Editing Chlorine (section)

==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 [[standard electrode potential]]s of the X<sub>2</sub>/X<sup>−</sup> couples (F, +2.866&nbsp;&thinsp;V; Cl, +1.395&nbsp;V; Br, +1.087&nbsp;&thinsp;V; I, +0.615&nbsp;V; [[Astatine|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 [[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 [[hydrogen chloride]], HCl, a major chemical in industry as well as in the laboratory, both as a gas and dissolved in water as [[hydrochloric acid]]. It is often produced by burning hydrogen gas in chlorine gas, or as a byproduct of chlorinating [[hydrocarbon]]s. Another approach is to treat [[sodium chloride]] with concentrated [[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 [[benzoyl chloride]] with [[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 [[hydrogen fluoride]], since hydrogen cannot form strong [[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-chlorine bonds 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 [[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 [[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 [[caesium|Cs<sup>+</sup>]] and [[quaternary ammonium cation|{{chem|NR|4|+}}]] (R = [[methyl group|Me]], [[ethyl group|Et]], [[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 [[nitrosyl chloride]] and [[phenol]], or salts with very low [[lattice energy|lattice energies]] such as tetraalkylammonium halides. It readily protonates [[nucleophile]]s containing lone-pairs or π bonds. [[Solvolysis]], [[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 [[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 [[noble gas]]es, with the exception of [[xenon]] in the highly unstable [[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 [[bismuth]]); and having an electronegativity higher than chlorine's ([[oxygen]] and [[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 [[nitrogen]] in NCl<sub>3</sub> is bearing a negative charge, the compound is usually called [[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 [[molybdenum(V) chloride|MoCl<sub>5</sub>]] and [[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, [[carbon tetrachloride]], or an organic chloride. For instance, [[zirconium dioxide]] reacts with chlorine at standard conditions to produce [[zirconium tetrachloride]], and [[uranium trioxide]] reacts with [[hexachloropropene]] when heated under [[reflux]] to give [[uranium tetrachloride]]. The second example also involves a reduction in [[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. [[scandium chloride]] is mostly ionic, but [[aluminium chloride]] is not). [[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 [[arsenic pentafluoride]]. The trichloride anion, {{chem2|[Cl3]-}}, has also been characterised; it is analogous to [[triiodide]].<ref name="Greenwood824" />

===Chlorine fluorides===
The three fluorides of chlorine form a subset of the [[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 [[pseudohalogen|pseudohalides]] of chlorine are also known, such as [[cyanogen chloride]] (ClCN, linear), chlorine [[cyanate]] (ClNCO), chlorine [[thiocyanate]] (ClSCN, unlike its oxygen counterpart), and chlorine [[azide]] (ClN<sub>3</sub>).<ref name="Greenwood824" />

[[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 reaction of its elements at 225&nbsp;°C, though it must then be separated and purified from [[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 [[carbon monoxide]] to form carbonyl chlorofluoride, COFCl. It will react analogously with [[hexafluoroacetone]], (CF<sub>3</sub>)<sub>2</sub>CO, with a [[potassium fluoride]] catalyst to produce heptafluoroisopropyl hypochlorite, (CF<sub>3</sub>)<sub>2</sub>CFOCl; with [[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

[[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 [[hydrogen]], [[potassium]], [[phosphorus]], [[arsenic]], [[antimony]], [[sulfur]], [[selenium]], [[tellurium]], [[bromine]], [[iodine]], and powdered [[molybdenum]], [[tungsten]], [[rhodium]], [[iridium]], and [[iron]].  It will also ignite water, along with many substances which in ordinary circumstances would be considered chemically inert such as [[asbestos]], concrete, glass, and sand. When heated, it will even corrode [[noble metal]]s as [[palladium]], [[platinum]], and [[gold]], and even the [[noble gas]]es [[xenon]] and [[radon]] do not escape fluorination. An impermeable fluoride layer is formed by [[sodium]], [[magnesium]], [[aluminium]], [[zinc]], [[tin]], and [[silver]], which may be removed by heating. [[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 [[hydrazine]] to form hydrogen fluoride, nitrogen, and chlorine gases was used in experimental rocket engine, but has problems largely stemming from its extreme [[Hypergolic propellant|hypergolicity]] resulting in ignition without any measurable delay. Today, it is mostly used in nuclear fuel processing, to oxidise [[uranium]] to [[uranium hexafluoride]] for its enriching and to separate it from [[plutonium]], as well as in the semiconductor industry, where it is used to clean [[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>

[[Chlorine pentafluoride]] (ClF<sub>5</sub>) is made on a large scale by direct fluorination of chlorine with excess [[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;°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. [[Arsenic pentafluoride]] and [[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, [[chloryl fluoride]], is one of the five known chlorine oxide fluorides. These range from the thermally unstable FClO to the chemically unreactive [[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 [[chlorine dioxide]] (ClO<sub>2</sub>) gas above a solution of hydrochloric acid and sodium chlorite in water, also containing dissolved chlorine dioxide]]
[[File:Dichlorine-heptoxide-3D-balls.png|thumb|right|Structure of [[dichlorine heptoxide]], Cl<sub>2</sub>O<sub>7</sub>, the most stable of the chlorine oxides]]
The [[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 [[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>

[[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 [[mercury(II) oxide]]. It is very soluble in water, in which it is in equilibrium with [[hypochlorous acid]] (HOCl), of which it is the anhydride. It is thus an effective bleach and is mostly used to make [[hypochlorite]]s. It explodes on heating or sparking or in the presence of ammonia gas.<ref name="Greenwood844" />

[[Chlorine dioxide]] (ClO<sub>2</sub>) was the first chlorine oxide to be discovered in 1811 by [[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&nbsp;°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 [[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 [[sulfur]], [[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&nbsp;°C: it is a dark brown solid that explodes below 0&nbsp;°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>

[[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 [[dichlorine hexoxide]] (Cl<sub>2</sub>O<sub>6</sub>).<ref name="Greenwood844" /> Chlorine perchlorate may also be considered a chlorine derivative of [[perchloric acid]] (HOClO<sub>3</sub>), similar to the thermally unstable chlorine derivatives of other oxoacids: examples include [[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 [[hydrogen fluoride]] does not proceed to completion.<ref name="Greenwood844" />

[[Dichlorine heptoxide]] (Cl<sub>2</sub>O<sub>7</sub>) is the anhydride of [[perchloric acid]] (HClO<sub>4</sub>) and can readily be obtained from it by dehydrating it with [[phosphoric acid]] at −10&nbsp;°C and then distilling the product at −35&nbsp;°C and 1&nbsp;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: [[hypochlorous acid]] (HOCl), [[chlorous acid]] (HOClO), [[chloric acid]] (HOClO<sub>2</sub>), and [[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, [[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 [[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 [[sodium perchlorate]] or [[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 [[hydrogen iodide]] and [[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" /> The Table below presents typical oxidation states for chlorine element as given in the secondary schools or colleges. There are more complex chemical compounds, the structure of which can only be explained using modern quantum chemical methods, for example, cluster technetium chloride [(CH<sub>3</sub>)<sub>4</sub>N]<sub>3</sub>[Tc<sub>6</sub>Cl<sub>14</sub>], in which 6 of the 14 chlorine atoms are formally divalent, and oxidation states are fractional.<ref>{{Cite journal |first1=Konstantin E. |last1=German |first2=Kryutchkov |last2=S.V. |first3=A.F. |last3=Kuzina |first4=V.I. |last4=Spitsyn |title=Synthesis and properties of new chloride technetium clusters |journal=Doklady Chemistry |year=1986 |volume=288 |issue=2 |pages=381–384}}</ref><ref>{{Cite journal |last1=Wheeler |first1=Ralph A. |last2=Hoffmann |first2=Roald. |date=October 1986 |title=A new magic cluster electron count and metal-metal multiple bonding |url=https://pubs.acs.org/doi/abs/10.1021/ja00281a025 |journal=Journal of the American Chemical Society |language=en |volume=108 |issue=21 |pages=6605–6610 |doi=10.1021/ja00281a025 |bibcode=1986JAChS.108.6605W |issn=0002-7863 |access-date=2023-11-08 |archive-date=2023-03-10 |archive-url=https://web.archive.org/web/20230310044010/https://pubs.acs.org/doi/abs/10.1021/ja00281a025 |url-status=live }}</ref> In addition, all the above chemical regularities are valid for "normal" or close to normal conditions, while at ultra-high pressures (for example, in the cores of large planets), chlorine can form a Na3Cl compound with sodium, which does not fit into traditional concepts of chemistry.<ref>{{Cite journal |last1=Zhang |first1=Weiwei |last2=Oganov |first2=Artem R. |last3=Goncharov |first3=Alexander F. |last4=Zhu |first4=Qiang |last5=Boulfelfel |first5=Salah Eddine |last6=Lyakhov |first6=Andriy O. |last7=Stavrou |first7=Elissaios |last8=Somayazulu |first8=Maddury |last9=Prakapenka |first9=Vitali B. |last10=Konôpková |first10=Zuzana |date=2013-12-20 |title=Unexpected Stable Stoichiometries of Sodium Chlorides |url=https://www.science.org/doi/10.1126/science.1244989 |journal=Science |language=en |volume=342 |issue=6165 |pages=1502–1505 |doi=10.1126/science.1244989 |pmid=24357316 |issn=0036-8075 |access-date=2023-11-08 |archive-date=2023-09-07 |archive-url=https://web.archive.org/web/20230907071303/https://www.science.org/doi/10.1126/science.1244989 |url-status=live |arxiv=1211.3644 |bibcode=2013Sci...342.1502Z }}</ref>

{| class="wikitable"
|-
! Chlorine oxidation state
| −1
| +1
| +3
| +5
| +7
|-
! Name
| [[chloride]]
| [[hypochlorite]]
| [[chlorite]]
| [[chlorate]]
| [[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|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 [[acyl chloride]]]]
Like the other carbon–halogen bonds, the C–Cl bond is a common functional group that forms part of core [[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 reaction|Chlorination]] modifies the physical properties of hydrocarbons in several ways: chlorocarbons are typically denser than [[water]] due to the higher atomic weight of chlorine versus hydrogen, and aliphatic organochlorides are [[alkylating agent]]s because chloride is a [[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 [[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 [[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 [[Lewis acid]] catalyst.<ref name="Ullmann" /> The [[haloform reaction]], using chlorine and [[sodium hydroxide]], is also able to generate alkyl halides from methyl ketones, and related compounds. Chlorine adds to the multiple bonds on [[alkene]]s and [[alkyne]]s as well, giving di- or tetrachloro 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 [[phosphorus pentachloride]] (PCl<sub>5</sub>) or [[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 = [[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 [[alkaloid]]s, [[terpene]]s, [[amino acid]]s, [[flavonoid]]s, [[steroid]]s, and [[fatty acid]]s.<ref name="Gribble" /><ref>{{cite journal | author = Kjeld C. Engvild | title = Chlorine-Containing Natural Compounds in Higher Plants | journal = [[Phytochemistry (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 [[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 [[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 [[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 [[insecticide]]s, such as [[DDT]], are [[persistent organic pollutant]]s 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=etal}}</ref> Due to the ready homolytic fission of the C–Cl bond to create chlorine radicals in the upper atmosphere, [[chlorofluorocarbon]]s have been phased out due to the harm they do to the ozone layer.<ref name="Greenwood844" />