Fluorine

Fluorine, ₉F

Liquid fluorine (F2 at extremely low temperature)
Fluorine
Pronunciation	/ˈflʊəriːn/ /ˈflɔːriːn/ (FLOR-een)
Allotropes	alpha, beta (see Allotropes of fluorine)
Appearance	gas: very pale yellow liquid: bright yellow solid: alpha is opaque, beta is transparent
Standard atomic weight Ar°(F)
	18.998403162±0.000000005 18.998±0.001 (abridged)
Fluorine in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson – ↑ F ↓ Cl oxygen ← fluorine → neon
Atomic number (Z)	9
Group	group 17 (halogens)
Period	period 2
Block	p-block
Electron configuration	[He] 2s2 2p5
Electrons per shell	2, 7
Physical properties
Phase at STP	gas
Melting point	(F2) 53.48 K ​(−219.67 °C, ​−363.41 °F)
Boiling point	(F2) 85.03 K ​(−188.11 °C, ​−306.60 °F)
Density (at STP)	1.696 g/L
when liquid (at b.p.)	1.505 g/cm3
Triple point	53.48 K, ​.252 kPa
Critical point	144.41 K, 5.1724 MPa
Heat of vaporization	6.51 kJ/mol
Molar heat capacity	Cp: 31 J/(mol·K) (at 21.1 °C) Cv: 23 J/(mol·K) (at 21.1 °C)
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, 0 (oxidizes oxygen)
Electronegativity	Pauling scale: 3.98
Ionization energies	1st: 1681 kJ/mol 2nd: 3374 kJ/mol 3rd: 6147 kJ/mol (more)
Covalent radius	64 pm
Van der Waals radius	135 pm
Spectral lines of fluorine
Other properties
Natural occurrence	primordial
Crystal structure	​cubic
Thermal conductivity	0.02591 W/(m⋅K)
Magnetic ordering	diamagnetic (−1.2×10−4)
CAS Number	7782-41-4
History
Naming	after the mineral fluorite, itself named after Latin fluo (to flow, in smelting)
Discovery	André-Marie Ampère (1810)
First isolation	Henri Moissan (June 26, 1886)
Named by	Humphry Davy André-Marie Ampère
Isotopes of fluorine
Main isotopes Decay abun­dance half-life (t1/2) mode pro­duct 18F trace 109.734 min β+ 18O 19F 100% stable

Fluorine is the chemical element with atomic number 9, represented by the symbol F. It is the lightest member of the halogen column of the periodic table and has a single stable isotope, fluorine-19. At standard pressure and temperature, fluorine is a pale yellow gas composed of diatomic molecules, F
₂. In stars, fluorine is relatively rare compared to other light elements. On Earth, fluorine is more common; it is the 13th most abundant element in the crust.

Fluorine's most important mineral, fluorite, was first formally described in 1530, in the context of metal smelting. The mineral's name derives from the Latin verb fluo, which means "stream" or "flow," because fluorite 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 of fluorine gas. The main use of elemental fluorine, uranium enrichment, was developed during the Manhattan Project. The vast majority, however, of commercial fluorine is never reduced to the element.

Fluorine has the highest electron affinity of any element but chlorine, and for this reason it is a very strong oxidizing agent. Fluorine forms stable compounds, fluorides, with all elements, except for helium and neon, for which the reaction has been attempted. Hydrofluoric acid, in contrast to other haloacids such as hydrochloric acid, is only a weak acid in water, but it is nonetheless extremely corrosive. Fluorides of lighter metal elements are ionic compounds (salts), which are usually water soluble. Heavier metal elements such as uranium can form volatile coordination compounds (separate molecules with several fluorine atoms surrounding a metal atom). Organic fluorine compounds tend to have high chemical and thermal stability and water-repellent properties. Several have large-scale commercial application, such as the fluorinated plastic polytetrafluoroethylene ("Teflon") used in cookware. Another major application is fluorinated refrigerants. Here, traditional chlorofluorocarbons ("Freons"), which cause ozone depletion, have been largely replaced by hydrofluorocarbons.

Although it helps prevent tooth decay, fluorine is not an essential mineral for mammals. Some organofluorine compounds are synthesized in microorganisms and plants. Several fluorine compounds, as well as elemental fluorine itself, are dangerously toxic. Nevertheless, an increasing number of pharmaceuticals (about 10% of new drugs) contain fluorine.

Characteristics

Electronic structure

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. Therefore, they experience a relatively high effective nuclear charge of +7. Fluorine 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 for any other element except neon and helium. The 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 that of any other element except chlorine. Fluorine has a relatively small covalent radius, on average about 60 picometers, which slightly exceeds the numbers of neon but is surpassed by those of oxygen.

Isotopes

Main article: Isotopes of fluorine

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 decay via beta plus decay (positron emission).

To date, only one nuclear isomer, fluorine-18m, has been characterized. 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.

Physical description

Fluorine atoms form diatomic molecules that are gaseous at room temperature. Though sometimes cited as yellow-green, fluorine gas is a very pale yellow; its color can only be observed in concentrated fluorine gas kept in long glass tubes. The element has a "pungent" odor that is noticeable even in extremely low concentrations. Fluorine gas's density is 1.696 grams per liter at 100 kilopascal and 0 °C, about 1.3 times as dense as air.

Fluorine condenses to a bright yellow liquid at −188.1 °C (−306.6 °F), a comparable temperature to the boiling points of oxygen and nitrogen. Fluorine solidifies at −219.6 °C (−363.3 °F) into a cubic structure, called beta-fluorine. This phase is transparent and soft, with significant disorder of the molecules. At −227.5 °C (−377.5 °F) fluorine undergoes a solid–solid phase transition into a monoclinic 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.

Elemental fluorine

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  Three circles: left is same color as background (old paper somewhat tanned); middle is very light yellow; right is more yellow.
• Alpha-fluorine-unit-cell-B-3D-vdW.png
  Alpha-fluorine-unit-cell-B-3D-vdW.png
• Solid alpha-fluorine's crystal structure: fluorine molecules lie in shingled layers.
  Solid alpha-fluorine's crystal structure: fluorine molecules lie in shingled layers.

Chemical reactivity

Halogen bond energies (kJ/mol)

X	XX	HX	BX3	AlX3	CX4
F	159	574	645	582	456
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 element. The removal of an electron from a fluorine atom requires so much energy that no known oxidant can oxidize fluorine to any positive oxidation state.

The high direct reactivity of fluorine gas results from the relative weakness of the fluorine–fluorine bond 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, due to the weak bonding between fluorine atoms.

Reactions between fluorine and other elements are often sudden or explosive. Fluorine is so reactive that water, halogens, and most other substances, even generally nonreactive ones such as radon, burn with a bright flame in a jet of fluorine gas. It can even oxidize elemental nitrogen, which is extremely nonreactive due to its triple bonds, to give nitrogen trifluoride, though this occurs only when activated by electric discharge.

External videos
Fluorine gas, impinging on several elements, causes bright flames even without a spark.

All metals react with fluorine to form fluorides, but different conditions are required for the reaction depending on the metal. Most frequently, the metals must be in powder forms, because many metals form layers of fluoride on their surfaces that resist further oxidation. Alkali metals react with fluorine violently and form fluorides with formula MF; alkaline earth metals react at room temperature as well, but such reactions are not so exothermic. The metals ruthenium, rhodium, palladium, platinum and gold react least readily with fluorine, and are oxidized by the halogen only in atmospheres of pure fluorine at temperatures of 300–450 °C (575–850 °F). 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 or any 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 reacts with hydrogen fluoride to form argon fluorohydride at low temperatures.

Origin and occurrence

Abundance in the Solar System

Atomic number	Element	Relative amount
6	Carbon	4,800
7	Nitrogen	1,500
8	Oxygen	8,800
9	Fluorine	1
10	Neon	1,400
11	Sodium	24
12	Magnesium	430

From the perspective of cosmology, fluorine is relatively rare with 400 ppb in the universe because of its tendency to undergo nuclear fusion with hydrogen to form oxygen and helium, or with helium to become neon and hydrogen, at solar core temperatures. Most fluorine forms either in Type II supernovae when a neutrino hits an atom of neon, in asymptotic giant branch stars, or in blue Wolf-Rayet stars with masses over 40 solar masses (in which stellar winds blow 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. Three minerals exist on earth that contain enough fluorine to be mined and used as industrial resources. The most important is fluorite, which is used in smelting, construction, and the manufacture of hydrogen fluoride. Fluorapatite is mined along with other apatites for its phosphate content, and is used mostly for production of phosphate fertilizers. The hexafluorosilicates produced as by-product phosphoric acid are mostly disposed of as waste. Cryolite is the least abundant of the three and is directly used for the production of aluminium. The latter two minerals originate from meteoric water; cryolite has also been found in magmatic water. Fluorocarbon-containing chlorofluorocarbons and tetrafluoromethane have been reported in rocks, presumably having formed without action of living organisms. They are not a commercially or environmentally important source of fluorine.

Notable fluorine-containing minerals

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  pink globular mass with crystal facets
• Fluorapatite
  Fluorapatite
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  A white rock on a black background

History

Henri Moissan, the first to isolate elemental fluorine

"Fluorine" is a word that ultimately derives from the Latin verb fluo, meaning "flow." 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.

Andreas Sigismund Marggraf made the first recorded preparation of "fluoric acid" (hydrofluoric acid in modern nomenclature) in 1764, when he 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 the name 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. Greek, Russian, and several other languages use the ftor or deratives, which comes from Greek φθόριος, meaning "destructive." The new Latin name (fluorum) gave the element its current symbol F, although the symbol Fl is 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 slow because it could only be prepared electrolytically and even then under stringent conditions, since the gas reacts with most materials. The generation of elemental fluorine from hydrofluoric acid proved to be exceptionally dangerous, killing or blinding several people who attempted early experiments on this halogen. Jean Dussaud referred to these scientists as "fluorine martyrs," a term still used. In 1886, French chemist Henri Moissan reported the isolation of elemental fluorine, after almost 74 years of effort by other chemists. Moissan received the 1906 Nobel Prize in chemistry for the feat.

The two most prominent developments of organofluorine compounds are chlorofluorocarbon refrigerants such as Freon-12, and Teflon. Both were associated with the DuPont company. Chlorofluorocarbons are now being replaced by hydrofluorocarbons.

The K-25 gaseous diffusion plant in Oak Ridge, Tennessee, where uranium hexafluoride was used to separate uranium-235 for the Manhattan Project

Large-scale productions of elemental fluorine began during World War II. Germany used high-temperature electrolysis to produce tons of chlorine trifluoride, a compound planned to be used as an incendiary. The Manhattan project in the United States produced even more fluorine for use in uranium separation. Gaseous uranium hexafluoride, was used to separate uranium-235, an important nuclear explosive, from the heavier uranium-238 in centrifuges and diffusion plants. Because uranium hexafluoride releases small quantities of corrosive fluorine, the separation plants were built with special materials. All pipes were coated with nickel, which forms a protective fluoride layer on its surface after exposure to fluorine. Joints and flexible parts were fabricated from Teflon.

Production

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

Electrolytic synthesis

Fluorine cell room at F2 Chemicals Ltd, Preston, UK

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

A mixture with the approximate composition KF•2HF melts at 70 °C (158 °F) and is electrolyzed between 70 °C and 130 °C (158–266 °F). Potassium bifluoride increases the electrical conductivity of the solution and provides the bifluoride anion, which is oxidized to form fluorine at the anode. When HF is electrolyzed, hydrogen forms at the cathode and the fluoride ions remain in solution. After electrolysis, potassium fluoride remains in solution.

2 HF₂^– → H₂↑ + F₂↑ + 2 F^–

Henri Moissan first pioneered this method of electrolysis. Moissan used platinum group metal electrodes and carved fluorite containers, but the modern process uses the steel cells that act as cathodes, while blocks of carbon are used as anodes (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.

Pure fluorine gas may be stored in steel cylinders, where the inside surface becomes passivated by 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 of fluorine gas. It involved the reaction of potassium hexafluoromanganate and antimony pentafluoride at 150 °C, in an atmosphere of hydrogen fluoride:

2 K
₂MnF
₆ + 4 SbF
₅ → 4 KSbF
₆ + 2 MnF
₃ + F
₂↑

This synthetic route is a rare chemical preparation of elemental fluorine, a reaction not previously thought possible. The manganese(IV) fluoride has to be prepared by reaction with fluorine gas itself, or with krypton difluoride, which is synthesized by reaction with elemental fluorine. This reaction is therefore not an industrially viable way to produce fluorine. Fluorine can by synthesized by the reaction of hexafluoronickelate ion NiF²⁻
₆ with the fluorides of the heavier noble gases (krypton and xenon), which can only be produced by a reaction of krypton or xenon with fluorine gas.

Economic aspects

With the exception of countries with planned economics, about 17,000 tonnes of fluorine are produced per year by 11 companies in G7 countries. Fluorine is relatively inexpensive, costing about $5–8 per kilogram when sold as uranium hexafluoride or sulfur hexafluoride. Because of difficulties in storage and handling, the price of pure fluorine gas is much higher.

Compounds

Fluorine exists in the −1 oxidation state in all compounds except for elemental fluorine, 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. Higher bonding can occur, for example boron monofluoride features a triple bond. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding.

Inorganic compounds

Inorganic acids

Hydrofluoric acid dip tanks for etching of silicon wafers

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. The acid's weak acidity in water is due to the hydrogen bonding between the fluoride and hydronium ions and the decrease in entropy that occur when hydrofluoric acid dissociates. 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, hydrogen fluoride is very corrosive, attacking glass. Due to the basicity of the fluoride ion, soluble fluorides give basic water solutions.

Perfluoroacids, which are acids that contain only hydrogen, fluorine and atoms of one other element in the center of the acid's anion, are generally very strong. Fluoroantimonic acid, one such acid, is a "superacid" and the strongest acid known. It has an extremely low pK_a of −31.3 and is 20 quintillion (2×10¹⁹) 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 the antimony atom and the hydron. By occupying all antimony's valence electrons, fluorine atoms prevent the hydron from bonding to it.

Metal fluorides

Metal fluorides have similarities to other metal halides and to metal oxides, but the ionic character is stronger in metal fluorides than in the corresponding chlorides or oxides. The solubity of ionic fluorides varies greatly, but tends to decrease as the number of fluorides increases. Alkali metal fluorides often resemble the chlorides in terms of structure (all having the sodium chloride structure) and solubilities. Because the fluoride anion is highly basic, many alkali metal fluorides form bifluorides with the formula MHF₂; this is a well-known process for sodium and potassium in chemical industry. Among other monofluorides, only silver(I) and thallium(I) fluorides are well-characterized; both are very soluble, unlike other corresponding halides. Unlike the chlorides, alkaline earth metals (except for beryllium) form fluorides that are only sparingly soluble. Several other difluorides, such as those of copper(II) and nickel(II), are soluble. No trifluoride is soluble in water, but several may be soluble in other solvents.

The fluorides of later first row transition metals

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  white powder in a tube and on a spoon
• Iron trifluoride
  Iron trifluoride
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  pink powder in a tube and on a spoon
• Nickel difluoride
  Nickel difluoride
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  light blue powder in a tube and on a spoon

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 partially caused by its small radius. Only rhenium is known to form bonds with seven fluorides, forming rhenium heptafluoride, which holds the record for number of charged ligands for a metal compound. The compounds shares the pentagonal bipyramid molecular geometry with iodine heptafluoride, the only well-studied analogous nonmetal compound. Metal hexafluorides and higher fluorides are oxidants: for example, platinum hexafluoride was the first compound to oxidize molecular oxygen and xenon (see below). Metal fluorides can be volatile solids, liquids, or gases at room temperature.

Nonmetal fluorides

The nonmetal fluorides are volatile. Period 2 elements (with the exception of boron, which forms a trifluoride) form fluorides that follow the octet rule: carbon tetrafluoride, nitrogen trifluoride, and oxygen difluoride. Lower 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. Silicon tetrafluoride is a weaker acid and less thermally stable, but carbon tetrafluoride is relatively chemically stable. Among pnictogens, reactivity and acidity of pentafluorides increases down the group; bismuth is an exception, because bismuth pentafluoride is not as acidic as antimony pentafluoride 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, NF⁺
₄, 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. In halogens, unlike previous groups, not all elements form fluorides of their 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 astatine fluoride has not been produced, even though this should be possible.

Noble gas compounds

Main article: Noble gas compound

Xenon tetrafluoride crystals, 1962

The noble gases are generally non-reactive because they all have fully filled electronic shells, which are extremely stable. The ability of fluorine-containing platinum hexafluoride to react with xenon was first reported by Neil Bartlett in 1962. He called the compound he prepared xenon hexafluoroplatinate, but since then the product has been revealed to be mixture of different chemicals. 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 include 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 indicate 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^•, which, because of its metastability, is used in the argon fluoride laser. Even though the reactivity of neon is the lowest of all the elements, the element forms a metastable chemical compound, neon monofluoride, NeF^•.

Ununoctium, the last currently known group 18 element, is predicted to form ununoctium difluoride, UuoF
₂, and ununoctium tetrafluoride, UuoF
₄, which is likely to have the tetrahedral T_d configuration. However, only a few atoms of ununoctium have been synthesized, and its chemical properties have not been examined yet (as of 2011).

Comparison between the highest oxidation states of oxides and fluorides

For groups 1—6 and 13—16 the highest oxidation states of oxides and fluorides are always equal, and differences are 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 oxidation states; for example, no binary oxide is known for krypton, but krypton difluoride is well-studied. 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 well-studied.

Later transition metals

Square planar structure of mercury(IV) fluoride

With the exceptions of the +7 and +8 oxidation states, 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. Gold(V) is only known in the hexafluoroaurate(V) ion, which can be synthesized indirectly under extreme conditions, and the gold(V) fluoride, which is obtained during hexafluoroaurate(V) decomposition. The high oxidizing potential of fluorine has led to the claim of the gold(VII) existence in gold heptafluoride, but current calculations show that the claimed AuF₇ molecule was AuF₅·F₂. The great oxidizing power of fluorine is also illustrated by the fluorine-containing complexes of copper(IV), silver(IV), nickel(IV), iridium(VI), and others. It is possible that the element 113, ununtrium, will be the first element in boron group to form a species in the +5 oxidation state, the fluorine-based hexafluoroununtrate(V), UutF⁻
₆; the possibility of a +5 oxygen-based species is not known to be calculated.

Fluorine is a generally stronger oxidizer than oxygen; however, this strength does not apply for every case. Dinitrogen pentoxide, with nitrogen in the oxidation state of +5, is known; but creating nitrogen pentafluoride would need to squeeze five fluorine atoms attached to the central atom. This is hard to perform, as a nitrogen atom is smaller than most other atoms. It is not known whether the molecule is possible to produce or not, and if possible, whether it is stable or not. 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 the form of a fluoride, calculations provided by Dementyev et al. in 1997 show that the element may be oxidized up to gold(IX), in the form of the tetroxoaurlyl(IX) ion, [AuO₄]⁺, but not as a fluorine-based compound or ion. Similarly, this and another (Rother et al., 1969) calculations revealed that oxygen-based complexes that contain iridium(IX), platinum(X), and mercury(VIII) might be possible. However, these species were denied by the University of Würzburg in a 2006 paper; it expects platinum(VI), gold(V), and mercury(IV), known in binary fluorides, to be the highest for the elements. It has been shown osmium and iridium may form heptafluorides; for osmium, even an octafluoride may be possible.

Halogens and noble gases

Among halogens, chlorine and bromine form perchlorates and perbromates, both oxygen-based and with the representative halogen in +7 state; however, chlorine, unlike bromine, forms a binary heptoxide. Out of their stable fluorinated species, pentafluorides are the species in the highest oxidation state achieved; however, bromine hexafluoride, BrF₆^•, is known as well. Iodine shows the reverse picture: no heptoxide is known, unlike heptafluoride, a well-known stable compound; 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 is not known nor expected to be stable, but octafluoroxenate(VI), XeF²⁻
₈, 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

Main article: Organofluorine chemistry

Organofluorine compounds are chemical compounds that contain 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, the molecular size is not changed significantly. The range of organofluorine compounds is thus diverse; consequently, the research in the area and its uses are 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 dangerously reactive, or reaction with fluorinating reagents such as sulfur tetrafluoride.

The most industrially important compounds of fluorine include Polytetrafluoroethylene (also called PTFE or Teflon) and hydrofluorocarbons, the main properties of which are affected by the carbon–fluorine bonds in them. The slippery nature of PTFE is the result of chemical stability and repulsion of highly charged fluorine atoms in polymeric chains. Its resistance to van der Waals forces makes PTFE the only known surface to which a gecko cannot stick. Properties of the chlorofluorocarbons and hydrochlorofluorocarbons depend on the number and identity of the halogen atoms. The volatility of these compounds is lower than in most organic compounds because of the strength of the carbon–fluorine bond and carbon–chlorine bond, as well as the molecular polarity induced by the halides and the polarity of halides themselves, which cause intermolecular interactions. The large difference between chlorine and fluorine atomic radii makes chlorofluorocarbons asymmetric, 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 the released halides quench the free radicals that sustain flames.

The large inductive effect of the trifluoromethyl group results in the high acid strength of many fluorinated organic acids, which may be comparable to mineral acids. In these compounds, the cation's affinity for the acid proton is decreased by the cation's fluorine content, which increases 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.

Some organofluorine compounds: (A) fluoromethane (methyl fluoride), used in semiconductor processing; (B) dichlorodifluoromethane (R-12 refrigerant), a CFC; (C) tetrafluoroethane (R-134a), an HFC; (D) trifluoromethanesulfonic acid (triflic acid), a superacid; (E) isoflurane, an ether-derivative used as an anesthetic; (F) perfluorooctanesulfonic acid (main component of Scotchgard), a fluorosurfactant; (G) a section of polytetrafluoroethylene (Teflon) polymer; (H) fluorouracil (Adrucil), an anti-cancer drug; (I) fluoxetine (Prozac), an anti-depression drug

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 application for elemental fluorine is preparation of 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. The second largest application for fluorine gas is sulfur hexafluoride, which is used as an inert dielectric medium in high voltage switching. Sulfur hexafluoride may be produced without using fluorine gas, but the reaction between pure sulfur and pure fluorine gas is the most commonly used in industry.

Elemental fluorine is used for production of tetrafluoromethane, which is utilized for plasma etching in semiconductor manufacturing, flat panel display production, and microelectromechanical systems fabrication. These and other 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.

Isotope applications

Whole-body PET scan using fluorine-18

Natural fluorine is monoisotopic, consisting solely of fluorine-19. Fluorine compounds are highly amenable to nuclear magnetic resonance, because fluorine-19 has a nuclear spin of ½, a high nuclear magnetic moment, and a high magnetogyric ratio, which allows it to make measurements quickly, comparable with a similar effect based on hydrogen-1. Although it is not one of basic NMR active nuclei used in science and medicine, fluorine-19 is still commonly used in nuclear magnetic resonance, especially in the study of protein structures and conformational changes. Natural fluorine's monoisotopic occurrence makes it useful in uranium enrichment, because uranium hexafluoride molecules differ in mass only due to mass differences between uranium-235 and uranium-238. These mass differences are used to separate uranium-235 and uranium-238 via diffusion and gas centrifugation.

Compounds containing fluorine-18, a radioactive isotope that emits positrons, are often used in 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 by tissues with a high need for glucose, such as the brain and most types of malignant tumors. Tomography can thus be used for diagnosis, staging, and monitoring treatment of cancers, particularly in Hodgkin's disease, lung cancer, breast cancer, and many others.

Uses of compounds

Inorganic fluorides and organofluorine compounds find use in a variety of materials and chemicals, including important pharmaceuticals, agrochemicals, lubricants, and textiles.

Hydrofluoric acid and certain fluoride-containing salts are useful for etching glass, light bulbs. Laboratory-produced sodium hexafluoroaluminate, better known as synthetic cryolite, the mineral composed mostly out of this chemical, is used in the electrolysis of aluminium and its purification metallurgy to lower the melting point of aluminium oxide; the compounds also act as a powerful flux for glass.

Perfluorooctanoic acid and tetrafluoroethylene are used in water resistant coatings and in the production of low friction plastics such as PTFE, or Teflon. 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 were once heavily used in the production of haloalkanes such as chlorofluorocarbons, which are used extensively in air conditioning and in refrigeration. CFCs have been banned for these applications because they contribute to ozone destruction, and have therefore been partly replaced with hydrofluorocarbons that contain no halogen other than fluorine, such as 1,1,1,2-tetrafluoroethane and 2,3,3,3-tetrafluoropropene. Another haloalkane, bromotrifluoromethane ("Halon") is still widely used in ship and aircraft gaseous fire suppression systems. Because Halon production has been banned since 1994, systems are dependent on the pre-ban stores and on recycling.

Organofluorine applications

• 
  Teflon tape, being wrapped on a pipe thread.
• A water droplet DWR-coated surface2 edit1.jpg
  A water droplet DWR-coated surface2 edit1.jpg
• Fluorine-containing durable water repellent makes a fabric water-resistant.
  Fluorine-containing durable water repellent makes a fabric water-resistant.
• 
  a sailer in coveralls wipes down red gas bottles
• File:A Halon fire suppression system in a ship's machinery room

Biological roles

Living organisms

Gifblaar is noted for making fluoroacetate. The plant kills many cattle in southern Africa.

See also: Fluoride deficiency

Fluoride is not considered an essential mineral element for mammals and humans, though 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 been controversial since its inception in 1945. Small amounts of fluoride may be beneficial for bone strength, but this is an issue only in the formulation of artificial diets.

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

Pharmaceuticals, agricultural chemicals, and poisons

Diflunisal has a half-life of 13 hours, which is much longer than most other non-steroidal anti-inflammatory drugs.

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 activation. For example, an aromatic ring may add to prevent the metabolism of a drug, but this presents a safety problem, because enzymes in the body metabolize some aromatic compounds into poisonous epoxides. Substituting a fluorine into a para position, however, protects the aromatic ring and prevents the epoxide from being produced. Adding fluorine to biologically active organics increases their lipophilicity, because the carbon–fluorine bond is even more hydrophobic than the carbon–hydrogen bond. 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, although the potential of the fluorine to be released as a fluoride leaving group is heavily dependent on the its position in the molecule.

Of drugs that have been commercialized in the past 50 years, 5–15% contain fluorine, and the percentage of currently available fluorine-containing drugs 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, also contain fluorine. Examples include sevoflurane, desflurane, and isoflurane, which are hydrofluorocarbon derivatives.

Many SSRI antidepressants are fluorinated organics, including citalopram, escitalopram, fluoxetine, fluvoxamine, and paroxetine. Fluoroquinolones are a commonly used family of broad-spectrum antibiotics. Because biological systems do not metabolize fluorinated molecules easily, 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. Because of these, water from agricultural sites contaminates 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

Main article: Fluorine absorption dating

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 the amount of fluoride with another object found in the same area. It is important as a separation technique in intra-site chronological analysis and inter-site comparisons.

However, if no actual age of any object is known, the ages can only be expressed in terms of one of the objects being older or younger than the other. The fluctuating amount of fluoride found in groundwater means the objects being compared 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

NASA projection of stratospheric ozone, in Dobson units, if chlorofluorocarbons had not been banned.

Chlorofluorocarbons and bromofluorocarbons have 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 halocarbons containing chlorine and bromine. Hydrofluorocarbons and perfluorocarbons are greenhouse gases about 4,000 to 10,000 times as potent as carbon dioxide. Sulfur hexafluoride exhibits an even stronger effect, having 20,000 times the global warming potential of carbon dioxide.

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 unclear; 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, as it is absorbed easily via the gastrointestinal tract in rats. PFOA has been shown not to metabolize in the body, and, unlike chlorinated hydrocarbons, it is neither genotoxic nor lipophilic. It binds to serum albumin and is excreted primarily from the kidney.

Precautions

Elementary state, fluoride ion, and fluoroacetate

See also: Fluoride poisoning

U.S. hazard signs for commercially transported fluorine

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 cause ignition at higher 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 the eyes, respiratory tract, lungs, liver and kidneys. At a concentration of 100 ppm, human eyes and noses are irritated and seriously damaged.

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 fluoride. The fluoride ion is somewhat toxic, in part because of its ability to form, by equilibration, small amounts of hydrogen fluoride in water. This mobile uncharged species diffuses across cell membranes to attack intracellular calcium. The 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 with calcium and interferes with various enzymes. Excess fluoride consumption can lead to skeletal fluorosis, which currently affects millions of people.

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"). 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, which affected nearly 300 people and killed one person.

Hydrofluoric acid

Topical treatment for hydrogen fluoride burns

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 a 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

• Halogens—the periodic group to which fluorine belongs
• Caesium—the least electronegative element, fluorine's "opposite"

Bibliography

• Aigueperse, Jean; Mollard, Paul; Devilliers, Didier; Chemla, Marius; Faron, Robert; Romano, Renée; Cuer, Jean Pierre (2005). Ullmann (ed.). Encyclopedia of Industrial Chemistry. Wiley-VCH. p. 35. doi:10.1002/14356007. ISBN 9783527306732.
• Dean, John A. (1999). Lange's handbook of chemistry (Fifteenth edition). McGraw-Hill, Inc. ISBN 0070161909. {{cite book}}: Invalid |ref=harv (help)
• Greenwood, N. N.; Earnshaw, A. (1998). Chemistry of the Elements (second edition). Butterworth Heinemann. ISBN 0750633654. {{cite book}}: Invalid |ref=harv (help)
• Lewars, Errol G. (2008). Modelling Marvels. Springer. ISBN 1402069723. Retrieved 7 May 2011. {{cite book}}: Invalid |ref=harv (help)
• Lide, David R. (1992). Handbook of Chemistry and Physics (73rd edition, 1992–1993, special student edition). CRC Press. ISBN 0849305667. {{cite book}}: Invalid |ref=harv (help)
• Template:Ru icon Lidin, P. A.; Molochko, V. A.; Andreeva, L. L. (2000). Knimicheskiye svoystva neorganicheskikh veshchestv (Chemical properties of inorganic substances). Khimiya. ISBN 5-7245-1163-0. {{cite book}}: Invalid |ref=harv (help)
• Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. ISBN 9780123526519. Retrieved 3 March 2011. {{cite book}}: Invalid |ref=harv (help)
• Yaws, Carl L.; Braker, William (2001). "Fluorine". Matheson gas data book, Book 2001. McGraw-Hill Professional. ISBN 9780071358545. {{cite book}}: Invalid |ref=harv (help)

External links

• Template:PeriodicVideo