From Academic Kids
Template:Elementbox header Template:Elementbox series Template:Elementbox groupperiodblock Template:Elementbox appearance img Template:Elementbox atomicmass gpm Template:Elementbox econfig Template:Elementbox epershell Template:Elementbox section physicalprop Template:Elementbox phase Template:Elementbox meltingpoint Template:Elementbox boilingpoint Template:Elementbox heatfusion kjpmol Template:Elementbox heatvaporiz kjpmol Template:Elementbox heatcapacity jpmolkat25 Template:Elementbox vaporpressure katpa Template:Elementbox section atomicprop Template:Elementbox crystalstruct Template:Elementbox oxistates Template:Elementbox electroneg pauling Template:Elementbox ionizationenergies4 Template:Elementbox atomicradius pm Template:Elementbox atomicradiuscalc pm Template:Elementbox covalentradius pm Template:Elementbox vanderwaalsrad pm Template:Elementbox section miscellaneous Template:Elementbox magnetic Template:Elementbox thermalcond wpmkat300k Template:Elementbox cas number Template:Elementbox isotopes begin Template:Elementbox isotopes stable Template:Elementbox isotopes end Template:Elementbox footer
Fluorine (from L. fluere, meaning "to flow"), is the chemical element in the periodic table that has the symbol F and atomic number 9. It is a poisonous pale yellow-green, univalent gaseous halogen that is the most chemically reactive and electronegative of all the elements. In its pure form, it is highly dangerous, causing severe chemical burns on contact with skin.
Pure fluorine is a corrosive pale yellow gas that is a powerful oxidizing agent. It is the most reactive and electronegative of all the elements, and readily forms compounds with most other elements. Fluorine even combines with the noble gases krypton, xenon, and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. In a jet of fluorine gas, glass, metals, water and other substances burn with a bright flame. It is far too reactive to be found in elemental form and has such an affinity for most elements, including silicon, that it can neither be prepared nor should be kept in glass vessels. In moist air it attacks water to form the equally dangerous hydrofluoric acid.
In aqueous solution, fluorine commonly occurs as the fluoride ion F-. Other forms are fluoro-complexes (such as [FeF4]-) or H2F+.
- Hydrofluoric acid (chemical formula HF) is used to etch glass in light bulbs and other products.
- Monatomic fluorine is used for plasma ashing in semiconductor manufacturing.
- Along with its compounds, fluorine is used in the production of uranium (from the hexafluoride) and in more than 100 different commercial fluorochemicals, including many high-temperature plastics.
- Fluorochlorohydrocarbons are used extensively in air conditioning and in refrigeration. Chlorofluorocarbons have been banned for these applications because they are suspected to contribute to the ozone hole. Sulfur hexafluoride is an extremely inert and (unusually for a fluorine compound) nontoxic gas. These classes of compounds are potent greenhouse gases.
- Potassium hexafluoroaluminate, the so-called cryolite, is used in electrolysis of aluminium.
- Sodium fluoride has been used as an insecticide, especially against cockroaches.
- Some other fluorides are often added to toothpaste and, somewhat controversially, to municipal water supplies to prevent dental cavities.
- It has been used in the past to help molten metal flow, hence the name.
- Fluorine-18, a radioactive isotope that emits positrons, is often used in positron emission tomography because of its half-life of 110 minutes.
Some researchers - including US space scientists in the early 1960s have studied elemental fluorine gas as a possible rocket propellant due to its exceptionally high specific impulse. Experiments failed since fluorine was so hard to handle.
Fluorine (L fluere meaning flow or flux) in the form of fluorspar (calcium fluoride) was described in 1529 by Georgius Agricola for its use as a flux, which is a substance that is used to promote the fusion of metals or minerals. In 1670 Schwandhard found that glass was etched when it was exposed to fluorspar that was treated with acid. Karl Scheele and many later researchers, including Humphry Davy, Gay-Lussac, Antoine Lavoisier, and Louis Thenard all would experiment with hydrofluoric acid, easily obtained by treating calcium fluoride (fluorspar) with concentrated sulfuric acid.
It was eventually realized that hydrofluoric acid contained a previously unknown element. This element was not isolated for many years after this due to its extreme reactivity - it is separated from its compounds only with difficulty and then it immediately attacks the remaining materials of the compound. Finally in 1886 fluorine was isolated by Henri Moissan after almost 74 years of continuous effort. It was an effort which cost several researchers their health or even their lives, and for Moissan, it earned him the 1906 Nobel Prize in chemistry.
The first commercial production of fluorine was needed for the atomic bomb Manhattan project in World War II where the compound uranium hexafluoride (UF6) was used to separate the U-235 and U-238 isotopes of uranium. Today both the gaseous diffusion process and the gas centrifuge process use gaseous (UF6) to produce enriched uranium for nuclear power applications.
Fluorine can often be substituted for hydrogen when it occurs in organic compounds. Through this mechanism, fluorine can have a very large number of compounds. Fluorine compounds involving noble gases were first synthesised by Howard Claassen, Henry Selig, and John Malm in 1962 - xenon tetrafluoride being the first. Fluorides of krypton and radon have also been prepared. This element is recovered from fluorite, cryolite, and fluorapatite.
See also: Fluorocarbon
Both elemental fluorine and fluoride ions are highly toxic. When it is a free element, fluorine has a characteristic pungent odor that is detectable in concentrations as low as 20 nL/L. It is recommended that the maximum allowable concentration for a daily 8-hour time-weighted exposure is 1 µL/L (part per million by volume (lower than e.g. hydrogen cyanide)
- Los Alamos National Laboratory – Fluorine (http://periodic.lanl.gov/elements/9.html)