Template:Infobox fluorine

Fluorine is the chemical element with atomic number 9, represented by the symbol F. Fluorine forms a single bond with itself in elemental form, resulting in the diatomic F2 molecule. F2 is a supremely reactive, poisonous, pale, yellowish brown gas. Elemental fluorine is the most chemically reactive and electronegative of all the elements. For example, it will readily "burn" hydrocarbons at room temperature, in contrast to the combustion of hydrocarbons by oxygen, which requires an input of energy with a spark. Therefore, molecular fluorine is highly dangerous, more so than other halogens such as the poisonous chlorine gas.

Fluorine's highest electronegativity and small atomic radius give unique properties to many of its compounds. For example, the enrichment of 235U, the principal nuclear fuel, relies on the volatility of UF6. Also, the carbon–fluorine bond is one of the strongest bonds in organic chemistry. This contributes to the stability and persistence of fluoroalkane based organofluorine compounds, such as PTFE/(Teflon) and PFOS. The carbon–fluorine bond's inductive effects result in the strength of many fluorinated acids, such as triflic acid and trifluoroacetic acid. Drugs are often fluorinated at biologically reactive positions, to prevent their metabolism and prolong their half-lives.


F2 is a corrosive pale yellow or brown[1] gas that is a powerful oxidizing agent. It is the most reactive and most electronegative of all the elements on the classic Pauling scale (4.0), and readily forms compounds with most other elements. It has an oxidation number -1, except when bonded to another fluorine in F2 which gives it an oxidation number of 0. Fluorine combines with the noble gases argon, krypton, xenon, and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. The reaction with hydrogen can occur at extremely low temperatures, using liquid hydrogen and solid fluorine. It is so reactive that metals, water, as well as most other substances, burn with a bright flame in a jet of fluorine gas. In moist air, it reacts with water to form the also dangerous hydrofluoric acid.

Fluorides are compounds that combine fluorine with some positively charged counterpart. They often consist of crystalline ionic salts. Fluorine compounds with metals are among the most stable of salts.

Hydrogen fluoride is a weak acid when dissolved in water, but is still very corrosive and attacks glass. Consequently, fluorides of alkali metals produce basic solutions. For example, a 1 M solution of NaF in water has a pH of 8.59 compared to a 1 M solution of NaOH, a strong base, which has a pH of 14.00.[2]


Main article: Isotopes of fluorine

Although fluorine (F) has multiple isotopes, only one of these isotopes (F-19) is stable, and the others have short half-lives and are not found in nature. Fluorine is thus a mononuclidic element.

The nuclide 18F is the radionuclide of fluorine with the longest half life (about 110 minutes), and commercially is an important source of positrons, finding its major use in positron emission tomography scanning.


Elemental fluorine, F2, is mainly used for the production of two compounds of commercial interest, uranium hexafluoride and sulfur hexafluoride.[3]

Industrial use of fluorine-containing compoundsEdit

File:Fluorite crystals 270x444.jpg

Dental and medical usesEdit

Chemistry of fluorineEdit

Fluorine forms a variety of very different compounds, owing to its small atomic size and covalent behavior. Elemental fluorine is a dangerously powerful oxidant, reflecting the extreme electronegativity of fluorine. Hydrofluoric acid is extremely dangerous, whereas in synthetic drugs incorporating an aromatic ring (e.g. flumazenil), fluorine is used to help prevent toxication or to delay metabolism[citation needed].

The fluoride ion is basic, therefore hydrofluoric acid is a weak acid in water solution. However, water is not an inert solvent in this case: when less basic solvents such as anhydrous acetic acid are used, hydrofluoric acid is the strongest of the hydrohalogenic acids. Also, owing to the basicity of the fluoride ion, soluble fluorides give basic water solutions. The fluoride ion is a Lewis base, and has a high affinity to certain elements such as calcium and silicon. For example, deprotection of silicon protecting groups is achieved with a fluoride. The fluoride ion is poisonous.

Fluorine as a freely reacting oxidant gives the strongest oxidants known.

The reactivity of fluorine toward the noble gas xenon was first reported by Neil Bartlett in 1962. Fluorides of krypton and radon have also been prepared. Argon fluorohydride has been observed at cryogenic temperatures.

The carbon-fluoride bond is covalent and very stable. The use of a fluorocarbon polymer, poly(tetrafluoroethene) or Teflon, is an example: it is thermostable and waterproof enough to be used in frying pans. Organofluorines may be safely used in applications such as drugs, without the risk of release of toxic fluoride. In synthetic drugs, toxication can be prevented. For example, an aromatic ring is useful but presents a safety problem: enzymes in the body metabolize some of them into poisonous epoxides. When the para position is substituted with fluorine, the aromatic ring is protected and epoxide is no longer produced.

The substitution of fluorine for hydrogen in organic compounds offers a very large number of compounds. An estimated fifth of pharmaceutical compounds and 30% of agrochemical compounds contain fluorine.[8] The -CF3 and -OCF3 moieties provide further variation, and more recently the -SF5 group.[9]


Fluorine cell room

Fluorine cell room at F2 Chemicals Ltd, Preston, UK

Industrial production of fluorine entails the electrolysis of hydrogen fluoride in the presence of potassium fluoride. This method is based on the pioneering studies by Moissan (see below). Fluorine gas forms at the anode, and hydrogen gas at the cathode. Under these conditions, the potassium fluoride (KF) converts to potassium bifluoride (KHF2), which is the actual electrolyte. This potassium bifluoride aids electrolysis by greatly increasing the electrical conductivity of the solution.

HF + KF → KHF2
2 KHF2 → 2 KF + H2 + F2

The HF required for the electrolysis is obtained as a byproduct of the production of phosphoric acid. Phosphate-containing minerals contain significant amounts of calcium fluorides, such as fluorite. Upon treatment with sulfuric acid, these minerals release hydrogen fluoride:

CaF2 + H2SO4 → 2 HF + CaSO4

In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl Christe discovered a purely chemical preparation involving the reaction of solutions in anhydrous HF, K2MnF6, and SbF5 at 150 °C:[10]

2 K2MnF6 + 4 SbF5 → 4 KSbF6 + 2 MnF3 + F2

Though not a practical synthesis on the large scale, this report demonstrates that electrolysis is not the sole route to the element.


The mineral fluorspar (also called fluorite), consisting mainly of calcium fluoride, was described in 1530 by Georgius Agricola for its use as a flux.[11] Fluxes are used to promote the fusion of metals or minerals. The etymology of the element's name reflects its history: Fluorine is pronounced /ˈflʊəriːn/, /ˈflʊərɨn/, or commonly /ˈflɔr-/; from Latin: fluere, meaning "to flow". In 1670 Schwanhard found that glass was etched when it was exposed to fluorspar that had been treated with acid. Carl Wilhelm Scheele and many later researchers, including Humphry Davy, Caroline Menard, Gay-Lussac, Antoine Lavoisier, and Louis Thenard all would experiment with hydrofluoric acid, easily obtained by treating fluorite with concentrated sulfuric acid.

Owing to its extreme reactivity, elemental fluorine was not isolated until many years after the characterization of fluorite. Progress in isolating elemental fluorine was slowed because it could only be prepared electrolytically and even then under stringent conditions since the gas attacks many materials. In 1886, the isolation of elemental fluorine was reported by Henri Moissan after almost 74 years of effort by other chemists.[12] The generation of elemental fluorine from hydrofluoric acid is exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. These individuals came to be referred to as "fluorine martyrs".[13] For Moissan, it earned him the 1906 Nobel Prize in chemistry.[14]

The first large-scale production of fluorine was undertaken in support of the Manhattan project, where the compound uranium hexafluoride (UF6) had been selected as the form of uranium that would allow separation of its 235U and 238U isotopes. Today both the gaseous diffusion process and the gas centrifuge process use gaseous UF6 to produce enriched uranium for nuclear power applications. In the Manhattan Project, it was found that UF6 decomposed into UF4 and F2. The corrosion problem due to the F2 was eventually solved by electrolytically coating all UF6 carrying piping with nickel metal, which forms a nickel difluoride that is not attacked by fluorine. Joints and flexible parts were made from teflon, then a very recently discovered fluorocarbon plastic which is also not attacked by F2.

Biological roleEdit

Though F2 is too reactive to have any natural biological role, fluorine is incorporated into compounds with biological activity. Naturally occurring organofluorine compounds are rare, the most notable example is fluoroacetate, which functions as a plant defence against herbivores in at least 40 plants in Australia, Brazil and Africa.[15] The enzyme adenosyl-fluoride synthase catalyzes the formation of 5'-deoxy-5'-fluoroadenosine. Fluorine is not an essential nutrient, but its importance in preventing tooth decay is well-recognized.[16] The effect is predominantly topical, although prior to 1981 it was considered primarily systemic (occurring through ingestion).[17]


Elemental fluorineEdit

Elemental fluorine (fluorine gas) is a highly toxic, corrosive oxidant, which can cause ignition of organic material. Fluorine gas has a characteristic pungent odor that is detectable in concentrations as low as 20 ppb. As it is so reactive, all materials of construction must be carefully selected and metal surfaces must be passivated.

Fluoride ionEdit

Main article: Fluoride poisoning

Fluoride ions are toxic: the lethal dose of sodium fluoride for a 70 kg human is estimated at 5–10 g.[18]

Hydrogen fluoride and hydrofluoric acidEdit

Main article: Hydrofluoric acid

Hydrogen fluoride and hydrofluoric acid are dangerous, far more so than the related hydrochloric acid, because undissociated molecular HF penetrates the skin and biological membranes, causing deep and painless burns. The free fluoride, once released from HF in dissociation, also is capable of chelating calcium ion to the point of causing death by cardiac dysrhythmia. Burns with areas larger than 25 square inches (160 cm2) have the potential to cause serious systemic toxicity.[19]


Organofluorines are naturally rare compounds. They can be nontoxic (perflubron and perfluorodecalin) or highly toxic (perfluoroisobutylene and fluoroacetic acid). Many pharmacueticals are organofluorines, such as the anti-cancer fluorouracil. Perfluorooctanesulfonic acid (PFOS) is a persistent organic pollutant.

See alsoEdit


  1. Theodore Gray. "Real visible fluorine". The Wooden Periodic Table.
  2. "pKa's of Inorganic and Oxo-Acids". Evans Group. Retrieved on 2008-11-29.
  3. M. Jaccaud, R. Faron, D. Devilliers, R. Romano (2005). Fluorine, in Ullmann’s Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. ISBN 3527310975. 
  4. Leonel R Arana, Nuria de Mas, Raymond Schmidt, Aleksander J Franz, Martin A Schmidt and Klavs F Jensen (2007). "Isotropic etching of silicon in fluorine gas for MEMS micromachining". J. Micromech. Microeng. 17: 384. doi:10.1088/0960-1317/17/2/026. 
  5. "Class I Ozone-Depleting Substances". Ozone Depletion. U.S. Environmental Protection Agency.
  6. compiled by Alexander Senning. (2007). Elsevier's dictionary of chemoetymology : the whies and whences of chemical nomenclature and terminology. Amsterdam: Elsevier. p. 149. ISBN 9780444522399, 
  7. Steve S Lim. "eMedicine - Corticosteroid-Induced Myopathy". 
  8. "Fluorine's treasure trove". ICIS news (2006-10-02). Retrieved on 2008-11-29.
  9. Bernhard Stump, Christian Eberle, W. Bernd Schweizer, Marcel Kaiser, Reto Brun, R. Luise Krauth-Siegel, Dieter Lentz, François Diederich (2009). "Pentafluorosulfanyl as a Novel Building Block for Enzyme Inhibitors: Trypanothione Reductase Inhibition and Antiprotozoal Activities of Diarylamines". ChemBioChem 10 (1): 79. doi:10.1002/cbic.200800565. PMID 19058274. 
  10. K. Christe (1986). "Chemical synthesis of elemental fluorine". Inorg. Chem. 25: 3721–3724. doi:10.1021/ic00241a001. 
  11. "Discovery of fluorine". Fluoride History.
  12. H. Moissan (1886). "Action d'un courant électrique sur l'acide fluorhydrique anhydre". Comptes rendus hebdomadaires des séances de l'Académie des sciences 102: 1543–1544, 
  13. Richard D. Duncan. (2008). Elements of faith : faith facts and learning lessons from the periodic table. Green Forest, Ark.: Master Books. p. 22. ISBN 9780890515471, 
  14. "The Nobel Prize in Chemistry 1906". Retrieved on 2009-07-07.
  15. Proudfoot AT, Bradberry SM, Vale JA (2006). "Sodium fluoroacetate poisoning". Toxicol Rev 25 (4): 213–9. doi:10.2165/00139709-200625040-00002. PMID 17288493. 
  16. Olivares M and Uauy R (2004). "Essential nutrients in drinking-water (Draft)". WHO. Retrieved on 2008-12-30.
  17. Pizzo G, Piscopo MR, Pizzo I, Giuliana G (September 2007). "Community water fluoridation and caries prevention: a critical review". Clin Oral Investig 11 (3): 189–93. doi:10.1007/s00784-007-0111-6. PMID 17333303. 
  18. Aigueperse, Jean; Paul Mollard, Didier Devilliers, Marius Chemla, Robert Faron, Renée Romano, Jean Pierre Cuer (2005). "Fluorine Compounds, Inorganic". in Ullmann. Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. 
  19. "Recommended Medical Treatment for Hydrofluoric Acid Exposure" (PDF). Honeywell Specialty Materials. Retrieved on 2009-05-06.

External linksEdit


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