Fluorine
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. In its elemental form 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 fluo 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. 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 that help deter predation.
Characteristics
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. Some solid nonmetals 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 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. It has a characteristic halogen-like pungent and biting odor detectable at 20 ppb. Fluorine condenses into a bright yellow liquid at, a transition temperature similar to those of oxygen and nitrogen.Fluorine has two solid forms, α- and β-fluorine. The latter crystallizes at 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 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
Natural fluorine consists entirely of the stable isotope. It has a high magnetogyric ratio and exceptional sensitivity to magnetic fields; because it is also the only isotope, it is highly suited for NMR and use 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. 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. Two metastable isomers of fluorine are known,, with a half-life of 162 nanoseconds, and, with a half-life of 2.2 milliseconds.Occurrence
Universe
Among the lighter elements, fluorine's abundance value of 400 ppb – 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
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.
History
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 for fluorite rocks. The name later evolved into fluorspar 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. 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. The New Latin name fluorum gave the element its current symbol F; Fl was used in early papers.