Sciencemadness Discussion Board

Removal of Iron from Manganese

Chemistry Alchemist - 25-8-2011 at 22:42

Ive just done a thermite reaction of MnO2 that also has Iron Oxide contaminated in, how do i clean away the Iron Metal from the Manganese metal with out it destroying the Manganese?

unionised - 26-8-2011 at 03:27

First you get a time machine, then you go back to before you did the reaction and you dissolve the Mn and Fe oxides in acid. Then you recrystallise the Mn salt to remove the Fe. Then you dissolve the pure Mn salt in water and re precipitate the Mn as the carbonate or hydroxide. You calcine that o get the pure oxide and then you repeat the thermite reaction.
Hypothetically you could vacuum distil the metals but that's not very practical either.

Chemistry Alchemist - 26-8-2011 at 04:01

Would you be able to write it out as a equation if possible... ive tried to crystallize the chloride but i was unsure on what one crystallized first...

Megamarko94 - 26-8-2011 at 05:29

Quote: Originally posted by Chemistry Alchemist  
ive tried to crystallize the chloride but i was unsure on what one crystallized first...

if there is no much iron... than its manganesse chloride because there is more of it also if it has slight pink color than its manganesse chloride...

Chemistry Alchemist - 26-8-2011 at 05:34

The solution was still really brown from the iron 3 chloride, but there were crystals at the bottom... Would they of been manganese chloride

blogfast25 - 26-8-2011 at 05:36

Quote: Originally posted by Chemistry Alchemist  
Ive just done a thermite reaction of MnO2 that also has Iron Oxide contaminated in, how do i clean away the Iron Metal from the Manganese metal with out it destroying the Manganese?


Hmmm... I thought you didn't have any aluminium? Also obtaining manganese metal from thermite is quite hard, I know this from personal experience. Show us your metal.

The iron needs to be removed prior to the reduction reaction. Small amounts of Fe can be removed by recrystallising the MnCl2 or MnSO4, larger quantities require more drastic measures like Nurdrage's method.

Also: try not to start too many separate but closely regards threads. This matter is closely related to your other manganese related thread...

[Edited on 26-8-2011 by blogfast25]

Chemistry Alchemist - 26-8-2011 at 05:58

I'll post a picture n it either tomorrow or the next day... Wen ever I get around to taking pictures of it...

not_important - 26-8-2011 at 06:20

Once you have the metal, electro-refining is the best route. A bit tricky, requires careful control of conditions, and is also the best way to get high purity Mn metal.

When starting from Mn compounds you adjust the pH and add an oxidiser to convert Fe(II) to (FeIII) which precipitates out as the hydrated oxide at the selected pH. If starting from MnO2 use the method I've posted here before - add enough con H2SO4 to make a thin paste of the MnO2, _slowly_ heat in a ceramic container until you drive off excess H2SO4, then continue heating up to red heat - 700 C. After allowing the reaction mix to cool, extract with warm water to get MnSO4. Iron compounds will have been converted to the oxides, ditto for several other metals that might be in there. If you want something other than MnSO4, add aqueous NH3 to ppt out the hydroxide, if thermite is the target then bubble air through the solution during the addition of NH3 and for some time afterwards, the Mn(OH)2 converts to Mn(III) hydrated oxide; wash, dry, roast in air.



blogfast25 - 26-8-2011 at 06:35

Quote: Originally posted by not_important  


[snip]

, if thermite is the target then bubble air through the solution during the addition of NH3 and for some time afterwards, the Mn(OH)2 converts to Mn(III) hydrated oxide; wash, dry, roast in air.



No need for the air, in my experience, bleach (hypochlorite) works instantly and completely. Filter, wash cake profusely with clean water, dry and lightly calcine the MnO2.

But MnO2 thermites aren't easy at all. The main problem is that the BP of manganese metal is very close to the MP of alumina. And since as to get decent separation of metal and slag you need to get above the MP of alumina, inevitably quite a bit of metal is simply evaporated. I've rarely obtained yields of more than 30 %.

You also need CaF2 (or Cryolite) as a slag fluidiser/coolant (heat sink) to get decent metal. W/o a flux the MnO2/Al mixture tends to behave like a flash powder. I'll see if I can post a photo of my manganese later on.

not_important - 26-8-2011 at 06:57

If you do use NaOCl, then don't use NH3 8-)

I prefer using ammonium salts because A) they tend not to bind to the ppts as strongly, B) they're easier to remove, C) for Mn they will complex other d-block metals without grabbing the Mn. But for thermite use some sodium contamination isn't going to be harmful, it'll boil off and/or function as a flux.

blogfast25 - 26-8-2011 at 07:39

Quote: Originally posted by not_important  


[snip]

C) for Mn they will complex other d-block metals without grabbing the Mn. But for thermite use some sodium contamination isn't going to be harmful, it'll boil off and/or function as a flux.



You mean complexing Zn2+ and Cu2+ as ammonia complexes? Are there any others?

I just find strong, clean reasonably priced NH3 solution not so easy to find. The weaker ones are a rip off.



Photo: Manganese reguli from MnO2 aluminothermic reduction (compared to a 1 p coin). In the irregular shape of these nuggets lay clues: the melt never got to settle well because of evaporating Mn.


[Edited on 26-8-2011 by blogfast25]

not_important - 26-8-2011 at 08:29

When I was doing this I had access to cheap slaked lime and lye with some NaCl content; both did fine for recovering & recycling the NH3 content of the salts. Plus the pottery place I was buying my MnO2 at was sloppy and got dust from one material spread about and into others. As most of their customers seemed to use MnO2 for heavy manganese with or without copper coloured glazes, a little iron, copper, or whatever, wasn't a problem for those folks.

Cu and Zn are p-block, but yes - along with Cd they form fairly strong NH3 complexes. But a number of other transition metals, as well as Ca and Mg, form weaker complexes that increase the solubility of otherwise low solubility compounds of them. I've posted the effect of ammonium slats on the solubility of CaSO4 several times, pointing out that the Ca(NO3)2 + (NH4)2SO4 route to AN needs a bit more work to get a clean product.


All of the thermal routes to metallic Mn suck, and the electrolytic means aren't as simple as those for some other metals.

blogfast25 - 26-8-2011 at 09:48

Quite a reputable book on pyrometallurgy (I forget which one right now) claims that metallic manganese (rarely needed, truth be told) is made by aluminothermic reduction of Mn3O4, in open crucibles. Apparently once reaction gets going the crucible is then continuously topped with more oxide/Al mixture and the crusible then gradually fills with liquid Mn. Hmmm...

An early lab method consisted in adding Mg powder (slowly!) to a molten KCl/anh. MnCl2 mixture. After reaction is complete, the crucible is then heated to above the MP of Mn for some time, then cooled to RT. Chemically pure Mn resulted, or so it was claimed...

My pottery MnO2 contained about 25 % acid insolubles. Some silicates, I think...

[Edited on 26-8-2011 by blogfast25]

Arthur Dent - 26-8-2011 at 12:35

I had some Manganese Carbonate from the pottery store I bought a while ago and decided to test it out.

I remember Peach mentioning that his Manganese Carbonate purchased on eBay was full of crap, and very impure. So I wanted to check mine if it was pure. First off it's the right color, a light tan, very fine powder with no signs of other impurities.

So I carefully and slowly added some concentrated hydrochloric acid to it, expecting to obtain a pink solution. Ugh! :mad:

The resulting mixture bubbled a bit, resulting in a dark coffee-colored solution, more reminiscent of ferric chloride! Wow! Even the crud from carbon zinc batteries has a lower iron content than that! I can't believe that this stuff contains so much iron! Or could it be something else? I really don't see what, though.

I succesfully removed Ferric Chloride impurities from Manganese Chloride months ago, using the chemical process of Ferric Hydroxide precipitation which resulted in a light pink solution, which yielded beautiful pink crystals unpon a long crystallization process. But this stuff I made this morning is so dark that I don't know if the Ferric Hydroxide trick will work with this crap. And the HCl is technical grade so it's not the culprit.

I know I shouldn't expect reagent-grade purity from pottery store stuff, but come on! that's just plain wrong!

Robert


[Edited on 26-8-2011 by Arthur Dent]

blogfast25 - 26-8-2011 at 13:07

Quote: Originally posted by Arthur Dent  
I know I shouldn't expect reagent-grade purity from pottery store stuff, but come on! that's just plain wrong!

Robert
[Edited on 26-8-2011 by Arthur Dent]


Precisely because they don't specify any degree of purity they can get away with this legally. And because these impurities probably don't affect performance in glazes the sellers will rarely face a justified complaint.

There is some high grade MnO2 available too on eBay, it's not all 'pottery crap'. I bought some quality MnO2 there and it was really quite pure.

Even quite large amounts of iron should be removable with the method you used though...

unionised - 27-8-2011 at 05:09

If you expose Mn++ to the air under alkaline conditions it oxidises so the oxides and carbonates are likley to contain Mn(III) and Mn(IV).
If you dissolve those in HCl you get mixed oxidation state complexes of Mn. Those complexes are dark green/ brown.
Their other characteristic is that they smell of chlorine because they are not stable.
Does the brown stuff you get smell of chlorine?
If it does then leaving it a while or heating it will help decolourise it.

If it doesn't then it's probably got Fe in it. You can extract that from a solution in HCl with ether

blogfast25 - 27-8-2011 at 06:49

Quote: Originally posted by unionised  
If you dissolve those in HCl you get mixed oxidation state complexes of Mn. Those complexes are dark green/ brown.
Their other characteristic is that they smell of chlorine because they are not stable.
Does the brown stuff you get smell of chlorine?
If it does then leaving it a while or heating it will help decolourise it.

If it doesn't then it's probably got Fe in it. You can extract that from a solution in HCl with ether


My experience is that Mn3+ in the presence of chloride breaks down incredibly quickly:

Mn3+ + Cl- === > Mn2+ + 1/2 Cl2

In fact, from what I observed, it appears that Mn3+ in conjuction with Cl- forms a very dark coloured, semi-covalent compound, MnCl3, which is even more unstable in water than MnCl4. The decomposition of MnCl3 is accompanied by great evolution of heat and proceeds very quickly.

Cheaper would be extraction of FeCl3 with acetone, in which it is greatly soluble. But I'm not sure about MnCl2's solubility in acetone.

[Edited on 27-8-2011 by blogfast25]

m1tanker78 - 27-8-2011 at 09:36

Quote: Originally posted by Arthur Dent  
I can't believe that this stuff contains so much iron! Or could it be something else?


Robert, even a tiny bit of iron impurity will yield an apparent impressive quantity of fluff when precipitated as hydroxide. I guess one should be thankful that iron is the main impurity? Easy to clean up in this case. ;)

Tank

Bezaleel - 28-8-2011 at 16:16

Quote: Originally posted by unionised  
(...)

If it doesn't then it's probably got Fe in it. You can extract that from a solution in HCl with ether


You say that FeCl3 will be extracted by the ether, while MnCl2 will remain in solution? Sounds like a handy extraction method.
Quote: Originally posted by blogfast25  

(...)

Cheaper would be extraction of FeCl3 with acetone, in which it is greatly soluble. But I'm not sure about MnCl2's solubility in acetone.

My experience is that MnCl2 also has an appreciable solubility in acetone. However, if you crystallise the MnCl2 contaminated with FeCl3, the FeCl3 will crystallise in the latter part of the crystallisation proces, meaning that it will reside on the outside of the MnCl2 crystals. You can quickly rince with acetone to remove the FeCl3. Since the soubility of FeCl3 in acetone is high, only use a small amount of acetone.


Another method of removing the iron from the manganese is by simply boiling off the water. Thereby the FeCl3 will largely decompose and, the remaining Fe2O3 can be filtered off after redissolving the product. This will remove about 90% of the Fe, is my experience.

Besides, a small amount of an emerald green substance is formed as a very finely divided by product, which I believe to be MnO. I believe this despite its unusually high stability.

not_important - 29-8-2011 at 06:13

Ah, but with ether you can extract from aqueous solution, while with acetone you need to work with the solids, which generally trap some of the material to be extracted.

AJKOER - 30-8-2011 at 08:23

Many have mentioned a path using sulfates, however, not mentioned is the enormous difference in solubility of the corresponding sulfate salts.

In fact, Ferric sulfate is 8 times more soluble at 20 C than the MnSO4.

Use this to your advantage in separating the Fe from the Mn.

blogfast25 - 30-8-2011 at 09:40

Quote: Originally posted by not_important  
Ah, but with ether you can extract from aqueous solution, while with acetone you need to work with the solids, which generally trap some of the material to be extracted.



Good point.

AJKOER:

The methods described above were presented precisely because separating Mn and Fe based on solubilities alone is very difficult.

Chemistry Alchemist - 31-8-2011 at 00:00

and if you were working with the solubility method, wouldn't you need to have a good amount of the stuff to get a good precipitate? ive been working with small amounts, not enough to separate via solubility... i think ill go with the Acetone method or the Crystallize MnCl2 out of solution method :)

Arthur Dent - 25-9-2011 at 12:52

After many-a-months in the dessicator, my solution of MnCl<sub>2</sub> is almost completely crystallized, so I decided to harvest my Manganese Chloride crystals which have grown to rather respectacle sizes, from 1 to 3 cm each!

I opened the dessicator and upon the rush of air, promptly emitted a puff of suffocating white HCl smoke...



The crystals have been deposited on a paper towel to dry a bit, but I'm wondering if these are hygroscopic and will absorb water from the air, Or I could just let them to dry out in the air for a while.

Also, they still smell of chlorine a bit, so I thought to wash them with some solvent. Is there a solvent in which MnCl<sub>2</sub> is insoluble? I have 99% isopropanol, anhydrous methanol and 99% acetone in stock...

Robert

blogfast25 - 26-9-2011 at 04:02

Yummy crystals, Robert! I would just let then dry in air now... MnCl2 isn't hygroscopic at all, in my experience. Chlorine will evaporate (it's couldn't in your dessicator). Everything will be fine: halleluyah! :cool:

Are they really that red or is that camera bias?

[Edited on 26-9-2011 by blogfast25]

Arthur Dent - 26-9-2011 at 04:26

@ blogfast25 : Yes sir this is the accurate color, they really look like hard candy! This morning however, I noticed that the edges of the beautiful crystals above have lost their shine and have started to turn a very light opaque pink, which is the color of anhydrous MnCl<sub>2</sub>, so it reacts a bit like Copper Sulphate crystals, who slowly turns white and anhydrous if left in a dry atmosphere.

It's very dry in my house, so I can imagine the manganese Chloride crystals have lost some of their hydration water. They do smell a bit less stingy.

The crystals have grown from a 150 ml solution down to 10 ml of liquid left when I pulled them out of the dessicator. Took a few months (3 or 4) in the dessicator at low vacuum (0.9 atm).

Robert

Random - 26-9-2011 at 05:03

Quote: Originally posted by blogfast25  
Everything will be fine: halleluyah! :cool:
[Edited on 26-9-2011 by blogfast25]



blogfast25 - 26-9-2011 at 07:46

@ Roger:

Next time perhaps boil in your soluion until boiling behaviour starts to change. On cooling and icing you should get a first crop of decent crystals of MnCl2.4H2O (depending on solubility - temp. function), the rest can be obtained in dessicator.

Watch for a U2U.

Chemistry Alchemist - 26-9-2011 at 07:49

when MnCl2 Crystals form, what shape do grow?

Arthur Dent - 26-9-2011 at 11:29

Manganese Chloride forms crystals with rhombohedral symmetry, much like Cadmium Chloride. A typical crystal looks like a cube skewed at a 45 degree angle.

blogfast25 - 26-9-2011 at 11:31

Quote: Originally posted by Chemistry Alchemist  
when MnCl2 Crystals form, what shape do grow?


The shape on the photos, LOL. Seriously, look up in Wiki about various crystalline structures. There should be pics and diagrams to your heart's content if you drill down a bit...

sxl168 - 16-11-2011 at 11:49

I was removing some small amounts of dissolved Iron today from MnCl2 prepared from alkaline batteries and I noticed this batch precipitated something else along with the Iron. The resulting solids are a moderate grey color and looked like a dirty off white when it first precipitated. My solution is a little acidic yet, so I wouldn't think any Mn had precipitated. Some of it is probably colloidal carbon that made it through the filter paper, but not the quantities observed.

I was wondering what additives are added to alkaline batteries if anyone knows, which this stuff might be composed of. The quantity that I have obtained so far is quite small <100 mg. Do they use Cd or Pb? Those additives would probably precipitate out in the conditions I have used. Could it also be Hg as I have used a few rather old cells (circa late 1980's) in this batch? I was wondering if anyone else has seen this in their precipitation reactions.

blogfast25 - 17-11-2011 at 09:43

Quote: Originally posted by sxl168  
I was removing some small amounts of dissolved Iron today from MnCl2 prepared from alkaline batteries and I noticed this batch precipitated something else along with the Iron.


Clues would lie in how exactly you removed the iron...

Chemistry Alchemist - 17-11-2011 at 10:02

So really the best way to remove the iron would be from the method NurdRage made a video about?

blogfast25 - 17-11-2011 at 10:30

Quote: Originally posted by Chemistry Alchemist  
So really the best way to remove the iron would be from the method NurdRage made a video about?


IMHO, yes.

See for instance, here:

http://www.sciencemadness.org/talk/viewthread.php?tid=17931#...

[Edited on 17-11-2011 by blogfast25]

sxl168 - 17-11-2011 at 11:14

I use HCl as my acid as it's the only acid I can get ahold of here dirt cheap, the following may behave differently with other acids. To remove the Iron, I oxidize it from Ferrous to Ferric with bleach or H2O2, doing this while the leaching solution is still highly acidic. Then I adjust the pH to about 5 with baking soda or Sodium Carbonate. I do not have a pH meter, I use my nose and 5 is just a best guess. The solution will only have a faint acidic smell to it and no alkaline smell. Zn and Mn will not drop out yet if done properly. While doing the pH adjustment, I have the solution heated so that it is at least 50C. Ferric drops out readily when heated (but really friggin' slowly at room temp, if at all). It might take 2 tries at the process if the solution is heavily contaminated with Iron. I get nice faint pink solutions once its done with no hints of Ferric colorization and no additional colorization upon standing if you use enough oxidizer. I've leached whole AAAA 9V cells, casings and all and was able to extract clean Mn/Zn with this method.

The hardest part IMO, is oxidizing the Iron but not the Mn. This is where having the acidic leaching solution works well. If MnO2 is formed while adding oxidizer, which it usually does, it will redissolve with stirring and oxidize the Iron. I usually know the Iron is finished being oxidized if I get a persistent bleach smell from the solution and MnO2 that won't redissolve. Just takes a few practice runs to get it right, IMO.

I saw NurdRage's video, but I don't have chemicals for his method. I thought everyone knew about method listed above as it's listed in many patents and battery recycling papers I have come across. It's a somewhat slow process, but works.

This other stuff that drops out seems to come when adjusting the pH higher to kick out the last remaining bits of Iron, and just before Zn/Mn start precipitating.

blogfast25 - 17-11-2011 at 12:32

Yes, your method is a variant of 'Nurdrage's': it relies on selective precipitation of Fe3+ at pH =<5, but not Mn, Zn and some others.

Perhaps there's another variant in there. Add copious bleach to the battery gunge, this will oxidise all to Fe (III) and Mn (IV). Wash superficially to get rid of the bleach. Dissolve washed gunge in HCl: MnCl2, FeCl3, ZnCl2 etc are leached out. Then treat the leachate like you do. That would eliminate the tricky bit of oxidising Fe (II) but not Mn (II).

Nurdrage's method has one small advantage: once the pH is right and you filter, the filtrate is free of cations like Na+, provided the precipitate that was used to 'buffer' the leachate had been washed properly.

You say you don't have the chemicals but for alkali you could use household ammonia, instead of NaOH. 'Washing soda' (Na2CO3) would also do.

[Edited on 17-11-2011 by blogfast25]

sxl168 - 18-11-2011 at 10:03

My bad, I was thinking of his video where he extracts the Mn from battery paste via burning sulfur. I forgot about his other video for purification. Yes indeed that method is the one that I use as a reference, but I just found it a bit lacking when I was dissolving whole cells with the casing (I only dissolve whole cells as they exist in 9 Volt packs, other size cells I do material separation). That results in a lot of iron contamination and it takes forever for it all to oxidize by exposure to air, which is why I was doing the oxidizer route. As for removing the alkali's, I precipitate everything out as carbonates once the iron has been removed and wash thoroughly. I keep the iron precipitate for use later as Ferric Chloride.

jsc - 18-11-2011 at 10:41

That pic is pure chemical porn. :P

Chemistry Alchemist - 23-11-2011 at 22:53

So its been epically hot recently and the solution which was some what hard to crystallize has now crystallized and i have nice long crystals of my guess Manganese(II) Chloride, so would i just poor off the rest of the solution and then wash with really cold water (~0 degrees C) or would the Manganese dissolve in the water even at that temperature? is there a solvent that can wash the remaining iron out and leave the manganese untouched? did you want me to post pics of the crystals?

blogfast25 - 24-11-2011 at 05:55

Wash with small amounts of iced water or with clean alcohol or acetone. Acetone would dissolve FeCl3 but only if the material is finely crushed. Pic(s) please!

Chemistry Alchemist - 24-11-2011 at 06:48

i tried washing with a few crystals but it kinda looks like the Iron(III) Chloride has fused in with the crystals... and the crystals were dissolving too fast to clean... the crystals are needle like as people said they should be if they are Manganese(II) Chloride

DSC01652.JPG - 108kB

They are a bit damp still, but you can still see some of the needle crystal structure off the the left side of it, the colour looks like it has a lot of iron in it

blogfast25 - 24-11-2011 at 06:57

Actually they 'look' fairly pure to me. Did you crystallise them from a highly acidic solution or from a fairly neutral one? FeCl3 doesn't survive a pH close to 7, or even 5 to 6: it hydrolyses and precipitates. If FeCl3 is present it is locked into the crustalline structure of the MnCl2.

Chemistry Alchemist - 24-11-2011 at 07:02

i started to slowly evaporate it straight away from the reaction of MnO2 and HCl after filtering... so fairly acetic but wouldn't the HCl evaporate leaving neutral solution...? fairly puree MnCl2?

blogfast25 - 24-11-2011 at 08:56

No. HCl forms an azeotrope with water at about 20 % HCl. No matter what strength HCl you started off from, by evaporation you always end up with roughly 20 %. I tell a small lie: if you started from very weak HCl you'd probably not reach azeotropic composition.

Chemistry Alchemist - 24-11-2011 at 09:28

I geared from somewhere because it's just HCl gas dissolved in water, the whole thing will just evaporate to nothing... So do I have reasonably pure MnCl2?

blogfast25 - 24-11-2011 at 11:02

Quote: Originally posted by Chemistry Alchemist  
I geared from somewhere because it's just HCl gas dissolved in water, the whole thing will just evaporate to nothing... So do I have reasonably pure MnCl2?


Yes because the H2O/HCl azeotrope completely eveporates, of course.