Fluorine

Fluorine
9F
- ↑ F ↓ Cl oxygen ← fluorine → neon
Fluorine in the periodic table
Appearance
gas: very pale yellow liquid: bright yellow solid: transparent (beta), opaque (alpha) Liquid fluorine at cryogenic temperatures
General properties
Name, symbol, number	fluorine, F, 9
Pronunciation	/ˈflʊəriːn/ FLUU-reen, /ˈflʊərɪn/, /ˈflɔəriːn/
Element category	diatomic nonmetal
Group, period, block	17 (halogens), 2, p
Standard atomic weight	18.998403163(6)
Electron configuration	[He] 2s2 2p5[23] 2, 7
History
Naming	after the mineral fluorite, itself named after Latin fluo (to flow, in smelting)
Discovery	André-Marie Ampère (1810)
First isolation	Henri Moissan[23] (June 26, 1886)
Named by	Humphry Davy
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa) 1.696[24] g/L
Liquid density at b.p.	1.505[25] g·cm−3
Melting point	53.48 K, −219.67 °C, −363.41[26] °F
Boiling point	85.03 K, −188.11 °C, −306.60[26] °F
Triple point	53.48 K, 90[26] kPa
Critical point	144.41 K, 5.1724[26] MPa
Heat of vaporization	6.51[24] kJ·mol−1
Molar heat capacity	(Cp) (21.1 °C) 31[25] J·mol−1·K−1 (Cv) (21.1 °C) 23[25] 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[23] (Pauling scale)
Ionization energies (more)	1st: 1681[27] kJ·mol−1
	2nd: 3374[27] kJ·mol−1
	3rd: 6147[27] kJ·mol−1
Covalent radius	64[28] pm
Van der Waals radius	135[29] pm
Miscellanea
Crystal structure	monoclinic alpha state (low-temperature)[30]
Magnetic ordering	diamagnetic, −1.2×10−4 (SI)[31][32]
Thermal conductivity	0.02591[33] W·m−1·K−1
CAS registry number	7782-41-4[23]
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[34]
v t e

Fluorine is an extremely reactive and poisonous chemical element. At room temperature, it is a pale yellow gas composed of diatomic molecules. It has an atomic number of 9 and is the lightest of the halogens and the most electronegative of the elements. Fluorine is the 24th most abundant element in the known universe and the 13th most abundant within the Earth's crust. It has a rich chemistry, forming compounds with nearly all other elements, including some of the noble gases.

The primary mineral source of fluorine, fluorite (calcium difluoride, CaF
2), was first described in 1529. At that time the Latin verb fluo, meaning "flow", became associated with fluorite rocks because they were added to metal ores to lower their melting points during smelting. First suggested as a chemical element in 1811, fluorine proved to be difficult and dangerous to separate from its compounds; several early experimenters were killed or badly hurt in their attempts at doing so. In 1886, French chemist Henri Moissan succeeded in isolating elemental fluorine using low temperature electrolysis, a process still used for the modern industrial production of fluorine.

Because of the expense of refining the pure element, nearly all commercially used fluorine remains in compound form throughout its processing. About half of mined fluorite is used directly in steel-making. The other half is converted to hydrogen fluoride, a dangerous acid that is the precursor to many fluorochemicals. The main use of hydrogen fluoride is in the synthesis of various organic fluorides and in cryolite, an inorganic material critical to aluminium refining. Organic fluorides have very high chemical and thermal stability; their largest market segments are in refrigerant gases and—in the form of polytetrafluoroethylene (Teflon)—electrical insulation and cookware. Modern pharmaceuticals such as atorvastatin (Lipitor) and fluoxetine (Prozac) contain fluorine. The fluoride ion, when directly applied to teeth, reduces decay; for this reason it is used in toothpaste and water fluoridation. The largest current end use of free fluorine, uranium enrichment, began in World II during the Manhattan Project. Global fluorochemical sales amount to over US$15 billion a year.

Fluorocarbon gases are greenhouse gases with warming potentials 4,000 to 10,000 times that of carbon dioxide. Sulfur hexafluoride–an insulating gas used in electricity plant–exhibits an even stronger effect, at about 20,000 times the global warming potential of carbon dioxide. Organofluorines endure in the environment due to the strength of the carbon–fluorine bond; the potential health impact of the most persistent of these compounds is unclear. While a few plants and bacteria synthesize organofluorine poisons for defense against herbivores, fluorine has no metabolic role in mammals.

Characteristics

Chemical reactivity

The difluorine bond is relatively weak, with a bond energy much less than that found in Cl
2 or Br
2 and similar to the easily cleaved oxygen–oxygen bonds of peroxides.^[35][36] For this reason, elemental fluorine easily dissociates to react with other atoms. On the other hand, bonds to non-fluorine atoms are very strong because of fluorine's high electronegativity. Both the easy dissolution of difluorine and the strong bonding to other atoms make fluorine extremely reactive.^[36] Many substances that are generally regarded as unreactive—such as powdered steel, glass fragments, and asbestos fibers—are readily consumed by cold fluorine gas. Wood and even water burn with flames when subjected to a jet of fluorine, without the need for a spark.^[37][38]

External video

Bright flames during fluorine reactions

Fluorine reacting with caesium

Reactions of elemental fluorine with metals require different conditions that depend on the metal. Many—such as aluminium, iron, and copper—must be powdered to overcome passivation (protective metal fluoride layers formed in the initial exposure).^[36] The alkali metals, such as sodium, react explosively. The alkaline earth metals, such as calcium, react somewhat less dramatically. The noble metals (gold, platinum, etc.) react least readily, requiring pure fluorine gas at 300–450 °C.^[39]

Fluorine reacts explosively with hydrogen in a manner similar to that of the alkali metals.^[40] The other halogens react readily with fluorine gas,^[41] as does the heavy noble gas radon.^[42] The lighter noble gases xenon and krypton react directly with fluorine under special conditions.^[43] Oxygen and fluorine gases do not normally react, but they can be combined under electric discharge at low pressures and temperatures. The products can be unstable and tend to separate back into fluorine and oxygen when heated.^[44][45][46] Nitrogen, with its very strong triple bonds, requires an electric discharge and very high temperatures to react with fluorine even though the end product, NF₃, is quite stable.^[47] Ammonia (NH₃) can react explosively with fluorine.^[48][49]

Carbon, in graphite or diamond form, does not react with fluorine gas at room temperature. Above 400°C, graphite reacts with fluorine to make a varying-composition solid called "carbon monofluoride". At higher temperatures, gaseous fluorocarbons start to be produced and the reaction can become explosive.^[50] Carbon dioxide and carbon monoxide react with fluorine at room temperature or just above.^[51] Organic chemicals, such as paraffins, react strongly when exposed to fluorine.^[52] Even fully halogenated organic molecules, such as the normally incombustible carbon tetrachloride, can explode.^[53]

The other solid nonmetals or metalloids (boron, silicon, arsenic, sulfur, germanium, phosphorus, selenium, tellurium) burn with a flame in room temperature fluorine.^[54] Hydrogen sulfide and sulfur dioxide combine readily with fluorine; the latter reaction can be explosive. Sulfuric acid reacts much more sluggishly.^[54]

Toxicity

Fluorine is highly toxic. The immediately dangerous to life or health concentration of fluorine is 25 ppm (hydrogen cyanide, in contrast, is 50 ppm).^[55] Above a concentration of 25 ppm, it causes significant irritation while attacking the eyes, airways, and lungs and affecting the liver and kidneys. At a concentration of 100 ppm, human eyes and noses are seriously damaged.^[56] Inhalation of 1,000 ppm fluorine will cause death in minutes^[57] (whereas for hydrogen cyanide a concentration of only 270 ppm is required).^[58]

Phases

At room temperature, fluorine is a gas of diatomic molecules.^[37] Though sometimes described as yellow-green, pure fluorine is actually a very pale yellow.^[59] The gas has a characteristic pungent odor that is noticeable in concentrations as low as 20 ppb.^[60] Fluorine condenses to a bright yellow liquid at −188 °C, which is near the condensation temperatures of oxygen and nitrogen.^[61]

Fluorine solidifies at −220 °C into a cubic structure called beta-fluorine.^[61] This phase is transparent and soft with a significant disordering of its molecular structure.^{[note 1]} With further cooling to −228 °C, fluorine undergoes a solid–solid phase transition into a monoclinic structure called alpha-fluorine. This phase is opaque and hard, with close-packed, shingled layers of molecules. The solid state phase change releases more energy than the melting point transition, and can be violent.^{[note 2][62][63]} In general, the solid state form of fluorine is structurally closer to that of oxygen than to those of the other halogens.^[62][63]

Electron arrangement

A neutral fluorine atom has nine electrons, one fewer than neon. The electronic configuration is 1s²2s²2p⁵: a filled inner shell of two electrons and an unfilled outer shell containing seven (one short of being filled). The outer electrons do not offer much shielding from the nucleus. They therefore experience a high effective nuclear charge of seven (nine minus two), which affects the physical properties of the atom.^[64]

Removal of electrons from neutral atoms is very difficult; fluorine's ionization energy (the energy required to remove an electron) is higher than that of any other element except neon and helium.^[65] Instead, fluorine exhibits a very strong preference for capturing one more electron to achieve the filled-shell electron configuration of the noble gas neon: 1s²2s²2p⁶.^[64] Fluorine is the element with the highest electronegativity (a relative measure of electron attraction by atoms).^[66] Its electron affinity (the energy released by adding an electron) is higher than that of any element except chlorine.^[67] Fluorine atoms have a small covalent radius of around 60 picometers; this is similar to the radii of oxygen and neon, its left and right hand periodic table neighbors.^{[note 3][68][69]}

Isotopes

One stable isotope of fluorine occurs naturally: fluorine-19, which contains ten neutrons.^[70] Fluorine is thus monoisotopic (having a single stable isotope) and mononuclidic (being found on Earth in only one isotope). Its monoisotopic occurrence makes it useful in uranium enrichment because UF₆ molecules differ only in mass due to mass differences between U-235 and U-238. These mass differences are used to separate U-235 and U-238 via diffusion and gas centrifugation.^[71] Fluorine's mononuclidic (100%) abundance make it well suited to magnetic resonance imaging, since it also has a high nuclear magnetogyric ratio^{[note 4]} (which translates to exceptional magnetic field sensitivity).^[72]

Seventeen radioisotopes have been synthesized, having mass numbers 14–18 and 20–31.^[73] The lightest fluorine isotopes, those with mass numbers of 14–16, decay via electron capture. Fluorine-17 and -18 undergo beta plus decay (emission of a positron). All isotopes heavier than fluorine-19 decay by beta minus mode (emission of an electron); some of these also decay by neutron emission. Fluorine-18 is the most stable radioisotope, with a half-life of 109.77 minutes before decaying to oxygen-18.^[73]

Occurrence

Universe

Abundance in the Solar System^[74]

Atomic number	Element	Relative amount
6	Carbon	4,800
7	Nitrogen	1,500
8	Oxygen	8,800
9	Fluorine	1
10	Neon	1,400
11	Sodium	24
12	Magnesium	430

For the ninth lightest element, fluorine is unusually rare given the lighter elements tend to be the more common ones. All of the elements from atomic number 6 (carbon) to atomic number 12 (magnesium) are hundreds to thousands of times more common, with the exception of sodium (which is tens of times more common).^[75] At 400 ppb, fluorine is estimated to be the 24th most common element in the universe.

Fluorine's rarity results from both a low stellar birth rate and rapid destruction. The main fusion reaction sequences of stars (stellar nucleosynthesis) which produce oxygen, carbon, and neon, bypass fluorine. Any fluorine which is nonetheless created is a large target (has a high nuclear cross section) for further fusion—either with hydrogen to form oxygen and helium, or with helium to form neon and hydrogen.^[75][76]

The presence of fluorine at all—outside its fleeting existence in stars—is somewhat of a mystery given these fluorine-eliminating reactions.^[75][77] Three theoretical solutions exist. In type II supernovae, atoms of neon could be 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 could blow the fluorine out of the star before hydrogen or helium can destroy it. In asymptotic giant branch stars (a type of red giant), pulses of fusion reactions could result in convection currents lifting fluorine out of the inner star.^[75][77]

Earth

Fluorine is the thirteenth most common element in the Earth's crust, comprising between 600 and 700 ppm by mass.^[78] Being so reactive, any free fluorine in the atmosphere of early Earth is likely to have been removed by being bound to surface rocks.^[79][80] It is essentially therefore found only in mineral compounds. The three most industrially significant are fluorite, fluorapatite, and cryolite.^[81][82]

Fluorite (CaF
2), also called fluorspar, is the main source of commercial fluorine. It is colorful, common, and can be found worldwide. China supplies over half the world's demand; Mexico is second. The U.S. produced most global fluorite in the early 20th century but its last mine, in Illinois, closed in 1995.^{[82][83][84][85][86]}

Fluorapatite (Ca₅(PO₄)₃F) and other apatites are mined in high volumes to produce phosphates for fertilizers. Most of the Earth's fluorine is bound in fluorapatite, but because the fluorine fraction is low (3.5%), it is discarded as waste. Only in the U.S. is there significant recovery: byproducts are used to supply water fluoridation.^[82]

Cryolite (Na
3AlF
6), the least abundant of the three major fluorine-containing minerals, is a concentrated source of fluorine. It was formerly used directly in aluminium production. The main commercial mine, on the west coast of Greenland, closed in 1987.^[82]

Major fluorine-containing minerals
Fluorite	Fluorapatite	Cryolite

Several other minerals, such as the gemstone topaz, include fluorine among their constituents. Fluorine is not significant in seawater or brines, unlike the other halides, because the alkaline earth fluorides (for example, CaF₂, MgF₂) precipitate out of water.^[82] Commercially insignificant quantities of organofluorines have been observed in volcanic eruptions and in geothermal springs. Their ultimate origin, whether from biological sources or geological formation, is unclear.^[87]

The possibility of small amounts of gaseous fluorine within crystals has been debated for many years. When crushed, one form of fluorite (antozonite) has a smell suggestive of fluorine.^[88][89] In 2012, a study reported detection of trace quantities (0.04% by weight) of diatomic fluorine in antozonite. It was suggested that radiation from small amounts of uranium within the crystals had caused the free fluorine defects.^[89]

Compounds

Fluorine has a rich chemistry including compounds formed with hydrogen, metals, main group nonmetals, and even noble gases, as well as a diverse set of organic compounds; ^{[note 5][90]} Its common oxidation state is −1.^{[note 6]} Because of its attraction for electrons, fluorine forms many ionic compounds. Covalent bonds involving fluorine are polar and are almost always, within molecules, single bonds.^{[note 7][91][92]}

Hydrogen fluoride

Fluorine combines with hydrogen to make a compound called hydrogen fluoride (HF). HF molecules cluster together weakly via hydrogen bonds. Because of this, hydrogen fluoride behaves more like water than HCl (hydrochloric acid).^[93][94][95] Hydrogen fluoride boils at a much higher temperature than the heavier hydrogen halides. HF is also fully miscible with water (that is, it dissolves in any proportion), unlike HCl, HBr, or HI.^[96]

Water solutions of hydrogen fluoride are called hydrofluoric acid. This is a chemically weak acid, unlike the other hydrohalic acids (such as hydrochloric) which are all strong.^{[97][note 8]} Although hydrofluoric acid is weak, it is very corrosive, even attacking glass.^[98]

Metal fluorides

The alkali metals form monofluorides that, like the alkali metal chlorides, are very ionic and soluble. They have the same atomic arrangement—the rock salt crystal structure—as sodium chloride.^[99][100] The difluorides of the alkaline earths are also very ionic but are generally very insoluble.^[101] Beryllium difluoride is an exception: it exhibits some covalent character, is water soluble, and has a structure similar to SiO₂ (quartz).^[102] Trifluorides are formed by many metals, particularly the rare earths, and are generally ionic.^{[103][104][105]}

The tetrafluorides represent a transition from ionic to covalent bonding. Zirconium and hafnium,^[106][7] along with several actinides,^[107] form high melting, ionic tetrafluorides.^{[108][note 9]} On the other hand, the tetrafluorides of titanium,^[109] vanadium,^[110] and niobium are polymeric.^[111] They melt or decompose below about 350 °C.^[112] The pentafluorides are even more covalently bonded, forming low dimensionality polymers or oligomeric molecules (clusters).^{[113][114][115]}

A total of thirteen metal hexafluorides have been characterized, all forming octahedredral molecules.^{[note 10]} All are volatile solids except molybdenum hexafluoride and rhenium hexafluoride (which are liquids) and tungsten hexafluoride (a gas).^{[116][117][118]} The only definite metal heptafluoride, that of rhenium, is a low-melting molecular solid. Its structure is a distorted pentagonal bipyramid.^[119] The higher metal fluorides are very reactive.^[120]

Progression of structure type with metal charge in the metal fluorides
Sodium fluoride, ionic	Bismuth pentafluoride, polymeric	Rhenium heptafluoride, molecular

Nonmetal fluorides

The binary fluorides of the main group metalloids and nonmetals are generally volatile, covalently bonded molecules that vary greatly in their reactivities. Nonmetals from the third row of the periodic table and below can form fluorides which are hypervalent (that is, they have more bonds than normal).^[121]

Boron trifluoride is a planar molecule in which the boron atom has an incomplete octet (with fewer bonds than normal). It is a weak Lewis acid and readily accepts a Lewis base, such as ammonia, thereby forming adducts (combinations).

The simplest binary compound with carbon is carbon tetrafluoride, an inert tetrahedral molecule.^{[note 11]} The atoms below carbon, silicon and germanium, also form tetrahedral tetrafluorides;^[122] these are Lewis acids.^[123][124]

Nitrogen trifluoride is stable against hydrolysis and is not a Lewis base.^[125] The atoms below nitrogen form trifluorides that are weak Lewis bases and are more reactive as the atom becomes heavier.^[125] Pentafluorides are formed by phosphorus, arsenic, and antimony. They are even more reactive than the respective trifluorides; antimony pentafluoride is the strongest Lewis acid of all charge-neutral compounds.^{[126][113][127]}

The chalcogens (elements making up the oxygen group) form a variety of fluorides. Unstable difluorides are known for oxygen (the only compound where oxygen is at a formal oxidation state +2) as well as for sulfur and selenium. Tetrafluorides and hexafluorides are known for sulfur, selenium, and tellurium. They tend to be more stable with more fluorine atoms and a lighter central atom: SF₆ is extremely inert.^[128][129]

The well-characterized heavier halogens (chlorine, bromine, and iodine) all form mono-, tri-, and pentafluorides: XF, XF₃, and XF₅. For XF₇, only iodine heptafluoride is known.^[130] Many of the halogen fluorides are powerful fluorinators (sources of fluorine atoms). Chlorine trifluoride readily fluorinates asbestos and refractory oxides; its use in industry requires precautions similar to those for fluorine gas.^[131][132]

Noble gas compounds

The noble gases are generally non-reactive because they have complete electron shells. Until the 1960s, no chemical bond with a noble gas was known. In 1962, Neil Bartlett reported the first chemical compound of xenon, xenon hexafluoroplatinate.^[133] Since then, xenon difluoride, xenon tetrafluoride and xenon hexafluoride have been isolated, as well as various oxyfluorides.^[134][135] Krypton, the lighter homolog of xenon, forms a difluoride and a few more complicated fluorine containing compounds.^[136] Radon, the heavier homolog of xenon, reacts readily with fluorine to form a solid, which is thought to be radon difluoride.^[137][138]

Some of the lightest noble gases form binary fluorides of only exceptionally limited stability. Argon reacts in extreme conditions with hydrogen fluoride to form argon fluorohydride.^[43] Helium and neon do not form any time-stable fluorides, but helium fluorohydride has been observed for milliseconds at extremely high pressure and low temperature.^[139] Neon is considered even less reactive than helium, and no fluorides have been even momentarily observed.^[140]

Organic compounds

The carbon–fluorine chemical bond is the strongest bond in organic chemistry;^[141] organofluorines are subsequently very stable.^[142] With a few exceptions, the C–F bond does not exist in nature, meaning the entire field is essentially "handmade",^[143] with research in particular areas tending to be driven by the commercial value of applications. The range of organofluorine compounds is diverse, reflecting the inherent complexity of organic chemistry.^[144]

Small molecules

Monofluoroalkanes (alkanes with one hydrogen replaced with fluorine) have properties similar to unfluorinated alkanes. They are soluble in many nonpolar solvents and have some chemical and thermal instability. As more fluorine atoms are substituted for hydrogen atoms, the properties change: solubility in hydrocarbons decreases and stability increases. Also, melting and boiling points decrease, while density goes up.^[145]

When all hydrogen atoms are replaced with fluorine atoms to make perfluorocarbons ("per" meaning maximum),^{[note 12]} a great difference is apparent: such compounds are extremely stable, and they are only subject to attack at standard conditions by sodium in liquid ammonia. They are also very insoluble, with few organic solvents capable of dissolving them.^[145]

Perfluorinated compound is a term for hydrocarbons that are fully fluorinated but which also have a functional group (a small non-hydrocarbon part of the molecule).^{[note 13][146]} Often this is a carboxylic acid (-CO₂H) group. Perfluorinated compounds exhibit many perfluorocarbon properties such as inertness, stability, non-wetting by water and oils, and slipperiness.^[147] The functional group is however available for reactions. It may also help the molecule to adhere to surfaces or behave as a surfactant (a soap-like mixture).^[148] Fluorosurfactants can lower the surface tension of water below that achievable with hydrocarbon-based surfactants. Industry practice is to also regard fluorotelomers‒perfluorinated compounds with fluorinated backbones, but also a few non-fluorinated carbons (typically two) near the functional group‒as perfluorinated.^[147]

Polymers

As with small molecules, replacing hydrogen with fluorine in a polymer increases chemical stability and reduces flammability. Melting points are typically much higher than in the corresponding hydrocarbon polymers.^[149]

The simplest fluoroplastic is PTFE (polytetrafluoroethylene, DuPont brand Teflon), which is a simple linear chain polymer with the repeating structural unit: –CF₂–. It has no hydrogen atoms and is the perfluoro analog of PE (polyethylene, structural unit: –CH₂–). PTFE, as expected for a perfluorocarbon, has a much higher chemical and thermal stability than that of polyethylene. However, its very high melting point makes it difficult to fashion into parts.^[150]

Various PTFE derivatives have lower maximum usage temperatures but have the benefit of being more melt-processable. FEP (fluorinated ethylene propylene) is structurally similar to PTFE but has some fluorine atoms replaced with –CF₃ groups. PFA (perfluoroalkoxy) has some fluorine atoms replaced with –OCF₃ groups.^[150] Nafion is a structurally complicated polymer. It has a PTFE-like backbone, but also contains side chains of perfluoro ether that end in sulfonic acid (–SO₂OH) groups.^[151][152]

There are other fluoroplastics that are not perfluorinated; these contain some C-H groups. PVDF (that is, polyvinylidene fluoride, with the structural unit  –CF₂CH₂–) has half the fluorine atoms of PTFE. PVF (polyvinyl fluoride, structural unit: –CH₂CHF–) has even less. Despite this, it still has many of the properties of fully fluorinated polymers.^[153]

History

Early discoveries

Fluorite, the main source mineral of fluorine, was described in 1529 by Georgius Agricola, who related its use as a flux—an additive that helps lower melting temperature during smelting.^{[154][155][note 14]} Agricola, the "father of mineralogy", invented several hundred new terms in his Latin works describing 16th-century industry. For fluorite rocks (schöne Flüsse in the German of the time), he created the Latin noun fluorés, from fluo (flow). The name for the mineral later evolved to fluorspar (still commonly used) and then to fluorite.^{[156][83][157]}

Hydrofluoric acid was used as a glass-etching agent from the 1720s, perhaps as early as 1670.^{[note 15]} Andreas Sigismund Marggraf made the first scientific report on its preparation in 1764 when he heated fluorite with sulfuric acid; the resulting solution corroded its glass container.^[158][159] Swedish chemist Carl Wilhelm Scheele repeated this reaction in 1771, recognizing the product as an acid, which he called "fluss-spats-syran" (fluor-spar-acid).^[159][160]

In 1810, French physicist André-Marie Ampère suggested that hydrofluoric acid was a compound of hydrogen with an unknown element, analogous to chlorine.^[161] Fluorite was then shown to be mostly composed of calcium fluoride.^[12] 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, although Greek, Russian, and some others (following Ampère's suggestion) use the name ftor or derivatives, from the Greek φθόριος (phthorios), meaning "destructive".^[162][163] The New Latin name (fluorum) gave the element its current symbol, F, although the symbol Fl was used in early papers.^{[101][note 16]}

Isolation

Progress in isolating the element was slowed by the exceptional dangers of generating fluorine; several 19th-century experimenters—the "fluorine martyrs"—were killed or badly hurt while working with hydrofluoric acid.^{[note 17]} Initial attempts to isolate the element were hindered by problems obtaining a suitable conducting liquid for electrolysis as well as by the extreme corrosiveness of hydrogen fluoride and of fluorine gas.^[12][13]

Edmond Frémy postulated that passing electric current through pure hydrofluoric acid (dry HF) might allow the element to be isolated. 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). He discovered, however, that pure hydrogen fluoride would not transmit an electric current.^{[12][13][164]}

French chemist Henri Moissan, formerly one of Frémy's students, continued the search. After trying many different approaches, he combined potassium difluoride and dry hydrogen fluoride. The mixture proved capable of conducting electricity, making electrolysis possible. However, rapid destruction of the platinum metal in his electrochemical cells stymied the quest. To continue, Moissan devised a strategy of cooling the reaction to extremely low temperatures, in a special bath, so as to slow the rate of corrosion. Moissan also made equipment that was more corrosion-resistant: containers crafted from a mixture of platinum and iridium (more chemically resistant than pure platinum) with fluorite stoppers.^[165][13] In 1886, Moissan crowned 74 years of effort by many chemists when he isolated elemental fluorine.^[164][166]

In 1906, two months before his death, Moissan received the Nobel Prize in chemistry.^{[note 18][167]} The award read:^[13]

...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.

Moissan's apparatus, 1887 publication	Henri Moissan, Nobel Prize photo

Application development

In the late 1920s, chlorofluorocarbon refrigerants were tested by researchers from the Frigidaire division of General Motors. In 1930, GM and DuPont formed a joint venture under the name Kinetic Chemicals, with a view to commercializing one such chlorofluorocarbon: Freon-12 (CCl
2F
2). It proved to be a marketplace success, rapidly replacing earlier more toxic refrigerants, and contributing to the growth of the overall market for kitchen refrigerators. By 1949, DuPont had bought out the joint venture and marketed several other Freon molecules.^{[144][159][168][169]}

In 1938, Teflon ((C₂F₄)_n) was accidentally discovered by a recently hired Kinetic chemist, Roy J. Plunkett. Undertaking research on the possible use of tetrafluoroethylene as a refrigerant, he encountered a mystery. Gas left in a cylinder overnight could not be released the next morning, but the weight of the container had not changed (indicating the gas had not leaked out). Cutting the cylinder open, he found white flakes of an unknown substance. Tests showed it to be polytetrafluoroethylene ('poly-' meaning 'many'). The new polymer was more resistant to corrosion and more stable at high temperatures than any other plastic. By 1941, it was being produced in significant quantities as a result of an accelerated commercialization program.^{[144][159][168]}

Large-scale production of elemental fluorine began during World War II. Germany used high-temperature electrolysis to produce tons of chlorine trifluoride, a compound planned for use as an incendiary.^[170] The Manhattan Project in the U.S. used even more elemental fluorine to make uranium hexafluoride for use in uranium enrichment plants. Because UF₆ is about as corrosive as fluorine itself, gaseous diffusion separation plants had to be built with special materials. Nickel was used for the membranes; fluoropolymers such as Teflon were used for seals; and liquid fluorocarbons were used as coolants and lubricants. After the war, the burgeoning nuclear weapons industry drove further development of fluorochemical compounds.^[171]

Industry and applications

Fluorite mining, the main source of fluorine, was a growing industry up to 1989 when it peaked at 5.6 million metric tons of ore extracted in that year. Environmental restrictions on the use of chlorofluorocarbons subsequently reduced production, down to 3.6 million tons in 1994. Production has steadily risen since that time. In 2003, it was estimated at 4.5 million tons with a corresponding revenue of $550 million. Subsequent market research reports estimated 2011 global fluorochemical sales at $15 billion and have predicted production figures over the period 2016‒18 in the range of 3.5 to about 5.9 million tons, with revenue of $20 billion, or more.^{[159][172][173][174][175]}

Mined fluorite is concentrated by flotation separation into two main grades, with about equal production of each. Metspar (60–85% purity) is used almost exclusively for iron smelting. Acidspar (97%+ purity) is mainly converted to hydrogen fluoride, the primary chemical intermediate for the fluorochemical industry.^{[84][159][176]}

The fluorine industry's supply chain, based on mass flows: click for links to related articles.

Inorganic fluorides

About 3 kg of metspar-grade fluorite are added to each metric ton of steel. The fluoride ions from CaF₂ lower the melt's temperature and viscosity, making it runnier. Metspar is similarly used to produce cast iron and other iron alloys.^[84][177]

Most acidspar is reacted with sulfuric acid to make hydrofluoric acid. Significant direct uses of HF include pickling (cleaning) steel, etching glass, and cracking alkanes in the petrochemical industry.^[84] Acidspar-grade fluorite is also added to ceramics, enamels, glass fibers, clouded glass, cement, and the outer coating of welding rods.^[84]

One third of HF (one sixth of mined fluorine) is used to make synthetic cryolite (sodium hexafluoroaluminate) and aluminium trifluoride. These compounds are used in the electrolysis of aluminium by the Hall–Héroult process. The fluorides are not reactants in the smelting process, but fluxes that lower the temperature of the melt. They are not consumed in the process and remain available to support smelting. However, over time, small amounts are lost through side reactions with the smelting apparatus and new fluorides must be added. About 23 kg are required for every metric ton of aluminium.^[84][178]

Fluorosilicates are the next most significant inorganic fluorides formed from HF. Sodium fluorosilicate is used for water fluoridation; as an intermediate for synthetic cryolite and silicon tetrafluoride; and for the treatment of effluents in laundries.^[179] 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); ammonium fluoride (various uses); and magnesium fluoride (antireflective optical coatings).^[84] Sodium and potassium difluorides are significant^why? to the chemical industry.^[100][180]

Organic fluorochemicals

Organofluoride production consumes over 20% of all mined fluorite and over 40% of hydrofluoric acid. Refrigerant gases are the dominant segment. Fluoropolymers represent less than one quarter the consumption of refrigerant gases in terms of fluorine usage but are growing faster.^[84][181] Fluorosurfactants are molecules used to make clothing and other items water-resistant. They make up a small market segment but generate over $1 billion in yearly revenue.^[182]

Industrially, production of fluorocarbons relies on indirect methods because the direct reaction of hydrocarbons with fluorine gas can be dangerous at temperatures above −150 °C. Many fluorochemicals are made by halogen exchange reactions: chlorinated hydrocarbons react with hydrogen fluoride to switch out chlorine for fluorine. The reactions are catalyzed, for example, by antimony halides in "Swarts fluorination". Another method is electrochemical fluorination, in which hydrocarbons are electrolyzed in hydrogen fluoride. In the Fowler process, hydrocarbons are reacted with solid carriers of fluorine, notably cobalt trifluoride.^{[183][144][184]}

Refrigerant gases

Halogenated molecules used in refrigeration are identified by the R-number system, which explains the amount of fluorine, chlorine, carbon, and hydrogen in each molecule.^[185] The DuPont brand Freon has been colloquially used for these compounds, but brand-neutral terminology uses "R" ("refrigerant") as the prefix.^{[note 19][84]}

Traditionally, chlorofluorocarbons (CFCs) were the predominant class of fluorinated organic chemical. Prominent CFCs included R-11 (trichlorofluoromethane), R-12 (dichlorodifluoromethane), and R-114 (1,2-dichlorotetrafluoroethane). Production of CFCs grew strongly through the 1980s, primarily for refrigeration and air conditioning but also for propellants and solvents. Production fell to less than 10% of the mid-1980s peak, by the early years of the 21st century, after most countries banned the end use of these materials.^[84]

Hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) serve as replacements for CFC refrigerants; few were commercially manufactured before 1990. More than 90% of fluorine used for organics goes into these two classes, in roughly 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).^[84]

Fluoropolymers

As of about 2006–2007, fluoropolymer volume was estimated at over 180,000 metric tons per year. The corresponding revenue estimate was over $3.5 billion.^[186] The 2011 global fluoropolymer market was estimated at slightly under $6 billion in revenue and predicted to grow 6.5% per year through 2016.^[187] Fluoropolymers are formed by polymerizing free radicals; other hydrocarbon polymerization techniques do not work.^[149]

Polytetrafluoroethylene (PTFE) is 60–80% of the world's fluoropolymer production on a weight basis.^[186] The DuPont brand Teflon is sometimes used generically for this substance.^[188] The largest application is in electrical insulation as PTFE is an excellent dielectric. It is also used in the chemical industry where corrosion resistance is needed: in coating pipes, tubing, and gaskets. Another major use is as architectural fabric in the form of PTFE-coated fiberglass cloth used for stadium roofs. The major consumer application is for non-stick cookware.^[188]

When stretched with a sudden jerk, PTFE film becomes a fine-pored membrane in the form of expanded PTFE (ePTFE). The term "Gore-Tex" is sometimes used generically although this is a specific brand name. ePTFE is used in rainwear, protective apparel, and liquids and gas filters. PTFE can also be formed into fibers which are used in pump packing seals and bag house filters.^[188]

Other fluoropolymers tend to have similar properties to PTFE, which leads to their use in electrical insulation and the chemical process industry. Unlike PTFE, these other fluoropolymers can be melt-processed. This makes them easier to work with as they can be formed into a complex shape, but they are also 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.^{[188][189][190]}

Fluorinated ionomers (polymers that include charged fragments) are expensive, chemically resistant materials used as membranes in electrochemical cells. Nafion, developed in the 1960s, was the first example and remains the most prominent material in the class. The initial application was as fuel cell material in spacecraft. Since then, this polymer has been transforming the 55 million tons per year chloralkali industry; it is replacing hazardous mercury-based cells with membrane cells. Recently, the fuel cell application has reemerged with efforts to instaill proton exchange membrane (PEM) fuel cells into automobiles.^{[191][192][193]}

Fluoroelastomers are rubber-like substances that are composed of crosslinked mixtures of fluoropolymers. Chemical-resistant O-rings are the primary application; Viton is a prominent brand.^[188]

Surfactants

Fluorinated surfactants are small organofluorine molecules, principally used for water and stain resistance. As of 2006, yearly revenues for this segment were over $1 billion. Fluorosurfactants are expensive chemicals, comparable to pharmaceutical chemicals: $200–2000 per kilogram ($90–900 per pound). Scotchgard has been a prominent brand; revenues in 2000 were over $300 million.^{[182][194][195]}

Fluorosurfactants make up a very small part of the overall surfactant market, most of which is hydrocarbon-based and much cheaper. Some potential applications (for example, low cost paints) are unable to use fluorosurfactants because of the price impact of compounding in even small amounts of fluorosurfactant. Usage in paints was only about $100 million as of 2006.^[182]

Fluorine gas

For countries with available data, about 17,000 metric tons of fluorine are produced per year. Fluorine is relatively inexpensive, costing about $5–8 per kilogram when sold as uranium hexafluoride or sulfur hexafluoride. Because of difficulties in storage and handling, the price of fluorine gas is much higher. Processes demanding large amounts of fluorine gas are generally vertically integrated and entail on-site production.^[196]

The largest application for elemental fluorine (up to 7,000 metric tons per year) is in the preparation of uranium hexafluoride, which is used in the production of nuclear fuels. 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.^[196] Fluorine's monoisotopic natural occurrence makes it useful in uranium enrichment, in which uranium-235 is separated by diffusion or centrifugation from uranium-238. The difference in mass between uranium hexafluoride molecules arises entirely from the different masses of the two uranium isotopes.^[84][196]

The second largest application for fluorine gas (about 6,000 metric tons per year) is in the production of sulfur hexafluoride, which is used as a dielectric medium in high voltage transformers and circuit breakers. SF₆ gas has a much higher dielectric strength than air and is extremely chemically inert. Switchgear using SF₆ has no hazardous polychlorinated biphenyls (PCBs), in contrast to traditional oil-filled devices.^[197]

Several compounds made from elemental fluorine serve the electronics industry. Rhenium and tungsten hexafluorides are used for chemical vapor deposition of thin metal films. Tetrafluoromethane is used for plasma etching.^{[198][199][200]} Nitrogen trifluoride is used for cleaning equipment.^[84]

Some organic fluorides are prepared from elemental fluorine rather than from HF. However, because direct fluorination is usually too hard to control, intermediate strength fluorinators are made from fluorine gas. The halogen fluorides ClF
3, BrF
3, and IF
5 provide gentler fluorination, with a series of strengths, and are easier to handle. Sulfur tetrafluoride is used for making fluorinated pharmaceuticals.^[84]

Production of fluorine gas

Industrial

Modern industrial production of elemental fluorine uses Moissan's process of electrolyzing potassium fluoride/hydrogen fluoride mixtures, but with an apparatus made of different materials. A steel container acts as the negative electrode, attracting H⁺ ions and releasing hydrogen gas. A carbon block (similar to that used in aluminium production) acts as the positive electrode, attracting F⁻ ions and releasing fluorine gas. The voltage difference between the electrodes is 8–12 volts.^[84][201]

Commercial temperatures are now higher than what Moissan used. A mixture with the approximate composition KF•2HF melts at 70 °C (158 °F) and is electrolyzed at 70–130 °C. Because HF alone cannot be electrolyzed, the presence of some KF is critical even though it is not consumed in the cell.^{[159][202][203]}

Pure fluorine gas may be stored in steel cylinders where the inside surface is passivated, as long as the temperature is kept below 200 °C. Above that temperature, nickel is required.^[159][202] Regulator valves are made of nickel. Fluorine piping is generally made of nickel or Monel (a nickel-copper alloy).^[204] Care must be taken to passivate all surfaces frequently and to exclude any water or greases. In the laboratory, fluorine gas may be used in glass tubing provided the pressure is low and moisture is excluded.^[204] Some sources instead recommend systems made of nickel, Monel, and PTFE.^[205]

Chemical

In 1986, when he was preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl O. Christe realized that the chemical generation of fluorine ought to be feasible. The main idea was that some metal fluoride anions either did not have neutral counterparts (or had such counterparts that were very unstable) and that acidifying these anions could result in chemical oxidation, rather than formation of the expected molecules. The method involved, which resulted in a high yield of fluorine at atmospheric pressures, was:^[206]

2 KMnO₄ + 2 KF + 10 HF + 3 H₂O₂ → 2 K₂MnF₆ + 8 H₂O + 3 O₂↑

2 K₂MnF₆ + 4 SbF₅ → 4 KSbF₆ + 2 MnF₃ + F₂↑

Christe subsequently commented that the reactants involved 'had been known for more than 100 years and even Moissan could have come up with this scheme.'^[207] Up until at least 2006, references could still be found in the literature asserting that fluorine was too reactive to be able to be separated from its compounds by chemical means.^[208]

Environmental concerns

Atmosphere

Because they deplete the ozone layer, chlorofluorocarbons (CFCs) and bromofluorocarbons (BFCs) have been strictly regulated via a series of international agreements called the Montreal Protocol. It is the chlorine and bromine from these molecules that cause harm, not the fluorine. Because of the inherent stability of these fully halogenated molecules (which makes them so nonflammable and useful), they are able to attain the upper reaches of the atmosphere before decomposing. At high altitudes, they release chlorine and bromine atoms which attack ozone molecules.^[209] Predictions are that even after the CFC ban, several generations will be required for these molecules to leave the atmosphere and for the ozone layer to fully recover. Early indications are that the CFC ban is working: ozone depletion has stopped, and recovery has started.^[210][211]

Hydrochlorofluorocarbons (HCFCs) are current replacements for CFCs, with about one tenth the ozone damaging potential (ODP).^[212] They were themselves originally scheduled for elimination by 2030 in developed nations and 2040 in undeveloped nations, with replacement by hydrofluorocarbons (HFCs) which have no chlorine and thus zero ODP. In 2003, the U.S. Environmental Protection Agency prohibited production of one HCFC and capped the production of the two others.^[213] In 2007, a new treaty was signed by almost all nations to move the HCFC phaseout up to 2020.^[214]

Fluorocarbon gases of all sorts (CFCs, HFCs, and the like) are greenhouse gases with about 4,000 to 10,000 times the potency of carbon dioxide. Sulfur hexafluoride exhibits an even stronger effect, at about 20,000 times the global warming potential of carbon dioxide.^[215]

Biopersistance

Organofluorines endure in the environment due to the strength of the carbon–fluorine bond. Perfluoroalkyl acids (PFAAs) have attracted particular attention as persistent global contaminants. Because of their acid group, PFAAs are water soluble in low concentrations.^[216] While there are other PFAAs, the lion's share of environmental research has been done on the two most well-known: perfluorooctanesulfonic acid (PFOS) and perfluorooctanoic acid (PFOA).^{[217][218][219]}

Trace quantities of PFAAs have been detected worldwide, from polar bears in the Arctic to the global human population. Both PFOS and PFOA have been detected in breast milk and the blood of newborns. A 2013 review showed widely varying amounts of PFAA in different soils and groundwater, with generally higher amounts in areas of more human activity. There was no clear pattern of one chemical dominating, and higher amounts of PFOS were correlated to higher amounts of PFOA.^{[217][218][220]}

In the body, PFAAs bind to proteins such as serum albumin. Unlike chlorinated hydrocarbons, PFAAs are not lipophilic (stored in fat). Their tissue distribution in humans is unknown, but studies in rats suggest they are present mostly in the liver, kidneys, and blood. They are not metabolized by the body but are excreted by the kidneys. Dwell time in the body varies greatly by species. Rodents have half-lives of days, while in humans they remain for years.^{[217][218][221]}

The potential health impact of PFAAs is unclear. Both PFOA and PFOS in high doses cause cancer and the death of newborns in rodents. However, studies on humans have not been able to establish an impact at current exposure levels.^{[217][218][221]}

Less fluorinated chemicals (not perfluorinated compounds) are also detectable in the environment. 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.^[222] Fluorine-containing agrichemicals are measurable in farmland runoff and nearby rivers.^[223]

Biological aspects

Natural biochemistry

Fluorine is not considered to be an essential mineral element for mammals or humans. Small amounts may be beneficial for bone strength, but this has not been definitively established. As there are many environmental sources of trace fluorine, the possibility of "fluorine deficiency" only pertains to artificial diets.^[224][225]

Biologically synthesized organofluorines have been found in microorganisms and plants,^[87] but not in animals.^[226] The most common example is fluoroacetate. It is used as a defense against herbivores by at least 40 green plants in Australia, Brazil, and Africa.^[227] Other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate.^[226] The enzyme adenosyl-fluoride synthase, which makes the carbon–fluorine bond, was isolated in bacteria in 2002.^[228]

Medicine

Dental care

Since the mid-20th century, population studies have shown that fluoride reduces tooth decay. The initial hypothesis was that fluoride helped by converting tooth enamel from the mineral hydroxyapatite to the more acid-resistant mineral fluorapatite. However, recent studies showed no difference in the frequency of caries (cavities) amongst teeth that were pre-fluoridated to different degrees. Current thinking is that fluoride prevents cavities primarily by helping teeth–which are in the very early stages of tooth decay–to regrow tooth enamel. In any case, it is only the fluoride that is directly present in the mouth (topical treatment) that prevents cavities; fluoride ions that are swallowed confer no such benefit.^[229]

Water fluoridation is the controlled addition of fluoride to a public water supply to reduce tooth decay.^[230] It began in the 1940s following studies of children in a region where water was naturally fluoridated. It is now used for about two thirds of the U.S. population on public water systems and for about 6% of people worldwide.^[231][232] Although the best available evidence shows no association with adverse effects other than dental fluorosis, most of which is mild,^[233] water fluoridation has been contentious for ethical, safety, and efficacy reasons.^[232] Opposition to water fluoridation exists despite its support by public health organizations.^[234] 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.^[235] Reviews of the scholarly literature in 2000 and 2007 associated water fluoridation with a significant reduction of tooth decay in children.^[236]

Toothpaste may contain fluorine in the form of sodium fluoride, tin difluoride, or (most commonly) sodium monofluorophosphate. The first fluoride toothpaste was introduced in 1955 in the U.S. Almost all toothpaste in developed countries is now fluoridated, as are many prescription and non-prescription mouthwashes. Fluoride may also be applied to teeth in gels, foams, or varnishes.^[235][237]

Pharmaceuticals

About 20% of modern pharmaceuticals contain fluorine.^[238] One of these, the cholesterol-reducer atorvastatin (Lipitor), was the number one money-making drug for nearly a decade, at least up until 2011.^[239] The branded asthma medication Seretide, a top-ten revenue drug as of the mid-2000s, contains two active ingredients, one of which—fluticasone−is fluorinated.^[240]

Even a single atom of fluorine added to a drug molecule can greatly change its chemical properties and thus how it interacts with the body. Because of the considerable stability of the carbon–fluorine bond, many drugs are fluorinated in order to delay their metabolism and elimination. This allows longer times between doses.^[241] Also, adding fluorine to organics increases their lipophilicity (ability to dissolve in fats) because the carbon–fluorine bond is even more hydrophobic than the carbon–hydrogen bond. This effect often increases a drug's bioavailability because of increased cell membrane penetration.^[240]

Many modern antidepressants are fluorinated molecules that selectively limit the body's binding of serotonin (with low serotonin availability in brain cells being a cause of depression). Prior to the 1980s, traditional antidepressants, such as the tricyclics, altered not only serotonin uptake but also affected several other neurotransmitters. This non-selective interaction caused many side effects. One of the first medications to alter only serotonin uptake—and be free of most side effects of previous pharmaceuticals—was the fluorine-containing drug fluoxetine (Prozac). It became the best-selling antidepressant. Some other selective serotonin reuptake inhibitor (SSRI) antidepressants that are fluorinated are citalopram (Celexa) and its isomer escitalopram (Lexapro), fluvoxamine (Luvox), and paroxetine (Paxil).^[242][243]

Quinolones are artificial compounds that are broad-spectrum antibiotics. Most of the currently used quinolones are fluorinated to make the drugs more powerful. Prominent examples include ciprofloxacin (Cipro) and levofloxacin (Levaquin). The latter was the highest selling U.S. antibiotic in 2010.^{[244][245][246][247]}

Fluorine also finds use in many steroidal drugs.^[248] Fludrocortisone (Florinef) is a mineralocorticoid (a compound used to retain sodium and water and thus raise blood pressure).^[249] Triamcinolone and dexamethasone are potent glucocorticoids (anti-inflammatories).^[249]

Several inhaled anesthetics, including the most common ones, are heavily fluorinated. The first fluorinated anesthetic, halothane, proved to be much safer (neither explosive nor flammable) and longer-lasting than those previously used. Modern fluorinated anesthetics are effective for even longer periods, and are almost insoluble in blood, allowing the patient to awaken more quickly. Examples include sevoflurane, desflurane, enflurane, and isoflurane, all of which are fluorinated ethers.^[250][251]

PET scanning

Fluorine-18 is routinely employed by the radiopharmaceutical industry, in particle accelerators, to produce radioactive tracers for positron emission tomography (PET) scanning. The radioisotope's half-life of almost two hours is long enough to allow its transportation from the production facility to the imaging center for radiation exposure to patients. The most widely used radiopharmaceutical is fluorodeoxyglucose (FDG).^[252] After injection into the blood, FDG is taken up by tissues with a high need for glucose, such as the brain and most types of malignant tumors.^[253] Computer assisted tomography (CAT) can then be used for detailed imaging.^[254]

Oxygen carrier research

Liquid fluorocarbons have a very high capacity for holding gas in solution. They can hold more oxygen or carbon dioxide than blood does and for that reason, have attracted ongoing interest as to their possible application in artificial blood, or liquid breathing.^[255]

Blood substitutes are the subject of research because the demand for blood transfusions grows faster than donations. In some scenarios, artificial blood may be more convenient or safe. Because fluorocarbons do not normally mix with water, they must be mixed into emulsions (small droplets of perfluorocarbon suspended in water) to be used as blood.^[256][257] One such product, Oxycyte, has been through initial clinical trials.^[258][259] PFC doping has the potential to aid endurance athletes and is therefore banned from sports. One cyclist's mysterious near death in 1998 prompted an investigation for PFC abuse.^[260][261]

Possible medical uses of liquid breathing (which uses pure perfluorocarbon liquid, not a water emulsion) involve assistance for premature babies or for burn victims (because the normal lung function is compromised). Both partial filling of the lungs and complete filling of the lungs have been considered, although only the former has any significant tests in humans. Several animal tests, and some human partial liquid ventilation trials, have been conducted.^[262] One effort, by Alliance Pharmaceuticals reached clinical trials but was abandoned because the results were not better than normal therapies.^[263]

Agrichemicals and poisons

An estimated 30% of agrichemical compounds contain fluorine.^[264] Most of them are herbicides and fungicides, but a few regulate crop growth. Fluorine substitution (usually of just a single atom or at most a trifluoromethyl group) is a powerful tool for new molecule design. The molecular effects—increasing biological stay time; membrane crossing; altering molecular recognition—are similar to those seen in fluorinated pharmaceuticals.^[265] Trifluralin is a prominent example, used widely in the U.S. as a weedkiller.^[265][266] Its suspected carcinogenic properties have caused many European countries to ban it.^[267]

Sodium monofluoroacetate (brand name 1080) is a commercial mammalian poison. The molecule is similar to the acetic acid molecule in vinegar but with a hydrogen atom changed out for fluorine atom (and another hydrogen atom changed out for a sodium atom). Fluoroacetate deprives cells of energy by replacing acetate in the Krebs cycle, halting a key part of cell metabolism. It was first synthesized in the late 19th century and then recognized as an insecticide in the early 20th century. Later, 1080 was widely used to control rats and other mammals. New Zealand is the largest consumer, using it to suppress the invasive Australian common brushtail possum, which threatens the indigenous kiwi.^[268] This particular poison is now banned in Europe and the U.S.^{[note 20][227][269]}

Fluorine-related hazards

Hydrofluoric acid

Hydrofluoric acid, the water solution of hydrogen fluoride, is a contact poison. Even though it is chemically only a weak acid, it is far more dangerous than the conventional strong mineral acids, such as nitric acid, sulfuric acid, or hydrochloric acid. Owing to its lesser chemical dissociation in water (thereby remaining a neutral molecule), hydrogen fluoride penetrates tissue more quickly than typical acids. Poisoning can occur readily through the skin or eyes or when inhaled or swallowed. From 1984 to 1994, at least nine U.S. workers died from accidents with HF.^[270]

Once in the blood, hydrogen fluoride reacts with calcium and magnesium, resulting in electrolyte imbalance (and potentially hypocalcemia). The consequent effect on the heart (cardiac arrhythmia) may be fatal.^[270] Formation of insoluble calcium fluoride also causes strong pain.^[271] Burns with areas larger than 160 cm² (about the size of a person's hand) can cause serious systemic toxicity.^[272]

Symptoms of exposure to hydrofluoric acid may not be immediately evident, with an 8-hour delay for 50% HF and up to 24 hours for lower concentrations. Hydrogen fluoride interferes with nerve function, meaning that burns may not initially be painful.

If the burn has been initially noticed, then HF should be washed off with a forceful stream of water for ten to fifteen minutes to prevent its further penetration into the body. Clothing used by the person burned may also present a danger.^[273] Hydrofluoric acid exposure is often treated with calcium gluconate, a source of calcium ions (Ca²⁺) that bind with the fluoride ions. Skin burns can be treated with a water wash and 2.5% calcium gluconate gel or special rinsing solutions.^{[274][275][276]} Because HF is absorbed, further medical treatment is necessary. Calcium gluconate may be injected or administered intravenously. Use of calcium chloride—a common laboratory reagent—in lieu of calcium gluconate is contraindicated, and may lead to severe complications. Sometimes surgical excision of tissue, or amputation, is required.^[272][277]

Fluoride ion

Soluble fluorides are moderately toxic. For sodium fluoride, the lethal dose for adults is 5–10 g, which is equivalent to 32–64 mg of elemental fluoride per kilogram of body weight.^[278] One fifth of the lethal dose may result in adverse health effects.^[279] Chronic excess fluoride consumption can lead to skeletal fluorosis, a disease of the bones that affects millions in Asia and Africa.^[279][280]

The fluoride ion is readily absorbed by the stomach and intestines. Ingested fluoride forms hydrofluoric acid in the stomach. In this form, it crosses cell membranes and then binds with calcium and interferes with various enzymes. The fluoride ion is excreted through urine. Exposure limits are based on urine testing, which determine the human body's capacity for ridding itself of fluoride ions.^[279][281]

Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluorides.^[282] Most current calls to poison control centers for possible fluoride poisoning come from the ingestion of fluoride-containing toothpaste.^[279] Malfunction of water fluoridation equipment has occurred several times, including an Alaskan incident that sickened nearly 300 people and killed one.^[283]

Dangers from toothpaste are more serious for small children—the U.S. Centers for Disease Control and Prevention recommend children under 6 years of age be supervised when brushing their teeth so that they do not swallow toothpaste.^[284] One regional study examined a year of fluoride poisoning reports for pre-teens. Of 87 cases, there was one fatality (from insecticide). The other 86 were all from dental fluoride. Most had no symptoms, but about 30% had stomach pains, with symptoms more likely when more fluoride was consumed.^[282] A larger study of American fluoride poisoning reports showed similar findings; 80% of the reports were related to children under 6. Few were serious, although several hundred cases were treated at health facilities each year.^[285]

See also

• Fluoride selective electrode, a device for measuring fluoride ion concentration
• Fluorine dating, a relative dating technique for buried organic matter
• Fluorine NMR, an analytical technique for fluorine-containing molecules
• Fluorous chemistry, a process of separating reagents from organic solvents
• Krypton fluoride laser and argon fluoride laser, two ultraviolet lasers

Notes

Sources

Citations

Indexed references

v t e Periodic table
	1	2															3	4	5	6	7	8	9	10	11	12	13	14	15	16	17	18
1	H																															He
2	Li	Be																									B	C	N	O	F	Ne
3	Na	Mg																									Al	Si	P	S	Cl	Ar
4	K	Ca															Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn	Ga	Ge	As	Se	Br	Kr
5	Rb	Sr															Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd	In	Sn	Sb	Te	I	Xe
6	Cs	Ba	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg	Tl	Pb	Bi	Po	At	Rn
7	Fr	Ra	Ac	Th	Pa	U	Np	Pu	Am	Cm	Bk	Cf	Es	Fm	Md	No	Lr	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Cn	Uut	Fl	Uup	Lv	Uus	Uuo
Alkali metal Alkaline earth metal Lan­thanide Actinide Transition metal Poor metal Metalloid Polyatomic nonmetal Diatomic nonmetal Noble gas Unknown chemical properties
Large version