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Chlorine perchlorate
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<h2 id="properties">Properties</h2> <p>Chlorine perchlorate is a pale greenish liquid. It is less stable than ClO2 (chlorine dioxide) and decomposes at room temperature to give O2 (<a href="/facts/Oxygen/acyn6o8r">oxygen</a>), Cl2 (<a href="/facts/Chlorine/y27UwYBM">chlorine</a>) and Cl2O6 (<a href="/facts/Dichlorine_hexoxide/KPCD0V7X">dichlorine hexoxide</a>): </p> 2 ClOClO3 → O2 + Cl2 + Cl2O6 <p>Chlorine perchlorate reacts with metal chlorides to form chlorine and the corresponding anhydrous perchlorate: </p> <a href="/facts/CrO2Cl2/RDvsSKyP">CrO2Cl2</a> + 2 ClOClO3 → 2 Cl2 + CrO2(ClO4)2 <a href="/facts/TiCl4/bGS32SsD">TiCl4</a> + 4 ClOClO3 → 4 Cl2 + Ti(ClO4)4 2 <a href="/facts/AgCl/qSbcrlqq">AgCl</a> + 2 ClOClO3 → 2 <a href="/facts/AgClO4/YX0v6ETL">AgClO4</a> + Cl2 <h2 id="reactions">Reactions</h2> <table><tbody><tr><th>Reactant</th><th>Conditions</th><th>Products</th></tr><tr><td>—</td><td>Heat</td><td><a href="/facts/Dichlorine_hexoxide/KPCD0V7X">dichlorine hexoxide</a> (80%), <a href="/facts/Chlorine_dioxide/1K7ybEdW">chlorine dioxide</a>, chlorine, oxygen</td></tr><tr><td>—</td><td><a href="/facts/Ultraviolet_light/kncBUqn5">Ultraviolet light</a></td><td><a href="/facts/Dichlorine_heptoxide/YqfYHlzE">dichlorine heptoxide</a>, chlorine, oxygen<a class="footnote-ref" id="fnref:4" href="#fn:4"><sup>4</sup></a></td></tr><tr><td><a href="/facts/Caesium_iodide/GX2FrsVy">caesium iodide</a></td><td>−45 °C</td><td>caesium tetraperchloratoiodate(III) Cs+[I(OClO3)4]−<a class="footnote-ref" id="fnref:5" href="#fn:5"><sup>5</sup></a></td></tr><tr><td>ClOSO2F or <a href="/facts/Chlorine_monofluoride/lbQjEEJu">ClF</a></td><td>—</td><td>M+ClO−4 (M = <a href="/facts/Caesium/dXKN1S5k">Cs</a> or <a href="/facts/Nitronium/ATVyweb5">[NO2]</a>)<a class="footnote-ref" id="fnref:6" href="#fn:6"><sup>6</sup></a></td></tr><tr><td><a href="/facts/Bromine/oLfj3C7D">bromine</a></td><td>−45 °C</td><td><a href="/facts/Bromine_perchlorate/Nd9Nk11x">bromine perchlorate</a> (BrOClO3)<a class="footnote-ref" id="fnref:7" href="#fn:7"><sup>7</sup></a></td></tr><tr><td><a href="/facts/Iodine/JlfMecMf">iodine</a>(0.33 <a href="/facts/Mole_(unit)/kCMFwCc0">mol</a>)</td><td>−50 °C</td><td>iodine(III) perchlorate I(OClO3)3<a class="footnote-ref" id="fnref:8" href="#fn:8"><sup>8</sup></a></td></tr></tbody></table> <h2 id="notes">Notes</h2> <h2 id="references">References</h2> <ol> <li id="fn:1"><p>A. J. Schell-Sorokin; D. S. Bethune; J. R. Lankard; M. M. T. Loy; P. P. Sorokin (1982). "Chlorine perchlorate a major photolysis product of chlorine dioxide". J. Phys. Chem. 86 (24): 4653–4655. doi:10.1021/j100221a001. <a href="/wiki/Doi_(identifier)" target="_blank">/wiki/Doi_(identifier)</a> <a href="#fnref:1" class="footnote-back-ref">↩</a></p></li> <li id="fn:2"><p>M. I. Lopez; J. E. Sicre (1988). "Ultraviolet spectrum of chlorine perchlorate". J. Phys. Chem. 92 (2): 563–564. doi:10.1021/j100313a062. <a href="/wiki/Doi_(identifier)" target="_blank">/wiki/Doi_(identifier)</a> <a href="#fnref:2" class="footnote-back-ref">↩</a></p></li> <li id="fn:3"><p>Rao, Balaji; Anderson, Todd A.; Redder, Aaron; Jackson, W. Andrew (2010-04-15). "Perchlorate Formation by Ozone Oxidation of Aqueous Chlorine/Oxy-Chlorine Species: Role of ClxOy Radicals". Environmental Science & Technology. 44 (8): 2961–2967. Bibcode:2010EnST...44.2961R. doi:10.1021/es903065f. ISSN 0013-936X. PMID 20345093. <a href="/wiki/Bibcode_(identifier)" target="_blank">/wiki/Bibcode_(identifier)</a> <a href="#fnref:3" class="footnote-back-ref">↩</a></p></li> <li id="fn:4"><p>Rao, Balaji; Anderson, Todd A.; Redder, Aaron; Jackson, W. Andrew (2010-04-15). "Perchlorate Formation by Ozone Oxidation of Aqueous Chlorine/Oxy-Chlorine Species: Role of ClxOy Radicals". Environmental Science & Technology. 44 (8): 2961–2967. Bibcode:2010EnST...44.2961R. doi:10.1021/es903065f. ISSN 0013-936X. PMID 20345093. <a href="/wiki/Bibcode_(identifier)" target="_blank">/wiki/Bibcode_(identifier)</a> <a href="#fnref:4" class="footnote-back-ref">↩</a></p></li> <li id="fn:5"><p>Cs+[I(OClO3)4]− is a pale yellow salt which is stable at room temperature. It has a square IO4 unit. <a href="/wiki/Salt_(chemistry)" target="_blank">/wiki/Salt_(chemistry)</a> <a href="#fnref:5" class="footnote-back-ref">↩</a></p></li> <li id="fn:6"><p>M+ClO−4 (M = Cs or [NO2]) reacts with BrOSO2F at −20 °C and produces bromine perchlorate (BrOClO3). Bromine perchlorate then reacts with hydrogen bromide (HBr) at −70 °C and produces elemental bromine (Br2) and perchloric acid (HClO4). <a href="/wiki/Hydrogen_bromide" target="_blank">/wiki/Hydrogen_bromide</a> <a href="#fnref:6" class="footnote-back-ref">↩</a></p></li> <li id="fn:7"><p>M+ClO−4 (M = Cs or [NO2]) reacts with BrOSO2F at −20 °C and produces bromine perchlorate (BrOClO3). Bromine perchlorate then reacts with hydrogen bromide (HBr) at −70 °C and produces elemental bromine (Br2) and perchloric acid (HClO4). <a href="/wiki/Hydrogen_bromide" target="_blank">/wiki/Hydrogen_bromide</a> <a href="#fnref:7" class="footnote-back-ref">↩</a></p></li> <li id="fn:8"><p>The last[6] attempt to form iodine monoperchlorate (IOClO3) occurred in 1972,[7] and even at low temperatures yielded instead the triperchlorate. On warming, the latter then decomposes to iodate. <a href="#fnref:8" class="footnote-back-ref">↩</a></p></li> </ol>