Fluorine (from L. Fluere, meaning "to flow"), is a chemical element in the periodic table that has the symbol F and atomic number 9. It is a poisonous pale yellow, 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.
|Name, Symbol, Number||Fluorine, F, 9|
|Group, Period, Block||17 (VIIA), 2 , p|
|Density, Hardness||1.696 kg/m3 (273 K), NA|
|Appearance||pale greenish-yellow gas|
|Atomic weight||18.9984 amu|
|Atomic radius (calc.)||50 (42) pm|
|Covalent radius||71 pm|
|van der Waals radius||147 pm|
|Electron configuration||[He]2s2s2 2p5|
|e- 's per energy level||2, 7|
|Oxidation states (Oxide)||-1 (strong acid)|
|State of matter||Gas (nonmagnetic)|
|Melting point||53.53 K (-363.32 ðF)|
|Boiling point||85.03 K (-306.62 ðF)|
|Molar volume||11.20 ×1010-6 m3/mol|
|Heat of vaporization||3.2698 kJ/mol|
|Heat of fusion||0.2552 kJ/mol|
|Vapor pressure||no data|
|Speed of sound||no data|
|Electronegativity||3.98 (Pauling scale)|
|Specific heat capacity||824 J/(kg*K)|
|Electrical conductivity||no data|
|Thermal conductivity||0.0279 W/(m*K)|
|1st ionization potential||1681.0 kJ/mol|
|2nd ionization potential||3374.2 kJ/mol|
|3rd ionization potential||6050.4 kJ/mol|
|4th ionization potential||8407.7 kJ/mol|
|5th ionization potential||11022.7 kJ/mol|
|6th ionization potential||15164.1 kJ/mol|
|7th ionization potential||17868 kJ/mol|
|8th ionization potential||92038.1 kJ/mol|
|9th ionization potential||106434.3 kJ/mol|
|Most stable isotopes|
|SI units & STP are used except where noted.|
|Table of contents|
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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 forms compounds with most other elements, including the noble gases 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 always occurs combined and has such an affinity for most elements, especially silicon, that it can neither be prepared nor should be kept in glass vessels.
In aqueous solution, fluorine commonly occurs as the fluoride ion F-. Other forms are fluoro-complexes (such as [FeF4]-) or H2F+.
Fluorine is used in the production of low friction plastics such as Teflon, and in halons such as Freon. Other uses:
- Hydrofluoric acid (chemical formula HFF) 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.
- Fluorochlorohydrocarbonss 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. Sulfurhexafluoride is an extremely inert and nontoxic gas. These classes of compounds are potent greenhouse gases.
- Potassiumhexafluoroaluminate, 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.
Fluorine (L fluere meaning flow or flux) in the form of fluorspar was described in 1529 by Georigius 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 (some experiments would end in tragedy).
This element was not isolated for many years after this due to the fact that when it is separated from one of its compounds 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.
The first commercial production of fluorine was for the atomic bomb Manhattan project in World War II where the compound uranium hexafluoride (UF6) was used to separate isotopes of uranium. This process is still is use today in 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 compoundss. Fluorine compounds involving noble gases have been confirmed with fluorides of krypton, radon, and xenon.
This element is recovered from fluorite, cryolite, and fluorapatite.
See also: Fluorocarbon
Fluorine and HF must be handled with great care and any contact with skin and eyes should be strictly avoided.
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 ppb. It is recommended that the maximum allowable concentration for a daily 8-hour time-weighted exposure is 1 ppm (lower than e.g. hydrogen cyanide)
However, safe handling procedures enable the transport of liquid fluorine by the ton.