; (C) dichlorodifluoromethane (R-12 refrigerant), a CFC; (D) trifluoromethanesulfonic acid (triflic acid), a superacid; (E) perfluorooctanesulfonic acid (main component of Scotchgard), a fluorosurfactant; (F) a section of polytetrafluoroethylene (Teflon) polymer

Small molecules[link]

Perfluorodecalin, a perfluorocarbon that is a liquid at room temperature. It boils at a lower temperature than its hydrocarbon analog, decalin.

Organic compounds with all C-H bonds changed to C-F are called fluorocarbons or perfluorocarbons (the "per" prefix signifying presence of as much fluorine as chemically possible). The very different molecular interaction potentials of perfluorocarbons causes these substances to be extremely insoluble in water and also not miscible with hydrocarbons, thus forming three layers of liquid (perfluorocarbon most dense) if liquid perfluorocarbons are mixed with liquid oil and water.

The perfluoro derivatives of alkanes (hydrocarbons with single bonds) have higher density and often higher melting and boiling points, and are also more thermally and chemically stable than their hydrocarbon analogs. However, the very small intermolecular interaction energies typical of perfluorinated molecules causes perfluoroalkanes to have much lower melting and boiling points than hydrocarbons of similar molecular weight (sometimes even lower than the hydrocarbon analogs).

In contrast, the perfluorination of hydrocarbons that 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.

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.

Partially fluorinated alkanes exist as well, the hydrofluorocarbons (HFCs). 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.^[288] As with fluorocarbons, chlorofluorocarbons and bromofluorocarbons are not flammable: they do not have carbon–hydrogen bonds to react and released halides quench flames.^[288]

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.^[289]

Polymers[link]

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.^[290]

The complex unit structure of Nafion

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). 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. Three fluoropolymers that contain hydrogen are polyvinylidene fluoride (PVDF, structural unit: –CF₂CH₂–), polyvinyl fluoride (PVF, structural unit: –CH₂CHF–), and ethylene tetrafluoroethylene (ETFE, structural unit: –CF₂CF₂CH₂CH₂–).

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[link]

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

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Notes[link]

1. ^ Density of air at 100 kilopascal and 0 °C is 1.2724 grams per liter.^[14]
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 is included for comparison among nonmetals, even though it is closer to metals chemically than to nonmetals. The noble gases are treated separately from nonmetals; hydrogen is discussed in the Hydrogen fluoride section and carbon in the Organic compounds 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, the yellow one is carbon, and the light gray elements have unknown properties.
3. ^ See citation for more detailed explanation.^[200]
4. ^ 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.^[271]
5. ^ 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 considered to be "high" here, but for larger palladium and platinum +6 is still considered to be "low."

References[link]