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Author: Subject: Trying to get potassium permanganate out of solution, keeps turning to MnO2
Junk_Enginerd
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[*] posted on 16-4-2021 at 22:36
Trying to get potassium permanganate out of solution, keeps turning to MnO2


So I've been trying to make potassium permanganate from carbon zinc batteries. I start off by rinsing it to get rid of salts, then dry it and set it on fire with supplemental oxygen gas to burn off some carbon, final carbon removal is done with heating and addition of KNO3 until it doesn't visibly react anymore. Then addition of KOH and heating until melting, pouring this off to get rid of gunk and chunks left behind, to get potassium manganate. This is dissolved in KOH solution and pH lowered with sodium bicarbonate until it turns to permanganate.

This all goes well to the point where I have a solution of beautiful deep purple permanganate and want to get rid of the water. It just won't cooperate. The more water i try to boil off, the more it reverts to MnO2. Is it the boiling that's wrong? Is the pH wrong?
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Sulaiman
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[*] posted on 17-4-2021 at 05:08


I recently had a go at recrystalising commercial potassium permanganate,
boiling the solution will definitely decompose some of your permanganate.

I had to leave the solution at RT in a jar covered by tissue paper
(to allow slow evaporation and keep dust out of the solution)
I got a few small batches of rather ugly crystals over the period of a couple of weeks,
which, when left to dry in open air deteriorated. https://www.sciencemadness.org/whisper/viewthread.php?tid=84...

I intend to revisit this problem soon, so any new info will be appreciated.....





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[*] posted on 17-4-2021 at 05:21


Does carbon dioxide destroy potassium permanganate forming potassium carbonate and manganese dioxide if so maybe dessicate off the solution (under vaccum if you can but not necessary) that way there is limited amount of carbon dioxide.

[Edited on 17-4-2021 by symboom]




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[*] posted on 17-4-2021 at 11:38


Ah. Okay so it may be a temperature issue. I wasn't sure whether it was the loss of water itself that was triggering it, or the heat or the bicarb decomposing from the heat...

So maybe a dehumidifier solution or low pressure evaporation is the way to go then.


Does CO2 destroy permanganate though? I experimented with turning manganate into permanganate by acidifying the solution with just carbonated water which worked great.
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clearly_not_atara
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[*] posted on 17-4-2021 at 14:24


The solubility of KMnO4 declines precipitously with temperature:

https://pubs.acs.org/doi/abs/10.1021/ja01976a004

It's about 25% at 65 C, 7% at 25 C and 3% at 0 C, with much higher solubility above that. In the face of this data, I'm not sure why anyone would want to precipitate by boiling. If the solution is not concentrated enough, reduce volume and then chill.

Furthermore, HMnO4 is strong in water. It doesn't make sense to blame bicarbonate, and in fact, I'd consider bubbling CO2 as an acidification technique.

In principle, the following reaction happens first:

3 MnO4(2-) + 4 H+ >> MnO2 (s) + 2 MnO4- (aq) + 2 H2O

So you should add acid to a warm solution of K2MnO4 and filter out MnO2 (which always precipitates) to obtain a concentrated solution of permanganate. Then cooling to 0 C on an ice bath should give KMnO4 crystals. Acetone may be an effective anti-solvent (or it might explode; consult appropriate literature before proceeding).

I don't think making KMnO4 by evaporation really works unless you use a vacuum; the solutions are unstable and release O2 as they evaporate, and boiling will destroy it.




[Edited on 04-20-1969 by clearly_not_atara]
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S.C. Wack
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[*] posted on 17-4-2021 at 18:43


Quote: Originally posted by Sulaiman  
I recently had a go at recrystalising commercial potassium permanganate,
boiling the solution will definitely decompose some of your permanganate.


This rather depends on the quality of your material...that which is OTC here is totally stable to water. One form of Mn in contact with another form pretty much always causes problems?

PS there will be no precipitation at all with chlorine...and apparently people here have really shitty permanganate...

[Edited on 18-4-2021 by S.C. Wack]




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[*] posted on 17-4-2021 at 19:06


Concentrated forms of KMnO4 will almost always decompose (commercial quality or otherwise). If you don't believe me try making a saturated solution and letting it sit for half a day or so. I would try to find a source, rather than to make it yourself.
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[*] posted on 17-4-2021 at 19:34


Quote: Originally posted by Deathunter88  
Concentrated forms of KMnO4 will almost always decompose (commercial quality or otherwise). If you don't believe me try making a saturated solution and letting it sit for half a day or so. I would try to find a source, rather than to make it yourself.


Yeah I did this, boiled it for hours this afternoon. Not so much as stains on the glass. Y'all using shit grade. And not following directions.

[Edited on 18-4-2021 by S.C. Wack]




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[*] posted on 18-4-2021 at 04:17


Well, I could make some, but from what I have followed on YouTube, it isn't totally simple.

Since I can buy all that I could possibly want, dirt cheap; I decided to take the path least resistance.

P.S. S.C. Wack, thanks for Library posting of Preparative Organic Chemistry, By Hilgetag and A. Martini.

Just spent a day or two reading through it. Pretty good book!

[Edited on 18-4-2021 by zed]

[Edited on 18-4-2021 by zed]
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[*] posted on 18-4-2021 at 04:40


Quote: Originally posted by S.C. Wack  
Yeah I did this, boiled it for hours this afternoon. Not so much as stains on the glass. Y'all using shit grade. And not following directions.
I don't know. I'm a bit skeptical of this claim, because I've had the same experience as others who have described the decomposition in this thread, and I have ACS grade permanganate from Fisher. So I guess that's considered "shit grade" in your book...



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[*] posted on 18-4-2021 at 15:08


BTW I've done this too (with the carbon, without the nitrate) and am familiar with the joy of manganate chemistry and both types of batteries.

Quote: Originally posted by zed  
P.S. S.C. Wack, thanks for Library posting of Preparative Organic Chemistry, By Hilgetag and A. Martini.

Just spent a day or two reading through it. Pretty good book!


Despite the thickness, this book was suited for scanning well; most have several problems. The 1945 version is also good. The original 1938 edition is better except for being in German, it's quite different from the translation. I just dumped my scans (and earlier, Gallica's) out there, Grandmaster P decided to host some of them.

Note that I scanned something like 40 others!

Quote: Originally posted by Texium (zts16)  
I'm a bit skeptical of this claim.


A match head was boiled in an open test tube above a blowtorch, maintaining a half full level, with a ceramic (unglazed boat) boiling stone, for 4-5 hours. What can I say...y'all pay more for less, have dirtier air, water, or glassware, haven't read preparations of KMnO4 from MnO2? Yes, failing is exceptionally easy; that's self-evident and well documented on this site. I don't recall the brand and the store doesn't exist in the west, but you should be able to find solutions for sale, similar if not identical OTC product at a farm store to test yourself (with clean glass and water!!!), and a J Chem Ed article that says the same thing I did.




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[*] posted on 21-4-2021 at 06:03


I hope this is some valuable input and might help you get some percipitation:

I think there are 3 problems:
1) the near neutral pH will increase solubility of unwanted MnO2,
2) the pH drives the reaction towards forming more MnO2 and
3) there is no way to tell if the MnO2 is from the wanted disproportionation or from the unwanted decomposition of KMnO4. Let me explain in detail:

1) I dont think that you really are percipitating out MnO2, most textbooks say that and in general this is not completely false, but what you are really percipitating ist the hydrated from, MnO(OH)2. To make MnO2, you actually have to heat the hydrated form to temperatures beyond 100 °C or let it sit for days (This is easily proven by adding some sulfuric acid to a percipitate of "MnO2", wich should not be soluable at RT in sulfuric acid). In general, MnO2, depending on ist surface and how old it is, is not soluable in most acids (except e.g. HCl , wich it decomposes). The hydrate however, wich you are forming, is soluable under acidic conditions and since your solution is not alkaline, you might have some ions in solution (byproducts Mn4+/Mn3+ and even Mn2+) without even knowing how much. Boiling down the Solution might result in forming insoluable carbonates/hydroxides before the permanganete will percipitate (Wich means your product might not even decompose yet).

Way to solve the Problem: An alkaline solution would give you all the unwanted insoluable hydroxides, wich can be filtered out (Glass wool or KMnO4 will attack you paper and yield will be lower) and probably wont make more problems.

2) Another factor is that the decomposition of permanganate is acutally „releasing“ electrons (Lewis base) as follows:
10 KMnO4 -> 3 K2MnO4 + 7 MnO4 + 2 O2- + 4 K+ + 6 O2
What actually is formed is K2O, I just wrote it as ions to make it more clear.
If we assume that this reaction is also partly taking place in solution, we can see that for 10 eq of KMnO4 we release 7 eq of MnO(OH)2 and 4 eq KOH. This means the less alkaline our solution is, the more prone our reaction mixture ist to decompose into KOH itself. Furthermore by having a slightly acidic mixture we might even remove the hydrate and neutralize the KOH and therefore shitfting the equilibrium, supporting the decomposition of the wanted product.
Way to solve the Problem: Add more KOH so the equilibrium is on your side, and less KMnO4 decomposes.

3) The last part is in my opinion the tricky one: There really is no way to tell if you really completed the reaction. The color of the permanganate is too intense to see the rest of the manganate. My suggestion: Give it time.

So how I would try to get the Permanganate:
Dissolve your manganate in KOH as you already did, bring the pH to neutral as before, heat it gently for a couple of ours (no boiling) and let it sit for 3 days. Then add KOH until you get a pH of about 11, stir for an hour and then filter the soulution, glass wool. After this, your solution should only consist of KOH and KMnO4. Add a solution of KOH (hard to figure out the concentration, I would start with 9.6 M, depends on your starting material) with cooling. This is to increase K+ concentration in solutuion and therefore lowering KMnO4 solubility. (To percipitate K2FeO4 for example, I found a preparation wich suggest to add approximatly 40 times the molar eq of KOH to a soultion of K2FeO4 to percipitate out K2FeO4.) Then cool the solution and wait for it to percipitate out.
Disclaimer: I have not tried this by myself yet, this is all theoretical.
Good Luck!


PS. This is my first post, I am open to suggestions and I am happy to be corrected if anything I stated is incorrect!
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[*] posted on 21-4-2021 at 07:47


Which is exactly why I figure to buy mine in a bottle. Trying to make it, messes with my absolutist ideals... I'm kind of a bottom feeder; if it appears that the project might prove difficult to successfully accomplish, and there is no big pay-off, I quit immediately. Plenty of really interesting projects I can fail at, without embarking on a project, I know in advance, is star-crossed.

There's a Song about that.

https://www.youtube.com/watch?v=2Py37G9qsfY





[Edited on 21-4-2021 by zed]
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