Sciencemadness Discussion Board - Manganese (II) Chloride Sep. from impurities.

Manganese (II) Chloride Sep. from impurities.

The Volatile Chemist - 9-11-2014 at 14:29

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

DraconicAcid - 9-11-2014 at 14:32

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.

The Volatile Chemist - 9-11-2014 at 14:37

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

DraconicAcid - 9-11-2014 at 14:46

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.

j_sum1 - 9-11-2014 at 14:46

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.

The Volatile Chemist - 9-11-2014 at 14:53

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.

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!

Texium - 9-11-2014 at 14:59

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.

DraconicAcid - 9-11-2014 at 15:01

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.

Freshly precipitated MnO2 will be much easier to dissolve/reduce than the original stuff you started with.

The Volatile Chemist - 9-11-2014 at 15:19

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.

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).

unionised - 10-11-2014 at 01:00

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.

blogfast25 - 10-11-2014 at 05:58

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]

elementcollector1 - 10-11-2014 at 06:54

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.

What if you refluxed the mixture so that the acetone reached a higher temperature?
(Unless it starts dissolving MnCl₂ then as well...)

unionised - 10-11-2014 at 07:35

It would be interesting to recrystallise the mixed salts from a mixture of acetone and water.

blogfast25 - 10-11-2014 at 09:21

@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?

The Volatile Chemist - 10-11-2014 at 12:34

Interesting ideas! I'll post when I've done one, rather busy at the moment.

AJKOER - 11-11-2014 at 04:58

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]

blogfast25 - 11-11-2014 at 09:46

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.

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 10-13 (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 10-38 !)

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]

elementcollector1 - 11-11-2014 at 12:37

Sounds like a good method of concentrating ammonia solution as a useful byproduct...

blogfast25 - 11-11-2014 at 14:06

Quote: Originally posted by elementcollector1

Sounds like a good method of concentrating ammonia solution as a useful byproduct...

Not really: if you want to do that just use a cheap, strong alkali like NaOH.

j_sum1 - 16-11-2014 at 00:11

Quote: Originally posted by The Volatile Chemist

Quote: Originally posted by DraconicAcid

Quote: Originally posted by j_sum1

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?

blogfast25 - 16-11-2014 at 04:37

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.

The Volatile Chemist - 16-11-2014 at 11:29

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.

j_sum1 - 16-11-2014 at 12:38

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?

The Volatile Chemist - 16-11-2014 at 12:46

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.

blogfast25 - 16-11-2014 at 13:11

Quote: Originally posted by The Volatile Chemist

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.

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 10-11

and Fe(OH)3 (Ks = 4 x 10-34, 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 10-13 (than Mn(OH)2) but still about 1021 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]

j_sum1 - 16-11-2014 at 15:53

I think I will use NaOH to remove the Fe. I like the idea of ammonia but my source of ammonia for my home lab is floor cleaner and so it doesn't make practical sense.

As an aside, I think this might make a good project for my students next year. They are required to do an EEI (extended experimental investigation). They get to choose but I get to make suggestions. Start with a lantern battery and isolate as many elements as possible. Lots of opportunities to learn some good techniques including filtration, washing, drying, selective precipitation, crystallisation, even pyrolysis of the case. I'm looking at one in front of me. It has:
pvc case
zinc cups
copper wires and terminals
graphite electrodes with copper wire in the centre
lead/zinc solder
mixture of MnO2/graphite powder wet with a solution of ZnCl2 (and perhaps some NH4Cl2)
cardboard/paper

There is a lot that students could investigate here and plenty to learn. Maybe even calcining the cardboard and casting the zinc.

blogfast25 - 17-11-2014 at 05:38

j_sum1:

Surely if you're a teacher you have access to clean NH3? Having said that, of course it works with NaOH too.

The only problem with your project is that battery crud is real messy: the graphite conductor that's part of the dry electrolyte stains everything! Some parents might not be so pleased if their beautiful broods come back with black fingernails and stained clothes. Make sure you do a dry run, just you, to evaluate that.

j_sum1 - 17-11-2014 at 12:57

I don't mix work and play. School chemicals and equipment stay at school.

[edited to add]
It seems to me to be not sensible to blur the lines between professional and private too much. I don't order chemicals or equipment through the school accounts. I don't do personal projects at school. I don't want my workplace in any way held accountable for what I do at home -- not good from a legal and risk assessment perspective.

Knowledge however is transferable. So if I am learning practical details about separation of Fe from Mn then my students may as well benefit from that. As for the mess from battery gunk: that's what lab-coats are for.

(I have been fortunate to have gotten hold of some chemicals that would otherwise have been dumped. And I have been able to acquire some broken glassware and re-purpose it. It is stupid to waste stuff that I can use and would probably buy anyway. It is also stupid to pay for disposal if I can do it.)

[Edited on 17-11-2014 by j_sum1]

DraconicAcid - 17-11-2014 at 17:23

Ammonium manganese phosphate is fairly insoluble; I don't know if iron phosphate is as insoluble...

(ETA: No- all metals other than alkali metals will interfere with this precipitation.)

[Edited on 18-11-2014 by DraconicAcid]