Fluorine

Fluorine
9F
- ↑ F ↓ Cl Periodic table oxygen ← fluorine → neon
Appearance
gas: very pale yellow liquid: bright yellow solid: transparent (beta), opaque (alpha) Liquid fluorine at cryogenic temperatures
General properties
Name, symbol, number	fluorine, F, 9
Pronunciation	/ˈflʊəriːn/ FLUU-reen, /ˈflʊərɪn/, /ˈflɔəriːn/
Element category	halogen
Group, period, block	17 (halogens), 2, p
Standard atomic weight	18.9984032(5)[1]
Electron configuration	[He] 2s2 2p5[2]
History
Naming	after the mineral fluorite, itself named after Latin fluo (to flow, in smelting)
Discovery	André-Marie Ampère (1810)
First isolation	Henri Moissan[2] (June 26, 1886)
Named by	Humphry Davy
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa) 1.696[3] g/L
Liquid density at b.p.	1.505[4] g·cm−3
Melting point	53.53 K, −219.62 °C, −363.32[5] °F
Boiling point	85.03 K, −188.12 °C, −306.62[5] °F
Critical point	144.4 K, 5.215[4] MPa
Heat of vaporization	6.51[3] kJ·mol−1
Molar heat capacity	(Cp) (21.1 °C) 825[4] J·mol−1·K−1 (Cv) (21.1 °C) 610[4] J·mol−1·K−1
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 38 44 50 58 69 85
Atomic properties
Oxidation states	−1 (oxidizes oxygen)
Electronegativity	3.98[2] (Pauling scale)
Ionization energies (more)	1st: 1,681[6] kJ·mol−1
	2nd: 3,374[6] kJ·mol−1
	3rd: 6,147[6] kJ·mol−1
Covalent radius	64[7] pm
Van der Waals radius	135[8] pm
Miscellanea
Crystal structure	monoclinic alpha state (low-temperature)[9]
	cubic beta state (high-temperature)
Magnetic ordering	diamagnetic, −1.2×10−4 (SI)[10][11]
Thermal conductivity	0.02591[12] W·m−1·K−1
CAS registry number	7782-41-4[2]
Most stable isotopes
Main article: Isotopes of fluorine
iso NA half-life DM DE (MeV) DP 18F trace 109.77 min+ β+ (96.9%) 0.634 18O ε (3.1%) 1.656 18O 19F 100% 19F is stable with 10 neutrons reference[13]
v t e · r

Fluorine is the chemical element with symbol F and atomic number 9. It is the lightest halogen and has a single stable isotope, fluorine-19. At standard pressure and temperature, fluorine is a pale yellow gas composed of diatomic molecules, F₂. Fluorine is the most electronegative element and is extremely reactive, requiring great care in handling. The compounds of fluorine are called fluorides.

In stars, fluorine is rare compared to other light elements. In Earth's crust, fluorine is the thirteenth-most abundant element. Fluorine's most important mineral, fluorite, was first formally described in 1529 in the context of smelting. The mineral's name derives from the Latin verb fluo, meaning "flow", because fluorite was added to metal ores to lower their melting points. Suggested as a chemical element in 1811, fluorine was named after the source mineral. The dangerous element resisted many attempts to isolate it, but in 1886, French chemist Henri Moissan succeeded. His method of electrolysis remains the industrial production method for fluorine gas. The largest use of elemental fluorine, uranium enrichment, was developed during the Manhattan Project.

Because of the difficulty in making elemental fluorine, most fluorine used in commerce is never converted to the free element. Instead, hydrofluoric acid is the key intermediate for the US$16 billion-per-year, global fluorochemical industry. The largest uses of inorganic fluorides are steel making and aluminium refining. Organofluorides tend to have high chemical and thermal stability. The largest commercial use is in refrigerant gases; even though traditional chlorofluorocarbons are banned, the replacements still contain fluorine. Polytetrafluoroethylene (Teflon) is the most important fluoropolymer and is used in electrical insulation and cookware.

While a few plants and bacteria synthesize organofluorine poisons, fluorine has no metabolic role in mammals. The fluoride ion, when directly applied to teeth, reduces decay and for this reason is used in toothpaste and water fluoridation. A growing fraction of modern pharmaceuticals contain fluorine; Lipitor and Prozac are prominent examples.

Characteristics[edit]

Phases[edit]

Fluorine forms diatomic molecules (F₂) that are gaseous at room temperature.^[14] Though sometimes cited as yellow-green, pure fluorine gas is actually a very pale yellow. The color can only be observed in concentrated fluorine gas when looking down the axis of long tubes as it appears transparent when observed from the side in normal tubes or if allowed to escape into the atmosphere.^[15] The element has a "pungent" characteristic odor that is noticeable in concentrations as low as 20 ppb.^[16]

Fluorine condenses to a bright yellow liquid at −188 °C (−307 °F),^[17] which is near the condensation temperatures of oxygen and nitrogen. It solidifies at −220 °C (−363 °F)^[17] into a cubic structure called beta-fluorine. This phase is transparent and soft with significant disorder of the molecules. At −228 °C (−378 °F) fluorine undergoes a solid–solid phase transition into a monoclinic structure called alpha-fluorine. This phase is opaque and hard with close-packed, shingled, 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's than to the other halogens'.^[9][18]

Low-temperature fluorine phases	Alpha-fluorine crystal structure

Atomic structure[edit]

A fluorine atom has nine protons and nine electrons, one fewer than neon, arranged in the electronic configuration [He] 2s² 2p⁵.^[19] Fluorine's outer electrons are relatively separate from each other, and thus they do not shield each other from the nucleus. Therefore, they experience a high effective nuclear charge. Fluorine has a relatively small covalent radius, around 60 picometers. This makes fluorine atoms similar in size to those of oxygen and neon.^{[note 1][20][21]}

Structure of the fluorine atom

Fluorine is reluctant to ionize. Instead, it exhibits a very strong preference for one more electron to achieve the extremely stable neon-like electron configuration.^[19] Fluorine's ionization energy (energy required to remove an electron to form F⁺) is higher than that of any element except neon or helium.^[22] Fluorine's electron affinity (energy released by adding an electron to form F^–) is higher than that of any element except chlorine.^[23]

Molecular structure[edit]

While an individual fluorine atom has one unpaired electron, molecular fluorine has all the electrons paired. This makes it diamagnetic (slightly repelled by magnets). In contrast, neighboring element oxygen, with two unpaired electrons per molecule, is paramagnetic (attracted to magnets).^[10][11]

The fluorine–fluorine bond of the difluorine molecule is relatively weak. The bond energy is much less than those of Cl₂ or Br₂ molecules and similar to the easily cleaved oxygen–oxygen bonds of peroxides or nitrogen–nitrogen bonds of hydrazines.^[24][25]

Chemical reactivity[edit]

Fluorine's chemistry is dominated by its strong tendency to gain an electron. It is the most electronegative element, and elemental fluorine is a strong oxidant. The removal of an electron from a fluorine atom requires so much energy that no reagents oxidize fluorine to a positive oxidation state.^[26]

External video

Bright flames during fluorine reactions

Fluorine reacting with caesium

Reactions with elemental fluorine are often sudden or explosive. Many substances that are generally regarded as unreactive—such as powdered steel, glass fragments, and asbestos fibers—are readily consumed by cold fluorine gas. Wood and even water burn with flames when subjected to a jet of fluorine, without the need for a spark.^[14][27]

Reactions of elemental fluorine with metals require different conditions that depend on the metal. Many — such as aluminium, iron, and copper — must be powdered to overcome the protective metal fluoride layers formed by passivation.^[25] The alkali metals, such as sodium, react with a bang. The alkaline earth metals, such as calcium, react not quite as dramatically. The noble metals gold, platinum, palladium, rhodium, and ruthenium react least readily, requiring pure fluorine gas at 300–450 °C (575–850 °F).^[28]

Fluorine reacts explosively with hydrogen in a manner similar to that of alkali metals.^[29] The other halogens react readily with fluorine gas^[30] as does the heavy noble gas radon.^[31] The lighter noble gases xenon and krypton can be made to react with fluorine under special conditions, while argon will undergo chemical trasformations only with hydrogen fluoride.^[32] Nitrogen, with its very stable triple bonds, requires electric discharge and high temperatures to combine with fluorine directly.^[33]

Isotopes[edit]

Fluorine occurs naturally on Earth exclusively in the form of its only stable isotope, fluorine-19,^[34] which makes the element monoisotopic and mononuclidic. These characteristics are a key in fluorine's usefulness in uranium enrichment and in nuclear magnetic resonance.

Seventeen radioisotopes have been synthesized: mass numbers 14–18 and 20–31.^[35] Fluorine-18 is the most stable radioisotope of fluorine, with a half-life of 109.77 minutes. It is the lightest unstable nuclide with equal odd numbers of protons and neutrons, 9 of each. (See also the "magic numbers" discussion of nuclide stability.)^[36]

The lightest fluorine isotopes, those with mass numbers of 14–16, decay via electron capture. ¹⁷F and ¹⁸F undergo beta plus decay (positron emission). All isotopes heavier than the stable fluorine-19 decay by beta minus mode (electron emission), while some also decay by neutron emission.^[35]

Origin and occurrence[edit]

In the universe[edit]

From the perspective of cosmology, fluorine is relatively rare with 400 ppb in the universe. Within stars, any fluorine that is created is rapidly eliminated through nuclear fusion: either with hydrogen to form oxygen and helium, or with helium to make neon and hydrogen. The presence of fluorine at all—outside of temporary existence in stars—is somewhat of a mystery because of the need to escape these fluorine-destroying reactions.^[38][39]

Three theoretical solutions to the mystery exist. In type II supernovae, atoms of neon could be hit by neutrinos during the explosion and converted to fluorine. In Wolf-Rayet stars (blue stars over 40 times heavier than the Sun), a strong solar wind could blow the fluorine out of the star before hydrogen or helium can destroy it. In asymptotic giant branch (a type of red giant) stars, pulses of fusion reactions could allow convection to lift fluorine out of the inner star. Only the red giant hypothesis has supporting evidence from observations.^[38][39]

On Earth[edit]

Fluorine is the thirteenth most common element in Earth's crust, comprising between 600 and 700 ppm of the crust by mass. Because of its reactivity, it is essentially only found in compounds. Three minerals exist that are industrially relevant sources: fluorite, fluorapatite, and cryolite.^[40][41]

• Fluorite (CaF₂), also called fluorspar, is the main source of commercial fluorine. Fluorite is a colorful mineral associated with hydrothermal deposits. It is common and found worldwide. China supplies more than half of the world's demand; Mexico is the second-largest producer. The United States produced most of the world's fluorite in the early 20th century, but its last mine, in Illinois, shut down in 1995.^{[41][42][43][44]}

• Fluorapatite (Ca₅(PO₄)₃F) is mined along with other apatites for its phosphate content and is used mostly for production of fertilizers. Most of the Earth's fluorine is bound in this mineral, but because the percentage within the mineral is low (3.5%), the fluorine is discarded as waste. Only in the United States is there significant recovery. There the hexafluorosilicates produced as byproducts are used to supply water fluoridation.^[41]

• Cryolite (Na₃AlF₆) is the least abundant of the three but is a concentrated source of fluorine. It was formerly used directly in aluminium production. However, the main commercial mine, on the west coast of Greenland, closed in 1987.^[41]

Major fluorine-containing minerals
Fluorite	Fluorapatite	Cryolite

History[edit]

Early discoveries[edit]

Steelmaking illustration, Agricola text

The word "fluorine" derives from the Latin stem of the main source mineral, fluorite, which was first mentioned in 1529 by Georgius Agricola, who described it as a flux—an additive that helps lower the melt temperature during smelting.^[45][46] Agricola, the "father of minerology", invented several hundred new terms in his Latin works describing 16th century industry. For fluorite rocks (schöne Flüsse in the German of the time), he created the Latin noun fluorés, from fluo (flow). The name for the mineral later evolved to fluorspar (still commonly used)^[47] and then to fluorite.^[44][48]

Hydrofluoric acid was known as a glass-etching agent from the 1720s, perhaps as early as 1670.^{[note 2]} Andreas Sigismund Marggraf made the first scientific report on its preparation in 1764 when he heated fluorite with sulfuric acid in glass, which was greatly corroded by the product.^[51][52] Swedish chemist Carl Wilhelm Scheele repeated this reaction in 1771,^[52][53] recognizing the product as an acid, which he called "fluss-spats-syran" (fluor-spar-acid); in English, it was known as "fluoric acid".

In 1810, French physicist André-Marie Ampère suggested that the acid was a compound of hydrogen with an unknown element, analogous to chlorine.^[54] Fluorite was then shown to be mostly composed of calcium fluoride.^[50] 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, although Greek, Russian, and some others (following Ampère's suggestion) use the name ftor or derivatives, from the Greek φθόριος (phthorios), meaning "destructive".)^[55][56] The New Latin name (fluorum) gave the element its current symbol, F, although the symbol Fl was used in early papers.^{[57][note 3]}

Isolation of the element[edit]

Progress in isolating the element was slowed by the exceptional dangers of generating fluorine: several 19th century experimenters, the "fluorine martyrs", were killed or blinded.^{[note 4]} Initial attempts to isolate the element were also hindered by material difficulties: the extreme corrosiveness and reactivity of hydrogen fluoride (and of fluorine gas) as well as problems getting a suitable conducting liquid for electrolysis.^[50][59]

Edmond Frémy thought that passing electric current through pure hydrofluoric acid (dry HF) might work. Previously, hydrogen fluoride was only available in a water solution. Frémy therefore devised a method for producing dry hydrogen fluoride by acidifying potassium bifluoride (KHF₂). Unfortunately, pure hydrogen fluoride did not pass an electric current.^[60][50][59]

French chemist Henri Moissan, formerly one of Frémy's students, continued the search. After trying many different approaches, he combined potassium bifluoride and dry hydrogen fluoride. The resultant solution conducted electricity. Moissan also constructed especially corrosion-resistant equipment: containers crafted from a mixture of platinum and iridium (more chemically resistant than pure platinum) with fluorite stoppers.^[61][59]

After 74 years of effort by many chemists, on 26 June 1886, Moissan isolated elemental fluorine.^[60][62] In 1906, two months before his death, Moissan received the Nobel Prize in chemistry.^[63][64][65] The citation:^{[59][note 5]}

...in recognition of the great services rendered by him in his investigation and isolation of the element fluorine...The whole world has admired the great experimental skill with which you have studied that savage beast among the elements.

Moissan's cell, 1887 publication	Henri Moissan, Nobel Prize photo

Later developments[edit]

During the 1930s and 1940s, the DuPont company commercialized organofluorine compounds at large scales. Following trials of chlorofluorcarbons as refrigerants by researchers at General Motors, DuPont developed large-scale production of Freon-12. Freon proved to be a marketplace hit, rapidly replacing earlier, more toxic, refrigerants and growing the overall market for kitchen refrigerators. In 1938, polytetrafluoroethylene (Teflon) was discovered by accident by a recently-hired DuPont PhD, Roy J. Plunkett. While working with a cylinder of tetrafluoroethylene, he was unable to release the gas although the weight had not changed. Scraping down the container, he found white flakes of a polymer new to the world. Tests showed the substance was resistant to corrosion from most substances and had better high temperature stability than any other plastic. In 1941, a crash program was making significant quantities of Teflon.^{[52][66][67][68]}

The Manhattan Project's uranium separation plant K-25

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.^[69] 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. Because uranium hexafluoride releases small quantities of corrosive fluorine, the separation plants were built with special materials. All pipes were coated with nickel; joints and flexible parts were fabricated from Teflon.^[66]

In 1958, a DuPont research manager in the Teflon business, Bill Gore, left the company because of its unwillingness to develop Teflon as wire-coating insulation. Gore's son Robert found a method for solving the wire-coating problem and the company W. L. Gore and Associates was born.^[70] In 1969, Robert Gore developed an expanded polytetrafluoroethylene (ePTFE) membrane which led to the large Gore-tex business in breathable rainwear. The company developed many other uses of PTFE.^[71]

In the 1970s and 1980s, concerns developed over the role chlorofluorocarbons play in damaging the ozone layer. By 1996, almost all nations had banned chlorofluorocarbon refrigerants and commercial production ceased. Fluorine continued to play a role in refrigeration though: hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) were commercialized as replacement refrigerants.^[72][73]

Industry and applications[edit]

An estimate pegged the fluorochemicals global market size at over $15 billion in 2011 and predicted 5% growth to 2018, when the market would exceed $20 billion.^[74] A different report said that, in 2016, global fluorochemicals production would be almost $20 billion on 3.5 million metric tons per year volume.^[75]

The main source of fluorine, fluorite mining, was estimated in 2003 to be a $550 million industry, extracting 4.5 million tons per year.^[52] Mined fluorite is concentrated by flotation separation into two main grades, with about equal production of each. Metspar (60–85% purity) is used almost exclusively for iron smelting. Acidspar (97%+ purity) is primarily converted to hydrofluoric acid (by reaction with sulfuric acid). The resultant HF is mostly used to produce organofluorides and synthetic cryolite.^[52][76][77]

Fluorine industry supply chain: Percentages indicate approximate mass flows of the element. Click for links to related articles.

Inorganic fluorides[edit]

About 3 kg (6.5 lb) of metspar grade fluorite are added to make each metric ton of steel. The fluoride ions from CaF₂ lower the melt's temperature and viscosity, making it runnier. Metspar is similarly used to produce cast iron and for other iron alloys.^[77][78]

Acidspar grade fluorite is added to ceramics, enamels, glass fibers, clouded glass, cement, and the outer coating of welding rods.^[77] Most acidspar is used to make hydrofluoric acid, a chemical intermediate for most fluorine-containing compounds. Significant direct uses of HF include pickling (cleaning) steel, etching glass, and cracking alkanes in the petrochemical industry.^[77]

Aluminium smelting process: cryolite (a fluoride) is required to dissolve alumnium oxide.

One third of HF (one sixth of mined fluorine) is used to make synthetic cryolite (sodium hexafluoroaluminate) and aluminium trifluoride. These compounds are used in the electrolysis of aluminium by the Hall–Héroult process. About 23 kg (51 lb) are required for every metric ton of aluminium.^[77] Although small amounts are lost over time through various side reaction, the fluorides are not reactants in the smelting process, but fluxes that lower the temperature of the melt.^[79]

Fluorosilicates are the next most significant inorganic fluorides formed from HF. Sodium fluorosilicate is used for water fluoridation, as an intermediate for synthetic cryolite and silicon tetrafluoride, and for treatment of effluents in laundries.^[80]

Other inorganic fluorides made in large quantities include cobalt difluoride (for organofluorine synthesis), nickel difluoride (electronics), lithium fluoride (a flux), sodium fluoride (water fluoridation), potassium fluoride (flux), ammonium fluoride (various uses), and magnesium fluoride (antireflection optical coatings).^[77] Sodium and potassium bifluorides are significant to the chemical industry.^[81][82]

Fluorocarbons[edit]

Making organic fluorides is the main use for hydrofluoric acid, consuming over 40% of it (over 20% of all mined fluorite). Within organofluorides, refrigerant gases are still the dominant segment, consuming about 80% of HF. Fluoropolymers are less than one quarter the size of refrigerant gases in terms of fluorine usage but are growing faster.^[77][83] Fluorosurfactants (materials like Scotchgard, used in durable water repellents) are a small segment in mass but are significant economically—over $1 billion yearly revenue.^[84]

Refrigerant gases[edit]

Traditionally chlorofluorocarbons (CFCs) were the predominant class of fluorinated organic chemical. CFCs are identified by a system of numbering (the R-number system) that explains the amount of fluorine, chlorine, carbon and hydrogen in the molecules.^[85] The DuPont brand Freon has been colloquially used for CFCs and similar halogenated molecules; brand-neutral terminology uses "R" ("refrigerant") as the prefix. Prominent CFCs included R-11 (trichlorofluoromethane), R-12 (dichlorodifluoromethane), and R-114 (1,2-dichlorotetrafluoroethane).^[77]

Production of CFCs grew strongly through the 1980s, primarily for refrigeration and air conditioning but also for propellants and solvents. Since the end use of these materials is now banned in most countries, this industry has shrunk dramatically. By the early 21st century, production of CFCs was less than 10% of the mid-1980s peak.^[77]

Hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) now serve as replacements for CFC refrigerants; few were commercially manufactured before 1990. Currently more than 90% of fluorine used for organics goes into these two classes, in about equal amounts. Prominent HCFCs include R-22 (chlorodifluoromethane) and R-141b (1,1-dichloro-1-fluoroethane). The main HFC is R-134a (1,1,1,2-tetrafluoroethane).^[77]

Fluoropolymers[edit]

The 2011 global fluoropolymer market was estimated at slightly under $6 billion revenue and predicted to grow 6.5% per year until 2016.^[86] As of about 2006–2007, fluoropolymer volume was estimated at over 180,000 metric tons per year. The corresponding revenue estimate was over $3.5 billion.^[87]

Polytetrafluoroethylene (PTFE) is 60–80% of the world's fluoropolymer production on a weight basis.^[87] The DuPont brand Teflon is sometimes used generically for the substance.^[88] The largest application for PTFE is in electrical insulation. It is an excellent dielectric. It is also used extensively in the chemical process industry where corrosion resistance is needed: in coating pipes, in tubing, and gaskets. Another major use is architectural fabric (PTFE-coated fiberglass cloth used for stadium roofs and such). The major consumer application is non-stick cookware.^[88]

Major PTFE applications
PTFE dielectric separating core and outer metal in a specialty coaxial cable	First Teflon branded frying pan, 1961	The interior of the Tokyo Dome. The roof is PTFE-coated fiberglass and air-supported.[89]

When stretched with a jerk, a PTFE film makes a fine-pored membrane: expanded PTFE (ePTFE). The term "Gore-tex" is sometimes used generically for this material, but that is a specific brand name. ePTFE is used in rainwear, protective apparel, and liquids and gas filters. PTFE can also be formed into fibers which are used in pump packing seals and bag house filters for industries with corrosive exhausts.^[88]

Other fluoropolymers tend to have similar properties to PTFE—high chemical resistance and good dielectric properties—which leads to use in the chemical process industry and electrical insulation. They are easier to work with (to form into complex shapes), but are more expensive than PTFE and have lower thermal stability. Films from two different fluoropolymers serve as glass-replacements in solar cells. Fluorinated ethylene propylene (FEP) is the second most produced fluoropolymer.^[88][90][91]

Fluorinated ionomers (polymers that include charged fragments) are expensive, chemically resistant materials used as membranes in electrochemical cells. Nafion, developed in the 1960s, was the first example and remains the most prominent material in the class. The initial application was as a fuel cell material in spacecraft. Since then, the material has been transforming the 55 million tons per year chloralkali industry; it is replacing hazardous mercury-based cells with membrane cells. Recently, the fuel cell application has reemerged with efforts to get proton exchange membrane (PEM) fuel cells into automobiles.^[92][93][94]

Fluoroelastomers are rubber-like substances that are composed of crosslinked mixtures of fluoropolymers. Viton is a prominent example. Chemical-resistant O-rings are the primary application.^[88]

Fluorine gas[edit]

For countries with available data, about 17,000 metric tons of fluorine are produced per year. Fluorine is relatively inexpensive, costing about $5–8 per kilogram ($2–4 per pound) when sold as uranium hexafluoride or sulfur hexafluoride. Because of difficulties in storage and handling, the price of fluorine gas is much higher. Processes demanding large amounts of fluorine gas generally vertically integrate and produce the gas onsite.^[95]

SF₆ transformers at a Russian railway

The largest application for elemental fluorine (up to 7,000 metric tons per year) is the preparation of uranium hexafluoride, which is used in the production of nuclear fuels. To obtain the compound, uranium dioxide is first treated with hydrofluoric acid, to produce uranium tetrafluoride. This compound is then further fluorinated by direct exposure to fluorine gas to make the hexafluoride.^[95] Fluorine's monoisotopic natural occurrence makes it useful in uranium enrichment because uranium hexafluoride molecules will differ in mass only because of mass differences between uranium-235 and uranium-238. These mass differences are used to concentrate uranium-235 by diffusion or centrifugation.^[77][95]

The second largest application for fluorine gas (about 6,000 metric tons per year) is in sulfur hexafluoride, which is used as a dielectric medium in high voltage transformers and circuit breakers. SF₆ gas has a much higher dielectric strength than air and is extremely chemically inert. Switchgear using SF₆ has no hazardous polychlorinated biphenyls (PCBs), in contrast to traditional oil-filled devices.^[96]

Several compounds made from elemental fluorine serve the electronics industry. Rhenium and tungsten hexafluorides are used for chemical vapor deposition of thin metal films onto semiconductors. Tetrafluoromethane is used for plasma etching in semiconductor and flat panel display manufacturing.^[97][98][99] Nitrogen trifluoride is used for cleaning equipment at display manufacturing plants.^[77]

Some organic fluorides are prepared from elemental fluorine rather than from HF. However, because direct fluorination is usually too hard to control, intermediate strength fluorinators are made from fluorine gas. The halogen fluorides ClF₃, BrF₃, and IF₅ provide gentler fluorination, with a series of strengths, and are easier to handle. Sulfur tetrafluoride is used for making fluorinated pharmaceuticals.^[77]

Production of fluorine gas[edit]

Industrial fluorine cells at F2 Chemicals, Preston, England.

Electrolytic synthesis[edit]

Today, elemental fluorine is produced using Moissan's process of electrolyzing potassium fluoride/hydrogen fluoride mixtures, but with an apparatus made of different materials: the steel container acts as the cathode, attracting H⁺ ions and releasing hydrogen gas; a carbon block (similar to that used in aluminum production) acts as the anode, attracting F^- ions and releasing fluorine gas. The voltage is 8–12 volts.^[77][100]

When combined, potassium fluoride and hydrogen fluoride spontaneously produce potassium bifluoride (KHF₂), which increases the conductivity of the solution. A mixture with the approximate composition KF•2HF melts at 70 °C (158 °F) and is electrolyzed at 70–130 °C (160–265 °F). Because HF alone cannot be electrolyzed, the presence of some KF is critical, even though it is not consumed in the cell and remains in solution.^{[52][101][102]}

Chemical routes[edit]

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. However, he stated in his work that the basics were known 50 years before the actual reaction.^[103] The main idea is that some metal fluoride anions do not have a neutral counterpart (or those are very unstable) and their acidifying would result in chemical oxidation, rather than formation of the expected molecules. Christe lists the following reactions as a possible way:^[104]

2 KMnO₄ + 2 KF + 10 HF + 3 H₂O₂ → 2 K₂MnF₆ + 8 H₂O + 3 O₂↑

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

Handling[edit]

Pure fluorine gas may be stored in steel cylinders, where the inside surface is passivated, as long as the temperature is kept below 200 °C (400 °F). Above that temperature, nickel is required.^[52][101] Regulator valves are made of nickel. Fluorine piping is generally made of nickel or Monel (nickel-copper alloy).^[105] Care must be taken to passivate all surfaces frequently and to exclude any water or greases. In the laboratory, fluorine gas may be used in glass tubing provided the pressure is low and moisture is excluded,^[105] but some sources recommend systems made of nickel, Monel, and PTFE.^[106]

Laboratory setup for fluorine reactions, University of Freiburg. The stainless steel enclosure on the left surrounds a small cyclinder of fluorine. The line's piping, fittings and valves are Monel. The rig is in a fume hood; many controls are external, allowing reactions with the door closed.^[107]

Biological aspects[edit]

South Africa's gifblaar is one of the few organisms that makes fluorine compounds.

Fluoride is not considered an essential mineral element for mammals or humans.^[108] Small amounts of fluoride may be beneficial for bone strength, but fluoride deficiency is an issue only in the formulation of artificial diets.^[109]

Biologically synthesized organofluorines have been found in microorganisms and plants,^[110] but not in animals.^[111] The most common example is fluoroacetate, with an active poison molecule identical to commercial "1080". It is used as a defense against herbivores by at least 40 green plants in Australia, Brazil, and Africa.^[112] Other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate.^[111]

The enzyme adenosyl-fluoride synthase, which makes the carbon–fluorine bond, was isolated in bacteria in 2002. The discovery was touted as possibly leading to biological routes for organofluorine synthesis.^[113]

Dental care[edit]

Since the mid-20th century, it has been understood (from population studies) that fluoride reduces tooth decay, somehow. Initially, researchers hypothesized that fluoride helped by converting tooth enamel from the mineral hydroxyapatite to the more acid-resistant mineral fluorapatite. However, recent studies showed no difference in the frequency of caries (cavities) amongst teeth that were pre-fluoridated to different degrees. Current thinking is that fluoride prevent cavities primarily by helping teeth that are in the very early stages of tooth decay to regrow tooth enamel. In any case, it is only the fluoride that is directly present in the mouth (topical treatment) that prevents cavities. Fluoride ions that are swallowed do not benefit the teeth.^[114]

Topical fluoride treatment in Panama

Water fluoridation is the controlled addition of fluoride to a public water supply to reduce tooth decay.^[115] Its use began in the 1940s, following studies of children in a region where water was naturally fluoridated. It is now used for about two thirds of the U.S. population on public water systems^[116] and for about 6% of people worldwide.^[117] Although the best available evidence shows no association with adverse effects other than dental fluorosis, most of which is mild,^[118] water fluoridation has been contentious for ethical, safety, and efficacy reasons.^[117] Opposition to water fluoridation exists despite its support by public health organizations.^[119] The benefits of water fluoridation have lessened recently—presumably because of the availability of fluoride in other forms—but are still measurable, particularly for low income groups.^[120] Reviews of the scholarly literature in 2000 and 2007 showed significant reduction of cavities in children associated with water fluoridatation.^[121]

Toothpaste may contain fluorine in the form of sodium fluoride, tin difluoride, or (most commonly) sodium monofluorophosphate. The first fluoride toothpaste was introduced in 1955 in the United States. Now, almost all toothpaste in developed countries is fluoridated.^[120] Fluoride may also be applied to teeth in gels, foams, or varnishes. It is also contained in prescription and non-prescription mouthwashes.^[122]

Pharmaceuticals[edit]

About 20% of modern pharmaceuticals contain fluorine, including commercially significant drugs in many different pharmaceutical classes.^[123] One of these, the cholesterol-reducer atorvastatin (Lipitor), was the number one money-making drug for nearly a decade. The branded asthma medication Serevent (Advair), a top-ten revenue drug as of the mid-2000s, also contains a fluorinated molecule: fluticasone.^[124][125]

Even a single atom of fluorine, added to a drug molecule, can greatly change its chemical properties and thus how it interacts with the body. For instance, because of the considerable stability of the carbon-fluorine bond, many drugs are fluorinated to delay their metabolism and elimination by the body. This allows longer times between doses.^[126] Also, adding fluorine to organics increases their lipophilicity (ability to dissolve in fats) because the carbon–fluorine bond is even more hydrophobic than the carbon–hydrogen bond. This effect often increases a drug's bioavailability because of increased cell membrane penetration.^[125]

Prozac (Fluoxetine) capsules

Many modern antidepressants are fluorinated molecules that selectively limit the body's binding of serotonin (low serotonin availability in brain cells is a cause of depression). Prior to the 1980s, traditional antidepressants, such as the tricyclics, altered not only serotonin uptake but also affected several other neurotransmitters. This non-selective interaction caused many side effects. One of the first drugs to alter only serotonin uptake—and be free of most side effects of previous drugs—was fluorine-containing Prozac (fluoxetine). It became the best-selling antidepressant and prompted the popular book Listening to Prozac. Some other selective serotonin reuptake inhibitor (SSRI) antidepressants that are fluorinated are Celexa (and its isomer Lexapro), Luvox, and Paxil.^[127][128]

Quinolones are artificial compounds that are broad-spectrum antimicrobials. Most of the recent quinolone drugs (and those in current common use) are fluorinated to make the drugs more powerful. Prominent examples of such fluoroquinolones include ciprofloxacin (Cipro) and levofloxacin (Levaquin). The latter was the highest selling U.S. antibiotic in 2010.^{[129][130][131][132]}

Fluorine also finds use in many steroidal drugs.^[133] Fludrocortisone (Florinef) is a synthetic mineralocorticoid (a compound used to retain sodium and water and thus raise blood pressure).^[134] Dexamethasone (Decadron) and triamcinolone (Kenalog) are potent synthetic glucocorticoids (anti-inflammatories).^[134]

Several inhaled anesthetics, including the most common ones, are heavily fluorinated. The first fluorinated anesthetic, halothane, proved to be much safer (neither explosive nor flammable) and longer-lasting than those previously used. Modern fluorinated anesthetics are longer-lasting still and almost insoluble in blood, which accelerates the awakening. Examples include sevoflurane, desflurane, enflurane, and isoflurane, all fluorinated ethers.^[135][136]

Agrichemicals and poisons[edit]

Sign warning of poisonous sodium fluoroacetate baits

An estimated 30% of agrichemical compounds contain fluorine.^[137] Most of them are poisons, but a few stimulate growth instead.

Sodium monofluoroacetate, the same acid as in vinegar but with one hydrogen changed out for fluorine, is a powerful poison. It was first synthesized in the late 19th century and was recognized as an insecticide in the early 20th century. Later, it became widely used to control mammalian pests (e.g. rats). The name "1080" refers to the catalogue number of the poison, which became its brand name. The compound is now banned in the European community and the United States^{[note 6]} but is still used in Australia and some other countries. Fluoroacetate deprives cells of energy by replacing acetate in the Krebs cycle, halting a key part of cell metabolism.^[112][138] Several insecticides contain sodium fluoride, which is much less toxic than fluoroacetate.^[139]

Trifluralin was once a very important weedkiller; for example, in 1998 over a half of U.S. cotton field area was coated with the chemical.^[140] However, its suspected carcinogenic properties caused some Northern European countries to ban it in 1993.^[141] Although currently banned throughout the whole European Union, this decision has recently been contested.^[142]

The currently used agrichemicals utilize another tactic: instead of directly poisoning organisms by interfering with metabolism, they transform the metabolic processes so that the organism itself produces poisonous compounds. For example, insects fed 29-fluorostigmasterol (a fluorinated steroid) produce fluoroacetates. Because vertebrates do not break down the fluorosteroid, the insecticide is more safe than 1080. If a fluorine is transferred to a body cell, it blocks metabolism at the position occupied.^[143]

Use of fluorine agrichemicals may increase if the cost of synthesis decreases and if green chemistry accentuates the differences between these agrichemicals and other less-environmentally-friendly alternatives.^[144]

Scanning[edit]

Whole-body PET scan using fluorine-18

Compounds containing fluorine-18, a radioactive isotope that emits positrons, are often used in positron emission tomography (PET) scanning, because the isotope's half-life of about 110 minutes is long by positron-emitter standards. One such radiopharmaceutical is 2-deoxy-2-(¹⁸F)fluoro-D-glucose (generically referred to as fludeoxyglucose), commonly abbreviated as ¹⁸F-FDG, or simply FDG.^[145] In PET imaging, FDG can be used for assessing glucose metabolism in the brain and for imaging cancer tumors. After injection into the blood, FDG is taken up by "FDG-avid" tissues with a high need for glucose, such as the brain and most types of malignant tumors.^[146] Tomography, often assisted by a computer to form a PET/CT (CT stands for "computer tomography") machine, can then be used to diagnose or monitor treatment of cancers; especially Hodgkin's lymphoma, lung cancer, and breast cancer.^[147]

Natural fluorine is monoisotopic, consisting solely of fluorine-19. Fluorine compounds are highly amenable to nuclear magnetic resonance (NMR), because fluorine-19 has a nuclear spin of ½, a high nuclear magnetic moment, and a high magnetogyric ratio. Fluorine compounds typically have a fast NMR relaxation, which enables the use of fast averaging to obtain a signal-to-noise ratio similar to hydrogen-1 NMR spectra.^[148] Fluorine-19 is commonly used in NMR study of metabolism, protein structures and conformational changes.^[149] In addition, inert fluorinated gases have the potential to be a cheap and efficient tool for imaging lung ventilation.^[150]

Hazards[edit]

Fluorine gas[edit]

Elemental fluorine is highly toxic. Above a concentration of 25 ppm, fluorine causes significant irritation while attacking the eyes, respiratory tract, lungs, liver and kidneys. At a concentration of 100 ppm, human eyes and noses are seriously damaged.^[151]

U.S. hazard signs for commercially transported fluorine^[152]

Hydrofluoric acid[edit]

Hydrogen fluoride is a toxic gas. Its water solution, hydrofluoric acid, is a contact poison and must be handled with extreme care, far beyond that accorded to other mineral acids. From 1984-1994, at least 9 U.S. workers died from accidents with HF.

HF burns

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 the skin or eyes or when inhaled or swallowed. Once in the blood, hydrogen fluoride reacts with calcium and magnesium, resulting in electrolyte imbalance (potentially hypocalcemia). The consequent cardiac arrhythmia may be fatal.^[153] Formation of insoluble calcium fluoride also causes strong pain.^[154] Burns with areas larger than 160 cm² (25 in²), about the size of a man's hand, can cause serious systemic toxicity.^[155]

Symptoms of exposure to hydrofluoric acid may not be immediately evident, with 8-hour delay for 50% HF and up to 24-hour for lower concentrations. Hydrogen fluoride interferes with nerve function, meaning that burns may not initially be painful. Accidental exposures can go unnoticed, delaying treatment and increasing injury.

If the burn has been initially noticed, then HF should be washed off with a forceful stream of water for ten to fifteen minutes, to prevent its further penetration into the body. Clothing used by the person burned may also exhibit danger.^[156] Hydrofluoric acid exposure is often treated with calcium gluconate, a source of Ca²⁺ that binds with the fluoride ions. Skin burns can be treated with a water wash and 2.5% calcium gluconate gel^[157][158] or special rinsing solutions.^[159]

Because HF is absorbed, medical treatment is necessary. Calcium gluconate may be injected or administered by IV. Use of calcium chloride is contra-indicated and may lead to severe complications. Sometimes surgical excision of tissue or amputation may be required.^[155][160]

Fluoride ion[edit]

Geographical areas with groundwater having over 1.5 mg/L of naturally occurring fluoride, which is above recommended levels.^[121]

Soluble fluorides are moderately toxic. For sodium fluoride, the lethal dose for adults is 5–10 g, which is equivalent to 32–64 mg of elemental fluoride per kilogram of body weight.^[161] The dose that may lead to adverse health effects is about one fifth the lethal dose.^[162] Chronic excess fluoride consumption can lead to skeletal fluorosis, a disease of the bones that affects millions in Asia and Africa.^[162][163]

The fluoride ion is readily absorbed by the stomach and intestines. Ingested fluoride forms hydrofluoric acid in the stomach. In this form, fluoride crosses cell membranes and then binds with calcium and interferes with various enzymes.

Fluoride is excreted through urine. Fluoride exposure limits are based on urine testing which has determined the human body's capacity for ridding itself of fluoride.^[162][164]

Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluoride,^[165] Currently, most calls to poison control centers for possible fluoride poisoning come from the ingestion of fluoride-containing toothpaste.^[162] Malfunction of water fluoridation equipment has occurred several times, including an Alaskan incident which sickened nearly 300 people and killed one.^[166]

Environmental concerns[edit]

Atmosphere[edit]

Because they deplete the ozone layer, Chlorofluorocarbons (CFCs) and bromofluorocarbons (BFCs) have been strictly regulated via a series of international agreements called the Montreal Protocol. It is the chlorine and bromine from these molecules that cause harm, not fluorine. Because of the inherent stability of these fully halogenated molecules (which makes them so nonflammable and useful), they are able to reach the upper reaches of the atmosphere before decomposing. Then at those altitudes, they release chlorine and bromine to attack the ozone.^[167] Predictions are that generations will be required, even after the CFC ban, for these molecules to leave the atmosphere and for the ozone layer to recover. Early indications are that the CFC ban is working—ozone depletion has stopped and recovery is underway.^[168][169]

NASA projection of stratospheric ozone over North America if CFCs had not been banned

Hydrochlorofluorocarbons (HCFCs) are current replacements for CFCs, with about one tenth the ozone damaging potential (ODP).^[170] They were originally scheduled for elimination by 2030 in developed nations (2040 in undeveloped). In 2003, the U.S. Environmental Protection Agency prohibited production of one HCFC and capped the production of the two others.^[171] In 2007, a new treaty was signed by almost all nations to move HCFC phaseout up to 2020 because HFCs, which have no chlorine and thus zero ODP, are available.^[172]

Fluorocarbon gases of all sorts (CFCs, HFCs, etc.) are greenhouse gases about 4,000 to 10,000 times as potent as carbon dioxide. Sulfur hexafluoride exhibits an even stronger effect, about 20,000 times the global warming potential of carbon dioxide.^[173][174]

Biopersistance[edit]

Because of the strength of the carbon–fluorine bond, organofluorines endure in the environment. Perfluorooctanoic acid (PFOA) and perfluorooctanesulfonic acid (PFOS), used in waterproofing sprays, are persistent global contaminants. Trace quantities of these substances have been detected worldwide, from polar bears in the Arctic to the global human population. PFOA has been detected in breast milk and the blood of newborns. One study indicates that PFOS levels in wildlife are starting to go down because of reducing production.^[175][176]

In the body, PFOA binds to a protein, serum albumin. PFOA's tissue distribution in humans is unknown, but studies in rats suggest it is likely to be present primarily in the liver, kidney, and blood. PFOA is not metabolized by the body, but is excreted by the kidneys.^[175][176]

The potential health effects of PFOA are unclear. Unlike chlorinated hydrocarbons, PFOA is not lipophilic (stored in fat), nor is it genotoxic (damaging genes). While both PFOA and PFOS cause cancer in high quantities in animals, studies on exposed humans have not been able to prove an impact at current exposures. Bottlenose dolphins have some of the highest PFOS concentrations of any wildlife studied; one study suggests an impact on their immune systems.^[175][176]

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.^[177] Fluorine-containing agrichemicals are measurable in farmland runoff and nearby rivers.^[178]

Compounds[edit]

Fluorine's common oxidation state is −1.^{[note 7]} With other atoms, fluorine forms either ionic bonds or polar covalent bonds. Covalent bonds involving fluorine atoms are almost always single bonds.^{[note 8][181]} Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding.^[182] Fluorine has a rich chemistry including compounds formed with hydrogen, metals, main group nonmetals, and even noble gases, as well as a diverse set of organic compounds.^{[note 9][183]}

Hydrogen fluoride[edit]

Boiling points of the hydrogen halides (blue) and hydrogen chalcogenides (red): HF and H₂O break trends.

Fluorine combines with hydrogen to make a compound called hydrogen fluoride (HF) or, in the context of water solutions, hydrofluoric acid. HF molecules cluster together weakly via hydrogen bonds. Because of this, hydrogen fluoride behaves more like water than like HCl.^{[184][185][186]} Hydrogen fluoride boils at a much higher temperature than the heavier hydrogen halides. HF is also fully miscible with water (dissolves in any proportion), unlike HCl, HBr, or HI.^[187]

In water solution, hydrogen fluoride is a weak acid, having lower dissociation constant compared to the other hydrohalic acids.^{[188][note 10]} Hydrofluoric acid is non-ideal: instead of having a constant acid dissociation constant (PKa), HF's inherent acidity increases at higher concentrations through a phenomenon called homoconjugation.^[190][191] Although hydrofluoric acid is weak, it is very corrosive, even attacking glass.^[190]

Metal fluorides[edit]

Metal fluorides share some similarities with other metal halides but are more ionic, often more like oxides in their bonding and crystal structures.^[192] However, there is a trend of the binary metal fluorides becoming more covalently-bonded and volatile as the amount of fluorination increases. For example, sodium monofluoride boils at 1704 °C (3099 °F),^[193] while rhenium heptafluoride has a boiling point below water's.^[194] Similar to fluorine itself, the higher metal fluorides are very reactive. Platinum hexafluoride was the first compound to oxidize molecular oxygen^[195] and xenon.^[196] The solubility of fluorides varies greatly but tends to decrease with more fluorine.^[197]

The alkali metals form monofluorides that, like the alkali metal chlorides, are very ionic and soluble. They have the same atomic arrangement—the rock salt crystal structure—as sodium chloride.^[198][81] Other metal monofluorides show clear difference versus their chlorides. For example, thallium^[199] and silver^[57] monofluorides are soluble, while the chlorides are not.

The difluorides of the alkaline earths are also very ionic, but are generally very insoluble. In contrast, the alkali earth dichlorides are fairly soluble.^[57] Beryllium difluoride differs from those of the other alkaline earths: the bonding has some covalent character, the structures are similar to SiO₂ (quartz), and the compound is water soluble.^[200]

Many metals form trifluorides, such as iron, bismuth, the rare earths, some actinides, and the metals in the aluminium and scandium columns of the periodic table.^{[citation needed]} Gold forms a trifluoride that has a polymer structure, with helical chains.^[201] In contrast, gold trichloride is molecular.^[201]

The tetrafluorides show a mixture of ionic and covalent bonding. Zirconium^[202] and hafnium,^{[citation needed]} along with several actinides,^[203] form tetrafluorides with an ionic structure based on 8-coordinate metal cations.^[204] Melting points are around 1000 °C (1850 °F).^{[note 11]} On the other hand, the tetrafluorides of titanium,^[207], vanadium,^[208] niobium^[209] are polymeric. The first melts at 284 °C (543 °F);^[210] the latter two decompose at around 350 °C (650 °F).^[211]

The pentafluorides are even more covalently bonded, forming low dimensionality polymers or oligomeric molecules.^[212] Bismuth and uranium pentafluorides share the same structure: straight chains of connected octahedra.^{[213][214][215]} Gold's pentafluoride is a dimer (doubled molecule), Au₂F₁₀.^[212] Niobium and tantalum pentafluorides are tetrameric molecules (fourfold clusters).^[212] The former melts at 80 °C (176 °F),^[216] the latter at 95 °C (203 °F).^[217]

There are thirteen metal hexafluorides.^{[note 12]} All are octahedral molecules.^{[citation needed]} At room temperature tungsten hexafluoride is a gas.^[206] Molybdenum hexafluoride and rhenium hexafluoride are liquids.^[218] The rest are volatile solids.^{[citation needed]}

The only definite metal heptafluoride, rhenium heptafluoride, is a low-melting molecular solid. Its structure is pentagonal bipyramidal.^{[192][note 13]}

Progression of structure type with metal charge in the metal fluorides
Sodium fluoride, ionic	Bismuth pentafluoride, polymeric	Rhenium heptafluoride, molecular

Nonmetal fluorides[edit]

The binary fluorides of the main group nonmetals and metalloids are generally volatile, covalently bonded molecules. The fluorides of those elements in the second row of the period table (carbon, nitrogen, and oxygen) follow the octet rule. Boron is an exception—its fluoride has an incomplete octet (less bonds than normal). Nonmetals from the third row of the period table and below form fluorides which are hypervalent (more bonds than normal).^[220] The discussion is organized by periodic table column.

Boron trifluoride is a planar molecule. It is a weak Lewis acid and readily accepts a Lewis base, forming adducts (combinations). The lone pair of a Lewis base, such as ammonia, allows the boron to complete its octet. With another fluoride ion (F^- is a Lewis base), BF₃ reacts to form the relatively unreactive BF₄^- anion.^[221]

Silicon tetrafluoride and germanium tetrafluoride adopt a tetrahedral structure^[222] and are Lewis acids.^[223][224] Germanium forms a difluoride, less stable than its tetrafluoride.^[225]

The pnictogens (nitrogen's column) form trifluorides that are weak Lewis bases and are more reactive as the pnictogen becomes heavier.^[226] In comparison, the pnictogen pentafluorides are much more reactive,^[227][228] with antimony pentafluoride being the strongest Lewis acid of all charge-neutral compounds.^[228] PF₅ and AsF₅ are trigonal bipramidal in all phases, as is gaseous SbF₅. Liquid and solid SbF₅ have more complicated structures with octahedral antimony and bridging fluorines.^[229] Nitrogen is different than other pnictogens as it forms a trifluoride that is stable against hydrolysis and is not a Lewis base.^[226] Also, nitrogen pentafluoride does not exist.^{[citation needed]}

The chalcogens (oxygen's column) form a variety of fluorides. Unstable difluorides are known for oxygen (the only compound where oxygen is at oxidation state +2) as well as sulfur and selenium.^{[citation needed]} Sulfur and selenium tetrafluorides are molecular while TeF₄ is a polymer. The tetrafluorides are thermally unstable and hydrolyze.^[230] The hexafluorides are the result of direct fluorination of sulfur, selenium, and tellurium, while other hexahalides of the elements do not even exist. SF₆ is extremely inert, while SeF₆ and TeF₆ show increasingly higher reactivity.^[230]

The well-characterized heavier halogens (chlorine, bromine, and iodine) all form mono-, tri-, and pentafluorides: XF, XF₃, and XF₅. Of the neutral +7 species, only iodine heptafluoride is known.^[231] The corresponding cations ClF₆⁺ and BrF₆⁺, are known and are extremely strong oxidizers.^[232] For the radioactive astatine only the non-volatile astatine monofluoride has been studied,^[233] but its existence is debated.^[234] Many of the halogen fluorides are powerful fluorinators (sources of fluorine atoms). ClF₃ readily fluorinates asbestos and refractory oxides, and industrial use require precautions similar to those for fluorine gas.^[235][236]

Notable nonmetal fluorides
Boron trifluoride (BF3)	Silicon tetrafluoride (SiF4)	Antimony pentafluoride (SbF5)	Sulfur hexafluoride (SF6)	Chlorine trifluoride (ClF3)

Noble gas compounds[edit]

Xenon tetrafluoride, 1962

The noble gases are generally non-reactive because they have complete electronic shells. Until the 1960s, no chemical bond with a noble gas was known. In 1962, Neil Bartlett reported the first chemical compound of xenon, xenon hexafluoroplatinate.^[237] Later in 1962, xenon was reported to react directly with fluorine to form the di- and tetrafluorides. Since then, xenon hexafluoride, various oxyfluorides, and their derivatives have been prepared.^[238][239]

Krypton, xenon's lighter homolog, also forms difluoride and a few more complicated fluorine-containing compounds.^[240] The possibility of a tetrafluoride^[241] or a hexafluoride has been debated.^[242]

Radon, xenon's heavier homolog has been shown to readily react with fluorine to form a solid compound. It is generally thought to be radon difluoride, but the exact formula is not known.^[233] If radon were not so radioactive and difficult to collect, its chemistry might be at least as extensive as xenon's.^[243]

The lightest noble gases do not form stable binary fluorides. Argon, however, reacts in extreme conditions with hydrogen fluoride to form argon fluorohydride.^[244] Helium and neon do not form any stable chemical compounds at all, but helium fluorohydride has been observed and it is unstable in gas phase, but it may be stable under enormous pressure.^[245] Neon is considered less reactive than helium and is not expected to form a stable compound capable of synthesis.^[246]

Organic compounds[edit]

The carbon–fluorine chemical bond is the strongest bond in organic chemistry.^[247] This C–F bond strength, along with the low polarizability of molecules containing C-F, makes organofluorines very stable.^[248] Fluorinated organics have similar sizes to corresponding unfluorinated molecules because of the small van der Waals radius of fluorine.^[248]

The range of organofluorine compounds is diverse, reflecting the inherent complexity of organic chemistry. A vast number of small molecules can exist with varying amounts of fluorine substitution, as well as many polymers. Research in particular areas is driven by the commercial value of applications.^[66]

Small molecules[edit]

Partially fluorinated alkanes are hydrofluorocarbons (HFCs). Monofluoroalkanes (alkanes with one hydrogen replaced with fluorine) have properties similar to unfluorinated alkanes. They are soluble in many nonpolar solvents and have some chemical and thermal instability. As more fluorines are substituted for hydrogens, the properties change. Solubility in hydrocarbons decreases and stability increases. Also, melting and boiling points decrease, while density goes up.^[249]

When all hydrogens are replaced with fluorines to make perfluorocarbons (the "per" means maximum), a great difference is revealed. Such compounds are extremely stable, and only sodium in liquid ammonia attacks them at standard conditions. They are also very insoluble, with few organic solvents capable of dissolving them.^[249]

Perfluorodecalin (trans isomer shown) is a liquid at room temperature. It boils at 142 °C (288 °F).[250] Its hydrocarbon analog, decalin, has a higher boiling point: 187 °C (369 °F).[251]	Demonstration of perfluorocarbon density: water with blue food dye (upper layer), perfluoroheptane (lower layer). The goldfish cannot swim down into the fluorocarbon. The crab floats at the liquid boundary. Animals were rescued after the photo.

Perfluorinated compounds, hydrocarbons that are fully fluorinated except for one functional group have a combination of properties.^[252] (See "Perfluorinated compound".) They exhibit many perfluorocarbon properties (e.g. inertness, stability, non-wetting by water and oils, slipperiness).^[253] However, the functional group is available for reactions or may make the molecule behave as a surfactant.^{[citation needed]} If a perfluorinated compound has a fluorinated tail but also a few non-fluorinated carbons (typically two) near the functional group, it is called a fluorotelomer. Industrially, such compounds are treated as perfluorinated.^[253]

Fluorinating an organic acid raises its acidity. For example, acetic acid and its mono-, di-, and trifluoroacetic derivatives show a trend of lowering pK_a: 4.74, 2.66, 1.24, 0.23 (thus increasing acidity).^[254] This happens because of fluorine's inductive effect (it stabilizes anions by spreading negative charge, allowing the hydrogen to be released).^[255] Similarly, the acidity is greatly increased for other perfluorocarboxyl acids, as well as the amines (which are not acids but become less basic if fluorinated).^[248] The perfluoroalkanesulfonic acids are also very notable for their acidity. Trifluoromethanesulfonic acid, is comparable to strong mineral acids.^[256][256]

Polymers[edit]

The complex unit structure of Nafion

As with small molecules, replacing hydrogen with fluorine in a polymer increases chemical stability and reduces flammability. Melting points are typically much higher than in the corresponding hydrocarbon polymers.^[257]

The simplest fluoroplastic is polytetrafluoroethylene (PTFE, DuPont brand Teflon), which is a simple linear chain polymer with the repeating structural unit: –CF₂–. It has no hydrogens and is the perfluoro analog of polyethylene (structural unit: –CH₂–). PTFE has high chemical and thermal stability, as expected for a perfluorocarbon, much stronger than polyethylene. However, its very high melting point makes it difficult to fashion into parts.^[258]

Various PTFE derivatives have lower maximum usage temperatures but have the benefit of being more melt-processable. FEP (fluorinated ethylene propylene is structurally similar to PTFE but has some fluorines replaced with the –CF₃ groups). PFA (perfluoroalkoxy has some fluorines replaced with –OCF₃).^[258]

There are other fluoroplastics that are not perfluorinated (contain some C-H). Polyvinylidene fluoride (PVDF, structural unit: –CF₂CH₂–), is an analog of PTFE with half the fluorines. PVF (polyvinyl fluoride, structural unit: –CH₂CHF–) contains one one-fourth the fluorines of PTFE. Despite this, it still has many properties of more fluorinated compounds.^[259]

Nafion is a structurally complicated polymer. It has a PTFE-like backbone, but also contains side chains of perfluoro ether that end in sulfonic acid (–SO₂OH) groups. Because of the polar side chains, it is very hydroscopic. With the addition of water and cations such as Na⁺, Nafion can be made into an ionic conductor.^[260][261]

See also[edit]

• Industrial gas

Notes[edit]

Periodic table colored to show how elements are treated in this article: dark gray elements are metals, green ones are nonmetals, blue ones are the noble gases, the purple one is hydrogen, the yellow one is carbon, and the light gray ones are those with indeterminate properties.

1. ^ Exact comparison of the sizes of fluorine, oxygen and neon atoms is not possible because of conflicting estimates from different sources.
2. ^ See the differing analysis of Partington^[49] and Weeks.^[50]
3. ^ Since 2012, the symbol Fl has been used for flerovium (element 114), a man-made transuranic element.^[58]
4. ^ The injured included Davy, Joseph Louis Gay-Lussac, Louis Jacques Thénard, and Irish chemists Thomas and George Knox. Belgian chemist Paulin Louyet and French chemist Jerome Nickles died. Moissan also experienced HF poisoning.^[50][59]
5. ^ Moissan's Nobel also honored his invention of the electric arc furnace
6. ^ A minor allowed use is in collars of sheep and cattle to kill predators like coyotes.
7. ^ It differs from this value in elemental fluorine, where the atoms are bonded to each other and thus at oxidation state 0. The very unstable anions F₂^- and F₃^- with intermediate oxidation states exist at very low temperatues, decomposing at around 40 K.^[179] Also, the F₄⁺ cation and a few related species have been predicted to be stable.^[180]
8. ^ Boron monofluoride and nitrogen monofluoride are metastable compounds with higher than single bonded fluorine. Boron monofluoride is stable as a ligand in some metal complexes.
9. ^ In this article, metalloids are lumped with the definite main group nonmetals because the fluoride chemistry is similar. The noble gases are treated separately. Hydrogen is discussed in the Hydrogen fluoride section; carbon in the Organic compounds section. P-block period 7 elements have not been studied and thus are not included. This is illustrated by the image to the right: the dark gray elements are metals, the green ones are nonmetals, the light blue ones are the noble gases, the purple one is hydrogen, the yellow one is carbon, and the light gray elements have unknown properties.
10. ^ For more detail, see explanation by Jim Clark.^[189]
11. ^ ZrF₄ melts at 932 °C (1764 °F).^[205] HfF₄ sublimes at 968 °C.^{[citation needed]} UF₄ melts at 1036 °C (1897 °F).^[206]
12. ^ IrF₆, MoF₆, OsF₆, NpF₆, PoF₆, PuF₆, PtF₆, ReF₆, RhF₆, RuF₆, TcF₆, UF₆, and WF₆
13. ^ In 1966, osmium heptafluoride was reported from daunting synthesis conditions (400 atmospheres of fluorine gas at 600 °C). The product was not structurally characterized. In 2006, the synthesis was re-attempted but only OsF₆ was observed.^[219]

Citations[edit]

Indexed references[edit]

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