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Bohrium (transition metal)
Hassium (transition metal)
Meitnerium (unknown chemical properties)
Darmstadtium (unknown chemical properties)
Roentgenium (unknown chemical properties)
Copernicium (transition metal)
Ununtrium (unknown chemical properties)
Flerovium (unknown chemical properties)
Ununpentium (unknown chemical properties)
Livermorium (unknown chemical properties)
Ununseptium (unknown chemical properties)
Ununoctium (unknown chemical properties)


Fluorine in the periodic table
gas: very pale yellow
liquid: bright yellow
solid: transparent (beta), opaque (alpha)
Small sample of pale yellow liquid fluorine condensed in liquid nitrogen
Liquid fluorine at cryogenic temperatures
General properties
Name, symbol, number fluorine, F, 9
Pronunciation /ˈflʊərn/ FLUU-reen, /ˈflʊərɪn/, /ˈflɔərn/
Element category diatomic nonmetal
Group, period, block 17 (halogens), 2, p
Standard atomic weight 18.998403163(6)
Electron configuration [He] 2s2 2p5[1]
2, 7
Physical properties
Phase gas
Density (0 °C, 101.325 kPa)
1.696[2] g/L
Liquid density at b.p. 1.505[3] g·cm−3
Melting point 53.48 K, −219.67 °C, −363.41[4] °F
Boiling point 85.03 K, −188.11 °C, −306.60[4] °F
Triple point 53.48 K, 90[4] kPa
Critical point 144.41 K, 5.1724[4] MPa
Heat of vaporization 6.51[2] kJ·mol−1
Molar heat capacity (Cp) (21.1 °C) 31[3] J·mol−1·K−1
(Cv) (21.1 °C) 23[3] 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[1] (Pauling scale)
Ionization energies
1st: 1681[5] kJ·mol−1
2nd: 3374[5] kJ·mol−1
3rd: 6147[5] kJ·mol−1
Covalent radius 64[6] pm
Van der Waals radius 135[7] pm
Crystal structure monoclinic
Fluorine has a monoclinic base-centered crystal structure

alpha (low-temperature)[8]
Magnetic ordering diamagnetic (−1.2×10−4)[9][10]
Thermal conductivity 0.02591[11] W·m−1·K−1
CAS registry number 7782-41-4[1]
Naming after the mineral fluorite, itself named after Latin fluo (to flow, in smelting)
Discovery André-Marie Ampère (1810)
First isolation Henri Moissan[1] (June 26, 1886)
Named by Humphry Davy
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
· references

Fluorine is a chemical element with symbol F and atomic number 9. The lightest halogen and most electronegative element–thus making fluorine extremely reactive–it exists as a pale yellow diatomic gas at standard conditions. Almost all other elements, including some noble gases, form compounds with fluorine. Being the most reactive of all halogens, elemental fluorine at high concentrations is extremely dangerous and poisonous for all living organisms.

Fluorite, the primary mineral source of fluorine, was first described in 1529; as it was added to metal ores to lower their melting points for smelting, the Latin verb fluo meaning "flow" became associated with it. Proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, with several early experimenters dying or sustaining injuries from their attempts. In 1886, French chemist Henri Moissan succeeded in isolating elemental fluorine using low-temperature electrolysis, a process still employed for modern production.

The element is 24th in universal abundance and 13th in terrestrial abundance. Due to the expense of refining pure fluorine, nearly all commercial applications handle it in compounds. About half of mined fluorite goes into steelmaking, the rest converted into corrosive hydrogen fluoride, a precursor to various organic fluorides and the critical aluminium refining flux cryolite. Organic fluorides have very high chemical and thermal stability, their major uses being refrigerants, electrical insulation and cookware, the last as PTFE. Pharmaceuticals such as atorvastatin and fluoxetine also contain fluorine, and the fluoride ion inhibits dental cavities, thus finding use in toothpaste and water fluoridation. Uranium enrichment, the largest application for free fluorine, began during the Manhattan Project in World War II. Global fluorochemical sales amount to over US$15 billion a year.

Fluorocarbon gases are generally greenhouse gases, with global-warming potentials 100 to 20,000 times that of carbon dioxide. Organofluorines persist in the environment due to the carbon–fluorine bond's strength, but the potential health impact of the most persistent such compounds is unclear. A few plants and bacteria synthesize organofluorine poisons for defense against herbivores, but fluorine has no known metabolic role in mammals.



Two concentric rings showing valence and non-valence electron shells
Simplified structure of the fluorine atom

Fluorine atoms have nine electrons, one fewer than neon, and electron configuration 1s22s22p5: two electrons in a filled inner shell and seven in an outer shell one short of completion. The outer electrons are ineffective at nuclear shielding, thereby experiencing a high effective nuclear charge of 9 − 2 = 7; this affects the atom's physical properties.[1]

Fluorine's first ionization energy is third-highest among all elements, behind helium and neon,[13] so removing electrons from neutral fluorine atoms is very difficult. Fluorine has a high electron affinity, second only to chlorine,[14] preferring to capture an electron and become isoelectronic with the noble gas neon;[1] it has the highest electronegativity of any element.[15] Fluorine atoms have a small covalent radius of around 60 picometers, similar to those of its period neighbors oxygen and neon.[16][17][note 1]


Fluorine has high reactivity because compared with Cl
and Br
, difluorine's bond energy is much lower, similar to a weak peroxide bond,[18][19] allowing elemental fluorine to dissociate easily. Bonds to other atoms are very strong because of its high electronegativity. Unreactive substances like powdered steel, glass fragments and asbestos fibers react quickly with cold fluorine gas; wood and water spontaneously combust under a fluorine jet.[2][20]

External video
Bright flames during fluorine reactions
Fluorine reacting with caesium

Reactions of elemental fluorine with metals require varying conditions: alkali metals cause explosions and alkaline earth metals display vigorous activity in bulk, but most other metals such as aluminium and iron must be powdered to prevent metal fluoride layers from passivating,[18] and noble metals require pure fluorine gas at 300–450 °C.[21] Metalloids and some solid nonmetals (sulfur, phosphorus, and selenium) burn with a flame in room temperature fluorine.[22][23] Hydrogen sulfide and sulfur dioxide combine readily with fluorine, the latter sometimes explosively, but sulfuric acid exhibits much less activity.[22]

Hydrogen, like the alkali metals, reacts explosively with fluorine.[24] Carbon, as lamp black, reacts at room temperature to yield fluoromethane. Graphite combines with fluorine above 400 °C, producing non-stoichiometric carbon monofluoride; higher temperatures generate gaseous fluorocarbons, sometimes with explosions.[25] Carbon dioxide and carbon monoxide react at room and slightly higher temperatures,[26] whereas paraffins and other organic chemicals beget strong reactions:[27] even fully substituted haloalkanes such as carbon tetrachloride, normally incombustible, may explode.[28] Although nitrogen trifluoride is stable, nitrogen requires electric discharge and elevated temperatures for reaction due to its very strong triple bonds;[29] ammonia's reaction is potentially explosive.[30][31] Oxygen does not combine under ambient conditions, but can be made to using electric discharge at low temperatures and pressures, the products tending to disintegrate into their constituent elements when heated.[32][33][34] Heavier halogens[35] and radon[36] react readily with fluorine; the lighter noble gases xenon and krypton require special conditions.[37]


4 diagonal placards with warnings, poison, corrosive, inhalant, oxidant
U.S. hazard signs for commercially transported fluorine.[38]

Fluorine is highly toxic, effects starting at lower concentrations than hydrogen cyanide's 50 ppm[39], and similar to chlorine[40]: significant irritation of the eyes and respiratory system as well as liver and kidney damage occur above 25 ppm. Human eyes and noses are seriously damaged at 100 ppm,[41] and inhalation of 1,000 ppm fluorine will cause death in minutes,[42] compared to 270 ppm for hydrogen cyanide.[43]


Main article: Phases of fluorine
Cube with spherical shapes on the corners and center and spinning molecules in planes in faces
Crystal structure of β-fluorine. Spheres indicate F
molecules that may assume any angle. Other molecules are constrained to planes.

At room temperature, fluorine is a gas of diatomic molecules,[2] pale yellow when pure (but sometimes described as yellow-green)[44] and possessing a characteristic pungent odor detectable at 20 ppb.[45] Fluorine condenses into a bright yellow liquid at −188 °C, a similar transition temperature to those of oxygen and nitrogen.[46]

Fluorine has two solid forms, α- and β-fluorine. The latter crystallizes at −220 °C and is transparent and soft, sharing the same disordered cubic structure of freshly crystallized solid oxygen[46][note 2] unlike the orthorhombic systems of other solid halogens.[8][50] Further cooling to −228 °C induces a phase transition into opaque, hard α-fluorine, which has a monoclinic structure featuring dense, angled layers of molecules. Transitioning from β- to α-fluorine is a more exothermic process than condensation of fluorine and can be violent.[8][50][note 3]


Main article: Isotopes of fluorine

Only one isotope of fluorine occurs naturally, the stable 19
containing ten neutrons.[51] This isotope has a high magnetogyric ratio[note 4] and therefore exceptional sensitivity to magnetic fields; because it is also the only stable isotope, it is used in magnetic resonance imaging.[53] Seventeen radioisotopes with mass numbers from 14 to 31 have been synthesized, of which 18
is most stable with a half-life of 109.77 minutes. Other radioisotopes have half-lives less than 70 seconds; most decay in less than half a second.[54] The isotopes 17
and 18
undergo β+ decay, lighter isotopes decay via electron capture, and those heavier than 19
undergo β decay or neutron emission.[54] One metastable isomer of fluorine is known, 18m
with a half-life of 234 nanoseconds.[55]



Solar System abundances[56]
Element Relative
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

Given that lighter elements show greater abundances, fluorine's value of 400 ppb – 24th among elements in the universe – is exceptional: all other elements from carbon to magnesium are at least twenty times more common.[57] This is because stellar nucleosynthesis processes bypass fluorine and any such atoms otherwise created have high nuclear cross sections, allowing further fusion with hydrogen or helium to generate oxygen or neon respectively.[57][58]

Beyond this transient existence, fluorine's presence is not well understood and three explanatory theories have been proposed:[57][59]


Fluorine is the thirteenth most common element in Earth's crust at 600–700 ppm by mass.[60] Free elemental fluorine in Earth's atmosphere would easily react with rocks, precluding its natural occurrence;[61][62] it is found only in minerals, of which fluorite, fluorapatite and cryolite are most industrially significant:[63][64]

  • Fluorite or fluorspar (CaF
    ), colorful and abundant worldwide, is fluorine's primary source with China and Mexico as its major suppliers. The U.S. led production in the early 20th century but ceased mining in 1995.[64][65][66][67][68]
  • Fluorapatite (Ca5(PO4)3F) and other apatites are much exploited for fertilizer phosphates. Although it contains most terrestrial fluorine, the low mass fraction of 3.5% leads to its disposal, except in the U.S. where it is used for water fluoridation.[64]
  • Cryolite (Na
    ), once used directly in aluminium production, is the rarest and most concentrated among these three minerals. The main commercial mine on Greenland's west coast closed in 1987.[64]
Major fluorine-containing minerals
pink globular mass with crystal facets Long prism-like crystal, without luster, at an angle coming out of aggregate-like rock A parallelogram-shaped outline with space-filling diatomic molecules (joined circles) arranged in two layers
Fluorite Fluorapatite Cryolite

Other minerals such as topaz also contain fluorine. Unlike heavier halides, fluorine precipitates out of water in alkaline earth fluorides rather than remaining in aqueous form.[64] Trace quantities of organofluorines with uncertain origin have been detected in volcanic eruptions and geothermal springs.[69] The existence of gaseous fluorine in crystals, suggested by crushed antozonite's smell, is contentious;[70][71] a 2012 study reported the presence of 0.04% F
by weight in antozonite, attributing these inclusions to radiation from its uranium atoms.[71]


Main article: Compounds of fluorine

Fluorine has a rich chemistry, encompassing both organic and inorganic domains, and combining with metals, nonmetals and even noble gases,[72][note 5] usually assuming an oxidation state of −1.[note 6] Its high electron affinity leads to its preference for ionic bonding, whereas its covalent bonds are polar and almost always single.[76][77][note 7]


Graph showing water and hydrogen fluoride breaking the trend of lower boiling points for lighter molecules
Boiling points of hydrogen halides and chalcogenides, showing the unusually high values for hydrogen fluoride and water.

Hydrogen and fluorine combine into hydrogen fluoride, where discrete molecules cluster through hydrogen bonds. The compound thus behaves more like water than hydrogen chloride,[78][79][80] boiling at much higher temperatures and being fully miscible with water unlike heavier hydrogen halides.[81] Hydrofluoric acid, aqueous hydrogen fluoride, is a weak acid unlike the other strong hydrohalic acids,[82][note 8] but corrosive enough to attack glass.[84]


Alkali metals form ionic and highly soluble monofluorides with the same cubic arrangement of sodium chloride, analogous to their corresponding chlorides.[85][86] Alkaline earth difluorides possess strong ionic bonds as well but are insoluble[87] save for beryllium difluoride, which also exhibits some covalent character and a quartz-like structure.[88] Rare earth elements and many other metals form mostly ionic trifluorides.[89][90][91]

Covalent bonding first comes to prominence in tetrafluorides: those of zirconium, hafnium[92][93] and several actinides[94] are ionic with high melting points,[95][note 9] but those of titanium,[98] vanadium[99] and niobium are polymeric,[100] melting or decomposing around or below 350 °C.[101] Pentafluorides continue this trend with their unbranched polymers and oligomeric complexes.[102][103][104] Thirteen metal hexafluorides are known,[note 10] all octahedral, and all volatile solids except for three: liquid MoF
and ReF
, and gaseous WF
.[105][106][107] Rhenium heptafluoride, the only characterized metal heptafluoride, is a low-melting molecular solid with pentagonal bipyramidal molecular geometry.[108] Metal fluorides with more fluorine atoms are particularly reactive.[109]

Structural progression of metal fluorides
Checkerboard-like lattice of small blue and large yellow balls, going in three dimensions so that each ball has 6 nearest neighbors of opposite type Straight chain of alternating balls, violet and yellow, with violet ones also linked to four more yellow perpendicularly to the chain and each other Ball and stick drawing showing central violet ball with a yellow one directly above and below and then an equatorial belt of 5 surrounding yellow balls
Sodium fluoride, ionic Bismuth pentafluoride, polymeric Rhenium heptafluoride, molecular

Other reactive nonmetals[edit]

Chlorine trifluoride, whose corrosive potential ignites asbestos, concrete, sand and other fire retardants.[110]

Binary fluorides of metalloids and p-block nonmetals are generally covalent and volatile, with varying reactivities. Period 3 and heavier nonmetals can form hypervalent fluorides.[111]

Boron trifluoride is planar and possesses an incomplete octet, thereby functioning as a Lewis acid and combining with Lewis bases like ammonia to form adducts.[112] Carbon tetrafluoride is tetrahedral and inert;[note 11] its group analogues, silicon and germanium tetrafluoride, are also tetrahedral[113] but behave as Lewis acids.[114][115] The pnictogens form trifluorides increasing in reactivity and basicity with higher molecular weight, although nitrogen trifluoride resists hydrolysis and is not basic.[116] Phosphorus, arsenic and antimony form pentafluorides more reactive than their respective trifluorides, with antimony pentafluoride the strongest neutral Lewis acid known.[102][117][118]

Chalcogens have diverse fluorides: unstable difluorides have been reported for oxygen (the only known compound with oxygen at oxidation state +2), sulfur and selenium, and tetrafluorides and hexafluorides exist for sulfur, selenium and tellurium. They are stabilized by more fluorine atoms and lighter central atoms, so sulfur hexafluoride is especially inert.[119][120] Chlorine, bromine and iodine can each form mono-, tri- and pentafluorides, but only iodine heptafluoride has been characterized among possible interhalogen heptafluorides.[121] Many of them are powerful sources of fluorine atoms, and chlorine trifluoride's industrial applications require similar precautions to those for fluorine gas.[122][123]

Noble gases[edit]

Main article: Noble gas compound
Black-and-white photo showing transparent crystals in a dish
Xenon tetrafluoride crystals photographed in 1962. Its synthesis, as with xenon hexafluoroplatinate, surprised many chemists.[124]

Noble gases, having complete electron shells, defied reaction with other elements until 1962 when Neil Bartlett reported synthesis of xenon hexafluoroplatinate;[125] xenon difluoride, tetrafluoride, hexafluoride and multiple oxyfluorides have been isolated since then.[126][127] Among other noble gases, krypton forms a difluoride,[128] and radon and fluorine generate a solid suspected to be radon difluoride.[129][130] Binary fluorides of lighter noble gases are exceptionally unstable: argon and hydrogen fluoride combine under extreme conditions to argon fluorohydride.[37] Helium and neon have no long-lived fluorides,[131] and no neon fluoride has ever been observed;[132] helium fluorohydride has been detected for milliseconds at high pressures and low temperatures.[131]

Organic compounds[edit]

Beaker with two layers of liquid, goldfish and crab in top, coin sunk in the bottom
Immiscible layers of colored water (top) and much denser perfluoroheptane (bottom) in a beaker; a goldfish and crab cannot penetrate the boundary. Quarters rest at the base.
Skeletal chemical formula
Chemical structure of Nafion, a fluoropolymer used in fuel cells and many other applications.[133]

The carbon–fluorine bond is organic chemistry's strongest,[134] giving stability to organofluorines.[135] It is almost non-existent in nature, manifesting instead in artificial compounds whose research is usually driven by commercial applications;[136] these compounds are diverse and reflect the complexity inherent in organic chemistry.[137]

Discrete molecules[edit]

Substitution of fluorine for more and more hydrogen atoms starting from an alkane gradually alters several properties: melting and boiling points are lowered, density increases, solubility in hydrocarbons decreases and overall stability increases. Perfluorocarbons,[note 12] where all hydrogen atoms are substituted, are therefore insoluble in most organic solvents, reacting at ambient conditions only with sodium in liquid ammonia.[138]

The term perfluorinated compound is used for what would otherwise be perfluorocarbons if not for a functional group,[139][note 13] usually that of carboxylic acids. These compounds share many properties with perfluorocarbons such as stability and hydrophobicity,[141] while the functional group augments their reactivity, enabling them to adhere to surfaces or act as a surfactant;[142] in particular, fluorosurfactants can lower water's surface tension more than hydrocarbon-based analogues. Fluorotelomers, which have some carbon atoms near the functional group unfluorinated, are also regarded as perfluorinated.[141]


Polymers exhibit the same stability increases afforded by fluorine substitution for hydrogen as discrete molecules, but their melting points are generally increased as well.[143] Polytetrafluoroethylene (PTFE), the simplest fluoropolymer and perfluoro analogue of polyethylene with structural unitCF
–, demonstrates this change as expected, but its very high melting point makes molding difficult.[144] Various PTFE derivatives are less temperature-tolerant but easier to mold: fluorinated ethylene propylene replaces some fluorine atoms with trifluoromethyl groups, perfluoroalkoxy alkanes do the same with trifluoromethoxy groups[144] and Nafion contains perfluoroether side chains capped with sulfonic acid groups.[145][146] Other fluoropolymers retain some hydrogen atoms; polyvinylidene fluoride has half the fluorine atoms of PTFE and polyvinyl fluoride has a quarter, but both behave much like perfluorinated polymers.[147]


Main article: History of fluorine

Early discoveries[edit]

Woodcut image showing man at open hearth with tongs and machine bellows to the side in background, man at water-operated hammer with quenching sluice nearby in foreground
Steelmaking illustration from De re metallica.

Georgius Agricola described fluorite in 1529 as a flux that lowered the melting point of metals during smelting,[148][149][note 14] creating the Latin fluorés from fluo (flow) for fluorite rocks as one of many Latin neologisms about 16th century industry. The mineral's name later evolved into fluorspar and then fluorite,[65][153][154] but its composition as calcium fluoride was only elucidated much later.[155]

Hydrofluoric acid was used in glass etching from 1720 onwards, perhaps as early as 1670;[note 15] Andreas Sigismund Marggraf reported its first characterization in 1764 when he heated fluorite with sulfuric acid, the resulting solution corroding its glass container.[157][158] Swedish chemist Carl Wilhelm Scheele repeated this in 1771, naming the acidic product fluss-spats-syran (fluorspar acid).[158][159] In 1810, French physicist André-Marie Ampère suggested that hydrogen and an element analogous to chlorine constituted hydrofluoric acid.[160] Sir Humphry Davy proposed naming this then-unknown substance fluorine from fluoric acid and the -ine suffix of other halogens, and the word for fluorine in most European languages is either of a similar etymology or derived from the Greek φθόριος (phthorios, destructive) following Ampère's suggestion.[161] The New Latin name fluorum gave the element its current symbol F, whereas Fl was used in early papers.[87][note 16]


Initial studies on fluorine were so dangerous that several 19th century experimenters were deemed "fluorine martyrs" on account of their misfortunes with hydrofluoric acid.[note 17] Isolation of elemental fluorine was hindered by the extreme corrosiveness of both it and hydrogen fluoride, as well as the lack of a simple and suitable electrolyte.[155][162] Edmond Frémy postulated that electrolysis of pure hydrofluoric acid to generate fluorine was feasible and devised a method to produce anhydrous samples from acidifying potassium bifluoride; instead, he discovered that the aqueous solution did not conduct.[155][162][163] Frémy's former student Henri Moissan persevered, and after much trial and error found that a mixture of potassium bifluoride and dry hydrogen fluoride was a conductor, enabling electrolysis. To prevent rapid corrosion of the platinum in his electrochemical cells, he cooled the reaction to extremely low temperatures in a special bath and forged cells from a more resistant mixture of platinum and iridium, using fluorite stoppers.[162][164] In 1886, after 74 years of effort, Moissan isolated elemental fluorine.[163][165]

In 1906, two months before his death, Moissan received the Nobel Prize in Chemistry[166] with citation:[162]

... in recognition of the great services rendered by him in his investigation and isolation of the element fluorine ... The whole world has admired the great experimental skill with which you have studied that savage beast among the elements.[note 18]

1887 drawing of Moissan's apparatus Nobel Prize photo of Moissan

Later uses[edit]

An ampoule of uranium hexafluoride or hex

The Frigidaire division of General Motors experimented with chlorofluorocarbon refrigerants in the late 1920s, and Kinetic Chemicals was formed as a joint venture between GM and DuPont in 1930 hoping to market CCl
as one such refrigerant. Replacing earlier and more toxic compounds, as well as fostering demand for kitchen refrigerators, it became profitable; by 1949 DuPont had bought out Kinetic and marketed several other Freon molecules.[137][158][167][168] Polytetrafluoroethylene was serendipitously discovered in 1938 by Roy J. Plunkett while working at Kinetic on tetrafluoroethylene as a refrigerant, and its superlative chemical and thermal resistance lent it to accelerated commercialization and mass production by 1941.[137][158][167]

Elemental fluorine's large-scale synthesis began during World War II, with Germany using high-temperature electrolysis for making tons of the planned incendiary chlorine trifluoride[169] and the Manhattan Project consuming greater quantities to produce uranium hexafluoride en route to enriching uranium. Since UF
is as corrosive as fluorine, gaseous diffusion plants required special materials: nickel for membranes, fluoropolymers for seals and liquid fluorocarbons as coolants and lubricants. This burgeoning nuclear industry later drove post-war fluorochemical development.[170]

Industry and applications[edit]

Fluorite mining, which supplies most global fluorine, peaked in 1989 when 5.6 million metric tons of ore were extracted. Chlorofluorocarbon restrictions lowered this value to 3.6 million tons in 1994, but production has been increasing since; around 4.5 million tons of ore and US$550 million revenue were generated in 2003, and later reports estimated 2011 global fluorochemical sales at $15 billion and predicted 2016–18 production figures of 3.5 to 5.9 million tons with at least a $20 billion revenue.[158][171][172][173][174] Froth flotation concentrates mined fluorite into two main grades of equal proportion: 60–85% pure metspar is almost all used in iron smelting whereas 97%+ pure acidspar is mainly converted to key industrial intermediate hydrogen fluoride.[66][158][175]

Fluorite Fluorapatite Hydrogen fluoride Metal smelting Glass production Fluorocarbons Sodium hexafluoroaluminate Pickling (metal) Fluorosilicic acid Alkane cracking Hydrofluorocarbon Hydrochlorofluorocarbons Chlorofluorocarbon Teflon Water fluoridation Uranium enrichment Sulfur hexafluoride Tungsten hexafluoride Phosphogypsum
Clickable diagram of the fluorochemical industry according to mass flows.

Inorganic fluorides[edit]

Aluminium extraction depends critically on cryolite.

As with other iron alloys, around 3 kg metspar is added to each metric ton of steel, the fluoride ions provided lowering its melting point and viscosity.[66][176] Alongside its role as an additive to materials like enamels and welding rod coats, most acidspar is reacted with sulfuric acid to form hydrofluoric acid, used straight in steel pickling, glass etching and alkane cracking.[66] One-third of HF goes into synthesizing cryolite and aluminium trifluoride, both fluxes in the Hall–Héroult process for aluminium extraction; although they are not consumed here, replenishment is necessitated by their occasional reactions with the smelting apparatus and each metric ton of aluminium requires about 23 kg of flux.[66][177] Fluorosilicates consume the second largest portion, with sodium fluorosilicate used in water fluoridation and laundry effluent treatment, as well as an intermediate en route to cryolite and silicon tetrafluoride.[178] Other important inorganic fluorides include those of cobalt, nickel and ammonium.[66][86][179]

Organic fluorides[edit]

Organofluorides consume over 20% of mined fluorite and over 40% of hydrofluoric acid, with refrigerant gases dominating and fluoropolymers increasing their market share.[66][180] Surfactants are a minor application but generate over US$1 billion in annual revenue.[181] Owing to how dangerous direct hydrocarbon–fluorine reactions are above −150 °C, industrial fluorocarbon production is indirect, mostly through halogen exchange reactions such as Swarts fluorination where chlorocarbon chlorines are substituted for fluorines by hydrogen fluoride under catalysts. Electrochemical fluorination subjects hydrocarbons to electrolysis in hydrogen fluoride, while the Fowler process treats them with solid fluorine carriers like cobalt trifluoride.[137][182]

Refrigerant gases[edit]

See also: Refrigerant

Halogenated refrigerants, termed Freons in informal contexts,[note 19] are identified by R-numbers which have an R prefix before stating the amount of fluorine, chlorine, carbon and hydrogen present.[66][183] Chlorofluorocarbons (CFCs) like R-11, R-12 and R-114 once dominated organofluorines, their production peaking in the 1980s and directed at air conditioning systems as well as propellant and solvent use; by the early 2000s, after their widespread international prohibition, production was below one-tenth of this peak.[66] Hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) are designed as replacements, their equally proportioned synthesis consuming more than 90% of fluorine in the organic industry. Important HCFCs include R-22, chlorodifluoromethane, and R-141b. The main HFC is R-134a[66] with HFO-1234yf coming to prominence owing to its global warming potential of less than 1% that of HFC-134a.[184]


SEM image with islands of white and interconnecting strings
Gore-Tex electron microscope image showing polymer islands connected by strands.[185]
Shiny spherical drop of water on blue cloth
Fluorosurfactant-treated fabrics are often hydrophobic.
Main article: Fluoropolymer

About 180,000 metric tons of fluoropolymers and over US$3.5 billion revenue per year were made in 2006–2007,[186] and the corresponding global market was estimated at just under $6 billion in 2011. It was predicted to grow 6.5% per year up to 2016.[187] Fluoropolymers can only be formed by polymerizing free radicals.[143]

Polytetrafluoroethylene (PTFE), sometimes called by its DuPont name Teflon,[188] represents 60–80% by mass of the world's fluoropolymer production.[186] Its dielectric properties enable its largest application of electrical insulation, whereas its inertness allows an additional use lining industrial pipes, tubing and gaskets requiring this property as well as coating fibreglass cloth for stadium roofs; non-stick cookware is a major consumer application.[188] Jerked PTFE film becomes expanded PTFE (ePTFE), a fine-pored membrane sometimes referred to by the brand name Gore-Tex and used for rainwear, protective apparel and filters; fibres may be fashioned into seals and dust filters.[188] Other fluoropolymers, including the second-most produced fluorinated ethylene propylene, mimic PTFE's properties and can substitute for it but are more moldable, albeit with higher prices and lower thermal stability. Unlike fluoropolymer film pairs replace glass in solar cells.[188][189]

The expensive yet resistant fluorinated ionomers are used as electrochemical cell membranes, of which the first and most prominent example is Nafion. Developed in the 1960s, it was first deployed in spacecraft and then replaced mercury-based chloralkali process cells, with use in proton exchange membrane fuel cells for vehicles following suit.[190][191][192] Fluoroelastomers such as Viton are crosslinked fluoropolymer mixtures mainly used in O-rings;[188] perfluorobutane (C4F10) is used as a fire-extinguishing agent.[193]


Fluorosurfactants are small organofluorine molecules used for repelling water and stains. Although expensive (US$200–2000 per kilogram) compared to pharmaceuticals, they produced over $1 billion in annual revenue by 2006; Scotchgard generated over $300 million alone in 2000.[181][194][195] They are a minority within the entire and much cheaper hydrocarbon-predominant market as possible applications like paints are burdened by compounding costs, this latter use valued at only $100 million in 2006.[181]

Elemental gas[edit]

Minaret-like electrical devices with wires running around them, thicker at at the bottom
transformers at a Russian railway.
See also: Industrial gas

At least 17,000 metric tons of fluorine are produced per year. It costs only US$5–8 per kilogram as uranium or sulfur hexafluoride, but handling challenges multiply its price as an element, most processes using the latter in large amounts employing in situ generation under vertical integration.[196]

The gas's largest application, consuming up to 7,000 metric tons annually, is in preparing UF
for the nuclear fuel cycle where it fluorinates uranium tetrafluoride formed from reacting uranium dioxide with hydrofluoric acid.[196] Its monoisotopic nature puts any mass differences between UF
molecules down to the presence of 235
or 238
, enabling uranium enrichment via diffusion or centrifuge.[2][66] Following that, about 6,000 metric tons per year go into producing the excellent yet inert dielectric SF
for high-voltage transformers and circuit breakers, eliminating hazardous polychlorinated biphenyls associated with oil-filled devices.[197] Electronics employs several compounds made from elemental fluorine: rhenium and tungsten hexafluoride in chemical vapor deposition, tetrafluoromethane in plasma etching[198][199][200] and nitrogen trifluoride in cleaning equipment.[66] Some organic fluorides' syntheses also use F
, but its reactivity often predicates conversion first to the gentler ClF
, BrF
and IF
which together allow calibrated fluorination, and fluorinated pharmaceuticals use sulfur tetrafluoride instead.[66]

Production of the free element[edit]


A machine room
Industrial fluorine cells at Preston.

Modern industrial F
production performs the same potassium fluoride/hydrogen fluoride electrolysis of Moissan with different apparatus: hydrogen and fluoride ions are reduced and oxidized at the steel container cathode and carbon block anode under 8–12 volts to hydrogen and fluorine gas respectively.[66][201] Temperatures are also elevated, KF•2HF melting at 70 °C (158 °F) and electrolyzed at 70–130 °C; because pure HF cannot be electrolyzed, KF is essential though catalytic.[158][202][203] F
may be stored in steel cylinders with passivated interiors below 200 °C, whereas nickel is required otherwise.[158][204] Regulator valves and piping are made of nickel, the latter possibly using Monel instead,[205] and frequent passivation along with strict exclusion of water and greases must be undertaken. In the laboratory, glassware may carry fluorine gas under low pressure and anhydrous conditions[205] but some sources recommend nickel-Monel-PTFE systems instead.[206]


While preparing for a 1986 conference celebrating the centennial of Moissan's achievement, Karl O. Christe reasoned that chemical fluorine generation should be feasible since some metal fluoride anions have no stable neutral counterparts, their acidification potentially triggering oxidation instead. He devised the following method which evolves fluorine at high yield and atmospheric pressure:[207]

2 KMnO4 + 2 KF + 10 HF + 3 H2O2 → 2 K2MnF6 + 8 H2O + 3 O2
2 K2MnF6 + 4 SbF5 → 4 KSbF6 + 2 MnF3 + F2

Christe later commented that the reactants "had been known for more than 100 years and even Moissan could have come up with this scheme."[208] As late as 2006, some references still asserted that fluorine was too reactive for any chemical isolation.[209]

Environmental concerns[edit]


Animation showing colored representation of ozone distribution by year above North America in 6 steps. It starts with a lot of ozone but by 2060 is all gone.
NASA projection of stratospheric ozone over North America without the Montreal Protocol.[210]

The Montreal Protocol set strict regulations on chlorofluorocarbons (CFCs) and bromofluorocarbons owing to their ozone damaging potential (ODP): their high stability which suited them to their original applications also meant that they decomposed at high altitudes where liberated chlorine and bromine atoms attacked ozone molecules.[211] Even with this and early indications of its efficiacy, predictions warned that several generations would pass before full recovery.[212][213] With one-tenth the ODP of CFCs, hydrochlorofluorocarbons (HCFCs) are current replacements,[214] themselves first scheduled for substitution by 2030–2040 with hydrofluorocarbons (HFCs) possessing no chlorine and zero ODP.[215] This date was brought forward to 2020 in 2007;[216] the Environmental Protection Agency had already prohibited one HCFC's production and capped those of two others in 2003.[215] Fluorocarbon gases are generally greenhouse gases with global-warming potentials (GWPs) of about 100 to 10,000, sulfur hexafluoride having a value around 20,000;[217] an outlier is HFO-1234yf which has attracted global demand due to its GWP of 4 compared with 1,430 for the current refrigerant standard HFC-134a.[184]


Perfluorooctanesulfonic acid, a key Scotchgard component until 2000.[218]

Organofluorines exhibit biopersistance due to the carbon–fluorine bond's strength and perfluoroalkyl acids (PFAAs), sparingly water-soluble owing to their acidic functional groups, are noted persistent organic pollutants[219] of which perfluorooctanesulfonic acid (PFOS) and perfluorooctanoic acid (PFOA) are most often researched.[220][221][222] PFAAs have been found in trace quantities worldwide from polar bears to humans, with PFOS and PFOA known to reside in breast milk and newborns' blood. A 2013 review showed slight correlation between groundwater and soil PFAA levels and human activity, but no chemical dominated and PFOS and PFOA were correlated themselves.[220][221][223] They bind in vivo to proteins such as serum albumin but are lipophobic unlike chlorocarbons and suspected to concentrate unmetabolized within humans in the liver and blood before excretion through the kidneys; their half-life depends much on species, from days in rodents to years in humans.[220][221][224] High doses of PFOS and PFOA cause cancer and death in newborn rodents but human studies have not established an effect at current exposure levels.[220][221][224] Other fluorochemicals also persist, like antibiotics and antidepressants in treated sewage[225] and agrichemicals in farmland runoff and nearby rivers,[226] since their biological metabolisms are hard.[225]

Biological role[edit]

Natural biochemistry[edit]

The gifblaar is one of the few organofluorine-synthesizing organisms.

Fluorine is not essential for mammals and fluorine deficiency pertains only to artificial diets as many environmental sources supply trace fluorine, though evidence suggests that trace amounts can strengthen bones.[227][228] Natural organofluorines have been found in microorganisms and plants[69] but not animals,[229] of which the most common is fluoroacetate, used to defend against herbivores by at least 40 plants in Africa, Australia and Brazil.[230] Other examples include ω-fluoro fatty acids, fluoroacetone and 2-fluorocitrate.[229] Adenosyl-fluoride synthase, which binds fluorine to carbon, was isolated from bacteria in 2002.[231]


Dental care[edit]

White man holding plastic tray with brown goop in it and sticking a small stick into a black boy's open mouth
Topical fluoride treatment in Panama

Population studies from the mid-20th century onwards show topical fluoride reducing dental caries. This was first explained as converting tooth enamel hydroxyapatite into less corrodable fluorapatite, but studies on pre-fluoridated teeth refuted that and current theories involve fluoride aiding enamel growth in small caries.[232] Following studies of children living where it was natural, controlled public water supply fluoridation against tooth decay[233] began in the 1940s and is now applied to water supplying 6% of the global population, including two-thirds of Americans on public water.[234][235] Despite best evidence indicating no side effects other than mostly benign dental fluorosis,[236] 2000 and 2007 meta-analyses demonstrating significant effects[237] and public endorsement, opposition still exists[238] on ethical and safety grounds;[235] the benefits have lessened, possibly due to other fluoride sources, but are still measurable in poorer people.[239] Sodium monofluorophosphate and sometimes sodium or tin(II) fluoride are often found in fluoride toothpastes, first introduced in the U.S. in 1955 and now ubiquitous in developed countries alongside fluoridated mouthwashes, gels, foams and varnishes.[239][240]


Capsules with "Prozac" and "DISTA" visible
Fluoxetine. A 2008 meta-analysis of this and three other antidepressants concluded that their efficacy, though statistically significant, was clinically insignificant "for any but the most severely depressed patients".[241]

20% of modern pharmaceuticals contain fluorine, which as a substituent can improve induced fit with target proteins or extend half-lives by protecting against metabolic enzymes.[242] Atorvastatin made more revenue than any other drug until it became generic in 2011[243] and the combination asthma prescription Seretide contains fluticasone among its two active ingredients.[244] One fluorine atom can alter pharmacokinetics, and many drugs are fluorinated to delay inactivation and lengthen dosage periods because the carbon–fluorine bond is very stable.[245] Fluorination also increases lipophilicity because the bond is more hydrophobic than the carbon–hydrogen bond, often helping in cell membrane penetration and hence bioavailability.[244]

Tricyclics and other pre-1980s antidepressants had several side effects due to their nonselective interference with neurotransmitters other than the serotonin target; the fluorinated fluoxetine was selective and one of the first avoiding this problem. Many current antidepressants receive this same treatment, including the selective serotonin reuptake inhibitors citalopram, its isomer escitalopram, fluvoxamine and paroxetine.[246][247] Quinolones are artificial broad-spectrum antibiotics that are often fluorinated to enhance their effects, prominent examples including ciprofloxacin and levofloxacin.[248][249][250][251] Fluorine also finds use in steroids:[252] fludrocortisone is a blood pressure-raising mineralocorticoid; triamcinolone and dexamethasone are strong glucocorticoids.[253] The majority of inhaled anesthetics are heavily fluorinated, with the prototype halothane proving much more inert and potent than its contemporaries. Later compounds such as the fluorinated ethers sevoflurane and desflurane improve on halothane and are almost insoluble in blood, allowing faster waking times.[254][255]

PET scanning[edit]

Rotating transparent image of a human figure with targeted organs highlighted
A full-body 18
PET scan.

Fluorine-18 is often found in radioactive tracers for positron emission tomography, as its half-life of almost two hours allows its transport from production facilities to imaging centers.[256] The most common tracer is fluorodeoxyglucose[256] which after intravenous injection is taken up by glucose-requiring tissues such as the brain and most malignant tumors,[257] whereby computer-assisted tomography can then be used for detailed imaging.[258]

Oxygen carriers[edit]

Liquid fluorocarbons can hold large volumes of oxygen or carbon dioxide, more so than blood, thereby attracting attention as respiratory media.[259] As blood transfusion demand exceeds supply, such less cumbersome substitutes have been studied, although they must often be emulsified for this use owing to immiscibility;[260][261] the substitute Oxycyte has passed initial clinical trials.[262][263] These substances can aid endurance athletes and are therefore banned from sports, with one cyclist's 1998 near death prompting investigation into their abuse.[264][265] Applications of pure perfluorocarbon liquid breathing include assisting burn victims and premature babies with deficient lungs. Complete and partial lung filling have been considered, though only animal tests and human trials of the former paradigm have been done thoroughly.[266] An Alliance Pharmaceuticals effort reached clinical trials but was abandoned due to inconclusive results.[267]


A 1080 warning in New Zealand.

About 30% of agrochemicals contain fluorine,[268] most of them herbicides and fungicides with a few crop regulators. Though often no more complex than trifluoromethylation, fluorination is a robust modification with effects analogous to fluorinated pharmaceuticals;[269] an example is trifluralin, with large-scale use against weeds in the U.S.,[269][270] but it is a suspected carcinogen and many European countries have banned it.[271] Sodium monofluoroacetate (1080) is a mammalian poison where two acetic acid hydrogens are replaced with fluorine and sodium, halting cell metabolism by replacing citric acid cycle acetate. First synthesized in the late 19th century, it was recognized as an insecticide in the early 20th and later deployed in its current use, with its largest consumer New Zealand using it to protect kiwis from the invasive Australian common brushtail possum.[272] Europe and the U.S. have banned 1080.[230][273][note 20]


Hydrofluoric acid[edit]

left and right hands, two views, burned index fingers
Hydrofluoric acid burns may not be evident for a day, after which calcium treatments are less effective.[274]
See also: Chemical burn

Hydrofluoric acid is a contact poison with greater hazards than many strong acids like sulfuric acid even though it is weak: it remains neutral in aqueous solution and thus penetrates tissue faster, whether through inhalation, ingestion or the skin, and at least nine U.S. workers died in such accidents from 1984 to 1994. It reacts with calcium and magnesium in the blood and imbalances them, leading to hypocalcemia and possible death through cardiac arrhythmia.[275] Insoluble calcium fluoride formation triggers strong pain[276] and burns larger than 160 cm2 can cause serious systemic toxicity.[277]

Exposure may not be evident for 8 hours for 50% HF, rising up to 24 hours for lower concentrations, and a burn may initially be painless as hydrogen fluoride affects nerve function. If it has been noticed, it should first be rinsed under a jet of water for 10–15 minutes to prevent further damage and any clothing worn should be removed.[278] Calcium gluconate is often applied next, providing calcium ions to bind with fluoride; skin burns can be treated with 2.5% calcium gluconate gel or special rinsing solutions.[279][280][281] Such HF absorption requires further medical treatment and calcium gluconate may also be injected, but using calcium chloride in its place is contraindicated and may cause severe complications. Excision or amputation of affected parts may be needed.[277][282]

Fluoride ion[edit]

Soluble fluorides are moderately toxic with 5–10 g sodium fluoride, or 32–64 mg fluoride ions per kilogram of body mass, a lethal dose for adults.[283] One-fifth of this can cause adverse health effects[284] and chronic excess consumption may lead to skeletal fluorosis which affects millions in Asia and Africa.[284][285] Ingested fluoride forms hydrofluoric acid in the stomach which is then easily absorbed by the intestines, where it crosses cell membranes, binds with calcium and interferes with various enzymes before its urinary excretion. Urine testing of the body's ability to clear fluoride ions sets exposure limits.[284][286]

A former major cause of fluoride poisoning was the accidental ingestion of insecticides with inorganic fluorides,[287] most current cases involving fluoride toothpaste swallowing instead.[284] Malfunctioning water fluoridation equipment is another cause, one incident in Alaska affecting almost 300 people and killing one.[288] Dangers from toothpaste are aggravated for small children, and the Centers for Disease Control and Prevention recommends supervising children below six brushing their teeth so they do not swallow toothpaste.[289] One regional study examined a year of pre-teen fluoride poisoning reports totalling 87 cases, including one death from ingesting insecticide. The rest were all from dental fluoride and most had no symptoms, but about 30% had stomach pains whose likeliness correlated with fluoride consumption.[287] A larger, similar study across the U.S. agreed; 80% of cases involved children under six and few were serious, though several hundred each year required special treatment.[290]

See also[edit]


  1. ^ Sources disagree on the radii of oxygen, fluorine and neon atoms. Precise comparison is thus impossible.
  2. ^ α-Fluorine has a regular pattern of molecules and is therefore a crystalline solid, but its molecules do not have a specific orientation. β-Fluorine's molecules have fixed location and minimal rotational uncertainty. For further detail on α-fluorine, see the 1970 structure by Pauling.[47] For further detail on the concept of disorder in crystals, see the referenced general reviews.[48][49]
  3. ^ A loud click is heard. Samples may shatter and sample windows blow out.
  4. ^ The ratio of the angular momentum to magnetic moment is called the gyromagnetic ratio. "Certain nuclei can for many purposes be thought of as spinning round an axis like the Earth or like a top. In general the spin endows them with angular momentum and with a magnetic moment; the first because of their mass, the second because all or part of their electric charge may be rotating with the mass."[52]
  5. ^
    118 squares with numbers and letters in them, mostly colored gray and green, with a few numbers and words outside the boxes
    Note how elements are categorized in this article: for example, antimony is included among nonmetals (colored green) even though its chemical properties are closer to metals (colored dark gray). Separate sections are dedicated to the noble gases (light blue), hydrogen (purple) and carbon (yellow). Period 7 p-block elements, colored light gray, have not been studied and are thus excluded.
  6. ^ Fluorine in F
    is defined to have oxidation state 0. With respect to unstable species, F
    and F
    have intermediate oxidation states, decomposing at around 40 K;[73] F+
    and related species are predicted to be stable.[74] Experimental evidence for the fluoronium ion, where fluorine has oxidation state +1, was reported in 2013.[75]
  7. ^ The metastable boron and nitrogen monofluoride have higher-order fluorine bonds, and some metal complexes use it as a bridging ligand. Hydrogen bonding is another possibility.
  8. ^ See also the explanation by Clark.[83]
  9. ^ ZrF
    melts at 932 °C,[96] HfF
    sublimes at 968 °C[93] and UF
    melts at 1036 °C.[97]
  10. ^ These thirteen are those of molybdenum, technetium, ruthenium, rhodium, tungsten, rhenium, osmium, iridium, platinum, polonium, uranium, neptunium and plutonium.
  11. ^ Carbon tetrafluoride is formally organic, but is included here rather than in the organofluorine chemistry section – where more complex carbon-fluorine compounds are discussed – for comparison with SiF
    and GeF
  12. ^ Perfluorocarbon and fluorocarbon are IUPAC synonyms for molecules of carbon and fluorine only, but in colloquial and commercial contexts the latter term may refer to any carbon- and fluorine-containing molecule, possibly with other elements.
  13. ^ This terminology is imprecise, and perfluorinated substance is also used.[140]
  14. ^ Basilius Valentinus supposedly described fluorite in the late 1400s, but because his writings were uncovered 200 years later, this work's veracity is doubtful.[150][151][152]
  15. ^ Partington[156] and Weeks[155] give differing accounts.
  16. ^ Fl, since 2012, is now used for flerovium.
  17. ^ Davy, Gay-Lussac, Thénard and the Irish chemists Thomas and George Knox were injured. Belgian chemist Paulin Louyet and French chemist Jerome Nickles died. Moissan also experienced serious hydrogen fluoride poisoning.[155][162]
  18. ^ Also honored was his invention of the electric arc furnace.
  19. ^ This DuPont trademark is sometimes further misused for CFCs, HFCs or HCFCs.
  20. ^ American sheep and cattle collars may use 1080 against predators like coyotes.


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