Upsilon
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Isolating manganese metal from potassium permanganate
I want to obtain a sample of manganese metal, and I want to know how to go about doing so. I have a theory in mind but I'm not sure if it will work.
First I would pass chlorine gas through a solution of potassium permanganate, which I hope causes the MnO4- ion to dissociate into MnO2 and O2:
2KMnO4 + Cl2 -> 2MnO2 + 2KCl + 2O2
I would then filter off the insoluble MnO2, and reduce it with aluminum powder:
3MnO2 + 4Al -> 2Al2O3 + 3Mn
Any ideas on this?
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DraconicAcid
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I'm sure you can find a better reducing agent than chlorine gas.
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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Upsilon
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Any suggestions? Chlorine gas really isn't hard to generate and work with, since it can be easily made with common household chemicals. If I reduce it
with any kind of acid, it would create manganese heptoxide, which is a liquid that decomposes in contact with water and is very explosive.
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DraconicAcid
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It takes very concentrated acid to make Mn2O7. Take your KMnO4, and add aqueous methanol or isopropanol to it, with a small amount of acid.
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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Upsilon
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Quote: Originally posted by DraconicAcid  | | It takes very concentrated acid to make Mn2O7. Take your KMnO4, and add aqueous methanol or isopropanol to it, with a small amount of acid.
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Yes it would, but if I didn't use concentrated acid, it would make permanganic acid which only exists in aqueous solution. Either way I can't make a
thermite reaction work with liquids.
Anyway, I'm not too big on organic chemistry; could you please explain how that reaction would work?
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DraconicAcid
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No, you can't do the thermite reaction with liquids. What you will do is to react potassium permanganate with a weak reducing agent such as an
alcohol (in aqueous solution) to get a precipitate of MnO2, which you can collect by filtration, dry, and then use for the thermite reaction.
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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Upsilon
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Interesting, could you please post the equation for this reaction?
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halogen
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Don't waste KMnO4. Use batteries. Can't find batteries, ship anyone of us here permanganate and we'd send you the batteries.
Thermite can be messy if you want a "sample" and MnO2 is known for violence andrapidity. Coarse particles slow that down, and how much simpler than Al
and MnO2. But... you might be able to prepare a Mn salt eutectic and electralyze it. That too sounds fun, road less traveled. Doing it in water is
possible but apparently difficult.
[Edited on 8-11-2013 by halogen]
[Edited on 8-11-2013 by halogen]
[Edited on 8-11-2013 by halogen]
[Edited on 8-11-2013 by halogen]
[Edited on 8-11-2013 by halogen]
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Upsilon
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KMnO4 really isn't that expensive, from the right person you can get 500g for around $20. I'd rather buy it then have to deal with tearing apart
batteries. If you don't recommend a thermite reaction, I could at least use a better method than electrolysis. It would take far too long. If I add
HCl to the MnO2, it would make water-soluble MnCl2 and water, correct? Using a solution of this MnCl2, I could throw in some aluminum foil and it will
make manganese metal over time.
[Edited on 8-11-2013 by Upsilon]
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halogen
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Couldn't hurt. Technically Mn reacts a little bit with water, check the activity series, so i don't know but sure, I'd give it a whack.
Well what how much you trying to make, butchering one/two battery no big deal, plus carbon electrodes if you ever need 'em and a bit of zinc.
Thermite IS recommended, just if you're looking for a big erudite block you know? You get a messy lump, and scattered bits. As is common knowledge
larger particles slower burn, plus, a lower oxide like Mn2O3 would be easier to control.
You'll do fine. Just don't get lots in you - manganism.
[Edited on 8-11-2013 by halogen]
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MrHomeScientist
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| Quote: | | If I add HCl to the MnO2, it would make water-soluble MnCl2 and water, correct? Using a solution of this MnCl2, I could throw in some aluminum foil
and it will make manganese metal over time. |
I don't think this will work. I've never heard of manganese being produced with aqueous chemistry; I think it will react with the water as it is
produced.
If you do try it, remember that dissolving MnO2 in HCl goes according to the equation:
MnO2 + 4HCl --> MnCl2 + 2H2O + Cl2
And note that this produces chlorine gas, so do this outside or in a fume hood. This reaction tends to be a lot harder than you'd think, at least when
dealing with commercial grade dioxide. Freshly precipitate MnO2 should react much better.
The 'usual' way to make manganese metal is via thermite, but it is actually extremely hard to get decent quality metal from this. The boiling point of
Mn is less than the melting point of aluminum oxide, so much of your product metal will boil away in the thermite reaction. Any you recover will be in
small nodules and likely contaminated with alumina. I've done this myself - the recommendations to use a lower oxide and coarse particle size are good
ones. Blogfast25 is the resident thermite expert, and I know he's posted here a number of times about the troubles with Mn thermites, so a search
along those lines would be valuable to you. Try google with the term "site:sciencemadness.org" if the local search engine fails you.
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elementcollector1
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I've managed to electrodeposit manganese on copper in an aqueous solution of MnCl2. If anyone's interested, it was a concentrated solution,
a blank PCB, a lead-tin anode straight out of a 60/40 solder tube, and 14V/2A power supply (car battery charger). Got a thin layer of bright silvery
metal that had many dark spots, and required brushing frequently. This eventually darkened slightly. The only problem is removing copper from the
equation to get a manganese foil - perhaps this might show some promising results? It doesn't affect zinc, which is more reactive than copper, and manganese is in the same boat.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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DraconicAcid
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| Quote: Originally posted by MrHomeScientist |
The 'usual' way to make manganese metal is via thermite, but it is actually extremely hard to get decent quality metal from this. The boiling point of
Mn is less than the melting point of aluminum oxide, so much of your product metal will boil away in the thermite reaction. Any you recover will be in
small nodules and likely contaminated with alumina. |
Manganese that is near boiling, exposed to air, will also oxidize.
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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Pyro
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reflux toluene with the KMnO4 to get MnO2, that way you can also make benzoic acid.
then make a thermite, but cover it immediately with something to prevent oxidation of the Mn
all above information is intellectual property of Pyro.  |
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Upsilon
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Would a CO2 or N2 filled environment prevent any oxidation of the Mn? Magnesium will still be able to burn in either to start the thermite reaction,
and they are less reactive than oxygen.
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Pyro
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You could initiate the reaction with KMnO4+glycerol, it's like a delayed fuse. as soon as you drip on the glycerol cover it.
you should see the youtube video on isolating B, he uses thermite.
all above information is intellectual property of Pyro. |
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Upsilon
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I was shopping around and found that it's far cheaper just to buy MnO2 on its own.
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Zephyr
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Upsilon is right, its far cheaper to just buy Manganese dioxide on its, Seattle pottery supply sells a pound for $4.
http://www.seattlepotterysupply.com/Merchant2/merchant.mvc?S...
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blogfast25
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A few points.
Apparently most Mn metal is now produced by electrodepositing, probably similarly to what EC1 did.
Pottery grades of MnO2 are almost always contaminated with Fe2O3. The separation of Fe and Mn is tedious. For quite pure MnO2 start from a reagent
grade Mn(II) salt, dissolve it, precipitate as Mn(OH)2 with NH3 solution and oxidise the hydroxide carefully with hypochlorite (commercial thin
bleach). Filter, wash and semi-calcine the resulting MnO2.
MnO2 thermites will not work very well unless you know what you’re doing. The main problem is that the boiling point of Mn (2061 C) and the melting
point of alumina (2071 C) are very close together (the end-temperature of the post-reaction mix has to be considerably higher than the latter value).
This results in much of the Mn being boiled off during reaction and in some cases to ‘empty crucible syndrome’. At a very minimum include fluorite
or lime in your formulation to cool things down a bit. Also, don't waste money on the finest Al powder money can buy: coarser is better here because
the mixture will burn slower and cooler.
It would also be better to start from the lower oxide Mn2O3 because less heat is generated during the reduction but Mn2O3 is harder to get.
Using permanganate as a source of Mn (as the title of the thread suggests) is really like putting the horse before the cart, considering just how hard
producing commercial permanganate actually is (see sticky thread in general chemistry).
[Edited on 28-12-2013 by blogfast25]
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eidolonicaurum
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There is a very easy way of doing this, I have done it myself. If you heat potassium permanganate to about 300C, it might be hotter, it decomposes
into potassium oxide, manganese dioxide and oxygen. The oxygen is very pure, it is difficult to get oxygen that is much purer, and by dumping the
solid products left over into water and filtering, good quality manganese dioxide is obtained. This you can then do a thermite with.
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bfesser
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Thread Pruned
3-1-2014 at 05:49 |
Oscilllator
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One other way you can do this is with a solution of hydrolysed sugar (sucrose -> glucose + fructose). Just make a concentrated solution of sugar,
add a few ml of concentrated HCl and leave it somewhere hot for a while. My sugar solution has turned yellow over time. This is highly effective and
cheap, since sugar is $1/kg for me at the supermarket.
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