Fluorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Fluorine[2/19/2010 12:15:30 AM]

navigation

Main pageContentsFeatured contentCurrent eventsRandom article

interaction

About WikipediaCommunity portalRecent changesContact WikipediaDonate to WikipediaHelp

toolbox

What links hereRelated changesUpload fileSpecial pagesPrintable versionPermanent linkCite this page

discussion view source historyTry Beta Log in / create account

oxygen ← fluorine → neon

-↑F↓Cl

Appearance

Yellowish gas

General properties

Name, symbol,number

fluorine, F, 9

Element category halogen

Group, period,block

17, 2, p

Standard atomicweight

18.9984032(5) g·mol−1

Electronconfiguration

1s2 2s2 2p5

Electrons per shell 2, 7 (Image)

Physical properties

Phase gas

Density (0 °C, 101.325 kPa)

1.7 g/L

Melting point 53.53 K, −219.62 °C,

−363.32 °F

Boiling point 85.03 K, −188.12 °C,

−306.62 °F

Critical point 144.13 K, 5.172 MPa

Heat of fusion (F2) 0.510 kJ·mol−1

FluorineFrom Wikipedia, the free encyclopedia

Fluorine is the chemical element with atomic number 9,represented by the symbol F. Fluorine forms a single bond withitself in elemental form, resulting in the diatomic F2 molecule.F2 is a supremely reactive, poisonous, pale, yellowish browngas. Elemental fluorine is the most chemically reactive andelectronegative of all the elements. For example, it will readily"burn" hydrocarbons at room temperature, in contrast to thecombustion of hydrocarbons by oxygen, which requires an inputof energy with a spark. Therefore, molecular fluorine is highlydangerous, more so than other halogens such as the poisonouschlorine gas.

Fluorine's highest electronegativity and small atomic radius giveunique properties to many of its compounds. For example, theenrichment of 235U, the principal nuclear fuel, relies on thevolatility of UF6. Also, the carbon–fluorine bond is one of thestrongest bonds in organic chemistry. This contributes to thestability and persistence of fluoroalkane based organofluorinecompounds, such as PTFE/(Teflon) and PFOS. The carbon–fluorine bond's inductive effects result in the strength of manyfluorinated acids, such as triflic acid and trifluoroacetic acid.Drugs are often fluorinated at biologically reactive positions, toprevent their metabolism and prolong their half-lives.

Contents [hide]

1 Characteristics2 Isotopes3 Applications

3.1 Industrial use of fluorine-containing compounds3.2 Dental and medical uses

4 Chemistry of fluorine5 Production6 History7 Biological role8 Precautions

8.1 Elemental fluorine8.2 Fluoride ion8.3 Hydrogen fluoride and hydrofluoric acid8.4 Organofluorines

9 See also10 References11 External links

Characteristics

F2 is a corrosive pale yellow or brown[1] gas that is a powerfuloxidizing agent. It is the most reactive and most electronegativeof all the elements on the classic Pauling scale (4.0), and

Periodic table

9F

article

search

languages

Afrikaans

AsturianuAzərbaycan

Bân-lâm-gúBosanskiБългарскиCatalàЧ вашлаČeskyCorsuCymraegDanskDeutschEestiΕλληνικάEspañolEsperanto

Fluorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Fluorine[2/19/2010 12:15:30 AM]

Heat ofvaporization

(F2) 6.62 kJ·mol−1

Specific heatcapacity

(25 °C) (F2)

31.304 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

(Weaklyacidic oxide)

Electronegativity 3.98 (Pauling scale)

Ionization energies(more)

1st: 1681.0 kJ·mol−1

2nd: 3374.2 kJ·mol−1

3rd: 6050.4 kJ·mol−1

Covalent radius 57±3 pm(see covalent radius of fluorine)

Van der Waalsradius

147 pm

Miscellanea

Crystal structure cubic

Magnetic ordering nonmagnetic

Thermalconductivity

(300 K) 27.7 m W·m−1·K−1

CAS registrynumber

7782-41-4

Most stable isotopes

Main article: Isotopes of fluorine

iso NA half-life DM DE (MeV) DP

18F syn 109.77 min β+ (97%) 0.64 18O

ε (3%) 1.656 18O

19F 100% 19F is stable with 10 neutrons

This box: view • talk • edit

readily forms compounds with most other elements. It has anoxidation number -1, except when bonded to another fluorine inF2 which gives it an oxidation number of 0. Fluorine combineswith the noble gases argon, krypton, xenon, and radon. Even indark, cool conditions, fluorine reacts explosively with hydrogen.The reaction with hydrogen can occur at extremely lowtemperatures, using liquid hydrogen and solid fluorine. It is soreactive that metals, water, as well as most other substances,burn with a bright flame in a jet of fluorine gas. In moist air, itreacts with water to form the also dangerous hydrofluoric acid.

Fluorides are compounds that combine fluorine with somepositively charged counterpart. They often consist of crystallineionic salts. Fluorine compounds with metals are among the moststable of salts.

Hydrogen fluoride is a weak acid when dissolved in water, butis still very corrosive and attacks glass. Consequently, fluoridesof alkali metals produce basic solutions. For example, a 1 Msolution of NaF in water has a pH of 8.59 compared to a 1 Msolution of NaOH, a strong base, which has a pH of 14.00.[2]

IsotopesMain article: Isotopes of fluorine

Although fluorine (F) has multiple isotopes, only one of theseisotopes (F-19) is stable, and the others have short half-livesand are not found in nature. Fluorine is thus a mononuclidicelement.

The nuclide 18F is the radionuclide of fluorine with the longesthalf life (about 110 minutes), and commercially is an importantsource of positrons, finding its major use in positron emissiontomography scanning.

ApplicationsElemental fluorine, F2, is mainly used for the production of twocompounds of commercial interest, uranium hexafluoride andsulfur hexafluoride.[3]

Industrial use of fluorine-containingcompounds

Atomic fluorine and molecular fluorine are used for plasmaetching in semiconductor manufacturing, flat panel displayproduction and MEMS (microelectromechanical systems)fabrication.[4] Xenon difluoride is also used for this last purpose.Hydrofluoric acid (chemical formula HF) is used to etch glass inlight bulbs and other products.Tetrafluoroethylene and perfluorooctanoic acid (PFOA) aredirectly used in the production of low friction plastics such asTeflon (or polytetrafluoroethylene).Fluorine is used indirectly in the production of halons such asfreon.Along with some of its compounds, fluorine is used in the

Euskara

FrançaisFryskFurlanGaeilgeGaelgGalegoHak-kâ-fa

Hawai`i

HrvatskiIdoBahasa IndonesiaÍslenskaItalianoעברית

KiswahiliKreyòl ayisyenKurdî / LatinaLatviešuLëtzebuergeschLietuviųLíguruLojbanMagyarМакедонски

Māori

Bahasa MelayuМонголNederlands

Norsk (bokmål)Norsk (nynorsk)NovialOccitanO'zbek

PlattdüütschPolskiPortuguêsRomânăRuna SimiРусскийSeelterskShqipSicilianuSimple EnglishSlovenčina

Fluorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Fluorine[2/19/2010 12:15:30 AM]

production of pure uranium from uranium hexafluoride and in thesynthesis of numerous commercial fluorochemicals, includingvitally important pharmaceuticals, agrochemical compounds,lubricants, and textiles.Fluorochlorohydrocarbons are used extensively in airconditioning and in refrigeration. Chlorofluorocarbons have beenbanned for these applications because they contribute to ozonedestruction and the ozone hole. Interestingly, since it is chlorineand bromine radicals which harm the ozone layer, not fluorine,compounds which do not contain chlorine or bromine but containonly fluorine, carbon and hydrogen (called hydrofluorocarbons)are not on the United States Environmental Protection Agency list of ozone-depleting substances,[5] andhave been widely used as replacements for the chlorine- and bromine-containing fluorocarbons.Hydrofluorocarbons do have a greenhouse effect, but a small one compared with carbon dioxide andmethane.Sodium hexafluoroaluminate (cryolite), is used in the electrolysis of aluminium.In much higher concentrations, sodium fluoride has been used as an insecticide, especially againstcockroaches.Fluorides have been used in the past to help molten metal flow. Hence the name, which derives fromLatin verb fluere, meaning to flow.[6]

Some researchers including US space scientists in the early 1960s have studied elemental fluorine gas asa possible rocket propellant due to its exceptionally high specific impulse. The experiments failed becausefluorine proved difficult to handle, and its combustion products proved extremely toxic and corrosive.Compounds of fluorine such as fluoropolymers, potassium fluoride and cryolite are utilized in applicationssuch as anti-reflective coatings and dichroic mirrors on account of their unusually low refractive index.

Dental and medical usesInorganic compounds of fluoride, including sodium fluoride (NaF), stannous fluoride (SnF2) and sodiumMFP, are used in toothpaste to prevent dental cavities. These or related compounds are also added tosome municipal water supplies, a process called water fluoridation, although the practice has remainedcontroversial since its beginnings in 1945.Many important agents for general anesthesia such as sevoflurane, desflurane, and isoflurane arehydrofluorocarbon derivatives.The fluorinated antiinflammatories dexamethasone and triamcinolone are among the most potent of thesynthetic corticosteroids class of drugs.[7]

Fludrocortisone ("Florinef") is one of the most common mineralocorticoids, a class of drugs which mimicsthe actions of aldosterone.Fluconazole is a triazole antifungal drug used in the treatment and prevention of superficial and systemicfungal infections.Fluoroquinolones are a family of broad-spectrum antibiotics.SSRI antidepressants, except in a few instances, are fluorinated molecules. These include citalopram,escitalopram oxalate, fluoxetine, fluvoxamine maleate, and paroxetine. A notable exception is sertraline.

Fluorite (CaF2) crystals

SlovenščinaСрпски / SrpskiSrpskohrvatski /СрпскохрватскиSuomiSvenska

То икTürkçeУкраїнськаUyghurche / Tiếng ViệtWinarayYorùbá

Fluorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Fluorine[2/19/2010 12:15:30 AM]

Because of the difficulty of biological systems in dealing with metabolism of fluorinated molecules,fluorinated antibiotics and antidepressants are among the major fluorinated organics found in treated citysewage and wastewater.

Compounds containing 18F, a radioactive isotope that emits positrons, are often used in positron emissiontomography, because its half-life of 110 minutes is long by the standards of positron-emitters. One suchspecies is fluorodeoxyglucose.

Chemistry of fluorineFluorine 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[citation needed].

The fluoride ion is basic, therefore hydrofluoric acid is a weak acid in water solution. However, water is not aninert solvent in this case: when less basic solvents such as anhydrous acetic acid are used, hydrofluoric acidis the strongest of the hydrohalogenic acids. Also, owing to the basicity of the fluoride ion, soluble fluoridesgive basic water solutions. The fluoride ion is a Lewis base, and has a high affinity to certain elements suchas calcium and silicon. For example, deprotection of silicon protecting groups is achieved with a fluoride. Thefluoride 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 ofkrypton and radon have also been prepared. Argon fluorohydride has been observed at cryogenictemperatures.

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 fryingpans. Organofluorines may be safely used in applications such as drugs, without the risk of release of toxicfluoride. In synthetic drugs, toxication can be prevented. For example, an aromatic ring is useful but presentsa safety problem: enzymes in the body metabolize some of them into poisonous epoxides. When the paraposition is substituted with fluorine, the aromatic ring is protected and epoxide is no longer produced.

The substitution of fluorine for hydrogen in organic compounds offers a very large number of compounds. Anestimated fifth of pharmaceutical compounds and 30% of agrochemical compounds contain fluorine.[8] The -CF3 and -OCF3 moieties provide further variation, and more recently the -SF5 group.[9]

See also: Category:fluorine compounds

ProductionIndustrial production of fluorine entails the electrolysis of hydrogenfluoride in the presence of potassium fluoride. This method isbased on the pioneering studies by Moissan (see below). Fluorinegas forms at the anode, and hydrogen gas at the cathode. Underthese conditions, the potassium fluoride (KF) converts to potassiumbifluoride (KHF2), which is the actual electrolyte. This potassiumbifluoride aids electrolysis by greatly increasing the electricalconductivity of the solution.

HF + KF → KHF22 KHF2 → 2 KF + H2 + F2

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

Fluorine cell room at F2 Chemicals Ltd,Preston, UK

Fluorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Fluorine[2/19/2010 12:15:30 AM]

CaF2 + H2SO4 → 2 HF + CaSO4

In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, KarlChriste discovered a purely chemical preparation involving the reaction of solutions in anhydrous HF,K2MnF6, and SbF5 at 150 °C:[10]

2 K2MnF6 + 4 SbF5 → 4 KSbF6 + 2 MnF3 + F2

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

HistoryThe mineral fluorspar (also called fluorite), consisting mainly of calcium fluoride, was described in 1530 byGeorgius Agricola for its use as a flux.[11] Fluxes are used to promote the fusion of metals or minerals. Theetymology of the element's name reflects its history: Fluorine is pronounced /'fl əri n/, /'fl ər n/, orcommonly /'fl r-/; from Latin: fluere, meaning "to flow". In 1670 Schwanhard found that glass was etchedwhen it was exposed to fluorspar that had been treated with acid. Carl Wilhelm Scheele and many laterresearchers, including Humphry Davy, Caroline Menard, Gay-Lussac, Antoine Lavoisier, and Louis Thenardall 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 characterizationof fluorite. Progress in isolating elemental fluorine was slowed because it could only be preparedelectrolytically and even then under stringent conditions since the gas attacks many materials. In 1886, theisolation of elemental fluorine was reported by Henri Moissan after almost 74 years of effort by otherchemists.[12] The generation of elemental fluorine from hydrofluoric acid is exceptionally dangerous, killing orblinding several scientists who attempted early experiments on this halogen. These individuals came to bereferred to as "fluorine martyrs".[13] For Moissan, it earned him the 1906 Nobel Prize in chemistry.[14]

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

Biological roleThough F2 is too reactive to have any natural biological role, fluorine is incorporated into compounds withbiological activity. Naturally occurring organofluorine compounds are rare, the most notable example isfluoroacetate, which functions as a plant defence against herbivores in at least 40 plants in Australia, Braziland Africa.[15] The enzyme adenosyl-fluoride synthase catalyzes the formation of 5'-deoxy-5'-fluoroadenosine. Fluorine is not an essential nutrient, but its importance in preventing tooth decay is well-recognized.[16] The effect is predominantly topical, although prior to 1981 it was considered primarilysystemic (occurring through ingestion).[17]

Precautions

Elemental fluorineElemental fluorine (fluorine gas) is a highly toxic, corrosive oxidant, which can cause ignition of organicmaterial. 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 bepassivated.

Fluorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Fluorine[2/19/2010 12:15:30 AM]

Fluoride ionMain article: Fluoride poisoning

Fluoride ions are toxic: the lethal dose of sodium fluoride for a 70 kg human is estimated at 5–10 g.[18]

Hydrogen fluoride and hydrofluoric acidMain article: Hydrofluoric acid

Hydrogen fluoride and hydrofluoric acid are dangerous, far more so than the related hydrochloric acid,because undissociated molecular HF penetrates the skin and biological membranes, causing deep andpainless burns. The free fluoride, once released from HF in dissociation, also is capable of chelating calciumion to the point of causing death by cardiac dysrhythmia. Burns with areas larger than 25 square inches(160 cm2) have the potential to cause serious systemic toxicity.[19]

OrganofluorinesOrganofluorines are naturally rare compounds. They can be nontoxic (perflubron and perfluorodecalin) orhighly toxic (perfluoroisobutylene and fluoroacetic acid). Many pharmacueticals are organofluorines, such asthe anti-cancer fluorouracil. Perfluorooctanesulfonic acid (PFOS) is a persistent organic pollutant.

See alsoFluorocarbonIsotopes of fluorineHalide mineralsWater fluoridation

References1. ^ Theodore Gray. "Real visible fluorine" . The Wooden Periodic Table.2. ^ "pKa's of Inorganic and Oxo-Acids" . Evans Group. Retrieved 2008-11-29.3. ^ M. Jaccaud, R. Faron, D. Devilliers, R. Romano (2005). Fluorine, in Ullmann’s Encyclopedia of Industrial Chemistry.

Weinheim: Wiley-VCH. ISBN 3527310975.4. ^ 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 .

5. ^ "Class I Ozone-Depleting Substances" . Ozone Depletion. U.S. Environmental Protection Agency.6. ^ compiled by Alexander Senning. (2007). Elsevier's dictionary of chemoetymology : the whies and whences of chemical

nomenclature and terminology . Amsterdam: Elsevier. p. 149. ISBN 9780444522399.7. ^ Steve S Lim. "eMedicine - Corticosteroid-Induced Myopathy" .8. ^ "Fluorine's treasure trove" . ICIS news. 2006-10-02. Retrieved 2008-11-29.9. ^ 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 ReductaseInhibition and Antiprotozoal Activities of Diarylamines". ChemBioChem 10 (1): 79. doi:10.1002/cbic.200800565 . PMID19058274 .

10. ^ K. Christe (1986). "Chemical synthesis of elemental fluorine". Inorg. Chem. 25: 3721–3724. doi:10.1021/ic00241a001 .11. ^ "Discovery of fluorine" . Fluoride History.12. ^ H. Moissan (1886). "Action d'un courant électrique sur l'acide fluorhydrique anhydre" . Comptes rendus hebdomadaires

des séances de l'Académie des sciences 102: 1543–1544.13. ^ Richard D. Duncan. (2008). Elements of faith : faith facts and learning lessons from the periodic table . Green Forest,

Ark.: Master Books. p. 22. ISBN 9780890515471.14. ^ "The Nobel Prize in Chemistry 1906" . Nobelprize.org. Retrieved 2009-07-07.15. ^ Proudfoot AT, Bradberry SM, Vale JA (2006). "Sodium fluoroacetate poisoning". Toxicol Rev 25 (4): 213–9.

doi:10.2165/00139709-200625040-00002 . PMID 17288493 .16. ^ Olivares M and Uauy R (2004). "Essential nutrients in drinking-water (Draft)" . WHO. Retrieved 2008-12-30.17. ^ Pizzo G, Piscopo MR, Pizzo I, Giuliana G (September 2007). "Community water fluoridation and caries prevention: a

critical review". Clin Oral Investig 11 (3): 189–93. doi:10.1007/s00784-007-0111-6 . PMID 17333303 .18. ^ Aigueperse, Jean; Paul Mollard, Didier Devilliers, Marius Chemla, Robert Faron, Renée Romano, Jean Pierre Cuer

Fluorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Fluorine[2/19/2010 12:15:30 AM]

This page was last modified on 11 February 2010 at 21:25.Text is available under the Creative Commons Attribution-ShareAlike License; additional terms may apply. See Terms of

Use for details.Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a non-profit organization.

Contact us Privacy policy About Wikipedia Disclaimers

v • d • e

v • d • e

Wikimedia Commons has mediarelated to: Fluorine

Look up fluorine in Wiktionary,the free dictionary.

(2005). "Fluorine Compounds, Inorganic". in Ullmann. Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH.19. ^ "Recommended Medical Treatment for Hydrofluoric Acid Exposure" (PDF). Honeywell Specialty Materials. Retrieved

2009-05-06.

External linksWebElements.com – FluorineIt's Elemental – FluorinePicture of liquid fluorine – chemie-master.deChemsoc.org

Diatomic chemical elements

Hydrogen H2 | Nitrogen N2 | Oxygen O2 | Fluorine F2 | Chlorine Cl2 | Bromine Br2 | Iodine I2 | Astatine At2 |

Periodic table

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo

Uue Ubn

Alkali metalsAlkaline earth

metalsLanthanoids Actinoids

Transition

metalsOther metals Metalloids

Other

nonmetalsHalogens Noble gases

Categories: Chemical elements | Halogens | Fluorine | Fluorinating agents | Biology and pharmacology ofchemical elements

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

navigation

Main pageContentsFeatured contentCurrent eventsRandom article

interaction

About WikipediaCommunity portalRecent changesContact WikipediaDonate to WikipediaHelp

toolbox

What links hereRelated changesUpload fileSpecial pagesPrintable versionPermanent linkCite this page

discussion view source historyTry Beta Log in / create account

sulfur ← chlorine → argon

F↑Cl↓Br

Appearance

pale yellow-green gas

General properties

Name, symbol,number

chlorine, Cl, 17

Element category Halogen

Group, period,block

17, 3, p

Standard atomicweight

35.453(2) g·mol−1

Electronconfiguration

[Ne] 3s2 3p5

Electrons per shell 2, 8, 7 (Image)

Physical properties

Phase gas

Density (0 °C, 101.325 kPa)

3.2 g/L

Melting point 171.6 K, -101.5 °C, -

150.7 °F

Boiling point 239.11 K, -34.04 °C, -

29.27 °F

Critical point 416.9 K, 7.991 MPa

Heat of fusion (Cl2) 6.406 kJ·mol−1

Heat of vaporization (Cl2) 20.41 kJ·mol−1

Specific heatcapacity

(25 °C) (Cl2)

33.949 J·mol−1·K−1

Vapor pressure

ChlorineFrom Wikipedia, the free encyclopedia

This article is about the chemical element. For the bleach, see Sodium hypochlorite.

Chlorine (pronounced /'kl əri n/ KLOR-een, from the Greekword 'χλωρóς' (khlôros, meaning 'pale green'), is the chemicalelement with atomic number 17 and symbol Cl. It is a halogen,found in the periodic table in group 17 (formerly VII, VIIa, orVIIb). As the chloride ion, which is part of common salt andother compounds, it is abundant in nature and necessary tomost forms of life, including humans. In its elemental form (Cl2or "dichlorine") under standard conditions, chlorine is a powerfuloxidant and is used in bleaching and disinfectants, as well as anessential reagent in the chemical industry. As a commondisinfectant, chlorine compounds are used in swimming pools tokeep them clean and sanitary. In the upper atmosphere,chlorine-containing molecules such as chlorofluorocarbons havebeen implicated in the destruction of the ozone layer.

Contents [hide]

1 Characteristics1.1 Isotopes1.2 Occurrence

2 History3 Production

3.1 Electrolysis3.2 Other methods3.3 Industrial production

4 Compounds4.1 Oxidation states

5 Applications and uses5.1 Production of industrial and consumer products5.2 Purification and disinfection5.3 Chemistry5.4 Use as a weapon5.5 Chlorine cracking5.6 Other uses

6 Health effects7 See also8 References9 External links

CharacteristicsAt standard temperature and pressure, two chlorine atoms formthe diatomic molecule Cl2. This is a pale yellow-green gas thathas its distinctive strong smell, the smell of bleach. The bondingbetween the two atoms is relatively weak (only of 242.580±0.004 kJ/mol) which makes the Cl2 molecule highly reactive.

Along with fluorine, bromine, iodine and astatine, chlorine is a

Periodic table

17Cl

article

search

languages

Afrikaans

AsturianuAzərbaycan

БеларускаяBosanskiБългарскиCatalàЧ вашлаČeskyCorsuCymraegDanskDeutsch

EestiΕλληνικάEspañol

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

P/Pa 1 10 100 1 k 10 k 100 k

at T/K 128 139 153 170 197 239

Atomic properties

Oxidation states 7, 6, 5, 4, 3, 2, 1, -1(strongly acidic oxide)

Electronegativity 3.16 (Pauling scale)

Ionization energies(more)

1st: 1251.2 kJ·mol−1

2nd: 2298 kJ·mol−1

3rd: 3822 kJ·mol−1

Covalent radius 102±4 pm

Van der Waalsradius

175 pm

Miscellanea

Crystal structure orthorhombic

Magnetic ordering diamagnetic[1]

Electrical resistivity (20 °C) > 10 Ω·m

Thermalconductivity

(300 K) 8.9x10-

3 W·m−1·K−1

Speed of sound (gas, 0 °C) 206 m/s

CAS registrynumber

7782-50-5

Most stable isotopes

Main article: Isotopes of chlorine

iso NA half-life DM DE (MeV) DP

35Cl 75.77% 35Cl is stable with 18 neutrons

36Cl trace 3.01×105 y β− 0.709 36Ar

ε - 36S

37Cl 24.23% 37Cl is stable with 20 neutrons

This box: view • talk • edit

member of the halogen series that forms the group 17 of theperiodic table—the most reactive group of elements. Itcombines readily with nearly all elements.

Compounds with oxygen, nitrogen, xenon, and krypton areknown, but do not form by direct reaction of the elements.[2]

Chlorine, though very reactive, is not as extremely reactive asfluorine. Pure chlorine gas does, however, support combustionof organic compounds such as hydrocarbons, although thecarbon component tends to burn incompletely, with much of itremaining as soot.[3] At 10 °C and atmospheric pressure, oneliter of water dissolves 3.10 L of gaseous chlorine, and at 30 °C(86 °F), 1 L of water dissolves only 1.77 liters of chlorine.[4]

Chlorine is a member of the salt-forming halogen series and isextracted from chlorides through oxidation often by electrolysis.With metals, it forms salts called chlorides. As the chloride ion,Cl−, it is also the most abundant dissolved ion in ocean water.

IsotopesMain article: Isotopes of chlorine

Chlorine has a wide range of isotopes, the two principal stableisotopes being 35Cl (75.77%) and 37Cl (24.23%); they givechlorine atoms an apparent atomic weight of 35.4527 g/mol.

Trace amounts of radioactive 36Cl exist in the environment, in aratio of about 7x10−13 to 1 with stable isotopes. 36Cl isproduced in the atmosphere by spallation of 36Ar by interactionswith cosmic ray protons. In the subsurface environment, 36Cl isgenerated primarily as a result of neutron capture by 35Cl ormuon capture by 40Ca. 36Cl decays to 36S and to 36Ar, with acombined half-life of 308,000 years. The half-life of thishydrophilic nonreactive isotope makes it suitable for geologicdating in the range of 60,000 to 1 million years. Additionally,large amounts of 36Cl were produced by irradiation of seawaterduring atmospheric detonations of nuclear weapons between1952 and 1958. The residence time of 36Cl in the atmosphere isabout 1 week. Thus, as an event marker of 1950s water in soiland ground water, 36Cl is also useful for dating waters less than50 years before the present. 36Cl has seen use in other areasof the geological sciences, including dating ice and sediments.

OccurrenceSee also: Category:Halide minerals

In nature, chlorine is found primarily as the chloride ion, a component of the salt that is deposited in the earthor dissolved in the oceans — about 1.9% of the mass of seawater is chloride ions. Even higherconcentrations of chloride are found in the Dead Sea and in underground brine deposits. Most chloride saltsare soluble in water, thus, chloride-containing minerals are usually only found in abundance in dry climates ordeep underground. Common chloride minerals include halite (sodium chloride), sylvite (potassium chloride),and carnallite (potassium magnesium chloride hexahydrate). Over 2000 naturally-occurring organic chlorinecompounds are known.[5]

Industrially, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water.Along with chlorine, this chloralkali process yields hydrogen gas and sodium hydroxide, according to the

EsperantoEuskara

FrançaisFurlanGaeilgeGaelgGalegoHak-kâ-fa

Hawai`i

HrvatskiIdoBahasa IndonesiaÍslenskaItalianoעברית

KiswahiliKreyòl ayisyenLatinaLatviešuLëtzebuergeschLietuviųLíguruLingálaLojbanMagyarМакедонски

Māori

Bahasa MelayuNederlands

Norsk (bokmål)Norsk (nynorsk)NovialOccitanO'zbek

PlattdüütschPolskiPortuguêsRomânăRuna SimiРусскийSeelterskShqipSimple EnglishSlovenčinaSlovenščina

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

following chemical equation:

2 NaCl + 2 H2O → Cl2 + H2 + 2 NaOH

HistoryThe most common compound of chlorine, sodium chloride, has been knownsince ancient times; archaeologists have found evidence that rock salt wasused as early as 3000 BC and brine as early as 6000 BC.[6] Hydrochloricacid was probably known to alchemist Abu Musa Jabir ibn Hayyan (Geber)around 800 AD.[7] Before 1400 AD, aqua regia (a mixture of nitric acid andhydrochloric acid) began to be used to dissolve gold,[8] and today this isstill one of the few reagents that will dissolve gold. Upon dissolving gold inaqua regia, chlorine gas is released along with other nauseating andirritating gases, but this wasn't known until much more recently.

Chlorine was first prepared and studied in 1774 by Swedish chemist CarlWilhelm Scheele, and therefore he is credited for its discovery.[9] He calledit "dephlogisticated muriatic acid air" since it was a gas (then called "airs")and it came from hydrochloric acid (then known as "muriatic acid").[9]

However, he failed to establish chlorine as an element, mistakenly thinkingthat it was the oxide obtained from the hydrochloric acid (see phlogistontheory).[9] He named the new element within this oxide as muriaticum.[9]

Regardless of what he thought, Scheele did isolate chlorine by reactingMnO2 (as the mineral pyrolusite) with HCl:

4 HCl + MnO2 → MnCl2 + 2 H2O + Cl2

Scheele observed several of the properties of chlorine: the bleaching effect on litmus, the deadly effect oninsects, the yellow green color, and the smell similar to aqua regia.

A number of chemists, including Claude Berthollet, suggested that Scheele's dephlogisticated muriatic acid airmust be a combination of oxygen and an undiscovered element, muriaticum.[10]

In 1809 Joseph Louis Gay-Lussac and Louis-Jacques Thénard tried to decompose dephlogisticated muriaticacid air by reacting it with charcoal to release the free element muriaticum (and carbon dioxide).[9] They didnot succeed and published a report in which they considered the possibility that dephlogisticated muriatic acidair is an element, but were not convinced.[11]

In 1810, Sir Humphry Davy tried the same experiment again, and concluded that it was an element, and not acompound.[9] He named this new element as chlorine, from the Greek word χλωρος (chlōros), meaninggreen-yellow.[12] The name halogen, meaning salt producer, was originally defined for chlorine (in 1811 byJohann Salomo Christoph Schweigger), and it was later applied to the rest of the elements in this family.[13]

In 1822, Michael Faraday liquefied chlorine for the first time.[14]

Chlorine was first used to bleach textiles in 1785.[15] In 1826, silver chloride was used to producephotographic images for the first time.[16] Chloroform was first used as an anesthetic in 1847.[16] A chlorinesolution in lime-water (hypochlorite) was first used as a germicide to prevent the spread of puerperal fever inthe maternity wards of Vienna General Hospital in Austria in 1847,[17] and in 1850 by John Snow to disinfectthe water supply in London after an outbreak of cholera. The US Department of Treasury called for all waterto be disinfected with chlorine by 1918.[16] Polyvinylchloride (PVC) was invented in 1912, initially without apurpose.[16] Chlorine gas was first introduced as a weapon on April 22, 1915 at Ypres by the GermanArmy,[18][19] and the results of this weapon were disastrous because gas masks had not yet been invented.

ProductionMain article: Chlorine production

Electrolysis

Liquid chlorine

Српски / SrpskiSrpskohrvatski /СрпскохрватскиSuomiSvenska

То икTürkçeУкраїнськаTiếng ViệtWinarayYorùbá

Žemaitėška

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

Chlorine can be manufactured by electrolysis of a sodium chloride solution(brine). The production of chlorine results in the co-products caustic soda(sodium hydroxide, NaOH) and hydrogen gas (H2). These two products, aswell as chlorine itself, are highly reactive. Chlorine can also be produced bythe electrolysis of a solution of potassium chloride, in which case the co-products are hydrogen and caustic potash (potassium hydroxide). There arethree industrial methods for the extraction of chlorine by electrolysis ofchloride solutions, all proceeding according to the following equations:

Cathode: 2 H+ (aq) + 2 e− → H2 (g)Anode: 2 Cl− (aq) → Cl2 (g) + 2 e−

Overall process: 2 NaCl (or KCl) + 2 H2O → Cl2 + H2 + 2 NaOH (or KOH)

Mercury cell electrolysis

Mercury cell electrolysis, also known as the Castner-Kellner process, wasthe first method used at the end of the nineteenth century to producechlorine on an industrial scale.[20][21] The "rocking" cells used have beenimproved over the years.[22] Today, in the "primary cell", titanium anodes(formerly graphite ones) are placed in a sodium (or potassium) chloridesolution flowing over a liquid mercury cathode. When a potential difference is applied and current flows,chlorine is released at the titanium anode and sodium (or potassium) dissolves in the mercury cathodeforming an amalgam. This flows continuously into a separate reactor ("denuder" or "secondary cell"), where itis usually converted back to mercury by reaction with water, producing hydrogen and sodium (or potassium)hydroxide at a commercially useful concentration (50% by weight). The mercury is then recycled to theprimary cell.

The mercury process is the least energy-efficient of the three main technologies (mercury, diaphragm andmembrane) and there are also concerns about mercury emissions.

It is estimated that there are still around 100 mercury-cell plants operating worldwide. In Japan, mercury-based chloralkali production was virtually phased out by 1987 (except for the last two potassium chlorideunits shut down in 2003). In the United States, there will be only five mercury plants remaining in operationby the end of 2008. In Europe, mercury cells accounted for 43% of capacity in 2006 and Western Europeanproducers have committed to closing or converting all remaining chloralkali mercury plants by 2020.[23]

Diaphragm cell electrolysis

In diaphragm cell electrolysis, an asbestos (or polymer-fiber) diaphragm separates a cathode and an anode,preventing the chlorine forming at the anode from re-mixing with the sodium hydroxide and the hydrogenformed at the cathode.[24] This technology was also developed at the end of the nineteenth century. Thereare several variants of this process: the Le Sueur cell (1893), the Hargreaves-Bird cell (1901), the Gibbs cell(1908), and the Townsend cell (1904).[25][26] The cells vary in construction and placement of the diaphragm,with some having the diaphragm in direct contact with the cathode.

The salt solution (brine) is continuously fed to the anode compartment and flows through the diaphragm tothe cathode compartment, where the caustic alkali is produced and the brine is partially depleted.

As a result, diaphragm methods produce alkali that is quite dilute (about 12%) and of lower purity than domercury cell methods. But diaphragm cells are not burdened with the problem of preventing mercurydischarge into the environment. They also operate at a lower voltage, resulting in an energy savings over themercury cell method,[26] but large amounts of steam are required if the caustic has to be evaporated to thecommercial concentration of 50%.

Membrane cell electrolysis

Development of this technology began in the 1970s. The electrolysis cell is divided into two "rooms" by acation permeable membrane acting as an ion exchanger. Saturated sodium (or potassium) chloride solution ispassed through the anode compartment, leaving at a lower concentration.[27] Sodium (or potassium)hydroxide solution is circulated through the cathode compartment, exiting at a higher concentration. A portion

Chlorine gas

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

of the concentrated sodium hydroxide solution leaving the cell is diverted as product, while the remainder isdiluted with deionized water and passed through the electrolysis apparatus again.

This method is more efficient than the diaphragm cell and produces very pure sodium (or potassium)hydroxide at about 32% concentration, but requires very pure brine.

Other electrolytic processes

Although a much lower production scale is involved, electrolytic diaphragm and membrane technologies arealso used industrially to recover chlorine from hydrochloric acid solutions, producing hydrogen (but no causticalkali) as a co-product.

Furthermore, electrolysis of fused chloride salts (Downs process) also enables chlorine to be produced, in thiscase as a by-product of the manufacture of metallic sodium or magnesium.

Other methodsBefore electrolytic methods were used for chlorine production, the direct oxidation of hydrogen chloride withoxygen or air was used in the Deacon process:

4 HCl + O2 → 2 Cl2 + 2 H2O

This reaction is accomplished with the use of copper(II) chloride (CuCl2) as a catalyst and is performed athigh temperature (about 400 °C). The amount of extracted chlorine is approximately 80%. Due to theextremely corrosive reaction mixture, industrial use of this method is difficult and several pilot trials failed inthe past. Nevertheless, recent developments are promising. Recently Sumitomo patented a catalyst for theDeacon process using ruthenium(IV) oxide (RuO2).[28]

Another earlier process to produce chlorine was to heat brine with acid and manganese dioxide.

2 NaCl + 2 H2SO4 + MnO2 → Na2SO4 + MnSO4 + 2 H2O + Cl2

Using this process, chemist Carl Wilhelm Scheele was the first to isolate chlorine in a laboratory. Themanganese can be recovered by the Weldon process.[29]

In the latter half of the 19th century, prior to the adoption of electrolytic methods of chlorine production, therewas substantial production of chlorine by these reactions to meet demand for bleach and bleaching powderfor use by textile industries; by the 1880s the UK, as well as supporting its own (then not inconsiderable)domestic textile production was exporting 70,000 tons per year of bleaching powder.[30] This demand wasmet by capturing hydrochloric acid driven off as a gas during the production of alkali by the Leblanc process,oxidizing this to chlorine (originally by reaction with manganese dioxide), later by direct oxidation by air usingthe Deacon process (in which case impurities capable of poisoning the catalyst had first to be removed), andsubsequently absorbing the chlorine onto lime.

Small amounts of chlorine gas can be made in the laboratory by putting concentrated hydrochloric acid in aflask with a side arm and rubber tubing attached. Manganese dioxide is then added and the flask stoppered.The reaction is not greatly exothermic. As chlorine is denser than air, it can be collected by placing the tubeinside a flask where it will displace the air. Once full, the collecting flask can be stoppered.

Another method for producing small amounts of chlorine gas in a lab is by adding concentrated hydrochloricacid (typically about 5M) to sodium hypochlorite or sodium chlorate solution.

Industrial productionLarge-scale production of chlorine involves several steps and many piecesof equipment. The description below is typical of a membrane plant. Theplant also simultaneously produces sodium hydroxide (caustic soda) andhydrogen gas. A typical plant consists of brine production/treatment, celloperations, chlorine cooling & drying, chlorine compression & liquefaction,liquid chlorine storage & loading, caustic handling, evaporation, storage &loading and hydrogen handling.

Brine

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

Key to the production of chlorine is the operation of the brinesaturation/treatment system. Maintaining a properly saturated solution withthe correct purity is vital, especially for membrane cells. Many plants have asalt pile which is sprayed with recycled brine. Others have slurry tanks thatare fed raw salt.

The raw brine is partially or totally treated with sodium hydroxide, sodiumcarbonate and a flocculant to reduce calcium, magnesium and otherimpurities. The brine proceeds to a large clarifier or a filter where theimpurities are removed. The total brine is additionally filtered before enteringion exchangers to further remove impurities. At several points in thisprocess, the brine is tested for hardness and strength.

After the ion exchangers, the brine is considered pure, and is transferred tostorage tanks to be pumped into the cell room. Brine, fed to the cell line, is heated to the correct temperatureto control exit brine temperatures according to the electrical load. Brine exiting the cell room must be treatedto remove residual chlorine and control pH levels before being returned to the saturation stage. This can beaccomplished via dechlorination towers with acid and sodium bisulfite addition. Failure to remove chlorine canresult in damage to the cells. Brine should be monitored for accumulation of both chlorate anions and sulfateanions, and either have a treatment system in place, or purging of the brine loop to maintain safe levels,since chlorate anions can diffuse through the membranes and contaminate the caustic, while sulfate anionscan damage the anode surface coating.

Cell room

The building that houses many electrolytic cells is usually called a cell room or cell house, although someplants are built outdoors. This building contains support structures for the cells, connections for supplyingelectrical power to the cells and piping for the fluids. Monitoring and control of the temperatures of the feedcaustic and brine is done to control exit temperatures. Also monitored are the voltages of each cell which varywith the electrical load on the cell room that is used to control the rate of production. Monitoring and controlof the pressures in the chlorine and hydrogen headers is also done via pressure control valves.

Direct current is supplied via a rectified power source. Plant load is controlled by varying the current to thecells. As the current is increased, flow rates for brine and caustic and deionized water are increased, whilelowering the feed temperatures.

Cooling and drying

Chlorine gas exiting the cell line must be cooled and dried since the exit gas can be over 80 °C (176 °F) andcontains moisture that allows chlorine gas to be corrosive to iron piping. Cooling the gas allows for a largeamount of moisture from the brine to condense out of the gas stream. This reduces both the coolingrequirements and feed flow of sulfuric acid required in the drying towers. Cooling also improves the efficiencyof both the compression and the liquefaction stage that follows. Chlorine exiting is ideally between 18 °C(64 °F) and 25 °C (77 °F). After cooling the gas stream passes through a series of towers with counterflowing sulfuric acid. The sulfuric acid is fed into the final tower at 98% and the first tower typically has astrength between 66% and 76% depending on materials of construction. These towers progressively removeany remaining moisture from the chlorine gas. After exiting the drying towers the chlorine is filtered to removeany remaining sulfuric acid.

Compression and liquefaction

Liquid Chlorine Analysis

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

Several methods of compression may be used: liquid ring, reciprocating, or centrifugal. The chlorine gas iscompressed at this stage and may be further cooled by inter- and after-coolers. After compression it flows tothe liquefiers, where it is cooled enough to liquefy. Non condensable gases and remaining chlorine gas arevented off as part of the pressure control of the liquefaction systems. These gases are routed to a gasscrubber, producing sodium hypochlorite, or used in the production of hydrochloric acid (by combustion withhydrogen) or ethylene dichloride (by reaction with ethylene).

Storage and loading

Liquid chlorine is typically gravity-fed to storage tanks. It can be loaded into rail or road tankers via pumps orpadded with compressed dry gas.

Caustic handling, evaporation, storage and loading

Caustic, fed to the cell room flows in a loop that is simultaneously bled off to storage with a part diluted withdeionized water and returned to the cell line for strengthening within the cells. The caustic exiting the cell linemust be monitored for strength, to maintain safe concentrations. Too strong or too weak a solution maydamage the membranes. Membrane cells typically produce caustic in the range of 30% to 33% by weight.The feed caustic flow is heated at low electrical loads to control its exit temperature. Higher loads require thecaustic to be cooled, to maintain correct exit temperatures. The caustic exiting to storage is pulled from astorage tank and may be diluted for sale to customers who require weak caustic or for use on site. Anotherstream may be pumped into a multiple effect evaporator set to produce commercial 50% caustic. Rail carsand tanker trucks are loaded at loading stations via pumps.

Hydrogen handling

Hydrogen produced may be vented unprocessed directly to the atmosphere or cooled, compressed and driedfor use in other processes on site or sold to a customer via pipeline, cylinders or trucks. Some possible usesinclude the manufacture of hydrochloric acid or hydrogen peroxide, as well as desulfurization of petroleumoils, or use as a fuel in boilers or fuel cells. In Porsgrunn the byproduct is used for the hydrogen fuelingstation at Hynor.

Energy consumption

Production of chlorine consumes a large amount of energy.[31] Energy consumption per unit weight of productis not far below that for iron and steel manufacture[32] and greater than for the production of glass[33] orcement.[34]

The amount of electrical energy required to produce a given amount of chlorine is fixed by the nature of theelectrochemical reaction. Any energy savings, therefore, can only be made by improving the efficiency of theprocess and reducing ancillary energy use.

CompoundsSee also: Category:Chlorine compounds

For general references to the chloride ion (Cl−), including references to specific chlorides, see chloride. Forother chlorine compounds see chlorate (ClO−

3), chlorite (ClO−2), hypochlorite (ClO−), and perchlorate (ClO−

4),and chloramine (NH2Cl).[35]

Other chlorine-containing compounds include:

Fluorides: chlorine monofluoride (ClF), chlorine trifluoride (ClF3), chlorine pentafluoride (ClF5)Oxides: chlorine dioxide (ClO2), dichlorine monoxide (Cl2O), dichlorine heptoxide (Cl2O7)Acids: hydrochloric acid (HCl), hypochlorous acid (HOCl), chloric acid (HClO3) and perchloric acid.

Oxidation states

Oxidationstate

Name Formula Example compounds

−1 chlorides Cl− ionic chlorides, organic chlorides, hydrochloric acid

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

0 chlorine Cl2 elemental chlorine

+1 hypochlorites ClO− sodium hypochlorite, calcium hypochlorite

+3 chlorites ClO−2 sodium chlorite

+4 chlorine dioxide ClO2

+5 chlorates ClO−3 sodium chlorate, potassium chlorate, chloric acid

+7 perchlorates ClO−4

potassium perchlorate, perchloric acid, magnesium perchlorateorganic perchlorates, ammonium perchlorate, dichlorine heptoxide

Chlorine exists in all odd numbered oxidation states from −1 to +7, as well as the elemental state of zero.Progressing through the states, hydrochloric acid can be oxidized using manganese dioxide, or hydrogenchloride gas oxidized catalytically by air to form elemental chlorine gas. The solubility of chlorine in water isincreased if the water contains dissolved alkali hydroxide. This is due to disproportionation:

Cl2 + 2 OH− → Cl− + ClO− + H2O

In hot concentrated alkali solution disproportionation continues:

2 ClO− → Cl− + ClO−2

ClO− + ClO−2 → Cl− + ClO−

3

Sodium chlorate and potassium chlorate can be crystallized from solutions formed by the above reactions. Iftheir crystals are heated, they undergo the final disproportionation step.

4 ClO−3 → Cl− + 3 ClO−

4

This same progression from chloride to perchlorate can be accomplished by electrolysis. The anode reactionprogression is:[36]

Reaction Electrodepotential

Cl− + 2 OH− → ClO− + H2O + 2 e− +0.89 volts

ClO− + 2 OH− → ClO−2 + H2O + 2 e− +0.67 volts

ClO−2 + 2 OH− → ClO−

3 + H2O + 2 e− +0.33 volts

ClO−3 + 2 OH− → ClO−

4 + H2O + 2 e− +0.35 volts

Each step is accompanied at the cathode by

2 H2O + 2 e− → 2 OH− + H2 (−0.83 volts)

Applications and uses

Production of industrial and consumer productsChlorine's principal applications are in the production of a wide range of industrial and consumerproducts.[37][38] For example, it is used in making plastics, solvents for dry cleaning and metal degreasing,textiles, agrochemicals and pharmaceuticals, insecticides, dyestuffs, household cleaning products, etc.

Purification and disinfectionChlorine is an important chemical for water purification (such as water treatment plants), in disinfectants, andin bleach. Chlorine in water is more than three times more effective as a disinfectant against Escherichia colithan an equivalent concentration of bromine, and is more than six times more effective than an equivalentconcentration of iodine.[39]

Chlorine is usually used (in the form of hypochlorous acid) to kill bacteria and other microbes in drinkingwater supplies and public swimming pools. In most private swimming pools chlorine itself is not used, butrather sodium hypochlorite, formed from chlorine and sodium hydroxide, or solid tablets of chlorinatedisocyanurates. Even small water supplies are now routinely chlorinated.[3] (See also chlorination)

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

It is often impractical to store and use poisonous chlorine gas for water treatment, so alternative methods ofadding chlorine are used. These include hypochlorite solutions, which gradually release chlorine into thewater, and compounds like sodium dichloro-s-triazinetrione (dihydrate or anhydrous), sometimes referred toas "dichlor", and trichloro-s-triazinetrione, sometimes referred to as "trichlor". These compounds are stablewhile solid and may be used in powdered, granular, or tablet form. When added in small amounts to poolwater or industrial water systems, the chlorine atoms hydrolyze from the rest of the molecule forminghypochlorous acid (HOCl) which acts as a general biocide killing germs, micro-organisms, algae, and so on.

ChemistryElemental chlorine is an oxidizer. It undergoes halogen substitution reactions with lower halide salts. Forexample, chlorine gas bubbled through a solution of bromide or iodide anions oxidizes them to bromine andiodine respectively.

Like the other halogens, chlorine participates in free-radical substitution reactions with hydrogen-containingorganic compounds. This reaction is often—but not invariably—non-regioselective, and hence, may result in amixture of isomeric products. It is often difficult to control the degree of substitution as well, so multiplesubstitutions are common. If the different reaction products are easily separated, e.g. by distillation,substitutive free-radical chlorination (in some cases accompanied by concurrent thermal dehydrochlorination)may be a useful synthetic route. Industrial examples of this are the production of methyl chloride, methylenechloride, chloroform and carbon tetrachloride from methane, allyl chloride from propylene, andtrichloroethylene and tetrachloroethylene from 1,2-dichloroethane.

Like the other halides, chlorine undergoes electrophilic additions reactions, most notably, the chlorination ofalkenes and aromatic compounds with a Lewis acid catalyst. Organic chlorine compounds tend to be lessreactive in nucleophilic substitution reactions than the corresponding bromine or iodine derivatives, but theytend to be cheaper. They may be activated for reaction by substituting with a tosylate group, or by the use ofa catalytic amount of sodium iodide.

Chlorine is used extensively in organic and inorganic chemistry as an oxidizing agent and in substitutionreactions because chlorine often imparts many desired properties to an organic compound, due to itselectronegativity.

Chlorine compounds are used as intermediates in the production of a number of important commercialproducts that do not contain chlorine. Examples are: polycarbonates, polyurethanes, silicones,polytetrafluoroethylene, carboxymethyl cellulose and propylene oxide.

Use as a weaponWorld War I

Main article: Poison gas in World War I

Chlorine gas, also known as bertholite, was first used as a weapon in World War I by Germany on April 22,1915 in the Second Battle of Ypres. As described by the soldiers it had a distinctive smell of a mixturebetween pepper and pineapple. It also tasted metallic and stung the back of the throat and chest. Chlorinecan react with water in the mucosa of the lungs to form hydrochloric acid, an irritant which can be lethal. Thedamage done by chlorine gas can be prevented by a gas mask, or other filtration method, which makes theoverall chance of death by chlorine gas much lower than those of other chemical weapons. It was pioneeredby a German scientist later to be a Nobel laureate, Fritz Haber of the Kaiser Wilhelm Institute in Berlin, incollaboration with the German chemical conglomerate IG Farben, who developed methods for dischargingchlorine gas against an entrenched enemy. It is alleged that Haber's role in the use of chlorine as a deadlyweapon drove his wife, Clara Immerwahr, to suicide. After its first use, chlorine was utilized by both sides asa chemical weapon, but it was soon replaced by the more deadly gases phosgene and mustard gas.[40]

Iraq War

Main article: 2007 chlorine bombings in Iraq

Chlorine gas has also been used by insurgents against the local population and coalition forces in the IraqWar in the form of chlorine bombs. On March 17, 2007, for example, three chlorine filled trucks were

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

NFPA 704

detonated in the Anbar province killing two and sickening over 350.[41] Other chlorine bomb attacks resultedin higher death tolls, with more than 30 deaths on two separate occasions.[42] Most of the deaths werecaused by the force of the explosions rather than the effects of chlorine, since the toxic gas is readilydispersed and diluted in the atmosphere by the blast. The Iraqi authorities have tightened up security forchlorine, which is essential for providing safe drinking water for the population.

Chlorine crackingThe element is widely used for purifying water owing to its powerfuloxidizing properties, especially potable water supplies and waterused in swimming pools. Several catastrophic collapses ofswimming pool ceilings have occurred owing to stress corrosioncracking of stainless steel rods used to suspend them.[43] Somepolymers are also sensitive to attack, including acetal resin andpolybutene. Both materials were used in hot and cold waterdomestic supplies, and stress corrosion cracking causedwidespread failures in the USA in the 1980s and 90s. One exampleshows an acetal joint in a water supply system, which when itfractured, caused substantial physical damage to computers in thelabs below the supply. The cracks started at injection moldingdefects in the joint and grew slowly until finally triggered. The fracture surface shows iron and calcium saltswhich were deposited in the leaking joint from the water supply before failure.

Other usesChlorine is used in the manufacture of numerous organic chlorine compounds, the most significant of which interms of production volume are 1,2-dichloroethane and vinyl chloride, intermediates in the production of PVC.Other particularly important organochlorines are methyl chloride, methylene chloride, chloroform, vinylidenechloride, trichloroethylene, perchloroethylene, allyl chloride, epichlorohydrin, chlorobenzene, dichlorobenzenesand trichlorobenzenes.

Chlorine is also used in the production of chlorates and in bromine extraction.

Health effectsChlorine is a toxic gas that irritates the respiratory system. Because it is heavier thanair, it tends to accumulate at the bottom of poorly ventilated spaces. Chlorine gas is astrong oxidizer, which may react with flammable materials.[44]

Chlorine is detectable in concentrations of as low as 0.2 ppm. Coughing and vomiting mayoccur at 30 ppm and lung damage at 60 ppm. About 1000 ppm can be fatal after a few deepbreaths of the gas.[4] Breathing lower concentrations can aggravate the respiratory system,and exposure to the gas can irritate the eyes.[45]

Chlorine's toxicity comes from its oxidizing power. When chlorine is inhaled at concentrations above 30ppm itbegins to react with water and cells which change it into hydrochloric acid (HCl) and hypochlorous acid(HClO).

When used at specified levels for water disinfection, although chlorine reaction with water itself usually doesn'trepresent a major concern for human health, other materials present in the water can generate disinfectionby-products that can damage human health.[46][47]

See alsoPolymer degradation

References1. ^ Magnetic susceptibility of the elements and inorganic compounds , in Handbook of Chemistry and Physics 81st edition,

Chlorine "attack" of an acetal resinplumbing joint.

03 0OX,₩

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

CRC press.2. ^ Windholz, Martha et al., ed (1976). Merck Index of Chemicals and Drugs, 9th ed.. Rahway, N.J.: Merck & Co.. ISBN

0911910263.3. ^ a b Hammond, C. R. (2000). The Elements, in Handbook of Chemistry and Physics 81st edition. CRC press. ISBN

0849304814.4. ^ a b "WebElements.com – Chlorine" . Mark Winter [The University of Sheffield and WebElements Ltd, UK]. Retrieved

2007-03-17.5. ^ "Risk assessment and the cycling of natural organochlorines" . Euro Chlor. Retrieved 2007-08-12.6. ^ "The earliest salt production in the world: an early Neolithic exploitation in Poiana Slatinei-Lunca, Romania" . Retrieved

2008-07-10.7. ^ Pereira, Jonathan (1854). The elements of materia medica and therapeutics, Volume 1 . Longman, Brown, Green, and

Longmans. p. 387.8. ^ Hoover, Herbert Clark (2003). Georgius Agricola de Re Metallica . Kessinger Publishing. p. 354. ISBN 0766131971.9. ^ a b c d e f "17 Chlorine" . Elements.vanderkrogt.net. Retrieved 2008-09-12.

10. ^ Ihde, Aaron John (1984). The development of modern chemistry . Courier Dover Publications. p. 158. ISBN0486642356.

11. ^ Gay-Lussac, Joseph Louis; Thénard, Louis-Jacques (1809). "On the nature and the properties of muriatic acid and ofoxygenated muriatic acid" . Mémoires de Physique et de Chimie de la Société d'Arcueil 2: 339–358.

12. ^ Sir Humphry Davy (1811). "On a Combination of Oxymuriatic Gas and Oxygene Gas" . Philosophical Transactions ofthe Royal Society 101: 155–162. doi:10.1098/rstl.1811.0008 .

13. ^ Snelders, H. A. M. (1971). "J. S. C. Schweigger: His Romanticism and His Crystal Electrical Theory of Matter" . Isis 62(3): 328.

14. ^ "Discovery of Chlorine" . Retrieved 2008-07-10.15. ^ "History of Chlorine" . Retrieved 2008-07-10.16. ^ a b c d Jacqueline Brazin. "Chlorine & its Consequences" . Retrieved 2008-07-10.17. ^ "Chlorine Story" . americanchemistry. Retrieved 2008-07-10.18. ^ "Chlorine - History" . Retrieved 2008-07-10.19. ^ "Weaponry: Use of Chlorine Gas Cylinders in World War I" . historynet.com. Retrieved 2008-07-10.20. ^ Pauling, Linus (1970). General Chemistry. Dover publications. ISBN 0-486-65622-5.21. ^ "Electrolytic Processes for Chlorine and Caustic Soda" . Lenntech Water treatment & air purification Holding B.V.,

Rotterdamseweg 402 M, 2629 HH Delft, The Netherlands. Retrieved 2007-03-17.22. ^ "Mercury cell" . Euro Chlor. Retrieved 2007-08-15.23. ^ "Regional Awareness-raising Workshop on Mercury Pollution" . UNEP. Retrieved 2007-10-28.24. ^ "Diaphragm cell" . Euro Chlor. Retrieved 2007-08-15.25. ^ "The Electrolysis of Brine" . Salt Manufacturers' Association. Retrieved 2007-03-17.26. ^ a b Kiefer, David M.. "When the Industry Charged Ahead" . Chemistry Chronicles. Retrieved 2007-03-17.27. ^ "Membrane cell" . Euro Chlor. Retrieved 2007-08-15.28. ^ Lopez, N; Gomezsegura, J; Marin, R; Perezramirez, J (2008). "Mechanism of HCl oxidation (Deacon process) over

RuO2". Journal of Catalysis 255: 29. doi:10.1016/j.jcat.2008.01.020 .29. ^ "The Chlorine Industry" . Lenntech Water treatment & air purification Holding B.V., Rotterdamseweg 402 M, 2629 HH

Delft, The Netherlands. Retrieved 2007-03-17.30. ^ Reader W J (1970 SBN 19 215937 2). Imperial Chemical Industries; A History. Volume 1. The Forerunners 1870-1926.

Oxford University Press. p. 102. citing Haber L F (1958). The Chemical Industry during the Nineteenth Century. Oxford:Clarendon Press.

31. ^ "Integrated Pollution Prevention and Control (IPPC) - Reference Document on Best Available Techniques in the Chlor-Alkali Manufacturing Industry" . European Commission. Retrieved 2007-09-02.

32. ^ "Integrated Pollution Prevention and Control (IPPC) - Best Available Techniques Reference Document on the Productionof Iron and Steel" . European Commission. Retrieved 2007-09-02.

33. ^ "Integrated Pollution Prevention and Control (IPPC) - Reference Document on Best Available Techniques in the GlassManufacturing Industry" . European Commission. Retrieved 2007-09-02.

34. ^ "Integrated Pollution Prevention and Control (IPPC) - Reference Document on Best Available Techniques in the Cementand Lime Manufacturing Industries" . European Commission. Retrieved 2007-09-02.

35. ^ "Chlorine compounds of the month" . Euro Chlor. Retrieved 2007-08-29.36. ^ Cotton, F. Albert and Wilkinson, Geoffrey (1966). Advanced Inorganic Chemistry, 2nd ed.. John Wiley & sons. p. 568.37. ^ "Uses" . Euro Chlor. Retrieved 2007-08-20.38. ^ "Chlorine Tree" . Chlorine Tree. Retrieved 2007-08-20.39. ^ Koski TA, Stuart LS, Ortenzio LF (1966). "Comparison of chlorine, bromine, iodine as disinfectants for swimming pool

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

v • d • e

v • d • e

v • d • e

Wikimedia Commons has mediarelated to: Chlorine

Look up chlorine in Wiktionary,the free dictionary.

water" . Applied Microbiology 14 (2): 276–279. PMID 4959984 . PMC 546668 .40. ^ "Weapons of War: Poison Gas" . First World War.com. Retrieved 2007-08-12.41. ^ Mahdi, Basim (2007-03-17). "Iraq gas attack makes hundreds ill" . CNN. Retrieved 2007-03-17.42. ^ "'Chlorine bomb' hits Iraq village" . BBC News. 2007-05-17. Retrieved 2007-05-17.43. ^ Bertolini, Luca; Elsener, Bernhard; Pedeferri, Pietro; Polder, Rob B. (2004). Corrosion of steel in concrete: prevention,

diagnosis, repair . Wiley-VCH. p. 148. ISBN 3527308008.44. ^ "Chlorine MSDS" . October 23, 1997 (Revised November 1999.45. ^ Winder, Chris (2001). "The Toxicology of Chlorine". Environmental Research 85 (2): 105–114.

doi:10.1006/enrs.2000.4110 . PMID 11161660 .46. ^ "What's in your Water?: Disinfectants Create Toxic By-products" . ACES News. College of Agricultural, Consumer and

Environmental Sciences - University of Illinois at Urbana-Champaign. 2009-03-31. Retrieved 2009-03-31.47. ^ Richardson, D.; Plewa, J.; Wagner, D.; Schoeny, R.; Demarini, M. (Nov 2007). "Occurrence, genotoxicity, and

carcinogenicity of regulated and emerging disinfection by-products in drinking water: a review and roadmap for research".Mutation research 636 (1-3): 178–242. doi:10.1016/j.mrrev.2007.09.001 . ISSN 0027-5107 . PMID 17980649 . edit

External linksElectrolytic productionComputational Chemistry WikiChlorine Production Using Mercury, EnvironmentalConsiderations and AlternativesNational Pollutant Inventory - ChlorineNational Institute for Occupational Safety and Health - ChlorinePageWebElements.com — ChlorineChlorine Institute - Trade association and lobby group representing the interests of the chlorine industryChlorine Online - Chlorine Online is an information resource produced by Eurochlor - the businessassociation of the European chlor-alkali industry

Chlorine compounds

ClF · ClF3 · ClF5 · ClFO3 · ClNO3 · ClO · ClO2 · ClO2F · Cl2O · Cl2O4 · Cl2O6 · Cl2O7

Diatomic chemical elements

Hydrogen H2 | Nitrogen N2 | Oxygen O2 | Fluorine F2 | Chlorine Cl2 | Bromine Br2 | Iodine I2 | Astatine At2 |

United States chemical weapons program

Agents and chemicals3-Quinuclidinyl benzilate (BZ) · Chlorine · Methylphosphonyl difluoride (DF) · Phosgene · QL ·Sarin (GB) · Sulfur mustard (HD) · VX

Weapons

Bigeye bomb · M1 chemical mine · M104 155mm Cartridge · M110 155mm Cartridge ·M121/A1 155mm Cartridge · M125 bomblet · M134 bomblet · M138 bomblet · M139 bomblet ·M2 mortar · M23 chemical mine · M34 cluster bomb · M360 105mm Cartridge ·M426 8-inch shell · M43 BZ cluster bomb · M44 generator cluster · M55 rocket ·M60 105mm Cartridge · M687 155mm Cartridge · XM-736 8-inch projectile · MC-1 bomb ·M47 bomb · Weteye bomb

Operations and testingDugway sheep incident · Edgewood Arsenal experiments · MKULTRA · Operation CHASE ·Operation Geranium · Operation LAC · Operation Red Hat · Operation Steel Box ·Operation Ranch Hand · Operation Top Hat · Project 112 · Project SHAD

Facilities

Anniston Army Depot · Anniston Chemical Activity · Blue Grass Army Depot ·Deseret Chemical Depot · Edgewood Chemical Activity · Hawthorne Army Depot ·Johnston Atoll Chemical Agent Disposal System · Newport Chemical Depot ·Pine Bluff Chemical Activity · Pueblo Chemical Depot · Tooele Chemical Agent Disposal Facility ·Umatilla Chemical Depot

Units and formations 1st Gas Regiment · U.S. Army Chemical Corps · Chemical mortar battalion

Equipment Chemical Agent Identification Set · M93 Fox · MOPP · People sniffer

Chlorine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Chlorine[2/19/2010 12:25:18 AM]

This page was last modified on 16 February 2010 at 12:18.Text is available under the Creative Commons Attribution-ShareAlike License; additional terms may apply. See Terms of

Use for details.Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a non-profit organization.

Contact us Privacy policy About Wikipedia Disclaimers

v • d • e

v • d • e

v • d • e

Related topicsAl-Shifa pharmaceutical factory · Chlorine bombings in Iraq · Herbicidal warfare · List of topics ·Poison gas in World War I · Tyler poison gas plot

Chemical agents

Blood Cyanogen chloride (CK) · Hydrogen cyanide (AC)

BlisterEthyldichloroarsine (ED) · Methyldichloroarsine (MD) · Phenyldichloroarsine (PD) · Lewisite (L) ·Sulfur mustard gas (HD · H · HT · HL · HQ) · Nitrogen mustard gas (HN1 · HN2 · HN3)

Nerve

G-agents Tabun (GA) · Sarin (GB) · Soman (GD) · Cyclosarin (GF) · GV

V-agents EA-3148 · VE · VG · VM · VR · VX

Novichok agents

Pulmonary Chlorine · Chloropicrin (PS) · Phosgene (CG) · Diphosgene (DP)

Incapacitating Agent 15 (BZ) · EA-3167 · Kolokol-1

Riot control Pepper spray (OC) · CS gas · CN gas (mace) · CR gas

E numbers

Colours (E100–199) • Preservatives (E200–299) • Antioxidants & Acidity regulators (E300–399) • Thickeners, stabilisers &emulsifiers (E400–499) • pH regulators & anti-caking agents (E500–599) • Flavour enhancers (E600–699) • Miscellaneous

(E900–999) • Additional chemicals (E1100–1599)

Waxes (E900–909) • Synthetic glazes (E910–919) • Improving agents (E920–929) • Packaging gases (E930–949) •Sweeteners (E950–969) • Foaming agents (E990–999)

L-cysteine (E920) • L-cystine (E921) • Potassium persulfate (E922) • Ammonium persulfate (E923) • Potassium bromate(E924) • Chlorine (E925) • Chlorine dioxide (E926) • Azodicarbonamide (E927) • Carbamide (E927b) • Benzoyl peroxide (E928)

Periodic table

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo

Uue Ubn

Alkali metalsAlkaline earth

metalsLanthanoids Actinoids

Transition

metalsOther metals Metalloids

Other

nonmetalsHalogens Noble gases

Categories: Chemical elements | Chlorine | Halogens | Hazardous air pollutants | Highly HazardousChemicals | Occupational safety and health | Pulmonary agents | Swimming pool equipment

Bromine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Bromine[2/19/2010 12:32:28 AM]

navigation

Main pageContentsFeatured contentCurrent eventsRandom article

interaction

About WikipediaCommunity portalRecent changesContact WikipediaDonate to WikipediaHelp

toolbox

What links hereRelated changesUpload fileSpecial pagesPrintable versionPermanent linkCite this page

discussion edit this page historyTry Beta Log in / create account

[edit]

selenium ← Bromine → krypton

Cl↑

Br↓I

Appearance

gas/liquid: red-brown

solid: metallic luster

General properties

Name, symbol,number

Bromine, Br, 35

Element category halogen

Group, period, block 17, 4, p

Standard atomicweight

79.904(1) g·mol−1

Electronconfiguration

[Ar] 4s2 3d10 4p5

Electrons per shell 2, 8, 18, 7 (Image)

Physical properties

Phase liquid

Density (near r.t.) (Br2, liquid) 3.1028

g·cm−3

Melting point 265.8 K, -7.2 °C, 19 °F

Boiling point 332.0 K, 58.8 °C,

137.8 °F

Critical point 588 K, 10.34 MPa

Heat of fusion (Br2) 10.571 kJ·mol−1

Heat of vaporization (Br2) 29.96 kJ·mol−1

Specific heat capacity (25 °C) (Br2)

75.69 J·mol−1·K−1

Vapor pressure

BromineFrom Wikipedia, the free encyclopedia

"Bromo" redirects here. For other uses, see Bromo (disambiguation).

Bromine (pronounced /'bro mi n/ BROH-meen or /'bro m n/BROH-min, from Greek: βρῶμος, brómos, meaning "stench (ofhe-goats)"),[2] is a chemical element with the symbol Br andatomic number 35. A halogen element, bromine is a reddish-brown volatile liquid at standard room temperature that isintermediate in reactivity between chlorine and iodine. Brominevapors are corrosive and toxic. Approximately 556,000 metrictons were produced in 2007.[3] The main applications forbromine are in fire retardants and fine chemicals.

Contents [hide]

1 History2 Characteristics

2.1 Isotopes2.2 Allotropes

3 Occurrence and production4 Compounds

4.1 Organic chemistry4.2 Inorganic chemistry

5 Applications5.1 Flame retardant5.2 Gasoline additive5.3 Pesticide5.4 Medical and veterinary5.5 Other uses

6 Biological role7 Safety8 References9 External links

HistoryBromine wasdiscoveredindependently by twochemists AntoineBalard[4] and CarlJacob Löwig[5] in1825 and 1826.[6]

Balard found bromidechemicals in the ashof sea weed from thesalt marshes ofMontpellier in 1826.The sea weed wasused to produce

Illustrative and secure bromine samplefor teaching

Periodic table

35Br

article

search

languages

Afrikaans

Azərbaycan

БеларускаяBosanskiБългарскиCatalàЧ вашлаČeskyCorsuCymraegDanskDeutschEestiΕλληνικάEspañolEsperantoEuskara

Bromine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Bromine[2/19/2010 12:32:28 AM]

[edit]

P/Pa 1 10 100 1 k 10 k 100 k

at T/K 185 201 220 244 276 332

Atomic properties

Oxidation states 7, 5, 4, 3, 1, -1(strongly acidic oxide)

Electronegativity 2.96 (Pauling scale)

Ionization energies 1st: 1139.9 kJ·mol−1

2nd: 2103 kJ·mol−1

3rd: 3470 kJ·mol−1

Atomic radius 120 pm

Covalent radius 120±3 pm

Van der Waals radius 185 pm

Miscellanea

Crystal structure orthorhombic

Magnetic ordering diamagnetic[1]

Electrical resistivity (20 °C) 7.8×1010Ω·m

Thermal conductivity (300 K) 0.122 W·m−1·K−1

Speed of sound (20°C) 206 m/s

CAS registry number 7726-95-6

Most stable isotopes

Main article: Isotopes of Bromine

iso NA half-life DM DE (MeV) DP

79Br 50.69% 79Br is stable with 44 neutrons

81Br 49.31% 81Br is stable with 46 neutrons

This box: view • talk • edit

iodine, but alsocontained bromine. Balard distilled the bromine from a solutionof seaweed ash saturated with chlorine. The properties of theresulting substance resembled that of an intermediate ofchlorine and iodine; with those results he tried to prove that thesubstance was iodine monochloride (ICl), but after failing to doso he was sure that he had found a new element and named itmuride, derived from the Latin word muria for brine.[4]

Carl Jacob Löwig isolated bromine from a mineral water springfrom his hometown Bad Kreuznach in 1825. Löwig used asolution of the mineral salt saturated with chlorine and extractedthe bromine with diethylether. After evaporation of the ether abrown liquid remained. With this liquid as a sample for his workhe applied for a position in the laboratory of Leopold Gmelin inHeidelberg. The publication of the results was delayed andBalard published his results first.[5]

After the French chemists Louis Nicolas Vauquelin, LouisJacques Thénard, and Joseph-Louis Gay-Lussac approved theexperiments of the young pharmacist Balard, the results werepresented at a lecture of the Académie des Sciences andpublished in Annales de Chimie et Physique.[7] In hispublication Balard states that he changed the name from murideto brôme on the proposal of M. Anglada. Other sources claimthat the French chemist and physicist Joseph-Louis Gay-Lussacsuggested the name brôme for the characteristic smell of thevapors.[8] Bromine was not produced in large quantities until1860.

The first commercial use, besides some minor medicalapplications, was the use of bromine for the daguerreotype. In1840 it was discovered that bromine had some advantages overthe previously used iodine vapor to create the light sensitivesilver halide layer used for daguerreotypy.[9]

Potassium bromide and sodium bromide were used as anticonvulsants and sedatives in the late 19th andearly 20th centuries, until they were gradually superseded by chloral hydrate and then the barbiturates.[10]

CharacteristicsBromine is the only liquid nonmetallic element at room temperature, and one of only two elements on theperiodic table that are liquid at room temperature. The melting point of bromine is −7.2 °C and the boilingpoint 58.8 °C (138 °F). The pure chemical element has the physical form of a diatomic molecule, Br2. It is adense, mobile, reddish-brown liquid, that evaporates easily at standard temperature and pressures to give ared vapor (its color resembles nitrogen dioxide) that has a strong disagreeable odor resembling that ofchlorine. Bromine is a halogen, and is less reactive than chlorine and more reactive than iodine. Bromine isslightly soluble in water, and highly soluble in carbon disulfide, aliphatic alcohols (such as methanol), andacetic acid. It bonds easily with many elements and has a strong bleaching action. Bromine, like chlorine, isalso used in maintenance of swimming pools.

Certain bromine-related compounds have been evaluated to have an ozone depletion potential orbioaccumulate in living organisms. As a result many industrial bromine compounds are no longermanufactured, are being restricted, or scheduled for phasing out. The Montreal Protocol mentions severalorganobromine compounds for this phase out.

Bromine is a powerful oxidizing agent. It reacts vigorously with metals, especially in the presence of water, as

FrançaisFurlanGaeilgeGaelgGalegoHak-kâ-fa

HrvatskiIdoBahasa IndonesiaÍslenskaItalianoעבריתBasa Jawa

KiswahiliKreyòl ayisyenLatinaLatviešuLëtzebuergeschLietuviųLíguruLojbanMagyarМакедонски

Māori

Bahasa MelayuNederlands

Norsk (bokmål)Norsk (nynorsk)OccitanO'zbek

PolskiPortuguêsRomânăRuna SimiРусскийSeelterskSimple EnglishSlovenčinaSlovenščinaСрпски / SrpskiSrpskohrvatski /СрпскохрватскиSuomiSvenska

То ик

Bromine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Bromine[2/19/2010 12:32:28 AM]

[edit]

[edit]

[edit]

well as most organic compounds, especially upon illumination.

IsotopesMain article: Isotopes of bromine

Bromine has 2 stable isotopes: 79Br (50.69 %) and 81Br (49.31%). At least another 23 radioisotopes areknown to exist.[11] Many of the bromine isotopes are fission products. Several of the heavier bromineisotopes from fission are delayed neutron emitters. All of the radioactive bromine isotopes are relatively shortlived. The longest half life is the neutron deficient 77Br at 2.376 days. The longest half life on the neutron richside is 82Br at 1.471 days. A number of the bromine isotopes exhibit metastable isomers. Stable 79Br exhibitsa radioactive isomer, with a half life of 4.86 seconds. It decays by isomeric transition to the stable groundstate.[12]

AllotropesAt a pressure of 55 GPa bromine converts to a metal. At 75 GPa it converts to a face centered orthorhombicstructure. At 100 GPa it converts to a body centered orthorhombic monoatomic form.[13]

Occurrence and productionSee also: Category:Halide minerals and Halide mineral

The diatomicelement Br2 doesnot occur naturally.Instead, bromineexists exclusively asbromide salts indiffuse amounts incrustal rock. Due toleaching, bromidesalts haveaccumulated in sea

water (65 ppm),[14] but at a lower concentration than chloride.Bromine may be economically recovered from bromide-richbrine wells and from the Dead Sea waters (up to 50000ppm).[15][16]

Approximately 556,000 metric tonnes (worth around US$2.5billion) of bromine are produced per year (2007) worldwidewith the United States, Israel, and China being the primaryproducers.[17][18][19] Bromine production has increased sixfold since the 1960s. The largest bromine reservein the United States is located in Columbia and Union County, Arkansas, U.S.[20] China's bromine reservesare located in the Shandong Province and Israel's bromine reserves are contained in the waters of the DeadSea. The bromide-rich brines are treated with chlorine gas, flushing through with air. In this treatment,bromide anions are oxidized to bromine by the chlorine gas.

2 Br− + Cl2 → 2 Cl− + Br2

Because of its commercial availability and long shelf-life, bromine is not typically prepared. Small amounts ofbromine can however be generated through the reaction of solid sodium bromide with concentrated sulfuricacid (H2SO4). The first stage is formation of hydrogen bromide (HBr), which is a gas, but under the reactionconditions some of the HBr is oxidized further by the sulfuric acid to form bromine (Br2) and sulfur dioxide(SO2).

NaBr (s) + H2SO4 (aq) → HBr (aq) + NaHSO4 (aq)2 HBr (aq) + H2SO4 (aq) → Br2 (g) + SO2 (g) + 2 H2O (l)

World bromine production trend

View of salt evaporation pans on the DeadSea, where Jordan (right) and Israel (left) producesalt and bromine

TürkçeУкраїнськаUyghurche / Tiếng ViệtWinarayYorùbá

Bromine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Bromine[2/19/2010 12:32:28 AM]

[edit]

[edit]

[edit]

Oxidation statesof bromine

-1 HBr

0 Br2+1 BrCl

+3 BrF3

+5 BrF5

+5 BrO−3

+7 BrO−4

Similar alternatives, such as the use of dilute hydrochloric acid with sodium hypochlorite, are also available.The most important thing is that the anion of the acid (in the above examples, sulfate and chloride,respectively) be more electronegative than bromine, allowing the substitution reaction to occur.

Reaction involving a strong oxidizing agent, such as potassium permanganate, on bromide ions in thepresence of an acid also gives bromine. An acidic solution of bromate ions and bromide ions will alsodisproportionate slowly to give bromine.

Bromine is only slightly soluble in water. But the solubility can be increased by the presence of bromide ions.However, concentrated solutions of bromine are rarely prepared in the lab as they will continually give offtoxic red-brown bromine gas due to its very high vapor pressure. Sodium thiosulphate is an excellent reagentfor getting rid of bromine completely including the stains and odor.

CompoundsSee also: Category:Bromine compounds

Organic chemistryOrganic compounds are brominated by either addition or substitution reactions. Bromineundergoes electrophilic addition to the double-bonds of alkenes, via a cyclic bromoniumintermediate. In non-aqueous solvents such as carbon disulfide, this affords the di-bromo product. For example, reaction with ethylene will produce 1,2-dibromoethane.Bromine also undergoes electrophilic addition to phenols and anilines. When used asbromine water, a small amount of the corresponding bromohydrin is formed as well asthe dibromo compound. So reliable is the reactivity of bromine that bromine water isemployed as a reagent to test for the presence of alkenes, phenols, and anilines. Likethe other halogens, bromine participates in free radical reactions. For example hydrocarbons are brominatedupon treatment with bromine in the presence of light.

Bromine, sometimes with a catalytic amount of phosphorus, easily brominates carboxylic acids at the α-position. This method, the Hell-Volhard-Zelinsky reaction, is the basis of the commercial route to bromoaceticacid. N-Bromosuccinimide is commonly used as a substitute for elemental bromine, being easier to handle,and reacting more mildly and thus more selectively. Organic bromides are often preferable relative to the lessreactive chlorides and more expensive iodide-containing reagents. Thus, Grignard and organolithiumcompound are most often generated from the corresponding bromides.

Inorganic chemistryBromine is an oxidizer, and it will oxidize iodide ions to iodine, being itself reduced tobromide:

Br2 + 2 I− → 2 Br− + I2

Bromine will also oxidize metals and metalloids to the corresponding bromides.Anhydrous bromine is less reactive toward many metals than hydrated bromine,however. Dry bromine reacts vigorously with aluminium, titanium, mercury as well asalkaline earths and alkali metals.

If bromine is dissolved in hydroxide containing water not only bromide (Br−) is formed,but also the hypobromite (OBr−). This hypobromite is responsible for the bleachingabilities of bromide solutions. In warm solutions the disproportion reaction of thehypobromite is quantitative. The resulting bromate is a strong oxidising agent and verysimilar to the chlorate.

3 BrO− → BrO−3 + 2 Br−

The perbromates are not accessible through electrolysis like the perchlorates, but only by reacting bromatesolutions with fluorine or ozone.

BrO3− + H2O + F2 → BrO−

4 + 2 HF− −

N-Bromosuccinimide

Bromine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Bromine[2/19/2010 12:32:28 AM]

[edit]

[edit]

[edit]

[edit]

[edit]

[edit]

BrO3 + O3 → BrO4 + O2

ApplicationsA wide variety of organobromine compounds are used in industry. Some are prepared from bromine andothers are prepared from hydrogen bromide, which is obtained by burning hydrogen in bromine.[3]

Illustrative of the addition reaction[21] is the preparation of 1,2-Dibromoethane, the organobromine compoundproduced in the largest amounts:

C2H4 + Br2 → CH2BrCH2Br

Flame retardantBrominated flame retardants represent a commodity of growingimportance. If the material burns the flame retardants producehydrobromic acid which interferes in the radical chain reaction ofthe oxidation reaction of the fire. The highly reactive hydrogenoxygen and hydroxy radicals react with hydrobromic acid and formless reactive bromine radicals.[22][23] The bromine containingcompounds can be placed in the polymers either duringpolymerization if a small amount of brominated monomer is addedor the bromine containing compound is added after polymerization.Tetrabromobisphenol A can be added to produce polyesters or epoxy resins. Epoxy used in printed circuitboards (PCB) are normally made from flame retardant resins, indicated by the FR in the abbreviation of theproducts (FR-4 and FR-2. Vinyl bromide can be used in the production of polyethylene, polyvinylchloride orpolypropylene. Decabromodiphenyl ether can be added to the final polymers.[24]

Gasoline additiveEthylene bromide was an additive in gasolines containing lead anti-engine knocking agents. It scavenges leadby forming volatile lead bromide, which is exhausted from the engine. This application accounted for 77% ofthe bromine uses in 1966 in the US. This application has declined since the 1970s due to environmentalregulations.[25] Ethylene bromide is also used as a fumigant, but again this application is declining.[19]

PesticideMethyl bromide was widely used as pesticide to fumigate soil. The Montreal Protocol onSubstances that Deplete the Ozone scheduled the phase out for the ozone depletingchemical until 2005. In 1991, an estimated 35,000 metric tonnes of the chemical wereused to control nematodes, fungi, weeds and other soil-borne diseases.[26][27]

Medical and veterinaryBromide compounds, especially potassium bromide, were frequently used assedatives in the 19th and early 20th century. Bromides in the form of simple salts are still used asanticonvulsants in both veterinary and human medicine.

Other usesThe bromides of calcium, sodium, and zinc account for a sizable part of the brominemarket. These salts form dense solutions in water that are used as drilling fluidssometimes called clear brine fluids.[19][28]

Bromine is also used in the production of brominated vegetable oil, which is used as anemulsifier in many citrus-flavored soft drinks (e.g. Mountain Dew). After the introductionin the 1940s the compound was extensively used until the UK and the US limited itsuse in the mid 1970s and alternative emulsifiers were developed.[29]

Soft drinks containing brominated vegetable oil are still sold in the US (2009).[30]

Several dyes, agrichemicals, and pharmaceuticals are

Tetrabromobisphenol A

Methyl bromide(bromomethane)

Orangefluoresces of DNAEthidium bromideintercalate

Bromine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Bromine[2/19/2010 12:32:28 AM]

[edit]

[edit]

[edit]

organobromine compounds. 1-Bromo-3-chloropropane, 1-bromoethylbenzene, and 1-bromoalkanes are prepared by the antimarkovnikov addition of HBr to alkenes.Ethidium bromide, EtBr, is used as a DNA stain in gel electrophoresis.High refractive index compoundsWater purification compounds, disinfectants and insecticides, such as tralomethrin(C22H19Br4NO3).[19]

Potassium bromide is used in some photographic developers to inhibit the formation of fog (undesiredreduction of silver).Vapor is used as the second step in sensitizing daguerreotype plates to be developed under Mercury (Hg)vapor. Bromine acts as an accelerator to the light sensitivity of the previously iodized plate.Bromine is also used to reduce mercury pollution from coal-fired power plants. This can be achieved eitherby treating activated carbon with bromine or by injecting bromine compounds onto the coal prior tocombustion.

Biological roleBromine has no known essential role in human or mammalian health, but inorganicbromine and organobromine compounds do occur naturally, and some may be of use tohigher organisms in dealing with parasites. For example, in the presence of H2O2formed by the eosinophil, and either chloride or bromide ions, eosinophil peroxidaseprovides a potent mechanism by which eosinophils kill multicellular parasites (such as,for example, the nematode worms involved in filariasis); and also certain bacteria (suchas tuberculosis bacteria). Eosinophil peroxidase is a haloperoxidase that preferentiallyuses bromide over chloride for this purpose, generating hypobromite (hypobromousacid).[31]

Marine organisms are the main source of organobromine compounds. Over 1600 compounds were identifiedby 1999. The most abundant one is methyl bromide with estimated 56,000 metric tonnes produced by marinealgae.[32] The essential oil of the Hawaiian alga Asparagopsis taxiformis consists of 80% methyl bromide.[33]

A famous example of a bromine-containing organic compound that has been used by humans for a long timeis Tyrian purple.[32][34] The brominated indigo is produced by a medium-sized predatory sea snail, themarine gastropod Murex brandaris. It took until 1909 before the organobromine nature of the compound wasdiscovered by Paul Friedländer.[35] Most organobromine compounds in nature arise via the action ofvanadium bromoperoxidase.[36]

SafetySee also: List of highly toxic gases

Elemental bromine is toxic and causes burns. As an oxidizing agent, it is incompatible with most organic andinorganic compounds. Care needs to be taken when transporting bromine; it is commonly carried in steeltanks lined with lead, supported by strong metal frames.

When certain ionic compounds containing bromine are mixed with potassium permanganate (KMnO4) and anacidic substance, they will form a pale brown cloud of bromine gas. This gas smells like bleach and is veryirritating to the mucus membranes. Upon exposure, one should move to fresh air immediately. If symptoms ofbromine poisoning arise, medical attention is needed.

References1. ^ Magnetic susceptibility of the elements and inorganic compounds , in Handbook of Chemistry and Physics 81st edition,

CRC press.2. ^ Gemoll W, Vretska K (1997). Griechisch-Deutsches Schul- und Handwörterbuch ("Greek-German dictionary"), 9th ed..

öbvhpt. ISBN 3-209-00108-1.3. ^ a b Jack F. Mills (2002). Bromine: in Ullmann's Encyclopedia of Chemical Technology. Weinheim: Wiley-VCH Verlag.

Tralomethrin

Tyrian purple

Bromine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Bromine[2/19/2010 12:32:28 AM]

doi:10.1002/14356007.a04_391 .4. ^ a b Balard, Antoine (1826). "Memoire of a peculire Substance contained in Sea Water" . Annals of Philosophy: 387–

and 411–.5. ^ a b Landolt, Hans Heinrich (1890). "Nekrolog: Carl Löwig" . Berichte der deutschen chemischen Gesellschaft 23: 905.

doi:10.1002/cber.18900230395 .6. ^ Weeks, Mary Elvira (1932). "The discovery of the elements: XVII. The halogen family.". Journal of Chemical Education 9:

1915.7. ^ Balard, A.J. (1826). Annales de Chimie et Physique 32: 337.8. ^ Wisniak, Jaime (2004). "Antoine-Jerôme Balard. The discoverer of bromine" . Revista CENIC Ciencias Químicas 35.9. ^ Barger, M. Susan; White, William Blaine (2000). "Technological Practice of Daguerreotypy". The Daguerreotype:

Nineteenth-century Technology and Modern Science. JHU Press. pp. 31–35. ISBN 9780801864582.10. ^ Shorter, Edward (1997). A History of Psychiatry: From the Era of the Asylum to the Age of Prozac. John Wiley and Sons.

p. 200. ISBN 9780471245315.11. ^ GE (1989). Chart of the Nuclides, 14th Edition. Nuclear Energy.12. ^ Audi, Georges (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass

Data Center) 729: 3. doi:10.1016/j.nuclphysa.2003.11.001 .13. ^ Duan, Defang (2007-09-26). "Ab initio studies of solid bromine under high pressure". Pysical Review B 76.

doi:10.1103/PhysRevB.76.104113 .14. ^ Tallmadge, John A.; Butt, John B.; Solomon Herman J. (1964). "Minerals From Sea Salt". Ind. Eng. Chem. 56: 44.

doi:10.1021/ie50655a008 .15. ^ Oumeish, Oumeish Youssef (1996). "Climatotherapy at the Dead Sea in Jordan". Clinics in Dermatology 14: 659.

doi:10.1016/S0738-081X(96)00101-0 .16. ^ Al-Weshah, Radwan A. (2008). "The water balance of the Dead Sea: an integrated approach". Hydrological Processes

14: 145. doi:10.1002/(SICI)1099-1085(200001)14:1<145::AID-HYP916>3.0.CO;2-N .17. ^ Emsley, John (2001). "Bromine". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford

University Press. pp. 69–73. ISBN 0198503407.18. ^ Lyday, Phyllis A.. "Comodity Report 2007: Bromine" . United States Geological Survey. Retrieved 2008-09-03.19. ^ a b c d Lyday, Phyllis A.. "Mineral Yearbook 2007: Bromine" . United States Geological Survey. Retrieved 2008-09-03.20. ^ "Bromine:An Important Arkansas Industry" .21. ^ N. A. Khan, F. E. Deatherage, and J. B. Brown (1963), "Stearolic Acid" , Org. Synth.; Coll. Vol. 4: 85122. ^ Green, Joseph (1996). "Mechanisms for Flame Retardancy and Smoke suppression -A Review". Journal of Fire Sciences

14: 426. doi:10.1177/073490419601400602 .23. ^ Kaspersma, Jelle; Doumena, Cindy; Munrob Sheilaand; Prinsa, Anne-Marie (2002). "Fire retardant mechanism of

aliphatic bromine compounds in polystyrene and polypropylene". Polymer Degradation and Stability 77: 325.doi:10.1016/S0141-3910(02)00067-8 .

24. ^ Weil, Edward D.; Levchik, Sergei (2004). "A Review of Current Flame Retardant Systems for Epoxy Resins". Journal ofFire Sciences 22: 25. doi:10.1177/0734904104038107 .

25. ^ Alaeea, Mehran; Ariasb, Pedro; Sjödinc, Andreas; Bergman, Åke (2003). "An overview of commercially used brominatedflame retardants, their applications, their use patterns in different countries/regions and possible modes of release".Environment International 29: 683. doi:10.1016/S0160-4120(03)00121-1 .

26. ^ Messenger, Belinda; Braun, Adolf (2000). "Alternatives to Methyl Bromide for the Control of Soil-Borne Diseases andPests in California" . Pest Management Analysis and Planning Program. Retrieved 2008-11-17.

27. ^ Decanio, Stephen J.; Norman, Catherine S. (2008). "Economics of the "Critical Use" of Methyl bromide under theMontreal Protocol". Contemporary Economic Policy 23: 376. doi:10.1093/cep/byi028 .

28. ^ Darley, H. C. H.; Gray, George Robert (1988). Composition and Properties of Drilling and Completion Fluids. GulfProfessional Publishing. pp. 61–62. ISBN 9780872011472.

29. ^ Kaufman, Vered R.; Garti, Nissim (1984). "Effect of cloudy agents on the stability and opacity of cloudy emulsions for softdrinks". International Journal of Food Science & Technology 19: 255. doi:10.1111/j.1365-2621.1984.tb00348.x .

30. ^ Horowitz, B. Zane (1997). "Bromism from Excessive Cola Consumption',Clinical Toxicology". Clinical Toxicology 35: 315.doi:10.3109/15563659709001219 .

31. ^ Mayeno AN, Curran AJ, Roberts RL, Foote CS (April 1989). "Eosinophils preferentially use bromide to generatehalogenating agents". J. Biol. Chem. 264 (10): 5660–8. PMID 2538427 .

32. ^ a b Gordon W. Gribble (1999). "The diversity of naturally occurring organobromine compounds". Chemical SocietyReviews 28: 335. doi:10.1039/a900201d .

33. ^ Burreson, B. Jay; Moore, Richard E.; Roller, Peter P. (1976). "Volatile halogen compounds in the alga Asparagopsistaxiformis (Rhodophyta)". Journal of Agricultural snd Food Chemistry 24: 856. doi:10.1021/jf60206a040 .

34. ^ Gordon W. Gribble (1998). "Naturally Occurring Organohalogen Compounds". Acc. Chem. Res. 31: 141.doi:10.1021/ar9701777 .

Bromine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Bromine[2/19/2010 12:32:28 AM]

This page was last modified on 17 February 2010 at 13:06.Text is available under the Creative Commons Attribution-ShareAlike License; additional terms may apply. See Terms of

Use for details.Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a non-profit organization.

Contact us Privacy policy About Wikipedia Disclaimers

[edit]

v • d • e

v • d • e

Wikimedia Commons has mediarelated to: Bromine

Look up bromine in Wiktionary,the free dictionary.

35. ^ Friedländer, P. (1909). "Über den Farbstoff des antiken Purpurs aus murex brandaris". Berichte der deutschenchemischen Gesellschaft 42: 765. doi:10.1002/cber.190904201122 .

36. ^ Butler, Alison; Carter-Franklin, Jayme N. (2004). "The role of vanadium bromoperoxidase in the biosynthesis ofhalogenated marine natural products". Natural Product Reports 21: 180. doi:10.1039/b302337k .

External linksWebElements.com – BromineTheodoregray.com – BromineUSGS Minerals Information: BromineBromine Science and Environmental Forum (BSEF)Thermal Conductivity of BROMINEViscosity of BromineChemistry in its element podcast (MP3) from the Royal Society of Chemistry's Chemistry World:Bromine

Diatomic chemical elements

Hydrogen H2 | Nitrogen N2 | Oxygen O2 | Fluorine F2 | Chlorine Cl2 | Bromine Br2 | Iodine I2 | Astatine At2 |

Periodic table

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo

Uue Ubn

Alkali metalsAlkaline earth

metalsLanthanoids Actinoids

Transition

metalsOther metals Metalloids

Other

nonmetalsHalogens Noble gases

Categories: Chemical elements | Halogens | Bromine | Highly Hazardous Chemicals

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

navigation

Main pageContentsFeatured contentCurrent eventsRandom article

interaction

About WikipediaCommunity portalRecent changesContact WikipediaDonate to WikipediaHelp

toolbox

What links hereRelated changesUpload fileSpecial pagesPrintable versionPermanent linkCite this page

discussion edit this page historyTry Beta Log in / create account

tellurium ← iodine → xenon

Br↑I↓At

Appearance

Lustrous metallic gray, violet as a gas

General properties

Name, symbol,number

iodine, I, 53

Element category halogen

Group, period,block

17, 5, p

Standard atomicweight

126.90447 g·mol−1

Electronconfiguration

[Kr] 4d10 5s2 5p5

Electrons per shell 2, 8, 18, 18, 7 (Image)

Physical properties

Phase solid

Density (near r.t.) 4.933 g·cm−3

Melting point 386.85 K, 113.7 °C,

236.66 °F

Boiling point 457.4 K, 184.3 °C,

363.7 °F

Triple point 386.65 K (113°C), 12.1 kPa

Critical point 819 K, 11.7 MPa

Heat of fusion (I2) 15.52 kJ·mol−1

Heat of vaporization (I2) 41.57 kJ·mol−1

Specific heatcapacity

(25 °C) (I2) 54.44

J·mol−1·K−1

Vapor pressure (rhombic)

IodineFrom Wikipedia, the free encyclopedia

For other uses, see Iodine (disambiguation).

Iodine (pronounced /'a . da n/ EYE-o-dyne, /'a . d n/ EYE-o-dən, or in chemistry /'a . di n/ EYE-o-deen; from Greek:ιώδης iodes meaning violet (or purple), is a chemical elementthat has the symbol I and the atomic number 53.

Chemically, iodine is the second least reactive of the halogens,and the second most electropositive halogen, trailing behindastatine in both of these categories. However, the element doesnot occur in the free state in nature. As with all other halogens(members of Group 17 in the periodic table), when freed fromits compounds iodine forms diatomic molecules (I2).

Iodine and its compounds are primarily used in medicine,photography, and dyes. Iodine is rare in the solar system andEarth's crust; however, the iodides are very soluble in water andthe element concentrates in seawater, where it occurs in farhigher amounts than in rocks. This mechanism helps to explainhow the element came to be required in trace amounts by allanimals and some plants, being the heaviest element commonlyused by living organisms (only tungsten, used in enzymes by afew bacteria, is heavier[2][3]).

Contents [hide]

1 Characteristics2 Occurrence3 Structure4 Production5 Isotopes6 History7 Applications

7.1 Disinfectant and water treatment7.2 Staining7.3 Radiocontrast agent7.4 Radioiodine

8 Iodine compounds8.1 Organic compounds

9 Chemistry9.1 Organic synthesis9.2 Clandestine synthetic chemical use

10 Biological role11 Extrathyroidal iodine

11.1 Iodine and the development of cancer11.2 Iodine and immunity11.3 Iodine in salivary glands and oral health11.4 Human dietary intake11.5 Deficiency

12 Radioiodine in biology12.1 Radioiodine and the thyroid

Periodic table

53I

article

search

languages

Afrikaans

БеларускаяBosanskiБългарскиCatalàЧ вашлаČeskyCorsuCymraegDanskDeutsch

EestiΕλληνικάEspañolEsperantoEuskara

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

[edit]

[edit]

P/Pa 1 10 100 1 k 10 k 100 k

at T/K 260 282 309 342 381 457

Atomic properties

Oxidation states 7, 5, 3, 1, -1

(strongly acidic oxide)

Electronegativity 2.66 (Pauling scale)

Ionization energies 1st: 1008.4 kJ·mol−1

2nd: 1845.9 kJ·mol−1

3rd: 3180 kJ·mol−1

Atomic radius 140 pm

Covalent radius 139±3 pm

Van der Waalsradius

198 pm

Miscellanea

Crystal structure orthorhombic

Magnetic ordering diamagnetic[1]

Electrical resistivity (0 °C) 1.3×107Ω·m

Thermalconductivity

(300 K) 0.449 W·m−1·K−1

Bulk modulus 7.7 GPa

CAS registrynumber

7553-56-2

Most stable isotopes

Main article: Isotopes of iodine

iso NA half-life DM DE (MeV) DP

123I syn 13 h ε, γ 0.16 123Te

127I 100% 127I is stable with 74 neutrons

129I trace 15.7×106 y β− 0.194 129Xe

131I syn 8.02070 d β−, γ 0.971 131Xe

This box: view • talk • edit

12.2 Iodine 12512.3 Iodine 129

12.3.1 Radioiodine and the kidney13 Precautions and toxicity of elemental iodine

13.1 Toxicity of iodide ion13.2 Iodine sensitivity

14 See also15 References16 External links

CharacteristicsIodine under standard conditions is a shiny grey solid. It can beseen apparently sublimating at standard temperatures into aviolet-pink gas that has an irritating odor. This halogen formscompounds with many elements, but is less reactive than theother members of its Group VII (halogens) and has somemetallic light reflectance.

Elemental iodine dissolves easily in mostorganic solvents such as hexane orchloroform due to its lack of polarity, but isonly slightly soluble in water. However, thesolubility of elemental iodine in water canbe increased by the addition of potassiumiodide. The molecular iodine reactsreversibly with the negative ion, generatingthe triiodide anion I3− in equilibrium, whichis soluble in water. This is also theformulation of some types of medicinal

(antiseptic) iodine, although tincture of iodine classicallydissolves the element in aqueous ethanol.

Solutions of elemental iodine have the unique property ofexhibiting dramatically different colors depending on the polarityof the solvent. When dissolved in nonpolar solvents like hexane,the solution appears deep violet; in moderately polardichloromethane the solution is dark crimson, and in stronglypolar solvents like acetone or ethanol, it appears dark orange orbrown. This is due to ligand field interactions of solventmolecules with the d-orbitals of iodine, which is the onlyhalogen with a sufficiently occupied electronic configuration toallow such interactions. This same property allows the formation of hypervalent iodine compounds, whichhave expanded bonding orbitals beyond the generally allowed octet rule.

Students who have seen the classroom demonstration in which iodine crystals are gently heated in a testtube to violet vapor may gain the impression that liquid iodine does not exist at atmospheric pressure. Thismisconception arises because the vapor produced has such a deep colour that the liquid appears not to form.In fact, if iodine crystals are heated carefully to just above their melting point of 113.7 °C, the crystals meltinto a liquid which is present under a dense blanket of the vapor.

When iodine is encapsulated into carbon nanotubes it forms atomic chains, whose structure depends on thenanotube diameter.[4]

Occurrence

In the gasphase iodineshows its violetcolor.

FrançaisFurlanGaeilgeGaelgGalegoHak-kâ-fa

HrvatskiIdoBahasa IndonesiaИронауÍslenskaItalianoעבריתBasa JawaKiswahiliKreyòl ayisyenLatinaLatviešuLëtzebuergeschLietuviųLíguruLojbanMagyarМакедонски

Nederlands

Norsk (bokmål)Norsk (nynorsk)OccitanO'zbek

PlattdüütschPolskiPortuguêsRomânăRuna SimiРусскийSeelterskShqipSicilianuSimple EnglishSlovenčinaSlovenščinaСрпски / SrpskiSrpskohrvatski /СрпскохрватскиSuomiSvenskaTagalog

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

[edit]

[edit]

Iodine naturally occurs in the environment chiefly as a dissolved iodide inseawater, although it is also found in some minerals and soils.[5] This elementalso exists in small amounts in the mineral caliche, found in Chile, between theAndes and the sea. A type of seaweed, kelp, tends to be high in iodine as well.

Organoiodine compounds are produced by marine life forms, the most notablebeing iodomethane (commonly called methyl iodide). The total iodomethanethat is produced by the marine environment, by microbial activity in rice paddiesand by the burning of biological material is estimated to be 214 kilotonnes.[6]

The volatile iodomethane is broken up by oxidation reactions in the atmosphereand a global iodine cycle is established.[5][6] Although the element is actually

quite rare, kelp and certain plants and other algae have some ability to concentrate iodine, which helpsintroduce the element into the food chain.

StructureIodine crystallizes in the orthorhombic space group Cmca No 64,Pearson symbol oS8, the same as black phosphorus. In the solidstate, I2 molecules are still represented by a short I-I bond of270 pm.

ProductionFrom the several places in which iodine occurs in nature only twoare used as source for iodine: the caliche, found in Chile and theiodine containing brines of gas and oil fields, especially in Japanand the United States.

The caliche, found in Chile contains sodium nitrate, which is the main product of the mining activities andsmall amounts of sodium iodate and sodium iodide. During leaching and production of pure sodium nitrate thesodium iodate and iodide is extracted.[7] The high concentration of iodine in the caliche and the extensivemining made Chile the largest producer of iodine in 2007.

Most other producers use natural occurring brine for the productionof iodine. The Japanese Minami Kanto gas field east of Tokyo andthe American Anadarko Basin gas field in northwest Oklahoma arethe two largest sources for iodine from brine. The brine has atemperature of over 60°C due to the depth of the source. The brineis first purified and acidified using sulfuric acid, then the iodidepresent is oxidized to iodine with chlorine. An iodine solution isproduced, but is dilute and must be concentrated. Air is blown intothe solution, causing the iodine to evaporate, then it is passed into an absorbing tower containing acid wheresulfur dioxide is added to reduce the iodine. The hydrogen iodide (HI) is reacted with chlorine to precipitatethe iodine. After filtering and purification the iodine is packed.[7][8]

2 HI + Cl2 → I2↑ + 2 HClI2 + 2 H2O + SO2 → 2 HI + H2SO42 HI + Cl2 → I2↓ + 2 HCl

The production of iodine from seawater via electrolysis is not used due to the sufficient abundance of iodine-rich brine. Another source of iodine was kelp, used in the 18th and 19th centuries, but it is no longereconomically viable.

Commercial samples often contain a large amount of impurities; they may be removed by sublimation. Theelement may also be prepared in an ultrapure form through the reaction of potassium iodide with copper(II)sulfate, which gives copper(II) iodide initially. That decomposes spontaneously to copper(I) iodide and iodine:

Cu2+ + 2 I– → CuI2

Iodomethane

Structure of solid iodine

Iodine output in 2005

То икTürkçeУкраїнська

Tiếng ViệtWinarayYorùbá

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

[edit]

[edit]

[edit]

[edit]

2 CuI2 → 2 CuI + I2

There are also other methods of isolating this element in the laboratory, for example the method used toisolate other halogens: oxidation of the iodide in hydrogen iodide (often made in situ with an iodide andsulfuric acid) by manganese dioxide (see below in Descriptive chemistry).

IsotopesMain article: Isotopes of iodine

There are 37 known (characterized) isotopes of iodine, but only one, 127I, is stable.

In many ways, 129I is similar to 36Cl. It is a soluble halogen, fairly non-reactive, exists mainly as a non-sorbing anion, and is produced by cosmogenic, thermonuclear, and in-situ reactions. In hydrologic studies,129I concentrations are usually reported as the ratio of 129I to total I (which is virtually all 127I). As is the casewith 36Cl/Cl, 129I/I ratios in nature are quite small, 10−14 to 10−10 (peak thermonuclear 129I/I during the1960s and 1970s reached about 10−7). 129I differs from 36Cl in that its half-life is longer (15.7 vs. 0.301million years), it is highly biophilic, and occurs in multiple ionic forms (commonly, I− and IO3

−) which havedifferent chemical behaviors. This makes it fairly easy for 129I to enter the biosphere as it becomesincorporated into vegetation, soil, milk, animal tissue, etc.

Excesses of stable 129Xe in meteorites have been shown to result from decay of "primordial" iodine-129produced newly by the supernovas which created the dust and gas from which the solar system formed. 129Iwas the first extinct radionuclide to be identified as present in the early solar system. Its decay is the basis ofthe I-Xe Iodine-xenon radiometric dating scheme, which covers the first 85 million years of solar systemevolution.

Effects of various radioiodine isotopes in biology are discussed below.

History

Iodine was discovered by Bernard Courtois in 1811.[9][10] He was born to a manufacturer of saltpeter (a vitalpart of gunpowder). At the time of the Napoleonic Wars, France was at war and saltpeter was in greatdemand. Saltpeter produced from French niter beds required sodium carbonate, which could be isolated fromseaweed washed up on the coasts of Normandy and Brittany. To isolate the sodium carbonate, seaweed wasburned and the ash then washed with water. The remaining waste was destroyed by adding sulfuric acid.One day Courtois added too much sulfuric acid and a cloud of purple vapor rose. Courtois noted that thevapor crystallized on cold surfaces making dark crystals. Courtois suspected that this was a new element butlacked the money to pursue his observations.

However he gave samples to his friends, Charles Bernard Desormes (1777–1862) and Nicolas Clément(1779–1841), to continue research. He also gave some of the substance to Joseph Louis Gay-Lussac (1778–1850), a well-known chemist at that time, and to physicist André-Marie Ampère (1775–1836). On 29November 1813, Dersormes and Clément made public Courtois’s discovery. They described the substance toa meeting of the Imperial Institute of France. On December 6, Gay-Lussac announced that the newsubstance was either an element or a compound of oxygen.[11][12][13] Ampère had given some of his sampleto Humphry Davy (1778–1829). Davy did some experiments on the substance and noted its similarity tochlorine.[14] Davy sent a letter dated December 10 to the Royal Society of London stating that he hadidentified a new element.[15] A large argument erupted between Davy and Gay-Lussac over who identifiediodine first but both scientists acknowledged Courtois as the first to isolate the chemical element.

Applications

Disinfectant and water treatmentElemental iodine is used as a disinfectant in various forms. The iodine exists as the element, or as the watersoluble triiodide anion I3- generated in situ by adding iodide to poorly-soluble iodine (the reverse chemicalreaction makes some free elemental iodine available for antisepsis). Alternatively, iodine may come from

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

[edit]

[edit]

[edit]

[edit]

iodophors, which contain iodine complexed with a solubilizing agent (iodide ion may be thought of loosely asthe iodophor in triiodide water solutions). Examples of such preparations include:[16]

Tincture of iodine (iodine in ethanol, or iodine and sodium iodide in a mixture of ethanol and water)Lugol's iodine (iodine and iodide in water, forming mostly triiodide)Povidone iodine (an iodophor)

StainingSee also: Staining

Iodine is a common general stain used in thin-layerchromatography. It is also used in the Gram stain as a mordant,after the sample is treated with crystal violet.

In particular, iodine forms an intense blue complex with the glucosepolymers starch and glycogen. Many applications rely on thisproperty:

Iodometry. The concentration of an oxidant can be determinedby adding it to an excess of iodide with a little free iodine, todestroy elemental iodine/triiodide as a result of oxidation by theoxidant. A starch indicator is then used as the indicator close tothe end-point, in order to increase the visual contrast (dark bluebecomes colorless, instead of the yellow of dilute triiodidebecoming colorless).An Iodine test may be used to test a sample substance for the presence of starch.The Iodine clock reaction is an extension of the techniques in iodometry.Iodine solutions are used in counterfeit banknote detection pens; the premise being that counterfeitbanknotes made using commercially available paper contain starch.Starch-iodide paper are used to test for the presence of oxidants such as peroxides. The oxidants convertiodide to iodine, which shows up as blue. A solution of starch and iodide can perform the samefunction.[17]

During colposcopy, Lugol's iodine is applied to the vagina and cervix. Normal vaginal tissue stains browndue to its high glycogen content (a color-reaction similar to that with starch), while abnormal tissuesuspicious for cancer does not stain, and thus appears pale compared to the surrounding tissue. Biopsy ofsuspicious tissue can then be performed. This is called a Schiller's Test.

Radiocontrast agentIodine, as a heavy element, is quite radio-opaque. Organiccompounds of a certain type (typically iodine-substituted benzenederivatives) are thus used in medicine as X-ray radiocontrastagents for intravenous injection. This is often in conjunction withadvanced X-ray techniques such as angiography and CT scanning

RadioiodineSome radioactive iodine isotopes can be used to treat thyroidcancer. The body accumulates iodine in the thyroid, thusradioactive iodine can selectively damage growing thyroid cancercells while the radioactive dose to the rest of the body remainssmall.

Iodine compoundsSee also: Category:Iodine compounds

Iodine forms many compounds. Potassium iodide is the most commercially significant iodine compound. It is a

Testing a seed for starch with a solutionof iodine

Diatrizoic acid, a radiocontrast

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

[edit]

[edit]

convenient source of the iodide anion; it is easier to handle than sodium iodide because it is not hygroscopic.Sodium iodide is especially useful in the Finkelstein reaction, because it is soluble in acetone, whilepotassium iodide is poorly so. In this reaction, an alkyl chloride is converted to an alkyl iodide. This relies onthe insolubility of sodium chloride in acetone to drive the reaction:

R-Cl (acetone) + NaI (acetone) → R-I (acetone) + NaCl (s)

Iodic acid (HIO3) and its salts are strong oxidizers. Periodic acid (HIO4) cleaves vicinal diols along the C-Cbond to give aldehyde fragments. 2-Iodoxybenzoic acid and Dess-Martin periodinane are hypervalent iodineoxidants used to specifically oxidize alcohols to ketones or aldehydes. Iodine pentoxide is a strong oxidant aswell.

Interhalogen compounds are well known; examples include iodine monochloride and trichloride; iodinepentafluoride and heptafluoride.

HI He

LiI BeI2 BI3 CI4 NI3

I2O4,I2O5,I4O9

IF,IF3,IF5,IF7

Ne

NaI MgI2 AlI3 SiI4PI3,P2I4

SICl,ICl3

Ar

KI CaI2 Sc TiI4 VI3 Cr MnI2 Fe CoI2 NiI2 CuI ZnI2 Ga2I6GeI2,GeI4

AsI3 Se IBr Kr

RbI SrI2 Y ZrI4 Nb Mo Tc Ru Rh Pd AgI CdI2 InI3SnI4,SnI2

SbI3 TeI4 I Xe

CsI BaI2 Hf Ta W Re Os Ir Pt AuIHg2I2,HgI2

TlI PbI2 Bi Po At Rn

Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Uub Uut Uuq Uup Uuh Uus Uuo

↓

La Ce Pr Nd Pm SmI2 Eu Gd TbI3 Dy Ho Er Tm Yb Lu

Ac ThI4 Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

Organic compoundsSee also: Organoiodine compound

Many organoiodine compounds exist, the simplest is iodomethane, approved as a soil fumigant. Iodinatedorganics are used as synthetic reagents, and also radiocontrast agents.

Biologically active substances like the thyroid hormones are naturally occurring organoiodine compounds.[18]

ChemistryElemental iodine is poorly soluble in water, with one gram dissolving in 3450 ml at 20 °C and 1280 ml at50 °C. By contrast with chlorine, the formation of the hypohalite ion (IO–) in neutral aqueous solutions ofiodine is negligible.

I2+ H2O H+ + I− + HIO (K = 2.0×10−13)[19]

Solubility in water is greatly improved if the solution contains dissolved iodides such as hydroiodic acid,potassium iodide, or sodium iodide; this extra solubility results from the high solubility of the I3− ion. Dissolvedbromides also improve water solubility of iodine. Iodine is soluble in a number of organic solvents, includingethanol (20.5 g/100 ml at 15 °C, 21.43 g/100 ml at 25 °C), diethyl ether (20.6 g/100 ml at 17 °C,25.20 g/100 ml at 25 °C), chloroform, acetic acid, glycerol, benzene (14.09 g/100 ml at 25 °C), carbontetrachloride (2.603 g/100 ml at 35 °C), and carbon disulfide (16.47 g/100 ml at 25 °C).[20] Aqueous and

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

[edit]

ethanol solutions are brown. Solutions in chloroform, carbon tetrachloride, and carbon disulfide are violet.

Molecular iodine can be prepared by oxidizing iodides with chlorine:

2 I− + Cl2 → I2 + 2 Cl−

or with manganese dioxide in acid solution:[19]

2 I− + 4 H+ + MnO2 → I2 + 2 H2O + Mn2+

Iodine is reduced to hydroiodic acid by hydrogen sulfide:[21]

I2 + H2S → 2 HI + S↓

or by hydrazine:

2 I2 + N2H4 → 4 HI + N2

Iodine is oxidized to iodate by nitric acid:[22]

I2 + 10 HNO3 → 2 HIO3 + 10 NO2 + 4 H2O

or by chlorates:[22]

I2 + 2 ClO3− → 2 IO3

− + Cl2

Iodine is converted in a two stage reaction to iodide and iodate in solutions of alkali hydroxides (such assodium hydroxide):[19]

I2 + 2 OH− → I− + IO− + H2O (K = 30)

3 IO− → 2 I− + IO3− (K = 1020)

Despite having the lowest electronegativity of the common halogens, iodine reacts violently with some metals,such as aluminum:

3 I2 + 2 Al → 2 AlI3

This reaction produces 314 kJ per mole of aluminum, comparable to thermite's 425 kJ. Yet the reactioninitiates spontaneously, and if unconfined, causes a cloud of gaseous iodine due to the high heat.

When dissolved in fuming sulfuric acid (65% oleum), iodine forms an intense blue solution. This has beenshown to be due to the formation of the I+2 cation, the result of iodine being oxidised by SO3:[23]

2 I2 + 2 SO3 + H2SO4 → 2 I+2 + SO2 + 2 HSO−4

The I+2 cation is also formed in the oxidation of iodine by SbF5 or TaF5. The resulting I+2Sb2F−11 or I+2Ta2F−

11can be isolated as deep blue crystals. The solutions of these salts turn red when cooled below −60 °C, due tothe formation of the I2+

4 cation:[23]

2 I+2 I2+4

Under slightly more alkaline conditions, I2+4 disproportionates into I+3 and an iodine(III) compound. Excess

iodine can then react with I+3 to form I+5 (green) and I3+15 (black).[23]

Organic synthesisWith phosphorus, iodine is able to replace hydroxyl groups on alcohols with iodide. For example, thesynthesis of methyl iodide from methanol, red phosphorus, and iodine.[24] The iodinating reagent isphosphorus triiodide that is formed in situ:

3 CH3OH + PI3 → 3 CH3I + H3PO3

Phosphorous acid is formed as a side-product.

The iodoform test uses an alkaline solution of iodine to react with methyl ketones to give the labiletriiodomethide leaving group, forming iodoform which precipitates.

Iodine is sometimes used to activate magnesium when preparing Grignard reagents; aryl and alkyl iodidesboth form Grignard reagents. Alkyl iodides such as iodomethane are good alkylating agents. Some drawbacksto use of iodo-organics in chemical synthesis are:

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

[edit]

[edit]

[edit]

[edit]

iodine compounds tend to be more expensive than the corresponding bromides and chlorides, in thatorderiodides tend to be much stronger alkylating agents, and so are more toxic (e.g. methyl iodide is very toxic(T+)[25]

low molecular weight iodides tend to have a much higher equivalent weight, compared with otheralkylating agents (e.g. methyl iodide versus dimethyl carbonate), due to the atomic mass of iodine.

Clandestine synthetic chemical useIn the United States, the Drug Enforcement Agency (DEA) regards iodine and compounds containing iodine(ionic iodides, iodoform, ethyl iodide, and so on) as reagents useful for the clandestine manufacture ofmethamphetamine. Persons who attempt to purchase significant quantities of such chemicals withoutestablishing a legitimate use are likely to find themselves the target of a DEA investigation. Persons sellingsuch compounds without doing due diligence to establish that the materials are not being diverted toclandestine use may be subject to stiff penalties, such as expensive fines or even imprisonment.[26][27]

Biological roleMain article: Iodine in biology

Iodine is an essential trace element for life, the heaviest element commonly needed by living organisms, andthe second-heaviest known to be used by any form of life (only tungsten, a component of a few bacterialenzymes, has a higher atomic number and atomic weight). Iodine's main role in animal biology is asconstituents of the thyroid hormones, thyroxine (T4) and triiodothyronine (T3). These are made from additioncondensation products of the amino acid tyrosine, and are stored prior to release in an iodine-containingprotein called thyroglobulin. T4 and T3 contain four and three atoms of iodine per molecule, respectively. Thethyroid gland actively absorbs iodide from the blood to make and release these hormones into the blood,actions which are regulated by a second hormone TSH from the pituitary. Thyroid hormones arephylogenetically very old molecules which are synthesized by most multicellular organisms, and which evenhave some effect on unicellular organisms.

Thyroid hormones play a basic role in biology, acting on gene transcription to regulate the basal metabolicrate.[citation needed] The total deficiency of thyroid hormones can reduce basal metabolic rate up to 50%, whilein excessive production of thyroid hormones the basal metabolic rate can be increased by100%.[citation needed] T4 acts largely as a precursor to T3, which is (with minor exceptions) the biologicallyactive hormone.

Extrathyroidal iodineIodine accounts for 65% of the molecular weight of T4 and 59% of the T3. 15–20 mg of iodine isconcentrated in thyroid tissue and hormones, but 70% of the body's iodine is distributed in other tissues,including mammary glands, eyes, gastric mucosa, the cervix, and salivary glands. In the cells of these tissuesiodide enters directly by sodium-iodide symporter (NIS). Its role in mammary tissue is related to fetal andneonatal development, but its role in the other tissues is unknown.[28] It has been shown to act as anantioxidant in these tissues.[28]

Iodine may have a relationship with selenium, and iodine supplementation in selenium-deficient populationsmay pose risks for thyroid function.[28]

The US Food and Nutrition Board and Institute of Medicine recommended daily allowance of iodine rangesfrom 150 micrograms /day for adult humans to 290 micrograms /day for lactating mothers. However, thethyroid gland needs no more than 70 micrograms /day to synthesize the requisite daily amounts of T4 and T3.These higher recommended daily allowance levels of iodine seem necessary for optimal function of a numberof body systems, including lactating breast, gastric mucosa, salivary glands, oral mucosa, thymus, epidermis,choroid plexus, etc.[29][30][31]

Iodine and the development of cancer

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

[edit]

[edit]

[edit]

[edit]

*Breast cancer. It is known that a diet lacking in iodine is connected with adverse health effects collectivelyreferred as iodine deficiency diseases or disorders. Studies also indicate that iodine deficiency, either dietaryor pharmacologic, can lead to breast atypia and increased incidence of malignancy in animal models, whileiodine treatment can reverse dysplasia.[32][33][34] Laboratory evidences demonstrate that the effect of iodineon breast cancer is in part independent of thyroid function and that iodine inhibit cancer promotion throughmodulation of the estrogen pathway. Gene array profiling of estrogen responsive breast cancer cell line showsthat the combination of iodine and iodide alters gene expression and inhibits the estrogen response throughup-regulating proteins involved in estrogen metabolism. Whether iodine/iodide will be useful as an adjuvanttherapy in the pharmacologic manipulation of the estrogen pathway in women with breast cancer has notbeen determined clinically.[32]

*Iodine and stomach cancer. Some researchers have found an epidemiologic correlation between iodinedeficiency, iodine-deficient goitre and gastric cancer;[35] [36][37] a decrease of the incidence of death ratefrom stomach cancer after implementation of the effective iodine-prophylaxis was reported too.[38] Theproposed mechanism of action is that iodide ion can function in gastric mucosa as an antioxidant reducingspecies that can detoxify poisonous reactive oxygen species, such as hydrogen peroxide.

Iodine and immunityIodine has important actions in the immune system. The high iodide-concentration of thymus suggests ananatomical rationale for this role of iodine in immune system.[39][40][41][42][43][44]

Iodine in salivary glands and oral healthThe trophic, antioxidant and apoptosis-inductor actions and the presumed anti-tumour activity of iodides mightalso be important for prevention of oral and salivary glands diseases.[45][46][47][48][49][50]

Human dietary intakeThe United States Recommended Daily Allowance (RDA) is 150 micrograms per day (μg/day) for both menand women, with a Tolerable Upper Intake Level (UL) for adults is 1,100 μg/day (1.1 mg/day).[51] Thetolerable upper limit was assessed by analyzing the effect of supplementation on thyroid-stimulatinghormone.[28]

Natural sources of iodine include sea life, such as kelp and certain seafood, as well as plants grown oniodine-rich soil.[52][53] Iodized salt is fortified with iodine.[53]

As of 2000, the median intake of iodine from food in the United States was 240 to 300 μg/day for men and190 to 210 μg/day for women.[51] In Japan, consumption is much higher due to the frequent consumption ofseaweed or kombu kelp.[28]

After iodine fortification programs (e.g. iodized salt) have been implemented, some cases of iodine-inducedhyperthyroidism have been observed (so called Jod-Basedow disease). The condition mainly seems to occurin people over forty, and the risk appears higher when iodine deficiency is severe and the initial rise in iodineintake is high.[54]

DeficiencyMain article: Iodine deficiency

In areas where there is little iodine in the diet, typically remote inland areas and semi-arid equatorial climateswhere no marine foods are eaten, iodine deficiency gives rise to hypothyroidism, symptoms of which areextreme fatigue, goitre, mental slowing, depression, weight gain, and low basal body temperatures.[55]

Iodine deficiency is the leading cause of preventable mental retardation, a result which occurs primarily whenbabies or small children are rendered hypothyroidic by a lack of the element. The addition of iodine to tablesalt has largely eliminated this problem in the wealthier nations, but as of March 2006, iodine deficiencyremained a serious public health problem in the developing world.[56] Iodine deficiency is also a problem incertain areas of Europe. In Germany it has been estimated to cause a billion dollars in health care costs peryear.[28]

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

[edit]

[edit]

[edit]

[edit]

[edit]

[edit]

Radioiodine in biology

Radioiodine and the thyroidThe most common compounds of iodine are the iodides of sodium (NaI) and potassium (KI) and the iodates(KIO3), as elemental iodine is mildly toxic to all living things. Normal iodine is an essential precursor for themanufacture of thyroid hormone.

Due to preferential uptake of iodine by the thyroid, isotopes with short half lives such as I131 can be used forthyroid ablation, a procedure in which radioactive iodine is administered intravenously or orally following adiagnostic scan. This procedure is generally performed on patients with thyroid cancer or hyperfunctioningthyroid tissue. After uptake, the iodine undergoes degeneration via beta decay, destroying its associatedthyroid tissue. Normally thyroidectomy is performed prior to ablation to avoid side effects of epilation andradiation toxicity. The purpose of radioablation is to destroy remnant tissue that was unable to be removedwith surgery.

Lower energy isotopes such as iodine-123, and less commonly iodine-125, are used as tracers to evaluatethe anatomic and physiologic function of the thyroid. Abnormal results may be caused by disorders such asGraves' Disease or Hashimoto's thyroiditis.

Potassium iodide has been distributed to populations exposed to nuclear fission accidents such as theChernobyl disaster. The iodide solution SSKI, a saturated solution of potassium (K) iodide in water, hasbeen used to block absorption of the radioiodine (it has no effect on other radioisotopes from fission). Tabletscontaining potassium iodide are now also manufactured and stocked in central disaster sites by thegovernments for this purpose. In theory, many harmful late-cancer effects of nuclear fallout might beprevented in this way, since an excess of thyroid cancers, presumably due to radioiodine uptake, is the onlyproven radioisotope contamination effect after a fission accident, or from contamination by fallout from anatomic bomb (prompt radiation from the bomb also cases other cancers, such as leukemias, directly). Takinglarge amounts of thyroid saturates iodide receptors prevents uptake of most radioactive iodine-131 that maybe present from fission product exposure (although it does not protect from other radioisotopes, nor from anyother form of direct radiation). The protective effect of KI lasts approximately 24 hours, so must be dosed dailyuntil a risk of significant exposure to radioiodines from fission products no longer exists.[57][58] Iodine-131(the most common radioiodine contaminant in fallout) also decays relatively rapidly with a half-life of 8 days,so that 99.95% of the original radioiodine is gone after three months.

Iodine 125Iodine-125 is also commonly used by radiation oncologists in low dose rate brachytherapy in the treatment ofcancer at sites other than the thyroid, especially in prostate cancer. The radioiodine is encapsulated intitanium seeds and implanted in the area of tumor involvement. In contrast to the blood-borne disseminationof radioiodine used in the thyroid, the radioiodine acts only locally in the area where it is implanted.

Iodine 129Iodine-129 (129I; half-life 15.7 million years) is a product of cosmic ray spallation on various isotopes ofxenon in the atmosphere, in cosmic ray muon interaction with tellurium-130, and also uranium and plutoniumfission, both in subsurface rocks and nuclear reactors. Artificial nuclear processes, in particular nuclear fuelreprocessing and atmospheric nuclear weapons tests, have now swamped the natural signal for this isotope.Nevertheless, it now serves as a groundwater tracer as indicator of nuclear waste dispersion into the naturalenvironment. In a similar fashion, 129I was used in rainwater studies to track fission products following theChernobyl disaster.

Radioiodine and the kidneyIn the 1970s imaging techniques were developed to employ radioiodine in diagnostics for renal hypertension;however methods using other chemical compounds, such as DMSA, are more popular in clinics nowadays.

Precautions and toxicity of elemental iodine

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

[edit]

[edit]

[edit]

[edit]

Elemental iodine is an oxidizing irritant and direct contact with skin can cause lesions, so iodine crystalsshould be handled with care. Solutions with high elemental iodine concentration such as tincture of iodine arecapable of causing tissue damage if use for cleaning and antisepsis is prolonged.

Elemental iodine (I2) is poisonous if taken orally in larger amounts; 2–3 grams of it is a lethal dose for anadult human.

Iodine vapor is very irritating to the eye, to mucous membranes, and in the respiratory tract. Concentration ofiodine in the air should not exceed 1 mg/m³ (eight-hour time-weighted average).

When mixed with ammonia and water, elemental iodine forms nitrogen triiodide which is extremely shocksensitive and can explode unexpectedly.

Toxicity of iodide ionExcess iodine has symptoms similar to those of iodine deficiency. Commonly encountered symptoms areabnormal growth of the thyroid gland and disorders in functioning and growth of the organism as a whole.Iodides are similar in toxicity to bromides.[citation needed]

Iodine sensitivitySome people develop a sensitivity to iodine. Application of tincture of iodine can cause a rash. Some cases ofreaction to Povidone-iodine (Betadine) have been documented to be a chemical burn.[59] Eating iodine-containing foods can cause hives. Medical use of iodine (i.e. as a contrast agent, see above) can causeanaphylactic shock in highly iodine sensitive patients. Some cases of sensitivity to iodine can be formallyclassified as iodine allergies. Iodine sensitivity is rare but has a considerable effect given the extremelywidespread use of iodine-based contrast media[60].

See alsoIodide as an antioxidantChemical oxygen iodine laserNutrition facts labelStarch indicator

References1. ^ Magnetic susceptibility of the elements and inorganic compounds , in Handbook of Chemistry and Physics 81st edition,

CRC press.2. ^ J McMaster and John H Enemark (1998). "The active sites of molybdenum- and tungsten-containing enzymes". Current

Opinion in Chemical Biology 2: 201. doi:10.1016/S1367-5931(98)80061-6 .3. ^ Russ Hille (2002). "Molybdenum and tungsten in biology". Trends in Biochemical Sciences 27: 360. doi:10.1016/S0968-

0004(02)02107-2 .4. ^ Guan, L; Suenaga, K; Shi, Z; Gu, Z; Iijima, S (Jun 2007). "Polymorphic structures of iodine and their phase transition in

confined nanospace.". Nano letters 7 (6): 1532. doi:10.1021/nl070313t . PMID 17477579 .5. ^ a b Dissanayake, C. B.; Chandrajith, Rohana; Tobschall, H. J. (1999). "The iodine cycle in the tropical environment —

implications on iodine deficiency disorders". International Journal of Environmental Studies 56: 357.doi:10.1080/00207239908711210 .

6. ^ a b N. Bell, L. Hsu, D. J. Jacob, M. G. Schultz, D. R. Blake, J. H. Butler, D. B. King, J. M. Lobert, and E. Maier-Reimer(2002). "Methyl iodide: Atmospheric budget and use as a tracer of marine convection in global models". Journal ofGeophysicalResearch 107: 4340. doi:10.1029/2001JD001151 .

7. ^ a b Jessica Elzea Kogel, Nikhil C. Trivedi, James M. Barker, Stanley T. Krukowski (2006). Industrial Minerals & Rocks:Commodities, Markets, and Uses . SME. pp. 541–552. ISBN 9780873352338.

8. ^ Tatsuo Maekawa, Shun-Ichiro Igari and Nobuyuki Kaneko (2006). "Chemical and isotopic compositions of brines fromdissolved-in-water type natural gas fields in Chiba, Japan". Geochemical Journal 40: 475. doi:10.2343/geochemj.40.475 .

9. ^ Bernard Courtois (1813). "Découverte d'une substance nouvelle dans le Vareck". Annales de chimie 88: 304. In French,seaweed that had been washed onto the shore was called "varec", "varech", or "vareck", whence the English word "wrack".Later, "varec" also referred to the ashes of such seaweed: the ashes were used as a source of iodine and salts of sodiumand potassium.

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

10. ^ Patricia A. Swain (2005). "Bernard Courtois (1777-1838) famed for discovering iodine (1811), and his life in Paris from1798" . Bulletin for the History of Chemistry 30 (2): 103.

11. ^ J. Gay-Lussac (1813). "Sur un nouvel acide formé avec la substance décourverte par M. Courtois". Annales de chimie88: 311.

12. ^ J. Gay-Lussac (1813). "Sur la combination de l'iode avec d'oxigène". Annales de chimie 88: 319.13. ^ J. Gay-Lussac (1814). "Mémoire sur l'iode". Annales de chimie 91: 5.14. ^ H. Davy (1813). "Sur la nouvelle substance découverte par M. Courtois, dans le sel de Vareck". Annales de chemie 88:

322.15. ^ Humphry Davy (January 1, 1814). "Some Experiments and Observations on a New Substance Which Becomes a Violet

Coloured Gas by Heat" . Phil. Trans. R. Soc. Lond. 104: 74. doi:10.1098/rstl.1814.0007 .16. ^ Block, Seymour Stanton (2001). Disinfection, sterilization, and preservation. Hagerstwon, MD: Lippincott Williams &

Wilkins. p. 159. ISBN 0-683-30740-1.17. ^ R. Toreki. "Peroxide" . The MSDS HyperGlossary.18. ^ Gribble, G. W. (1996). "Naturally occurring organohalogen compounds - A comprehensive survey". Progress in the

Chemistry of Organic Natural Products 68: 1–423. PMID 8795309 .19. ^ a b c F. A. Cotton and G. Wilkinson (1988). Advanced Inorganic Chemistry, 5th ed.. John Wiley & Sons. ISBN

0471849979.20. ^ Martha Windholz, editor ; Susan Budavari, associate editor ; Lorraine Y. Stroumtsos, assistant editor ; Margaret Noether

Fertig, assistant editor. (1976). Merck Index of Chemicals and Drugs, 9th ed. S.l.: J A Majors Company. ISBN 0911910263.21. ^ N.L. Glinka (1981). General Chemistry (volume 2). Mir Publishing.22. ^ a b Linus Pauling (1988). General Chemistry. Dover Publications. ISBN 0486656225.23. ^ a b c Egon Wiberg; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. p. 419-420.

ISBN 0123526515.24. ^ King, C. S.; Hartman, W. W. (1943), "Methyl Iodide" , Org. Synth.; Coll. Vol. 2: 39925. ^ "Safety data for iodomethane" . Oxford University.26. ^ 21 "USC Sec. 872 01/22/02" .27. ^ "Chemical Supplier Convicted of Diversion of Iodine" .28. ^ a b c d e f Patrick L (2008). "Iodine: deficiency and therapeutic considerations" . Altern Med Rev 13 (2): 116. PMID

18590348 .29. ^ Brown-Grant, K. (1961). "Extrathyroidal iodide concentrating mechanisms" . Physiol Rev. 41: 189.30. ^ Spitzweg, C., Joba, W., Eisenmenger, W. and Heufelder, A.E. (1998). "Analysis of human sodium iodide symporter gene

expression in extrathyroidal tissues and cloning of its complementary deoxyribonucleic acid from salivary gland, mammarygland, and gastric mucosa". J Clin Endocrinol Metab. 83: 1746. doi:10.1210/jc.83.5.1746 .

31. ^ Banerjee, R.K., Bose, A.K., Chakraborty, t.K., de, S.K. and datta, A.G. (1985). "Peroxidase catalysed iodotyrosineformation in dispersed cells of mouse extrathyroidal tissues". J Endocrinol. 2: 159.

32. ^ a b Stoddard II, F. R.; Brooks, A. D.; Eskin, B. A.; Johannes, G. J. (2008). "Iodine Alters Gene Expression in the MCF7Breast Cancer Cell Line: Evidence for an Anti-Estrogen Effect of Iodine" . International Journal of Medical Science 5 (4):189. PMID 18645607 . PMC 2452979 .

33. ^ Eskin, B. A.; Grotkowski, C. E.; Connolly, C. P.; Ghent W. R.; (1995). "Different tissue responses for iodine and iodide inrat thyroid and mammary glands". Bioligal Trace Elements Research 49 (5): 9. doi:10.1007/BF02788999 . PMID14965610 .

34. ^ Venturi, S.; Grotkowski, CE; Connolly, CP; Ghent, WR (2001). "Is there a role for iodine in breast diseases?". The Breast10 (1): 379. doi:10.1054/brst.2000.0267 . PMID 7577324 .

35. ^ Josefssson M, Ekblad E. (2009). Victor R. Preedy, Gerard N. Burrow MD, Ronald Watson. ed. Sodium Iodide Symporter(NIS) in Gastric Mucosa: Gastric Iodide Secretion. In: Comprehensive Handbook of Iodine: Nutritional, Biochemical,Pathological and Therapeutic Aspects.

36. ^ Abnet CC, Fan JH, Kamangar F, Sun XD, Taylor PR, Ren JS, Mark SD, Zhao P, Fraumeni JF Jr, Qiao YL, Dawsey SM(2006). Self-reported goiter is associated with a significantly increased risk of gastric noncardia adenocarcinoma in a largepopulation-based Chinese cohort.. 119. p. 1508.

37. ^ Behrouzian R, Aghdami N. (2004). East Mediterr Health J.. 10. p. 921.38. ^ Golkowski F, Szybinski Z, Rachtan J, Sokolowski A, Buziak-Bereza M, Trofimiuk M, Hubalewska-Dydejczyk A, Przybylik-

Mazurek E, Huszno B. (2007). "Iodine prophylaxis--the protective factor against stomach cancer in iodine deficient areas".Eur J Nutr. 46: 251. doi:10.1007/s00394-007-0657-8 .

39. ^ Venturi S, Venturi M (September 2009). "Iodine, thymus, and immunity". Nutrition 25 (9): 977–9.doi:10.1016/j.nut.2009.06.002 . PMID 19647627 .

40. ^ Venturi S.; Venturi A, Cimini D, Arduini C, Venturi M, Guidi A. (1993). "A new hypothesis: iodine and gastric cancer.".Europ. J. Cancer. Prev. 2: 17.

41. ^ Marani L; Venturi S, Masala R (1985). "Role of iodine in delayed immune response.". Isr. J. Med. Sci. 21: 864.

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

[edit]

v • d • e

v • d • e

v • d • e

Wikimedia Commons has mediarelated to: Iodine

Look up iodine in Wiktionary,the free dictionary.

42. ^ Ma F; Zhao W, Kudo M, Aoki K, Misumi J. (2002). "Inhibition of vacuolation toxin activity of Helicobacter pylori by iodine,nitrite and potentiation by sodium chloride, sterigmatocystin and fluoride.". Toxicol in Vitro 16: 531. doi:10.1016/S0887-2333(02)00045-0 .

43. ^ Klebanoff S.J. (1967). "Iodination of bacteria: A bacterial mechanism.". J Exp Med 126: 1063.doi:10.1084/jem.126.6.1063 .

44. ^ "Iodine enhances ig-G-synthesis by human peripheral blood Iyphocytes in vitro.". Acta Endocr 103: 103. 1983.45. ^ Venturi S.; Venturi M. (2009). "Iodine in evolution of salivary glands and in oral health". Nutrition and Health 20 (2): 119–

134. PMID 19835108 .46. ^ Bahar, G.; Feinmesser, R.; Shpitzer, T.; Popovtzer, A.; Nagler, R.M. (2007). "Salivary analysis in oral cancer patients:

DNA and protein oxidation, reactive nitrogen species, and antioxidant profile". Cancer 109 (1): 54–59.doi:10.1002/cncr.22386 . PMID 17099862 .

47. ^ Banerjee, R.K.; Bose, A.K.; Chakraborty, T.K.; De, S.K.; Datta, A.G. (1985). "Peroxidase-catalysed iodotyrosine formationin dispersed cells of mouse extrathyroidal tissues". J Endocrinol 2: 159–165.

48. ^ Banerjee, R.K.; Datta, A.G. (1986). "Salivary peroxidases". Mol Cell Biochem 70 (1): 21–29. PMID 3520291 .49. ^ Bartelstone, H. J. (1951). "Radioiodine penetration through intact enamel with uptake by bloodstream and thyroid gland".

J Dent Res 5: 728–733.50. ^ Bartelstone, H.J.; Mandel, I.D.; Oshry, E.; Seidlin, S.M. (1947). "Use of radioactive iodine as a tracer in the Study of the

Physiology of teeth". Science 106: 132.51. ^ a b United States National Research Council (2000). Dietary Reference Intakes for Vitamin A, Vitamin K, Arsenic, Boron,

Chromium, Copper, Iodine, Iron, Manganese, Molybdenum, Nickel, Silicon, Vanadium, and Zinc . National AcademiesPress. pp. 258–259.

52. ^ "Sources of iodine" . International Council for the Control of Iodine Deficiency Disorders.53. ^ a b "MedlinePlus Medical Encyclopedia: Iodine in diet" .54. ^ Wu T, Liu GJ, Li P, Clar C (2002). "Iodised salt for preventing iodine deficiency disorders". Cochrane Database Syst Rev

(3): CD003204. doi:10.1002/14651858.CD003204 . PMID 12137681 .55. ^ Felig, Philip; Frohman, Lawrence A. (2001). "Endemic Goiter" . Endocrinology & metabolism. McGraw-Hill Professional.

ISBN 9780070220010.56. ^ "Micronutrients - Iodine, Iron and Vitamin A" . UNICEF.57. ^ "Frequently Asked Questions on Potassium Iodide" . Food and Drug Administration. Retrieved 2009-06-06.58. ^ "Potassium Iodide as a Thyroid Blocking Agent in Radiation Emergencies" . Food and Drug Administration. Retrieved

2009-06-06.59. ^ D. O. Lowe, S. R. Knowles, E. A. Weber, C. J. Railton, and N. H. Shear (2006). "Povidone-iodine-induced burn: case

report and review of the literature". Pharmacotherapy 26: 1641-5. doi:10.1592/phco.26.11.1641 .60. ^ Katelaris, Constance (2009). "'Iodine Allergy' label is misleading". Australian Prescriber, Vol. 32, 125-128. Available at

http://www.australianprescriber.com/magazine/32/5/125/8/

External links"Micronutrient Research for Optimum Health", Linus PaulingInstitute, OSU Oregon State UniversityATSDR - CSEM: Radiation Exposure from Iodine 131 U.S.Department of Health and Human Services (public domain)ChemicalElements.com - Iodinewho.int, WHO Global Database on Iodine DeficiencyOxidizing Agents > IodineWebElements.com – Iodine

Diatomic chemical elements

Hydrogen H2 | Nitrogen N2 | Oxygen O2 | Fluorine F2 | Chlorine Cl2 | Bromine Br2 | Iodine I2 | Astatine At2 |

Iodine compounds

ICN · ICl · ICl3 · IF3 · IF5 · IF7 · I2O5

Periodic table

H He

Li Be B C N O F Ne

Iodine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Iodine[2/19/2010 12:34:01 AM]

This page was last modified on 18 February 2010 at 19:04.Text is available under the Creative Commons Attribution-ShareAlike License; additional terms may apply. See Terms of

Use for details.Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a non-profit organization.

Contact us Privacy policy About Wikipedia Disclaimers

Na Mg Al Si P S Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo

Uue Ubn

Alkali metalsAlkaline earth

metalsLanthanoids Actinoids

Transition

metalsOther metals Metalloids

Other

nonmetalsHalogens Noble gases

Categories: Dietary minerals | Iodine | Halogens | DEA List I chemicals | Biology and pharmacology ofchemical elements | Chemical elements

Astatine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Astatine[2/19/2010 12:36:14 AM]

navigation

Main pageContentsFeatured contentCurrent eventsRandom article

interaction

About WikipediaCommunity portalRecent changesContact WikipediaDonate to WikipediaHelp

toolbox

What links hereRelated changesUpload fileSpecial pagesPrintable versionPermanent linkCite this page

discussion edit this page historyTry Beta Log in / create account

[edit]

polonium ← astatine → radon

I↑

At↓

Uus

Appearance

black solid (presumed)

General properties

Name, symbol,number

astatine, At, 85

Element category halogens

Group, period, block 17, 6, p

Standard atomicweight

(210) g·mol−1

Electron configuration [Xe] 4f14 5d10 6s2 6p5

Electrons per shell 2, 8, 18, 32, 18, 7

(Image)

Physical properties

Phase solid

Melting point 575 K, 302 °C, 576 °F

Boiling point 610 K, 337 °C, 639 °F

Heat of vaporization 40 kJ·mol−1

Vapor pressure

P/Pa 1 10 100 1 k 10 k 100 k

at T/K 361 392 429 475 531 607

Atomic properties

Oxidation states ±1, 3, 5, 7

Electronegativity 2.2 (Pauling scale)

Ionization energies 1st: 890±40 kJ·mol−1

Covalent radius 150 pm

Van der Waals radius 202 pm

Miscellanea

Magnetic ordering no data

Thermal conductivity (300 K) 1.7 W·m−1·K−1

CAS registry number 7440-68-8

Most stable isotopes

Main article: Isotopes of astatine

AstatineFrom Wikipedia, the free encyclopedia

Astatine (pronounced /'æstəti n/ AS-tə-teen or /'æstət n/ AS-tət-in) is a radioactive chemical element with the symbol Atand atomic number 85. It is the heaviest of the discoveredhalogens. Although astatine is produced by radioactive decay innature, due to its short half life it is found only in minuteamounts. Astatine was first produced by Dale R. Corson,Kenneth Ross MacKenzie, and Emilio Segrè in 1940. Threeyears passed before traces of astatine were also found innatural minerals. Until recently most of the physical andchemical characteristics of astatine were inferred fromcomparison with other elements. Some astatine isotopes areused as alpha-particle emitters in science applications, andmedical applications for astatine 211 have been tested. Astatineis currently the rarest naturally-occurring element, with lessthan 30g estimated to be contained in the entire Earth'scrust.[1]

Contents [hide]

1 Characteristics2 History3 Occurrence

3.1 Production4 Compounds5 Isotopes6 Applications7 Precautions8 References9 External links

CharacteristicsThis highly radioactive element has been confirmed by massspectrometers to behave chemically much like other halogens,especially iodine (it would probably accumulate in the thyroidgland like iodine), though astatine is thought to be more metallicthan iodine. Researchers at the Brookhaven NationalLaboratory have performed experiments that have identifiedand measured elementary reactions that involve astatine;[2]

however, chemical research into astatine is limited by itsextreme rarity, which is a consequence of its extremely shorthalf-life. Its most stable isotope has a half-life of around 8.3hours. The final products of the decay of astatine are isotopesof lead. The halogens get darker in color with increasingmolecular weight and atomic number. Thus, following the trend,astatine would be expected to be a nearly black solid, which,when heated, sublimes into a dark, purplish vapor (darker thaniodine). Astatine is expected to form ionic bonds with metals

Periodic table

85At

article

search

languages

Azərbaycan

БеларускаяBosanskiБългарскиCatalàЧ вашлаČeskyCorsuDanskDeutschEestiΕλληνικάEspañolEsperantoEuskara

Français

Astatine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Astatine[2/19/2010 12:36:14 AM]

[edit]

[edit]

[edit]

[edit]

iso NA half-life DM DE (MeV) DP

210At trace 8.1 h ε, β+ 3.981 210Po

α 5.631 206Bi

211At syn 7.2 h

This box: view • talk • edit

such as sodium, like the other halogens, but it can be displacedfrom the salts by lighter, more reactive halogens. Astatine canalso react with hydrogen to form hydrogen astatide, which whendissolved in water, forms the exceptionally strong hydroastaticacid. Astatine is the least reactive of the halogens, being lessreactive than iodine.[3]

HistoryThe existence of "eka-iodine" had been predicted by DmitriMendeleev. Astatine (after Greek αστατος astatos, meaning "unstable") was first synthesized in 1940 by DaleR. Corson, Kenneth Ross MacKenzie, and Emilio Segrè at the University of California, Berkeley bybombarding bismuth with alpha particles.[4]

As the periodic table of elements was long known, several scientists tried to find the element following iodinein the halogen group. The unknown substance was called Eka-iodine before its discovery because the nameof the element was to be suggested by the discoverer. The claimed discovery in 1931 at the AlabamaPolytechnic Institute (now Auburn University) by Fred Allison and associates, led to the spurious name for theelement as alabamine (Ab) for a few years.[5][6][7] This discovery was later shown to be an erroneous one.

The name dakin was proposed for this element in 1937 by the chemist Rajendralal De working in Dhaka,Bangladesh.[8]

The name helvetium was chosen by the Swiss chemist Walter Minder, when he announced the discovery ofelement 85 in 1940, but changed his suggested name to anglohelvetium in 1942.[9]

It took three years before astatine was found as product of the natural decay processes. The short-livedelement was found by the two scientists Berta Karlik and Traude Bernert.[10][11]

OccurrenceAstatine occurs naturally in three natural radioactive decay series, but because of its short half-life is foundonly in minute amounts. Astatine-218 (218At) is found in the uranium series, 216At is in the thorium series,and 215At as well as 219At are in the actinium series[12]. The most long-lived of these naturally-occurringastatine isotopes is 219At with a half-life of 56 seconds.

Astatine is the rarest naturally-occurring element, with the total amount in Earth's crust estimated to be lessthan 1 oz (28 g) at any given time. This amounts to less than one teaspoon of the element. Guinness WorldRecords has dubbed the element the rarest on Earth, stating: "Only around 0.9 oz (25 g) of the elementastatine (At) occurring naturally". Isaac Asimov, in a 1957 essay on large numbers, scientific notation, and thesize of the atom, wrote that in "all of North and South America to a depth of ten miles", the number ofastatine-215 atoms at any time is "only a trillion".[13]

ProductionAstatine is produced by bombarding bismuth with energetic alpha particles to obtain the relatively long-livedisotopes 209At through 211At, which can then be distilled from the target by heating in the presence of air.The energy of the alpha particles determine which isotopes are produced:

Reaction Energy of alpha particle20983Bi + 42α → 211

85At + 2 10n 26 MeV[14]

20983Bi + 42α → 210

85At + 3 10n 40 MeV[14]

20983Bi + 42α → 209

85At + 4 10n 60 MeV[15]

CompoundsMultiple compounds of astatine have been synthesized in microscopic amounts and studied as intensively as

FurlanGaeilgeGaelgGalegoHak-kâ-fa

HrvatskiIdoBahasa IndonesiaÍslenskaItalianoעבריתBasa Jawa

ҚазақшаKiswahiliLatinaLatviešuLëtzebuergeschLietuviųLíguruLojbanMagyar

МонголNederlands

Norsk (bokmål)Norsk (nynorsk)Occitan

PolskiPortuguêsRomânăRuna SimiРусскийSeelterskSicilianuSimple EnglishSlovenčinaSlovenščinaСрпски / SrpskiSrpskohrvatski /СрпскохрватскиSuomiSvenskaTagalog

TürkçeУкраїнська

Uyghurche / Tiếng ViệtWinaray

Astatine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Astatine[2/19/2010 12:36:14 AM]

[edit]

[edit]

[edit]

[edit]

possible before their inevitable radioactive disintegration. The reactions are normally tested with dilutesolutions of astatine mixed with larger amounts of iodine. The iodine acts as a carrier, ensuring that there issufficient material for laboratory techniques such as filtration and precipitation to work.[14][16]

While these compounds are primarily of theoretical interest, they are being studied for potential use in nuclearmedicine.[17] Astatine is expected to form ionic bonds with metals such as sodium, like the other halogens,but it can be displaced from the salts by lighter, more reactive halogens. Astatine can also react withhydrogen to form hydrogen astatide (HAt), which when dissolved in water, forms hydroastatic acid.

Some examples of astatic compounds are:

Sodium astatide (NaAt)Magnesium astatide (MgAt2)Carbon tetraastatide (CAt4)

IsotopesMain article: isotopes of astatine

Astatine has 33 known isotopes, all of which are radioactive; the range of their mass numbers is from 191 to223. There exist also 23 metastable excited states. The longest-lived isotope is 210At, which has a half-life of8.1 hours; the shortest-lived known isotope is 213At, which has a half-life of 125 nanoseconds.[18]

ApplicationsThe least stable isotopes of astatine have no practical applications other than scientific study due to theirextremely short life, but heavier isotopes have medical uses. Astatine-211 is an alpha emitter with a physicalhalf-life of 7.2 h. These features have led to its use in radiation therapy.[19] An investigation of the efficacy ofastatine-211–tellurium colloid for the treatment of experimental malignant ascites in mice reveals that thisalpha-emitting radiocolloid can be curative without causing undue toxicity to normal tissue. By comparison,beta-emitting phosphorus-32 as colloidal chromic phosphate had no antineoplastic activity. The mostcompelling explanation for this striking difference is the dense ionization and short range of action associatedwith alpha-emission. These results have important implications for the development and use of alpha-emittersas radiocolloid therapy for the treatment of human tumors.[20]

PrecautionsSince astatine is extremely radioactive, it should be handled with extreme care. Because of its extreme rarity,it is not likely that the general public will be exposed.

Astatine is a halogen, and standard precautions apply. It is reactive, sharing similar chemical characteristicswith iodine.

There are toxicologic studies of astatine-211 on mice indicating that radioactive poisoning is the major effecton living organisms.[21]

References1. ^ Close, Frank. Particle Physics: A Very Short Introduction. Oxford University Press: New York, 2004. Page 2.2. ^ C. R. Hammond (2004). The Elements, in Handbook of Chemistry and Physics 81st edition. CRC press. ISBN

0849304857.3. ^ Anders, E. (1959). "Technetium and Astatine Chemistry". Annual Review of Nuclear Science 9: 203–220.

doi:10.1146/annurev.ns.09.120159.001223 .4. ^ D. R. Corson, K. R. MacKenzie, and E. Segrè (1940). "Artificially Radioactive Element 85". Phys. Rev. 58: 672–678.

doi:10.1103/PhysRev.58.672 .5. ^ Fred Allison, Edgar J. Murphy, Edna R. Bishop, and Anna L. Sommer (1931). "Evidence of the Detection of Element 85 in

Certain Substances". Phys. Rev. 37: 1178–1180. doi:10.1103/PhysRev.37.1178 .6. ^ "Alabamine & Virginium" . time. Retrieved 2008-07-10.7. ^ Trimble, R. F. (1975). "What happened to alabamine, virginium, and illinium?". J. Chem. Educ. 52: 585.

Yorùbá

Astatine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Astatine[2/19/2010 12:36:14 AM]

[edit]

v • d • e

v • d • e

Wikimedia Commons has mediarelated to: Astatine

Look up astatine in Wiktionary,the free dictionary.

8. ^ 85 Astatine9. ^ Alice Leigh-Smith, Walter Minder (1942). "Experimental Evidence of the Existence of Element 85 in the Thorium Family".

Nature 150: 767–768. doi:10.1038/150767a0 .10. ^ Karlik, Berta; Bernert, Traude (1943). "Eine neue natürliche α-Strahlung". Naturwissenschaften 31 (25–26): 298–299.

doi:10.1007/BF01475613 .11. ^ Karlik, Berta; Bernert, Trande (1943). "Das Element 85 in den natürlichen Zerfallsreihen". Zeitschrift für Physik 123 (1–2):

51–72. doi:10.1007/BF01375144 .12. ^ "astatine (At)" . Encyclopedia Britannica online. Retrieved 2008-06-22.13. ^ http://ia331335.us.archive.org/1/items/onlyatrillion017765mbp/onlyatrillion017765mbp_djvu.txt14. ^ a b c Nefedov, V. D.; Norseev, Yu V; Toropova, M A; Khalkin, Vladimir A (1968). "Astatine". Russ. Chem. Rev. 37: 87–

98. doi:10.1070/RC1968v037n02ABEH001603 .15. ^ Barton, G. W.; Ghiorso, A.; Perlman, I. (1951). "Radioactivity of Astatine Isotopes". Physical Reviews 82 (1): 13–19.

doi:10.1103/PhysRev.82.13 .16. ^ Aten Jun., A. H. W.; Doorgeest, T.; Hollstein, U.; Moeken, H. P. (1952). "Section 5: radiochemical methods. Analytical

chemistry of astatine". Analyst 77: 774–777. doi:10.1039/AN9527700774 .17. ^ Boyd, Jade (1007-08-27). "Nuclear Nanocapsules, The New Cancer Weapon" . Medical News Today. Retrieved 2008-

11-05.18. ^ Audi, Georges (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass

Data Center) 729: 3–128. doi:10.1016/j.nuclphysa.2003.11.001 .19. ^ Wilbur, D. S. (1 October 2001). "Overcoming the Obstacles to Clinical Evaluation of 211At-Labeled

Radiopharmaceuticals." . Journal Nuclear Medicine 42 (10): 1516–1518.20. ^ Bloomer, W. D.; McLaughlin, WH; Neirinckx, RD; Adelstein, SJ; Gordon, PR; Ruth, TJ; Wolf, AP (1981). "Astatine-211--

tellurium radiocolloid cures experimental malignant ascites" . Science 212 (4492): 340–341.doi:10.1126/science.7209534 . PMID 7209534 .

21. ^ Cobb, L.M.; Harrison, A.; Butler, S.A. (1988). "Toxicity of Astatine-211 in the Mouse". Human & Experimental Toxicology7: 529. doi:10.1177/096032718800700602 .

External linksWebElements.com - AstatineDoc Brown's Chemistry Clinic - Group 7 The HalogensChemistry in its element podcast (MP3) from the RoyalSociety of Chemistry's Chemistry World: Astatine

Diatomic chemical elements

Hydrogen H2 | Nitrogen N2 | Oxygen O2 | Fluorine F2 | Chlorine Cl2 | Bromine Br2 | Iodine I2 | Astatine At2 |

Periodic table

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo

Uue Ubn

Alkali metalsAlkaline earth

metalsLanthanoids Actinoids

Transition

metalsOther metals Metalloids

Other

nonmetalsHalogens Noble gases

Categories: Chemical elements | Halogens | Astatine

Astatine - Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Astatine[2/19/2010 12:36:14 AM]

This page was last modified on 17 February 2010 at 18:00.Text is available under the Creative Commons Attribution-ShareAlike License; additional terms may apply. See Terms of

Use for details.Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a non-profit organization.

Contact us Privacy policy About Wikipedia Disclaimers