Fluorine

Oxygen - Fluorine - Neon   F Cl       Full table
General
Name, Symbol, Number	Fluorine, F, 9
Series	Halogens
Group, Period, Block	17 (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]2s2 2p5
e- 's per energy level	2, 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 ×10-3 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
iso NA half-life DM DE MeV DP 19F 100% 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 kept in glass vessels.

In solution, fluorine commonly occurs as the fluoride ion F-. Fluorides are compounds that combine this fluoride ion with some positively charged radical.

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 HF) 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.
• 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. Both of these classes of compounds are potent greenhouse gases.
• 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

A certain hypothesis held by some says that fluorine can be substituted for hydrogen when it occurs in organic compounds. Through this mechanism it is thought that fluorine can have a very large number of compounds. 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.

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

External Links

• WebElements.com - Fluorine (http://www.webelements.com/webelements/elements/text/F/index)
• EnvironmentalChemistry.com - Fluorine (http://environmentalchemistry.com/yogi/periodic/F)
• It's Elemental - Fluorine (http://education.jlab.org/itselemental/ele009)