{{Short description|Anion and term for chemical compounds containing it}}{{Chembox| Name = Chlorate| ImageFile = Chlorate-2D.png| ImageName = Structure and bonding in the chlorate ion| ImageFile1 = Chlorate-3D-vdW.png| 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 = ClO3âˆ’| MolarMass = 83.4512 | Appearance = | Density = | Solubility = | pKa =

}}| Section3 = {hide}Chembox Structure! Common name !! Stock name !! Oxidation state !! Formula

Trigonal pyramidal molecular geometry>Trigonal pyramidal\| Dipole = {edih}\| Section7 = {{Chembox Hazards\| MainHazards = oxidation agent }}\| Section8 = {hide}Chembox Related	bromate\|iodate{edih}	chloride	hypochlorite>chlorite\|perchlorate}} }} }}Chlorate is the common name of the {{chem\|ClO\|3\|âˆ’}} anion, whose chlorine atom is in the +5 oxidation state. The term can also refer to chemical compounds containing this anion, with chlorates being the salts of chloric acid. Other 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\|âˆ’}} ion commonly called perchlorate can also be called chlorate(VII).As predicted by valence shell electron pair repulsion theory, chlorate anions have trigonal pyramidal structures.Chlorates are powerful 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 perchlorates instead. Structure and bonding The chlorate ion cannot be satisfactorily represented by just one Lewis structure, since all the Clâ€“O bonds are the same length (1.49 Ã… in potassium chlorate), and the chlorine atom is hypervalent. Instead, it is often thought of as a hybrid of multiple 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 hydroxides like KOH: 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 Clâˆ’ (oxidation number âˆ’1) and chlorate(V) {{chem\|ClO\|3\|âˆ’}} (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 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 50â€“70 Â°C is performed by the 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.JOURNAL, Rao, B., Hatzinger, P. B., BÃ¶hlke, J. K., Sturchio, N. C., Andraski, B. J., F. D., Eckardt, Jackson, W., Natural Chlorate in the Environment: Application of a New IC-ESI/MS/MS Method with a Cl18O3âˆ’ Internal Standard, Environ. Sci. Technol., 2010, 44, 22, 8429â€“8434, 10.1021/es1024228, 20968289, 2010EnST...44.8429R, 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.JOURNAL, Coates, J. D., Achenbach, L. A., Microbial perchlorate reduction: rocket-fuelled metabolism, Nature Reviews Microbiology, 2004, 2, July, 569â€“580, 10.1038/nrmicro926, 15197392, 21600794, 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 â€“ eventually â€“ is present in the pore-filling solutions.WEB,www.mindat.org/, Home, mindat.org, In 2011, a study of the Georgia Institute of Technology unveiled the presence of magnesium chlorate on the planet Mars.NEWS,www.letemps.ch/sciences/2015/09/28/eau-liquide-reperee-pentes-martiennes, De l’EAU liquide rÃ©pÃ©rÃ©e sur les pentes martiennes, Le Temps, 28 September 2015, Compounds (salts) {{See also\|Category:Chlorates}}Examples of chlorates include potassium chlorate, KClO3 sodium chlorate, NaClO3 magnesium chlorate, Mg(ClO3)2 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%;”
Hypochlorite >		\| ClOâˆ’
Chlorite >			ClO	âˆ’}}
			ClO	âˆ’}}
Perchlorate >			ClO	âˆ’}}

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

External links

EB1911, Chlorates, 6, 254, 1,

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chlorate

Structure and bonding

Preparation

Laboratory

Industrial

Natural occurrence

Compounds (salts)

Other oxyanions

Toxicity

References

External links