View source for Fluorine

{{distinguish|Florin|Fluorene|Fluoride|Fluorone|Florine}}
{{Infobox fluorine}}
'''Fluorine''' is a [[chemical element]]; it has [[Chemical symbol|symbol]] '''F''' and [[atomic number]] 9. It is the lightest [[halogen]]{{refn|Assuming that [[hydrogen]] is not considered a halogen.||group="note"}} and exists at [[Standard temperature and pressure|standard conditions]] as pale yellow [[Diatomic molecule|diatomic]] gas. Fluorine is extremely [[Reactivity (chemistry)|reactive]] as it reacts with all other [[Periodic table|elements]] except for the light [[Noble gas|inert gases]]. It is highly [[toxicity|toxic]].

Among the elements, fluorine ranks [[Abundance of the chemical elements|24th in cosmic abundance and 13th in crustal abundance]]. [[Fluorite]], the primary mineral source of fluorine, which gave the element its name, was first described in 1529; as it was added to metal [[ore]]s to lower their melting points for [[smelting]], the Latin verb {{lang|la|fluo}} meaning {{gloss|to flow}} gave the mineral its name. Proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, and several early experimenters died or sustained injuries from their attempts. Only in 1886 did French chemist [[Henri Moissan]] isolate elemental fluorine using low-temperature [[electrolysis]], a process still employed for modern production. Industrial production of fluorine gas for [[Enriched uranium#Enrichment methods|uranium enrichment]], its largest application, began during the [[Manhattan Project]] in [[World War&nbsp;II]].

Owing to the expense of refining pure fluorine, most commercial applications use fluorine compounds, with about half of mined fluorite used in [[steelmaking]]. The rest of the fluorite is converted into [[hydrogen fluoride]] en route to various organic fluorides, or into [[cryolite]], which plays a key role in [[Hall–Héroult process|aluminium refining]]. The [[carbon–fluorine bond]] is usually very stable. Organofluorine compounds are widely used as [[refrigerant]]s, electrical insulation, and [[Polytetrafluoroethylene|PTFE]] (Teflon). Pharmaceuticals such as [[atorvastatin]] and [[fluoxetine]] contain C−F bonds. The [[Fluoride|fluoride ion]] from dissolved fluoride salts inhibits dental cavities and so finds use in [[toothpaste]] and [[water fluoridation]]. [[Fluorochemical industry|Global fluorochemical]] sales amount to more than [[United States dollar|US$]]15&nbsp;billion a year.

[[Fluorocarbon]] gases are generally [[greenhouse gas]]es with [[Global warming potential|global-warming potentials]] 100 to 23,500 times that of [[carbon dioxide]], and [[Sulfur hexafluoride|SF<sub>6</sub>]] has the highest global warming potential of any known substance. [[Organofluorine chemistry|Organofluorine compounds]] often persist in the environment due to the strength of the carbon–fluorine bond. Fluorine has no known metabolic role in mammals; a few [[plant]]s and [[Sponge|marine sponges]] synthesize organofluorine poisons (most often [[Fluoroacetic acid|monofluoroacetates]]) that help deter predation.{{sfn|Lee et al.|2014}}
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==Characteristics==

===Electron configuration===
Fluorine atoms have nine electrons, one fewer than [[neon]], and [[electron configuration]] 1s<sup>2</sup>2s<sup>2</sup>2p<sup>5</sup>: two electrons in a filled inner shell and seven in an outer shell requiring one more to be filled. The outer electrons are ineffective at nuclear [[Shielding effect|shielding]], and experience a high [[effective nuclear charge]] of 9&nbsp;− 2&nbsp;= 7; this affects the atom's physical properties.{{sfn|Jaccaud et al.|2000|p=381}}

Fluorine's [[ionization energy|first ionization energy]] is third-highest among all elements, behind helium and neon,{{sfn|Dean|1999|p=564}} which complicates the removal of electrons from neutral fluorine atoms. It also has a high [[electron affinity]], second only to [[chlorine]],{{sfn|Lide|2004|pp=10.137–10.138}} and tends to capture an electron to become [[isoelectronicity|isoelectronic]] with the noble gas neon;{{sfn|Jaccaud et al.|2000|p=381}} it has the highest [[electronegativity]] of any reactive element.<ref>{{harvnb|Moore|Stanitski|Jurs|2010|p=[https://books.google.com/books?id=ZOm8L9oCwLMC&pg=PA156 156]}}.</ref> Fluorine atoms have a small [[Covalent radius of fluorine|covalent radius]] of around 60&nbsp;[[picometer]]s, similar to those of its [[period 2 element|period]] neighbors oxygen and neon.<ref name="CSD">{{Harvnb|Cordero et al.|2008}}.</ref><ref>{{Harvnb|Pyykkö|Atsumi|2009}}.</ref>{{refn|Sources disagree on the radii of oxygen, fluorine, and neon atoms. Precise comparison is thus impossible.||group="note"}}
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===Reactivity===
{{external media|width=240px|float=right|video1=[https://www.youtube.com/watch?v=vtWp45Eewtw Bright flames] during fluorine reactions|video2 = Fluorine [https://www.youtube.com/watch?v=TLOFaWdPxB0 reacting with caesium]}}
[[File:Fluorine molecule ball.png|thumb|Fluorine 3D molecule]]
The [[bond energy]] of [[Compounds of fluorine#Difluorine|difluorine]] is much lower than that of either {{chem|Cl|2}} or {{chem|Br|2}} and similar to the easily cleaved [[peroxide]] bond; this, along with high electronegativity, accounts for fluorine's easy [[Dissociation (chemistry)|dissociation]], high reactivity, and strong bonds to non-fluorine atoms.{{sfn|Greenwood|Earnshaw|1998|p=804}}<ref>{{harvnb|Macomber|1996|p=[https://books.google.com/books?id=HyuogOtzoaYC&pg=PA230 230]}}</ref> Conversely, bonds to other atoms are very strong because of fluorine's high electronegativity. Unreactive substances like [[Powder coating|powdered steel]], glass fragments, and [[asbestos]] fibers react quickly with cold fluorine gas; wood and water spontaneously combust under a fluorine jet.{{sfn|Jaccaud et al.|2000|p=382}}<ref name="pop mechanics F reactivity">{{harvnb|Nelson|1947}}.</ref>

Reactions of elemental fluorine with metals require varying conditions. [[Alkali metal]]s cause explosions and [[alkaline earth metal]]s display vigorous activity in bulk; to prevent [[Passivation (chemistry)|passivation]] from the formation of metal fluoride layers, most other metals such as aluminium and iron must be powdered,{{sfn|Greenwood|Earnshaw|1998|p=804}} and [[noble metal]]s require pure fluorine gas at {{convert|300|–|450|C|F}}.{{sfn|Lidin|Molochko|Andreeva|2000|pp=442–455}} Some solid nonmetals (sulfur, phosphorus) react vigorously in liquid fluorine.{{sfn|Wiberg|Wiberg|Holleman|2001|p=404}} [[Hydrogen sulfide]]{{sfn|Wiberg|Wiberg|Holleman|2001|p=404}} and [[sulfur dioxide]]{{sfn|Patnaik|2007|p=472}} combine readily with fluorine, the latter sometimes explosively; [[sulfuric acid]] exhibits much less activity, requiring elevated temperatures.{{sfn|Aigueperse et al.|2000|p=400}}

[[Hydrogen]], like some of the alkali metals, reacts explosively with fluorine.{{sfn|Greenwood|Earnshaw|1998|pp=76, 804}} [[Carbon]], as [[lamp black]], reacts at room temperature to yield [[tetrafluoromethane]]. Graphite combines with fluorine above {{convert|400|°C|°F|abbr=on}} to produce [[non-stoichiometric compound|non-stoichiometric]] [[carbon monofluoride]]; higher temperatures generate gaseous [[fluorocarbon]]s, sometimes with explosions.<ref name="Kuriakose, A. K.">{{harvnb|Kuriakose|Margrave|1965}}.</ref> Carbon dioxide and carbon monoxide react at or just above room temperature,<ref name="Hasegawa, Yasuo">{{harvnb|Hasegawa et al.|2007}}.</ref> whereas [[paraffin wax|paraffins]] and other organic chemicals generate strong reactions:<ref>{{harvnb|Lagow|1970|pp=64–78}}.</ref> even completely substituted [[haloalkane]]s such as [[carbon tetrachloride]], normally incombustible, may explode.<ref name="Navarrini, Walter">{{harvnb|Navarrini et al.|2012}}.</ref> Although [[nitrogen trifluoride]] is stable, nitrogen requires an [[Electric discharge in gases|electric discharge]] at elevated temperatures for reaction with fluorine to occur, due to the very strong [[triple bond]] in elemental nitrogen;{{sfn|Lidin|Molochko|Andreeva|2000|p=252}} ammonia may react explosively.<ref name="Anhydrous ammonia MSDS">{{Harvnb|Tanner Industries|2011}}.</ref><ref name="Scott I. Morrow">{{harvnb|Morrow|Perry|Cohen|1959}}.</ref> [[Oxygen]] does not combine with fluorine under ambient conditions, but can be made to react using electric discharge at low temperatures and pressures; the products tend to disintegrate into their constituent elements when heated.<ref name="Emeléus Sharpe 1974 111">{{harvnb|Emeléus|Sharpe|1974|p=[https://books.google.com/books?id=qi0vujzNSz4C&pg=PA111 111]}}.</ref>{{sfn|Wiberg|Wiberg|Holleman|2001|p=457}}<ref>{{harvnb|Brantley|1949|p=[https://books.google.com/books?id=Yo7r_1qJG2gC&pg=PA26-IA83 26]}}.</ref> Heavier halogens{{sfn|Jaccaud et al.|2000|p=383}} react readily with fluorine as does the noble gas [[radon]];<ref>{{harvnb|Pitzer|1975}}.</ref> of the other noble gases, only [[xenon]] and [[krypton]] react, and only under special conditions.<ref name="Khriachtchev et al. 2000">{{harvnb|Khriachtchev et al.|2000}}.</ref> [[Argon]] does not react with fluorine gas; however, it does form a compound with fluorine, [[argon fluorohydride]].
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===Phases===
[[File:Beta fluorine unit cell.svg|right|thumb|Crystal structure of β-fluorine. Spheres indicate {{chem|F|2}} molecules that may assume any angle. Other molecules are constrained to planes.|alt=Cube with spherical shapes on the corners and center and spinning molecules in planes in faces]]
{{main|Phases of fluorine}}
[[File:Beta-fluorine crystal structure.gif|thumb|Animation showing the crystal structure of beta-fluorine. Molecules on the faces of the unit cell have rotations constrained to a plane.]]
At room temperature, fluorine is a gas of [[diatomic molecule]]s,{{sfn|Jaccaud et al.|2000|p=382}} pale yellow when pure (sometimes described as yellow-green).<ref>{{harvnb|Burdon|Emson|Edwards|1987}}.</ref> It has a characteristic halogen-like pungent and biting odor detectable at 20&nbsp;[[parts per billion|ppb]].{{sfn|Lide|2004|p=4.12}} Fluorine condenses into a bright yellow liquid at {{convert|-188|°C|°F|abbr=on}}, a transition temperature similar to those of oxygen and nitrogen.{{sfn|Dean|1999|p=523}}

Fluorine has two solid forms, α- and β-fluorine. The latter crystallizes at {{convert|-220|°C|°F|abbr=on}} and is transparent and soft, with the same disordered [[Cubic crystal system|cubic]] structure of freshly crystallized solid oxygen,{{sfn|Dean|1999|p=523}}{{refn|α-Fluorine has a regular pattern of molecules and is a crystalline solid, but its molecules do not have a specific orientation. β-Fluorine's molecules have fixed locations and minimal rotational uncertainty.<ref>{{harvnb|Pauling|Keaveny|Robinson|1970}}.</ref>||group="note"}} unlike the [[orthorhombic crystal system|orthorhombic]] systems of other solid halogens.{{sfn|Young|1975|p=10}}<ref name="Barrett 1967" /> Further cooling to {{convert|-228|°C|°F|abbr=on}} induces a [[phase transition]] into opaque and hard α-fluorine, which has a [[monoclinic crystal system|monoclinic]] structure with dense, angled layers of molecules. The transition from β- to α-fluorine is more [[exothermic]] than the condensation of fluorine, and can be violent.{{sfn|Young|1975|p=10}}<ref name="Barrett 1967">{{harvnb|Barrett|Meyer|Wasserman|1967}}.</ref>

===Isotopes===
{{Main|Isotopes of fluorine}}
Only one [[isotope]] of fluorine occurs naturally in abundance, the stable isotope {{chem|19|F}}.<ref>{{Harvnb|National Nuclear Data Center|NuDat 2.1|loc=[http://www.nndc.bnl.gov/nudat2/reCenter.jsp?z=9&n=10 Fluorine-19]}}.</ref> It has a high [[magnetogyric ratio]]{{refn|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."<ref>{{harvnb|Vigoureux|1961}}.</ref>|group="note"}} and exceptional sensitivity to magnetic fields; because it is also [[monoisotopic element|the only stable isotope]], it is [[Fluorine-19 nuclear magnetic resonance|used]] in [[magnetic resonance imaging]].<ref>{{harvnb|Meusinger|Chippendale|Fairhurst|2012|pp=752, 754}}.</ref> Eighteen [[radioisotope]]s with [[mass number]]s 13–31 have been synthesized, of which [[Fluorine-18|{{chem|18|F}}]] is the most stable with a [[half-life]] of 109.734&nbsp;minutes.{{NUBASE2020|ref}}<ref>{{cite journal | last=Kahlbow | first=J. | display-authors=etal | collaboration=SAMURAI21-NeuLAND Collaboration | title=Magicity versus Superfluidity around <sup>28</sup>O viewed from the Study of <sup>30</sup>F | journal=Physical Review Letters | volume=133 | issue=8 | date=2024-08-23 | page=082501 | issn=0031-9007 | doi=10.1103/PhysRevLett.133.082501| doi-access=free | arxiv=2407.19303 }}</ref> {{chem|18|F}} is a natural [[trace radioisotope]] produced by [[cosmic ray spallation]] of atmospheric [[argon]] as well as by reaction of protons with natural oxygen: <sup>18</sup>O + p → <sup>18</sup>F + n.<ref name=SCOPE50>
[http://www.scopenvironment.org/downloadpubs/scope50 SCOPE 50 - Radioecology after Chernobyl] {{webarchive|url=https://web.archive.org/web/20140513065145/http://www.scopenvironment.org/downloadpubs/scope50/ |date=2014-05-13 }}, the [[Scientific Committee on Problems of the Environment]] (SCOPE), 1993. See table 1.9 in Section 1.4.5.2.</ref> Other radioisotopes have half-lives less than 70&nbsp;seconds; most decay in less than half a second.<ref name="isotopes">{{Harvnb|National Nuclear Data Center|NuDat 2.1}}.</ref> The isotopes {{chem|17|F}} and {{chem|18|F}} undergo [[positron emission|β<sup>+</sup> decay]] and [[electron capture]], lighter isotopes decay by [[proton emission]], and those heavier than {{chem|19|F}} undergo [[beta minus decay|β<sup>−</sup> decay]] (the heaviest ones with delayed [[neutron emission]]).<ref name="isotopes" /><ref>{{harvnb|''NUBASE''|2016|pp=030001-23–030001-27}}.</ref> Two [[nuclear isomer|metastable isomers]] of fluorine are known, {{chem|18m|F}}, with a half-life of 162(7)&nbsp;nanoseconds, and {{chem|26m|F}}, with a half-life of 2.2(1)&nbsp;milliseconds.<ref>{{harvnb|''NUBASE''|2016|pp=030001–24}}.</ref>

==Occurrence==
{{Main|Origin and occurrence of fluorine}}

===Universe===
{| class="wikitable" style="float:right; margin-left:15px; margin-top:0; font-size:10pt; line-height:11pt;"
|+ style="margin-bottom: 5px;" | Solar System abundances<ref name="Cameron">{{harvnb|Cameron|1973}}.</ref>
! style="text-align:center;" | Atomic <br /> number
! style="width:45%;"| Element
! style="padding-right: 5px; padding-left: 10px;" | Relative <br /> amount
|-
| style="text-align:center;" | 6
| style="text-align:center;"| Carbon
| style="padding-right:5px; text-align:right;"| 4,800
|-
| style="text-align:center;" | 7
| style="text-align:center;"| Nitrogen
| style="padding-right:5px; text-align:right;"|1,500
|-
| style="text-align:center;" | 8
| style="text-align:center;"| Oxygen
| style="padding-right:5px; text-align:right;"|8,800
|- style="background:#ff9;"
| style="text-align:center;"| 9
| style="text-align:center; "| Fluorine
| style="padding-right:5px; text-align:right;"|1
|-
| style="text-align:center;" | 10
| style="text-align:center;"| Neon
| style="padding-right:5px; text-align:right;"|1,400
|-
| style="text-align:center;" | 11
| style="text-align:center;"| Sodium
| style="padding-right:5px; text-align:right;"|24
|-
| style="text-align:center;" | 12
| style="text-align:center;"| Magnesium
| style="padding-right:5px; text-align:right;"|430
|}
Among the lighter elements, fluorine's abundance value of 400&nbsp;[[parts-per notation|ppb]] (parts per billion) – 24th among elements in the universe – is exceptionally low: other elements from carbon to magnesium are twenty or more times as common.<ref name="Crosswell" /> This is because [[stellar nucleosynthesis]] processes bypass fluorine, and any fluorine atoms otherwise created have high [[nuclear cross section]]s, allowing collisions with hydrogen or helium to generate oxygen or neon respectively.<ref name="Crosswell" /><ref>{{Harvnb|Clayton|2003|pp=[https://books.google.com/books?id=fXcdHyLUVnEC&pg=PA101 101–104]}}.</ref>

Beyond this transient existence, three explanations have been proposed for the presence of fluorine:<ref name="Crosswell">{{Harvnb|Croswell|2003}}.</ref><ref name="Milky Way">{{harvnb|Renda et al.|2004}}.</ref>
* during [[type II supernova]]e, bombardment of neon atoms by [[neutrino]]s could transmute them to fluorine;
* the solar wind of [[Wolf–Rayet star]]s could blow fluorine away from any hydrogen or helium atoms; or
* fluorine is borne out on convection currents arising from fusion in [[asymptotic giant branch]] stars.

===Earth===
{{See also|List of countries by fluorite production}}
Fluorine is the thirteenth most [[Abundance of elements in Earth's crust|common element in Earth's crust]] at 600–700&nbsp;ppm (parts per million) by mass.{{sfn|Jaccaud et al.|2000|p=384}} Though believed not to occur naturally, elemental fluorine has been shown to be present as an occlusion in antozonite, a variant of fluorite.{{sfn|Schmedt|Mangstl|Kraus|2012}} Most fluorine exists as fluoride-containing minerals. [[Fluorite]], [[fluorapatite]] and [[cryolite]] are the most industrially significant.{{sfn|Jaccaud et al.|2000|p=384}}{{sfn|Greenwood|Earnshaw|1998|p=795}} Fluorite ({{chem|Ca|F|2}}), also known as fluorspar, abundant worldwide, is the main source of fluoride, and hence fluorine. China and Mexico are the major suppliers.{{sfn|Greenwood|Earnshaw|1998|p=795}}<ref name="KGS fluorite terminology" /><ref name="AcountF">{{harvnb|Villalba|Ayres|Schroder|2008}}.</ref><ref name="historical fluorspar statistics">{{Harvnb|Kelly|Miller|2005}}.</ref><ref name="UK fluorspar">{{Harvnb|Lusty et al.|2008}}.</ref> Fluorapatite (Ca<sub>5</sub>(PO<sub>4</sub>)<sub>3</sub>F), which contains most of the world's fluoride, is an inadvertent source of fluoride as a byproduct of fertilizer production.{{sfn|Greenwood|Earnshaw|1998|p=795}} Cryolite ({{chem|Na|3|Al|F|6}}), used in the production of aluminium, is the most fluorine-rich mineral. Economically viable natural sources of cryolite have been exhausted, and most is now synthesised commercially.{{sfn|Greenwood|Earnshaw|1998|p=795}}

<gallery mode="packed" heights="200px">
File:Fluorite-270246.jpg|Fluorite: Pink globular mass with crystal facets
File:Apatite Canada.jpg|Fluorapatite: Long, prismatic crystal, dull in [[Lustre (mineralogy)|lustre]], protruding, at an angle, from [[Matrix (geology)|matrix]] of aggregate-like rock
File:Ivigtut cryolite edit.jpg|Cryolite: A parallelogram-shaped outline with diatomic molecules arranged in two layers
</gallery>

Other minerals such as [[topaz]] contain fluorine. Fluorides, unlike other halides, are insoluble and do not occur in commercially favorable concentrations in saline waters.{{sfn|Greenwood|Earnshaw|1998|p=795}} Trace quantities of organofluorines of uncertain origin have been detected in volcanic eruptions and geothermal springs.{{sfn|Gribble|2002}} The existence of gaseous fluorine in crystals, suggested by the smell of crushed [[antozonite]], is contentious;{{sfn|Richter|Hahn|Fuchs|2001|p=3}}{{sfn|Schmedt|Mangstl|Kraus|2012}} a 2012 study reported the presence of 0.04% {{chem|F|2}} by weight in antozonite, attributing these [[inclusion (mineral)|inclusions]] to radiation from the presence of tiny amounts of [[uranium]].{{sfn|Schmedt|Mangstl|Kraus|2012}}

==History==
{{Main|History of fluorine}}

===Early discoveries===
[[File:Book9-25.gif|thumb|upright|alt=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]]'']]
In 1529, [[Georgius Agricola]] described fluorite as an additive used to lower the melting point of metals during [[smelting]].{{sfn|Greenwood|Earnshaw|1998|p=790}}<ref name="assassinated">{{harvnb|Senning|2007|p=[https://books.google.com/books?id=Fl4sdCYrq3cC&pg=PA149 149]}}.</ref>{{refn|[[Basilius Valentinus]] supposedly described fluorite in the late 15th century, but because his writings were uncovered 200 years later, this work's veracity is doubtful.<ref name="PopSci">{{Harvnb|Stillman|1912}}.</ref><ref>{{harvnb|Principe|2012|pp=140, 145}}.</ref><ref>{{Harvnb|Agricola|Hoover|Hoover|1912|loc=footnotes and commentary, pp. xxx, 38, 409, 430, 461, 608}}.</ref>||group="note"}} He penned the Latin word ''fluorēs'' (''fluor,'' flow) for fluorite rocks. The name later evolved into ''fluorspar'' (still commonly used) and then ''fluorite''.<ref name="KGS fluorite terminology">{{harvnb|Norwood|Fohs|1907|p=[https://archive.org/stream/bulletin01kentgoog#page/n94/mode/2up 52]}}.</ref>{{sfn|Greenwood|Earnshaw|1998|p=109}}<ref>{{Harvnb|Agricola|Hoover|Hoover|1912|loc=preface, pp.&nbsp;[https://archive.org/stream/georgiusagricola00agririch#page/380/mode/2up 380–381]}}.</ref> The composition of fluorite was later determined to be [[calcium difluoride]].<ref name="Weeks">{{harvnb|Weeks|1932}}.</ref>

[[Hydrofluoric acid]] was used in [[glass etching]] from 1720 onward.{{refn|Or perhaps from as early as 1670 onwards; Partington<ref name="Partington">{{harvnb|Partington|1923}}.</ref> and Weeks<ref name="Weeks" /> give differing accounts.||group="note"}} [[Andreas Sigismund Marggraf]] first characterized it in 1764 when he heated fluorite with sulfuric acid, and the resulting solution corroded its glass container.<ref>{{harvnb|Marggraf|1770}}.</ref><ref name="ModFcchem" /> Swedish chemist [[Carl Wilhelm Scheele]] repeated the experiment in 1771, and named the acidic product ''fluss-spats-syran'' (fluorspar acid).<ref name="ModFcchem" /><ref>{{harvnb|Scheele|1771}}.</ref> In 1810, the French physicist [[André-Marie Ampère]] suggested that hydrogen and an element analogous to chlorine constituted hydrofluoric acid.<ref>{{harvnb|Ampère|1816}}.</ref> He also proposed in a letter to [[Humphry Davy|Sir Humphry Davy]] dated August 26, 1812 that this then-unknown substance may be named ''fluorine'' from fluoric acid and the ''-ine'' suffix of other halogens.<ref>{{Cite book|url=https://books.google.com/books?id=kslaDwAAQBAJ&pg=PA3|title = Fluorine: A Paradoxical Element|isbn = 9780128129913|last1 = Tressaud|first1 = Alain|date = 6 October 2018| publisher=Academic Press }}</ref>{{sfn|Davy|1813|p=[https://books.google.com/books?id=dhdGAAAAMAAJ&pg=PA278 278]}} This word, often with modifications, is used in most European languages; however, Greek, Russian, and some others, following Ampère's later suggestion, use the name ''ftor'' or derivatives, from the Greek φθόριος (''phthorios'', destructive).{{sfn|Banks|1986|p=11}} The New Latin name ''fluorum'' gave the element its current symbol '''F'''; Fl was used in early papers.<ref name="dissolve" />{{refn|Fl, since 2012, is used for [[flerovium]].||group="note"}}

===Isolation===
[[file:Recherches sur l’isolement du fluor, Fig. 5.PNG|thumb|upright|left|1887 drawing of Moissan's apparatus]]
Initial studies on fluorine were so dangerous that several 19th-century experimenters were deemed "fluorine martyrs" after misfortunes with hydrofluoric acid.{{refn|[[Humphry Davy|Davy]], [[Joseph Louis Gay-Lussac|Gay-Lussac]], [[Louis Jacques Thénard|Thénard]], and the Irish chemists Thomas and George Knox were injured. Belgian chemist [[Paulin Louyet]] and French chemist {{ill|Jérôme Nicklès|de}} died. Moissan also experienced serious hydrogen fluoride poisoning.<ref name="Weeks" /><ref name="Toon" />||group="note"}} Isolation of elemental fluorine was hindered by the extreme corrosiveness of both elemental fluorine itself and hydrogen fluoride, as well as the lack of a simple and suitable [[electrolyte]].<ref name="Weeks" /><ref name = "Toon" /> [[Edmond Frémy]] postulated that [[electrolysis]] of pure hydrogen fluoride to generate fluorine was feasible and devised a method to produce anhydrous samples from acidified [[potassium bifluoride]]; instead, he discovered that the resulting (dry) hydrogen fluoride [[insulator (electricity)|did not conduct]] electricity.<ref name="Weeks" /><ref name="Toon" /><ref name="Asimov">{{harvnb|Asimov|1966|p=162}}.</ref> 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 cell]]s, he cooled the reaction to extremely low temperatures in a special bath and forged cells from a more resistant mixture of platinum and [[iridium]], and used fluorite stoppers.<ref name="Toon" />{{sfn|Greenwood|Earnshaw|1998|pp=789–791}} In 1886, after 74 years of effort by many chemists, Moissan isolated elemental fluorine.<ref name="Asimov" /><ref name="'qwen'">{{harvnb|Moissan|1886}}.</ref>

In 1906, two months before his death, Moissan received the [[Nobel Prize in Chemistry]],<ref name="Chem Nobel laureates">{{harvnb|Viel|Goldwhite|1993|p=[https://books.google.com/books?id=jEy67gEvIuMC&pg=PA35 35]}}.</ref> with the following citation:<ref name="Toon">{{harvnb|Toon|2011}}.</ref>
{{blockquote|1=[I]n recognition of the great services rendered by him in his investigation and isolation of the element fluorine&nbsp;... The whole world has admired the great experimental skill with which you have studied that savage beast among the elements.{{refn|Also honored was his invention of the [[electric arc furnace]].||group="note"}}}}

===Later uses===
[[File:Uranium hexafluoride crystals sealed in an ampoule.jpg|thumb|upright=1.15|right|An ampoule of [[uranium hexafluoride]]]]
The [[Frigidaire]] division of [[General Motors]] (GM) 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 Freon-12 ({{chem|link=dichlorodifluoromethane|CCl|2|F|2}}) as one such [[refrigerant]]. It replaced earlier and more toxic compounds, increased demand for kitchen refrigerators, and became profitable; by 1949 DuPont had bought out Kinetic and marketed several other [[Freon]] compounds.<ref name="ModFcchem" /><ref name="Jstg">{{harvnb|Okazoe|2009}}.</ref>{{sfn|Hounshell|Smith|1988|pp=156–157}}<ref name="DD Freon history">{{Harvnb|DuPont|2013a}}.</ref> [[Polytetrafluoroethylene]] (Teflon) was serendipitously discovered in 1938 by [[Roy J. Plunkett]] while working on refrigerants at Kinetic, and its superlative chemical and thermal resistance lent it to accelerated commercialization and mass production by 1941.<ref name="ModFcchem" /><ref name="Jstg" />{{sfn|Hounshell|Smith|1988|pp=156–157}}

Large-scale production of elemental fluorine began during World War II. Germany used high-temperature electrolysis to make tons of the planned incendiary [[chlorine trifluoride]]<ref>{{harvnb|Meyer|1977|p=111}}.</ref> and the [[Manhattan Project]] used huge quantities to produce [[uranium hexafluoride]] for uranium enrichment. Since {{chem|UF|6}} 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.<ref>{{harvnb|Kirsch|2004|pp=[https://books.google.com/books?id=cqnEZSTA6cUC&pg=PA61 60–66]}}.</ref>

==Compounds==
{{main|Fluorine compounds}}
Fluorine has a rich chemistry, encompassing organic and inorganic domains. It combines with metals, nonmetals, [[metalloid]]s, and most noble gases.<ref name="Riedel Kaupp 2009">{{harvnb|Riedel|Kaupp|2009}}.</ref> Fluorine's high electron affinity results in a preference for [[ionic bonding]]; when it forms [[covalent bond]]s, these are polar, and almost always [[single bond|single]].<ref name="NF charge">{{harvnb|Harbison|2002}}.</ref><ref>{{harvnb|Edwards|1994|p=[https://books.google.com/books?id=TRMRDXHWX7EC&pg=PA515 515]}}.</ref><ref group="note">The metastable [[boron monofluoride|boron]] and [[nitrogen monofluoride]] have higher-order fluorine bonds, and some [[metal complex]]es use it as a [[bridging ligand]]. [[Hydrogen bonding]] is another possibility.</ref>

===Oxidation states===
In compounds, fluorine almost exclusively assumes an [[oxidation state]] of −1. Fluorine in {{chem|F|2}} is defined to have oxidation state 0. The unstable species {{chem|F|2|-}} and {{chem|F|3|-}}, which decompose at around 40&nbsp;K, have intermediate oxidation states;{{sfn|Wiberg|Wiberg|Holleman|2001|p=422}} {{chem|F|4|+}} and a few related species are predicted to be stable.<ref>{{harvnb|Schlöder|Riedel|2012}}.</ref>

===Metals===
{{See also|Fluoride volatility}}
Alkali metals form ionic and highly soluble [[monofluoride]]s; these have the [[Cubic crystal system#Rock-salt structure|cubic arrangement of sodium chloride]] and analogous chlorides.<ref>{{harvnb|Katakuse et al.|1999|p=[https://books.google.com/books?id=pjFTI5KgMFkC&pg=PA267 267]}}.</ref>{{sfn|Aigueperse et al.|2000|pp=420–422}} Alkaline earth [[difluoride]]s possess strong ionic bonds but are insoluble in water,<ref name="dissolve">{{harvnb|Storer|1864|pp=[https://archive.org/details/b21901053/page/278 278–280]}}.</ref> with the exception of [[beryllium difluoride]], which also exhibits some covalent character and has a [[silicon dioxide|quartz]]-like structure.<ref name="Beryllium chemistry">{{harvnb|Walsh|2009|pp=[https://books.google.com/books?id=3-GbhmSfyeYC&pg=PA99 99–102], [https://books.google.com/books?id=3-GbhmSfyeYC&pg=PA118 118–119]}}.</ref> [[Rare earth element]]s and many other metals form mostly ionic [[trifluoride]]s.{{sfn|Emeléus|Sharpe|1983|pp=89–97}}<ref name="Crystal chemistry of fluorides">{{harvnb|Babel|Tressaud|1985|pp=[https://books.google.com/books?id=SaWLysMGF3cC&pg=PA91 91–96]}}.</ref><ref>{{harvnb|Einstein et al.|1967}}.</ref>

Covalent bonding first comes to prominence in the [[tetrafluoride]]s: those of [[zirconium]], [[hafnium]]<ref>{{harvnb|Brown et al.|2005|p=[https://books.google.com/books?id=DvqwTdVhjMEC&pg=PA144 144]}}.</ref><ref name="Perry 2011 193">{{harvnb|Perry|2011|p=[https://books.google.com/books?id=SFD30BvPBhoC&pg=PA193 193]}}.</ref> and several [[actinides]]<ref>{{harvnb|Kern et al.|1994}}.</ref> are ionic with high melting points,{{sfn|Lide|2004|pp=4.60, 4.76, 4.92, 4.96}}{{refn|{{chem|Zr|F|4}} melts at {{convert|932|°C|°F|abbr=on}},{{sfn|Lide|2004|p=4.96}} {{chem|Hf|F|4}} sublimes at {{convert|968|°C|°F|abbr=on}},<ref name="Perry 2011 193" /> and {{chem|U|F|4}} melts at {{convert|1036|°C|°F|abbr=on}}.{{sfn|Lide|2004|p=4.92}}|group="note"}} while those of [[titanium]],{{sfn|Greenwood|Earnshaw|1998|p=964}} [[vanadium]],<ref>{{harvnb|Becker|Müller|1990}}.</ref> and [[niobium]] are polymeric,{{sfn|Greenwood|Earnshaw|1998|p=990}} melting or decomposing at no more than {{convert|350|°C|°F|abbr=on}}.{{sfn|Lide|2004|pp=4.72, 4.91, 4.93}} [[Pentafluoride]]s continue this trend with their linear polymers and [[oligomer]]ic complexes.{{sfn|Greenwood|Earnshaw|1998|pp=561–563}}{{sfn|Emeléus|Sharpe|1983|pp=256–277}}{{sfn|Mackay|Mackay|Henderson|2002|pp=355–356}} Thirteen metal [[hexafluoride]]s are known,{{refn|These thirteen are those of molybdenum, technetium, ruthenium, rhodium, tungsten, rhenium, osmium, iridium, platinum, polonium, uranium, neptunium, and plutonium.|group="note"}} all octahedral, and are mostly volatile solids but for liquid {{chem|link=molybdenum hexafluoride|Mo|F|6}} and {{chem|link=rhenium hexafluoride|Re|F|6}}, and gaseous {{chem|link=tungsten hexafluoride|W|F|6}}.{{sfn|Greenwood|Earnshaw|1998|loc=(various pages, by metal in respective chapters)}}{{sfn|Lide|2004|pp=4.71, 4.78, 4.92}}<ref name="D_BLOCK_XF6">{{harvnb|Drews et al.|2006}}.</ref> [[Rhenium heptafluoride]], the only characterized metal [[heptafluoride]], is a low-melting molecular solid with [[pentagonal bipyramidal molecular geometry]].{{sfn|Greenwood|Earnshaw|1998|p=819}} Metal fluorides with more fluorine atoms are particularly reactive.<ref>{{harvnb|Bartlett|1962}}.</ref>

{| cellpadding="3" style="margin:1em auto 1em auto; text-align:center;"
|-
| colspan="3" style="text-align:center;"|'''Structural progression of metal fluorides'''
|-
| [[File:Sodium-fluoride-unit-cell-3D.png|132px|alt=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|Sodium fluoride: cubic lattice of alternating sodium and fluorine atoms with no distinct molecules]]
| [[File:Bismuth-pentafluoride-chain-from-xtal-1971-3D-balls.png|500px|alt=Straight chain of alternating balls, violet and yellow, with violet ones also linked to four more yellow perpendicularly to the chain and each other|Bismuth pentafluoride: arbitrarily long, straight chain of atoms]]
| [[File:Rhenium-heptafluoride-3D-balls.png|121px|alt=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|Rhenium heptafluoride: discrete small molecule]]
|- style="text-align:center;"
| [[Sodium fluoride]], ionic
| [[Bismuth pentafluoride]], polymeric
| [[Rhenium heptafluoride]], molecular
|}

===Hydrogen===
{{Main|Hydrogen fluoride|hydrofluoric acid}}
[[File:Boiling-points Chalcogen-Halogen.svg|alt=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|thumb|right]]
Hydrogen and fluorine combine to yield hydrogen fluoride, in which discrete molecules form clusters by hydrogen bonding, resembling water more than [[hydrogen chloride]].<ref name="Pauling HF hydrogen bonds">{{harvnb|Pauling|1960|pp=[https://archive.org/details/natureofchemical0000paul_3ed/page/454 454–464]}}.</ref><ref name="Atkins HF">{{harvnb|Atkins|Jones|2007|pp=184–185}}.</ref><ref name="New Scientist HF">{{harvnb|Emsley|1981}}.</ref> It boils at a much higher temperature than heavier hydrogen halides and unlike them is [[miscible]] with water.{{sfn|Greenwood|Earnshaw|1998|pp=812–816}} Hydrogen fluoride readily hydrates on contact with water to form aqueous hydrogen fluoride, also known as hydrofluoric acid. Unlike the other hydrohalic acids, which are [[strong acids|strong]], hydrofluoric acid is a [[weak acid]] at low concentrations.{{sfn|Wiberg|Wiberg|Holleman|2001|p=425}}<ref>{{Harvnb|Clark|2002}}.</ref> However, it can attack glass, something the other acids cannot do.<ref name="H+">{{harvnb|Chambers|Holliday|1975|pp=328–329}}.</ref>

===Other reactive nonmetals===
[[File:Chlorine-trifluoride-3D-balls.png|thumb|[[Chlorine trifluoride]], whose corrosive potential ignites asbestos, concrete, sand and other fire retardants{{sfn|Air Products and Chemicals|2004|p=1}}]]
Binary fluorides of metalloids and p-block nonmetals are generally covalent and volatile, with varying reactivities. [[Period 3]] and heavier nonmetals can form [[hypervalent molecule|hypervalent]] fluorides.<ref name="Noury">{{harvnb|Noury|Silvi|Gillespie|2002}}.</ref>

[[Boron trifluoride]] is planar and possesses an incomplete octet. It functions as a [[Lewis acids and bases|Lewis acid]] and combines with Lewis bases like ammonia to form [[adduct]]s.<ref>{{harvnb|Chang|Goldsby|2013|p=706}}.</ref> [[Carbon tetrafluoride]] is tetrahedral and inert;{{refn|Carbon tetrafluoride is formally organic, but is included here rather than in the [[#Organic compounds|organofluorine chemistry section]] – where more complex carbon-fluorine compounds are discussed – for comparison with {{chem|Si|F|4}} and {{chem|Ge|F|4}}.|group="note"}} [[carbon group|its group]] analogues, silicon and germanium tetrafluoride, are also tetrahedral<ref>{{harvnb|Ellis|2001|p=69}}.</ref> but behave as Lewis acids.{{sfn|Aigueperse et al.|2000|p=423}}{{sfn|Wiberg|Wiberg|Holleman|2001|p=897}} The [[pnictogen]]s form trifluorides that increase in reactivity and basicity with higher molecular weight, although [[nitrogen trifluoride]] resists hydrolysis and is not basic.<ref name="PnFx">{{harvnb|Raghavan|1998|pp=[https://books.google.com/books?id=pBiS0jc-kWIC&pg=PA164 164–165]}}.</ref> The pentafluorides of phosphorus, arsenic, and antimony are more reactive than their respective trifluorides, with [[antimony pentafluoride]] the strongest neutral Lewis acid known, only behind [[gold pentafluoride]].{{sfn|Greenwood|Earnshaw|1998|pp=561–563}}<ref name="Arsenic">{{harvnb|Godfrey et al.|1998|p=[https://books.google.com/books?id=vVhpurkfeN4C&pg=PA98 98]}}.</ref>{{sfn|Aigueperse et al.|2000|p=432}}

[[Chalcogen]]s have diverse fluorides: unstable difluorides have been reported for oxygen (the only known compound with oxygen in an oxidation state of +2), sulfur, and selenium; tetrafluorides and hexafluorides exist for sulfur, selenium, and tellurium. The latter are stabilized by more fluorine atoms and lighter central atoms, so [[sulfur hexafluoride]] is especially inert.<ref name="(O,S,Se,Te)Fn">{{harvnb|Murthy|Mehdi Ali|Ashok|1995|pp=[https://books.google.com/books?id=0xl17YU8WzQC&pg=PA180 180–182], [https://books.google.com/books?id=0xl17YU8WzQC&pg=PA206 206–208]}}.</ref>{{sfn|Greenwood|Earnshaw|1998|pp=638–640, 683–689, 767–778}} Chlorine, bromine, and iodine can each form mono-, tri-, and pentafluorides, but only [[iodine heptafluoride]] has been characterized among possible [[interhalogen]] heptafluorides.{{sfn|Wiberg|Wiberg|Holleman|2001|pp=435–436}} Many of them are powerful sources of fluorine atoms, and industrial applications using chlorine trifluoride require precautions similar to those using fluorine.{{sfn|Greenwood|Earnshaw|1998|pp=828–830}}<ref name="Patnaik">{{harvnb|Patnaik|2007|pp=[https://books.google.com/books?id=-CRRJBVv5d0C&pg=PA478 478–479]}}.</ref>

===Noble gases===
{{main|Noble gas compound}}
[[File:Xenon tetrafluoride crop.gif|thumb|left|upright|alt=Black-and-white photo showing transparent crystals in a dish|These xenon tetrafluoride crystals were photographed in 1962. The compound's synthesis, as with xenon hexafluoroplatinate, surprised many chemists.{{sfn|Moeller|Bailar|Kleinberg|1980|p=236}}]]
[[Noble gas]]es, having complete electron shells, defied reaction with other elements until 1962 when [[Neil Bartlett (chemist)|Neil Bartlett]] reported synthesis of [[xenon hexafluoroplatinate]];{{sfn|Wiberg|Wiberg|Holleman|2001|pp=392–393}} [[xenon difluoride]], [[xenon tetrafluoride|tetrafluoride]], [[xenon hexafluoride|hexafluoride]], and multiple oxyfluorides have been isolated since then.{{sfn|Wiberg|Wiberg|Holleman|2001|p=395–397, 400}} Among other noble gases, krypton forms a [[krypton difluoride|difluoride]],{{sfn|Lewars|2008|p=68}} and radon and fluorine generate a solid suspected to be [[radon difluoride]].<ref name="astatineradon">{{harvnb|Pitzer|1993|p=[https://books.google.com/books?id=CdUVspeIJAYC&pg=PA111 111]}}.</ref>{{sfn|Lewars|2008|p=67}} Binary fluorides of lighter noble gases are exceptionally unstable: argon and hydrogen fluoride combine under extreme conditions to give [[argon fluorohydride]].<ref name="Khriachtchev et al. 2000" /> Helium has no long-lived fluorides,<ref name="HeNeCompounds">{{harvnb|Bihary|Chaban|Gerber|2002}}.</ref> and no neon fluoride has ever been observed;{{sfn|Lewars|2008|p=71}} helium fluorohydride has been detected for milliseconds at high pressures and low temperatures.<ref name="HeNeCompounds" />

===Organic compounds===
[[File:FluorocarbonCrabFish.JPG|thumb|alt=Beaker with two layers of liquid, goldfish and crab in top, coin sunk in the bottom|[[miscibility|Immiscible]] layers of colored water (top) and much denser [[perfluoroheptane]] (bottom) in a beaker; a goldfish and crab cannot penetrate the boundary; [[Quarter (United States coin)|quarters]] rest at the bottom.]]
{{main|Organofluorine chemistry}}

[[File:Nafion2.svg|upright=1.35|thumb|alt=Skeletal chemical formula|Chemical structure of [[Nafion]], a fluoropolymer used in fuel cells and many other applications{{sfn|Hoogers|2002|pp=4–12}}]]
The [[carbon–fluorine bond]] is [[organic chemistry]]'s strongest,<ref name="hagan">{{harvnb|O'Hagan|2008}}.</ref> and gives stability to organofluorines.{{sfn|Siegemund et al.|2005|p=444}} It is almost non-existent in nature, but is used in artificial compounds. Research in this area is usually driven by commercial applications;<ref>{{Harvnb|Sandford|2000|p=455}}.</ref> the compounds involved are diverse and reflect the complexity inherent in organic chemistry.<ref name="Jstg" />

====Discrete molecules====
{{main|Fluorocarbon|Perfluorinated compound}}
The substitution of hydrogen atoms in an [[alkane]] by progressively more fluorine atoms gradually alters several properties: melting and boiling points are lowered, density increases, solubility in hydrocarbons decreases and overall stability increases. [[Perfluorocarbon]]s,{{refn|''Perfluorocarbon'' and ''fluorocarbon'' are [[IUPAC]] synonyms for molecules containing 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.|group="note"}} in which all hydrogen atoms are substituted, are insoluble in most organic solvents, reacting at ambient conditions only with sodium in liquid ammonia.{{sfn|Siegemund et al.|2005|pp=451–452}}

The term ''[[perfluorinated compound]]'' is used for what would otherwise be a perfluorocarbon if not for the presence of a [[functional group]],<ref>{{harvnb|Barbee|McCormack|Vartanian|2000|p=[https://books.google.com/books?id=EdzO0aUxDJkC&pg=PA116 116]}}.</ref>{{refn|This terminology is imprecise, and ''perfluorinated substance'' is also used.<ref>{{harvnb|Posner et al.|2013|pp=[https://books.google.com/books?id=dthjVAGP-XIC&dq=%22Perfluorinated+compounds,+hydrocarbons+that+are+fully+fluorinated+except+for+one+functional+group%22&pg=PA187 187–190]}}.</ref>|group="note"}} often a [[carboxylic acid]]. These compounds share many properties with perfluorocarbons such as stability and [[hydrophobicity]],<ref name="fluorotelomer" /> while the functional group augments their reactivity, enabling them to adhere to surfaces or act as [[surfactant]]s.<ref name="Salager2002">{{Harvnb|Salager|2002|p=45}}.</ref> [[Fluorosurfactant]]s, in particular, can lower the [[surface tension]] of water more than their hydrocarbon-based analogues. [[Fluorotelomer]]s, which have some unfluorinated carbon atoms near the functional group, are also regarded as perfluorinated.<ref name="fluorotelomer">{{harvnb|Posner|2011|p=27}}.</ref>

====Polymers====
Polymers exhibit the same stability increases afforded by fluorine substitution (for hydrogen) in discrete molecules; their melting points generally increase too.{{sfn|Carlson|Schmiegel|2000|p=3}} [[Polytetrafluoroethylene]] (PTFE), the simplest fluoropolymer and perfluoro analogue of [[polyethylene]] with [[structural unit]] –{{chem|C|F|2}}–, demonstrates this change as expected, but its very high melting point makes it difficult to mold.{{sfn|Carlson|Schmiegel|2000|pp=3–4}} Various PTFE derivatives are less temperature-tolerant but easier to mold: [[fluorinated ethylene propylene]] replaces some fluorine atoms with [[trifluoromethyl]] groups, [[perfluoroalkoxy alkane]]s do the same with [[trifluoromethoxy]] groups,{{sfn|Carlson|Schmiegel|2000|pp=3–4}} and [[Nafion]] contains perfluoroether side chains capped with [[sulfonic acid]] groups.<ref>{{harvnb|Rhoades|2008|p=[https://books.google.com/books?id=k3AoGn0upf0C&pg=PA2 2]}}.</ref><ref>{{harvnb|Okada et al.|1998}}.</ref> 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.{{sfn|Carlson|Schmiegel|2000|p=4}}

==Production==
Elemental fluorine and virtually all fluorine compounds are produced from [[hydrogen fluoride]] or its aqueous solution, [[hydrofluoric acid]]. Hydrogen fluoride is produced in [[kiln]]s by the [[endothermic reaction]] of [[fluorite]] (CaF<sub>2</sub>) with sulfuric acid:{{sfn|Aigueperse et al.|2000}}

:CaF<sub>2</sub>  +  H<sub>2</sub>SO<sub>4</sub>  →  2 HF(g)  +  CaSO<sub>4</sub>

The gaseous HF can then be absorbed in water or liquefied.<ref>{{cite book |last1=Norris Shreve |last2=Joseph Brink, Jr. |title=Chemical Process Industries |date=1977 |isbn=0070571457 |page=321 |publisher=McGraw-Hill |edition=4}}</ref>

About 20% of manufactured HF is a byproduct of fertilizer production, which produces [[hexafluorosilicic acid]] (H<sub>2</sub>SiF<sub>6</sub>), which can be degraded to release HF thermally and by hydrolysis:
:H<sub>2</sub>SiF<sub>6</sub>  →  2 HF  +  SiF<sub>4</sub>
:SiF<sub>4</sub>  +  2 H<sub>2</sub>O   →  4 HF  +  SiO<sub>2</sub>

===Industrial routes to F<sub>2</sub>===
[[File:Fluorine cell room.jpg|thumb|Industrial fluorine cells at [[City of Preston, Lancashire|Preston]]|alt=A machine room]]
Moissan's method is used to produce industrial quantities of fluorine, via the electrolysis of a [[potassium bifluoride]]/[[hydrogen fluoride]] mixture: hydrogen ions are reduced at a steel container [[cathode]] and fluoride ions are oxidized at a carbon block [[anode]], under 8–12&nbsp;volts, to generate hydrogen and fluorine gas respectively.<ref name="AcountF" />{{sfn|Jaccaud et al.|2000|p=386}} Temperatures are elevated, KF•2HF melting at {{convert|70|Celsius}} and being electrolyzed at {{convert|70|–|130|Celsius}}. KF, which acts to provide electrical conductivity, is essential since pure HF cannot be electrolyzed because it is virtually non-conductive.<ref name="ModFcchem" />{{sfn|Jaccaud et al.|2000|pp=384–285}}{{sfn|Greenwood|Earnshaw|1998|pp=796–797}} Fluorine can be stored in steel cylinders that have passivated interiors, at temperatures below {{convert|200|Celsius}}; otherwise nickel can be used.<ref name="ModFcchem" />{{sfn|Jaccaud et al.|2000|pp=384–385}} Regulator valves and pipework are made of nickel, the latter possibly using [[Monel]] instead.{{sfn|Jaccaud et al.|2000|pp=390–391}} Frequent passivation, along with the strict exclusion of water and greases, must be undertaken. In the laboratory, glassware may carry fluorine gas under low pressure and anhydrous conditions;{{sfn|Jaccaud et al.|2000|pp=390–391}} some sources instead recommend nickel-Monel-PTFE systems.<ref name="Shriver Inorganic">{{harvnb|Shriver|Atkins|2010|p=427}}.</ref>

===Laboratory routes===
While preparing for a 1986 conference to celebrate 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 triggers oxidation instead. He devised a method which evolves fluorine at high yield and atmospheric pressure:<ref name="'year'">{{harvnb|Christe|1986}}.</ref>

:2 [[Potassium permanganate|KMnO<sub>4</sub>]] + 2 [[Potassium fluoride|KF]] + 10 [[Hydrogen fluoride|HF]] + 3 [[Hydrogen peroxide|H<sub>2</sub>O<sub>2</sub>]] → 2 K<sub>2</sub>MnF<sub>6</sub> + 8 [[Water|H<sub>2</sub>O]] + 3 [[Molecular oxygen|O<sub>2</sub>]]↑
:2 K<sub>2</sub>MnF<sub>6</sub> + 4 [[Antimony pentafluoride|SbF<sub>5</sub>]] → 4 KSbF<sub>6</sub> + 2 [[Manganese(III) fluoride|MnF<sub>3</sub>]] + F<sub>2</sub>↑

Christe later commented that the reactants "had been known for more than 100 years and even Moissan could have come up with this scheme."<ref>{{Harvnb|Christe Research Group|n.d.}}</ref> As late as 2008, some references still asserted that fluorine was too reactive for any chemical isolation.<ref>{{harvnb|Carey|2008|page=173}}.</ref>

==Industrial applications==
{{Main|Fluorochemical industry}}
Fluorite mining, which supplies most global fluorine, peaked in 1989 when 5.6&nbsp;million [[metric ton]]s of ore were extracted. Chlorofluorocarbon restrictions lowered this to 3.6&nbsp;million tons in 1994; production has since been increasing. Around 4.5&nbsp;million tons of ore and revenue of [[United States dollar|US$]]550&nbsp;million were generated in 2003; later reports estimated 2011 global fluorochemical sales at $15&nbsp;billion and predicted 2016–18 production figures of 3.5 to 5.9&nbsp;million tons, and revenue of at least $20&nbsp;billion.<ref name="ModFcchem">{{harvnb|Kirsch|2004|pp=3–10}}.</ref><ref>{{Harvnb|Miller|2003b}}.</ref><ref>{{Harvnb|PRWeb|2012}}.</ref><ref>{{Harvnb|Bombourg|2012}}.</ref><ref>{{Harvnb|TMR|2013}}.</ref> [[Froth flotation]] separates mined fluorite into two main metallurgical grades of equal proportion: 60–85% pure metspar is almost all used in iron smelting whereas 97%+ pure acidspar is mainly converted to the key industrial [[chemical intermediate|intermediate]] hydrogen fluoride.<ref name="AcountF" /><ref name="ModFcchem" /><ref name="SME mining 2006">{{harvnb|Fulton|Miller|2006|p=[https://books.google.com/books?id=zNicdkuulE4C&pg=PA471 471]}}.</ref>
<imagemap>
Image:The fluorine economy.svg|thumb|675px|center|Clickable diagram of the fluorochemical industry according to mass flows
rect 9 6 81 34 [[Fluorite]]
rect 9 172 81 199 [[Fluorapatite]]
rect 142 5 244 34 [[Hydrogen fluoride]]
rect 142 65 245 97 [[Smelting|Metal smelting]]
rect 142 121 244 154 [[Glass production]]
rect 309 5 411 33 [[Organofluorine|Fluorocarbons]]
rect 310 63 413 92 [[Sodium hexafluoroaluminate]]
rect 311 121 414 154 [[Pickling (metal)]]
rect 310 171 412 200 [[Hexafluorosilicic acid|Fluorosilicic acid]]
rect 309 211 412 243 [[Cracking (chemistry)|Alkane cracking]]
rect 483 6 585 34 [[Hydrofluorocarbon]]s
rect 484 47 585 76 [[Chlorofluorocarbon#Development of alternatives for CFCs|Hydrochlorofluorocarbons]]
rect 483 88 586 116 [[Chlorofluorocarbon]]
rect 483 128 585 160 [[Polytetrafluoroethylene|Teflon]]
rect 484 170 586 200 [[Water fluoridation]]
rect 483 210 586 238 [[Enriched uranium|Uranium enrichment]]
rect 484 258 586 287 [[Sulfur hexafluoride]]
rect 484 297 585 357 [[Tungsten hexafluoride]]
rect 28 246 177 293 [[Phosphogypsum]]
desc bottom-left
</imagemap>

[[File:SF6 current transformer TGFM-110 Russia.jpg|thumb|alt=Minaret-like electrical devices with wires around them, thicker at the bottom|{{chem|SF|6}} current transformers at a Russian railway]]
{{see also|Industrial gas}}
At least 17,000&nbsp;metric&nbsp;tons of fluorine are produced each year. It costs only $5–8 per kilogram as uranium or sulfur hexafluoride, but many times more as an element because of handling challenges. Most processes using free fluorine in large amounts employ ''in situ'' generation under [[vertical integration]].{{sfn|Jaccaud et al.|2000|p=392}}

The largest application of fluorine gas, consuming up to 7,000&nbsp;metric&nbsp;tons annually, is in the preparation of {{chem|UF|6}} for the [[nuclear fuel cycle]]. Fluorine is used to fluorinate [[uranium tetrafluoride]], itself formed from uranium dioxide and hydrofluoric acid.{{sfn|Jaccaud et al.|2000|p=392}} Fluorine is monoisotopic, so any mass differences between {{chem|UF|6}} molecules are due to the presence of {{chem|235|U}} or {{chem|238|U}}, enabling uranium enrichment via gaseous diffusion or [[gas centrifuge]].{{sfn|Jaccaud et al.|2000|p=382}}<ref name="AcountF" /> About 6,000&nbsp;metric&nbsp;tons per year go into producing the inert [[dielectric]] {{chem|SF|6}} for high-voltage transformers and circuit breakers, eliminating the need for hazardous [[polychlorinated biphenyl]]s associated with {{nowrap|oil-filled}} devices.{{sfn|Aigueperse et al.|2000|p=430}} Several fluorine compounds are used in electronics: rhenium and tungsten hexafluoride in [[chemical vapor deposition]], [[tetrafluoromethane]] in [[plasma etching]]{{sfn|Jaccaud et al.|2000|pp=391–392}}<ref>{{harvnb|El-Kareh|1994|p=[https://books.google.com/books?id=Wg1SnT6Cb0QC&pg=PA317 317]}}.</ref><ref>{{harvnb|Arana et al.|2007}}.</ref> and [[nitrogen trifluoride]] in cleaning equipment.<ref name="AcountF" /> Fluorine is also used in the synthesis of organic fluorides, but its reactivity often necessitates conversion first to the gentler {{chem|Cl|F|3}}, {{chem|Br|F|3}}, or {{chem|I|F|5}}, which together allow calibrated fluorination. Fluorinated pharmaceuticals use [[sulfur tetrafluoride]] instead.<ref name="AcountF" />

===Inorganic fluorides===
[[File:Hall-Heroult cell schematic.svg|thumb|upright=1.35|Aluminium extraction depends critically on cryolite]]
As with other iron alloys, around {{convert|3|kg|lb|abbr=on}} metspar is added to each metric ton of steel; the fluoride ions lower its melting point and [[viscosity]].<ref name="AcountF" /><ref name="USGS">{{Harvnb|Miller|2003a}}.</ref> Alongside its role as an additive in materials like enamels and welding rod coats, most acidspar is reacted with sulfuric acid to form hydrofluoric acid, which is used in steel [[Pickling (metal)|pickling]], glass etching and [[cracking (chemistry)|alkane cracking]].<ref name="AcountF" /> One-third of HF goes into synthesizing [[sodium hexafluoroaluminate|cryolite]] and [[aluminium trifluoride]], both fluxes in the [[Hall–Héroult process]] for aluminium extraction; replenishment is necessitated by their occasional reactions with the smelting apparatus. Each metric ton of aluminium requires about {{convert|23|kg|lb|abbr=on}} of flux.<ref name="AcountF" /><ref>{{Harvnb|Energetics, Inc.|1997|pp=41, 50}}.</ref> Fluorosilicates consume the second largest portion, with [[sodium fluorosilicate]] used in water fluoridation and laundry effluent treatment, and as an intermediate en route to cryolite and silicon tetrafluoride.{{sfn|Aigueperse et al.|2000|p=428}} Other important inorganic fluorides include those of [[cobalt difluoride|cobalt]], [[nickel difluoride|nickel]], and [[ammonium fluoride|ammonium]].<ref name="AcountF" />{{sfn|Aigueperse et al.|2000|pp=420–422}}<ref>{{harvnb|Willey|2007|p=[https://books.google.com/books?id=PrKemPXyXNIC&pg=PA113 113]}}.</ref>

===Organic fluorides===
Organofluorides consume over 20% of mined fluorite and over 40% of hydrofluoric acid, with [[#Refrigerant gases|refrigerant gases]] dominating and [[#Fluoropolymers|fluoropolymers]] increasing their market share.<ref name="AcountF" /><ref name="GIA f-chem report description">{{Harvnb|PRWeb|2010}}.</ref> [[#Surfactants|Surfactants]] are a minor application but generate over $1&nbsp;billion in annual revenue.<ref name="fluorosurfactant market">{{harvnb|Renner|2006}}.</ref> Due to the danger from direct hydrocarbon–fluorine reactions above {{convert|-150|°C|°F|abbr=on}}, industrial fluorocarbon production is indirect, mostly through [[Finkelstein reaction|halogen exchange reactions]] such as [[Swarts fluorination]], in which chlorocarbon chlorines are substituted for fluorines by hydrogen fluoride under catalysts. [[Electrochemical fluorination]] subjects hydrocarbons to electrolysis in hydrogen fluoride, and the [[Fowler process]] treats them with solid fluorine carriers like [[cobalt trifluoride]].<ref name="Jstg" /><ref>{{Harvnb|Green|Slinn|Simpson|Woytek|1994|pp=[https://books.google.com/books?id=TRMRDXHWX7EC&pg=PA91 91–93]}}.</ref>

====Refrigerant gases====
{{See also|Refrigerant}}
Halogenated refrigerants, termed Freons in informal contexts,{{refn|This DuPont trademark is sometimes further misused for CFCs, HFCs, or HCFCs.||group="note"}} are identified by [[Refrigerant#Classification of refrigerants|R-numbers]] that denote the amount of fluorine, chlorine, carbon, and hydrogen present.<ref name="AcountF" /><ref>{{Harvnb|DuPont|2013b}}.</ref> [[Chlorofluorocarbons]] (CFCs) like [[trichlorofluoromethane|R-11]], [[dichlorodifluoromethane|R-12]], and [[1,2-dichlorotetrafluoroethane|R-114]] once dominated organofluorines, peaking in production in the 1980s. Used for air conditioning systems, propellants and solvents, their production was below one-tenth of this peak by the early 2000s, after widespread international prohibition.<ref name="AcountF" /> Hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) were designed as replacements; their synthesis consumes more than 90% of the fluorine in the organic industry. Important HCFCs include R-22, [[chlorodifluoromethane]], and [[1,1-Dichloro-1-fluoroethane|R-141b]]. The main HFC is [[1,1,1,2-Tetrafluoroethane|R-134a]]<ref name="AcountF" /> with a new type of molecule [[2,3,3,3-Tetrafluoropropene|HFO-1234yf]], a [[Hydrofluoroolefin]] (HFO) coming to prominence owing to its [[global warming potential]] of less than 1% that of HFC-134a.<ref name="HFO">{{harvnb|Walter|2013}}.</ref>

====Polymers====
[[File:A water droplet DWR-coated surface2 edit1.jpg|thumb|right|alt=Shiny spherical drop of water on blue cloth|Fluorosurfactant-treated fabrics are often [[hydrophobe|hydrophobic]].]]
{{Main|Fluoropolymer}}
About 180,000&nbsp;metric&nbsp;tons of fluoropolymers were produced in 2006 and 2007, generating over $3.5&nbsp;billion revenue per year.<ref name="yearly FP production">{{harvnb|Buznik|2009}}.</ref> The global market was estimated at just under $6&nbsp;billion in 2011.<ref>{{Harvnb|PRWeb|2013}}.</ref> Fluoropolymers can only be formed by [[polymerization|polymerizing]] free radicals.{{sfn|Carlson|Schmiegel|2000|p=3}}

Polytetrafluoroethylene (PTFE), sometimes called by its DuPont name Teflon,<ref name="encyc of structure" /> represents 60–80% by mass of the world's fluoropolymer production.<ref name="yearly FP production" /> The largest application is in [[electrical insulation]] since 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 in PFTE-coated [[fabric structure|fiberglass cloth]] for stadium roofs. The major consumer application is for [[non-stick surface|non-stick cookware]].<ref name="encyc of structure">{{Harvnb|Martin|2007|pp=[https://books.google.com/books?id=xv420pEC2qMC&pg=PA189 187–194]}}.</ref> Jerked PTFE film becomes expanded PTFE (ePTFE), a fine-pored [[synthetic membrane|membrane]] sometimes referred to by the brand name [[Gore-Tex]] and used for rainwear, [[personal protective equipment|protective apparel]], and [[filtration|filters]]; [[PTFE fiber|ePTFE fibers]] may be made into [[seal (mechanical)|seals]] and [[Dust collector|dust filters]].<ref name="encyc of structure" /> Other fluoropolymers, including [[fluorinated ethylene propylene]], mimic PTFE's properties and can substitute for it; they are more moldable, but also more costly and have lower thermal stability. Films from two different fluoropolymers replace glass in solar cells.<ref name="encyc of structure" /><ref name="DeBergalis">{{harvnb|DeBergalis|2004}}.</ref>

The chemically resistant (but expensive) fluorinated [[ionomer]]s are used as electrochemical cell membranes, of which the first and most prominent example is [[Nafion]]. Developed in the 1960s, it was initially deployed as fuel cell material in spacecraft and then replaced mercury-based [[chloralkali process]] cells. Recently, the fuel cell application has reemerged with efforts to install [[proton exchange membrane]] fuel cells into automobiles.<ref name="Fluorinated Ionomers">{{harvnb|Grot|2011|pp=[https://books.google.com/books?id=E8H1Hwd5GXUC&pg=PA6 1–10]}}.</ref><ref name="Functional Materials">{{harvnb|Ramkumar|2012|p=[https://books.google.com/books?id=ep7U5G4O3mQC&pg=PA567 567]}}.</ref><ref name="Chlor-alkali and chlorate technology">{{harvnb|Burney|1999|p=[https://books.google.com/books?id=rtm_bjQd8SUC&pg=PA111 111]}}.</ref> [[Fluoroelastomer]]s such as [[Viton]] are [[crosslink]]ed fluoropolymer mixtures mainly used in [[O-ring]]s;<ref name="encyc of structure" /> [[perfluorobutane]] (C<sub>4</sub>F<sub>10</sub>) is used as a fire-extinguishing agent.{{sfn|Slye|2012|p=10}}

====Surfactants====
{{Main|Fluorinated surfactant|Durable water repellent}}
Fluorosurfactants are small organofluorine molecules used for repelling water and stains. Although expensive (comparable to pharmaceuticals at $200–2000 per kilogram), they yielded over $1&nbsp;billion in annual revenues by 2006; [[Scotchgard]] alone generated over $300&nbsp;million in 2000.<ref name="fluorosurfactant market" /><ref name="Fluorinated surfactants">{{Harvnb|Kissa|2001|pp=[https://books.google.com/books?id=iAmE8v3bFnUC&pg=PA522 516–551]}}.</ref><ref name="Ullman fibers">{{Harvnb|Ullmann|2008|pp=538, 543–547}}.</ref> Fluorosurfactants are a minority in the overall surfactant market, most of which is taken up by much cheaper hydrocarbon-based products. Applications in [[surfactants in paint|paints]] are burdened by [[compounding]] costs; this use was valued at only $100&nbsp;million in 2006.<ref name="fluorosurfactant market" />

====Agrichemicals====
About 30% of [[agrichemical]]s contain fluorine,<ref>{{Harvnb|ICIS|2006}}.</ref> most of them [[herbicide]]s and [[fungicide]]s with a few [[plant hormone|crop regulators]]. Fluorine substitution, usually of a single atom or at most a [[trifluoromethyl]] group, is a robust modification with effects analogous to fluorinated pharmaceuticals: increased biological stay time, membrane crossing, and altering of molecular recognition.<ref name="Theodoridis" /> [[Trifluralin]] is a prominent example, with large-scale use in the U.S. as a weedkiller,<ref name="Theodoridis">{{harvnb|Theodoridis|2006}}.</ref><ref name="trifluralin">{{Harvnb|EPA|1996}}.</ref> but it is a suspected carcinogen and has been banned in many European countries.<ref>{{Harvnb|DG Environment|2007}}.</ref> [[Sodium monofluoroacetate]] (1080) is a mammalian poison in which one [[sodium acetate]] hydrogen is replaced with fluorine; it disrupts cell metabolism by replacing acetate in the [[citric acid cycle]]. First synthesized in the late 19th century, it was recognized as an insecticide in the early 20th century, and was later deployed in its current use. New Zealand, the largest consumer of 1080, uses it to protect [[Kiwi (bird)|kiwi]]s from the invasive Australian [[common brushtail possum]].<ref name="Beasley">{{Harvnb|Beasley|2002}}.</ref> Europe and the U.S. have banned 1080.{{sfn|Proudfoot|Bradberry|Vale|2006}}{{sfn|Eisler|1995}}{{refn|American sheep and cattle collars may use 1080 against predators like coyotes.||group="note"}}

==Medicinal applications==

===Dental care===
[[File:US Navy 090526-F-1333S-023 A service member embarked aboard the Military Sealift Command hospital ship USNS Comfort (T-AH 20) gives a Fluoride treatment to a patient during a Continuing Promise 2009 medical civil service projec.jpg|thumb|alt=Man holding plastic tray with brown material in it and sticking a small stick into a boy's open mouth|Topical fluoride treatment in Panama]]
{{Main|Fluoride therapy|Water fluoridation|Water fluoridation controversy}}
Population studies from the mid-20th century onwards show [[topical medication|topical]] fluoride reduces [[dental caries]]. This was first attributed to the conversion of tooth enamel [[hydroxyapatite]] into the more durable fluorapatite, but studies on pre-fluoridated teeth refuted this hypothesis, and current theories involve fluoride aiding enamel growth in small caries.<ref name="Pizzo">{{harvnb|Pizzo et al.|2007}}.</ref> After studies of children in areas where fluoride was naturally present in drinking water, controlled [[water supply|public water supply]] fluoridation to fight tooth decay<ref>{{Harvnb|CDC|2001}}.</ref> began in the 1940s and is now applied to water supplying 6 percent of the global population, including two-thirds of Americans.<ref name="Ripa">{{harvnb|Ripa|1993}}.</ref><ref name="Cheng Chalmers Sheldon 2007">{{harvnb|Cheng|Chalmers|Sheldon|2007}}.</ref> Reviews of the scholarly literature in 2000 and 2007 associated water fluoridation with a significant reduction of tooth decay in children.<ref name="NHMRC">{{harvnb|NHMRC|2007}}; see {{harvnb|Yeung|2008}} for a summary.</ref> Despite such endorsements and evidence of no adverse effects other than mostly benign [[dental fluorosis]],<ref>{{harvnb|Marya|2011|p=[https://books.google.com/books?id=NKYy0w1OjJoC&pg=PA343 343]}}.</ref> [[water fluoridation controversy|opposition]] still exists on ethical and safety grounds.<ref name="Cheng Chalmers Sheldon 2007" /><ref>{{harvnb|Armfield|2007}}.</ref> The benefits of fluoridation have lessened, possibly due to other fluoride sources, but are still measurable in low-income groups.<ref name="Dental caries" /> [[Sodium monofluorophosphate]] and sometimes sodium or [[tin(II) fluoride]] are often found in fluoride [[toothpaste]]s, first introduced in the U.S. in 1955 and now ubiquitous in developed countries, alongside fluoridated mouthwashes, gels, foams, and varnishes.<ref name="Dental caries">{{harvnb|Baelum|Sheiham|Burt|2008|p=[https://books.google.com/books?id=fZfXWhSmG1UC&pg=PA518 518]}}.</ref><ref name="Crest continuing ed">{{Harvnb|Cracher|2012|p=12}}.</ref>

===Pharmaceuticals===
[[File:Prozac pills.jpg|thumb|alt=Capsules with "Prozac" and "DISTA" visible|[[Fluoxetine]] capsules]]
Twenty percent of modern pharmaceuticals contain fluorine.<ref name="Emsley 2011 178">{{harvnb|Emsley|2011|p=178}}.</ref> One of these, the cholesterol-reducer [[atorvastatin]] (Lipitor), made more revenue than any other drug until it became generic in 2011.<ref name="url_Crain's_2011">{{Harvnb|Johnson|2011}}.</ref> The combination asthma prescription [[Seretide]], a top-ten revenue drug in the mid-2000s, contains two active ingredients, one of which – [[fluticasone]] – is fluorinated.<ref name="Swinson" /> Many drugs are fluorinated to delay inactivation and lengthen dosage periods because the carbon–fluorine bond is very stable.<ref name="Hagmann">{{harvnb|Hagmann|2008}}.</ref> Fluorination also increases [[lipophilicity]] because the bond is more hydrophobic than the [[carbon–hydrogen bond]], and this often helps in cell membrane penetration and hence [[bioavailability]].<ref name="Swinson">{{harvnb|Swinson|2005}}.</ref>

[[Tricyclic antidepressant|Tricyclics]] and other pre-1980s [[antidepressant]]s had several side effects due to their non-selective interference with [[neurotransmitter]]s other than the [[serotonin]] target; the fluorinated [[fluoxetine]] was selective and one of the first to avoid this problem. Many current antidepressants receive this same treatment, including the [[selective serotonin reuptake inhibitor]]s: [[citalopram]], its [[enantiomer]] [[escitalopram]], and [[fluvoxamine]] and [[paroxetine]].<ref>{{Harvnb|Mitchell|2004|pp=[https://books.google.com/books?id=Duy5_XAD-vgC&pg=PA37 37–39]}}.</ref><ref>{{Harvnb|Preskorn|1996|loc=[http://www.preskorn.com/books/ssri_s2.html chap. 2]}}.</ref> [[Quinolones]] are artificial [[broad-spectrum antibiotic]]s that are often fluorinated to enhance their effects. These include [[ciprofloxacin]] and [[levofloxacin]].<ref>{{harvnb|Werner et al.|2011}}.</ref><ref>{{Harvnb|Brody|2012}}.</ref><ref name="pmid17342653">{{harvnb|Nelson et al.|2007}}.</ref><ref>{{harvnb|King|Malone|Lilley|2000}}.</ref> Fluorine also finds use in steroids:<ref>{{harvnb|Parente|2001|p=[https://books.google.com/books?id=fsfnKeFflUoC&pg=PA40 40]}}.</ref> [[fludrocortisone]] is a blood pressure-raising [[mineralocorticoid]], and [[triamcinolone]] and [[dexamethasone]] are strong [[glucocorticoid]]s.<ref name="dontburn">{{harvnb|Raj|Erdine|2012|p=[https://books.google.com/books?id=SQDge93GucMC&pg=PA58 58]}}.</ref> The majority of inhaled [[anesthetic]]s are heavily fluorinated; the prototype [[halothane]] is much more inert and potent than its contemporaries. Later compounds such as the fluorinated [[ether]]s [[sevoflurane]] and [[desflurane]] are better than halothane and are almost insoluble in blood, allowing faster waking times.<ref name="fut">{{harvnb|Filler|Saha|2009}}.</ref><ref name="anestetics">{{harvnb|Bégué|Bonnet-Delpon|2008|pp=[https://books.google.com/books?id=QMVSvZ-R7I0C&pg=PA335 335–336]}}.</ref>

===PET scanning===
{{main|Positron emission tomography}}
[[File:PET-MIPS-anim.gif|thumb|A full-body {{chem|18|F}} PET scan with glucose tagged with radioactive fluorine-18. The normal brain and kidneys take up enough glucose to be imaged. A malignant tumor is seen in the upper abdomen. Radioactive fluorine is seen in urine in the bladder.|alt=Rotating transparent image of a human figure with targeted organs highlighted]]
Fluorine-18 is often found in [[radioactive tracer]]s for positron emission tomography, as its half-life of almost two hours is long enough to allow for its transport from production facilities to imaging centers.<ref name="18F">{{harvnb|Schmitz et al.|2000}}.</ref> The most common tracer is [[fluorodeoxyglucose]]<ref name="18F" /> which, after intravenous injection, is taken up by glucose-requiring tissues such as the brain and most malignant tumors;<ref>{{harvnb|Bustamante|Pedersen|1977}}.</ref> [[X-ray computed tomography|computer-assisted tomography]] can then be used for detailed imaging.<ref>{{harvnb|Alavi|Huang|2007|p=41}}.</ref>

===Oxygen carriers===
{{see also|Blood substitute|Liquid breathing}}
Liquid fluorocarbons can hold large volumes of oxygen or carbon dioxide, more so than blood, and have attracted attention for their possible uses in artificial blood and in liquid breathing.<ref name="Gabrielet">{{harvnb|Gabriel et al.|1996}}.</ref> 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.<ref>{{Harvnb|Sarkar|2008}}.</ref><ref name="Schimmeyer">{{Harvnb|Schimmeyer|2002}}.</ref> One such product, [[Oxycyte]], has been through initial clinical trials.<ref>{{Harvnb|Davis|2006}}.</ref> These substances can aid endurance athletes and are banned from sports; one cyclist's near death in 1998 prompted an investigation into their abuse.<ref>{{Harvnb|Gains|1998}}.</ref><ref>{{Harvnb|Taber|1999}}.</ref> Applications of pure perfluorocarbon liquid breathing (which uses pure perfluorocarbon liquid, not a water emulsion) include assisting burn victims and premature babies with deficient lungs. Partial and complete lung filling have been considered, though only the former has had any significant tests in humans.<ref name="ReferenceA">{{harvnb|Shaffer|Wolfson|Clark|1992|p=102}}.</ref> An Alliance Pharmaceuticals effort reached clinical trials but was abandoned because the results were not better than normal therapies.<ref name="Alliance stock drop">{{harvnb|Kacmarek et al.|2006}}.</ref>

==Biological role==
{{Main|Biological aspects of fluorine}}
Fluorine is not [[dietary mineral|essential]] for humans and other mammals, but small amounts are known to be beneficial for the strengthening of dental enamel (where the formation of fluorapatite makes the enamel more resistant to attack, from acids produced by bacterial fermentation of sugars). Small amounts of fluorine may be beneficial for bone strength, but the latter has not been definitively established.{{sfn|Nielsen|2009}} Both the WHO and the Institute of Medicine of the US National Academies publish recommended daily allowance (RDA) and upper tolerated intake of fluorine, which varies with age and gender.{{sfn|Olivares|Uauy|2004}}{{sfn|Food and Nutrition Board}}

Natural organofluorines have been found in microorganisms, plants{{sfn|Gribble|2002}} and, recently, animals.<ref>{{cite journal |last1=Xiao-Hua |first1=Xu |last2=Guang-Min |first2=Yao |last3=Yan-Ming |first3=Li |last4=Jian-Hua |first4=Lu |last5=Chang-Jiang |first5=Lin |last6=Xin |first6=Wang |last7=Chui-Hua |first7=Kong |title=5-Fluorouracil Derivatives from the Sponge Phakellia fusca |journal=J. Nat. Prod. |date=January 4, 2003 |volume=2 |issue=66 |pages=285–288 |doi=10.1021/np020034f|pmid=12608868 }}</ref> The most common is [[sodium fluoroacetate|fluoroacetate]], which is used as a [[plant defense against herbivory|defense against herbivores]] by at least 40 plants in Africa, Australia and Brazil.{{sfn|Proudfoot|Bradberry|Vale|2006}} Other examples include terminally fluorinated [[fatty acid]]s, [[fluoroacetone]], and [[2-fluorocitrate]].<ref name="Murphy2003">{{harvnb|Murphy|Schaffrath|O'Hagan|2003}}</ref> An enzyme that binds fluorine to carbon – [[adenosyl-fluoride synthase]] – was discovered in bacteria in 2002.{{sfn|O'Hagan et al.|2002}}

==Toxicity==
{{main|Biological aspects of fluorine#Hazards|l1=Fluorine-related hazards}}
Elemental fluorine is highly toxic to living organisms. Its effects in humans start at concentrations lower than [[hydrogen cyanide]]'s 50&nbsp;ppm<ref>{{harvnb|The National Institute for Occupational Safety and Health|1994a}}.</ref> and are similar to those of chlorine:<ref>{{harvnb|The National Institute for Occupational Safety and Health|1994b}}.</ref> significant irritation of the eyes and respiratory system as well as liver and kidney damage occur above 25&nbsp;ppm, which is the [[IDLH|immediately dangerous to life and health]] value for fluorine.<ref name="Keplinger">{{harvnb|Keplinger|Suissa|1968}}.</ref> The eyes and nose are seriously damaged at 100&nbsp;ppm,<ref name="Keplinger" /> and inhalation of 1,000&nbsp;ppm fluorine will cause death in minutes,<ref>{{harvnb|Emsley|2011|p=179}}.</ref> compared to 270&nbsp;ppm for hydrogen cyanide.<ref>{{harvnb|Biller|2007|p=939}}.</ref>
{{clear}}

===Hydrofluoric acid===
{{Chembox
| container_only = yes
| Section7       = {{Chembox Hazards
| ExternalSDS    =
| GHSPictograms = {{GHS03}}{{GHS05}}{{GHS06}}{{GHS07}}{{GHS08}}{{GHS09}}
| GHSSignalWord = Danger
| HPhrases = {{H-phrases|270|314|310+330|}}<ref>{{cite web|url=https://www.airgas.com/msds/001061.pdf|archive-url=https://web.archive.org/web/20150419162827/https://www.airgas.com/msds/001061.pdf|archive-date=19 April 2015|title=Fluorine. Safety data sheet|publisher=Airgas}}</ref>
| NFPA-H = 4
| NFPA-F = 0
| NFPA-R = 4
| NFPA-S = W+OX
| NFPA_ref = 
 }}
}}
[[File:HF burned hands.jpg|thumb|alt=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.<ref name="Eaton">{{Harvnb|Eaton|1997}}.</ref>]]
{{See also|Chemical burn}}
Hydrofluoric acid is the weakest of the [[hydrohalic acid]]s, having a [[Acid dissociation constant|pKa]] of 3.2 at 25&nbsp;°C.<ref>"Inorganic Chemistry" by Gary L. Miessler and Donald A. Tarr, 4th edition, Pearson</ref> Pure hydrogen fluoride is a volatile liquid due to the presence of hydrogen bonding, while the other hydrogen halides are gases. It is able to attack glass, concrete, metals, and organic matter.<ref>"Inorganic Chemistry" by Shriver, Weller, Overton, Rourke and Armstrong, 6th edition, Freeman</ref>

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 leading to [[hypocalcemia]] and possible death through [[cardiac arrhythmia]].<ref name="OSHA">{{harvnb|Blodgett|Suruda|Crouch|2001}}.</ref> Insoluble calcium fluoride formation triggers strong pain<ref name="Goldfrank">{{Harvnb|Hoffman et al.|2007|p=1333}}.</ref> and burns larger than 160&nbsp;cm<sup>2</sup> (25&nbsp;in<sup>2</sup>) can cause serious systemic toxicity.<ref name="RMTHFE">{{Harvnb|HSM|2006}}.</ref>

Exposure may not be evident for eight hours for 50% HF, rising to 24 hours for lower concentrations, and a burn may initially be painless as hydrogen fluoride affects nerve function. If skin has been exposed to HF, damage can be reduced by rinsing it under a jet of water for 10–15 minutes and removing contaminated clothing.<ref>{{harvnb|Fischman|2001|pp=[https://books.google.com/books?id=PyUSgdZUGr4C&pg=PA458 458–459]}}.</ref> [[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.<ref name="pmid2741315">{{harvnb|El Saadi et al.|1989}}.</ref><ref name="pmid17091088">{{harvnb|Roblin et al.|2006}}.</ref><ref name="pmid15461243">{{harvnb|Hultén et al.|2004}}.</ref> Hydrofluoric acid absorption requires further medical treatment; calcium gluconate may be injected or administered intravenously. Using calcium chloride – a common laboratory reagent – in lieu of calcium gluconate is contraindicated, and may lead to severe complications. Excision or amputation of affected parts may be required.<ref name="RMTHFE" /><ref>{{Harvnb|Zorich|1991|pp=[https://books.google.com/books?id=sdB3KjhWzFQC&pg=PA183 182–183]}}.</ref>

===Fluoride ion===

{{See also|Fluoride toxicity}}
Soluble fluorides are moderately toxic: 5–10&nbsp;g sodium fluoride, or 32–64&nbsp;mg fluoride ions per kilogram of body mass, represents a lethal dose for adults.<ref>{{Harvnb|Liteplo et al.|2002|p=100}}.</ref> One-fifth of the lethal dose can cause adverse health effects,<ref name="Nochimson" /> and chronic excess consumption may lead to [[skeletal fluorosis]], which affects millions in Asia and Africa, and, in children, to reduced intelligence.<ref name="Nochimson" /><ref name="pmid19305069">{{harvnb|Reddy|2009}}.</ref> Ingested fluoride forms hydrofluoric acid in the stomach which is easily absorbed by the intestines, where it crosses cell membranes, binds with calcium and interferes with various enzymes, before urinary [[excretion]]. Exposure limits are determined by urine testing of the body's ability to clear fluoride ions.<ref name="Nochimson">{{Harvnb|Shin|Silverberg|2013}}.</ref><ref>{{harvnb|Baez|Baez|Marthaler|2000}}.</ref>

Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluorides.<ref name="Augenstein et al. 1991">{{Harvnb|Augenstein et al.|1991}}.</ref> Most current calls to poison control centers for possible fluoride poisoning come from the ingestion of fluoride-containing toothpaste.<ref name="Nochimson" /> Malfunctioning water fluoridation equipment is another cause: one incident in Alaska affected almost 300 people and killed one person.<ref name="Gessner94">{{harvnb|Gessner et al.|1994}}.</ref> 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 that they do not swallow toothpaste.<ref>{{Harvnb|CDC|2013}}.</ref> One regional study examined a year of pre-teen fluoride poisoning reports totaling 87 cases, including one death from ingesting insecticide. Most had no symptoms, but about 30% had stomach pains.<ref name="Augenstein et al. 1991" /> A larger study across the U.S. had similar findings: 80% of cases involved children under six, and there were few serious cases.<ref>{{Harvnb|Shulman|Wells|1997}}.</ref>

==Environmental concerns==

===Atmosphere===
[[File:Future ozone layer concentrations.gif|thumb|alt=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<ref>{{Harvnb|Beck|Newman|Schindler|Perkins|2011}}.</ref>]]
{{See also|Ozone depletion|global warming}}
The [[Montreal Protocol]], signed in 1987, set strict regulations on chlorofluorocarbons (CFCs) and [[bromofluorocarbon]]s due to their ozone damaging potential (ODP). The high stability which suited them to their original applications also meant that they were not decomposing until they reached higher altitudes, where liberated chlorine and bromine atoms attacked ozone molecules.<ref name="UNEP FAQ Ozone">{{Harvnb|Aucamp|Björn|2010|pp=4–6, 41, 46–47}}.</ref> Even with the ban, and early indications of its efficacy, predictions warned that several generations would pass before full recovery.<ref>{{harvnb|Mitchell Crow|2011}}.</ref><ref>{{Harvnb|Barry|Phillips|2006}}.</ref> With one-tenth the ODP of CFCs, hydrochlorofluorocarbons (HCFCs) are the current replacements,<ref name="EPA OLP Science">{{Harvnb|EPA|2013a}}.</ref> and are themselves scheduled for substitution by 2030–2040 by hydrofluorocarbons (HFCs) with no chlorine and zero ODP.<ref name="HCFC0">{{Harvnb|EPA|2013b}}.</ref> In 2007 this date was brought forward to 2020 for developed countries;<ref name="HCFC Elim speedup">{{harvnb|McCoy|2007}}.</ref> the [[Environmental Protection Agency]] had already prohibited one HCFC's production and capped those of two others in 2003.<ref name="HCFC0" /> Fluorocarbon gases are generally [[greenhouse gas]]es with [[global-warming potential]]s (GWPs) of about 100 to 10,000; sulfur hexafluoride has a value of around 20,000.{{sfn|Forster et al.|2007|pp=212–213}} An outlier is [[HFO-1234yf]] which is a new type of refrigerant called a [[Hydrofluoroolefin]] (HFO) and has attracted global demand due to its GWP of less than 1 compared to 1,430 for the current refrigerant standard [[HFC-134a]].<ref name="HFO" />

===Biopersistence===
[[File:PFOS-3D-vdW.png|left|[[Perfluorooctanesulfonic acid]], a key [[Scotchgard]] component until 2000{{sfn|Schwarcz|2004|p=37}}|thumb]]
{{main|Biological aspects of fluorine#Biopersistence|l1=Biopersistence of fluorinated organics}}
Organofluorines exhibit biopersistence due to the strength of the carbon–fluorine bond. [[Perfluoroalkyl acid]]s (PFAAs), which are sparingly water-soluble owing to their acidic functional groups, are noted [[persistent organic pollutant]]s;<ref>{{Harvnb|Giesy|Kannan|2002}}.</ref> [[perfluorooctanesulfonic acid]] (PFOS) and [[perfluorooctanoic acid]] (PFOA) are most often researched.<ref name="PFOA" /><ref name="PFOA evidence" /><ref>{{Harvnb|EPA|2012}}.</ref> PFAAs have been found in trace quantities worldwide from polar bears to humans, with PFOS and PFOA known to reside in breast milk and the blood of newborn babies. A 2013 review showed a slight correlation between groundwater and soil PFAA levels and human activity; there was no clear pattern of one chemical dominating, and higher amounts of PFOS were correlated to higher amounts of PFOA.<ref name="PFOA" /><ref name="PFOA evidence" /><ref>{{Harvnb|Zareitalabad|Siemens|Hamer|Amelung|2013}}.</ref> In the body, PFAAs bind to proteins such as [[serum albumin]]; they tend to concentrate within humans in the liver and blood before excretion through the kidneys. Dwell time in the body varies greatly by species, with half-lives of days in rodents, and years in humans.<ref name="PFOA">{{harvnb|Steenland|Fletcher|Savitz|2010}}.</ref><ref name="PFOA evidence">{{harvnb|Betts|2007}}.</ref><ref name="tox">{{harvnb|Lau et al.|2007}}.</ref> 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.<ref name="PFOA" /><ref name="PFOA evidence" /><ref name="tox" />

==See also==
{{Portal|Chemistry}}
<!-- Please keep entries in alphabetical order & add a short description [[WP:SEEALSO]] -->
{{div col|colwidth=18em}}
* [[Argon fluoride laser]]
* [[Electrophilic fluorination]]
* [[Fluoride selective electrode]], which measures fluoride concentration
* [[Fluorine absorption dating]]
* [[Fluorous chemistry]], a process used to separate reagents from organic solvents
* [[Krypton fluoride laser]]
* [[Radical fluorination]]
{{div col end}}
<!-- please keep entries in alphabetical order -->

==Notes==
{{reflist|group="note"}}

==Sources==

=== Citations ===
{{Reflist|21em}}

===Indexed references===
{{sfn whitelist |CITEREFJaccaud_et_al.2000 |CITEREFAigueperse_et_al.2000 |CITEREFMeusingerChippendaleFairhurst2012 |CITEREFNUBASE2016 |CITEREFSiegemund_et_al.2005 |CITEREFCarlsonSchmiegel2000}}
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{{Refend}}

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[[Category:Fluorine| ]]
[[Category:Chemical elements]]
[[Category:Halogens]]
[[Category:Reactive nonmetals]]
[[Category:Diatomic nonmetals]]
[[Category:Fluorinating agents]]
[[Category:Oxidizing agents]]
[[Category:Industrial gases]]
[[Category:Gases with color]]