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)
Electron configuration	[He] 2s2 2p5 2, 7
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
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa) 1.696 g/L
Liquid density at b.p.	1.505 g·cm−3
Melting point	53.53 K, −219.62 °C, −363.32 °F
Boiling point	85.03 K, −188.12 °C, −306.62 °F
Critical point	144.4 K, 5.215 MPa
Heat of vaporization	6.51 kJ·mol−1
Molar heat capacity	(Cp) (21.1 °C) 825 J·mol−1·K−1 (Cv) (21.1 °C) 610 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 (Pauling scale)
Ionization energies (more)	1st: 1,681 kJ·mol−1
	2nd: 3,374 kJ·mol−1
	3rd: 6,147 kJ·mol−1
Covalent radius	64 pm
Van der Waals radius	135 pm
Miscellanea
Crystal structure	cubic alpha state: just below the melting point, at 1 atm
	monoclinic beta state
Magnetic ordering	diamagnetic, −1.2×10−4 (SI)
Thermal conductivity	0.02591 W·m−1·K−1
CAS registry number	7782-41-4
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

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.

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 1530 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 free fluorine. Instead, hydrofluoric acid is the key intermediate for the US$16 billion-per-year, global fluorochemical industry. The fluorides of low-charged metals are ionic compounds (salts); those of high-charged metals are volatile molecular compounds. The largest uses of inorganic fluoride compounds are steel making and aluminium refining. Organic fluorine compounds tend to have high chemical and thermal stability. The largest commercial use is in refrigerant gases (Freon); 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

Phases

Main article: Phases of fluorine

Fluorine forms diatomic molecules (F₂) that are gaseous at room temperature. 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. The element 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 (−307 °F), which is near the condensation temperatures of oxygen and nitrogen. It solidifies at −220 °C (−363 °F) 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 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'.

Low-temperature fluorine phases	Alpha-fluorine crystal structure

Atomic structure

A fluorine atom has nine protons and nine electrons, one fewer than neon, arranged in the electronic configuration [He] 2s² 2p⁵. 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.

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. Fluorine's ionization energy (energy required to remove an electron to form F⁺) is higher than that of any element except neon or helium. Fluorine's electron affinity (energy released by adding an electron to form F^–) is higher than that of any element except chlorine.

Molecular structure

Main article: Difluorine

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).

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.

Chemical reactivity

Main article: Chemical characteristics of fluorine

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.

External video

Cold fluorine gas impinging on several substances causes bright flames

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.

Reactions of elemental fluorine with metals require different conditions that depend on the metal. Often, the metal (such as aluminium, iron, or copper) must be powdered because many metals passivate by forming protective layers of the metal fluoride that resist further fluoridation. The alkali metals react with fluorine with a bang; the alkaline earth metals react not quite as aggressively. The noble metals ruthenium, rhodium, palladium, platinum, and gold react least readily, requiring pure fluorine gas at 300–450 °C (575–850 °F).

Fluorine reacts explosively with hydrogen in a manner similar to that of alkali metals. The halogens react readily with fluorine gas as does the heavy noble gas radon. 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. Nitrogen, with its very stable triple bonds, requires electric discharge and high temperatures to combine with fluorine directly.

Isotopes

Main article: Isotopes of fluorine

Fluorine occurs naturally on Earth exclusively in the form of its only stable isotope, fluorine-19, which makes the element monoisotopic and mononuclidic. Seventeen radioisotopes have been synthesized: mass numbers 14–18 and 20–31. Fluorine-18 is the most stable radioisotope of fluorine, with a half-life of 109.77 minutes. It is also the lightest unstable nuclide with equal odd numbers of protons and neutrons.

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.

Origin and occurrence

Main article: Origin and occurrence of fluorine

In the 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

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.

Three theoretical solutions to the mystery exist. In type II supernovae, atoms of neon are 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 blows the fluorine out of the star before hydrogen or helium can destroy it. In asymptotic giant branch (a type of red giant) stars, fusion reactions occur in pulses and convection lifts fluorine out of the inner star. Only the red giant hypothesis has supporting evidence from observations.

On 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 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.

• 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 the last mine, in Illinois, shut down in 1995.
• 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.
• 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.

Major fluorine-containing minerals
Fluorite	Fluorapatite	Cryolite

History

Smelting illustration from Agricola's De re metallica, where fluorite was first described

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 melt ores and slags during smelting. Fluorite stones were called schone flusse in the German of the time. Agricola, writing in Latin but describing 16th century industry, invented several hundred new Latin terms. For the schone flusse stones, he used the Latin noun fluores, "fluxes", because they made metal ores flow when in a fire. After Agricola, the name for the mineral evolved to fluorspar (still commonly used) and then to fluorite.

Some sources claim that the first production of hydrofluoric acid was by Heinrich Schwanhard, a German glass cutter, in 1670. A peer-reviewed study of Schwanhard's writings, though, showed no specific mention of fluorite and only discussion of an extremely strong acid. It was hypothesized that this was probably nitric acid or aqua regia, both capable of etching soft glass. Andreas Sigismund Marggraf made the first recorded preparation of hydrofluoric acid in 1764 when he heated fluorite with sulfuric acid in glass, which was greatly corroded by the product. In 1771, Swedish chemist Carl Wilhelm Scheele repeated this reaction. Scheele recognized the product of the reaction 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. Fluorite was then shown to be mostly composed of calcium fluoride.

Sir Humphry Davy originally suggested the name fluorine, taking the root from the name of "fluoric acid" and the -ine suffix, similarly to other halogens. This name, with modifications, came to most European languages. (Greek, Russian, and several other languages use the name ftor or derivatives, which was suggested by Ampère and comes from the Greek φθόριος (phthorios), meaning "destructive".) The New Latin name (fluorum) gave the element its current symbol, F, although the symbol Fl has been used in early papers. The symbol Fl is now used for the super-heavy element flerovium.

Owing to its extreme reactivity, elemental fluorine was not isolated until many years after the characterization of fluorite. Progress in isolating elemental fluorine was slow because its electrolytic preparation was hard to perform and because the gas reacted with most materials. The generation of elemental fluorine proved to be exceptionally dangerous, killing or blinding several early experimenters who became known as the "fluorine martyrs".

Edmond Frémy thought that passing electric current through pure hydrofluoric acid 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.

Henri Moissan, fluorine discover

French chemist Henri Moissan, formerly one of Frémy's students, continued the search. After trying many different approaches, he built on Frémy's earlier attempt by combining potassium bifluoride and 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. After 74 years of effort by many chemists, on 26 June 1886, Moissan isolated elemental fluorine. Moissan's report to the French Academy of making fluorine showed appreciation for the feat:

One can indeed make various hypotheses on the nature of the liberated gas; the simplest would be that we are in the presence of fluorine

Moissan later devised a less expensive apparatus for making fluorine: copper equipment coated with copper fluoride. In 1906, two months before his death, Moissan received the Nobel Prize in chemistry for his fluorine isolation as well as the invention of the electric arc furnace.

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.

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

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 developed as replacement refrigerants.

Industry and applications

Main article: Fluorochemical industry

The global market for fluorochemicals was about US$16 billion per year as of 2006 and was predicted to reach 2.6 million metric tons per year by 2015. Fluorite mining (the main source of fluorine) was estimated in 2003 to be a $550 million industry, extracting 4.5 million tons per year.

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.

Fluorine industry supply chain: major sources, intermediates and applications. Click for links to related articles.

Inorganic fluorides

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

About 3 kg (6.5 lb) of metspar grade fluorite, added directly to the batch, are used for every metric ton of steel made. The fluoride ions from CaF₂ lower the melt's temperature and viscosity (make the liquid runnier). Metspar is similarly used in cast iron production and for other iron-containing alloys.

Fluorite of the acidspar grade is used directly as an additive to ceramics and enamels, glass fibers and clouded glass, and cement, as well as in the outer coating of welding rods. Acidspar is primarily used for making hydrofluoric acid, which is a chemical intermediate for most fluorine-containing compounds. Significant direct uses of HF include pickling (cleaning) of steel, cracking of alkanes in the petrochemical industry, and etching of glass.

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.

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), magnesium fluoride (antireflection optical coatings), and calcium fluoride (lenses). Sodium and potassium bifluorides are significant to the chemical industry.

Fluorocarbons

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. Fluorosurfactant (materials like Scotchgard, used in durable water repellents) are a small segment in mass but are significant economically--over $1 billion yearly revenue.

Refrigerant gases

See also: Refrigerant

Traditionally chlorofluorocarbons (CFCs) were the predominant fluorinated organic chemical. CFCs are identified by a system of numbering that explains the amount of fluorine, chlorine, carbon and hydrogen in the molecules. The term Freon has been colloquially used for CFCs and similar halogenated molecules though strictly speaking this is just a DuPont brand name. Brand neutral terminology is to use "R" as the prefix. 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 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) 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).

Fluoropolymers

Main article: Fluoropolymer

As of about 2006–2007, estimates of the global fluoropolymer production varied from over 100,000 to 180,000 metric tons per year. Yearly revenue estimates ranged from over $2.5 billion to over $3.5 billion.

Polytetrafluoroethylene (PTFE) is 60–80% of the world's fluoropolymer production on a weight basis. The term Teflon is sometimes used generically for the substance, but is a DuPont brand name. The largest application for PTFE is in electrical insulation. It is an excellent dielectric and very chemically stable. 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.

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.

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.

Microscopic structure of ePTFE ("Gore-Tex")

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. 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 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.

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.

Fluorine 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 ($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.

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 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.

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 switching stations. SF₆ gas has a much higher dielectric strength than air. It is extremely inert and, compared to oil-filled switchgear, has no hazardous polychlorinated biphenyls (PCBs).

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. Nitrogen trifluoride is used for cleaning equipment at display manufacturing plants.

For making niche organofluorines and fluorine-containing pharmaceuticals, direct fluorination is usually too hard to control. Preparation of intermediate strength fluorinators from fluorine gas solves this problem. 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.

Production of fluorine gas

Fluorine cell room at F2 Chemicals, Preston, England.

Commercial producers of fluorine gas continue to use the method of electrolysis pioneered by Moissan, with some modifications in the cell design.

Potassium bifluoride forms spontaneously from potassium fluoride and the hydrogen fluoride:

HF + KF → KHF₂

A mixture with the approximate composition KF•2HF melts at 70 °C (158 °F) and is electrolyzed between 70 °C and 130 °C (160–265 °F). Potassium bifluoride increases the electrical conductivity of the solution and provides the bifluoride anion, which releases fluorine at the anode (negative part of the cell). If HF alone is electrolyzed, hydrogen forms at the cathode (positive part of the cell) and the fluoride ions remain in solution. After electrolysis, potassium fluoride remains in solution.

2 HF₂^– → H₂↑ + F₂↑ + 2 F^–

The modern version of the process uses steel containers as cathodes, while blocks of carbon are used as anodes. The carbon electrodes are similar to those used in the electrolysis of aluminium. The voltage for the electrolysis is between 8 and 12 volts.

Biological aspects

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

Main article: Biological aspects of fluorine

Fluoride is not considered an essential mineral element for mammals and humans. Small amounts of fluoride may be beneficial for bone strength, but this is an issue only in the formulation of artificial diets. See also "Fluoride deficiency".

Biologically synthesized organofluorines have been found in microorganisms and plants, but not in animals. 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; other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate. In bacteria, the enzyme adenosyl-fluoride synthase, which makes the carbon–fluorine bond, was isolated. The discovery was touted as possibly leading to biological routes for organofluorine synthesis.

Dental care

Main articles: Fluoride therapy and Water fluoridation

Fluoride ions in contact with teeth have been long thought to limit cavities by turning the forming the hydroxyapatite of teeth into less soluble fluorapatite. The more up-to-date studies show no difference in caries levels between teeth with enamel fluoridated to different degrees, while low levels of fluoride in plaque fluid and saliva definitely do help to fight the early caries. The process of fluoride absorption works only by direct contact (topical treatment). 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. Its use began in the 1940s, following studies of children in a region where water is naturally fluoridated. It is now used for about two-thirds of the U.S. population on public water systems and for about 5.7% of people worldwide. Although the best available evidence shows no association with adverse effects other than fluorosis (dental and, in worse cases, skeletal), most of which is mild, water fluoridation has been contentious for ethical, safety, and efficacy reasons, and 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. Systematic reviews in 2000 and 2007 showed significant reduction of cavities in children associated with water fluoridation.

Sodium fluoride, tin difluoride, and, most commonly, sodium monofluorophosphate, are used in toothpaste. In 1955, the first fluoride toothpaste was introduced, in the United States. Now, almost all toothpaste in developed countries is fluoridated. For example, 95% of European toothpaste contains fluoride. Gels and foams are often advised for special patient groups, particularly those undergoing radiation therapy to the head (cancer patients). The patient receives a four-minute application of a high amount of fluoride. Varnishes exist that perform a similar function, but are more quickly applied. Fluoride is also contained in prescription and non-prescription mouthwashes and is a trace component of foods manufactured from fluoridated water supplies.

Pharmaceuticals

About 20% of pharmaceuticals contain fluorine, including important drugs in many different pharmaceutical classes. One of these is Atorvastatin (Lipitor), the cholesterol-reducing drug which topped the list of bestselling drugs in the world for nearly a decade.

Fluorine is added to drug molecules as even a single atom of it can greatly change the chemical properties of the molecule in ways that are desirable. 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.^{[citation needed]} Adding fluorine to biologically active 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.

Prozac

Fluorine finds use in many steroidal drugs. Fludrocortisone (Florinef) is a mineralocorticoid (used to retain sodium and water and raise blood pressure). Dexamethasone (Decadron) and triamcinolone (Kenalog) are potent glucocorticoids (anti-inflammatories).

Prior to 1980s, antidepressants altered not only the serotonin uptake (lack of serotonin is a reason for a depression), but also altered norepinephrine uptake which caused many side effects. The first drug to only alter the serotonin uptake was fluorine-containing Prozac. It became the best-selling antidepressant, prompted the writing of the popular book Listening to Prozac and gave birth to the extensive selective serotonin reuptake inhibitor (SSRI) class of antidepressants. Many other SSRI antidepressants are fluorinated organics, including Celexa, Luvox, and Lexapro.

Quinolones are compounds that are broad-spectrum antibiotics. Most of the more recent versions (and those in current common use) are fluorinated which makes the drugs more powerful. Prominent examples include ciprofloxacin (Cipro) and levofloxacin (Levaquin). The latter was the highest selling U.S. antibiotic in 2010.

Several inhaled general anesthetic agents, including the most common ones, also contain fluorine. The first fluorinated anesthetic agent, 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.

Agrichemicals and poisons

Sign warning of poisonous sodium fluoroacetate baits

An estimated 30% of agrichemical compounds contain fluorine. Most of them are poisons, but a few stimulate the growth instead. It is expected that how often the fluorine agrichemicals will be used depends on two factors: if the synthesis reaction will be improved (to reduce the prices) and if green chemistry will be taken in account to a larger scale (fluorochemicals are more environment-friendly).

Synthetic sodium fluoroacetate has been used as an insecticide but is especially effective against mammalian pests. The name "1080" refers to the catalogue number of the poison, which became its brand name. Fluoroacetate is similar to acetate, which has a pivotal role in the Krebs cycle (a key part of cell metabolism). Fluoroacetate halts the cycle and causes cells to be deprived of energy. Several other insecticides contain sodium fluoride, which is much less toxic than fluoroacetate. Currently the compound is banned.

Another important agrichemcial is Trifluralin. It was once very important (for example, in 1998 over a half of U.S. cotton field area was coated with the chemical); however, its suspected carcinogenic properties caused some Northern European countries to ban it in 1993. Currently, the whole European Union has it banned, although there was a case intended to cancel the decision.

The currently used agrichemicals utilize another tactic: instead of being poisonous themselves, e.g., by directly affecting metabolism, they transform the metabolism so the organism produces poisonous compounds. For example, insects fed 29-fluorostigmasterol produce the fluoroacetates from it. If a fluorine is transferred to a body cell, it blocks metabolism at the position occupied.

Scanning

Main articles: Positron emission tomography and Fludeoxyglucose (18F)

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. 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. 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.

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. Fluorine-19 is commonly used in NMR study of metabolism, protein structures and conformational changes. In addition, inert fluorinated gases have the potential to be a cheap and efficient tool for imaging lung ventilation.

Hazards

Fluorine gas

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.

U.S. hazard signs for commercially transported fluorine

Hydrofluoric acid

See also: Chemical burn

Hydrogen fluoride is a gas, but upon contact with moisture, including tissue, hydrogen fluoride immediately converts to hydrofluoric acid, which is a contact poison and must be handled with extreme care, far beyond that accorded to other mineral acids. Owing to its lesser chemical dissociation in water (remaining a neutral molecule), hydrogen fluoride penetrates tissue more quickly than typical acids.

HF burns

Poisoning can occur readily through exposure of skin or eyes, or when inhaled or swallowed. Symptoms of exposure to hydrofluoric acid may not be immediately evident, with 8-hour delay for 50% HF and up to 24-hour if the concentration is smaller. Hydrogen fluoride interferes with nerve function, meaning that burns may not initially be painful. Accidental exposures can go unnoticed, delaying treatment and increasing the extent and seriousness of the injury. 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.

Once in the blood, hydrogen fluoride reacts with calcium and magnesium, resulting in electrolyte unbalance, cardiac arrhythmia, and potentially, death. Formation of insoluble calcium fluoride possibly causes both a fall in calcium serum and the strong pain associated with tissue toxicity. In some cases, exposures can lead to hypocalcemia. Burns with areas larger than 160 cm² (25 in²) can cause serious systemic toxicity from interference with blood and tissue calcium levels.

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

Fluoride ion

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

See also: Fluoride poisoning

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, Currently, most 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, which affected nearly 300 people and killed one.

Environmental concerns

Atmosphere

See also: Ozone depletion and global warming

Chlorofluorocarbons (CFCs) and bromofluorocarbons (BFCs) have been strictly regulated via a series of international agreements, the Montreal Protocol, because they deplete the ozone layer. 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, and then release chlorine and bromine to attack the ozone at those altitudes. 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.

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

Hydrochlorofluorocarbons (HCFCs) are current replacements for CFCs; HCFCs have about one tenth the ozone damaging potential (ODP) of CFCs. They were originally scheduled for elimination by 2030 in developed nations (2040 in undeveloped). In 2007, a new treaty was signed by almost all nations to move that phaseout up by ten years because HFCs, which have no chlorine and thus zero ODP, are available. Meanwhile, individual HCFCs with the highest anti-ozone depleting potential are being phased out first. For example, in 2003, HCFC-141b was phased out in the U.S. by Environmental Protection Agency regulation. Many of the other HCFCs are now being produced at a fraction of their previous production rates.

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

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. One study indicates that PFOS levels in wildlife are starting to go down because of the recent reduced production of that chemical.

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. In the body, PFOA binds to a protein, serum albumin; it has been detected in breast milk and the blood of newborns. PFOA is not metabolized by the body, but is excreted by the kidneys.

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.

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.

Compounds

Main article: Compounds of fluorine

Fluorine as straight (a) or bent (b) bridging ligands.

Fluorine's common oxidation state is −1. With other atoms, fluorine forms either ionic bonds or polar covalent bonds. Covalent bonds involving fluorine atoms are almost always single bonds, although at least two examples of a higher order bond exist.. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine has a rich chemistry including inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds.

Inorganic

Hydrogen fluoride

Main articles: Hydrogen fluoride and Hydrofluoric acid

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, especially in the context of water solutions, hydrofluoric acid. The HF molecules interact weakly through hydrogen bonds, thus creating extra clustering associations with other HF molecules. Because of this, hydrogen flouride behaves more like water than like other hydrogen halides, such as HCl. Hydrogen fluoride does not boil until 20 °C in contrast to the heavier hydrogen halides with much lower boiling points. Hydrogen fluoride is fully miscible with water (dissolves in any proportion).

Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in water solution. 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. Although hydrofluoric acid is weak, it is very corrosive, even attacking glass.

Metal fluorides

Metal fluorides share some similarities with other metal halides but are more ionic. Sometimes they are more like oxides in their bonding and crystal structures. The solubility of fluorides varies greatly but tends to decrease as the charge on the metal ion increases.

Alkali metal fluorides are very ionic and soluble and have the same structure as rock salt (NaCl). Other ionic metal fluorides show clear difference versus their chlorides. For example, thallium and silver monofluorides are soluble, while the chlorides are not. The alkaline earth metals form very ionic difluorides, such as CaF₂. However, in contrast to the chlorides, they are water insoluble. Aluminium and gold form ionic trifluorides while the trichlorides are covalent.

At larger oxidation numbers, the metal fluorides exhibit covalent character: titanium, niobium, and vanadium tetrafluorides are polymeric. They normally have low melting points or decompose easily. Beryllium fluoride (BeF₂) also exhibits significant covalent character, and forms similar structures to SiO₂ (quartz), but unlike the other alkaline earth salts, it is very soluble in water.

Higher fluorides are normally discrete molecules, which contrasts the behavior of corresponding oxides. While oxygen forms discrete molecules with only five metals (manganese heptoxide, technetium heptoxide, ruthenium tetroxide, osmium tetroxide, and iridium tetroxide), fluorine forms molecules with fifteen metals. This is because its small size and single charge as an ion allows surrounding metal atoms with more fluorines than oxygen can. These compounds are highly reactive, acting like acids. For example, platinum hexafluoride was the first compound to oxidize molecular oxygen and xenon.

Structure types in the metal fluorides
Sodium fluoride, ionic	Bismuth pentafluoride, polymeric	Rhenium heptafluoride, discrete molecule

Nonmetal fluorides

The nonmetal binary fluorides are volatile compounds. Nonmetals from period 3 and below form fluorides which are hypervalent (have more bonds than normal).

Boron trifluoride is a planar molecule, where the central boron atom has only six electrons (and thus an incomplete octet). It readily accepts a Lewis base, forming adducts with molecules that have lone-pairs such as ammonia. With another fluoride ion it completes its octet to form the relatively unreactive BF−
4 anion. Silicon tetrafluoride adopts a molecular tetrahedral structure and is weakly acidic.

Pnictogens (nitrogen's periodic table column) fluorides become more reactive as the pnictogen becomes heavier and are weak Lewis bases. The pentafluorides are much more reactive than the trifluorides, with antimony pentafluoride being the strongest Lewis acid of all charge-neutral compounds. Nitrogen is different than other pnictogens, as it forms a triflouride that stable against hydrolysis and is not a Lewis base, and does not form pentafluoride.

The chalcogens (oxygen's periodic table 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]} The tetafluorides are thermally unstable and hydrolyze. They form adducts with other (acidic) fluorides through their lone pair. Sulfur and selenium tetrafluorides are molecular while TeF₄ is a polymer. 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. In addition, several chalcogen fluorides occur which have more than chalcogen (O₂F₂, S₂F₁₀, etc.).^{[citation needed]}

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. The corresponding cations ClF+
6 and BrF+
6, are known and are extremely strong oxidizers. For the radioactive astatine only the non-volatile astatine monofluoride has been studied, but its existence is debated. 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.

Hypervalent nonmetal fluorides: some distinctive structures
The trigonal bipyramid of PF5	The see-saw of SF4	The T-shape of ClF3	The square pyramid of BrF5

Noble gas compounds

Main article: Noble gas compound

Xenon tetrafluoride, 1962

The noble gases are generally non-reactive because they have complete electronic shells, and 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. 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. Krypton, xenon's lighter homolog, also forms difluoride and a few more complicated fluorine-containing compounds; the possibility of the existence of tetrafluoride and hexafluoride has been debated. Radon, xenon's heavier homolog has been shown to readily react with fluorine to form a solid compound, generally thought to be radon difluoride, but its exact structure has not been clearly established; if radon were not as radioactive and difficult to collect, its chemistry could be at least as extensive as xenon's.

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 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. Neon is considered to be even less reactive than helium, and is not expected to form a stable compound capable of synthesis.

Organic

Main article: Organofluorine chemistry

The carbon–fluorine chemical bond in organofluorine compounds is the strongest bond in organic chemistry. This C–F bond stability along with the low polarizability of molecules containing C-F make organofluorines very stable. Fluorinated organics have similar sizes to corresponding unfluorinated molecules because of the small van der Waals radius of fluorine.

The range of organofluorine compounds is diverse, reflecting the inherent complexity of organic chemistry. A vast number of small molecules exist with varying amounts of fluorine substitution, as well as many polymers—research into particular areas is driven by the commercial value of applications.

Small molecules

Perfluorodecalin a perfluorocarbon that is a liquid at room temperature. It boils at a lower temperature than its hydrocarbon analog, decalin.

Monofluoroalkanes (alkanes with one hydrogen replaced with fluorine) may be chemically and thermally unstable, yet are soluble in many solvents; but as more fluorines are in instead of hydrogens, the stability increases, while melting and boiling points, and solubility decrease. While the densities and viscosities are increased, the dielectric constants, surface tensions, and refractive indices fall.

Partially fluorinated alkanes are the hydrofluorocarbons (HFCs). Substituting other halogens in combination with fluorine gives rise to chlorofluorocarbons (CFCs) or bromofluorocarbons (BFCs) (if some hydrogen is retained, HCFCs and the like). Properties depend on the number and identity of the halogen atoms. In general, the boiling points are elevated by combination of halogen atoms because the varying size and charge of different halogens allows more intermolecular attractions. As with fluorocarbons, chlorofluorocarbons and bromofluorocarbons are not flammable: they do not have carbon–hydrogen bonds to react and released halides quench flames.

When all hydrogens are replaced with fluorine to achieve perfluoroalkanes, 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 compounds, as opposed to perfluorocarbons, is the term used for molecules that would be perfluorocarbons—only carbon and fluorine atoms—except for having an extra functional group (even though another definition exists). They exhibit many perfluorocarbon properties (e.g. inertness, stability, non-wetting by water and oils, slipperiness). 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.

Fluorinated organic acids have higher acidity than their hydrocarbon analogues. As fluorines are added to the acid, its strength grows: consider acetic acid and its mono-, di-, and trifluoroacetic derivatives and their pK_a values (4.74, 2.66, 1.24, and 0.23). This happens because of the inductive effect of fluorines (they stabilize anions by spreading negative charge, allowing the hydrogen to be released). 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). The perfluoroalkanesulfonic acids are also very notable for their acidity. The sulfonic acid derivative, trifluoromethanesulfonic acid, is comparable to strong mineral acids.

Polymers

The complex unit structure of Nafion polymer

Fluoropolymers are similar in many regards with smaller molecules; adding fluorine to a polymer increases chemical stability and reduces flammability. Melting points are typically much higher than in the corresponding hydrocarbon polymers.

The simplest fluoroplastic is polytetrafluoroethylene (PTFE, DuPont brand Teflon), which is a simple linear chain polymer with the repeating structural unit: –CF₂–. PTFE has no hydrogens and can be thought of as 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.

Various PTFE derivatives have lower maximum usage temperatures, but have the benefit of being 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).

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.

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

See also

• Industrial gas

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 indeterminate properties.

Indexed references

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• Greenwood, N. N.; Earnshaw, A. (1998). Chemistry of the elements (2nd ed.). Butterworth Heinemann. ISBN 0-7506-3365-4. 
• Hounshell, David A.; Smith, John Kelly (1988). Science and corporate strategy: DuPont R&D, 1902–1980. Cambridge University Press. ISBN 0-521-32767-9. 
• Lewars, Errol G. (2008). Modeling marvels: Computational anticipation of novel molecules. Springer. ISBN 1-4020-6972-3. 
• Lide, David R. (2004). Handbook of chemistry and physics (84th ed.). CRC Press. ISBN 0-8493-0566-7. 
• (Russian) Lidin, P. A.; Molochko, V. A.; Andreeva, L. L. (2000). Химические свойства неорганических веществ [Chemical properties of inorganic substances]. Khimiya. ISBN 5-7245-1163-0. 
• Mackay, Kenneth Malcolm; Mackay, Rosemary Ann; Henderson, W. (2002). Introduction to modern inorganic chemistry (6th ed.). CRC Press. ISBN 0-7487-6420-8. 
• Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. ISBN 978-0-12-352651-9. Retrieved 3 March 2011. 
• Yaws, Carl L.; Braker, William (2001). "Fluorine". Matheson gas data book, Book 2001. McGraw-Hill Professional. ISBN 978-0-07-135854-5. 
• Ullmann, Franz, ed. (2005). Encyclopedia of Industrial Chemistry. Wiley-VCH. ISBN 978-3-527-30673-2. 
  • Aigueperse, Jean; Mollard, Paul; Devilliers, Didier; Chemla, Marius; Faron, Robert; Romano, Renée; Cuer, Jean Pierre. "Fluorine Compounds, Inorganic". doi:10.1002/14356007. ISBN 978-3-527-30673-2. 
  • Siegemund, Günter; Schwertweger, Werner; Feiring, Andrew; Smart, Bruce; Behr, Fred; Vogel, Herward; McKusick, Blain. "Fluorine Compounds, Organic". doi:10.1002/14356007.a11_349. 
  • Jassaud, Michael; Faron, Robert; Devilliers, Didier; Romano, René. "Fluorine". doi:10.1002/14356007.a11_293. 
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External links

• Fluorine at The Periodic Table of Videos (University of Nottingham)