Editing Perchlorate (section)

==Chemical properties==
The perchlorate ion is the least [[redox]] reactive of the generalized [[:Category:chlorates|chlorates]]. Perchlorate contains [[chlorine]] in its highest [[oxidation state|oxidation number]] (+7). A table of [[reduction potential]]s of the four [[:Category:chlorates|chlorates]] shows that, contrary to expectation, perchlorate in [[aqueous solution]] is the weakest [[Oxidizing agent|oxidant]] among the four.<ref>{{Cotton&Wilkinson5th|page=564}}</ref>

{|class="wikitable"
|-
!Ion !!Acidic reaction !!''E''° (V) !!Neutral/basic reaction !!''E''° (V)
|-
|align="center" |[[Hypochlorite]] ||{{chem2|2 H+ + 2 HOCl + 2 e− → Cl2 (''g'') + 2 H2O}} ||align="center" |1.63||{{chem2|ClO− + H2O + 2 e− → Cl− + 2 OH−}} ||align="center" |0.89
|-
|align="center" |[[Chlorite]] ||{{chem2|6 H+ + 2 HOClO + 6 e− → Cl2 (''g'') + 4 H2O}} ||align="center" |1.64||{{chem2|ClO2− + 2 H2O + 4 e− → Cl− + 4 OH−}} ||align="center" |0.78
|-
|align="center" |[[Chlorate]] ||{{chem2|12 H+ + 2 ClO3− + 10 e− → Cl2 (''g'') + 6 H2O}} ||align="center" |1.47||{{chem2|ClO3− + 3 H2O + 6 e− → Cl− + 6 OH−}} ||align="center" |0.63
|-
|align="center" |Perchlorate ||{{chem2|16 H+ + 2 ClO4− + 14 e− → Cl2 (''g'') + 8 H2O}} ||align="center" |1.42||{{chem2|ClO4− + 4 H2O + 8 e− → Cl− + 8 OH−}} ||align="center" |0.56
|}

These data show that the perchlorate and chlorate are stronger oxidizers in acidic conditions than in basic conditions.

Gas phase measurements of heats of reaction (which allow computation of Δ<sub>f</sub>''H''°) of various chlorine oxides do follow the expected trend wherein [[Dichlorine heptoxide|{{chem2|Cl2O7}}]] exhibits the largest endothermic value of Δ<sub>f</sub>''H''° (238.1&nbsp;kJ/mol) while [[Dichlorine monoxide|{{chem2|Cl2O}}]] exhibits the lowest endothermic value of Δ<sub>f</sub>''H''° (80.3&nbsp;kJ/mol).<ref>Wagman, D. D.; Evans, W. H.; Parker, V. P.; Schumm, R. H.; Halow, I.; Bailey, S. M.; Churney, K. L.; Nuttall, R. L. ''J. Phys. Chem. Ref. Data'' Vol. 11(2); 1982, American Chemical Society and the American Institute of Physics.</ref>

===Weak base and weak coordinating anion===
As [[perchloric acid]] is one of the strongest mineral acids, perchlorate is a very weak [[Base (chemistry)|base]] in the sense of [[Brønsted–Lowry acid–base theory]]. 
As it is also generally a [[weakly coordinating anion]], perchlorate is commonly used as a [[supporting electrolyte|background]], or [[supporting electrolyte|supporting, electrolyte]].

===Weak oxidant in aqueous solution due to kinetic limitations===
{{Main articles|Inner sphere electron transfer|Outer sphere electron transfer}}
Perchlorate compounds oxidize organic compounds, especially when the mixture is heated. The explosive decomposition of [[ammonium perchlorate]] is catalyzed by metals and heat.<ref name="Housecroft2018">{{cite book | last1=Housecroft | first1=C.E. | last2=Sharpe | first2=A.G. | year=2018 | title=Inorganic Chemistry. 5th edition | publisher=Pearson | isbn=978-1-292-13414-7 | url=https://books.google.com/books?id=8VyjtAEACAAJ | access-date=2024-09-02 | page=1298}}</ref>

As perchlorate is a weak [[Lewis base]] (''i.e.'', a weak electron pair donor) and a weak [[Nucleophile|nucleophilic]] anion, it is also a very weakly [[Coordination complex|coordinating]] [[anion]].<ref name="Housecroft2018" /> This is why it is often used as a [[supporting electrolyte]] to study the [[Coordination complex|complexation]] and the [[Chemical species|chemical speciation]] of many [[cation]]s in [[aqueous solution]] or in [[electroanalytical methods]] ([[voltammetry]], [[electrophoresis]]…).<ref name="Housecroft2018" /> Although the perchlorate reduction is [[thermodynamics|thermodynamically]] favorable {{Nowrap|(∆''G'' < 0; ''E''° > 0)}}, and that {{chem2|ClO4-}} is expected to be a strong [[oxidant]], most often in aqueous solution, it is practically an inert species behaving as an extremely slow [[oxidant]] because of severe [[Chemical kinetics|kinetics limitations]].<ref name="Taube1953">{{cite journal | last1=Taube | first1=Henry | last2=Myers | first2=Howard | last3=Rich | first3=Ronald L. | title=Observations on the mechanism of electron transfer in solution | journal=Journal of the American Chemical Society | volume=75 | issue=16 | date=1953 | issn=0002-7863 | doi=10.1021/ja01112a546 | pages=4118–4119}}</ref><ref name="Brown2006">{{cite book | last1=Brown | first1=Gilbert M. | last2=Gu | first2=Baohua | title=Perchlorate | chapter=The Chemistry of Perchlorate in the Environment | publisher=Kluwer Academic Publishers | publication-place=Boston, MA | date=2006 | isbn=978-0-387-31114-2 | doi=10.1007/0-387-31113-0_2 | pages=17–47}}</ref> The [[metastable]] character of perchlorate in the presence of [[reducing agent|reducing]] cations such as {{Chem2|Fe(2+)}} in solution is due to the difficulty to form an [[activated complex]] facilitating the [[electron transfer]] and the exchange of oxo groups in the opposite direction. These strongly hydrated cations cannot form a sufficiently stable coordination bridge with one of the four oxo groups of the perchlorate anion. Although thermodynamically a mild reductant, {{Chem2|Fe(2+)}} ion exhibits a stronger trend to remain coordinated by water molecules to form the corresponding hexa-aquo complex in solution. The high [[activation energy]] of the cation binding with perchlorate to form a transient [[inner sphere complex]] more favourable to [[electron transfer]] considerably hinders the [[redox]] reaction.<ref name="Marcus1992_NobelLecture">{{cite web | last1=Marcus | first1=Rudolph A. | title=Electron transfer reactions in chemistry: Theory and experiment | url=https://www.nobelprize.org/uploads/2018/06/marcus-lecture.pdf | access-date=2024-09-02}}</ref> The redox reaction rate is limited by the formation of a favorable [[activated complex]] involving an oxo-bridge between the perchlorate anion and the metallic cation.<ref name="Taube1954">{{cite journal | last1=Taube | first1=Henry | last2=Myers | first2=Howard | title=Evidence for a bridged activated complex for electron transfer reactions | journal=Journal of the American Chemical Society | volume=76 | issue=8 | date=1954 | issn=0002-7863 | doi=10.1021/ja01637a020 | pages=2103–2111}}</ref> It depends on the [[molecular orbital]] rearrangement ([[HOMO and LUMO]] [[Orbital hybridisation|orbitals]]) necessary for a fast [[Transition metal oxo complex#Oxygen-atom transfer|oxygen atom transfer]] (OAT)<ref name="OAT">{{cite web | last1=Bakhtchadjian | first1=Robert | last2=Rajeev | first2=Anjana | last3=Liao | first3=Guangjian | last4=Yin | first4=Guochuan | last5=Sankaralingam | first5=Muniyandi | year=2023 | title=Oxygen Atom Transfer Reactions | publisher=Bentham Science Publishers | url=https://benthambooks.com/book/9789815050929/ | isbn=9789815050929 | access-date=2024-09-17}}</ref> and the associated electron transfer as studied experimentally by [[Henry Taube]] (1983 Nobel Prize in Chemistry)<ref name="Taube1983">{{cite web|title=Press Release: The 1983 Nobel Prize in Chemistry| url=http://nobelprize.org/nobel_prizes/chemistry/laureates/1983/press.html|publisher=NobelPrize.org The Official Website of the Nobel Prize|access-date=2024-09-02}}</ref><ref name="Taube1984">{{cite journal | last1=Taube | first1=Henry | title=Electron transfer between metal complexes: Retrospective | journal=Science | volume=226 | issue=4678 | date=1984-11-30 | issn=0036-8075 | doi=10.1126/science.6494920 | pages=1028–1036| pmid=6494920 | bibcode=1984Sci...226.1028T }}</ref> and theoretically by [[Rudolph A. Marcus]] (1992 Nobel Prize in Chemistry),<ref name="Marcus1992">{{cite web | title=The Nobel Prize in Chemistry 1992 | website=NobelPrize.org | date=1992 | url=https://www.nobelprize.org/prizes/chemistry/1992/marcus/facts/ | access-date=2024-09-02}}</ref> both awarded for their respective works on the mechanisms of electron-transfer reactions with metal complexes and in chemical systems. 

In contrast to the {{Chem2|Fe(2+)}} cations which remain unoxidized in deaerated perchlorate aqueous solutions free of dissolved oxygen, other cations such as Ru(II) and Ti(III) can form a more stable bridge between the metal centre and one of the oxo groups of {{chem2|ClO4-}}. In the [[inner sphere electron transfer]] mechanism to observe the perchlorate reduction, the {{chem2|ClO4-}} anion must quickly transfer an oxygen atom to the reducing cation.<ref name="Taube1982">{{cite book | last1=Taube | first1=Henry | editor1=Rorabacher, D. B. | editor2=Endicott, J. F. | title=Observations on Atom-Transfer Reactions. In: Mechanistic Aspects of Inorganic Reactions | publisher=American Chemical Society | publication-place=Washington, D. C. | volume=198 | date=1982-09-27 | isbn=978-0-8412-0734-9 | doi=10.1021/bk-1982-0198.ch007 | page=151}}</ref><ref name="Bakac2010">{{cite book | last1=Bakac | first1=Andreja | year=2010 | title=Physical Inorganic Chemistry: Reactions, Processes, and Applications | publisher=Wiley | pages=620 | isbn=978-0-470-60255-3 | url=https://books.google.com/books?id=dl7z0JscRTQC&pg=PR7 | access-date=2024-09-02}}</ref> When it is the case, metallic cations can readily reduce perchlorate in solution.<ref name="Taube1983" /> Ru(II) can reduce {{chem2|ClO4-}} to {{chem2|ClO3-}}, while V(II), V(III), Mo(III), Cr(II) and Ti(III) can reduce {{chem2|ClO4-}} to {{chem2|Cl-}}.<ref>Urbansky, Edward T. (1998). [https://clu-in.org/download/contaminantfocus/perchlorate/urbansky2.pdf Perchlorate Chemistry: Implications for Analysis and Remediation] {{Webarchive|url=https://web.archive.org/web/20220129071859/https://clu-in.org/download/contaminantfocus/perchlorate/urbansky2.pdf |date=29 January 2022}}</ref>
 
Some metal complexes, especially those of [[rhenium]], and some metalloenzymes can [[Catalyst|catalyze]] the reduction of perchlorate under mild conditions.<ref>{{cite journal|doi=10.1002/1521-3773(20001201)39:23<4310::AID-ANIE4310>3.0.CO;2-D |date=2000 |volume=39 |issue=23 |last1=Abu-Omar |first1=Mahdi M. |last2=McPherson |first2=Lee D. |last3=Arias |first3=Joachin |last4=Béreau |first4=Virginie M. |title=Clean and Efficient Catalytic Reduction of Perchlorate |journal=Angewandte Chemie |pages=4310–4313 |pmid=29711910 |bibcode=2000AngCh..39.4310A }}</ref> [[Perchlorate reductase]] (see below), a [[Molybdenum in biology|molybdoenzyme]], also catalyzes the reduction of perchlorate.<ref>{{cite journal |doi=10.1074/jbc.M116.714618|doi-access=free |date=2016 |volume=291 |issue=17 |last1=Youngblut |first1=Matthew D. |last2=Tsai |first2=Chi-Lin |last3=Clark |first3=Iain C. |last4=Carlson |first4=Hans K. |last5=Maglaqui |first5=Adrian P. |last6=Gau-Pan |first6=Phonchien S. |last7=Redford |first7=Steven A. |last8=Wong |first8=Alan |last9=Tainer |first9=John A. |last10=Coates |first10=John D. |title=Perchlorate Reductase is Distinguished by Active Site Aromatic Gate Residues |journal=Journal of Biological Chemistry |pages=9190–9302 |pmid=26940877 |pmc=4861485 }}</ref>  Both the Re- and Mo-based [[catalyst]]s operate via metal-oxo intermediates.

===Microbiology===
Over 40 phylogenetically and metabolically diverse microorganisms capable of growth using perchlorate as an electron acceptor<ref>{{cite journal | pmid = 19921177 | doi=10.1007/s00253-009-2336-6 | volume=86 | issue=1 | title=Description of the novel perchlorate-reducing bacteria ''Dechlorobacter hydrogenophilus'' gen. nov., sp. nov. and ''Propionivibrio militaris'', sp. nov | pmc=2822220 | year=2010 | journal=Appl Microbiol Biotechnol | pages=335–43 | author=Thrash JC, Pollock J, Torok T, Coates JD}}</ref> have been isolated since 1996. Most originate from the [[Pseudomonadota]], but others include the [[Bacillota]], ''[[Moorella perchloratireducens]]'' and ''Sporomusa'' sp., and the [[archaeon]] ''[[Archaeoglobus|Archaeoglobus fulgidus]]''.<ref name=Coates>{{cite journal | title = Microbial perchlorate reduction: rocket-fuelled metabolism |author1=John D. Coates |author2=Laurie A. Achenbach | journal = [[Nature Reviews Microbiology]] | volume = 2 | issue = 7 | pages = 569–580 | year = 2004 | pmid = 15197392| doi = 10.1038/nrmicro926|s2cid=21600794 }}</ref><ref>{{cite journal|journal=Science|date=5 April 2013|volume=340|issue=6128|pages=85–87| doi=10.1126/science.1233957|title=Archaeal (Per)Chlorate Reduction at High Temperature: An Interplay of Biotic and Abiotic Reactions|author=Martin G. Liebensteiner, Martijn W. H. Pinkse, Peter J. Schaap, Alfons J. M. Stams, Bart P. Lomans|pmid=23559251|bibcode=2013Sci...340...85L|s2cid=32634949}}</ref> With the exception of ''A. fulgidus'', microbes that grow via perchlorate reduction utilize the enzymes [[perchlorate reductase]] and [[chlorite dismutase]], which collectively take perchlorate to chloride.<ref name=Coates/> In the process, free [[oxygen]] ({{chem2|O2}}) is generated.<ref name=Coates/>