Fluorine is a chemical element with the symbol F and atomic number 9. It is the lightest halogen and exists at standard conditions as a highly toxic, pale yellow diatomic gas. As the most electronegative element, it is extremely reactive, as it reacts with all other elements except for argon, neon, and helium.

Fluorine, 9F
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)
  • 18.998403162±0.000000005
  • 18.998±0.001 (abridged)[1]
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


Atomic number (Z)9
Groupgroup 17 (halogens)
Periodperiod 2
Block  p-block
Electron configuration[He] 2s2 2p5[2]
Electrons per shell2, 7
Physical properties
Phase at STPgas
Melting point(F2) 53.48 K (−219.67 °C, −363.41 °F)[3]
Boiling point(F2) 85.03 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]
Spectral lines of fluorine
Other properties
Natural occurrenceprimordial
Crystal structure cubic
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 by
Main isotopes of fluorine[13]
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct
18F trace 109.8 min β+ (97%) 18O
ε (3%) 18O
19F 100% stable
 Category: Fluorine
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Among the elements, fluorine ranks 24th in universal abundance and 13th in terrestrial abundance. Fluorite, the primary mineral source of fluorine which gave the element its name, was first described in 1529; as it was added to metal ores to lower their melting points for smelting, the Latin verb fluo meaning 'flow' gave the mineral its name. Proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, and several early experimenters died or sustained injuries from their attempts. Only in 1886 did French chemist Henri Moissan isolate elemental fluorine using low-temperature electrolysis, a process still employed for modern production. Industrial production of fluorine gas for uranium enrichment, its largest application, began during the Manhattan Project in World War II.

Owing to the expense of refining pure fluorine, most commercial applications use fluorine compounds, with about half of mined fluorite used in steelmaking. The rest of the fluorite is converted into corrosive hydrogen fluoride en route to various organic fluorides, or into cryolite, which plays a key role in aluminium refining. Molecules containing a carbon–fluorine bond often have very high chemical and thermal stability; their major uses are as refrigerants, electrical insulation and cookware, the last as PTFE (Teflon). Pharmaceuticals such as atorvastatin and fluoxetine contain C−F bonds. The fluoride ion from dissolved fluoride salts inhibits dental cavities, and so finds use in toothpaste and water fluoridation. Global fluorochemical sales amount to more than US$69 billion a year.

Fluorocarbon gases are generally greenhouse gases with global-warming potentials 100 to 23,500 times that of carbon dioxide, and SF6 has the highest global warming potential of any known substance. Organofluorine compounds often persist in the environment due to the strength of the carbon–fluorine bond. Fluorine has no known metabolic role in mammals; a few plants and sea sponges synthesize organofluorine poisons (most often monofluoroacetates) that help deter predation.[14]

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