Barium permanganate is a chemical compound, with the formula Ba(MnO4)2.[2] It forms violet to brown crystals that are sparingly soluble in water.
| Identifiers | |
|---|---|
3D model (JSmol)
|
|
| ChemSpider | |
| ECHA InfoCard | 100.029.191 |
| EC Number |
|
PubChem CID
|
|
| UNII | |
CompTox Dashboard (EPA)
|
|
| |
| |
| Properties | |
| Ba(MnO4)2[1] | |
| Molar mass | 375.198 g/mol |
| Appearance | dark violet to brown crystals |
| Odor | odorless |
| Density | 3.77 g/cm3 |
| Melting point | 200 °C (392 °F; 473 K) (decomposes) |
| 62.5 g/100 mL (29 °C) | |
| Solubility | decomposes in alcohol |
| Structure | |
| rhombic | |
| Hazards | |
| NFPA 704 (fire diamond) | |
| Related compounds | |
Other cations
|
Magnesium permanganate Strontium permanganate |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
| |
Preparation
editBarium permanganate may be produced by disproportionation of barium manganate in a mildly acidic solution,[3] including solutions carbon dioxide or sulfuric acid:[4]
- 3 BaMnO4 + 2 CO2 → Ba(MnO4)2 + 2 BaCO3 + MnO2
- 3 BaMnO4 + 2 H2SO4 → Ba(MnO4)2 + 2 BaSO4 + MnO2 + 2 H2O
It can also be prepared by oxidation of barium manganate with strong oxidants. Preparations relying on aqueous reactions of barium manganate are extremely slow process due to the low solubility of the manganate.[3]
Another way to synthesize barium permanganate is by the reaction between silver permanganate and barium chloride. Highly pure samples can be obtained from the similar reaction between potassium permanganate and aluminium sulfate to form aluminium permanganate, which is then reacted with a stoichiometric amount of barium hydroxide.[5]
Reactions
editBarium permanganate is a strong oxidizer. It is thermally stable up to 180 °C, above which it decomposes in two stages between 180–350 and 500–700 °C.[5]
- 2 Ba(MnO4)2 → 2 BaMnO3 + 2 MnO2 + 3 O2
- 4 BaMnO3 → 4 BaO + 2 Mn2O3 + O2
The decomposition has been shown to proceed at slow rates above 160 °C, and that irradiation with UV or X-rays lowers this temperature.[6] Crystal defects and impurities play a role in the mechanism.
Permanganic acid can be prepared by the reaction of dilute sulfuric acid with a solution barium permanganate, the insoluble barium sulfate byproduct being removed by filtering:[3]
- Ba(MnO4)2 + H2SO4 → 2 HMnO4 + BaSO4
The sulfuric acid used must be dilute; reactions of permanganates with concentrated sulfuric acid yield the anhydride, manganese heptoxide.
References
edit- ^ Lide, David R., ed. (2009). CRC Handbook of Chemistry and Physics (90th ed.). Boca Raton, Florida: CRC Press. p. 4-50. ISBN 978-1-4200-9084-0.
- ^ PubChem
- ^ a b c Olsen, J. C. (1900). Permanganic Acid by Electrolysys. Easton, PA: The Chemical Publishing Company.
- ^ Brauer, Georg (1981). Handbuch der präparativen anorganischen Chemie. Band III (3rd ed.). Stuttgart: Enke. ISBN 3-432-87823-0. OCLC 310719495.
- ^ a b Kotai, Laszlo; Gacs, Istvan; Sajo, Istvan E.; Sharma, Pradeep K.; Banerji, Kalyan K. (2011-03-29). "ChemInform Abstract: Beliefs and Facts in Permanganate Chemistry - An Overview on the Synthesis and the Reactivity of Simple and Complex Permanganates". ChemInform. 42 (13): no. doi:10.1002/chin.201113233.
- ^ Prout, E. G.; Herley, P. J. (1961). "The Thermal Decomposition of Barium Permanganate". The Journal of Physical Chemistry. 65 (2): 208–212. doi:10.1021/j100820a005. ISSN 0022-3654.
 | This inorganic compound–related article is a stub. You can help Wikipedia by expanding it. |