Chlorate

From Wikipedia, the free encyclopedia
Jump to: navigation, search
The chlorate ion
CAS # 14866-68-3[1]
Structure and bonding in the chlorate ion

The chlorate anion has the formula ClO3. In this case, the chlorine atom is in the +5 oxidation state. "Chlorate" can also refer to chemical compounds containing this anion; chlorates are the salts of chloric acid. "Chlorate", when followed by a Roman numeral in parentheses, e.g. chlorate(VII), refers to a particular oxyanion of chlorine.

As predicted by VSEPR, 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 which formerly used chlorates in the past now use the more stable perchlorates instead.

Structure and bonding[edit]

The chlorate ion cannot be satisfactorily represented by just one Lewis structure, since all the Cl-O bonds are the same length (1.49 Å in potassium chlorate[2]), and the chlorine atom is hypervalent. Instead, it is often thought of as a hybrid of multiple resonance structures:

Resonance structures of the chlorate ion

Preparation[edit]

Laboratory[edit]

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 Cl (oxidation number −1) and chlorate(V) ClO
3
(oxidation number +5). Reaction of cold aqueous metal hydroxides with chlorine produces the chloride and hypochlorite (oxidation number +1) instead.

Industrial[edit]

The industrial scale synthesis for sodium chlorate starts from aqueous sodium chloride solution (brine) rather than chlorine gas. If equipment for electrolysis allows mixing of the chlorine and the sodium hydroxide, then the disproportionation reaction described above occurs. The heating of the reactants to 50-70°C is performed by the electrical power used for electrolysis.[citation needed]

Natural Occurrence[edit]

A recent study has discovered the presence of natural chlorate deposits around the world with relatively high concentrations found in arid and hyper-arid regions.[3] The chlorate was also measured in rainfall samples with the amount of chlorate similar to perchlorate. It is suspected that both chlorate and perchlorate may share a common natural formation mechanism(s) 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 micro-organisms 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.[4]

Compounds (salts)[edit]

Examples of chlorates include

Other oxyanions[edit]

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

Common name Stock name Oxidation state Formula
Hypochlorite Chlorate(I) +1 ClO
Chlorite Chlorate(III) +3 ClO2
Chlorate Chlorate(V) +5 ClO3
Perchlorate Chlorate(VII) +7 ClO4

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

Toxicity[edit]

Chlorates are relatively toxic, though they form generally harmless chlorides upon reduction.

References[edit]

  1. ^ "ChemIndustry". Retrieved 9 April 2014. 
  2. ^ J. Danielsen, A. Hazell, F. K. Larsen (1981). "The structure of potassium chlorate at 77 and 298 K". Acta Cryst. B37: 913–915. doi:10.1107/S0567740881004573. 
  3. ^ Rao, B.; Hatzinger, P. B.; Böhlke, J. K.; Sturchio, N. C.; Andraski, B. J.; Eckardt, F. D.; Jackson, W. (2010). "Natural Chlorate in the Environment: Application of a New IC-ESI/MS/MS Method with a Cl18O3 Internal Standard". Environ. Sci. Technol. 44: 8429–8434. doi:10.1021/es1024228. PMID 20968289. 
  4. ^ Coates, J. D.; Achenbach, L. A. (2004). "Microbial perchlorate reduction: rocket-fuelled metabolism". Nature Reviews Microbiology 2 (July): 569–580. doi:10.1038/nrmicro926. PMID 15197392.