Fluorine

Fluorine is the chemical element with symbol F and atomic number 9. At room temperature, the element is a pale yellow gas composed of diatomic molecules, F
₂. Fluorine is the lightest halogen and the most electronegative element. It requires great care in handling as it is extremely reactive and poisonous.

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	common: −1
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 24th most abundant element in the universe and the 13th most abundant within the Earth's crust. The main source mineral, fluorite, was first scientifically described in 1529. At that time, the Latin verb fluo, meaning "flow", became associated with fluorite rocks because they were added to metal ores to lower their melting points during smelting. First suggested as a chemical element in 1811, fluorine proved difficult and dangerous to separate from its compounds. In 1886, French chemist Henri Moissan succeeded in isolating elemental fluorine using low temperature electrolysis. In the 1940s, the largest current end use of free fluorine, uranium enrichment, began during World War II's Manhattan Project.

Compounds of fluorine are called fluorides, and global fluorochemical sales are over US$15 billion per year. Because of the expense of refining the pure element, 99% of commercially used fluorine remains in compound form throughout its processing. About half of all mined fluorite is used directly in steel-making. The other half is converted to hydrogen fluoride, a dangerous acid which is the precursor to many fluorochemicals. About one third of created hydrogen fluoride ends up in synthetic cryolite, an inorganic material critical to aluminium refining. Even more hydrogen fluoride is used to make organic fluorides. These compounds have very high chemical and thermal stability; their largest market segment is in refrigerant gases ("Freon"). (Even though traditional chlorofluorocarbons are now mostly prohibited, the replacement molecules still contain fluorine.) The predominant fluoropolymer is polytetrafluoroethylene (Teflon), which 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; for this reason, it is used in toothpaste and water fluoridation. A significant fraction of modern pharmaceuticals—such as atorvastatin (Lipitor) and fluoxetine (Prozac)—contain fluorine.

History

Main article: History of fluorine

Early discoveries

Steelmaking illustration, Agricola text

The word "fluorine" derives from the Latin stem of the main source mineral, fluorite. The stone was described in 1529 by Georgius Agricola, who related its use as a flux—an additive that helps lower the melt temperature during smelting. Agricola, the "father of mineralogy", 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) and then to fluorite.

Hydrofluoric acid was known as a glass-etching agent from the 1720s, perhaps as early as 1670. Andreas Sigismund Marggraf made the first scientific report on its preparation in 1764 when he heated fluorite with sulfuric acid; the resulting solution corroded its glass container. Swedish chemist Carl Wilhelm Scheele repeated this reaction in 1771, recognizing the product as an acid, which he called "fluss-spats-syran" (fluor-spar-acid).

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, although Greek, Russian, and some others (following Ampère's suggestion) use the name ftor or derivatives, from the Greek φθόριος (phthorios), meaning "destructive". The New Latin name (fluorum) gave the element its current symbol, F, although the symbol Fl was used in early papers.

Isolation

Progress in isolating the element was slowed by the exceptional dangers of generating fluorine: several 19th-century experimenters, the "fluorine martyrs", were killed or badly hurt while working with hydrofluoric acid. Initial attempts to isolate the element were hindered by problems getting a suitable conducting liquid for electrolysis as well as by the extreme corrosiveness of hydrogen fluoride and of fluorine gas.

Edmond Frémy thought that passing electric current through pure hydrofluoric acid (dry HF) might allow the element to be isolated. 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 would not pass an electric current.

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 mixture proved to conduct electricity, making electrolysis possible. However, rapid destruction of the platinum metal in his electrochemical cells still stymied the quest. To continue, Moissan devised a strategy of cooling the reaction in a special bath to extremely low temperatures, which slowed the rate of corrosion. Moissan also constructed more corrosion-resistant equipment: containers crafted from a mixture of platinum and iridium (more chemically resistant than pure platinum) with fluorite stoppers. In 1886, Moissan crowned 74 years of effort by many chemists when he isolated elemental fluorine.

In 1906, two months before his death, Moissan received the Nobel Prize in chemistry. The citation:

...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 apparatus, 1887 publication	Henri Moissan, Nobel Prize photo

Application development

In the late 1920s, chlorofluorocarbon refrigerants were tested by researchers from the Frigidaire division of General Motors. In 1930, GM formed a joint venture with DuPont, Kinetic Chemicals, to commercialize Freon-12 (CCl
₂F
₂). It proved to be a marketplace hit, rapidly replacing earlier, more toxic, refrigerants and growing the overall market for kitchen refrigerators. By 1949, DuPont had bought out the joint venture and marketed several other Freon molecules.

Uranium hexafluoride, U.S. Department of Energy

In 1938, Teflon was discovered by accident by a recently hired Kinetic Ph.D., Roy J. Plunkett. Undertaking research on possible use of tetrafluoroethylene as a refrigerant, he encountered a mystery. Gas left in a cylinder overnight could not be released the next morning, but the weight of the container had not changed (indicating the gas had not leaked out). Cutting the cylinder open, he found white flakes of a polymer new to the world. Tests showed the substance was polytetrafluoroethylene (poly meaning many). The polymer was more resistant to corrosion and had better high temperature stability than did any other plastic. By 1941, an accelerated commercialization program was making significant quantities.

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 used even more elemental fluorine to make uranium hexafluoride for use in uranium enrichment plants. Because UF₆ is corrosive in a similar manner to fluorine itself, gaseous diffusion separation plants were built with special materials. Nickel was used for the membranes; fluoropolymers such as Teflon were used for seals; and liquid fluorocarbons were used as coolants and lubricants. After the war, the nuclear weapons industry drove further development of fluorochemicals.

Characteristics

Phases

Solid fluorine's beta crystal structure: the spheres indicate F₂ molecules that are disordered by rotations to any angle; other molecules are disordered in planes.

Main article: Phases of fluorine

At room temperature, fluorine is a gas of diatomic molecules. Though sometimes described as yellow-green, pure fluorine is actually a very pale yellow. The gas has a pungent characteristic odor that is noticeable in concentrations as low as 20 ppb. Fluorine condenses to a bright yellow liquid at −188 °C, which is near the condensation temperatures of oxygen and nitrogen.

Fluorine solidifies at −220 °C into a cubic structure called beta-fluorine. This phase is transparent and soft with significant disorder of the molecules. With further cooling to −228 °C, 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 releases more energy than the melting point transition, and can be violent. In general, fluorine's solid state form is more similar to oxygen's than to those of the other halogens.

Isotopes

Main article: Isotopes of fluorine

Fluorine is defined by its atomic number, nine, which indicates how many protons it has. One stable isotope occurs naturally: fluorine-19, which contains ten neutrons. The element is both monoisotopic (having a single stable isotope) and mononuclidic (being found on Earth in only one isotope). These attributes make it useful in uranium enrichment and in nuclear magnetic resonance.

Seventeen radioisotopes have been synthesized: mass numbers 14–18 and 20–31. The lightest fluorine isotopes, those with mass numbers of 14–16, decay via electron capture. Fluorine-17 and -18 undergo beta plus decay (emission of a positron). All isotopes heavier than fluorine-19 decay by beta minus mode (emission of an electron); some of these also decay by neutron emission. Fluorine-18 is the most stable radioisotope, with a half-life of 109.77 minutes before decaying to oxygen-18.

Electron arrangement

Simplified structure of the fluorine atom

The neutral atom has nine electrons, one fewer than neon. The electronic configuration is 1s²2s²2p⁵: a filled inner shell of two electrons and an unfilled outer shell containing seven (one short of being filled). The outer electrons do not shield each other much from the nucleus. They therefore experience a high effective nuclear charge of seven (nine minus two), which affects the physical properties of the atom.

Removal of electrons from neutral atoms is very difficult; fluorine's ionization energy (the energy required to remove an electron) is higher than that of any other element except neon and helium. Instead, fluorine exhibits a very strong preference for capturing one more electron to achieve the filled-shell electron configuration of noble gas neon: 1s²2s²2p⁶. Fluorine is the element with the highest electronegativity (a relative measure of electron attraction by atoms). Its electron affinity (the energy released by adding an electron) is higher than that of any element except chlorine. Fluorine atoms have a small covalent radius of around 60 picometers. This is similar to the radii of oxygen and neon, the left and right hand neighbors of the periodic table.

Chemical reactivity

Main article: Chemical characteristics of fluorine

The difluorine bond is relatively weak, with a bond energy much less than that in Cl
₂ or Br
₂ and similar to the easily cleaved oxygen–oxygen bonds of peroxides. For this reason, elemental fluorine easily dissociates to react with other atoms. On the other hand, bonds to non-fluorine atoms are very strong because of atomic fluorine's electron attraction. Both the easy breaking apart of difluorine and the strong bonding to other atoms make fluorine extremely reactive. 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.

External videos
Bright flames during fluorine reactions
Fluorine reacting with caesium

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 passivation (protective metal fluoride layers formed with initial exposure). The alkali metals, such as sodium, react explosively. The alkaline earth metals, such as calcium, react somewhat less dramatically. The noble metals (gold, platinum, etc.) react least readily, requiring pure fluorine gas at 300–450 °C.

Fluorine reacts explosively with hydrogen in a manner similar to that of alkali metals. The other halogens react readily with fluorine gas, as does the heavy noble gas radon. The lighter noble gases xenon and krypton react directly with fluorine under special conditions. Oxygen and fluorine gases do not normally react, but they can be combined under electric discharge at low pressures and temperatures. The products can be unstable and tend to separate back into fluorine and oxygen when heated. Nitrogen, with its very strong triple bonds, requires electric discharge and very high temperatures to react with fluorine even though the end product, NF₃, is quite stable. Ammonia (NH₃) can react explosively with fluorine.

Carbon, in graphite or diamond form, is impervious to fluorine gas at room temperature. Above 400°C, graphite reacts with fluorine to make a varying-composition solid called "carbon monofluoride". At higher temperatures, gaseous fluorocarbons start to be produced; the reaction can become explosive. Carbon dioxide and carbon monoxide react with fluorine at room temperature or a little above. Organic chemicals, such as paraffins, react strongly when exposed to fluorine. Even fully halogenated organic molecules, such as normally incombustible carbon tetrachloride, can explode.

The other solid nonmetals or metalloids (boron, silicon, arsenic, sulfur, germanium, phosphorus, selenium, tellurium) burn with a flame in room temperature fluorine. Hydrogen sulfide and sulfur dioxide combine readily with fluorine; the latter reaction can be explosive. Sulfuric acid reacts much more sluggishly.

Occurrence

Main article: Origin and occurrence of fluorine

Universe

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

At 400 ppb, fluorine is estimated to be the 24th most common element in the Universe. For the ninth lightest element, it is unusually rare (the lighter elements tend to be the more common ones). All of the elements from atomic number 6 (carbon) to atomic number 12 (magnesium) are hundreds to thousands of times more common except for sodium (tens of times more common).

Fluorine's rarety results from both a low birth rate and rapid destruction. The main fusion reaction sequences of stars (stellar nucleosynthesis) which produce oxygen, carbon, and neon bypass fluorine. Any fluorine which is nonetheless created is a large target (has a high nuclear cross section) for further fusion—either with hydrogen to form oxygen and helium, or with helium to make neon and hydrogen.

The presence of fluorine at all—outside its fleeting existence in stars—is somewhat of a mystery because of these fluorine-eliminating reactions. Three theoretical solutions 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.

Earth

See also: List of countries by fluorite production

Fluorine is the thirteenth most common element in Earth's crust, comprising between 600 and 700 ppm by mass. (Rocky planets like Earth tend to concentrate elements that form non-volatile compounds.) On Earth, fluorine, because of its reactivity, is found essentially only in mineral compounds. The three most industrially significant are fluorite, fluorapatite, and cryolite:

• Fluorite (CaF
  ₂), also called fluorspar, is the main source of commercial fluorine. It is colorful, common, and found worldwide. China supplies over half the world's demand; Mexico is second. The United States produced most global fluorite in the early 20th century, but its last mine, in Illinois, closed in 1995.

• Fluorapatite (Ca₅(PO₄)₃F) and other apatites are mined in high volumes to produce phosphates for fertilizers. Most of the Earth's fluorine is bound in fluorapatite, but because the fluorine fraction is low (3.5%), it is discarded as waste. Only in the United States is there significant recovery: byproducts are used to supply water fluoridation.

• Cryolite (Na
  ₃AlF
  ₆), the least abundant of the three major fluorine-containing minerals, is a concentrated source of fluorine. It was formerly used directly in aluminium production. The main commercial mine, on the west coast of Greenland, closed in 1987.

Major fluorine-containing minerals
Fluorite	Fluorapatite	Cryolite

Several other minerals, such as the gemstone topaz, contain fluoride. Fluoride is not significant in seawater or brines, unlike the other halides, because the alkaline earth fluorides (e.g. CaF₂, MgF₂) precipitate out of water. Commercially insignificant quantities of organofluorines have been observed in volcanic eruptions and in geothermal springs. Their ultimate origin (from biological sources or geological formation) is unclear.

The possibility of small amounts of gaseous fluorine within crystals has been debated for many years. One form of fluorite, antozonite, has an smell suggestive of fluorine when crushed. In 2012, a study reported detection of trace quantities (0.04% by weight) of diatomic fluorine in antozonite. It was suggested that radiation from small amounts of uranium within the crystals had caused the free fluorine defects.

Industry and applications

Main article: Fluorochemical industry

The main source of fluorine, fluorite mining, was a growing industry through 1989 when it reached a high of 5.6 million metric tons per year. Environmental restrictions on a main end market (chlorofluorocarbons) led to reduced production, down to 3.6 million tons in 1994. Since then production has steady risen. In 2003, it was estimated at 4.5 million tons and a corresponding revenue of $550 million. A 2012 market research report estimated that production would hit about 5.9 million tons by 2017.

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 hydrogen fluoride, which is the primary chemical intermediate for the industry.

A market survey estimated 2011 global fluorochemical sales at $15 billion. It predicted 5% growth to 2018, when the market would exceed $20 billion. A different report said that, in 2016, global fluorochemical production would be almost $20 billion with a volume of 3.5 million metric tons per year.

The fluorine industry's supply chain, based on mass flows: click for links to related articles.

Inorganic fluorides

About 3 kg 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 other iron alloys.

Cryolite (Na₃AlF₆) use in aluminium smelting

Acidspar grade fluorite is added to ceramics, enamels, glass fibers, clouded glass, cement, and the outer coating of welding rods. Most acidspar is reacted with sulfuric acid to make hydrofluoric acid. Significant direct uses of HF include pickling (cleaning) steel, etching glass, and cracking alkanes in the petrochemical industry.

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. The fluorides are not reactants in the smelting process but fluxes that lower the temperature of the melt. They are not consumed in the process and remain available to support smelting. However, over time small amounts are lost through side reactions with the smelting apparatus, and new fluorides must be added. About 23 kg are required for every metric ton of aluminium.

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. 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). Sodium and potassium bifluorides are significant to the chemical industry.

Organic fluorochemicals

Organofluoride production consumes over 40% of hydrofluoric acid (over 20% of all mined fluorite). Within organofluorides, refrigerant gases are the dominant segment, about 80% on a fluorine basis. Fluoropolymers are less than one quarter the size of refrigerant gases in terms of fluorine usage but are growing faster. Fluorosurfactants are molecules used to make clothing and other items water resistant. They are a small segment in size but are over $1 billion yearly revenue.

Industrially, production of fluorocarbons relies on indirect methods because the direct reaction of hydrocarbons with fluorine gas can be dangerous at temperatures above −150 °C. Many fluorochemicals are made by halogen exchange reactions: chlorinated hydrocarbons react with hydrogen fluoride to switch out chlorine for fluorine. The reactions are catalyzed, for example by antimony halides in "Swarts fluorination". Another method is electrochemical fluorination, in which hydrocarbons are electrolyzed in hydrogen fluoride. In the Fowler process, hydrocarbons are reacted with solid carriers of fluorine, notably cobalt trifluoride.

Refrigerant gases

See also: Refrigerant

Halogenated molecules used in refrigeration are identified by a system of numbering (the R-number system) that explains the amount of fluorine, chlorine, carbon, and hydrogen in the molecules. The DuPont brand Freon has been colloquially used for these compounds, but brand-neutral terminology uses "R" ("refrigerant") as the prefix.

Traditionally chlorofluorocarbons (CFCs) were the predominant class of fluorinated organic chemical. Prominent CFCs included R-11 (trichlorofluoromethane), R-12 (dichlorodifluoromethane), and R-114 (1,2-dichlorotetrafluoroethane). 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.

Hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) serve as replacements for CFC refrigerants; few were commercially manufactured before 1990. 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).

Fluoropolymers

Main article: Fluoropolymer

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. The 2011 global fluoropolymer market was estimated at slightly under $6 billion in revenue and predicted to grow 6.5% per year through 2016. Fluoropolymers are formed by polymerizing free radicals; other hydrocarbon polymerization techniques do not work.

"Gore-Tex" electron microscope image: the membrane's microstructure has islands of polymer with connecting strands.

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

When jerk-stretched, PTFE film becomes a fine-pored membrane: expanded PTFE (ePTFE). The term "Gore-Tex" is sometimes used generically although 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.

Other fluoropolymers tend to have similar properties to PTFE, which leads to use in electrical insulation and the chemical process industry. Unlike PTFE, the other fluoropolymers can be melt processed. This makes them easier to work with—they can be formed into complex shapes—but they are more expensive than PTFE and have lower thermal stability. Fluorinated ethylene propylene (FEP) is the second most produced fluoropolymer. Films from two different fluoropolymers serve as glass-replacements in solar cells.

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 polymer 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.

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

Surfactants

Fabric treated with fluorosurfactant

Main articles: Fluorinated surfactant and Durable water repellent

Fluorinated surfactants are small organofluorine molecules, principally used for water and stain resistance. As of 2006, yearly revenues for this segment were over $1 billion. Fluorosurfactants are expensive chemicals, comparable to pharmaceutical chemicals: $200–2000 per kilogram ($90–900 per pound). Scotchgard has been a prominent brand; in 2000 the revenues were over $300 million.

Fluorosurfactants make a very small part of the overall surfactant market, most of which is hydrocarbon based and much cheaper. Some potential applications (e.g. low cost paints) are unable to use fluorosurfactants because of the price impact of compounding in even small amounts of fluorosurfactant. Use in paints was only about $100 million as of 2006.

Fluorine gas

See also: Industrial gas

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 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.

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. Fluorine's monoisotopic natural occurrence makes it useful in uranium enrichment, in which uranium-235 is separated by diffusion or centrifugation from uranium-238. The difference in mass between uranium hexafluoride molecules arises entirely from the uranium isotopes.

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.

Several compounds made from elemental fluorine serve the electronics industry. Rhenium and tungsten hexafluorides are used for chemical vapor deposition of thin metal films. Tetrafluoromethane is used for plasma etching. Nitrogen trifluoride is used for cleaning equipment.

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 ClF
₃, BrF
₃, and IF
₅ provide gentler fluorination, with a series of strengths, and are easier to handle. Sulfur tetrafluoride is used for making fluorinated pharmaceuticals.

Production of fluorine gas

Industrial fluorine cells

Modern industrial production of elemental fluorine uses Moissan's process of electrolyzing potassium fluoride/hydrogen fluoride mixtures but with an apparatus made of different materials. The steel container acts as the negative electrode, attracting H⁺ ions and releasing hydrogen gas. A carbon block (similar to that used in aluminium production) acts as the positive electrode, attracting F⁻ ions and releasing fluorine gas. The voltage difference between the electrodes is 8–12 volts.

Commercial temperatures are now higher than what Moissan used. A mixture with the approximate composition KF•2HF melts at 70 °C (158 °F) and is electrolyzed at 70–130 °C. Because HF alone cannot be electrolyzed, the presence of some KF is critical even though it is not consumed in the cell.

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. Above that temperature, nickel is required. Regulator valves are made of nickel. Fluorine piping is generally made of nickel or Monel (nickel-copper alloy). 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, but some sources recommend systems made of nickel, Monel, and PTFE.

Chemical synthesis of the element

In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl O. 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. The main idea was that some metal fluoride anions did not have a neutral counterpart (or had ones that are very unstable) and their acidifying could result in chemical oxidation, rather than formation of the expected molecules. Christe's reactions:

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

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

Biological aspects

Main article: Biological aspects of fluorine

Natural biochemistry

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. Small amounts of fluoride may be beneficial for bone strength, but this has not been definitively established. As there are many environmental sources of trace fluorine, the possibility of "fluorine deficiency" only pertains to artificial diets.

Biologically synthesized organofluorines have been found in microorganisms and plants but not in animals. The most common example is fluoroacetate. It is used as a defense against herbivores by at least 40 green plants in Australia, Brazil, and Africa. Other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate. The enzyme adenosyl-fluoride synthase, which makes the carbon–fluorine bond, was isolated in bacteria in 2002.

Medicine

Dental care

Main articles: Fluoride therapy and Water fluoridation

Since the mid-20th century, population studies have shown that fluoride reduces tooth decay. The initial hypothesis was 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.

Topical fluoride treatment in Panama

Water fluoridation is the controlled addition of fluoride to a public water supply to reduce tooth decay. It 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 and for about 6% of people worldwide. Although the best available evidence shows no association with adverse effects other than dental fluorosis, most of which is mild, water fluoridation has been contentious for ethical, safety, and efficacy reasons. Opposition to water fluoridation exists despite its support by public health organizations. 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. Reviews of the scholarly literature in 2000 and 2007 showed significant reduction of cavities in children associated with water fluoridation.

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. Almost all toothpaste in developed countries is now fluoridated, as are many prescription and non-prescription mouthwashes. Fluoride may also be applied to teeth in gels, foams, or varnishes.

Pharmaceuticals

About 20% of modern pharmaceuticals contain fluorine, including commercially significant drugs in many different pharmaceutical classes. One of these, the cholesterol-reducer atorvastatin (Lipitor), was the number one money-making drug for nearly a decade. The branded asthma medication Seretide, a top-ten revenue drug as of the mid-2000s, contains two active ingredients, one of which is fluorinated: fluticasone.

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. Because of the considerable stability of the carbon–fluorine bond, many drugs are fluorinated to delay their metabolism and elimination. This allows longer times between doses. 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.

Fluoxetine (Prozac)

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 fluoxetine (Prozac). It became the best-selling antidepressant. Some other selective serotonin reuptake inhibitor (SSRI) antidepressants that are fluorinated are citalopram (Celexa) and its isomer escitalopram (Lexapro), fluvoxamine (Luvox), and paroxetine (Paxil).

Quinolones are artificial compounds that are broad-spectrum antibiotics. Most of the currently used quinolones are fluorinated to make the drugs more powerful. Prominent examples include ciprofloxacin (Cipro) and levofloxacin (Levaquin). The latter was the highest selling U.S. antibiotic in 2010.

Fluorine also finds use in many steroidal drugs. Fludrocortisone (Florinef) is a mineralocorticoid (a compound used to retain sodium and water and thus raise blood pressure). Triamcinolone and dexamethasone are potent glucocorticoids (anti-inflammatories).

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 effective for even longer periods, and are almost insoluble in blood, allowing the patient to awaken more quickly. Examples include sevoflurane, desflurane, enflurane, and isoflurane, all fluorinated ethers.

PET scanning

Whole-body PET scan using fluorine-18

Main article: Positron emission tomography

Fluorine-18 is being routinely produced by the radiopharmaceutical industry using particle accelerators to produce radioactive tracers for positron emission tomography (PET) scanning. The radioisotope's half-life of almost two hours is long enough to allow its transportation from the production facility to the imaging center while keeping minimal the patient radiation exposure. The most widely used radiopharmaceutical is fluorodeoxyglucose (FDG). After injection into the blood, FDG is taken up by tissues with a high need for glucose, such as the brain and most types of malignant tumors. Computer assisted tomography (CAT) can then be used for detailed imaging.

Oxygen carrier research

See also: Blood substitute and Liquid breathing

Liquid fluorocarbons have a very high capacity for holding gas in solution. They can hold more oxygen or carbon dioxide than blood does and for that reason, they have attracted ongoing interest related to the possibility of artificial blood or of liquid breathing.

Blood substitutes are the subject of research because the demand for blood transfusions grows faster than donations. In some scenarios, artificial blood may be more convenient or safe. Because fluorocarbons do not normally mix with water, they must be mixed into emulsions (small droplets of perfluorocarbon suspended in water) to be used as blood. One such product, Oxycyte, has been through initial clinical trials. PFC doping has the potential to aid endurance athletes and is therefore banned from sports. One cyclist's mysterious near death in 1998 prompted an investigation for PFC abuse.

Possible medical uses of liquid breathing (which uses pure perfluorocarbon liquid, not a water emulsion) involve assistance for premature babies or for burn victims (because the normal lung function is compromised). Both partial filling of the lungs and complete filling of the lungs have been considered, although only the former has any significant tests in humans. Several animal tests have been performed and some human partial liquid ventilation trials. One effort, by Alliance Pharmaceuticals reached clinical trials but was abandoned because the results were not better than normal therapies.

Agrichemicals and poisons

1080 use in New Zealand

An estimated 30% of agrichemical compounds contain fluorine. Most of them are herbicides and fungicides, but a few regulate crop growth. Fluorine substitution (usually of just a single atom or at most a trifluoromethyl group) is a powerful tool for new molecule design. The molecular effects—increasing biological stay time, membrane crossing, altering molecular recognition—are similar to fluorinated pharmaceuticals. Trifluralin is a prominent example, used widely in the United States as a weedkiller. Its suspected carcinogenic properties have caused many European countries to ban it.

Sodium monofluoroacetate (brand name 1080) is a commercial mammalian poison. The molecule is similar to the one in vinegar but with a hydrogen changed out for fluorine. It was first synthesized in the late 19th century and then recognized as an insecticide in the early 20th century. Later, 1080 was widely used to control rats and other mammals. The compound is now banned in Europe and the United States, but it 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. Several insecticides contain sodium fluoride, which is much less toxic to humans than fluoroacetate.

Compounds

Main article: Compounds of fluorine

Fluorine's common oxidation state is −1. Because of its attraction for electrons, fluorine forms many ionic compounds. Covalent bonds involving fluorine are polar and almost always, within molecules, are single bonds. 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.

Hydrogen fluoride

Main articles: Hydrogen fluoride and Hydrofluoric acid

Hydrogen fluoride shows similar trend breaking as water: boiling points of the hydrogen halides and hydrogen chalcogenides.

Fluorine combines with hydrogen to make a compound called hydrogen fluoride (HF). HF molecules cluster together weakly via hydrogen bonds. Because of this, hydrogen fluoride behaves more like water than like HCl (hydrochloric acid). 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.

Water solutions of hydrogen fluoride are called hydrofluoric acid. This is a chemically weak acid, unlike the other hydrohalic acids (hydrochloric, etc.) which are all strong. Although hydrofluoric acid is weak, it is very corrosive, even attacking glass.

Metal fluorides

See also: Fluoride volatility

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. The difluorides of the alkaline earths are also very ionic but are generally very insoluble. Beryllium difluoride is an exception: it exhibits some covalent character, is water soluble, and has a structure similar to SiO₂ (quartz). Trifluorides are formed by many metals, particularly the rare earths, and are in general ionic.

The tetrafluorides are a transition region from ionic to covalent bonding. Zirconium and hafnium, along with several actinides, form high melting, ionic tetrafluorides. On the other hand, the tetrafluorides of titanium, vanadium, niobium are polymeric. They melt or decompose below about 350 °C. The pentafluorides are even more covalently bonded, forming low dimensionality polymers or oligomeric molecules (clusters).

A total of thirteen metal hexafluorides have been characterized, all octahedredral molecules. All are volatile solids except molybdenum hexafluoride and rhenium hexafluoride (liquids) and tungsten hexafluoride (a gas). The only definite metal heptafluoride, that of rhenium, is a low-melting molecular solid. Its structure is a distorted pentagonal bipyramid. The higher metal fluorides are very reactive.

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

Nonmetal fluorides

The binary fluorides of the main group nonmetals and metalloids are generally volatile, covalently bonded molecules but vary greatly in their reactivities. Nonmetals from the third row of the periodic table and below can form fluorides which are hypervalent (more bonds than normal):

• Boron trifluoride is a planar molecule where the boron atom has an incomplete octet (less bonds than normal). It is a weak Lewis acid and readily accepts a Lewis base, such as ammonia, forming adducts (combinations).

• The simplest binary compound with carbon is carbon tetrafluoride, an inert tetrahedral molecule. The atoms below carbon, silicon and germanium, also form tetrahedral tetrafluorides. But the compounds are Lewis acids.

Chlorine trifluoride

• The atoms below nitrogen form trifluorides that are weak Lewis bases and are more reactive as the atom becomes heavier. (Nitrogen's trifluoride is stable against hydrolysis and is not a Lewis base.) Pentafluorides are formed by phosphorus, arsenic and antimony. They are even more reactive than the respective trifluorides; antimony pentafluoride is the strongest Lewis acid of all charge-neutral compounds.

• The chalcogens (oxygen's column) form a variety of fluorides. Unstable difluorides are known for oxygen (the only compound where oxygen is at formal oxidation state +2) as well as sulfur and selenium. Tetrafluorides and hexafluorides are known for sulfur, selenium, and tellurium. They tend to be more stable with more fluorines and with a lighter central atom: SF₆ is extremely inert.

• The well-characterized heavier halogens (chlorine, bromine, and iodine) all form mono-, tri-, and pentafluorides: XF, XF₃, and XF₅. For XF₇, only iodine heptafluoride is known. Many of the halogen fluorides are powerful fluorinators (sources of fluorine atoms). Chlorine trifluoride readily fluorinates asbestos and refractory oxides, and industrial use requires precautions similar to those for fluorine gas.

Noble gas compounds

Main article: Noble gas compound

Xenon tetrafluoride, 1962

The noble gases are generally non-reactive because they have complete electron 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. Since then, xenon's difluoride, tetrafluoride, and hexafluoride have been isolated, as well as various oxyfluorides. Krypton, xenon's lighter homolog forms a difluoride and a few more complicated fluorine-containing compounds. Radon, xenon's heavier homolog, reacts readily with fluorine to form a solid, thought to be radon difluoride.

The lightest noble gases do not form stable binary fluorides. Argon, however, reacts in extreme conditions with hydrogen fluoride to form argon fluorohydride. Helium and neon do not form any time-stable fluorides, but helium fluorohydride has been observed for milliseconds at extremely high pressure and low temperature. Neon is considered even less reactive than helium, and no fluorides have been even momentarily observed.

Organic compounds

Main article: Organofluorine chemistry

The carbon–fluorine chemical bond is the strongest bond in organic chemistry, and organofluorines are very stable. Save a few exceptions, the C–F bond does not exist in nature, meaning the entire field is essentially "man-made"; research in particular areas is driven by the commercial value of applications. The range of organofluorine compounds is diverse, reflecting the inherent complexity of organic chemistry.

Small molecules

Perfluorocarbon density demonstration (animals were rescued after the photo).

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.

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.

Perfluorinated compound is a term for hydrocarbons that are fully fluorinated but which also have a functional group (a small non-hydrocarbon part of the molecule). Often the functional group is a carboxylic acid (-CO₂H) group. Perfluorinated compounds exhibit many perfluorocarbon properties (e.g. inertness, stability, non-wetting by water and oils, slipperiness). However, the functional group is available for reactions. It may also help the molecule to adhere to surfaces or behave as a surfactant (a soap-like mixture). Fluorosurfactants can lower the surface tension of water below that achievable with hydrocarbon-based surfactants. 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.

Polymers

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.

The complex unit structure of Nafion

The simplest fluoroplastic is PTFE (polytetrafluoroethylene, 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 PE (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.

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). 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.

There are other fluoroplastics that are not perfluorinated (contain some C-H). PVDF (polyvinylidene fluoride, structural unit: –CF₂CH₂–) has half the fluorines of PTFE. PVF (polyvinyl fluoride, structural unit: –CH₂CHF–) has even less. Despite this, it still has many properties of fully fluorinated polymers.

Hazards

The U.S. hazard signs for commercially transported fluorine

Main article: Fluorine-related hazards

Fluorine gas

Elemental fluorine is highly toxic. Above a concentration of 25 ppm, it causes significant irritation while attacking the eyes, airways and lungs and affecting the liver and kidneys. At a concentration of 100 ppm, human eyes and noses are seriously damaged.

Hydrofluoric acid

See also: Chemical burn

Hydrofluoric acid, the water solution of hydrogen fluoride, is a contact poison. Even though it is chemically only a weak acid, it is far more dangerous than the conventional strong mineral acids, such as nitric acid, sulfuric acid, or hydrochloric acid. 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. From 1984 to 1994, at least nine U.S. workers died from accidents with HF.

Typical HF burns: the outward signs may not be evident for a day, at which point calcium treatments are less effective.

Once in the blood, hydrogen fluoride reacts with calcium and magnesium, resulting in electrolyte imbalance (potentially hypocalcemia). The consequent effect on the heart (cardiac arrhythmia) may be fatal. Formation of insoluble calcium fluoride also causes strong pain. Burns with areas larger than 160 cm², about the size of a man's hand, can cause serious systemic toxicity.

Symptoms of exposure to hydrofluoric acid may not be immediately evident, with 8-hour delay for 50% HF and up to 24 hours for lower concentrations. Hydrogen fluoride interferes with nerve function, meaning that burns may not initially be painful.

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 present a danger. 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 or special rinsing solutions. Because HF is absorbed, further medical treatment is necessary. Calcium gluconate may be injected or administered intravenously. Use of calcium chloride is contraindicated and may lead to severe complications. Sometimes surgical excision of tissue or amputation is required.

Fluoride ion

See also: Fluoride toxicity

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. The dose that may lead to adverse health effects is about one fifth the lethal dose. Chronic excess fluoride consumption can lead to skeletal fluorosis, a disease of the bones that affects millions in Asia and Africa.

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.

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

Dangers from toothpaste are more serious for small children—the U.S. CDC recommends children under 6 years of age be supervised when brushing their teeth so they don't swallow toothpaste. One regional study examined a year of fluoride poisoning reports for pre-teens. Amongst 87 cases, there was one fatality (from insecticide). The other 86 were all from dental fluoride. Most had no symptoms, but about 30% had stomach pains, with symptoms more likely when more fluoride was consumed. A larger study of American fluoride poisoning reports showed similar findings. 80% of the reports were related to children under 6. Few were serious, although several hundred cases were treated at health facilities yearly.

Environmental concerns

Atmosphere

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

See also: Ozone depletion and global warming

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 the fluorine. Because of the inherent stability of these fully halogenated molecules (which makes them so nonflammable and useful), they are able to attain the upper reaches of the atmosphere before decomposing. At high altitudes, they release chlorine and bromine atoms which attack ozone molecules. 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 fully recover. Early indications are that the CFC ban is working: ozone depletion has stopped, and recovery has started.

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

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.

Biopersistance

Main article: Biopersistance of fluorinated organics

Because of the strength of the carbon–fluorine bond, organofluorines endure in the environment. Perfluoroalkyl acids (PFAAs) have attracted particular attention as persistent global contaminants. Because of the acid group, PFAAs are water soluble in low concentrations. While there are other PFAAs, the lion's share of environmental research has been done on the two most well-known: perfluorooctanesulfonic acid (PFOS) and perfluorooctanoic acid (PFOA).

Trace quantities of PFAAs have been detected worldwide, from polar bears in the Arctic to the global human population. Both PFOS and PFOA has been detected in breast milk and the blood of newborns. A 2013 review showed widely varying amounts of PFAA in different soils and groundwater, with generally higher amounts where there was more human activity. There was no clear pattern of one chemical dominating, and higher amounts of PFOS were correlated to higher amounts of PFOA.

The PFOS molecule

In the body, PFAAs bind to proteins such as serum albumin. Unlike chlorinated hydrocarbons, PFAAs are not lipophilic (stored in fat). Their tissue distribution in humans is unknown, but studies in rats suggest they are present mostly in the liver, kidney, and blood. They are not metabolized by the body but are excreted by the kidneys. Dwell time in the body varies greatly by species. Rodents have half-lives of days, while in humans they remain for years.

The potential health impact of PFAAs is unclear. Both PFOA and PFOS in high doses cause cancer and the death of newborns in rodents. However, studies on humans have not been able to prove an impact at current exposures.

Less fluorinated chemicals (not perfluorinated compounds) are also detectable in the environment. 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. Fluorine-containing agrichemicals are measurable in farmland runoff and nearby rivers.

See also

• Fluoride selective electrode, the device for measuring fluoride ion concentration
• Fluorine dating, a relative dating technique for buried organic matter
• Fluorine NMR, an analytical technique for fluorine-containing molecules
• Fluorous chemistry, perfluorocarbon solvent extractions
• Krypton fluoride laser and argon fluoride laser, ultraviolet lasers

Notes

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 unknown properties.

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Fischman, Michael L. (2001). "Semiconductor Manufacturing Hazards In John B. Sullivan and Gary R. Krieger, eds., Clinical Environmental Health and Toxic Exposures (pp. 431–465). 2nd ed.". Philadelphia, PA: Lippincott Williams & Wilkins. ISBN 978-0-683-08027-8. {{cite book}}: Invalid |ref=harv (help); Missing or empty |title= (help)

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