View source for Fluoride

{{Short description|Ion of fluorine}}
{{About|the fluoride ion|a review of fluorine compounds|Compounds of fluorine|the fluoride additive used in toothpaste|Fluoride therapy}}
{{Distinguish|Floride (disambiguation){{!}}Floride|Fluorite}}
{{Use dmy dates|date=July 2022}}
{{Chembox
| IUPACName = Fluoride<ref>{{cite web|url=https://pubchem.ncbi.nlm.nih.gov/compound/28179|title=Fluorides – PubChem Public Chemical Database|work=The PubChem Project|location=USA|publisher=National Center for Biotechnology Information|at=Identification}}</ref>
| ImageFileL1 = F- crop.svg
| ImageSizeL1 = 100px
| ImageFileR1 = Fluoride_ion.svg
| ImageSizeR1 = 100px
|Section1={{Chembox Identifiers
| CASNo = 16984-48-8
| CASNo_Ref = {{cascite|correct|CAS}}
| UNII_Ref = {{fdacite|correct|FDA}}
| UNII = Q80VPU408O
| PubChem = 28179
| ChemSpiderID = 26214
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| KEGG = C00742
| KEGG_Ref = {{keggcite|correct|kegg}}
| MeSHName = Fluoride
| ChEBI = 17051
| ChEMBL = 1362
| ChEMBL_Ref = {{ebicite|correct|EBI}}
| Gmelin = 14905
| SMILES = [F-]
| StdInChI = 1S/FH/h1H/p-1
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| StdInChIKey = KRHYYFGTRYWZRS-UHFFFAOYSA-M
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
}}
|Section2={{Chembox Properties
| Formula = {{Chem|F|-}}
| F=1 
| ConjugateAcid = [[Hydrogen fluoride]]
}}
|Section3={{Chembox Thermochemistry
| DeltaHf = −333 kJ mol<sup>−1</sup>
| Entropy = 145.58 J/mol K (gaseous)<ref>{{cite web |url=http://webbook.nist.gov/cgi/cbook.cgi?ID=C16984488&Mask=1#Thermo-Gas |title=Fluorine anion |pages=1–1951 |publisher=NIST |access-date=July 4, 2012|year=1998 |last1=Chase |first1=M. W. }}</ref>
}}
|Section4={{Chembox Related
| OtherAnions = {{unbulleted list|[[Chloride]]|[[Bromide]]|[[Iodide]]}}
}}
}}

'''Fluoride''' ({{IPAc-en|ˈ|f|l|ʊər|aɪ|d|,_|ˈ|f|l|ɔr|-}})<ref name="wells313">{{cite book|last1=Wells|first1=J.C.|title=Longman pronunciation dictionary|date=2008|publisher=Pearson Education Limited/Longman|location=Harlow, England|isbn=978-1-4058-8118-0|page=313|edition=3rd}}. According to this source, {{IPAc-en|ˈ|f|l|uː|ə|r|aɪ|d}} is a possible pronunciation in British English.</ref> is an [[Inorganic chemistry|inorganic]], [[Monatomic ion|monatomic]] [[Ion#Anions and cations|anion]] of [[fluorine]], with the [[chemical formula]] {{Chem|F|−}} (also written {{Chem|[F]|−}}), whose salts are typically white or colorless. Fluoride salts typically have distinctive bitter tastes, and are odorless. Its salts and minerals are important [[Reagent|chemical reagents]] and industrial chemicals, mainly used in the production of [[hydrogen fluoride]] for [[fluorocarbon]]s. Fluoride is classified as a weak base since it only partially associates in solution, but concentrated fluoride is corrosive and can attack the skin.

Fluoride is the simplest fluorine [[anion]]. In terms of charge and size, the fluoride [[ion]] resembles the [[hydroxide]] ion. Fluoride ions occur on [[Earth]] in several minerals, particularly [[fluorite]], but are present only in trace quantities in bodies of water in nature.

== Nomenclature ==
Fluorides include compounds that contain ionic fluoride and those in which fluoride does not dissociate. The nomenclature does not distinguish these situations. For example, [[sulfur hexafluoride]] and [[carbon tetrafluoride]] are not sources of fluoride ions under ordinary conditions.

The systematic name ''fluoride'', the valid [[IUPAC]] name, is determined according to the additive nomenclature. However, the name ''fluoride'' is also used in compositional IUPAC nomenclature which does not take the nature of bonding involved into account.
''Fluoride'' is also used non-systematically, to describe compounds which release fluoride upon dissolving. Hydrogen fluoride is itself an example of a non-systematic name of this nature. However, it is also a [[trivial name]], and the [[preferred IUPAC name]] for ''fluorane''.{{citation needed|date = September 2017}}

==Occurrence==
[[File:Calcite sur fluorine (USA) 1.JPG|right|thumb|Fluorite crystals]]
[[Fluorine]] is estimated to be the 13th-most [[Abundance of elements in Earth's crust|abundant element]] in Earth's crust and is widely dispersed in nature, entirely in the form of fluorides. The vast majority is held in [[:Category:Fluorine minerals|mineral deposits]], the most commercially important of which is [[fluorite]] (CaF<sub>2</sub>).<ref name=Aigueperse/> Natural weathering of some kinds of rocks,<ref>{{cite journal |last1=Derakhshani |first1=R |last2=Raoof |first2=A |last3=Mahvi |first3=AH |last4=Chatrouz |first4=H |title=Similarities in the Fingerprints of Coal Mining Activities, High Ground Water Fluoride, and Dental Fluorosis in Zarand District, Kerman Province, Iran |journal=Fluoride |date=2020 |volume=53 |issue=2 |pages=257–267}}</ref><ref>{{cite journal |last1=Derakhshani |first1=R |last2=Tavallaie |first2=M |last3=Malek Mohammad |first3=T |last4=Abbasnejad |first4=A |last5=Haghdoost |first5=A |title=Occurrence of fluoride in groundwater of Zarand region, Kerman province, Iran |journal=Fluoride |date=2014 |volume=47 |issue=2 |pages=133–138}}</ref> as well as human activities, releases fluorides into the [[biosphere]] through what is sometimes called the [[fluorine cycle]].

===In water===
Fluoride is naturally present in groundwater, [[Fresh water|fresh]] and [[Saline water|saltwater]] sources, as well as in rainwater, particularly in urban areas exposed to air pollution.<ref>{{cite web |url= https://www.atsdr.cdc.gov/phs/phs.asp?id=210&tid=38 |title= Public Health Statement for Fluorides, Hydrogen Fluoride, and Fluorine |publisher= [[ATSDR]] |date= September 2003 }}</ref> Seawater fluoride levels are usually in the range of 0.86 to 1.4&nbsp;mg/L, and average 1.1&nbsp;mg/L<ref>{{cite web|url=http://www.env.gov.bc.ca/wat/wq/BCguidelines/fluoride/fluoridetoo-01.html|title=Ambient Water Quality Criteria for Fluoride|publisher=Government of British Columbia|access-date=8 October 2014|archive-date=24 September 2015|archive-url=https://web.archive.org/web/20150924035152/http://www.env.gov.bc.ca/wat/wq/BCguidelines/fluoride/fluoridetoo-01.html}}</ref> (milligrams per [[litre]]). For comparison, [[chloride]] concentration in seawater is about 19&nbsp;g/L. The low concentration of fluoride reflects the insolubility of the [[Alkaline earth metal|alkaline earth]] fluorides, e.g., CaF<sub>2</sub>.

Concentrations in fresh water vary more significantly. [[Surface water]] such as rivers or lakes generally contains between 0.01 and 0.3&nbsp;mg/L.<ref>{{cite book|last1=Liteplo|first1=Dr R.|last2=Gomes|first2=R.|last3=Howe|first3=P.|last4=Malcolm|first4=Heath|title=FLUORIDES - Environmental Health Criteria 227: 1st draft|date=2002|publisher=World Health Organization|location=Geneva|isbn=978-92-4-157227-9|url=http://www.inchem.org/documents/ehc/ehc/ehc227.htm#1.4}}</ref> [[Groundwater]] (well water) concentrations vary even more, depending on the presence of local fluoride-containing minerals. For example, natural levels of under 0.05&nbsp;mg/L have been detected in parts of Canada but up to 8&nbsp;mg/L in parts of China; in general levels rarely exceed 10&nbsp;mg/litre<ref name=Fawell>{{cite web|author1=Fawell, J.K.|title=Fluoride in Drinking-water Background document for development of WHO Guidelines for Drinking-water Quality|url=https://www.who.int/water_sanitation_health/dwq/chemicals/fluoride.pdf|publisher=World Health Organization|access-date=6 May 2016|ref=WHO/SDE/WSH/03.04/96|language=en|display-authors=etal}}</ref> 
* In parts of Asia the groundwater can contain dangerously high levels of fluoride, leading to serious [[Fluoride toxicity#Chronic toxicity|health problems]].<ref>{{cite journal |last1=Yadav |first1=Krishna Kumar |last2=Kumar |first2=Sandeep |last3=Pham |first3=Quoc Bao |last4=Gupta |first4=Neha |last5=Rezania |first5=Shahabaldin |last6=Kamyab |first6=Hesam |last7=Yadav |first7=Shalini |last8=Vymazal |first8=Jan |last9=Kumar |first9=Vinit |last10=Tri |first10=Doan Quang |last11=Talaiekhozani |first11=Amirreza |last12=Prasad |first12=Shiv |last13=Reece |first13=Lisa M. |last14=Singh |first14=Neeraja |last15=Maurya |first15=Pradip Kumar |last16=Cho |first16=Jinwoo |title=Fluoride contamination, health problems and remediation methods in Asian groundwater: A comprehensive review |journal=Ecotoxicology and Environmental Safety |date=October 2019 |volume=182 |article-number=109362 |doi=10.1016/j.ecoenv.2019.06.045|pmid=31254856 |bibcode=2019EcoES.18209362Y |s2cid=195764865 }}</ref> 
* Worldwide, 50 million people receive water from water supplies that naturally have close to the "optimal level".<ref>{{cite web|last1=Tiemann|first1=Mary|title=Fluoride in Drinking Water: A Review of Fluoridation and Regulation Issues|url=https://www.fas.org/sgp/crs/misc/RL33280.pdf|publisher=Congressional Research Service|access-date=6 May 2016|ref=7-5700 www.crs.gov RL33280|page=3|date=April 5, 2013}}</ref>
* In other locations the level of fluoride is very low, sometimes leading to [[Water fluoridation|fluoridation]] of public water supplies to bring the level to around 0.7–1.2 ppm.
*[[Mining]] can increase local fluoride levels<ref>{{cite journal |last1=Chandio |first1=Tasawar Ali |last2=Khan |first2=Muhammad Nasiruddin |last3=Muhammad |first3=Maria Taj |last4=Yalcinkaya |first4=Ozcan |last5=Wasim |first5=Agha Arslan |last6=Kayis |first6=Ahmet Furkan |title=Fluoride and arsenic contamination in drinking water due to mining activities and its impact on local area population |journal=Environmental Science and Pollution Research |date=January 2021 |volume=28 |issue=2 |pages=2355–2368 |doi=10.1007/s11356-020-10575-9|pmid=32880840 |bibcode=2021ESPR...28.2355C |s2cid=221463681 |doi-access=free }}</ref>

Fluoride can be present in rain, with its concentration increasing significantly upon exposure to volcanic activity<ref>{{cite journal |last1=Bellomo |first1=Sergio |last2=Aiuppa |first2=Alessandro |last3=D'Alessandro |first3=Walter |last4=Parello |first4=Francesco |title=Environmental impact of magmatic fluorine emission in the Mt. Etna area |journal=Journal of Volcanology and Geothermal Research |date=August 2007 |volume=165 |issue=1–2 |pages=87–101 |doi=10.1016/j.jvolgeores.2007.04.013|bibcode=2007JVGR..165...87B }}</ref> or atmospheric pollution derived from burning fossil fuels or other sorts of industry,<ref>{{cite journal|last1=Smith|first1=Frank A.|last2=Hodge|first2=Harold C.|last3=Dinman|first3=B. D.|title=Airborne fluorides and man: Part I|journal=CRC Critical Reviews in Environmental Control|date=9 January 2009|volume=8|issue=1–4|pages=293–371|doi=10.1080/10643387709381665}}</ref><ref>{{cite journal|last1=Smith|first1=Frank A.|last2=Hodge|first2=Harold C.|last3=Dinman|first3=B. D.|title=Airborne fluorides and man: Part II|journal=CRC Critical Reviews in Environmental Control|date=9 January 2009|volume=9|issue=1|pages=1–25|doi=10.1080/10643387909381666}}</ref> particularly [[aluminium smelter]]s.<ref>{{cite journal |last1=Arnesen |first1=A.K.M. |last2=Abrahamsen |first2=G. |last3=Sandvik |first3=G. |last4=Krogstad |first4=T. |title=Aluminium-smelters and fluoride pollution of soil and soil solution in Norway |journal=Science of the Total Environment |date=February 1995 |volume=163 |issue=1–3 |pages=39–53 |doi=10.1016/0048-9697(95)04479-K|bibcode=1995ScTEn.163...39A }}</ref>

===In plants===
All vegetation contains some fluoride, which is absorbed from soil and water.<ref name=Fawell /> Some plants concentrate fluoride from their environment more than others. All tea leaves contain fluoride; however, mature leaves contain as much as 10 to 20 times the fluoride levels of young leaves from the same plant.<ref>{{cite journal |vauthors=Wong MH, Fung KF, Carr HP |title=Aluminium and fluoride contents of tea, with emphasis on brick tea and their health implications |journal=Toxicology Letters |volume=137 |issue=1–2 |pages=111–20 |year=2003 |pmid=12505437 |doi=10.1016/S0378-4274(02)00385-5 }}</ref><ref name="pmid18078704">{{cite journal |vauthors=Malinowska E, Inkielewicz I, Czarnowski W, Szefer P |title=Assessment of fluoride concentration and daily intake by human from tea and herbal infusions |journal=Food Chem. Toxicol. |volume=46 |issue=3 |pages=1055–61 |year=2008 |pmid=18078704 |doi=10.1016/j.fct.2007.10.039}}</ref><ref>{{cite journal |vauthors=Gardner EJ, Ruxton CH, Leeds AR |title=Black tea&mdash;helpful or harmful? A review of the evidence |journal=European Journal of Clinical Nutrition |volume=61 |issue=1 |pages=3–18 |year=2007 |pmid=16855537 |doi=10.1038/sj.ejcn.1602489 |doi-access= }}</ref>

== Chemical properties ==

=== Basicity ===
Fluoride can act as a [[Base (chemistry)|base]]. It can combine with a [[proton]] ({{H+}}):

:{{chem2 | F- + H+ -> HF }}

This neutralization reaction forms [[hydrogen fluoride]] (HF), the [[conjugate acid]] of fluoride.

In aqueous solution, fluoride has a [[acid dissociation constant|p''K''<sub>b</sub>]] value of 10.8. It is therefore a [[weak base]], and tends to remain as the fluoride ion rather than generating a substantial amount of hydrogen fluoride. That is, the following equilibrium favours the left-hand side in water:

:{{chem2 | F- + H2O <-> HF + HO- }}

However, upon prolonged contact with moisture, soluble fluoride salts will decompose to their respective hydroxides or oxides, as the hydrogen fluoride escapes. Fluoride is distinct in this regard among the halides. The identity of the solvent can have a dramatic effect on the equilibrium shifting it to the right-hand side, greatly increasing the rate of decomposition.

=== Structure of fluoride salts===
Salts containing fluoride are numerous and adopt myriad structures. Typically the fluoride anion is surrounded by four or six cations, as is typical for other halides. [[Sodium fluoride]] and [[sodium chloride]] adopt the same structure. For compounds containing more than one fluoride per cation, the structures often deviate from those of the chlorides, as illustrated by the main fluoride mineral [[fluorite]] (CaF<sub>2</sub>) where the Ca<sup>2+</sup> ions are surrounded by eight F<sup>−</sup> centers. In CaCl<sub>2</sub>, each Ca<sup>2+</sup> ion is surrounded by six Cl<sup>−</sup> centers. The difluorides of the transition metals often adopt the [[rutile]] structure whereas the dichlorides have [[cadmium chloride]] structures.

===Inorganic chemistry===
Upon treatment with a standard acid, fluoride salts convert to [[hydrogen fluoride]] and metal [[salt (chemistry)|salt]]s. With very strong acids, it can be doubly protonated to give [[fluoronium|{{Chem|H|2|F|+}}]]. Oxidation of fluoride gives fluorine. Solutions of inorganic fluorides in water contain F<sup>−</sup> and [[bifluoride]] {{chem|HF|2|-}}.<ref>{{cite book|last1=Wiberg|last2=Holleman|first2=A.F.|title=Inorganic chemistry|date=2001|publisher=Academic Press, W. de Gruyter.|location=San Diego, Calif. : Berlin|isbn=978-0-12-352651-9|edition=1st English ed., [edited] by Nils Wiberg.}}</ref> Few inorganic fluorides are soluble in water without undergoing significant hydrolysis. In terms of its reactivity, fluoride differs significantly from [[chloride]] and other halides, and is more strongly solvated in [[protic solvent]]s due to its smaller radius/charge ratio. Its closest chemical relative is [[hydroxide]], since both have similar geometries.

===Naked fluoride===
Most fluoride salts dissolve to give the bifluoride ({{chem|HF|2|-}}) anion. Sources of true F<sup>−</sup> anions are rare because the highly basic fluoride anion abstracts protons from many, even adventitious, sources. Relative [[solvation|unsolvated]] fluoride, which does exist in aprotic solvents, is called "naked". '''Naked fluoride''' is a strong [[Lewis base]],<ref>{{cite journal |last1=Schwesinger |first1=Reinhard |last2=Link |first2=Reinhard |last3=Wenzl |first3=Peter |last4=Kossek |first4=Sebastian |title=Anhydrous Phosphazenium Fluorides as Sources for Extremely Reactive Fluoride Ions in Solution |journal=Chemistry: A European Journal |volume=12 |issue=2 |pages=438–45 |year=2005 |pmid=16196062 |doi=10.1002/chem.200500838 }}</ref> and a powerful nucleophile. Some quaternary ammonium salts of naked fluoride include [[tetramethylammonium fluoride]]  and [[tetrabutylammonium fluoride]].<ref name=dimagno>{{cite journal |author1=Haoran Sun |author2=Stephen G. DiMagno |name-list-style=amp | title= Anhydrous Tetrabutylammonium Fluoride | journal= [[Journal of the American Chemical Society]] | year= 2005 | volume= 127 | pages= 2050–1| doi=10.1021/ja0440497 | pmid= 15713075 | issue= 7}}</ref> [[Cobaltocenium]] fluoride is another example.<ref>{{cite journal |doi=10.1021/ja00103a045|title=Cobaltocenium Fluoride: A Novel Source of "Naked" Fluoride Formed by Carbon-Fluorine Bond Activation in a Saturated Perfluorocarbon|year=1994|last1=Bennett|first1=Brian K.|last2=Harrison|first2=Roger G.|last3=Richmond|first3=Thomas G.|journal=Journal of the American Chemical Society|volume=116|issue=24|pages=11165–11166}}</ref> However, they all lack structural characterization in aprotic solvents. Because of their high basicity, many so-called naked fluoride sources are in fact bifluoride salts. In late 2016 [[imidazolium]] fluoride was synthesized that is the closest approximation of a thermodynamically stable and structurally characterized example of a "naked" fluoride source in an aprotic solvent (acetonitrile).<ref>{{cite journal | last1 = Alič | first1 = B. | last2 = Tavčar | first2 = G. | year = 2016 | title = Reaction of N-heterocyclic carbene (NHC) with different HF sources and ratios – A free fluoride reagent based on imidazolium fluoride | journal = J. Fluorine Chem. | volume = 192 | pages = 141–146 | doi = 10.1016/j.jfluchem.2016.11.004 }}</ref> The sterically demanding imidazolium cation stabilizes the discrete anions and protects them from polymerization.<ref>{{cite journal | last1 = Alič | first1 = B. | last2 = Tramšek | first2 = M. | last3 = Kokalj | first3 = A. | last4 = Tavčar | first4 = G. | year = 2017| title = Discrete GeF5– Anion Structurally Characterized with a Readily Synthesized Imidazolium Based Naked Fluoride Reagent | journal = Inorg. Chem. | volume = 56 | issue = 16| pages = 10070–10077 | doi = 10.1021/acs.inorgchem.7b01606 | pmid = 28792216 }}</ref><ref>{{cite journal | last1 = Zupanek | first1 = Ž. | last2 = Tramšek | first2 = M. | last3 = Kokalj | first3 = A. | last4 = Tavčar | first4 = G. | year = 2018| title = Reactivity of VOF3 with N-Heterocyclic Carbene and Imidazolium Fluoride: Analysis of Ligand–VOF3 Bonding with Evidence of a Minute π Back-Donation of Fluoride | journal = Inorg. Chem. | volume = 57 | issue = 21| pages = 13866–13879 | doi = 10.1021/acs.inorgchem.8b02377 | pmid = 30353729 | s2cid = 53031199 }}</ref>

=== Biochemistry ===
At physiological pHs, [[hydrogen fluoride]] is usually fully ionised to fluoride. In [[biochemistry]], fluoride and hydrogen fluoride are equivalent. Fluorine, in the form of fluoride, is considered to be a [[micronutrient]] for human health, necessary to prevent dental cavities, and to promote healthy bone growth.<ref name="who.int">{{cite web|last1=Fawell|first1=J. |title=Fluoride in Drinking-water|url=https://www.who.int/water_sanitation_health/dwq/chemicals/fluoride.pdf|publisher=World Health Organization|access-date=10 March 2016}}</ref> The tea plant (''[[Camellia sinensis]]'' L.) is a known accumulator of fluorine compounds, released upon forming infusions such as the common beverage. The fluorine compounds decompose into products including fluoride ions. Fluoride is the most bioavailable form of fluorine, and as such, tea is potentially a vehicle for fluoride dosing.<ref name=Chan2013>{{Cite journal|last=Chan|first=Laura|author2=Mehra, Aradhana |author3=Saikat, Sohel |author4= Lynch, Paul |title=Human exposure assessment of fluoride from tea (''Camellia sinensis'' L.): A UK based issue?|journal=Food Research International|date=May 2013|volume=51|issue=2|pages=564–570|doi=10.1016/j.foodres.2013.01.025}}</ref> Approximately, 50% of absorbed fluoride is excreted renally with a twenty-four-hour period. The remainder can be retained in the oral cavity, and lower digestive tract. Fasting dramatically increases the rate of fluoride absorption to near 100%, from a 60% to 80% when taken with food.<ref name=Chan2013 /> Per a 2013 study, it was found that consumption of one litre of tea a day, can potentially supply the daily recommended intake of 4&nbsp;mg per day. Some lower quality brands can supply up to a 120% of this amount. Fasting can increase this to 150%. The study indicates that tea drinking communities are at an increased risk of [[dental fluorosis|dental]] and [[skeletal fluorosis]], in the case where water fluoridation is in effect.<ref name=Chan2013 /> Fluoride ion in low doses in the mouth reduces tooth decay.<ref>{{Cite web | url=http://oradyne.net/fluoride-free-toothpaste/ | title=Fluoride Free Toothpaste – Fluoride (Finally!) Explained| date=2016-06-27}}</ref> For this reason, it is used in toothpaste and water fluoridation. At much higher doses and frequent exposure, fluoride causes health complications and can be toxic.

==Applications==
{{See also|Fluorochemical industry|Biological aspects of fluorine|Fluorine}}
Fluoride salts and hydrofluoric acid are the main fluorides of industrial value.

===Organofluorine chemistry===
{{main|Organofluorine chemistry}}
Organofluorine compounds are pervasive. Many drugs, many polymers, refrigerants, and many inorganic compounds are made from fluoride-containing reagents. Often fluorides are converted to hydrogen fluoride, which is a major reagent and precursor to reagents. Hydrofluoric acid and its anhydrous form, [[hydrogen fluoride]], are particularly important.<ref name="Aigueperse">{{cite encyclopedia|last1=Aigueperse|first1=Jean|encyclopedia=Ullmann's Encyclopedia of Industrial Chemistry|last2=Mollard|first2=Paul|last3=Devilliers|first3=Didier|last4=Chemla|first4=Marius|last5=Faron|first5=Robert|last6=Romano|first6=René|last7=Cuer|first7=Jean Pierre|year=2000|doi=10.1002/14356007.a11_307|isbn=978-3-527-30673-2|chapter=Fluorine Compounds, Inorganic}}</ref>

===Production of metals and their compounds===
The main uses of fluoride, in terms of volume, are in the production of cryolite, Na<sub>3</sub>AlF<sub>6</sub>. It is used in [[aluminium smelting]]. Formerly, it was mined, but now it is derived from hydrogen fluoride. Fluorite is used on a large scale to separate slag in steel-making. Mined [[fluorite]] (CaF<sub>2</sub>) is a commodity chemical used in steel-making.  [[Uranium hexafluoride]] is employed in the purification of uranium isotopes.

===Cavity prevention===
{{main|Fluoride therapy|Water fluoridation}}
[[Image:Sodium fluoride tablets.jpg|thumb|right|Sodium fluoride sold in tablets for cavity prevention.]]
Fluoride-containing compounds, such as [[sodium fluoride]] or [[sodium monofluorophosphate]] are used in topical and systemic [[fluoride therapy]] for preventing [[Dental caries|tooth decay]]. They are used for [[water fluoridation]] and in many products associated with [[oral hygiene]].<ref name="mcdonagh2000">{{cite journal|author1=McDonagh M. S. |author2=Whiting P. F. |author3=Wilson P. M. |author4=Sutton A. J. |author5=Chestnutt I. |author6=Cooper J. |author7=Misso K. |author8=Bradley M. |author9=Treasure E. |author10=Kleijnen J. |year=2000 |title=Systematic review of water fluoridation |journal=[[British Medical Journal]] |volume=321 |issue=7265| pages=855–859|doi=10.1136/bmj.321.7265.855 |pmid=11021861|pmc=27492}}</ref> Originally, sodium fluoride was used to fluoridate water; [[hexafluorosilicic acid]] (H<sub>2</sub>SiF<sub>6</sub>) and its salt [[sodium hexafluorosilicate]] (Na<sub>2</sub>SiF<sub>6</sub>) are more commonly used additives, especially in the United States. The fluoridation of water is known to prevent tooth decay<ref>{{cite journal |vauthors=Griffin SO, Regnier E, Griffin PM, Huntley V |title=Effectiveness of fluoride in preventing caries in adults |journal=J. Dent. Res. |volume=86 |issue=5 |pages=410–5 |year=2007 |pmid=17452559 |doi=10.1177/154405910708600504|hdl=10945/60693 |s2cid=58958881 |hdl-access=free }}</ref><ref>{{cite journal |author1=Winston A. E. |author2=Bhaskar S. N. |title=Caries prevention in the 21st century |journal=J. Am. Dent. Assoc. |volume=129 |issue=11 |pages=1579–87 |date=1 November 1998 |pmid=9818575 |url=http://jada.ada.org/cgi/pmidlookup?view=long&pmid=9818575 |archive-url=https://archive.today/20120715143950/http://jada.ada.org/cgi/pmidlookup?view=long&pmid=9818575 |archive-date=15 July 2012 |doi=10.14219/jada.archive.1998.0104 |url-access=subscription }}</ref> and is considered by the U.S. [[Centers for Disease Control and Prevention]] to be "one of 10 great public health achievements of the 20th century".<ref>{{cite web|title=Community Water Fluoridation|url=https://www.cdc.gov/fluoridation/|publisher=Centers for Disease Control and Prevention|access-date=10 March 2016}}</ref><ref>{{cite web|title=Ten Great Public Health Achievements in the 20th Century|url=https://www.cdc.gov/about/history/tengpha.htm|publisher=Centers for Disease Control and Prevention|access-date=10 March 2016|archive-url=https://web.archive.org/web/20160313072852/http://www.cdc.gov/about/history/tengpha.htm|archive-date=2016-03-13}}</ref> In some countries where large, centralized water systems are uncommon, fluoride is delivered to the populace by fluoridating table salt. Despite medical consensus as to its safety, fluoridation of water has its critics {{crossreference|(see [[Water fluoridation controversy]])}}.<ref>{{cite journal |author=Newbrun E |title=The fluoridation war: a scientific dispute or a religious argument? |journal=Journal of Public Health Dentistry |volume=56 |issue=5 Spec No |pages=246–52 |year=1996 |pmid=9034969 |doi=10.1111/j.1752-7325.1996.tb02447.x}}</ref> Fluoridated [[toothpaste]] is in common use. [[Meta-analysis]] show the efficacy of 500 ppm fluoride in toothpastes.<ref>{{Cite journal|last1=Walsh|first1=Tanya|last2=Worthington|first2=Helen V.|last3=Glenny|first3=Anne-Marie|last4=Marinho|first4=Valeria Cc|last5=Jeroncic|first5=Ana|date=March 4, 2019|title=Fluoride toothpastes of different concentrations for preventing dental caries|journal=Cochrane Database of Systematic Reviews|volume=3|issue=3|article-number=CD007868|doi=10.1002/14651858.CD007868.pub3|issn=1469-493X|pmc=6398117|pmid=30829399}}</ref><ref>{{Cite web|title=Remineralization of initial carious lesions in deciduous enamel after application of dentifrices of different fluoride concentrations|url=https://www.springermedizin.de/remineralization-of-initial-carious-lesions-in-deciduous-enamel-/8679072|access-date=2021-02-24|website=springermedizin.de|language=de}}</ref> However, no beneficial effect can be detected when more than one fluoride source is used for daily oral care.<ref>{{Cite journal|last1=Hausen|first1=H.|last2=Kärkkäinen|first2=S.|last3=Seppä|first3=L.|date=February 2000|title=Application of the high-risk strategy to control dental caries|journal=Community Dentistry and Oral Epidemiology|volume=28|issue=1|pages=26–34|doi=10.1034/j.1600-0528.2000.280104.x|issn=0301-5661|pmid=10634681}}</ref>{{request quotation|date=August 2021}}

===Laboratory reagent===
Fluoride salts are commonly used in biological assay processing to [[enzyme inhibitor|inhibit]] the activity of [[phosphatases]], such as [[serine]]/[[threonine]] phosphatases.<ref>{{cite journal |vauthors=Nakai C, Thomas JA |title=Properties of a phosphoprotein phosphatase from bovine heart with activity on glycogen synthase, phosphorylase, and histone |journal=J. Biol. Chem. |volume=249 |issue=20 |pages=6459–67 |year=1974 |doi=10.1016/S0021-9258(19)42179-0 |pmid=4370977 |doi-access=free }}</ref> Fluoride mimics the [[nucleophile|nucleophilic]] [[hydroxide]] ion in these enzymes' active sites.<ref>{{cite journal |vauthors=Schenk G, Elliott TW, Leung E |title=Crystal structures of a purple acid phosphatase, representing different steps of this enzyme's catalytic cycle |journal=BMC Struct. Biol. |volume=8 |page=6 |year=2008 |pmid=18234116 |doi=10.1186/1472-6807-8-6 |pmc=2267794|display-authors=etal |doi-access=free }}</ref> [[Beryllium fluoride]] and [[aluminium fluoride]] are also used as phosphatase inhibitors, since these compounds are structural mimics of the [[phosphate]] group and can act as analogues of the [[transition state]] of the reaction.<ref>{{cite journal |vauthors=Wang W, Cho HS, Kim R |title=Structural characterization of the reaction pathway in phosphoserine phosphatase: crystallographic "snapshots" of intermediate states |journal=J. Mol. Biol. |volume=319 |issue=2 |pages=421–31 |year=2002 |pmid=12051918 |doi=10.1016/S0022-2836(02)00324-8 |display-authors=etal}}</ref><ref>{{cite journal |vauthors=Cho H, Wang W, Kim R |title=BeF(3)(-) acts as a phosphate analog in proteins phosphorylated on aspartate: structure of a BeF(3)(-) complex with phosphoserine phosphatase |journal=Proc. Natl. Acad. Sci. U.S.A. |volume=98 |issue=15 |pages=8525–30 |year=2001 |pmid=11438683 |doi=10.1073/pnas.131213698 |pmc=37469|bibcode = 2001PNAS...98.8525C |display-authors=etal|doi-access=free }}</ref>

==Dietary recommendations==
The U.S. Institute of Medicine (IOM) updated Estimated Average Requirements (EARs) and Recommended Dietary Allowances (RDAs) for some minerals in 1997. Where there was not sufficient information to establish EARs and RDAs, an estimate designated Adequate Intake (AI) was used instead. AIs are typically matched to actual average consumption, with the assumption that there appears to be a need, and that need is met by what people consume. The current AI for women 19 years and older is 3.0&nbsp;mg/day (includes pregnancy and lactation). The AI for men is 4.0&nbsp;mg/day. The AI for children ages 1–18 increases from 0.7 to 3.0&nbsp;mg/day. The major known risk of [[Fluorine deficiency|fluoride deficiency]] appears to be an increased risk of bacteria-caused tooth cavities. As for safety, the IOM sets tolerable upper intake levels (ULs) for vitamins and minerals when evidence is sufficient. In the case of fluoride the UL is 10&nbsp;mg/day. Collectively the EARs, RDAs, AIs and ULs are referred to as [[Dietary Reference Intake]]s (DRIs).<ref name="DRItext">{{cite book | last1 = Institute of Medicine | title = Dietary Reference Intakes for Calcium, Phosphorus, Magnesium, Vitamin D and Fluoride | chapter = Fluoride | publisher = The National Academies Press | year = 1997 | location = Washington, DC | pages = 288–313 | doi = 10.17226/5776 | pmid = 23115811 | isbn = 978-0-309-06403-3 | chapter-url = https://www.nap.edu/read/5776/chapter/10| author1-link = Institute of Medicine }}</ref>

The [[European Food Safety Authority]] (EFSA) refers to the collective set of information as Dietary Reference Values, with Population Reference Intake (PRI) instead of RDA, and Average Requirement instead of EAR. AI and UL are defined the same as in the United States. For women ages 18 and older the AI is set at 2.9&nbsp;mg/day (including pregnancy and lactation). For men, the value is 3.4&nbsp;mg/day. For children ages 1–17 years, the AIs increase with age from 0.6 to 3.2&nbsp;mg/day. These AIs are comparable to the U.S. AIs.<ref>{{cite web | title = Overview on Dietary Reference Values for the EU population as derived by the EFSA Panel on Dietetic Products, Nutrition and Allergies| year = 2017| url = https://www.efsa.europa.eu/sites/default/files/assets/DRV_Summary_tables_jan_17.pdf}}</ref> The EFSA reviewed safety evidence and set an adult UL at 7.0&nbsp;mg/day (lower for children).<ref name="EFSAtext">{{citation| title = Tolerable Upper Intake Levels For Vitamins And Minerals| publisher = European Food Safety Authority| year = 2006| url = http://www.efsa.europa.eu/sites/default/files/efsa_rep/blobserver_assets/ndatolerableuil.pdf}}</ref>

For U.S. food and dietary supplement labeling purposes, the amount of a vitamin or mineral in a serving is expressed as a percent of Daily Value (%DV). Although there is information to set Adequate Intake, fluoride does not have a Daily Value and is not required to be shown on food labels.<ref name="FedReg">{{cite web|url=https://www.gpo.gov/fdsys/pkg/FR-2016-05-27/pdf/2016-11867.pdf |title=Federal Register May 27, 2016 Food Labeling: Revision of the Nutrition and Supplement Facts Labels. FR page 33982.}}</ref>

==Estimated daily intake==
Daily intakes of fluoride can vary significantly according to the various sources of exposure. Values ranging from 0.46 to 3.6–5.4&nbsp;mg/day have been reported in several studies (IPCS, 1984).<ref name="who.int" /> In areas where water is [[water fluoridation|fluoridated]] this can be expected to be a significant source of fluoride, however fluoride is also naturally present in virtually all foods and beverages at a wide range of concentrations.<ref>{{cite web |url=http://ndb.nal.usda.gov/ndb/nutrients/report?nutrient1=313&nutrient2=&nutrient3=&fg=&max=25&subset=0&offset=0&sort=c&totCount=563&measureby=g |archive-url=https://web.archive.org/web/20140526022117/http://ndb.nal.usda.gov/ndb/nutrients/report?nutrient1=313&nutrient2=&nutrient3=&fg=&max=25&subset=0&offset=0&sort=c&totCount=563&measureby=g |archive-date=26 May 2014 |title=Nutrient Lists|publisher=Agricultural Research Service United States Department of Agriculture|access-date=25 May 2014}}</ref> The maximum safe daily consumption of fluoride is 10&nbsp;mg/day for an adult (U.S.) or 7&nbsp;mg/day (European Union).<ref name="DRItext" /><ref name="EFSAtext" />

The upper limit of fluoride intake from all sources (fluoridated water, food, beverages, fluoride dental products and dietary fluoride supplements) is set at 0.10&nbsp;mg/kg/day for infants, toddlers, and children through to 8 years old. For older children and adults, who are no longer at risk for dental fluorosis, the upper limit of fluoride is set at 10&nbsp;mg/day regardless of weight.<ref>{{cite journal | doi = 10.1111/j.1752-7325.1999.tb03272.x | volume=59 | title=Total Fluoride Intake and Implications for Dietary Fluoride Supplementation | year=1999 | journal=Journal of Public Health Dentistry | pages=211–223  | last1 = Levy | first1 = Steven M. | last2 = Guha-Chowdhury | first2 = Nupur| issue=4 | pmid=10682326 }}</ref>

{| class="wikitable"
|+ Examples of fluoride content
|-
! Food/Drink !! Fluoride<br />(mg per 1000g/ppm) !! Portion !! Fluoride<br />(mg per portion)
|-
| Black tea (brewed) || style="text-align: right;"|3.73 ||  1 cup, 240 g (8 fl oz) || style="text-align: right;"|0.884
|-
| Raisins, seedless || style="text-align: right;"|2.34 || small box, {{convert|43|g|oz|abbr=on}} || style="text-align: right;"|0.101
|-
| Table wine || style="text-align: right;"|1.53 || Bottle, {{convert|750|mL|impoz|abbr=on}} || style="text-align: right;"|1.150
|-
| Municipal tap-water,<br />(Fluoridated) || style="text-align: right;"|0.81 || Recommended daily intake,<br /> {{convert|3|L|USgal}} || style="text-align: right;"|2.433
|-
| Baked potatoes, Russet || style="text-align: right;"|0.45 || Medium potato, {{convert|140|g|lb|abbr=on}} || style="text-align: right;"|0.078
|-
| Lamb || style="text-align: right;"|0.32 || Chop, {{convert|170|g|oz|abbr=on}} || style="text-align: right;"|0.054
|-
| Carrots || style="text-align: right;"|0.03 || 1 large carrot, {{convert|72|g|oz|abbr=on}} || style="text-align: right;"|0.002
|-
| colspan="4" style="text-align: center;" | Source: Data taken from United States Department of Agriculture, [https://fdc.nal.usda.gov/fdc-app.html#/?component=0 National Nutrient Database] {{Webarchive|url=https://web.archive.org/web/20140301101828/http://ndb.nal.usda.gov/ndb/nutrients/index |date=2014-03-01 }}<ref>{{Cite web |title=Food Composition Databases: Food Search: Fluoride |url=https://ndb.nal.usda.gov/ndb/nutrients/report/nutrientsfrm?max=25&offset=0&totCount=0&nutrient1=313&nutrient2=&subset=0&sort=c&measureby=g |archive-url=https://web.archive.org/web/20181205193315/https://ndb.nal.usda.gov/ndb/nutrients/report/nutrientsfrm?max=25&offset=0&totCount=0&nutrient1=313&nutrient2=&subset=0&sort=c&measureby=g |archive-date=5 December 2018 |publisher=[[Agricultural Research Service]], [[United States Department of Agriculture]] |access-date=5 December 2018}}</ref>
|}

== Safety ==
{{main|Fluoride toxicity}}

=== Ingestion ===
According to the U.S. Department of Agriculture, the Dietary Reference Intakes, which is the "highest level of daily nutrient intake that is likely to pose no risk of adverse health effects" specify 10&nbsp;mg/day for most people, corresponding to 10 L of fluoridated water with no risk. For young children the values are smaller, ranging from 0.7&nbsp;mg/d to 2.2&nbsp;mg/d for infants.<ref>{{cite web|title=Dietary Reference Intakes: EAR, RDA, AI, Acceptable Macronutrient Distribution Ranges, and UL |url=http://fnic.nal.usda.gov/dietary-guidance/dietary-reference-intakes/dri-tables-and-application-reports|archive-url=https://web.archive.org/web/20150213140728/http://fnic.nal.usda.gov/dietary-guidance/dietary-reference-intakes/dri-tables-and-application-reports|archive-date=13 February 2015|publisher=United States Department of Agriculture|access-date=9 September 2017}}</ref> Water and food sources of fluoride include community water fluoridation, seafood, tea, and gelatin.<ref>{{cite web|title=Fluoride in diet|url=https://www.medlineplus.gov/ency/article/002420.htm|publisher=U.S. National Library of Medicine|access-date=10 March 2016}}</ref>

Soluble fluoride salts, of which [[sodium fluoride]] is the most common, are toxic, and have resulted in both accidental and self-inflicted deaths from [[Acute toxicity|acute poisoning]].<ref name=Aigueperse/> The lethal dose for most adult humans is estimated at 5 to 10&nbsp;g (which is equivalent to 32 to 64&nbsp;mg elemental fluoride per kg body weight).<ref>{{cite book | last = Gosselin | first = RE |author2=Smith RP |author3=Hodge HC | title = Clinical toxicology of commercial products | publisher = Williams & Wilkins| location = Baltimore (MD) | year = 1984 | pages =III–185–93| isbn =978-0-683-03632-9 }}</ref><ref>{{cite book |last=Baselt |first=RC |url=https://archive.org/details/dispositionoftox0000base_v7n5/page/636/mode/2up |title=Disposition of toxic drugs and chemicals in man |publisher=Biomedical Publications |year=2008 |isbn=978-0-9626523-7-0 |location=Foster City (CA) |pages=636–40 |url-access=registration}}</ref><ref>{{cite book | last = IPCS | title = Environmental health criteria 227 (Fluoride) | publisher = International Programme on Chemical Safety, World Health Organization| location = Geneva | year = 2002 | page =100| isbn =978-92-4-157227-9}}</ref> A case of a fatal poisoning of an adult with 4&nbsp;grams of sodium fluoride is documented,<ref name=acute/> and a dose of 120&nbsp;g sodium fluoride has been survived.<ref>{{cite journal |vauthors=Abukurah AR, ((Moser AM Jr)), Baird CL, ((Randall RE Jr)), Setter JG, Blanke RV | title=Acute sodium fluoride poisoning | journal=JAMA | year=1972 | pages=816–7 | volume=222 | issue=7 | pmid=4677934 | doi=10.1001/jama.1972.03210070046014}}</ref> For [[sodium fluorosilicate]] (Na<sub>2</sub>SiF<sub>6</sub>), the [[median lethal dose]] (LD<sub>50</sub>) orally in rats is 125&nbsp;mg/kg, corresponding to 12.5&nbsp;g for a 100&nbsp;kg adult.<ref name="merck">The Merck Index, 12th edition, Merck & Co., Inc., 1996</ref>

Treatment may involve oral administration of dilute [[calcium hydroxide]] or [[calcium chloride]] to prevent further absorption, and injection of [[calcium gluconate]] to increase the calcium levels in the blood.<ref name=acute>{{cite journal|pmid=20323400|pmc=1581810|year=1945|last1=Rabinowitch|first1=IM|title=Acute Fluoride Poisoning|volume=52|issue=4|pages=345–9|journal=Canadian Medical Association Journal}}</ref> [[Hydrogen fluoride]] is more dangerous than salts such as NaF because it is corrosive and volatile, and can result in fatal exposure through inhalation or upon contact with the skin; calcium gluconate gel is the usual antidote.<ref>{{cite journal |vauthors=Muriale L, Lee E, Genovese J, Trend S | year = 1996 | title = Fatality due to acute fluoride poisoning following dermal contact with hydrofluoric acid in a palynology laboratory | journal = Ann. Occup. Hyg. | volume = 40 | issue = 6| pages = 705–710 | pmid = 8958774 | doi = 10.1016/S0003-4878(96)00010-5 }}</ref>

In the higher doses used to treat [[osteoporosis]], sodium fluoride can cause pain in the legs and incomplete stress fractures when the doses are too high; it also irritates the stomach, sometimes so severely as to cause ulcers. Slow-release and [[enteric]]-coated versions of sodium fluoride do not have gastric side effects in any significant way, and have milder and less frequent complications in the bones.<ref>{{cite journal |vauthors=Murray TM, Ste-Marie LG |title=Prevention and management of osteoporosis: consensus statements from the Scientific Advisory Board of the Osteoporosis Society of Canada. 7. Fluoride therapy for osteoporosis |journal=CMAJ |volume=155 |issue=7 |pages=949–54 |year=1996 |pmid=8837545 |pmc=1335460}}</ref> In the lower doses used for [[water fluoridation]], the only clear adverse effect is [[dental fluorosis]], which can alter the appearance of children's teeth during [[Human tooth development|tooth development]]; this is mostly mild and is unlikely to represent any real effect on aesthetic appearance or on public health.<ref>{{cite book |url=http://nhmrc.gov.au/_files_nhmrc/file/publications/synopses/Eh41_Flouridation_PART_A.pdf |year=2007 |title=A systematic review of the efficacy and safety of fluoridation |author=National Health and Medical Research Council (Australia) |isbn=978-1-86496-415-8 |access-date=2010-02-21 |archive-url=https://web.archive.org/web/20091014191758/http://www.nhmrc.gov.au/_files_nhmrc/file/publications/synopses/Eh41_Flouridation_PART_A.pdf |archive-date=2009-10-14 }} Summary: {{cite journal |author= Yeung CA |title= A systematic review of the efficacy and safety of fluoridation |journal= Evid.-Based Dent. |volume=9 |issue=2 |pages=39–43 |year=2008 |pmid=18584000 |doi=10.1038/sj.ebd.6400578 |doi-access=free }}</ref> Fluoride was known to enhance bone mineral density at the lumbar spine, but it was not effective for vertebral fractures and provoked more nonvertebral fractures.<ref>{{cite journal|last=Haguenauer|first=D|author2=Welch, V |author3=Shea, B |author4=Tugwell, P |author5=Adachi, JD |author6= Wells, G |title=Fluoride for the treatment of postmenopausal osteoporotic fractures: a meta-analysis.|journal=Osteoporosis International |date=2000|volume=11|issue=9|pages=727–38|pmid=11148800 |doi=10.1007/s001980070051|s2cid=538666}}</ref> In areas that have naturally occurring high levels of fluoride in [[groundwater]] which is used for [[drinking water]], both [[dental fluorosis|dental]] and [[skeletal fluorosis]] can be prevalent and severe.<ref>{{cite web |author=World Health Organization |year=2004 |title=Fluoride in drinking-water |url=https://www.who.int/water_sanitation_health/dwq/chemicals/en/fluoride.pdf |archive-url=https://web.archive.org/web/20160304082148/http://www.who.int/water_sanitation_health/dwq/chemicals/en/fluoride.pdf |archive-date=2016-03-04 |access-date=2014-02-13}}</ref>

==== Hazard maps for fluoride in groundwater ====

Around one-third of the human population drinks water from groundwater resources. Of this, about 10%, approximately 300 million people, obtain water from groundwater resources that are heavily contaminated with arsenic or fluoride.<ref>Eawag (2015) Geogenic Contamination Handbook – Addressing Arsenic and Fluoride in Drinking Water. C.A. Johnson, A. Bretzler (Eds.), Swiss Federal Institute of Aquatic Science and Technology (Eawag), Duebendorf, Switzerland. (download: www.eawag.ch/en/research/humanwelfare/drinkingwater/wrq/geogenic-contamination-handbook/)</ref>  These trace elements derive mainly from minerals.<ref>{{cite journal | last1 = Rodríguez-Lado | first1 = L. | last2 = Sun | first2 = G. | last3 = Berg | first3 = M. | last4 = Zhang | first4 = Q. | last5 = Xue | first5 = H. | last6 = Zheng | first6 = Q. | last7 = Johnson | first7 = C.A. | year = 2013 | title = Groundwater arsenic contamination throughout China | url = https://www.dora.lib4ri.ch/eawag/islandora/object/eawag%3A7346| journal = Science | volume = 341 | issue = 6148| pages = 866–868 | doi = 10.1126/science.1237484 | pmid = 23970694 | bibcode = 2013Sci...341..866R | s2cid = 206548777 }}</ref> Maps locating potential problematic wells are available.<ref>[http://www.gapmaps.org Groundwater Assessment Platform]</ref>

=== Topical ===
Concentrated fluoride solutions are corrosive.<ref>{{cite journal |vauthors=Nakagawa M, Matsuya S, Shiraishi T, Ohta M |title=Effect of fluoride concentration and pH on corrosion behavior of titanium for dental use |journal=Journal of Dental Research |volume=78 |issue=9 |pages=1568–72 |year=1999 |pmid=10512392 |doi=10.1177/00220345990780091201|s2cid=32650790 }}</ref> Gloves made of [[nitrile rubber]] are worn when handling fluoride compounds. The hazards of solutions of fluoride salts depend on the concentration. In the presence of [[strong acid]]s, fluoride salts release [[hydrogen fluoride]], which is corrosive, especially toward glass.<ref name=Aigueperse/>

== Other derivatives ==
Organic and inorganic anions are produced from fluoride, including:
*[[Bifluoride]], used as an etchant for glass<ref>{{Cite web|url=http://www.eng.chimko.com/item29/|title=Ammonium bifluoride in the glass industry - Chimex Ltd}}</ref>
*[[Tetrafluoroberyllate]]
*[[Hexafluoroplatinate]]
*[[Tetrafluoroborate]] used in organometallic synthesis
*[[Hexafluorophosphate]] used as an electrolyte in commercial secondary batteries.
*[[Trifluoromethanesulfonate]]

== See also ==
{{Portal|Medicine}}
* [[Per- and polyfluoroalkyl substances]]
* [[Fluorine-19 nuclear magnetic resonance spectroscopy]]
* [[Fluoride deficiency]]
* [[Fluoride selective electrode]]
* [[Fluoride therapy]]
* [[Sodium monofluorophosphate]]

== References ==
{{Reflist|35em}}

== External links ==
{{Commons category|Fluorides}}
* [https://fas.org/sgp/crs/misc/RL33280.pdf "Fluoride in Drinking Water: A Review of Fluoridation and Regulation Issues"], [[Congressional Research Service]]
* [https://web.archive.org/web/20150203111150/http://apps.nccd.cdc.gov/MWF/index.asp U.S. government site for checking status of local water fluoridation]

{{Fluorides}}
{{Monatomic anion compounds}}

{{Authority control}}

[[Category:Fluorides| ]]
[[Category:Anions]]
[[Category:Biology and pharmacology of chemical elements]]
[[Category:Nephrotoxins]]