Sodium fluoride

Sodium fluoride (NaF) is an inorganic compound with the formula NaF. It is used in trace amounts in the fluoridation of drinking water, in toothpaste, in metallurgy, and as a flux. It is a colorless or white solid that is readily soluble in water. It is a common source of fluoride in the production of pharmaceuticals and is used to prevent dental cavities.

Sodium fluoride
Names
Pronunciation /ˌsoʊdiəm ˈflʊəraɪd/[1]
IUPAC name
Sodium fluoride
Other names
Florocid
Identifiers
CAS Number
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.028.789
EC Number
  • 231-667-8
KEGG
PubChem CID
RTECS number
  • WB0350000
UNII
UN number 1690
InChI
  • InChI=1S/FH.Na/h1H;/q;+1/p-1 Y
    Key: PUZPDOWCWNUUKD-UHFFFAOYSA-M Y
  • InChI=1/FH.Na/h1H;/q;+1/p-1
    Key: PUZPDOWCWNUUKD-REWHXWOFAH
Properties
Chemical formula
 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)
Solubility in water
 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]
Magnetic susceptibility (χ)
 −16.4·10−6 cm3/mol
Refractive index (nD)
 1.3252
Structure
Crystal structure
 Cubic
Lattice constant
a = 462 pm
Molecular shape
 Octahedral
Thermochemistry
Heat capacity (C)
 46.82 J/(mol K)
Std molar
entropy (S⦵298)
 51.3 J/(mol K)
Std enthalpy of
formation (ΔfH⦵298)
 -573.6 kJ/mol
Gibbs free energy (ΔfG⦵)
 -543.3 kJ/mol
Pharmacology
A01AA01 (WHO) A12CD01 (WHO),
V09IX06 (WHO) (18F)
Hazards
GHS labelling:
Pictograms
Signal word
 Danger
Hazard statements
 H301, H315, H319, H335[4]
NFPA 704 (fire diamond)
3
0
0
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
52–130 mg/kg (oral in rats, mice, rabbits)[5]
NIOSH (US health exposure limits):
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]
Safety data sheet (SDS) [4]
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 YN ?)
Infobox references

In 2017, it was the 247th most commonly prescribed medication in the United States, with more than one million prescriptions.[7][8]

Uses
Sodium fluoride is sold in tablets for cavity prevention

Dental caries
See also: Fluoride therapy and 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.[9][10][11] Although sodium fluoride is used to fluoridate water and 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 United States.[12]

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.[13][14]

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.[15] 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.[16] Additionally the annihilation photons produced by decay of 18F have a high energy of 511 keV compared to the 140 keV photons of 99mTc.[17]

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.[18] 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 to 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").[19]

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

Over a century ago, sodium fluoride was used as a stomach poison for plant-feeding insects.[20] 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.[21]

Safety
See also: Fluoride poisoning

The lethal dose for a 70 kg (154 lb) human is estimated at 5–10 g.[19]

Fluorides, particularly aqueous solutions of sodium fluoride, are rapidly and quite extensively absorbed by the human body.[22]

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 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.[23]

Sodium fluoride is classed as toxic by both inhalation (of dusts or aerosols) and ingestion.[24] 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.[25]

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-coated versions of sodium fluoride do not have significant gastric side effects, and have milder and less frequent complications in the bones.[26] 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.[27] 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.[22]

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;[28][29] 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.[30]

Production

NaF is prepared by neutralizing hydrofluoric acid or hexafluorosilicic acid (H2SiF6), both 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.[19]

See also
Wikimedia Commons has media related to Sodium fluoride.

References
  1. Wells, John C. (2008), Longman Pronunciation Dictionary (3rd ed.), Longman, pp. 313 and 755, ISBN 978-1-4058-8118-0. 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 978-1-4398-5511-9.
  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. Sigma-Aldrich Co., Sodium Fluoride.
  5. Martel, B.; Cassidy, K. (2004), Chemical Risk Analysis: A Practical Handbook, Butterworth–Heinemann, p. 363, ISBN 978-1-903996-65-2
  6. NIOSH Pocket Guide to Chemical Hazards. "#0563". National Institute for Occupational Safety and Health (NIOSH).
  7. "The Top 300 of 2020". ClinCalc. Retrieved 11 April 2020.
  8. "Sodium Fluoride – Drug Usage Statistics". ClinCalc. Retrieved 11 April 2020.
  9. Bourne, volume editor, Geoffrey H. (1986). Dietary research and guidance in health and disease. Basel: Karger. p. 153. ISBN 978-3-8055-4341-5. {{cite book}}: |first1= has generic name (help)
  10. Klein, Cornelis (1999). Hurlbut, Cornelius S. (ed.). Manual of Mineralogy (after James D. Dana) (21st ed., rev. ed.). New York: J. Wiley. ISBN 978-0-471-31266-6.
  11. 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. S2CID 204616785.
  12. 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)
  13. Haguenauer, D; Welch, V; Shea, B; Tugwell, P; Wells, G (2000). "Fluoride for treating postmenopausal osteoporosis". The Cochrane Database of Systematic Reviews. 2010 (4): CD002825. doi:10.1002/14651858.CD002825. PMC 8453489. PMID 11034769.
  14. 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. 19 (3): 257–68. doi:10.1007/s00198-007-0437-6. PMID 17701094. S2CID 25890845.
  15. 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. PMID 5059349.
  16. 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. PMID 25750032.
  17. 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.
  18. Halpern, D.F. (2001), "Sodium Fluoride", Encyclopedia of Reagents for Organic Synthesis, John Wiley & Sons, doi:10.1002/047084289X.rs071, ISBN 978-0-471-93623-7
  19. Aigueperse, Jean; Mollard, Paul; Devilliers, Didier; Chemla, Marius; Faron, Robert; Romano, René; Cuer, Jean Pierre (2000). "Fluorine Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a11_307.
  20. House, James E.; House, Kathleen A. (2015-09-10). Descriptive Inorganic Chemistry. Academic Press. p. 397. ISBN 978-0-12-802979-4.
  21. Metcalf, Robert L. (2007), "Insect Control", Ullmann's Encyclopedia of Industrial Chemistry (7th ed.), Wiley, p. 9
  22. Kapp, Robert (2005), "Fluorine", Encyclopedia of Toxicology, vol. 2 (2nd ed.), Elsevier, pp. 343–346
  23. Greene Shepherd (2005), "Fluoride", Encyclopedia of Toxicology, vol. 2 (2nd ed.), Elsevier, pp. 342–343
  24. NaF MSDS. hazard.com
  25. CDC – NIOSH Pocket Guide to Chemical Hazards
  26. Murray TM, Ste-Marie LG (1996). "Prevention and management of osteoporosis: consensus statements from the Scientific Advisory Board of the Osteoporosis Society of Canada. 7. Fluoride therapy for osteoporosis". CMAJ. 155 (7): 949–54. PMC 1335460. PMID 8837545.
  27. National Health and Medical Research Council (Australia) (2007). A systematic review of the efficacy and safety of fluoridation (PDF). ISBN 978-1-86496-415-8. Summary: Yeung CA (2008). "A systematic review of the efficacy and safety of fluoridation". Evid Based Dent. 9 (2): 39–43. doi:10.1038/sj.ebd.6400578. PMID 18584000.
  28. Wells, A.F. (1984), Structural Inorganic Chemistry, Oxford: Clarendon Press, ISBN 978-0-19-855370-0
  29. "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
  30. Mineral Handbook (PDF), Mineral Data Publishing, 2005.
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