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Author: Subject: Permanganates
blogfast25
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[*] posted on 27-8-2011 at 04:32


Quote: Originally posted by Chemistry Alchemist  
i want manganese metal for some time now, is it possible to add HCl to the KMnO4 = KCl + MnCl2 + Cl2 and then with the KCl and MnCl2 solution, add aluminum to it to precipitate the manganese? is this a good method to use?


Again, this doesn't really belong in this thread.

To answer your question:

the reduction reaction 3 Mn2+(aq) + 2 Al(s) === > 3 Mn(s) + 2 Al3+(aq) can indeed proceed. Unfortunately manganese is a very reactive metal and any Mn formed would immediately be oxidised completely by water, first to Mn(OH)2, which would then in turn be oxidised again (by air oxygen) to MnO2. No manganese metal can be obtained this way. That's in contrast to for instance copper, which can be obtained that way but copper is far more 'noble' than manganese...

The correct equation for the reaction you were trying to write is in fact:

KMnO4 + 8 HCl === > KCl + MnCl2 + 5/2 Cl2 + 4 H2O






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[*] posted on 22-9-2011 at 22:52
Potassium Manganate disporporting


I done the casual reaction of KMnO4 + Break fluid, this Produces small amounts of Potassium Manganate when the potassium Permanganate decomposes due to the heat, i washed the remains in water to clean the Manganese oxides, and left nice deep green solution of manganate, now, i wanted to dry it to a powder so i poored it into a evaporating dish and left it in the sun, an hour later i came back and the solution had turned clear with a brown precipitate, what has happened here? is the brown precipitate dissolved manganese salt being oxidized in the air to form the Insoluble oxide and left a solution of potassium ____? please explain if possible :)



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[*] posted on 23-9-2011 at 04:35


Quote: Originally posted by Chemistry Alchemist  
I done the casual reaction of KMnO4 + Break fluid, this Produces small amounts of Potassium Manganate when the potassium Permanganate decomposes due to the heat, i washed the remains in water to clean the Manganese oxides, and left nice deep green solution of manganate, now, i wanted to dry it to a powder so i poored it into a evaporating dish and left it in the sun, an hour later i came back and the solution had turned clear with a brown precipitate, what has happened here? is the brown precipitate dissolved manganese salt being oxidized in the air to form the Insoluble oxide and left a solution of potassium ____? please explain if possible :)


Manganates (Mn [+VI] or MnO<sub>4</sub><sup>2-</sup>;) are inherently unstable, cannot be isolated and on attempting to isolate, disproportionate: Mn (VI) === > Mn (IV) + Mn (II), so brown MnO2 precipitates...

[Edited on 23-9-2011 by blogfast25]




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[*] posted on 23-9-2011 at 04:39


Thanks :) what would be the potassium salt?



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[*] posted on 23-9-2011 at 05:23


K2MnO4. But as a solid this is almost impossible to synthesize. In solution it's emerald green (depending concentration, obviously...)

[Edited on 23-9-2011 by blogfast25]




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[*] posted on 23-9-2011 at 05:25


but isnt K2MnO4 a green solution?



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[*] posted on 28-9-2011 at 04:31


Quote: Originally posted by blogfast25  

disproportionate: Mn (VI) === > Mn (IV) + Mn (II), so brown MnO2 precipitates...

[Edited on 23-9-2011 by blogfast25]


They disproportionate like:
Mn(VI) = Mn(IV) + Mn(VII)

3MnO<sub>4</sub><sup>2-</sup> + 2H<sub>2</sub>O = MnO<sub>2</sub> + 2MnO<sub>4</sub><sup>-</sup> + 4OH<sup>-</sup>




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[*] posted on 28-9-2011 at 08:45


Quote: Originally posted by rstar  
They disproportionate like:
Mn(VI) = Mn(IV) + Mn(VII)

3MnO<sub>4</sub><sup>2-</sup> + 2H<sub>2</sub>O = MnO<sub>2</sub> + 2MnO<sub>4</sub><sup>-</sup> + 4OH<sup>-</sup>


I only once made some K2MnO4 solution and didn't see it form any purple KMnO4.




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[*] posted on 29-9-2011 at 02:26


Quote: Originally posted by blogfast25  

I only once made some K2MnO4 solution and didn't see it form any purple KMnO4.


Your solution might be quite basic, in which green K2MnO4 is stable, but as your solution will become acidic it will gradually change to purple KMnO4.

Try adding some acids ;)




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[*] posted on 29-9-2011 at 05:02


Quote: Originally posted by rstar  
Your solution might be quite basic, in which green K2MnO4 is stable, but as your solution will become acidic it will gradually change to purple KMnO4.

Try adding some acids ;)


My solution was very alkaline: MnO2 + KOH + KClO3 fusion product, leached. But I haven't got any right now...




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[*] posted on 4-1-2012 at 00:53


This is possible to use air oxygen instead of potassium nitrate in first reaction?I want to use 50% potassium Hydroxide solution and MnO2 .

I decide to enter air oxygen in this boiling solution by air pump.Does it possible?
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[*] posted on 4-1-2012 at 09:27


Quote: Originally posted by Waffles SS  
This is possible to use air oxygen instead of potassium nitrate in first reaction?I want to use 50% potassium Hydroxide solution and MnO2 .

I decide to enter air oxygen in this boiling solution by air pump.Does it possible?


I doubt that Mn (IV) can be oxidised to Mn (VI) in solution by air oxygen, In fact I'm pretty sure it isn't possible. It requires fusion at least, in which case vigourous aireation might do the trick.




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[*] posted on 4-1-2012 at 11:37


Thanks dear blogfast25,
What about pure oxygen or ozone(O3)?
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[*] posted on 5-1-2012 at 10:42


Ozone? Where are you gonna get that from?

Nah, good old fusing of MnO2/KOH with KNO3 or KClO3 does it nicely.




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[*] posted on 5-1-2012 at 12:56


Quote: Originally posted by blogfast25  
Ozone? Where are you gonna get that from?

I have a ozone generator(its for swimming pool) :D
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[*] posted on 5-1-2012 at 13:12


I have always found it interesting that ozone oxidizes chlorine, but not carbon tetrachloride.
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[*] posted on 6-1-2012 at 08:12


@ Waffles SS
Quote:
This is possible to use air oxygen instead of potassium nitrate in first reaction?I want to use 50% potassium Hydroxide solution and MnO2 .

I decide to enter air oxygen in this boiling solution by air pump.Does it possible?


Yes but not 50% solution. Commercial methods sparge air through MnO2/KOH mixtures but they use molten KOH at highish temperatures.

See the article I wrote and posted as http://www.sciencemadness.org/scipics/MnOXY.doc

Page 12ff.

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[*] posted on 29-4-2012 at 13:03


Is it possible to synthesize in a relatively easy way calcium manganate? Or any other alkali earth, whose chemistry is probably the same...



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[*] posted on 30-4-2012 at 09:13


If manganate (MnO<sub>4</sub><sup>2-</sup>;) - Mn [VI] - is what you mean then no. It's not very stable at all. Heating MnO2 with KOH and a strong oxidiser like chlorate or nitrate yields it but it decays quickly.

[Edited on 30-4-2012 by blogfast25]




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[*] posted on 30-4-2012 at 12:15


Yes. Manganese, technetium and rhenium have always been quite interesting metals (at least to me). Oh, also chromium.



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[*] posted on 18-7-2012 at 11:07


Has anyone tried preparing calcium permanganate? Maybe it's easier to isolate, anyone knows its properties?

By the way, I have found that sodium permanganate is available as about 30% solution, so it should be definitelly possible to make it.

A very small amount is produced with the reaction of mno2 and bleach with sodium hydroxide. So I am thinking, if we would put MnO2 into hypochlorite cell and electrolyse it, maybe the continuous slow supply of Cl2 and hypochlorite in a basic solution would over time make a bigger amount of permanganate.
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[*] posted on 18-7-2012 at 11:24


Quote: Originally posted by Random  
Has anyone tried preparing calcium permanganate? Maybe it's easier to isolate, anyone knows its properties?

By the way, I have found that sodium permanganate is available as about 30% solution, so it should be definitelly possible to make it.

A very small amount is produced with the reaction of mno2 and bleach with sodium hydroxide. So I am thinking, if we would put MnO2 into hypochlorite cell and electrolyse it, maybe the continuous slow supply of Cl2 and hypochlorite in a basic solution would over time make a bigger amount of permanganate.


You need strong mixing of the MnO2 in the satured solution of sodium chloride, and it work, I already make some with the same process.




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[*] posted on 18-7-2012 at 11:28


Quote: Originally posted by plante1999  
Quote: Originally posted by Random  
Has anyone tried preparing calcium permanganate? Maybe it's easier to isolate, anyone knows its properties?

By the way, I have found that sodium permanganate is available as about 30% solution, so it should be definitelly possible to make it.

A very small amount is produced with the reaction of mno2 and bleach with sodium hydroxide. So I am thinking, if we would put MnO2 into hypochlorite cell and electrolyse it, maybe the continuous slow supply of Cl2 and hypochlorite in a basic solution would over time make a bigger amount of permanganate.


You need strong mixing of the MnO2 in the satured solution of sodium chloride, and it work, I already make some with the same process.


Yeah, I thought stirring would be useful there. What concentration was it? Could it be useful solution for example for oxidation of alcohols?
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[*] posted on 17-8-2012 at 10:57


I don't remember perfectly but yes it could be used to oxidize alcohols. Reaction with alcohol is fast and exothermic.



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[*] posted on 19-8-2012 at 19:57


I noticed that most (if not all) of this thread is about alkali metal permanganates. Does anyone have information on the permanganates of transition metals, poor metals, etc.? I imagine CuMnO4 would be an amazing oxidizer :D
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