View source for Chlorate

{{Short description|Anion and term for chemical compounds containing it}}
{{Chembox
|   Name = Chlorate
|   ImageFile = Chlorate-2D.png
|   ImageClass = skin-invert-image
|   ImageName = Structure and bonding in the chlorate ion
|   ImageFile1 = Chlorate-3D-vdW.png
|   ImageClass1 = bg-transparent
|   ImageName1 = The chlorate ion
|   OtherNames = Chlorate(V)
| Section1 = {{Chembox Identifiers
| CASNo = 14866-68-3
| CASNo_Ref = 
| ChEBI = 49709
| ChEBI_Comment = ion
| ChEBI1 = 23109
| ChEBI1_Comment = class of salts
| ChemSpiderID  = 94578
| Gmelin = 1491
| PubChem  =  104770
|  EC_number  =  
| UNNumber  =  1461
| UNII  =  08Z8093742
| SMILES  =  O=Cl(=O)[O-]
| StdInChI=1S/ClHO3/c2-1(3)4/h(H,2,3,4)/p-1
| StdInChIKey = XTEGARKTQYYJKE-UHFFFAOYSA-M
}}
| Section2 = {{Chembox Properties
|   Formula = ClO<sub>3</sub><sup>−</sup>
|   MolarMass = 83.4512 
|   Appearance = 
|   Density = 
|   Solubility = 
|   pKa = 
  }}
| Section3 = {{Chembox Structure
|   MolShape = [[Trigonal pyramidal molecular geometry|Trigonal pyramidal]]
|   Dipole =
  }}
| Section7 = {{Chembox Hazards
|   MainHazards = oxidation agent
  }}
| Section8 = {{Chembox Related
|   OtherAnions = {{ubl|[[bromate]]|[[iodate]]}}
|   OtherCompounds = {{ubl|[[chloride]]|[[hypochlorite]]|[[chlorite]]|[[perchlorate]]}}
  }}
}}

'''Chlorate''' is the common name of the {{chem|ClO|3|−}} anion, whose [[chlorine]] atom is in the +5 [[oxidation state]]. The term can also refer to [[chemical compound]]s containing this anion, with [[Chlorate#Compounds (salts)|chlorates]] being the [[salt (chemistry)|salt]]s of [[chloric acid]]. Other [[Chlorate#Other oxyanions|oxyanions of chlorine]] can be named "chlorate" followed by a Roman numeral in parentheses denoting the oxidation state of chlorine: e.g., the {{chem|ClO|4|−}} ion commonly called [[perchlorate]] can also be called chlorate(VII).

As predicted by [[VSEPR|valence shell electron pair repulsion theory]], chlorate anions have [[trigonal pyramid (chemistry)|trigonal pyramidal structures]].

Chlorates are powerful [[oxidizing agent|oxidizers]] and should be kept away from organics or easily oxidized materials. Mixtures of chlorate salts with virtually any combustible material (sugar, [[sawdust]], charcoal, [[organic solvents]], metals, etc.) will readily [[deflagrate]]. Chlorates were once widely used in [[pyrotechnics]] for this reason, though their use has fallen due to their instability. Most pyrotechnic applications that formerly used chlorates now use the more stable [[perchlorate]]s instead.

==Structure and bonding==
The chlorate ion cannot be satisfactorily represented by just one [[Lewis structure]], since all the Cl–O bonds are the same length (1.49&nbsp;Å in [[potassium chlorate]]<ref>{{ cite journal | journal = Acta Crystallogr. B | year = 1981 | volume = 37 | pages =913–915 | title = The structure of potassium chlorate at 77 and 298 K |author1=J. Danielsen |author2=A. Hazell |author3=F. K. Larsen | issue = 4 | doi = 10.1107/S0567740881004573 }}</ref>), and the chlorine atom is [[Hypervalent molecule|hypervalent]]. Instead, it is often thought of as a hybrid of multiple [[Resonance (chemistry)|resonance structures]]:

[[File:Chlorate-resonance-2D.svg|400px|Resonance structures of the chlorate ion]]

==Preparation==

===Laboratory===
Metal chlorates can be prepared by adding [[chlorine]] to hot metal [[hydroxide]]s like [[Potassium hydroxide|KOH]]:

:{{chem2 | 3 Cl2 + 6 KOH -> 5 KCl + KClO3 + 3 H2O }}

In this reaction, chlorine undergoes [[disproportionation]], both reduction and oxidation. Chlorine, oxidation number 0, forms chloride (Cl<sup>−</sup>; oxidation number −1) and chlorate(V) ({{chem2|ClO3−}}; oxidation number +5). The reaction of cold aqueous metal hydroxides with chlorine produces the chloride and [[hypochlorite]] (oxidation number +1) instead.{{citation needed|date=July 2018}}

===Industrial===
{{Main|Chloralkali process}}

The industrial-scale synthesis for [[sodium chlorate]] starts from an aqueous [[sodium chloride]] solution (brine) rather than chlorine gas. If the electrolysis equipment allows for the mixing of the [[chlorine]] and the [[sodium hydroxide]], then the disproportionation reaction described above occurs. The heating of the reactants to 50–70&nbsp;°C is performed by the [[electric power|electrical power]] used for [[electrolysis]].{{Citation needed|date=December 2007}}

==Natural occurrence==
A 2010 study has discovered the presence of natural chlorate deposits around the world, with relatively high concentrations found in arid and hyper-arid regions.<ref>{{cite journal |last1=Rao |first1=B. |last2=Hatzinger |first2=P. B. |last3=Böhlke |first3=J. K. |last4=Sturchio |first4=N. C. |last5=Andraski |first5=B. J. |first6=F. D. |last6=Eckardt |last7=Jackson |first7=W. |title=Natural Chlorate in the Environment: Application of a New IC-ESI/MS/MS Method with a Cl<sup>18</sup>O<sub>3</sub><sup>−</sup> Internal Standard |journal=Environ. Sci. Technol. |year=2010 |volume=44 |issue=22 |pages=8429–8434 |doi=10.1021/es1024228 |pmid=20968289|bibcode=2010EnST...44.8429R}}</ref> The chlorate was also measured in rainfall samples with the amount of chlorate similar to [[perchlorate]]. It is suspected that chlorate and perchlorate may share a common natural formation mechanism and could be a part of the chlorine biogeochemistry cycle. From a microbial standpoint, the presence of natural chlorate could also explain why there is a variety of microorganisms capable of reducing chlorate to chloride. Further, the evolution of chlorate reduction may be an ancient phenomenon as all perchlorate reducing bacteria described to date also utilize chlorate as a terminal electron acceptor.<ref>{{cite journal |last1=Coates |first1=J. D. |last2=Achenbach |first2=L. A. |title=Microbial perchlorate reduction: rocket-fuelled metabolism |journal=Nature Reviews Microbiology |year=2004 |volume=2 |issue=July |pages=569–580 |doi=10.1038/nrmicro926 |pmid=15197392|s2cid=21600794 }}</ref> It should be clearly stated, that currently no chlorate-dominant minerals are known. This means that the chlorate anion exists only as a substitution in the known mineral species, or – eventually – is present in the pore-filling solutions.<ref>{{cite web |url=http://www.mindat.org/ |title=Home |website=mindat.org}}</ref>

In 2011, a study by the [[Georgia Tech|Georgia Institute of Technology]] unveiled the presence of [[magnesium chlorate]] on the planet Mars.<ref>{{Cite news|url=https://www.letemps.ch/sciences/2015/09/28/eau-liquide-reperee-pentes-martiennes|title = De l'EAU liquide répérée sur les pentes martiennes|newspaper = Le Temps|date = 28 September 2015}}</ref>

==Compounds (salts)==
{{See also|Category:Chlorates}}

Examples of chlorates include
* [[potassium chlorate]], KClO<sub>3</sub>
* [[sodium chlorate]], NaClO<sub>3</sub>
* [[magnesium chlorate]], Mg(ClO<sub>3</sub>)<sub>2</sub>

==Other oxyanions==
If a Roman numeral in brackets follows the word "chlorate", this indicates the [[oxyanion]] contains chlorine in the indicated oxidation state, namely:

{| class="wikitable" style="font-size: 100%;"
|-
! Common name !! Stock name !! Oxidation state !! Formula
|-
| [[Hypochlorite]] || Chlorate(I) || +1 || ClO<sup>−</sup>
|-
| [[Chlorite]] || Chlorate(III) || +3 || {{chem|ClO|2|−}}
|-
| Chlorate || Chlorate(V) || +5 || {{chem|ClO|3|−}}
|-
| [[Perchlorate]] || Chlorate(VII) || +7 || {{chem|ClO|4|−}}
|}

Using this convention, "chlorate" means ''any'' chlorine oxyanion. Usually, "chlorate" refers only to chlorine in the +5 oxidation state.

==Toxicity==
Chlorates are relatively toxic, though they form generally harmless chlorides on reduction.

==References==
{{Reflist}}

==External links==
{{Commons category|Chlorate ion}}
* {{cite EB1911|wstitle=Chlorates|volume=6|page=254|short=1}}

{{Chlorates}}
{{Authority control}}

[[Category:Chlorates| ]]
[[Category:Chlorine oxides]]