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Fluorine, 00F
Small sample of pale yellow liquid fluorine condensed in liquid nitrogen
Liquid fluorine (at extremely low temperatures)
Allotropesalpha, beta (see Allotropes of fluorine)
Appearancegas: very pale yellow
liquid: bright yellow
solid: alpha is opaque, beta is transparent
Standard atomic weight Ar°(F)
Fluorine in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson


Groupgroup 17 (halogens)
Periodperiod 2
Block  p-block
Electron configuration[He] 2s2 2p5[2]
Electrons per shell2, 7
Physical properties
Phase at STPgas
Melting point53.48 K ​(−219.67 °C, ​−363.41 °F)[3]
Boiling point85.03 K ​(−188.11 °C, ​−306.60 °F)[3]
Density (at STP)1.696 g/L[4]
when liquid (at b.p.)1.505 g/cm3[5]
Triple point53.48 K, ​90 kPa[3]
Critical point144.41 K, 5.1724 MPa[3]
Heat of vaporization6.51 kJ/mol[4]
Molar heat capacityCp: 31 J/(mol·K)[5] (at 21.1 °C)
Cv: 23 J/(mol·K)[5] (at 21.1 °C)
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 38 44 50 58 69 85
Atomic properties
Oxidation states−1, 0[6] (oxidizes oxygen)
ElectronegativityPauling scale: 3.98[2]
Ionization energies
  • 1st: 1681 kJ/mol
  • 2nd: 3374 kJ/mol
  • 3rd: 6147 kJ/mol
  • (more)[7]
Covalent radius64 pm[8]
Van der Waals radius135 pm[9]
Color lines in a spectral range
Spectral lines of fluorine
Other properties
Natural occurrenceprimordial
Crystal structurecubic
Cubic crystal structure for fluorine
Thermal conductivity0.02591 W/(m⋅K)[10]
Magnetic orderingdiamagnetic (−1.2×10−4)[11][12]
CAS Number7782-41-4[2]
Namingafter the mineral fluorite, itself named after Latin fluo (to flow, in smelting)
DiscoveryAndré-Marie Ampère (1810)
First isolationHenri Moissan[2] (June 26, 1886)
Named byHumphry Davy
Isotopes of fluorine
Main isotopes Decay
abun­dance half-life (t1/2) mode pro­duct
18F trace 109.734 min β+ 18O
19F 100% stable
 Category: Fluorine
| references
A more real picture 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 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]

Chemical compounds containing fluorine ions are called fluorides. Fluorine only exists in one oxidation state: -1.

Occurrence[change | change source]

Fluorite crystals, the "ore" of fluorine

Fluorine is not found as an element on the earth becase it is 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]

  1. "Standard Atomic Weights: Fluorine". CIAAW. 2013.
  2. 2.0 2.1 2.2 2.3 2.4 Jaccaud et al. 2000, p. 381.
  3. 3.0 3.1 3.2 3.3 Haynes 2011, p. 4.121.
  4. 4.0 4.1 Jaccaud et al. 2000, p. 382.
  5. 5.0 5.1 5.2 Compressed Gas Association 1999, p. 365.
  6. Himmel, D.; Riedel, S. (2007). "After 20 Years, Theoretical Evidence That 'AuF7' Is Actually AuF5·F2". Inorganic Chemistry. 46 (13). 5338–5342. doi:10.1021/ic700431s.
  7. Dean 1999, p. 4.6.
  8. Dean 1999, p. 4.35.
  9. Matsui 2006, p. 257.
  10. Yaws & Braker 2001, p. 385.
  11. Mackay, Mackay & Henderson 2002, p. 72.
  12. Cheng et al. 1999.
  13. 13.0 13.1 Chisté & Bé 2011.