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Topic: Reducing MnO2  (Read 25821 times)

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Offline science2000

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Reducing MnO2
« on: December 21, 2005, 11:20:56 PM »
Does anyone know of a way to reduce magnanese(IV) dioxide to manganese(II), besides adding HCl to it? What would it do if I attempted to dissolve it in dilute acetic acid?

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Offline lemonoman

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Re:Reducing MnO2
« Reply #1 on: December 22, 2005, 01:10:03 AM »
"MnO2 particles are unstable in anoxic environments, and will be reduced to dissolved Mn2+" (from Joceline Boucher's Ocean Studies 212 Lectures).

If you can obtain an 'anoxic environment' - i.e. one with no oxygen...as in water with 0.00% dissolved oxygen - then that should do the trick nicely.

Two notes, is that anoxic water is probably hard to come by.  I know you can get rid of most dissolved gases in liquids by freezing then melting, refreezing and melting, over and over...

Second thing is that "Mn2+ is thermodynamically unstable and can eventual oxidize to MnO2(s) in the presence of oxygen" (Same source).  So if you have an anoxic aqueous environment that is open to the atmosphere, be prepared for an amount of reoxidation at the surface of the solution.

Best of luck!
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Offline Borek

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Re:Reducing MnO2
« Reply #2 on: December 22, 2005, 03:28:50 AM »
Acidify and add some reducing agent. Acetic acid will be probably too weak, I would rather try sulphuric with reducing agent like Fe2+, or some organic - like ascorbic acid.
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Offline Borek

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Re:Reducing MnO2
« Reply #3 on: December 22, 2005, 06:26:31 AM »
Quote from: lemonoman on December 22, 2005, 01:10:03 AM
"MnO2 particles are unstable in anoxic environments, and will be reduced to dissolved Mn2+" (from Joceline Boucher's Ocean Studies 212 Lectures).

If you can obtain an 'anoxic environment' - i.e. one with no oxygen...as in water with 0.00% dissolved oxygen - then that should do the trick nicely.

Not questioning the lectures, but are you sure it read the information correctly? Lack of oxygen is not enough for the reduction to take place - you need some kind of reductor to do the trick. I believe sea water contains lots of organic reducing agents (lots meaning many different, not very large quantities).
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Offline lemonoman

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Re:Reducing MnO2
« Reply #4 on: December 26, 2005, 12:29:49 AM »
Quote from: Borek on December 22, 2005, 06:26:31 AM
Not questioning the lectures, but are you sure it read the information correctly? Lack of oxygen is not enough for the reduction to take place - you need some kind of reductor to do the trick. I believe sea water contains lots of organic reducing agents (lots meaning many different, not very large quantities).

Noted; thanks for the clarification!
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metalriffzz

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Re:Reducing MnO2
« Reply #5 on: December 27, 2005, 12:46:30 AM »
Theres nothing complicated here. Just add your MnO2 to a little hydrogen peroxide and any acid. Vinegar works fine. I found this out trying to get rid of those nasty stains KMnO4 leaves behind.
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Offline science2000

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Re:Reducing MnO2
« Reply #6 on: December 30, 2005, 10:07:35 PM »
Wouldn't the oxygen produced from the decomposing peroxide oxidize the water and prevent the reduction? I don't know if I understand.
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metalriffzz

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Re:Reducing MnO2
« Reply #7 on: January 07, 2006, 06:35:04 PM »
An acidified solution of H2O2 will cause the Mn to go all the way to the +2 oxidation state from +4. Make sure the acid is added first and little or no oxygen will evolve from the solution
« Last Edit: January 07, 2006, 06:38:18 PM by metalriffzz »
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Offline woelen

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Re:Reducing MnO2
« Reply #8 on: January 12, 2006, 07:11:11 AM »
Keep in mind, that not all MnO2 reacts in this way easily. Strongly calcined MnO2 (very fine crystalline stuff) hardly dissolves and even in concentrated acids it hardly is capable to react. This is a problem with many oxides. When they are calcined and become crystalline, then they dissolve with great difficulty.

Freshly precipitated amorphous MnO2 indeed quickly dissolves in an acidified solution of H2O2, forming colorless Mn(2+) ions.
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Offline science2000

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Re:Reducing MnO2
« Reply #9 on: January 12, 2006, 09:25:36 PM »
What about MnO2 from a dry-cell battery? That's the stuff I'd use.
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Offline woelen

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Re:Reducing MnO2
« Reply #10 on: January 13, 2006, 06:15:36 PM »
Quote from: science2000 on January 12, 2006, 09:25:36 PM
What about MnO2 from a dry-cell battery? That's the stuff I'd use.
Probably that dissolves quite well, because it also needs to react easily in the battery. The only problem is that this MnO2 is terribly impure. Mostly it is mixed with carbon as far as I remember and it also contains a lot of other stuff (it is humid, due to some electrolyte).
A good source of MnO2 are ceramics and pottery supplies. These also have many other interesting metal salts (e.g. salt of Co, Fe, Cu, Ni, Bi, Cr).
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Offline Borek

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Re:Reducing MnO2
« Reply #11 on: January 13, 2006, 06:33:01 PM »
It must contain electrolyte to conduct electric current.
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Offline jdurg

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Re:Reducing MnO2
« Reply #12 on: January 13, 2006, 07:56:18 PM »
Quote from: Borek on January 13, 2006, 06:33:01 PM
It must contain electrolyte to conduct electric current.

Many times that electrolyte is sodium hydroxide.
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Offline science2000

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Re:Reducing MnO2
« Reply #13 on: January 13, 2006, 08:28:04 PM »
What are some impurities in battery MnO2? Any heavy metals or things I should be warned about?
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Offline Alberto_Kravina

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Re:Reducing MnO2
« Reply #14 on: January 14, 2006, 04:36:01 AM »
Quote from: science2000 on January 13, 2006, 08:28:04 PM
What are some impurities in battery MnO2? Any heavy metals or things I should be warned about?
An impurity could be Zn2+
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