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=                              Fluorine                              =
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                            Introduction
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Fluorine is a chemical element; it has symbol F and atomic number 9.
It is the lightest halogen and exists at standard conditions as pale
yellow diatomic gas. Fluorine is extremely reactive as it reacts with
all other elements except for the light noble gases. It is highly
toxic.

Among the elements, fluorine ranks 24th in cosmic abundance and 13th
in crustal abundance. Fluorite, the primary mineral source of
fluorine, which gave the element its name, was first described in
1529; as it was added to metal ores to lower their melting points for
smelting, the Latin verb  meaning  gave the mineral its name. Proposed
as an element in 1810, fluorine proved difficult and dangerous to
separate from its compounds, and several early experimenters died or
sustained injuries from their attempts. Only in 1886 did French
chemist Henri Moissan isolate elemental fluorine using low-temperature
electrolysis, a process still employed for modern production.
Industrial production of fluorine gas for uranium enrichment, its
largest application, began during the Manhattan Project in World War
II.

Owing to the expense of refining pure fluorine, most commercial
applications use fluorine compounds, with about half of mined fluorite
used in steelmaking. The rest of the fluorite is converted into
hydrogen fluoride en route to various organic fluorides, or into
cryolite, which plays a key role in aluminium refining. The
carbon-fluorine bond is usually very stable. Organofluorine compounds
are widely used as refrigerants, electrical insulation, and PTFE
(Teflon). Pharmaceuticals such as atorvastatin and fluoxetine contain
C−F bonds. The fluoride ion from dissolved fluoride salts inhibits
dental cavities and so finds use in toothpaste and water fluoridation.
Global fluorochemical sales amount to more than US$15 billion a year.

Fluorocarbon gases are generally greenhouse gases with global-warming
potentials 100 to 23,500 times that of carbon dioxide, and SF6 has the
highest global warming potential of any known substance.
Organofluorine compounds often persist in the environment due to the
strength of the carbon-fluorine bond. Fluorine has no known metabolic
role in mammals; a few plants and marine sponges synthesize
organofluorine poisons (most often monofluoroacetates) that help deter
predation.


Electron configuration
========================
Fluorine atoms have nine electrons, one fewer than neon, and electron
configuration 1s22s22p5: two electrons in a filled inner shell and
seven in an outer shell requiring one more to be filled. The outer
electrons are ineffective at nuclear shielding, and experience a high
effective nuclear charge of 9 − 2 = 7; this affects the atom's
physical properties.

Fluorine's first ionization energy is third-highest among all
elements, behind helium and neon, which complicates the removal of
electrons from neutral fluorine atoms. It also has a high electron
affinity, second only to chlorine, and tends to capture an electron to
become isoelectronic with the noble gas neon; it has the highest
electronegativity of any reactive element. Fluorine atoms have a small
covalent radius of around 60 picometers, similar to those of its
period neighbors oxygen and neon.


Reactivity
============
The bond energy of difluorine is much lower than that of either  or
and similar to the easily cleaved peroxide bond; this, along with high
electronegativity, accounts for fluorine's easy dissociation, high
reactivity, and strong bonds to non-fluorine atoms. Conversely, bonds
to other atoms are very strong because of fluorine's high
electronegativity. Unreactive substances like powdered steel, glass
fragments, and asbestos fibers react quickly with cold fluorine gas;
wood and water spontaneously combust under a fluorine jet.

Reactions of elemental fluorine with metals require varying
conditions. Alkali metals cause explosions and alkaline earth metals
display vigorous activity in bulk; to prevent passivation from the
formation of metal fluoride layers, most other metals such as
aluminium and iron must be powdered, and noble metals require pure
fluorine gas at 300 -. Some solid nonmetals (sulfur, phosphorus) react
vigorously in liquid fluorine. Hydrogen sulfide and sulfur dioxide
combine readily with fluorine, the latter sometimes explosively;
sulfuric acid exhibits much less activity, requiring elevated
temperatures.

Hydrogen, like some of the alkali metals, reacts explosively with
fluorine. Carbon, as lamp black, reacts at room temperature to yield
tetrafluoromethane. Graphite combines with fluorine above 400 °C to
produce non-stoichiometric carbon monofluoride; higher temperatures
generate gaseous fluorocarbons, sometimes with explosions. Carbon
dioxide and carbon monoxide react at or just above room temperature,
whereas paraffins and other organic chemicals generate strong
reactions: even completely substituted haloalkanes such as carbon
tetrachloride, normally incombustible, may explode. Although nitrogen
trifluoride is stable, nitrogen requires an electric discharge at
elevated temperatures for reaction with fluorine to occur, due to the
very strong triple bond in elemental nitrogen; ammonia may react
explosively. Oxygen does not combine with fluorine under ambient
conditions, but can be made to react using electric discharge at low
temperatures and pressures; the products tend to disintegrate into
their constituent elements when heated. Heavier halogens react readily
with fluorine as does the noble gas radon; of the other noble gases,
only xenon and krypton react, and only under special conditions. Argon
does not react with fluorine gas; however, it does form a compound
with fluorine, argon fluorohydride.


Phases
========
At room temperature, fluorine is a gas of diatomic molecules, pale
yellow when pure (sometimes described as yellow-green). It has a
characteristic halogen-like pungent and biting odor detectable at 20
ppb. Fluorine condenses into a bright yellow liquid at -188 °C, a
transition temperature similar to those of oxygen and nitrogen.

Fluorine has two solid forms, α- and β-fluorine. The latter
crystallizes at -220 °C and is transparent and soft, with the same
disordered cubic structure of freshly crystallized solid oxygen,
unlike the orthorhombic systems of other solid halogens. Further
cooling to -228 °C induces a phase transition into opaque and hard
α-fluorine, which has a monoclinic structure with dense, angled layers
of molecules. The transition from β- to α-fluorine is more exothermic
than the condensation of fluorine, and can be violent.


Isotopes
==========
Only one isotope of fluorine occurs naturally in abundance, the stable
isotope . It has a high magnetogyric ratio and exceptional sensitivity
to magnetic fields; because it is also the only stable isotope, it is
used in magnetic resonance imaging. Eighteen radioisotopes with mass
numbers 13-31 have been synthesized, of which Fluorine-18 is the most
stable with a half-life of 109.734 minutes.  is a natural trace
radioisotope produced by cosmic ray spallation of atmospheric argon as
well as by reaction of protons with natural oxygen: 18O + p → 18F + n.
[http://www.scopenvironment.org/downloadpubs/scope50 SCOPE 50 -
Radioecology after Chernobyl] , the Scientific Committee on Problems
of the Environment (SCOPE), 1993. See table 1.9 in Section 1.4.5.2.
Other radioisotopes have half-lives less than 70 seconds; most decay
in less than half a second. The isotopes  and  undergo β+ decay and
electron capture, lighter isotopes decay by proton emission, and those
heavier than  undergo β− decay (the heaviest ones with delayed neutron
emission). Two metastable isomers of fluorine are known, , with a
half-life of 162(7) nanoseconds, and , with a half-life of 2.2(1)
milliseconds.


Universe
==========
Solar System abundances  Atomic   number         Element         Relative   amount
6       Carbon  4,800
7       Nitrogen        |1,500
8       Oxygen  |8,800  9       Fluorine        |1
10      Neon    |1,400
11      Sodium  |24
12      Magnesium       |430
Among the lighter elements, fluorine's abundance value of 400 ppb
(parts per billion) - 24th among elements in the universe - is
exceptionally low: other elements from carbon to magnesium are twenty
or more times as common. This is because stellar nucleosynthesis
processes bypass fluorine, and any fluorine atoms otherwise created
have high nuclear cross sections, allowing collisions with hydrogen or
helium to generate oxygen or neon respectively.

Beyond this transient existence, three explanations have been proposed
for the presence of fluorine:
* during type II supernovae, bombardment of neon atoms by neutrinos
could transmute them to fluorine;
* the solar wind of Wolf-Rayet stars could blow fluorine away from any
hydrogen or helium atoms; or
* fluorine is borne out on convection currents arising from fusion in
asymptotic giant branch stars.


Earth
=======
Fluorine is the 13th most abundant element in Earth's crust at 600-700
ppm (parts per million) by mass. Though believed not to occur
naturally, elemental fluorine has been shown to be present as an
occlusion in antozonite, a variant of fluorite. Most fluorine exists
as fluoride-containing minerals. Fluorite, fluorapatite and cryolite
are the most industrially significant. Fluorite (), also known as
fluorspar, abundant worldwide, is the main source of fluoride, and
hence fluorine. China and Mexico are the major suppliers. Fluorapatite
(Ca5(PO4)3F), which contains most of the world's fluoride, is an
inadvertent source of fluoride as a byproduct of fertilizer
production. Cryolite (), used in the production of aluminium, is the
most fluorine-rich mineral. Economically viable natural sources of
cryolite have been exhausted, and most is now synthesised
commercially.


File:Fluorite-270246.jpg|Fluorite: Pink globular mass with crystal
facets
File:Apatite Canada.jpg|Fluorapatite: Long, prismatic crystal, dull in
lustre, protruding, at an angle, from matrix of aggregate-like rock
File:Ivigtut cryolite edit.jpg|Cryolite: A parallelogram-shaped
outline with diatomic molecules arranged in two layers


Other minerals such as topaz contain fluorine. Fluorides, unlike other
halides, are insoluble and do not occur in commercially favorable
concentrations in saline waters. Trace quantities of organofluorines
of uncertain origin have been detected in volcanic eruptions and
geothermal springs. The existence of gaseous fluorine in crystals,
suggested by the smell of crushed antozonite, is contentious; a 2012
study reported the presence of 0.04%  by weight in antozonite,
attributing these inclusions to radiation from the presence of tiny
amounts of uranium.


Early discoveries
===================
In 1529, Georgius Agricola described fluorite as an additive used to
lower the melting point of metals during smelting. He penned the Latin
word 'fluorēs' ('fluor,' flow) for fluorite rocks. The name later
evolved into 'fluorspar' (still commonly used) and then 'fluorite'.
The composition of fluorite was later determined to be calcium
difluoride.

Hydrofluoric acid was used in glass etching from 1720 onward. Andreas
Sigismund Marggraf first characterized it in 1764 when he heated
fluorite with sulfuric acid, and the resulting solution corroded its
glass container. Swedish chemist Carl Wilhelm Scheele repeated the
experiment in 1771, and named the acidic product 'fluss-spats-syran'
(fluorspar acid). In 1810, the French physicist André-Marie Ampère
suggested that hydrogen and an element analogous to chlorine
constituted hydrofluoric acid. He also proposed in a letter to Sir
Humphry Davy dated August 26, 1812 that this then-unknown substance
may be named 'fluorine' from fluoric compounds and the '-ine' suffix
of other halogens. This word, often with modifications, is used in
most European languages; however, Greek, Russian, and some others,
following Ampère's later suggestion, use the name 'ftor' or
derivatives, from the Greek φθόριος ('phthorios', destructive). The
New Latin name 'fluorum' gave the element its current symbol F; Fl was
used in early papers.


Isolation
===========
1887 drawing of Moissan's apparatus
Initial studies on fluorine were so dangerous that several
19th-century experimenters were deemed "fluorine martyrs" after
misfortunes with hydrofluoric acid. Isolation of elemental fluorine
was hindered by the extreme corrosiveness of both elemental fluorine
itself and hydrogen fluoride, as well as the lack of a simple and
suitable electrolyte. Edmond Frémy postulated that electrolysis of
pure hydrogen fluoride to generate fluorine was feasible and devised a
method to produce anhydrous samples from acidified potassium
bifluoride; instead, he discovered that the resulting (dry) hydrogen
fluoride did not conduct electricity. Frémy's former student Henri
Moissan persevered, and after much trial and error found that a
mixture of potassium bifluoride and dry hydrogen fluoride was a
conductor, enabling electrolysis. To prevent rapid corrosion of the
platinum in his electrochemical cells, he cooled the reaction to
extremely low temperatures in a special bath and forged cells from a
more resistant mixture of platinum and iridium, and used fluorite
stoppers. In 1886, after 74 years of effort by many chemists, Moissan
isolated elemental fluorine.

In 1906, two months before his death, Moissan received the Nobel Prize
in Chemistry, with the following citation:


Later uses
============
The Frigidaire division of General Motors (GM) experimented with
chlorofluorocarbon refrigerants in the late 1920s, and Kinetic
Chemicals was formed as a joint venture between GM and DuPont in 1930
hoping to market Freon-12 () as one such refrigerant. It replaced
earlier and more toxic compounds, increased demand for kitchen
refrigerators, and became profitable; by 1949 DuPont had bought out
Kinetic and marketed several other Freon compounds.
Polytetrafluoroethylene (Teflon) was serendipitously discovered in
1938 by Roy J. Plunkett while working on refrigerants at Kinetic, and
its superlative chemical and thermal resistance lent it to accelerated
commercialization and mass production by 1941.

Large-scale production of elemental fluorine began during World War
II. Germany used high-temperature electrolysis to make tons of the
planned incendiary chlorine trifluoride and the Manhattan Project used
huge quantities to produce uranium hexafluoride for uranium
enrichment. Since  is as corrosive as fluorine, gaseous diffusion
plants required special materials: nickel for membranes,
fluoropolymers for seals, and liquid fluorocarbons as coolants and
lubricants. This burgeoning nuclear industry later drove post-war
fluorochemical development.


                             Compounds
======================================================================
Fluorine has a rich chemistry, encompassing organic and inorganic
domains. It combines with metals, nonmetals, metalloids, and most
noble gases. Fluorine's high electron affinity results in a preference
for ionic bonding; when it forms covalent bonds, these are polar, and
almost always single.


Oxidation states
==================
In compounds, fluorine almost exclusively assumes an oxidation state
of −1. Fluorine in  is defined to have oxidation state 0. The unstable
species  and , which decompose at around 40 K, have intermediate
oxidation states;  and a few related species are predicted to be
stable.


Metals
========
Alkali metals form ionic and highly soluble monofluorides; these have
the cubic arrangement of sodium chloride and analogous chlorides.
Alkaline earth difluorides possess strong ionic bonds but are
insoluble in water, with the exception of beryllium difluoride, which
also exhibits some covalent character and has a quartz-like structure.
Rare earth elements and many other metals form mostly ionic
trifluorides.

Covalent bonding first comes to prominence in the tetrafluorides:
those of zirconium, hafnium and several actinides are ionic with high
melting points, while those of titanium, vanadium, and niobium are
polymeric, melting or decomposing at no more than 350 °C.
Pentafluorides continue this trend with their linear polymers and
oligomeric complexes. Thirteen metal hexafluorides are known, all
octahedral, and are mostly volatile solids but for liquid  and , and
gaseous . Rhenium heptafluoride, the only characterized metal
heptafluoride, is a low-melting molecular solid with pentagonal
bipyramidal molecular geometry. Gold heptafluoride is a
low-temperature complex of molecular F2 with AuF5, with NPA
calculations indicating that the fluorine in the F2 ligand is nearly
neutral while those in the AuF5 portion of the molecule have strong
negative partial charges. This is consistent with the F2 ligand
representing fluorine in the zero oxidation state. Metal fluorides
with more fluorine atoms are particularly reactive.

|**Structural progression of metal fluorides**
Sodium fluoride: cubic lattice of alternating sodium and fluorine
atoms with no distinct molecules        Bismuth pentafluoride: arbitrarily
long, straight chain of atoms   Rhenium heptafluoride: discrete small
molecule        Sodium fluoride, ionic  Bismuth pentafluoride, polymeric
Rhenium heptafluoride, molecular


Hydrogen
==========
Hydrogen and fluorine combine to yield hydrogen fluoride, in which
discrete molecules form clusters by hydrogen bonding, resembling water
more than hydrogen chloride. It boils at a much higher temperature
than heavier hydrogen halides and unlike them is miscible with water.
Hydrogen fluoride readily hydrates on contact with water to form
aqueous hydrogen fluoride, also known as hydrofluoric acid. Unlike the
other hydrohalic acids, which are strong, hydrofluoric acid is a weak
acid at low concentrations. However, it can attack glass, something
the other acids cannot do.


Other reactive nonmetals
==========================
Binary fluorides of metalloids and p-block nonmetals are generally
covalent and volatile, with varying reactivities. Period 3 and heavier
nonmetals can form hypervalent fluorides.

Boron trifluoride is planar and possesses an incomplete octet. It
functions as a Lewis acid and combines with Lewis bases like ammonia
to form adducts. Carbon tetrafluoride is tetrahedral and inert; its
group analogues, silicon and germanium tetrafluoride, are also
tetrahedral but behave as Lewis acids. The pnictogens form
trifluorides that increase in reactivity and basicity with higher
molecular weight, although nitrogen trifluoride resists hydrolysis and
is not basic. The pentafluorides of phosphorus, arsenic, and antimony
are more reactive than their respective trifluorides, with antimony
pentafluoride the strongest neutral Lewis acid known, only behind gold
pentafluoride.

Chalcogens have diverse fluorides: unstable difluorides have been
reported for oxygen (the only known compound with oxygen in an
oxidation state of +2), sulfur, and selenium; tetrafluorides and
hexafluorides exist for sulfur, selenium, and tellurium. The latter
are stabilized by more fluorine atoms and lighter central atoms, so
sulfur hexafluoride is especially inert. Chlorine, bromine, and iodine
can each form mono-, tri-, and pentafluorides, but only iodine
heptafluoride has been characterized among possible interhalogen
heptafluorides. Many of them are powerful sources of fluorine atoms,
and industrial applications using chlorine trifluoride require
precautions similar to those using fluorine.


Noble gases
=============
Noble gases, having complete electron shells, defied reaction with
other elements until 1962 when Neil Bartlett reported synthesis of
xenon hexafluoroplatinate; xenon difluoride, tetrafluoride,
hexafluoride, and multiple oxyfluorides have been isolated since then.
Among other noble gases, krypton forms a difluoride, and radon and
fluorine generate a solid suspected to be radon difluoride. Binary
fluorides of lighter noble gases are exceptionally unstable: argon and
hydrogen fluoride combine under extreme conditions to give argon
fluorohydride. Helium has no long-lived fluorides, and no neon
fluoride has ever been observed; helium fluorohydride has been
detected for milliseconds at high pressures and low temperatures.


Organic compounds
===================
The carbon-fluorine bond is organic chemistry's strongest, and gives
stability to organofluorines. It is almost non-existent in nature, but
is used in artificial compounds. Research in this area is usually
driven by commercial applications; the compounds involved are diverse
and reflect the complexity inherent in organic chemistry.


Discrete molecules
====================
The substitution of hydrogen atoms in an alkane by progressively more
fluorine atoms gradually alters several properties: melting and
boiling points are lowered, density increases, solubility in
hydrocarbons decreases and overall stability increases.
Perfluorocarbons, in which all hydrogen atoms are substituted, are
insoluble in most organic solvents, reacting at ambient conditions
only with sodium in liquid ammonia.

The term 'perfluorinated compound' is used for what would otherwise be
a perfluorocarbon if not for the presence of a functional group, often
a carboxylic acid. These compounds share many properties with
perfluorocarbons such as stability and hydrophobicity, while the
functional group augments their reactivity, enabling them to adhere to
surfaces or act as surfactants. Fluorosurfactants, in particular, can
lower the surface tension of water more than their hydrocarbon-based
analogues. Fluorotelomers, which have some unfluorinated carbon atoms
near the functional group, are also regarded as perfluorinated.


Polymers
==========
Polymers exhibit the same stability increases afforded by fluorine
substitution (for hydrogen) in discrete molecules; their melting
points generally increase too. Polytetrafluoroethylene (PTFE), the
simplest fluoropolymer and perfluoro analogue of polyethylene with
structural unit --, demonstrates this change as expected, but its very
high melting point makes it difficult to mold. Various PTFE
derivatives are less temperature-tolerant but easier to mold:
fluorinated ethylene propylene replaces some fluorine atoms with
trifluoromethyl groups, perfluoroalkoxy alkanes do the same with
trifluoromethoxy groups, and Nafion contains perfluoroether side
chains capped with sulfonic acid groups. Other fluoropolymers retain
some hydrogen atoms; polyvinylidene fluoride has half the fluorine
atoms of PTFE and polyvinyl fluoride has a quarter, but both behave
much like perfluorinated polymers.


About 180,000 metric tons of fluoropolymers were produced in 2006 and
2007, generating over $3.5 billion revenue per year. The global market
was estimated at just under $6 billion in 2011. Fluoropolymers can
only be formed by polymerizing free radicals.

Polytetrafluoroethylene (PTFE), sometimes called by its DuPont name
Teflon, represents 60-80% by mass of the world's fluoropolymer
production. 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 in PFTE-coated fiberglass
cloth for stadium roofs. The major consumer application is for
non-stick cookware. Jerked PTFE film becomes expanded PTFE (ePTFE), a
fine-pored membrane sometimes referred to by the brand name Gore-Tex
and used for rainwear, protective apparel, and filters; ePTFE fibers
may be made into seals and dust filters. Other fluoropolymers,
including fluorinated ethylene propylene, mimic PTFE's properties and
can substitute for it; they are more moldable, but also more costly
and have lower thermal stability. Films from two different
fluoropolymers replace glass in solar cells.

The chemically resistant (but expensive) fluorinated ionomers are used
as electrochemical cell membranes, of which the first and most
prominent example is Nafion. Developed in the 1960s, it was initially
deployed as fuel cell material in spacecraft and then replaced
mercury-based chloralkali process cells. Recently, the fuel cell
application has reemerged with efforts to install proton exchange
membrane fuel cells into automobiles. Fluoroelastomers such as Viton
are crosslinked fluoropolymer mixtures mainly used in O-rings;
perfluorobutane (C4F10) is used as a fire-extinguishing agent.


                             Production
======================================================================
Elemental fluorine and virtually all fluorine compounds are produced
from hydrogen fluoride or its aqueous solution, hydrofluoric acid.
Hydrogen fluoride is produced in kilns by the endothermic reaction of
fluorite (CaF2) with sulfuric acid:

:CaF2  +  H2SO4  →  2 HF(g)  +  CaSO4

The gaseous HF can then be absorbed in water or liquefied.

About 20% of manufactured HF is a byproduct of fertilizer production,
which produces hexafluorosilicic acid (H2SiF6), which can be degraded
to release HF thermally and by hydrolysis:
:H2SiF6  →  2 HF  +  SiF4
:SiF4  +  2 H2O   →  4 HF  +  SiO2


Industrial routes to F₂
====================================
Moissan's method is used to produce industrial quantities of fluorine,
via the electrolysis of a potassium bifluoride/hydrogen fluoride
mixture: hydrogen ions are reduced at a steel container cathode and
fluoride ions are oxidized at a carbon block anode, under 8-12 volts,
to generate hydrogen and fluorine gas respectively. Temperatures are
elevated, KF•2HF melting at 70 Celsius and being electrolyzed at 70 -.
KF, which acts to provide electrical conductivity, is essential since
pure HF cannot be electrolyzed because it is virtually non-conductive.
Fluorine can be stored in steel cylinders that have passivated
interiors, at temperatures below 200 Celsius; otherwise nickel can be
used. Regulator valves and pipework are made of nickel, the latter
possibly using Monel instead. Frequent passivation, along with the
strict exclusion of water and greases, must be undertaken. In the
laboratory, glassware may carry fluorine gas under low pressure and
anhydrous conditions; some sources instead recommend nickel-Monel-PTFE
systems.


Laboratory routes
===================
While preparing for a 1986 conference to celebrate the centennial of
Moissan's achievement, Karl O. Christe reasoned that chemical fluorine
generation should be feasible since some metal fluoride anions have no
stable neutral counterparts; their acidification potentially triggers
oxidation instead. He devised a method which evolves fluorine at high
yield and atmospheric pressure:

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

Christe later commented that the reactants "had been known for more
than 100 years and even Moissan could have come up with this scheme."
As late as 2008, some references still asserted that fluorine was too
reactive for any chemical isolation.


                      Industrial applications
======================================================================
Fluorite mining, which supplies most global fluorine, peaked in 1989
when 5.6 million metric tons of ore were extracted. Chlorofluorocarbon
restrictions lowered this to 3.6 million tons in 1994; production has
since been increasing. Around 4.5 million tons of ore and revenue of
US$550 million were generated in 2003; later reports estimated 2011
global fluorochemical sales at $15 billion and predicted 2016-18
production figures of 3.5 to 5.9 million tons, and revenue of at least
$20 billion. Froth flotation separates mined fluorite into two main
metallurgical grades of equal proportion: 60-85% pure metspar is
almost all used in iron smelting whereas 97%+ pure acidspar is mainly
converted to the key industrial intermediate hydrogen fluoride.

Image:The fluorine economy.svg|thumb|675px|center|Clickable diagram of
the fluorochemical industry according to mass flows
rect 9 6 81 34 Fluorite
rect 9 172 81 199 Fluorapatite
rect 142 5 244 34 Hydrogen fluoride
rect 142 65 245 97 Metal smelting
rect 142 121 244 154 Glass production
rect 309 5 411 33 Fluorocarbons
rect 310 63 413 92 Sodium hexafluoroaluminate
rect 311 121 414 154 Pickling (metal)
rect 310 171 412 200 Fluorosilicic acid
rect 309 211 412 243 Alkane cracking
rect 483 6 585 34 Hydrofluorocarbons
rect 484 47 585 76 Hydrochlorofluorocarbons
rect 483 88 586 116 Chlorofluorocarbon
rect 483 128 585 160 Teflon
rect 484 170 586 200 Water fluoridation
rect 483 210 586 238 Uranium enrichment
rect 484 258 586 287 Sulfur hexafluoride
rect 484 297 585 357 Tungsten hexafluoride
rect 28 246 177 293 Phosphogypsum
desc bottom-left



At least 17,000 metric tons of fluorine are produced each year. It
costs only $5-8 per kilogram as uranium or sulfur hexafluoride, but
many times more as an element because of handling challenges. Most
processes using free fluorine in large amounts employ 'in situ'
generation under vertical integration.

The largest application of fluorine gas, consuming up to 7,000 metric
tons annually, is in the preparation of  for the nuclear fuel cycle.
Fluorine is used to fluorinate uranium tetrafluoride, itself formed
from uranium dioxide and hydrofluoric acid. Fluorine is monoisotopic,
so any mass differences between  molecules are due to the presence of
or , enabling uranium enrichment via gaseous diffusion or gas
centrifuge. About 6,000 metric tons per year go into producing the
inert dielectric  for high-voltage transformers and circuit breakers,
eliminating the need for hazardous polychlorinated biphenyls
associated with  devices. Several fluorine compounds are used in
electronics: rhenium and tungsten hexafluoride in chemical vapor
deposition, tetrafluoromethane in plasma etching and nitrogen
trifluoride in cleaning equipment. Fluorine is also used in the
synthesis of organic fluorides, but its reactivity often necessitates
conversion first to the gentler , , or , which together allow
calibrated fluorination. Fluorinated pharmaceuticals use sulfur
tetrafluoride instead.


Inorganic fluorides
=====================
As with other iron alloys, around 3 kg metspar is added to each metric
ton of steel; the fluoride ions lower its melting point and viscosity.
Alongside its role as an additive in materials like enamels and
welding rod coats, most acidspar is reacted with sulfuric acid to form
hydrofluoric acid, which is used in steel pickling, glass etching and
alkane cracking. One-third of HF goes into synthesizing cryolite and
aluminium trifluoride, both fluxes in the Hall-Héroult process for
aluminium extraction; replenishment is necessitated by their
occasional reactions with the smelting apparatus. Each metric ton of
aluminium requires about 23 kg of flux. Fluorosilicates consume the
second largest portion, with sodium fluorosilicate used in water
fluoridation and laundry effluent treatment, and as an intermediate en
route to cryolite and silicon tetrafluoride. Other important inorganic
fluorides include those of cobalt, nickel, and ammonium.


Organic fluorides
===================
Organofluorides consume over 20% of mined fluorite and over 40% of
hydrofluoric acid, with refrigerant gases dominating and
fluoropolymers increasing their market share. Surfactants are a minor
application but generate over $1 billion in annual revenue. Due to the
danger from direct hydrocarbon-fluorine reactions above -150 °C,
industrial fluorocarbon production is indirect, mostly through halogen
exchange reactions such as Swarts fluorination, in which chlorocarbon
chlorines are substituted for fluorines by hydrogen fluoride under
catalysts. Electrochemical fluorination subjects hydrocarbons to
electrolysis in hydrogen fluoride, and the Fowler process treats them
with solid fluorine carriers like cobalt trifluoride.


Refrigerant gases
===================
Halogenated refrigerants, termed Freons in informal contexts, are
identified by R-numbers that denote the amount of fluorine, chlorine,
carbon, and hydrogen present. Chlorofluorocarbons (CFCs) like R-11,
R-12, and R-114 once dominated organofluorines, peaking in production
in the 1980s. Used for air conditioning systems, propellants and
solvents, their production was below one-tenth of this peak by the
early 2000s, after widespread international prohibition.
Hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) were
designed as replacements; their synthesis consumes more than 90% of
the fluorine in the organic industry. Important HCFCs include R-22,
chlorodifluoromethane, and R-141b. The main HFC is R-134a with a new
type of molecule HFO-1234yf, a Hydrofluoroolefin (HFO) coming to
prominence owing to its global warming potential of less than 1% that
of HFC-134a.


Surfactants
=============
Fluorosurfactants are small organofluorine molecules used for
repelling water and stains. Although expensive (comparable to
pharmaceuticals at $200-2000 per kilogram), they yielded over $1
billion in annual revenues by 2006; Scotchgard alone generated over
$300 million in 2000. Fluorosurfactants are a minority in the overall
surfactant market, most of which is taken up by much cheaper
hydrocarbon-based products. Applications in paints are burdened by
compounding costs; this use was valued at only $100 million in 2006.


Agrichemicals
===============
About 30% of agrichemicals contain fluorine, most of them herbicides
and fungicides with a few crop regulators. Fluorine substitution,
usually of a single atom or at most a trifluoromethyl group, is a
robust modification with effects analogous to fluorinated
pharmaceuticals: increased biological stay time, membrane crossing,
and altering of molecular recognition. Trifluralin is a prominent
example, with large-scale use in the U.S. as a weedkiller, but it is a
suspected carcinogen and has been banned in many European countries.
Sodium monofluoroacetate (1080) is a mammalian poison in which one
sodium acetate hydrogen is replaced with fluorine; it disrupts cell
metabolism by replacing acetate in the citric acid cycle. First
synthesized in the late 19th century, it was recognized as an
insecticide in the early 20th century, and was later deployed in its
current use. New Zealand, the largest consumer of 1080, uses it to
protect kiwis from the invasive Australian common brushtail possum.
Europe and the U.S. have banned 1080.


Dental care
=============
Population studies from the mid-20th century onwards show topical
fluoride reduces dental caries. This was first attributed to the
conversion of tooth enamel hydroxyapatite into the more durable
fluorapatite, but studies on pre-fluoridated teeth refuted this
hypothesis, and current theories involve fluoride aiding enamel growth
in small caries. After studies of children in areas where fluoride was
naturally present in drinking water, controlled public water supply
fluoridation to fight tooth decay began in the 1940s and is now
applied to water supplying 6 percent of the global population,
including two-thirds of Americans. Reviews of the scholarly literature
in 2000 and 2007 associated water fluoridation with a significant
reduction of tooth decay in children. Despite such endorsements and
evidence of no adverse effects other than mostly benign dental
fluorosis, opposition still exists on ethical and safety grounds. The
benefits of fluoridation have lessened, possibly due to other fluoride
sources, but are still measurable in low-income groups. Sodium
monofluorophosphate and sometimes sodium or tin(II) fluoride are often
found in fluoride toothpastes, first introduced in the U.S. in 1955
and now ubiquitous in developed countries, alongside fluoridated
mouthwashes, gels, foams, and varnishes.


Pharmaceuticals
=================
Twenty percent of modern pharmaceuticals contain fluorine. One of
these, the cholesterol-reducer atorvastatin (Lipitor), made more
revenue than any other drug until it became generic in 2011. The
combination asthma prescription Seretide, a top-ten revenue drug in
the mid-2000s, contains two active ingredients, one of which -
fluticasone - is fluorinated. Many drugs are fluorinated to delay
inactivation and lengthen dosage periods because the carbon-fluorine
bond is very stable. Fluorination also increases lipophilicity because
the bond is more hydrophobic than the carbon-hydrogen bond, and this
often helps in cell membrane penetration and hence bioavailability.

Tricyclics and other pre-1980s antidepressants had several side
effects due to their non-selective interference with neurotransmitters
other than the serotonin target; the fluorinated fluoxetine was
selective and one of the first to avoid this problem. Many current
antidepressants receive this same treatment, including the selective
serotonin reuptake inhibitors: citalopram, its enantiomer
escitalopram, and fluvoxamine and paroxetine. Quinolones are
artificial broad-spectrum antibiotics that are often fluorinated to
enhance their effects. These include ciprofloxacin and levofloxacin.
Fluorine also finds use in steroids: fludrocortisone is a blood
pressure-raising mineralocorticoid, and triamcinolone and
dexamethasone are strong glucocorticoids. The majority of inhaled
anesthetics are heavily fluorinated; the prototype halothane is much
more inert and potent than its contemporaries. Later compounds such as
the fluorinated ethers sevoflurane and desflurane are better than
halothane and are almost insoluble in blood, allowing faster waking
times.


PET scanning
==============
Fluorine-18 is often found in radioactive tracers for positron
emission tomography, as its half-life of almost two hours is long
enough to allow for its transport from production facilities to
imaging centers. The most common tracer is fluorodeoxyglucose which,
after intravenous injection, is taken up by glucose-requiring tissues
such as the brain and most malignant tumors; computer-assisted
tomography can then be used for detailed imaging.


Oxygen carriers
=================
Liquid fluorocarbons can hold large volumes of oxygen or carbon
dioxide, more so than blood, and have attracted attention for their
possible uses in artificial blood and in liquid breathing. Because
fluorocarbons do not normally mix with water, they must be mixed into
emulsions (small droplets of perfluorocarbon suspended in water) to be
used as blood. One such product, Oxycyte, has been through initial
clinical trials. These substances can aid endurance athletes and are
banned from sports; one cyclist's near death in 1998 prompted an
investigation into their abuse. Applications of pure perfluorocarbon
liquid breathing (which uses pure perfluorocarbon liquid, not a water
emulsion) include assisting burn victims and premature babies with
deficient lungs. Partial and complete lung filling have been
considered, though only the former has had any significant tests in
humans. An Alliance Pharmaceuticals effort reached clinical trials but
was abandoned because the results were not better than normal
therapies.


                          Biological role
======================================================================
Fluorine is not essential for humans and other mammals, but small
amounts are known to be beneficial for the strengthening of dental
enamel (where the formation of fluorapatite makes the enamel more
resistant to attack, from acids produced by bacterial fermentation of
sugars). Small amounts of fluorine may be beneficial for bone
strength, but the latter has not been definitively established. Both
the WHO and the Institute of Medicine of the US National Academies
publish recommended daily allowance (RDA) and upper tolerated intake
of fluorine, which varies with age and gender.

Natural organofluorines have been found in microorganisms, plants and,
recently, animals. The most common is fluoroacetate, which is used as
a defense against herbivores by at least 40 plants in Africa,
Australia and Brazil. Other examples include terminally fluorinated
fatty acids, fluoroacetone, and 2-fluorocitrate. An enzyme that binds
fluorine to carbon - adenosyl-fluoride synthase - was discovered in
bacteria in 2002.


                              Toxicity
======================================================================
Elemental fluorine is highly toxic to living organisms. Its effects in
humans start at concentrations lower than hydrogen cyanide's 50 ppm
and are similar to those of chlorine: significant irritation of the
eyes and respiratory system as well as liver and kidney damage occur
above 25 ppm, which is the immediately dangerous to life and health
value for fluorine. The eyes and nose are seriously damaged at 100
ppm, and inhalation of 1,000 ppm fluorine will cause death in minutes,
compared to 270 ppm for hydrogen cyanide.


Hydrofluoric acid
===================
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Hydrofluoric acid is the weakest of the hydrohalic acids, having a pKa
of 3.2 at 25 °C. Pure hydrogen fluoride is a volatile liquid due to
the presence of hydrogen bonding, while the other hydrogen halides are
gases. It is able to attack glass, concrete, metals, and organic
matter.

Hydrofluoric acid is a contact poison with greater hazards than many
strong acids like sulfuric acid even though it is weak: it remains
neutral in aqueous solution and thus penetrates tissue faster, whether
through inhalation, ingestion or the skin, and at least nine U.S.
workers died in such accidents from 1984 to 1994. It reacts with
calcium and magnesium in the blood leading to hypocalcemia and
possible death through cardiac arrhythmia. Insoluble calcium fluoride
formation triggers strong pain and burns larger than 160 cm2 (25 in2)
can cause serious systemic toxicity.

Exposure may not be evident for eight hours for 50% HF, rising to 24
hours for lower concentrations, and a burn may initially be painless
as hydrogen fluoride affects nerve function. If skin has been exposed
to HF, damage can be reduced by rinsing it under a jet of water for
10-15 minutes and removing contaminated clothing. Calcium gluconate is
often applied next, providing calcium ions to bind with fluoride; skin
burns can be treated with 2.5% calcium gluconate gel or special
rinsing solutions. Hydrofluoric acid absorption requires further
medical treatment; calcium gluconate may be injected or administered
intravenously. Using calcium chloride - a common laboratory reagent -
in lieu of calcium gluconate is contraindicated, and may lead to
severe complications. Excision or amputation of affected parts may be
required.


Fluoride ion
==============
Soluble fluorides are moderately toxic: 5-10 g sodium fluoride, or
32-64 mg fluoride ions per kilogram of body mass, represents a lethal
dose for adults. One-fifth of the lethal dose can cause adverse health
effects, and chronic excess consumption may lead to skeletal
fluorosis, which affects millions in Asia and Africa, and, in
children, to reduced intelligence. Ingested fluoride forms
hydrofluoric acid in the stomach which is easily absorbed by the
intestines, where it crosses cell membranes, binds with calcium and
interferes with various enzymes, before urinary excretion. Exposure
limits are determined by urine testing of the body's ability to clear
fluoride ions.

Historically, most cases of fluoride poisoning have been caused by
accidental ingestion of insecticides containing inorganic fluorides.
Most current calls to poison control centers for possible fluoride
poisoning come from the ingestion of fluoride-containing toothpaste.
Malfunctioning water fluoridation equipment is another cause: one
incident in Alaska affected almost 300 people and killed one person.
Dangers from toothpaste are aggravated for small children, and the
Centers for Disease Control and Prevention recommends supervising
children below six brushing their teeth so that they do not swallow
toothpaste. One regional study examined a year of pre-teen fluoride
poisoning reports totaling 87 cases, including one death from
ingesting insecticide. Most had no symptoms, but about 30% had stomach
pains. A larger study across the U.S. had similar findings: 80% of
cases involved children under six, and there were few serious cases.


Atmosphere
============
The Montreal Protocol, signed in 1987, set strict regulations on
chlorofluorocarbons (CFCs) and bromofluorocarbons due to their ozone
damaging potential (ODP). The high stability which suited them to
their original applications also meant that they were not decomposing
until they reached higher altitudes, where liberated chlorine and
bromine atoms attacked ozone molecules. Even with the ban, and early
indications of its efficacy, predictions warned that several
generations would pass before full recovery. With one-tenth the ODP of
CFCs, hydrochlorofluorocarbons (HCFCs) are the current replacements,
and are themselves scheduled for substitution by 2030-2040 by
hydrofluorocarbons (HFCs) with no chlorine and zero ODP. In 2007 this
date was brought forward to 2020 for developed countries; the
Environmental Protection Agency had already prohibited one HCFC's
production and capped those of two others in 2003. Fluorocarbon gases
are generally greenhouse gases with global-warming potentials (GWPs)
of about 100 to 10,000; sulfur hexafluoride has a value of around
20,000. An outlier is HFO-1234yf which is a new type of refrigerant
called a Hydrofluoroolefin (HFO) and has attracted global demand due
to its GWP of less than 1 compared to 1,430 for the current
refrigerant standard HFC-134a.


Biopersistence
================
Organofluorines exhibit biopersistence due to the strength of the
carbon-fluorine bond. Perfluoroalkyl acids (PFAAs), which are
sparingly water-soluble owing to their acidic functional groups, are
noted persistent organic pollutants; perfluorooctanesulfonic acid
(PFOS) and perfluorooctanoic acid (PFOA) are most often researched.
PFAAs have been found in trace quantities worldwide from polar bears
to humans, with PFOS and PFOA known to reside in breast milk and the
blood of newborn babies. A 2013 review showed a slight correlation
between groundwater and soil PFAA levels and human activity; there was
no clear pattern of one chemical dominating, and higher amounts of
PFOS were correlated to higher amounts of PFOA. In the body, PFAAs
bind to proteins such as serum albumin; they tend to concentrate
within humans in the liver and blood before excretion through the
kidneys. Dwell time in the body varies greatly by species, with
half-lives of days in rodents, and years in humans. High doses of PFOS
and PFOA cause cancer and death in newborn rodents but human studies
have not established an effect at current exposure levels.


                              See also
======================================================================
* Argon fluoride laser
* Electrophilic fluorination
* Fluoride selective electrode, which measures fluoride concentration
* Fluorine absorption dating
* Fluorous chemistry, a process used to separate reagents from organic
solvents
* Krypton fluoride laser
* Radical fluorination


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=========
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Original Article: http://en.wikipedia.org/wiki/Fluorine

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