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Encyclopedia > Flourine
oxygenfluorineneon
 
F
Cl  
 
 

General
Name, Symbol, Number Fluorine, F, 9
Series Halogens
Group, Period, Block 17 (VIIA), 2 , p
Density, Hardness 1.696 kg/m3, NA
Appearance pale greenish-yellow gas
Atomic properties
Atomic weight 18.9984 g/mol
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 cm3/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/(kgK)
Electrical conductivity no data
Thermal conductivity 0.0279 W/(mK)
1st ionization potential 1.681 MJ/mol
2nd ionization potential 3.374 MJ/mol
3rd ionization potential 6.050 MJ/mol
4th ionization potential 8.407 MJ/mol
5th ionization potential 11.02 MJ/mol
6th ionization potential 15.16 MJ/mol
7th ionization potential 17.87 MJ/mol
8th ionization potential 92.04 MJ/mol
9th ionization potential 106.4 MJ/mol
Most stable isotopes
iso NA half-life DM DE DP
19F 100% F is stable with 10 neutrons
SI units & STP are used except where noted.

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.

Contents

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 readily forms compounds with most other elements. Fluorine even combines with 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 is far too reactive to be found in elemental form and has such an affinity for most elements, especially 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+.


Fluorides are compounds that combine fluoride with some positively charged rest. They often consist of ions. Fluorine compounds with metals are among the most stable of salts.


Applications

Fluorine is used in the production of low friction plastics such as Teflon, and in halons such as Freon. Other uses:

Some researchers - including US space scientists in the early 1960s have studied elemental fluorine gas a possible rocket propellant due to its exceptionally high specific impulse. Experiments failed since fluorine was so hard to handle


History

Fluorine (L fluere meaning flow or flux) in the form of fluorspar (calcium fluoride) 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, easily obtained by treating calcium fluoride (fluorspar) with concentrated sulfuric acid.


It was eventually realised 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.


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 compounds. Fluorine compounds involving noble gases were first synthesised by Howard Claassen, Henry Selig, 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


Precautions

Both 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 nL/L. It is recommended that the maximum allowable concentration for a daily 8-hour time-weighted exposure is 1 L/L (part per million by volume (lower than e.g. hydrogen cyanide)


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


References

External links

  • WebElements.com – Fluorine (http://www.webelements.com/webelements/elements/text/F/index.html)
  • EnvironmentalChemistry.com – Fluorine (http://environmentalchemistry.com/yogi/periodic/F.html)
  • It's Elemental – Fluorine (http://education.jlab.org/itselemental/ele009.html)
  • Picture of liquid fluorine – chemie-master.de (http://www.chemie-master.de/pse/pse.php?modul=F)
  • Chemsoc.org (http://www.chemsoc.org/viselements/pages/fluorine.html)
  • Periodic Table of Elements (http://nautilus.fis.uc.pt/st2.5/index-en.html)

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