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Author: Subject: H2o2 + Kmno4
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[*] posted on 13-4-2005 at 17:47
H2o2 + Kmno4


if h2o2 and Kmno4 are aded will the reaction create manganeese oxide h2o +o2??
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[*] posted on 13-4-2005 at 17:57


Yea, the peroxide will reduce the permanganate to [!Mn<sup>2+</sup>] Mn<sup>4+</sup> in acidic conditions.

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Perhaps you could get a table of standard reduction potentials to help answer similar questions in future. If you don't know how to use this, then I would love to reccomend some great books on chemistry; perhaps from the FTP?

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[edit - brain fart: two times two is four, not two!]

[Edited on 21-11-2005 by Ramiel]




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[*] posted on 14-4-2005 at 15:24


Are you quite sure it does not reduce Mn(VII) to Mn(IV), MnO2, instead? Mn(II) is very easily oxidized by any source of free oxygen.
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[*] posted on 14-4-2005 at 16:27


it reduces to MnO2



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[*] posted on 15-4-2005 at 15:04


I suppose outcome will be pH dependent.



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[*] posted on 20-11-2005 at 08:10


Sorry for raising an old thread, but it is the only one about this specific reaction

According to the posts here, the following reaction should be

KMnO4 + H2O2 -> H2O + O2 + MnO2 + ?

What happens to the potassium ion? I'm pretty sure that KMnO4 isn't just a catalyst!?

I tried this with a 30% H2O2, adding KMnO4 produced quite warm (hot) oxygen and there was brownish precipitate. Seemed like MnO2 allright.
Tried googeling, but the only reaction I found was

H2O2 + KMnO4 -> KMnO4 + H2O + O2 + MnO2, which is impossible to balance.




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[*] posted on 20-11-2005 at 10:51


2MnO4- + 8H+ +5H2O2 --> 2 Mn2+ + 8H2O +5O2 +10H+
And
O2 +4H+ + 2Mn(2+) --> 2MnO2 +8H+ +2H2O

Potassium is just a spectator ion here. And remember, all commercial peroxide contains (I think) phosphoric acid as a stabilizer, and only traces of acid are needed to get this reaction going as more is produced as it progresses.




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[*] posted on 20-11-2005 at 12:20


Quote:
Originally posted by rogue chemist
And remember, all commercial peroxide contains (I think) phosphoric acid as a stabilizer, and only traces of acid are needed to get this reaction going as more is produced as it progresses.


Um? As MnO4- decomposes, the solution will go basic, no?

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[*] posted on 20-11-2005 at 12:47


Under near-neutral conditions KMnO4 is reduced as follows:

2KMnO4 + H2O2 --> 2MnO2 + 2K(+) + 2OH(-) + 2O2

In ionic form:

2MnO4(-) + H2O2 --> 2MnO2 + 2OH(-) + 2O2

The liquid indeed becomes basic.

The MnO2 causes more H2O2 to be decomposed catalytically:

2H2O2 --> 2H2O + O2


The reduction to Mn(2+) only occurs at low pH.

[Edited on 20-11-2005 by woelen]




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[*] posted on 20-11-2005 at 12:59


Yes, but all commercial peroxide is stabilized with an acid, so would not the reaction I posted be more accurate?
Yes, that would happen with a neutral pH, but how many of us have unstabilized peroxide?:o.

I suppose the redox table I have pinned to the wall in front of me could just be lacking the basic half reaction...



EDIT: I was wrong, I just tried it. pH of peroxide ~4.5, pH after reaction ~8.5. I guess my redox table is just incorrect, or the acidic permanganete half reaction does not apply here.

[Edited on 20-11-2005 by rogue chemist]




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[*] posted on 20-11-2005 at 15:02


woelen is quite correct, as he already knows.

The amount of stabilizer is very small. Pure 35% H2O2 has a pH of 4.6, says K-O. BTW, some popular US OTC sources for 10, 20, 30, and 40% H2O2 use sodium phosphate as the stabilizer, the original manufacturer mixes TSP with H3PO4.
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