Sodium fluoride

Sodium fluoride
Names
Pronunciation /ˌsoʊdiəm ˈflʊəraɪd/[1]
IUPAC name
Sodium fluoride
Other names
Florocid
Identifiers
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.028.789
EC Number 231-667-8
KEGG
RTECS number WB0350000
UNII
UN number 1690
Properties
NaF
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
1.3252
Structure
Cubic
a = 462 pm
Octahedral
Thermochemistry
46.82 J/mol K
51.3 J/mol K
-573.6 kJ/mol
-543.3 kJ/mol
Pharmacology
A01AA01 (WHO) A12CD01 (WHO),
V09IX06 (WHO) (18F)
Hazards
Safety data sheet [4]
GHS pictograms
GHS signal word Danger
H301, H315, H319, H335[4]
NFPA 704
Flammability code 0: Will not burn. E.g., waterHealth code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gasReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogenSpecial hazards (white): no codeNFPA 704 four-colored diamond
0
3
0
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
52–200 mg/kg (oral in rats, mice, rabbits)[5]
US health exposure limits (NIOSH):
PEL (Permissible)
 TWA 2.5 mg/m3[6]
REL (Recommended)
 TWA 2.5 mg/m3[6]
IDLH (Immediate danger)
 250 mg/m3 (as F)[6]
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).
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Infobox references

Sodium fluoride (NaF) is an inorganic compound with the formula NaF. It is colorless or white solid that is readily soluble in water. It is a common source of fluoride in the production of pharmaceuticals and it is used to prevent cavities.

Production

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 sodium bifluoride (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]


Uses

Sodium fluoride is sold in tablets for cavity prevention

Dental caries

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.[8][9][10] 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.[11]

Osteoporosis

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.[12][13]

Medical imaging

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.[14] 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.[15] Additionally the annihilation photons produced by decay of 18F have a high energy of 511-keV compared to 140-keV photons of 99mTc.[16]

Chemistry

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.[17] 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

Sodium fluoride is used as a cleaning agent (e.g., as a "laundry sour").[7] 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.[18]

Safety

Fluorides, particularly aqueous solutions of sodium fluoride, are rapidly and quite extensively absorbed by the human body.[19]

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.[20]

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.[21] 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.[22]

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.[23] 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.[24] 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.[19]

As of 2014 there have been only three reported cases of fluoride toxicity associated with the ingestion of fluoride-containing toothpaste. As an example, one of these involved a 45 year old woman who came to her doctor complaining of unusual swelling and pain in her fingers. Tests showed elevated levels of fluoride in her urine and blood. When questioned about this, the woman admitted to the regular ingestion of large amounts of toothpaste, consuming a tube of it every two days and swallowing 68.5 mg of fluoride every day, because she "liked the taste".[25] When asked to switch to a non-fluoride form of toothpaste, her fluoride levels dropped dramatically and her condition subsided.[25]

Chemical structure

Sodium fluoride is an inorganic ionic compound, dissolving in water 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;[26][27] its lattice spacing, approximately 462 pm, is somewhat smaller than that of sodium chloride.

Occurrence

The mineral form of NaF, villiaumite, is moderately rare. It is known from plutonic nepheline syenite rocks.[28]

See also

References

  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 /ˈflɔːraɪd/ and, in the UK, also /ˈfluːəraɪd/.
  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. 1 2 Sigma-Aldrich Co., Sodium Fluoride. Retrieved on 2015-03-17.
  5. ↑ Martel, B.; Cassidy, K. (2004), Chemical Risk Analysis: A Practical Handbook, Butterworth–Heinemann, p. 363, ISBN 1-903996-65-1
  6. 1 2 3 "NIOSH Pocket Guide to Chemical Hazards #0563". National Institute for Occupational Safety and Health (NIOSH).
  7. 1 2 3 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.
  8. ↑ Bourne, volume editor, Geoffrey H. (1986). Dietary research and guidance in health and disease. Basel: Karger. p. 153. ISBN 3-8055-4341-7.
  9. ↑ 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.
  10. ↑ 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.
  11. ↑ Division of Oral Health, National Center for Prevention Services, CDC (1993), Fluoridation census 1992 (PDF), retrieved 2008-12-29.
  12. ↑ 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.
  13. ↑ 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.
  14. ↑ 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.
  15. ↑ 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. PMC 4561601.
  16. ↑ 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.
  17. ↑ Halpern, D.F. (2001), "Sodium Fluoride", Encyclopedia of Reagents for Organic Synthesis, John Wiley & Sons, doi:10.1002/047084289X.rs071, ISBN 0471936235
  18. ↑ Metcalf, Robert L. (2007), "Insect Control", Ullmann's Encyclopedia of Industrial Chemistry (7th ed.), Wiley, p. 9
  19. 1 2 Kapp, Robert (2005), "Fluorine", Encyclopedia of Toxicology, 2 (2nd ed.), Elsevier, pp. 343–346
  20. ↑ Greene Shepherd (2005), "Fluoride", Encyclopedia of Toxicology, 2 (2nd ed.), Elsevier, pp. 342–343
  21. ↑ NaF MSDS. hazard.com
  22. ↑ CDC – NIOSH Pocket Guide to Chemical Hazards
  23. ↑ 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.
  24. ↑ 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.
  25. 1 2 "Fluoride toothpaste–what are the dangers of chronic ingestion in adults? Q&A 154.4" (PDF). May 2014.
  26. ↑ Wells, A.F. (1984), Structural Inorganic Chemistry, Oxford: Clarendon Press, ISBN 0-19-855370-6
  27. ↑ "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
  28. ↑ Mineral Handbook (PDF) (version 1), Mineral Data Publishing, 2005.
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