Sodium fluoride

Sodium fluoride
IUPAC name Sodium fluoride
Other names Florocid
Identifiers
CAS number	7681-49-4 Y
PubChem	5235
ChemSpider	5045 Y
UNII	8ZYQ1474W7 Y
EC number	231-667-8
UN number	1690
KEGG	C08142 N
ChEBI	CHEBI:28741 Y
ChEMBL	CHEMBL1528 Y
RTECS number	WB0350000
ATC code	A01AA01
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
Molecular formula	NaF
Molar mass	41.988713 g/mol
Appearance	White solid
Odor	odorless
Density	2.558 g/cm3
Melting point	993 °C
Boiling point	1695 °C
Solubility in water	0.5 M (20 °C)
Solubility	soluble in HF insoluble in ethanol
Vapor pressure	1 mmHg @ 1077 C°[1]
Hazards
MSDS	Sodium fluoride MSDS
EU Index	009-004-00-7
EU classification	Toxic (T) Irritant (Xi)
R-phrases	R25, R32, R36/38
S-phrases	(S1/2), S22, S36, S45
NFPA 704	0 3 0
Flash point	Non-flammable
LD50	52–200 mg/kg (oral in rats, mice, rabbits)[2]
Related compounds
Other anions	Sodium chloride Sodium bromide Sodium iodide
Other cations	Lithium fluoride Potassium fluoride Rubidium fluoride Caesium fluoride
Related compounds	TASF reagent
N (verify) (what is: Y/N?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

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.^[3][4]

The mineral form of NaF, villiaumite, is moderately rare. It is known from plutonic nepheline syenite rocks.^[5]

Production

NaF is prepared by neutralizing hydrofluoric acid or hexafluorosilicic acid (H₂SiF₆), byproducts of 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.^[6]

Applications

Sodium fluoride is sold in tablets for cavity prevention.

Fluoride salts are used to enhance the strength of teeth by the formation of fluorapatite, a naturally occurring component of tooth enamel.^[7][8] 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.^[9] Toothpaste often contains sodium fluoride to prevent cavities.^[10] Alternatively, sodium fluoride is used as a cleaning agent, e.g. as a "laundry scour".^[6] 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.^[11] Like other fluorides, sodium fluoride finds use in desilylation in organic synthesis. The fluoride is the reagent for the synthesis of fluorocarbons.

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.

Safety

See also: Fluoride poisoning

The lethal dose for a 70 kg (154 lb) human is estimated at 5–10 g.^[6] Sodium fluoride is classed as toxic by both inhalation (of dusts or aerosols) and ingestion.^[12] In high enough doses, it has been shown to affect the heart and circulatory system.

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.^[13] 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.^[14]

See also

• Fluorine
• Cryolite

References

1. ^ 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
2. ^ Martel, B.; Cassidy, K. (2004), Chemical Risk Analysis: A Practical Handbook, Butterworth–Heinemann, p. 363, ISBN 1903996651 
3. ^ Wells, A.F. (1984), Structural Inorganic Chemistry, Oxford: Clarendon Press, ISBN 0-19-855370-6 
4. ^ "Chemical and physical information" (PDF), Toxicological profile for fluorides, hydrogen fluoride, and fluorine, Agency for Toxic Substances and Disease Registry (ATDSR), September 2003, pp. 187, http://www.atsdr.cdc.gov/toxprofiles/tp11.pdf, retrieved 2008-11-01 
5. ^ (PDF) Mineral Handbook, Mineral Data Publishing, 2005, http://rruff.geo.arizona.edu/doclib/hom/villiaumite.pdf. 
6. ^ ^{a b c} Aigueperse, Jean; Paul Mollard, Didier Devilliers, Marius Chemla, Robert Faron, Renée Romano, Jean Pierre Cuer (2005), "Fluorine Compounds, Inorganic", in Ullmann, Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH, doi:10.1002/14356007.a11_307 
7. ^ Bourne, Geoffrey Howard (1986), Dietary research and guidance in health and disease, Karger, p. 153, ISBN 3-805-5434-17, http://books.google.com/?id=OW0gAAAAMAAJ , Snippet view from page 153
8. ^ Klein, Cornelis; Hurlbut, Cornelius Searle; Dana, James Dwight (1999), Manual of Mineralogy (21 ed.), Wiley, ISBN 0-471-31266-5 
9. ^ Division of Oral Health, National Center for Prevention Services, CDC (1993) (PDF), Fluoridation census 1992, http://cdc.gov/fluoridation/pdf/statistics/1992.pdf, retrieved 2008-12-29. 
10. ^ "Sodium fluoride, Molecule of the week". American Chemical Society. 2008-02-19. http://portal.acs.org/portal/acs/corg/content?_nfpb=true&_pageLabel=PP_ARTICLEMAIN&node_id=841&content_id=WPCP_008239&use_sec=true&sec_url_var=region1. Retrieved 2008-11-01. 
11. ^ Halpern, D.F. (2001), "Sodium Fluoride", Encyclopedia of Reagents for Organic Synthesis, John Wiley & Sons, doi:10.1002/047084289X.rs071 
12. ^ http://hazard.com/msds/mf/baker/baker/files/s3722.htm NaF MSDS
13. ^ 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.
14. ^ National Health and Medical Research Council (Australia). A systematic review of the efficacy and safety of fluoridation [PDF]. 2007. ISBN 1864964154. 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.