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Name , symbol , atomic number Fluorine, F, 9
Element category Halogens
Group , period , block 17 , 2 , p
Look pale, yellowish gas
CAS number 7782-41-4
EC number 231-954-8
ECHA InfoCard 100,029,049
Mass fraction of the earth's envelope 0.028%
Atomic mass 18.998403163 (6) and
Atomic radius (calculated) 50 (42) pm
Covalent radius 71 pm
Van der Waals radius 147 pm
Electron configuration [ He ] 2 s 2 2 p 5
1. Ionization energy 17th.42282 (5) eV1 681.05 kJ / mol
2. Ionization energy 34.97081 (12) eV3 374.17 kJ / mol
3. Ionization energy 62.70798 (25) eV6 050.4 kJ / mol
4. Ionization energy 87.175 (17) eV8 411.11 kJ / mol
5. Ionization energy 114.249 (6) eV11 023.3 kJ / mol
6. Ionization energy 157.16311 (25) eV15 163.9 kJ / mol
7. Ionization energy 185.1868 (6) eV17 867.8 kJ / mol
Physical state gaseous (F 2 )
density 1.6965 kg / m 3 at 273 K.
magnetism diamagnetic
Melting point 53.53 K (−219.62 ° C)
boiling point 85.15 K (−188 ° C)
Molar volume (solid) 11.20 · 10 −6 m 3 · mol −1
Heat of vaporization 6.32 kJ / mol
Heat of fusion 0.2552 kJ mol −1
Thermal conductivity 0.0279 W m −1 K −1
Oxidation states −1
Normal potential 2.87 V (F + e - → F - )
Electronegativity 3.98 ( Pauling scale )
isotope NH t 1/2 ZA ZE (M eV ) ZP
17 F. {syn.} 64.49 s β + 2.761 17 O
18 F {syn.} 109.77 min β + , ε 1.656 18 O
19 F. 100% Stable
20 F {syn.} 11.00 s β - 7.025 20 Ne
21 F. {syn.} 4.158 s β - 5.684 21 Ne
For other isotopes, see the list of isotopes
NMR properties
number I
γ in
rad · T −1 · s −1
E r  ( 1 H) f L at
B = 4.7 T
in MHz
19 F. 1/2 25.17 · 10 7 0.832 94.077 (at 2.3488 T)
safety instructions
GHS hazard labeling from  Regulation (EC) No. 1272/2008 (CLP) , expanded if necessary
03 - Oxidising 04 - gas bottle 06 - Toxic or very toxic 05 - Corrosive


H and P phrases H: 270-280-330-314
EUH: 071
P: 260-280-244-220-304 + 340-303 + 361 + 353-305 + 351 + 338-315-370 + 376-405-403

Switzerland: 0.1 ml m −3 or 0.15 mg m −3

As far as possible and customary, SI units are used.
Unless otherwise noted, the data given apply to standard conditions .

Fluorine [ ˈfluːoːɐ̯ ] is a chemical element with the symbol F and the atomic number 9. In the periodic table it is in the 7th  main group and thus belongs to the 17th  IUPAC group , the halogens , of which it is the lightest. Under normal conditions it is in the form of the diatomic molecule F 2 in gaseous form and is the most reactive of all elements. It reacts with all elements with the exception of the noble gases helium and neon . Fluorine is colorless and appears pale yellow in a highly compressed form. It is the most electronegative of all elements and always has the oxidation state −1 in connection with other elements .

The name of the element is derived from the Latin fluorescent . This term referred to the most important naturally occurring mineral fluorite (fluorspar) , which is used in metallurgy as a flux to lower the melting point of ores (in the original context : lapides igni liquescentes (fluores) ).

Elemental fluorine is very poisonous and extremely corrosive . Its penetrating odor can be noticed even in low concentrations. Its salts ( fluoride and various fluorocomplex salts such as sodium monofluorophosphate ) are also very toxic in higher concentrations, but traces are administered for the prophylaxis of dental caries . Because of this (fluorine or fluoride are involved in the formation of bones and teeth), they are sometimes added to drinking water or table salt ( fluoridation ).


The first fluorine salt described was naturally occurring calcium fluoride (fluorspar). It was described by Georgius Agricola in 1530 and mentioned by him in 1556 as an aid for smelting ores. It makes ore smelting and slag thinner and allows them to flow (flux).

From 1771 Carl Wilhelm Scheele dealt with fluorspar and its properties for the first time, as well as the hydrofluoric acid formed from it during acid treatment . He researched the reactions when hydrofluoric acid acts on glass with the formation of silicon tetrafluoride and fluorosilicic acid . Another property he discovered in fluorspar was fluorescence , which is named after the mineral.

Henri Moissan

In a letter to the editor to Philosophical Magazine in 1808, which is only signed with the abbreviation "EB", the writer complained about what he believed to be an inconsistent approach to naming new elements. In a supplement he hit for the hydrofluoric acid (ger .: fluoric acid called bound) raw material fluorine ago. In a letter to Humphry Davy dated August 25, 1812, André-Marie Ampère expressed the idea that, as in hydrochloric acid, in hydrofluoric acid the radical (English: fluorine, occasionally also fluorine , analogous to chlorine for chlorine) is bound to hydrogen be. Afterwards, many chemists tried to isolate the element, but this was difficult because of its reactivity and toxicity. On June 26, 1886, Henri Moissan succeeded for the first time in producing and characterizing elemental fluorine. He obtained it by electrolysis of a solution of potassium hydrogen difluoride in liquid hydrogen fluoride at low temperatures in a specially developed apparatus (partly from fluorspar). For this achievement Moissan received the Nobel Prize in Chemistry in 1906 .

Fluorine production took off in the Second World War , on the one hand through the development of nuclear weapons in the USA ( Manhattan Project ), since the isotope enrichment of 235 uranium takes place via gaseous uranium hexafluoride (UF 6 ), which is produced with the help of elemental fluorine. On the other hand, IG Farben operated a fluorine electrolysis cell in Gottow at that time, the product of which was supposedly only used to manufacture a new incendiary agent ( chlorine trifluoride ) for incendiary bombs . Whether it would have been possible in Germany at that time to enrich uranium 235 with the help of this fluorine production has been discussed controversially.


Fluorite crystals

In the earth's crust , fluorine is a relatively common element at 525  ppm . Due to its reactivity, it occurs only very rarely in elemental form in nature, but almost exclusively as fluoride in the form of some minerals . An exception is stinkspar (a uranium-containing fluorite variety) and others. from Wölsendorf , as well as Villiaumit, in which small amounts of elemental fluorine were created by radiolysis , which causes a strong odor due to released fluorine during mechanical processing. Seawater contains little dissolved fluoride, since the presence of calcium limits the solubility due to the formation of poorly soluble calcium fluoride. The most common fluorine minerals are the fluorite CaF 2 and the fluoroapatite Ca 5 (PO 4 ) 3 F. Most of the fluorite is bound in fluoroapatite, but this contains only a small mass fraction of fluorine of approx. 3.8%. Therefore, fluorapatite is not degraded because of its fluorine content, but primarily as a source of phosphate. The main source for the extraction of fluorine and fluorine compounds is fluorite. Larger fluorite deposits exist in Mexico , China , South Africa , Spain and Russia . Fluorite is also found in Germany , for example in the Wölsendorf mentioned above.

Another naturally occurring fluorine mineral is cryolite Na 3 AlF 6 . The originally significant cryolite deposits near Ivigtut on Greenland have been exploited. The cryolite required in aluminum production is now produced chemically.

Fluoride ions also occur in some rare minerals, in which they replace the hydroxide groups. Examples are asbestos and the gemstone topaz Al 2 SiO 4 (OH, F) 2 , sellaite MgF 2 and bastnäsite (La, Ce) (CO 3 ) F. The category: fluorine mineral provides an overview .

A few organisms can incorporate fluoride into organofluorine compounds. The South African bush Gifblaar and other plant species of the genus Dichapetalum can synthesize fluoroacetic acid and store it in their leaves. This serves to ward off predators, for which fluoroacetic acid is lethal. The poison effect is triggered by interrupting the citric acid cycle .

Extraction and presentation

Liquid fluorine

The starting material for the extraction of elemental fluorine and other fluorine compounds is predominantly fluorite (CaF 2 ). Hydrogen fluoride is obtained from this by reaction with concentrated sulfuric acid .

Reaction of calcium fluoride with sulfuric acid.

Another source of hydrofluoric acid is phosphate production , where hydrofluoric acid is a waste product from the processing of fluorapatite.

Most of the hydrofluoric acid produced is used to manufacture fluorinated compounds. Where the reactivity of the hydrofluoric acid is insufficient, elemental fluorine is used, which is obtained from a smaller part of the HF production.

Since fluorine is one of the strongest oxidizing agents , it can only be obtained chemically with great difficulty and not economically. Instead, an electrochemical process is used. The gross response is as follows:

The process is named after Henri Moissan . No pure hydrogen fluoride is used for the electrolysis , but a mixture of potassium fluoride and hydrogen fluoride in a ratio of 1: 2 to 1: 2.2. The main reason for using this mixture is that the conductivity of the melt is greatly increased compared to pure hydrogen fluoride, which, like pure water, conducts electricity only very little. For the electrolysis it is important that the melt is completely free of water, otherwise oxygen instead of fluorine would be produced during the electrolysis .

Scheme of a fluorine electrolysis cell

Technically, the so-called medium- temperature process with temperatures of 70 to 130 ° C and a potassium fluoride-hydrogen fluoride mixture of 1: 2 is used. At higher hydrogen fluoride contents, a higher vapor pressure would arise, so that work would have to be carried out at low temperatures and complex cooling. At lower contents (about 1: 1) the melting temperatures are higher (1: 1 ratio: 225 ° C), which makes handling considerably more difficult and promotes corrosion . The electrolysis takes place with graphite - electrodes in cells from steel or Monel instead, the additional iron plates for the separation of anode - and cathode compartment containing, in order to prevent mixing of the gases produced. A voltage of approximately 8-12 volts is applied to the electrodes  . The hydrogen fluoride consumed in the electrolysis is continuously replaced.

The raw fluorine that leaves the electrolysis cell is contaminated with hydrogen fluoride, but also contains oxygen , tetrafluoromethane (CF 4 ) and other perfluorocarbons that are produced by the reaction of fluorine and the electrode material. These impurities can be removed by freezing out and adsorbing hydrogen fluoride on sodium fluoride.

In the laboratory, fluorine can be represented by the decomposition of manganese tetrafluoride (MnF 4 ). For this purpose, K 2 MnF 6 is first mixed with SbF 5 , with the unstable blue-violet MnF 4 being released. This manganese tetrafluoride decomposes at temperatures above 150 ° C in F 2 and MnF 3 .


Physical Properties

Fluorine is a pale yellow, pungent smelling gas at room temperature . The color depends on the thickness of the layer; below a meter thickness the gas appears colorless, only above it is it yellow. Below −188 ° C, fluorine is liquid and has a "canary yellow" color. The melting point of fluorine is −219.52 ° C. Two modifications of solid fluorine are known. Between −227.6 ° C and the melting point, fluorine is present in a cubic crystal structure with a lattice parameter a = 667 pm (β-fluorine). The monoclinic α-modification with the lattice parameters a = 550 pm, b = 328 pm, c = 728 pm and β = 102.17 ° is stable below −227.6 ° C. With a density of 1.6959 kg / m³ at 0 ° C and 1013 hPa, fluorine is heavier than air. The critical point is at a pressure of 52.2 bar and a temperature of 144.2 K (−129 ° C).

Molecular properties

Molecular orbital scheme of fluorine

In its elemental state, fluorine, like the other halogens, is in the form of diatomic molecules . The bond length in the fluorine molecule is 144 pm shorter than the single bonds in other elements (for example, carbon -carbon bond: 154 pm). Despite this short bond, the dissociation energy of the fluorine-fluorine bond is low compared to other bonds at 158 ​​kJ / mol and roughly corresponds to that of the iodine molecule with a bond length of 266 pm. The main reasons for the low dissociation energy are that the lone pairs of electrons of the fluorine atoms come close together and repulsions occur. This weak bond causes the high reactivity of fluorine.

The bond in the fluorine molecule can also be explained by the molecular orbital theory . The s and p atomic orbitals of the individual atoms are combined to form binding and antibonding molecular orbitals. The 1s and 2s orbitals of the fluorine atoms become σ s and σ s * - bonding and antibonding molecular orbitals , respectively . Since these orbitals are completely filled with electrons , they do not add anything to the bond. The 2p orbitals become a total of six molecular orbitals with different energies. These are the bonding σ p , π y and π z , as well as the corresponding antibonding σ p *, π y * and π z * molecular orbitals. The π orbitals have the same energy. If electrons are distributed in the molecular orbitals, it happens that both all binding and antibonding π * orbitals are completely occupied. This results in a bond order of (6–4) / 2 = 1 and a diamagnetic behavior, which is also observed.

Chemical properties

Fluorine is one of the strongest oxidants that are stable at room temperature . It is the most electronegative element and reacts with all elements except helium and neon. The reactions are usually violent. In contrast to all other halogens without light activation , fluorine reacts explosively with hydrogen as a solid at −200 ° C to form hydrogen fluoride . Fluorine is the only element that reacts directly with the noble gases krypton , xenon and radon . So formed at 400 ° C of xenon and fluorine Xenon (II) fluoride .

Many other substances also react vigorously with fluorine, including many hydrogen compounds such as water , ammonia , monosilane , propane or organic solvents . Fluorine reacts differently with water under different conditions. If small amounts of fluorine are poured into cold water, hydrogen peroxide and hydrofluoric acid are formed .

When an excess of fluorine reacts with small amounts of water, ice or hydroxides, on the other hand, the main products are oxygen and oxygen difluoride .

With solid materials, however, fluorine reacts more slowly and in a more controlled manner because of the smaller attack surface. In the case of many metals, the reaction with elemental fluorine leads to the formation of a passivation layer on the metal surface, which protects the metal from further attack by the gas. Since the layer is not tight at high temperatures or fluorine pressures, a further reaction of fluorine and metal can occur, which leads to the melting of the material. Since fresh metal is constantly exposed during melting, which is then ready to react with fluorine, it can ultimately even lead to an uncontrolled course of the reaction (so-called fluorine fire ).

Even plastics usually react in a very controlled manner with elemental fluorine at room temperature. As with metals, the reaction with fluorine also leads to the formation of a fluorinated surface layer in plastic.

Glass is inert to fluorine free from hydrogen fluoride at room temperature. At higher temperatures, however, a more or less rapid reaction is observed. Responsible for this are fluorine atoms, which are formed by the thermal dissociation of the molecular fluorine and are therefore particularly reactive. The product of the reaction is gaseous silicon tetrafluoride . On the other hand, traces of hydrogen fluoride lead to a rapid reaction even without heating.


Fluorine is one of 22 pure elements . Naturally occurring fluorine consists 100% of the isotope 19 F. In addition, another 16 artificial isotopes from 14 F to 31 F and the isomer 18m F are known. Except for the isotope 18 F, which has a half-life of 109.77 minutes, all other artificial isotopes decay within zeptoseconds (10 −21 s) to a little over a minute.

18 F is used in cancer diagnostics in the form of fluorodeoxyglucose , fluoroethylcholine , fluoroethyl tyrosine or 18 F fluoride as a radionuclide in positron emission tomography (PET).

See also: List of Fluorine Isotopes


Due to the high reactivity and the difficult handling of fluorine, elemental fluorine can only be used to a limited extent. It is mainly processed into fluorinated compounds that cannot be produced in any other way. Most of the fluorine produced is required for the production of uranium hexafluoride , which, due to its volatility, enables the enrichment of 235 U with gas centrifuges or the gas diffusion process. This isotope is important for nuclear fission . A second important product that can only be made with the help of elemental fluorine is sulfur hexafluoride . This serves as a gaseous insulator, for example in high-voltage switches and gas-insulated pipelines .

Fluorine is also used to fluorinate the surface of plastics . This is used, among other things, in fuel tanks in automobiles , whereby a fluorinated barrier layer is formed which, among other things, causes lower gasoline permeability. This application of fluorination competes with coextrusion technologies and the metal tank. A second effect of fluorination is that the surface energy of many plastics can be increased. This is mainly used where paints , varnishes or adhesives are to be applied to otherwise hydrophobic plastic surfaces ( polyolefins ). The advantages of fluorinating plastic surfaces are that bodies with pronounced 3D structures and cavities can be treated. In addition, small parts can be treated as bulk goods and the effect is retained for a long time. Fluorination is used when more common and less expensive methods, such as. B. the flame , cannot be used. Other possible effects that can be achieved by fluorinating plastic surfaces are: improved fiber-matrix adhesion , reduced friction and improved selectivities in membrane technology .

If fluorine and graphite are heated together, graphite fluoride is produced , which can be used as a dry lubricant and electrode material.


There are several types of evidence for fluoride ions . In the so-called creep test , a fluoride-containing substance is mixed with concentrated sulfuric acid in a test tube made of glass .

Fluoride ions react with sulfuric acid to form sulfate ions and hydrogen fluoride.

Hydrogen fluoride vapors rise and etch the glass. At the same time, due to the change in the surface, the sulfuric acid is no longer able to wet the glass .

Water drop sample

A second possibility of detection is the so-called water drop test . The fluoride-containing substance is brought together with silica and sulfuric acid. Gaseous silicon tetrafluoride is produced . A drop of water is held over the vessel with the sample. The reaction of silicon tetrafluoride with the water forms silicon dioxide , which crystallizes as a characteristic white border around the drop.

Formation of silicon tetrafluoride.
Reaction in a drop of water

See also: Evidence for Fluoride

In modern analysis , especially for organic fluorine compounds, NMR spectroscopy plays a major role. Fluorine has the advantage of being 100% an isotope ( pure element ) that can be detected by NMR spectroscopy.

Biological importance

It is controversial whether fluorine is an essential trace element for the human organism . According to an opinion published in 2013 by the European Food Safety Authority (EFSA), fluoride is not an essential nutrient, as it is neither used for growth processes nor for tooth development and signs of a fluoride deficiency could not be identified.

The body contains around 5 g of fluoride (for a body weight of 70 kg). It is very unevenly distributed, the vast majority of which is contained in the bones and teeth .

Fluoride can protect against dental caries and harden tooth enamel . The incorporation of fluoride ions instead of hydroxide ions into the hydroxyapatite of the teeth results in fluorapatite . Due to the low solubility of fluorapatite , fluoride has a remineralizing effect , in that the apatite dissolved by acids is precipitated again in the presence of fluoride . In addition, fluoride has an inhibitory effect on certain enzymes and causes an interruption of glycolysis in caries-causing bacteria , which inhibits their growth.

Fluoride is usually absorbed naturally through drinking water or food. If children ingest too much fluoride during tooth development, dental fluorosis can develop . This causes punctiform to spotty brown discolorations ( mottled teeth or mottled enamel ) on the tooth surface, and the tooth is more fragile and less resistant. The maximum recommended maximum amount of fluoride that a person should ingest daily is 0.7 mg for infants up to six months of age and 0.9 mg for children up to three years of age is 0.9 mg. Children aged four to eight should not consume more than 2.2 mg of fluoride per day. When the tooth development is then completed, the person can tolerate higher doses of up to 10 mg fluoride per day.

In Germany, drinking water fluoridation is not permitted; in Switzerland it was carried out in Basel until 2003 . For this reason, no salt containing fluoride was allowed to be sold in Basel until the year 2000.

Since fluoride, similar to selenium , has a toxic effect in larger quantities, there is only a small area in which fluoride can occur in the body without being toxic.


Fluorine and many fluorine compounds are very toxic to humans and other living beings, the lethal dose (LD 50 , one hour) for elemental fluorine is 150–185  ppm . Elemental fluorine has a strong burning and corrosive effect on the lungs , skin and especially the eyes . Even a five-minute contact with 25 ppm fluorine can cause considerable irritation to the eyes. At the same time, the reaction with water (air humidity, skin surface) creates the likewise toxic hydrogen fluoride. Acute fluorine poisoning manifests itself with different symptoms, depending on the route, the compound and the dose in which the fluorine entered the body. A gastrointestinal resulting acute poisoning with fluoride leads to mucosal burns, nausea, initially mucous, later bloody vomiting, unquenchable thirst, severe abdominal pain and bloody diarrhea. Some of those affected die. If hydrogen fluoride and dust-like fluorides are absorbed with the breath, tears, sneezing, coughing, shortness of breath, pulmonary edema and even death with convulsions follow . If the skin is poisoned with hydrogen fluoride (also in acidic solutions of fluoride), it leads to profound necrosis and poorly healing ulcers .

As a weakly dissociated molecule, hydrogen fluoride is easily absorbed through the skin. It leads to painful inflammation, later to stubborn, poorly healing ulcers. In addition, HF forms strong hydrogen bonds and is thus able to change the tertiary structure of proteins. With aluminum ions, fluoride forms fluoridoaluminate complexes, which have a phosphate-like structure and thus contribute to the deregulation of G proteins . The result is an intervention in the receptor-coupled signal transmission and - via signal-dependent phosphorylation / dephosphorylation - in the activity of many enzymes. The best-known example of enzyme inhibition by fluoride is enolase , an enzyme in the glycolysis chain . This inhibition is used when measuring blood sugar : The sodium fluoride placed in the sampling tube inhibits the in-vitro breakdown of glucose after the blood sample, so that the value measured later comes as close as possible to the actual in-vivo value.

The highly toxic fluoroacetates and fluoroacetamide are metabolized to fluorocitrate after absorption . This connection leads to the blockade of the enzyme aconitase, which is important for the citric acid cycle . This causes an accumulation of citrate in the blood, which in turn cuts off the body's cells from the energy supply. Perfluorinated alkanes , which are being tested as blood substitutes , and the commercially available fluorocarbons such as PTFE (Teflon) , PVDF or PFA are considered non-toxic.

The sparingly soluble calcium fluoride , which is formed when it reacts with calcium - for example in the bones - was previously thought to be inert and therefore harmless. At least the dusts of calcium fluoride have proven to be toxic in animal experiments as well as in humans. Whether poorly soluble calcium fluoride is actually formed in vivo in acute fluoride poisoning, as is so often assumed, could not be proven in the context of specific investigations.

Ingesting more than 20 mg of fluoride per day leads to chronic fluoride poisoning, also known as fluorosis . Symptoms are cough, expectoration, shortness of breath, dental fluorosis with changes in the structure and color of the tooth enamel , fluorosteopathy and, in some cases, fluorocachexia. The fluorosteopathy leads to a loss of elasticity and increased bone fragility ( osteosclerosis ) up to the complete stiffening of joints or even the spine due to the increase in bone tissue . Since bone growth can be stimulated at the same time with the help of high doses of fluoride, fluoride is used to treat various forms of osteoporosis .

Previous studies examined a possible link between fluoride intake from drinking water and the occurrence of osteosarcomas , a type of cancer. The amount of fluoride ingested was estimated based on the preferred drinking water source. A statistical analysis of this found a positive correlation between fluoride intake and cancer rate, but only in men, not women. An independent commentary on this analysis notes that the connection was only detectable in the first group of people examined, but not in a later second group. The authors come to the conclusion that no increased cancer risk can be derived from this. In its assessment in 1982, the International Agency for Cancer Research came to the conclusion that there was no evidence of a carcinogenic effect of inorganic fluorides in connection with the fluoridation of drinking water . A connection is now disputed.

Damage caused by working with fluorides, such as skeletal fluorosis, lung damage, irritation of the gastrointestinal tract or chemical burns are recognized as occupational diseases . In the occupational diseases system, they are recorded under Bk No. 13 08.

safety instructions

Due to its high reactivity, fluorine must be stored in special containers. The materials must be designed in such a way that they form a passivation layer when they come into contact with fluorine and thus prevent further reaction. Examples of suitable materials are steel or nickel - copper - alloy Monel . For example, glass , which is attacked by the hydrogen fluoride formed, or aluminum are not suitable . Flammable substances such as fat must also not come into contact with fluorine, as they burn with a violent reaction.

Fluorine does not burn itself, but it has a fire-promoting effect . Fires in the presence of fluorine cannot be extinguished with extinguishing agents ; further entry of fluorine must first be prevented.


As the most electronegative of all elements, fluorine occurs in compounds almost exclusively in the −I oxidation state. Fluorine compounds are known of all elements except helium and neon.

Hydrogen fluoride

Hydrogen fluoride is a very corrosive, poisonous gas. The aqueous solution of hydrogen fluoride is called hydrofluoric acid . While anhydrous, liquid hydrogen fluoride is one of the strongest acids, the so-called super acids , hydrofluoric acid is only moderately strong. Hydrogen fluoride is one of the few substances that react directly with glass . Accordingly, the use as an etching solution in the glass industry is an application of hydrofluoric acid. In addition, hydrogen fluoride is the starting material for the production of fluorine and many other fluorine compounds.


Calcium fluoride

Fluorides are the salts of hydrogen fluoride. They are the most important and most common fluorine salts. The sparingly soluble calcium fluoride CaF 2 occurs in nature in the form of the mineral fluorite . Technically, other fluorides also play a role. Examples are the under use of said uranium hexafluoride or sodium fluoride , which among other things, as a wood preservative is used and when about 100 years ago rat poison and insecticide marketed.

A fluoride frequently used in organic chemistry is tetrabutylammonium fluoride (TBAF). Since TBAF is soluble in organic solvents and the fluoride ion is not influenced by cations (so-called "naked fluoride") it is used as a fluoride source in organic reactions. Another important reaction of tetrabutylammonium fluoride is the cleavage of silyl ethers , which are used as protective groups for alcohols .

Organic fluorine compounds

A number of organic fluorine compounds exist . One of the best-known groups of substances containing fluorine are chlorofluorocarbons (CFCs). The low molecular weight CFCs with one or two carbon atoms are gaseous substances and used to be used as refrigerants in refrigerators and propellants for spray cans . Since these substances increase ozone depletion and thus damage the ozone layer , their production and use have been severely restricted by the Montreal Protocol . In contrast, fluorocarbons are harmless to the ozone layer. Another environmentally harmful impact of fluorine-containing organic compounds is their absorption capacity for infrared radiation . Therefore, they act as greenhouse gases .

An organic fluorine compound known from everyday life is polytetrafluoroethene (PTFE), which is used as a coating on frying pans under the trade name Teflon® . Perfluorinated surfactants , which are used in the manufacture of PTFE, among other things, and other perfluorinated compounds have an extremely stable carbon-fluorine bond. This bond gives the substances a high chemical and thermal resistance, which also means that the substances are persistent in the environment and are hardly broken down.

See also category: organofluorine compounds

Other fluorine compounds

With the other halogens, fluorine forms a number of interhalogen compounds . An important example of this is chlorine trifluoride , a poisonous gas that is mainly used as a fluorinating agent .

Fluorine is more electronegative than oxygen, which is why the compounds between fluorine and oxygen are not referred to as halogen oxides like the other halogen-oxygen compounds , but rather as oxygen fluorides .

In contrast to the heavier halogens, there is only one fluoric acid, the hypofluoric acid HOF. The reason for this is that fluorine does not form three-center-four-electron bonds.

Fluorine also forms some compounds with the noble gases krypton , xenon and argon, such as xenon (II) fluoride . Krypton difluoride , the only known krypton compound, is the most powerful oxidizing agent known . Other noble gas compounds of fluorine often also contain atoms of other elements, such as argon hydrogen fluoride (HArF), the only known argon compound.


Web links

Wiktionary: Fluorine  - explanations of meanings, word origins, synonyms, translations
Commons : Fluor  - album with pictures, videos and audio files

Individual evidence

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This version was added to the list of articles worth reading on January 21, 2008 .