|gas: very pale yellow
liquid: bright yellow
solid: transparent (beta), opaque (alpha)
Liquid fluorine at very cold temperatures
|Name, symbol, number||fluorine, F, 9|
|Pronunciation||//, //, //|
|Group, period, block||17, 2, p|
|Standard atomic weight||18.9984032(5) g/mol|
|Electron configuration||[He] 2s2 2p5|
|Electrons per shell||2, 7 (Image)|
|Density||(0 °C, 101.325 kPa)
|Liquid density at b.p.||1.505 g/cm3|
|Melting point||53.53 K, −219.62 °C, −363.32 °F|
|Boiling point||85.03 K, −188.12 °C, −306.62 °F|
|Critical point||144.4 K, 5.215 MPa|
|Heat of vaporization||6.51 kJ/mol|
|Specific heat capacity||(25 °C) (Cp) (21.1 °C) 825 J·mol−1·K−1
(Cv) (21.1 °C) 610 J/(mol·K)
|Electronegativity||3.98 (Pauling scale)|
||1st: 1,681 kJ/mol|
|2nd: 3,374 kJ/mol|
|3rd: 6,147 kJ/mol|
|Covalent radius||64 pm|
|Van der Waals radius||135 pm|
|Crystal structure note||the structure shows solid fluorine, just under the melting point, 1 atm|
|Thermal conductivity||(300 K) 0.02591 W/(m·K)|
|CAS registry number||7782-41-4|
|Most stable isotopes|
|Main article: Isotopes of fluorine|
Fluorine (symbol F) is a chemical element that is very poisonous. Its atomic number (which is the number of protons in it) is 9, and its atomic mass is 19. It is part of the Group 7 (halogens) on the periodic table of elements.
Properties[change | change source]
Fluorine is a light yellow diatomic gas. It is very reactive gas, which exists as diatomic molecules. It is actually the most reactive element. Fluorine has a very high attraction for electrons, because it is missing one. This makes it the most powerful oxidizing agent. It can rip electrons from water (making oxygen) and ignite propane on contact. It does not need a spark. Metals can catch on fire when placed in a stream of fluorine. After it is reduced by reacting with other things, it forms the stable fluoride ion. Fluorine is very poisonous. Fluorine bonds very strongly with carbon. It can react with the unreactive noble gases. It explodes when mixed with hydrogen. The melting point of fluorine is -363.33°F (-219.62°C), the boiling point is -306.62°F (-188.12°C).
Chemical compounds[change | change source]
- Aluminium fluoride
- Antimony trifluoride
- Antimony pentafluoride
- Arsenic trifluoride
- Arsenic pentafluoride
- Bismuth(III) fluoride
- Bismuth(V) fluoride
- Bromine trifluoride
- Bromine pentafluoride
- Chlorine monofluoride
- Chlorine trifluoride
- Cobalt(II) fluoride
- Cobalt(III) fluoride
- Disulfur decafluoride
- Hydrofluoric acid, a solution of hydrogen fluoride in water
- Hydrogen fluoride
- Iodine trifluoride
- Iodine pentafluoride
- Iodine heptafluoride
- Manganese(II) fluoride
- Manganese(III) fluoride
- Manganese(IV) fluoride
- Potassium fluoride
- Selenium tetrafluoride
- Selenium hexafluoride
- Silver(I) fluoride, brown-yellow
- Silver(II) fluoride, highly reactive, white or gray
- Sodium aluminium fluoride, cryolite
- Sodium fluoride
- Sulfur hexafluoride
- Sulfur tetrafluoride
- Tellurium(IV) fluoride
- Tellurium(VI) fluoride
- Thallium(I) fluoride
- Thallium(III) fluoride
- Tin(II) fluoride
- Tin(IV) fluoride
- Zinc fluoride
Occurrence[change | change source]
Fluorine is not found as an element on the earth; it is much too reactive. Several fluorides are found in the earth, though. When calcium phosphate is reacted with sulfuric acid to make phosphoric acid, some hydrofluoric acid is produced. Also, fluorite can be reacted with sulfuric acid to make hydrofluoric acid. Fluorite naturally occurs on the earths' crust in rocks, coal and clay.
Preparation[change | change source]
Fluorine is normally made by electrolysis. Hydrogen fluoride is dissolved in potassium fluoride. This mixture is melted and an electric current is passed through it. This is electrolysis. Hydrogen is produced at one side and fluorine at the other side. If the sides are not separated, the cell may explode.
Someone made fluorine in 1986 without using electrolysis. They produced manganese(IV) fluoride by using various chemical compounds, which released fluorine gas.
Uses[change | change source]
Fluorine is used to enrich uranium for nuclear weapons. It is also used to make sulfur hexafluoride. Sulfur hexafluoride is used to propel stuff out of an aerosol can. It is also used to make integrated circuits. Fluorine compounds have many uses. Fluoride ions are in fluorine compounds. Fluoride ions can be in toothpaste. Some are used in nonstick coatings. Freons contain fluorine.
Safety[change | change source]
Fluorine as an element is extremely reactive and toxic. It can react with almost everything, even glass. Fluorine is also poisonous.
Fluoride ions are somewhat toxic. If too much toothpaste containing fluoride is eaten then fluoride poisoning may occur. Fluoride is not reactive, though.
Related pages[change | change source]
Sources[change | change source]
- Wieser, Michael E.; Coplen, Tyler B. (2010). "Atomic weights of the elements 2009 (IUPAC Technical Report)". Pure and Applied Chemistry 83: 359–396. doi:10.1351/PAC-REP-10-09-14.
- Aigueperse et al. 2005, "Fluorine", p. 1.
- Aigueperse et al. 2005, "Fluorine", p. 2.
- Compressed Gas Association (1999). Handbook of compressed gases. Springer. p. 365. ISBN 9780412782305.
- Dean 1999, p. 3.29.
- Dean 1999, p. 4.6.
- Dean 1999, p. 4.35.
- Kim, Sung-Hoon (2006), Functional dyes, Elsevier, p. 257, ISBN 9780444521763
- Young, David A. (2011-06-10) Phase Diagrams of the Elements . Springer, 10. Report.
- Mackay, Mackay & Henderson 2002, p. 72.
- Yaws & Braker 2001, p. 385.
- Chiste, V.; Be, M. M. (2006). "F-18". Table de radionucleides. Laboratoire National Henri Becquerel. http://www.nucleide.org/DDEP_WG/Nuclides/F-18_tables.pdf. Retrieved 15 June 2011.