Fluorine (, , or ) is the chemical element with atomic number 9, represented by the symbol F. Under normal conditions, elemental fluorine is a pale yellow gas of diatomic molecules, . In stars, fluorine is relatively rare especially for such a light element. In the Earth's crust, fluorine is more common: the 13th most abundant element. The lightest halogen, fluorine has a single stable isotope, fluorine-19.

Fluorine has the highest electron affinity of any element but chlorine and for this reason is one of the strongest oxidizing agents known. Fluorine forms stable compounds, fluorides, with all other elements for which the reaction has been attempted, except for helium and neon. In contrast to hydrochloric acid, hydrogen fluoride is only a weak acid in water, but is nonetheless extremely corrosive. Metal fluorides are often ionic salts, usually water soluble. Heavier metal elements, such as uranium can also form volatile coordination compound molecules with several fluorine atoms surrounding each metal atom. Organic fluorine compounds tend to have high chemical and thermal stability, water-repellent properties, and low melting and boiling points. Several have large-scale commercial application, such as hydrofluorocarbon gases in refrigeration or the fluorinated plastic polytetrafluoroethylene ("Teflon") in cookware. Production of chlorofluorocarbons and a number of other previously important fluorinated refrigerants that contain halogens other than fluorine, have been banned by international accord since the mid-1990s, as causes of ozone depletion.

Because of its high reactivity, fluorine is not found as a free element on Earth. Fluorine's most important mineral fluorite, calcium fluoride, was first formally described in 1530, in the context of metal smelting. The Latin noun ''fluo'', which means "stream" or "current", gave the name to the mineral, because it was added to metal ores to lower their melting points. Suggested to be a chemical element in 1811, fluorine was named after the source mineral. It was not until 1886 that elemental fluorine was obtained by French chemist Henri Moissan, whose method of electrolysis remains the only industrial production method.

Despite helping prevent tooth decay, fluorine is not considered an essential mineral element for mammals, including humans. However, some organofluorine compounds are synthesized in microorganisms and plants. Several fluorine compounds and elemental fluorine itself are dangerously toxic. Nevertheless, an increasing number of pharmaceuticals (about 10% of new drugs) contain fluorine.

Characteristics

Physical properties

A fluorine atom has nine protons and thus nine electrons, arranged in electronic configuration [He]2s²2p⁵, one fewer than neon. Fluorine's outer electrons are relatively separate from each other, and do not shield each other from the nucleus, thus they experience a relatively high effective nuclear charge of +7. Fluorine both tightly holds its own electrons and has an attraction for one more electron to achieve the extremely stable neon arrangement.

Fluorine's first ionization energy (energy required to remove an electron to form F⁺) is 1,681 kilojoules per mole, which is higher than any other element's except for neon and helium. Second and third ionization energies of fluorine are 3,374 and 6,147 kilojoules per mole, respectively. Fluorine's electron affinity (energy released by adding an electron to form F^–) is 328 kilojoules per mole, which is higher than any other element except chlorine. Fluorine has a relatively small covalent radius, on average about 60 picometers, slightly larger than neon but smaller than oxygen.

Fluorine atoms form diatomic molecules with the chemical formula that are gaseous at room temperature. Though sometimes cited as yellow-green, the gas is a very pale yellow, visible only in long tubes. Fluorine gas's density is 1.696 grams per liter at 100 kilopascal and 0 °C,, about 1.3 times as dense as air.|group="note"}} Fluorine liquefies at −188.1 °C (−306.6 °F), comparable to oxygen and nitrogen, and is then bright yellow.

Fluorine solidifies at −219.6 °C (−363.3 °F) into a cubic (high symmetry) structure, called beta-fluorine. This phase is transparent and soft, with significant disorder of the molecules. At −227.5 °C fluorine undergoes a solid–solid phase transition into a monoclinic (low symmetry) structure called alpha-fluorine. This phase is opaque and hard with close-packed layers of molecules. The solid state phase change requires more energy than the melting point transition and can be violent, shattering samples and blowing out sample holder windows. In general, fluorine's solid state is more similar to oxygen (which has a monoclinic low-temperature structure and similar temperature alpha-beta transition) than to the other halogens.

Isotopes

Fluorine occurs naturally on Earth exclusively in the form of its only stable isotope, fluorine-19, which makes the element both monoisotopic and mononuclidic. In total, at least 17 radioisotopes have been synthesized, ranging in mass number from 14 to 31.

Fluorine-18 is the most stable radioisotope of fluorine, with a half-life of 109.77 minutes, and the lightest unstable nuclide with equal odd numbers of protons and neutrons. All isotopes heavier than the stable fluorine-19 decay via beta minus decay (electron emission), for some isotopes possibly together with neutron emission. Isotopes lighter than the stable fluorine-19 undergo electron capture, while fluorine-17 and fluorine-18 also decay via beta plus decay (positron emission).

To date, only one nuclear isomer has been characterized: fluorine-18m. Its half-life before gamma ray emission is approximately 160 nanoseconds, which is less than that of the ground states of the isotopes from fluorine-17 to fluorine-30, except for fluorine-28.

Chemical reactivity

! X	! XX	! HX	! BX3	! AlX3	! CX4
''F''
Cl	243	428	444	427	327
Br	193	363	368	360	272
I	151	294	272	285	239

Fluorine's chemistry is dominated by its tendency to gain an electron. It is the most electronegative of the elements. Oxidation of fluorine by the extraction of an electron requires so much energy that no known oxidant can oxidize fluorine to any positive oxidation state.

The high direct reactivity of fluorine gas with other elements is also a result of the relatively weak fluorine–fluorine bonds in elemental fluorine. The bond energy is similar to the easily cleaved oxygen–oxygen bonds of peroxides or nitrogen–nitrogen bonds of hydrazines and significantly weaker than those of dichlorine or dibromine molecules. The covalent radius of fluorine in difluorine molecules, about 71 picometers, is significantly larger than that in other compounds, a demonstration of the weak bonding between fluorine atoms.

Reactions between fluorine and other elements (possible for all elements except helium, neon and argon) are often sudden or explosive. It is so reactive that water, halogens, as well as most other substances, even the generally nonreactive ones like the noble gas radon, burn with a bright flame in a jet of fluorine gas. It can even oxidize elemental nitrogen, which is extremely nonreactive due to formation of triple bonds, to give nitrogen trifluoride, but this reaction can occur only at electric discharge.

All metals react with fluorine, forming fluorides, but different conditions are required depending on the metal. Most frequently, the metals must be powdered, because many metals form fluorides layers that resist further oxidation. Alkali metals react with it violently, forming fluorides with formula MF; alkaline earth metals react at room temperature as well, but such reactions are not so exothermic. Ruthenium, rhodium, palladium, platinum and gold, the most nonreactive metals to fluorine, react in atmosphere of pure fluorine only at 300–450 °C. Fluorine reacts explosively with hydrogen in a manner similar to that of alkali metals.

Fluorine is known to form compounds with all elements up to einsteinium, element 99, except for helium, neon, astatine and francium; it is also known to form compounds with rutherfordium, element 104, and seaborgium, element 106. No attempt has been made to oxidize astatine, francium, four later actinides, dubnium and all elements above seaborgium with fluorine, due to the radioactive instability of these elements, though such oxidations are possible in theory. Computational studies have suggested that helium could form a bond with fluorine, and excited states containing neon–fluorine bonds have been observed in a mixture of neon and fluorine irradiated with electrons. Argon forms argon fluorohydride at low temperatures.

Origin and occurrence

	Atomic number	Element	Relative amount

From the perspective of cosmology, fluorine is relatively rare with 400 ppb in the universe. Such a low occurrence is due to quick fusion with hydrogen to form oxygen and helium, or with helium to become neon and hydrogen at solar temperatures. Most fluorine is created either in Type II supernovae when a neutrino hits an atom of neon, in asymptotic giant branch stars, or when a blue Wolf-Rayet star with a mass over 40 solar masses has a stellar wind blowing the fluorine out of the star before hydrogen or helium can destroy it. Even though fluorine, due to its chemical activity, does not exist in its elementary state on Earth, it can be found in the interstellar medium, and fluorine cations exist in stars and planetary nebulae.

Fluorine is the thirteenth most common element in Earth's crust, making up between 600 and 700 ppm of the crust by mass. There are three minerals that are mined and contain enough fluorine to be used as industrial resources. One source for fluoride is fluorite, which is widespread and the most important. It is used in smelting, construction, and the manufacture of hydrogen fluoride. Fluorapatite is mined along with other apatites because of its phosphate content, mostly for production of phosphate fertilizers. The hexafluorosilicates produced as by-product of the phosphoric acid production are mostly disposed of as waste. Cryolite is the least abundant of the three and it is directly used for the production of aluminium. The latter two minerals originate from meteoric water; cryolite has also been found in magmatic waters. Fluorocarbon-containing chlorofluorocarbons and tetrafluoromethane have been reported in rocks, presumably having formed without action of living organisms. However, they are not a commercially or environmentally important source of fluorine.

{{Image gallery |title=Notable fluorine-containing minerals |lines=1 |width=200 |height=250 |align=center |Fluorite-270246.jpg|pink globular mass with crystal facets|

Fluorite |Apatite Canada.jpg|Long prism-like crystal, without shinyness, at an angle coming out of aggregate-like rock|Fluorapatite |Ivigtut cryolite edit.jpg|A white rock on a black background|Cryolite }}

History

"Fluorine" is a word that ultimately derives from the Latin noun ''fluo'', meaning stream. The mineral fluorite, a natural form of calcium fluoride, was first mentioned in 1529 by Georgius Agricola, who named it after its use as a "flux"—an additive that helps melt ores and slags during smelting. Agricola first named the mineral "fluorspar" as a latinization of the German ''Flußspat''. Since then, the mineral has been renamed "fluorite," although "fluorspar" is still sometimes used.

The first recorded preparation of "fluoric acid" (hydrofluoric acid in modern nomenclature) was in 1764 by Andreas Sigismund Marggraf, who heated fluorite with sulfuric acid in glass, which was greatly corroded by the product. In 1771, Swedish chemist Carl Wilhelm Scheele repeated this reaction. In 1810, French physicist André-Marie Ampère suggested that the acid was a compound of hydrogen with an unknown element, analogous to chlorine; fluorite was then shown to be mostly composed of calcium fluoride. Sir Humphry Davy originally suggested to call the element ''fluorine'', taking the root from the name of "fluoric acid" and the -ine suffix, similarly to other halogens; this name, with modifications, came to most European languages; however, in Greek, Russian and several other languages the name ''ftor'' and deratives are in use, coming from Greek ''φθόριος'', meaning "destructive". The new Latin name (''fluorum'') gave the element its current symbol F, even though the symbol Fl is also seen in pre-Moissan papers.

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 all but certain exotic materials. The generation of elemental fluorine from hydrofluoric acid proved to be exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. Jean Dussaud referred to them as "fluorine martyrs", a term still used. In 1886, the isolation of elemental fluorine was reported by French chemist Henri Moissan after almost 74 years of effort by other chemists. For Moissan, the feat helped earn the 1906 Nobel Prize in chemistry.|group="note"}}

The two most prominent developments of organofluorine compounds, chlorofluorocarbon refrigerants (for example Freon-12, introduced in the late 1920s), and Teflon (invented 1938) were both associated with the DuPont company. Chlorofluorocarbons are now being replaced by hydrofluorocarbons.

The first large-scale productions of elemental fluorine were by opposing sides during World War II. Germany initiated high-temperature electrolysis production of fluorine to produce tons of chlorine trifluoride, a compound planned as an incendiary. The United States's Manhattan project produced even more fluorine for use in uranium separation. Gaseous uranium hexafluoride, was used to separate uranium-235, key for nuclear explosions, from the more atomically inert uranium-238 in centrifuges and diffusion plants.

Production

Industrially, fluorine is used either directly as the mined mineral fluorite or as hydrogen fluoride (from reaction of sulfuric acid with fluorite). Only a very small fraction of industrial fluorine is ever electrolyzed to molecular fluorine, F₂. Most fluorine contained in organofluorines is not from molecular fluorine, but is derived from hydrogen fluoride.

Electrolytic synthesis

Several thousand tons of elemental fluorine are produced annually by electrolysis of potassium bifluoride in hydrogen fluoride. Potassium bifluoride forms spontaneously from potassium fluoride and the hydrogen fluoride:

:HF + KF → KHF₂

The mixture with the approximate composition KF•2HF melts at 70 °C and is electrolyzed between 70 °C and 130 °C. Potassium bifluoride increases the electrical conductivity of the solution and provides the bifluoride anion that is oxidized to form fluorine at the anode, while hydrogen forms at the cathode. Formed fluoride ions remain in solution, with soluted compound after electrolysis being potassium fluoride. :2 HF₂^– → H₂↑ + F₂↑ + 2 F^– These electrolytes and the electrolytic method using them are essentially these first pioneered by Henri Moissan. However, improvements have since been made in electrodes and containment: while Moissan used platinum group metal electrodes and carved fluorite containers, the modern process uses the steel cell itself as cathode, while blocks of carbon are used as anode (the Söderberg carbon electrodes are similar to those used in the electrolysis of aluminium). The voltage for the electrolysis varies between 8 and 12 volts.

If the fluorine gas is cleaned of hydrogen fluoride and oxygen impurities, it may be stored in steel cylinders, where the inside surface becomes passivated with a metal fluoride layer that resists further attack.

Chemical routes

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 potassium hexafluoromanganate and antimony pentafluoride at 150 °C, conducted in hydrogen fluoride: :2 + 4 → 4 + 2 + ↑ This synthetic route is conceptually interesting as a rare chemical preparation of elemental fluorine – a reaction not previously thought possible. However, the manganese(IV) fluoride is only known to be prepared via reaction with fluorine gas itself, or reaction with krypton difluoride, also obtained upon reaction with elemental fluorine and therefore this is not a industrially suitable way to produce fluorine; other perfluoromangantes(IV) also require fluorine gas. Other fluorine-forming reactions are known with hexafluoronickelate ion, , and the heavier noble gases (krypton and xenon) fluorides, which also require elemental fluorine to be produced.

Economic aspects

With the exception of countries with planned economics, about 17,000 tonnes of fluorine are produced per year, all coming from 11 companies from G7 countries. Price of fluorine is relatively small, at the level of 5–8 United States dollars per kilogram, if units for fluorine gas production are integrated into those of uranium hexafluoride or sulfur hexafluoride. However, because of difficulties in storage and handling, the price is much higher when fluorine is supplied in cylinders.

Compounds

Fluorine atoms are in the −1 oxidation state in all molecules except for difluorine, where the atoms are bonded to each other and thus at oxidation state 0. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds; however, higher bonding is also known, for example, boron monofluoride features a triple bond. Fluoride may be a linear bridging atom between two metals in some complex molecules. However, molecules containing fluorine may exhibit hydrogen bonding.

Inorganic compounds

Inorganic acids

Unlike other hydrohalic acids, such as hydrochloric acid, hydrofluoric acid is only a weak acid in water solution, with acid dissociation constant (pK_a) equal to 3.18. However, water is not an inert solvent for hydrogen fluoride. When less basic solvents such as dry acetic acid are used, hydrofluoric acid is the strongest of the hydrohalic acids. Despite its weakness as an acid in water, HF is inherently very corrosive, attacking glass. Owing to the basicity of the fluoride ion, soluble fluorides give basic water solutions.

Perfluoroacids, i. e. acids that contain only hydrogen, fluorine and atoms of one more element, named central as it is in the center of the acid's anion, are usually very strong. Fluoroantimonic acid, one such an acid, is a "superacid" and the strongest acid known. It has an extremely low pK_a of −31.3 and is 20 quintillion (2) times stronger than pure sulfuric acid, which has pK_a of −12. This happens because fluorine atoms are univalent and thus cannot form strong chemical bonds to both antimony atom and the hydron. Occupying all antimony's valence electrons, fluorine atoms do not let the hydron to be bonded to it.

Metal fluorides

Metal fluorides have similarities to other metal halides and to metal oxides but in them, the ionic character is stronger than in corresponding chlorides or oxides. Ionic fluorides vary in solubility greatly, but with a trend to decrease solubility as the number of fluorides grows up: the alkali metal fluorides often resemble the chlorides in terms of structure (all having sodium chloride structure) and solubilities. Reflecting the high basicity of fluoride anion, many alkali metal fluorides form bifluorides with the formula MHF₂; this is a well-known process for sodium and potassium. Out of other monofluorides, only silver(I) and thallium(I) fluorides are well-characterized, both very soluble, unlike other corresponding halides. Unlike the chlorides, alkaline earth metals (except for beryllium) fluorides are sparingly soluble. However, several other difluorides, such as those of copper(II) and nickel(II), are soluble. No trifluoride is soluble in water, but several ones may be soluble in other solvents.

{{Image gallery |title=The fluorides of later first row transition metals |lines=2 |width=120 |height=250 |align=center |Fluorid manganatý rotated.PNG|white powder in a tube and on a spoon|

Manganese difluoride |Fluorid železitý rotated.PNG|brown powder in a tube and on a spoon|Iron trifluoride |Fluorid kobaltnatý rotated.PNG|pink powder in a tube and on a spoon|Cobalt difluoride |Fluorid nikelnatý rotated.PNG|light green powder in a tube and on a spoon|Nickel difluoride |Fluorid měďnatý rotated.PNG|light blue powder in a tube and on a spoon|Copper difluoride }}

While metal tri- and lower fluorides are ionic solids, metal penta- and higher fluorides are molecular and volatile. Tetrafluorides are the borderline: for example, zirconium tetrafluoride is an ionic solid, but germanium tetrafluoride is a molecular gas. This property of the fluoride ion is caused by its small radius. Only one metal is known to form bonds with seven fluorides—rhenium, forming rhenium heptafluoride, which hold the record for number of charged ligands for a metal compound; the compounds shares the pentagonal bipyramidal with iodine heptafluoride (the only analogous nonmetal compound, which is well-known). Metal hexafluorides (and higher fluorides, existing or possible) are oxidants; for example, platinum hexafluoride was the first compound to oxidize oxygen molecule and xenon (see below). Such compounds vary in their phases, as volatile solid, liquid and gaseous fluorides are known.

Nonmetal fluorides

Nonmetal fluorides are all volatile. Period 2 elements (with the exception of boron, which forms a trifluoride) form fluorides that follow octet rule: carbon tetrafluoride, nitrogen trifluoride and oxygen difluoride, as they need some of theirs electrons to be transferred to 3d-orbitals to break the rule, which requires too much energy. The following periods may form fluorides that are hypervalent molecules, such as phosphorus pentafluoride. The reactivity of such species varies greatly: sulfur hexafluoride is inert, while chlorine fluorides are oxidants.

Boron trifluoride is a Lewis acid, while silicon tetrafluoride is a weaker acid and less stable thermally; tetrafluoromethane (carbon tetrafluoride) is relatively stable chemically. Among pnictogens, reactivity of pentafluorides increases down the group, as well as the acidity; bismuth is an exception, because its pentafluoride is not as acidic as its antimony analog due to its polymeric structure and its trifluoride is ionic. Nitrogen is another special case because it is not known to form a pentafluoride, but tetrafluoroammonium ion, , with nitrogen in the formal oxidation state of +5, is known. Chalcogens show a similar trend: hexafluorides increase in acidity and reactivity down the group; oxygen is not known to be oxidized above difluoride. Halogens, unlike previous groups, not all form highest fluorides of highest oxidation states; chlorine and bromine form pentafluorides, both strong fluorinators; iodine may be oxidized up to iodine heptafluoride. Astatine is not well-studied, and unlike its other halides, no astatine fluoride has been produced,

Noble gas compounds

The noble gases, helium, neon, argon, krypton, xenon, and radon are generally non-reactive because they all have fully filled electronic shells, which are extremely stable. The reactivity of fluorine-containing platinum hexafluoride toward the noble gas xenon was first reported by Neil Bartlett only in 1962. He called the compound he prepared xenon hexafluoroplatinate, but since then that has been revealed to be mixture of different chemicals.|group="note"}} Later that year, xenon was oxidized directly with fluorine, to form xenon difluoride. Today, only xenon and krypton have well-characterized binary noble gas–fluorine compounds, which are xenon difluoride, krypton difluoride, xenon tetrafluoride, krypton tetrafluoride, xenon hexafluoride and their deratives. Several oxyfluorides and oxyfluoroxenates are known, including xenon oxytetrafluoride, XeOF₄.

Radon readily reacts with fluorine to form a solid compound, which is generally thought to be radon difluoride. However, it decomposes on attempted vaporization and its exact composition is uncertain. Calculations have shown that radon difluoride can be ionic, unlike all other binary noble gas fluorides.

Argon can react at extreme conditions with hydrogen fluoride, to form its only stable compound—argon fluoride hydride. Helium can form an analogous helium fluoride hydride but it is metastable, with a lifetime of at most 14 nanoseconds. Argon forms binary argon monofluoride, ArF^•, but it is metastable; because of its metastability, the compound found its use in the argon fluoride laser. Even though the reactivity of neon is the lowest of all the elements,|group="note"}} the element also forms a chemical compound, neon monofluoride, NeF^•, which is metastable.

Ununoctium, the last currently known group 18 element, is predicted to form ununoctium difluoride, , and ununoctium tetrafluoride, , which is likely to have the tetrahedral T_d configuration. However, only four atoms of ununoctium have been synthesized, and its chemical properties have not been examined yet.

Comparison between the highest oxidation states of oxides and fluorides

For groups 1—6 and 13—16 highest oxidation states of oxides and fluorides are always equal; the difference is only seen in groups 7—11, mercury, halogens, and the noble gases. The general trend is fluorination allows to achieve relatively low but hardly achievable for an element compounds; for example, no binary oxide is known for krypton, but krypton difluoride is well-known. However, very high oxidation states of several elements are known for oxygen only; for example, none has shown that existence of ruthenium octafluoride is possible yet, while ruthenium tetroxide is known.

Taking relatively low but hard-to-achieve oxidation states of metals, fluorine is the key in achieving many rare high oxidation states of the transition metals. For instance, direct reaction of the respective metals with fluorine gives rise to palladium(VI) and platinum(VI). The only occurrence of mercury(IV) is binary mercury(IV) fluoride, synthesized at temperatures close to absolute zero. Fluorine-containing complexes of copper(IV), silver(IV), nickel(IV), iridium(VI), and others that are examples of element oxidation state fluorine-containing compounds, rarely occurring in any other compounds, are known. Gold(V) is only known in hexafluoroaurate(V) ion, which can be synthesized indirectly on extreme conditions, and gold(V) fluoride, which is obtained during hexafluoroaurate(V) decomposition. The high oxidizing potential of fluorine has led to claim of gold(VII) existence in gold heptafluoride, but current calculations show that the claimed AuF₇ molecule was only AuF₅·F₂. It is also possible that the element 113, ununtrium, will be the first element in boron group to form a species in +5 oxidation state, the fluorine-based hexafluoroununtrate(V), ; possibility of +5 oxygen-based species is not known to be calculated.

Even though fluorine is a generally stronger oxidizer than oxygen, creating, for example, nitrogen pentafluoride would need to squeeze five fluorine atoms attached to the small central atom, which is hard to perform, and the resulting molecule may not be stable at all; however, existence of this compound cannot be denied at all. Similarly, the highest oxidation states of several late transition metals may be achieved in oxides only: for example, even though only gold(V) is known now, and only in form of a fluoride, calculations show that the element may be oxidized up to gold(IX), if form of tetroxoaurlyl(IX) ion, [AuO₄]⁺, but not a fluorine-based compound or ion. platinum(X), however, a later work denies these species and predict that platinum(VI), gold(V), and mercury(IV), all known in binary fluorides, are the highest for the elements. It also shows osmium and iridium may form heptafluorides; for osmium, even an octafluoride may be possible.

Among halogens, chlorine and bromine form perchlorates and perbromates, both oxygen-based and with halogen in +7 state; however, chlorine, unlike bromine, also forms a binary heptoxide. For their stable fluorinated species, pentafluorides are the highest species achieved; however, bromine hexafluoride, BrF₆^•, is known as well. Iodine shows the reverse picture: no heptoxide is known, unlike heptafluoride, a well-known stable compounds; however, periodic acid, containing iodine(VII), is known as well.

Noble gases do not show a trend as well: as noted above, krypton has no known binary oxides, but has a well-studied difluoride. Xenon forms a tetroxide of oxygen-based species, but only a hexafluoride of fluorine-based ones. Neutral xenon octafluoride ion is not known nor expected to be stable, but octafluoroxenate(VI), , has been synthesized. Contradictory data is known about fluorides and especially oxides of radon; no binary fluoride or oxide of lighter noble gases are known.

Organic compounds

Organofluorine compounds are chemical compounds containing a carbon–fluorine chemical bond. This bond is the strongest covalent bond in organic chemistry and is very stable. Fluorine replaces hydrogen in hydrocarbons even at room temperature; after the reaction, molecular size is not changed significantly. The range of organofluorine compounds is thus diverse, in part because the area is driven by commercial value of such compounds in materials science and pharmaceutical chemistry. Organofluorine compounds are synthesized via both direct reaction with fluorine gas, which can be dangerous due to the reactivity of the elemental fluorine, and reaction with fluorinating reagents such as sulfur tetrafluoride.

The most industrially important compounds of fluorine include Teflon and hydrofluorocarbons, the main properties of which are affected by carbon–fluorine bonds in them. The slippery nature of Teflon is the result of chemical stability and repulsion of highly charged fluorine atoms in polymeric chains. Resistance of the chemical to van der Waals forces means that it is the only known surface to which a gecko cannot stick. Properties of the chlorofluorocarbons and hydrochlorofluorocarbons are relative to the number and identity of the halogen atoms. Volatility of these compounds is lower than in most organic compounds due to strength of carbon–fluorine bond and carbon–chlorine bond. This is caused by the molecular polarity induced by the halides and the polarity of halides, which induces intermolecular interactions; due to large difference between chlorine and fluorine atom radii, the compounds molecules do not have symmetry, which increases the polarity in the molecules; these effects lead to high solubility potential and higher boiling points of chlorofluorocarbons compared to those of parent hydrocarbons. Chlorofluorocarbons are far less flammable than methane, in part because they contain fewer carbon–hydrogen bonds and in part because, in the case of the chlorides and bromides, the released halides quench the free radicals that sustain flames.

The large inductive (electron-withdrawing) effect of the trifluoromethyl group results in the high strength of many fluorinated organic acids, which may be comparable to mineral acids. In these compounds, the affinity of the acid cation for the acid proton is decreased by the cation's fluorine content, which increase its affinity for the extra electron left when the acidic proton leaves. For example, acetic acid is a weak acid, with pK_a equal to 4.76, while its fluorinated derivative, trifluoroacetic acid has pK_a of −0.23, giving it 33,000 times greater formal acidic potential.

Applications

Approximately half of mined fluorite is used to help molten metal flow, especially in iron smelting. The other half is converted to hydrofluoric acid, which is mostly used to produce organofluorides or synthetic cryolite.

Uses of fluorine gas

Elemental fluorine is occasionally used as a fluorinating agent in industrial processes. The largest use for elemental fluorine is preparing uranium hexafluoride, used in the production of nuclear fuels. To obtain the compound, uranium dioxide is treated with hydrofluoric acid, to produce uranium tetrafluoride, which is oxidized by fluorine to give uranium hexafluoride. Unlike uranium hexafluoride, sulfur hexafluoride, used as an inert dielectric medium in high voltage switching and the second largest use for fluorine gas, does not necessarily require fluorine to be produced, but most often, it is produced in reaction between sulfur and fluorine.

Other than those, elemental fluorine is used for production tetrafluoromethane, plasma etching in semiconductor manufacturing, flat panel display production, and microelectromechanical systems fabrication. These and other uses are said to require up to 2,000 tonnes annually.

United States and Soviet space scientists in the early 1960s studied elemental fluorine as a possible rocket propellant, due to its exceptionally high specific impulse when used as an oxidizer. The experiments failed because fluorine proved difficult to handle, and its combustion product (typically hydrogen fluoride) was extremely toxic and corrosive.

Uses of compounds

Inorganic fluorides and organofluorine compounds, a fraction of which are prepared from elemental fluorine, find use in a variety of materials and chemicals, including important pharmaceuticals, agrochemicals, lubricants, and textiles.

The first recorded use of fluorides was to help molten metal flow by breaking the oxide coating of metals and offering a lower-friction surface. For this reason, the name of fluorite is derived from Latin verb ''fluere'', meaning to flow.

Hydrofluoric acid and certain fluoride-containing salts are useful etchants for glass, e.g. for light bulbs. Laboratory-produced sodium hexafluoroaluminate, better known as synthetic cryolite, the mineral composed mostly out of this chemical, is used in aluminium metallurgy. In the electrolysis of the metal and its purification, it acts to lower the melting point of aluminium oxide and acts like a powerful flux for glass.

Perfluorooctanoic acid and tetrafluoroethylene are directly used in water resistant coatings and in the production of low friction plastics such as Teflon, or PTFE. The low van der Waals forces in solid Teflon give it unusual antiadhesive properties. Nafion, a strongly acidic fluorinated polymer, is a component of fuel cells.

Other fluorine-based compounds are used in the production of haloalkanes such as chlorofluorocarbons, which are used extensively in air conditioning and in refrigeration. They have been banned for these applications because they contribute to ozone destruction.

Isotope applications

Natural fluorine is monoisotopic, consisting of fluorine-19. Fluorine compounds are highly amenable to nuclear magnetic resonance, because fluorine-19 has a nuclear spin of ½, high nuclear magnetic moment and a high magnetogyric ratio, making measurements very fast, comparable with similar effect based on hydrogen-1. Fluorine-19 nuclear magnetic resonance is not one of basic magnetic resonance spectrometers used in science and medicine, but using this isotope is quite common. It has found uses in studies of protein structures and conformational changes. The monoisotopic occurrence of fluorine assists in its use in uranium enrichment, as uranium hexafluoride molecules differ in mass only due to uranium atom isotope mass differences, i.e. uranium-235 and uranium-238. Molecules with different masses are separated via diffusion and gas centrifugation.

Compounds containing fluorine-18, a radioactive isotope that emits positrons, are often used in positron emission tomography ("PET scanning"), because its half-life of about 110 minutes is long by the standards of positron-emitters. One such species is 2-deoxy-2-(¹⁸F)fluoro-D-glucose, commonly abbreviated as ¹⁸F-FDG. In PET imaging, ¹⁸F-FDG can be used for the assessment of glucose metabolism in the brain and for imaging tumors in oncology. This radiopharmaceutical is retained by cells and is taken up many tissues with a high need for glucose, such as the brain and most types of malignant tumors. As a result, the tomography can be used for diagnosis, staging, and monitoring treatment of cancers, particularly in Hodgkin's disease, lung cancer, breast cancer, and many others.

Biological roles

Living organisms

Fluoride is not considered an essential mineral element for mammals and humans, in the sense of being necessary for life, but its role in prevention of tooth decay is well-established. Sodium fluoride, tin(II) fluoride, and, most commonly, sodium monofluorophosphate, are used in toothpaste. These or related compounds, such as fluorosilicates are added to many municipal water supplies, a process called water fluoridation, which has had controversy since its beginnings in 1945. Small amounts of fluoride may be beneficial for bone strength, but this is an issue only for the formulation of artificial diets.

Biologically synthesized organofluorines have been found in microorganisms and plants, but not animals. The most common example is fluoroacetate, which occurs as a plant defense against herbivores in at least 40 plants in Australia, Brazil and Africa. Other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate, which are all believed to be biosynthesized in biochemical pathways from the intermediate fluoroacetaldehyde. Adenosyl-fluoride synthase is an enzyme capable of biologically synthesizing the carbon–fluorine bond.

Pharmaceuticals, agricultural chemicals, and poisons

Several important pharmaceuticals contain fluorine. Because of the considerable stability of the carbon-fluorine bond, many drugs are fluorinated to prevent their metabolism and prolong their half-lives, allowing for longer times between dosing and application. For example, an aromatic ring may be a useful drug addition to prevent metabolism, but presents a safety problem: enzymes in the body metabolize some of them into poisonous epoxides. When the ''para'' position is substituted with fluorine, however, the aromatic ring is protected and epoxide is no longer produced. Another reason for fluorination of biologically active organics is that many organic compounds show increased lipophilicity due to addition of fluorine (the carbon–fluorine bond is even more hydrophobic than the carbon–hydrogen bond), and this effect often increases a drug's bioavailability due to increased cell membrane penetration. Since the carbon–fluorine bond is strong, organofluorides are generally very stable towards metabolism, although the potential of the fluorine to be released as a fluoride leaving group is heavily dependent on the its position in the molecule.

Out of drugs that were commercialized in the past 50 years, 5–15% contain fluorine, and this percentage is increasing. For example, fludrocortisone is one of the most common mineralocorticoids, a class of drugs that mimics the actions of aldosterone. The anti-inflammatories dexamethasone and triamcinolone, which are among the most potent of the synthetic corticosteroids class of drugs, contain fluorine. Several inhaled general anesthetic agents, including the most commonly used inhaled agents, contain fluorine. Examples are sevoflurane, desflurane, and isoflurane, which are hydrofluorocarbon derivatives.

Many SSRI antidepressants are fluorinated organics, such as citalopram, escitalopram, fluoxetine, fluvoxamine, and paroxetine. Fluoroquinolones are a commonly used family of broad-spectrum antibiotics. Because of the difficulty of biological systems in dealing with metabolism of fluorinated molecules, fluorinated pharmaceuticals (often antibiotics and antidepressants) are among the major fluorinated organics found in treated city sewage and wastewater.

In addition to pharmaceuticals, an estimated 30% of agrochemical compounds contain fluorine. Although this water is usually not treated along with sewage, it does contaminate rivers with runoff organofluorines. Synthetic sodium fluoroacetate has been used as an insecticide, especially against cockroaches, and is effective as a bait-poison against mammalian pests. Several other insecticides contain sodium fluoride, which is much less toxic than fluoroacetate.

Chronology

Because groundwater contains fluorine ions, organic items such as bone that are buried in soil will absorb those ions over time. As such, it is possible to determine the relative age of an object by comparing it with the amount of fluoride in another object found in the same area. It is important in separation technique in intra-site chronological analysis and inter-site comparisons.

However, if no real age of any object is known, ages can only be expressed in older than or younger than between the two objects. The fluctuating amount of fluoride found in groundwater means the objects in comparison must be in the same local area in order for the comparisons to be accurate. This technique is not always reliable, given that not all objects absorb fluorine at the same rates.

Environmental concerns

Chlorofluorocarbons and bromofluorocarbons have recently come under strict environmental regulation due to their long residence times in the atmosphere, and their contribution to ozone depletion. Since it is specifically chlorine and bromine radicals that harm the ozone layer, not fluorine, compounds that 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, and have been widely used as replacements for the chlorine- and bromine-containing halocarbons. Hydrofluorocarbons and perfluorocarbons however, function as another type of pollutant: they are greenhouse gases about 4,000 to 10,000 times that of carbon dioxide. Sulfur hexafluoride exhibits an even stronger such an effect, exceeding that of carbon dioxide 20,000 times.

Because of the strength of the carbon–fluorine bond, many synthetic fluorocarbons and fluorocarbon-based compounds are persistent in the environment. The fluorosurfactants perfluorooctanesulfonic acid (PFOS) and perfluorooctanoic acid (PFOA), used in waterproofing sprays, and other related chemicals, are persistent global contaminants. PFOS is a persistent organic pollutant and may be harming the health of wildlife. The potential health effects of PFOA to humans are not well-known; its tissue distribution in humans is unknown, but studies in rats suggest it is likely to be present primarily in the liver, kidney, and blood, being absorbed easily via the gastrointestinal tract in rats. PFOA has been shown not to metabolize in the body, not to be genotoxic nor lipophilic, unlike chlorinated hydrocarbons; it binds to serum albumin and is excreted primarily from the kidney.

Precautions

Elementary state, fluoride ion, and fluoroacetate

Elemental fluorine is a highly toxic, corrosive oxidant, and is extremely reactive to organic material (except for perfluorinated substances) even at very low concentrations and can even cause ignition on larger ones. Fluorine gas has a pungent characteristic odor that is noticeable in concentrations as low as 20 ppb. Significant irritation to humans can be caused by concentration of fluorine of 25 ppm; at this and higher concentrations fluorine attacks eyes, respiratory tract, lungs, liver and kidneys. At concentration of 100 ppm, human eyes and noses are irritated and damaged seriously.

Soluble fluorides are moderately toxic. In the case of the simple salt sodium fluoride, 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). A toxic dose that may lead to adverse health effects is estimated at 3 to 5 mg/kg of elemental fluoride. The fluoride ion is somewhat toxic in biology, in part because of its ability to form, by equilibration, small amounts of hydrogen fluoride in water, and this mobile uncharged species diffuses across cell membranes to attack intracellular calcium. Fluoride ion is readily absorbed by the stomach, intestines and excreted through urine. Urine tests have been used to ascertain rates of excretion in order to set upper limits in exposure to fluoride compounds and associated detrimental health effects. Ingested fluoride initially acts locally on the intestinal mucosa, where it forms hydrofluoric acid in the stomach. Thereafter it binds calcium and interferes with various enzymes. Excess of fluoride consumption can lead to skeletal fluorosis, of which currently millions people are now are affected.

Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluoride, or (more rarely) rodenticides containing sodium fluoroacetate ("Compound 1080") containing organofluorine. Currently, most fluoride poisonings are due to the ingestion of fluoride-containing toothpaste. Malfunction of water fluoridation equipment has occurred several times, including a notable incident in Alaska, effected nearly 300 people and one person died.

Hydrofluoric acid

Hydrofluoric acid is a contact poison, and must be handled with extreme care far beyond that accorded to other mineral acids, even the analogous hydrochloric acid, HCl. Owing to its lesser chemical dissociation in water (remaining a neutral molecule), hydrogen fluoride penetrates tissue more quickly than typical acids. Poisoning can occur readily through exposure of skin or eyes, or when inhaled or swallowed. Symptoms of exposure to hydrofluoric acid may not be immediately evident. Hydrogen fluoride interferes with nerve function, meaning that burns may not initially be painful. Accidental exposures can go unnoticed, delaying treatment and increasing the extent and seriousness of the injury.

Once absorbed into blood through the skin, hydrogen fluoride reacts with blood calcium and may cause cardiac arrest. Formation of insoluble calcium fluoride possibly causes both fall in calcium serum and the strong pain associated with tissue toxicity. In some cases, exposures can lead to hypocalcemia. Burns with areas larger than 160 cm² (25 in²) can cause serious systemic toxicity from interference with blood and tissue calcium levels.

Hydrofluoric acid exposure is often treated with calcium gluconate, a source of Ca²⁺ that binds with the fluoride ions. Hydrogen fluoride chemical burns to the skin can be treated with a water wash and 2.5% calcium gluconate gel or special rinsing solutions. However, because it is absorbed, medical treatment is necessary; in some cases, amputation may be required.

See also

Caesium—the least electronegative element, fluorine's "opposite"

Notes

References

Bibliography

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External links

Category:Biology and pharmacology of chemical elements Category:Chemical elements Category:Fluorinating agents Category:Halogens Category:Oxidizing agents

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