From Wikipedia, the free encyclopedia
Jump to: navigation, search
Hydrogen (diatomic nonmetal)
Helium (noble gas)
Lithium (alkali metal)
Beryllium (alkaline earth metal)
Boron (metalloid)
Carbon (polyatomic nonmetal)
Nitrogen (diatomic nonmetal)
Oxygen (diatomic nonmetal)
Fluorine (diatomic nonmetal)
Neon (noble gas)
Sodium (alkali metal)
Magnesium (alkaline earth metal)
Aluminium (other metals)
Silicon (metalloid)
Phosphorus (polyatomic nonmetal)
Sulfur (polyatomic nonmetal)
Chlorine (diatomic nonmetal)
Argon (noble gas)
Potassium (alkali metal)
Calcium (alkaline earth metal)
Scandium (transition metal)
Titanium (transition metal)
Vanadium (transition metal)
Chromium (transition metal)
Manganese (transition metal)
Iron (transition metal)
Cobalt (transition metal)
Nickel (transition metal)
Copper (transition metal)
Zinc (transition metal)
Gallium (other metals)
Germanium (metalloid)
Arsenic (metalloid)
Selenium (polyatomic nonmetal)
Bromine (diatomic nonmetal)
Krypton (noble gas)
Rubidium (alkali metal)
Strontium (alkaline earth metal)
Yttrium (transition metal)
Zirconium (transition metal)
Niobium (transition metal)
Molybdenum (transition metal)
Technetium (transition metal)
Ruthenium (transition metal)
Rhodium (transition metal)
Palladium (transition metal)
Silver (transition metal)
Cadmium (transition metal)
Indium (other metals)
Tin (other metals)
Antimony (metalloid)
Tellurium (metalloid)
Iodine (diatomic nonmetal)
Xenon (noble gas)
Caesium (alkali metal)
Barium (alkaline earth metal)
Lanthanum (lanthanide)
Cerium (lanthanide)
Praseodymium (lanthanide)
Neodymium (lanthanide)
Promethium (lanthanide)
Samarium (lanthanide)
Europium (lanthanide)
Gadolinium (lanthanide)
Terbium (lanthanide)
Dysprosium (lanthanide)
Holmium (lanthanide)
Erbium (lanthanide)
Thulium (lanthanide)
Ytterbium (lanthanide)
Lutetium (lanthanide)
Hafnium (transition metal)
Tantalum (transition metal)
Tungsten (transition metal)
Rhenium (transition metal)
Osmium (transition metal)
Iridium (transition metal)
Platinum (transition metal)
Gold (transition metal)
Mercury (transition metal)
Thallium (other metals)
Lead (other metals)
Bismuth (other metals)
Polonium (other metals)
Astatine (metalloid)
Radon (noble gas)
Francium (alkali metal)
Radium (alkaline earth metal)
Actinium (actinide)
Thorium (actinide)
Protactinium (actinide)
Uranium (actinide)
Neptunium (actinide)
Plutonium (actinide)
Americium (actinide)
Curium (actinide)
Berkelium (actinide)
Californium (actinide)
Einsteinium (actinide)
Fermium (actinide)
Mendelevium (actinide)
Nobelium (actinide)
Lawrencium (actinide)
Rutherfordium (transition metal)
Dubnium (transition metal)
Seaborgium (transition metal)
Bohrium (transition metal)
Hassium (transition metal)
Meitnerium (unknown chemical properties)
Darmstadtium (unknown chemical properties)
Roentgenium (unknown chemical properties)
Copernicium (transition metal)
Ununtrium (unknown chemical properties)
Flerovium (unknown chemical properties)
Ununpentium (unknown chemical properties)
Livermorium (unknown chemical properties)
Ununseptium (unknown chemical properties)
Ununoctium (unknown chemical properties)


Fluorine in the periodic table
gas: very pale yellow
liquid: bright yellow
solid: transparent (beta), opaque (alpha)
Small sample of pale yellow liquid fluorine condensed in liquid nitrogen
Liquid fluorine at cryogenic temperatures
General properties
Name, symbol, number fluorine, F, 9
Pronunciation /ˈflʊərn/ FLUU-reen, /ˈflʊərɪn/, /ˈflɔərn/
Element category diatomic nonmetal
Group, period, block 17 (halogens), 2, p
Standard atomic weight 18.998403163(6)
Electron configuration [He] 2s2 2p51
2, 7
Physical properties
Phase gas
Density (0 °C, 101.325 kPa)
1.6962 g/L
Liquid density at b.p. 1.5053 g·cm−3
Melting point 53.48 K, −219.67 °C, −363.414 °F
Boiling point 85.03 K, −188.11 °C, −306.604 °F
Triple point 53.48 K, 904 kPa
Critical point 144.41 K, 5.17244 MPa
Heat of vaporization 6.512 kJ·mol−1
Molar heat capacity (Cp) (21.1 °C) 313 J·mol−1·K−1
(Cv) (21.1 °C) 233 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 3.981 (Pauling scale)
Ionization energies
1st: 16815 kJ·mol−1
2nd: 33745 kJ·mol−1
3rd: 61475 kJ·mol−1
Covalent radius 646 pm
Van der Waals radius 1357 pm
Crystal structure monoclinic
Fluorine has a monoclinic base-centered crystal structure

alpha state (low-temperature)8
Magnetic ordering diamagnetic (−1.2×10−4 (SI)910)
Thermal conductivity 0.0259111 W·m−1·K−1
CAS registry number 7782-41-41
Naming after the mineral fluorite, itself named after Latin fluo (to flow, in smelting)
Discovery André-Marie Ampère (1810)
First isolation Henri Moissan1 (June 26, 1886)
Named by Humphry Davy
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
· references

Fluorine is an extremely reactive and poisonous elemental gas with a pale yellow appearance. It is composed of diatomic molecules, has an atomic number of 9, and is the lightest of the halogens and most electronegative of the elements. Fluorine is the 24th most abundant element in the known universe and the 13th most abundant within the Earth's crust. It has a rich chemistry, forming compounds with nearly all other elements, including some of the noble gases.

The primary mineral source of fluorine, fluorite (calcium fluoride, CaF
), was first 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 to be difficult and dangerous to separate from its compounds; several early experimenters were killed or badly hurt in their attempts at doing so. In 1886, French chemist Henri Moissan succeeded in isolating elemental fluorine using low temperature electrolysis, a process still used for the modern industrial production of fluorine.

Because of the expense of refining the pure element, nearly all commercially used fluorine remains in compound form throughout its processing. About half of mined fluorite is used directly in steel-making. The other half is converted to hydrogen fluoride, a dangerous acid that is the precursor to many fluorochemicals. The main use of hydrogen fluoride is in the synthesis of various organic fluorides and in cryolite, an inorganic material critical to aluminium refining. Organic fluorides have very high chemical and thermal stability; their largest market segments are in refrigerant gases and—in the form of polytetrafluoroethylene (Teflon)—electrical insulation and cookware. Modern pharmaceuticals such as atorvastatin (Lipitor) and fluoxetine (Prozac) contain fluorine. The fluoride ion, when directly applied to teeth, reduces decay; for this reason it is used in toothpaste and water fluoridation. The largest current end use of free fluorine, uranium enrichment, began in World War II during the Manhattan Project. Global fluorochemical sales amount to over US$15 billion a year.

Fluorocarbon gases are generally greenhouse gases with warming potentials 100 to 10,000 times that of carbon dioxide. Sulfur hexafluoride–an insulating gas used in electrical plants–exhibits an even stronger effect, at about 20,000 times the global-warming potential of carbon dioxide. Organofluorines endure in the environment due to the strength of the carbon–fluorine bond; the potential health impact of the most persistent of these compounds is unclear. While a few plants and bacteria synthesize organofluorine poisons for defense against herbivores, fluorine has no metabolic role in mammals.


Chemical reactivity

The difluorine bond is relatively weak, with a bond energy much less than that found in Cl
or Br
and similar to the easily cleaved oxygen–oxygen bonds of peroxides.1314 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 fluorine's high electronegativity. Both the easy dissolution of difluorine and its strong bonding to other atoms make fluorine extremely reactive.14 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.215

External video
Bright flames during fluorine reactions
Fluorine reacting with caesium

Reactions of elemental fluorine with metals require varying conditions. The alkali metals, such as sodium, react explosively. The alkaline earth metals, such as calcium, react somewhat less dramatically. Most of the remaining metals—such as aluminium, iron, and copper—must be powdered to overcome passivation (protective metal fluoride layers formed in the initial exposure).14 The noble metals (gold, platinum, and the like) react least readily, requiring pure fluorine gas at 300–450 °C.16

The metalloids (boron, silicon, germanium, arsenic, antimony and tellurium), and some of the solid nonmetals (sulfur, phosphorus and selenium) burn with a flame in room temperature fluorine.1718 Hydrogen sulfide and sulfur dioxide combine readily with fluorine; the latter reaction can be explosive. Sulfuric acid reacts much more sluggishly.17

Fluorine reacts explosively with hydrogen in a manner similar to that of the alkali metals.19 Carbon, as lamp black, reacts with fluorine at room temperature to yield fluoromethane. 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 and the reaction can become explosive.20 Carbon dioxide and carbon monoxide react with fluorine at room temperature or just above.21 Organic chemicals, such as paraffins, react strongly when exposed to fluorine.22 Even fully halogenated organic molecules, such as the normally incombustible carbon tetrachloride, can explode.23 Nitrogen, with its very strong triple bonds, requires an electric discharge and very high temperatures to react with fluorine even though the end product, NF3, is quite stable.24 Ammonia (NH3) can react explosively with fluorine.2526 Oxygen does not normally react, but can be combined with fluorine under an electric discharge at low pressures and temperatures. The products tend to be unstable and separate back into fluorine and oxygen when heated.272829 The other halogens react readily with fluorine,30 as does the heavy noble gas radon.31 The lighter noble gases xenon and krypton react directly with fluorine under special conditions.32


4 diagonal placards with warnings, poison, corrosive, inhalant, oxidant
The U.S. hazard signs for commercially transported fluorine33

Fluorine is highly toxic. Its immediately dangerous to life or health concentration is 25 ppm (hydrogen cyanide, in contrast, is 50 ppm).34 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.35 Inhalation of 1,000 ppm fluorine will cause death in minutes36 (whereas for hydrogen cyanide a concentration of only 270 ppm is required).37


Cube with spherical shapes on the corners and center and spinning molecules in planes in faces
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.

At room temperature, fluorine is a gas composed of diatomic molecules.2 Though sometimes described as yellow-green, pure fluorine is actually a very pale yellow.38 It has a characteristic pungent odor that is noticeable in concentrations as low as 20 ppb.39 Fluorine condenses to a bright yellow liquid at −188 °C, which is near the condensation temperatures of oxygen and nitrogen.40

Fluorine has two solid forms: beta-fluorine and alpha-fluorine. The beta-phase crystallizes at −220 °C and is transparent and soft. It has a disordered cubic structure that is the same as that of solid oxygen when first crystallized.40note 1 With further cooling to −228 °C, fluorine undergoes a solid–solid phase transition into a form called alpha-fluorine, this being opaque and hard. It has a monoclinic structure featuring close-packed, shingled layers of molecules. The phase change from beta- to alpha-fluorine releases more energy than the melting point transition, and can be violent.note 2844 Solid fluorine is unlike the other halogens (chlorine, bromine, and iodine), all of which have orthorhombic crystalline structures.844

Electron arrangement

A bunch of rings showing valence and non-valence electron shells
Simplified structure of the fluorine atom

A neutral fluorine 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 offer much shielding from the nucleus. They therefore experience a high effective nuclear charge of seven (nine minus two), which affects the physical properties of the atom.1

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.45 Instead, fluorine exhibits a very strong preference for capturing one more electron to achieve the filled-shell electron configuration of the noble gas neon: 1s22s22p6.1 Fluorine has the highest electronegativity (a relative measure of electron attraction by atoms) among the elements.46 Its electron affinity (the energy released by adding an electron) is higher than that of any element except chlorine.47 Fluorine atoms have a small covalent radius of around 60 picometers; this is similar to the radii of oxygen and neon, its left and right hand periodic table neighbors.note 34849


One stable isotope of fluorine occurs naturally: fluorine-19, which contains ten neutrons.50 Fluorine is thus monoisotopic (having a single stable isotope) and mononuclidic (being found on Earth in only one isotope). Its monoisotopic occurrence makes it useful in uranium enrichment because UF6 molecules differ only in mass due to mass differences between U-235 and U-238. These mass differences are used to separate U-235 and U-238 via diffusion and gas centrifugation.51 Fluorine's mononuclidic (100%) abundance make it well suited to magnetic resonance imaging, since it also has a high nuclear magnetogyric rationote 4 (which translates to exceptional magnetic field sensitivity).53

Seventeen radioisotopes have been synthesized, having mass numbers 14–18 and 20–31.54 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-18.54



Abundance in the Solar System55
Element Relative
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

For the ninth lightest element, fluorine is unusually rare given 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, with the exception of sodium (which is tens of times more common).56 At 400 ppb, fluorine is estimated to be the 24th most common element in the universe.

Fluorine's rarity results from both a low stellar 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 form neon and hydrogen.5657

The presence of fluorine at all—outside its fleeting existence in stars—is somewhat of a mystery given these fluorine-eliminating reactions.5658 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 stars (a type of red giant), pulses of fusion reactions could result in convection currents lifting fluorine out of the inner star.5658


Fluorine is the thirteenth most common element in the Earth's crust, comprising between 600 and 700 ppm by mass.59 Being so reactive, any free fluorine in the atmosphere of early Earth is likely to have been removed by being bound to surface rocks.6061 It is essentially therefore found only in mineral compounds. Fluorite, fluorapatite, and cryolite are the three most industrially significant compounds.6263

Fluorite (CaF
), also called fluorspar, is the main source of commercial fluorine. It is colorful, common, and can be found worldwide. China supplies over half the world's demand; Mexico is second. The U.S. produced most global fluorite in the early 20th century but its last mine, in Illinois, closed in 1995.6364656667

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 U.S. is there significant recovery: byproducts are used to supply water fluoridation.63

Cryolite (Na
), 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 1987.63

Major fluorine-containing minerals
pink globular mass with crystal facets Long prism-like crystal, without luster, at an angle coming out of aggregate-like rock A parallelogram-shaped outline with space-filling diatomic molecules (joined circles) arranged in two layers
Fluorite Fluorapatite Cryolite

Several other minerals, such as the gemstone topaz, include fluorine among their constituents. Fluorine is not significant in seawater or brines, unlike the other halides, because the alkaline earth fluorides (for example, CaF2, MgF2) precipitate out of water.63 Commercially insignificant quantities of organofluorines have been observed in volcanic eruptions and in geothermal springs. Their ultimate origin, whether from biological sources or geological formation, is unclear.68

The possibility of small amounts of gaseous fluorine within crystals has been debated for many years. When crushed, one form of fluorite (antozonite) has a smell suggestive of fluorine.6970 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.70


Fluorine has a rich chemistry encompassing compounds formed with metals, nonmetals, and even noble gases, as well as a diverse set of organic compounds; note 571 Its common oxidation state is −1.note 6 Because of its attraction for electrons, fluorine forms many ionic compounds. Covalent bonds involving fluorine are polar and are almost always, within molecules, single bonds.note 77576

Metal fluorides

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.7778 The difluorides of the alkaline earths are also very ionic but are generally very insoluble.79 Beryllium difluoride is an exception: it exhibits some covalent character, is water soluble, and has a structure similar to SiO2 (quartz).80 Trifluorides are formed by many metals, particularly the rare earths, and are generally ionic.818283

The tetrafluorides represent a transition from ionic to covalent bonding. Zirconium and hafnium,8485 along with several actinides,86 form high melting, ionic tetrafluorides.87note 8 On the other hand, the tetrafluorides of titanium,90 vanadium,91 and niobium are polymeric.92 They melt or decompose below about 350 °C.93 The pentafluorides are even more covalently bonded, forming low dimensionality polymers or oligomeric molecules (clusters).949596

A total of thirteen metal hexafluorides have been characterized, all forming octahedredral molecules.note 9 All are volatile solids except molybdenum hexafluoride and rhenium hexafluoride (which are liquids) and tungsten hexafluoride (a gas).979899 The only definite metal heptafluoride, that of rhenium, is a low-melting molecular solid. Its structure is a distorted pentagonal bipyramid.100 The higher metal fluorides are very reactive.101

Progression of structure type with metal charge in the metal fluorides
checkerboard-like lattice of small blue and large yellow balls, going in three dimensions so that each ball has 6 nearest neighbors of opposite type a straight chain of alternating balls, violet and yellow, with violet ones linked additionally to four more yellow perpendicularly to the chain and each other Ball and stick drawing showing central violet ball with a yellow one directly above and below and then an equatorial belt of 5 surrounding yellow balls.
Sodium fluoride, ionic Bismuth pentafluoride, polymeric Rhenium heptafluoride, molecular

Hydrogen fluoride

graph showing trend-breaking water and HF boiling points: big jogs up versus a trend that is down with lower molecular weight for the other series members.
Boiling points of the hydrogen halides and hydrogen chalcogenides. Hydrogen fluoride shows a similar breaking trend to that of water.

Fluorine combines with hydrogen to make hydrogen fluoride (HF), a compound in which the HF molecules cluster weakly together via hydrogen bonds. Because of this, hydrogen fluoride behaves more like water than HCl (hydrochloric acid).102103104 It boils at a much higher temperature than the heavier hydrogen halides. HF is also fully miscible with water (that is, it dissolves in any proportion), unlike HCl, HBr, or HI.105

Water solutions of hydrogen fluoride are called hydrofluoric acid. This is a chemically weak acid, unlike the other hydrohalic acids (such as hydrochloric) which are all strong.106note 10 Although hydrofluoric acid is weak, it is very corrosive, even attacking glass.108

Nonmetal fluorides

A chlorine trifluoride molecule. The corrosivity of chlorine trifluoride is such that it can ignite asbestos, concrete, gravel, and sand, and other highly fire-retardant materials.109

The binary fluorides of the main group metalloids and nonmetals are generally volatile, covalently bonded molecules that vary greatly in their reactivities. Nonmetals from the third row of the periodic table and below can form fluorides which are hypervalent (that is, they have more bonds than normal).110

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

The simplest binary compound with carbon is carbon tetrafluoride, an inert tetrahedral molecule.note 11 The atoms below carbon, silicon and germanium, also form tetrahedral tetrafluorides:111 these are Lewis acids.112113

Nitrogen and its congeners form trifluorides of increasing reactivity and Lewis basicity. Nitrogen trifluoride is stable against hydrolysis and is not a Lewis base.114 The atoms below nitrogen form trifluorides that are weak Lewis bases and are more reactive as the atom becomes heavier.114 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.11594116

The chalcogens (elements making up the oxygen group) form a variety of fluorides. Unstable difluorides are known for oxygen (the only compound in which oxygen is at a formal oxidation state +2) as well as for sulfur and selenium. Tetrafluorides and hexafluorides are known for sulfur, selenium, and tellurium. They tend to be more stable with more fluorine atoms and a lighter central atom: sulfur hexafluoride (SF6) is extremely inert.117118

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.119 Many of the halogen fluorides are powerful fluorinators (sources of fluorine atoms). Chlorine trifluoride readily fluorinates asbestos and refractory oxides; its use in industry requires precautions similar to those for fluorine gas.120121

Noble gas compounds

black-and-white photo showing icy-looking crystals in a dish
Xenon tetrafluoride, 1962. Its synthesis, as with xenon hexafluoroplatinate, amazed many chemists.122

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 synthesis of xenon hexafluoroplatinate, the first noble gas compound.123 Since then, xenon difluoride, xenon tetrafluoride and xenon hexafluoride have been isolated, as well as various oxyfluorides.124125 Krypton, the lighter homolog of xenon, forms a difluoride and a few more complicated fluorine containing compounds.126 Radon, the heavier homolog of xenon, reacts readily with fluorine to form a solid, which is thought to be radon difluoride.127128

Some of the lightest noble gases form binary fluorides of only exceptionally limited stability. Argon reacts in extreme conditions with hydrogen fluoride to form argon fluorohydride.32 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.129 Neon is considered even less reactive than helium, and no fluorides have been even momentarily observed.130

Organic compounds

The carbon–fluorine chemical bond is the strongest bond in organic chemistry;131 organofluorines are subsequently very stable.132 With a few exceptions, the C–F bond does not exist in nature, meaning the entire field is essentially "handmade",133 with research in particular areas tending to be driven by the commercial value of applications. The range of organofluorine compounds is diverse, reflecting the inherent complexity of organic chemistry.134

Small molecules

beaker with two layers of liquid, goldfish and crab in top one. (uncaptioned, we also see coins at the bottom of the bottom layer)
Perfluorocarbon density demonstration. A beaker holds water with blue food dye (upper liquid layer) and a much more dense perfluoroheptane (a fluorocarbon) in its lower liquid layer. The two liquids cannot mix and the dye cannot dissolve in the fluorocarbon. A goldfish and a red clawed crab have been introduced into the water. The goldfish cannot penetrate the dense fluorocarbon. The crab floats at the liquid boundary with only parts of its legs penetrating the fluorocarbon fluid, unable to sink to the bottom of the beaker. Quarter coins rest at the bottom of the beaker. (The animals were rescued from their predicament after the photo was taken).

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 fluorine atoms are substituted for hydrogen atoms, the properties change: solubility in hydrocarbons decreases and stability increases. Also, melting and boiling points decrease, while density goes up.135

When all hydrogen atoms are replaced with fluorine atoms to make perfluorocarbons ("per" meaning maximum),note 12 a great difference becomes apparent: such compounds are extremely stable, and they are only subject to attack at standard conditions by sodium in liquid ammonia. They are also very insoluble, with few organic solvents capable of dissolving them.135

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).note 13137 Often this is a carboxylic acid (-CO2H) group. Perfluorinated compounds exhibit many perfluorocarbon properties such as inertness, stability, non-wetting by water and oils, and slipperiness.138 The functional group is however available for reactions. It may also help the molecule to adhere to surfaces or behave as a surfactant (a soap-like mixture).139 Fluorosurfactants can lower the surface tension of water below that achievable with hydrocarbon-based surfactants. Industry practice is to also regard fluorotelomers‒perfluorinated compounds with fluorinated backbones, but also a few non-fluorinated carbons (typically two) near the functional group‒as perfluorinated.138


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.140

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 hydrogen atoms and is the perfluoro analog of PE (polyethylene, structural unit: –CH2–). PTFE, as expected for a perfluorocarbon, has a much higher chemical and thermal stability than that of polyethylene. However, its very high melting point makes it difficult to fashion into parts.141

Skeletal chemical formula
The complex unit structure of Nafion, a fluoropolymer of high chemical inertness and with broad applications, including in fuel cells.142

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 fluorine atoms replaced with –CF3 groups. PFA (perfluoroalkoxy) has some fluorine atoms replaced with –OCF3 groups.141 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.143144

There are other fluoroplastics that are not perfluorinated; these contain some C-H groups. PVDF (that is, polyvinylidene fluoride, with the structural unit  –CF2CH2–) has half the fluorine atoms of PTFE. PVF (polyvinyl fluoride, structural unit: –CH2CHF–) has even less. Despite this, it still has many of the properties of fully fluorinated polymers.145


Early discoveries

woodcut image with top showing man at open hearth with tongs and machine bellows to the side. Bottom shows man at water-operated hammer with a sluice nearby for quenching.
Steelmaking illustration, Agricola text

Fluorite, the main source mineral of fluorine, was described in 1529 by Georgius Agricola, who related its use as a flux—an additive that helps lower melting temperature during smelting.146147note 14 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.15164152

Hydrofluoric acid was used as a glass-etching agent from the 1720s, and perhaps as early as 1670.note 15 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.155156 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).156157

In 1810, French physicist André-Marie Ampère suggested that hydrofluoric acid was a compound of hydrogen with an unknown element, analogous to chlorine.158 Fluorite was then shown to be mostly composed of calcium fluoride.154 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".159 The New Latin name (fluorum) gave the element its current symbol, F, although the symbol Fl was used in early papers.79note 16


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.note 17 Initial attempts to isolate the element were hindered by problems obtaining a suitable conducting liquid for electrolysis as well as by the extreme corrosiveness of hydrogen fluoride and of fluorine gas.154160

Edmond Frémy postulated that passing electric current through pure hydrofluoric acid (that is, hydrogen fluoride) 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). He discovered, however, that pure hydrogen fluoride would not transmit an electric current.154160161

French chemist Henri Moissan, formerly one of Frémy's students, continued the search. After trying many different approaches, he combined potassium difluoride and dry hydrogen fluoride. The mixture proved capable of conducting electricity, making electrolysis possible. However, rapid destruction of the platinum metal in his electrochemical cells stymied the quest. To continue, Moissan devised a strategy of cooling the reaction to extremely low temperatures, in a special bath, so as to slow the rate of corrosion. Moissan also made equipment that was more corrosion-resistant: containers crafted from a mixture of platinum and iridium (more chemically resistant than pure platinum) with fluorite stoppers.162160 In 1886, Moissan crowned 74 years of effort by many chemists when he isolated elemental fluorine.161163

In 1906, two months before his death, Moissan received the Nobel Prize in chemistry.note 18164 The award read:160

... 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 and DuPont formed a joint venture under the name Kinetic Chemicals, with a view to commercializing one such chlorofluorocarbon: Freon-12 (CCl
). It proved to be a marketplace success, rapidly replacing earlier more toxic refrigerants, and contributing to the growth of the overall market for kitchen refrigerators. By 1949, DuPont had bought out the joint venture and marketed several other Freon molecules.134156165166

In 1938, Teflon ((C2F4)n) was accidentally discovered by a recently hired Kinetic chemist, Roy J. Plunkett. Undertaking research on the 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 an unknown substance. Tests showed it to be polytetrafluoroethylene ('poly-' meaning 'many'). The new polymer was more resistant to corrosion and more stable at high temperatures than any other plastic. By 1941, it was being produced in significant quantities as a result of an accelerated commercialization program.134156165

Uranium hexafluoride, referred to as "hex" in the nuclear industry (U.S. Department of Energy)

Large-scale production of elemental fluorine began during World War II. Germany used high-temperature electrolysis to produce tons of chlorine trifluoride, a compound planned for use as an incendiary.167 The Manhattan Project in the U.S. used even more elemental fluorine to make uranium hexafluoride for use in uranium enrichment plants. Because UF6 is about as corrosive as fluorine itself, gaseous diffusion separation plants had to be 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 burgeoning nuclear weapons industry drove further development of fluorochemical compounds.168

Industry and applications

Fluorite mining, the main source of fluorine, was a growing industry up to 1989 when it peaked at 5.6 million metric tons of ore extracted in that year. Environmental restrictions on the use of chlorofluorocarbons subsequently reduced production, down to 3.6 million tons in 1994. Production has steadily risen since that time. In 2003, it was estimated at 4.5 million tons with a corresponding revenue of US$550 million. Subsequent market research reports estimated 2011 global fluorochemical sales at $15 billion and have predicted production figures over the period 2016‒18 in the range of 3.5 to about 5.9 million tons, with revenue of $20 billion, or more.156169170171172

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 mainly converted to hydrogen fluoride, the primary chemical intermediate for the fluorochemical industry.65156173

Fluorite Fluorapatite Hydrogen fluoride Metal smelting Glass production Fluorocarbons Sodium hexafluoroaluminate Pickling (metal) Fluorosilicic acid Alkane cracking Hydrofluorocarbon Hydrochlorofluorocarbons Chlorofluorocarbon Teflon Water fluoridation Uranium enrichment Sulfur hexafluoride Tungsten hexafluoride Phosphogypsum
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 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.65174

Cryolite (Na3AlF6) use in aluminium smelting

Most acidspar-grade fluorite is reacted with sulfuric acid to make hydrofluoric acid (HF). Significant direct uses of HF include pickling (cleaning) steel, etching glass, and cracking alkanes in the petrochemical industry.65 Acidspar-grade fluorite is also added to ceramics, enamels, glass fibers, clouded glass, cement, and the outer coating of welding rods.65

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.65175

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 the treatment of effluents in laundries.176 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 (antireflective optical coatings).6578177

Organic fluorochemicals

Organofluoride production consumes over 20% of all mined fluorite and over 40% of hydrofluoric acid. Refrigerant gases are the dominant segment. Fluoropolymers represent less than one quarter the consumption of refrigerant gases in terms of fluorine usage but are growing faster.65178 Fluorosurfactants are molecules used to make clothing and other items water-resistant. They make up a small market segment but generate over US$1 billion in yearly revenue.179

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.134180

Refrigerant gases

Halogenated molecules used in refrigeration are identified by the R-number system, which explains the amount of fluorine, chlorine, carbon, and hydrogen in each molecule.181 The DuPont brand Freon has been colloquially used for these compounds, but brand-neutral terminology uses "R" ("refrigerant") as the prefix.note 1965

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. By the early years of the 21st century, production had fallen to less than 10% of the mid-1980s peak, after most countries banned the end use of these chemicals.65

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 roughly 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).65 In response to greenhouse gas concerns, world-wide demand for HFO-1234yf, another HFC, is increasing as it has a global-warming potential that is less than one per cent that of HFC-134a.182


As of about 2006–2007, fluoropolymer volume was estimated at over 180,000 metric tons per year. The corresponding revenue estimate was over US$3.5 billion.183 The 2011 global fluoropolymer market was estimated at slightly under $6 billion in revenue and predicted to grow 6.5% per year through 2016.184 Fluoropolymers are formed by polymerizing free radicals; other hydrocarbon polymerization techniques do not work.140

SEM image with islands of white and interconnecting strings
"Gore-Tex" electron microscope image: the membrane's microstructure has islands of polymer with connecting strands.185

Polytetrafluoroethylene (PTFE) represents 60–80% of the world's fluoropolymer production on a weight basis.183 The DuPont brand Teflon is sometimes used generically for this substance.186 The largest application is in electrical insulation since PTFE is an excellent dielectric. It is also used in the chemical industry where corrosion resistance is needed: in coating pipes, tubing, and gaskets. Another major use is as architectural fabric in the form of PTFE-coated fiberglass cloth used for stadium roofs. The major consumer application is for non-stick cookware.186

When stretched with a sudden jerk, PTFE film becomes a fine-pored membrane in the form of expanded PTFE (ePTFE). The term "Gore-Tex" is sometimes used generically although this 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.186

Other fluoropolymers tend to have similar properties to PTFE, which leads to their use in electrical insulation and the chemical process industry. Unlike PTFE, these other fluoropolymers can be melt-processed. This makes them easier to work with as they can be formed into complex shapes, but they are also 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.186187

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 fuel cell material in spacecraft. Since then, this 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 instaill proton exchange membrane (PEM) fuel cells into automobiles.188189190

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.186

There are also fluorocarbon-based fire extinguishers.191


A shiny spherical drop of water on blue cloth
Fabric treated with fluorosurfactant

Fluorinated surfactants are small organofluorine molecules, principally used for water and stain resistance. As of 2006, yearly revenues for this segment were over US$1 billion. Fluorosurfactants are expensive chemicals, comparable to pharmaceutical chemicals: $200–2000 per kilogram ($90–900 per pound). Scotchgard has been a prominent brand; revenues in 2000 were over $300 million.179192193

Fluorosurfactants make up a very small part of the overall surfactant market, most of which is hydrocarbon-based and much cheaper. Some potential applications (for example, low cost paints) are unable to use fluorosurfactants because of the price impact of compounding in even small amounts of fluorosurfactant. Usage in paints represented a market of only about US$100 million as of 2006.179

Fluorine gas

For countries with available data, about 17,000 metric tons of fluorine are produced per year. Fluorine is relatively inexpensive, costing about US$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 are generally vertically integrated and entail on-site production.194

minaret-like electrical things with wires running around and fat at the bottom
SF6 transformers at a Russian railway

The largest application for elemental fluorine (up to 7,000 metric tons per year) is in the preparation of uranium hexafluoride, which is used in the production of nuclear fuels. 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.194 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 different masses of the two uranium isotopes.652

The second largest application for fluorine gas (about 6,000 metric tons per year) is in the production of 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.195

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.196197198 Nitrogen trifluoride is used for cleaning equipment.65

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
, BrF
, and IF
provide gentler fluorination, with a series of strengths, and are easier to handle. Sulfur tetrafluoride is used for making fluorinated pharmaceuticals.65

Production of fluorine gas


A machine room
Industrial fluorine cells, F2 Chemicals Ltd., Preston, England.

Modern industrial production of elemental fluorine uses Moissan's process of electrolyzing a mixture of potassium fluoride and hydrogen fluoride, but with an apparatus made of different materials. A 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.65199

Commercial temperatures are now higher than those used by Moissan. 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.156200201

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.156202 Regulator valves are made of nickel. Fluorine piping is generally made of nickel or Monel (a nickel-copper alloy).203 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.203 Some sources instead recommend systems made of nickel, Monel, and PTFE.204


In 1986, when he was preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl O. Christe realized that the chemical generation of fluorine ought to be feasible. The main idea was that some metal fluoride anions either did not have neutral counterparts (or had such counterparts that were very unstable) and that acidifying these anions could result in chemical oxidation, rather than formation of the expected molecules. The method involved, which resulted in a high yield of fluorine at atmospheric pressures, was:205

2 KMnO4 + 2 KF + 10 HF + 3 H2O2 → 2 K2MnF6 + 8 H2O + 3 O2
2 K2MnF6 + 4 SbF5 → 4 KSbF6 + 2 MnF3 + F2

Christe subsequently commented that the reactants involved 'had been known for more than 100 years and even Moissan could have come up with this scheme.'206 Up until at least 2006, references could still be found in the literature asserting that fluorine was too reactive to be able to be separated from its compounds by chemical means.207

Environmental concerns


An animation showing colored representation of ozone distribution by year, above North America, through 6 steps. It starts with a lot of ozone especially over Alaska and by 2060 is almost all gone from north to south.
NASA projection of stratospheric ozone levels over North America if CFCs had not been banned208

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.209 Predictions are that even after the CFC ban, several generations will be required 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.210211

Hydrochlorofluorocarbons (HCFCs) are current replacements for CFCs, with about one-tenth the ozone damaging potential (ODP).212 They were themselves originally scheduled for elimination by 2030 in developed nations and 2040 in undeveloped nations, 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.213 In 2007, a new treaty was signed by almost all nations to move the HCFC phaseout up to 2020.214

Fluorocarbon gases (CFCs, HFCs, and the like) are generally greenhouse gases with about 100 to 10,000 times the potency of carbon dioxide. Sulfur hexafluoride exhibits an even stronger effect, at about 20,000 times the global-warming potential of carbon dioxide.215 A notable outlier is HFO-1234yf, which has a global warming potential (GWP) of only four times that of carbon dioxide, compared with a GWP of 1,430 for HFC-134a, the current industry refrigerant standard. Global demand for, and interest it, HFO-1234yf is increasing in response to greenhouse gas concerns.182


Organofluorines endure in the environment due to the strength of the carbon–fluorine bond. Perfluoroalkyl acids (PFAAs) have attracted particular attention as persistent global contaminants. Because of their acid group, PFAAs are water soluble in low concentrations.216 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).217218219

Trace quantities of PFAAs have been detected worldwide, from polar bears in the Arctic to the global human population. Both PFOS and PFOA have 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 in areas of more human activity. There was no clear pattern of one chemical dominating, and higher amounts of PFOS were correlated to higher amounts of PFOA.217218220

The PFOS molecule. PFOS was a key ingredient in Scotchgard, until it was phased out from that product in 2000.221

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, kidneys, 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.217218222

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 establish an impact at current exposure levels.217218222

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.223 Fluorine-containing agrichemicals are measurable in farmland runoff and nearby rivers.224

Biological aspects

Natural biochemistry

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

Fluorine is not considered to be an essential mineral element for mammals or humans. Small amounts 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" pertains only to artificial diets.225226

Biologically synthesized organofluorines have been found in microorganisms and plants,68 but not in animals.227 The most common example is fluoroacetate. It is used as a defense against herbivores by at least 40 green plants in Africa, Australia and Brazil.228 Other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate.227 The enzyme adenosyl-fluoride synthase, which makes the carbon–fluorine bond, was isolated in bacteria in 2002.229


Dental care

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 prevents cavities primarily by helping teeth–which 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 confer no such benefit.230

white man holding plastic tray with brown goop in it and sticking a small stick into a black boy's open mouth
Topical fluoride treatment in Panama

Water fluoridation is the controlled addition of fluoride to a public water supply to reduce tooth decay.231 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.232233 Although the best available evidence shows no association with adverse effects other than dental fluorosis, most of which is mild,234 water fluoridation has been contentious for ethical, safety, and efficacy reasons.233 Opposition to water fluoridation exists despite its support by public health organizations.235 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.236 Reviews of the scholarly literature in 2000 and 2007 associated water fluoridation with a significant reduction of tooth decay in children.237

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 U.S. 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.236238


Fluorine atoms are present in 20% of modern pharmaceuticals.239 Replacing hydrogen with fluorine can protect drugs from degradation by metabolic enzymes and extend their active lifetimes in the body. The introduction of fluorine can alter a molecule's shape so that it binds better to its target protein.239 One pharmaceutical cholesterol-reducer atorvastatin (Lipitor), was the number one money-making drug for nearly a decade, at least up until 2011.240 The branded asthma medication Seretide, a top-ten revenue drug as of the mid-2000s, contains two active ingredients, one of which—fluticasone−is fluorinated.241

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.242 Replacing hydrogen with fluorine can protect drugs.243 Because of the considerable stability of the carbon–fluorine bond, many drugs are fluorinated in order to delay their metabolism and elimination. This allows longer times between doses.242 The introduction of fluorine can alter a molecule's shape so that it binds better to its target protein.243 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.241

large image of just a capsule with words Prozac and DISTA visible
Fluoxetine (Prozac). A 2008 meta-analysis of four antidepressants, including Prozac, concluded that their efficacy, although statistically significant, did not attain clinical significance "for any but the most severely depressed patients."244

Many modern antidepressants are fluorinated molecules that selectively limit the body's binding of serotonin (with low serotonin availability in brain cells being 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 medications to alter only serotonin uptake—and be free of most side effects of previous pharmaceuticals—was the fluorine-containing drug 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).245246

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 2010.247248249250

Fluorine also finds use in many steroidal drugs.251 Fludrocortisone (Florinef) is a mineralocorticoid (a compound used to retain sodium and water and thus raise blood pressure).252 Triamcinolone and dexamethasone are potent glucocorticoids (anti-inflammatories).252

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 of which are fluorinated ethers.253254

PET scanning

A rotating, transparent image of a human figure with targeted organs highlighted
Whole-body PET scan using fluorine-18

Fluorine-18 is routinely employed by the radiopharmaceutical industry, in particle accelerators, to produce radioactive tracers for positron emission tomography (PET) scanning. Its half-life of almost two hours is long enough to allow its transportation from the production facility to the imaging center for radiation exposure to patients. The most widely used radiopharmaceutical is fluorodeoxyglucose (FDG).255 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.256 Computer assisted tomography (CAT) can then be used for detailed imaging.257

Oxygen carrier research

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, have attracted ongoing interest as to their possible application in artificial blood, or liquid breathing.258

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.259260 One such product, Oxycyte, has been through initial clinical trials.261262 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.263264

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, and some human partial liquid ventilation trials, have been conducted.265 One effort, by Alliance Pharmaceuticals reached clinical trials but was abandoned because the results were not better than normal therapies.266

Agrichemicals and poisons

An estimated 30% of agrichemical compounds contain fluorine.267 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 those seen in fluorinated pharmaceuticals.268 Trifluralin is a prominent example, used widely in the U.S. as a weedkiller.268269 Its suspected carcinogenic properties have caused it to be banned in many European countries.270

Sodium monofluoroacetate (brand name 1080) is a commercial mammalian poison. The molecule is similar to the acetic acid molecule in vinegar but with a hydrogen atom changed out for fluorine atom (and another hydrogen atom changed out for a sodium atom). Fluoroacetate deprives cells of energy by replacing acetate in the Krebs cycle, halting a key part of cell metabolism. 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. New Zealand is the largest consumer, using it to suppress the invasive Australian common brushtail possum, which threatens the indigenous kiwi.271 This particular poison is now banned in Europe and the U.S.note 20228272

Fluorine-related hazards

Hydrofluoric acid

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 (thereby 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.273

left and right hands, two views, burned index fingers
Typical HF burns: the outward signs may not be evident for a day, at which point calcium treatments are less effective.274

Once in the blood, hydrogen fluoride reacts with calcium and magnesium, resulting in electrolyte imbalance (and potentially hypocalcemia). The consequent effect on the heart (cardiac arrhythmia) may be fatal.273 Formation of insoluble calcium fluoride also causes strong pain.275 Burns with areas larger than 160 cm2 (about the size of a person's hand) can cause serious systemic toxicity.276

Symptoms of exposure to hydrofluoric acid may not be immediately evident, with an 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, the 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.277 Hydrofluoric acid exposure is often treated with calcium gluconate, a source of calcium ions (Ca2+) that bind with the fluoride ions. Skin burns can be treated with a water wash and 2.5% calcium gluconate gel or special rinsing solutions.278279280 Because HF is absorbed, further medical treatment is necessary. Calcium gluconate may be injected or administered intravenously. Use of calcium chloride—a common laboratory reagent—in lieu of calcium gluconate is contraindicated, and may lead to severe complications. Sometimes surgical excision of tissue, or amputation, is required.276281

Fluoride ion

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.282 One fifth of the lethal dose may result in adverse health effects.283 Chronic excess fluoride consumption can lead to skeletal fluorosis, a disease of the bones that affects millions in Asia and Africa.283284

The fluoride ion is readily absorbed by the stomach and intestines. Ingested fluoride forms hydrofluoric acid in the stomach. In this form, it crosses cell membranes and then binds with calcium and interferes with various enzymes. The fluoride ion is excreted through urine. Exposure limits are based on urine testing, which determine the human body's capacity for ridding itself of fluoride ions.283285

Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluorides.286 Most current calls to poison control centers for possible fluoride poisoning come from the ingestion of fluoride-containing toothpaste.283 Malfunction of water fluoridation equipment has occurred several times, including an Alaskan incident that sickened nearly 300 people and killed one.287

Dangers from toothpaste are more serious for small children—the U.S. Centers for Disease Control and Prevention recommend children under six years of age be supervised when brushing their teeth so that they do not swallow toothpaste.288 One regional study examined a year of fluoride poisoning reports for pre-teens. Of 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.286 A larger study of American fluoride poisoning reports showed similar findings; 80% of the reports were related to children under six. Few were serious, although several hundred cases were treated at health facilities each year.289

See also


118 squares with numbers and letters in them, mostly colored gray and green, with a few numbers and words outside the boxes
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.
  1. ^ Alpha fluorine is solid and crystalline in that there is a regular pattern of repeated molecules. However, the diatomic molecules themselves are disordered within the crystal structure by random rotation. In contrast, in beta fluorine, the diatomic molecules are both fixed in location and have minimal rotational disorder. For further detail on alpha fluorine, see the 1970 structure by Pauling.41 For further detail on the concept of disorder in crystals, see the referenced general reviews.4243
  2. ^ A loud click is heard. Samples may shatter and sample windows blow out.
  3. ^ Exact comparison of the sizes of fluorine, oxygen and neon atoms is not possible because of conflicting estimates from different sources.
  4. ^ The ratio of the angular momentum to magnetic moment is called the gyromagnetic ratio: 'Certain nuclei can for many purposes be thought of as spinning round an axis like the Earth or like a top. In general the spin endows them with angular momentum and with a magnetic moment; the first because of their mass, the second because all or part of their electric charge may be rotating with the mass.'52
  5. ^ In this article, metalloids are lumped with the definite main group nonmetals because the fluoride chemistry is similar. The noble gases are treated separately. Hydrogen is discussed in the Hydrogen fluoride section; carbon in the Organic compounds section. The most recently created heavy elements have not been studied and thus are not included.
  6. ^ It differs from this value in elemental fluorine, where the atoms are bonded to each other and thus at oxidation state 0. The very unstable anions F2- and F3- with intermediate oxidation states exist at very low temperatures, decomposing at around 40 K.72 The F4+ cation and a few related species have been predicted to be stable.73 Theoretical and experimental evidence for the short-lived existence of a fluoronium ion, in which a fluorine atom bridges two carbon atoms, and is positively charged, was reported in 2013.74
  7. ^ Two cases of greater than single bonds are known: the metastable compounds boron monofluoride and nitrogen monofluoride. Fluorine may also act as a bridging ligand (less than single bond) between metals in some metal complexes. Molecules containing fluorine may also exhibit hydrogen bonding.
  8. ^ ZrF4 melts at 932 °C.88 HfF4 sublimes at 968 °C.85 UF4 melts at 1036 °C.89
  9. ^ IrF6, MoF6, OsF6, NpF6, PoF6, PuF6, PtF6, ReF6, RhF6, RuF6, TcF6, UF6, and WF6
  10. ^ For more detail, see the explanation by Clark.107
  11. ^ CF4 is formally an organic compound, but is noted here for structural comparison to SiF4 and GeF4. See the Organic compounds section for an overview of the vast number of fluorinated carbon-containing molecules.
  12. ^ The term "fluorocarbon" is defined by IUPAC as being the same as a perfluorocarbon (a molecule with only carbon and fluorine), but in regular practice the usage is blurred, so that fluorocarbon often used for fluorinated organic molecules including those not fully fluorinated, especially in a commercial context.
  13. ^ The term perfluorinated substance is also used for these molecules. In practice, terminology for classes of highly fluorinated molecules is imprecise. See the referenced summary of the terms used in the literature.136
  14. ^ Fluorite was also described by alchemist "Basilius Valentinus", supposedly in the late 1400s. However, it is alleged that "Valentinus" was a hoax as his writings were not known until about 1600.148149150
  15. ^ See the differing accounts of Partington153 and Weeks.154
  16. ^ Since 2012, the symbol Fl has been used for flerovium (element 114), an artificially synthesized transuranic element.
  17. ^ The injured included the English chemist Davy, the French chemists Joseph Louis Gay-Lussac and Louis Jacques Thénard, and the Irish chemists Thomas and George Knox. Belgian chemist Paulin Louyet and French chemist Jerome Nickles died. Moissan also experienced serious HF poisoning.154160
  18. ^ Moissan's Nobel also honored his invention of the electric arc furnace
  19. ^ The terminology is further muddled by the colloquial misuse of "Freon" to refer to now-banned CFCs, or to HFCs and HCFCs.
  20. ^ A minor allowed use in the U.S. is in the collars of sheep and cattle to kill predators such as coyotes.


  1. ^ a b c d e f Jaccaud et al. 2000, p. 381.
  2. ^ a b c d e Jaccaud et al. 2000, p. 382.
  3. ^ a b c Compressed Gas Association 1999, p. 365.
  4. ^ a b c d Haynes 2011, p. 4.121.
  5. ^ a b c Dean 1999, p. 4.6.
  6. ^ Dean 1999, p. 4.35.
  7. ^ Matsui 2006, p. 257.
  8. ^ a b c Young 1975, p. 10.
  9. ^ Mackay, Mackay & Henderson 2002, p. 72.
  10. ^ Cheng et al. 1999.
  11. ^ Yaws & Braker 2001, p. 385.
  12. ^ Chisté 2006.
  13. ^ Macomber 1996, p. 230.
  14. ^ a b c Greenwood & Earnshaw 1998, p. 804.
  15. ^ Nelson 1947.
  16. ^ Lidin, Molochko & Andreeva 2000, pp. 442–455.
  17. ^ a b Mellor 1922.
  18. ^ Sidgwick 1950, p. 701.
  19. ^ Greenwood & Earnshaw 1998, p. 844.
  20. ^ Kuriakose & Margrave 1965.
  21. ^ Hasegawa et al. 2007.
  22. ^ Lagow 1970.
  23. ^ Navarrini et al. 2012.
  24. ^ Lidin, Molochko & Andreeva 2000, p. 252.
  25. ^ Tanner Industries 2011.
  26. ^ Morrow, Perry & Cohen 1959.
  27. ^ Emeléus & Sharpe 1974, p. 111.
  28. ^ Wiberg, Wiberg & Holleman 2001, p. 457.
  29. ^ Brantley 1949, p. 26.
  30. ^ Jaccaud et al. 2000, p. 383.
  31. ^ Pitzer 1975.
  32. ^ a b Khriachtchev et al. 2000.
  33. ^ National Oceanic and Atmospheric Administration.
  34. ^ The National Institute for Occupational Safety and Health 1994.
  35. ^ Keplinger & Suissa 1968.
  36. ^ Emsley 2011, p. 179.
  37. ^ Biller 2007, p. 939.
  38. ^ Burdon, Emson & Edwards 1987.
  39. ^ Lide 2004, p. 4.12.
  40. ^ a b Dean 1999, p. 523.
  41. ^ Pauling, Keaveny & Robinson 1970.
  42. ^ Bürgi 2000.
  43. ^ Mueller 2009.
  44. ^ a b Barrett, Meyer & Wasserman 1967.
  45. ^ Dean 1999, p. 564.
  46. ^ Moore, Stanitski & Jurs 2010, p. 156.
  47. ^ Lide 2004, pp. 10.137–10.138.
  48. ^ Cordero et al. 2008.
  49. ^ Pyykkö & Atsumi 2009.
  50. ^ National Nuclear Data Center NuDat 2.1, Fluorine-19.
  51. ^ Koch et al. 2012, pp. 752, 754.
  52. ^ Vigoureux 1961.
  53. ^ Meusinger, Chippendale & Fairhurst 2012, pp. 752, 754.
  54. ^ a b National Nuclear Data Center NuDat 2.1.
  55. ^ Cameron 1973.
  56. ^ a b c d Croswell 2003.
  57. ^ Clayton 2003, pp. 101–104.
  58. ^ a b Renda et al. 2004.
  59. ^ Lodders 2003.
  60. ^ Schulze-Makuch & Irwin 2008, p. 121.
  61. ^ Haxel, Hedrick & Orris 2005.
  62. ^ Jaccaud et al. 2000, p. 384.
  63. ^ a b c d e Greenwood & Earnshaw 1998, p. 795.
  64. ^ a b Norwood & Fohs 1907, p. 52.
  65. ^ a b c d e f g h i j k l m n o Villalba, Ayres & Schroder 2008.
  66. ^ Kelly & Miller 2005.
  67. ^ Lusty et al. 2008.
  68. ^ a b Gribble 2002.
  69. ^ Richter, Hahn & Fuchs 2001, p. 3.
  70. ^ a b Schmedt, Mangstl & Kraus 2012.
  71. ^ Riedel & Kaupp 2009.
  72. ^ Wiberg, Wiberg & Holleman 2001, p. 422.
  73. ^ Schlöder & Riedel 2012.
  74. ^ Struble et al. 2013
  75. ^ Harbison 2002.
  76. ^ Edwards 1994, p. 515.
  77. ^ Katakuse et al. 1999, p. 267.
  78. ^ a b Aigueperse et al. 2000, pp. 420–422.
  79. ^ a b Storer 1864, pp. 278–280.
  80. ^ Walsh 2009, pp. 99–102, 118–119.
  81. ^ Emeléus & Sharpe 1983, pp. 89–97.
  82. ^ Babel & Tressaud 1985, pp. 91–96.
  83. ^ Einstein et al. 1967.
  84. ^ Brown et al. 2005, p. 144.
  85. ^ a b Perry 2011, p. 193.
  86. ^ Kern et al. 1994.
  87. ^ Lide 2004, pp. 4.60, 4.76, 4.92, 4.96.
  88. ^ Lide 2004, p. 4.96.
  89. ^ Lide 2004, p. 4.92.
  90. ^ Greenwood & Earnshaw 1998, p. 964.
  91. ^ Becker & Müller 1990.
  92. ^ Greenwood & Earnshaw 1998, p. 990.
  93. ^ Lide 2004, p. 4.72, 4.91, 4.93.
  94. ^ a b Greenwood & Earnshaw 1998, pp. 561–563.
  95. ^ Emeléus & Sharpe 1983, pp. 256–277.
  96. ^ Mackay, Mackay & Henderson 2002, pp. 243–244.
  97. ^ Greenwood & Earnshaw 1998, (various pages, by metal in respective chapter).
  98. ^ Lide 2004, p. 4.71, 4.78, 4.92.
  99. ^ Drews et al. 2006.
  100. ^ Greenwood & Earnshaw 1998, p. 819.
  101. ^ Bartlett 1962.
  102. ^ Pauling 1960, pp. 454–464.
  103. ^ Atkins & Jones 2007, pp. 184–185.
  104. ^ Emsley 1981.
  105. ^ Greenwood & Earnshaw 1998, pp. 812–816.
  106. ^ Wiberg, Wiberg & Holleman 2001, p. 425.
  107. ^ Clark 2002.
  108. ^ Chambers & Holliday 1975, pp. 328–329.
  109. ^ Air Products and Chemicals 2004, p. 1.
  110. ^ Noury, Silvi & Gillespie 2002.
  111. ^ Ellis 2001, p. 69.
  112. ^ Aigueperse et al. 2000, p. 423.
  113. ^ Wiberg, Wiberg & Holleman 2001, p. 897.
  114. ^ a b Raghavan 1998, pp. 164–5.
  115. ^ Godfrey et al. 1998, p. 98.
  116. ^ Aigueperse et al. 2000, p. 432.
  117. ^ Murthy, Mehdi Ali & Ashok 1995, pp. 180–182, 206–208.
  118. ^ Greenwood & Earnshaw 1998, pp. 638–640, 683–689, 767–778.
  119. ^ Wiberg, Wiberg & Holleman 2001, pp. 435-436.
  120. ^ Greenwood & Earnshaw 1998, pp. 828–830.
  121. ^ Patnaik 2007, pp. 478–479.
  122. ^ Moeller, Bailar & Kleinberg 1980, p. 236.
  123. ^ Wiberg, Wiberg & Holleman 2001, pp. 392–393.
  124. ^ Wiberg, Wiberg & Holleman 2001, p. 438.
  125. ^ Wiberg, Wiberg & Holleman 2001, p. 400.
  126. ^ Lewars 2008, p. 68.
  127. ^ Pitzer 1993, p. 111.
  128. ^ Lewars 2008, p. 67.
  129. ^ Bihary, Chaban & Gerber 2002.
  130. ^ Lewars 2008, p. 71.
  131. ^ O'Hagan 2008.
  132. ^ Siegemund et al. 2005, p. 444.
  133. ^ Sandford 2000, p. 455.
  134. ^ a b c d Okazoe 2009.
  135. ^ a b Siegemund et al. 2005, pp. 451–452.
  136. ^ Posner et al. 2013, pp. 187–190.
  137. ^ Barbee, McCormack & Vartanian 2000, p. 116.
  138. ^ a b Posner 2011, p. 27.
  139. ^ Salager 2002, p. 45.
  140. ^ a b Carlson & Scmiegel 2005, p. 3.
  141. ^ a b Carlson & Scmiegel 2005, pp. 3–4.
  142. ^ Hoogers 2004, p. 4-12.
  143. ^ Rhoades 2008, p. 2.
  144. ^ Okada et al. 1998.
  145. ^ Carlson & Scmiegel 2005, p. 4.
  146. ^ Greenwood & Earnshaw 1998, p. 790.
  147. ^ Senning 2007, p. 149.
  148. ^ Stillman 1912.
  149. ^ Principe 2012, pp. 140, 145.
  150. ^ Agricola, Hoover & Hoover 1912, footnotes and commentary, pp. xxx, 38, 409, 430, 461, 608.
  151. ^ Greenwood & Earnshaw 1998, p. 109.
  152. ^ Agricola, Hoover & Hoover 1912, preface, pp. 380–381.
  153. ^ Partington 1923.
  154. ^ a b c d e Weeks 1932.
  155. ^ Marggraf 1770.
  156. ^ a b c d e f g h Kirsch 2004, pp. 3–10.
  157. ^ Scheele 1771.
  158. ^ Ampère 1816.
  159. ^ Banks 1986, p. 11.
  160. ^ a b c d e Toon 2011.
  161. ^ a b Asimov 1966, p. 162.
  162. ^ Greenwood & Earnshaw 1998, pp. 789–791.
  163. ^ Moissan 1886.
  164. ^ Viel & Goldwhite 1993, p. 35.
  165. ^ a b Hounshell & Smith 1988, pp. 156–157.
  166. ^ DuPont 2013a.
  167. ^ Meyer 1977, p. 111.
  168. ^ Kirsch 2004, pp. 60–6.
  169. ^ Miller 2003b.
  170. ^ PRWeb 2012.
  171. ^ Bombourg 2012.
  172. ^ TMR 2013.
  173. ^ Fulton & Miller 2006, p. 471.
  174. ^ Miller 2003a.
  175. ^ Energetics, Inc. 1997, pp. 41, 50.
  176. ^ Aigueperse et al. 2000, p. 428.
  177. ^ Willey 2007, p. 113.
  178. ^ PRWeb 2010.
  179. ^ a b c Renner 2006.
  180. ^ Green et al. 1994, pp. 91–93.
  181. ^ DuPont 2013b.
  182. ^ a b Walter 2014
  183. ^ a b Buznik 2009.
  184. ^ PRWeb 2013.
  185. ^ Ebnesajjad 2000, pp. 176–7.
  186. ^ a b c d e Martin 2007, pp. 187–94.
  187. ^ DeBergalis 2004.
  188. ^ Grot 2011, pp. 1–10.
  189. ^ Ramkumar 2012, p. 567.
  190. ^ Burney 1999, pp. 111.
  191. ^ "Fluorine". Retrieved 18 February 2014. 
  192. ^ Kissa 2001, pp. 516–51.
  193. ^ Ullmann 2008, pp. 538, 543–547.
  194. ^ a b Jaccaud et al. 2000, p. 392.
  195. ^ Aigueperse et al. 2000, p. 430.
  196. ^ Jaccaud et al. 2000, pp. 391–392.
  197. ^ El-Kareh 1994, p. 317.
  198. ^ Arana et al. 2007.
  199. ^ Jaccaud et al. 2000, p. 386.
  200. ^ Jaccaud et al. 2000, pp. 384–285.
  201. ^ Greenwood & Earnshaw 1998, pp. 796–797.
  202. ^ Jaccaud et al. 2000, pp. 384–385.
  203. ^ a b Jaccaud et al. 2000, pp. 390–391.
  204. ^ Shriver & Atkins 2010, p. 427.
  205. ^ Christe 1986.
  206. ^ Christe Research Group n.d.
  207. ^ Carey 2006.
  208. ^ Beck et al. 2011.
  209. ^ Aucamp & Björn 2010, pp. 4–6, 41, 46–7.
  210. ^ Crow 2011.
  211. ^ Barry & Phillips 2006.
  212. ^ EPA 2013a.
  213. ^ EPA 2013b.
  214. ^ McCoy 2007.
  215. ^ Forster et al. 2007, pp. 212–213.
  216. ^ Giesy & Kannan 2002.
  217. ^ a b c d Steenland, Fletcher & Savitz 2010.
  218. ^ a b c d Betts 2007.
  219. ^ EPA 2012.
  220. ^ Zareitalabad et al. 2013.
  221. ^ Schwarcz 2004, p. 37.
  222. ^ a b Lau et al. 2007.
  223. ^ Lietz & Meyer 2006, pp. 7–8.
  224. ^ Ahrens 2011.
  225. ^ Nielsen 2009.
  226. ^ Olivares & Uauy 2004.
  227. ^ a b Murphy, Schaffrath & O'Hagan 2003
  228. ^ a b Proudfoot, Bradberry & Vale 2006.
  229. ^ O'Hagan et al. 2002.
  230. ^ Pizzo 2007.
  231. ^ CDC 2001.
  232. ^ Ripa 1993.
  233. ^ a b Cheng, Chalmers & Sheldon 2007.
  234. ^ Marya 2011, p. 343.
  235. ^ Armfield 2007.
  236. ^ a b Baelum, Sheiham & Burt 2008, p. 518.
  237. ^ NHMRC 2007; see Yeung 2008 for a summary.
  238. ^ Cracher 2012, p. 12.
  239. ^ a b Emsley 2011, p. 178.
  240. ^ Johnson 2011.
  241. ^ a b Swinson 2005.
  242. ^ a b Hagmann 2008.
  243. ^ a b Emsley 2011.
  244. ^ Kirsch, I.; Deacon, B. J.; Huedo-Medina, T. B.; Scoboria, A.; Moore, T. J.; Johnson, B. T. (February 2008). "Initial Severity and Antidepressant Benefits: A Meta-Analysis of Data Submitted to the Food and Drug Administration". PLoS Medicine. doi:10.1371/journal.pmed.0050045. Retrieved 16 February 2014. 
  245. ^ Mitchell 2004, pp. 37–39.
  246. ^ Preskorn 1996, chap. 2.
  247. ^ Werner et al. 2011.
  248. ^ Brody 2012.
  249. ^ Nelson et al. 2007.
  250. ^ King, Malone & Lilley 2000.
  251. ^ Parente 2001, p. 40.
  252. ^ a b Raj & Erdine 2012, p. 58.
  253. ^ Filler & Saha 2009.
  254. ^ Bégué & Bonnet-Delpon 2008, pp. 335–336.
  255. ^ Schmitz et al. 2000.
  256. ^ Bustamante & Pedersen 1977.
  257. ^ Alavi & Huang 2007, p. 41.
  258. ^ Gabriel et al. 1996.
  259. ^ Sarkar 2008.
  260. ^ Schimmeyer 2002.
  261. ^ Tasker 2008.
  262. ^ Davis 2006.
  263. ^ Gains 1998.
  264. ^ Taber 1999.
  265. ^ Shaffer, Wolfson & Clark 1992, p. 102.
  266. ^ Kacmarek et al. 2006.
  267. ^ ICIS 2006.
  268. ^ a b Theodoridis 2006.
  269. ^ EPA 1996.
  270. ^ DG Environment 2007.
  271. ^ Beasley 2002.
  272. ^ Eisler 1995.
  273. ^ a b Blodgett, Suruda & Crouch 2001.
  274. ^ Eaton 1997–
  275. ^ Hoffman et al. 2007, p. 1333.
  276. ^ a b HSM 2006.
  277. ^ Fischman 2001, pp. 458–459.
  278. ^ El Saadi et al. 1989.
  279. ^ Roblin et al. 2006.
  280. ^ Hultén et al. 2004.
  281. ^ Zorich 1991, pp. 182–3.
  282. ^ Liteplo et al. 2002, p. 100.
  283. ^ a b c d Shin & Silverberg 2013.
  284. ^ Reddy 2009.
  285. ^ Baez, Baez & Marthaler 2000.
  286. ^ a b Augenstein et al. 1991.
  287. ^ Gessner et al. 1994.
  288. ^ CDC 2013.
  289. ^ Shulman & Wells 1997.
Indexed references
Agricola, Georgius; Hoover, Herbert Clark; Hoover, Lou Henry (1912). De Re Metallica. London: The Mining Magazine. 
Ahrens, L. (2011). "Polyfluoroalkyl Compounds in the Aquatic Environment: A Review of Their Occurrence and Fate". Journal of Environmental Monitoring 13 (1): 20–31. doi:10.1039/c0em00373e. PMID 21031178.  
Aigueperse, J.; Mollard, P.; Devilliers, D.; Chemla, M.; Faron, R.; Romano, R. E.; Cue, J. P. (2000). "Fluorine Compounds, Inorganic". In Ullmann, Franz (ed.). Ullmann's Encyclopedia of Industrial Chemistry 15. Weinheim: Wiley-VCH. pp. 397–441. doi:10.1002/14356007. ISBN 3527306730.  
Air Products and Chemicals (2004). "Safetygram #39 Chlorine Triflouride". Allentown, PA: Air Products and Chemicals. Retrieved 16 February 2014. 
Alavi, Abbas; Huang, Steve S. (2007). "Positron Emission Tomography in Medicine: An Overview". In Hayat, M. A. (ed.). Cancer Imaging, Volume 1: Lung and Breast Carcinomas. Burlington, MA: Academic Press. pp. 39–44. ISBN 978-0-12-370468-9. 
Ampère, André-Marie (1816). "Suite d'une classification naturelle pour les corps simples". Annales de chimie et de physique (in French) 2: 1–5. 
Arana, L. R.; Mas, N.; Schmidt, R.; Franz, A. J.; Schmidt, M. A.; Jensen, K. F. (2007). "Isotropic Etching of Silicon in Fluorine Gas for MEMS Micromachining". Journal of Micromechanics and Microengineering 17 (2): 384. Bibcode:2007JMiMi..17..384A. doi:10.1088/0960-1317/17/2/026.  
Armfield, J. M. (2007). "When Public Action Undermines Public Health: A Critical Examination of Antifluoridationist Literature". Australia and New Zealand Health Policy 4: 25. doi:10.1186/1743-8462-4-25. PMC 2222595. PMID 18067684.  
Asimov, Isaac (1966). The Noble Gases. New York, NY: Basic Books. ISBN 978-0-465-05129-8. 
Atkins, Peter; Jones, Loretta (2007). Chemical Principles: The Quest for Insight (4th ed.). New York, NY: W. H. Freeman. ISBN 978-1-4292-0965-6. 
Aucamp, Pieter J.; Björn, Lars Olof (2010). "Questions and Answers about the Environmental Effects of the Ozone Layer Depletion and Climate Change: 2010 Update". United Nations Environmental Programme. Retrieved 14 October 2013. 
Augenstein, W. L. et al. (1991). "Fluoride Ingestion in Children: A Review of 87 cases". Pediatrics 88 (5): 907–912. PMID 1945630. 
Babel, Dietrich; Tressaud, Alain (1985). "Crystal Chemistry of Fluorides". In Hagenmuller, Paul (ed.). Inorganic Solid Fluorides: Chemistry And Physics. Orlando, FL: Academic Press. pp. 78–203. ISBN 978-0-12-412490-5. 
Baelum, Vibeke; Sheiham, Aubrey; Burt, Brian (2008). "Caries Control for Populations". In Fejerskov, Ole; Kidd, Edwina (eds.). Dental Caries: The Disease and Its Clinical Management (2nd ed.). Oxford: Blackwell Munksgaard. pp. 505–526. ISBN 978-1-4051-3889-5. 
Baez, Ramon J.; Baez, Martha X.; Marthaler, Thomas M. (2000). "Urinary Fluoride Excretion by Children 4–6 Years Old in a South Texas Community". Revista Panamericana de Salud Pública 7 (4): 242–248. doi:10.1590/S1020-49892000000400005. 
Banks, R. E. (1986). Journal of Fluorine Chemistry 33 (1–4). pp. 3–26. doi:10.1016/S0022-1139(00)85269-0. 
Bansal, Raj K. (2003). A Textbook of Organic Chemistry (4th ed.). New Delhi: New Age International Publishers. ISBN 81-224-1459-1. 
Barbee, K.; McCormack, K.; Vartanian, V. (2000). "EHS Concerns with Ozonated Water Spray Processing". In Mendicino, L. (ed.). Environmental Issues in the Electronics and Semiconductor Industries. Pennington, NJ: The Electrochemical Society. pp. 108–121. ISBN 978-1-56677-230-3. 
Barrett, C. S.; Meyer, L.; Wasserman, J. (1967). "Argon—Fluorine Phase Diagram". The Journal of Chemical Physics 47 (2): 740–743. Bibcode:1967JChPh..47..740B. doi:10.1063/1.1711946.  
Barry, Patrick L.; Phillips, Tony (26 May 2006). "Good News and a Puzzle". National Aeronautics and Space Administration. Retrieved 6 January 2012. 
Bartlett, N. (1962). "Xenon Hexafluoroplatinate (V) Xe+[PtF6". Proceedings of the Chemical Society (6): 218. doi:10.1039/PS9620000197.  
Bartlett, N.; Lohmann, D. H. (1962). "Dioxygenyl Hexafluoroplatinate (V), O2+[PtF6". Proceedings of the Chemical Society 3: 115. doi:10.1039/PS9620000097.  
Beasley, Michael (August 2002). Guidelines for the safe use of sodium fluoroacetate (1080). Wellington: Occupational Safety & Health Service, Department of Labour (New Zealand). ISBN 0-477-03664-3. Retrieved 11 November 2013. 
Beck, Jefferson; Newman, Paul; Schindler, Trent L.; Perkins, Lori (2011). "What Would have Happened to the Ozone Layer if Chlorofluorocarbons (CFCs) had not been Regulated?". National Aeronautics and Space Administration. Retrieved 15 October 2013. 
Becker, S.; Müller, B. G. (1990). "Vanadium Tetrafluoride". Angewandte Chemie International Edition in English 29 (4): 406. doi:10.1002/anie.199004061.  
Bégué, Jean-Pierre; Bonnet-Delpon, Danièle (2008). Bioorganic and Medicinal Chemistry of Fluorine. Hoboken, NJ: John Wiley & Sons. ISBN 978-0-470-27830-7. 
Berger, R.; Protsch, R. (1991). "Fluorine Dating". In Göksu, M; Oberhofer, H. Y.; Regulla, D. (eds.). Scientific Dating Methods. Dordrecht: Kluwer Academic Publishers. pp. 251–270. ISBN 978-0-7923-1461-5. 
Betts, K. S. (2007). "Perfluoroalkyl Acids: What is the Evidence Telling Us?". Environmental Health Perspectives 115 (5): A250–A256. doi:10.1289/ehp.115-a250. PMC 1867999. PMID 17520044.  
Bihary, Z.; Chaban, G. M.; Gerber, R. B. (2002). "Stability of a Chemically Bound Helium Compound in High-pressure Solid Helium". The Journal of Chemical Physics 117 (11): 5105–5108. Bibcode:2002JChPh.117.5105B. doi:10.1063/1.1506150.  
Biller, José (2007). Interface of Neurology and Internal Medicine (illustrated ed.). Philadelphia: Lippincott Williams & Wilkins. ISBN 0-7817-7906-5. 
Blodgett, D. W.; Suruda, A. J.; Crouch, B. I. (2001). "Fatal Unintentional Occupational Poisonings by Hydrofluoric Acid in the U.S". American Journal of Industrial Medicine 40 (2): 215–220. doi:10.1002/ajim.1090. 
Bombourg, Nicolas (4 July 2012). "World Fluorochemicals Market, Freedonia". Reporterlinker. Retrieved 20 October 2013. 
Brantley, L. R. (1949). "Fluorine, vol. 3, no. 1, pp. 11–18". In Roy Squires and Arthur C. Clarke, eds., Pacific Rockets: Journal of the Pacific Rocket Society. Volume 1, Number 1 to Volume 4, Number 2, Incl. Sawyer Publishing. ISBN 978-0-9794418-5-1. 
Brody, Jane E. (10 September 2012). "Popular Antibiotics May Carry Serious Side Effects". The New York Times Well blog. Retrieved 18 October 2013. 
Brown, Paul L.; Mompean, Federico J.; Perrone, Jane; Illemassène, Myriam (2005). Chemical Thermodynamics of Zirconium. Amsterdam: Elsevier B. V. ISBN 978-0-444-51803-3. 
Burdon, J.; Emson, B.; Edwards, A. J. (1987). "Is Fluorine Gas Really Yellow?". Journal of Fluorine Chemistry 34 (3–4): 471. doi:10.1016/S0022-1139(00)85188-X.  
Bürgi, H. B. (2000). "Motion and Disorder in Crystal Structure Analysis: Measuring and Distinguishing them". Annual Review of Physical Chemistry 51: 275–296. Bibcode:2000ARPC...51..275B. doi:10.1146/annurev.physchem.51.1.275. PMID 11031283.  
Burney, H. (1999). "Past, Present and Future of the Chlor-Alkali Industry". In Burney, H. S.; Furuya, N; Hine, F; Ota, K-I. Chlor-Alkali and Chlorate Technology: R. B. MacMullin Memorial Symposium. Pennington, NJ: The Electrochemical Society. pp. 105–126. ISBN 1-56677-244-3. 
Bustamante, E.; Pedersen, P. L. (1977). "High Aerobic Glycolysis of Rat Hepatoma Cells in Culture: Role of Mitochondrial Hexokinase". Proceedings of the National Academy of Sciences 74 (9): 3735–3739. Bibcode:1977PNAS...74.3735B. doi:10.1073/pnas.74.9.3735. PMC 431708. PMID 198801.  
Buznik, V. M. (2009). "Fluoropolymer Chemistry in Russia: Current Situation and Prospects". Russian Journal of General Chemistry 79 (3): 520–526. doi:10.1134/S1070363209030335.  
Cameron, A. G. W. (1973). "Abundance of the Elements in the Solar System". Space Science Review 15: 121–146. Bibcode:1973SSRv...15..121C. doi:10.1007/BF00172440.  
Carey, W. C. (2006). American Scientists. Norwell, MA: Infobase Publishing. ISBN 0-8160-5499-1. 
Carlson, D. P.; Schmiegel, W. (2000). "Fluoropolymers, Organic". In Ullmann, Franz (ed.). Ullmann's Encyclopedia of Industrial Chemistry 15. Weinheim: Wiley-VCH. pp. 495–533. doi:10.1002/14356007.a11_393. ISBN 3527306730.  
Centers for Disease Control and Prevention (2001). "Recommendations for Using Fluoride to Prevent and Control Dental Caries in the United States". MMWR Recommendations and Reports (Atlanta, GA) 50 (RR–14): 1–42. PMID 11521913. Retrieved 14 October 2013. 
Centers for Disease for Control and Prevention (10 July 2013). "Community Water Fluoridation". Retrieved 25 October 2013. 
Chambers, C.; Holliday, A. K. (1975). Modern Inorganic Chemistry: An Intermediate Text. London: Butterworth & Co. ISBN 978-0-408-70663-6. 
Cheng, H.; Fowler, D. E.; Henderson, P. B.; Hobbs, J. P.; Pascolini, M. R. (1999). "On the Magnetic Susceptibility of Fluorine". The Journal of Physical Chemistry A 103 (15): 2861–2866. doi:10.1021/jp9844720.  
Cheng, K. K.; Chalmers, I.; Sheldon, T. A. (2007). "Adding Fluoride to Water Supplies". BMJ 335: 699–702. doi:10.1136/bmj.39318.562951.BE. PMC 2001050. PMID 17916854.  
Chisté, V.; Bé, M. M. (2006). "F-18". Table de radionucléides. Gif-sur-Yvette, France: Laboratoire National Henri Becquerel. Retrieved 15 June 2011. 
Christe, Karl O. (1986). "Chemical Synthesis of Elemental Fluorine". Inorganic Chemistry (University of Southern California, Loas Angeles, CA) 25 (21): 3721–3722. doi:10.1021/ic00241a001. 
Christe Research Group (n.d.). "Chemical Synthesis of Elemental Fluorine:". Retrieved 12 January 2013. 
Clark, Jim (2002). "The Acidity of the Hydrogen Halides". chemguide.co.uk. Retrieved 15 October 2013. 
Clayton, Donald (2003). Handbook of Isotopes in the Cosmos: Hydrogen to Gallium. New York, NY: Cambridge University Press. ISBN 978-0-521-82381-4. 
Compressed Gas Association (1999). Handbook of Compressed Gases (4th ed.). Boston: Kluwer Academic Publishers. ISBN 978-0-412-78230-5. 
Cordero, B.; Gómez, V.; Platero-Prats, A. E.; Revés, M.; Echeverría, J.; Cremades, E.; Barragán, F.; Alvarez, S. (2008). "Covalent Radii Revisited". Dalton Transactions (21): 2832–2838. doi:10.1039/b801115j. 
Cracher, Connie M. (2012). "Current Concepts in Preventive Dentistry". dentalcare.com. Retrieved 14 October 2013. 
Croswell, Ken (2003). "Fluorine: An element–ary Mystery". Sky and Telescope (Berkeley, CA) (September). Retrieved 17 October 2013. 
Crow, James Mitchell (2011). "First Signs of Ozone-hole Recovery Spotted". Nature News. doi:10.1038/news.2011.293.  
Danielson, M. A.; Falke, J. J. (1996). "Use of 19F NMR to Probe Protein Structure and Conformational Changes". Annual Review of Biophysics and Biomolecular Structure 25: 163–195. doi:10.1146/annurev.bb.25.060196.001115. PMC 2899692. PMID 8800468.  
Davis, Nicole. "Better than blood". Popular Science (November 2006). Archived from the original on 4 June 2011. Retrieved 20 October 2013. 
Dean, John A. (1999). Lange's Handbook of Chemistry (15th ed.). New York: McGraw-Hill. ISBN 0-07-016190-9. 
Debergalis, M. (2004). "Fluoropolymer Films in the Photovoltaic Industry". Journal of Fluorine Chemistry 125 (8): 1255–1257. doi:10.1016/j.jfluchem.2004.05.013.  
European Commission, Directorate-General for the Environment (2007). Trifluralin (Report). Brussels: European Commission. http://www.unece.org/fileadmin/DAM/env/lrtap/TaskForce/popsxg/2008/Trifluralin_RA%20dossier_proposal%20for%20submission%20to%20the%20UNECE%20POP%20Protocol.pdf. Retrieved 14 October 2013.
Dixon, D. A.; Wang, T. H.; Grant, D. J.; Peterson, K. A.; Christe, K. O.; Schrobilgen, G. J. (2007). "Heats of Formation of Krypton Fluorides and Stability Predictions for KrF4 and KrF6 from High Level Electronic Structure Calculations". Inorganic Chemistry 46 (23): 10016–10021. doi:10.1021/ic701313h. PMID 17941630.  
Drews, T.; Supeł, J.; Hagenbach, A.; Seppelt, K. (2006). "Solid State Molecular Structures of Transition Metal Hexafluorides". Inorganic Chemistry 45 (9): 3782–3788. doi:10.1021/ic052029f. PMID 16634614.  
DuPont (2013a). "Freon". Retrieved 17 October 2013. 
DuPont (2013b). "Understanding the Refrigerant 'R' Nomenclature". Retrieved 17 October 2013. 
Eaton, Charles (1997–). "Figure hfl". E-Hand.com: The Electronic Textbook of Hand Surgery. The Hand Center (former practice of Dr. Eaton). Retrieved 28 September 2013. 
Eberle, S. H.; Berei, Klara; Vasáros, László (1985). "Chemical Behavior and Compounds of Astatine". In Kugler, Hans Karl; Keller, Cornelius (eds.). At – Astatine. Gmelin Handbook of Inorganic Chemistry (8th ed.). Berlin: Springer-Verlag. pp. 183–289. ISBN 978-3-662-05870-1. 
Ebnesajjad, Sina (2000). Fluoroplastics, Volume 1: Non-Melt Processible Fluoroplastics. Plastic Designs Library. Norwich, NY: William Andrew. ISBN 978-0-8155-1727-6. 
Edwards, Philip Neil (1994). "Use of Fluorine in Chemotherapy". In Banks, R. E.; Smart, B. E.; Tatlow, J. C. (eds.). Organofluorine Chemistry: Principles and Commercial Applications. New York, NY: Plenum Press. pp. 501–542. ISBN 978-0-306-44610-8. 
Einstein, F. W. B.; Rao, P. R.; Trotter, J.; Bartlett, N. (1967). "The Crystal Structure of Gold Trifluoride". Journal of the Chemical Society A: Inorganic, Physical, Theoretical 4: 478–482. doi:10.1039/J19670000478. 
Eisler, Ronald (1995). Sodium Monofluoroacetate (1080) Hazards to Fish, Wildlife and Invertebrates: A Synoptic Review (Report). Laurel, MD: Patuxent Environmental Science Center (U.S. National Biological Service). http://www.pwrc.usgs.gov/eisler/CHR_30_Sodium_monofluoroacetate.pdf. Retrieved 5 June 2011.
Ellis, Brian (2001). Scientific Essentialism. Cambridge: Cambridge University Press. ISBN 978-0-521-80094-5. 
El-Kareh, Badih (1994). Fundamentals of Semiconductor Processing Technology. Norwell, MA, and Dordrecht: Kluwer Academic Publishers. ISBN 978-0-7923-9534-8. 
El Saadi, M. S.; Hall, A. H.; Hall, P. K.; Riggs, B. S.; Augenstein, W. L.; Rumack, B. H. (1989). "Hydrofluoric Acid Dermal Exposure". Veterinary and Human Toxicology 31 (3): 243–247. PMID 2741315. 
Emeléus, H. J.; Sharpe, A. G. (1974). Advances in Inorganic Chemistry and Radiochemistry 16. New York, NY: Academic Press. ISBN 978-0-08-057865-1. 
Emeléus, H. J.; Sharpe, A. G. (1983). Advances in Inorganic Chemistry and Radiochemistry 27. Academic Press. ISBN 0-12-023627-3. 
Emsley, John (1981). "The Hidden Strength of Hydrogen". New Scientist 91 (1264): 291–292. 
Emsley, John (2011). Nature's Building Blocks: An A–Z Guide to the Elements (2nd ed.). Oxford: Oxford University Press. ISBN 978-0-19-960563-7. 
Energetics, Inc. (1997). Energy and Environmental Profile of the U.S. Aluminum Industry (Report). Columbia, MD. http://www1.eere.energy.gov/manufacturing/resources/aluminum/pdfs/aluminum.pdf. Retrieved 15 October 2013.
Filler, R.; Saha, R. (2009). "Fluorine in Medicinal Chemistry: A Century of Progress and a 60-year Retrospective of Selected Highlights". Future Medicinal Chemistry 1 (5): 777–791. doi:10.4155/fmc.09.65. PMID 21426080.  
Fischman, Michael L. (2001). "Semiconductor Manufacturing Hazards". In Sullivan, John B.; Krieger, Gary R. (eds.). Clinical Environmental Health and Toxic Exposures (2nd ed.). Philadelphia, PA: Lippincott Williams & Wilkins. pp. 431–465. ISBN 978-0-683-08027-8. 
Forster, P.; Ramaswamy, V; Artaxo, P.; Berntsen, T.; Betts, R.; Fahey, D. W.; Haywood, J.; Lean, J.; Lowe, D. C.; Myhre, G.; Nganga, J.; Prinn, R.; Raga, G.; Schulz, M.; Van Dorland, R. (2007). "Changes in Atmospheric Constituents and in Radiative Forcing". In Solomon, S.; Manning, M; Chen, Z.; Marquis, M; Averyt, K. B.; Tignor, M.; Miller, H. L. (eds.). Climate Change 2007: The Physical Science Basis. Contribution of Working Group I to the Fourth Assessment Report of the Intergovernmental Panel on Climate Change. Cambridge: Cambridge University. pp. 129–234. ISBN 978-0-521-70596-7. 
Fulton, Robert B.; Miller, M. Michael (2006). "Fluorspar". In Kogel, Jessica Elzea; Trivedi, Nikhil C.; Barker, James M.; Krukowski, Stanley T. Industrial Minerals & Rocks: Commodities, Markets, and Uses. Littleton, CO: Society for Mining, Metallurgy, and Exploration (U.S.). pp. 461–473. ISBN 978-0-87335-233-8. 
Gabriel, J. L.; Miller Jr, T. F.; Wolfson, M. R.; Shaffer, T. H. (1996). "Quantitative Structure-Activity Relationships of Perfluorinated Hetero-Hydrocarbons as Potential Respiratory Media". ASAIO Journal 42 (6): 968–973. doi:10.1097/00002480-199642060-00009. PMID 8959271.  
Gains, Paul (18 October 1998). "A New Threat in Blood Doping". The New York Times. Retrieved 18 October 2013. 
Gessner, B. D.; Beller, M.; Middaugh, J. P.; Whitford, G. M. (1994). "Acute Fluoride Poisoning from a Public Water System". New England Journal of Medicine 330 (2): 95–99. doi:10.1056/NEJM199401133300203. PMID 8259189.  
Giesy, J.P.; Kannan, K. (2002). "Perfluorochemical Surfactants in the Environment". Environmental Science & Technology 36 (7): 146A–152A. doi:10.1021/es022253t. 
Godfrey, S. M.; McAuliffe, C. A.; Mackie, A. G.; Pritchard, R. G. (1998). "Inorganic Derivatives of the Elements". In Norman, Nicholas C. (ed.). Chemistry of Arsenic, Antimony and Bismuth. London: Blackie Academic & Professional. pp. 67–158. ISBN 978-0-7514-0389-3. 
Green, S. W.; Slinn, D. S. L.; Simpson, R. N. F.; Woytek, A. J. (1994). "Perfluorocarbon Fluids". In Banks, R. E.; Smart, B. E.; Tatlow, J. C. (eds.). Organofluorine Chemistry: Principles and Applications. New York, NY: Plenum Press. pp. 89–119. ISBN 978-0-306-44610-8. 
Greenwood, N. N.; Earnshaw, A. (1998). Chemistry of the Elements (2nd ed.). Oxford: Butterworth Heinemann. ISBN 0-7506-3365-4. 
Gribble, G. W. (2002). "Naturally Occurring Organofluorines". In Neison (ed.), A. H. (ed.). Organofluorines. Berlin: Springer. pp. 121–136. doi:10.1007/10721878_5. ISBN 3-540-42064-9. 
Grosse, A. V.; Kirshenbaum, A. D.; Streng, A. G.; Streng, L. V. (1963). "Krypton Tetrafluoride: Preparation and Some Properties". Science 139 (3559): 1047–1048. Bibcode:1963Sci...139.1047G. doi:10.1126/science.139.3559.1047. PMID 17812982.  
Grot, Walter (2011). Fluorinated Ionomers (2nd ed.). Oxford and Waltham, MA: Elsevier. ISBN 978-1-4377-4457-6. 
Hagmann, W. K. (2008). "The Many Roles for Fluorine in Medicinal Chemistry". Journal of Medicinal Chemistry 51 (15): 4359–4369. doi:10.1021/jm800219f. PMID 18570365.  
Harbison, G. S. (2002). "The Electric Dipole Polarity of the Ground and Low-lying Metastable Excited States of NF". Journal of the American Chemical Society 124 (3): 366–367. PMID 11792193.  
Hasegawa, Y.; Otani, R.; Yonezawa, S.; Takashima, M. (2007). "Reaction Between Carbon Dioxide and Elementary Fluorine". Journal of Fluorine Chemistry 128 (1): 17–28. doi:10.1016/j.jfluchem.2006.09.002. 
Haxel, G. B.; Hedrick, J. B.; Orris, G. J. (2005). Rare Earth Elements—Critical Resources for High Technology, Fact Sheet 087-02 (Report). U.S. Geological Survey. http://pubs.usgs.gov/fs/2002/fs087-02/. Retrieved 31 January 2014.
Haynes, William M., ed. (2011). Handbook of Chemistry and Physics (92nd ed.). Boca Raton: CRC Press. ISBN 1-4398-5511-0. 
Hoffman, Robert; Nelson, Lewis; Howland, Mary; Lewin, Neal; Flomenbaum, Neal; Goldfrank, Lewis (2007). Goldfrank's Manual of Toxicologic Emergencies. New York, NY: McGraw-Hill Professional. ISBN 978-0-07-144310-4. 
Honeywell (2006). Recommended medical treatment for hydrofluoric acid exposure. Morristown, NJ: Honeywell International. Retrieved 9 January 2014. 
Hoogers, G. (2002). "Fuel Cell Components and Their Impact on Performance". In Hoogers, G. (ed.). Fuel Cell Technology Handbook. Boca Raton, FL: CRC Press. pp. 4–1–4–27. ISBN 0-8493-0877-1. 
Hounshell, David A.; Smith, John Kelly (1988). Science and Corporate Strategy: DuPont R & D, 1902–1980. Cambridge: Cambridge University Press. ISBN 0-521-32767-9. 
Hultén, P.; Höjer, J.; Ludwigs, U.; Janson, A. (2004). "Hexafluorine vs. Standard Decontamination to Reduce Systemic Toxicity After Dermal Exposure to Hydrofluoric Acid". Clinical Toxicology 42 (4): 355–361. doi:10.1081/CLT-120039541. PMID 15461243.  
Hwang, I.-C.; Seppelt, K. (2001). "Gold Pentafluoride: Structure and Fluoride Ion Affinity". Angewandte Chemie International Edition 40 (19): 3690–3693. doi:10.1002/1521-3773(20011001)40:19<3690::AID-ANIE3690>3.0.CO;2-5. 
ICIS (2 October 2006). "Fluorine's Treasure Trove". Surrey: Reed Business Information. Retrieved 24 October 2013. 
Jaccaud, M.; Faron, R.; Devilliers, D.; Romano, R. (2000). "Fluorine". In Ullmann, Franz (ed.). Ullmann's Encyclopedia of Industrial Chemistry 15. Weinheim: Wiley-VCH. pp. 381–395. doi:10.1002/14356007.a11_293. ISBN 3527306730.  
Johnson, Linda A. (28 December 2011). "Against Odds, Lipitor Became World's Top Seller". The Boston Globe. Retrieved 24 October 2013. 
Kacmarek, R. M.; Wiedemann, H. P.; Lavin, P. T.; Wedel, M. K.; Tütüncü, A. S.; Slutsky, A. S. (2006). "Partial Liquid Ventilation in Adult Patients with Acute Respiratory Distress Syndrome". American Journal of Respiratory and Critical Care Medicine 173 (8): 882–889. doi:10.1164/rccm.200508-1196OC. PMID 16254269.  
Katakuse, Itsuo; Ichihara, Toshio; Ito, Hiroyuki; Sakurai, Tohru; Matsuo, Takekiyo (1999). "SIMS Experiment". In Arai, T; Mihama, K; Yamamoto, K; Sugano, S. Mesoscopic Materials and Clusters: Their Physical and Chemical Properties. Tokoyo: Kodansha. pp. 259–273. ISBN 4-06-208635-2. 
Kelly, T. D.; Miller, M. M. (2005). "Historical Fluorspar Statistics". U.S. Geological Service. Retrieved 10 February 2014. 
Keplinger, M. L.; Suissa, L. W. (1968). "Toxicity of Fluorine Short-Term Inhalation". American Industrial Hygiene Association Journal 29 (1): 10–18. doi:10.1080/00028896809342975. PMID 5667185.  
Kern, S.; Hayward, J.; Roberts, S.; Richardson, J. W.; Rotella, F. J.; Soderholm, L.; Cort, B.; Tinkle, M.; West, M.; Hoisington, D.; Lander, G. A. (1994). "Temperature Variation of the Structural Parameters in Actinide Tetrafluorides". The Journal of Chemical Physics 101 (11): 9333–9337. Bibcode:1994JChPh.101.9333K. doi:10.1063/1.467963.  
Khriachtchev, L.; Pettersson, M.; Runeberg, N.; Lundell, J.; Räsänen, M. (2000). "A Stable Argon Compound". Nature 406 (6798): 874–876. doi:10.1038/35022551. PMID 10972285.  
King, D. E.; Malone, R.; Lilley, S. H. (2000). "New Classification and Update on the Quinolone Antibiotics". American Family Physician 61 (9): 2741–2748. PMID 10821154. Retrieved 8 October 2013. 
Kirsch, Peer (2004). Modern Fluoroorganic Chemistry: Synthesis, Reactivity, Applications. Weinheim: Wiley-VCH. ISBN 978-3-527-30691-6. 
Kissa, Erik (2001). Fluorinated Surfactants and Repellents (2nd ed.). New York, NY: Marcel Dekker. ISBN 978-0-8247-0472-8. 
Koch, Günter; Träger, Siegfried; Max, Arthur; Krebs, Wolf-Dieter; Stoll, Wolfgang; Heit, Werner; Warnecke, Ernst; Brennecke, Peter; Merz, Erich (2012). "Nuclear Technology, 3. Fuel Cycle". In Ullmann, Franz. Ullmann's Encyclopedia of Industrial Chemistry 24. Weinheim: Wiley-VCH. pp. 741–822. doi:10.1002/14356007.o17_o06. 
Kuethe, D. O.; Caprihan, A.; Fukushima, E.; Waggoner, R. A. (1998). "Imaging Lungs Using Inert Fluorinated Gases". Magnetic Resonance in Medicine 39 (1): 85–88. doi:10.1002/mrm.1910390114. PMID 9438441.  
Kuriakose, A. K.; Margrave, J. L. (1965). "Kinetics of the Reactions of Elemental Fluorine. IV. Fluorination of Graphite". Journal of Physical Chemistry 69 (8): 2772–2775. doi:10.1021/j100892a049.  
Kylstra, J. A. (1977). The Feasibility of Liquid Breathing in Man. Durham, NC: Duke University Medical Center. Retrieved 15 October 2013. 
Lagow, R. J. (1970). The Reactions of Elemental Fluorine; A New Approach to Fluorine Chemistry (PhD thesis, Rice University, TX). Ann Arbor, MI: UMI. 
Lau, C.; Anitole, K.; Hodes, C.; Lai, D.; Pfahles-Hutchens, A.; Seed, J. (2007). "Perfluoroalkyl Acids: A Review of Monitoring and Toxicological Findings". Toxicological Sciences 99 (2): 366–394. doi:10.1093/toxsci/kfm128. PMID 17519394.  
Lewars, Errol G. (2008). Modeling Marvels: Computational Anticipation of Novel Molecules. Dordrecht: Springer. ISBN 1-4020-6972-3. 
Lide, David R. (2004). Handbook of Chemistry and Physics (84th ed.). Boca Raton, Florida: CRC Press. ISBN 0-8493-0566-7. 
Lidin, R.; Molochko, V.A.; Andreeva, L.L. (2000). Химические свойства неорганических веществ Chemical Properties of Inorganic Substances (in Russian). Khimiya. ISBN 5-7245-1163-0. 
Lietz, A. C.; Meyer, Michael T. (2006). Evaluation of Emerging Contaminants of Concern at the South District Waste Water Treatment Plant Based on Seasonal Sampling Events, Miami-Dade Country, Florida, 2004 (Report). U.S. Geological Survey Scientific Investigations. http://pubs.usgs.gov/sir/2006/5240/pdf/sir2006-5240.pdf. Retrieved 6 June 2011.
Liteplo, R.; Gomes, R.; Howe, P.; Malcolm, H. (2002). Environmental Health Criteria 227 (Fluoride). Geneva: United Nations Environment Programme; International Labour Organization; World Health Organization. ISBN 92-4-157227-2. Retrieved 14 October 2013. 
Lodders, Katharina (2003). "Solar System Abundances and Condensation Temperatures of the Elements". The Astrophysical Journal 591 (2): 1220–1247. Bibcode:2003ApJ...591.1220L. doi:10.1086/375492. 
Lusty, P. A. J.; Brown, T. J.; Ward, J.; Bloomfield, S. (2008). "The Need for Indigenous Fluorspar Production in England". Nottingham: British Geological Survey. Retrieved 13 October 2013. 
Mackay, Kenneth Malcolm; Mackay, Rosemary Ann; Henderson, W. (2002). Introduction to Modern Inorganic Chemistry (6th ed.). Cheltenham: Nelson Thornes. ISBN 0-7487-6420-8. 
Macomber, Roger (1996). Organic chemistry 1. Sausalito, CA: University Science Books. ISBN 978-0-935702-90-3. 
Marggraf, Andreas Sigismun (1770). "Observation concernant une volatilisation remarquable d'une partie de l'espece de pierre, à laquelle on donne les noms de flosse, flüsse, flus-spaht, et aussi celui d'hesperos; laquelle volatilisation a été effectuée au moyen des acides" [Observation of a remarkable volatilization of part of a type of stone to which one gives the name flosse, flüsse, flus-spaht, as well as that of hesperos; which volatilization was effected by means of acids]. Mémoires de l'Académie royale des sciences et belles-lettres (in French) XXIV: 3–11. 
Martin, John W., ed. (2007). Concise Encyclopedia of the Structure of Materials. Oxford and Amsterdam: Elsevier. ISBN 978-0-08-045127-5. 
Marya, C. M. (2011). A Textbook of Public Health Dentistry. New Delhi: Jaypee Brothers Medical Publishers. ISBN 978-93-5025-216-1. 
Matsui, M (2006). "Fluorine-containing Dyes". In Kim, Sung-Hoon (ed.). Functional dyes. Orlando, FL: Academic Press. pp. 257–266. ISBN 978-0-12-412490-5. 
Moissan, Henri (1886). "Action d'un courant électrique sur l'acide fluorhydrique anhydre". Comptes rendus hebdomadaires des séances de l'Académie des sciences (in French) 102: 1543–1544. Retrieved 9 October 2013. 
McCoy, M. (2007). "SURVEY Market Challenges Dim the Confidence of the World's Chemical CEOs". Chemical & Engineering News 85 (23): 11. doi:10.1021/cen-v085n023.p011a.  
Mellor, J. W. (1922). A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Volume I. London and New York, NY: Longmans, Green and Co. 
Meusinger, Reinhard; Chippendale, A. Margaret; Fairhurst, Shirley A. (2012). "Nuclear Magnetic Resonance and Electron Spin Resonance Spectroscopy". In Ullmann, Franz. Ullmann's Encyclopedia of Industrial Chemistry 24. Weinheim: Wiley-VCH. pp. 609–660. doi:10.1002/14356007.b05_471. 
Meyer, Eugene (1977). Chemistry of Hazardous Materials. Englewood Cliffs, NJ New York: Prentice Hall. ISBN 978-0-13-129239-0. 
Miller, M. Michael (2003a). "Fluorspar". U.S. Geological Survey Minerals Yearbook. Reston, VA: U.S. Geological Survey. pp. 27.1–27.12. Retrieved 13 October 2013. 
Miller, M. Michael (2003b). "Mineral Resource of the Month, Fluorspar". U.S. Geological Survey. Retrieved 24 October 2013. 
Mitchell, E. Siobhan (2004). Antidepressants. New York, NY: Chelsea House Publishers. ISBN 978-1-4381-0192-7. 
Moeller, T.; Bailar, J. C.; Kleinberg (1980). Chemistry, with Inorganic Qualitative Analysis (3rd ed.). New York: Academic Press. ISBN 0-12-503350-8. 
Moore, John W.; Stanitski, Conrad L.; Jurs, Peter C. (2010). Principles of Chemistry: The Molecular Science. Belmont, CA: Brooks/Cole. ISBN 978-0-495-39079-4. 
Morrow, S. I.; Perry, D. D.; Cohen, M. S. (1959). "The Formation of Dinitrogen Tetrafluoride in the Reaction of Fluorine and Ammonia". Journal of the American Chemical Society 81 (23): 6338–6339. doi:10.1021/ja01532a066.  
Mueller, Peter (2009). 5.067 Crystal Structure Refinement. Cambridge, MA: MIT OpenCourseWare. Retrieved 13 October 2013. 
Murphy, C. D.; Schaffrath, C.; O'Hagan, D. (2003). "Fluorinated Natural Products: The Biosynthesis of Fluoroacetate and 4-Fluorothreonine in Streptomyces cattleya". Chemosphere 52 (2): 455–461. doi:10.1016/S0045-6535(03)00191-7. PMID 12738270.  
Murthy, C. Parameshwara; Mehdi Ali, S. F.; Ashok, D. (1995). University Chemistry, Volume I. New Delhi: New Age International. ISBN 978-81-224-0742-6. 
National Health and Medical Research Council (2007). "A Systematic Review of the Efficacy and Safety of Fluoridation, Part A: Review of Methodlogy and Results". Canberra: Australian Government. ISBN 1-86496-421-9. Retrieved 8 October 2013. 
The National Institute for Occupational Safety and Health (1994). "Fluorine". Documentation for Immediately Dangerous To Life or Health Concentrations (IDLHs). Washington, DC. Retrieved 15 January 2014. 
National Nuclear Data Center. "NuDat 2.1 Database". Upton, NY: Brookhaven National Laboratory. Retrieved 25 October 2013. 
National Oceanic and Atmospheric Administration. "UN/NA 1045 (United Nations/North America Fluorine Data Sheet)". Silver Spring, MD. Retrieved 15 October 2013. 
Navarrini, W.; Venturini, F.; Tortelli, V.; Basak, S.; Pimparkar, K. P.; Adamo, A.; Jensen, K. F. (2012). "Direct Fluorination of Carbon Monoxide in Microreactors". Journal of Fluorine Chemistry 142: 19–23. doi:10.1016/j.jfluchem.2012.06.006.  
Nelson, Eugene W. (1947). "'Bad Man' of The Elements". Popular Mechanics 88 (2): 106–108, 260. 
Nelson, John H. (2003). Nuclear Magnetic Resonance Spectroscopy. Upper Saddle River, NJ: Prentice Hall. ISBN 978-0-13-033451-0. 
Nelson, J. M.; Chiller, T. M.; Powers, J. H.; Angulo, F. J. (2007). "Food Safety: Fluoroquinolone‐ResistantCampylobacterSpecies and the Withdrawal of Fluoroquinolones from Use in Poultry: A Public Health Success Story". Clinical Infectious Diseases 44 (7): 977–980. doi:10.1086/512369. PMID 17342653.  
Nielsen, F. H. (2009). "Micronutrients in Parenteral Nutrition: Boron, Silicon, and Fluoride". Gastroenterology 137 (5): S55–S60. doi:10.1053/j.gastro.2009.07.072. PMID 19874950.  
Norwood, Charles J.; Fohs, F. Julius (1907). Kentucky Geological Survey, Bulletin No. 9: Fluorspar Deposits of Kentucky. Lexington, KY: Kentucky Geological Survey. 
Noury, S.; Silvi, B.; Gillespie, R. J. (2002). "Chemical Bonding in Hypervalent Molecules: Is the Octet Rule Relevant?". Inorganic Chemistry 41 (8): 2164–2172. doi:10.1021/ic011003v. PMID 11952370. Retrieved 23 May 2012.  
Okada, T.; Xie, G.; Gorseth, O.; Kjelstrup, S.; Nakamura, N.; Arimura, T. (1998). "Ion and Water Transport Characteristics of Nafion Membranes as Electrolytes". Electrochimica Acta 43 (24): 3741–3747. doi:10.1016/S0013-4686(98)00132-7.  
Okazoe, T. (2009). "Overview on the History of Organofluorine Chemistry from the Viewpoint of Material Industry". Proceedings of the Japan Academy, Series B 85 (8): 276–289. Bibcode:2009PJAB...85..276O. doi:10.2183/pjab.85.276.  
Olivares, M.; Uauy, R. (2004). Essential Nutrients in Drinking Water (Draft) (Report). Geneva: World Health Organization. http://www.who.int/water_sanitation_health/dwq/en/nutoverview.pdf. Retrieved 14 October 2013.
Oxtoby, David W.; Gillis, H. P.; Campion, Alan (2012). Principles of Modern Chemistry (7th ed.). Belmont, CA: Brooks/Cole. ISBN 978-0-8400-4931-5. 
O'Hagan, D. (2008). "Understanding Organofluorine Chemistry. An Introduction to the C–F Bond". Chemical Society Reviews 37 (2): 308–319. doi:10.1039/b711844a. PMID 18197347.  
O'Hagan, D.; Schaffrath, C.; Cobb, S. L.; Hamilton, J. T. G.; Murphy, C. D. (2002). "Biochemistry: Biosynthesis of an Organofluorine Molecule". Nature 416 (6878): 279. Bibcode:2002Natur.416..279O. doi:10.1038/416279a. PMID 11907567.  
Parente, Luca (2001). "The Development of Synthetic Glucocorticoids". In Goulding, Nicolas J.; Flower, Rod J. (ed.). Glucocorticoids. Basel: Birkhäuser. pp. 35–53. ISBN 978-3-7643-6059-7. 
Partington, J. R. (1923). "The early history of hydrofluoric acid". Memoirs and Proceedings of the Manchester Literary and Philosophical Society 67 (6): 73–87. 
Patnaik, Pradyot (2007). A Comprehensive Guide to the Hazardous Properties of Chemical Substances (3rd ed.). Hoboken, NJ: John Wiley & Sons. ISBN 978-0-471-71458-3. 
Pauling, Linus (1960). The Nature of the Chemical Bond (3rd ed.). Ithaca, NY: Cornell University Press. ISBN 978-0-8014-0333-0. 
Pauling, L.; Keaveny, I.; Robinson, A. B. (1970). "The Crystal Structure of α-Fluorine". Journal of Solid State Chemistry 2 (2): 225–227. Bibcode:1970JSSCh...2..225P. doi:10.1016/0022-4596(70)90074-5.  
Perry, Dale L. (2011). Handbook of Inorganic Compounds (2nd ed.). Boca Raton, FL: CRC Press. ISBN 978-1-4398-1461-1. 
Pitzer, K. S. (1975). "Fluorides of Radon and Element 118". Journal of the Chemical Society, Chemical Communications (18): 760b–761. doi:10.1039/C3975000760B.  
Pitzer, Kenneth S., ed. (1993). Molecular Structure and Statistical Thermodynamics: Selected Papers of Kenneth S. Pitzer. Singapore: World Scientific Publishing. ISBN 978-981-02-1439-5. 
Pizzo, G.; Piscopo, M. R.; Pizzo, I.; Giuliana, G. (2007). "Community Water Fluoridation and Caries Prevention: A Critical Review". Clinical Oral Investigations 11 (3): 189–193. doi:10.1007/s00784-007-0111-6. PMID 17333303.  
Posner, Stefan (2011). "Perfluorinated Compounds: Occurrence and Uses in Products". In Knepper, Thomas P.; Large, Frank T. (eds.). Polyfluorinated Chemicals and Transformation Products. Heidelberg: Springer Science+Business Media. pp. 25–40. ISBN 978-3-642-21871-2. 
Posner, Stefan, et al. (2013). Per- and Polyfluorinated Substances in the Nordic Countries: Use Occurrence and Toxicology. Copenhagen: Nordic Council of Ministers. doi:10.6027/TN2013-542. ISBN 978-92-893-2562-2. 
Preskorn, Sheldon H. (1996). Clinical Pharmacology of SSRI's. Caddo, OK: Professional Communications. ISBN 978-1-884735-08-0. 
Principe, Lawrence M. (2012). The Secrets of Alchemy. Chicago, IL: University of Chicago Press. ISBN 978-0-226-68295-2. 
Proudfoot, A. T.; Bradberry, S. M.; Vale, J. A. (2006). "Sodium Fluoroacetate Poisoning". Toxicological Reviews 25 (4): 213–219. doi:10.2165/00139709-200625040-00002. PMID 17288493.  
PRWeb (28 October 2010). "Global Fluorochemicals Market to Exceed 2.6 Million Tons by 2015, According to a New Report by Global Industry Analysts, Inc.". prweb.com. Retrieved 24 October 2013. 
PRWeb (23 February 2012). "Global Fluorspar Market to Reach 5.94 Million Metric Tons by 2017, According to New Report by Global Industry Analysts, Inc.". prweb.com. Retrieved 24 October 2013. 
PRWeb (7 April 2013). "Fluoropolymers Market is Poised to Grow at a CAGR of 6.5% & to Reach $9,446.0 Million by 2016 – New report by MarketsandMarkets". prweb.com. Retrieved 24 October 2013. 
Pyykkö, P.; Atsumi, M. (2009). "Molecular Double-Bond Covalent Radii for Elements Li–E112". Chemistry - A European Journal 15 (46): 12770–12779. doi:10.1002/chem.200901472.  
Raghavan, P. S. (1998). Concepts and Problems in Inorganic Chemistry. Delhi: Discovery Publishing House. ISBN 978-81-7141-418-5. 
Raj, P. Prithvi; Erdine, Serdar (2012). Pain-Relieving Procedures: The Illustrated Guide. Chichester: John Wiley & Sons. ISBN 978-0-470-67038-5. 
Ramkumar, Jayshree (2012). "Nafion Persulphonate Membrane: Unique Properties and Various Applications". In Banerjee, S.; Tyagi, A. K. (ed.). Functional Materials: Preparation, Processing and Applications. London and Waltham, MA: Elsevier. pp. 549–578. ISBN 978-0-12-385142-0. 
Reddy, D. (2009). "Neurology of Endemic Skeletal Fluorosis". Neurology India 57 (1): 7–12. doi:10.4103/0028-3886.48793. PMID 19305069.  
Reiche, I. (2006). "Fluorine and Its Relevance for Archaeological Studies". In Tressaud, Alain, ed. Fluorine and the Environment - Agrochemicals, Archaeology, Green Chemistry & Water. Advances in Fluorine Science 2. pp. 253–283. doi:10.1016/S1872-0358(06)02008-2. ISBN 9780444526724.  
Remy, Heinrich (1956). Treatise on Inorganic Chemistry: Introduction and Main Groups of the Periodic Table. Amsterdam: Elsevier. 
Renda, A.; Fenner, Y.; Gibson, B. K.; Karakas, A. I.; Lattanzio, J. C.; Campbell, S.; Chieffi, A.; Cunha, K.; Smith, V. V. (2004). "On the Origin of Fluorine in the Milky Way". Monthly Notices of the Royal Astronomical Society 354 (2): 575–581. arXiv:astro-ph/0410580. Bibcode:2004MNRAS.354..575R. doi:10.1111/j.1365-2966.2004.08215.x.  
Renner, R. (2006). "The Long and the Short of Perfluorinated Replacements". Environmental Science & Technology 40: 12–13. Bibcode:2006EnST...40...12R. doi:10.1021/es062612a. PMID 16433328.  
Rhoades, David Walter (2008). Broadband Dielectric Spectroscopy Studies of Nafion. PhD dissertation, University of Southern Mississippi, MS. Ann Arbor, MI: ProQuest. ISBN 978-0-549-78540-8. 
Richter, M.; Hahn, O.; Fuchs, R. (2001). "Purple Fluorite: A Little Known Artists' Pigment and Its Use in Late Gothic and Early Renaissance Painting in Northern Europe". Studies in Conservation 46 (1): 1–13. JSTOR 1506878. 
Riedel, S.; Kaupp, M. (2009). "The Highest Oxidation States of the Transition Metal Elements". Coordination Chemistry Reviews 253 (5–6): 606–624. doi:10.1016/j.ccr.2008.07.014.  
Ripa, L. W. (2008). "A Half-century of Community Water Fluoridation in the United States: Review and Commentary". Journal of Public Health Dentistry 53 (1): 17–44. doi:10.1111/j.1752-7325.1993.tb02666.x. PMID 8474047.  
Roblin, I.; Urban, M.; Flicoteau, D.; Martin, C.; Pradeau, D. (2006). "Topical Treatment of Experimental Hydrofluoric Acid Skin Burns by 2.5% Calcium Gluconate". Journal of Burn Care & Research 27 (6): 889–894. doi:10.1097/01.BCR.0000245767.54278.09. PMID 17091088.  
Salager, Jean-Louis (2002). Surfactants: Types and Uses. FIRP Booklet # 300-A. Bogotá, Colombia: Laboratory of Formulation, Interfaces, Rheology, and Processes, Universidad de los Andes. Retrieved 13 October 2013. 
Sandford, Graham (2000). "Organofluorine Chemistry". Philosophical Transactions 358: 455–471. doi:10.1098/rsta.2000.0541. 
Sarkar, S. (2008). "Artificial Blood". Indian Journal of Critical Care Medicine 12 (3): 140–144. doi:10.4103/0972-5229.43685. PMC 2738310. PMID 19742251.  
Scheele, Carl Wilhelm (1771). "Undersŏkning om fluss-spat och dess syra" [Investigation of Fluorite and Its Acid]. Kungliga Svenska Vetenskapsademiens Handlingar [Proceedings of the Royal Swedish Academy of Science] (in Swedish) 32: 129–138. 
Schimmeyer, S. (2002). "The Search for a Blood Substitute". Illumin (Columbia, SC: University of Southern Carolina) 15 (1). Retrieved 15 October 2013. 
Schlöder, T.; Riedel, S. (2012). "Investigation of Heterodimeric and Homodimeric Radical Cations of the Series: [F2O2+, [F2Cl2+, [Cl2O2+, [F4+, and [Cl4+". RSC Advances 2 (3): 876–881. doi:10.1039/C1RA00804H.  
Schmedt, J.; Mangstl, M.; Kraus, F. (2012). "Occurrence of Difluorine F2 in Nature—In Situ Proof and Quantification by NMR Spectroscopy". Angewandte Chemie International Edition 51 (31): 7847–7849. doi:10.1002/ange.201203515.  
Schmitz, A.; Kälicke, T.; Willkomm, P.; Grünwald, F.; Kandyba, J.; Schmitt, O. (2000). "Use of Fluorine-18 Fluoro-2-deoxy-D-glucose Positron Emission Tomography in Assessing the Process of Tuberculous Spondylitis". Journal of Spinal Disorders 13 (6): 541–544. doi:10.1097/00002517-200012000-00016. PMID 11132989. Retrieved 8 October 2013.  
Schulze-Makuch, D.; Irwin, L. N. (2008). Life in the Universe: Expectations and Constraints (2nd ed.). Berlin: Springer-Verlag. ISBN 978-3-540-76816-6. 
Schwarcz, Joseph A. (2004). The Fly in the Ointment: 70 Fascinating Commentaries on the Science of Everyday Life. Toronto: ECW Press. ISBN 1-55022-621-5. 
Senning, A. (2007). Elsevier's Dictionary of Chemoetymology: The Whies and Whences of Chemical Nomenclature and Terminology. Amsterdam and Oxford: Elsevier. ISBN 978-0-444-52239-9. 
Shaffer, T. H.; Wolfson, M. R.; Clark Jr, L. C. (1992). "Liquid Ventilation". Pediatric Pulmonology 14 (2): 102–109. doi:10.1002/ppul.1950140208. PMID 1437347.  
Shin, Richard D.; Silverberg, Mark A. (2013). "Fluoride Toxicity". New York: Medscape. Retrieved 15 October 2013. 
Shorafa, H.; Seppelt, K. (2006). "Osmium(VII) Fluorine Compounds". Inorganic Chemistry 45 (19): 7929–7934. doi:10.1021/ic0608290. PMID 16961386.  
Shriver, Duward; Atkins, Peter (2010). Solutions Manual for Inorganic Chemistry. New York, NY: W. H. Freeman. ISBN 978-1-4292-5255-3. 
Shulman, J. D.; Wells, L. M. (1997). "Acute Fluoride Toxicity from Ingesting Home-use Dental Products in Children, Birth to 6 Years of Age". Journal of Public Health Dentistry 57 (3): 150–158. doi:10.1111/j.1752-7325.1997.tb02966.x. PMID 9383753.  
Sidgwick, N. V. (1950). The chemical elements and their compounds 1. Oxford: Clarendon Press. p. 701. 
Siegemund, G. N.; Schwertfeger, W.; Feiring, A.; Smart, B.; Behr, F.; Vogel, H.; McKusick, B. (2000). "Fluorine Compounds, Organic". In Ullmann, Franz, ed. Ullmann's Encyclopedia of Industrial Chemistry 15. Weinheim: Wiley-VCH. pp. 443–494. doi:10.1002/14356007.a11_349. ISBN 3527306730.  
Snow, T. P.; Destree, J. D.; Jensen, A. G. (2007). "The Abundance of Interstellar Fluorine and Its Implications". The Astrophysical Journal 655: 285–298. arXiv:astro-ph/0611066. Bibcode:2007ApJ...655..285S. doi:10.1086/510187.  
Snow, T. P.; York, D. G. (1981). "The Detection of Interstellar Fluorine in the Line of Sight Toward Delta Scorpii". The Astrophysical Journal 247: L39–L41. Bibcode:1981ApJ...247L..39S. doi:10.1086/183585.  
Steenland, K.; Fletcher, T.; Savitz, D. A. (2010). "Epidemiologic Evidence on the Health Effects of Perfluorooctanoic Acid (PFOA)". Environmental Health Perspectives 118 (8): 1100–1108. doi:10.1289/ehp.0901827. PMC 2920088. PMID 20423814.  
Stillman, John Maxson (December 1912). "Basil Valentine, A Seventeenth Century Hoax". Popular Science Monthly 81. Retrieved 14 October 2013. 
Storer, Frank H. (1864). First Outlines of a Dictionary of Solubilities of Chemical Substances. Cambridge: Sever and Francis. 
Struble, M. D.; Scerba, M. T.; Siegler, M.; Lectka, T. (2013). "Evidence for a Symmetrical Fluoronium Ion in Solution". Science 340 (6128): 57–60. doi:10.1126/science.1231247. 
Swinson, Joel (2005). "Fluorine – A Vital Element in the Medicine Chest". PharmaChem (Pharmaceutical Chemistry) (June): 26–27. Retrieved 9 October 2013. 
Taber, Andrew (22 April 1999). "Dying to ride". Salon. Retrieved 18 October 2013. 
Tanner Industries (January 2011). "Anhydrous Ammonia: (MSDS) Material Safety Data Sheet". Southampton, MA: tannerind.com. Retrieved 24 October 2013. 
Tasker, Fred (19 March 2008). "Miami Herald: Artificial blood goes from science fiction to science fact". Miami Herald (archived via noblood.org). Retrieved 24 October 2013. 
Theodoridis, George (2006). "Fluorine-Containing Agrochemicals: An Overview of Recent Developments". In Tressaud, Alain (ed.). Fluorine and the Environment : Agrochemicals, Archaeology, Green Chemistry & Water. Amsterdam and Oxford: Elsevier. pp. 121–176. ISBN 978-0-444-52672-4. 
Transparency Market Research (17 May 2013). "Fluorochemicals Market is Expected to Reach USD 21.5 Billion Globally by 2018: Transparency Market Research". Transparency Market Research Blog. Retrieved 15 October 2013. 
Toon, Richard (2011). "Fluorine, An Obsession with a Tragic Past". Education in Chemistry 48 (5): 148–151. 
Ullmann, Fritz (2008). Ullmann's Fibers (2 volumes). Weinheim: Wiley-VCH. ISBN 978-3-527-31772-1. 
United States Environmental Protection Agency (1996). "R.E.D. Facts: Trifluralin". Washington, DC. Retrieved 17 October 2013. 
United States Environmental Protection Agency (2008). "Ozone depletion glossary". Washington, DC. Retrieved 15 October 2013. 
United States Environmental Protection Agency (2010). "Brief questions and answers on ozone depletion | Ozone layer protection". Washington, DC. Retrieved 15 October 2013. 
United States Environmental Protection Agency (2012). "Emerging Contaminants – Perfluorooctane Sulfonate (PFOS) and Perfluorooctanoic Acid (PFOA)". Washington, DC. Retrieved 4 November 2013. 
United States Environmental Protection Agency (2013a). "Class I Ozone-depleting Substances". Washington, DC. Retrieved 15 October 2013. 
United States Environmental Protection Agency (2013b). "Phaseout of HCFCs (Class II Ozone-Depleting Substances)". Washington, DC. Retrieved 15 October 2013. 
Viel, Claude; Goldwhite, Harold (1993). "1906 Nobel Laureate: Henri Moissan, 1852–1907". In Laylin, K. James (ed.). Nobel Laureates in Chemistry, 1901–1992. Washington, DC: American Chemical Society; Chemical Heritage Foundation. pp. 35–41. ISBN 978-0-8412-2690-6. 
Vigoureux, P. (1961). "The Gyromagnetic Ratio of the Proton". Contemporary Physics 2 (5): 360–366. doi:10.1080/00107516108205282. 
Villalba, G.; Ayres, R. U.; Schroder, H. (2008). "Accounting for Fluorine: Production, Use, and Loss". Journal of Industrial Ecology 11: 85–101. doi:10.1162/jiec.2007.1075.  
Walsh, Kenneth A. (2009). Beryllium Chemistry and Processing. Materials Park, OH: ASM International. ISBN 978-0-87170-721-5. 
Walter, P. (2014). "Honeywell Invests $300m in Green Refigerant". Chemistry World 11 (2): 17. doi:10.1021/es062612a. PMID 16433328. 
Weeks, M. E. (1932). "The Discovery of the Elements. XVII. The Halogen Family". Journal of Chemical Education 9 (11): 1915–1939. Bibcode:1932JChEd...9.1915W. doi:10.1021/ed009p1915.  
Weinstock, B.; Malm, J. G. (1958). "Osmium Hexafluoride and its Identity with the Previously Reported Octafluoride". Journal of the American Chemical Society 80 (17): 4466–4468. doi:10.1021/ja01550a007.  
Werner, N. L.; Hecker, M. T.; Sethi, A. K.; Donskey, C. J. (2011). "Unnecessary use of Fluoroquinolone Antibiotics in Hospitalized Patients". BMC Infectious Diseases 11: 187–193. doi:10.1186/1471-2334-11-187. PMC 3145580. PMID 21729289.  
Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 978-0-12-352651-9. 
Willey, Ronald R. (2007). Practical Equipment, Materials, and Processes for Optical Thin Films. Charlevoix, MI: Willey Optical. ISBN 978-0-615-14397-2. 
Yaws, Carl L.; Braker, William (2001). "Fluorine". Matheson Gas Data Book (7th ed.). Parsippany, NJ: Matheson Tri-Gas. ISBN 978-0-07-135854-5. 
Yeung, C. A. (2008). "A Systematic Review of the Efficacy and Safety of Fluoridation". Evidence-Based Dentistry 9 (2): 39–43. doi:10.1038/sj.ebd.6400578. PMID 18584000.  
Young, David A. (1975). Phase Diagrams of the Elements (Report). CA: Lawrence Livermore Laboratory. http://www.osti.gov/bridge/servlets/purl/4010212-0BbwUC/4010212.pdfaccess. Retrieved 10 June 2011.
Zareitalabad, P.; Siemens, J.; Hamer, M.; Amelung, W. (2013). "Perfluorooctanoic Acid (PFOA) and Perfluorooctanesulfonic Acid (PFOS) in Surface Waters, Sediments, Soils and Wastewater – A Review on Concentrations and Distribution Coefficients". Chemosphere 91 (6): 725–732. doi:10.1016/j.chemosphere.2013.02.024. PMID 23498059.  
Zhang, Y.; Liu, X. -W. (2005). "Fluorine Abundances in Planetary Nebulae". The Astrophysical Journal 631: L61–L64. arXiv:astro-ph/0508339. Bibcode:2005ApJ...631L..61Z. doi:10.1086/497113.  
Zorich, Robert (1991). Handbook of Quality Integrated Circuit Manufacturing. San Diego, CA: Academic Press. ISBN 978-0-323-14055-3.