saps
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H2o2 + Kmno4
if h2o2 and Kmno4 are aded will the reaction create manganeese oxide h2o +o2??
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Ramiel
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Yea, the peroxide will reduce the permanganate to [!Mn<sup>2+</sup>] Mn<sup>4+</sup> in acidic conditions.
Teach a man to fish...
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?
Sincerely
- D
ps. Moved.
[edit - brain fart: two times two is four, not two!]
[Edited on 21-11-2005 by Ramiel]
Caveat Orator
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JohnWW
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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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budullewraagh
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it reduces to MnO2
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Borek
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I suppose outcome will be pH dependent.
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Ashendale
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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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The_Davster
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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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12AX7
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| 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?
Tim
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woelen
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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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The_Davster
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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? .
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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S.C. Wack
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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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