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Fluorine

[it:Fluoro]]
Oxygen - Fluorine - Neon
 
F
Cl  
 
 

General
Name, Symbol, NumberFluorine, F, 9
Series Halogens
Group, Period, Block17 (VIIA), 2 , p
Density, Hardness 1.696 kg/m3 (273 K), NA
Appearance pale greenish-yellow gas
Atomic Properties
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 level2, 7
Oxidation states (Oxide) -1 (strong acid)
Crystal structure cubic
Physical Properties
State of matter Gas (nonmagnetic)
Melting point 53.53 K (-363.32 °F)
Boiling point 85.03 K (-306.62 °F)
Molar volume 11.20 ×1010-3 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
Miscellaneous
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
isoNAhalf-life DMDE MeVDP
19F100%F is stable with 10 neutrons
SI units & STP are used except where noted.
Fluorine 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.

Table of contents
1 Notable Characteristics
2 Applications
3 History
4 Compounds
5 Precautions
6 External Links

Notable Characteristics

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+.

Fluorides are compounds that combine fluoride with some positively charged rest. They often consist of ions.

Applications

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.
  • Monoatomic 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 untoxic gas. These classes of compounds are potent greenhouse gases.
  • Potassium hexafluoro aluminate, 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.

    • Some researchers have studied elemental fluorine gas a possible rocket propellant due to its exceptionally high specific impulse.

    History

    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.

    Compounds

    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 rare gases have been confirmed with fluorides of krypton, radon, and xenon. This element is recovered from fluorite, cryolite, and fluorapatite.

    See also: Fluorocarbon

    Precautions

    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. prussic acid)

    However, safe handling procedures enable the transport of liquid fluorine by the ton.

    External Links