Sodium fluoride
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Sodium fluoride
IUPAC name	Sodium fluoride
Identifiers
CAS number	[7681-49-4]
Properties
Molecular formula	NaF
Molar mass	41.99 g/mol
Appearance	White solid
Density	2.558 g/cm³, solid
Melting point	993 °C
Boiling point	1700 °C
Solubility in water	4.13 g/100 g at 25 °C
Hazards
EU classification	Toxic (T)
NFPA 704	0 2 0
R-phrases	R25, R32, R36, R38
S-phrases	S22, S36, S45
Flash point	Non-flammable.
Related compounds
Other anions	sodium chloride sodium bromide sodium iodide
Other cations	potassium fluoride calcium fluoride caesium fluoride
Related bases	None listed.
Related compounds	TASF reagent
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references

Sodium fluoride is the chemical compound with the formula NaF. This colourless solid is the main source of the fluoride ion in diverse applications. NaF is less expensive and less hygroscopic than potassium fluoride (KF).

Production

NaF is prepared by neutralizing waste hydrofluoric acid resulting from the production of superphosphate fertilizer. It is also generated by treating sodium hydroxide and sodium carbonate with hydrofluoric acid, followed by concentrating the resulting solutions, sometimes with the addition of alcohols 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.^[1]

Structure and basic properties

NaF crystallizes in the sodium chloride motif where both Na⁺ and F⁻ occupy octahedral coordination sites.^[2]

Sodium fluoride is sold in tablets for cavity prevention.

It is an ionic compound, dissolving to give separated Na⁺ and F^- ions.

Applications

NaF is used as a cleaning agent, often to remove iron stains. A variety of specialty chemical applications exist in synthesis and extractive metallurgy. NaF is a reagent for the synthesis of fluorocarbons. Representative substrates include electrophilic chlorides including acyl chlorides, sulfur chlorides, and phosphorus chloride.^[3] Like other fluorides, NaF finds use in desilylation in organic synthesis.

Fluoride salts were used widely to enhance the strength of teeth by the formation of fluoroapatite, a naturally occurring component of tooth enamel. In the US, NaF was once used to fluoridate drinking water but its use has been displaced by hexafluorosilicic acid (H₂SiF₆) or the related sodium salt (Na₂SiF₆). Toothpaste often contains sodium fluoride to prevent cavities.

Safety

See also: Water fluoridation and water fluoridation controversy

The lethal dose for a 70 kg human is estimated at 5 – 10 g.^[1]

See also

• Fluorine
• Water fluoridation
• Fluoride poisoning
• Caries
• Osteoporosis
• Insecticide
• Cryolite
• Light metals

References

1. ^ ^{a b} Jean Aigueperse, Paul Mollard, Didier Devilliers, Marius Chemla, Robert Faron, Renée Romano, Jean Pierre Cuer, “Fluorine Compounds, Inorganic” in Ullmann’s Encyclopedia of Industrial Chemistry 2005 Wiley-VCH, Weinheim. doi:10.1002/14356007.a11 307
2. ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
3. ^ Halpern, D. F. “Sodium Fluoride” Encyclopedia of Reagents for Organic Synthesis, 2001, John Wiley & Sons. DOI: 10.1002/047084289X.rs071.

External links

• Sodium Fluoride at chemistry.org (American Chemical Society)
• Chemical Profile for Sodium Fluoride (CAS Number: 7681-49-4) at Scorecard, the pollution information site