ThoughtsIControl
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Question - C and MnO2 separation
I've extracted some manganese dioxide from a carbon-zinc battery. Carbon is an impurity that I'd like to remove from the MnO2.
My question is:
How do you remove the carbon impurities from the gooey manganese dioxide found in batteries?
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clearly_not_atara
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Kind of a weird guess but maybe piranha solution would work? As long as it doesn't get reduced to MnSO4...
[Edited on 04-20-1969 by clearly_not_atara]
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DraconicAcid
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I think the best way to separate these two would be to reduce the MnO2 to a water-soluble Mn(II) salt, filter out the carbon, then reoxidize the
manganese to MnO2.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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njl
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If the mush isn't contaminated with anything too harmful, I would think hitting it with a torch in open air would be a good way to remove carbon.
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ThoughtsIControl
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Thank you guys.
I think that the piranha solution would favor the MnSO4 because of how electronegative it is compared to hydrogen peroxide.
I think I am going to take the advice of DraconicAcid here.
I'll probably end up using hydrochloric acid to dissolve it..
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j_sum1
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You need the wisdom of Nurdrage on this.
Here are some of his videos that I think you will find useful. (Total watch time <half hour.)
https://www.youtube.com/watch?v=knc1lSupAwQ
https://www.youtube.com/watch?v=2gXByJkg0iY
https://www.youtube.com/watch?v=BLJgBSrhZI8
https://www.youtube.com/watch?v=4JXbF4QgxJk
https://www.youtube.com/watch?v=nvMVlhBmv7M
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ThoughtsIControl
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Thanks j_sum1 haha.
Actually, I really enjoy watching the videos of Nurdrage.
I'm checking these videos out. I plan on playing around with the chemistry of manganese dioxide since I already have it. I'm going to write out a lab
procedure for each little experiment. Oh! While I'm thinking about it -->
I plan to make myself some nitric acid and perhaps even sulfuric acid using the method of having an electron permeable membrane separating the
mixtures. Maybe you've heard of it as "clay pot electrolysis"
Would this be a good idea in order to get a fair yield or should I just wait for the rest of the month for my glassware to come in the mail so I can
distill it?
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mysteriusbhoice
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I found something better than a clay pot!
mix concrete and sand with whatever saturated salt solution you want to use it for nitrate/sulfate/etc it will wont matter at first since it will act
as a true ion exchange composite!!
The resulting concrete cannot bear load and is quite brittle but we aint gonna use it for building anyway and its compact enough to keep its shape.
its resistant to both acids and strong alkali though initial bubbling may occur at first due to carbonates present but adding calcium sulfate seems to
make it fully acid resistant and again its fragile.
let the concrete cure in a pot and with an inner mold to create a chamber and it should look something like this.
I call this my ghetto style Ion exchange composite.
Just like any membrane it can suffer from fouling but when it does it will still be porous and its performance will just go down.
However i still advice against using it in solutions containing sticky organics like sugars phenolics etc..
[Edited on 20-5-2020 by mysteriusbhoice]
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DraconicAcid
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Make sure there's a reducing agent in solution, or you risk getting chlorine gas formed.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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j_sum1
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Let me concur with Draconic Acid here.
The first major incident in my lab was producing a cloud of Cl2 gas that I was not expecting -- doing exactly what you are doing now. It was less
than 20g of battery paste and a little HCl in a small flask. It foamed all over the concrete floor and produced enough gas that I held my breath and
abandoned the shed. I returned several hours later for clean up.
I would tend to go for sulfuric rather than hydrochloric.
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chornedsnorkack
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Sulphuric acid is not a reducing agent.
Of acids which are reducing:
H2SO3 is itself volatile and smelly.
H2C2O4 is not volatile and oxidises to CO2 which is feebly toxic compared to SO2 or
Cl2
H3PO2 is reducing, but phosphates are commonly poorly soluble. Also, while H3PO2 is not itself volatile,
it may dismute to PH3, which is volatile and toxic.
HI oxidizes to I2 which is much less volatile than Cl2. Especially in a solution rich in I-
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mysteriusbhoice
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new batteries contain MnO2 + junk and spent ones contain Mn3O4 + MnO2 + other junk
Mn3O4 can be dissolved in acetic or nitric acid then anodically oxidized to form MnO2 again
and yes i found that out due to it contaminating my PbO2 plating solution and the oxide that produced was MnO2..... instead
if you used nitric acid you could probably just burn it to give MnO2 but it wont be pure unlike electrolysing it with graphite at higher voltage to
get fine powder or flakes.
So here would be what I would do and would only work if the battery was spent.
mix the Mn3O4 with either acetic acid or nitric acid.
filter solution and electrolyze and get a crop of clean MnO2 from the anode sludge and it should be insoluble in acetic or nitric acid.
For new batteries use H2SO4 and do the same
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ThoughtsIControl
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Alright I'll probably whip up some oxalic acid before I start to use as a reducing agent.
How much of it am I supposed to use during the synthesis of sulfuric and/or nitric acid? Well, how do I go about calculating this..
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