Jump to content

Fluorine: Difference between revisions

Edit links
From Wikipedia, the free encyclopedia
Browse history interactively
← Previous editNext edit →
Content deleted Content added
VisualWikitext
Undid revision 279356293 by 193.188.46.64 (talk)
No edit summary
Line 1: Line 1:
{{Infobox fluorine}}
{{Infobox fluorine}}
'''Fluorine''' is a [[chemical element]], represented by the symbol '''F''', and the [[atomic number]] 9. Fluorine forms a single bond with itself in elemental form, resulting in the diatomic F<sub>2</sub> molecule. F<sub>2</sub> is a supremely reactive, [[poison]]ous, pale, yellowish brown gas. Elemental fluorine is the most chemically reactive and [[electronegativity|electronegative]] of all the elements. For example, it will readily "burn" [[hydrocarbons]] at room temperature, in contrast to the [[combustion]] of hydrocarbons by [[oxygen]], which requires an [[activation energy|input of energy]] with a spark. Therefore, molecular fluorine is highly dangerous, more so than other [[halogen]]s such as the poisonous [[chlorine]] gas.
'''Fluorine''' is a [[fucked up chemical element]], represented by the symbol '''F''', and the [[atomic number]] 9. Fluorine forms a single bond with itself in elemental form, resulting in the diatomic F<sub>2</sub> molecule. F<sub>2</sub> is a supremely reactive, [[poison]]ous, pale, yellowish brown gas. Elemental fluorine is the most chemically reactive and [[electronegativity|electronegative]] of all the elements. For example, it will readily "burn" [[hydrocarbons]] at room temperature, in contrast to the [[combustion]] of hydrocarbons by [[oxygen]], which requires an [[activation energy|input of energy]] with a spark. Therefore, molecular fluorine is highly dangerous, more so than other [[halogen]]s such as the poisonous [[chlorine]] gas.


Fluorine's highest electronegativity and small atomic radius give unique properties to many of its compounds. For example, the enrichment of <sup>235</sup>[[uranium|U]], the principal nuclear fuel, relies on the volatility of [[Uranium hexafluoride|UF<sub>6</sub>]]. Also, the [[carbon–fluorine bond]] is considered as the strongest bond in organic chemistry. This contributes to the stability and persistence of [[fluorocarbon|fluoroalkane]] based [[organofluorine compounds]], such as [[PTFE]]/(Teflon) and [[PFOS]]. The carbon–fluorine bond also sharply increases the efficacy of many [[pharmaceutical]]s and results in the strength of many [[superacid]]s.
Fluorine's highest electronegativity and small atomic radius give unique properties to many of its compounds. For example, the enrichment of <sup>235</sup>[[uranium|U]], the principal nuclear fuel, relies on the volatility of [[Uranium hexafluoride|UF<sub>6</sub>]]. Also, the [[carbon–fluorine bond]] is considered as the strongest bond in organic chemistry. This contributes to the stability and persistence of [[fluorocarbon|fluoroalkane]] based [[organofluorine compounds]], such as [[PTFE]]/(Teflon) and [[PFOS]]. The carbon–fluorine bond also sharply increases the efficacy of many [[pharmaceutical]]s and results in the strength of many [[superacid]]s.

Revision as of 13:28, 24 March 2009

Fluorine, 9F
Small sample of pale yellow liquid fluorine condensed in liquid nitrogen
Liquid fluorine (F2 at extremely low temperature)
Fluorine
Pronunciation
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
–
↑
F
↓
Cl
oxygen ← fluorine → neon
Atomic number (Z)9
Groupgroup 17 (halogens)
Periodperiod 2
Block  p-block
Electron configuration[He] 2s2 2p5[3]
Electrons per shell2, 7
Physical properties
Phase at STPgas
Melting point(F2) 53.48 K ​(−219.67 °C, ​−363.41 °F)[4]
Boiling point(F2) 85.03 K ​(−188.11 °C, ​−306.60 °F)[4]
Density (at STP)1.696 g/L[5]
when liquid (at b.p.)1.505 g/cm3[6]
Triple point53.48 K, ​.252 kPa[7]
Critical point144.41 K, 5.1724 MPa[4]
Heat of vaporization6.51 kJ/mol[5]
Molar heat capacityCp: 31 J/(mol·K)[6] (at 21.1 °C)
Cv: 23 J/(mol·K)[6] (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 statescommon: −1
ElectronegativityPauling scale: 3.98[3]
Ionization energies
  • 1st: 1681 kJ/mol
  • 2nd: 3374 kJ/mol
  • 3rd: 6147 kJ/mol
  • (more)[8]
Covalent radius64 pm[9]
Van der Waals radius135 pm[10]
Color lines in a spectral range
Spectral lines of fluorine
Other properties
Natural occurrenceprimordial
Crystal structure ​cubic
Cubic crystal structure for fluorine
Thermal conductivity0.02591 W/(m⋅K)[11]
Magnetic orderingdiamagnetic (−1.2×10−4)[12][13]
CAS Number7782-41-4[3]
History
Namingafter the mineral fluorite, itself named after Latin fluo (to flow, in smelting)
DiscoveryAndré-Marie Ampère (1810)
First isolationHenri Moissan[3] (June 26, 1886)
Named by
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

Fluorine is a fucked up chemical element, represented by the symbol F, and the atomic number 9. Fluorine forms a single bond with itself in elemental form, resulting in the diatomic F2 molecule. F2 is a supremely reactive, poisonous, pale, yellowish brown gas. Elemental fluorine is the most chemically reactive and electronegative of all the elements. For example, it will readily "burn" hydrocarbons at room temperature, in contrast to the combustion of hydrocarbons by oxygen, which requires an input of energy with a spark. Therefore, molecular fluorine is highly dangerous, more so than other halogens such as the poisonous chlorine gas.

Fluorine's highest electronegativity and small atomic radius give unique properties to many of its compounds. For example, the enrichment of 235U, the principal nuclear fuel, relies on the volatility of UF6. Also, the carbon–fluorine bond is considered as the strongest bond in organic chemistry. This contributes to the stability and persistence of fluoroalkane based organofluorine compounds, such as PTFE/(Teflon) and PFOS. The carbon–fluorine bond also sharply increases the efficacy of many pharmaceuticals and results in the strength of many superacids.

Characteristics

F2 is a corrosive pale yellow or brown[15] gas that is a powerful oxidizing agent. It is the most reactive and most electronegative of all the elements (4.0), and readily forms compounds with most other elements. It has an oxidation number -1, except when bonded to another fluorine in F2 which gives it an oxidation number of 0. Fluorine even combines with the noble gases argon, krypton, xenon, and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. The reaction with hydrogen occurs even at extremely low temperatures, using liquid hydrogen and solid fluorine. It is so reactive that metals, and even water, as well as other substances, burn with a bright flame in a jet of fluorine gas. In moist air it reacts with water to form also-dangerous hydrofluoric acid.

Fluorides are compounds that combine fluorine with some positively charged counterpart. They often consist of crystalline ionic salts. Fluorine compounds with metals are among the most stable of salts.

Hydrogen fluoride is a weak acid when dissolved in water. Consequently, fluorides of alkali metals produce basic solutions. For example, a 1 M solution of NaF in water has a pH of 8.59 compared to a 1 M solution of NaOH, a strong base, which has a pH of 14.00.[16]

Applications

Elemental fluorine, F2, is mainly used for the production of two compounds of commercial interest, uranium hexafluoride and sulfur hexafluoride.[17]

Industrial use of fluorine-containing compounds:

Dental and medical uses

Chemistry of fluorine

Fluorine forms a variety of very different compounds, owing to its small atomic size and covalent behavior. Elemental fluorine is a dangerously powerful oxidant, reflecting the extreme electronegativity of fluorine. Hydrofluoric acid is extremely dangerous, whereas in synthetic drugs incorporating an aromatic ring (e.g. flumazenil), fluorine is used to help prevent toxication or to delay metabolism.

The fluoride ion is basic, therefore hydrofluoric acid is a weak acid in water solution. However, water is not an inert solvent in this case: when less basic solvents such as anhydrous acetic acid are used, hydrofluoric acid is the strongest of the hydrohalogenic acids. Also, owing to the basicity of the fluoride ion, soluble fluorides give basic water solutions. The fluoride ion is a Lewis base, and has a high affinity to certain elements such as calcium and silicon. For example, deprotection of silicon protecting groups is achieved with a fluoride. The fluoride ion is poisonous.

Fluorine as a freely reacting oxidant gives the strongest oxidants known.

The reactivity of fluorine toward the noble gas xenon was first reported by Neil Bartlett in 1962. Fluorides of krypton and radon have also been prepared. Also argon fluorohydride has been observed at cryogenic temperatures.

The carbon-fluoride bond is covalent and very stable. The use of a fluorocarbon polymer, poly(tetrafluoroethene) or Teflon, is an example: it is thermostable and waterproof enough to be used in frying pans. Organofluorines may be safely used in applications such as drugs, without the risk of release of toxic fluoride. In synthetic drugs, toxication can be prevented. For example, an aromatic ring is useful but presents a safety problem: enzymes in the body metabolize some of them into poisonous epoxides. When the para position is substituted with fluorine, the aromatic ring is protected and epoxide is no longer produced.

The substitution of hydrogen for fluorine in organic compounds offers a very large number of compounds. An estimated fifth of pharmaceutical compounds and 30% of agrochemical compounds contain fluorine.[21] The -CF3 and -OCF3 moieties provide further variation, and more recently the -SF5 group.[22]


Fluorite (CaF2) crystals
See also: Category:Fluorine compounds

Production

Fluorine cell room at F2 Chemicals Ltd, Preston, UK

Industrial production of fluorine entails the electrolysis of hydrogen fluoride in the presence of potassium fluoride. This method is based on the pioneering studies by Moissan (see below). Fluorine gas forms at the anode, and hydrogen gas at the cathode. Under these conditions, the potassium fluoride (KF) converts to potassium bifluoride (KHF2), which is the actual electrolyte, This potassium bifluoride aids electrolysis by greatly increasing the electrical conductivity of the solution.

HF + KF → KHF2
2 KHF2 → 2 KF + H2 + F2

The HF required for the electrolysis is obtained as a byproduct of the production of phosphoric acid. Phosphate-containing minerals contain significant amounts of calcium fluorides, such as fluorite. Upon treatment with sulfuric acid, these minerals release hydrogen fluoride:

CaF2 + H2SO4 → 2 HF + CaSO4

In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl Christe discovered a purely chemical preparation involving the reaction of solutions in anhydrous HF, K2MnF6, and SbF5 at 150 °C:[23]

Template:PotassiumTemplate:ManganeseF6 + 2Template:AntimonyF5 → 2Template:PotassiumTemplate:AntimonyF6 + Template:ManganeseF3 + ½F2

Though not a practical synthesis on the large scale, this report demonstrates that electrolysis is not the sole route to the element.

History

The mineral fluorspar (also called fluorite), consisting mainly of calcium fluoride, was described in 1530 by Georgius Agricola for its use as a flux.[24] Fluxes are used to promote the fusion of metals or minerals. The etymology of the element's name reflects its history: Fluorine Template:Pron-en, /ˈflʊərɨn/, or commonly /ˈflɔr-/; from Template:Lang-la, meaning "to flow". In 1670 Schwanhard found that glass was etched when it was exposed to fluorspar that had been treated with acid. Carl Wilhelm Scheele and many later researchers, including Humphry Davy, Caroline Menard, Gay-Lussac, Antoine Lavoisier, and Louis Thenard all would experiment with hydrofluoric acid, easily obtained by treating fluorite with concentrated sulfuric acid.

Owing to its extreme reactivity, elemental fluorine was not isolated until many years after the characterization of fluorite. Progress in isolating elemental fluorine was slowed because it could only be prepared electrolytically and even then under stringent conditions since the gas attacks many materials. In 1886, the isolation of elemental fluorine was reported by Henri Moissan after almost 74 years of effort by other chemists.[25] The generation of elemental fluorine from hydrofluoric acid is exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. These individuals came to be referred to as "fluorine martyrs".[citation needed] For Moissan, it earned him the 1906 Nobel Prize in chemistry.

The first large-scale production of fluorine was undertaken in support of the Manhattan project, where the compound uranium hexafluoride (UF6) had been selected as the form of uranium that would allow separation of its 235U and 238U isotopes. Today both the gaseous diffusion process and the gas centrifuge process use gaseous UF6 to produce enriched uranium for nuclear power applications. In the Manhattan Project, it was found that UF6 decomposed into UF4 and F2. The corrosion problem due to the F2 was eventually solved by electrolytically coating all UF6 carrying piping with nickel metal, which forms a nickel difluoride that is not attacked by fluorine. Joints and flexible parts were made from teflon, then a very recently discovered fluorocarbon plastic which is also not attacked by F2.

Biological role

Main article: Fluorine deficiency

Though F2 is too reactive to have any natural biological role, fluorine is incorporated into compounds with biological activity. Naturally occurring organofluorine compounds are rare, the most notable example is fluoroacetate, which functions as a plant defence against herbivores in at least 40 plants in Australia, Brazil and Africa.[26] The enzyme adenosyl-fluoride synthase catalyzes the formation of 5'-deoxy-5'-fluoroadenosine. Additionally, fluoride might have a natural role in preventing tooth decay.[27]

Precautions

Elemental fluorine

Elemental fluorine (fluorine gas) is a highly toxic, corrosive oxidant, which can cause ignition of organic material. Fluorine gas has a characteristic pungent odor that is detectable in concentrations as low as 20 ppb. As it is so reactive, all materials of construction must be carefully selected and metal surfaces must be passivated.

Fluoride ion

Main article: Fluoride poisoning

Fluoride ions are toxic: the lethal dose of sodium fluoride for a 70 kg human is estimated at 5–10 g.[28]

Hydrogen fluoride and hydrofluoric acid

Main article: Hydrofluoric acid

Hydrogen fluoride and hydrofluoric acid are dangerous, far moreso than the related hydrochloric acid, because molecular HF penetrates the skin and biological membranes.

Organofluorines

Organofluorines are naturally rare compounds. They can be nontoxic (perflubron and perfluorodecalin) or highly toxic (perfluoroisobutylene and fluoroacetic acid). Many pharmacuticals are organofluorines, such as the anti-cancer fluorouracil. Perfluorooctanesulfonic acid (PFOS) is a persistent organic pollutant.

See also

Notes

  1. ^ "Standard Atomic Weights: Fluorine". CIAAW. 2021.
  2. ^ Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (2022-05-04). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.
  3. ^ a b c d Jaccaud et al. 2000, p. 381.
  4. ^ a b c Haynes 2011, p. 4.121.
  5. ^ a b Jaccaud et al. 2000, p. 382.
  6. ^ a b c Compressed Gas Association 1999, p. 365.
  7. ^ "Triple Point | The Elements Handbook at KnowledgeDoor". KnowledgeDoor.
  8. ^ Dean 1999, p. 4.6.
  9. ^ Dean 1999, p. 4.35.
  10. ^ Matsui 2006, p. 257.
  11. ^ Yaws & Braker 2001, p. 385.
  12. ^ Mackay, Mackay & Henderson 2002, p. 72.
  13. ^ Cheng et al. 1999.
  14. ^ Chisté & Bé 2011.
  15. ^ Theodore Gray. "Real visible fluorine". The Wooden Periodic Table.
  16. ^ "pKa's of Inorganic and Oxo-Acids" (PDF). Evans Group. Retrieved 2008-11-29.
  17. ^ M. Jaccaud, R. Faron, D. Devilliers, R. Romano “Fluorine” in Ullmann’s Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, 2005.
  18. ^ Leonel R Arana, Nuria de Mas, Raymond Schmidt, Aleksander J Franz, Martin A Schmidt and Klavs F Jensen (2007). "Isotropic etching of silicon in fluorine gas for MEMS micromachining". J. Micromech. Microeng. 17: 384. doi:10.1088/0960-1317/17/2/026.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  19. ^ "Class I Ozone-Depleting Substances". Ozone Depletion. U.S. Environmental Protection Agency.
  20. ^ eMedicine - Corticosteroid-Induced Myopathy : Article by Steve S Lim, MD
  21. ^ "Fluorine's treasure trove". ICIS news. 2006-10-02. Retrieved 2008-11-29.
  22. ^ Bernhard Stump, Christian Eberle, W. Bernd Schweizer, Marcel Kaiser, Reto Brun, R. Luise Krauth-Siegel, Dieter Lentz, François Diederich (2009). "Pentafluorosulfanyl as a Novel Building Block for Enzyme Inhibitors: Trypanothione Reductase Inhibition and Antiprotozoal Activities of Diarylamines". ChemBioChem. 10: 79. doi:10.1002/cbic.200800565.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  23. ^ K. Christe (1986). "Chemical synthesis of elemental fluorine". Inorg. Chem. 25: 3721–3724. doi:10.1021/ic00241a001. {{cite journal}}: Unknown parameter |years= ignored (help)
  24. ^ Fluoride History Discovery of fluorine
  25. ^ H. Moissan (1886). "Action d'un courant électrique sur l'acide fluorhydrique anhydre". Comptes rendus hebdomadaires des séances de l'Académie des sciences. 102: 1543–1544.
  26. ^ Proudfoot AT, Bradberry SM, Vale JA (2006). "Sodium fluoroacetate poisoning". Toxicol Rev. 25 (4): 213–9. doi:10.2165/00139709-200625040-00002. PMID 17288493.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  27. ^ Yeung CA (2008). "A systematic review of the efficacy and safety of fluoridation". Evid Based Dent. 9 (2): 39–43. doi:10.1038/sj.ebd.6400578. PMID 18584000.
  28. ^ Aigueperse, Jean (2005), "Fluorine Compounds, Inorganic", in Ullmann (ed.), Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH {{citation}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)

References

Wikimedia Commons has media related to Fluorine.
Look up fluorine in Wiktionary, the free dictionary.

Template:Link FA

Retrieved from "https://en.wikipedia.org/w/index.php?title=Fluorine&oldid=279356725"