Sodium fluoride
From Wikipedia, the free encyclopedia
| Sodium fluoride | |
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| 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 |
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Sodium fluoride is an ionic 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 KF.
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[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
Using an excess of HF gives the 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.[1]
[edit] Structure, properties, Uses
NaF crystallizes in the sodium chloride motif where both Na+ and F− occupy octahedral coordination sites.[2]
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 its sodium salt (Na2SiF6). Toothpaste often contains sodium fluoride to prevent cavities.
[edit] Safety
- See also: Water fluoridation and water fluoridation controversy
The lethal dose for a 70 kg human is estimated at 5 – 10 g.[1]
[edit] See also
- Fluorine
- Water fluoridation
- Fluoride poisoning
- Caries
- Osteoporosis
- Insecticide
- Cryolite
- Light metals
[edit] References
- ^ 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
- ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
- ^ Halpern, D. F. “Sodium Fluoride” Encyclopedia of Reagents for Organic Synthesis, 2001, John Wiley & Sons. DOI: 10.1002/047084289X.rs071.
[edit] 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
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