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I read this question and answer, and it seemed like a rather nice way to make sodium permanganate. Using the method with sodium hydroxide, I produced a solution. Once filtered and decanted, it was dark green.

At first this seemed somewhat strange, as I would expect to see a red color for sodium permanganate. Though, after a bit of research it seems this is actually sodium manganate. The color matches, and considering it only exists in very basic conditions, and since there is a lot of sodium hydroxide in the solution, this would make sense. To test this further, I added some sodium bicarbonate to, oddly enough, act as an acid. The sample of solution quickly turned the pink color I had expected.

So, I currently have some amount of sodium manganate, sodium hypochlorite, sodium hydroxide, and sodium chloride in this solution.

At first I hoped I might be able to extract this using organic solvents, but unfortunately wikipedia claims potassium permanganate decomposes in organic solvents and alcohol. I assume the sodium permanganate or manganate would do the same.

Upon boiling down the solution I would end up with a mess of crystals of various shades of blackish and whitish.

So, my real question here is, if there is a way or ways to remove the other substances (sodium hypochlorite, sodium hydroxide, sodium chloride) and attain more or less pure sodium permanganate crystals.

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    $\begingroup$ The easiest way to get some permanganate is to go with potassium salt. Potassium permanganate is fairly soluble in warm water, but only moderately soluble in cold water. No so luck with the sodium salt. Afafik, the sodium salt in clean state may be acquired by electrolysis of sodium hydroxide solution with manganese anode. $\endgroup$ – permeakra Jul 23 '16 at 10:06
  • $\begingroup$ Yes, permeakra is correct. Sodium permanganate has solubility of 900 g/L so it will always come last from the solution. And there is no simple way to separate it from other salts. Also, it's hygroscopic. BTW, use less concentrated sodium hydroxide for oxidation, so you will get directly sodium permanganate and not manganate. $\endgroup$ – vapid Jul 25 '16 at 13:32
  • $\begingroup$ Solid potassium permanganate isn't heat stable, so I wouldn't try to boil off all the water. $\endgroup$ – Zhe Oct 21 '16 at 19:55
  • $\begingroup$ KMnO4 is likely to react with organic solvents rather than react with them... $\endgroup$ – Eashaan Godbole Jul 19 '17 at 16:26
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You are correct. First step produces $\ce{K2MnO4}$, then you need to oxidize it.

My first guess was that adding $\ce{H2SO4}$ will cause disproportionation of $\ce{K2MnO4}$ into $\ce{KMnO4}$ and $\ce{MnO2 + K2SO4}$. Filtering removes $\ce{MnO2}$. Crystallization will give you pure product. Or you can use a mix of $\ce{KMnO4 + K2SO4}$ for many purposes.

Industry uses more efficient method: electrolysis of $\ce{K2MnO4}$.

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  • $\begingroup$ Wouldn't the addition of sulfuric acid produce manganese heptoxide? $\endgroup$ – ChemBird Jul 23 '16 at 20:03
  • $\begingroup$ It wouldn't. Mn$^{5+}$ (as in K**$_3$**MnO4) and Mn$^{6+}$ (as in K**$_2$**MnO4) are only stable under basic conditions. In acid they disproportionate to yield MnO2 and KMnO4. If you add too much H2SO4 you can get Mn2O7 - a very reactive and dangerous Mn$^{7+}$ oxide that inflames EtOH on contact. $\endgroup$ – sixtytrees Jul 23 '16 at 20:07

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