|Jmol-3D images||Image 1|
|Molar mass||138.55 g mol−1|
|Appearance||colourless/ white crystalline powder|
|Melting point||610 °C (1,130 °F; 883 K)
decomposes from 400 °C
|Solubility in water||0.76 g/100 mL (0 °C)
1.5 g/100 mL (25 °C)
4.76 g/100 mL (40 °C)
21.08 g/100 mL (100 °C)
|Solubility product, Ksp||1.05·10-2|
|Solubility||negligible in alcohol
insoluble in ether
|Solubility in ethanol||47 mg/kg (0 °C)
120 mg/kg (25 °C)
|Solubility in acetone||1.6 g/kg|
|Solubility in ethyl acetate||15 mg/kg|
|Refractive index (nD)||1.4724|
heat capacity C
|Std enthalpy of
|Gibbs free energy ΔG||-300.4 kJ/mol|
|GHS signal word||Danger|
|GHS hazard statements||H271, H302|
|GHS precautionary statements||P220|
|EU classification||O Xn|
|S-phrases||(S2), S13, S22, S27|
|Other anions||Potassium chloride
|Other cations||Ammonium perchlorate
|Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)|
|(what is: / ?)|
Potassium perchlorate is the inorganic salt with the chemical formula KClO4. Like other perchlorates, this salt is a strong oxidizer and potentially reacts with many organic substances. This usually obtained as a colorless, crystalline solid is a common oxidizer used in fireworks, ammunition percussion caps, explosive primers, and is used variously in propellants, flash compositions, stars, and sparklers. It has been used as a solid rocket propellant, although in that application it has mostly been replaced by the higher performance ammonium perchlorate. KClO4 has the lowest solubility of the alkali metal perchlorates (1.5 g in 100 mL of water at 25 °C).
KClO4 is prepared industrially by treating an aqueous solution of sodium perchlorate with KCl. This single precipitation reaction exploits the low solubility of KClO4, which is about 100 times less than the solubility of NaClO4 (209.6 g/100 mL at 25 °C).
KClO4 is an oxidizer in the sense that it exothermically transfers oxygen to combustible materials, greatly increasing their rate of combustion relative to that in air. Thus, with glucose it gives carbon dioxide:
- 3 KClO4 + C6H12O6 → 6 H2O + 6 CO2 + 3 KCl
The conversion of solid glucose into hot gaseous CO2 is the basis of the explosive force of this and other such mixtures. Even with cane sugar, KClO4 yields a low explosive, provided the necessary confinement. Otherwise such mixtures simply deflagrate with an intense purple flame characteristic of potassium. Flash compositions used in firecrackers usually consist of fine aluminium powder mixed with potassium perchlorate.
As an oxidizer, potassium perchlorate can be used safely in the presence of sulfur, whereas potassium chlorate cannot. The greater reactivity of chlorate is typical – perchlorates are kinetically poorer oxidants. Chlorate produces chloric acid, which is highly unstable and can lead to premature ignition of the composition. Correspondingly, perchloric acid is quite stable.
Potassium perchlorate can be used as an antithyroid agent used to treat hyperthyroidism, usually in combination with one other medication. This application exploits the similar ionic radii and hydrophilicity of perchlorate and iodide.
Since 2005, a cartridge with potassium perchlorate mixed with anthracene and sulfur is used for generating the black smoke signalling the failure of reaching a two-thirds majority needed for the election of new pope by a papal conclave.
- "Potassium Perchlorate MSDS". J.T. Baker. 2007-02-16. Retrieved 2007-12-10.
- Benenson, Walter; Stöcker, Horst. Handbook of Physics. Springer. p. 780. ISBN 0387952691.
- Sigma-Aldrich Co., Potassium perchlorate. Retrieved on 2014-05-27.
- Potassium perchlorate in Linstrom, P.J.; Mallard, W.G. (eds.) NIST Chemistry WebBook, NIST Standard Reference Database Number 69. National Institute of Standards and Technology, Gaithersburg MD. http://webbook.nist.gov (retrieved 2014-05-27)
- Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 0-618-94690-X.
- Helmut Vogt, Jan Balej, John E. Bennett, Peter Wintzer, Saeed Akbar Sheikh, Patrizio Gallone “Chlorine Oxides and Chlorine Oxygen Acids” in Ullmann's Encyclopedia of Industrial Chemistry 2002, Wiley-VCH. doi:10.1002/14356007.a06_483
- Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
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