View source for Chemistry:Chlorine

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{{Infobox chlorine|engvar=en-GB}}

'''Chlorine''' is a [[Chemistry:Chemical element|chemical element]] with the [[Chemistry:Symbol|symbol]] '''Cl''' and [[Physics:Atomic number|atomic number]] 17. The second-lightest of the [[Chemistry:Halogen|halogen]]s, it appears between [[Chemistry:Fluorine|fluorine]] and [[Chemistry:Bromine|bromine]] in the periodic table and its properties are mostly intermediate between them. Chlorine is a yellow-green gas at room temperature. It is an extremely reactive element and a strong [[Chemistry:Oxidizing agent|oxidising agent]]: among the elements, it has the highest [[Physics:Electron affinity|electron affinity]] and the third-highest [[Chemistry:Electronegativity|electronegativity]] on the revised [[Chemistry:Electronegativity#Pauling electronegativity|Pauling scale]], behind only [[Chemistry:Oxygen|oxygen]] and fluorine.

Chlorine played an important role in the experiments conducted by medieval [[Philosophy:Alchemy|alchemists]], which commonly involved the heating of chloride [[Chemistry:Salt|salts]] like [[Chemistry:Ammonium chloride|ammonium chloride]] (sal ammoniac) and [[Chemistry:Sodium chloride|sodium chloride]] (common salt), producing various chemical substances containing chlorine such as [[Chemistry:Hydrogen chloride|hydrogen chloride]], [[Chemistry:Mercury(II) chloride|mercury(II) chloride]] (corrosive sublimate), and [[Chemistry:Hydrochloric acid|hydrochloric acid]] (in the form of {{lang|la|[[Chemistry:Aqua regia|aqua regia]]}}). However, the nature of free chlorine gas as a separate substance was only recognised around 1630 by [[Biography:Jan Baptist van Helmont|Jan Baptist van Helmont]]. [[Biography:Carl Wilhelm Scheele|Carl Wilhelm Scheele]] wrote a description of chlorine gas in 1774, supposing it to be an [[Chemistry:Oxide|oxide]] of a new element. In 1809, chemists suggested that the gas might be a pure element, and this was confirmed by Sir Humphry Davy in 1810, who named it after the Ancient Greek {{lang|grc|χλωρός}} ({{transl|grc|khlōrós}}, "pale green") because of its colour.

Because of its great reactivity, all chlorine in the Earth's crust is in the form of [[Physics:Ion|ion]]ic [[Chemistry:Chloride|chloride]] compounds, which includes table salt. It is the [[Earth:Abundance of elements in Earth's crust|second-most abundant]] [[Chemistry:Halogen|halogen]] (after fluorine) and twenty-first most abundant chemical element in Earth's crust. These crustal deposits are nevertheless dwarfed by the huge reserves of chloride in seawater.

Elemental chlorine is commercially produced from [[Earth:Brine|brine]] by [[Chemistry:Electrolysis|electrolysis]], predominantly in the chlor-alkali process. The high oxidising potential of elemental chlorine led to the development of commercial bleaches and [[Chemistry:Disinfectant|disinfectant]]s, and a [[Biology:Reagent|reagent]] for many processes in the chemical industry. Chlorine is used in the manufacture of a wide range of consumer products, about two-thirds of them organic chemicals such as [[Chemistry:Polyvinyl chloride|polyvinyl chloride]] (PVC), many intermediates for the production of plastics, and other end products which do not contain the element. As a common disinfectant, elemental chlorine and chlorine-generating compounds are used more directly in [[Earth:Swimming pool|swimming pool]]s to keep them [[Chemistry:Swimming pool sanitation|sanitary]]. Elemental chlorine at high [[Chemistry:Concentration|concentration]] is extremely dangerous, and poisonous to most living organisms. As a [[Social:Chemical warfare|chemical warfare]] agent, chlorine was first used in World War&nbsp;I as a poison gas weapon.

In the form of chloride ions, chlorine is necessary to all known species of life. Other types of chlorine compounds are rare in living organisms, and artificially produced chlorinated organics range from inert to toxic. In the [[Earth:Upper atmosphere|upper atmosphere]], chlorine-containing organic molecules such as [[Earth:Chlorofluorocarbon|chlorofluorocarbon]]s have been implicated in [[Earth:Ozone depletion|ozone depletion]]. Small quantities of elemental chlorine are generated by oxidation of chloride ions in [[Biology:Neutrophil|neutrophil]]s as part of an [[Biology:Immune system|immune system]] response against bacteria.

==History==
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 [[Earth:Brine|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 |access-date=2008-07-10 |url-status=dead |archive-url=https://web.archive.org/web/20110430145935/http://antiquity.ac.uk/ProjGall/weller/ |archive-date=April 30, 2011 }}</ref>

===Early discoveries===
Around 900, the authors of the Arabic writings attributed to [[Biography:Jabir ibn Hayyan|Jabir ibn Hayyan]] (Latin: Geber) and the Persian physician and alchemist [[Biography:Abu Bakr al-Razi|Abu Bakr al-Razi]] ({{circa}} 865–925, Latin: Rhazes) were experimenting with sal ammoniac ([[Chemistry:Ammonium chloride|ammonium chloride]]), which when it was distilled together with vitriol (hydrated sulfates of various metals) produced [[Chemistry:Hydrogen chloride|hydrogen chloride]].<ref>{{Cite book|last=Kraus|first=Paul|year=1942–1943|title=Jâbir ibn Hayyân: Contribution à l'histoire des idées scientifiques dans l'Islam. I. Le corpus des écrits jâbiriens. II. Jâbir et la science grecque|publisher=[[Organization:Institut Français d'Archéologie Orientale|Institut Français d'Archéologie Orientale]]|location=Cairo|oclc=468740510|isbn=978-3487091150}} vol. II, pp. 41–42; {{cite book|last=Multhauf|first=Robert P.|author-link=Robert P. Multhauf|year=1966|title=The Origins of Chemistry|location=London|publisher=Oldbourne|isbn=}} pp. 141–42.</ref> However, it appears that in these early experiments with chloride [[Chemistry:Salt|salts]], the gaseous products were discarded, and hydrogen chloride may have been produced many times before it was discovered that it can be put to chemical use.<ref>{{harvnb|Multhauf|1966|loc=p. 142, note 79}}.</ref> One of the first such uses was the synthesis of [[Chemistry:Mercury(II) chloride|mercury(II) chloride]] (corrosive sublimate), whose production from the heating of [[Chemistry:Mercury (element)|mercury]] either with [[Chemistry:Alum|alum]] and ammonium chloride or with vitriol and sodium chloride was first described in the ''De aluminibus et salibus'' ("On Alums and Salts", an eleventh- or twelfth century Arabic text falsely attributed to Abu Bakr al-Razi and translated into Latin in the second half of the twelfth century by Gerard of Cremona, 1144–1187).<ref>{{harvnb|Multhauf|1966|pp=160–163}}.</ref> Another important development was the discovery by pseudo-Geber (in the ''De inventione veritatis'', "On the Discovery of Truth", after c. 1300) that by adding ammonium chloride to [[Chemistry:Nitric acid|nitric acid]], a strong solvent capable of dissolving gold (i.e., ''[[Chemistry:Aqua regia|aqua regia]]'') could be produced.<ref>{{Cite journal|last1=Karpenko|first1=Vladimír|last2=Norris|first2=John A.|year=2002|title=Vitriol in the History of Chemistry|journal=Chemické listy|volume=96|issue=12|pages=997–1005|url=http://www.chemicke-listy.cz/ojs3/index.php/chemicke-listy/article/view/2266}} p. 1002.</ref> Although ''aqua regia'' is an unstable mixture that continually gives off fumes containing free chlorine gas, this chlorine gas appears to have been ignored until c. 1630, when its nature as a separate gaseous substance was recognised by the Brabantian chemist and physician [[Biography:Jan Baptist van Helmont|Jan Baptist van Helmont]].<ref name="Greenwood789">{{harvnb|Greenwood|Earnshaw|1997|pp=789–92}}</ref><ref group=en>{{cite book|last1=van Helmont|first1=Joannis Baptistae|title=Opera omnia [All Works]|date=1682|publisher=Johann Just Erythropel|location=Frankfurt-am-Main, (Germany)|url=https://books.google.com/books?id=Qy5AAAAAcAAJ&pg=PP5|language=la}} From ''"Complexionum atque mistionum elementalium figmentum."'' (Formation of combinations and of mixtures of elements), §37, [https://books.google.com/books?id=Qy5AAAAAcAAJ&pg=PA105 p. 105:] ''"Accipe salis petrae, vitrioli, & alumnis partes aequas: exsiccato singula, & connexis simul, distilla aquam. Quae nil aliud est, quam merum sal volatile. Hujus accipe uncias quatuor, salis armeniaci unciam junge, in forti vitro, alembico, per caementum (ex cera, colophonia, & vitri pulverre) calidissime affusum, firmato; mox, etiam in frigore, Gas excitatur, & vas, utut forte, dissilit cum fragore."'' (Take equal parts of saltpeter [i.e., sodium nitrate], vitriol [i.e., concentrated sulfuric acid], and alum: dry each and combine simultaneously; distill off the water [i.e., liquid]. That [distillate] is nothing else than pure volatile salt [i.e., spirit of nitre, nitric acid]. Take four ounces of this [viz, nitric acid], add one ounce of Armenian salt [i.e., ammonium chloride], [place it] in a strong glass alembic sealed by cement ([made] from wax, rosin, and powdered glass) [that has been] poured very hot; soon, even in the cold, gas is stimulated, and the vessel, however strong, bursts into fragments.) From ''"De Flatibus"'' (On gases), [https://books.google.com/books?id=Qy5AAAAAcAAJ&pg=PA408 p. 408]: ''"Sal armeniacus enim, & aqua chrysulca, quae singula per se distillari, possunt, & pati calorem: sin autem jungantur, & intepescant, non possunt non, quin statim in Gas sylvestre, sive incoercibilem flatum transmutentur."'' (Truly Armenian salt [i.e., ammonium chloride] and nitric acid, each of which can be distilled by itself, and submitted to heat; but if, on the other hand, they be combined and become warm, they cannot but be changed immediately into carbon dioxide [note: van Helmont's identification of the gas is mistaken] or an incondensable gas.)<br>
See also:
* [https://www.encyclopedia.com/people/science-and-technology/chemistry-biographies/johannes-joan-baptista-van-helmont Helmont, Johannes (Joan) Baptista Van, Encyclopedia.Com]: "Others were chlorine gas from the reaction of nitric acid and sal ammoniac; … "
* Wisniak, Jaime (2009) "Carl Wilhelm Scheele," ''Revista CENIC Ciencias Químicas'', '''40''' (3): 165–73 ; see p. 168: "Early in the seventeenth century Johannes Baptiste van Helmont (1579–1644) mentioned that when sal marin (sodium chloride) or sal ammoniacus and aqua chrysulca (nitric acid) were mixed together, a flatus incoercible (non-condensable gas) was evolved."</ref>
[[File:PSM V31 D740 Carl Wilhelm Scheele.jpg|thumb|left|upright=0.65|Carl Wilhelm Scheele, discoverer of chlorine]]

===Isolation===
The element was first studied in detail in 1774 by Swedish chemist [[Biography:Carl Wilhelm Scheele|Carl Wilhelm Scheele]], and he is credited with the discovery.<ref>{{cite journal|last1=Scheele|first1=Carl Wilhelm|title=Om Brunsten, eller Magnesia, och dess Egenskaper|journal=Kongliga Vetenskaps Academiens Handlingar [Proceedings of the Royal Scientific Academy]|date=1774|volume=35|pages=89–116, 177–94|url=https://babel.hathitrust.org/cgi/pt?id=mdp.39015039452910;view=1up;seq=99|trans-title=On braunstein [i.e., pyrolusite, manganese dioxide], or magnesia, and its properties|language=sv}} [https://babel.hathitrust.org/cgi/pt?id=mdp.39015039452910;view=1up;seq=103 In section 6 on pp. 93–94 of his paper], Scheele described how chlorine was produced when a mixture of hydrochloric acid and manganese dioxide (''Brunsten'') was heated: ''"6) (a) På 1/2 uns fint rifven Brunsten slogs 1 uns ren Spiritus salis. … samt lukten fo̊rsvunnen."'' ( 6) (a) On one half ounce of finely ground Braunstein [pyrolusite] was poured one ounce of pure ''spiritus salis'' [spirit of salt, hydrogen chloride]. After this mixture had been standing in the cold for one hour, the acid had assumed a dark brown colour. One part of this solution was poured into a glass, which was placed over the fire. The solution gave off an odour like warm ''aqua regia'' and after one quarter's hour duration, it was as clear and colourless as water, and the smell had disappeared.) For an English translation of the relevant passages of this article, see: ''The Early History of Chlorine : Papers by Carl Wilhelm Scheele (1774), C. L. Berthollet (1785), Guyton de Morveau (1787), J. L. Gay-Lussac and L. J. Thenard (1809)'' (Edinburgh, Scotland: Alembic Club, 1912), [https://babel.hathitrust.org/cgi/pt?id=coo.31924012394379;view=1up;seq=9 pp. 5–10.]</ref><ref name="krogt">{{cite web |url=http://elements.vanderkrogt.net/element.php?sym=Cl |title=17 Chlorine |publisher=Elements.vanderkrogt.net |access-date=2008-09-12 |url-status=dead |archive-url=https://web.archive.org/web/20100123003013/http://elements.vanderkrogt.net/element.php?sym=Cl |archive-date=2010-01-23 }}</ref> Scheele produced chlorine by reacting [[Chemistry:Manganese dioxide|MnO<sub>2</sub>]] (as the mineral [[Chemistry:Pyrolusite|pyrolusite]]) with HCl:<ref name="Greenwood789" />

: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, the deadly effect on insects, the yellow-green color, and the smell similar to [[Chemistry:Aqua regia|aqua regia]].<ref name="Greenwood792">{{harvnb|Greenwood|Earnshaw|1997|pp=792–93}}</ref> He called it "''dephlogisticated muriatic acid air''" since it is a gas (then called "airs") and it came from [[Chemistry:Hydrochloric acid|hydrochloric acid]] (then known as "muriatic acid").<ref name="krogt" /> He failed to establish chlorine as an element.<ref name="krogt" />

Common chemical theory at that time held that an acid is a compound that contains oxygen (remnants of this survive in the German and Dutch names of [[Chemistry:Oxygen|oxygen]]: ''sauerstoff'' or ''zuurstof'', both translating into English as ''acid substance''), 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=https://books.google.com/books?id=34KwmkU4LG0C&pg=PA158|page=158|title=The development of modern chemistry|author=Ihde, Aaron John|publisher=Courier Dover Publications|date=1984|isbn=978-0-486-64235-2}}</ref><ref name="Weeks">{{cite journal|doi = 10.1021/ed009p1915|title = The discovery of the elements. XVII. The halogen family|date = 1932|last1 = Weeks|first1 = Mary Elvira|journal = Journal of Chemical Education|volume = 9|issue = 11|page = 1915|bibcode=1932JChEd...9.1915W}}</ref>

In 1809, [[Biography:Joseph Louis Gay-Lussac|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|last1=Gay-Lussac|last2=Thenard|title = Extrait des mémoires lus à l'Institut national, depuis le 7 mars 1808 jusqu'au 27 février 1809.|trans-title=Extracts from memoirs read at the national Institute, from 7 March 1808 to 27 February 1809|journal = Mémoires de Physique et de Chimie de la Société d'Arcueil|volume = 2|pages = 295–358|date = 1809|url = https://books.google.com/books?id=hnJKAAAAYAAJ&pg=PA295}} See: §
''De la nature et des propriétés de l'acide muriatique et de l'acide muriatique oxigéné'' (On the nature and properties of muriatic acid and of oxidized muriatic acid), pp. 339–58. From pp. 357–58: ''"Le gaz muriatique oxigéné n'est pas, en effect, décomposé … comme un corps composé."'' ("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.") For a full English translation of this section, see: [http://web.lemoyne.edu/~giunta/thenard.html Joseph Louis Gay-Lussac and Louis Jacques Thénard, "On the nature and the properties of muriatic acid and of oxygenated muriatic acid" (Lemoyne College, Syracuse, New York)]</ref>

In 1810, Sir Humphry Davy tried the same experiment again, and concluded that the substance was an element, and not a compound.<ref name="krogt" /> He announced his results to the Royal Society on 15 November that year.<ref name="Greenwood789" /> At that time, he named this new element "chlorine", from the Greek word χλωρος (''chlōros'', "green-yellow"), in reference to its color.<ref>{{cite journal|last1=Davy|first1=Humphry|title=The Bakerian Lecture. On some of the combinations of oxymuriatic gas and oxygene, and on the chemical relations of these principles, to inflammable bodies|journal=Philosophical Transactions of the Royal Society of London|date=1811|volume=101|pages=1–35|url=https://babel.hathitrust.org/cgi/pt?id=mdp.39015034564297;view=1up;seq=15|bibcode=1811RSPT..101....1D|doi=10.1098/rstl.1811.0001|doi-access=free}} Davy named chlorine on [https://babel.hathitrust.org/cgi/pt?id=mdp.39015034564297;view=1up;seq=46 p. 32:] "After consulting some of the most eminent chemical philosophers in this country, it has been judged most proper to suggest a name founded upon one of its obvious and characteristic properties – its colour, and to call it ''Chlorine'', or ''Chloric'' gas.* *From χλωρος."</ref> The name "[[Chemistry:Halogen|halogen]]", meaning "salt producer", was originally used for chlorine in 1811 by Johann Salomo Christoph Schweigger.<ref>{{cite journal|last1=Schweigger|first1=J.S.C.|title=Nachschreiben des Herausgebers, die neue Nomenclatur betreffend|journal=Journal für Chemie und Physik|date=1811|volume=3|issue=2|pages=249–55|url=https://babel.hathitrust.org/cgi/pt?id=njp.32101076802287;view=1up;seq=295|trans-title=Postscript of the editor concerning the new nomenclature|language=de}} On p. 251, Schweigger proposed the word "halogen": ''"Man sage dafür lieber mit richter Wortbildung ''Halogen'' (da schon in der Mineralogie durch ''Werner's'' Halit-Geschlecht dieses Wort nicht fremd ist) von αλς ''Salz'' und dem alten γενειν (dorisch γενεν) ''zeugen''."'' (One should say instead, with proper morphology, "halogen" (this word is not strange since [it's] already in mineralogy via Werner's "halite" species) from αλς [als] "salt" and the old γενειν [genein] (Doric γενεν) "to beget".)</ref> This term was later used as a generic term to describe all the elements in the chlorine family (fluorine, bromine, iodine), after a suggestion by Jöns Jakob Berzelius in 1826.<ref>In 1826, Berzelius coined the terms ''Saltbildare'' (salt-formers) and ''Corpora Halogenia'' (salt-making substances) for the elements chlorine, iodine, and fluorine. See: {{cite journal|last1=Berzelius|first1=Jacob|title=Årsberättelser om Framstegen i Physik och Chemie|journal=Arsb. Vetensk. Framsteg |trans-title=Annual Report on Progress in Physics and Chemistry|date=1826|publisher=P.A. Norstedt & Söner|location=Stockholm, Sweden|volume=6|page=187|url=https://babel.hathitrust.org/cgi/pt?id=hvd.32044092556919;view=1up;seq=195|language=sv}} From p. 187: ''"De förre af dessa, d. ä. ''de electronegativa'', dela sig i tre klasser: 1) den första innehåller kroppar, som förenade med de electropositiva, omedelbart frambringa salter, hvilka jag derför kallar ''Saltbildare'' (Corpora Halogenia). Desse utgöras af chlor, iod och fluor *)."'' (The first of them [i.e., elements], i.e., the electronegative [ones], are divided into three classes: 1) The first includes substances which, [when] united with electropositive [elements], immediately produce salts, and which I therefore name "salt-formers" (salt-producing substances). These are chlorine, iodine, and fluorine *).)</ref><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 |pages=328–38|date=1971| doi=10.1086/350763| jstor=229946|s2cid=170337569}}</ref> In 1823, [[Biography:Michael Faraday|Michael Faraday]] liquefied chlorine for the first time,<ref>{{cite journal|last1=Faraday|first1=M.|title=On fluid chlorine|journal=Philosophical Transactions of the Royal Society of London|date=1823|volume=113|pages=160–64|url=https://babel.hathitrust.org/cgi/pt?id=mdp.39015034564487;view=1up;seq=244|bibcode=1823RSPT..113..160F|doi=10.1098/rstl.1823.0016|doi-access=free}}</ref><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 |editor=Chodos, Alan |publisher=American Physical Society |access-date=2010-05-08 |url-status=dead |archive-url=https://web.archive.org/web/20100615142343/http://aps.org/publications/apsnews/200108/history.cfm |archive-date=June 15, 2010 }}</ref><ref>{{cite web |url=http://www-history.mcs.st-andrews.ac.uk/Biographies/Faraday.html |author=O'Connor J. J. |author2=Robertson E. F. |work=School of Mathematics and Statistics, University of St Andrews, Scotland |title=Michael Faraday |access-date=2010-05-08 |url-status=dead |archive-url=https://web.archive.org/web/20100220051522/http://www-history.mcs.st-andrews.ac.uk/Biographies/Faraday.html |archive-date=2010-02-20 }}</ref> and demonstrated that what was then known as "solid chlorine" had a structure of [[Physics:Clathrate hydrate|chlorine hydrate]] (Cl<sub>2</sub>·H<sub>2</sub>O).<ref name="Greenwood789" />

===Later uses===
Chlorine gas was first used by French chemist Claude Berthollet to bleach textiles in 1785.<ref name="encyc">{{cite encyclopedia |encyclopedia=Encyclopædia Britannica |title=Bleaching |url=http://www.1902encyclopedia.com/B/BLE/bleaching.html |access-date=2012-05-02 |edition=9th Edition (1875) and 10th Edition (1902) |url-status=dead |archive-url=https://web.archive.org/web/20120524111017/http://www.1902encyclopedia.com/B/BLE/bleaching.html |archive-date=2012-05-24 }}</ref><ref name="Cotton">{{Cite book| title=The Cotton Industry| last=Aspin| first=Chris| publisher=Shire Publications Ltd| date=1981| isbn=978-0-85263-545-2| page=[https://archive.org/details/cottonindustry0000aspi/page/24 24]| url=https://archive.org/details/cottonindustry0000aspi/page/24}}</ref> Modern bleaches resulted from further work by Berthollet, who first produced [[Chemistry:Sodium hypochlorite|sodium hypochlorite]] in 1789 in his laboratory in the town of Javel (now part of [[Place:Paris|Paris]], France), by passing chlorine gas through a solution of sodium carbonate. The resulting liquid, known as "''Eau de Javel''" ("Javel water"), was a weak solution of [[Chemistry:Sodium hypochlorite|sodium hypochlorite]]. This process was not very efficient, and alternative production methods were sought. Scottish chemist and industrialist Charles Tennant first produced a solution of [[Chemistry:Calcium hypochlorite|calcium hypochlorite]] ("chlorinated lime"), then solid calcium hypochlorite (bleaching powder).<ref name="encyc" /> These compounds produced low levels of elemental chlorine and could be more efficiently transported than sodium hypochlorite, which remained as dilute solutions because when purified to eliminate water, it became a dangerously powerful and unstable oxidizer. Near the end of the nineteenth century, E. S. Smith patented a method of sodium hypochlorite production involving electrolysis of [[Earth:Brine|brine]] to produce [[Chemistry:Sodium hydroxide|sodium hydroxide]] and chlorine gas, which then mixed to form sodium hypochlorite.<ref>{{cite web |url=http://www.chm.bris.ac.uk/motm/bleach/bleachh.htm |title=Bleach (Sodium Hypochlorite) |author=Paul May |publisher=University of Bristol |access-date=13 December 2016 |archive-url=https://web.archive.org/web/20161213230407/http://www.chm.bris.ac.uk/motm/bleach/bleachh.htm |archive-date=13 December 2016 |url-status=live}}</ref> This is known as the [[Physics:Chloralkali process|chloralkali process]], first introduced on an industrial scale in 1892, and now the source of most elemental chlorine and sodium hydroxide.<ref name="Greenwood798">{{harvnb|Greenwood|Earnshaw|1997|p=798}}</ref> In 1884 Chemischen Fabrik Griesheim of Germany developed another [[Physics:Chloralkali process|chloralkali process]] which entered commercial production in 1888.<ref>{{cite book|url=https://books.google.com/books?id=OI0fTJhydh4C&q=Chlorine-Alkali+Griesheim+1884&pg=PA220|title=History of Industrial Gases|isbn=978-0-306-47277-0|last1=Almqvist|first1=Ebbe|publisher=Springer Science & Business Media|date=2003|page=220}}</ref>

Elemental chlorine solutions dissolved in [[Chemistry:Base|chemically basic]] water (sodium and [[Chemistry:Calcium hypochlorite|calcium hypochlorite]]) were first used as anti-[[Biology:Putrefaction|putrefaction]] agents and [[Chemistry:Disinfectant|disinfectant]]s in the 1820s, in France, long before the establishment of the [[Biology:Germ theory of disease|germ theory of disease]]. This practice was pioneered by Antoine-Germain Labarraque, who adapted Berthollet's "Javel water" bleach and other chlorine preparations (for a more complete history, see below).<ref name="bouvet">{{cite journal |author=Bouvet, Maurice |language=fr |trans-title=The great pharmacists: Labarraque (1777–1850) |title=Les grands pharmaciens: Labarraque (1777–1850) |journal=Revue d'Histoire de la Pharmacie |date=1950 |volume=38 |issue=128 |pages=97–107 |doi=10.3406/pharm.1950.8662 }}</ref> Elemental chlorine has since served a continuous function in topical [[Chemistry:Antiseptic|antisepsis]] (wound irrigation solutions and the like) and public sanitation, particularly in swimming and drinking water.<ref name="Greenwood792" />

Chlorine gas was first used as a weapon on April 22, 1915, at [[Earth:Ypres|Ypres]] by the German Army.<ref>{{cite web|url=http://www.drcordas.com/education/weaponsmassd/Chlorine.pdf |title=Chlorine – History |access-date=2008-07-10 |url-status=dead |archive-url=https://web.archive.org/web/20070221140939/http://www.drcordas.com/education/weaponsmassd/Chlorine.pdf |archive-date=21 February 2007 }}</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|access-date = 2008-07-10|url-status = live|archive-url = https://web.archive.org/web/20080702044210/http://www.historynet.com/weaponry-use-of-chlorine-gas-cylinders-in-world-war-i.htm|archive-date = 2008-07-02|date = 2006-06-12}}</ref> The effect on the allies was devastating because the existing gas masks were difficult to deploy and had not been broadly distributed.<ref>{{Cite web
 |author      = Staff
 |date        = 29 July 2004
 |url         = http://www.vac-acc.gc.ca/remembers/sub.cfm?source=history/firstwar/canada/Canada4
 |title       = On the Western Front, Ypres 1915
 |publisher   = Veteran Affairs Canada
 |access-date  = 2008-04-08
 |url-status     = dead
 |archive-url  = https://web.archive.org/web/20081206024237/http://www.vac-acc.gc.ca/remembers/sub.cfm?source=history%2Ffirstwar%2Fcanada%2FCanada4
 |archive-date = 6 December 2008
}}</ref><ref>{{Cite book
 | first=Victor | last=Lefebure |author2=Wilson, Henry
 | title=The Riddle of the Rhine: Chemical Strategy in Peace and War
 | year=2004 | publisher=Kessinger Publishing | isbn=978-1-4179-3546-8 }}</ref>

==Properties==
[[File:Chlorine liquid in an ampoule.jpg|thumb|Chlorine, liquefied under a pressure of 7.4 bar at room temperature, displayed in a quartz ampule embedded in [[Chemistry:Poly(methyl methacrylate)|acrylic glass]]]]
[[File:Die chemischen elemente cl.jpg|thumb|Solid chlorine at −150&nbsp;°C]]
Chlorine is the second [[Chemistry:Halogen|halogen]], being a [[Chemistry:Nonmetal|nonmetal]] in group 17 of the periodic table. Its properties are thus similar to [[Chemistry:Fluorine|fluorine]], [[Chemistry:Bromine|bromine]], and [[Chemistry:Iodine|iodine]], and are largely intermediate between those of the first two. Chlorine has the electron configuration [Ne]3s<sup>2</sup>3p<sup>5</sup>, with the seven electrons in the third and outermost shell acting as its [[Physics:Valence electron|valence electron]]s. Like all halogens, it is thus one electron short of a full octet, and is hence a strong oxidising agent, reacting with many elements in order to complete its outer shell.<ref name="Greenwood800">{{harvnb|Greenwood|Earnshaw|1997|pp=800–804}}</ref> Corresponding to periodic trends, it is intermediate in [[Chemistry:Electronegativity|electronegativity]] between fluorine and bromine (F: 3.98, Cl: 3.16, Br: 2.96, I: 2.66), and is less reactive than fluorine and more reactive than bromine. It is also a weaker oxidising agent than fluorine, but a stronger one than bromine. Conversely, the [[Chemistry:Chloride|chloride]] ion is a weaker reducing agent than bromide, but a stronger one than fluoride.<ref name="Greenwood800" /> It is intermediate in [[Physics:Atomic radius|atomic radius]] between fluorine and bromine, and this leads to many of its atomic properties similarly continuing the trend from iodine to bromine upward, such as first ionisation energy, [[Physics:Electron affinity|electron affinity]], enthalpy of dissociation of the X<sub>2</sub> molecule (X = Cl, Br, I), ionic radius, and X–X bond length. (Fluorine is anomalous due to its small size.)<ref name="Greenwood800" />

All four stable halogens experience intermolecular [[Physics:Van der Waals force|van der Waals force]]s of attraction, and their strength increases together with the number of electrons among all homonuclear diatomic halogen molecules. Thus, the melting and boiling points of chlorine are intermediate between those of fluorine and bromine: chlorine melts at −101.0&nbsp;°C and boils at −34.0&nbsp;°C. As a result of the increasing molecular weight of the halogens down the group, the density and heats of fusion and vaporisation of chlorine are again intermediate between those of bromine and fluorine, although all their heats of vaporisation are fairly low (leading to high volatility) thanks to their diatomic molecular structure.<ref name="Greenwood800" /> The halogens darken in colour as the group is descended: thus, while fluorine is a pale yellow gas, chlorine is distinctly yellow-green. This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group.<ref name="Greenwood800" /> Specifically, the colour of a halogen, such as chlorine, results from the [[Physics:Atomic electron transition|electron transition]] between the highest occupied antibonding ''π<sub>g</sub>'' molecular orbital and the lowest vacant antibonding ''σ<sub>u</sub>'' molecular orbital.<ref name="Greenwood804">{{harvnb|Greenwood|Earnshaw|1997|pp=804–809}}</ref> The colour fades at low temperatures, so that solid chlorine at −195&nbsp;°C is almost colourless.<ref name="Greenwood800" />

Like solid bromine and iodine, solid chlorine crystallises in the [[Physics:Orthorhombic crystal system|orthorhombic crystal system]], in a layered lattice of Cl<sub>2</sub> molecules. The Cl–Cl distance is 198&nbsp;pm (close to the gaseous Cl–Cl distance of 199&nbsp;pm) and the Cl···Cl distance between molecules is 332&nbsp;pm within a layer and 382&nbsp;pm between layers (compare the van der Waals radius of chlorine, 180&nbsp;pm). This structure means that chlorine is a very poor conductor of electricity, and indeed its conductivity is so low as to be practically unmeasurable.<ref name="Greenwood800" />

===Isotopes===
{{main|Physics:Isotopes of chlorine}}
Chlorine has two stable isotopes, <sup>35</sup>Cl and <sup>37</sup>Cl. These are its only two natural isotopes occurring in quantity, with <sup>35</sup>Cl making up 76% of natural chlorine and <sup>37</sup>Cl making up the remaining 24%. Both are synthesised in stars in the [[Astronomy:Oxygen-burning process|oxygen-burning]] and [[Astronomy:Silicon-burning process|silicon-burning process]]es.<ref name="Cameron">{{cite journal|last1=Cameron |first1=A. G. W. |year=1973 |title=Abundance of the Elements in the Solar System |url=http://pubs.giss.nasa.gov/docs/1973/1973_Cameron_1.pdf |journal=Space Science Reviews |volume=15 |issue=1 |pages=121–46 |doi=10.1007/BF00172440 |bibcode=1973SSRv...15..121C |s2cid=120201972 |url-status=dead |archive-url=https://web.archive.org/web/20111021030549/http://pubs.giss.nasa.gov/docs/1973/1973_Cameron_1.pdf |archive-date=2011-10-21 }}</ref> Both have nuclear spin 3/2+ and thus may be used for [[Physics:Nuclear magnetic resonance|nuclear magnetic resonance]], although the spin magnitude being greater than 1/2 results in non-spherical nuclear charge distribution and thus resonance broadening as a result of a nonzero nuclear quadrupole moment and resultant quadrupolar relaxation. The other chlorine isotopes are all radioactive, with [[Physics:Half-life|half-lives]] too short to occur in nature primordially. Of these, the most commonly used in the laboratory are <sup>36</sup>Cl (''t''<sub>1/2</sub> = 3.0×10<sup>5</sup>&nbsp;y) and <sup>38</sup>Cl (''t''<sub>1/2</sub> = 37.2&nbsp;min), which may be produced from the [[Physics:Neutron activation|neutron activation]] of natural chlorine.<ref name="Greenwood800" />

The most stable chlorine radioisotope is <sup>36</sup>Cl. The primary decay mode of isotopes lighter than <sup>35</sup>Cl is [[Physics:Electron capture|electron capture]] to isotopes of [[Chemistry:Sulfur|sulfur]]; that of isotopes heavier than <sup>37</sup>Cl is [[Physics:Beta decay|beta decay]] to isotopes of [[Chemistry:Argon|argon]]; and <sup>36</sup>Cl may decay by either mode to stable <sup>36</sup>S or <sup>36</sup>Ar.<ref name="NUBASE">{{NUBASE 2003}}</ref> [[Physics:Chlorine-36|<sup>36</sup>Cl]] occurs in trace quantities in nature as a [[Physics:Cosmogenic nuclide|cosmogenic nuclide]] in a ratio of about (7–10)&nbsp;×&nbsp;10<sup>−13</sup> to 1 with stable chlorine isotopes: it is produced in the atmosphere by [[Physics:Spallation|spallation]] of <sup>36</sup>[[Chemistry:Argon|Ar]] by interactions with [[Astronomy:Cosmic ray|cosmic ray]] [[Physics:Proton|proton]]s. In the top meter of the lithosphere, <sup>36</sup>Cl is generated primarily by thermal neutron activation of <sup>35</sup>Cl and spallation of <sup>39</sup>[[Chemistry:Potassium|K]] and <sup>40</sup>[[Chemistry:Calcium|Ca]]. In the subsurface environment, [[Physics:Muon capture|muon capture]] by <sup>40</sup>[[Chemistry:Calcium|Ca]] becomes more important as a way to generate <sup>36</sup>Cl.<ref name="Zreda">{{cite journal | author = M. Zreda | year = 1991 | title = Cosmogenic chlorine-36 production rates in terrestrial rocks | journal = Earth and Planetary Science Letters | volume = 105 | issue = 1–3 | pages = 94–109|bibcode = 1991E&PSL.105...94Z |doi = 10.1016/0012-821X(91)90123-Y |display-authors=0}}</ref><ref>{{cite journal | author = M. Sheppard and M. Herod | year = 2012 | title = Variation in background concentrations and specific activities of 36Cl, 129I and U/Th-series radionuclides in surface waters | journal = Journal of Environmental Radioactivity | volume = 106 | pages = 27–34 | doi=10.1016/j.jenvrad.2011.10.015| pmid = 22304997 }}</ref>

==Chemistry and compounds==
{| class="wikitable" style="float:right; margin-top:0; margin-left:1em; text-align:center; font-size:10pt; line-height:11pt; width:25%;"
|+ style="margin-bottom: 5px;" | Halogen bond energies (kJ/mol)<ref name="Greenwood804" />
|-
! X
! XX
! HX
! BX<sub>3</sub>
! AlX<sub>3</sub>
! CX<sub>4</sub>
|-
! F
| 159
| 574
| 645
| 582
| 456
|-
! Cl
|243
|428
|444
|427
|327
|-
! Br
|193
|363
|368
|360
|272
|-
! I
|151
|294
|272
|285
|239
|}
Chlorine is intermediate in reactivity between fluorine and bromine, and is one of the most reactive elements. Chlorine is a weaker oxidising agent than fluorine but a stronger one than bromine or iodine. This can be seen from the [[Chemistry:Standard electrode potential|standard electrode potential]]s of the X<sub>2</sub>/X<sup>−</sup> couples (F, +2.866&thinsp;V; Cl, +1.395&thinsp;V; Br, +1.087&thinsp;V; I, +0.615&thinsp;V; At, approximately +0.3&nbsp;&thinsp;V). However, this trend is not shown in the bond energies because fluorine is singular due to its small size, low polarisability, and inability to show hypervalence. As another difference, chlorine has a significant chemistry in positive oxidation states while fluorine does not. Chlorination often leads to higher oxidation states than bromination or iodination but lower oxidation states than fluorination. Chlorine tends to react with compounds including M–M, M–H, or M–C bonds to form M–Cl bonds.<ref name="Greenwood804" />

Given that E°({{sfrac|1|2}}O<sub>2</sub>/H<sub>2</sub>O) = +1.229&nbsp;V, which is less than +1.395&nbsp;V, it would be expected that chlorine should be able to oxidise water to oxygen and hydrochloric acid. However, the kinetics of this reaction are unfavorable, and there is also a bubble [[Chemistry:Overpotential|overpotential]] effect to consider, so that electrolysis of aqueous chloride solutions evolves chlorine gas and not oxygen gas, a fact that is very useful for the industrial production of chlorine.<ref name="Greenwood853">{{harvnb|Greenwood|Earnshaw|1997|pp=853–856}}</ref>

===Hydrogen chloride===
[[File:DCl Neutron powder.png|thumb|upright=1.8|right|Structure of solid deuterium chloride, with D···Cl hydrogen bonds]]
The simplest chlorine compound is [[Chemistry:Hydrogen chloride|hydrogen chloride]], HCl, a major chemical in industry as well as in the laboratory, both as a gas and dissolved in water as [[Chemistry:Hydrochloric acid|hydrochloric acid]]. It is often produced by burning hydrogen gas in chlorine gas, or as a byproduct of chlorinating [[Chemistry:Hydrocarbon|hydrocarbon]]s. Another approach is to treat [[Chemistry:Sodium chloride|sodium chloride]] with concentrated [[Chemistry:Sulfuric acid|sulfuric acid]] to produce hydrochloric acid, also known as the "salt-cake" process:<ref name="Greenwood809" />
:NaCl + H<sub>2</sub>SO<sub>4</sub> {{overunderset|⟶|150&thinsp;°C|&thinsp;}} NaHSO<sub>4</sub> + HCl
:NaCl + NaHSO<sub>4</sub> {{overunderset|⟶|540–600&thinsp;°C|&thinsp;}} Na<sub>2</sub>SO<sub>4</sub> + HCl
In the laboratory, hydrogen chloride gas may be made by drying the acid with concentrated sulfuric acid. Deuterium chloride, DCl, may be produced by reacting [[Chemistry:Benzoyl chloride|benzoyl chloride]] with [[Chemistry:Heavy water|heavy water]] (D<sub>2</sub>O).<ref name="Greenwood809">{{harvnb|Greenwood|Earnshaw|1997|pp=809–812}}</ref>

At room temperature, hydrogen chloride is a colourless gas, like all the hydrogen halides apart from [[Chemistry:Hydrogen fluoride|hydrogen fluoride]], since hydrogen cannot form strong [[Chemistry:Hydrogen bond|hydrogen bond]]s to the larger electronegative chlorine atom; however, weak hydrogen bonding is present in solid crystalline hydrogen chloride at low temperatures, similar to the hydrogen fluoride structure, before disorder begins to prevail as the temperature is raised.<ref name="Greenwood809" /> Hydrochloric acid is a strong acid (p''K''<sub>a</sub> = −7) because the hydrogen bonds to chlorine are too weak to inhibit dissociation. The HCl/H<sub>2</sub>O system has many hydrates HCl·''n''H<sub>2</sub>O for ''n'' = 1, 2, 3, 4, and 6. Beyond a 1:1 mixture of HCl and H<sub>2</sub>O, the system separates completely into two separate liquid phases. Hydrochloric acid forms an [[Physics:Azeotrope|azeotrope]] with boiling point 108.58&nbsp;°C at 20.22&nbsp;g HCl per 100&nbsp;g solution; thus hydrochloric acid cannot be concentrated beyond this point by distillation.<ref name="Greenwood812">{{harvnb|Greenwood|Earnshaw|1997|pp=812–816}}</ref>

Unlike hydrogen fluoride, anhydrous liquid hydrogen chloride is difficult to work with as a solvent, because its boiling point is low, it has a small liquid range, its dielectric constant is low and it does not dissociate appreciably into H<sub>2</sub>Cl<sup>+</sup> and {{chem|HCl|2|-}} ions – the latter, in any case, are much less stable than the [[Chemistry:Bifluoride|bifluoride]] ions ({{chem|HF|2|-}}) due to the very weak hydrogen bonding between hydrogen and chlorine, though its salts with very large and weakly polarising cations such as [[Chemistry:Caesium|Cs<sup>+</sup>]] and [[Chemistry:Quaternary ammonium cation|{{chem|NR|4|+}}]] (R = [[Chemistry:Methyl group|Me]], [[Chemistry:Ethyl group|Et]], [[Chemistry:Butyl group|Bu<sup>''n''</sup>]]) may still be isolated. Anhydrous hydrogen chloride is a poor solvent, only able to dissolve small molecular compounds such as [[Chemistry:Nitrosyl chloride|nitrosyl chloride]] and [[Chemistry:Phenol|phenol]], or salts with very low [[Physics:Lattice energy|lattice energies]] such as tetraalkylammonium halides. It readily protonates [[Chemistry:Electrophile|electrophile]]s containing lone-pairs or π bonds. [[Chemistry:Solvolysis|Solvolysis]], [[Chemistry:Ligand|ligand]] replacement reactions, and oxidations are well-characterised in hydrogen chloride solution:<ref name="Greenwood818">{{harvnb|Greenwood|Earnshaw|1997|pp=818–819}}</ref>
:Ph<sub>3</sub>SnCl + HCl ⟶ Ph<sub>2</sub>SnCl<sub>2</sub> + PhH (solvolysis)
:Ph<sub>3</sub>COH + 3 HCl ⟶ {{chem|Ph|3|C|+|HCl|2|-}} + H<sub>3</sub>O<sup>+</sup>Cl<sup>−</sup> (solvolysis)
:{{chem|Me|4|N|+|HCl|2|-}} + BCl<sub>3</sub> ⟶ {{chem|Me|4|N|+|BCl|4|-}} + HCl (ligand replacement)
:PCl<sub>3</sub> + Cl<sub>2</sub> + HCl ⟶ {{chem|PCl|4|+|HCl|2|-}} (oxidation)

===Other binary chlorides===
[[File:Nickel(II)-chloride-hexahydrate-sample.jpg|thumb|right|Hydrated [[Chemistry:Nickel(II) chloride|nickel(II) chloride]], NiCl<sub>2</sub>(H<sub>2</sub>O)<sub>6</sub>]]
Nearly all elements in the periodic table form binary chlorides. The exceptions are decidedly in the minority and stem in each case from one of three causes: extreme inertness and reluctance to participate in chemical reactions (the [[Chemistry:Noble gas|noble gas]]es, with the exception of [[Chemistry:Xenon|xenon]] in the highly unstable [[Chemistry:Xenon dichloride|XeCl<sub>2</sub>]] and XeCl<sub>4</sub>); extreme nuclear instability hampering chemical investigation before decay and transmutation (many of the heaviest elements beyond [[Chemistry:Bismuth|bismuth]]); and having an electronegativity higher than chlorine's ([[Chemistry:Oxygen|oxygen]] and [[Chemistry:Fluorine|fluorine]]) so that the resultant binary compounds are formally not chlorides but rather oxides or fluorides of chlorine.<ref name="Greenwood821">{{harvnb|Greenwood|Earnshaw|1997|pp=821–844}}</ref> Even though [[Chemistry:Nitrogen|nitrogen]] in NCl<sub>3</sub> is bearing a negative charge, the compound is usually called [[Chemistry:Nitrogen trichloride|nitrogen trichloride]].

Chlorination of metals with Cl<sub>2</sub> usually leads to a higher oxidation state than bromination with Br<sub>2</sub> when multiple oxidation states are available, such as in [[Chemistry:Molybdenum(V) chloride|MoCl<sub>5</sub>]] and [[Chemistry:Molybdenum(III) bromide|MoBr<sub>3</sub>]]. Chlorides can be made by reaction of an element or its oxide, hydroxide, or carbonate with hydrochloric acid, and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen chloride gas. These methods work best when the chloride product is stable to hydrolysis; otherwise, the possibilities include high-temperature oxidative chlorination of the element with chlorine or hydrogen chloride, high-temperature chlorination of a metal oxide or other halide by chlorine, a volatile metal chloride, [[Chemistry:Carbon tetrachloride|carbon tetrachloride]], or an organic chloride. For instance, [[Chemistry:Zirconium dioxide|zirconium dioxide]] reacts with chlorine at standard conditions to produce zirconium tetrachloride, and [[Chemistry:Uranium trioxide|uranium trioxide]] reacts with [[Chemistry:Hexachloropropene|hexachloropropene]] when heated under [[Reflux|reflux]] to give [[Chemistry:Uranium tetrachloride|uranium tetrachloride]]. The second example also involves a reduction in [[Chemistry:Oxidation state|oxidation state]], which can also be achieved by reducing a higher chloride using hydrogen or a metal as a reducing agent. This may also be achieved by thermal decomposition or disproportionation as follows:<ref name="Greenwood821" />
: EuCl<sub>3</sub> + {{sfrac|1|2}} H<sub>2</sub> ⟶ EuCl<sub>2</sub> + HCl
: ReCl<sub>5</sub> {{overunderset|⟶|at "bp"|&thinsp;}} ReCl<sub>3</sub> + Cl<sub>2</sub>
: AuCl<sub>3</sub> {{overunderset|⟶|160&thinsp;°C|&thinsp;}} AuCl + Cl<sub>2</sub>

Most metal chlorides with the metal in low oxidation states (+1 to +3) are ionic. Nonmetals tend to form covalent molecular chlorides, as do metals in high oxidation states from +3 and above. Both ionic and covalent chlorides are known for metals in oxidation state +3 (e.g. [[Chemistry:Scandium chloride|scandium chloride]] is mostly ionic, but [[Chemistry:Aluminium chloride|aluminium chloride]] is not). [[Chemistry:Silver chloride|Silver chloride]] is very insoluble in water and is thus often used as a qualitative test for chlorine.<ref name="Greenwood821" />

===Polychlorine compounds===
Although dichlorine is a strong oxidising agent with a high first ionisation energy, it may be oxidised under extreme conditions to form the {{chem2|[Cl2]+}} cation. This is very unstable and has only been characterised by its electronic band spectrum when produced in a low-pressure discharge tube. The yellow {{chem2|[Cl3]+}} cation is more stable and may be produced as follows:<ref name="Greenwood842">{{harvnb|Greenwood|Earnshaw|1997|pp=842–844}}</ref>
:{{chem2|Cl2 + ClF + AsF5}} {{overset|−78 °C|⟶}} {{chem2|[Cl3]+[AsF6]-}}
This reaction is conducted in the oxidising solvent [[Chemistry:Arsenic pentafluoride|arsenic pentafluoride]]. The trichloride anion, {{chem2|[Cl3]-}}, has also been characterised; it is analogous to [[Chemistry:Triiodide|triiodide]].<ref name="Greenwood824" />

===Chlorine fluorides===
The three fluorides of chlorine form a subset of the [[Chemistry:Interhalogen|interhalogen]] compounds, all of which are diamagnetic.<ref name="Greenwood824">{{harvnb|Greenwood|Earnshaw|1997|pp=824–828}}</ref> Some cationic and anionic derivatives are known, such as {{chem|ClF|2|-}}, {{chem|ClF|4|-}}, {{chem|ClF|2|+}}, and Cl<sub>2</sub>F<sup>+</sup>.<ref name="Greenwood835">{{harvnb|Greenwood|Earnshaw|1997|pp=835–842}}</ref> Some [[Chemistry:Pseudohalogen|pseudohalides]] of chlorine are also known, such as [[Chemistry:Cyanogen chloride|cyanogen chloride]] (ClCN, linear), chlorine [[Chemistry:Cyanate|cyanate]] (ClNCO), chlorine [[Chemistry:Thiocyanate|thiocyanate]] (ClSCN, unlike its oxygen counterpart), and chlorine [[Chemistry:Azide|azide]] (ClN<sub>3</sub>).<ref name="Greenwood824" />

[[Chemistry:Chlorine monofluoride|Chlorine monofluoride]] (ClF) is extremely thermally stable, and is sold commercially in 500-gram steel lecture bottles. It is a colourless gas that melts at −155.6&nbsp;°C and boils at −100.1&nbsp;°C. It may be produced by the direction of its elements at 225&nbsp;°C, though it must then be separated and purified from [[Chemistry:Chlorine trifluoride|chlorine trifluoride]] and its reactants. Its properties are mostly intermediate between those of chlorine and fluorine. It will react with many metals and nonmetals from room temperature and above, fluorinating them and liberating chlorine. It will also act as a chlorofluorinating agent, adding chlorine and fluorine across a multiple bond or by oxidation: for example, it will attack [[Chemistry:Carbon monoxide|carbon monoxide]] to form carbonyl chlorofluoride, COFCl. It will react analogously with [[Chemistry:Hexafluoroacetone|hexafluoroacetone]], (CF<sub>3</sub>)<sub>2</sub>CO, with a [[Chemistry:Potassium fluoride|potassium fluoride]] catalyst to produce heptafluoroisopropyl hypochlorite, (CF<sub>3</sub>)<sub>2</sub>CFOCl; with [[Chemistry:Nitrile|nitrile]]s RCN to produce RCF<sub>2</sub>NCl<sub>2</sub>; and with the sulfur oxides SO<sub>2</sub> and SO<sub>3</sub> to produce ClSO<sub>2</sub>F and ClOSO<sub>2</sub>F respectively. It will also react exothermically with compounds containing –OH and –NH groups, such as water:<ref name="Greenwood824" />
:H<sub>2</sub>O + 2 ClF ⟶ 2 HF + Cl<sub>2</sub>O

[[Chemistry:Chlorine trifluoride|Chlorine trifluoride]] (ClF<sub>3</sub>) is a volatile colourless molecular liquid which melts at −76.3&nbsp;°C and boils at 11.8&nbsp;&thinsp;°C. It may be formed by directly fluorinating gaseous chlorine or chlorine monofluoride at 200–300&nbsp;°C. One of the most reactive chemical compounds known, the list of elements it sets on fire is diverse, containing [[Chemistry:Hydrogen|hydrogen]], [[Chemistry:Potassium|potassium]], [[Chemistry:Phosphorus|phosphorus]], [[Chemistry:Arsenic|arsenic]], [[Chemistry:Antimony|antimony]], [[Chemistry:Sulfur|sulfur]], [[Chemistry:Selenium|selenium]], [[Chemistry:Tellurium|tellurium]], [[Chemistry:Bromine|bromine]], [[Chemistry:Iodine|iodine]], and powdered [[Chemistry:Molybdenum|molybdenum]], [[Chemistry:Tungsten|tungsten]], [[Chemistry:Rhodium|rhodium]], [[Chemistry:Iridium|iridium]], and [[Chemistry:Iron|iron]].  It will also ignite water, along with many substances which in ordinary circumstances would be considered chemically inert such as [[Chemistry:Asbestos|asbestos]], concrete, glass, and sand. When heated, it will even corrode [[Chemistry:Noble metal|noble metal]]s as [[Chemistry:Palladium|palladium]], [[Chemistry:Platinum|platinum]], and [[Chemistry:Gold|gold]], and even the [[Chemistry:Noble gas|noble gas]]es [[Chemistry:Xenon|xenon]] and [[Chemistry:Radon|radon]] do not escape fluorination. An impermeable fluoride layer is formed by [[Chemistry:Sodium|sodium]], [[Chemistry:Magnesium|magnesium]], [[Chemistry:Aluminium|aluminium]], [[Chemistry:Zinc|zinc]], [[Chemistry:Tin|tin]], and [[Chemistry:Silver|silver]], which may be removed by heating. [[Chemistry:Nickel|Nickel]], copper, and steel containers are usually used due to their great resistance to attack by chlorine trifluoride, stemming from the formation of an unreactive layer of metal fluoride. Its reaction with [[Chemistry:Hydrazine|hydrazine]] to form hydrogen fluoride, nitrogen, and chlorine gases was used in experimental rocket engine, but has problems largely stemming from its extreme [[Chemistry:Hypergolic propellant|hypergolicity]] resulting in ignition without any measurable delay. Today, it is mostly used in nuclear fuel processing, to oxidise [[Chemistry:Uranium|uranium]] to [[Chemistry:Uranium hexafluoride|uranium hexafluoride]] for its enriching and to separate it from [[Chemistry:Plutonium|plutonium]], as well as in the semiconductor industry, where it is used to clean [[Chemistry:Chemical vapor deposition|chemical vapor deposition]] chambers.<ref name=SIcvd>{{cite news |title=In Situ Cleaning of CVD Chambers |newspaper=Semiconductor International |date=1 June 1999 |url=http://www.semiconductor.net/article/209105-In_Situ_Cleaning_of_CVD_Chambers.php }}{{dead link|date=August 2017 |bot=InternetArchiveBot |fix-attempted=yes }}</ref> It can act as a fluoride ion donor or acceptor (Lewis base or acid), although it does not dissociate appreciably into {{chem|ClF|2|+}} and {{chem|ClF|4|-}} ions.<ref name="Greenwood828">{{harvnb|Greenwood|Earnshaw|1997|pp=828–831}}</ref>

[[Chemistry:Chlorine pentafluoride|Chlorine pentafluoride]] (ClF<sub>5</sub>) is made on a large scale by direct fluorination of chlorine with excess [[Chemistry:Fluorine|fluorine]] gas at 350&nbsp;°C and 250&nbsp;atm, and on a small scale by reacting metal chlorides with fluorine gas at 100–300&nbsp;&thinsp;°C. It melts at −103&nbsp;°C and boils at −13.1&nbsp;°C. It is a very strong fluorinating agent, although it is still not as effective as chlorine trifluoride. Only a few specific stoichiometric reactions have been characterised. [[Chemistry:Arsenic pentafluoride|Arsenic pentafluoride]] and [[Chemistry:Antimony pentafluoride|antimony pentafluoride]] form ionic adducts of the form [ClF<sub>4</sub>]<sup>+</sup>[MF<sub>6</sub>]<sup>−</sup> (M = As, Sb) and water reacts vigorously as follows:<ref name="Greenwood832">{{harvnb|Greenwood|Earnshaw|1997|pp=832–835}}</ref>
:2 H<sub>2</sub>O + ClF<sub>5</sub> ⟶ 4 HF + FClO<sub>2</sub>

The product, [[Chemistry:Chloryl fluoride|chloryl fluoride]], is one of the five known chlorine oxide fluorides. These range from the thermally unstable FClO to the chemically unreactive [[Chemistry:Perchloryl fluoride|perchloryl fluoride]] (FClO<sub>3</sub>), the other three being FClO<sub>2</sub>, F<sub>3</sub>ClO, and F<sub>3</sub>ClO<sub>2</sub>. All five behave similarly to the chlorine fluorides, both structurally and chemically, and may act as Lewis acids or bases by gaining or losing fluoride ions respectively or as very strong oxidising and fluorinating agents.<ref name="Greenwood875">{{harvnb|Greenwood|Earnshaw|1997|pp=875–880}}</ref>

=== Chlorine oxides ===
[[File:Chlorine dioxide gas and solution.jpg|thumb|right|Yellow [[Chemistry:Chlorine dioxide|chlorine dioxide]] (ClO<sub>2</sub>) gas above a solution containing chlorine dioxide.{{What?|impurity changes dark green to yellow? does not make sense.|date=March 2022}}]]
[[File:Dichlorine-heptoxide-3D-balls.png|thumb|right|Structure of [[Chemistry:Dichlorine heptoxide|dichlorine heptoxide]], Cl<sub>2</sub>O<sub>7</sub>, the most stable of the chlorine oxides]]
The [[Chemistry:Chlorine oxide|chlorine oxide]]s are well-studied in spite of their instability (all of them are endothermic compounds). They are important because they are produced when [[Earth:Chlorofluorocarbon|chlorofluorocarbon]]s undergo photolysis in the upper atmosphere and cause the destruction of the ozone layer. None of them can be made from directly reacting the elements.<ref name="Greenwood844">{{harvnb|Greenwood|Earnshaw|1997|pp=844–850}}</ref>

[[Chemistry:Dichlorine monoxide|Dichlorine monoxide]] (Cl<sub>2</sub>O) is a brownish-yellow gas (red-brown when solid or liquid) which may be obtained by reacting chlorine gas with yellow [[Chemistry:Mercury(II) oxide|mercury(II) oxide]]. It is very soluble in water, in which it is in equilibrium with [[Chemistry:Hypochlorous acid|hypochlorous acid]] (HOCl), of which it is the anhydride. It is thus an effective bleach and is mostly used to make [[Chemistry:Hypochlorite|hypochlorite]]s. It explodes on heating or sparking or in the presence of ammonia gas.<ref name="Greenwood844" />

[[Chemistry:Chlorine dioxide|Chlorine dioxide]] (ClO<sub>2</sub>) was the first chlorine oxide to be discovered in 1811 by [[Biography:Humphry Davy|Humphry Davy]]. It is a yellow paramagnetic gas (deep-red as a solid or liquid), as expected from its having an odd number of electrons: it is stable towards dimerisation due to the delocalisation of the unpaired electron. It explodes above −40&thinsp;°C as a liquid and under pressure as a gas and therefore must be made at low concentrations for wood-pulp bleaching and water treatment. It is usually prepared by reducing a [[Chemistry:Chlorate|chlorate]] as follows:<ref name="Greenwood844" />
:{{chem|ClO|3|-}} + Cl<sup>−</sup> + 2 H<sup>+</sup> ⟶ ClO<sub>2</sub> + {{sfrac|1|2}} Cl<sub>2</sub> + H<sub>2</sub>O
Its production is thus intimately linked to the redox reactions of the chlorine oxoacids. It is a strong oxidising agent, reacting with [[Chemistry:Sulfur|sulfur]], [[Chemistry:Phosphorus|phosphorus]], phosphorus halides, and potassium borohydride. It dissolves exothermically in water to form dark-green solutions that very slowly decompose in the dark. Crystalline clathrate hydrates ClO<sub>2</sub>·''n''H<sub>2</sub>O (''n'' ≈ 6–10) separate out at low temperatures. However, in the presence of light, these solutions rapidly photodecompose to form a mixture of chloric and hydrochloric acids. Photolysis of individual ClO<sub>2</sub> molecules result in the radicals ClO and ClOO, while at room temperature mostly chlorine, oxygen, and some ClO<sub>3</sub> and Cl<sub>2</sub>O<sub>6</sub> are produced. Cl<sub>2</sub>O<sub>3</sub> is also produced when photolysing the solid at −78&thinsp;°C: it is a dark brown solid that explodes below 0&thinsp;°C. The ClO radical leads to the depletion of atmospheric ozone and is thus environmentally important as follows:<ref name="Greenwood844" />
:Cl• + O<sub>3</sub> ⟶ ClO• + O<sub>2</sub>
:ClO• + O• ⟶ Cl• + O<sub>2</sub>

[[Chemistry:Chlorine perchlorate|Chlorine perchlorate]] (ClOClO<sub>3</sub>) is a pale yellow liquid that is less stable than ClO<sub>2</sub> and decomposes at room temperature to form chlorine, oxygen, and [[Chemistry:Dichlorine hexoxide|dichlorine hexoxide]] (Cl<sub>2</sub>O<sub>6</sub>).<ref name="Greenwood844" /> Chlorine perchlorate may also be considered a chlorine derivative of [[Chemistry:Perchloric acid|perchloric acid]] (HOClO<sub>3</sub>), similar to the thermally unstable chlorine derivatives of other oxoacids: examples include [[Chemistry:Chlorine nitrate|chlorine nitrate]] (ClONO<sub>2</sub>, vigorously reactive and explosive), and chlorine fluorosulfate (ClOSO<sub>2</sub>F, more stable but still moisture-sensitive and highly reactive).<ref name="Greenwood883">{{harvnb|Greenwood|Earnshaw|1997|pp=883–885}}</ref> Dichlorine hexoxide is a dark-red liquid that freezes to form a solid which turns yellow at −180&thinsp;°C: it is usually made by reaction of chlorine dioxide with oxygen. Despite attempts to rationalise it as the dimer of ClO<sub>3</sub>, it reacts more as though it were chloryl perchlorate, [ClO<sub>2</sub>]<sup>+</sup>[ClO<sub>4</sub>]<sup>−</sup>, which has been confirmed to be the correct structure of the solid. It hydrolyses in water to give a mixture of chloric and perchloric acids: the analogous reaction with anhydrous [[Chemistry:Hydrogen fluoride|hydrogen fluoride]] does not proceed to completion.<ref name="Greenwood844" />

[[Chemistry:Dichlorine heptoxide|Dichlorine heptoxide]] (Cl<sub>2</sub>O<sub>7</sub>) is the anhydride of [[Chemistry:Perchloric acid|perchloric acid]] (HClO<sub>4</sub>) and can readily be obtained from it by dehydrating it with [[Chemistry:Phosphoric acid|phosphoric acid]] at −10&thinsp;°C and then distilling the product at −35&thinsp;°C and 1&thinsp;mmHg. It is a shock-sensitive, colourless oily liquid. It is the least reactive of the chlorine oxides, being the only one to not set organic materials on fire at room temperature. It may be dissolved in water to regenerate perchloric acid or in aqueous alkalis to regenerate perchlorates. However, it thermally decomposes explosively by breaking one of the central Cl–O bonds, producing the radicals ClO<sub>3</sub> and ClO<sub>4</sub> which immediately decompose to the elements through intermediate oxides.<ref name="Greenwood844" />

===Chlorine oxoacids and oxyanions===
{| class="wikitable" style="float:right; margin-top:0; margin-left:1em; text-align:center; font-size:10pt; line-height:11pt; width:25%;"
|+ Standard reduction potentials for aqueous Cl species<ref name="Greenwood853" />
! {{nowrap|E°(couple)}}!!{{nowrap|''a''(H<sup>+</sup>) {{=}} 1}}<br>(acid)!!{{nowrap|E°(couple)}}!!{{nowrap|''a''(OH<sup>−</sup>) {{=}} 1}}<br>(base)
|-
|Cl<sub>2</sub>/Cl<sup>−</sup>||+1.358|||Cl<sub>2</sub>/Cl<sup>−</sup>||+1.358
|-
|HOCl/Cl<sup>−</sup>||+1.484||ClO<sup>−</sup>/Cl<sup>−</sup>||+0.890
|-
|{{chem|ClO|3|-}}/Cl<sup>−</sup>||+1.459||&thinsp;||&thinsp;
|-
|HOCl/Cl<sub>2</sub>||+1.630||ClO<sup>−</sup>/Cl<sub>2</sub>||+0.421
|-
|HClO<sub>2</sub>/Cl<sub>2</sub>||+1.659||&thinsp;||&thinsp;
|-
|{{chem|ClO|3|-}}/Cl<sub>2</sub>||+1.468||&thinsp;||&thinsp;
|-
|{{chem|ClO|4|-}}/Cl<sub>2</sub>||+1.277||&thinsp;||&thinsp;
|-
|HClO<sub>2</sub>/HOCl||+1.701||{{chem|ClO|2|-}}/ClO<sup>−</sup>||+0.681
|-
|&thinsp;||&thinsp;||{{chem|ClO|3|-}}/ClO<sup>−</sup>||+0.488
|-
|{{chem|ClO|3|-}}/HClO<sub>2</sub>||+1.181||{{chem|ClO|3|-}}/{{chem|ClO|2|-}}||+0.295
|-
|{{chem|ClO|4|-}}/{{chem|ClO|3|-}}||+1.201||{{chem|ClO|4|-}}/{{chem|ClO|3|-}}||+0.374
|}

Chlorine forms four oxoacids: [[Chemistry:Hypochlorous acid|hypochlorous acid]] (HOCl), [[Chemistry:Chlorous acid|chlorous acid]] (HOClO), [[Chemistry:Chloric acid|chloric acid]] (HOClO<sub>2</sub>), and [[Chemistry:Perchloric acid|perchloric acid]] (HOClO<sub>3</sub>). As can be seen from the redox potentials given in the adjacent table, chlorine is much more stable towards disproportionation in acidic solutions than in alkaline solutions:<ref name="Greenwood853" />
:{|
|-
| Cl<sub>2</sub> + H<sub>2</sub>O || {{eqm}} HOCl + H<sup>+</sup> + Cl<sup>−</sup> || ''K''<sub>ac</sub> = 4.2 × 10<sup>−4</sup> mol<sup>2</sup> l<sup>−2</sup>
|-
| Cl<sub>2</sub> + 2 OH<sup>−</sup> || {{eqm}} OCl<sup>−</sup> + H<sub>2</sub>O + Cl<sup>−</sup> || ''K''<sub>alk</sub> = 7.5 × 10<sup>15</sup> mol<sup>−1</sup> l
|}

The hypochlorite ions also disproportionate further to produce chloride and chlorate (3 ClO<sup>−</sup> {{eqm}} 2 Cl<sup>−</sup> + {{chem|ClO|3|-}}) but this reaction is quite slow at temperatures below 70&nbsp;°C in spite of the very favourable equilibrium constant of 10<sup>27</sup>. The chlorate ions may themselves disproportionate to form chloride and perchlorate (4 {{chem|ClO|3|-}} {{eqm}} Cl<sup>−</sup> + 3 {{chem|ClO|4|-}}) but this is still very slow even at 100&nbsp;°C despite the very favourable equilibrium constant of 10<sup>20</sup>. The rates of reaction for the chlorine oxyanions increases as the oxidation state of chlorine decreases. The strengths of the chlorine oxyacids increase very quickly as the oxidation state of chlorine increases due to the increasing delocalisation of charge over more and more oxygen atoms in their conjugate bases.<ref name="Greenwood853" />

Most of the chlorine oxoacids may be produced by exploiting these disproportionation reactions. Hypochlorous acid (HOCl) is highly reactive and quite unstable; its salts are mostly used for their bleaching and sterilising abilities. They are very strong oxidising agents, transferring an oxygen atom to most inorganic species. Chlorous acid (HOClO) is even more unstable and cannot be isolated or concentrated without decomposition: it is known from the decomposition of aqueous chlorine dioxide. However, [[Chemistry:Sodium chlorite|sodium chlorite]] is a stable salt and is useful for bleaching and stripping textiles, as an oxidising agent, and as a source of chlorine dioxide. Chloric acid (HOClO<sub>2</sub>) is a strong acid that is quite stable in cold water up to 30% concentration, but on warming gives chlorine and chlorine dioxide. Evaporation under reduced pressure allows it to be concentrated further to about 40%, but then it decomposes to perchloric acid, chlorine, oxygen, water, and chlorine dioxide. Its most important salt is [[Chemistry:Sodium chlorate|sodium chlorate]], mostly used to make chlorine dioxide to bleach paper pulp. The decomposition of chlorate to chloride and oxygen is a common way to produce oxygen in the laboratory on a small scale. Chloride and chlorate may comproportionate to form chlorine as follows:<ref name="Greenwood856">{{harvnb|Greenwood|Earnshaw|1997|pp=856–870}}</ref>
:{{chem|ClO|3|-}} + 5 Cl<sup>−</sup> + 6 H<sup>+</sup> ⟶ 3 Cl<sub>2</sub> + 3 H<sub>2</sub>O

Perchlorates and perchloric acid (HOClO<sub>3</sub>) are the most stable oxo-compounds of chlorine, in keeping with the fact that chlorine compounds are most stable when the chlorine atom is in its lowest (−1) or highest (+7) possible oxidation states. Perchloric acid and aqueous perchlorates are vigorous and sometimes violent oxidising agents when heated, in stark contrast to their mostly inactive nature at room temperature due to the high activation energies for these reactions for kinetic reasons. Perchlorates are made by electrolytically oxidising sodium chlorate, and perchloric acid is made by reacting anhydrous [[Chemistry:Sodium perchlorate|sodium perchlorate]] or [[Chemistry:Barium perchlorate|barium perchlorate]] with concentrated hydrochloric acid, filtering away the chloride precipitated and distilling the filtrate to concentrate it. Anhydrous perchloric acid is a colourless mobile liquid that is sensitive to shock that explodes on contact with most organic compounds, sets [[Chemistry:Hydrogen iodide|hydrogen iodide]] and [[Chemistry:Thionyl chloride|thionyl chloride]] on fire and even oxidises silver and gold. Although it is a weak ligand, weaker than water, a few compounds involving coordinated {{chem|ClO|4|-}} are known.<ref name="Greenwood856" />

{| class="wikitable"
|-
! Chlorine oxidation state
| −1
| +1
| +3
| +5
| +7
|-
! Name
| [[Chemistry:Chloride|chloride]]
| [[Chemistry:Hypochlorite|hypochlorite]]
| [[Chemistry:Chlorite|chlorite]]
| [[Chemistry:Chlorate|chlorate]]
| [[Chemistry:Perchlorate|perchlorate]]
|-
! Formula
| Cl<sup>−</sup>
| ClO<sup>−</sup>
| {{chem|ClO|2|−}}
| {{chem|ClO|3|−}}
| {{chem|ClO|4|−}}
|-
! Structure
| [[File:Chloride-ion-3D-vdW.png|50px|The chloride ion]]
| [[File:Hypochlorite-3D-vdW.png|50px|The hypochlorite ion]]
| [[File:Chlorite-3D-vdW.png|50px|The chlorite ion]]
| [[File:Chlorate-3D-vdW.png|50px|The chlorate ion]]
| [[File:Perchlorate-3D-vdW.png|50px|The perchlorate ion]]
|}

===Organochlorine compounds===
{{main|Chemistry:Organochlorine compound}}
[[File:Phosphorus pentachloride mechanism.png|thumb|upright=2.25|right|Suggested mechanism for the chlorination of a carboxylic acid by phosphorus pentachloride to form an [[Chemistry:Acyl chloride|acyl chloride]]]]
Like the other carbon–halogen bonds, the C–Cl bond is a common functional group that forms part of core [[Chemistry:Organic chemistry|organic chemistry]]. Formally, compounds with this functional group may be considered organic derivatives of the chloride anion. Due to the difference of electronegativity between chlorine (3.16) and carbon (2.55), the carbon in a C–Cl bond is electron-deficient and thus electrophilic. Chlorination modifies the physical properties of hydrocarbons in several ways: chlorocarbons are typically denser than [[Chemistry:Water|water]] due to the higher atomic weight of chlorine versus hydrogen, and aliphatic [[Chemistry:Organochloride|organochloride]]s are [[Chemistry:Alkylating agent|alkylating agent]]s because chloride is a [[Chemistry:Leaving group|leaving group]].<ref name="Ullmann">M. Rossberg et al. "Chlorinated Hydrocarbons" in ''Ullmann's Encyclopedia of Industrial Chemistry'' 2006, Wiley-VCH, Weinheim. {{doi|10.1002/14356007.a06_233.pub2}}</ref>

Alkanes and [[Chemistry:Aryl|aryl]] alkanes may be chlorinated under free-radical conditions, with UV light. However, the extent of chlorination is difficult to control: the reaction is not [[Chemistry:Regioselectivity|regioselective]] and often results in a mixture of various isomers with different degrees of chlorination, though this may be permissible if the products are easily separated. Aryl chlorides may be prepared by the Friedel-Crafts halogenation, using chlorine and a [[Chemistry:Lewis acid|Lewis acid]] catalyst.<ref name="Ullmann" /> The [[Chemistry:Haloform reaction|haloform reaction]], using chlorine and [[Chemistry:Sodium hydroxide|sodium hydroxide]], is also able to generate alkyl halides from methyl ketones, and related compounds. Chlorine adds to the multiple bonds on alkenes and alkynes as well, giving di- or tetra-chloro compounds. However, due to the expense and reactivity of chlorine, organochlorine compounds are more commonly produced by using hydrogen chloride, or with chlorinating agents such as [[Chemistry:Phosphorus pentachloride|phosphorus pentachloride]] (PCl<sub>5</sub>) or [[Chemistry:Thionyl chloride|thionyl chloride]] (SOCl<sub>2</sub>). The last is very convenient in the laboratory because all side products are gaseous and do not have to be distilled out.<ref name="Ullmann" />

Many organochlorine compounds have been isolated from natural sources ranging from bacteria to humans.<ref name="Gribble">{{cite journal | title = Naturally Occurring Organohalogen Compounds | author = Gordon W. Gribble | journal = Acc. Chem. Res. | volume = 31 | issue = 3 | pages = 141–52 | year = 1998 | doi = 10.1021/ar9701777}}</ref><ref name="Gribble99">{{cite journal | title = The diversity of naturally occurring organobromine compounds | author = Gordon W. Gribble | journal = [[Organization:Chemical Society Reviews|Chemical Society Reviews]] | volume = 28 | issue = 5 | pages = 335–46| year = 1999 | doi = 10.1039/a900201d}}</ref> Chlorinated organic compounds are found in nearly every class of biomolecules including [[Chemistry:Alkaloid|alkaloid]]s, [[Biology:Terpene|terpene]]s, [[Chemistry:Amino acid|amino acid]]s, [[Chemistry:Flavonoid|flavonoid]]s, [[Biology:Steroid|steroid]]s, and [[Chemistry:Fatty acid|fatty acid]]s.<ref name="Gribble" /><ref>{{cite journal | author = Kjeld C. Engvild | title = Chlorine-Containing Natural Compounds in Higher Plants | journal = Phytochemistry | volume = 25 | issue = 4 | pages = 7891–91 | year = 1986 | doi=10.1016/0031-9422(86)80002-4| bibcode = 1986PChem..25..781E }}</ref> Organochlorides, including [[Chemistry:Polychlorinated dibenzodioxins|dioxins]], are produced in the high temperature environment of forest fires, and dioxins have been found in the preserved ashes of lightning-ignited fires that predate synthetic dioxins.<ref>{{cite journal|author=Gribble, G. W.|year=1994|title=The Natural production of chlorinated compounds|journal=Environmental Science and Technology|volume=28|pages=310A–319A|doi=10.1021/es00056a712|issue=7|pmid=22662801|bibcode=1994EnST...28..310G}}</ref> In addition, a variety of simple chlorinated hydrocarbons including dichloromethane, chloroform, and [[Chemistry:Carbon tetrachloride|carbon tetrachloride]] have been isolated from marine algae.<ref>{{cite journal | doi = 10.1021/np50088a001 | author = Gribble, G. W. | title = Naturally occurring organohalogen compounds – A comprehensive survey | journal = Progress in the Chemistry of Organic Natural Products | year = 1996 | volume = 68 | pages = 1–423 | pmid = 8795309 | issue = 10}}</ref> A majority of the [[Chemistry:Chloromethane|chloromethane]] in the environment is produced naturally by biological decomposition, forest fires, and volcanoes.<ref>[http://www.atsdr.cdc.gov/toxprofiles/tp106-c1-b.pdf Public Health Statement – Chloromethane] {{webarchive|url=https://web.archive.org/web/20070927203426/http://www.atsdr.cdc.gov/toxprofiles/tp106-c1-b.pdf |date=2007-09-27 }}, Centers for Disease Control, Agency for Toxic Substances and Disease Registry</ref>

Some types of organochlorides, though not all, have significant toxicity to plants or animals, including humans. Dioxins, produced when organic matter is burned in the presence of chlorine, and some insecticides, such as [[Chemistry:DDT|DDT]], are persistent organic pollutants which pose dangers when they are released into the environment. For example, DDT, which was widely used to control insects in the mid 20th century, also accumulates in food chains, and causes reproductive problems (e.g., eggshell thinning) in certain bird species.<ref>{{cite book | title=Introduction to Ecotoxicology | publisher=Blackwell Science | year=1999 | isbn=978-0-632-03852-7 | page=68 | author=Connell, D.|display-authors=0}}</ref> Due to the ready homolytic fission of the C–Cl bond to create chlorine radicals in the upper atmosphere, [[Earth:Chlorofluorocarbon|chlorofluorocarbon]]s have been phased out due to the harm they do to the ozone layer.<ref name="Greenwood844" />

==Occurrence and production==
{{Main|Chemistry:Chlorine production|Physics:Chloralkali process}}
[[File:Liquid chlorine in flask.jpg|upright|thumb|Liquid chlorine analysis]]
Chlorine is too reactive to occur as the free element in nature but is very abundant in the form of its chloride salts. It is the twenty-first most abundant element in Earth's crust and makes up 126&nbsp;[[Parts per million|parts per million]] of it, through the large deposits of chloride minerals, especially [[Chemistry:Sodium chloride|sodium chloride]], that have been evaporated from water bodies. All of these pale in comparison to the reserves of chloride ions in seawater: smaller amounts at higher concentrations occur in some inland seas and underground [[Earth:Brine|brine]] wells, such as the Great Salt Lake in Utah and the [[Earth:Dead Sea|Dead Sea]] in Israel.<ref name="Greenwood795">{{harvnb|Greenwood|Earnshaw|1997|p=795}}</ref>

Small batches of chlorine gas are prepared in the laboratory by combining hydrochloric acid and [[Chemistry:Manganese dioxide|manganese dioxide]], but the need rarely arises due to its ready availability. In industry, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water. This method, the [[Physics:Chloralkali process|chloralkali process]] industrialized in 1892, now provides most industrial chlorine gas.<ref name="Greenwood798" /> Along with chlorine, the method yields [[Chemistry:Hydrogen|hydrogen]] gas and [[Chemistry:Sodium hydroxide|sodium hydroxide]], which is the most valuable product. The process proceeds according to the following [[Chemistry:Chemical equation|chemical equation]]:<ref>{{Holleman&Wiberg|page=408}}</ref>

: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<sub>2</sub>O + 2 e<sup>−</sup> → H<sub>2</sub> + 2 OH<sup>−</sup>
:Anode: 2 Cl<sup>−</sup> → Cl<sub>2</sub> + 2 e<sup>−</sup>

In diaphragm cell electrolysis, an [[Chemistry:Asbestos|asbestos]] (or polymer-fiber) diaphragm separates a cathode and an [[Physics:Anode|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/the-chlorine-universe/how-is-chlorine-produced/the-diaphragm-cell-process.aspx|title=The diaphragm cell process|publisher=Euro Chlor|access-date=2007-08-15|url-status=dead|archive-url=https://web.archive.org/web/20111111191151/http://www.eurochlor.org/the-chlorine-universe/how-is-chlorine-produced/the-diaphragm-cell-process.aspx|archive-date=2011-11-11}}</ref> The salt solution (brine) is continuously fed to the anode compartment and flows through the diaphragm to the cathode compartment, where the caustic [[Chemistry:Alkali|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 [[Chemistry:Mercury (element)|mercury]] disposal and they are more energy efficient.<ref name="Greenwood798" />

Membrane cell electrolysis employs [[Physics:Semipermeable membrane|permeable membrane]] as an [[Chemistry:Ion exchange|ion exchange]]r. Saturated sodium (or potassium) chloride solution is passed through the anode compartment, leaving at a lower [[Chemistry:Concentration|concentration]]. This method also produces very pure sodium (or potassium) hydroxide but has the disadvantage of requiring very pure brine at high concentrations.<ref name="Euro Chlor4">{{cite web|url=http://www.eurochlor.org/the-chlorine-universe/how-is-chlorine-produced/the-membrane-cell-process.aspx|title=The membrane cell process|publisher=Euro Chlor|access-date=2007-08-15|url-status=dead|archive-url=https://web.archive.org/web/20111111190719/http://www.eurochlor.org/the-chlorine-universe/how-is-chlorine-produced/the-membrane-cell-process.aspx|archive-date=2011-11-11}}</ref>

[[File:Chloralkali membrane.svg|center|upright=2.75|thumb|Membrane cell process for chloralkali production]]

In the [[Chemistry:Deacon process|Deacon process]], hydrogen chloride recovered from the production of [[Chemistry:Organochlorine compound|organochlorine compound]]s is recovered as chlorine. The process relies on oxidation using oxygen:
: 4 HCl + O<sub>2</sub> → 2 Cl<sub>2</sub> + 2 H<sub>2</sub>O

The reaction requires a catalyst. As introduced by Deacon, early catalysts were based on copper. Commercial processes, such as the Mitsui MT-Chlorine Process, have switched to chromium and ruthenium-based catalysts.<ref name="chlorine">Schmittinger, Peter ''et al.'' (2006) "Chlorine" in ''Ullmann's Encyclopedia of Industrial Chemistry'', Wiley-VCH Verlag GmbH & Co., {{doi|10.1002/14356007.a06_399.pub2}}</ref> The chlorine produced is available in cylinders from sizes ranging from 450&nbsp;g to 70&nbsp;kg, as well as drums (865&nbsp;kg), tank wagons (15&nbsp;tonnes on roads; 27–90&nbsp;tonnes by rail), and barges (600–1200&nbsp;tonnes).<ref name="Greenwood796">{{harvnb|Greenwood|Earnshaw|1997|pp=796–800}}</ref>

==Applications==
Sodium chloride is the most common chlorine compound, and is the main source of chlorine for the demand by the chemical industry. About 15000 chlorine-containing compounds are commercially traded, including such diverse compounds as chlorinated [[Chemistry:Methane|methane]], [[Chemistry:Ethane|ethane]]s, [[Chemistry:Vinyl chloride|vinyl chloride]], [[Chemistry:Polyvinyl chloride|polyvinyl chloride]] (PVC), [[Chemistry:Aluminium trichloride|aluminium trichloride]] for [[Chemistry:Catalysis|catalysis]], the chlorides of [[Chemistry:Magnesium|magnesium]], [[Chemistry:Titanium|titanium]], [[Chemistry:Zirconium|zirconium]], and [[Chemistry:Hafnium|hafnium]] which are the precursors for producing the pure form of those elements.<ref name="Greenwood792" />

Quantitatively, of all elemental chlorine produced, about 63% is used in the manufacture of organic compounds, and 18% in the manufacture of inorganic chlorine compounds.{{sfn|Greenwood|Earnshaw|1997|p=798}} About 15,000 chlorine compounds are used commercially.{{sfn|Greenwood|Earnshaw|1997|p=793}} The remaining 19% of chlorine produced is used for bleaches and disinfection products.<ref name="Greenwood796" /> The most significant of organic compounds in terms of production volume are 1,2-dichloroethane and [[Chemistry:Vinyl chloride|vinyl chloride]], intermediates in the production of [[Chemistry:Polyvinyl chloride|PVC]]. Other particularly important organochlorines are methyl chloride, methylene chloride, [[Chemistry:Chloroform|chloroform]], vinylidene chloride, [[Chemistry:Trichloroethylene|trichloroethylene]], perchloroethylene, [[Chemistry:Allyl chloride|allyl chloride]], [[Chemistry:Epichlorohydrin|epichlorohydrin]], [[Chemistry:Chlorobenzene|chlorobenzene]], [[Chemistry:Dichlorobenzene|dichlorobenzene]]s, and [[Chemistry:Trichlorobenzene|trichlorobenzene]]s. The major inorganic compounds include HCl, Cl<sub>2</sub>O, HOCl, NaClO<sub>3</sub>, chlorinated isocyanurates, [[Chemistry:Aluminium chloride|AlCl<sub>3</sub>]], [[Chemistry:Silicon tetrachloride|SiCl<sub>4</sub>]], [[Chemistry:Tin(IV) chloride|SnCl<sub>4</sub>]], [[Chemistry:Phosphorus trichloride|PCl<sub>3</sub>]], [[Chemistry:Phosphorus pentachloride|PCl<sub>5</sub>]], [[Chemistry:Phosphoryl chloride|POCl<sub>3</sub>]], [[Chemistry:Arsenic trichloride|AsCl<sub>3</sub>]], [[Chemistry:Antimony trichloride|SbCl<sub>3</sub>]], [[Chemistry:Antimony pentachloride|SbCl<sub>5</sub>]], [[Chemistry:Bismuth chloride|BiCl<sub>3</sub>]], and [[Chemistry:Zinc chloride|ZnCl<sub>2</sub>]].<ref name="Greenwood796" />

=== Sanitation, disinfection, and antisepsis ===
{{Main|Chemistry:Water chlorination|Chemistry:Bleach}}

==== Combating putrefaction ====

In France (as elsewhere), animal intestines were processed to make musical instrument strings, [[Engineering:Goldbeater's skin|Goldbeater's skin]] and other products. This was done in "gut factories" (''boyauderies''), and it was an odiferous and unhealthy process. In or about 1820, the Société d'encouragement pour l'industrie nationale offered a prize for the discovery of a method, chemical or mechanical, for separating the peritoneal membrane of animal intestines without [[Biology:Putrefaction|putrefaction]].<ref name="nbu" /><ref name="knight">{{cite book |author=Knight, Charles |page=427 |title=Arts and sciences |volume=1 |publisher=Bradbury, Evans & Co. |date=1867}}</ref> The prize was won by Antoine-Germain Labarraque, a 44-year-old French chemist and pharmacist who had discovered that Berthollet's chlorinated bleaching solutions ("''Eau de Javel''") not only destroyed the smell of putrefaction of animal tissue decomposition, but also actually retarded the decomposition.<ref name="knight" /><ref name="bouvet" />

Labarraque's research resulted in the use of chlorides and hypochlorites of lime ([[Chemistry:Calcium hypochlorite|calcium hypochlorite]]) and of sodium ([[Chemistry:Sodium hypochlorite|sodium hypochlorite]]) in the ''boyauderies.'' The same chemicals were found to be useful in the routine [[Chemistry:Disinfectant|disinfection]] and deodorization of [[Engineering:Latrine|latrine]]s, [[Engineering:Sewerage|sewers]], markets, abattoirs, anatomical theatres, and morgues.<ref name="gedeon">{{cite book |author=Gédéon, Andras |url=https://books.google.com/books?id=vw9qq7Ghjp4C |title=Science and technology in medicine |publisher=Springer |date=2006 |pages=181–82 |isbn=978-0-387-27874-2 |url-status=live |archive-url=https://web.archive.org/web/20151231220856/https://books.google.com/books?id=vw9qq7Ghjp4C&printsec=frontcover |archive-date=2015-12-31 }}</ref> They were successful in hospitals, lazarets, prisons, infirmaries (both on land and at sea), [[Biology:Magnanery|magnaneries]], stables, cattle-sheds, etc.; and they were beneficial during exhumations,<ref>{{cite book |title=On the disinfecting properties of Labarraque's preparations of chlorine |author=Labarraque, Antoine Germain |others=Translated by James Scott |date=1828 |page=8 |url=https://books.google.com/books?id=pD0XAQAAMAAJ |url-status=live |archive-url=https://web.archive.org/web/20151231220856/https://books.google.com/books?id=pD0XAQAAMAAJ&printsec=frontcover |archive-date=2015-12-31 }}</ref> embalming, outbreaks of epidemic disease, fever, and [[Medicine:Blackleg (disease)|blackleg]] in cattle.<ref name="nbu">{{cite book |editor=Hoefer, Jean Chrétien Ferdinand |contribution=Labarraque, Antoine-Germain |title=Nouvelle biographie universelle |volume=28 |pages=323–24 |ol=24229911M }}</ref>

==== Disinfection ====
Labarraque's chlorinated lime and soda solutions have been advocated since 1828 to prevent infection (called "contagious infection", presumed to be transmitted by "miasmas"), and to treat [[Biology:Putrefaction|putrefaction]] of existing wounds, including septic wounds.<ref>Scott, James (trans.). ''[https://books.google.com/books?id=pD0XAQAAMAAJ On the disinfecting properties of Labarraque's preparations of chlorine] {{webarchive|url=https://web.archive.org/web/20151231220856/https://books.google.com/books?id=pD0XAQAAMAAJ&printsec=frontcover |date=2015-12-31 }}'' (S. Highley, 1828) Accessed Nov 1, 2011.</ref> In his 1828 work, Labarraque recommended that doctors breathe chlorine, wash their hands in chlorinated lime, and even sprinkle chlorinated lime about the patients' beds in cases of "contagious infection". In 1828, the contagion of infections was well known, even though the agency of the [[Biology:Microorganism|microbe]] was not discovered until more than half a century later.

During the Paris cholera outbreak of 1832, large quantities of so-called ''chloride of lime'' were used to disinfect the capital. This was not simply modern [[Chemistry:Calcium chloride|calcium chloride]], but chlorine gas dissolved in lime-water (dilute [[Chemistry:Calcium hydroxide|calcium hydroxide]]) to form [[Chemistry:Calcium hypochlorite|calcium hypochlorite]] (chlorinated lime). Labarraque's discovery helped to remove the terrible stench of decay from hospitals and dissecting rooms, and by doing so, effectively deodorised the Latin Quarter of Paris.<ref name="corbin">Corbin, Alain (1988). ''[https://books.google.com/books?id=LI1M4sLcvPAC The Foul and the Fragrant: Odor and the French Social Imagination] {{webarchive|url=https://web.archive.org/web/20151231220856/https://books.google.com/books?id=LI1M4sLcvPAC&printsec=frontcover |date=2015-12-31 }}''. Harvard University Press. pp. 121–22.</ref> These "putrid miasmas" were thought by many to cause the spread of "contagion" and "infection" – both words used before the germ theory of infection. Chloride of lime was used for destroying odors and "putrid matter". One source claims chloride of lime was used by Dr. John Snow to disinfect water from the cholera-contaminated well that was feeding the Broad Street pump in 1854 London,<ref>{{cite book|publisher=Wiley|year=2010|chapter-url=http://samples.sainsburysebooks.co.uk/9780470561324_sample_411014.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://samples.sainsburysebooks.co.uk/9780470561324_sample_411014.pdf |archive-date=2022-10-09 |url-status=live|title=White's Handbook of Chlorination and Alternative Disinfectants|place=Hoboken, NJ|page=452|author=Lewis, Kenneth A.|doi=10.1002/9780470561331.ch9|isbn=978-0-470-56133-1|chapter=Ch. 9 Hypochlorination – Sodium Hypochlorite}}{{dead link|date=December 2017 |bot=KAP03 |fix-attempted=yes }}</ref> though three other reputable sources that describe that famous cholera epidemic do not mention the incident.<ref name="Vinten">Vinten-Johansen, Peter, Howard Brody, Nigel Paneth, Stephen Rachman and Michael Rip. (2003). ''Cholera, Chloroform, and the Science of Medicine''. New York:Oxford University.</ref><ref>Hemphill, Sandra. (2007). ''The Strange Case of the Broad Street Pump: John Snow and the Mystery of Cholera''. Los Angeles:University of California</ref><ref>Johnson, Steven. (2006). ''The Ghost Map: The Story of London's Most Terrifying Epidemic and How It Changed Science, Cities, and the Modern World''. New York :Riverhead Books</ref> One reference makes it clear that chloride of lime was used to disinfect the offal and filth in the streets surrounding the Broad Street pump – a common practice in mid-nineteenth century England.<ref name="Vinten" />{{rp|296}}

==== Semmelweis and experiments with antisepsis ====
[[File:Semmelweis Ignác.jpg|thumb|Ignaz Semmelweis|left|upright=0.6 ]]
Perhaps the most famous application of Labarraque's chlorine and [[Chemistry:Base|chemical base]] solutions was in 1847, when Ignaz Semmelweis used chlorine-water (chlorine dissolved in pure water, which was cheaper than chlorinated lime solutions) to disinfect the hands of Austrian doctors, which Semmelweis noticed still carried the stench of decomposition from the dissection rooms to the patient examination rooms. Long before the germ theory of disease, Semmelweis theorized that "cadaveric particles" were transmitting decay from fresh medical cadavers to living patients, and he used the well-known "Labarraque's solutions" as the only known method to remove the smell of decay and tissue decomposition (which he found that soap did not). The solutions proved to be far more effective antiseptics than soap (Semmelweis was also aware of their greater efficacy, but not the reason), and this resulted in Semmelweis's celebrated success in stopping the transmission of childbed fever ("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 |access-date=2008-07-10 |url-status=bot: unknown |archive-url=https://web.archive.org/web/20110429214443/http://www.americanchemistry.com/s_chlorine/sec_content.asp?CID=1166&DID=4476&CTYPEID=109 |archive-date=2011-04-29 }}</ref>

Much later, during World War I in 1916, a standardized and diluted modification of Labarraque's solution containing hypochlorite (0.5%) and boric acid as an acidic stabilizer was developed by [[Biography:Henry Drysdale Dakin|Henry Drysdale Dakin]] (who gave full credit to Labarraque's prior work in this area). Called [[Chemistry:Dakin's solution|Dakin's solution]], the method of wound irrigation with chlorinated solutions allowed antiseptic treatment of a wide variety of open wounds, long before the modern antibiotic era. A modified version of this solution continues to be employed in wound irrigation in modern times, where it remains effective against bacteria that are resistant to multiple antibiotics (see Century Pharmaceuticals).<ref>{{Cite journal | last1 = Rezayat | first1 = C. | last2 = Widmann | first2 = W. D. | last3 = Hardy | first3 = M. A. | doi = 10.1016/j.cursur.2006.04.009 | title = Henry Drysdale Dakin: More Than His Solution | journal = Current Surgery | volume = 63 | issue = 3 | pages = 194–96 | year = 2006 | pmid = 16757372}}</ref>
{{clear}}

==== Public sanitation ====
[[File:Liquid Pool Chlorine.jpg|thumb|Liquid Pool Chlorine]]

The first continuous application of chlorination to drinking U.S. water was installed in Jersey City, New Jersey, in 1908.<ref>Joseph Cotruvo, Victor Kimm, Arden Calvert. [http://www.epaalumni.org/hcp/drinkingwater.pdf “Drinking Water: A Half Century of Progress.”] EPA Alumni Association. March 1, 2016.</ref> By 1918, the US Department of Treasury called for all drinking water to be disinfected with chlorine. Chlorine is presently an important chemical for [[Earth:Water purification|water purification]] (such as in water treatment plants), in [[Chemistry:Disinfectant|disinfectant]]s, and in [[Chemistry:Bleach|bleach]]. 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|edition = 81st| publisher =CRC press| date = 2000| isbn = 978-0-8493-0481-1}}</ref>

Chlorine is usually used (in the form of [[Chemistry:Hypochlorous acid|hypochlorous acid]]) to kill [[Biology:Bacteria|bacteria]] and other microbes in drinking water supplies and public swimming pools. In most private swimming pools, chlorine itself is not used, but rather [[Chemistry:Sodium hypochlorite|sodium hypochlorite]], formed from chlorine and [[Chemistry:Sodium hydroxide|sodium hydroxide]], or solid tablets of chlorinated isocyanurates. The drawback of using chlorine in swimming pools is that the chlorine reacts with the [[Chemistry:Amino acid|amino acid]]s in proteins in human hair and skin. Contrary to popular belief, the distinctive "chlorine aroma" associated with swimming pools is not the result of elemental chlorine itself, but of [[Chemistry:Monochloramine|chloramine]], a chemical compound produced by the reaction of free dissolved chlorine with amines in organic substances including those in urine and sweat.<ref>{{cite web|title=Chloramines & Pool Operation|url=https://www.cdc.gov/healthywater/swimming/aquatics-professionals/chloramines.html|access-date=13 March 2022|publisher=Centres for Disease Control and Prevention}}</ref> As a disinfectant in water, chlorine is more than three times as effective against ''[[Biology:Escherichia coli|Escherichia coli]]'' as [[Chemistry:Bromine|bromine]], and more than six times as effective as [[Chemistry:Iodine|iodine]].<ref>{{cite journal|author=Koski T. A. |author2=Stuart L. S. |author3=Ortenzio L. F.|title=Comparison of chlorine, bromine, iodine as disinfectants for swimming pool water|journal = Applied Microbiology|volume = 14|issue = 2|date = 1966|pages = 276–79|pmid=4959984|pmc=546668|doi=10.1128/AEM.14.2.276-279.1966 }}</ref> Increasingly, [[Chemistry:Monochloramine|monochloramine]] itself is being directly added to drinking water for purposes of disinfection, a process known as [[Chemistry:Chloramination|chloramination]].{{r|CDC-Chloramination}}

It is often impractical to store and use poisonous chlorine gas for water treatment, so alternative methods of adding chlorine are used. These include [[Chemistry:Hypochlorite|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 [[Biology:Biocide|biocide]], killing germs, microorganisms, algae, and so on.{{sfn|Greenwood|Earnshaw|1997|p=860}}<ref>{{Holleman&Wiberg|page=411}}</ref>

=== Use as a weapon ===

==== World War I ====

Chlorine gas, also known as bertholite, was first [[Social:Chemical warfare|used as a weapon]] in [[Social:World War I|World War I]] by Germany on April 22, 1915, in the Second Battle of Ypres.<ref>"Battle of Ypres" ''The Canadian Encyclopedia''</ref><ref name="CEN2015">{{cite journal | first = Sarah | last = Everts | url = http://chemicalweapons.cenmag.org/when-chemicals-became-weapons-of-war/ | title = When Chemicals Became Weapons of War | journal = Chemical & Engineering News | volume = 93 | issue = 8 | date = February 23, 2015 | url-status = live | archive-url = https://web.archive.org/web/20160330203223/http://chemicalweapons.cenmag.org/when-chemicals-became-weapons-of-war/ | archive-date = March 30, 2016 }}</ref> As described by the soldiers, it had the distinctive smell of a mixture of pepper and pineapple.{{cn|date=August 2022}} It also tasted metallic and stung the back of the throat and chest. Chlorine reacts with water in the [[Biology:Mucous membrane|mucosa]] of the lungs to form [[Chemistry:Hydrochloric acid|hydrochloric acid]], destructive to living tissue and potentially lethal. Human respiratory systems can be protected from chlorine gas by [[Engineering:Gas mask|gas mask]]s with activated charcoal or other filters, which makes chlorine gas much less lethal than other chemical weapons. It was pioneered by a German scientist later to be a Nobel laureate, [[Biography:Fritz Haber|Fritz Haber]] of the Kaiser Wilhelm Institute in Berlin, in collaboration with the German chemical conglomerate IG Farben, which developed methods for discharging chlorine gas against an [[Engineering:Trench|entrenched]] enemy.<ref>{{Cite book| url=https://books.google.com/books?id=G9FljcEASycC&pg=PA226| page=226| title=Enriching the Earth: Fritz Haber, Carl Bosch, and the Transformation of World Food Production| isbn=978-0-262-69313-4| author1=Smil, Vaclav| year=2000| url-status=live| archive-url=https://web.archive.org/web/20151231220856/https://books.google.com/books?id=G9FljcEASycC&pg=PA226| archive-date=2015-12-31}}</ref> After its first use, both sides in the conflict used chlorine as a chemical weapon, but it was soon replaced by the more deadly [[Chemistry:Phosgene|phosgene]] and [[Chemistry:Mustard gas|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|access-date=2007-08-12|url-status=live|archive-url=https://web.archive.org/web/20070821004525/http://www.firstworldwar.com/weaponry/gas.htm|archive-date=2007-08-21}}</ref>

==== Middle east ====

Chlorine gas was also used during the Iraq War in Anbar Province in 2007, with insurgents packing truck bombs with mortar shells and chlorine tanks. The attacks killed two people from the explosives and sickened more than 350. 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. In some bombings, over a hundred civilians were hospitalized due to breathing difficulties. The Iraqi authorities tightened security for elemental chlorine, which is essential for providing safe drinking water to the population.<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|access-date=2007-03-17|title=Iraq gas attack makes hundreds ill|url-status=live|archive-url=https://web.archive.org/web/20070317201051/http://www.cnn.com/2007/WORLD/meast/03/17/iraq.main/index.html|archive-date=2007-03-17}}</ref><ref name="Chlorine bomb BBC">{{cite news|url = http://news.bbc.co.uk/2/hi/middle_east/6660585.stm|work = BBC News|date = 2007-05-17|access-date = 2007-05-17|title = 'Chlorine bomb' hits Iraq village|url-status = dead|archive-url = https://web.archive.org/web/20070526010713/http://news.bbc.co.uk/2/hi/middle_east/6660585.stm|archive-date = 2007-05-26}}</ref>

On 23 October 2014, it was reported that the [[Religion:Islamic State of Iraq and the Levant|Islamic State of Iraq and the Levant]] had used chlorine gas in the town of Duluiyah, Iraq.<ref>{{cite news|url=https://www.washingtonpost.com/world/middle_east/islamic-state-militants-allegedly-used-chlorine-gas-against-iraqi-security-forces/2014/10/23/c865c943-1c93-4ac0-a7ed-033218f15cbb_story.html|title=Islamic State militants allegedly used chlorine gas against Iraqi security forces|newspaper={{wipe|The Washington Post}}|first=Loveday|last=Morris|date=2014-10-23|access-date=2021-06-08}}</ref> Laboratory analysis of clothing and soil samples confirmed the use of chlorine gas against Kurdish Peshmerga Forces in a vehicle-borne improvised explosive device attack on 23 January 2015 at the Highway 47 Kiske Junction near Mosul.<ref>{{cite news|url=https://www.dropbox.com/s/efilp3g8mxwiev4/Lab-Report-Kisic-Junction-Chlorine.pdf?dl=0|title=Lab report on chlorine gas usage|publisher=Kurdistan Region Security Council |date=March 14, 2015 }}</ref>

Another country in the middle east, [[Place:Syria|Syria]], has used chlorine as a [[Chemistry:Chemical weapon|chemical weapon]]<ref>{{cite news |url=https://www.nytimes.com/2017/02/13/world/middleeast/syrian-chlorine-bombs-aleppo-human-rights-watch.html?_r=0 |title=Syria Used Chlorine Bombs Systematically in Aleppo, Report Says |access-date=2017-05-10 |url-status=live |archive-url=https://web.archive.org/web/20170515042507/https://www.nytimes.com/2017/02/13/world/middleeast/syrian-chlorine-bombs-aleppo-human-rights-watch.html?_r=0 |archive-date=2017-05-15 |newspaper=The New York Times |date=2017-02-13 |last1=Gladstone |first1=Rick }}</ref> delivered from [[Engineering:Barrel bomb|barrel bomb]]s and rockets.<ref>{{cite news |url=https://www.bbc.com/news/world-middle-east-37291182 |title=Syrian forces 'drop chlorine' on Aleppo |access-date=2017-05-10 |url-status=live |archive-url=https://web.archive.org/web/20170513191745/http://www.bbc.com/news/world-middle-east-37291182 |archive-date=2017-05-13 |work=BBC News |date=2016-09-07 }}</ref><ref>{{cite web |url=http://www.foxnews.com/world/2017/03/06/ignoring-un-russia-and-assad-continue-syrian-chemical-weapons-and-bombing-attacks-labeled-war-crimes.html |title=Ignoring UN, Russia and Assad continue Syrian chemical weapons and bombing attacks labeled war crimes |website=Fox News |access-date=2017-05-11 |url-status=live |archive-url=https://web.archive.org/web/20170425092749/http://www.foxnews.com/world/2017/03/06/ignoring-un-russia-and-assad-continue-syrian-chemical-weapons-and-bombing-attacks-labeled-war-crimes.html |archive-date=2017-04-25 |date=2017-03-06 }}</ref> In 2016, the OPCW-UN Joint Investigative Mechanism concluded that the Syrian government used chlorine as a chemical weapon in three separate attacks.<ref>[https://www.reuters.com/article/us-mideast-crisis-syria-opcw-timeline-idUSKBN1HG1M7 "Timeline of investigations into Syria's chemical weapons"]. Reuters. April 9, 2018.</ref> Later investigations from the OPCW's Investigation and Identification Team concluded that the {{wipe|Syrian Air Force}} was responsible for chlorine attacks in 2017 and 2018.<ref>[https://www.bbc.com/news/world-middle-east-56721409 "Syrian air force behind 2018 chlorine attack on Saraqeb, OPCW finds"] BBC News. April 12, 2021.</ref>

==Biological role==
The [[Chemistry:Chloride|chloride]] anion is an [[Chemistry:Mineral (nutrient)|essential nutrient]] for metabolism. Chlorine is needed for the production of [[Chemistry:Hydrochloric acid|hydrochloric acid]] in the stomach and in cellular pump functions.<ref>{{cite web|url=http://www.med.umich.edu/1libr/aha/aha_schlorid_crs.htm |access-date=30 April 2010 |url-status=dead |archive-url=https://web.archive.org/web/20090331211149/http://med.umich.edu/1libr/aha/aha_schlorid_crs.htm |archive-date=31 March 2009 |title=Blood (Serum) Chloride Level Test}}</ref> The main dietary source is table salt, or sodium chloride. Overly low or high concentrations of chloride in the blood are examples of electrolyte disturbances. [[Medicine:Hypochloremia|Hypochloremia]] (having too little chloride) rarely occurs in the absence of other abnormalities. It is sometimes associated with hypoventilation.<ref name="pmid3764530">{{cite journal |last1=Lavie|first1=CJ| last2=Crocker|first2=EF|last3=Key|first3=KJ |last4=Ferguson|first4=TG |title=Marked hypochloremic metabolic alkalosis with severe compensatory hypoventilation |journal=South. Med. J. |volume=79 |issue=10 |pages=1296–99 |date=October 1986 |pmid=3764530 |doi= 10.1097/00007611-198610000-00025}}</ref> It can be associated with chronic respiratory acidosis.<ref name="pmid13611033">{{cite journal |last1=Levitin|first1=H |last2=Branscome|first2=W |last3=Epstein|first3=FH |title=The pathogenesis of hypochloremia in respiratory acidosis |journal=J. Clin. Invest. |volume=37 |issue=12 |pages=1667–75 |date=December 1958 |pmid=13611033 |pmc=1062852 |doi=10.1172/JCI103758}}</ref> [[Medicine:Hyperchloremia|Hyperchloremia]] (having too much chloride) usually does not produce symptoms. When symptoms do occur, they tend to resemble those of [[Medicine:Hypernatremia|hypernatremia]] (having too much [[Chemistry:Sodium|sodium]]). Reduction in blood chloride leads to cerebral dehydration; symptoms are most often caused by rapid rehydration which results in [[Medicine:Cerebral edema|cerebral edema]]. Hyperchloremia can affect oxygen transport.<ref name="pmid9760315">{{cite journal |last1=Cambier|first1=C |last2=Detry|first2=B |last3=Beerens|first3=D|last4=Florquin|display-authors=3 |title=Effects of hyperchloremia on blood oxygen binding in healthy calves |journal=J. Appl. Physiol. |volume=85 |issue=4 |pages=1267–72 |date=October 1998 |pmid=9760315 |doi= 10.1152/jappl.1998.85.4.1267}}</ref>

==Hazards==
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Chlorine is a toxic gas that attacks the respiratory system, eyes, and skin.<ref>{{Cite web|url=http://www.bt.cdc.gov/agent/chlorine/basics/facts.asp |title=Facts About Chlorine |website=www.bt.cdc.gov |access-date=2016-04-12 |url-status=bot: unknown |archive-url=https://web.archive.org/web/20160423193302/http://www.bt.cdc.gov/agent/chlorine/basics/facts.asp |archive-date=2016-04-23 }}</ref> Because it is denser 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 |archive-url=https://web.archive.org/web/20070926183417/http://www.westlake.com/datasheets/MSDS_Chlorine.pdf |archive-date= 2007-09-26 |title=Chlorine MSDS|date=1997-10-23}}</ref><ref>{{cite web|url=http://cameochemicals.noaa.gov/chemical/2862|title=Chlorine|author=NOAA Office of Response and Restoration, US GOV|work=noaa.gov|access-date=25 August 2015|url-status=live|archive-url=https://web.archive.org/web/20151015212302/http://cameochemicals.noaa.gov/chemical/2862|archive-date=15 October 2015}}</ref>

Chlorine is detectable with measuring devices in concentrations as low as 0.2 parts per million (ppm), and by smell at 3&nbsp;ppm. Coughing and vomiting may occur at 30&nbsp;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="Greenwood792" /> The IDLH (immediately dangerous to life and health) concentration is 10&nbsp;ppm.<ref name="NIOSH">{{PGCH|0115}}</ref> 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|date = 2001 |doi = 10.1006/enrs.2000.4110 |title = The Toxicology of Chlorine|author = Winder, Chris|pages = 105–14|pmid = 11161660|bibcode = 2001ER.....85..105W }}</ref> When chlorine is inhaled at concentrations greater than 30&nbsp;ppm, it reacts with water within the lungs, producing [[Chemistry:Hydrochloric acid|hydrochloric acid]] (HCl) and [[Chemistry:Hypochlorous acid|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. Other materials present in the water may generate disinfection by-products that are associated with negative effects on human health.<ref>{{cite web|url=http://news.aces.illinois.edu/news/whats-your-water-disinfectants-create-toxic-products|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|access-date=2009-03-31|url-status=live|archive-url=https://web.archive.org/web/20140903101530/http://news.aces.illinois.edu/news/whats-your-water-disinfectants-create-toxic-products|archive-date=2014-09-03}}</ref><ref>{{cite journal |doi=10.1016/j.mrrev.2007.09.001 |title=Occurrence, genotoxicity, and carcinogenicity of regulated and emerging disinfection by-products in drinking water: A review and roadmap for research |date=2007 |last1=Richardson |first1=Susan D. |last2=Plewa |first2=Michael J. |last3=Wagner |first3=Elizabeth D. |last4=Schoeny |first4=Rita |last5=DeMarini |first5=David M. |journal=Mutation Research/Reviews in Mutation Research |volume=636 |pages=178–242 |pmid=17980649 |issue=1–3}}</ref>

In the United States, the [[Organization:Occupational Safety and Health Administration|Occupational Safety and Health Administration]] (OSHA) has set the [[Earth:Permissible exposure limit|permissible exposure limit]] for elemental chlorine at 1&nbsp;ppm, or 3&nbsp;mg/m<sup>3</sup>. The [[Organization:National Institute for Occupational Safety and Health|National Institute for Occupational Safety and Health]] has designated a [[Earth:Recommended exposure limit|recommended exposure limit]] of 0.5&nbsp;ppm over 15 minutes.<ref name="NIOSH" />

In the home, accidents occur when hypochlorite bleach solutions come into contact with certain acidic drain-cleaners to produce chlorine gas.<ref>{{Cite web|url=https://www.forbes.com/sites/alexberezow/2013/11/04/why-you-should-never-mix-different-drain-cleaners/#1d08c8fc32c8|title=Why You Should Never Mix Different Drain Cleaners|last=Berezow|first=Alex|website=Forbes|access-date=2016-04-12|url-status=live|archive-url=https://web.archive.org/web/20160425184357/http://www.forbes.com/sites/alexberezow/2013/11/04/why-you-should-never-mix-different-drain-cleaners/#1d08c8fc32c8|archive-date=2016-04-25}}</ref> Hypochlorite bleach (a popular laundry additive) combined with [[Chemistry:Ammonia|ammonia]] (another popular laundry additive) produces chloramines, another toxic group of chemicals.<ref>{{Cite web|url=http://www.doh.wa.gov/YouandYourFamily/HealthyHome/Contaminants/BleachMixingDangers|title=Bleach Mixing Dangers : Washington State Dept. of Health|website=www.doh.wa.gov|access-date=2016-04-12|url-status=live|archive-url=https://web.archive.org/web/20160414074554/http://www.doh.wa.gov/YouandYourFamily/HealthyHome/Contaminants/BleachMixingDangers|archive-date=2016-04-14}}</ref>

=== Chlorine-induced cracking in structural materials ===
Chlorine is widely used for purifying water, especially potable water supplies and water used in swimming pools. Several catastrophic collapses of swimming pool ceilings have occurred from chlorine-induced [[Chemistry:Stress corrosion cracking|stress corrosion cracking]] of [[Chemistry:Stainless steel|stainless steel]] suspension rods.<ref>{{Cite book|url=https://books.google.com/books?id=cEmq232h1zcC&pg=PA148|page=148|title=Corrosion of steel in concrete: prevention, diagnosis, repair|author=Bertolini, Luca|author2=Elsener, Bernhard|author3=Pedeferri, Pietro|author4=Polder, Rob B.|publisher=Wiley-VCH|date=2004|isbn=978-3-527-30800-2}}</ref> Some [[Chemistry:Polymer|polymer]]s are also sensitive to attack, including acetal resin and [[Chemistry:Polybutene|polybutene]]. Both materials were used in hot and cold water domestic plumbing, and [[Chemistry:Stress corrosion cracking|stress corrosion cracking]] caused widespread failures in the US in the 1980s and 1990s.<ref>{{cite book|author=Lewis, P.R.|title=Polymer Product Failure|url=https://books.google.com/books?id=SFm8rKotSaUC&pg=PA19|access-date=2011-04-30|date=1 January 2000|publisher=iSmithers Rapra Publishing|isbn=978-1-85957-192-7|pages=19–|url-status=live|archive-url=https://web.archive.org/web/20130510041524/http://books.google.com/books?id=SFm8rKotSaUC&pg=PA19|archive-date=10 May 2013}}</ref>

<!-- Don't move this picture from this subsection or delete it, as this subsection refers to it. -->
[[File:Chlorine attack1.jpg|thumb|center|Chlorine "attack" on an acetal resin plumbing joint resulting from a fractured acetal joint in a water supply system which started at an injection molding defect in the joint and slowly grew until the part failed; the fracture surface shows iron and calcium salts that were deposited in the leaking joint from the water supply before failure and are the indirect result of the chlorine attack]]

=== Chlorine-iron fire ===
The element [[Chemistry:Iron|iron]] can combine with chlorine at high temperatures in a strong exothermic reaction, creating a ''chlorine-iron fire''.<ref name="bayerdatasheet2008">{{cite web|url=https://tecci.bayer.de/io-tra-pro/emea/de/docId-2857612/Chlorine.pdf |title=Chlorine: Product Datasheet |publisher=Bayer MaterialScience AG |date=2008-04-21 |access-date=2013-12-17 |url-status=dead |archive-url=https://web.archive.org/web/20120915114003/https://tecci.bayer.de/io-tra-pro/emea/de/docId-2857612/Chlorine.pdf |archive-date=September 15, 2012 }}</ref><ref name="sanders2004">{{Cite book | isbn = 978-0-7506-7749-3 | title = Chemical Process Safety: Learning from Case Histories, 3rd Revised edition | last1 = Sanders | first1 = Roy E. | year = 2004 | publisher = Elsevier Science & Technology | location = Oxford | page = 92 }}</ref> Chlorine-iron fires are a risk in chemical process plants, where much of the pipework that carries chlorine gas is made of steel.<ref name="bayerdatasheet2008" /><ref name="sanders2004" />

== See also ==
* [[Biology:Chlorine cycle|Chlorine cycle]]
* [[Physics:Chlorine gas poisoning|Chlorine gas poisoning]]
* [[Chemistry:Industrial gas|Industrial gas]]
* [[Chemistry:Polymer degradation|Polymer degradation]]
* [[Biology:Reductive dechlorination|Reductive dechlorination]]

== References ==
{{reflist | 30em | refs =

<ref name = CDC-Chloramination>
{{cite web
 | url           = https://www.cdc.gov/healthywater/drinking/public/chloramine-disinfection.html
 | title         = Disinfection with chloramine
 | website       = Centers for Disease Control and Prevention (CDC)
 | location      = Atlanta, Georgia
 | access-date   = 2019-01-20
 | archive-url   = https://web.archive.org/web/20190120152952/https://www.cdc.gov/healthywater/drinking/public/chloramine-disinfection.html
 | url-status      = live
 | archive-date  = 2019-01-20
}}</ref>

}}

== Explanatory notes ==
{{reflist|group=en}}

== General bibliography ==

== External links ==
* [http://www.periodicvideos.com/videos/017.htm Chlorine] at ''The Periodic Table of Videos'' (University of Nottingham)
* Agency for Toxic Substances and Disease Registry: [https://web.archive.org/web/20100607045216/http://www.atsdr.cdc.gov/substances/toxsubstance.asp?toxid=36 Chlorine]
* [https://web.archive.org/web/20090518141937/http://electrochem.cwru.edu/encycl/art-b01-brine.htm Electrolytic production]
* [https://web.archive.org/web/20100429205421/http://www.amazingrust.com/Experiments/how_to/Liquid_Cl2.html Production and liquefaction of chlorine]
* [https://web.archive.org/web/20080706020039/http://oceana.org/chlorine Chlorine Production Using Mercury, Environmental Considerations and Alternatives]
* [https://web.archive.org/web/20091029202733/http://www.npi.gov.au/database/substance-info/profiles/20.html National Pollutant Inventory – Chlorine]
* [https://www.cdc.gov/niosh/topics/chlorine/ National Institute for Occupational Safety and Health – Chlorine Page]
* [http://www.chlorineinstitute.org/ Chlorine Institute] – Trade association representing the chlorine industry
* [http://www.eurochlor.org/ Chlorine Online] – the web portal of Eurochlor – the business association of the European chlor-alkali industry

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[[Category:Chlorine| ]]
[[Category:Chemical elements]]
[[Category:Diatomic nonmetals]]
[[Category:Gases with color]]
[[Category:Halogens]]
[[Category:Hazardous air pollutants]]
[[Category:Industrial gases]]
[[Category:Oxidizing agents]]
[[Category:Pulmonary agents]]
[[Category:Reactive nonmetals]]
[[Category:World Health Organization essential medicines]]

{{Sourceattribution|Chlorine}}