Fluorine is a highly reactive chemical element with atomic symbol F. Having the atomic number nine, fluorine is the lightest halogen. Fluorine is a yellow-green gas which does not occur as a free, unreacted element in the natural environment. Under conditions of standard temperature and pressure, elemental fluorine forms a diatomic molecule with chemical formula F2. Chemically, fluorine is one of the strongest known oxidizing agents, and even more reactive and hazardous than chlorine. Its very high electron affinity causes fluorine to react directly with almost all other elements except for several of the Noble gases.
|Previous Element: Oxygen
Next Element: Neon
|Phase at Room Temp.||gas|
|Melting Point (K)||53.58|
|Boiling Point (K)||85.1|
|Heat of Fusion (kJ/mol)||1.0|
|Heat of Vaporization (kJ/mol)||6.5|
|Heat of Atomization (kJ/mol)||79|
|Thermal Conductivity (J/m sec K)||0.03|
|Electrical Conductivity (1/mohm cm)||0|
|Number of Isotopes||16(1 natural)|
|Electron Affinity (kJ/mol)||328|
|First Ionization Energy (kJ/mol)||1681|
|Second Ionization Energy (kJ/mol)||3374.1|
|Third Ionization Energy (kJ/mol)||6050.3|
|Atomic Volume (cm3/mol)||12.6|
|Ionic Radius2- (pm)||---|
|Ionic Radius1- (pm)||119|
|Atomic Radius (pm)||72|
|Ionic Radius1+ (pm)||---|
|Ionic Radius2+ (pm)||---|
|Ionic Radius3+ (pm)||---|
|Common Oxidation Numbers||-1|
|Other Oxid. Numbers||---|
|In Earth's Crust (mg/kg)||5.85×102|
|In Earth's Ocean (mg/L)||1.3×100|
|In Human Body (%)||0.004%|
|Regulatory / Health|
|OSHA Permissible Exposure Limit||1ppm=1.55mg/m3
|OSHA PEL Vacated 1989||TWA:0.1ppm|
|NIOSH Recommended Exposure Limit||TWA:0.1ppm|
Fluoride has been used for the last seven decades as a tooth decay preventative, although systematic side effects of tooth discoloration can occur for higher doses. A number of fluorocarbons have been produced as refrigerants and other industrial chemicals; some of these compounds are very effective greenhouse gases with a potency as high as 10,000 times that of carbon dioxide.
In the universe at large, fluorine is somewhat uncommon with an average concentration of 0.4 parts per million; this outcome is due to the solar temperatures needed to produce it also enable it to rapidly fuse with hydrogen to form oxygen and helium, or with helium to make neon and hydrogen. Most fluorine is created either during a supernova when a neutrino collides with a neon atom, or when a blue Wolf-Rayet star of more than 40 solar masses has a stellar wind blowing the fluorine out of the star sooner than hydrogen or helium can destroy the fluorine.
Fluorite, CaF2, also called fluorspar, is a common natural source for fluorine. Almost half of all mined fluorite is employed to assist in molten metal flow, particularly in iron smelting processes. Most of the balance is utilized in conversion to hydrofluoric acid, much of which is applied to produce organofluorides or synthetic cryolite.
Specialized uses of fluorine include in fine micro-etching for the semiconductor industry. The compound sulfur hexaflouride is quite non-reactive so is utilized in numerous applications where an inert tracer gas is needed for precision measurement and lack of interference in atmospheric chemical processes; for example, this gas was used in calibrating the first line source air pollution dispersion model to understand the dispersion of roadway pollutants.
The single stable isotope of fluorine is 19F, which exhibits a nuclear spin of one half and a high nuclear magnetic moment; correspondingly, fluorine compounds of the stable isotope are highly amenable to nuclear magnetic resonance spectroscopy. The monoisotopic occurance of fluorine assists in its use in uranium enrichment, as uranium hexafluoride molecules differ in mass only due to uranium atom isotope mass differences. The molecules with different masses due to uranium mass differences are separated via diffusion and gas centrifugation.
The concentration of fluorides in soils typically ranges between 200 and 300 parts per million; moreover, levels are higher in areas with notable fluoride-containing mineral deposits. Higher levels may also occur where phosphate fertilizers have been applied, where coal-fired power plants or fluoride-releasing industries are situated, or in the vicinity of certain hazardous waste sites. People may be exposed to fluorides through skin contact with these soils.
Use of fluoride to combat dental caries has been conducted systematically since the year 1945; however, side effects of tooth enamel mottling and discoloration are somewhat common. For example a 1997 study in the USA indicated the prevalence of over 29 percent of children had noticeable cosmetic side efffects of flouridation.
Accumulation of chlorofluorocarbons in the upper atmosphere has had an adverse impact upon the ozone layer, with formation of a dramatic hole in the Southern Hemisphere ozone layer, leading to measurable increases in incidence of skin cancer in humans.
- Jean Aigueperse, Paul Mollard, Didier Devilliers, Marius Chemla, Robert Faron, Renée Romano and Jean Pierre Cuer. 2005. Fluorine Compounds, Inorganic. in Ullmann. Encyclopedia of Industrial Chemistry. Wiley-VCH. Weinheim
- Agostino Renda, Yeshe Fenner, Brad K.Gibson, Amanda Karakas, John C.Lattanzio, Simon Campbell, Alessandro Chieffi, Katia Cunha et al. 2004. On the origin of fluorine in the Milky Way. Monthly Notices of the Royal Astronomical Society 354: 575–580.
- Paul Connett, James Beck and Spedding Micklem. 2010. The Case Against Fluoride: How Hazardous Waste Ended Up in Our Drinking Water and the Bad Science and Powerful Politics That Keep It There. Chelsea Green Publishing. 384 pages
- C.Michael Hogan, Richard Venti, Leda Patmore et al. 1972. Calibration of a roadway air pollutant dispersal model using sulfur hexafluoride. ESL Inc./U.S.Environmental Protection Agency
- Milton R. Beychok (2005), Fundamentals of Stack Gas Dispersion, 4th Edition, author-published, ISBN 0-9644588-0-2.
- Bradford D.Gessner, Michael Beller, John P.Middaugh, Gary M.Whitford. 1994. Acute fluoride poisoning from a public water system. New England Journal of Medicine 330 (2): 95–99.
- J.H.Nelson. 2003. Nuclear Magnetic Resonance Spectroscopy. Prentice Hall. ISBN 0-13-033451-0