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The Volatile Chemist
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Manganese (II) Chloride Sep. from impurities.
I produced some MnCl2 from Manganese (IV) Oxide, and realized there's a large amount of Iron impurities in the solution. I don't have
enough to do a lot of experimentation with, so I don't want to waste it. Right now the solution's yellow-green-ish, and acidified from excess HCl. My
separation process theories are so far:
1. Let the iron oxidize to Fe2O3 from the air
2. Turn both to acetates, and let the iron oxidize to Fe2O3 from the air.
3. add H2O2 to selectively precipitate the iron as oxide, and not the manganese.
4. Slowly add Na2CO3 and selectively precipitate out one salt as carbonate, as Wikipedia states this is how it was done
industrially.
Thanks,
Nathan
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DraconicAcid
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Quote: Originally posted by The Volatile Chemist  | I produced some MnCl2 from Manganese (IV) Oxide, and realized there's a large amount of Iron impurities in the solution. I don't have
enough to do a lot of experimentation with, so I don't want to waste it. Right now the solution's yellow-green-ish, and acidified from excess HCl. My
separation process theories are so far:
1. Let the iron oxidize to Fe2O3 from the air
2. Turn both to acetates, and let the iron oxidize to Fe2O3 from the air.
3. add H2O2 to selectively precipitate the iron as oxide, and not the manganese.
4. Slowly add Na2CO3 and selectively precipitate out one salt as carbonate, as Wikipedia states this is how it was done
industrially.
Thanks,
Nathan |
I'd go with the carbonate. You don't want it to get too basic, or you'll be reforming MnO2 or Mn(OH)3 on exposure to air.
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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The Volatile Chemist
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Thanks. Which should precipitate first? I'm pretty sure the iron will, but not positive. If I let the solution sit for a while after precipitation,
will the Iron oxidize (and the manganese not)?
Thanks!
I found this thread on MnCl2. After I finish collecting a manganese salt, I'd like to write a series of articles on manganese in the
SMWiki, hope I get to it first 
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DraconicAcid
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The iron will precipitate first. MnCO3 is surprisingly soluble. (Ksp is 3.5e-8; that of FeCO3 is 3.2e-11)
You can also separate them by oxidizing the manganese to MnO2 in acidic solution, so that it precipitates while the iron stays in solution.
(We do this for qualitative analysis at my college- the solution is made very acidic with conc. nitric acid, and potassium chlorate is added to the
hot solution *very* carefully.)
It's usually better to precipitate the stuff you want, and leave the unwanted stuff in solution, since the unwanted stuff is never *completely*
insoluble.
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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I have a pile of MnO2 salvaged from batteries just sitting there with all of the impurities it contains. I haven't used it since I also have a kilo
of pure MnO2. I might just convert it to MnCl2 and follow your lead here.
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The Volatile Chemist
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| Quote: Originally posted by DraconicAcid | The iron will precipitate first. MnCO3 is surprisingly soluble. (Ksp is 3.5e-8; that of FeCO3 is 3.2e-11)
You can also separate them by oxidizing the manganese to MnO2 in acidic solution, so that it precipitates while the iron stays in solution.
(We do this for qualitative analysis at my college- the solution is made very acidic with conc. nitric acid, and potassium chlorate is added to the
hot solution *very* carefully.)
It's usually better to precipitate the stuff you want, and leave the unwanted stuff in solution, since the unwanted stuff is never *completely*
insoluble. |
Thanks! I really don't want to have to precipitate the Manganese as MnO2 again, just to have to go through the same process of HCl and
MnO2. I'll probably precipitate the iron, then precipitate the manganese after filtering. Thanks for your help!
| Quote: Originally posted by j_sum1 | I have a pile of MnO2 salvaged from batteries just sitting there with all of the impurities it contains. I haven't used it since I also have a kilo
of pure MnO2. I might just convert it to MnCl2 and follow your lead here.
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Yes, it's a slightly dirty reaction, but if I get 50% yeild, I'll be happy. My main goal is not to waste more reagents on it. Just the
Na2CO3.
You should do it, though! I'd like to see your results!
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Texium
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I've tried it before. The results were depressingly bad. I used dilute hydroxide for selective precipitation. It brought about more headaches and
dirty glassware than anything else.
There is an old and relevant thread here though, that you might want to check out.
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DraconicAcid
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| Quote: Originally posted by The Volatile Chemist | | Quote: Originally posted by DraconicAcid |
It's usually better to precipitate the stuff you want, and leave the unwanted stuff in solution, since the unwanted stuff is never *completely*
insoluble. |
Thanks! I really don't want to have to precipitate the Manganese as MnO2 again, just to have to go through the same process of HCl and
MnO2. I'll probably precipitate the iron, then precipitate the manganese after filtering. Thanks for your help!
|
Freshly precipitated MnO2 will be much easier to dissolve/reduce than the original stuff you started with.
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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The Volatile Chemist
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| Quote: Originally posted by DraconicAcid | | Quote: Originally posted by The Volatile Chemist | | Quote: Originally posted by DraconicAcid |
It's usually better to precipitate the stuff you want, and leave the unwanted stuff in solution, since the unwanted stuff is never *completely*
insoluble. |
Thanks! I really don't want to have to precipitate the Manganese as MnO2 again, just to have to go through the same process of HCl and
MnO2. I'll probably precipitate the iron, then precipitate the manganese after filtering. Thanks for your help!
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Freshly precipitated MnO2 will be much easier to dissolve/reduce than the original stuff you started with. |
OK, thanks. I'll post when I've decided and done one of these procedures.
| Quote: Originally posted by zts16 | I've tried it before. The results were depressingly bad. I used dilute hydroxide for selective precipitation. It brought about more headaches and
dirty glassware than anything else.
There is an old and relevant thread here though, that you might want to check out. |
Yeah, it's definitely not one of my favorites. I saw that link (Posted it above).
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unionised
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The traditional approach is to dissolve the MnCl2 in 20% HCl then extract Fe(III) with ether.
Of course, that's only useful if you have access to ether.
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blogfast25
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Nurdrage on eliminating iron(III) from MnSO4 solution by means of selective precipitation:
https://www.youtube.com/watch?v=BLJgBSrhZI8
I've used variants (chloride instead of sulphate) of this several times with great success.
Other than ether, acetone is a also a good solvent for FeCl3 but would have to applied several times on very finely ground contaminated MnCl2 hydrate.
[Edited on 10-11-2014 by blogfast25]
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elementcollector1
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| Quote: Originally posted by blogfast25 |
Other than ether, acetone is a also a good solvent for FeCl3 but would have to applied several times on very finely ground contaminated MnCl2 hydrate.
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What if you refluxed the mixture so that the acetone reached a higher temperature?
(Unless it starts dissolving MnCl2 then as well...)
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unionised
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It would be interesting to recrystallise the mixed salts from a mixture of acetone and water.
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blogfast25
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@EC1:
Should work, as long as MnCl2 solubility remains low, as you wrote.
@unionised:
Or use the acetone as an anti-solvent for MnCl2 hydrate, thus precipitating it, leaving the FeCl3 in the acetone/water phase?
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The Volatile Chemist
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Interesting ideas! I'll post when I've done one, rather busy at the moment.
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AJKOER
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Here is a possibly interesting find. Per this site http://www.public.asu.edu/~jpbirk/qual/qualanal/manganese.ht... , to quote:
"Addition of aqueous ammonia precipitates white Mn(OH)2:
Mn2+(aq) + 2NH3(aq) + 2H2O(l) <==> Mn(OH)2(s) + 2NH4+(aq)
The precipitate does not dissolve in excess ammonia, but does dissolve in solutions containing ammonium salts. The precipitate is easily oxidized by
atmospheric oxygen to form Mn(III) or Mn(IV), which turns the precipitate a brownish color."
However, a similar possible solubility of a Fe(OH)2 impurity, in say (NH4)2SO4 (which an be readily prepared from the action of aqueous ammonia on
household Epsom salts, MgSO4, and filtering out the Mg(OH)2 precipitate), no such mention (see http://www.public.asu.edu/~jpbirk/qual/qualanal/iron.html ).
Wikipedia on Fe(OH)2 simply states on its preparation, to quote:
"Iron(II) hydroxide is poorly soluble (1.43 × 10−3 g/L). It precipitates from the reaction of iron(II) sulfate and hydroxide ions (from a soluble
compound containing hydroxide ion):
FeSO4 + 2 OH− → Fe(OH)2 + SO4 2− "
with no mention of ammonium salt solubility.
When I have to purify my battery's MnO2, I may investigate this less smelly and inexpensive route than employing acetone.
[Edit] I created some Fe(OH)2 from green aqueous FeCl2 plus household ammonia in excess. I then added some solid NH4Cl and stirred. No evident
dissolution of the Fe(OH)2. Perhaps on evaporation of any unreacted NH3 and a lowering of the pH, it may still occur. So far, dissolution of
MnCl2/FeCl2 in NH3/NH4Cl may work to separate the manganese ion from the iron.
[Edited on 11-11-2014 by AJKOER]
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blogfast25
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| Quote: Originally posted by AJKOER | The precipitate does not dissolve in excess ammonia, but does dissolve in solutions containing ammonium salts.
When I have to purify my battery's MnO2, I may investigate this less smelly and inexpensive route than employing acetone.
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If true, it has that in common with Mg(OH)2. And it's possible, the Ks of Mn(OH)2 is not very low: 1.9 x 10<sup>-13</sup> (http://bilbo.chm.uri.edu/CHM112/tables/KspTable.htm)
That dissolution would be caused by the modest acidity of ammonium salts (of strong acids), ammonium being the conjugated acid of a weak base, NH3.
Thus:
Mn(OH)2(s) + 2 NH4(+)(aq) < === > Mn2+(aq) + 2 NH3(aq) + 2 H2O(l)
You wouldn't have to worry about ferrous iron because if the Mn and Fe both come from some source of MnO2, the iron would ALWAYS be ferric because
MnO2 is such a powerful oxidiser. And Fe(OH)3 is of course extremely insoluble (Ks = 4 x 10<sup>-38</sup> !)
But, but, but: there is a snake in the grass. Look at the equilibrium equation: to push it fully to the right (and fully dissolve the Mn(OH)2) you'd
need 2 mol ammonium salt per mol of Mn(OH)2 and release 2 mol of ammonia into the bargain (so I hope you prefer the stench of ammonia to the citrussy
odour of acetone!  )
[Edited on 11-11-2014 by blogfast25]
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elementcollector1
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Sounds like a good method of concentrating ammonia solution as a useful byproduct...
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blogfast25
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Not really: if you want to do that just use a cheap, strong alkali like NaOH.
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j_sum1
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| Quote: Originally posted by The Volatile Chemist | | Quote: Originally posted by DraconicAcid | | Quote: Originally posted by j_sum1 | I have a pile of MnO2 salvaged from batteries just sitting there with all of the impurities it contains. I haven't used it since I also have a kilo
of pure MnO2. I might just convert it to MnCl2 and follow your lead here.
|
Yes, it's a slightly dirty reaction, but if I get 50% yeild, I'll be happy. My main goal is not to waste more reagents on it. Just the
Na2CO3.
You should do it, though! I'd like to see your results! |
You aren't kidding when you said that it's dirty. Well, everything with battery innards is dirty but add the fact that it is foaming with Cl2 and
overflowing the reaction vessel and you do have a bit of a mess.
Nurdrage uses NaOH to purify Fe from MnSO4 in this video https://www.youtube.com/watch?v=BLJgBSrhZI8.
I assume the same technique will work equally well with chlorides and sulfates.
What are the relative merits of hydroxides versus carbonates? |
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blogfast25
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| Quote: Originally posted by j_sum1 |
1. I assume the same technique will work equally well with chlorides and sulfates.
2. What are the relative merits of hydroxides versus carbonates? |
'Yes' to 1.
2. With carbonates you'll precipitate Fe(OH)3 (there is no ferric carbonate) and MnCO3. Since as the latter is also much more soluble than the former,
it probably makes little difference. But I would always recommend ammonia as the source of alkali, because ammonium salts appear easier to wash out.
Traces can also be removed by calcining your end product.
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The Volatile Chemist
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Forget what I said about writing the articles. I literally got NO MnCO3. I don't know what happened. The reaction was dirty enough that I
don't want to do it again.
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j_sum1
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| Quote: Originally posted by The Volatile Chemist | | Forget what I said about writing the articles. I literally got NO MnCO3. I don't know what happened. The reaction was dirty enough that I
don't want to do it again. |
If MnCO3 is "surprisingly soluble" might it have happened that you precipitated only the ferric compounds? Ie, manganese carbonate produced but not
isolated?
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The Volatile Chemist
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It's possible. Regardless, it's down the drain now. I really don't think MnCO3 could be THAT soluble, that I could not precipitate it.
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blogfast25
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You're misinterpreting solubility in this context. It's perfectly possible to precipitate MnCO3, I've done it several times.
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.
[Edited on 16-11-2014 by blogfast25]
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