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Author: Subject: Permanganates
DerAlte
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 Quote: Originally posted by blogfast25 Something I want to try shortly [big cough!] is to fuse 'MnOx≈2' with an excess of KNO3 and just a bit of KOH (to ensure alkalinity): MnO2 + 2 KNO3 ===> K2MnO4 + 2 NO2 I have reason to believe that with a 50 %w stoichiometric excess (of KNO3) the conversion [IV] to [VI] may be near 100 %. Then maybe leach with 1 M KOH and acidify to [VII] + [IV]? Maybe with acetic acid ≈ 2 M, slowly and with vigorous stirring? Sooner or later anyone always catches the 'permanganate virus'!

Once a transition metal gets hold of you and tangles you in its swarm of outer electrons you are doomed!

Love to try that idea.

Der Alte
blogfast25
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I carried out two attempts to prepare potassium manganate (VI) with:

MnO2 (s) + 2 KNO3 (l) === > K2MnO4 (l) + 2 NO2 (g)

… and both failed. 5.0 g o MnO2 (pottery), 17.4 g of KNO3 (a 50 % excess) and 1 g of KOH (to ensure some alkalinity to protect any K2MnO4 formed) were mixed, ground and then fused together in a nickel crucible for 20 minutes on a medium-high Bunsen heat. In neither instances were significant amount of the emerald potassium manganate obtained, despite vigorous bubbling of the melt (what are these bubbles, O2?)

I then attempted:

MnO2 (s) + KOH (l) + KNO3 (l) === > K2MnO4 (l) + NO (g) + ½ H2O (g)

… in the same conditions and rather stupidly not realising at that point this equation isn’t balanced: ½ O is missing on the right hand side (*)!

Formulation: MnO2: 5.0 g; KOH: 4.1 g (a 20 % excess); KNO3: 9.5 g (a 50 % excess). This too melted easily and started bubbling right away, the latter which subsided completely after only about 5 minutes of heating. Heating was then stopped.

The crucible content was then leached with about 100 ml of approx. 1 M KOH solution. It was clear that much potassium manganate (VI) had formed with relatively little unreacted MnO2 in the leachate. The leachate was a comforting very deep green colour.

And then something unexpected happened: on hot filtering (normal filter paper) the manganate started to decompose to MnO2, ON THE FILTER! Finely formed MnO2 found its way through the filter, creating a mess in the filtrate. Even though I could still see green droplets passing into the filtrate, they too seemed to immediately revert to MnO2. But what is being oxidised here, since as no potassium permanganate is being formed? The filter itself?

A little bit of the unfiltered K2MnO4 bearing liquor was set aside. Simple attempts at acidifying this with dilute H2SO4 and acetic acid, to provoke the disproportionation to (IV) + (VII) were also unsuccessful.

(*) The correctly balanced equation is:

3 MnO2 (s) + 4 KOH (l) + 2 KNO3 (l) === > 3 K2MnO4 (l) + 2 NO (l) + 2 H2O (l)

Or: MnO2 (s) + 1.333… KOH (l) + 0.666… KNO3 (l) === > K2MnO4 (l) + 0.666… NO (l) + 0.666… H2O (l), which isn't that far from the original unbalanced equation.

[Edited on 25-8-2013 by blogfast25]

Formatik
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 Quote: Originally posted by blogfast25 And the something unexpected happened: on hot filtering (normal filter paper) the manganate started to decompose to MnO2, ON THE FILTER! Finely formed MnO2 found its way through the filter, creating a mess in the filtrate. Even though I could still see green droplets passing into the filtrate, they too seemed to immediately revert to MnO2. But what is being oxidised here, since as no potassium permanganate is being formed? The filter itself?

One old preparation I read says distinctly not to use paper (Ausführliches Lehrbuch der pharmaceutischen Chemie, E.A. Schmidt, 848) to filter the potassium manganate solution, they used asbestos. Glass wool can be used instead. Pumice or fiberglass (which has the organic aspect destroyed) might also work for filtration.
blogfast25
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Thank you, Formatik, very useful. It makes sense, manganates being such fragile compounds. I'll also have a look at that used filter paper, for clues.

The German experimenter (Versuchschemie) only mentioned filtering, so I assumed he used paper. But one paper may not be equal to another...

[Edited on 25-8-2013 by blogfast25]

DerAlte
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@Formatik & blogfast25 et al.:

Filter paper is a very definite NO for any manganate! I should have mentioned that in the MnOXY document. Cellolose is attacked even at high pH. I use filter paper as a cheap chromatograph to determine whether manganate is present during the electrolytic conversion to permanganate. Put a drop of the converting electrolyte on to filter paper and watch it spread out. Manganate and permangante separate, green outer, magenta inner. But in a few moments, only brown crud remains.

I use glass wool in the funnel. It traps MnO2 particles but may require a second pass, pouring the filtrate back over the funnel. The presence of MnO2 traps most of any remaining, but filtration may be slow.

You said: Formulation: MnO2: 5.0 g; KOH: 4.1 g (a 20 % excess); KNO3: 9.5 g (a 50 % excess). This too melted easily and started bubbling right away, the latter which subsided completely after only about 5 minutes of heating. Heating was then stopped.

5 mins. sounds an excessively short time to expect the reaction to complete, even with activated MnO2 instead of pottery grade. Was the temp. 'bright red heat' or merely enough to well melt the mix? Activated contains xH2O and can be expected to fizzle for a bit. Even if x is very small, steam should be given off as part of the MnO2/KOH reaction.

Old references talk about "hard" filter paper more resistant to oxidation, but I have never seen any.

Congratulations on trying out your idea, blogger. Not only do you blog fast, but you experiment fast!

Regards, DA
blogfast25
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DerAlte:

Good to see you still keep an eye on this thread, more about that later.

An additional observation: no NOx was ever seen, or smelled during fusion/reaction, so I need to assume the overall reaction was:

2 MnO2 + 4 KOH + 2 KNO3 --> 2 K2MnO4 + 2 KNO2 + 2 H2O

… as in fact described in ‘t’Jap patent’ . If so, I was probably not using adequate reagent quantities and will need to repeat at least once, with the right quantities.

For obvious reasons, I don’t know what the conversion was and there certainly was unreacted crud left but much K2MnO4 had formed too. Perhaps about 50 %? The difference with the ‘KNO3 only’ test was truly striking.

I accept that 5 minutes is a very short time but what’s a boy to do when after that time ALL steam evolution stops other than to assume it’s ‘game over’? Initially gas evolution was surprisingly vigorous and I had to stir to prevent over boiling. This wasn't water that was already there, no, this was being formed in situ, no question about it. Of course it's possible that more steam would have slowly and invisibly bled off over the course of a longer heat.

Can you confirm/infirm that I read correctly/incorrectly in your document that for manganate (VI) formation LOW temperature is to be preferred?

DA, I’ve finally managed till the wee small hours last night to trawl through this entire thread and your excellent summary of prior art. Boy, permanganates are still a lot harder to prepare than I thought (and I never had any illusions about it being easy to begin with). Hats off to you and Xenoid, especially for his work on electrolytic oxidation of MnO2.

It appears to me that in terms of Free Energy (G), the series Mn II/III/IV/V/VI/VII is a valley with the black crud at the lowest point and everything else wanting to roll back to that point, given half a chance!

I will continue some limited experimentation but need to prepare some ‘activated’ MnO2 first. I’ll use the MnSO4/NaClO method for that.

I’ll also want to make some Na manganate (V), like Xenoid did, just to be able to say: ‘I saw that!’

Thanks for the tip on filtration media. I should have known better. The German guy actually mentioned ‘harder’ filter paper. I’ll try and use (very clean! ) glass frit.

Oh, and one last thing. The little amount of green solution I managed to syphon off without ever seeing the filter paper, when gently acidified, once with HOAc, once with 1 M H2SO4, never yielded permanganate, not even a whiff, only brownish crud and some bubbles. Oxidation of water?

[Edited on 26-8-2013 by blogfast25]

DerAlte
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@blogfast
 Quote: … as in fact described in ‘t’Jap patent’

I now believe most of what that says. I am always very sceptical about patents.
Exceptions are (1) yields – have you honestly got yields in any experimental process much greater than70-80%? (2) I have not yet successfully managed the electrolytic conversion of an alkaline suspension of ‘MnO2’ to MnO4- but do think it ought to be feasible, given a correct set of conditions and catalyst.
 Quote: Can you confirm/infirm that I read correctly/incorrectly in your document that for manganate (VI) formation LOW temperature is to be preferred?

{Aside: I am not sure your use of the word 'infirm' is the reverse of confirm! Infirm describes my current condition, unfortunately}.

Yes. Higher temperature favors hypomanganate with KOH. From the original sourece: (but this is with direct oxidation using air):

The ﬁrst step in roasting processes is the formation of K3MnO4 from MnO2 ore. This is promoted by high temperature and high KOH and low H2O concentration. The second step oxidizes Mn(V) to Mn(VI). A lower temperature and control of moisture in the air is used.

Note also, from the same source:

Aqueous potassium permanganate solutions are not perfectly thermodynamically stable at 25C, because MnO2, not MnO4-, is the thermodynamically stable form of manganese in water. Thus permanganate tends to oxidize water with the evolution of oxygen and the deposition of manganese dioxide, which acts to further catalyze the reaction.

Which is why black crud is often the result…
 Quote: It appears to me that in terms of Free Energy (G), the series Mn II/III/IV/V/VI/VII is a valley with the black crud at the lowest point and everything else wanting to roll back to that point, given half a chance!

Absolutely. Any experiment to produce Mn X seems to start with xMnO2.yMn2O3.zH2O and end up with the same, with only x,y,z changed!

One way to convert manganate to permanganate without disproportionation and the black plague is to use hypochlorite;
2MnO4-- + ClO- + 2H+  Cl- + H2O + 2MnO4-
Compare anodic oxidation:
MnO4--  MnO4- + e-

I’ve got a lot more but enough for now,
Regards,
Der Alte
blogfast25
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In the recesses of my chemical cavern I found some 'MnO2' I bought years ago, as a 'high grade' (definitely not 'pottery'). It's very brown (and very fine), rather than black and I believe it's a synthetic grade (it dissolves completely in HCl and appears also Fe free). So next week end I'll try some longer, relatively 'cool' fusions with KOH + KNO3 quantities in accordance with some experimenters here, using that material. Target: K2MnO4. Any specific recommendations welcome.

Thanks DA.

[Edited on 28-8-2013 by blogfast25]

DerAlte
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The brown stuff is what one usually gets when one prepares "MnO2" chemically. It contains xH2O, x order <2. Expect more fizzle! Also, try heating a small quantity carefully to reduce water content. If it turns black, or darker, then x is fairly large.

Electrolytic MnO2 is quoted as 98% MnO2. I have never tried making it, so I don't know what color it is.

Der Alte
blogfast25
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I'll have look at heating a bit of it tomorrow. A bit of water doesn't do this thing any harm though, does it?

Any recommended formulation for best chances with K2MnO4?

Ta.

[Edited on 29-8-2013 by blogfast25]

DerAlte
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Blogfast wrote:
 Quote: A bit of water doesn't do this thing any harm though, does it?

A source I have says
 Quote: liquid-phase oxidation: KOH, (O2)air, H2O , MnO2 ore -- > K2MnO4; electrolysis -- > KMnO4;

Of course this is with air oxidation but many industrial processes use very concentrated KOH solution :
 Quote: The USSR process (118–120) is discontinuous, uses turbine-agitated, low pressure reactors having a volume of 4 m3 each, and processes 2000–2500 L/ batch. Preconcentrated molten potassium hydroxide (70–80%) is added to the reactor with a quantity of 78–80% MnO2ore (<0.1mm particle size) resulting in a 1:5 molar ratio of MnO2: KOH. Air, or O2, is introduced below the liquid level by a sparging device at such a rate that a positive pressure of 186 – 216 kPa (1.9 – 2.2 atm) is maintained. The temperature is kept at 250 – 320C for the duration of the reaction, which requires approximately 4–6 h for completion. The reaction mixture, which reportedly remains ﬂuid during the entire time, is then emptied through a siphon.

From which I would guess that the presence of water is not deleterious – it is produced in the reaction anyrate. I have used highly concentrated NaOH and added nitrate and MnO2xH2O and heated, driving off water and the reaction seems to go as if you merely melted the hydroxide.

Der Alte
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As for buying permanganate in the Eastern US: http://www.lowes.com/pd_112505-677-PF65N_0__?Ntt=potassium&U... Lowe's sells cheap 5lbs containers of it...
Brain&Force
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I wonder where you would find it on the west coast - it's not available where I live. I have to resort to aquarium suppliers, which supply it only as a solution.

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vmelkon
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 Quote: Originally posted by The Volatile Chemist As for buying permanganate in the Eastern US: http://www.lowes.com/pd_112505-677-PF65N_0__?Ntt=potassium&U... Lowe's sells cheap 5lbs containers of it...

Lucky you.
I guess 5 lbs is something like 2.5 kg.
I had to buy mine from England and it was 70$for 800 g, shipping included to Canada. //EDIT: woops, the actual amount is 46.35$USD
26.50$(800 g of KMnO4) + 19.84$ shipping.

[Edited on 2-4-2014 by vmelkon]

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plante1999
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You gotta love shipping to Canada.

I never asked for this.
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Bubbling with CO2 a boiling solution of potassium manganate, should work :

3 K2MnO4 (aq) + 2 CO2 --(boiling > 100*C)--> 2 KMnO4 (aq) + 2 K2CO3 (aq) + MnO2(s)

The reason is that H2CO3 (carbonic acid) is a stronger acid than [H2MnO4] (the theoretic formula of manganic acid) is .
We know this because, K2MnO4 has stronger basic properties than K2CO3...

So, I think that the reaction mechanism is:

2K+ + MnO42- + [H2CO3](aq) --(activation energy > 100*C)--> 2K+ + CO32- + [H2MnO4](aq)

...("H2CO3" and "H2MnO4" are placed in brackets, because this are unstable compounds and exist only in aquous solutions with H3O+ anions and the respective cations associated)

[H2MnO4] + 2 H2O <--> 2H3O+ + MnO42- <--> 3 H2O + [MnO3]

"MnO3" is a very unstable oxide of manganese .

[MnO3] <--> MnO2 (solid) + [O]2-

The atomic oxygen [O] makes a nucleophilic attack on the an other manganate ion , a double oxigen-manganese bond is formed, but, a single O-Mn bond is breaked ...

[O]2- + MnO42- + 2H3O+ --> (MnO)O42- + 2H3O+ --> MnO4- + 3H2O

As soon the permanganate ion is formed, a potassium ion get closer ( there was 2 K+ before the oxidation of the manganate ion to the permanganate ion):

K+ + MnO4- = KMnO4 (aq)

I'm now thinking if I also can use calcium manganate instead of the potassium one .
Random
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How would you form calcium manganate?
Kagutsuchi
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Does anyone know anything about aluminum permanganate? I'd try to prepare some but I'm unsure if it will worth it or it will be completely useless or even decompose instantly.
I'd try the
2KMnO4+dilute H2SO4----->2HMnO4+K2So4
3HMnO4+Al----->Al(MnO4)3+1.5H2 method.

P.S.:I'd be happy to hear if you tried it in pyrotechnics.

[Edited on 22-6-2015 by Kagutsuchi]
softbeard
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 Quote: Originally posted by Kagutsuchi 2KMnO4+dilute H2SO4----->2HMnO4+K2So4 3HMnO4+Al----->Al(MnO4)3+1.5H2 method.

Hey Kagutsuchi, I don't know the specifics of a theoretical 'aluminum permanaganate' but I'm sure the 2nd equation you've written is nonsense. There is no way the MnO4- anion is going to survive with a reductant like Al metal around. Much less yielding H2 gas in the process.
You'll get instant reduction of the MnO4- to Mn3O4 or Mn++ by Al metal, depending on mainly the pH.
Furthermore, I really doubt Al(MnO4)3 exists as a compound. I would venture that even if it does exist, it would be impossible to prepare in an aqueous medium.
MolecularWorld
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I can only see it in snippets, but the Kirk-Othmer Encyclopedia of Chemical Technology seems to suggest aluminum permanganate exists, can be crystallized, is unstable above 80*C, and can be formed from the reaction of cold solutions of potassium permanganate and aluminum sulfate: aluminum permanganate stays in solution, potassium aluminum sulfate precipitates (I couldn't actually see the equation).

gatosgr
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Have you found an easy way to turn decomposed KMnO4 from exposure to air back to KMnO4?

[Edited on 9-3-2016 by gatosgr]

clearly_not_atara
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 Quote: Have you found an easy way to turn decomposed KMnO4 from exposure to air back to KMnO4?

From the first page:

 Quote: The industrial method of heating MnO2 with KOH and using air as an oxidizer is doomed to failure.

This reaction instead almost always produces potassium manganate, K2MnO4. Unfortunately Mn(VI) is more stable than Mn(VII), despite being far less interesting. Stupid quantum mechanics demons. On the other hand:

https://en.wikipedia.org/wiki/Potassium_ferrate

 Quote: Edmond Frémy (1814 – 1894) later discovered that fusion of potassium hydroxide and iron(III) oxide in air produced a compound that was soluble in water.

In the case of iron, Fe(VI) is more stable than Fe(V) which is more stable than Fe(IV), although none of these is particularly stable. As a result only potassium ferrate is produced when iron oxide is heated with KOH. But luckily, Fe(VI) is a stronger oxidizing agent than Mn(VII), which means that this should happen:

2FeO4(2-) + 6MnO4(2-) + 5K+ + 5 H2O >> Fe2O3 + 6MnO4- + 5 KOH + 5 OH-

Excess ferrate can be removed, presumably, by simply allowing it to decompose.
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 Quote: Originally posted by MadHatter Given the recent CPSC(assholes) victory against FireFox, I'm sure that any convenient method will be appreciated. KMnO4 sales by pyro suppliers are now limited to 1 LB a year. There's still some OTC sources for now but we don't know how long that'll last.

So dont go to a pyro supplier. I got mine at a water depot, 5 lbs for about 45\$. Farmers use it to take iron out of well water for cattle. They didnt even give me a second glance when I bought it.
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KMnO4 from manganese metal electrodes: A paper and my results.

While looking for a method to prepare permanganates, I came across a paper from 1921 titled "The Electrolytic Production of Sodium and Potassium Permanganates from Ferromanganese", by Wilson et al.. I could get the full-text access with my uni library account and found very interesting information.

I'll just give the main points of the paper and technical details that are relevant for those who want to try it. And then I'll outline my two runs with a similar setup and the results I got.

The paper

On the electrolytic cell the authors say:
"The electrolyses were carried
out in a cell which consisted of a cylindrical glass jar (12.5
cm. in diameter by 15 cm. high), containing within it a
porous porcelain cup (5 em. in diameter by 12.5 em. high)
which served as a diaphragm. In this cell were placed the
two electrodes, the cathode (of 16-gage sheet iron, 3 em. by
15 em.) being within the cup, and the ferromanganese anode
standing in the jar outside, and about 3 em. away from
the cup. The freshly cast anodes had the dimensions 2.5
em. square by 15 em. in height. Electrical contact with the
anode was made by clamping a strip of brass, carrying a
binding post, against a freshly ground surface near the top."

On the anolyte and catholyte they used:
"It was found that the carbonate electrolyte (as compared
with the hydroxide) gives the purest product-uncon-
taminated with manganate-at the best efficiencies and for
the least expenditure of power. Hence, this electrolyte is
recommended for technical operation."

"The anolyte was 12 per cent sodium carbonate solution
(previously found to give about the optimum results), and
the catholyte was 8 per cent sodium hydroxide solution."

On anode composition:
"An experiment with spiegel iron (about 40
per cent Mn) showed practically zero yield of per-
manganate."

They state the cathode was simply iron and the current density was around 5 to 10 A/sq. dm.

In total they did 180 runs with varying conditons so this paper is just pure gold. If some of you want more of it I can paste the whole of it here (7 pages) or upload it somewhere (where?)

My two test-runs

My setup consists of a 600 ml beaker with a terracotta flower pot (acts as a porous membrane) in it. The hole of the flower pot was plugged with rubber stopper so that the flower pot can hold the anolyte without leaks
The beaker contains the catholyte. The levels of the electrolytes were adjusted to be at the same height to prevent the electrolytes from mixing.

The anode was a pile of 99.9% manganese metal flakes (ordered on OnyxMet.com. Btw they sell Mn flakes for 16 bucks/kg and ferromanganese 86% for 8 bucks/kg) that I put in the terracotta pot. Electrical contact was made with a long manganese flake touching the pile of Mn and connected to a power supply (+). I didn't use ferromanganese because I didn't have it, but it appears to work fine with pure Mn.

The cathode was a stainless steel plate.

1) The first run was done with 1M KOH in both cathode and anode compartment. A voltage of about 3V with a current of 100 to 200 mA was applied to the cell for a few hours. The catholyte stayed crystal clear with bubbles of gas being evolved at the cathode (oxygen I guess). The anolyte turned a purple color at first but within a few minutes turned instead a deep green/brownish color which indicates that potassium manganate was produced. I didn't process this batch and insted changed the composition of the anolyte:

2) The second run was done in the same conditions except:
- The anolyte used was about 10-15% Na2CO3. Catholyte was still KOH soln. because I didn't want to throw it away and thought it should have no influence on the run.
- The current supplied was about 500 mA and the voltage 5V, because I wanted to produce NaMnO4 in high enough conc. to be able to precipitate KMnO4 upon adding KCl and I don't have time to wait a week for the run to "complete". The temperature of the anolyte was about 25°C for the whole run (room temp 16°C). Total run time was about 50 hours.
The anolyte turned very dark purple and a fine mist of permanganate is produced so I used plastic wrap to contain it.

Results
Concentration: I titrated a sample of the anolyte with Mohr's salt (a stable source of Fe2+ ions) solution by dripping the anolyte into acidified mohr's salt solution until the purple colors stays. The anolyte was found to be about 0.4M in conentration with an error of say 5 or 10 % because I don't have very precise titration instruments (such as a burette).

Extraction of the permanganate
I filtered the anolyte on a GLASS filter funnel (no paper) and added conc. KCl solution to the filtrate. the undissolved solids appear to be MnO2 but I'll have to check that. The filtrate is currently evaporating at about 80°C on a hotplate with a fan blowing over it. Dark needle-like crystals appear to be forming upon cooling. I'll report back when I have separated the crystals.

I plan to do a more carefully controlled test on a 1 liter anolyte scale but I cannot go ahead right now because I don't have enough manganese left (the pieces are too small to work with) so I'll have to order some more Mn. I might try to cast a manganese metal plate.

First picture: The cell. Second picture: Mn flakes I used piled up in the flower pot.

Update: The crystals that formed from the evaporated anolyte are highly impure as evidenced by the presence of white crystals and brown crud among the needles of KMnO4. The crystals were dissolved in a small amount of water and the solution was heated up to boiling. It was filtered while hot on a preheated filter funnel (to remove MnO2, again). The filtrate yielded 0.37 g of small (and pure) needles of KMnO4 upon cooling.

(More KMnO4 could be extracted from the anolyte, but this run was just to prove that the method works and I didn't try to be particularly efficient in the extraction.)

The product was tested by adding a few drops of conc. H2SO4 to a crystal to form Mn2O7 and the reaction with ethanol was as expected: it ignited!

[Edited on 1-9-2016 by Romain]

[Edited on 1-9-2016 by Romain]

[Edited on 1-9-2016 by Romain]
zts16
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