Potassium manganate

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Potassium manganate
Potassium-manganate-unit-cell-3D-balls.png
Identifiers
CAS number 10294-64-1 YesY
PubChem 160931
EC number 233-665-2
Jmol-3D images Image 1
Properties
Molecular formula K2MnO4
Molar mass 197.132 g/mol
Appearance dark green crystals
Density 2.78 g/cm3, solid
Melting point 190 °C (374 °F; 463 K)
Solubility in water decomposes
Acidity (pKa) 7.1
Structure
Crystal structure isomorphous with K2SO4
Coordination
geometry
tetrahedral anion
Hazards
R-phrases 8-36/37/38
S-phrases 17-26-36/37/39
Main hazards oxidizer
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 1: Exposure would cause irritation but only minor residual injury. E.g., turpentine Reactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calcium Special hazard OX: Oxidizer. E.g., potassium perchlorateNFPA 704 four-colored diamond
Related compounds
Related compounds KMnO4
MnO2
K2CrO4
K2FeO4
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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Infobox references

Potassium manganate is the inorganic compound with the formula K2MnO4. This green-colored salt is an intermediate in the industrial synthesis of potassium permanganate (KMnO4), a common chemical. Occasionally, potassium manganate and potassium permanganate are confused, but they are different compounds with distinctly different properties.

Structure[edit]

K2MnO4 is a salt, consisting of K+ cations and MnO42− anions. X-ray crystallography shows that the anion is tetrahedral, with Mn-O distances of 1.66 Å, ca. 0.03 Å longer than the Mn-O distances in KMnO4.[1] It is isostructural with potassium sulfate.

Synthesis[edit]

The industrial route entails treatment of MnO2 with air:

2 MnO2 + 4 KOH + O2 → 2 K2MnO4 + 2 H2O

The transformation gives a green-colored melt. In fact, one can test an unknown substance for the presence of manganese by heating the sample in strong KOH in air. The production of a green coloration indicates the presence of Mn. This green color results from an intense absorption at 610 nm.

In laboratory, K2MnO4 can be synthesized by heating a solution of KMnO4 in concentrated KOH solution followed by cooling to give green crystals:[2]

4 KMnO4 + 4 KOH → 4 K2MnO4 + O2 + 2 H2O

This reaction illustrates the relatively rare role of hydroxide as a reducing agent. Solutions of K2MnO4 are generated by allowing a solution of KMnO4 in 5-10 M KOH to stir for a day at room temperature followed by removal of MnO2, which is insoluble. The concentration of K2MnO4 in such solutions can be checked by measuring their absorbance at 610 nm.

The one-electron reduction of permanganate to manganate can also be effected using iodide as the reducing agent:

2 KMnO4 + 2 KI → 2 K2MnO4 + I2

The conversion is signaled by the color change from purple, characteristic of permanganate, to the green color of manganate. This reaction also illustrates the fact that manganate(VII) can serve as an electron acceptor in addition to its usual role as an oxygen-transfer reagent. Barium manganate, BaMnO4, is generated by the reduction of KMnO4 with iodide in the presence of barium chloride. Just like BaSO4, BaMnO4 exhibits low solubility in virtually all solvents.

An easy method for preparing potassium manganate in the laboratory involves heating crystals or powder of pure potassium permanganate. Potassium permanganate will decompose into potassium manganate, manganese dioxide and oxygen gas:

2KMnO4 → K2MnO4 + MnO2 + O2

This reaction is a laboratory method to prepare oxygen.

Reactions[edit]

See also: Manganate

At lower pH, the manganate ion will disproportionate to permanganate ion and manganese dioxide:

3 K2MnO4 + 2 H2O → 2 KMnO4 + MnO2 + 4 KOH

The colorful nature of this reaction has led the manganate/manganate(VII) pair to be referred to as a chemical chameleon. This disproportionation reaction, which becomes rapid when [OH
] < 1M, follows bimolecular kinetics.[1]

Literature cited[edit]

  1. ^ Palenik, G. J. (1967). "Crystal Structure of Potassium Manganate". Inorg. Chem. 6: 507–511. doi:10.1021/ic50049a01. 
  2. ^ Nyholm, R. S.; Woolliams, P. R. (1968). "Manganates(VI)". Inorg. Synth. Inorganic Syntheses 11: 56–61. doi:10.1002/9780470132425.ch11. ISBN 978-0-470-13242-5. 

Other references[edit]

  • Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.

See category for a list.

External links[edit]