Flouride Defined

OWLS™ Water Education: Fluoride Defined

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Fluoride /flʊərd/ is an inorganic anion of fluorine with the chemical formula F
. It contributes no color to fluoride salts. Fluorite is roughly 49% fluoride by mass, making fluoride the main component of fluorite apart from calcium ions. Fluoride contributes a distinctive bitter taste, but no odor to fluoride salts. Its salts are mainly mined as a precursor to hydrogen fluoride. Although it is classified as a weak base, concentrated fluoride solutions are corrosive and can attack the skin.

Fluoride is the simplest unary fluorine anion, the other being the tentatively investigated difluoridofluorate(1-) anion. Its salts are important chemical reagents and industrial chemicals, mainly used in the production of hydrogen fluoride for fluorocarbons. Structurally, and to some extent chemically, the fluoride ion resembles the hydroxide ion. Fluoride ions occur on earth in several minerals, particularly fluorite, but are only present in trace quantities in water.


The systematic name fluoride, the valid IUPAC name, is determined according to the additive nomenclature. However, the name fluoride is also used in compositional IUPAC nomenclature which does not take the nature of bonding involved. Examples of such naming are sulfur hexafluoride and beryllium fluoride, which contains no fluoride ions whatsoever.

Fluoride is also used non-systematically, to describe compounds which releases hydrogen fluoride upon acidification, or a compound that otherwise incorporates fluorine in some form, such as methyl fluoride and fluorosilicic acid. Hydrogen fluoride is itself an example of a non-systematic name of this nature. However, it is also a trivial name, and the preferred IUPAC name for fluorane.


Fluorite crystals

Many fluoride minerals are known, but of paramount commercial importance is fluorite.[3] It is composed of calcium fluoride, with small impurities. The soft, colorful mineral is found worldwide and is common.

Seawater fluoride levels are usually in the range of 0.86 to 1.4 mg/L, and average 1.1 mg/L.[4] For comparison, chloride concentration in seawater is about 19 mg/L. The low concentration of fluoride reflects the insolubility of the alkaline earth fluorides, e.g., CaF2.

Fluoride is found naturally in low concentration in drinking water and foods. Fresh water supplies generally contain between 0.01–0.3 ppm.[5][6] In some locations, the fresh water contains dangerously high levels of fluoride, leading to serious health problems.

Chemical properties


Fluoride can act as a base. It can combine with a proton (H+

F + H+ → HF

This neutralization reaction forms hydrogen fluoride (HF), the conjugate acid of fluoride.

In aqueous solution, fluoride has a pKb value of 10.8. It is therefore a weak base, and tends to remain as the fluoride ion rather than generating a substantial amount of hydrogen fluoride. That is, the following equilibrium favours the left-hand side in water:

F + H2O is in a disfavored equilibrium with HF + HO

However, upon prolonged contact with moisture, fluoride salts will decompose to their respective hydroxides or oxides, as the hydrogen fluoride escapes. Fluoride is distinct in this regard among the halides. The identity of the solvent can have a dramatic effect on the equilibrium shifting it to the right-hand side, greatly increasing the rate of decomposition.


Counter-intuitively, unlike the lower halides, the variety of possible salts containing true fluoride ions is rather restricted. Most metal fluorides are highly polar, covalently bonded molecular networks instead of ionic lattices. All of the alkali metal fluorides, and most of the alkali earth metal fluorides are true salts. In such true salts, fluoride typically assumes the primitive or face-centred cubic motifs. Fluoride is also found with weakly coordinating counter cations, such as in ammonium fluoride, in which it assumes the close-packed hexagonal motif. Under aqueous conditions, fluoride exists as a trigonal pyramidal-shaped hydrated complex, namely [F(H2O)3]. Fluoride has the smallest monatomic, crystal and effective, ionic radii: 199 and 133 pm, respectively.

Inorganic Chemistry

Upon treatment with a standard acid, fluoride salts convert to hydrogen fluoride and metal salts. With strong acids, it can be doubly protonated to give H
. Oxidation of fluoride gives fluorine. Solutions of inorganic fluorides in water contain F and bifluoride HF
.[7] Few inorganic fluorides are soluble in water without undergoing significant hydrolysis. In terms of its reactivity, fluoride differs significantly from chloride and other halides, and is more strongly solvated in protic solvents due to its smaller radius/charge ratio. Its closest chemical relative is hydroxide. When relatively unsolvated, for example in nonprotic solvents, fluoride anions are called “naked”. Naked fluoride is a very strong Lewis base,[8] it is easily reacted with Lewis acids, forming strong adducts. Fluoride is susceptible to extreme ultraviolet radiation, ejecting an electron to become highly reactive atomic fluorine. It has a standard electrode potential of 2.87 Volts.


At physiological pHs, hydrogen fluoride is usually fully ionised to fluoride. In biochemistry, fluoride and hydrogen fluoride are equivalent. Fluorine, in the form of fluoride, is considered to be a micronutrient for human health, necessary to prevent dental cavities, and to promote healthy bone growth.[5] The tea plant (Camellia sinensis L.) is a known accumulator of fluorine compounds, released upon forming infusions such as the common beverage. The fluorine compounds decompose into products including fluoride ions. Fluoride is the most bioavailable form of fluorine, and as such, tea is potentially a vehicle for fluoride dosing.[9] Approximately, fifty percent of absorbed fluoride is excreted renally with a twenty four hour period. The remainder can be retained in the oral cavity, and lower digestive tract. Fasting dramatically increases the rate of fluoride absorption to near hundred percent, from a sixty to eighty percent when taken with food.[9] Per a 2013 study, it was found that consumption of one litre of tea a day, can potentially supply the daily recommended intake of 4 mg per day. Some lower quality brands can supply up to a 120 percent of this amount. Fasting can increase this to 150 percent. The study indicates that tea drinking communities are at an increased risk of fluorosis, in the case where water fluoridation is in effect.[9] Fluoride ion in low doses in the mouth reduces tooth decay. For this reason, it is used in toothpaste and water fluoridation. At much higher doses, fluoride causes health complications and can be toxic.


Fluoride salts and hydrofluoric acid are the main fluorides of industrial value. Compounds with C-F bonds fall into the realm of organofluorine chemistry. The main uses of fluoride, in terms of volume, are in the production of cryolite, Na3AlF6. It is used in aluminium smelting. Formerly, it was mined, but now it is derived from hydrogen fluoride. Fluorite is used on a large scale to separate slag in steel-making. Mined fluorite (CaF2) is a commodity chemical used in steel-making.

Hydrofluoric acid and its anhydrous form, hydrogen fluoride, is also used in the production of fluorocarbons. Hydrofluoric acid has a variety of specialized applications, including its ability to dissolve glass.[3]

Cavity prevention

Fluoride is sold in tablets for cavity prevention.

Fluoride-containing compounds, such as sodium fluoride or sodium monofluorophosphate are used in topical and systemic fluoride therapy for preventing tooth decay. They are used for water fluoridation and in many products associated with oral hygiene.[10] Originally, sodium fluoride was used to fluoridate water; hexafluorosilicic acid (H2SiF6) and its salt sodium hexafluorosilicate (Na2SiF6) are more commonly used additives, especially in the United States. The fluoridation of water is known to prevent tooth decay[11][12] and is considered by the U.S. Centers for Disease Control and Prevention as “one of 10 great public health achievements of the 20th century”.[13][14] In some countries where large, centralized water systems are uncommon, fluoride is delivered to the populace by fluoridating table salt. For the method of action for cavity prevention (see Fluoride therapy). Fluoridation of water has its critics (see Water fluoridation controversy).[15]

Biochemical reagent

Fluoride salts are commonly used in biological assay processing to inhibit the activity of phosphatases, such as serine/threonine phosphatases.[16] Fluoride mimics the nucleophilic hydroxide ion in these enzymes’ active sites.[17] Beryllium fluoride and aluminium fluoride are also used as phosphatase inhibitors, since these compounds are structural mimics of the phosphate group and can act as analogues of the transition state of the reaction.[18][19]

Estimated daily intake

Daily intakes of fluoride can vary significantly according to the various sources of exposure. Values ranging from 0.46 to 3.6–5.4 mg/day have been reported in several studies (IPCS, 1984).[20] In areas where water is fluoridated this can be expected to be a significant source of fluoride, however fluoride is also naturally present in huge range of foods, in a wide range of concentrations.[21] The maximum safe daily consumption of fluoride is 10mg for an adult.

Examples of fluoride content
Food/Drink Fluoride
(mg per 100g)
Portion Fluoride
(mg per portion)
Black Tea (brewed) 0.373 1 cup, 240g (8 fl oz) 0.884
Raisins, seedless 0.234 small box, 43g (1.5 oz) 0.033
Table wine 0.153 Bottle, 750ml (26.4 fl oz) 1.150
Municipal tap-water,
0.081 Recommended daily intake,
3 litres (0.79 US gal)
Baked potatoes, Russet 0.045 Medium potato, 140g (0.3 lb) 0.078
Lamb 0.032 Chop, 170g (6 oz) 0.054
Carrots 0.003 1 large carrot, 72g (2.5 oz) 0.002

Data taken from United States Department of Agriculture, National Nutrient Database


Main article: Fluoride toxicity


According to the U.S. Department of Agriculture, the Dietary Reference Intakes, which is the “highest level of daily nutrient intake that is likely to pose no risk of adverse health effects” specify 10 mg/day for most people, corresponding to 10 L of fluoridated water with no risk. For infants and young children the values are smaller, ranging from 0.7 mg/d for infants to 2.2 mg/d.[22] Water and food sources of fluoride include community water fluoridation, seafood, tea, and gelatin.[23]

Soluble fluoride salts, of which sodium fluoride is the most common, are only mildly toxic, although they have resulted in both accidental and suicidal deaths from acute poisoning.[3] The lethal dose for most adult humans is estimated at 5 to 10 g (which is equivalent to 32 to 64 mg/kg elemental fluoride/kg body weight).[24][25][26] However, a case of a fatal poisoning of an adult with 4 grams of sodium fluoride is documented,[27] while a dose of 120 g sodium fluoride has been survived.[28] For Sodium fluorosilicate (Na2SiF6), the median lethal dose (LD50) orally in rats is 0.125 g/kg, corresponding to 12.5 g for a 100 kg adult.[29]

The fatal period ranges from 5 min to 12 hours.[27] The mechanism of toxicity involves the combination of the fluoride anion with the calcium ions in the blood to form insoluble calcium fluoride, resulting in hypocalcemia; calcium is indispensable for the function of the nervous system, and the condition can be fatal.

Treatment may involve oral administration of dilute calcium hydroxide or calcium chloride to prevent further absorption, and injection of calcium gluconate to increase the calcium levels in the blood.[27] Hydrogen fluoride is more dangerous than salts such as NaF because it is corrosive and volatile, and can result in fatal exposure through inhalation or upon contact with the skin; calcium gluconate gel is the usual antidote.[30]

In the higher doses used to treat osteoporosis, 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.[31] 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.[32] Fluoride was known to enhance the measurement of bone mineral density at the lumbar spine, but it was not effective for vertebral fractures and provoked more non vertebral fractures.[33]

In areas that have naturally occurring high levels of fluoride in groundwater both dental and skeletal fluorosis can be prevalent and severe.[34]

A popular urban myth claims that the Nazis used fluoride in concentration camps, but there’s no historical evidence for it.[2]


Concentrated fluoride solutions are corrosive. Gloves made of nitrile rubber are worn when handling fluoride compounds. The hazards of solutions of fluoride salts depend on the concentration. In the presence of strong acids, fluoride salts release hydrogen fluoride, which is highly corrosive.

Other derivatives

Organic and inorganic anions are produced from fluoride, including:

See also


  1. “Fluorides – PubChem Public Chemical Database”. The PubChem Project. USA: National Center for Biotechnology Information. Identification.
  2. “Fluorine anion”. NIST. Retrieved 7/4/2012.
  3. Aigueperse, Jean; Mollard, Paul; Devilliers, Didier; Chemla, Marius; Faron, Robert; Romano, Renée; Cuer, Jean Pierre (2005), “Fluorine Compounds, Inorganic”, Ullmann’s Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH, p. 307, doi:10.1002/14356007.a11_307
  4. “Ambient Water Quality Criteria for Fluoride”. Government of British Columbia. Retrieved 8 October 2014.
  5. Fluoride in Drinking-water: Background document for development of WHO Guidelines for Drinking-water Quality. World Health Organization, 2004, p. 2.
  6. Environmental Health Criteria 227: Fluorides. World Health Organization, 2002, p. 38.
  7. Holleman, A. F.; Wiberg, E. (2001) Inorganic Chemistry, Academic Press: San Diego, ISBN 0-12-352651-5.
  8. Schwesinger, Reinhard; Link, Reinhard; Wenzl, Peter; Kossek, Sebastian (2006). “Anhydrous phosphazenium fluorides as sources for extremely reactive fluoride ions in solution”. Chemistry – A European Journal 12 (2): 438. doi:10.1002/chem.200500838.
  9. Chan, Laura; Mehra, Aradhana; Saikat, Sohel; Lynch, Paul (May 2013). “Human exposure assessment of fluoride from tea (Camellia sinensis L.): A UK based issue?”. Food Research International 51 (2): 564–570. doi:10.1016/j.foodres.2013.01.025. Retrieved 9 August 2013.
  10. McDonagh M. S., Whiting P. F., Wilson P. M., Sutton A. J., Chestnutt I., Cooper J., Misso K., Bradley M., Treasure E., & Kleijnen J. (2000). “Systematic review of water fluoridation”. British Medical Journal 321 (7265): 855–859. doi:10.1136/bmj.321.7265.855. PMC 27492. PMID 11021861.
  11. Griffin SO, Regnier E, Griffin PM, Huntley V (2007). “Effectiveness of fluoride in preventing caries in adults”. J. Dent. Res. 86 (5): 410–5. doi:10.1177/154405910708600504. PMID 17452559.
  12. Winston A. E., Bhaskar S. N. (1 November 1998). “Caries prevention in the 21st century”. J. Am. Dent. Assoc. 129 (11): 1579–87. doi:10.14219/jada.archive.1998.0104. PMID 9818575.
  13. Community Water Fluoridation – Oral Health. Cdc.gov (2013-07-10). Retrieved on 2013-07-22.
  14. Ten Great Public Health Achievements in the 20th Century. US Centers for Disease Control and Prevention.
  15. Newbrun E (1996). “The fluoridation war: a scientific dispute or a religious argument?”. J. Public Health Dent. 56 (5 Spec No): 246–52. doi:10.1111/j.1752-7325.1996.tb02447.x. PMID 9034969.
  16. Nakai C, Thomas JA (1974). “Properties of a phosphoprotein phosphatase from bovine heart with activity on glycogen synthase, phosphorylase, and histone”. J. Biol. Chem. 249 (20): 6459–67. PMID 4370977.
  17. Schenk G, Elliott TW, Leung E, et al. (2008). “Crystal structures of a purple acid phosphatase, representing different steps of this enzyme’s catalytic cycle”. BMC Struct. Biol. 8: 6. doi:10.1186/1472-6807-8-6. PMC 2267794. PMID 18234116.
  18. Wang W, Cho HS, Kim R, et al. (2002). “Structural characterization of the reaction pathway in phosphoserine phosphatase: crystallographic “snapshots” of intermediate states”. J. Mol. Biol. 319 (2): 421–31. doi:10.1016/S0022-2836(02)00324-8. PMID 12051918.
  19. Cho H, Wang W, Kim R, et al. (2001). “BeF(3)(-) acts as a phosphate analog in proteins phosphorylated on aspartate: structure of a BeF(3)(-) complex with phosphoserine phosphatase”. Proc. Natl. Acad. Sci. U.S.A. 98 (15): 8525–30. Bibcode:2001PNAS…98.8525C. doi:10.1073/pnas.131213698. PMC 37469. PMID 11438683.
  20. http://www.who.int/water_sanitation_health/dwq/chemicals/fluoride.pdf
  21. “Nutrient Lists”. Agricultural Research Service United States Department of Agriculture. Retrieved 25 May 2014.
  22. [1]
  23. Fluoride in diet, U.S. National Library of Medicine
  24. Gosselin, RE; Smith RP; Hodge HC (1984). Clinical toxicology of commercial products. Baltimore (MD): Williams & Wilkins. pp. III–185–93. ISBN 0-683-03632-7.
  25. Baselt, RC (2008). Disposition of toxic drugs and chemicals in man. Foster City (CA): Biomedical Publications. pp. 636–40. ISBN 978-0-9626523-7-0.
  26. IPCS (2002). Environmental health criteria 227 (Fluoride). Geneva: International Programme on Chemical Safety, World Health Organization. p. 100. ISBN 92-4-157227-2.
  27. Rabinowitch, IM (1945). “Acute Fluoride Poisoning”. Canadian Medical Association journal 52 (4): 345–9. PMC 1581810. PMID 20323400.
  28. Abukurah AR, Moser AM Jr, Baird CL, Randall RE Jr, Setter JG, Blanke RV (1972). “Acute sodium fluoride poisoning”. JAMA 222 (7): 816–7. doi:10.1001/jama.1972.03210070046014. PMID 4677934.
  29. The Merck Index, 12th edition, Merck & Co., Inc., 1996
  30. Muriale L, Lee E, Genovese J, Trend S (1996). “Fatality due to acute fluoride poisoning following dermal contact with hydrofluoric acid in a palynology laboratory”. Ann Occup Hyg. 40 (6): 705–710. doi:10.1016/S0003-4878(96)00010-5. PMID 8958774.
  31. Murray TM, Ste-Marie LG. Prevention and management of osteoporosis: consensus statements from the Scientific Advisory Board of the Osteoporosis Society of Canada. 7. Fluoride therapy for osteoporosis. CMAJ. 1996;155(7):949–54. PMID 8837545.
  32. National Health and Medical Research Council (Australia). A systematic review of the efficacy and safety of fluoridation [PDF]. 2007. ISBN 1-86496-415-4. Summary: Yeung CA. A systematic review of the efficacy and safety of fluoridation. Evid Based Dent. 2008;9(2):39–43. doi:10.1038/sj.ebd.6400578. PMID 18584000. Lay summary: NHMRC, 2007.
  33. Haguenauer, D; Welch, V; Shea, B; Tugwell, P; Adachi, JD; Wells, G (2000). “Fluoride for the treatment of postmenopausal osteoporotic fractures: a meta-analysis.”. Osteoporosis international : a journal established as result of cooperation between the European Foundation for Osteoporosis and the National Osteoporosis Foundation of the USA 11 (9): 727–38. doi:10.1007/s001980070051. PMID 11148800.
  34. World Health Organization (2004). Fluoride in drinking-water (PDF).

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