Sodium fluoride
Sodium fluoride
IUPAC name
Sodium fluoride
Other names
ECHA InfoCard 100.028.789 Edit this at Wikidata
EC Number
  • 231-667-8
RTECS number
  • WB0350000
UN number 1690
  • InChI=1S/FH.Na/h1H;/q;+1/p-1 checkY
  • InChI=1/FH.Na/h1H;/q;+1/p-1
Molar mass 41.988173 g/mol
Appearance White to greenish solid
Odor odorless
Density 2.558 g/cm3
Melting point 993 °C (1,819 °F; 1,266 K)
Boiling point 1,704 °C (3,099 °F; 1,977 K)
36.4 (0 °C); 40.4 (20 °C); 50.5 (100 °C) g/L[1]
Solubility slightly soluble in HF, ammonia
negligible in alcohol, acetone, SO2, dimethylformamide
Vapor pressure 1 mmHg @ 1077 C°[2]
46.82 J/mol K
51.3 J/mol K
-573.6 kJ/mol
-543.3 kJ/mol
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
52–200 mg/kg (oral in rats, mice, rabbits)[3]
Related compounds
Other anions
Sodium chloride
Sodium bromide
Sodium iodide
Other cations
Lithium fluoride
Potassium fluoride
Rubidium fluoride
Caesium fluoride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Sodium fluoride is an inorganic chemical compound with the formula NaF. A colorless solid, it is a source of the fluoride ion in diverse applications. Sodium fluoride is less expensive and less hygroscopic than the related salt potassium fluoride.

Structure, general properties, occurrence

Sodium fluoride is an ionic compound, dissolving to give separated Na+ and F ions. Like sodium chloride, it crystallizes in a cubic motif where both Na+ and F occupy octahedral coordination sites;[4][5] its lattice spacing, approximately 462 pm, is somewhat smaller than that of sodium chloride.

The mineral form of NaF, villiaumite, is moderately rare. It is known from plutonic nepheline syenite rocks.[6]


NaF is prepared by neutralizing hydrofluoric acid or hexafluorosilicic acid (H2SiF6), byproducts of the reaction of fluorapatite (Ca5(PO4)3F) (from phosphate rock) from the production of superphosphate fertilizer. Neutralizing agents include sodium hydroxide and sodium carbonate. Alcohols are sometimes used to precipitate the NaF:

HF + NaOH → NaF + H2O

From solutions containing HF, sodium fluoride precipitates as the bifluoride salt NaHF2. Heating the latter releases HF and gives NaF.

HF + NaF ⇌ NaHF2

In a 1986 report, the annual worldwide consumption of NaF was estimated to be several million tonnes.[7]


See also: Fluoride therapy and Water fluoridation

Sodium fluoride is sold in tablets for cavity prevention.

Fluoride salts is often added to drinking water and some food products for dental health. Other fluoride sources are used as well, e.g., salts of hexafluorosilicate. The fluoride enhance the strength of teeth by the formation of fluorapatite, a naturally occurring component of tooth enamel.[8][9] Although sodium fluoride is also used to fluoridate water and, indeed, is the standard by which other water-fluoridation compounds are gauged, hexafluorosilicic acid (H2SiF6) and its salt sodium hexafluorosilicate (Na2SiF6) are more commonly used additives in the U.S.[10] Toothpaste often contains sodium fluoride to prevent cavities.[11]

Sodium fluoride is used as a cleaning agent (e.g., as a "laundry sour").[7]

A variety of specialty chemical applications exist in synthesis and extractive metallurgy. It reacts with electrophilic chlorides including acyl chlorides, sulfur chlorides, and phosphorus chloride.[12] Like other fluorides, sodium fluoride finds use in desilylation in organic synthesis. The fluoride is the reagent for the synthesis of fluorocarbons.[citation needed]

In medical imaging, fluorine-18-labelled sodium fluoride is used in positron emission tomography (PET). Relative to conventional bone scintigraphy carried out with gamma cameras or SPECT systems, PET offers more sensitivity and spatial resolution. A disadvantage of PET is that fluorine-18 labelled sodium fluoride is less widely available than conventional technetium-99m-labelled radiopharmaceuticals.[citation needed]

Sodium fluoride is used as a stomach poison for plant-feeding insects. Inorganic fluorides such as fluorosilicates and sodium fluoride complex magnesium ions as magnesium fluorophosphate. They inhibit enzymes such as enolase that require Mg2+ as a prosthetic group. Thus, fluoride poisoning prevents phosphate transfer in oxidative metabolism.[13]


See also: Fluoride poisoning

Fluorides, particularly aqueous solutions of sodium fluoride, are rapidly and quite extensively absorbed.[14]

Fluorides interfere with electron transport and calcium metabolism. Calcium is essential for maintaining cardiac membrane potentials and in regulating coagulation. Large ingestion of fluoride salts or hydrofluoric acid may result in fatal arrhythmias due to profound hypocalcemia. Recreational inhalation of fluoridated hydrocarbon refrigerants like Freon has been associated with "sudden sniffing death", which is thought to be a fatal arrhythmia caused by myocardial sensitization to catecholamines.[15]

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.[15]

The lethal dose for a 70 kg (154 lb) human is estimated at 5–10 g.[7] Sodium fluoride is classed as toxic by both inhalation (of dusts or aerosols) and ingestion.[16] 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 mg/m3 over an eight-hour time-weighted average.[17]

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 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.[18] 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 tooth development; this is mostly mild and is unlikely to represent any real effect on aesthetic appearance or on public health.[19] 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.[14]

See also


  1. ^ Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). CRC Press. p. 5.194. ISBN 1439855110.
  2. ^ 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
  3. ^ Martel, B.; Cassidy, K. (2004), Chemical Risk Analysis: A Practical Handbook, Butterworth–Heinemann, p. 363, ISBN 1-903996-65-1((citation)): CS1 maint: multiple names: authors list (link)
  4. ^ Wells, A.F. (1984), Structural Inorganic Chemistry, Oxford: Clarendon Press, ISBN 0-19-855370-6
  5. ^ "Chemical and physical information", Toxicological profile for fluorides, hydrogen fluoride, and fluorine (PDF), Agency for Toxic Substances and Disease Registry (ATDSR), September 2003, p. 187, retrieved 2008-11-01
  6. ^ Mineral Handbook (PDF), Mineral Data Publishing, 2005.
  7. ^ a b c Aigueperse, Jean (2005), "Fluorine Compounds, Inorganic", in Ullmann (ed.), Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH, doi:10.1002/14356007.a11_307 ((citation)): Unknown parameter |coauthors= ignored (|author= suggested) (help)
  8. ^ Bourne, Geoffrey Howard (1986), Dietary research and guidance in health and disease, Karger, p. 153, ISBN 3-8055-4341-7, Snippet view from page 153
  9. ^ Klein, Cornelis; Hurlbut, Cornelius Searle; Dana, James Dwight (1999), Manual of Mineralogy (21 ed.), Wiley, ISBN 0-471-31266-5
  10. ^ Division of Oral Health, National Center for Prevention Services, CDC (1993), Fluoridation census 1992 (PDF), retrieved 2008-12-29.((citation)): CS1 maint: multiple names: authors list (link)
  11. ^ "Sodium fluoride, Molecule of the week". American Chemical Society. 2008-02-19. Retrieved 2008-11-01.
  12. ^ Halpern, D.F. (2001), "Sodium Fluoride", Encyclopedia of Reagents for Organic Synthesis, John Wiley & Sons, doi:10.1002/047084289X.rs071
  13. ^ Robert L. Metcalf (2007), "Insect Control", Ullmann's Encyclopedia of Industrial Chemistry (7th ed.), Wiley, p. 9
  14. ^ a b Robert Kapp (2005), "Fluorine", Encyclopedia of Toxicology, vol. 2 (2nd ed.), Elsevier, pp. 343–346
  15. ^ a b Greene Shepherd (2005), "Fluoride", Encyclopedia of Toxicology, vol. 2 (2nd ed.), Elsevier, pp. 342–343
  16. ^ NaF MSDS.
  17. ^ CDC - NIOSH Pocket Guide to Chemical Hazards
  18. ^ Murray TM, Ste-Marie LG. Prevention and management of osteoporosis: consensus statements from the Scientific Advisory Board of the Osteoporosis Society of Canada. 7. Fluoride therapy for osteoporosis. CMAJ. 1996;155(7):949–54. PMID 8837545.
  19. ^ National Health and Medical Research Council (Australia). A systematic review of the efficacy and safety of fluoridation [PDF]. 2007. ISBN 1-86496-415-4. Summary: Yeung CA. A systematic review of the efficacy and safety of fluoridation. Evid Based Dent. 2008;9(2):39–43. doi:10.1038/sj.ebd.6400578. PMID 18584000.