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Sodium fluoride

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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. NaF is less expensive and less hygroscopic than potassium fluoride (KF).

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

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.

Structure and basic properties

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

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

Applications

Sodium fluoride (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. 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.