Editing Chlorine (section)

===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]]
|}