Fluorine: Wikis


Note: Many of our articles have direct quotes from sources you can cite, within the Wikipedia article! This article doesn't yet, but we're working on it! See more info or our list of citable articles.

Did you know ...

More interesting facts on Fluorine

Include this on your site/blog:


From Wikipedia, the free encyclopedia



Yellowish gas
General properties
Name, symbol, number fluorine, F, 9
Element category halogen
Group, period, block 172, p
Standard atomic weight 18.9984032(5)g·mol−1
Electron configuration 1s2 2s2 2p5
Electrons per shell 2, 7 (Image)
Physical properties
Phase gas
Density (0 °C, 101.325 kPa)
1.7 g/L
Melting point 53.53 K, −219.62 °C, −363.32 °F
Boiling point 85.03 K, −188.12 °C, −306.62 °F
Critical point 144.13 K, 5.172 MPa
Heat of fusion (F2) 0.510 kJ·mol−1
Heat of vaporization (F2) 6.62 kJ·mol−1
Specific heat capacity (25 °C) (F2)
31.304 J·mol−1·K−1
Vapor pressure
P/Pa 1 10 100 1 k 10 k 100 k
at T/K 38 44 50 58 69 85
Atomic properties
Oxidation states −1
(Weaklyacidic oxide)
Electronegativity 3.98 (Pauling scale)
Ionization energies
1st: 1681.0 kJ·mol−1
2nd: 3374.2 kJ·mol−1
3rd: 6050.4 kJ·mol−1
Covalent radius 57±3 pm
(see covalent radius of fluorine)
Van der Waals radius 147 pm
Crystal structure cubic
Magnetic ordering nonmagnetic
Thermal conductivity (300 K) 27.7 m W·m−1·K−1
CAS registry number 7782-41-4
Most stable isotopes
Main article: Isotopes of fluorine
iso NA half-life DM DE (MeV) DP
18F syn 109.77 min β+ (97%) 0.64 18O
ε (3%) 1.656 18O
19F 100% 19F is stable with 10 neutrons

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 is found in the -1 oxidation state, 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]


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 compounds

  • Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMS (microelectromechanical systems) fabrication.[4] Xenon difluoride is also used for this last purpose.
  • Hydrofluoric acid (chemical formula HF) is used to etch glass in light bulbs and other products.
  • Tetrafluoroethylene and perfluorooctanoic acid (PFOA) are directly used in the production of low friction plastics such as Teflon (or polytetrafluoroethylene).
  • Fluorine is used indirectly in the production of halons such as freon.
  • Along with some of its compounds, fluorine is used in the production of pure uranium from uranium hexafluoride and in the synthesis of numerous commercial fluorochemicals, including vitally important pharmaceuticals, agrochemical compounds, lubricants, and textiles.
  • Chlorofluorocarbons (CFCs) are used extensively in air conditioning and in refrigeration. CFCs have been banned for these applications because they contribute to ozone destruction and the ozone hole. Interestingly, since it is chlorine and bromine radicals which harm the ozone layer, not fluorine, compounds which do not contain chlorine or bromine but contain only fluorine, carbon and hydrogen (called hydrofluorocarbons) are not on the United States Environmental Protection Agency list of ozone-depleting substances,[5] and have been widely used as replacements for the chlorine- and bromine-containing fluorocarbons. Hydrofluorocarbons do have a greenhouse effect, but a small one compared with carbon dioxide and methane.
  • Sodium hexafluoroaluminate (cryolite), is used in the electrolysis of aluminium.
  • In much higher concentrations, sodium fluoride has been used as an insecticide, especially against cockroaches.
  • Fluorides have been used in the past to help molten metal flow. Hence the name, which derives from Latin verb fluere, meaning to flow.[6]
  • Some researchers including US space scientists in the early 1960s have studied elemental fluorine gas as a possible rocket propellant due to its exceptionally high specific impulse. The experiments failed because fluorine proved difficult to handle, and its combustion products proved extremely toxic and corrosive.
  • Compounds of fluorine such as fluoropolymers, potassium fluoride and cryolite are utilized in applications such as anti-reflective coatings and dichroic mirrors on account of their unusually low refractive index.
Fluorite (CaF2) crystals

Dental and medical uses

Chemistry of fluorine

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 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 role

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 fluorine

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 ion

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 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 also


  1. ^ Theodore Gray. "Real visible fluorine". The Wooden Periodic Table. http://theodoregray.com/PeriodicTable/Samples/009.5/index.s12.html. 
  2. ^ "pKa's of Inorganic and Oxo-Acids". Evans Group. http://www2.lsdiv.harvard.edu/labs/evans/pdf/evans_pKa_table.pdf. Retrieved 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. http://www.epa.gov/ozone/ods.html. 
  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. http://books.google.de/books?id=Fl4sdCYrq3cC&pg=PT158. 
  7. ^ Steve S Lim. "eMedicine - Corticosteroid-Induced Myopathy". http://www.emedicine.com/pmr/topic35.htm. 
  8. ^ "Fluorine's treasure trove". ICIS news. 2006-10-02. http://www.icis.com/Articles/2006/09/30/2016413/fluorines-treasure-trove.html. Retrieved 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. http://www.fluoride-history.de/fluorine.htm. 
  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. http://gallica.bnf.fr/ark:/12148/bpt6k3058f/f1541.chemindefer. 
  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. http://books.google.com/books?id=kgVAlzGXx6oC. 
  14. ^ "The Nobel Prize in Chemistry 1906". Nobelprize.org. http://nobelprize.org/nobel_prizes/chemistry/laureates/1906/. Retrieved 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. http://www.who.int/water_sanitation_health/dwq/en/nutoverview.pdf. Retrieved 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. http://www51.honeywell.com/sm/hfacid/common/documents/HF_medical_book.pdf. Retrieved 2009-05-06. 

External links

1911 encyclopedia

Up to date as of January 14, 2010

From LoveToKnow 1911

FLUORINE (symbol F, atomic weight iv), a chemical element of the halogen group. It is never found in the uncombined condition, but in combination with calcium as fluor-spar CaF2 it is widely distributed; it is also found in cryolite Na3A1F6, in fluor-apatite, CaF 2.3Ca 3 P 2 O 8, and in minute traces in seawater, in some mineral springs, and as a constituent of the enamel of the teeth. It was first isolated by H. Moissan in 1886 by the electrolysis of pure anhydrous hydrofluoric acid containing dissolved potassium fluoride. The U-shaped electrolytic vessel and the electrodes are made of an alloy of platinum-iridium, the limbs of the tube being closed by stoppers made of fluor-spar, and fitted with two lateral exit tubes for carrying off the gases evolved. Whilst the electrolysis is proceeding, the apparatus is kept at a constant temperature of - 23° C. by means of liquid methyl chloride. The fluorine, which is liberated as a gas at the anode, is passed through a well cooled platinum vessel, in order to free it from any acid fumes that may be carried over, and finally through two platinum tubes containing sodium fluoride to remove the last traces of hydrofluoric acid; it is then collected in a platinum tube closed with fluor-spar plates. B. Brauner (Jour. Chem. Soc., 18 94, 6 5, p. 393) obtained fluorine by heating potassium fluorplumbate 3KF HF PbF 4. At 200° C. this salt decomposes, giving off hydrofluoric acid, and between 230-250° C. fluorine is liberated.

Fluorine is a pale greenish-yellow gas with a very sharp smell; its specific gravity is 1.265 (H. Moissan); it has been liquefied, the liquid also being of a yellow colour and boiling at - 187° C. It is the most active of all the chemical elements; in contact with hydrogen combination takes place between the two gases with explosive violence, even in the dark, and at as low a temperature as - 210 C.; finely divided carbon burns in the gas, forming carbon tetrafluoride; water is decomposed even at ordinary temperatures, with the formation of hydrofluoric acid and "ozonised" oxygen; iodine, sulphur and phosphorus melt and then inflame in the gas; it liberates chlorine from chlorides, and combines with most metals instantaneously to form fluorides; it does not, however, combine with oxygen. Organic compounds are rapidly attacked by the gas.

Only one compound of hydrogen and fluorine is known, namely hydrofluoric acid, HF or H 2 F 2, which was first obtained by C. Scheele in 1771 by decomposing fluor-spar with concentrated sulphuric acid, a method still used for the commercial preparation of the aqueous solution of the acid, the mixture being distilled from leaden retorts and the acid stored in leaden or gutta-percha bottles. The perfectly anhydrous acid is a very volatile colourless liquid and is best obtained, according to G. Gore (Phil. Trans., 1869, p. 173) by decomposing the double fluoride of hydrogen and potassium, at a red heat in a platinum retort fitted with a platinum condenser surrounded by a freezing mixture, was having a platinum receiver luted on. It can also be prepared in the anhydrous condition by passing a current of hydrogen over dry silver fluoride. The pure acid thus obtained is a most dangerous substance to handle, its vapour even when highly diluted with air having an exceedingly injurious action on the respiratory organs, whilst inhalation of the pure vapour is followed by death. The anhydrous acid boils at 19 0.5 C. (H Moissan), and on cooling, sets to a solid mass at - 102°. 5 C., which melts at - 9 2° 3 C. (K. Olszewski, Monats. fiir Chemie, 1886, 7, p. 37 1). Potassium and sodium readily dissolve in the anhydrous acid with evolution of hydrogen and formation of x. 19 fluorides. The aqueous solution is strongly acid to litmus and dissolves most metals directly. Its most important property is that it rapidly attacks glass, reacting with the silica of the glass to form gaseous silicon fluoride, and consequently it is used for etching. T. E. Thorpe (Jour. Chem. Soc., 1889, 55, p. 163) determined the vapour density of hydrofluoric acid at different temperatures, and showed that there is no approach to a definite value below about 88° C. where it reaches the value 10.29 corresponding to the molecular formula HF; at temperatures below 88° C. the value increases rapidly, showing that the molecule is more complex in its structure. (For references see J. N. Friend, The Theory of Valency (1909), p. iii.) The aqueous solution behaves on concentration similarly to the other halogen acids; E. Deussen (Zeit. anorg. Chem., 1905, 44, pp. 300, 408; 1906, 49, p. 2 97) found the solution of constant boiling point to contain 43.2% HF and to boil at (750 mm.).

The salts of hydrofluoric acid are known as fluorides and are easily obtained by the action of the acid on metals or their oxides, hydroxides or carbonates. The fluorides of the alkali metals, of silver, and of most of the heavy metals are soluble in water; those of the alkaline earths are insoluble. A characteristic property of the alkaline fluorides is their power of combining with a molecule of hydrofluoric acid and with the fluorides of the more electro-negative elements to form double fluorides, a behaviour not shown by other metallic halides. Fluorides can be readily detected by their power of etching glass when warmed with sulphuric acid; or by warming them in a glass tube with concentrated sulphuric acid and holding a moistened glass rod in the mouth of the tube, the water apparently gelatinizes owing to the decomposition of the silicon fluoride formed. The atomic weight of fluorine has been determined by the conversion of calcium, sodium and potassium fluorides into the corresponding sulphates. J. Berzelius, by converting silver fluoride into silver chloride, obtained the value 19.44, and by analysing calcium fluoride the value 19.16; the more recent work of H. Moissan gives the value 19.05.

See H. Moissan, Le Fluor et ses composes (Paris, 1900).

<< Fluorescence

Fluor-spar >>

Simple English

[[File:|thumb|A colorful picture of fluorine. Fluorine is not really so bright green.]]

A more real picture of flourine

Fluorine (symbol F) is a chemical element. Its atomic number (which is the number of protons in it) is 9, and its atomic mass is 19. It is part of the Group 7 (halogens) on the periodic table of elements.



Fluorine is a light yellow diatomic gas. It is very reactive gas, which exists as diatomic molecules. It is actually the most reactive element. Fluorine has a very high attraction for electrons, because it is missing one. This makes it the most powerful oxidizing agent. It can rip electrons from water (making oxygen) and ignite propane on contact. It does not need a spark. Metals can catch on fire when placed in a stream of fluorine. After it is reduced by reacting with other things, it forms the stable fluoride ion. Fluorine is very poisonous. Fluorine bonds very strongly with carbon. It can react with the unreactive noble gases. It explodes when mixed with hydrogen.


Fluorine is not found as an element on the earth; it is much too reactive. Several fluorides are found in the earth, though. When calcium phosphate is reacted with sulfuric acid to make phosphoric acid, some hydrofluoric acid is produced. Also, fluorite can be reacted with sulfuric acid to make hydrofluoric acid.


Fluorine is normally made by electrolysis. Hydrogen fluoride is dissolved in potassium fluoride. This mixture is melted and an electric current is passed through it. This is electrolysis. Hydrogen is produced at one side and fluorine at the other side. If the sides are not separated, the cell may explode.

Someone made fluorine in 1986 without using electrolysis. They produced manganese(IV) fluoride by using various chemical compounds, which released fluorine gas.


Fluorine is used to enrich uranium for nuclear weapons. It is also used to make sulfur hexafluoride. Sulfur hexafluoride is used to propel stuff out of an aerosol can. It is also used to make integrated circuits. Fluorine compounds have many uses. Fluoride ions are in fluorine compounds. Fluoride ions can be in toothpaste. Some are used in nonstick coatings. Freons contain fluorine.


Fluorine as an element is extremely reactive and toxic. It can react with almost everything, even glass. Fluorine is also poisonous.

Fluoride ions are somewhat toxic. If too much toothpaste containing fluoride is eaten then fluoride poisoning may occur. Fluoride is not reactive, though.

Other pages


Got something to say? Make a comment.
Your name
Your email address