|Molar mass||41.988173 g/mol|
|Appearance||White to greenish solid|
|Melting point||993 °C (1,819 °F; 1,266 K)|
|Boiling point||1,704 °C (3,099 °F; 1,977 K)|
|36.4 g/L (0 °C);
40.4 g/L (20 °C);
50.5 g/L (100 °C)
|Solubility||slightly soluble in HF, ammonia
negligible in alcohol, acetone, SO2, dimethylformamide
|Vapor pressure||1 mmHg @ 1077 C°|
Refractive index (nD)
a = 462 pm
|46.82 J/mol K|
|51.3 J/mol K|
Std enthalpy of
Gibbs free energy (ΔfG˚)
|Safety data sheet|||
|H301, H315, H319|
|EU classification|| T – Toxic
Xi – Irritant
|R-phrases||R25, R32, R36/38|
|S-phrases||(S1/2), S22, S36, S45|
|Lethal dose or concentration (LD, LC):|
LD50 (Median dose)
|52–200 mg/kg (oral in rats, mice, rabbits)|
|US health exposure limits (NIOSH):|
|TWA 2.5 mg/m3|
|TWA 2.5 mg/m3|
IDLH (Immediate danger
|250 mg/m3 (as F)|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is: / ?)(|
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; its lattice spacing, approximately 462 pm, is somewhat smaller than that of sodium chloride.
NaF is prepared by neutralizing hydrofluoric acid or hexafluorosilicic acid (H2SiF6), byproducts of the reaction of fluorapatite (Ca5(PO4)3F) (from phosphate rock) from 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.
Treatment of osteoporosis
Sodium fluoride is used to treat osteoporosis, particularly in postmenopausal women. It stimulates bone formation and increases bone density in women with osteoporosis by 8 percent per year in the lumbar spine and by 4 percent in the proximal femur, but it decreases cortical bone density in the radius by 2 percent per year. Bone with excess fluoride content has an abnormal structure (possibly due to mineralization outstripping supply of calcium and vitamin D) and its fragility may be increased. Thus sodium fluoride therapy results in large increases in bone mineral density, but the effect on fracture rates is small. There is evidence that using smaller doses of slow release sodium fluoride may be more effective and also decreases the rate of side effects, which can include severe gastrointestinal irritation and joint pain.
In medical imaging, fluorine-18-labelled sodium fluoride (USP, sodium fluoride F18) is one of the oldest tracers used in positron emission tomography (PET), having been in use since the 1960s. Relative to conventional bone scintigraphy carried out with gamma cameras or SPECT systems, PET offers more sensitivity and spatial resolution. Fluorine-18 has a half-life of 110 min, which requires it to be used promptly once produced; this logistical limitation hampered its adoption in the face of the more convenient technetium-99m-labelled radiopharmaceuticals. However fluorine-18 is generally considered to be a superior radiopharmaceutical for skeletal imaging. In particular it has a high and rapid bone uptake accompanied by very rapid blood clearance, which results in a high bone-to-background ratio in a short time. Additionally the annihilation photons produced by decay of 18F have a high energy of 511-keV compared to 140-keV photons of 99mTc.
Fluoride salts are often added to drinking water (as well as certain food products in some countries) for the purposes of maintaining dental health. The fluoride enhances the strength of teeth by the formation of fluorapatite, a naturally occurring component of tooth enamel. Although sodium fluoride is 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. Toothpaste often contains sodium fluoride to prevent cavities, although tin(II) fluoride and sodium monofluorophosphate are generally considered superior for this application.
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. Like other fluorides, sodium fluoride finds use in desilylation in organic synthesis. Sodium fluoride can be used to produce fluorocarbons via the Finkelstein reaction; this process has the advantage of being simple to perform on a small scale but is rarely used on an industrial scale due the existence of more effective techniques (e.g. Electrofluorination, Fowler process).
Sodium fluoride is used as a cleaning agent (e.g., as a "laundry sour"). Sodium fluoride is used as a stomach poison for plant-feeding insects. Inorganic fluorides such as fluorosilicates and sodium fluoride complex magnesium ions as magnesium fluorophosphate. They inhibit enzymes such as enolase that require Mg2+ as a prosthetic group. Thus, fluoride poisoning prevents phosphate transfer in oxidative metabolism.
Fluorides, particularly aqueous solutions of sodium fluoride, are rapidly and quite extensively absorbed.
Fluorides interfere with electron transport and calcium metabolism. Calcium is essential for maintaining cardiac membrane potentials and in regulating coagulation. Large ingestion of fluoride salts or hydrofluoric acid may result in fatal arrhythmias due to profound hypocalcemia. Recreational inhalation of fluoridated hydrocarbon refrigerants like Freon has been associated with "sudden sniffing death", which is thought to be a fatal arrhythmia caused by myocardial sensitization to catecholamines.
Chronic over-absorption can cause hardening of bones, calcification of ligaments, and buildup on teeth. Fluoride can cause irritation or corrosion to eyes, skin, and nasal membranes.
The lethal dose for a 70 kg (154 lb) human is estimated at 5–10 g. Sodium fluoride is classed as toxic by both inhalation (of dusts or aerosols) and ingestion. In high enough doses, it has been shown to affect the heart and circulatory system. For occupational exposures, the Occupational Safety and Health Administration and the National Institute for Occupational Safety and Health have established occupational exposure limits at 2.5 mg/m3 over an eight-hour time-weighted average.
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. 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. A chronic fluoride ingestion of 1 ppm of fluoride in drinking water can cause mottling of the teeth (fluorosis) and an exposure of 1.7 ppm will produce mottling in 30–50 % of patients.
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