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Fluoride

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Fluoride is the element fluorine when it is bonded with another element.[1] [2] Both organic and inorganic compounds containing the element fluorine are sometimes 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. Fluorine-containing compounds range from potent toxins such as sarin to life-saving pharmaceuticals such as efavirenz, and from inert materials such as calcium fluoride to the highly reactive sulfur tetrafluoride. The range of fluorine-containing compounds is considerable as fluorine is capable of forming compounds with all the elements except helium and neon.[3][4]

Compounds containing fluoride anions and in many cases those containing covalent bonds to fluorine are called fluorides.

Occurrence

The mineral fluorite, a common mineral and chief source of fluoride for commercial applications (the fluorite is the purple-colored part of this specimen).

Solutions of inorganic fluorides in water contain F− and bifluoride HF2−.[5] Few inorganic fluorides are soluble in water without undergoing significant hydrolysis. Examples of inorganic fluorides include hydrofluoric acid (HF), sodium fluoride (NaF), and uranium hexafluoride (UF6). 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. The Si-F linkage is one of the strongest single bonds. In contrast, other silyl halides are easily hydrolyzed.

Natural occurrence

Many fluoride minerals are known, but of paramount commercial importance are fluorite and fluorapatite.[6] Fluoride is found naturally in low concentration in drinking water and foods. Water with underground sources is more likely to have higher levels of fluoride, whereas the concentration in seawater averages 1.3 parts per million (ppm).[7] Fresh water supplies generally contain between 0.01–0.3 ppm, whereas the ocean contains between 1.2 and 1.5 ppm.[8]

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.[6]

Organic synthesis

Fluoride reagents are significant in synthetic organic chemistry. 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).

Inorganic fluorides

Sulfur hexafluoride is an inert, nontoxic insulator that is used in electrical transformers. Uranium hexafluoride is used 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. The volatility of fluorides of uranium and other elements may also be used for nuclear fuel reprocessing.

PTFE is often used to coat non-stick frying pans as it is not water-wettable and possesses high heat resistance.

Fluoropolymers

Fluoropolymers such as polytetrafluoroethylene, Teflon, are used as chemically inert and biocompatible materials for a variety of applications, including as surgical implants such as coronary bypass grafts,[9] and a replacement for soft tissue in cosmetic and reconstructive surgery.[10] These compounds are also commonly used as non-stick surfaces in cookware and bakeware, and the fluoropolymer fabric Gore-Tex used in breathable garments for outdoor use.

Cavity prevention

Main article: 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.[11] Originally, sodium fluoride was used to fluoridate water; however, 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[12][13] 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".[14][15] In some countries where large, centralized water systems are uncommon, fluoride is delivered to the populace by fluoridating table salt. Fluoridation of water is not without critics, however (see Water fluoridation controversy).[16]

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.[17] 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.[18]

Fluoride salts are commonly used to inhibit the activity of phosphatases, such as serine/threonine phosphatases.[19] Fluoride mimics the nucleophilic hydroxyl ion in these enzymes' active sites.[20] 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.[21][22]

Toxicology

Main article: Fluoride poisoning
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 NaF is the most common, are mildly toxic but have resulted in both accidental and suicidal deaths from acute poisoning.[6] While the minimum fatal dose in humans is not known, a case of a fatal poisoning of an adult with 4 grams of NaF is documented.[23] 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.[24] The fatal period ranges from 5 min to 12 hours.[23] 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.[23] 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.[25]

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.[26] 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.[27]

See also

References

  1. ^ http://education.jlab.org/itselemental/ele009.html
  2. ^ http://www.fluoroseal.com/fluorine.html
  3. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. p. 804
  4. ^ Khriachtchev, Leonid (24 August 2000). "A stable argon compound". Nature. 406 (6798): 874–876. doi:10.1038/35022551. PMID 10972285. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
  5. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  6. ^ a b c Aigueperse, Jean (2005). "Fluorine Compounds, Inorganic". Ullmann’s Encyclopedia of Industrial Chemistry. Wiley-VCH, Weinheim. p. 307. doi:10.1002/14356007.a11. {{cite book}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
  7. ^ 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. ^ Environmental Health Criteria 227: Fluorides. World Health Organization, 2002, page 38. Page accessed on February 22, 2007.
  9. ^ Kannan RY, Salacinski HJ, Butler PE, Hamilton G, Seifalian AM (2005). "Current status of prosthetic bypass grafts: a review". J. Biomed. Mater. Res. Part B Appl. Biomater. 74 (1): 570–81. doi:10.1002/jbm.b.30247. PMID 15889440.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  10. ^ Singh S., Baker J. L. (2000). "Use of expanded polytetrafluoroethylene in aesthetic surgery of the face". Clin Plast Surg. 27 (4): 579–93. PMID 11039891.
  11. ^ 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)
  12. ^ 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)
  13. ^ 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.
  14. ^ Community Water Fluoridation - Oral Health
  15. ^ http://www.cdc.gov/about/history/tengpha.htm
  16. ^ 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.
  17. ^ 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)
  18. ^ 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)
  19. ^ 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.
  20. ^ 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 (1): 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)
  21. ^ 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)
  22. ^ 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. 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)
  23. ^ a b c I. M. Rabinowitch. Acute Fluoride Poisoning. Can Med Assoc J. 1945, 52, 345–349. [1]
  24. ^ The Merck Index, 12th edition, Merck & Co., Inc., 1996
  25. ^ Muriale L, Lee E, Genovese J, Trend S. Fatality due to acute fluoride poisoning following dermal contact with hydrofluoric acid in a palynology laboratory. Ann Occup Hyg. 1996 40, 705–710. PMID 8958774.
  26. ^ 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.
  27. ^ National Health and Medical Research Council (Australia). A systematic review of the efficacy and safety of fluoridation [PDF]. 2007. ISBN 1864964154. 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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