Sodium fluoride

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Sodium fluoride
Sodium fluoride
Pronunciation /ˌsdiəm ˈflʊərd/[1]
IUPAC name
Sodium fluoride
Other names
ECHA InfoCard 100.028.789
EC Number 231-667-8
RTECS number WB0350000
UN number 1690
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 g/L (0 °C);
40.4 g/L (20 °C);
50.5 g/L (100 °C)[2]
Solubility slightly soluble in HF, ammonia
negligible in alcohol, acetone, SO2, dimethylformamide
Vapor pressure 1 mmHg @ 1077 C°[3]
−16.4·10−6 cm3/mol
a = 462 pm
46.82 J/mol K
51.3 J/mol K
-573.6 kJ/mol
-543.3 kJ/mol
A01AA01 (WHO) A12CD01 (WHO),
V09IX06 (WHO) (18F)
Safety data sheet [4]
H301, H315, H319[4]
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
52–200 mg/kg (oral in rats, mice, rabbits)[6]
US health exposure limits (NIOSH):
PEL (Permissible)
TWA 2.5 mg/m3[5]
REL (Recommended)
TWA 2.5 mg/m3[5]
IDLH (Immediate danger)
250 mg/m3 (as F)[5]
Related compounds
Other anions
Sodium chloride
Sodium bromide
Sodium iodide
Sodium astatide
Other cations
Lithium fluoride
Potassium fluoride
Rubidium fluoride
Caesium fluoride
Francium fluoride
Related compounds
TASF reagent
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YesYN ?)
Infobox references

Sodium fluoride (NaF) is a chemical compound and medication.[7][8] As a medication it is primarily used to prevent tooth decay in children older than 6 months in areas where the drinking water is low in fluoride.[8] Sodium fluoride is used as a liquid, pill, or paste by mouth with this use being known as fluoride therapy.[9][10]

Normal doses may occasionally result in white marks on the teeth. Excessive doses can result in brown or yellow coloring of the teeth.[9] It is believed to work mostly through direct contact with the teeth.[8] Sodium fluoride is a colorless solid. It is most often made by mixing fluorosilicic acid with sodium hydroxide.[7]

Sodium fluoride came into use to prevent tooth decay in the 1940s.[11] It is on the World Health Organization's List of Essential Medicines, the most important medications needed in a basic health system.[12] In the United Kingdom a typical month supply costs the NHS about 36 pence.[9] It is also commonly used for water fluoridation and added to toothpaste.[7][9] An estimated 90 percent of toothpastes in the United States contains some type of fluoride.[13]


Sodium fluoride is sold in tablets for cavity prevention

Dental caries[edit]

Fluoride salts are often added to municipal drinking water (as well as certain food products in some countries) for the purposes of maintaining dental health. The fluoride enhances the strength of teeth by the formation of fluorapatite, a naturally occurring component of tooth enamel.[14][15][16] Although sodium fluoride is 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.[17] Toothpaste often contains sodium fluoride to prevent cavities, although tin(II) fluoride is generally considered superior for this application.[citation needed]


Fluoride supplementation has been extensively studied for the treatment of 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.[18][19]

Medical imaging[edit]

In medical imaging, fluorine-18-labelled sodium fluoride (USP, sodium fluoride F18) is one of the oldest tracers used in positron emission tomography (PET), having been in use since the 1960s.[20] Relative to conventional bone scintigraphy carried out with gamma cameras 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 radiopharmaceuticals. 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.[21] Additionally the annihilation photons produced by decay of 18F have a high energy of 511-keV compared to 140-keV photons of 99mTc.[22]


Sodium fluoride has a variety of specialty chemical applications in synthesis and extractive metallurgy. It reacts with electrophilic chlorides including acyl chlorides, sulfur chlorides, and phosphorus chloride.[23] Like other fluorides, sodium fluoride finds use in desilylation in organic synthesis. Sodium fluoride can be used to produce fluorocarbons 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 the existence of more effective techniques (e.g. Electrofluorination, Fowler process).

Other uses[edit]

Sodium fluoride is used as a cleaning agent (e.g., as a "laundry sour").[24] 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.[25]


Fluorides, particularly aqueous solutions of sodium fluoride, are rapidly and quite extensively absorbed.[26]

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. 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.[27]

The lethal dose for a 70 kg (154 lb) human is estimated at 5–10 g.[24] Sodium fluoride is classed as toxic by both inhalation (of dusts or aerosols) and ingestion.[28] 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.[29]

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.[30] 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.[31] 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.[26]

Chemical structure[edit]

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;[32][33] 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.[34]


NaF is prepared by neutralizing hydrofluoric acid or hexafluorosilicic acid (H2SiF6), byproducts of the reaction of fluorapatite (Ca5(PO4)3F) 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:

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.[24]

See also[edit]


  1. ^ Wells, John C. (2008), Longman Pronunciation Dictionary (3rd ed.), Longman, pp. 313 and 755, ISBN 9781405881180 . According to this source, an alternative pronunciation of the second word is /ˈflɔːrd/ and, in the UK, also /ˈflərd/.
  2. ^ Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). CRC Press. p. 5.194. ISBN 1439855110. 
  3. ^ 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
  4. ^ a b Sigma-Aldrich Co., Sodium Fluoride. Retrieved on 2015-03-17.
  5. ^ a b c "NIOSH Pocket Guide to Chemical Hazards #0563". National Institute for Occupational Safety and Health (NIOSH). 
  6. ^ Martel, B.; Cassidy, K. (2004), Chemical Risk Analysis: A Practical Handbook, Butterworth–Heinemann, p. 363, ISBN 1-903996-65-1 
  7. ^ a b c Spellman, Frank R. (2008). Handbook of Water and Wastewater Treatment Plant Operations, Second Edition (2 ed.). CRC Press. p. 131. ISBN 9781420075304. 
  8. ^ a b c WHO Model Formulary 2008 (PDF). World Health Organization. 2009. pp. 501–502. ISBN 9789241547659. Retrieved 8 January 2017. 
  9. ^ a b c d British national formulary : BNF 69 (69 ed.). British Medical Association. 2015. pp. 699–700. ISBN 9780857111562. 
  10. ^ Saha, Ashok Kumar (2016). Otology & Middle Ear Surgery. JP Medical Ltd. ISBN 9789352501229. 
  11. ^ Murray, John J.; Nunn, June H.; Steele, James G. (2003). The Prevention of Oral Disease. OUP Oxford. p. 53. ISBN 9780192632791. 
  12. ^ "WHO Model List of Essential Medicines (19th List)" (PDF). World Health Organization. April 2015. Retrieved 8 December 2016. 
  13. ^ Company, DIANE Publishing (1992-09-30). Review Of Fluoride: Benefits And Risks. Report Of The Ad Hoc Subcommittee On Fluoride. DIANE Publishing. p. 12. ISBN 9780788107290. 
  14. ^ Bourne, volume editor, Geoffrey H. (1986). Dietary research and guidance in health and disease. Basel: Karger. p. 153. ISBN 3-8055-4341-7. 
  15. ^ Jr, Cornelis Klein, Cornelius S. Hurlbut, (1999). Manual of mineralogy : (after James D. Dana) (21st ed., rev. ed.). New York: J. Wiley. ISBN 0-471-31266-5. 
  16. ^ Selwitz, Robert H; Ismail, Amid I; Pitts, Nigel B (January 2007). "Dental caries". The Lancet. 369 (9555): 51–59. doi:10.1016/S0140-6736(07)60031-2. PMID 17208642. 
  17. ^ Division of Oral Health, National Center for Prevention Services, CDC (1993), Fluoridation census 1992 (PDF), retrieved 2008-12-29. 
  18. ^ Haguenauer, D; Welch, V; Shea, B; Tugwell, P; Wells, G (2000). "Fluoride for treating postmenopausal osteoporosis". The Cochrane Database of Systematic Reviews (4): CD002825. doi:10.1002/14651858.CD002825. PMID 11034769. 
  19. ^ Vestergaard, P; Jorgensen, NR; Schwarz, P; Mosekilde, L (March 2008). "Effects of treatment with fluoride on bone mineral density and fracture risk—a meta-analysis". Osteoporosis international : a journal established as result of cooperation between the European Foundation for Osteoporosis and the National Osteoporosis Foundation of the USA. 19 (3): 257–68. doi:10.1007/s00198-007-0437-6. PMID 17701094. 
  20. ^ Blau, Monte; Ganatra, Ramanik; Bender, Merrill A. (January 1972). "18F-fluoride for bone imaging". Seminars in Nuclear Medicine. 2 (1): 31–37. doi:10.1016/S0001-2998(72)80005-9. 
  21. ^ Ordonez, A. A.; DeMarco, V. P.; Klunk, M. H.; Pokkali, S.; Jain, S.K. (October 2015). "Imaging Chronic Tuberculous Lesions Using Sodium [18F]Fluoride Positron Emission Tomography in Mice". Molecular Imaging and Biology. 17 (5): 609–614. doi:10.1007/s11307-015-0836-6. 
  22. ^ Grant, F. D.; Fahey, F. H.; Packard, A. B.; Davis, R. T.; Alavi, A.; Treves, S. T. (12 December 2007). "Skeletal PET with 18F-Fluoride: Applying New Technology to an Old Tracer". Journal of Nuclear Medicine. 49 (1): 68–78. doi:10.2967/jnumed.106.037200. PMID 18077529. 
  23. ^ Halpern, D.F. (2001), "Sodium Fluoride", Encyclopedia of Reagents for Organic Synthesis, John Wiley & Sons, doi:10.1002/047084289X.rs071, ISBN 0471936235 
  24. ^ a b c Aigueperse, Jean; Mollard, Paul; Devilliers, Didier; Chemla, Marius; Faron, Robert; Romano, Renée; Cuer, Jean Pierre (2005), "Fluorine Compounds, Inorganic", in Ullmann, Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH, doi:10.1002/14356007.a11_307, ISBN 3527306730 
  25. ^ Metcalf, Robert L. (2007), "Insect Control", Ullmann's Encyclopedia of Industrial Chemistry (7th ed.), Wiley, p. 9 
  26. ^ a b Kapp, Robert (2005), "Fluorine", Encyclopedia of Toxicology, 2 (2nd ed.), Elsevier, pp. 343–346 
  27. ^ Greene Shepherd (2005), "Fluoride", Encyclopedia of Toxicology, 2 (2nd ed.), Elsevier, pp. 342–343 
  28. ^ NaF MSDS.
  29. ^ CDC – NIOSH Pocket Guide to Chemical Hazards
  30. ^ 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.
  31. ^ 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. Lay summary: NHMRC, 2007.
  32. ^ Wells, A.F. (1984), Structural Inorganic Chemistry, Oxford: Clarendon Press, ISBN 0-19-855370-6 
  33. ^ "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 
  34. ^ Mineral Handbook (PDF) (version 1), Mineral Data Publishing, 2005.