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Fluoride

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Fluoride
Names
IUPAC name
Fluoride[1]
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
14905
KEGG
MeSH Fluoride
  • InChI=1S/FH/h1H/p-1 checkY
    Key: KRHYYFGTRYWZRS-UHFFFAOYSA-M checkY
  • [F-]
Properties
F−
Molar mass 18.9984032 g mol−1
Thermochemistry
145.58 J/mol K (gaseous)[2]
−333 kJ mol−1
Related compounds
Other anions
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Fluoride is the anion F−, the reduced form of fluorine when as an ion and when bonded to another element. Inorganic fluorine containing compounds are called fluorides. Fluoride, like other halides, is a monovalent ion (−1 charge). Its compounds often have properties that are distinct relative to other halides. Structurally, and to some extent chemically, the fluoride ion resembles the hydroxide ion.

Occurrence

The mineral fluorite, a common mineral and chief source of fluoride for commercial applications.[3][4]

Solutions of inorganic fluorides in water contain F− and bifluoride HF−
2.[5] 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 due to its smaller radius/charge ratio. Its closest chemical relative is hydroxide. When relatively unsolvated, fluoride anions are called "naked". Naked fluoride is a very strong Lewis base.[6] The presence of fluoride and its compounds can be detected by 19F NMR spectroscopy.

Natural occurrence

Many fluoride minerals are known, but of paramount commercial importance are fluorite and fluorapatite.[3]

Fluoride is usually found naturally in low concentration in drinking water and foods. The concentration in seawater averages 1.3 parts per million (ppm). Fresh water supplies generally contain between 0.01–0.3 ppm, whereas the ocean contains between 1.2 and 1.5 ppm.[7] In some locations, the fresh water contains dangerously high levels of fluoride, leading to serious health problems.

Applications

Fluorides are pervasive in modern technology. Hydrofluoric acid is the fluoride synthesized on the largest scale. It is produced by treating fluoride minerals with sulfuric acid. Hydrofluoric acid and its anhydrous form hydrogen fluoride are used in the production of fluorocarbons and aluminium fluorides. Hydrofluoric acid has a variety of specialized applications, including its ability to dissolve glass.[3]

Inorganic chemicals

Fluoride salts are used in the manufacture of many inorganic chemicals, many of which contain fluoride covalently bonded to the metal or nonmetal in question. Some examples of these are:

  • Cryolite (Na3AlF6) is a pesticide that can leave fluoride on agricultural commodities.[8][9] Cryolite was originally utilized in the preparation of aluminium.
  • Sulfuryl fluoride (SO2F2) is used as a pesticide and fumigant on agricultural crops. In 2010, the United States Environmental Protection Agency proposed to withdraw the use of sulfuryl fluoride on food. Sulfuryl fluoride releases fluoride when metabolized.[10][11]
  • Sulfur hexafluoride is an inert, nontoxic insulator gas that is used in electrical transformers and as a tracer gas in indoor air quality investigations.
  • Uranium hexafluoride, although not ionic, is prepared from fluoride reagents. It is utilized in the separation of isotopes of uranium between the fissile isotope U-235 and the non-fissile isotope U-238 in preparation of nuclear reactor fuel and atomic bombs. This is due to the volatility of fluorides of uranium.

    • Organic chemicals

    Main article: Organofluorine chemistry

    Fluoride reagents are significant in synthetic organic chemistry. Organofluorine chemistry has produced many useful compounds over the last 50 years. Included in this area are polytetrafluorethylene (Teflon), polychlorotrifluoroethylene (moisture barriers), efavirenz (pharmaceutical used for treatment of HIV), fluoxetine (an antidepressant), 5-fluorouracil (an anticancer drug), hydrochlorofluorocarbons and hydrofluorcarbons (refrigerants, blowing agents and propellants).

    Due to the affinity of silicon for fluoride, and the ability of silicon to expand its coordination number, silyl ether protecting groups can be easily removed by the fluoride sources such as sodium fluoride and tetra-n-butylammonium fluoride (TBAF). This is quite useful for organic synthesis and the production of fine chemicals. The Si-F linkage is one of the strongest single bonds. In contrast, other silyl halides are easily hydrolyzed.

    Cavity prevention

    Main articles: fluoride therapy and water fluoridation

    Fluoride-containing compounds 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.[12] 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[13][14] 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".[15][16] In some countries where large, centralized water systems are uncommon, fluoride is delivered to the populace by fluoridating table salt. Fluoridation of water has its critics (see Water fluoridation controversy).[17]

    Structure of halothane.

    Biomedical applications

    Positron emission tomography is commonly carried out using fluoride-containing pharmaceuticals such as fluorodeoxyglucose, which is labelled with the radioactive isotope fluorine-18, which emits positrons when it decays into 18O.

    Numerous drugs contain fluorine including antipsychotics such as fluphenazine, HIV protease inhibitors such as tipranavir, antibiotics such as ofloxacin and trovafloxacin, and anesthetics such as halothane.[18] Fluorine is incorporated in the drug structures to reduce drug metabolism, as the strong C-F bond resists deactivation in the liver by cytochrome P450 oxidases.[19]

    Fluoride salts are commonly used to inhibit the activity of phosphatases, such as serine/threonine phosphatases.[20] Fluoride mimics the nucleophilic hydroxyl ion in these enzymes' active sites.[21] 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.[22][23]

    Toxicology

    Main article: Fluoride toxicity
    Reaction of the irreversible inhibitor diisopropylfluorophosphate with a serine protease

    Fluoride-containing compounds are so diverse that it is not possible to generalize on their toxicity, which depends on their reactivity and structure, and in the case of salts, their solubility and ability to release fluoride ions.

    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]

    See also

    References

    1. ^ "Fluorides - PubChem Public Chemical Database". The PubChem Project. USA: National Center for Biotechnology Information. Identification.
    2. ^ "Fluorine anion". NIST. Retrieved 7/4/2012. {{cite web}}: Check date values in: |accessdate= (help)
    3. ^ a b c d Aigueperse, Jean (2005). "Fluorine Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. p. 307. doi:10.1002/14356007.a11_307. ISBN 978-3527306732. {{cite encyclopedia}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
    4. ^ Villalba, Gara; Ayres, Robert U.; Schroder, Hans (2008). "Accounting for Fluorine: Production, Use, and Loss". Journal of Industrial Ecology. 11: 85–101. doi:10.1162/jiec.2007.1075.
    5. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
    6. ^ Schwesinger, Reinhard; Link, Reinhard; Wenzl, Peter; Kossek, Sebastian "Anhydrous phosphazenium fluorides as sources for extremely reactive fluoride ions in solution" Chemistry—A European Journal (2006), Volume Date2005, 12(2), 438-445. doi:10.1002/chem.200500838
    7. ^ Environmental Health Criteria 227: Fluorides. World Health Organization, 2002, page 38. Page accessed on February 22, 2007.Fluoride in Drinking-water: Background document for development of WHO Guidelines for Drinking-water Quality. World Health Organization, 2004, page 2. Page accessed on February 22, 2007.
    8. ^ US EPA (1996). "EPA R.E.D. FACTS Cryolite" (PDF). Retrieved 7/5/2012. {{cite web}}: Check date values in: |accessdate= (help); Unknown parameter |month= ignored (help)
    9. ^ "Cryolite (Kryocide) Pesticide Tolerance and Feed Additive Regulation". 1996. Retrieved 7/5/2012. {{cite web}}: Check date values in: |accessdate= (help); Unknown parameter |month= ignored (help)
    10. ^ US EPA (5/9/2012). "Questions and Answers about EPA's Sulfuryl Fluoride Actions". Retrieved 7/5/2012. {{cite web}}: Check date values in: |accessdate= and |date= (help)
    11. ^ US EPA (5/9/2012). "Sulfuryl Fluoride". Retrieved 7/5/2012. {{cite web}}: Check date values in: |accessdate= and |date= (help)
    12. ^ 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.{{cite journal}}: CS1 maint: multiple names: authors list (link)
    13. ^ 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.{{cite journal}}: CS1 maint: multiple names: authors list (link)
    14. ^ Winston A. E., Bhaskar S. N. (1 November 1998). "Caries prevention in the 21st century". J. Am. Dent. Assoc. 129 (11): 1579–87. PMID 9818575.
    15. ^ Community Water Fluoridation – Oral Health
    16. ^ Ten Great Public Health Achievements in the 20th Century – CDC
    17. ^ 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.
    18. ^ Park BK, Kitteringham NR, O'Neill PM (2001). "Metabolism of fluorine-containing drugs". Annu. Rev. Pharmacol. Toxicol. 41: 443–70. doi:10.1146/annurev.pharmtox.41.1.443. PMID 11264465.{{cite journal}}: CS1 maint: multiple names: authors list (link)
    19. ^ Fisher MB, Henne KR, Boer J (2006). "The complexities inherent in attempts to decrease drug clearance by blocking sites of CYP-mediated metabolism". Curr. Opin. Drug Discov. Devel. 9 (1): 101–9. PMID 16445122.{{cite journal}}: CS1 maint: multiple names: authors list (link)
    20. ^ 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.
    21. ^ 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. {{cite journal}}: Explicit use of et al. in: |author= (help)CS1 maint: multiple names: authors list (link) CS1 maint: unflagged free DOI (link)
    22. ^ 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. {{cite journal}}: Explicit use of et al. in: |author= (help)CS1 maint: multiple names: authors list (link)
    23. ^ 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. {{cite journal}}: Explicit use of et al. in: |author= (help)CS1 maint: multiple names: authors list (link)
    24. ^ Gosselin, RE (1984). Clinical toxicology of commercial products. Baltimore (MD): Williams & Wilkins. pp. III–185–93. ISBN 0-683-03632-7. {{cite book}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
    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. {{cite book}}: Cite has empty unknown parameter: |coauthors= (help)
    26. ^ IPCS (2002). Environmental health criteria 227 (Fluoride). Geneva: International Programme on Chemical Safety, World Health Organization. p. 100. ISBN 92-4-157227-2. {{cite book}}: Cite has empty unknown parameter: |coauthors= (help)
    27. ^ a b c I. M. Rabinowitch. Acute Fluoride Poisoning. Can Med Assoc J. 1945, 52, 345–349. [1]
    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.{{cite journal}}: CS1 maint: multiple names: authors list (link)
    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. PMID 8958774.{{cite journal}}: CS1 maint: multiple names: authors list (link)
    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.
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