View source for Sodium fluoride

{{Short description|Ionic compound (NaF)}}
{{Use dmy dates|date=November 2024}}
{{cs1 config |name-list-style=vanc |display-authors=6}}
{{Chembox
| Verifiedfields = changed
| Watchedfields = changed
| verifiedrevid = 464362396
| Name = Sodium fluoride
| ImageFile = Sodium-fluoride-3D-ionic.png
| ImageName = Sodium fluoride
| ImageFile2 = Sodium fluoride.jpg
| ImageName2 = Sample of sodium fluoride, AR grade
| pronounce ={{IPAc-en|ˌ|s|oʊ|d|i|ə|m|_|ˈ|f|l|ʊər|aɪ|d}}<ref>{{citation|last=Wells|first=John C.|year=2008|title=Longman Pronunciation Dictionary|edition=3rd|publisher=Longman|pages=313 and 755|isbn=978-1-4058-8118-0}}. According to this source, an alternative pronunciation of the second word is {{IPAc-en|ˈ|f|l|ɔr|aɪ|d}} and, in the UK, also {{IPAc-en|ˈ|f|l|uː|ə|r|aɪ|d}}.</ref>
| IUPACName = Sodium fluoride
| OtherNames = Florocid
|Section1={{Chembox Identifiers
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| ChemSpiderID = 5045
| UNII_Ref = {{fdacite|correct|FDA}}
| UNII = 8ZYQ1474W7
| InChI = 1/FH.Na/h1H;/q;+1/p-1
| InChIKey = PUZPDOWCWNUUKD-REWHXWOFAH
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| StdInChIKey = PUZPDOWCWNUUKD-UHFFFAOYSA-M
| ChEMBL_Ref = {{ebicite|correct|EBI}}
| ChEMBL = 1528
| CASNo = 7681-49-4
| CASNo_Ref = {{cascite|correct|CAS}}
| RTECS = WB0350000
| EINECS = 231-667-8
| UNNumber = 1690
| ChEBI_Ref = {{ebicite|correct|EBI}}
| ChEBI = 28741
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| StdInChI = 1S/FH.Na/h1H;/q;+1/p-1
| PubChem = 5235
| KEGG_Ref = {{keggcite|changed|kegg}}
| KEGG = C08142
 }}
|Section2={{Chembox Properties
| Formula = NaF
| MolarMass = 41.988173 g/mol
| Appearance = White to greenish solid
| Odor = odorless
| Density = 2.558 g/cm<sup>3</sup>
| Solubility = 36.4 g/L (0 °C);<br /> 40.4 g/L (20 °C);<br /> 50.5 g/L (100 °C)<ref>{{cite book | editor= Haynes, William M. | year = 2011 | title = CRC Handbook of Chemistry and Physics | edition = 92nd
| publisher = [[CRC Press]] | isbn = 978-1-4398-5511-9|page=5.194| url=https://archive.org/details/handbookofchemis00crcp/page/5/mode/2up }}</ref>
| SolubleOther = slightly soluble in [[hydrofluoric acid|HF]], [[ammonia]] <br /> negligible in [[ethanol|alcohol]], [[acetone]], [[sulfur dioxide|SO<sub>2</sub>]], [[dimethylformamide]]
| MeltingPtC = 993
| BoilingPtC = 1704
| pKa = 3,20 (weak base, see [[hydrofluoric acid|HF]])
| VaporPressure = 1 mmHg @ 1077 °C<ref>Lewis, R.J. Sax's Dangerous Properties of Industrial Materials. 10th ed. Volumes 1–3 New York, NY: John Wiley & Sons Inc., 1999., p. 3248</ref>
| RefractIndex = 1.3252
| MagSus = −16.4·10<sup>−6</sup> cm<sup>3</sup>/mol
 }}
|Section3={{Chembox Structure
| MolShape = Octahedral
| CrystalStruct = Cubic
| LattConst_a = 462 pm
 }}
|Section5={{Chembox Thermochemistry
| DeltaHf = -573.6 kJ/mol
| DeltaGf = -543.3 kJ/mol
| Entropy = 51.3 J/(mol K)
| HeatCapacity = 46.82 J/(mol K)
 }}
|Section6={{Chembox Pharmacology
| ATCCode_prefix = A01
| ATCCode_suffix = AA01
| ATC_Supplemental = {{ATC|A12|CD01}},<br />{{ATC|V09|IX06}} ([[Fluorine-18|<sup>18</sup>F]])
}}
|Section7={{Chembox Hazards
| ExternalSDS = <ref name="sigma">{{Sigma-Aldrich|id=450022|name=Sodium Fluoride|access-date=17 March 2015}}</ref>
| HPhrases = {{H-phrases|301|315|319|335}}<ref name="sigma" />
| NFPA-H = 3
| NFPA-F = 0
| NFPA-R = 0
| FlashPt = Non-flammable
| GHSPictograms = {{GHS06|Acute Toxicity}} {{GHS07}} {{GHS08}}
| GHSSignalWord = Danger
| LD50 = 52–130 mg/kg (oral in rats, mice, rabbits)<ref>{{Citation |author1=Martel, B. |author2=Cassidy, K. |title=Chemical Risk Analysis: A Practical Handbook |publisher=Butterworth–Heinemann |year=2004 |isbn=978-1-903996-65-2 |page=363}}</ref>
| PEL = TWA 2.5 mg/m<sup>3</sup><ref name=PGCH>{{PGCH|0563}}</ref>
| IDLH = 250 mg/m<sup>3</sup> (as F)<ref name=PGCH />
| REL = TWA 2.5 mg/m<sup>3</sup><ref name=PGCH />
 }}
|Section8={{Chembox Related
| OtherAnions = [[Sodium chloride]]<br />[[Sodium bromide]]<br />[[Sodium iodide]]<br />[[Sodium astatide]]
| OtherCations = [[Lithium fluoride]]<br />[[Potassium fluoride]]<br />[[Rubidium fluoride]]<br />[[Caesium fluoride]]<br />[[Francium fluoride]]
| OtherCompounds = [[TASF reagent]]
  }}
}}
<!-- Definition and medical uses -->
'''Sodium fluoride''' ('''NaF''') is an [[inorganic compound]] with the formula {{chem2|auto=1|NaF}}. It is a colorless or white solid that is readily soluble in water. It is used in trace amounts in the [[fluoridation]] of drinking water to prevent [[tooth decay]], and in [[toothpaste]]s and topical pharmaceuticals for the same purpose. In 2022, it was the 221st most commonly prescribed medication in the United States, with more than 1{{nbsp}}million prescriptions.<ref>{{cite web | title=The Top 300 of 2022 | url=https://clincalc.com/DrugStats/Top300Drugs.aspx | website=ClinCalc | access-date=30 August 2024 | archive-date=30 August 2024 | archive-url=https://web.archive.org/web/20240830202410/https://clincalc.com/DrugStats/Top300Drugs.aspx | url-status=live }}</ref><ref>{{cite web | title = Sodium Fluoride Drug Usage Statistics, United States, 2013 - 2022 | website = ClinCalc | url = https://clincalc.com/DrugStats/Drugs/SodiumFluoride | access-date = 30 August 2024 }}</ref> It is also used in [[metallurgy]] and in [[medical imaging]].

== Uses ==
[[File:Sodium fluoride tablets.jpg|thumb|left|Sodium fluoride is sold in tablets for cavity prevention]]

=== Dental caries ===
{{See also|Fluoride therapy|Water fluoridation}}
Fluoride salts are often added to municipal drinking water (as well as to certain food products in some countries) for the purpose of maintaining dental health. The fluoride enhances the strength of teeth by the formation of [[fluorapatite]], a naturally occurring component of [[tooth enamel]].<ref>{{cite book|last1=Bourne|first1= Geoffrey H.|title=Dietary research and guidance in health and disease|date=1986|publisher=Karger|location=Basel|isbn=978-3-8055-4341-5|page=153}}</ref><ref>{{cite book|last1=Klein|first1=Cornelis|editor-first=Cornelius S.|editor-last=Hurlbut|title=Manual of Mineralogy (after James D. Dana)|date=1999|publisher=J. Wiley|location=New York|isbn=978-0-471-31266-6|edition=21st ed., rev.}}</ref><ref>{{cite journal|last1=Selwitz|first1=Robert H|last2=Ismail|first2=Amid I|last3=Pitts|first3=Nigel B|title=Dental caries|journal=The Lancet|date=January 2007|volume=369|issue=9555|pages=51–59|doi=10.1016/S0140-6736(07)60031-2|pmid=17208642|s2cid=204616785}}</ref> Although sodium fluoride is used to [[Water fluoridation|fluoridate water]] and is the standard by which other water-fluoridation compounds are gauged, [[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 in the United States.<ref>{{Citation |url=http://cdc.gov/fluoridation/pdf/statistics/1992.pdf |title= Fluoridation census 1992 |author= Division of Oral Health, National Center for Prevention Services, CDC |year=1993 |access-date=29 December 2008 |postscript=.}}</ref>

=== Osteoporosis ===
Fluoride supplementation has been extensively studied for the treatment of [[osteoporosis|postmenopausal osteoporosis]]. This supplementation does not appear to be effective; even though sodium fluoride increases bone density, it does not decrease the risk of fractures.<ref>{{cite journal|last1=Haguenauer|first1=D|last2=Welch|first2=V|last3=Shea|first3=B|last4=Tugwell|first4=P|last5=Wells|first5=G|title=Fluoride for treating postmenopausal osteoporosis.|journal=The Cochrane Database of Systematic Reviews|date=2000|volume=2010|issue=4|pages=CD002825|pmid=11034769|doi=10.1002/14651858.CD002825|pmc=8453489}}</ref><ref>{{cite journal|last1=Vestergaard|first1=P|last2=Jorgensen|first2=NR|last3=Schwarz|first3=P|last4=Mosekilde|first4=L|title=Effects of treatment with fluoride on bone mineral density and fracture risk—a meta-analysis.|journal=Osteoporosis International|date=March 2008|volume=19|issue=3|pages=257–68|pmid=17701094|doi=10.1007/s00198-007-0437-6|s2cid=25890845}}</ref>

=== Medical imaging ===
In medical imaging, [[fluorine-18]]-labelled sodium fluoride ([[United States Pharmacopeia|USP]], sodium fluoride Na<sup>18</sup>F) is one of the oldest tracers used in [[positron emission tomography]] (PET), having been in use since the 1960s.<ref>{{cite journal|last1=Blau|first1=Monte|last2=Ganatra|first2=Ramanik|last3=Bender|first3=Merrill A.|title=18F-fluoride for bone imaging|journal=Seminars in Nuclear Medicine|date=January 1972|volume=2|issue=1|pages=31–37|doi=10.1016/S0001-2998(72)80005-9|pmid=5059349}}</ref> Relative to conventional bone [[scintigraphy]] carried out with [[gamma camera]]s or [[SPECT]] systems, PET offers more sensitivity and spatial resolution. Fluorine-18 has a [[half-life]] of 110 min, which requires it to be used promptly once produced; this logistical limitation hampered its adoption in the face of the more convenient [[technetium-99m]]-labelled [[radiopharmaceutical]]s. However, fluorine-18 is generally considered to be a superior [[radiopharmaceutical]] for skeletal imaging. In particular it has a high and rapid bone uptake accompanied by very rapid blood clearance, which results in a high bone-to-background ratio in a short time.<ref>{{cite journal|last1=Ordonez|first1=A. A.|last2=DeMarco|first2=V. P.|last3=Klunk|first3=M. H.|last4=Pokkali|first4=S.|last5=Jain|first5=S.K.|title=Imaging Chronic Tuberculous Lesions Using Sodium [18F]Fluoride Positron Emission Tomography in Mice.|journal=Molecular Imaging and Biology|date=October 2015|volume=17|issue=5|pages=609–614|doi=10.1007/s11307-015-0836-6|pmid=25750032|pmc=4561601}}</ref> Additionally the annihilation photons produced by decay of <sup>18</sup>F have a high energy of 511 keV  compared to the 140 keV photons of <sup>99m</sup>Tc.<ref>{{cite journal|last1=Grant|first1=F. D.|last2=Fahey|first2=F. H.|last3=Packard|first3=A. B.|last4=Davis|first4=R. T.|last5=Alavi|first5=A.|last6=Treves|first6=S. T.|title=Skeletal PET with 18F-Fluoride: Applying New Technology to an Old Tracer|journal=Journal of Nuclear Medicine|date=12 December 2007|volume=49|issue=1|pages=68–78|doi=10.2967/jnumed.106.037200|pmid=18077529|doi-access=free}}</ref>

=== Chemistry ===
Sodium fluoride has a variety of specialty chemical applications in synthesis and extractive [[metallurgy]]. It reacts with electrophilic chlorides including [[acyl chloride]]s, sulfur chlorides, and phosphorus chloride.<ref name=halpern>{{citation | last = Halpern | first = D.F. | contribution = Sodium Fluoride | title = Encyclopedia of Reagents for Organic Synthesis | year = 2001 | publisher = [[John Wiley & Sons]] | doi = 10.1002/047084289X.rs071| isbn = 978-0-471-93623-7 }}</ref> Like other fluorides, sodium fluoride finds use in [[desilylation]] in [[organic synthesis]]. Sodium fluoride can be used to produce [[fluorocarbon]]s via the [[Finkelstein reaction]]; this process has the advantage of being simple to perform on a small scale but is rarely used on an industrial scale due to the existence of more effective techniques (e.g. [[Electrofluorination]], [[Fowler process]]).

=== Biology ===
Sodium fluoride is sometimes added at relatively high concentrations (~20 mM) to protein lysis buffers in order to inhibit endogenous [[phosphatases]] and thereby protect phosphorylated protein sites.<ref name="Somerville Wang 1988 pp. 118–129">{{cite journal | last1=Somerville | first1=Laura L. | last2=Wang | first2=Kuan | title=Sarcomere matrix of striated muscle: In vivo phosphorylation of titin and nebulin in mouse diaphragm muscle | journal=Archives of Biochemistry and Biophysics | publisher=Elsevier BV | volume=262 | issue=1 | year=1988 | issn=0003-9861 | doi=10.1016/0003-9861(88)90174-9 | pages=118–129| pmid=3355162 }}</ref> [[Sodium pyrophosphate]] and [[Sodium orthovanadate]] are also used for this purpose.<ref name="Thermo Fisher Scientific 2017">{{cite web | title=Overview of Protease and Phosphatase Inhibition for Protein Preparation - US | website=Thermo Fisher Scientific | date=10 May 2017 | url=https://www.thermofisher.com/us/en/home/life-science/protein-biology/protein-biology-learning-center/protein-biology-resource-library/pierce-protein-methods/protease-phosphatase-inhibitors.html | access-date=3 February 2023}}</ref>

=== Insecticide ===
Inorganic fluorides such as [[fluorosilicate]]s and sodium fluoride [[chelation|complex]] magnesium ions as magnesium [[monofluorophosphate|fluorophosphate]]. They inhibit enzymes such as [[enolase]] that require Mg<sup>2+</sup> as a prosthetic group. Thus, fluoride poisoning prevents phosphate transfer in [[oxidative metabolism]].<ref>{{citation | author=Metcalf, Robert L. | contribution=Insect Control | title=Ullmann's Encyclopedia of Industrial Chemistry | edition=7th | publisher=Wiley | year=2007 | page=9| title-link=Ullmann's Encyclopedia of Industrial Chemistry }}</ref> Sodium fluoride, patented as an insecticide in 1896, was commonly used through the 1970s on ants and other domestic pests,<ref name="Fluorides in the Environment">{{cite book|last1=Weinstein|first1=L.H.|url=https://books.google.com/books?id=hll2t6FVmgkC|title=Fluorides in the Environment: Effects on Plants and Animals|last2=Davison|first2=A.|date=2004|publisher=CABI Publishing|isbn=9780851998725|page=76|language=en}}</ref> and as a stomach poison for plant-feeding insects.<ref>{{cite book|last1=House|first1=James E.|url=https://books.google.com/books?id=pu3HBQAAQBAJ&q=Descriptive+Inorganic+Chemistry|title=Descriptive Inorganic Chemistry|last2=House|first2=Kathleen A.|date=10 September 2015|publisher=Academic Press|isbn=978-0-12-802979-4|page=397|language=en}}</ref> 

Its use, along with that of [[sodium fluorosilicate]], declined over the 20th century as the products were banned or restricted due to the possibility of poisoning, intentional or accidental.<ref name="Fluorides in the Environment" /> In 1942, for instance, [[Oregon State Hospital#1901–2000: Expansion and overcrowding|47 inmates at the Oregon State Hospital died]] after consuming scrambled eggs which had been inadvertently prepared with sodium fluoride; while assisting the cooks, another inmate had confused a container of insecticide—used by the hospital to control cockroaches—with powdered milk, which was stored nearby.<ref>{{cite news|url=https://www.nytimes.com/1983/10/04/obituaries/unsuspecting-poisoner-of-47-at-a-hospital-in-1942-is-dead.html|title=Unsuspecting Poisoner of 47 At a Hospital in 1942 Is Dead|work=[[The New York Times]]|date=4 October 1983|access-date=4 August 2016|archive-date=20 August 2016|archive-url=https://web.archive.org/web/20160820204404/http://www.nytimes.com/1983/10/04/obituaries/unsuspecting-poisoner-of-47-at-a-hospital-in-1942-is-dead.html|url-status=live}}</ref>

=== Other uses ===
Sodium fluoride is used as a cleaning agent (e.g., as a "[[laundry sour]]").<ref name=Aigueperse />

Sodium fluoride can be used in a nuclear [[molten salt reactor]].

== Safety ==
{{See also|Fluoride poisoning}}
The lethal dose for a 70&nbsp;kg (154&nbsp;lb) human is estimated at 5–10&nbsp;g.<ref name=Aigueperse />

Fluorides, particularly aqueous solutions of sodium fluoride, are rapidly and quite extensively absorbed by the human body.<ref name="fluorine">{{citation | author=Kapp, Robert | contribution=Fluorine | title=Encyclopedia of Toxicology | edition=2nd | volume=2 | publisher=Elsevier | year=2005 | pages=343–346}}</ref>

Fluorides interfere with electron transport and [[calcium metabolism]]. [[Calcium]] is essential for maintaining [[cardiac]] [[membrane]] potentials and in regulating [[coagulation]]. High ingestion of fluoride salts or hydrofluoric acid may result in fatal [[Heart arrhythmia|arrhythmias]] due to profound [[hypocalcemia]]. Chronic over-absorption can cause hardening of bones, calcification of ligaments, and buildup on teeth. Fluoride can cause irritation or corrosion to eyes, skin, and nasal membranes.<ref name="fluoride">{{citation | author=Greene Shepherd | contribution=Fluoride | title=Encyclopedia of Toxicology | edition=2nd | volume=2 | publisher=Elsevier | year=2005 | pages=342–343}}</ref>

Sodium fluoride is classed as toxic by both [[inhalation]] (of dusts or aerosols) and [[ingestion]].<ref>[http://hazard.com/msds/mf/baker/baker/files/s3722.htm NaF MSDS]. hazard.com</ref>  In high enough doses, it has been shown to affect the heart and circulatory system. For occupational exposures, the [[Occupational Safety and Health Administration]] and the [[National Institute for Occupational Safety and Health]] have established occupational exposure limits at 2.5&nbsp;mg/m<sup>3</sup> over an eight-hour time-weighted average.<ref>[https://www.cdc.gov/niosh/npg/npgd0563.html CDC – NIOSH Pocket Guide to Chemical Hazards]</ref>

In the higher doses used to treat [[osteoporosis]], plain 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 [[peptic ulcer disease]]. Slow-release and [[Enteric coating|enteric-coated]] versions of sodium fluoride do not have significant gastric side effects, and have milder and less frequent complications in the bones.<ref>{{cite journal |author=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]]. A chronic fluoride ingestion of 1 ppm of fluoride in drinking water can cause [[mottling]] of the teeth (fluorosis) and an exposure of 1.7 ppm will produce mottling in 30%–50% of patients.<ref name="fluorine" /> Studies have shown that dental fluorosis negatively impacts the self-esteem and self-image of adolescents.<ref>{{cite journal |last1=Molina-Frechero |first1=N |last2=Nevarez-Rascón |first2=M |last3=Nevarez-Rascón |first3=A |last4=González-González |first4=R |last5=Irigoyen-Camacho |first5=ME |last6=Sánchez-Pérez |first6=L |last7=López-Verdin |first7=S |last8=Bologna-Molina |first8=R |title=Impact of Dental Fluorosis, Socioeconomic Status and Self-Perception in Adolescents Exposed to a High Level of Fluoride in Water. |journal=International Journal of Environmental Research and Public Health |date=12 January 2017 |volume=14 |issue=1 |page=73 |doi=10.3390/ijerph14010073 |doi-access=free |pmid=28085102 |pmc=5295324 }}</ref> <ref>{{cite journal |last1=Nilchian |first1=F |last2=Asgary |first2=I |last3=Mastan |first3=F |title=The Effect of Dental Fluorosis on the Quality of Life of Female High School and Precollege Students of High Fluoride-Concentrated Area. |journal=Journal of International Society of Preventive & Community Dentistry |date=July 2018 |volume=8 |issue=4 |pages=314–319 |doi=10.4103/jispcd.JISPCD_94_18 |doi-access=free |pmid=30123763 |pmc=6071349 }}</ref>

== Chemical structure ==
Sodium fluoride is an inorganic [[ionic compound]], dissolving in water to give separated Na<sup>+</sup> and F<sup>−</sup> ions. Like [[sodium chloride]], it crystallizes in a cubic motif where both Na<sup>+</sup> and F<sup>−</sup> occupy [[Octahedral molecular geometry|octahedral coordination sites]];<ref name=wells>{{Citation | last = Wells | first = A.F. | year = 1984 | title = Structural Inorganic Chemistry | place = Oxford | publisher = [[Clarendon Press]] | isbn = 978-0-19-855370-0}}</ref><ref name=tp>{{citation | url = http://www.atsdr.cdc.gov/toxprofiles/tp11.pdf | contribution = Chemical and physical information | title = Toxicological profile for fluorides, hydrogen fluoride, and fluorine | pages = 187 | publisher = Agency for Toxic Substances and Disease Registry (ATDSR) | date = September 2003 | access-date = 1 November 2008}}</ref> its lattice spacing, approximately 462 [[picometre|pm]], is smaller than that of sodium chloride (564 pm).

== Occurrence ==
The [[mineral]] form of NaF, [[villiaumite]], is moderately rare. It is known from plutonic [[nepheline syenite]] rocks.<ref name=minhandbook>{{Citation | url = http://rruff.geo.arizona.edu/doclib/hom/villiaumite.pdf | title = Mineral Handbook
| year = 2005 | publisher = Mineral Data Publishing | issue = version 1 | postscript = .}}</ref>

== Production ==
NaF is prepared by neutralizing [[hydrofluoric acid]] or [[hexafluorosilicic acid]] (H<sub>2</sub>SiF<sub>6</sub>), both byproducts of the reaction of fluorapatite (Ca<sub>5</sub>(PO<sub>4</sub>)<sub>3</sub>F) from [[phosphate rock]] during the production of [[superphosphate]] fertilizer.  Neutralizing agents include [[sodium hydroxide]] and [[sodium carbonate]].  Alcohols are sometimes used to precipitate the NaF:<ref name=Aigueperse/>
:HF  +  NaOH  →   NaF  +  H<sub>2</sub>O
From solutions containing HF, sodium fluoride precipitates as the [[bifluoride]] salt [[sodium bifluoride]] (NaHF<sub>2</sub>).  Heating the latter releases HF and gives NaF.
:HF  +  NaF  ⇌   NaHF<sub>2</sub>

In a 1986 report, the annual worldwide consumption of NaF was estimated to be several million tonnes.<ref name=Aigueperse>{{Ullmann|doi=10.1002/14356007.a11_307|title=Fluorine Compounds, Inorganic|year=2000|last1=Aigueperse|first1=Jean|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|isbn=3-527-30673-0}}</ref>

== See also ==
{{Commons category}}

== References ==
{{reflist}}

{{Sodium compounds}}
{{Stomatological preparations}}
{{Fluorides}}
{{Portal bar | Medicine}}
{{Authority control}}

[[Category:Fluorides]]
[[Category:Metal halides]]
[[Category:Alkali metal fluorides]]
[[Category:Sodium compounds]]
[[Category:Dental drugs]]
[[Category:Rock salt crystal structure]]