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Author: Subject: reduction of MnO2 with metabisulfite
sternman318
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[*] posted on 30-7-2011 at 12:26
reduction of MnO2 with metabisulfite


So I attempted this,in a very crude manner, using MnO2 from batteries and stump remover. I added a solution of NaHSO4 ( I dont have sulfuric acid) to some MnO2, then added some sodium metabisulfite. Solution got very hot, blah blah.

Let the solution settle and filtered it, and got a weird yellowish brown syrup. Letting it cool, I got the formation of yellow-brown crystals, which I assumed would have been NaSO4 ( I added a lot in the process), but are insoluble in water and dissolve with the addition of NaHSO4. Not sure how to go about testing this precip.

But now I have a slightly cloudy, pink solution, which hopefully has some Mn2+ in it.

My question: should I have used an acidified solution? I know this releases SO2, which directly reduces the MnO2, but because it's all in solution, would a neutral solution work better, or not at all?

[Edited on 30-7-2011 by sternman318]
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kryss
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[*] posted on 30-7-2011 at 12:39


Did it get hot before you added the metabisulphite to the bisulphate mixture?
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sternman318
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[*] posted on 1-8-2011 at 05:50


I believe so, probably because of the neutralization of some electrolyte ( I did not wait to rinse it, I merely wanted to see if results were possible).

However, I do have about 50 mls now of a nice rose pink solution, so it did work !

my question remains still, on whether sodium metabisulfite needs acidic conditions in order to reduce?
I assume yes, atleast for this reaction because the MnO2, going by its std. electrode potential, it needs acidic conditions. ( however nurdrage reduced MnO2 by pumping SO2 into an un-acidified solution of MnO2). Also, this will release SO2, which appears soluble enough in water that I shouldnt be loosing a whole bunch out of the solution when adding some acid.

I need to read up on my redox reactions, huh?
I will try a smalll scale experiment of reducing Cu2+ later.
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UnintentionalChaos
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[*] posted on 4-8-2011 at 14:54


Manganese (II) sulfite trihydrate is poorly soluble in water and represents a significant loss as solid.



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sternman318
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[*] posted on 5-8-2011 at 15:48


Quote: Originally posted by UnintentionalChaos  
Manganese (II) sulfite trihydrate is poorly soluble in water and represents a significant loss as solid.


Ah, thank you! I did not even think about that. Some reading reveals that most sulfites are insoluble ( unless in an acidic environment). I seem to be having a hard time reading about reduction with either metabisulfite or sodium sulfite ( woelen has mentioned that the metabisulfite anion essentially forms two bisulfite ions in solution). So I guess that answer to my question is that it does need acidic conditions, from what I can tell.
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[*] posted on 5-8-2011 at 17:34


Quote: Originally posted by sternman318  

my question remains still, on whether sodium metabisulfite needs acidic conditions in order to reduce?
I assume yes, atleast for this reaction because the MnO2, going by its std. electrode potential, it needs acidic conditions. ( however nurdrage reduced MnO2 by pumping SO2 into an un-acidified solution of MnO2).


Bubbling SO2 into water creates an acidic solution, albeit weak.

SO2 + H2O <---> H2SO3




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sternman318
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[*] posted on 5-8-2011 at 17:43


With that said, what is the active species? SO2, or sulfurous acid/bisulfite/sulfite? ( yes I know H2SO3 doesnt really exist in solution)
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