View source for Fluorine

{{Infobox fluorine}}
'''Fluorine''' is the [[chemical element]] with symbol '''F''' and [[atomic number]] 9. The lightest [[halogen]], it has a single stable [[isotope]], fluorine‑19. At [[standard conditions for temperature and pressure|standard pressure and temperature]], the element is a pale yellow gas composed of [[diatomic molecule]]s, F<sub>2</sub>. Fluorine is the most [[electronegative]] element and also the most reactive, requiring great care in handling. [[Compounds of fluorine]] are called fluorides.

In stars, fluorine is rare, for a light element, because it is consumed in fusion reactions. Within Earth's crust, fluorine is the thirteenth‑most abundant element. The main source mineral, [[fluorite]], was first formally described in 1529. The mineral's name derived from the Latin verb ''{{lang|la|fluo}}'', meaning "flow", because it was added to metal ores to lower their melting points. Suggested as a chemical element in 1811, the dangerous element injured many experimenters who tried to isolate it. In 1886, French chemist [[Henri Moissan]] succeeded. His method of [[electrolysis]] remains the industrial production method for fluorine gas. The largest use of elemental fluorine, [[enriched uranium|uranium enrichment]], began during the [[Manhattan Project]].

[[fluorochemical industry|Global fluorochemical]] sales are over US$15 billion per year. Because of the difficulty in making elemental fluorine, ninety-nine percent of commercially used fluorine is never converted to the free element. About half of all mined fluorite is used directly in [[steel-making]]. The other half is converted to [[hydrofluoric acid]], a precursor to other chemicals. The most important synthetic inorganic fluoride is [[cryolite]], a commodity that is critical for [[Hall–Héroult process|aluminium refining]]. Organic fluorides have very high chemical and thermal stability. Their largest market segment is in [[refrigerant]] gases. Even though traditional [[chlorofluorocarbons|CFCs]] are now mostly banned, the replacement molecules still contain fluorine. The predominant [[fluoropolymer]] is [[PTFE|Teflon]], which is used in electrical insulation and cookware.

While a few plants and bacteria synthesize organofluorine poisons, fluorine has no metabolic role in mammals. The [[fluoride|fluoride ion]], when directly applied to teeth, reduces decay. For this reason, it is used in toothpaste and [[water fluoridation]]. A growing fraction of modern pharmaceuticals contain fluorine; [[Lipitor]] and [[Prozac]] are prominent examples.
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==Characteristics==

[[File:Fluorine with significant periodic table neighbors landscape.svg|400px|center]]

===Isotopes===

{{Main|Isotopes of fluorine}}

The [[atomic nucleus|nucleus]] of a fluorine atom contains nine [[proton]]s. One stable [[isotope]] occurs naturally: fluorine-19, which contains ten [[neutrons]].<ref>{{cite web |author=National Nuclear Data Center |title=NuDat 2.1 database – Fluorine-19 |url=http://www.nndc.bnl.gov/nudat2/reCenter.jsp?z=9&n=10 |publisher=[[Brookhaven National Laboratory]] |accessdate=31 September 2005}}</ref> This makes the element both [[monoisotopic element|monoisotopic]] and [[mononuclidic element|mononuclidic]], characteristics that enable its use in [[uranium enrichment]] and in [[nuclear magnetic resonance]].

Seventeen [[radioisotope]]s have been synthesized: [[mass number]]s 14–18 and 20–31.<ref name="isotopes">{{cite web |author=National Nuclear Data Center |title=NuDat 2.1 database |url=http://www.nndc.bnl.gov/nudat2/ |publisher=Brookhaven National Laboratory |accessdate=31 September 2005}}</ref> [[Fluorine-18]] is the most stable, with a [[half-life]] of 109.77&nbsp;minutes. The lightest fluorine isotopes, those with mass numbers of 14–16, decay via [[electron capture]]. Fluorine-17 and -18 undergo [[positron emission|beta plus decay]] (emission of a [[positron]]). All isotopes heavier than fluorine-19 decay by [[beta minus decay|beta minus mode]] (emission of an [[electron]]); some of these also decay by [[neutron emission]].<ref name="isotopes"/>

===Atomic structure===

The neutral fluorine atom has nine electrons, one fewer than [[neon]], arranged in the [[electronic configuration]] [He] 2s<sup>2</sup> 2p<sup>5</sup>.{{sfn|Jaccaud et al.|2005|p=1}} Fluorine's outer electrons are relatively separate from each other. Hence, they do not shield each other from the nucleus, so they experience a high [[effective nuclear charge]]. Fluorine has a small [[Covalent radius of fluorine|covalent radius]], around 60&nbsp;[[picometer]]s, which is similar to [[oxygen]] or neon.{{#tag:ref|Exact comparison of the sizes of fluorine, oxygen and neon atoms is not possible because of conflicting estimates from different sources.||group="note"}}<ref name="CSD">{{cite journal|author= Cordero, Beatriz; Gómez, Verónica; Platero-Prats, Ana E.; Revés, Marc; Echeverría, Jorge; Cremades, Eduard; Barragán, Flavia; Alvarez, Santiago|title=Covalent radii revisited | journal=Dalton Translations|year=2008|pages=2832–2838|doi=10.1039/b801115j|issue=21}}</ref><ref>{{cite doi|10.1002/chem.200901472}}</ref>

[[File:Fluorine shielding.svg|upright=1.8|alt=A bunch of rings showing valence and non-valence electron shells|Structure of the fluorine atom|center|thumb]]

Fluorine is reluctant to [[ionization|ionize]]. Its [[ionization energy]] (the energy required to remove an electron) is higher than that of any element except neon or [[helium]].{{sfn|Dean|1999|p=564}} Instead of ionizing, it exhibits a very strong preference for one more electron to achieve the extremely stable neon‑like electron configuration.{{sfn|Jaccaud et al.|2005|p=1}} Fluorine is the element with the highest [[electronegativity]] (a relative measure of electron attraction by atoms).<ref>{{cite book|url = http://books.google.com/?id=ZOm8L9oCwLMC&pg=PA156|year = 2010| page =156 |title = Principles of chemistry: The molecular Science|last1 = Moore|first1 = John W.|last2 = Stanitski|first2 = Conrad L.|last3 = Jurs|first3 = Peter C.|publisher=Cengage Learning| isbn = 978-0-495-39079-4}}</ref> Its [[electron affinity]] (the energy released by adding an electron) is higher than that of any element except [[chlorine]].{{sfn|Lide|2004|pp=10.137–10.138}}

===Molecular structure===

{{Main|Compounds of fluorine#Difluorine|l1=Difluorine}}

Fluorine forms [[diatomic molecules]], F<sub>2</sub>. The difluorine bond is relatively weak, with a [[bond energy]] much less than that in Cl<sub>2</sub> or Br<sub>2</sub>. It is similar to I<sub>2</sub> or to the easily cleaved oxygen–oxygen bonds of [[peroxide]]s or nitrogen–nitrogen bonds of [[hydrazine]]s.<ref>{{cite book|year = 1996|url=http://books.google.com/books?id=HyuogOtzoaYC&pg=PA230| page =230 |title = Organic chemistry|volume=1|last1 = Macomber|first1 = Roger S.|publisher=University Science Books| isbn = 0-935702-90-3}}</ref>{{sfn|Greenwood|Earnshaw|1998|p=804}}

While an individual fluorine atom has one unpaired electron, molecular fluorine has all the electrons paired. This makes it [[diamagnetic]] (slightly repelled by magnets). In contrast, neighboring element oxygen, with two unpaired electrons per molecule, is [[paramagnetic]] (attracted to magnets).{{sfn|Mackay|Mackay|Henderson|2002|p=72}}<ref name = Magsusc>{{cite journal|journal=Journal of Physical Chemistry A | year = 1999|volume = 103|pages = 2861–2866|first1 = H.|last1 = Cheng|first2 = D.E.|last2 = Fowler|first3 = P.B.|last3 = Henderson|first4 = J.P. |last4 = Hobbs|first5 = M.R.|last5 = Pascaloni|title = On the magnetic susceptibility of fluorine|doi=10.1021/jp9844720|issue=15}}</ref>

===Phases===

{{Main|Phases of fluorine}}

Fluorine is gaseous at room temperature.{{sfn|Jaccaud et al.|2005|p=2}}  Though sometimes cited as yellow-green, pure fluorine is actually a very pale yellow.<ref>{{cite journal|journal=Journal of Fluorine Chemistry | year = 1987|volume = 34|pages = 471–474|first1 = J.|issue = 3–4|last1 = Burdon|first2 = B.|last2 = Emson|first3 = A.J.|last3 = Edwards|title = Is fluorine really yellow?|doi = 10.1016/S0022-1139(00)85188-X}}</ref> The gas has a "pungent" characteristic odor that is noticeable in concentrations as low as 20&nbsp;[[parts per billion|ppb]].{{sfn|Lide|2004|p=4.12}} Fluorine condenses to a bright yellow liquid at −188&nbsp;°C (−307&nbsp;°F),{{sfn|Dean|1999|p=523}} which is near the condensation temperature of oxygen or [[nitrogen]].

Fluorine solidifies at −220&nbsp;°C (−363&nbsp;°F){{sfn|Dean|1999|p=523}} into a [[Cubic crystal system|cubic]] structure called beta-fluorine. This [[phase (matter)|phase]] is transparent and soft with significant disorder of the molecules. With further cooling to −228&nbsp;°C (−378&nbsp;°F), fluorine undergoes a solid–solid [[phase transition]] into a [[monoclinic crystal system|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, shattering samples and blowing out sample holder windows. In general, fluorine's solid state phases are more similar to oxygen's than to those of its fellow halogens.<ref name="phases"/><ref>{{Cite doi|10.1063/1.1711946|noedit}}</ref>

{| style="margin:1em auto 1em auto; text-align:left; width:500px;" cellpadding="0" cellspacing="0"
|- style="vertical-align:top;"
| style="width:300px; padding-right:15px;" | [[File:Phase diagram of fluorine (1975).png|300px|alt=angled lines showing linear pressure temp relations of the lower phase boundaries]]
| style="width:290px; padding-left:15px;" | [[File:Alpha-fluorine-unit-cell-B-3D-vdW.png|300px|alt=A parallelogram-shaped outline with space-filling diatomic molecules (joined circles) arranged in two layers]]
|- style="text-align:center;"
| style="width:300px; padding-right:15px; "|Low-temperature fluorine phases
| style="width:290px; padding-left:15px; "|Alpha-fluorine crystal structure
|}

===Chemical reactivity===

{{Main|Compounds of fluorine#Chemical characteristics of fluorine|l1=Chemical characteristics of fluorine}}

Reactions with elemental fluorine are often sudden or explosive. 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.{{sfn|Jaccaud et al.|2005|p=2}}<ref name="pop mechanics F reactivity">{{cite journal|last=Nelson|first=Eugene W.|title='Bad man' of the elements|journal=Popular Mechanics|year=1947|volume=88|issue=2|pages=106–108, 260|url=http://books.google.com/books?id=1t8DAAAAMBAJ&pg=PA106}}</ref>

{{external media
| width = 240px
| float = right
| video1 = [http://www.youtube.com/watch?v=vtWp45Eewtw Bright flames] during fluorine reactions
| video2 = Fluorine [http://www.youtube.com/watch?v=TLOFaWdPxB0 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 the protective metal fluoride layers formed by [[Passivation (chemistry)|passivation]].{{sfn|Greenwood|Earnshaw|1998|p=804}} The [[alkali metal]]s, such as sodium, react explosively. The [[alkaline earth metal]]s, such as calcium, react somewhat less dramatically. The [[noble metal]]s gold, platinum, palladium, rhodium, and ruthenium react least readily, requiring pure fluorine gas at 300–450&nbsp;°C (575–850&nbsp;°F).{{sfn|Lidin|Molochko|Andreeva|2000|pp=442–455}}

Fluorine reacts explosively with [[hydrogen]] in a manner similar to that of alkali metals.{{sfn|Greenwood|Earnshaw|1998|p=844}} The other halogens react readily with fluorine gas,{{sfn|Jaccaud et al.|2005|p=3}} as does the heavy noble gas [[radon]].<ref>{{cite journal|journal=Journal of the Chemical Society: Chemical communications|year=1975|pages=760b–761|title=Fluorides of radon and element 118|author=Pitzer, Kenneth S.|doi=10.1039/C3975000760B|issue=18}}</ref> The lighter noble gases [[xenon]] and [[krypton]] react with fluorine under special conditions; [[argon]] will react only with [[hydrogen fluoride]].<ref name="Rasanen2000">{{cite journal|last1 = Khriachtchev|first1 = Leonid|first2 = Mika|last2 =Pettersson| first3 = Nino|last3 =Runeberg| first4 = Jan|last4 =Lundell|first5 =Markku|last5 =Räsänen|year = 2000|title = A stable argon compound|journal= Nature|volume = 406|pages = 874–876|doi = 10.1038/35022551|accessdate =29 April 2011|url = http://www.nature.com/nature/journal/v406/n6798/abs/406874a0.html|pmid = 10972285|issue = 6798}}</ref> Oxygen and fluorine gases do not normally react, but they can be combined under special conditions ([[Electric discharge in gases|electric discharge]] and low pressure/temperature). The resultant products can be unstable and tend to separate back into fluorine and oxygen when heated.<ref name="EmeleusSharpe1974">{{cite book|author=Emeleus, H.J.; Sharpe, A.G.|title=Advances in inorganic chemistry and radiochemistry|url=http://books.google.com/books?id=qi0vujzNSz4C&pg=PA111|year=1974|publisher=Academic Press|isbn=978-0-08-057865-1|page=111}}</ref>{{sfn|Wiberg|Wiberg|Holleman|2001|p=457}}<ref name="Clarke">{{cite book|author=Squires, Roy; Clarke, Arthur C. |title=Journal of the Pacific Rocket Society|url=http://books.google.com/books?id=Yo7r_1qJG2gC&pg=PA26-IA83|publisher=Sawyer Publishing|isbn=978-0-9794418-5-1|page=26}}</ref> Nitrogen, with its very stable [[triple bond]]s, requires electric discharge and very high temperatures to react with fluorine.{{sfn|Lidin|Molochko|Andreeva|2000|p=252}} However, ammonia can react explosively with fluorine.<ref name="Anhydrous ammonia MSDS">{{cite web|last=|first=|date=|title=Anhydrous ammonia: (MSDS) Material Safety Data Sheet|url=http://www.tannerind.com/anhydrous-msds.html|publisher=Tanner Industries|accessdate=14 September 2013}}</ref><ref name="Scott I. Morrow">{{Cite doi|10.1021/ja01532a066}}</ref>

[[Carbon]], in graphite or diamond form, is impervious to fluorine gas at room temperature. Above 400°C, graphite reacts with fluorine to make a fluorinated solid, "[[carbon monofluoride]]". At even higher temperatures, gaseous fluorocarbons start to be produced; the reaction can become explosive.<ref name="Kuriakose, A. K.">{{Cite doi|10.1021/j100892a049}}</ref> Carbon dioxide and carbon monoxide react with fluorine at room temperature or a little above. The reactions make mixtures of products, are heat-generating, and are difficult to control.<ref name="Hasegawa, Yasuo">{{Cite doi|10.1016/j.jfluchem.2006.09.002}}</ref> Organic chemicals, such as [[alkane|paraffin]]s, react strongly when exposed to fluorine.<ref>{{cite thesis|author=Lagow, James|title=The reactions of elemental fluorine; a new approach to fluorine chemistry|type=Ph.D. thesis|year=1970|publisher=Rice University}}</ref> Even fully halogenated organic molecules, such as the normally incombustible solvent [[carbon tetrachloride]], can explode.<ref name="Navarrini, Walter">{{Cite doi|10.1016/j.jfluchem.2012.06.006}}</ref>

The other solid [[nonmetal]]s or [[metalloid]]s (boron, silicon, arsenic, sulfur, germanium, phosphorus, selenium, tellurium) burn with a flame in room temperature fluorine.<ref name=Mellor>{{cite book|last=Mellor|first=J.W. |year=1927|title=A Comprehensive Treatise on Inorganic and Theoretical Chemistry Volume I |publisher= Lonmans, Green and Co|isbn=}}</ref> [[Hydrogen sulfide]] and [[sulfur dioxide]] combine readily with fluorine; the latter reaction can be explosive. [[Sulfuric acid]] reacts much more sluggishly.<ref name=Mellor/>

==Origin and occurrence==
{{See also|List of countries by fluorite production}}

{| class="wikitable" style="float:right; margin-left:15px; margin-top:0; font-size:10pt; line-height:11pt;"
|+ style="margin-bottom: 5px;" | Abundance in the Solar System<ref name="Cameron">{{cite journal|journal= Space Science Review | url = http://pubs.giss.nasa.gov/docs/1973/1973_Cameron_1.pdf | year = 1973 |volume = 15 | pages = 121–146 | first1 = A.G.W. | last1 = Cameron | title = Abundance of the elements in the Solar System|bibcode = 1973SSRv...15..121C |doi = 10.1007/BF00172440}}</ref>
! style="padding-right: 10px; padding-left: 5px;" | Atomic <br> number
! style="width:45%;"| Element
! style="padding-right: 5px; padding-left: 10px;" | Relative <br> amount
|-
| style="padding-left: 5px;" | 6
| style="text-align:center;"| Carbon
| style="padding-right:5px; text-align:right;"| 4,800
|-
| style="padding-left: 5px;" | 7
| style="text-align:center;"| Nitrogen
| style="padding-right:5px; text-align:right;"|1,500
|-
| style="padding-left: 5px;" | 8
| style="text-align:center;"| Oxygen
| style="padding-right:5px; text-align:right;"|8,800
|- style="background:#ff9;"
| style="padding-left:5px; "| ''9''
| style="text-align:center; "| ''Fluorine''
| style="padding-right:5px; text-align:right;"|''1''
|-
| style="padding-left: 5px;" | 10
| style="text-align:center;"| Neon
| style="padding-right:5px; text-align:right;"|1,400
|-
| style="padding-left: 5px;" | 11
| style="text-align:center;"| Sodium
| style="padding-right:5px; text-align:right;"|24
|-
| style="padding-left: 5px;" | 12
| style="text-align:center;"| Magnesium
| style="padding-right:5px; text-align:right;"|430
|}

From the perspective of cosmology, fluorine is relatively rare with 400&nbsp;ppb in the universe. Within stars, any fluorine that is [[Stellar nucleosynthesis|created]] is rapidly eliminated through [[nuclear fusion]]: either with hydrogen to form oxygen and helium, or with helium to make neon and hydrogen. The presence of fluorine at all—outside of temporary existence in stars—is somewhat of a mystery because of the need to escape these fluorine-destroying reactions.<ref name=Crosswell/><ref name="Milky Way"/> Three theoretical solutions exist. In [[type II supernova]]e, atoms of neon could be hit by [[neutrino]]s 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]] (a type of red giant) stars, pulses of fusion reactions could allow [[convection]] to lift fluorine out of the inner star. Only the red giant hypothesis has supporting evidence from observations.<ref name=Crosswell>{{cite journal|last1=Croswell|first1=Ken|title=Fluorine: An element–ary mystery| journal=Sky and Telescope| year = 2003|url=http://kencroswell.com/fluorine.html|accessdate=3 May 2011}}</ref><ref name="Milky Way">{{Cite journal|last1=Renda|first1=A.|first2=Y.|last2=Fenner|first3=B.K.|last3=Gibson|first4=A.I. |last4=Karakas|first5=J.C. |last5=Lattanzio|first6=S.|last6=Campbell|first7=A.|last7=Chieffi
|first8=K. |last8=Cunha|first9=V.V. |last9=Smith|displayauthors=9|url=http://cdsweb.cern.ch/record/800212/files/0410580.pdf?version=1|title =On the origin of fluorine in the Milky Way|journal=Monthly Notices of the Royal Astronomical Society|volume =354|pages =575–580|year =2004|doi =10.1111/j.1365-2966.2004.08215.x|bibcode=2004MNRAS.354..575R|issue =2|arxiv = astro-ph/0410580}}</ref>

In space, fluorine commonly combines with hydrogen to form hydrogen fluoride. (This compound has been suggested as a tracer to enable tracking reservoirs of hydrogen in the universe.)<ref>{{cite web|author=Neufeld, David; Bergin, Edwin; Gerin, Maryvonne |year=2010|publisher=[[European Space Agency]]|title=Tracing the Milky Way's hidden reservoirs of gas|url=http://www.esa.int/Our_Activities/Space_Science/Herschel/Tracing_the_Milky_Way_s_hidden_reservoirs_of_gas}}</ref> In addition to HF, monatomic fluorine has been observed in the [[interstellar medium]].<ref>{{cite journal|last1=Snow|first1=T.P., Jr.|last2=York|first2=D.G.|title=The detection of interstellar fluorine in the line of sight toward Delta Scorpii|year = 1981|bibcode=1981ApJ...247L..39S|volume=247|page=L39|journal=Astrophysical Journal|doi=10.1086/183585}}</ref><ref name="Snow 2007">{{cite journal|last=Snow|first=Theodore, P.|coauthors=Destree, Joshua D.; Jensen, Adam G.|title=The abundance of interstellar fluorine and its implications|journal=The Astrophysical Journal|year=2007|volume=655|issue=1|pages=285–298 |doi = 10.1086/510187|arxiv = astro-ph/0611066 |bibcode = 2007ApJ...655..285S }}</ref> Fluorine [[cation]]s have been seen in planetary nebulae and in stars, including our Sun.<ref>{{cite journal|journal= The Astrophysical Journal |last1=Zhang|first1=Y.|last2=Liu|first2=X.-W.|title=Fluorine abundances in planetary nebulae|year = 2005| pages = L61–L63| doi = 10.1086/497113|volume = 631|arxiv = astro-ph/0508339 |bibcode = 2005ApJ...631L..61Z|issue=1}}</ref>

Fluorine is the thirteenth most [[Abundance of elements in Earth's crust|common element in Earth's crust]], comprising between 600 and 700&nbsp;[[parts per million|ppm]] of the crust by mass. Because of its reactivity, it is usually found as fluoridated compounds. Three minerals exist that are industrially relevant sources: fluorite, [[fluorapatite]], and cryolite.{{sfn|Jaccaud et al.|2005|p=4}}{{sfn|Greenwood|Earnshaw|1998|p=795}}

{| cellpadding="3" style="margin:1em auto 1em auto; text-align:center;"
|-
| colspan="3" style="text-align:center;"| '''Major fluorine-containing minerals'''
|-
| [[File:Fluorite-270246.jpg|200px|alt=pink globular mass with crystal facets]]
| [[File:Apatite Canada.jpg|200px|alt=Long prism-like crystal, without luster, at an angle coming out of aggregate-like rock|Solid alpha-fluorine's crystal structure: fluorine molecules lie in shingled layers.]]
| [[File:Ivigtut cryolite edit.jpg|200px|alt=A parallelogram-shaped outline with space-filling diatomic molecules (joined circles) arranged in two layers]]
|- style="text-align:center;"
| Fluorite
| Fluorapatite
| Cryolite
|}

* '''Fluorite''' (CaF<sub>2</sub>), also called fluorspar, is the main source of commercial fluorine. Fluorite is colorful, common, and found worldwide. China supplies more than half of the world's demand; Mexico is second. The United States produced most global fluorite in the early 20th century, but its last mine, in Illinois, shut down in 1995.{{sfn|Greenwood|Earnshaw|1998|p=795}}<ref name="AcountF">{{cite journal|last1 = Villalba|first1 = Gara|last2 = Ayres|first2 = Robert U.|last3 = Schroder|first3 = Hans|title = Accounting for fluorine: production, use, and loss|journal= Journal of Industrial Ecology|volume = 11|pages = 85–101|year = 2008|doi =10.1162/jiec.2007.1075}}</ref><ref name="historical fluorspar statistics">{{cite web|last=Kelly|first=T.D|title=Historical fluorspar statistics|url=http://minerals.usgs.gov/ds/2005/140/fluorspar.pdf||publisher=United States Geological Service|accessdate=25 January 2012}}</ref><ref name="UK fluorspar">{{cite web|last=Lusty|first=P.A.J.|title=The need for indigenous fluorspar production in England|url=https://www.bgs.ac.uk/downloads/start.cfm?id=1328|publisher=British Geological Survey|accessdate=6 October 2013|coauthors=Brown, T.J.; Ward, J; Bloomfeld, S.|year=2008}}</ref><ref name="KGS fluorite terminology"/>}}</ref>

* '''Fluorapatite''' (Ca<sub>5</sub>(PO<sub>4</sub>)<sub>3</sub>F) is mined, along with other [[apatite]]s, in high volumes for its [[phosphate]] content, which is used in fertilizer. 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 United States is there significant recovery. There, the [[hexafluorosilicate]]s produced as byproducts are used to supply water fluoridation.{{sfn|Greenwood|Earnshaw|1998|p=795}}

* '''Cryolite''' (Na<sub>3</sub>AlF<sub>6</sub>) is the least abundant of the three major fluorine-containing minerals but 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.{{sfn|Greenwood|Earnshaw|1998|p=795}}

Several other minerals, such as the gemstone [[topaz]], contain fluoride. Fluoride is not significant in seawater or brines, unlike the other [[halide]]s, because the alkaline earth fluorides precipitate out of water.{{sfn|Greenwood|Earnshaw|1998|p=795}} Commercially insignificant quantities of organofluorines have been observed in volcanic eruptions and in geothermal springs. Their ultimate origin (from biological sources or geological formation) is unclear.<ref name=Gribble>{{cite journal|last1 =Gribble|first1 =Gordon W.|title =Naturally occurring organofluorines|volume =3N|pages =121–136|year =2002|doi =10.1007/10721878_5|journal= The Handbook of Environmental Chemistry|isbn =3-540-42064-9}}</ref>

The possibility of small amounts of gaseous fluorine within crystals has been debated for many years. One form of fluorite, [[antozonite]], has an ozone smell (suggestive of fluorine) when crushed. The mineral also has a dark black color, perhaps from free calcium. 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 [[Crystallographic defect|defects]].<ref name="Elementar">{{cite journal|author=Schmedt, Jörn; Mangst, Martin; Kraus, Florian|title=Elementares Fluor F<sub>2</sub> in der Natur – In-situ-Nachweis und Quantifizierung durch NMR-Spektroskopie|journal=Angewandte Chemie|year=2012|volume=124|issue=31|pages=7968–7971|doi= 10.1002/ange.201203515|language=German}}</ref>

==History==

{{Main|History of fluorine}}

===Early discoveries===

[[File:Book9-25.gif|thumb|right|alt=woodcut image with top showing man at open hearth with tongs and machine bellows to the side. Bottom shows man at water-operated hammer with a sluice nearby for quenching.|Steelmaking illustration, Agricola text]]

The word "fluorine" derives from the Latin stem of the main source mineral, fluorite. The stone was described in 1529 by [[Georgius Agricola]], who related its use as a [[flux (metallurgy)|flux]]—an additive that helps lower the melt temperature during [[smelting]].{{sfn|Greenwood|Earnshaw|1998|p=790}}<ref name="assassinated">{{cite book|url = http://books.google.com/?id=Fl4sdCYrq3cC&pg=PT149|isbn = 978-0-444-52239-9|page = 149|first = Alexander|last = Senning|year = 2007|publisher = Elsevier|title = Elsevier's dictionary of chemoetymology: The whies and whences of chemical nomenclature and terminology}}</ref> 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){{sfn|Greenwood|Earnshaw|1998|p=109}} and then to fluorite.<ref name="KGS fluorite terminology">{{cite book|last=Norwood|first=Charles J.|title=Kentucky geological survey Bulletin 9: Fluorspar deposits of Kentucky|year=1907|publisher=Globe Printing Company|url=http://books.google.com/books?id=CGRMAAAAYAAJ&pg=PA34|coauthors=Fohs, Julius F.|page=52|chapter=Fluorspar and its Occurrence}}</ref><ref>{{cite book|url = http://www.farlang.com/gemstones/agricola-metallica/page_003 |pages = preface, 380–381 (3, 416–417 in linked viewer)|first = Georgius |last=Agricola |title=De Re Metallica|others =trans. Hoover, Herbert Clark; Hoover, Henry Lou|year = 1912}}</ref>

Hydrofluoric acid was known as a glass-etching agent from the 1720s, perhaps as early as 1670.{{#tag:ref|See the differing accounts of Partington<ref name=Partington>{{cite journal|last=Partington|first=J.R.|title=The early history of hydrofluoric acid|journal=Memoirs and Proceedings of the Manchester Literary and Philosophical Society|year=1923|volume=67|issue=6|pages=73–87}}</ref> and Weeks.<ref name="Weeks"/>||group="note"}} [[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.<ref>{{{cite journal|author=Marggraf, Andreas Sigismun|year=1770|url=http://books.google.com/books?id=f28pAXS5SGIC&pg=PA3|title=Observation concernant une volatilisation remarquable d'une partie de l'espece de pierre, à laquelle on donne les noms de flosse, flüsse, flus-spaht, et aussi celui d'hesperos; laquelle volatilisation a été effectuée au moyen des acides|trans_title=Observation of a remarkable volatilization of part of a type of stone to which one gives the name flosse, flüsse, flus-spaht, as well as that of hesperos; which volatilization was effected by means of acids|journal=Mémoires de l'Académie royale des sciences et belles-lettres|pages=3–11|language=French}}</ref><ref name="ModFcchem"/> 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).<ref name="ModFcchem"/><ref>{{cite journal|author=Scheele, Carl Wilhelm|year=1771|url=http://books.google.com/books?id=d984AAAAMAAJ&pg=PA120|title=Undersŏkning om fluss-spat och dess syra|trans_title=Investigation of fluorite and its acid|journal=Kungliga Svenska Vetenskapsademiens Handlingar (Proceedings of the Royal Swedish Academy of Science)|volume=32|pages=129–138|language=Swedish}}</ref>

In 1810, French physicist [[André-Marie Ampère]] suggested that the acid was a compound of hydrogen with an unknown element, analogous to chlorine.<ref>{{cite journal|last = Ampere|first = André-Marie|title = Suite d'une classification naturelle pour les corps simples|url = http://books.google.com/?id=4jEFAAAAQAAJ&pg=RA1-PA5|journal= Annales de chimie et de physique|year = 1816|volume = 2|pages = 1–5|language=French}}</ref> Fluorite was then shown to be mostly composed of [[calcium fluoride]].<ref name="Weeks">{{cite journal|last1 = Weeks|first1 = Mary Elvira|authorlink1=Mary Elvira Weeks|title = The discovery of the elements. XVII. The halogen family|journal= Journal of Chemical Education|volume = 9|pages = 1915–1939|year = 1932|doi = 10.1021/ed009p1915|issue = 11|bibcode = 1932JChEd...9.1915W}}</ref> [[Humphry Davy|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".<ref>{{cite web|url=http://elements.vanderkrogt.net/element.php?sym=F |title=09 Fluorine |publisher=elements.vanderkrogt.net |accessdate=24 January 2012}}</ref><ref name=LSJ>{{LSJ|fqo/rios|φθόριος|ref}}</ref> The New Latin name (''fluorum'') gave the element its current symbol, F, although the symbol Fl was used in early papers.<ref name="dissolve"/>{{#tag:ref|Since 2012, the symbol Fl has been used for [[flerovium]] (element 114), a man-made [[transuranic element]].<ref>{{cite web |url=http://www.iupac.org/news/news-detail/article/element-114-is-named-flerovium-and-element-116-is-named-livermorium.html |title=Element 114 is named flerovium and element 116 is named livermorium |date=30 May 2012 |publisher=[[International Union of Pure and Applied Chemistry|IUPAC]] |accessdate=1 June 2012}}</ref> ||group="note"}}

===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 incapacitated while working with hydrofluoric acid.{{#tag:ref|The injured included Davy, [[Joseph Louis Gay-Lussac]], [[Louis Jacques Thénard]], and Irish chemists [[Thomas and George Knox]]. Belgian chemist [[Paulin Louyet]] and French chemist [[Jerome Nickles]] died. Moissan also experienced serious HF poisoning.<ref name="Weeks"/><ref name = "Toon"/>||group="note"}} Initial attempts to isolate the element were also hindered by material difficulties: problems getting a suitable conducting liquid for [[electrolysis]] as well as the extreme corrosiveness of hydrogen fluoride (and of fluorine gas).<ref name="Weeks"/><ref name = "Toon"/>

[[Edmond Frémy]] thought that passing electric current through pure hydrofluoric acid (dry HF) might work. Previously, hydrogen fluoride was only available in a water solution. Frémy therefore devised a method for producing dry hydrogen fluoride by acidifying [[potassium bifluoride]] (KHF<sub>2</sub>). Unfortunately, pure hydrogen fluoride did not pass an electric current.<ref name="Weeks"/><ref name = "Toon"/><ref name=Asimov>{{cite book|last=Asimov|first=Isaac|year=1966|title=The noble gases|isbn=978-0-465-05129-8|publisher=Basic Books|page=162}}</ref>

French chemist [[Henri Moissan]], formerly one of Frémy's students, continued the search. After trying many different approaches, he combined potassium bifluoride and dry hydrogen fluoride. The mixture conducted electricity. Moissan also constructed especially corrosion-resistant equipment: containers crafted from a mixture of [[platinum]] and [[iridium]] (more chemically resistant than pure platinum) with fluorite stoppers.{{sfn|Greenwood|Earnshaw|1998|pp=789–791}}<ref name = "Toon"/> After 74 years of effort by many chemists, on 26 June 1886, Moissan isolated elemental fluorine.<ref name=Asimov/><ref name='qwen'>{{cite journal|title = Action d'un courant électrique sur l'acide fluorhydrique anhydre| last1 = Moissan| first1 = Henry| journal= Comptes rendus hebdomadaires des séances de l'Académie des sciences| year = 1886| volume = 102| pages = 1543–1544| url = http://gallica.bnf.fr/ark:/12148/bpt6k3058f/f1541.chemindefer|accessdate=7 May 2011|language=French}}</ref> In 1906, two months before his death, Moissan received the Nobel Prize in chemistry.<ref>{{cite web|accessdate=7 July 2009|url=http://nobelprize.org/nobel_prizes/chemistry/laureates/1906/ |publisher=nobelprize.org|title=The Nobel Prize in chemistry 1906}}</ref>{{sfn|Greenwood|Earnshaw|1998|p=791}}<ref name="Chem Nobel laureates">{{cite book|last=James|first=Laylin K.|title=Nobel laureates in chemistry, 1901–1992|year=1993|publisher=Chemical Heritage Foundation|isbn=978-0-8412-2690-6|url=http://books.google.com/books?id=jEy67gEvIuMC&pg=PA35|page=35}}</ref> The citation:<ref name = "Toon">{{cite journal|title=Fluorine, an obsession with a tragic past|author=Toon, Richard|journal= Education in Chemistry|date=September 2011|pages=148–151|url= http://www.rsc.org/images/EiC_Sept2011_fluorine_tcm18-207109.pdf}}</ref>{{#tag:ref|Moissan's Nobel also honored his invention of the [[electric arc furnace]]||group="note"}}

<blockquote>...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.</blockquote>

{| style="margin:1em auto 1em auto; text-align:left; width:463px;" cellpadding="0" cellspacing="0"
|- style="vertical-align:top;"
| style="width:243px; padding-right:20px;" | [[File:Recherches sur l’isolement du fluor, Fig. 5.PNG|223px|alt=]]
| style="width:220px; padding-left:20px;" | [[File:Henri Moissan HiRes.jpg|200px|alt=]]
|- style="text-align:center; vertical-align:top;"
| style="width:243px; padding-right:20px; "|Moissan's cell, 1887 publication
| style="width:220px; padding-left:20px; "|Henri Moissan, Nobel Prize photo
|}

===Later developments===

During the 1930s and 1940s, the [[DuPont]] company commercialized organofluorine compounds at large scales. Following trials of chlorofluorcarbons as refrigerants by [[General Motors]], DuPont developed large-scale production of [[Freon-12]] in 1930. It proved to be a marketplace hit, rapidly replacing earlier, more toxic, refrigerants and growing the overall market for kitchen refrigerators. In 1938, Teflon was discovered by accident by a recently hired DuPont Ph.D., [[Roy J. Plunkett]]. While working with a cylinder of [[tetrafluoroethylene]], he was unable to release the gas although the weight had not changed. Scraping down the container, he found white flakes of a polymer new to the world. Tests showed the substance was more resistant to corrosion and had better high temperature stability than any other plastic. By 1941, an accelerated commercialization program was making significant quantities.<ref name="ModFcchem"/><ref name="Jstg">{{cite journal|last1 = Okazoe|first1 = Takashi|title = Overview on the history of organofluorine chemistry from the viewpoint of material industry|journal= Proceedings of the Japan Academy, Series B|volume = 85|pages = 276–289|year = 2009|doi = 10.2183/pjab.85.276|issue = 8|bibcode = 2009PJAB...85..276O|format=PDF|url=https://www.jstage.jst.go.jp/article/pjab/85/8/85_8_276/_pdf}}</ref>{{sfn|Hounshell|Smith|1988|pp=156–157}}<ref name="DD Freon history">{{cite web|title=Freon|url=http://www2.dupont.com/Phoenix_Heritage/en_US/1930_d_detail.html|publisher=dupont.com|accessdate=10 November 2012}}</ref>

Large-scale productions of elemental fluorine began during World War II. Germany used high-temperature electrolysis to produce tons of [[chlorine trifluoride]], a compound planned to be used as an incendiary.<ref>{{cite book|last1 = Meyer|first1 = Eugene|title = Chemistry of hazardous materials|journal= Prentice Hall|edition = 1st|page = 111|year = 1977|isbn=9780131292390}}</ref> The [[Manhattan Project]] in the United States used even more fluorine. Uranium was converted to gaseous [[uranium hexafluoride]] for [[gaseous diffusion]] processing to produce [[enriched uranium]] by separating the [[fissile]] isotope [[uranium-235]] from the heavier [[uranium-238]]. Because of uranium hexafluoride is corrosive,<ref name="Jstg"/> the separation plants were built with special materials. All pipes were coated with [[nickel]]; joints and flexible parts were fabricated from Teflon.<ref name="Jstg"/>{{failed verification|date=October 2013}}

In the 1970s, concern developed that chlorofluorocarbons were damaging the [[ozone layer]]. By 1996, almost all nations had banned CFCs, and commercial production ceased. Fluorine continued to play a role in refrigeration though: [[hydrochlorofluorocarbons]] (HCFCs) and [[hydrofluorocarbons]] (HFCs) were the replacement refrigerants.<ref>{{cite web|url=http://www.epa.gov/ozone/defns.html#hcfc|title=Ozone depletion glossary|accessdate=3 September 2008|author=U.S. Environmental Protection Agency|year=2008}}</ref><ref>{{cite web|url=http://www.epa.gov/ozone/science/q_a.html |title=Brief questions and answers on ozone depletion &#124; Ozone layer protection |author= U.S. Environmental Protection Agency |year=2006 |accessdate=8 November 2011}}</ref>

==Industry and applications==

{{Main|Fluorochemical industry}}

A market survey estimated 2011 global fluorochemical sales at $15 billion. It predicted 5% growth to 2018, when the market would exceed $20 billion.<ref>{{cite web|title=Fluorochemicals market is expected to reach USD 21.5 billion globally by 2018: Transparency Market Research |url=http://www.tmrblog.com/2013/05/fluorochemicals-market-is-expected-to.html|publisher=Transparency Market Research Blog|author=Transparency Market Research|year=2013|accessdate=3 June 2013}}</ref> A different report said that, in 2016, global fluorochemical production would be almost $20 billion with a volume of 3.5 million metric tons per year.<ref>{{cite web|url=http://www.prnewswire.com/news-releases/world-fluorochemicals-market-freedonia-161338935.html|title=World Fluorochemicals Market, Freedonia|year=2012|accessdate=25 May 2013|publisher=PR Newswire<!--author not stated-->}}</ref> The main source of fluorine, fluorite mining, was estimated in 2003 to be a $550-million industry, extracting 4.5&nbsp;million&nbsp;tons per year.<ref name="ModFcchem">{{cite book|chapter = Fluorine|url = http://books.google.com/?id=ycSf3EK9i2gC&pg=PA5|pages = 3–10|title = Modern fluoroorganic chemistry: Synthesis, reactivity, applications|isbn = 978-3-527-30691-6|author1 = Kirsch, Peer|year = 2004}}</ref>

Mined fluorite is concentrated by [[Froth flotation|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 primarily converted to hydrogen fluoride, which is the primary [[chemical intermediate]] for the industry.<ref name="AcountF"/><ref name="ModFcchem"/><ref name="SME mining 2006">{{cite book|title=Industrial minerals & rocks: Commodities, markets, and uses|year=2006|publisher=Society for Mining, Metallurgy, and Exploration (U.S.)|isbn=978-0-87335-233-8|pages=461–473|url=http://books.google.com/books?id=zNicdkuulE4C&pg=PA472|last1=Kogel|first2= Nikhil C. |last2=Trivedi|first3=James M. |last3=Barker}}</ref>

<center>
<imagemap>
Image:The fluorine economy.svg|thumb|675px|center|Fluorine industry supply chain: Percentages indicate approximate mass flows of the element. Click for links to related articles.
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>
</center>

===Inorganic fluorides===

About 3&nbsp;kg (6.5&nbsp;lb) of metspar-grade fluorite are added to make each metric ton of steel. The fluoride ions from CaF<sub>2</sub> lower the melt's temperature and [[viscosity]], making it runnier. Metspar is similarly used to produce cast iron and other iron alloys.<ref name="AcountF"/><ref name="USGS">{{cite book|chapter=Fluorspar|author=Miller, M. Michael|publisher=U.S. Geological Survey|year=2003|pages=27.1–27.12|title=U.S. Geological Survey Minerals Yearbook|url=http://minerals.usgs.gov/minerals/pubs/commodity/fluorspar/fluormyb03.pdf}}</ref>

Acidspar grade fluorite is added to ceramics, enamels, glass fibers, clouded glass, cement, and the outer coating of welding rods.<ref name="AcountF"/> Most acidspar is reacted with sulfuric acid to make hydrofluoric acid. Significant direct uses of HF include [[Pickling (metal)|pickling]] (cleaning) steel, etching glass, and [[cracking (chemistry)|cracking alkanes]] in the petrochemical industry.<ref name="AcountF"/>

One third of HF (one sixth of mined fluorine) is used to make synthetic cryolite ([[sodium hexafluoroaluminate]]) and [[aluminium fluoride|aluminium trifluoride]]. These compounds are used in the electrolysis of aluminium by the [[Hall–Héroult process]]. About 23&nbsp;kg (51&nbsp;lb) are required for every metric ton of aluminium.<ref name="AcountF"/> Although small amounts are lost through side reactions, the fluorides are not reactants in the smelting process, but fluxes that lower the temperature of the melt.<ref>{{cite report|url=https://www1.eere.energy.gov/manufacturing/resources/aluminum/pdfs/aluminum.pdf|title=Energy and Environmental Profile of the U.S. Aluminum Industry|year=1997|author=Energetics, Inc.|pages=41, 50}}</ref>

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 treatment of effluents in laundries.{{sfn|Aigueperse et al.|2005|p=33}}

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]] (antireflection optical coatings).<ref name="AcountF"/> Sodium and potassium [[bifluoride]]s are significant to the chemical industry.{{sfn|Aigueperse et al.|2005|pp=25–27}}<ref>{{cite book|title=Practical equipment, materials, and processes for optical thin films|author=Willey, Ronald R.|page=113|isbn=9780615143972|publisher=Willey Optical|year=2007}}</ref>

===Fluorocarbons===

Organofluoride production consumes over 40% of hydrofluoric acid (over 20% of all mined fluorite). Within organofluorides, refrigerant gases are still the dominant segment, about 80% on a fluorine basis. Fluoropolymers are less than one quarter the size of refrigerant gases in terms of fluorine usage but are growing faster.<ref name="AcountF"/><ref name="GIA f-chem report description">{{cite web|title=Global fluorochemicals Market to exceed 2.6 million tons by 2015, according to a new report by Global Industry Analysts, Inc.|url=http://www.prweb.com/releases/fluorochemicals/organic_inorganic/prweb4708534.htm|publisher=Global Industry Analysts (via PRWeb)|accessdate=26 January 2012|year=2010}}</ref> [[Fluorosurfactant]]s (materials like [[Scotchgard]], used in [[durable water repellent]]s) are a small segment in mass but are over $1 billion yearly revenue.<ref name="fluorosurfactant market">{{cite journal |author=Renner, R. |title=The long and the short of perfluorinated replacements |journal=Environmental Science & Technology |volume=40 |issue=1 |year=2006 |pmid=16433328 |doi=10.1021/es062612a |pages=12–3|bibcode = 2006EnST...40...12R }}</ref>

====Refrigerant gases====

{{See also|Refrigerant}}

Traditionally chlorofluorocarbons (CFCs) were the predominant class of fluorinated organic chemical. CFCs are identified by a system of numbering ([[Refrigerant#Refrigerants by class and by R-number|the R-number system]]) that explains the amount of fluorine, chlorine, carbon and hydrogen in the molecules.<ref>{{cite web|url=http://www2.dupont.com/Refrigerants/en_CA/products/understanding.html|title=Understanding the refrigerant "R" nomenclature|publisher=DuPoint|accessdate=14 June 2013|year=2013}}</ref> The DuPont brand Freon has been colloquially used for CFCs and similar halogenated molecules; brand-neutral terminology uses "R" ("refrigerant") as the prefix. Prominent CFCs included R-11 ([[trichlorofluoromethane]]), R-12 ([[dichlorodifluoromethane]]), and R-114 ([[1,2-dichlorotetrafluoroethane]]).<ref name="AcountF"/>

Production of CFCs grew strongly through the 1980s, primarily for refrigeration and air conditioning but also for propellants and solvents. Since the end use of these materials is now banned in most countries, this industry has shrunk dramatically. By the early 21st century, production of CFCs was less than 10% of the mid-1980s peak.<ref name="AcountF"/>

Hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) now serve as replacements for CFC refrigerants; few were commercially manufactured before 1990. Currently more than 90% of fluorine used for organics goes into these two classes, in about equal amounts. Prominent HCFCs include R-22 ([[chlorodifluoromethane]]) and R-141b ([[1,1-Dichloro-1-fluoroethane|1,1-dichloro-1-fluoroethane]]). The main HFC is R-134a ([[1,1,1,2-Tetrafluoroethane|1,1,1,2-tetrafluoroethane]]).<ref name="AcountF"/>

====Fluoropolymers====

{{Main|Fluoropolymer}}

The 2011 global fluoropolymer market was estimated at slightly under $6 billion in revenue and predicted to grow 6.5% per year until 2016.<ref>{{cite web|url=http://www.prweb.com/releases/fluoropolymers/market/prweb10608563.htm|title=Fluoropolymers market is poised to grow at a CAGR of 6.5% & to reach $9,446.0 Million by 2016 – New report by MarketsandMarkets|author=MarketsandMarkets|year=2013}}</ref> As of about 2006–2007, fluoropolymer volume was estimated at over 180,000&nbsp;metric&nbsp;tons per year. The corresponding revenue estimate was over $3.5 billion.<ref name="yearly FP production">{{cite doi|10.1134/S1070363209030335}}</ref>

Polytetrafluoroethylene (PTFE) is 60–80% of the world's fluoropolymer production on a weight basis.<ref name="yearly FP production"/> The DuPont brand Teflon is sometimes used generically for the substance.<ref name="encyc of structure"/> 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, in tubing, and in gaskets. Another major use is [[Fabric structure|architectural fabric]] (PTFE-coated fiberglass cloth used for stadium roofs). The major consumer application is [[Non-stick surface|non-stick]] cookware.<ref name="encyc of structure">{{cite book|last=Martin|first=John Wilson|title=Concise encyclopedia of the structure of materials|year=2007|publisher=Elsevier|isbn=978-0-08-045127-5|pages=187–194|url=http://books.google.com/books?id=xv420pEC2qMC&pg=PA189}}</ref>

{| style="margin:1em auto 1em auto; text-align:left; width:691px;" cellpadding="0" cellspacing="0"
|-
| colspan="3" style="text-align:center; padding-bottom:8px;"| '''Major PTFE applications'''
|-
| style="width:257px;" | [[File:Coaxial cable cutaway PTFE.svg|257px|alt=cutaway of a coax cable (plastic around metal around PTFE around metal core)]]
| style="width:169px; padding-left:30px; padding-right:30px;" | [[File:Happy Pan Poster.jpg|109px|alt=advertisement for the "Happy Pan" with DuPont Teflon]]
| style="width:265px;" | [[File:Tokyo Dome 2007-1.jpg|265px|alt=vast expanse with no columns, gray roof with sections, bright green playing field]]
|-
| PTFE dielectric separating core and outer metal in a specialty coaxial cable
| style="padding-left:30px; padding-right:30px;" | First Teflon branded frying pan, 1961
| The interior of the [[Tokyo Dome]]. The roof is PTFE-coated fiberglass and air-supported.<ref name=Nakagawa>{{cite web|last=Nakagawa|first=Ulara|title=15 sights that make Tokyo so fascinating|url=http://www.cnngo.com/tokyo/life/15-reasons-why-tokyo-rules-583723|publisher=CNN|accessdate=31 December 2011|year=2011}}</ref>
|}
{{-}}

When jerk-stretched, PTFE film becomes a fine-pored [[Artificial membrane|membrane]]: expanded PTFE (ePTFE). The term "[[Gore-Tex]]" is sometimes used generically although that is a specific brand name. ePTFE is used in rainwear, [[Personal protective equipment|protective apparel]], and [[filtration|liquids and gas filters]]. PTFE can also be formed into [[PTFE fiber|fibers]] which are used in [[seal (mechanical)|pump packing seals]] and [[Dust collector|bag house filters]].<ref name="encyc of structure"/>

[[File:Goretex photo.png|alt=SEM image with islands of white and interconnecting strings|Electron microscope image of ePTFE ("Gore-Tex"): the structure consists of islands of polymer with connecting strands.<ref>{{cite book|title=Fluoroplastics, Volume 1: Non-Melt Processible Fluoroplastics|author=Ebnesajjad, Sina|year=2000|pages=176–177|publisher=William Andrew|isbn=9780815517276|url=http://books.google.com/books?id=-H4yIE71aGMC&pg=PA176}}</ref>|thumb|right|upright=1.1]]

Other fluoropolymers tend to have similar properties to PTFE, which leads to use in electrical insulation and the chemical process industry. Unlike PTFE, the other fluoropolymers can be melt processed. This makes them easier to work with (to form into complex shapes). But they are more expensive than PTFE and have lower thermal stability. [[Fluorinated ethylene propylene]] (FEP) is the second most produced fluoropolymer. Films from two different fluoropolymers serve as glass-replacements in solar cells.<ref name="encyc of structure"/><ref name=Altenergy>{{cite web|last=Bhiwankar|first=Nikhil|title=Weathering the storm: Fluoropolymer films protect solar modules and ensure performance|publisher=altenergymag.com|year=2011|url=http://www.altenergymag.com/emagazine/2011/10/weathering-the-storm-fluoropolymer-films-protect-solar-modules-and-ensure-performance/1800|accessdate=31 December 2011}}</ref><ref name=DeBergalis>{{cite journal|last=DeBergalis|first=Michael|title=Fluoropolymer films in the photovoltaic industry|journal=Journal of Fluorine Chemistry|year=2004|volume=125|pages=1255–1257|url=http://www2.dupont.com/Photovoltaics/en_US/assets/downloads/pdf/DeBergalis_flpr_films_wp.pdf|doi=10.1016/j.jfluchem.2004.05.013|issue=8}}</ref>

Fluorinated [[ionomer]]s (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 a fuel cell material in spacecraft. Since then, the polymer has been transforming the 55&nbsp;million&nbsp;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 get [[proton exchange membrane]] (PEM) fuel cells into automobiles.<ref name="Fluorinated Ionomers">{{cite book|last=Grot|first=Walter|title=Fluorinated ionomers|year=2011|publisher=Elsevier|isbn=978-1-4377-4457-6|pages=1–10|url=http://books.google.com/books?id=E8H1Hwd5GXUC&pg=PA6}}</ref><ref name="Functional Materials">{{cite book|last=Ramkumar|first=Jayshreee|title=Functional materials: Preparation, processing and applications|year=2012|publisher=Elsevier|isbn=978-0-12-385142-0|pages=549–578|url=http://books.google.com/books?id=ep7U5G4O3mQC&pg=PA567|editor=Banerjee, S|chapter=Nafion Persulphonate Membrane: Unique Properties and Various Applications}}</ref><ref name="Chlor-alkali and chlorate technology">{{cite book|last=Burney|first=H.S.|title=Chlor-alkali and chlorate technology: R.B. MacMullin memorial symposium|year=1999|publisher=The Electrochemical Society|isbn=1-56677-244-3|pages=105–126|url=http://books.google.com/books?id=rtm_bjQd8SUC&pg=PA111|chapter=Past, Present, and Future of the Chlor-Alkali Industry}}</ref>

[[Fluoroelastomer]]s are rubber-like substances that are composed of [[crosslink]]ed mixtures of fluoropolymers. Chemical-resistant [[O-ring]]s are the primary application. [[Viton]] is a prominent brand.<ref name="encyc of structure"/>

===Fluorine gas===

{{see also|Industrial gas}}

For countries with available data, about 17,000&nbsp;metric&nbsp;tons of fluorine are produced per year. Fluorine is relatively inexpensive, costing about $5–8 per kilogram ($2–4 per pound) when sold as uranium hexafluoride or sulfur hexafluoride. Because of difficulties in storage and handling, the price of fluorine gas is much higher. Processes demanding large amounts of fluorine gas generally [[vertical integration|vertically integrate]] and produce the gas onsite.{{sfn|Jaccaud et al.|2005|p=12}}

[[File:SF6 current transformer TGFM-110 Russia.jpg|thumb|upright=1.2|alt=minaret-like electrical things with wires running around and fat at the bottom|SF<sub>6</sub> transformers at a Russian railway]]

The largest application for elemental fluorine (up to 7,000&nbsp;metric&nbsp;tons per year) is the preparation of uranium hexafluoride, which is used in the production of nuclear fuels. To obtain the compound, [[uranium dioxide]] is first treated with hydrofluoric acid, to produce [[uranium tetrafluoride]]. This compound is then further fluorinated by direct exposure to fluorine gas to make the hexafluoride.{{sfn|Jaccaud et al.|2005|p=12}} Fluorine's monoisotopic natural occurrence makes it useful in uranium enrichment because uranium hexafluoride molecules will differ in mass only because of mass differences between uranium-235 and uranium-238. These mass differences are used to concentrate uranium-235 by diffusion or centrifugation.<ref name="AcountF"/>{{sfn|Jaccaud et al.|2005|p=12}}

The second largest application for fluorine gas (about 6,000&nbsp;metric&nbsp;tons per year) is in [[sulfur hexafluoride]], which is used as a dielectric medium in high voltage transformers and circuit breakers. SF<sub>6</sub> gas has a much higher dielectric strength than air and is extremely chemically inert. Switchgear using SF<sub>6</sub> has no hazardous [[polychlorinated biphenyl]]s (PCBs), in contrast to traditional {{nowrap|oil-filled}} devices.{{sfn|Aigueperse et al.|2005|p=35}}

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]].{{sfn|Jaccaud et al.|2005|pp=11–12}}<ref>{{cite book|url = http://books.google.com/?id=Wg1SnT6Cb0QC&pg=PA317|page = 317|chapter = Fluorine-based plasmas|title = Fundamentals of semiconductor processing technology|isbn = 978-0-7923-9534-8|last = El-Kareh|first = Badih|year = 1994|accessdate=7 May 2011}}</ref><ref>{{cite journal|first1 = Leonel R.|last1 = Arana|first2 = Nuria|last2 = de Mas|first3 = Aleksander J.|last3 = Schmidt|last4 = Franz|first4 = Martin A. |first5 =F. |last5 =Schmidt |first6 = Klaus F. |last6 =Jensen|doi = 10.1088/0960-1317/17/2/026|title = Isotropic etching of silicon in fluorine gas for MEMS micromachining|journal= Journal Micromechanical Microenergy|year = 2007|volume = 17|page = 384|issue = 2|bibcode = 2007JMiMi..17..384A}}</ref> [[Nitrogen trifluoride]] is used for cleaning equipment.<ref name="AcountF"/>

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<sub>3</sub>, BrF<sub>3</sub>, and IF<sub>5</sub> provide gentler fluorination, with a series of strengths, and are easier to handle. Sulfur tetrafluoride is used for making fluorinated pharmaceuticals.<ref name="AcountF"/>

==Production of fluorine gas==

[[File:Fluorine cell room.jpg|thumb|upright=1.5|Industrial fluorine cells at F2 Chemicals, [[City of Preston, Lancashire|Preston, England]]|alt=A machine room]]

Today, elemental fluorine is produced using Moissan's process of electrolyzing potassium fluoride/hydrogen fluoride mixtures, but with an apparatus made of different materials. The steel container acts as the cathode, attracting H<sup>+</sup> ions and releasing hydrogen gas. A carbon block (similar to that used in aluminium production) acts as the anode, attracting F<sup>-</sup> ions and releasing fluorine gas. The voltage is 8–12&nbsp;volts.<ref name="AcountF"/>{{sfn|Jaccaud et al.|2005|p=6}}

When combined, potassium fluoride and hydrogen fluoride spontaneously produce potassium bifluoride (KHF<sub>2</sub>), which increases the conductivity of the solution. A mixture with the approximate composition KF•2HF melts at 70&nbsp;°C (158&nbsp;°F) and is electrolyzed at 70–130&nbsp;°C (160–265&nbsp;°F). Because HF alone cannot be electrolyzed, the presence of some KF is critical even though it is not consumed in the cell.<ref name="ModFcchem"/>{{sfn|Jaccaud et al.|2005|pp=4–5}}{{sfn|Greenwood|Earnshaw|1998|pp=796–797}}

In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, [[Karl O. Christe]] discovered a purely chemical preparation of fluorine gas. However, he stated in his work that the basics were known 50 years before the actual reaction.<ref name=Kirsch2004>{{Cite book | title = Modern fluoroorganic chemistry: Synthesis, reactivity, applications | url = http://books.google.com/books?id=aZoQP2YFsakC&pg=PA7 | year = 2004 | author = Kirsch, Peer | page = 7 | isbn = 3-527-30691-9 | publisher = John Wiley & Sons}}</ref> The main idea is that some metal fluoride anions do not have a neutral counterpart (or those are very unstable) and their acidifying would result in chemical oxidation, rather than formation of the expected molecules. Christe lists the following reactions as a possible way:<ref name='year'>{{cite journal|author = Christe, K.|journal= [[Inorganic Chemistry]]|volume = 25|pages = 3721–3724|doi = 10.1021/ic00241a001|title = Chemical synthesis of elemental fluorine|year = 1986|issue = 21}}</ref>

:2&nbsp;KMnO<sub>4</sub> + 2&nbsp;KF + 10&nbsp;HF + 3&nbsp;H<sub>2</sub>O<sub>2</sub> → 2&nbsp;K<sub>2</sub>MnF<sub>6</sub> + 8&nbsp;H<sub>2</sub>O + 3&nbsp;O<sub>2</sub>↑

:2&nbsp;K<sub>2</sub>MnF<sub>6</sub> + 4&nbsp;SbF<sub>5</sub> → 4&nbsp;KSbF<sub>6</sub> + 2&nbsp;MnF<sub>3</sub> + F<sub>2</sub>↑

Pure fluorine gas may be stored in steel cylinders, where the inside surface is passivated, as long as the temperature is kept below 200&nbsp;°C (400&nbsp;°F). Above that temperature, nickel is required.<ref name="ModFcchem"/>{{sfn|Jaccaud et al.|2005|pp=4–5}} Regulator valves are made of nickel. Fluorine piping is generally made of nickel or [[Monel]] (nickel-copper alloy).{{sfn|Jaccaud et al.|2005|pp=10–11}} 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,{{sfn|Jaccaud et al.|2005|pp=10–11}} but some sources recommend systems made of nickel, Monel, and PTFE.<ref name="Shriver Inorganic">{{cite book|last1=Shriver|first1=Duward|last2=Atkins|first2=Peter|title=Solutions manual for inorganic chemistry|year=2010|publisher=Macmillan|isbn=978-1-4292-5255-3|url=http://books.google.com/books?id=ZWKuqaeYXSEC&pg=PA427|page=427}}</ref>

[[File:Laboratory apparatus for reactions with fluorine gas.jpg|thumb|center|upright=2|Laboratory setup for fluorine reactions, [[University of Freiburg]]. The stainless steel enclosure on the left surrounds a small cylinder of fluorine. The line's piping, fittings and valves are Monel. The rig is in a fume hood; many controls are external, allowing reactions with the door closed.<ref>{{cite web|url=http://portal.uni-freiburg.de/molchem/research/f2-lab/f2-lab-intro#|title=Fluorine Laboratory|author=Lehrstuhl für Molekül- und Koordinationschemie|publisher=University of Freiburg|year=2013|accessdate=28 May 2013}}</ref>]]

==Biological aspects==

{{Main|Biological aspects of fluorine}}

===Natural biochemistry===

[[File:Gifblaar.jpg|thumb|South Africa's [[gifblaar]] is one of the few organisms that makes fluorine compounds.]]

Fluoride is not considered an [[micronutrient|essential]] mineral element for mammals or humans. Small amounts of fluoride 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.<ref name="gastroenterology1">{{cite journal|last1 = Nielsen|first1 = Forrest H.|doi = 10.1053/j.gastro.2009.07.072|title = Micronutrients in parenteral nutrition: Boron, silicon, and fluoride|pmid = 19874950|year = 2009|journal= Gastroenterology|volume = 137|issue = 5 Suppl|pages =S55–S60}}</ref><ref name="quott">{{cite report|last1= Olivares |first1=M. |last2= Uauy |first2=R.|year = 2004| url=http://www.who.int/water_sanitation_health/dwq/en/nutoverview.pdf|title=Essential nutrients in drinking water (Draft)|publisher=[[World Health Organization|WHO]]|accessdate=30 December 2008}}</ref>

Biologically synthesized organofluorines have been found in microorganisms and plants,<ref name=Gribble>{{cite journal|last1 =Gribble|first1 =Gordon W.|title =Naturally occurring organofluorines|volume =3N|pages =121–136|year =2002|doi =10.1007/10721878_5|journal= The Handbook of Environmental Chemistry|series =The Handbook of Environmental Chemistry|isbn =3-540-42064-9}}</ref> but not in animals.<ref name="Murphy2003">{{Cite journal| last1 = Murphy |pmid=12738270| first1 = C.| last2 = Schaffrath| first2 = C.| last3 = O'Hagan| first3 = D.| title = Fluorinated natural products: The biosynthesis of fluoroacetate and 4-fluorothreonine in ''Streptomyces cattleya''| doi = 10.1016/S0045-6535(03)00191-7| year = 2003| journal= Chemosphere| volume = 52| issue = 2| pages =455–461}}</ref> The most common example is [[sodium fluoroacetate|fluoroacetate]]. It is used as a [[plant defense against herbivores|defense against herbivores]] by at least 40 green plants in Australia, Brazil, and Africa.<ref name="Proudfoot">{{cite journal |last1=Proudfoot|first1 = A.T.|last2= Bradberry|first2= S.M. |last3= Vale |first3= J.A. |title=Sodium fluoroacetate poisoning |journal=Toxicological Reviews |volume=25 |issue=4 |pages=213–219 |year=2006 |pmid=17288493 |doi=10.2165/00139709-200625040-00002}}</ref> Other biologically synthesized organofluorines include ω-fluoro [[fatty acid]]s, fluoroacetone, and 2-fluorocitrate.<ref name="Murphy2003"/> The enzyme [[adenosyl-fluoride synthase]], which makes the carbon–fluorine bond, was isolated in bacteria in 2002. The discovery was touted as possibly leading to biological routes for organofluorine synthesis.<ref>{{cite journal|last1= O'Hagan |first1=D.|last2= Schaffrath |first2=C. |last3=Cobb |first3=S.L. |last4=Hamilton |first4=J.T. |last5=Murphy |first5=C.D.|year = 2002|title = Biochemistry: Biosynthesis of an organofluorine molecule|journal= Nature |volume = 416|page = 279 |pmid = 11907567|doi = 10.1038/416279a|issue = 6878|bibcode = 2002Natur.416..279O }}</ref>

===Dental care===

{{Main|Fluoride therapy|Water fluoridation}}

Since the mid-20th century, researchers have noted in population studies that fluoride somehow 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 prevent cavities primarily by helping teeth that 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 medication|topical treatment]]) that prevents cavities; fluoride ions that are swallowed do not benefit the teeth.<ref name=Pizzo>{{cite journal |author= Pizzo, G.; Piscopo, M.R.; Pizzo, I.; Giuliana, G. |title= Community water fluoridation and caries prevention: a critical review |journal= Clinical Oral Investigation |volume=11 |issue=3 |pages=189–193 |year=2007 |pmid=17333303 |doi=10.1007/s00784-007-0111-6|url=http://www.newmediaexplorer.org/chris/Pizzo-2007.pdf}}</ref>

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

Water fluoridation is the controlled addition of fluoride to a [[public water supply]] to reduce tooth decay.<ref>{{cite journal |title= Recommendations for using fluoride to prevent and control dental caries in the United States |author= Centers for Disease Control and Prevention |journal= MMWR Recommendations and Reports |volume=50 |issue=RR–14 |pages=1–42 |year=2001 |pmid=11521913 |url=http://cdc.gov/mmwr/preview/mmwrhtml/rr5014a1.htm}}</ref> 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<ref name=Ripa>{{cite journal |author= Ripa, L.W. |url=http://aaphd.org/docs/position%20papers/A%20Half-Century%20of%20Community%20Water1993.pdf |archiveurl=http://web.archive.org/web/20090304021822/http://aaphd.org/docs/position%20papers/A%20Half-Century%20of%20Community%20Water1993.pdf |archivedate=2009-03-04 |title= A half-century of community water fluoridation in the United States: review and commentary |journal= Journal of Public Health Dentistry |volume=53 |issue=1 |pages=17–44 |year=1993 |pmid=8474047 |doi= 10.1111/j.1752-7325.1993.tb02666.x}}</ref> and for about 6% of people worldwide.<ref name=Cheng2007>{{cite journal |author= Cheng, K.K.; Chalmers, I.; Sheldon, T.A. |title= Adding fluoride to water supplies |url=http://www.appgaf.org.uk/data/433-water-fluoridation.pdf| journal=BMJ |volume=335 |issue=7622 |pages=699–702 |year=2007 |pmid=17916854 |pmc=2001050 |doi=10.1136/bmj.39318.562951.BE}}</ref> Although the best available evidence shows no association with adverse effects other than [[dental fluorosis]], most of which is mild,<ref>{{cite book|isbn=9789350252161 |year=2011 |title=A textbook of public health dentistry|author= Marya, C.M.|page=343|publisher=JP Medical Limited}}</ref> water fluoridation has been contentious for ethical, safety, and efficacy reasons.<ref name=Cheng2007/> [[Opposition to water fluoridation]] exists despite its support by public health organizations.<ref>{{cite journal |last=Armfield|first=J.M.|title= When public action undermines public health: A critical examination of antifluoridationist literature |journal= Australia and New Zealand Health Policy |volume=4 |page=25 |year=2007 |pmid=18067684 |pmc=2222595 |doi=10.1186/1743-8462-4-25|issue=1}}</ref> 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.<ref name="Dental caries"/> Reviews of the scholarly literature in 2000 and 2007 showed significant reduction of cavities in children associated with water fluoridatation.<ref name=NHMRC>{{cite web |url=http://www.nhmrc.gov.au/_files_nhmrc/publications/attachments/eh41_1.pdf |accessdate=24 February 2009 |year=2007 |title= A systematic review of the efficacy and safety of fluoridation |author= National Health and Medical Research Council (Australia)}} Summary: {{cite journal |author= Yeung, C.A. |title= A systematic review of the efficacy and safety of fluoridation |journal= Evidence-Based Dentistry |volume=9 |issue=2 |pages=39–43 |year=2008 |pmid=18584000 |doi=10.1038/sj.ebd.6400578 |laydate=2007 |laysource=NHMRC}}</ref>

[[Toothpaste]] may contain fluorine in the form of sodium fluoride, [[tin(II) fluoride|tin difluoride]], or (most commonly) [[sodium monofluorophosphate]]. The first fluoride toothpaste was introduced in 1955 in the United States. 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.<ref name="Crest continuing ed">{{cite web|last=Cracher|first=Connie Myers|title=Current concepts in preventive dentistry|url=http://www.dentalcare.com/media/en-US/education/ce334/ce334.pdf|publisher=dentalcare.com|accessdate=20 January 2012|page=12|year=2009}}</ref><ref name="Dental caries">{{cite book|title=Dental caries: The disease and its clinical management|year=2008|publisher=John Wiley & Sons|isbn=978-1-4051-3889-5|url=http://books.google.com/books?id=fZfXWhSmG1UC&pg=PA518|author=Fejerskov, Ole|coauthors=Kidd, Edwina;|page=518}}</ref>

===Pharmaceuticals===

About 20% of modern pharmaceuticals contain fluorine, including commercially significant drugs in many different pharmaceutical classes.<ref>{{cite book |last=Emsley |first=John |title=Nature's building blocks: An A–Z guide to the elements |edition=2nd |year=2011 |publisher=Oxford University Press | isbn=978-0-19-960563-7 |page=178}}</ref> One of these, the cholesterol-reducer [[atorvastatin]] (Lipitor), was the number one money-making drug for nearly a decade. The branded asthma medication [[Serevent]] (Advair), a top-ten revenue drug as of the mid-2000s, also contains a fluorinated molecule: [[fluticasone]].<ref name="url_Crain's_2011">{{cite web | url = http://www.crainsnewyork.com/article/20111228/HEALTH_CARE/111229902 | title = Lipitor becomes world's top-selling drug | date = 28 December 2011 | work = Crain's New York Business |accessdate=6 October 2013}}</ref><ref name = Swinson/>

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 to delay their [[drug metabolism|metabolism]] and elimination by the body. This allows longer times between doses.<ref name="Hagmann">{{cite journal |author=Hagmann, W.K. |title=The many roles for fluorine in medicinal chemistry |journal=Journal of Medicinal Chemistry |volume=51 |issue=15 |pages=4359–69 |year=2008 |month=August|doi=10.1021/jm800219f |pmid=18570365}}</ref> 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.<ref name = Swinson>{{cite journal |last=Swinson |first=Joel |journal=PharmaChem |title = Fluorine – A vital element in the medicine chest |year=2005 |pages=26–27 |accessdate=26 August 2010 |url=http://www.halocarbon.com/halocarbon_media/swinson_109.pdf |publisher=Pharmaceutical Chemistry}}</ref>

[[File:Prozac pills.jpg|thumb|alt=large image of just a capsule with words Prozac and DISTA visible|Prozac (fluoxetine)]]

Many modern [[antidepressant]]s are fluorinated molecules that selectively limit the body's binding of [[serotonin]] (low serotonin availability in brain cells is a cause of depression). Prior to the 1980s, traditional antidepressants, such as the [[tricyclic antidepressant|tricyclics]], altered not only serotonin uptake but also affected several other [[neurotransmitters]]. This non-selective interaction caused many side effects. One of the first drugs to alter only serotonin uptake—and be free of most side effects of previous drugs—was fluorine-containing [[Prozac]] (fluoxetine). It became the {{nowrap|best-selling}} antidepressant and prompted the popular book ''[[Listening to Prozac]]''. Some other [[selective serotonin reuptake inhibitor]] (SSRI) antidepressants that are fluorinated are [[citalopram|Celexa]] (and its isomer [[escitalopram|Lexapro]]), [[fluvoxamine|Luvox]], and [[paroxetine|Paxil]].<ref>{{cite book|title = Antidepressants|year = 2004|pages=37–39|last1 = Mitchell|first1 = E. Siobhan|last2 = Triggle|first2 = D.J.|publisher = Infobase Publishing|isbn=978-1-4381-0192-7}}</ref><ref>{{cite book|url=http://www.preskorn.com/books/ssri_s2.html|chapter=2: Rational Drug Discovery and SSRIs|title=Clinical Pharmacology of SSRI's|edition=1st|last=Preskorn|first=Sheldon H.|year=1996|publisher=Professional Communications, Inc.|isbn=1-884735-08-8}}</ref>

[[Quinolones]] are artificial compounds that are [[broad-spectrum antibiotic]]s. 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.<ref>{{cite doi|10.1186/1471-2334-11-187}}</ref><ref>{{cite web|url=http://well.blogs.nytimes.com/2012/09/10/popular-antibiotics-may-carry-serious-side-effects/|title=Popular antibiotics may carry serious side effects|last=Brody|first=Jane E.|year=2012|work=The New York Times|accessdate=3 June 2013}}</ref><ref name="pmid17342653">{{Cite journal|author=Nelson, J.M.; Chiller, T.M.; Powers, J.H.; Angulo, F.J.|title=Fluoroquinolone-resistant Campylobacter species and the withdrawal of fluoroquinolones from use in poultry: A public health success story |url=http://www.cdc.gov/narms/pdf/JNelson_FluoroquinoloneRCampy_CID.pdf|journal=Clinical Infectious Diseases |volume=44 |issue=7 |pages=977–980 |year=2007 |pmid=17342653 |doi=10.1086/512369}}</ref><ref>{{cite journal|url=http://www.aafp.org/afp/2000/0501/p2741.html|last1=King|first1=Dana E.|last2=Malone|first2=Robb|last3=Lilley|first3=Sandra H.|journal=American Family Physican|year=2000|volume=61|issue=9|pages=2741–2748|title=New classification and update on the quinolone antibiotics}}</ref>

Fluorine also finds use in many steroidal drugs.<ref>{{cite book|title=Glucocorticoids|author=Goulding, Nicolas J.; Flower, Rod J.|page=40|publisher=Springer|year=2001|isbn=9783764360597}}</ref> [[Florinef]] (fludrocortisone) is a [[mineralocorticoid]] (a compound used to retain sodium and water and thus raise blood pressure).<ref name="dontburn"/> [[Kenalog]] (triamcinolone) and [[dexamethasone]] are potent [[glucocorticoid]]s (anti-inflammatories).<ref name="dontburn">{{cite book |title=Pain-relieving procedures: The illustrated guide |page=58 |publisher=John Wiley & Sons |year=2012 |last1=Raj |first1=P. Prithvi |last2=Erdine |first2=Serdar |isbn=9781118300459}}</ref>

Several inhaled [[anesthetic]]s, 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 longer-lasting still and almost insoluble in blood, which accelerates the awakening. Examples include [[sevoflurane]], [[desflurane]], [[enflurane]], and [[isoflurane]], all fluorinated [[ether]]s.<ref name="fut">{{cite journal |last1=Filler |first1=R. |last2=Saha |first2=R. |url=http://www.future-science-group.com/_img/pics/fluorine_in_medicinal_chemistry.....pdf |title=Fluorine in medicinal chemistry: A century of progress and a 60-year retrospective of selected highlights |journal=Future Medicinal Chemistry |volume=1 |issue=5 |pages=777–791 |year=2009 |pmid=21426080 |doi=10.4155/fmc.09.65}}</ref><ref name="anestetics">{{cite book |title=Bioorganic and Medicinal Chemistry of Fluorine |pages=335–336 |publisher= John Wiley & Sons |year=2008 |last1= Bégué |first1=Jean-Pierre |last2=Bonnet-Delpon |first2=Daniele |isbn=9780470281871}}</ref>

===Agrichemicals and poisons===

[[Image:1080PoisonWarning gobeirne.jpg|thumb|right|upright=1.15|[[1080 usage in New Zealand|1080 use in New Zealand]]]]

An estimated 30% of [[agrochemical|agrichemical compounds]] contain fluorine.<ref>{{cite web|title = Fluorine's treasure trove|accessdate =20 February 2011|url = http://www.icis.com/Articles/2006/09/30/2016413/fluorines-treasure-trove.html|work = ICIS news|date = 2 October 2006}}</ref> Most of them are [[herbicide]]s and [[fungicide]]s, but a few [[Plant hormone|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 fluorinated pharmaceuticals.<ref name=Theodoridis/> [[Trifluralin]] is a prominent example, used widely in the United States as a weedkiller.<ref name=Theodoridis>{{cite book|pages=121–176|author=Theodoridis, George|year=2006|isbn=9780444526724|publisher=Elseiver|title=Fluorine and the Environment: Agrochemicals, Archaeology, Green Chemistry & Water}}</ref><ref name="trifluralin">{{cite journal|journal=Pesticides News|volume=52|year=2001|pages=20–21|title=Fact sheet: Trifluralin}}<!--|Author not stated--></ref> Its suspected carcinogenic properties have caused many European countries to ban it.<ref>{{cite report|author = European Commission|title = Trifluralin |year = 2007|url = http://www.unece.org/fileadmin/DAM/env/lrtap/TaskForce/popsxg/2008/Trifluralin_RA%20dossier_proposal%20for%20submission%20to%20the%20UNECE%20POP%20Protocol.pdf}}</ref>

[[Sodium monofluoroacetate]] (brand name 1080) is a commercial mammalian [[poison]]. The molecule is similar to the one in vinegar but with a hydrogen changed out for fluorine.{{#tag:ref|Vinegar is diluted [[acetic acid]]. A derivative salt is [[sodium acetate]].||group="note"}} 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. The compound is now banned in the European community and the United States,{{#tag:ref|A minor allowed use in the US is in collars of sheep and cattle to kill predators such as coyotes.||group="note"}} but it is still used in Australia and some other countries. Fluoroacetate deprives cells of energy by replacing acetate in the [[Krebs cycle]], halting a key part of cell metabolism.<ref name="Proudfoot">{{cite journal |author=Proudfoot, A.T.; Bradberry, S.M.; Vale, J.A. |title=Sodium fluoroacetate poisoning |journal=Toxicology Review |volume=25 |issue=4 |pages=213–219 |year=2006 |pmid=17288493 |doi=10.2165/00139709-200625040-00002}}</ref><ref>{{cite report|url = http://www.pwrc.usgs.gov/infobase/eisler/CHR_30_Sodium_monofluoroacetate.pdf|last = Eisler|first = Ronald|title = Biological report 27: Sodium monofluoroacetate (1080) Hazards to fish, wildlife and invertebrates: A synoptic review|publisher = Patuxent Environmental Science Center (U.S. National Biological Service)|year = 1995|accessdate =5 June 2011}}</ref> Several insecticides contain sodium fluoride, which is much less toxic than fluoroacetate.<ref>{{cite web|publisher = Scorecard|title = Class I ozone-depleting substances|work = Sodium fluoride – pesticidal uses|url = http://scorecard.goodguide.com/chemical-profiles/pesticides.tcl?edf_substance_id=7681-49-4| accessdate =20 February 2011}}</ref>

===Scanning===

[[File:PET-MIPS-anim.gif|thumb|upright|Whole-body PET scan using fluorine-18|alt=A rotating, transparent image of a human figure with targeted organs highlighted]]

{{main|Positron emission tomography|Fluorine NMR}}

Compounds containing the manmade isotope fluorine-18 are often used in positron emission tomography (PET) scanning because the isotope's half-life of 110 minutes is relatively long for a positron emitter. One such [[radiopharmaceutical]] is [[fluorodeoxyglucose]] (FDG).<ref>{{cite journal |doi = 10.1097/00002517-200012000-00016 |last1 = Schmitz |pmid = 11132989 |first1 = A. |year = 2000 |pages = 541–544 |issue = 6 |last2 = Kälicke |volume = 13 |journal= Journal of spinal disorders |first2 = T. |last3 = Willkomm |first3 = P. |last4 = Grünwald |first4 = F. |last5 = Kandyba |first5 = J. |last6 = Schmitt |first6 = O.|title=Use of fluorine-18 fluoro-2-deoxy-D-glucose positron emission tomography in assessing the process of tuberculous spondylitis|url=http://www.ecios.org/Synapse/Data/PDFData/0157CIOS/cios-2-167.pdf}}</ref> 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.<ref>{{cite journal|url=http://www.pnas.org/content/74/9/3735.full.pdf+html |title=High aerobic glycolysis of rat hepatoma cells in culture: Role of mitochondrial hexokinase|first1= Ernesto|last1 =Bustamante| first2= Peter L.|last2 =Pedersen|volume = 74|issue = 9|pages = 3735–3739|journal= Proceedings of the National Academy of Sciences|year=1977|pmid=198801|pmc=431708|doi=10.1073/pnas.74.9.3735|bibcode = 1977PNAS...74.3735B}}</ref> [[computed tomography|Computer assisted tomagraphy]] can then be used for detailed imaging.<ref>{{cite book|title=Cancer imaging: Lung and breast carcinomas|publisher=Academic Press|page=41|author=Hayat, M.A.|year=2007|isbn=9780123742124}}</ref>

[[Nuclear magnetic resonance]] (NMR) of fluorine within compounds is a common analysis. Natural fluorine has a single isotope (F-19) and a [[nuclear spin]] of ½. The isotope has a high [[magnetogyric ratio]] along with fast [[Relaxation (NMR)|NMR relaxation]] (in compounds). This translates to fluorine NMR spectra with [[signal-to-noise ratio]]s similar to [[proton NMR]].<ref>{{cite book|title = Nuclear magnetic resonance spectroscopy|first = J.H.|last = Nelson|publisher = Prentice Hall|year = 2003|url = http://books.google.com/?ei=5mlhTambF4fHswbu3_S1CA|isbn = 0-13-033451-0|pages=129–139}}</ref> Fluorine is substituted into proteins so that fluorine NMR can be used to study structure, conformation changes, and metabolism.<ref>{{cite journal|doi =10.1146/annurev.bb.25.060196.001115|title = Use of <sup>19</sup>F NMR to probe protein structure and conformational changes|first1 = Mark A.|last1 = Danielson|first2 = Joseph J.|last2 = Falke|year =1996|journal=Annual Review of Biophysics and Biomolecular Structure|volume =25|pages =163–195|pmid =8800468|pmc =2899692}}</ref> Fluorinated gases have the potential for imaging lung ventilation.<ref>{{cite journal|doi= 10.1002/mrm.1910390114|title = Imaging lungs using inert fluorinated gases|first1 = Dean O.|last1 = Kuethe|first2 = Arvind|last2 = Caprihan|first3 = Eiichi|last3 = Fukushima|first4 = R. Allen|last4 = Waggoner|year =2005|journal=Magnetic Resonance in Medicine|volume =39|pages =85–88|pmid= 9438441|issue= 1}}</ref>

===Other biological aspects===

[[blood substitute|Artificial blood]] and [[liquid breathing]] research make use of perfluorocarbons (PFCs) because they can hold more oxygen or carbon dioxide than blood does.<ref name="Gabrielet">{{cite journal|author=Gabriel, J.L.; Miller, T.F.; Wolfson, M.R. Jr; Shaffer, T.H.|title= Quantitative structure-activity relationships of perfluorinated hetro-hydrocarbons as potential respiratory media. Application to oxygen solubility, partition coefficient, viscosity, vapor pressure, and density|issue=6|volume=42|pages=968–973|year=1996|journal=ASAIO Journal|doi=10.1097/00002480-199642060-00009|pmid=8959271}}</ref> A blood substitute, [[Oxycyte]], has been through initial clinical trials.<ref>{{cite journal| author=Tasker, Fred| date=19 March 2008| url=http://www.miamiherald.com/1057/story/459521.html |title=Miami Herald: Artificial blood goes from science fiction to science fact| archiveurl=http://www.noblood.org/news-hot-topics-such-hepatitis-c-sars-aids/4407-artificial-blood-goes-science-fiction-science-fact.html |archivedate=19 March 2008| journal=Miami Herald ''(at noblood.org)'' }}</ref><ref>{{cite journal|first=Nicole|last=Davis|year=2006|url=http://www.popsci.com/scitech/article/2006-11/better-blood|archiveurl=http://web.archive.org/web/20110604181554/http://www.popsci.com/scitech/article/2006-11/better-blood|archivedate=2011-06-04| title=Better than blood|journal=[[Popular Science]]|accessdate=30 September 2012}}</ref> PFCs have the potential to aid endurance athletes and are therefore banned from sports; the near death of cyclist [[Mauro Gianetti]] was investigated because PFC use was suspected.<ref>{{cite news|url=http://www.nytimes.com/1998/10/18/sports/a-new-threat-in-blood-doping.html|date=18 October 1998|work=The New York Times|title=A New Threat in Blood Doping|author=Gains, Paul}}</ref><ref>{{cite news|url=http://www.salon.com/1999/04/21/cycling/|work=Salon|title=Dying to ride|author=Taber, A.|date=21 April 1999}}</ref> Several [[Liquid breathing#Examples in fiction|fictional treatments of PFC breathing]] exist; the 1989 film ''[[The Abyss]]'' faked [[deep diving|deep sea divers]] breathing PFC but showed a real rat surviving 30 minutes immersion.<ref name=kylstra>{{cite book |title=The feasibility of liquid breathing in man |author=Kylstra, J.A. |year=1977 |publisher=Duke University |url=http://archive.rubicon-foundation.org/4257 |accessdate=5 May 2008}}</ref><ref>{{cite web |url=http://www.experiencefestival.com/liquid_breathing_-_space_travel |title=Liquid breathing – Space travel |accessdate=17 May 2008| publisher=experiencefestival.com |author=The Global Oneness Commitment}}</ref><ref>{{cite journal|author= Harmetz, Aljean|year=1989|url=http://query.nytimes.com/gst/fullpage.html?res=950DE0D7133DF935A3575BC0A96F948260&sec=&spon=&pagewanted=all|title= FILM; 'The Abyss': A foray into deep waters |journal=The New York Times|accessdate=2 October 2012}}</ref> A liquid breathing effort (but with only partial filling of the lungs) by Alliance Pharmaceuticals reached clinical trials but was abandoned.<ref name="Alliance stock drop">{{cite doi|10.1164/rccm.200508-1196OC}}</ref>

[[Fluorine absorption dating]] is a method used to see how long an organic object has been buried. Because groundwater contains fluorine ions, items such as bone that are buried in soil will absorb those ions over time. Therefore, it is possible to determine the [[relative dating|relative age]] of an object by comparing the amount of fluoride with another object found in the same area.<ref>{{cite book|last1=Göksu Murphy|first1 = H.Y.|last2 = Oberhofer|first2 = Martin|last3 = Regulla|first3 = D.F.|title = Scientific Dating Methods|publisher = Kluver Academic Publishers|url = http://books.google.com/?id=maZIlXibg5YC&pg=PA267|year = 1991|page = 267|isbn=978-0-7923-1461-5}}</ref> But if no actual age of any object is known, the ages can only be expressed in terms of one of the objects being older or younger than the other. The fluctuating amount of fluoride found in groundwater means the objects being compared must be in the same local area in order for the comparisons to be accurate. The technique is not always reliable, given that not all objects absorb fluoride at the same rates.<ref>{{cite journal|last1 = Reiche|first1 = I.|title = Fluorine and its relevance for archaeological studies|volume = 2|pages = 253–83|year = 2006|doi = 10.1016/S1872-0358(06)02008-2}}</ref>

==Hazards==

===Fluorine gas===

[[File:DOT hazmat signs - Fluorine.svg|thumb|right|upright=1.6|alt=4 diagonal placards with warnings, poison, corrosive, inhalant, oxidant|The U.S. hazard signs for commercially transported fluorine<ref name="NOAA data sheet">{{cite web|title=UN/NA 1045 (United Nations/North America fluorine data sheet)|url=http://cameochemicals.noaa.gov/unna/1045|publisher=National Oceanic and Atmospheric Administration|accessdate=17 June 2011}}</ref>]]

Elemental fluorine is highly toxic. Above a concentration of 25&nbsp;ppm, it causes significant irritation while attacking the eyes, airways and lungs, and affecting the liver and kidneys. At a concentration of 100&nbsp;ppm, human eyes and noses are seriously damaged.<ref>{{cite journal|last1 = Keplinger|first1 = M.L.|last2 = Suissa|first2 = L.W.|title = Toxicity of fluorine short-term inhalation|journal= American Industrial Hygiene Association Journal|volume = 29|year = 1968|issue = 1|pages = 10–18|pmid = 5667185|doi=10.1080/00028896809342975}}</ref>

===Hydrofluoric acid===

{{See also|Chemical burn}}

Hydrofluoric acid, the water solution of hydrogen fluoride, is a contact poison. It must be handled with extreme care, far beyond that accorded to other dangerous [[mineral acids]], such as [[nitric acid]]. From 1984 to 1994, at least nine U.S. workers died from accidents with HF. Owing to its lesser chemical dissociation in water (remaining a neutral molecule), hydrogen fluoride penetrates tissue more quickly than typical acids. Poisoning can occur readily through the skin or eyes or when inhaled or swallowed.<ref name="OSHA"/>

Once in the blood, hydrogen fluoride reacts with calcium and magnesium, resulting in electrolyte imbalance (potentially [[hypocalcemia]]). The consequent [[cardiac arrhythmia]] may be fatal.<ref name="OSHA">{{cite journal|doi = 10.1002/ajim.1090|title = Fatal unintentional occupational poisonings by hydrofluoric acid in the U.S|year = 2001|last1 = Blodgett|first1 = David W.|last2 = Suruda|first2 = Anthony J.|last3 = Crouch|first3 = Barbara Insley|journal= American Journal of Industrial Medicine|volume = 40|issue = 2|pages = 215–220|pmid = 11494350|url=http://www.chem.purdue.edu/chemsafety/Equip/HFfacts12.pdf}}</ref> Formation of insoluble [[calcium fluoride]] also causes strong pain.<ref name=Goldfrank>{{cite book|title = Goldfrank's manual of toxicologic emergencies|publisher = McGraw-Hill Professional|year = 2007|isbn = 978-0-07-144310-4|page = 1333|url= http://books.google.com/books?id=gWvDBZrTA8MC}}</ref> Burns with areas larger than 160&nbsp;cm<sup>2</sup> (25&nbsp;in<sup>2</sup>), about the size of a man's hand, can cause serious systemic toxicity.<ref name='RMTHFE'>{{cite web|url=http://www51.honeywell.com/sm/hfacid/common/documents/HF_medical_book.pdf |title=Recommended medical treatment for hydrofluoric acid exposure |accessdate=6 May 2009 |pages=2–12 |publisher=Honeywell Specialty Materials}}</ref>

[[File:HF burned hands.jpg|thumb|upright=1.2|alt=left and right hands, two views, burned index fingers|Typical HF burns: the outward signs may not be evident for a day, at which point calcium treatments are less effective.<ref name=Eaton>{{cite web|last=Eaton|first=Charles|title=Figure hfl|url=http://www.eatonhand.com/complic/figures/hfl.htm|work=E-Hand.com: the electronic textbook of hand surgery|publisher=The Hand Center (former practice of Dr. Eaton)|accessdate=28 September 2013}}</ref>]]

Symptoms of exposure to hydrofluoric acid may not be immediately evident, with 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.<ref>{{cite book|title=Clinical environmental health and toxic exposures|publisher= Lippincott Williams & Wilkins|pages=458–459|author= Sullivan, John Burke; Krieger, Gary R.|year=2001|isbn=9780683080278|url=http://books.google.com/books?id=PyUSgdZUGr4C&pg=PA459}}</ref> Hydrofluoric acid exposure is often treated with [[calcium gluconate]], a source of Ca<sup>2+</sup> that binds with the fluoride ions. Skin burns can be treated with a water wash and 2.5% calcium gluconate gel<ref name="pmid2741315">{{cite journal |last1=El Saadi |first1=M.S. |last2=Hall |first2=A.H. |last3=Hall |first3=P.K.|last4=Riggs |first4=B.S. |last5=Augenstein |first5=W.L. |last6=Rumack |first6=B.H. |title=Hydrofluoric acid dermal exposure |journal=Veterinary and human toxicology |volume=31 |issue=3 |pages=243–47 |year=1989 |pmid=2741315}}</ref><ref name="pmid17091088">{{cite journal |last1=Roblin |first1=Isabelle |last2=Urban|first2=Martine |last3=Flicoteau |first3=Domitille |last4=Martin |first4=Chantel |last5=Pradeau |first5=Dominique |title=Topical treatment of experimental hydrofluoric acid skin burns by 2.5% calcium gluconate |journal=Journal of Burn Care & Research |volume=27 |issue=6 |pages=889–894 |year=2006 |pmid=17091088|doi=10.1097/01.BCR.0000245767.54278.09}}</ref> or special rinsing solutions.<ref name="pmid15461243">{{cite journal |last1=Hultén |first1=Peter |last2=Höjer|first2=J. |last3=Ludwigs |first3=U. |last4=Janson |first4=A. |title=Hexafluorine vs. standard decontamination to reduce systemic toxicity after dermal exposure to hydrofluoric acid |journal=Journal of Toxicology – Clinical Toxicology |volume=42 |issue=4 |pages=355–361 |year=2004 |pmid=15461243 |doi=10.1081/CLT-120039541}}</ref>

Because HF is absorbed, further medical treatment is necessary. Calcium gluconate may be injected or administered intravenously. Use of calcium chloride is contraindicated and may lead to severe complications. Sometimes surgical excision of tissue or amputation is required.<ref name="RMTHFE"/><ref>{{cite book|title=Handbook of Quality Integrated Circuit Manufacturing|last=Zorich|first=Robert|pages=182–183|url=http://books.google.com/books?id=sdB3KjhWzFQC&pg=PA183|publisher=Academic Press|year=1991|isbn=9780323140553}}</ref>

===Fluoride ion===

{{See also|Fluoride toxicity}}

Soluble fluorides are moderately toxic. For sodium fluoride, the lethal dose for adults is 5–10&nbsp;g, which is equivalent to 32–64&nbsp;mg of elemental fluoride per kilogram of body weight.<ref name="zzzzzzz23">{{cite book|url=http://www.inchem.org/documents/ehc/ehc/ehc227.htm|author = Liteplo, R.; Gomes R.; Howe, P.; Malcolm, H.|title = Environmental health criteria 227 (Fluoride)|publisher = United Nations Environment Programme; International Labour Organization; World Health Organization|year = 2002|page =100| isbn =92-4-157227-2}}</ref> The dose that may lead to adverse health effects is about one fifth the lethal dose.<ref name="Nochimson"/> Chronic excess fluoride consumption can lead to [[skeletal fluorosis]], a disease of the bones that affects millions in Asia and Africa.<ref name="Nochimson"/><ref name="pmid19305069">{{cite journal |author=Reddy, D.R. |title=Neurology of endemic skeletal fluorosis |journal=Neurology India |volume=57 |issue=1 |pages=7–12 |year=2009 |pmid=19305069 |doi= 10.4103/0028-3886.48793}}</ref>

The fluoride ion is readily absorbed by the stomach and intestines. Ingested fluoride forms hydrofluoric acid in the stomach. In this form, fluoride crosses cell membranes and then binds with calcium and interferes with various enzymes. Fluoride is [[excretion|excreted]] through urine. Fluoride exposure limits are based on urine testing which has determined the human body's capacity for ridding itself of fluoride.<ref name="Nochimson">{{cite web|last =Nochimson|first =Geofrey |title =Fluoride toxicity|publisher = [[emedicine]]|year = 2011|url =http://emedicine.medscape.com/article/814774-overview|accessdate =19 May 2011}}</ref><ref>{{cite journal|last1 = Baez|first1 = Ramon J.|last2 = Baez|first2 = Martha X.|last3 = Marthaler|first3 = Thomas M.|title = Urinary fluoride excretion by children 4–6 years old in a south Texas community| journal= Revista Panamericana de Salud Pública|volume = 7|year = 2000|doi = 10.1590/S1020-49892000000400005|issue = 4|pages = 242–248}}</ref>

Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluoride,<ref>{{cite journal|last1=Augenstein|first1=W.L. 
|first2=David G.|last2=Spoerke|first3=Ken W.|last3=Kulig|displayauthors=2 |title = Fluoride ingestion in children: a review of 87 cases|url = http://pediatrics.aappublications.org/cgi/content/abstract/88/5/907|year = 1991|journal= Pediatrics |pages = 907–912|volume = 88|issue = 5|pmid = 1945630}}</ref> Currently, most calls to poison control centers for possible fluoride poisoning come from the ingestion of fluoride-containing toothpaste.<ref name="Nochimson"/> Malfunction of water fluoridation equipment has occurred several times, including an Alaskan incident which sickened nearly 300 people and killed one.<ref name=Gessner94>{{cite journal|last1 = Gessner|first1 = Bradford D.|last2 = Beller|first2 = Michael|last3 = Middaugh|first3 = John P.|last4 = Whitford|first4 = Gary M.|title = Acute fluoride poisoning from a public water system|journal= [[New England Journal of Medicine]]|year=1994|volume = 330|issue = 2|pages = 95–99|pmid = 8259189|doi = 10.1056/NEJM199401133300203}}</ref>

==Environmental concerns==

===Atmosphere===

{{See also|Ozone depletion|global warming}}

Because they deplete the ozone layer, chlorofluorocarbons (CFCs) and [[bromofluorocarbon]]s (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 then release chlorine and bromine atoms which attack ozone molecules.<ref name="UNEP FAQ Ozone">{{cite web|author=Aucamp, Pieter J.; Björn, Lars Olof|title=Questions and answers about the environmental effects of the ozone layer depletion and climate change: 2010 update|url=http://ozone.unep.org/Assessment_Panels/EEAP/eeap-report2010-FAQ.pdf|publisher=United Nations Environmental Programme|accessdate=6 January 2012|pages=4–6, 41, 46–47}}</ref> Predictions are that generations will be required, even after the CFC ban, for these molecules to leave the atmosphere and for the ozone layer to fully recover. Early indications are that the CFC ban is working: ozone depletion has stopped, and recovery has started.<ref>{{cite doi|10.1038/news.2011.293}}</ref><ref>{{cite web|title=Good news and a puzzle|url=http://science.nasa.gov/science-news/science-at-nasa/2006/26may_ozone/|accessdate=6 January 2012|author1=Barry, Patrick L.|author2=Phillips, Tony|publisher=NASA|date=2006-05-26}}</ref>

[[File:Future ozone layer concentrations.gif|thumb|upright=1.4|Animation|alt=An animation showing colored representation of ozone distribution by year, above North America, through 6 steps. It starts with a lot of ozone especially over Alaska and by 2060 is almost all gone from north to south.|NASA projection of stratospheric ozone levels over North America if CFCs had not been banned]]

Hydrochlorofluorocarbons (HCFCs) are current replacements for CFCs, with about one tenth the ozone damaging potential (ODP).<ref name="EPA OLP Science">{{cite web|title=Class I Ozone-depleting Substances|url=http://www.epa.gov/ozone/science/ods/classone.html|publisher=U.S. Environmental Protection Agency|accessdate=6 January 2012}}</ref> They were originally scheduled for elimination by 2030 in developed nations (2040 in undeveloped). In 2003, the U.S. [[Environmental Protection Agency]] prohibited production of one HCFC and capped the production of the two others.<ref>{{cite web|url=http://www.epa.gov/ozone/title6/phaseout/classtwo.html|title=Phaseout of HCFCs (Class II Ozone-Depleting Substances)|publisher=U.S. Environmental Protection Agency|accessdate=15 August 2012|year=2012}}</ref> In 2007, a new treaty was signed by almost all nations to move HCFC phaseout up to 2020 because HFCs, which have no chlorine and thus zero ODP, are available.<ref name="HCFC Elim speedup">{{cite journal|last=McCoy|first=Michael|title=Nations speed up HCFC elimination|journal=Chemical & Engineering News|year=2007|volume=85|issue=40|page=11|url=http://pubs.acs.org/cen/news/85/i40/8540notw8.html|accessdate=25 December 2012|doi=10.1021/cen-v085n023.p011a}}</ref>

Fluorocarbon gases of all sorts (CFCs, HFCs, etc.) are [[greenhouse gases]] about 4,000 to 10,000&nbsp;times as potent as carbon dioxide. Sulfur hexafluoride exhibits an even stronger effect, about 20,000 times the global warming potential of carbon dioxide.<ref name="greenhouse">{{cite book|chapter = Changes in Atmospheric Constituents and in Radiative Forcing.|title = Climate change 2007: The physical science basis. Contribution of Working Group I to the fourth assessment report of the Intergovernmental Panel on Climate|isbn = 978-0-521-70596-7|publisher = Cambridge University|author = [[Intergovernmental Panel on Climate Change]]|url = http://www.ipcc.ch/pdf/assessment-report/ar4/wg1/ar4-wg1-chapter2.pdf|pages = 208&ndash;216}}</ref>

===Biopersistance===

Because of the strength of the carbon–fluorine bond, organofluorines endure in the environment. [[Perfluorooctanoic acid]] (PFOA) and [[perfluorooctanesulfonic acid]] (PFOS), used in waterproofing sprays, are [[persistent organic pollutant|persistent global contaminants]]. Trace quantities of these substances have been detected worldwide, from polar bears in the Arctic to the global human population. PFOA has been detected in breast milk and the blood of newborns. One study indicates that PFOS levels in wildlife are starting to go down because of reducing production.<ref name="PFOA"/><ref name="PFOA evidence"/>

In the body, PFOA binds to a protein, [[serum albumin]]. PFOA's tissue distribution in humans is unknown, but studies in rats suggest it is present mostly in the liver, kidney, and blood. PFOA is not metabolized by the body, but is excreted by the kidneys.<ref name="PFOA"/><ref name="PFOA evidence"/>

The potential health effects of PFOA are unclear. Unlike chlorinated hydrocarbons, PFOA is not [[lipophilic]] (stored in fat), nor is it [[Genotoxicity|genotoxic]] (damaging genes). While both PFOA and PFOS cause cancer in high quantities in animals, studies on exposed humans have not been able to prove an impact at current exposures. [[Bottlenose dolphin]]s have some of the highest PFOS concentrations of any wildlife studied; one study suggests an impact on their immune systems.<ref name="PFOA">{{cite journal |last1 = Steenland |first1 = Kyle |last2 = Fletcher |first2 = Tony |last3 = Savitz |first3 = David A. |title = Epidemiologic evidence on the health effects of perfluorooctanoic acid (PFOA) |doi = 10.1289/ehp.0901827 |pmc = 2920088 |pmid = 20423814 |year = 2010 |issue = 8 |volume = 118 |journal= Environmental Health Perspectives|url = http://ehp03.niehs.nih.gov/article/fetchArticle.action?articleURI=info%3Adoi%2F10.1289%2Fehp.0901827|accessdate=3 May 2011 |pages = 1100–8}}</ref><ref name="PFOA evidence">{{cite journal|last=Betts|first=Kellyn S.|title=Perfluoroalkyl acids: What is the evidence telling us?|journal=Environmental Health Perspectives|year=2007|volume=115|issue=5|pages=A250–A256|doi=10.1289/ehp.115-a250|url=http://ehp03.niehs.nih.gov/article/info%3Adoi%2F10.1289%2Fehp.115-a250|accessdate=5 September 2011|pmid=17520044|pmc=1867999}}</ref>

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.<ref>{{cite report|title = Evaluation of emerging contaminants of concern at the south district wastewater treatment plant based on seasonal sampling events, Miami-Dade Country, Florida, 2004|author=Lietz, A.C.; Meyer, Michael T.|publisher = U.S. Geological Survey Scientific Investigations|year = 2006|accessdate =6 June 2011|url=http://pubs.usgs.gov/sir/2006/5240/pdf/sir2006-5240.pdf|pages=7–8|url=http://pubs.usgs.gov/sir/2006/5240/pdf/sir2006-5240.pdf}}</ref> Fluorine-containing agrichemicals are measurable in farmland [[Surface runoff|runoff]] and nearby rivers.<ref>{{cite doi|10.1039/c0em00373e}}</ref>
{{-}}

==Compounds==

{{main|Compounds of fluorine}}

Fluorine's common [[oxidation state]] is −1.{{#tag:ref|It differs from this value in elemental fluorine, where the atoms are bonded to each other and thus at oxidation state 0. The very unstable anions F<sub>2</sub><sup style="margin-left:-0.75ex;">-</sup> and F<sub>3</sub style="margin-left:-1ex;"><sup>-</sup> with intermediate oxidation states exist at very low temperatures, decomposing at around 40&nbsp;K.{{sfn|Wiberg|Wiberg|Holleman|2001|p=422}} The F<sub>4</sub><sup style="margin-left:-0.75ex;">+</sup> cation and a few related species have been predicted to be stable.<ref>{{cite doi|10.1039/C1RA00804H}}</ref>|group="note"}} Because of its attraction for electrons, fluorine forms polar covalent bonds or ionic bonds to other atoms. Covalent bonds involving fluorine atoms are almost always [[single bond]]s,<ref group="note">[[Boron monofluoride]] and [[nitrogen monofluoride]] are metastable compounds with higher than single bonded fluorine. Boron monofluoride is stable as a ligand in some coordination compounds.</ref><ref name="NF charge">{{cite pmid|11792193}}</ref> although fluorine may act as a [[bridging ligand]] between metals in some [[metal complex]]es. Molecules containing fluorine may also exhibit [[hydrogen bonding]].<ref>{{cite book|title=Organofluorine chemistry: Principles and commercial applications|page=515|publisher=Springer|author=Smart, Bruce E.; Tatlow, J.C.|year=1994|isbn=0306446103}}</ref> 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.<ref group="note" name="metal">In this article, metalloids are lumped with the definite main group nonmetals because the fluoride chemistry is similar. The noble gases are treated [[#Noble gas compounds|separately]]. Hydrogen is discussed in the [[#Hydrogen fluoride|Hydrogen fluoride]] section; carbon in the [[#Organic compounds|Organic compounds]] section. The most recently created heavy elements have not been studied and thus are not included. This is illustrated by the image to the right: the dark gray elements are metals, the green ones are nonmetals, the light blue ones are the noble gases, the purple one is hydrogen, the yellow one is carbon, and the light gray elements have unknown properties.</ref><ref>{{cite doi|10.1016/j.ccr.2008.07.014}}</ref>

===Hydrogen fluoride===

{{Main|Hydrogen fluoride|Hydrofluoric acid}}

[[File:Boiling-points Chalcogen-Halogen.svg|upright=1.4|alt=graph showing trend-breaking water and HF boiling points: big jogs up versus a trend that is down with lower molecular weight for the other series members.|Boiling points of the hydrogen halides (blue) and hydrogen chalcogenides (red): HF and H<sub>2</sub>O break trends.|thumb|right]]

Fluorine combines with hydrogen to make a compound called hydrogen fluoride (HF) or, in the context of water solutions, hydrofluoric acid. HF molecules cluster together weakly via hydrogen bonds. Because of this, hydrogen fluoride behaves more like [[water]] than like [[hydrogen chloride|HCl]] (hydrochloric acid).<ref name="Pauling HF hydrogen bonds">{{cite book|last=Pauling|first=Linus A.|title=The nature of the chemical bond and the structure of molecules and crystals: An introduction to modern structural chemistry|year=1960|publisher=Cornell University Press|isbn=978-0-8014-0333-0|pages=454–464|url=http://books.google.com/books?id=L-1K9HmKmUUC&pg=PA454}}</ref><ref name="Atkins HF">{{cite book|last=Atkins|first=Peter|title=Chemical principles: The quest for insight|year=2008|publisher=W. H. Freeman & Co|isbn=978-1-4292-0965-6|pages=184–185|url=http://books.google.com/books?id=4R6hb1OIMRUC&pg=PA184|last2=Jones|first2=Loretta}}</ref><ref name="New Scientist HF">{{cite journal|last=Emsley|first=John|title=The hidden strength of hydrogen|journal=New Scientist|year=1981|volume=91|issue=1264|pages=291–292|url=http://books.google.com/books?id=ZbthaZCUXy4C&pg=PA292}}</ref> Hydrogen fluoride boils at a much higher temperature than the heavier hydrogen halides. HF is also fully [[miscible]] with water (dissolves in any proportion), unlike HCl, [[Hydrogen bromide|HBr]], or [[Hydrogen iodide|HI]].{{sfn|Greenwood|Earnshaw|1998|pp=812–816}}

In water solution, hydrogen fluoride is a [[weak acid]]; the other hydrohalic acids are all [[strong acid|strong]].{{sfn|Wiberg|Wiberg|Holleman|2001|p=425}}{{#tag:ref|For more detail, see explanation by Jim Clark.<ref>{{cite web|last=Clark|first=Jim|title=The acidity of the hydrogen halides| url=http://www.chemguide.co.uk/inorganic/group7/acidityhx.html|accessdate=4 September 2011|year=2002}}</ref>|group="note"}} Hydrofluoric acid is non-ideal: instead of having a constant [[acid dissociation constant]] (PKa), HF's inherent acidity increases at higher concentrations through a phenomenon called [[homoassociation|homoconjugation]].<ref name="H+">{{cite book|author=Chambers, C.; Holliday, A.K.|title=Modern inorganic chemistry (An intermediate text)|year=1975|publisher=The Butterworth Group|pages=328–329|url=http://files.rushim.ru/books/neorganika/Chambers.pdf}}</ref><ref>{{cite book|author=Hannan, Henry J.|title=Course in chemistry for IIT-JEE 2011|year=2010|publisher=Tata McGraw Hill Education Private Limited|isbn=9780070703360|pages=15–22|url=http://books.google.com/books?id=wNYMUeCcaSEC&pg=SA15-PA22}}</ref> Although hydrofluoric acid is weak, it is very corrosive, even attacking glass.<ref name="H+"/>

===Metal fluorides===

{{See also|Fluoride volatility}}

Metal fluorides share some similarities with other [[metal halide]]s but are more ionic, often more like [[metal oxide|oxides]] in their bonding and crystal structures.{{sfn|Greenwood|Earnshaw|1998|p=819}} However, there is a trend of the binary metal fluorides becoming more covalently bonded and volatile as the amount of fluorination increases.<!--we don't need a cn tag here, it's a generalization supported by many examples below--> For example, [[sodium monofluoride]] boils at 1704&nbsp;°C (3099&nbsp;°F),{{sfn|Lide|2004|p=4.84}} whereas [[rhenium heptafluoride]] has a boiling point below water's.{{sfn|Lide|2004|p=4.79, 4.91}} Similar to fluorine itself, the higher metal fluorides are very reactive. [[Platinum hexafluoride]] was the first compound to oxidize molecular oxygen<ref>{{cite journal| title=Dioxygenyl hexafluoroplatinate (V), O<sub>2</sub><sup style="margin-left:-0.75ex;">+</sup>[PtF<sub>6</sub>]<sup>−</sup>| author= Bartlett, Neil; Lohmann, D. H. | journal= Proceedings of the Chemical Society | publisher = Chemical Society | issue = 3 | page = 115 | year = 1962 | doi = 10.1039/PS9620000097}}</ref> and xenon.<ref>{{cite journal | title = Xenon hexafluoroplatinate (V) Xe<sup>+</sup>[PtF<sub>6</sub>]<sup>−</sup> | author = Bartlett, Neil | journal= Proceedings of the Chemical Society | publisher = Chemical Society | issue = 6 | page = 218 | year = 1962 | doi = 10.1039/PS9620000197}}</ref> The solubility of fluorides varies greatly but tends to decrease with more fluorine.<ref>{{cite book|title=Principle of modern chemistry|author=Oxtoby, David W.; Gillis, H. Pat; Campion, Alan|page=693|publisher= Cengage learning|year=2012|isbn=9780840049315}}</ref>

The alkali metals form [[monofluoride]]s that, like the alkali metal chlorides, are very ionic and soluble. They have the same atomic arrangement—[[Cubic crystal system#Rock-salt structure|the rock salt crystal structure]]—as sodium chloride.<ref>{{cite book|title=Mesoscopic materials and clusters: Their physical and chemical properties|url=http://books.google.com/books?id=pjFTI5KgMFkC&pg=PA267|page=267|publisher=Springer|author=Arai, Toshihiro|year=1999|isbn=9783540648840}}</ref>{{sfn|Aigueperse et al.|2005|pp=25–27}} Other metal monofluorides show clear difference versus their chlorides. For example, thallium<ref>{{cite book| last = Remy | first = Heinrich | year = 1956 | title = Treatise on inorganic chemistry: Introduction and main groups of the periodic table | publisher = Elsevier | page = 383}}</ref> and silver<ref name="dissolve"/> monofluorides are soluble, while the chlorides are not.

The [[difluoride]]s of the alkaline earths are also very ionic, but are generally very insoluble. In contrast, the alkali earth dichlorides are fairly soluble.<ref name="dissolve">{{cite book|pages=278–280|title=First outlines of a dictionary of solubilities of chemical substances| author= Storer, Frank Humphreys| publisher= Cambridge University Press| year=1864| isbn=978-1-176-62256-2}}</ref> [[Beryllium difluoride|Beryllium's difluoride]] differs from those of the other alkaline earths: the bonding has some covalent character, the structures are similar to [[silicon dioxide|SiO<sub>2</sub>]] (quartz), and the compound is water soluble.<ref name="Beryllium chemistry">{{cite book|last=Walsh|first=Kenneth A|title=Beryllium chemistry and processing|publisher=ASM International|isbn=978-0-87170-721-5|pages=99–102, 118–119|url=http://books.google.com/books?id=3-GbhmSfyeYC&pg=PA119|year=2009}}</ref>

[[File:AuF3structure-2.png|thumb|right|Spiral chains in gold trifluoride|upright=1.2]]

Many metals form [[trifluoride]]s, including iron, bismuth, the [[rare earth element|rare earths]], some [[actinide]]s, and the metals in the aluminium and scandium columns of the periodic table. In general, the metal trifluorides have ionic bonding structures with octahedral coordination of the metal cations. Some show the same structure as [[rhenium trioxide]]. The larger metal atoms can have dense structures with more than six surrounding fluorines. In contrast, [[gold trifluoride|gold's trifluoride]] has an open structure with helical chains of [[square planar]] AuF<sub>4</sub>.{{sfn|Emeléus|Sharpe|1983|pp=89–97}}<ref name="Crystal chemistry of fluorides">{{cite book|last1=Babel|first1=Dietrich|last2=Alain|first2=Tressaud|title=Inorganic solid fluorides: chemistry and physics|publisher=Academic Press Inc.|editor=Poul Hagenmuller|location=Orlando|chapter=Crystal chemistry of fluorides|pages=91–96|year=1985}}</ref><ref>{{cite journal | author = Einstein, F.W.B.; Rao, P.R.; Trotter, James; Bartlett, Neil | title = The crystal structure of gold trifluoride | year = 1967 | journal = Journal of the Chemical Society A: Inorganic, Physical, Theoretical | volume = 4 | pages = 478–482 | doi = 10.1039/J19670000478 }}</ref>

The [[tetrafluoride]]s show a mixture of ionic and covalent bonding. Zirconium<ref>{{cite book|page=144| url=http://books.google.com/?id=DvqwTdVhjMEC&pg=PA144| title=Chemical Thermodynamics of Zirconium| author= Brown, Paul L.; Mompean, Federico J.; Perrone, Jane; Illemassène, Myriam| publisher= Gulf Professional Publishing| year=2005| isbn=0-444-51803-7}}</ref> and hafnium,<ref>{{cite book |last=Perry |first=D.L. |title=Handbook of Inorganic Compounds|year=2011 |publisher=CRC Press |edition=2nd |isbn=9781439814611 |page=193}}</ref> along with several actinides,<ref>{{cite doi|10.1063/1.467963}}</ref> form tetrafluorides with an ionic structure based on 8-coordinate metal cations.{{sfn|Lide|2004|pp=4.60, 4.76, 4.92, 4.96}} Melting points are around 1000&nbsp;°C (1850&nbsp;°F).{{#tag:ref|ZrF<sub>4</sub> melts at 932&nbsp;°C (1764&nbsp;°F).{{sfn|Lide|2004|p=4.96}} HfF<sub>4</sub> sublimes at 968&nbsp;°C.<ref>{{cite book|title=Handbook of Inorganic Compounds|edition=2nd|publisher=CRC Press|author=Perry, Dale|page=193}}</ref> UF<sub>4</sub> melts at 1036&nbsp;°C (1897&nbsp;°F).{{sfn|Lide|2004|p=4.92}}|group="note"}} On the other hand, the tetrafluorides of titanium,{{sfn|Greenwood|Earnshaw|1998|p=964}}, vanadium,<ref>{{cite doi|10.1002/anie.199004061}}</ref> niobium{{sfn|Greenwood|Earnshaw|1998|p=990}} are polymeric. The first melts at 284&nbsp;°C (543&nbsp;°F);{{sfn|Lide|2004|p=4.91}} the latter two decompose at around 350&nbsp;°C (650&nbsp;°F).{{sfn|Lide|2004|p=4.72, 4.93}}

The [[pentafluoride]]s are even more covalently bonded, forming low dimensionality polymers or [[oligomer]]ic molecules.<ref name=AuF5/> [[Bismuth pentafluoride|Bismuth]] and [[uranium pentafluoride]]s share the same structure: straight chains of connected octahedra.{{sfn|Greenwood|Earnshaw|1998|pp=561–563}}{{sfn|Emeléus|Sharpe|1983|pp=256–277}}{{sfn|Mackay|Mackay|Henderson|2002|pp=243–244}} [[Gold(V) fluoride|Gold's pentafluoride]] is a [[dimer (chemistry)|dimer]] (doubled molecule), Au<sub>2</sub>F<sub>10</sub>.<ref name=AuF5>{{cite doi|10.1002/1521-3773(20011001)40:19<3690::AID-ANIE3690>3.0.CO;2–5}}</ref> [[Niobium pentafluoride|Niobium]] and [[tantalum pentafluoride]]s are [[tetramer]]ic molecules (fourfold clusters).<ref name=AuF5/> The former melts at 80&nbsp;°C (176&nbsp;°F),{{sfn|Lide|2004|p=4.72}} the latter at 95&nbsp;°C (203&nbsp;°F).{{sfn|Lide|2004|p=4.88}}

There are thirteen metal [[hexafluoride]]s.{{#tag:ref|[[Iridium hexafluoride|IrF<sub>6</sub>]], [[molybdenum hexafluoride|MoF<sub>6</sub>]], [[osmium hexafluoride|OsF<sub>6</sub>]], [[Neptunium hexafluoride|NpF<sub>6</sub>]], [[polonium hexafluoride|PoF<sub>6</sub>]], [[plutonium hexafluoride|PuF<sub>6</sub>]], [[platinum hexafluoride|PtF<sub>6</sub>]], [[Rhenium hexafluoride|ReF<sub>6</sub>]], [[rhodium hexafluoride|RhF<sub>6</sub>]], [[ruthenium hexafluoride|RuF<sub>6</sub>]], [[technetium hexafluoride|TcF<sub>6</sub>]], [[uranium hexafluoride|UF<sub>6</sub>]], and [[tungsten hexafluoride|WF<sub>6</sub>]]|group="note"}} All are octahedral molecules.<ref name="D_BLOCK_XF6">{{cite journal |author= Drews, T.; Supeł, J.; Hagenbach, A.; Seppelt, K. |title= Solid state molecular structures of transition metal hexafluorides |journal= Inorganic Chemistry |year= 2006 |volume= 45 |issue= 9 |pages= 3782–3788 |doi= 10.1021/ic052029f |pmid= 16634614 }}</ref> At room temperature [[tungsten hexafluoride]] is a gas.{{sfn|Lide|2004|p=4.92}} [[Molybdenum hexafluoride]] and [[rhenium hexafluoride]] are liquids.{{sfn|Lide|2004|p=4.71, 4.78}} The rest are volatile solids.{{sfn|Greenwood|Earnshaw|1998|loc=(various pages, by individual metal in respective chapter)}}

The only definite metal [[heptafluoride]], that of rhenium, is a low-melting molecular solid. The structure is a distorted [[pentagonal bipyramid]]: the five equatorial fluorines are buckled slightly above and below the plane.{{sfn|Greenwood|Earnshaw|1998|p=819}} In 1966, osmium heptafluoride was reported from daunting synthesis conditions (400 atmospheres of fluorine gas at 600&nbsp;°C), but the product was not structurally characterized. In 2006, the synthesis was re-attempted but only osmium hexafluoride was observed.<ref>{{cite doi|10.1021/ic0608290}}</ref> Previously, a 1913 report of an [[octafluoride]], that of osmium, was shown in 1958 to be a mistaken identification of the same compound, osmium hexafluoride.<ref name="fluorides">{{ cite journal | last1 = Riedel | first1 = S. | last2 = Kaupp | first2 = M. | title = The highest oxidation states of the transition metal elements | journal = Coordination Chemistry Reviews | year = 2009 | volume = 253 | issue = 5–6 | doi = 10.1016/j.ccr.2008.07.014 | pages = 606–624 | url = http://144.206.159.178/ft/243/588116/14862785.pdf }}</ref><ref name="Osmium octafluoride and hexafluoride">{{cite journal|last=Weinstock|first=Bernard|coauthors=Malm, John G.|title=Osmium Hexafluoride and its Identity with the Previously Reported Octafluoride|journal=Journal of the American Chemical Society|year=1958|month=September|volume=80|issue=17|pages=4466–4468|url=http://pubs.acs.org/doi/abs/10.1021/ja01550a007|doi=10.1021/ja01550a007}}</ref>

{| cellpadding="3" style="margin:1em auto 1em auto; text-align:center;"
|-
| colspan="3" style="text-align:center;"| '''Progression of structure type with metal charge in the 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|Sodiun fluoride: checkerboard-like lattice of small sodium and large fluorine atoms, going in three dimensions so that no "molecule" can be isolated]]
| [[File:Bismuth-pentafluoride-chain-from-xtal-1971-3D-balls.png|500px|alt=a straight chain of alternating balls, violet and yellow, with violet ones linked additionally to four more yellow perpendicularly to the chain and each other|Bismuth pentafluoride: a straight chain of atoms forming an infinitely long molecule]]
| [[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: a discrete small molecule]]
|- style="text-align:center;"
| Sodium fluoride, ionic
| Bismuth pentafluoride, polymeric
| Rhenium heptafluoride, molecular
|}

===Nonmetal fluorides===

The binary fluorides of the main group nonmetals and metalloids are generally volatile, covalently bonded molecules. The fluorides of those elements in the [[Period 2|second row of the periodic table]] (carbon, nitrogen, and oxygen) follow the [[octet rule]]. Boron is an exception—its fluoride has an [[incomplete octet]] (less bonds than normal). Nonmetals from [[period 3|the third row of the periodic table]] and below can form fluorides which are [[hypervalent molecule|hypervalent]] (more bonds than normal).<ref name=Noury>{{cite journal | url = http://alpha.chem.umb.edu/chemistry/Seminar/06-09%20WQE/InorgI.Carter.pdf | title = Chemical bonding in hypervalent molecules: Is the octet rule relevant? | journal= Inorganic Chemistry | year = 2002 | volume = 41 | pages = 2164–2172 | accessdate = 23 May 2012| author = Noury, Stephane; Silvi, Bernard; Gillespie, Ronald J. | doi = 10.1021/ic011003v | issue = 8 | pmid = 11952370}}</ref> The discussion is organized by [[Group (periodic table)|periodic table column]].

[[Boron trifluoride]] is a planar molecule. It is a weak [[Lewis acid]] and readily accepts a [[Lewis base]], forming [[adduct]]s (combinations). The lone pair of a Lewis base, such as ammonia, allows the boron to complete its octet. With another fluoride ion (F<sup>−</sup> is a Lewis base), BF<sub>3</sub> reacts to form the relatively unreactive {{chem|BF|4|-}} anion.{{sfn|Greenwood|Earnshaw|1998|pp=375–376}}

[[File:Selenium-tetrafluoride-gas-3D-balls.png|thumb|The unusual seesaw shape of sulfur and selenium tetrafluoride]]

[[Silicon tetrafluoride]] and [[germanium tetrafluoride]] adopt a [[tetrahedral]] structure<ref>{{cite book|title=Scientific essentialism|page=69|author=Ellis, Bryan David|publisher=Cambridge University Press|year=2001|isbn=0521800943}}</ref> and are Lewis acids.{{sfn|Aigueperse et al.|2005|p=28}}{{sfn|Wiberg|Wiberg|Holleman|2001|p=897}} Silicon forms a variety of fluorides (more diversity than the [[silanes]]). With excess silicon, gaseous [[silicon difluoride|SiF<sub>2</sub>]] is made, which condenses into a polymer. Forms of Si<sub>n</sub>F<sub>2n+2</sub> are liquids and solids and are known up to Si<sub>14</sub>F<sub>30</sub>.{{sfn|Greenwood|Earnshaw|1998|pp=376–377}} Reaction of excess germanium with its tetrafluoride also creates a [[Germanium difluoride|difluoride]]. The compound is a low-melting solid with spiral chains of linked GeF<sub>3</sub> units. Germanium also forms a [[Inner sphere electron transfer#Mixed valence compounds|mixed valence compound]] (i.e. with both +4 and +2 Ge atoms): [(GeF<sub>2</sub>)<sub>4</sub>GeF<sub>4</sub>].{{sfn|Greenwood|Earnshaw|1998|pp=198–199}}

The [[pnictogen]]s (nitrogen's column) form trifluorides that are weak Lewis bases and are more reactive as the pnictogen becomes heavier.<ref name="PnFx"/> However, nitrogen trifluoride is stable against hydrolysis and is not a Lewis base.<ref name="PnFx">{{cite book |title= Concepts and problems in inorganic Chemistry |last1= Raghavan|first1=P.S. |year=1998 |publisher=Discovery Publishing House |isbn=9788171414185 |pages=164–165 |url=http://books.google.com/books?id=pBiS0jc-kWIC&pg=PA164}}</ref> In comparison to the trifluorides, the pnictogen pentafluorides are even more reactive.{{sfn|Aigueperse et al.|2005|p=37}}<ref name=Arsenic/> Antimony pentafluoride is the strongest Lewis acid of all charge-neutral compounds.<ref name=Arsenic>{{cite book |title= Chemistry of arsenic, antimony and bismuth |last1= Norman|first1=Nicholas C. |year=1998 |publisher=Springer |isbn=075140389X |page=97 |url=http://books.google.com/books?id=vVhpurkfeN4C&pg=PA98}}</ref> The trifluorides have a [[Trigonal pyramidal molecular geometry|trigonal pyramid structure]] (ammonia's shape). The pentafluorides are generally [[Trigonal bipyramidal molecular geometry|trigonal bipramid]]al. However, the liquid and solid phases of SbF<sub>5</sub> have more complicated structures with octahedral antimony and bridging fluorines.{{sfn|Greenwood|Earnshaw|1998|pp=561–563}}  

The [[chalcogen]]s (oxygen's column) form a variety of fluorides. Unstable difluorides are known for oxygen (the only compound where oxygen is at oxidation state +2) as well as sulfur and selenium. The geometry is [[bent molecular geometry|bent]] (similar to water). Sulfur and selenium tetrafluorides show the rarece [[seesaw molecular geometry]], but solid TeF<sub>4</sub> is a polymer. All the tetrafluorides are thermally unstable and hydrolyze. The hexafluorides are the result of direct fluorination of sulfur, selenium, and tellurium and are octahedral molecules. SF<sub>6</sub> is extremely inert, while SeF<sub>6</sub> and TeF<sub>6</sub> show increasingly higher reactivity. Various chalcofluoride molecules exist with two chalcogen atoms also: O<sub>2</sub>F<sub>2</sub> ("FOOF"), S<sub>2</sub>F<sub>10</sub>, etc.<ref name="(O,S,Se,Te)Fn">{{cite book|title=University chemistry, Vol. 1|publisher=New Age International|pages=180–182, 206–208|isbn=8122407420|author=Murthy, C. Parameshwara}}</ref>{{sfn|Greenwood|Earnshaw|1998|pp=638–640, 683–689, 767–778}} 

The well-characterized heavier halogens (chlorine, bromine, and iodine) all form mono-, tri-, and pentafluorides: XF, XF<sub>3</sub>, and XF<sub>5</sub>. Of the neutral +7 species, only [[iodine heptafluoride]] is known.{{sfn|Wiberg|Wiberg|Holleman|2001|p=435}} The corresponding cations, {{chem|ClF|6|+}} and {{chem|BrF|6|+}}, are known and are extremely strong oxidizers.{{sfn|Wiberg|Wiberg|Holleman|2001|p=436}} For the [[radioactive]] [[astatine]] only the non-volatile astatine monofluoride has been studied,<ref name=astatineradon>{{cite book | title = Molecular structure and statistical thermodynamics: Selected papers of Kenneth S. Pitzer, Vol. 1 | publisher = World Scientific | year = 1993 | page = 111 | isbn = 9810214391| editor = Pitzer, Kenneth Sanborn}}</ref> but its existence is debated.<ref>{{cite book|title=Gmelin handbook of inorganic chemistry: At—Astatine|edition=8th|publisher=Springer-Verlag|page=224|isbn=9783540935162|author=Gmelin, Leopold}}</ref> Many of the halogen fluorides are powerful fluorinators (sources of fluorine atoms). ClF<sub>3</sub> readily fluorinates asbestos and refractory oxides, and industrial use requires precautions similar to those for fluorine gas.{{sfn|Greenwood|Earnshaw|1998|pp=828–830}}<ref name=Patnaik>{{cite book|last=Patnaik|first=Pradyot|title=A comprehensive guide to the hazardous properties of chemical substances|year=2007|publisher=John Wiley & Sons|isbn=9780471714583|pages=478–479|url=http://books.google.com/books?id=-CRRJBVv5d0C&pg=PA478}}</ref>

{| cellpadding="3" style="margin:1em auto 1em auto; text-align:center;"
|-
| colspan="5" style="text-align:center;"| '''Notable nonmetal fluorides, across the periodic table columns'''
|-
| [[File:Boron-trifluoride-3D-balls.png|150px]]
| [[File:Silicon-tetrafluoride-3D-balls.png|150px]]
| [[File:Antimony-pentafluoride-monomer-3D-balls.png|150px]]
| [[File:Sulfur-hexafluoride-3D-balls.png|150px]]
| [[File:Chlorine-trifluoride-3D-balls.png|150px]]
|- style="text-align:center;"
| Boron trifluoride 
| Silicon tetrafluoride 
| Antimony pentafluoride 
| Sulfur hexafluoride 
| Chlorine trifluoride 
|}

===Noble gas compounds===

{{main|Noble gas compound}}

The [[noble gas]]es are generally non-reactive because they have complete electron shells. Until the 1960s, no chemical bond with a noble gas was known. 

[[File:Xenon tetrafluoride crop.gif|thumb|right|alt=black-and-white photo showing icy-looking crystals in a dish|Xenon tetrafluoride, 1962]]

In 1962, [[Neil Bartlett (chemist)|Neil Bartlett]] reported the first chemical compound of xenon, [[xenon hexafluoroplatinate]].{{sfn|Wiberg|Wiberg|Holleman|2001|pp=392–393}} Later in 1962, xenon was reported to react directly with fluorine to form the [[xenon difluoride|di-]] and [[xenon tetrafluoride|tetrafluorides]]. Since then, [[xenon hexafluoride]], various oxyfluorides, and their derivatives have been prepared.{{sfn|Wiberg|Wiberg|Holleman|2001|p=438}}{{sfn|Wiberg|Wiberg|Holleman|2001|p=400}}

Krypton, xenon's lighter homolog, also forms a [[krypton difluoride|difluoride]] and a few more complicated fluorine-containing compounds.{{sfn|Lewars|2008|p=68}} The possibility of a tetrafluoride<ref>{{cite doi|10.1126/science.139.3559.1047}}</ref> or a hexafluoride has been debated.<ref>{{cite doi|10.1021/ic701313h}}</ref>

Radon, xenon's heavier homolog, has been shown to readily react with fluorine to form a solid compound. It is generally thought to be [[radon difluoride]], but the exact formula is not known.<ref name=astatineradon/> If radon were not so radioactive and difficult to collect, its chemistry might be at least as extensive as xenon's.{{sfn|Lewars|2008|p=67}}

The lightest noble gases do not form stable binary fluorides. Argon, however, reacts in extreme conditions with hydrogen fluoride to form [[argon fluorohydride]].<ref>{{cite doi|10.1038/35022551}}</ref> 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.<ref>{{cite doi|10.1063/1.1506150}}</ref> Neon is considered even less reactive than helium and no fluorides have been even momentarily observed.{{sfn|Lewars|2008|p=71}}

===Organic compounds===

{{main|Organofluorine chemistry}}

The [[carbon–fluorine bond|carbon–fluorine chemical bond]] is the strongest bond in [[organic chemistry]].<ref name=hagan>{{cite journal |author=O'Hagan, D. |title=Understanding organofluorine chemistry. An introduction to the C–F bond |journal=Chemical Society Reviews|volume=37 |issue=2 |pages=308–319 |year=2008 |pmid=18197347 |doi=10.1039/b711844a}}</ref> This C–F bond strength, along with the low polarizability of molecules containing C-F, makes organofluorines very stable.{{sfn|Siegemund et al.|2005|p=2}} Fluorinated organics have similar sizes to corresponding unfluorinated molecules because of the small van der Waals radius of fluorine.{{sfn|Siegemund et al.|2005|p=2}}

The range of organofluorine compounds is diverse, reflecting the inherent complexity of organic chemistry. A vast number of small molecules can exist with varying amounts of fluorine substitution, as well as many polymers. Research in particular areas is driven by the commercial value of applications.<ref name="Jstg"/>

====Small molecules====

Partially fluorinated alkanes are [[hydrofluorocarbon]]s (HFCs). 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 fluorines are substituted for hydrogens, the properties change. Solubility in hydrocarbons decreases and stability increases. Also, melting and boiling points decrease, while density goes up.{{sfn|Siegemund et al.|2005|pp=7–8}}

When all hydrogens are replaced with fluorines to make ''perfluorocarbons'' (the "per" means maximum), a great difference is revealed. Such compounds are extremely stable, and only sodium in liquid ammonia attacks them at standard conditions. They are also very insoluble, with few organic solvents capable of dissolving them.{{sfn|Siegemund et al.|2005|pp=7–8}}

{| style="margin:1em auto 1em auto; text-align:left; width:645px;" cellpadding="0" cellspacing="0"
|- style="vertical-align:top;"
| style="width:295px; padding-right:15px;" | [[File:Perfluorodecalin-3D-balls.png|295px|alt=two six-membered rings joined on a side, all F substituted]]
| style="width:320px; padding-left:15px;" | [[File:FluorocarbonCrabFish.JPG|320px|alt=beaker with two layers of liquid, goldfish and crab in top one. (uncaptioned, we also see coins at the bottom of the bottom layer)]]
|- style="text-align:left;"
|- style="vertical-align:top;"
| style="width:295px; padding-right:15px; "|[[Perfluorodecalin]] (trans isomer shown) is a liquid at room temperature. It boils at 142&nbsp;°C (288&nbsp;°F).<ref>{{cite web|url=http://www.f2chemicals.com/pdf/sds/Perfluorodecalin%20-%20SDS20109%20-%20ENG.pdf|title=Safety data sheet: Perfluorodecalin|year=2011|accessdate=14 June 2013|author=F2 Chemicals Limited}}</ref> Its hydrocarbon analog, [[decalin]], has a higher boiling point: 187&nbsp;°C (369&nbsp;°F).<ref>{{cite book|author=Mackay, Donald; Shiu, Wan-Ying; Ma, Kuo-Ching; Lee, Sum Chi|title=Handbook of Physical-Chemical Properties and Environmental Fate for Organic Chemicals|edition=2nd|year=2010|page=263|publisher=CRC Press|isbn=9781420044393|url=http://books.google.com/books?id=wjd-nEugVskC&pg=PA263}}</ref>
| style="width:320px; padding-left:15px; "|Perfluorocarbon density demonstration: upper layer is water with blue dye; lower layer is [[perfluoroheptane]]. The goldfish cannot swim down. The crab floats at the boundary. (Animals were rescued after the photo.)
|}

[[Perfluorinated compound]]s, hydrocarbons that are fully fluorinated except for one functional group, have a combination of properties.<ref>{{cite book|isbn=9781566772303|author=Mendicino, L.|page=116|year=1999|title=Environmental issues in the electronics and semiconductor industries|publisher=The Electrochemical Society}}</ref> They exhibit many perfluorocarbon properties (e.g. inertness, stability, non-wetting by water and oils, slipperiness).<ref name="fluorotelomer"/> However, the functional group is available for reactions or may make the molecule behave as a surfactant.<ref name=Salager2002>{{cite journal |last=Salager |first=Jean-Louis |title=Surfactants-Types and Uses |version=FIRP Booklet # 300-A |pages=45 |publisher=Universidad de los Andes Laboratory of Formulation, Interfaces Rheology, and Processes |year=2002 |url=http://nanoparticles.org/pdf/Salager-E300A.pdf |format=PDF |accessdate=7 September 2008}}</ref> If a perfluorinated compound has a fluorinated tail but also a few non-fluorinated carbons (typically two) near the functional group, it is called a [[fluorotelomer]]. Industrially, such compounds are treated as perfluorinated.<ref name="fluorotelomer">{{cite book |page=27 |year=2011 |publisher=Springer |author=Knepper, Thomas P.; Lange, Frank T. |isbn=9783642218712 |title=Polyfluorinated Chemicals and Transformation Products}}</ref>

Fluorinating an organic acid raises its acidity. For example, [[acetic acid]] and its [[fluoroacetic acid|mono-]], di-, and [[trifluoroacetic acid|trifluoroacetic]] derivatives show a trend of lowering pK<sub>a</sub>: 4.74, 2.66, 1.24, 0.23 (thus increasing acidity).{{sfn|Siegemund et al.|2005|p=29}} This happens because of fluorine's [[inductive effect]] (it stabilizes anions by spreading negative charge, allowing the hydrogen to be released).<ref>{{cite book|url=http://books.google.ca/books?id=1B6ijcTkD5EC&pg=PA52|author=Bansal, Raj|page=52|title=A Textbook of Organic Chemistry|publisher=New Age International|year=2003|isbn=9788122414592}}</ref> Similarly, the acidity is greatly increased for other perfluorocarboxyl acids, as well as the amines (which are not acids but become less basic if fluorinated).{{sfn|Siegemund et al.|2005|p=2}} The perfluoroalkanesulfonic acids are also very notable for their acidity. [[Trifluoromethanesulfonic acid]], is comparable to strong [[mineral acid]]s.{{sfn|Siegemund et al.|2005|p=32}}{{sfn|Siegemund et al.|2005|p=32}}

====Polymers====

[[File:Nafion2.svg|upright=1.4|thumb|alt=Skeletal chemical formula|The complex unit structure of Nafion]]

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.{{sfn|Carlson|Scmiegel|2005|p=3}}

The simplest fluoroplastic is polytetrafluoroethylene (PTFE, DuPont brand Teflon), which is a simple linear chain polymer with the repeating [[structural unit]]: –CF<sub>2</sub>–. It has no hydrogens and is the perfluoro analog of [[polyethylene]] (structural unit: –CH<sub>2</sub>–). PTFE has high chemical and thermal stability, as expected for a perfluorocarbon, much stronger than polyethylene. However, its very high melting point makes it difficult to fashion into parts.{{sfn|Carlson|Scmiegel|2005|pp=3–4}}

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 fluorines replaced with the –CF<sub>3</sub> groups). PFA ([[perfluoroalkoxy]] has some fluorines replaced with –OCF<sub>3</sub>).{{sfn|Carlson|Scmiegel|2005|pp=3–4}} [[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]] (–SO<sub>2</sub>OH) groups.<ref>{{cite book|title=Broadband dielectric spectroscopy studies of Nafion|page=2|publisher=ProQuest|year=2008|author=Rhoades, David Walter|isbn=9780549785408}}</ref><ref>{{cite doi|10.1016/S0013-4686(98)00132-7}}</ref>

There are other fluoroplastics that are not perfluorinated (contain some C-H). [[Polyvinylidene fluoride]] (PVDF, structural unit: –CF<sub>2</sub>CH<sub>2</sub>–), is an analog of PTFE with half the fluorines. PVF ([[polyvinyl fluoride]], structural unit: –CH<sub>2</sub>CHF–), contains one one-fourth the fluorines of PTFE. Despite this, it still has many properties of fully fluorinated polymers.{{sfn|Carlson|Scmiegel|2005|p=4}}

==Notes==

[[File:Periodic table fluorine.svg|thumb|alt=118 squares with numbers and letters in them, mostly colored gray and green, with a few numbers and words outside the boxes|Periodic table colored to show how elements are treated in this article: dark gray elements are metals, green ones are nonmetals, blue ones are the noble gases, the purple one is hydrogen, the yellow one is carbon, and the light gray ones are those with unknown properties.]]
<references group="note" />

==Sources==

===Citations===

{{Reflist|colwidth=30em}}

===Indexed references===
{{refbegin}}
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* {{cite book | last = Dean | first = John A. | title = Lange's Handbook of Chemistry | edition = 15th | publisher = McGraw-Hill | year = 1999 | isbn = 0-07-016190-9 | ref = harv}}
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{{refend}}

{{diatomicelements}}
{{Compact periodic table}}

{{Good article}}
{{Use dmy dates|date=November 2012}}

[[Category:Fluorine| ]]
[[Category:Chemical elements]]
[[Category:Halogens]]
[[Category:Diatomic nonmetals]]
[[Category:Biology and pharmacology of chemical elements]]
[[Category:Fluorinating agents]]
[[Category:Oxidizing agents]]
[[Category:Industrial gases]]

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