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{{About|the chemical element|the bleach|Sodium hypochlorite|the upcoming film|Chlorine (film)}}
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{{Infobox chlorine}}
'''Chlorine''' ({{IPAc-en|icon|ˈ|k|l|ɔər|iː|n}} {{Respell|KLOHR|een}}; from the [[Greek language|Greek]] word 'χλωρóς' (''khlôros'', meaning 'pale green')) is the  [[chemical element]] with [[atomic number]] 17 and symbol '''Cl'''. It is a [[halogen]], found in the [[periodic table]] in [[group 17]] (formerly VII, VIIa, or VIIb). As the [[chloride]] ion, which is part of [[common salt]] and other compounds, it is abundant in nature and necessary to most forms of life, including [[human]]s. In its elemental form (Cl<sub>2</sub> or "dichlorine") under [[standard conditions]], chlorine is a powerful [[oxidizing agent|oxidant]] and is used in [[Bleach (chemical)|bleaching]] and disinfectants, as well as an essential reagent in the chemical industry. As a common disinfectant, chlorine compounds are used in [[swimming pool]]s to keep them clean and [[swimming pool sanitation|sanitary]]. In the [[upper atmosphere]], chlorine-containing molecules such as [[chlorofluorocarbons]] have been implicated in the [[ozone depletion|destruction]] of the [[ozone layer]].

== Characteristics ==
===Physical characteristics===
At [[standard temperature and pressure]], two chlorine atoms form the [[diatomic molecule]] {{chem|Cl|2}}. This is a pale yellow-green gas that has its distinctive strong smell, the smell of bleach. The bonding between the two atoms is relatively weak (only 242.580 ±0.004&nbsp;kJ/mol), which makes the {{chem|Cl|2}} molecule highly reactive.

===Chemical characteristics===
Along with [[fluorine]], [[bromine]], [[iodine]], and [[astatine]], chlorine is a member of the [[halogen]] series that forms the group 17 of the periodic table. Chlorine forms compounds with almost all of the elements to give compounds that are usually called chlorides.  Chlorine gas reacts with most organic compounds, and will even sluggishly support the combustion of [[hydrocarbon]]s.<ref name=CRC/> 

=====Hydrolysis=====
At 10&nbsp;°[[Celsius|C]] and atmospheric pressure, one [[litre|liter]] of [[water]] dissolves 3.10&nbsp;L of gaseous chlorine,<ref name="webelement">{{cite web|url=http://www.webelements.com/webelements/elements/text/Cl/index.html|title=WebElements.com – Chlorine|accessdate=2007-03-17|publisher=Mark Winter [The University of Sheffield and WebElements Ltd, UK]}}</ref> Solutions of chlorine in water contain chlorine (Cl<sub>2</sub>), hydrochloric acid, and hypochlorous acid:
:Cl<sub>2</sub>  +  H<sub>2</sub>O   <math>\overrightarrow{\leftarrow}</math>   HCl + HClO
This conversion to the right is called [[disproportionation]], because the ingredient chlorine both increases and decreases in formal oxidation state.  The solubility of chlorine in water is increased if the water contains dissolved alkali hydroxide, and in this way, [[chlorine bleach]] is produced.
:Cl<sub>2</sub> + 2 OH<sup>-</sup> → ClO<sup>-</sup> + Cl<sup>-</sup> + H<sub>2</sub>O
Chlorine gas only exists in a neutral or acidic solution.

==== Compounds====
{{See also|Category:Chlorine compounds}}
Chlorine exists in all odd numbered [[oxidation states]] from −1 to +7, as well as the elemental state of zero (see Table). Progressing through the states, hydrochloric acid can be oxidized using manganese dioxide, or [[hydrogen chloride]] gas oxidized catalytically by air to form elemental chlorine gas. 

{| class="wikitable"
! Oxidation<br />state !! Name !! Formula !! Illustrative compounds
|-
| align="center"|−1 || chlorides || align="center"|Cl<sup>−</sup> || ionic chlorides, organic chlorides, hydrochloric acid
|-
| align="center"|0 || chlorine || align="center"|Cl<sub>2</sub> || elemental chlorine
|-
| align="center"|+1 || hypochlorites || align="center"|ClO<sup>−</sup> || [[sodium hypochlorite]], [[calcium hypochlorite]]
|-
| align="center"|+3 || [[chlorite]]s || align="center"|{{chem|ClO|2|-}} || [[sodium chlorite]]
|-
| align="center"|+4 || [[chlorine dioxide]] || align="center"|{{chem|ClO|2}} ||
|-
| align="center"|+5 || [[chlorate]]s || align="center"|{{chem|ClO|3|-}} || [[sodium chlorate]], [[potassium chlorate]], [[chloric acid]]
|-
| align="center"|+7 || [[perchlorate]]s || align="center"|{{chem|ClO|4|-}} ||  [[perchloric acid]], perchlorate salts such as [[magnesium perchlorate]], [[dichlorine heptoxide]]
|}
=====Interhalogen compounds=====
Chlorine oxidizes bromide and iodide salts to bromine and iodine, respectively.  But it cannot oxidize fluoride to fluorine.  It makes are variety of "interhalogen compounds" such as the chlorine fluorides, [[chlorine monofluoride]] (ClF), [[chlorine trifluoride]] (ClF<sub>3</sub>), [[chlorine pentafluoride]] (ClF<sub>5</sub>). Chlorides of bromine and iodine are also known.

=====Organo chlorine compounds=====
{{main|organochloride}}
Chlorine is used extensively in [[organic chemistry|organic]] in [[substitution (chemistry)|substitution]] and addition reactions. Chlorine often imparts many desired properties to an [[organic compound]], in part due to its electronegativity. Organochlorine compounds are also serious pollutants, either as side products of industrial processes or as persistent pesticidess.

Many important industrial products are produced via organochlorine intermediates. Examples include [[polycarbonate]]s, [[polyurethane]]s, [[silicone]]s, [[polytetrafluoroethylene]], [[carboxymethyl cellulose]], and [[propylene oxide]].  Like the other halogens, chlorine participates in [[free-radical substitution reaction]]s with hydrogen-containing organic compounds. When applied to organic substrates, reaction is often—but not invariably—non-regioselective, and, hence, may result in a mixture of isomeric products. It is often difficult to control the degree of substitution as well, so multiple substitutions 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]], [[methylene chloride]], [[chloroform]], and [[carbon tetrachloride]] from methane, [[allyl chloride]] from propylene, and [[trichloroethylene]], and [[tetrachloroethylene]] from [[ethylene dichloride|1,2-dichloroethane]].

Like the other halides, chlorine undergoes electrophilic additions reactions, the most notable one being the chlorination of alkenes and aromatic compounds with a Lewis acid catalyst. Organic chlorine compounds tend to be less reactive in nucleophilic substitution reactions than the corresponding bromine or iodine derivatives, but they tend to be cheaper. They may be activated for reaction by substituting with a tosylate group, or by the use of a catalytic amount of [[sodium iodide]].

=====Chlorides=====
{{main|chloride}}
Chlorine combines with almost all elements to give chlorides.  Compounds with [[oxygen]], [[nitrogen]], [[xenon]], and [[krypton]] are known, but do not form by direct reaction of the elements.<ref>{{cite book|title=Merck Index of Chemicals and Drugs, 9th ed.|publisher=Merck & Co.|location = Rahway, N.J.|year=1976|isbn=0911910263|editor=Windholz, Martha ''et al.''}}</ref> Chloride is one of the most common anions in nature. [[Hydrogen chloride]] and its aqueous solution, [[hydrochloric acid]], are produced on megaton scale annually both as valued intermediates but sometimes as undesirable pollutants.

=====Chlorine oxides=====
Chlorine forms a variety of oxides: [[chlorine dioxide]] (ClO<sub>2</sub>), [[dichlorine monoxide]] (Cl<sub>2</sub>O), [[dichlorine heptoxide]] (Cl<sub>2</sub>O<sub>7</sub>).  The anionic derivatives of these same oxides are also well known including [[chlorate]] ({{chem|ClO|3|-}}), [[chlorite]] ({{chem|ClO|2|-}}), [[hypochlorite]] (ClO<sup>−</sup>), and [[perchlorate]] ({{chem|ClO|4|-}}), and [[chloramine]] (NH<sub>2</sub>Cl).<ref name="Euro Chlor6">{{cite web|url=http://www.eurochlor.org/index.asp?page=678|title=Chlorine compounds of the month|publisher=Euro Chlor|accessdate=2007-08-29}}</ref> The acid derivatives of these anions are [[hypochlorous acid]] (HOCl), [[chloric acid]] (HClO<sub>3</sub>) and [[perchloric acid]] (HClO<sub>4</sub>).

In hot concentrated alkali solution [[hypochlorite]] disproportionates:
:2 ClO<sup>−</sup> → Cl<sup>−</sup> + {{chem|ClO|2|-}}
:ClO<sup>−</sup> + {{chem|ClO|2|-}} → Cl<sup>−</sup> + {{chem|ClO|3|-}}
[[Sodium chlorate]] and [[potassium chlorate]] can be crystallized from solutions formed by the above reactions. If their crystals are heated, they undergo a further, final disproportionation:
:4 {{chem|ClO|3|-}} → Cl<sup>−</sup> + 3 {{chem|ClO|4|-}}

This same progression from chloride to perchlorate can be accomplished by electrolysis. The anode reaction progression is:<ref name="c_w">{{cite book|author=Cotton, F. Albert and Wilkinson, Geoffrey|title=Advanced Inorganic Chemistry, 2nd ed. |publisher=John Wiley & sons| page= 568|year=1966|isbn=}}</ref> 

:{| class="wikitable"
! Reaction !! Electrode<br />potential
|-
| align="center"|Cl<sup>−</sup> + 2 OH<sup>−</sup> → ClO<sup>−</sup> + H<sub>2</sub>O + 2 e<sup>−</sup> || +0.89 volts
|-
| align="center"|ClO<sup>−</sup> + 2 OH<sup>−</sup> → {{chem|ClO|2|-}} + H<sub>2</sub>O + 2 e<sup>−</sup> || +0.67 volts
|-
| align="center"|{{chem|ClO|2|-}} + 2 OH<sup>−</sup> → {{chem|ClO|3|-}} + H<sub>2</sub>O + 2 e<sup>−</sup> || +0.33 volts
|-
| align="center"|{{chem|ClO|3|-}} + 2 OH<sup>−</sup> → {{chem|ClO|4|-}} + H<sub>2</sub>O + 2 e<sup>−</sup> || +0.35 volts
|}
Each step is accompanied at the cathode by
:2 H<sub>2</sub>O  +  2 e<sup>−</sup>   →   2 OH<sup>−</sup>  +  H<sub>2</sub>  (−0.83 volts)

===Occurrence===
{{See also category|Halide minerals}}
In nature, chlorine is found primarily as the [[chloride ion]], a component of the salt that is deposited in the earth or dissolved in the [[ocean]]s — about 1.9% of the mass of seawater is chloride ions. Even higher concentrations of chloride are found in the [[Dead Sea]] and in underground [[brine]] deposits. Most chloride salts are soluble in water, thus, chloride-containing minerals are usually only found in abundance in dry climates or deep underground. Common chloride minerals include ''[[halite]]'' ([[sodium chloride]]), ''[[sylvite]]'' ([[potassium chloride]]), and ''[[carnallite]]'' (potassium magnesium chloride hexahydrate). Over 2000 naturally occurring organic chlorine compounds are known.<ref name="Euro Chlor">{{cite web|url=http://www.eurochlor.org/upload/documents/document236.pdf|title=Risk assessment and the cycling of natural organochlorines|publisher=Euro Chlor|accessdate=2007-08-12}}</ref>

In the interstellar medium, chlorine is produced in [[supernova]]e via the [[r-process]].<ref>{{cite journal|author= A.G.W. Cameron|title=Stellar Evolution, Nuclear Astrophysics, and Nucleogenesis| journal=CRL-41 | month=June| year=1957}} "http://www.fas.org/sgp/eprint/CRL-41.pdf"</ref>

=== Isotopes ===
{{Main|Isotopes of chlorine}}
Chlorine has a wide range of [[Isotopes of chlorine|isotopes]], the two principal stable [[isotope]]s being <sup>35</sup>Cl (75.77%) and <sup>37</sup>Cl (24.23%); they give chlorine [[atom]]s an apparent atomic weight of 35.4527&nbsp;g/mol.

Trace amounts of [[radioactive]] [[Chlorine-36|<sup>36</sup>Cl]] exist in the environment, in a ratio of about 7x10<sup>−13</sup> to 1 with stable isotopes. <sup>36</sup>Cl is produced in the atmosphere by [[spallation]] of <sup>36</sup>[[argon|Ar]] by interactions with [[cosmic ray]] [[proton]]s. In the subsurface environment, <sup>36</sup>Cl is generated primarily as a result of [[neutron capture]] by <sup>35</sup>Cl or [[muon capture]] by <sup>40</sup>[[calcium|Ca]]. <sup>36</sup>Cl decays to <sup>36</sup>[[sulfur|S]] and to <sup>36</sup>[[argon|Ar]], with a combined [[half-life]] of 308,000 years. The half-life of this [[hydrophilic]] nonreactive isotope makes it suitable for [[geologic dating]] in the range of 60,000 to 1&nbsp;million years. Additionally, large amounts of <sup>36</sup>Cl were produced by irradiation of [[seawater]] during atmospheric detonations of [[nuclear weapon]]s between 1952 and 1958. The residence time of <sup>36</sup>Cl in the atmosphere is about 1 week. Thus, as an event marker of 1950s water in [[soil]] and [[ground water]], <sup>36</sup>Cl is also useful for dating waters less than 50 years before the present. <sup>36</sup>Cl has seen use in other areas of the geological sciences, including dating ice and sediments.

== History ==
[[File:Chlorine liquid in an ampoule.jpg|thumb|left|upright|Liquid chlorine]]
The most common compound of chlorine, sodium chloride, has been known since ancient times; archaeologists have found evidence that rock salt was used as early as 3000 BC and brine as early as 6000 BC.<ref>{{cite web |url = http://antiquity.ac.uk/ProjGall/weller/|title = The earliest salt production in the world: an early Neolithic exploitation in Poiana Slatinei-Lunca, Romania |accessdate=2008-07-10}}</ref> 

Elemental chlorine was first prepared and studied in 1774 by Swedish chemist [[Carl Wilhelm Scheele]], and, therefore, he is credited for its discovery.<ref name=krogt>{{cite web|url=http://elements.vanderkrogt.net/elem/cl.html |title=17 Chlorine |publisher=Elements.vanderkrogt.net |accessdate=2008-09-12}}</ref> He called it "''dephlogisticated muriatic acid air''" since it is a gas (then called "airs") and it came from [[hydrochloric acid]] (then known as "muriatic acid").<ref name=krogt/> However, he failed to establish chlorine as an element, mistakenly thinking that it was the [[oxide]] obtained from the hydrochloric acid (see [[phlogiston theory]]).<ref name=krogt/> He named the new element within this oxide as ''muriaticum''.<ref name=krogt/> Regardless of what he thought, Scheele did isolate chlorine by reacting [[Manganese dioxide|MnO<sub>2</sub>]] (as the mineral [[pyrolusite]]) with HCl:

:4 HCl + MnO<sub>2</sub> →  MnCl<sub>2</sub> + 2 H<sub>2</sub>O + Cl<sub>2</sub>

Scheele observed several of the properties of chlorine: the bleaching effect on [[Litmus test (chemistry)|litmus]], the deadly effect on insects, the yellow green color, and the smell similar to [[aqua regia]].

At the time, common chemical theory was: Any acid is a compound that contains oxygen (still sounding in the German and Dutch names of [[oxygen]]: ''sauerstoff'' or ''zuurstof'', both translating into English as ''acid stuff''), so a number of chemists, including [[Claude Berthollet]], suggested that Scheele's ''dephlogisticated muriatic acid air''  must be a combination of oxygen and the yet undiscovered element, ''muriaticum''.<ref>{{Cite book|url=http://books.google.com/?id=34KwmkU4LG0C&pg=PA158|page=158|title=The development of modern chemistry|author=Ihde, Aaron John|publisher=Courier Dover Publications|year=1984|isbn=0486642356}}</ref>

In 1809, [[Joseph Louis Gay-Lussac]] and [[Louis-Jacques Thénard]] tried to decompose ''dephlogisticated muriatic acid air'' by reacting it with charcoal to release the free element ''muriaticum'' (and carbon dioxide).<ref name="krogt"/> They did not succeed and published a report in which they considered the possibility that ''dephlogisticated muriatic acid air'' is an element, but were not convinced.<ref>{{cite journal|title = On the nature and the properties of muriatic acid and of oxygenated muriatic acid|journal = Mémoires de Physique et de Chimie de la Société d'Arcueil|volume = 2|pages = 339–358|year = 1809|url = http://web.lemoyne.edu/~giunta/thenard.html|text = "In fact, oxygenated muriatic acid is not decomposed by charcoal, and it might be supposed, from this fact and those that are communicated in this Memoir, that this gas is a simple body. The phenomena that it presents can be explained well enough on this hypothesis; we shall not seek to defend it, however, as it appears to us that they are still better explained by regarding oxygenated muriatic acid as a compound body."|author = Gay-Lussac, Joseph Louis; Thénard, Louis-Jacques}}</ref>

In 1810, [[Sir Humphry Davy]] tried the same experiment again, and concluded that it is an element, and not a compound.<ref name=krogt/> He named this new element as chlorine, from the Greek word χλωρος (''chlōros''), meaning green-yellow.<ref>{{cite journal|author = Sir Humphry Davy |title = On a Combination of Oxymuriatic Gas and Oxygene Gas|journal = Philosophical Transactions of the Royal Society|volume = 101|pages = 155–162|year = 1811|doi = 10.1098/rstl.1811.0008|url=http://www.chemteam.info/Chem-History/Davy-Chlorine-1811.html }}</ref> The name halogen, meaning salt producer, was originally defined for chlorine (in 1811 by [[Johann Salomo Christoph Schweigger]]), and it was later applied to the rest of the elements in this family.<ref>{{Cite journal|title=J. S. C. Schweigger: His Romanticism and His Crystal Electrical Theory of Matter|author=Snelders, H. A. M.|journal=Isis|volume=62|issue=3 |year=1971|pages=328|url=http://www.jstor.org/pss/229946|doi=10.1086/350763}}</ref> In 1823, [[Michael Faraday]] liquefied chlorine for the first time.<ref>{{cite web |url = http://www.aps.org/publications/apsnews/200108/history.cfm |title =
This Month in Physics History September 4, 1821 and August 29, 1831: Faraday and Electromagnetism|accessdate = 2010-05-08}}</ref><ref>{{cite web |url = http://www-history.mcs.st-andrews.ac.uk/Biographies/Faraday.html |title =
Michael Faraday|accessdate = 2010-05-08}}</ref>

Chlorine was first used to bleach textiles in 1785.<ref>{{cite web |url = http://members.aol.com/manbio999/chlorine.htm|title = History of Chlorine |accessdate = 2008-07-10}}</ref>  In 1826, [[silver chloride]] was used to produce photographic images for the first time.<ref name=history>{{ cite web|url = http://ocw.mit.edu/NR/rdonlyres/Earth--Atmospheric--and-Planetary-Sciences/12-091January--IAP--2006/0EF9264B-3205-44A3-8306-8E8364917DF0/0/brazin.pdf|title = Chlorine & its Consequences|author = Jacqueline Brazin|accessdate=2008-07-10|archiveurl = http://web.archive.org/web/20060918145109/http://ocw.mit.edu/NR/rdonlyres/Earth--Atmospheric--and-Planetary-Sciences/12-091January--IAP--2006/0EF9264B-3205-44A3-8306-8E8364917DF0/0/brazin.pdf |archivedate = September 18, 2006|deadurl=yes}}</ref> Chloroform was first used as an anesthetic in 1847.<ref name=history/> A chlorine solution in lime-water ([[hypochlorite]]) was first used as a [[germicide]] to prevent the spread of [[puerperal fever]] in the maternity wards of [[Vienna General Hospital]] in Austria in 1847,<ref>{{cite web|url = http://www.americanchemistry.com/s_chlorine/sec_content.asp?CID=1166&DID=4476&CTYPEID=109|title = Chlorine Story|publisher = americanchemistry|accessdate=2008-07-10}}</ref> and in 1850 by [[John Snow (physician)|John Snow]] to disinfect the water supply in London after an outbreak of [[cholera]]. The US [[Department of Treasury]] called for all water to be disinfected with chlorine by 1918.<ref name=history/> [[Polyvinylchloride]] (PVC) was invented in 1912, initially without a purpose.<ref name=history/> Chlorine gas was first introduced as a weapon on April 22, 1915 at [[Ypres]] by the [[German Army]],<ref>{{cite web |url = http://www.drcordas.com/education/weaponsmassd/Chlorine.pdf|title = Chlorine - History|accessdate=2008-07-10}}</ref><ref>{{cite web |title = Weaponry: Use of Chlorine Gas Cylinders in World War I|url = http://www.historynet.com/weaponry-use-of-chlorine-gas-cylinders-in-world-war-i.htm|publisher = historynet.com|accessdate = 2008-07-10}}</ref> and the results of this weapon were disastrous because [[gas mask]]s had not yet been invented.<!-- was it the first gas used as a weapon? --><!-- pvc uses -->

== Production ==
{{Main|Chlorine production}}
In industry, 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 following [[chemical equation]]:

:2 NaCl + 2 H<sub>2</sub>O → Cl<sub>2</sub> + H<sub>2</sub> + 2 NaOH
The electrolysis of chloride solutions all proceed according to the following equations:

:Cathode: 2 H<sup>+</sup> (aq) + 2 e<sup>−</sup> → H<sub>2</sub> (g)
:Anode: 2 Cl<sup>−</sup> (aq) → Cl<sub>2</sub> (g) + 2 e<sup>−</sup>
Overall process: 2 NaCl (or KCl) + 2 H<sub>2</sub>O → Cl<sub>2</sub> + H<sub>2</sub> + 2 NaOH (or KOH)
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 hydrogen formed at the cathode.<ref name="Euro Chlor3">{{cite web|url=http://www.eurochlor.org/animations/diaphragm-cell.asp|title=Diaphragm cell|publisher=Euro Chlor|accessdate=2007-08-15}}</ref> The salt solution (brine) is continuously fed to the anode compartment and flows through the diaphragm to the cathode compartment, where the [[Causticity|caustic]] [[alkali]] is produced and the brine is partially depleted. Diaphragm methods produce dilute and slightly impure alkali but they are not burdened with the problem of preventing mercury discharge into the environment and they more energy efficient.  Membrane cell electrolysis employ [[Semipermeable membrane|permeable membrane]] as an [[ion exchange]]r. Saturated sodium (or potassium) chloride solution is passed through the anode compartment, leaving at a lower [[concentration]].<ref name="Euro Chlor4">{{cite web|url=http://www.eurochlor.org/animations/membrane-cell.asp|title=Membrane cell|publisher=Euro Chlor|accessdate=2007-08-15}}</ref> 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.
[[Image:Liquid chlorine in flask.jpg|upright|thumb|thumb|Liquid Chlorine Analysis]]

===Laboratory methods===
Small amounts of chlorine gas can be made in the laboratory by combining hydrochloric acid  and manganese dioxide.  Alternatively a strong acid such as sulfuric acid or hydrochloric acid reacts with [[sodium hypochlorite]] or [[sodium chlorate]] solution to release chlorine gas.  In the home, accidents occur when hypochlorite bleach solutions are combined with certain acidic drain-cleaners.

== Applications ==
=== Production of industrial and consumer products ===
Chlorine's principal applications are in the production of a wide range of industrial and consumer products.<ref name="Euro Chlor5">{{cite web|url=http://www.eurochlor.org/uses|title=Uses|publisher=Euro Chlor|accessdate=2007-08-20}}</ref><ref name="Chlorine Tree">{{cite web|url=http://www.chlorinetree.org|title=Chlorine Tree|publisher=Chlorine Tree|accessdate=2007-08-20}}</ref> 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 disinfection ===
Chlorine is an important chemical for [[water purification]] (such as water treatment plants), in [[disinfectant]]s, and in [[bleach]]. Chlorine in water is more than three times more effective as a disinfectant against ''[[Escherichia coli]]'' than an equivalent concentration of [[bromine]], and is more than six times more effective than an equivalent concentration of [[iodine]].<ref>{{cite journal|author=Koski TA, Stuart LS, Ortenzio LF|title=Comparison of chlorine, bromine, iodine as disinfectants for swimming pool water|journal = Applied Microbiology|volume = 14|issue = 2|year = 1966|pages = 276–279|url = http://www.pubmedcentral.nih.gov/articlerender.fcgi?tool=pubmed&pubmedid=4959984 |pmid=4959984|pmc=546668 }}</ref>

Chlorine is usually used (in the form of [[hypochlorous acid]]) to kill [[bacterium|bacteria]] and other microbes in [[drinking water]] supplies and public swimming pools. In most private swimming pools, chlorine itself is not used, but rather [[sodium hypochlorite]], formed from chlorine and [[sodium hydroxide]], or solid tablets of chlorinated isocyanurates. Even small water supplies are now routinely chlorinated.<ref name=CRC>{{cite book| author = Hammond, C. R.|title = The Elements, in Handbook of Chemistry and Physics 81st edition| publisher =CRC press| year = 2000| isbn = 0849304814}}</ref> (''See also'' [[chlorination]])

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

=== Use as a weapon ===
* '''World War I '''
{{Main|Poison gas in World War I}}

Chlorine gas, also known as '''bertholite''', was first [[chemical warfare|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 mixture between pepper and pineapple. It also tasted metallic and stung the back of the throat and chest. Chlorine can react with water in the mucosa of the lungs to form [[hydrochloric acid]], an irritant that can be lethal. The damage done by chlorine gas can be prevented by a gas mask, or other filtration method, which makes the overall chance of death by chlorine gas much lower than those of other chemical weapons. It was pioneered by a German scientist later to be a Nobel laureate, [[Fritz Haber]] of the Kaiser Wilhelm Institute in Berlin, in collaboration with the German chemical conglomerate [[IG Farben]], who developed methods for discharging chlorine gas against an [[trench|entrenched]] enemy. It is alleged that Haber's role in the use of chlorine as a deadly weapon drove his wife, [[Clara Immerwahr]], to suicide.{{Citation needed|date=August 2010}}<!--The Immerwahr article says there is source for the motive of here suicide--> After its first use, chlorine was utilized by both sides as a chemical weapon, but it was soon replaced by the more deadly gases [[phosgene]] and [[mustard gas]].<ref name="First World War">{{cite web|url=http://www.firstworldwar.com/weaponry/gas.htm|title=Weapons of War: Poison Gas|publisher=First World War.com|accessdate=2007-08-12}}</ref>

* '''Iraq War '''
{{Main|2007 chlorine bombings in Iraq}}

Chlorine gas has also been used by insurgents against the local population and coalition forces in the [[Iraq War]] in the form of [[chlorine bomb]]s. On March 17, 2007, for example, three chlorine filled trucks were detonated in the Anbar province killing two and sickening over 350.<ref name="cnnchlorine">{{cite news|url=http://www.cnn.com/2007/WORLD/meast/03/17/iraq.main/index.html|last=Mahdi|first=Basim|publisher=CNN|date=2007-03-17|accessdate=2007-03-17|title= Iraq gas attack makes hundreds ill}}</ref> Other chlorine bomb attacks resulted in higher death tolls, with more than 30 deaths on two separate occasions.<ref name="Chlorine bomb BBC">{{cite news|url= http://news.bbc.co.uk/2/hi/middle_east/6660585.stm|last = |first= |publisher = BBC News|date=2007-05-17|accessdate = 2007-05-17|title = 'Chlorine bomb' hits Iraq village}}</ref> Most of the deaths were caused by the force of the explosions rather than the effects of chlorine, since the toxic gas is readily dispersed and diluted in the atmosphere by the blast. The Iraqi authorities have tightened up security for chlorine, which is essential for providing safe drinking water for the population.

=== Chlorine cracking ===
[[Image:Chlorine attack1.jpg|thumb|right|Chlorine "attack" on an acetal resin plumbing joint.]]
The element is widely used for purifying water owing to its powerful oxidizing properties, especially potable water supplies and water used in swimming pools. Several catastrophic collapses of swimming pool ceilings have occurred owing to [[stress corrosion cracking]] of [[stainless steel]] rods used to suspend them.<ref>{{Cite book|url=http://books.google.com/?id=cEmq232h1zcC&pg=PA148|page=148|title=Corrosion of steel in concrete: prevention, diagnosis, repair|author=Bertolini, Luca; Elsener, Bernhard; Pedeferri, Pietro; Polder, Rob B.|publisher=Wiley-VCH|year=2004|isbn=3527308008}}</ref> Some [[polymer]]s are also sensitive to attack, including [[acetal resin]] and [[polybutene]]. Both materials were used in hot and cold water domestic supplies, and [[stress corrosion cracking]] caused widespread failures in the USA in the 1980s and 1990s. One example shows an acetal joint in a water supply system, which, when it fractured, caused substantial physical damage to computers in the labs below the supply. The cracks started at [[injection molding]] defects in the joint and grew slowly until finally triggered. The fracture surface shows iron and calcium salts that were deposited in the leaking joint from the water supply before failure.

=== Other uses ===
Chlorine is used in the manufacture of numerous organic chlorine compounds, the most significant of which in terms 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]], [[vinylidene chloride]], [[trichloroethylene]], [[perchloroethylene]], [[allyl chloride]], [[epichlorohydrin]], [[chlorobenzene]], dichlorobenzenes, and [[trichlorobenzene]]s.

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

== Mapping of industrial releases in the United States ==

One tool that maps the most recent release information of chlorine [http://toxmap.nlm.nih.gov/toxmap/tri/mapIt.do?chemicalName=chlorine] to particular locations in the United States<ref>{{cite web|url=http://toxmap.nlm.nih.gov/toxmap/|title=TRI Releases Map |publisher=Toxmap.nlm.nih.gov |accessdate=2010-03-23}}</ref> and also provides additional information about such releases is [[TOXMAP]]. TOXMAP is a Geographic Information System (GIS) from the Division of Specialized Information Services of the [[United States National Library of Medicine]] (NLM) that uses maps of the United States to help users visually explore data from the [[United States Environmental Protection Agency]]'s (EPA) [[Toxics Release Inventory]] and [[Superfund Basic Research Program]]s. TOXMAP is a resource funded by the US Federal Government. TOXMAP's chemical and environmental health information is taken from NLM's Toxicology Data Network (TOXNET)<ref>[http://toxnet.nlm.nih.gov TOXNET - Databases on toxicology, hazardous chemicals, environmental health, and toxic releases]</ref> and [[PubMed]], and from other authoritative sources.

== Health effects ==
[[Image:Skull and crossbones.svg|left|40px]]
{| class="wikitable" style="float:right; margin:0em 0em 0em 1em"
! style="background:#FF9900"|[[NFPA 704]]
|-
| align="left"|{{NFPA 704|Health = 3|Flammability = 0|Reactivity = 0|Other = OX}}
|-
|}
Chlorine is a toxic gas that irritates the respiratory system. Because it is heavier than air, it tends to accumulate at the bottom of poorly ventilated spaces. Chlorine gas is a strong oxidizer, which may react with flammable materials.<ref name="msds-cl">{{cite web|url=http://www.westlake.com/datasheets/MSDS_Chlorine.pdf |title=Chlorine MSDS|date=October 23, 1997 (Revised November 1999}}</ref>

Chlorine is detectable in concentrations of as low as 0.2 ppm. Coughing and vomiting may occur at 30 ppm and lung damage at 60&nbsp;ppm. About 1000&nbsp;ppm can be fatal after a few deep breaths of the gas.<ref name="webelement"/>  Breathing lower concentrations can aggravate the respiratory system, and exposure to the gas can irritate the eyes.<ref name=tox>{{cite journal|journal = Environmental Research|volume = 85|issue = 2|year = 2001 |doi = 10.1006/enrs.2000.4110 |title = The Toxicology of Chlorine|author = Winder, Chris|pages = 105–114|pmid = 11161660}}</ref>
The toxicity of chlorine comes from its oxidizing power. When chlorine is inhaled at concentrations above 30&nbsp;ppm, it begins 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, the reaction of chlorine with water is not a major concern for human health. However, other materials present in the water may generate disinfection by-products that can damage human health.<ref>{{cite web|url=http://www.aces.uiuc.edu/news/stories/news4724.html|title=What's in your Water?: Disinfectants Create Toxic By-products|date=2009-03-31|work=ACES News|publisher=College of Agricultural, Consumer and Environmental Sciences - University of Illinois at Urbana-Champaign|accessdate=2009-03-31}}</ref><ref>{{cite pmid|17980649}}</ref>

== See also ==
* [[Polymer degradation]]

== References ==
{{Reflist|colwidth=30em}}

== External links ==
{{Commons}}
{{Wiktionary|chlorine}}
*[[Agency for Toxic Substances and Disease Registry]]: [http://www.atsdr.cdc.gov/substances/toxsubstance.asp?toxid=36 Chlorine]
* [http://electrochem.cwru.edu/encycl/art-b01-brine.htm Electrolytic production]
* [http://www.amazingrust.com/Experiments/how_to/Liquid_Cl2.html Production and liquefaction of chlorine]
* [http://www.oceana.org/chlorine Chlorine Production Using Mercury, Environmental Considerations and Alternatives]
* [http://www.npi.gov.au/database/substance-info/profiles/20.html National Pollutant Inventory - Chlorine]
* [http://www.cdc.gov/niosh/topics/chlorine/ National Institute for Occupational Safety and Health - Chlorine Page]
* [http://www.webelements.com/webelements/elements/text/Cl/index.html WebElements.com — Chlorine]
* [http://www.chlorineinstitute.org/ Chlorine Institute] - Trade association and lobby group representing the interests of the chlorine industry
* [http://www.eurochlor.org/ Chlorine Online] - Chlorine Online is an information resource produced by Eurochlor - the business association of the European chlor-alkali industry
{{Chlorine compounds}}
{{diatomicelements}}
{{U.S. chemical weapons}}
{{Chemical warfare}}
{{E number infobox 920-929}}
{{Compact periodic table}}

{{Good article}}

[[Category:Chemical elements]]
[[Category:Chlorine|*]]
[[Category:Halogens]]
[[Category:Hazardous air pollutants]]
[[Category:Occupational safety and health]]
[[Category:Pulmonary agents]]
[[Category:Swimming pool equipment]]
[[Category:Oxidizing agents]]

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