Fluorine From Wikipedia, the free encyclopedia page is in the middle of an expansion or major revamping This article or section is in the process of an expansion or major restructuring. You are welcome to assist in its construction by editing it as well. If this article or section has not been edited in several days, please remove this template. This article was last edited by ClueBot NG (talk | contribs) 26 seconds ago. (Purge) Fluorine 9F ↑ F ↓ Cl oxygen ← fluorine → neon Fluorine in the periodic table Appearance gas: very pale yellow liquid: bright yellow solid: transparent (beta), opaque (alpha) Small sample of pale yellow liquid fluorine condensed in liquid nitrogen Liquid fluorine at cryogenic temperatures General properties Name, symbol, number fluorine, F, 9 Pronunciation /ˈflʊəriːn/ fluu-reen, /ˈflʊərɪn/, /ˈflɔəriːn/ Element category diatomic nonmetal Group, period, block 17 (halogens), 2, p Standard atomic weight 18.998403163(6) Electron configuration [He] 2s2 2p5[1] 2, 7 Physical properties Phase gas Density (0 °C, 101.325 kPa) 1.696[2] g/L Liquid density at b.p. 1.505[3] g·cm−3 Melting point 53.48 K, −219.67 °C, −363.41[4] °F Boiling point 85.03 K, −188.11 °C, −306.60[4] °F Triple point 53.48 K, 90[4] kPa Critical point 144.41 K, 5.1724[4] MPa Heat of vaporization 6.51[2] kJ·mol−1 Molar heat capacity (Cp) (21.1 °C) 31[3] J·mol−1·K−1 (Cv) (21.1 °C) 23[3] J·mol−1·K−1 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 (oxidizes oxygen) Electronegativity 3.98[1] (Pauling scale) Ionization energies (more) 1st: 1681[5] kJ·mol−1 2nd: 3374[5] kJ·mol−1 3rd: 6147[5] kJ·mol−1 Covalent radius 64[6] pm Van der Waals radius 135[7] pm Miscellanea Crystal structure monoclinic Fluorine has a monoclinic base-centered crystal structure alpha state (low-temperature)[8] Magnetic ordering diamagnetic (−1.2×10−4 (SI)[9][10]) Thermal conductivity 0.02591[11] W·m−1·K−1 CAS registry number 7782-41-4[1] History Naming after the mineral fluorite, itself named after Latin fluo (to flow, in smelting) Discovery André-Marie Ampère (1810) First isolation Henri Moissan[1] (June 26, 1886) Named by Humphry Davy Most stable isotopes Main article: Isotopes of fluorine iso NA half-life DM DE (MeV) DP 18F trace 109.77 min β+ (96.9%) 0.634 18O ε (3.1%) 1.656 18O 19F 100% 19F is stable with 10 neutrons reference[12] v t e · references Fluorine is an extremely reactive and poisonous chemical element with atomic number 9. The lightest halogen and most electronegative element, it exists as a pale yellow diatomic gas at standard conditions. Almost all other elements, including some noble gases, form compounds with fluorine. Fluorite (calcium fluoride, CaF 2), the primary mineral source of fluorine, was first described in 1529; trace amounts of F2 lie embedded within it. Contemporarily the Latin verb fluo, meaning "flow", became associated with fluorite rocks as an additive to metal ores which lowered melting points for smelting. Proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, with several early experimenters dying or sustaining injuries from their attempts. In 1886, French chemist Henri Moissan succeeded in isolating elemental fluorine using low-temperature electrolysis, a process still employed for modern production. The element is 24th in universal abundance and 13th in terrestrial abundance. Due to the expense of refining pure fluorine, nearly all commercial applications handle it bound to compounds. About half of mined fluorite goes into steelmaking, the rest converted into corrosive hydrogen fluoride, a precursor to various organic fluorides and the critical aluminium refining material cryolite. Organic fluorides have very high chemical and thermal stability, their major uses being refrigerants and – as PTFE – cookware, as well as electrical insulation. Pharmaceuticals such as atorvastatin and fluoxetine also contain fluorine, while the fluoride ion inhibits dental cavities, thus finding use in toothpaste and water fluoridation. Uranium enrichment, the largest application for free fluorine, began in World War II during the Manhattan Project. Global fluorochemical sales amount to over US$15 billion a year. Fluorocarbon gases are generally greenhouse gases, with global-warming potentials 100 to 20,000 (for sulfur hexafluoride) times that of carbon dioxide. Organofluorines persist in the environment due to the carbon–fluorine bond's strength, but the potential health impact of the most persistent such compounds is unclear. While a few plants and bacteria synthesise organofluorine poisons for defence against herbivores, fluorine has no metabolic role in mammals. Contents [hide] 1 Characteristics 1.1 Chemical reactivity 1.2 Toxicity 1.3 Phases 1.4 Electron arrangement 1.5 Isotopes 2 Occurrence 2.1 Universe 2.2 Earth 3 Compounds 3.1 Metal fluorides 3.2 Hydrogen fluoride 3.3 Nonmetal fluorides 3.4 Noble gas compounds 3.5 Organic compounds 4 History 4.1 Early discoveries and etymology 4.2 Isolation 4.3 Application development 5 Industry and applications 5.1 Inorganic fluorides 5.2 Organic fluorochemicals 5.3 Fluorine gas 6 Production of fluorine gas 6.1 Industrial 6.2 Chemical 7 Environmental concerns 7.1 Atmosphere 7.2 Biopersistance 8 Biological aspects 8.1 Natural biochemistry 8.2 Medicine 8.3 Agrichemicals and poisons 9 Fluorine-related hazards 9.1 Hydrofluoric acid 9.2 Fluoride ion 10 See also 11 Notes 12 Sources Characteristics[edit] Chemical reactivity[edit] Main article: Chemical characteristics of fluorine Fluorine's superlative reactivity stems from two reasons.[13] Compared with Cl 2 and Br 2, difluorine's bond energy is much lower, similar to those of weak peroxide bonds;[14][13] elemental fluorine thus dissociates easily in reactions. Conversely, bonds to other atoms are very strong because of fluorine's high electronegativity. Otherwise inert substances like powdered steel, glass fragments and asbestos fibres react quickly with cold fluorine gas, while wood and water spontaneously combust under a fluorine jet.[2][15] External video Bright flames during fluorine reactions Fluorine reacting with caesium Reactions of elemental fluorine with metals require varying conditions: alkali metals cause explosions and alkaline earth metals display vigorous activity in bulk, but most other metals such as aluminium and iron must be powdered to prevent protective metal fluoride layers from passivating,[13] and noble metals require pure fluorine gas at 300–450 °C.[16] Metalloids and some solid nonmetals (sulfur, phosphorus, selenium) burn with a flame in room temperature fluorine.[17][18] Hydrogen sulfide and sulfur dioxide combine readily with fluorine, the latter with chances of exploding, but sulfuric acid exhibits much less activity.[17] Hydrogen, analogous to alkali metals, reacts explosively with fluori