Sodium fluoride
From Wikipedia, the free encyclopedia
| Sodium fluoride | |
|---|---|
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| IUPAC name | Sodium fluoride |
| Other names | Florocid |
| Identifiers | |
| CAS number | 7681-49-4 |
| EINECS number | |
| RTECS number | WB0350000 |
| 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 | |
| MSDS | Sodium fluoride MSDS |
| 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 references |
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Sodium fluoride is the chemical compound with the formula NaF. This colourless solid is the main source of the fluoride ion in diverse applications. Sodium fluoride is less expensive and less hygroscopic than potassium fluoride.
The mineral form of NaF, villiaumite, is moderately rare. It is known from plutonic nepheline syenite rocks.[1]
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[edit] Structure and basic properties
Sodium fluoride is an ionic compound, dissolving to give separated Na+ and F− ions. It crystallizes in the cubic (sodium chloride) motif where both Na+ and F− occupy octahedral coordination sites.[2][3]
[edit] 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.[4]
[edit] Applications
Fluoride salts are used to enhance the strength of teeth by the formation of fluorapatite, a naturally occurring component of tooth enamel. In the US, sodium fluoride was once used to fluoridate drinking water but this has been displaced by hexafluorosilicic acid (H2SiF6) or the related sodium salt (Na2SiF6). Toothpaste often contains sodium fluoride to prevent cavities.[5]
Alternatively, sodium fluoride is used as a cleaning agent, often to remove iron stains. A variety of specialty chemical applications exist in synthesis and extractive metallurgy. The fluoride is the reagent for the synthesis of fluorocarbons. Representative substrates include electrophilic chlorides including acyl chlorides, sulfur chlorides, and phosphorus chloride.[6] Like other fluorides, sodium fluoride finds use in desilylation in organic synthesis.
In medicial 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.
[edit] Safety
- See also: Water fluoridation and Fluoride poisoning
Sodium fluoride is classed as toxic by both inhalation and ingestion.[7] In high enough doses, it has been shown to affect the heart and circulatory system, and the lethal dose for a 70 kg human is estimated at 5–10 g.[4]
[edit] See also
[edit] References
- ^ "Mineral Handbook" (PDF). Mineral Data Publishing (2005).
- ^ Wells, A.F. (1984). Structural Inorganic Chemistry. Clarendon Press. ISBN 0-19-855370-6.
- ^ "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 on 1 November 2008
- ^ a b 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
- ^ "Sodium fluoride, Molecule of the week". American Chemical Society (2008-02-19). Retrieved on 2008-11-01.
- ^ Halpern, D.F. (2001), "Sodium Fluoride", Encyclopedia of Reagents for Organic Synthesis, John Wiley & Sons, doi:
- ^ http://www.jtbaker.com/msds/englishhtml/S3722.htm NaF MSDS
[edit] External links
- Chemical profile for sodium fluoride at Scorecard, the pollution information site
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