underground
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purification of manganese dioxide from batteries
As the subject says, is there any way to purify manganese dioxide from batteries ?
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j_sum1
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yep.
I can't search right now but there is plenty on the board I think.
Nurdrage also has a procedure for purifying MnSO4 that is appropriate for purifying battery gunk.
I can give more details later if no one else does.
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j_sum1
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Right. Back.
Relevant thread here http://www.sciencemadness.org/talk/viewthread.php?tid=44590#...
There are probably many more. This one dates from shortly after my joining the board and attempting a similar project. My goal was to start with
battery gunk and finish with as many elements as possible.
The first thing to realise is that the stuff in batteries is a horrible mixture to work with and notoriously messy. It is not just MnO2 but a paste
typically composed of a ZnCl2(aq) solution or NH4Cl2(aq) solution, MnO2 powder and graphite powder.
My procedure is as follows:
Mix contents of one D cell with water and filter. Filtrate contains ZnCl2. The zinc can be reduced via electrolysis.
Digest the filter cake in concentrate HCl. This liberates copious quantities of Cl2 gas. If you are not able to deal with this then you should
choose another acid. I added the HCl dripwise to the battery gunk in a flask using a separatory funnel and bubbled the evolved gas through a bucket
of NaOH solution.
Filter the contents of the flask. The residue should be mostly graphite powder. The filtrate contains MnCl2. It will also have significant
impurities of iron.
The following part is a procedure I found from Nurdrage: https://www.youtube.com/watch?v=BLJgBSrhZI8
Separate out one third of the MnCl2 solution. Add to this third an alkaline solution to form a precipitate. I used NaOH. Ammonia and Na2CO3 are
also suggested and have some advantages. (Refer to blogfast's comments in the link above.
Filter and wash with water. The filtrate contains Na+ solution and can be discarded. The filter cake contains Mn(OH)2 or MnCO3 depending on the
alkali used.
Add the washed filter cake to the remaining two thirds of solution. It should mostly redissolve. However it is critical that some precipitation
remains undissolved and that the solution appears cloudy. The Fe impurities should be part of this suspended solid.
Filter and retain the filtrate. This should be relatively pure MnCl2 (Or MnSO4 if you used sulfuric acid in the first step.)
Recrystallise this solution to retrieve pure MnCl2 (or the sulfate). You should have lovely light pink crystals.
To convert these crystals to MnO2, first dissolve in water. Then add alkali again. NaOH, ammonia or Na2CO3. Filter out the precipitate formed and
wash with water.
Calcine the filter cake to form MnO2.
Finally, reduce to Mn via a thermite reaction using Mg powder of Al powder.
I admit to not having actually completed all these steps from beginning to end. I have done some and also played a little with pottery grade MnO2.
This is a fun project and something of a rite of passage I think. But it is time-consuming and messy. Yields are low since Mn is thrown out at
several steps.
If your target is MnO2 for some other purpose then I think you would be better served to order some from a pottery supplier. If you want to learn
some chemistry techniques then you could do a lot worse than disassembling a few batteries.
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blogfast25
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underground:
If you search for it, 'DerAlte' put together a handy guide on the subject of Mn recovery from battery electrolyte, on this forum (as a *.doc file,
IIRW)
Personally I'm now more in favour of using the oxalic acid (very OTC) route (to reduce the MnO2/Mn2O3 mixture) than the HCl one, because of the
chlorine issue. Nurdrage also has a video on the oxalic acid route.
| Quote: | | Finally, reduce to Mn via a thermite reaction using Mg powder of Al powder. |
Although possible, the MnO2 thermite is one of the trickier ones. Get it wrong and yield can be very, very low. Very fast deflagrations of MnO2/Al
mixtures are fairly common due to the low BP of Mn. It's not w/o danger for inexperienced experimenters. Careful formulation required!
[Edited on 25-4-2015 by blogfast25]
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j_sum1
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Surprised to hear you say that about Mn/Al thermite blogfast. I have done a few and have not had any significant flare-ups. (Now, CuO is
fast.) I agree on the low yield though. The product of the thermite has never in my experience been a nice ingot but instead has been a porous
mixture with presumably Al2O3. Rather brittle too. But it does show that brownish surface that you get on a piece of Mn.
I actually popped back to make sure credit was given where it was due.
Blogfast's helpful knowledge of the work done on this:
Quote: Originally posted by blogfast25  |
But the method (Nurdrage's for argument's sake, but it's actually based on a Russian patent first presented here by 'peach') relies on the immense
difference in solubility products between Mn(OH)2 (Ks = 1.8 x 10<sup>-11</sup> and Fe(OH)3 (Ks = 4 x 10<sup>-34</sup>, http://bilbo.chm.uri.edu/CHM112/tables/KspTable.htm ).
One can show mathematically that the conditions (the lowest pH) in which Fe(OH)3 still precipitates are such that Mn(OH)2 does not. Both are
'insoluble' but Mn(OH)2 is far, far more soluble than Fe(OH)3.
Now, MnCO3 does have a slightly lower Ks = 1.9 x 10<sup>-13</sup> (than Mn(OH)2) but still about 10<sup>21</sup> larger than
that of Fe(OH)3. Solubility is relative, as shown by the concept of solubility products. |
Of course, now you'll have me looking up the oxalic acid route. I haven't properly finished the last one. My immediate excuse is that I don't have
oxalic acid at the moment. (It is on my shopping list, but a few hundred dollars from the top.) Of course that is just an excuse. Damn you
blogfast. You've made more work for me.
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blogfast25
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| Quote: Originally posted by j_sum1 | | Of course, now you'll have me looking up the oxalic acid route. I haven't properly finished the last one. My immediate excuse is that I don't have
oxalic acid at the moment. (It is on my shopping list, but a few hundred dollars from the top.) Of course that is just an excuse. Damn you
blogfast. You've made more work for me. |
In essence, replace HCl with an oxalic acid solution, after preliminary purification of the battery crud (removal of water soluble compounds).
Oxalic acid is reduced to CO2, with sulphuric acid to provide acidity, so much fizzing occurs. It's strongly exothermic, so the oxalic acid solution
has to be added slowly with constant stirring and with intermittent cooling if needed (to avoid boiling over). Reaction is over when fizzing
subsides.
HOOC-COOH === > 2 CO2 + 2 H<sup>+</sup> + 2 e
MnO2 + 4 H<sup>+</sup> + 2 e === > Mn<sup>2+</sup> + 2 H2O
After filtration, one obtains an MnSO4 solution. Some Fe is also dissolved, if present in the source.
Forget about Mn ingots from the MnO2 thermite: part of the Mn evaporates, causing a very irregular, often porous regulus. It works but quite poorly
compared to many other thermites. My Mn yield has never exceeded about 30 %.
[Edited on 25-4-2015 by blogfast25]
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j_sum1
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I just spent some time searching for that link by DerAlte that you mentioned. I didn't find it. If you have a link easily available, i would love to
see it.
It's late here at right now. Heading to bed.
J.
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blogfast25
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| Quote: Originally posted by j_sum1 | I just spent some time searching for that link by DerAlte that you mentioned. I didn't find it. If you have a link easily available, i would love to
see it.
|
Unfortunately it's buried in one of the many and often lengthy threads by DerAlte on the subject of manganese. I'll have a look.
Ok, gotcha: DerAlte's remarkably complete monography on Mn oxides:
http://www.sciencemadness.org/talk/viewthread.php?tid=8480&a...
Nurdrage on the oxalic acid method:
https://www.youtube.com/watch?v=2gXByJkg0iY
[Edited on 25-4-2015 by blogfast25]
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byko3y
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2MnSO4 + 8H2O + 5K2S2O6(O2) = 2KMnO4 + 4K2SO4 + 8H2SO4
This is a good way to make permanganate, btw.
I tried to perform the reaction, and at first I thought I failed, but after few hours solution turned purple. But after standing for a few days I have
almost clear solution, because MnO2 catalyzes KMnO4 decomposition to MnO2. So, basically, I've got a pure MnO2 precipitate. You can make it faster by
mixing the KMnO4 and MnSO4. This way you can make ultra pure γ-MnO2 with high catalytic activity.
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underground
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| Quote: Originally posted by byko3y | 2MnSO4 + 8H2O + 5K2S2O6(O2) = 2KMnO4 + 4K2SO4 + 8H2SO4
This is a good way to make permanganate, btw.
I tried to perform the reaction, and at first I thought I failed, but after few hours solution turned purple. But after standing for a few days I have
almost clear solution, because MnO2 catalyzes KMnO4 decomposition to MnO2. So, basically, I've got a pure MnO2 precipitate. You can make it faster by
mixing the KMnO4 and MnSO4. This way you can make ultra pure γ-MnO2 with high catalytic activity. |
What is K2S2O6(O2) ?
Can you give more details about the above process ? With this way you can purify mno2 from batteries ?
P.S. I read also somewere that mno2 from alkaline batteries is better than mno2 from carbon-zinc batteries.
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blogfast25
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| Quote: Originally posted by underground | What is K2S2O6(O2) ?
Can you give more details about the above process ? With this way you can purify mno2 from batteries ?
P.S. I read also somewere that mno2 from alkaline batteries is better than mno2 from carbon-zinc batteries. |
He meant potassium persulphate (K<sub>2</sub>S<sub>2</sub>O<sub>8</sub>, a very powerful oxidiser. It's also quite hard to get.
It's been claimed that the filler from alkaline batteries contains no graphite (because the KOH solution acts as the electrolyte, instead of the
graphite conductor). In that case careful washing of the alkaline battery filler should yield fairly pure MnO<sub>2</sub>, without the
chemical kerfuffle that's needed to extract the manganese from dry batteries. Whether the 'no graphite in alkalines' claim is true I do not know.
There's only one way to find out!
Having said that, in spent alkaline batteries the Mn will be no longer as MnO2 but also as lower oxides because that's what using the
batteries does to the initial MnO2. So you would have to sacrifice new batteries.
[Edited on 25-4-2015 by blogfast25]
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Metacelsus
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When I needed some manganese(ii) sulfate, I used dead batteries because first, I had no other use for them, and second, they needed less reducing
agent to get to Mn(ii). They were alkaline batteries, but there was some inert material (graphite, I think, but I didn't test it).
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byko3y
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Yes, I was talking about peroxydisulfate. In my country you can easily obtain ammonium persulfate from electronics store (I don't know how it should
be called correctly, the one where you can buy solder, chips, multimeters, etc.), it is used for etching.
Sorry, I don't have any more information on that reaction (maybe AgNO3 as a catalyst will help you, though I don't know, didn't use it).
Also, here MnSO4 is dirt chip, like 30$ for 25 kg bag. Do you have some regulations on selling of manganese compounds?
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j_sum1
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Thanks for the link blogfast. I will take a read of it.
BTW, the oxalic acid route is far from otc where I live. I haven't been able to find any products that contain a significant quantity of oxalic acid.
Not even pottery suppliers. Rust stain remover is lactic acid around here. I suppose there is always tea and rhubarb leaves. Also H2SO4 is not
OTC. I do have some but I wouldn't vouch for its purity. I don't suppose it would matter too much in this instance.
Would another organic acid work instead? I am thinking maybe citric, lactic or tartaric?
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blogfast25
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| Quote: Originally posted by j_sum1 | | BTW, the oxalic acid route is far from otc where I live. Would another organic acid work instead? I am thinking maybe citric, lactic or tartaric?
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Wow. No wood cleaner (oxalic acid) in Oz? The streets of eBay UK are lined with OA sellers! Funny old world...
MnO2 is of course a very powerful oxidiser: HCl, H2O2, sulphites, oxalic acid, it's all easily oxidised by it in acidic conditions.
But other carboxylic acids? Not that I know of.
Metabisulphate (wine making) should work though. See nurdrage's use of SO2.
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j_sum1
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Wood cleaner? I dunno. I will have to take a look. But I haven't come across it before.
And Oz is a big place. What is available in one location may not necessarily be available elsewhere. Automotive parts places stopped selling battery
acid in my town about five years ago. It might be different in other places.
Yep. It is a funny old world. There is no consistency on what is considered injurious to the public good.
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blogfast25
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There's a thread by Magpie (IIRW) on preparing OA from sucrose and nitric acid (believe it or not). I tried it and it works. But preparing such a
cheap compound that way isn't a very practical way to obtain it, is it?
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byko3y
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j_sum1, I bought all the battery acid in my town, so now there's no acid sold here.
However, I can order it online. I'm pretty sure you can order it too http://www.chemicals.co.uk/battery-acid .
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j_sum1
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Oh. I can get it if needed. I have made my own H2SO4. I have access to a range of acids at my place of work but I prefer to not mix work and home
projects. I am compiling a shopping list for argentscientific who looks like being a good supplier for Aus customers and I am sure will put together
an acid pack for me. The problem is that I have another shopping list for a few elements and reagents at three other places. And I also have my eye
on some decent glassware which might have to include a heating mantle. My bank balance can't take too many hits in a short space of time.
At last count I had 16 incompleted projects in my lab manual too. I really shouldn't be in too much of a hurry.
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