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12AX7
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Manganese Chloride Crystals
-- Such a tease, I don't have crystals yet. I have a solution though.
Probably belongs in the book of complicated syntheses, since the MnO2/Mn2O3 came from hydrolysis of permanganate experiments.
Anyways, I took the sudges from those experiments and dissolved in HCl, outdoors.  Filtered more MnO2 from it, and I've been evaporating the solution seen below.
I'm guessing the color is a rather strong ferric chloride impurity. I'll probably recrystallize this once or twice.
Tim
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chloric1
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I have had astoundingly pure rose pink manganous chloride by dissolving electrolytic Manganese in about 15% HCl. The reaction puts Alka-Seltzer to
shame!
Fellow molecular manipulator
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12AX7
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It oughta! I was dissolving zinc the other day and had to dump it into a larger jar, too much foam...
Further observations: shit, I think the above solution is turning to gel. 
Tim
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Arthur Dent
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Sorry to revive this antique thread but i've been attempting to make some Manganese chloride out of MnO2 that's been thoroughly washed and cleaned.
So I put a jar filled with about 200 ml of slightly aqueous MnO2 sludge (I let it deposit and removed most of the water) in a water bath and then I
poured about 200 ml of conc. hydrochloric acid (100 ml at a time 30 min. apart). This is done outside at a temp. of 3 deg. C.
An ominous cloud of greenish yellow chlorine was hovering above the solution as expected and the jar warmed up a bit, but several hours later, my
solution is still stinkin' of Cl and is still a dark black sludge. Just how much HCl do I need to add? And how long is this reaction going to take
before it stabilizes and does not evolve any more chlorine?
I'm asking because until this has stabilized, I can't take this stuff inside, and the weather up north is calling for temps of -6 by Tuesday, so not
the ideal situation for glassware filled with solutions...
Thanks for any advice.
Robert
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Xenoid
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Hmmm.... at 3 oC, probably forever! You need to apply some heat to get rid of the chlorine. You'll need to concentrate the solution somewhat anyway,
before you get any MnCl2.4H2O crystallising, it is quite soluble, 74g/100g H2O at 20 oC. and still 63g/100g at 0 oC.
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gsd
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As Xenoid has said you need to heat (boil) the reaction mixture in order to drive the reaction to completion. Lot of unreacted HCl will remain in the
the solution. This is not a problem if you are not fussy about the purity of final product. Just filter off the sluge and boil the clear filtrate to
concentrate it enough to yield MnCl2.4H2O on cooling.
If you are looking for a purer final product the work-up becomes more elaborate. MnO2 ore is always associated with Fe impurities which will get into
solution as can be seen by reddish/yellow tinge to the othewise pink solution of MnCl2. The iron can be precipitated out by bringing the pH of
filtrate to about 4 to 4.5. This is best done by neutralizing the excess HCl with MnCO3 (or MnO), and then refiltering and concentration.
Gsd
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Arthur Dent
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Thanks, that's what I was afraid of, well I guess i'll have to pull out the ol' hotplate. I wish I didn't because even done outside, the amount of Cl
generated even at 4oC was quite impressive (and noxious).
And a fumehood is not in the plans this year (I just don't have the room, even though this would be way cool!).
Robert
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Xenoid
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Yes! I realise this doesn't help your current situation, but most pottery suppliers sell manganese carbonate as well as the oxide. Although somewhat
more expensive, the carbonate reaction with HCl is a little more benign 
Edit: See Doktor Klawonn's thread on making manganese carbonate from old batteries, in the "Reagents and Apparatus Acquisition" section.
[Edited on 13-12-2010 by Xenoid]
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blogfast25
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MnCl2 crystallises quite easily. I converted my stockpile a while ago to anhydrous MnCl2. Bu I'll be making 'pure pink MnCl2 hydrate shortly again...
If it's FeCL3 that's bothering you, try wahing with dry acetone: FeCl3 is highly soluble in it, not sure about MnCl2 though...
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Arthur Dent
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I finally got to the last step of my synthesis of MnCl2 (thanks to the holiday break), so I filtered-off the remaining liquor that I had prepared a
few weeks ago with MnO2 and HCl and indeed, instead of a pink-ish liquid, what I got was a crystal-clear but very dark solution, similar to 12AX7's
photo in the first post.
The solution looks like very dark tea, so I imagine that there are lots of iron impurities in this solution. Now I know that I could precipitate the
ferric chloride impurities with manganese carbonate but unfortunately, I don't have any.
But if I were to add some sodium carbonate, would this neutralise the solution enough to precipitate the iron? Or would the manganese precipitate into
a carbonate too?
I'm starting to regret making the MnCl2 first, I should have prepared the carbonate first as per Doktor Klawonn's thread. But for now, I'm stuck with
my contaminated Manganese Chloride and I'd love to be able to purify it in a simple way.
I do have some acetone, but do I need to completely crystallize my solution to use that process?
Robert
--- Art is making something out of nothing and selling it. - Frank Zappa ---
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S.C. Wack
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Quote: Originally posted by Arthur Dent  | | But if I were to add some sodium carbonate, would this neutralise the solution enough to precipitate the iron? Or would the manganese precipitate into
a carbonate too? |
Yes.
It would be sensible to precipitate some carbonate and use that. In a perfect world, an amount equal to the amount of iron.
[Edited on 3-1-2011 by S.C. Wack]
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UnintentionalChaos
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With a large amount of sulfite around, manganese will precipitate as MnSO3*3H2O. Any idea if iron (II) sulfite is similarly poor in solubility?
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
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blogfast25
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| Quote: Originally posted by Arthur Dent | I'm starting to regret making the MnCl2 first, I should have prepared the carbonate first as per Doktor Klawonn's thread. But for now, I'm stuck with
my contaminated Manganese Chloride and I'd love to be able to purify it in a simple way.
I do have some acetone, but do I need to completely crystallize my solution to use that process?
Robert
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Here’s what I did in a similar situation (Mn2+/Fe3+ mixture, this from pottery grade MnO2).
Precipitate everything with alkalised commercial bleach (NaOCl solution): the Mn2+ is oxidised to MnO2 and precipitates, the Fe goes down as
Fe(OH)3.nH2O. Re-acidify carefully with diluted acid to about neutral. Now wash the precipitate repeatedly with vinegar (very dilute sulphuric or even
very hydrochloric will also do, pH 3 to 4). Freshly precipitated Fe(OH)3 is soluble in weakly acidic solutions but MnO2 is not (with dilute HCl
you’ll get some reaction with the MnO2 but not much). This is perhaps best done as a series of decantations (not on the filter): add fresh vinegar
to the slurry, mix well, allow to stand, then decant off the supernatant liquid, repeat untill you can't find any Fe in the wash vinegar.
Then redissolve the washed MnO2 in strong HCl.
You can prove you’ve separated the Fe and the Mn by realkalising the vinegar wash that contains the Fe3+ (as acetate), Fe(OH)3 then precipitates
again..
Alternatively crystallise the mixture of chorides and treat with acetone, in which FeCl3 is highly soluble. This may not work brilliantly if your iron
contamination is fairly high because FeCl3.6H2O is a bit of a barstool to crystallise… Going by your description that may be the case...
Alternatively you may be able to take advantage of an annoying property of Fe(OH)3.nH2O: peptisation. When you wash a fresh precipitate of
Fe(OH)3 with distilled or deionised water what happens is that when the wash water (filtrate) gets below a certain ionic strength the Fe(OH)3 is
transformed from a precipitate into a colloidal solution, usually all at once! In plain English: the stuff just runs right through your filter! I’ve
not seen it with Fe(OH)3 yet but I have with Sn(OH)4 and it’s very frustrating. But I cannot vouch for it happening with Fe(OH)3 as a co-precipitate
with MnO2…
[Edited on 2-1-2011 by blogfast25]
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Arthur Dent
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Ugh. Back to square one as they say! I definitely will attempt the acetone method because frankly, I don't want to see Manganese Dioxide ever again!
LOL It's so dirty and hard to wash off. I'll concentrate the solution as much as I can by heat and then decant the resulting syrup with dry CaCl2.
I have a brand new can of acetone, so i'll just pour some along with the salt/concentrate in a glass flask and shake, hoping that some of the MnCl2
will precipitate a bit, then rinse off the ppt with fresh acetone.
Out of curiosity, could I use an electrolysis process to isolate/separate the metals?
Robert
--- Art is making something out of nothing and selling it. - Frank Zappa ---
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blogfast25
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| Quote: Originally posted by Arthur Dent | Ugh. Back to square one as they say! I definitely will attempt the acetone method because frankly, I don't want to see Manganese Dioxide ever again!
LOL It's so dirty and hard to wash off. I'll concentrate the solution as much as I can by heat and then decant the resulting syrup with dry CaCl2.
I have a brand new can of acetone, so i'll just pour some along with the salt/concentrate in a glass flask and shake, hoping that some of the MnCl2
will precipitate a bit, then rinse off the ppt with fresh acetone.
Out of curiosity, could I use an electrolysis process to isolate/separate the metals?
Robert |
Hmmm…I’m assuming you started from battery gunge, right? Well, it’s really the graphite in there (some 30 % of it) that makes that stuff so
black, dirty and hard to remove from ANYTHING! Now that you’ve got a relatively pure manganese salt, precipitating the oxides and dealing with them
is much, much less hassle, without the sticky, smeary graphite…
The trouble with crystallysing is that both chlorides are very soluble but you could try this. Reduce solution by simmering gently until first
crystals form. Cool and ice and isolate the crystals. Possibly these are MnCl2, obviously contaminated with FeCl3. What with several small aliquots of
acetone, until wash acetone is colourless. If it works, you’re likely to get low yield though…
Electrolytically? Fractionated electroplating may be possible but fresh manganese is very reactive: fresh powder react with water at RT, more reactive
than Mg powder… Fe3+ plates out when you add aluminium flakes to a ferric solution byt that introduces Al3+ into it, that then needs to be separated
out too!
[Edited on 4-1-2011 by blogfast25]
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peach
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ContactZ
If you like this discussion and want to see some pictures, send me a PM for my facebook name & msn
John
I have been working on this recently as well, in an attempt to make electroluminescent paint (Zinc sulphide doped with a fraction of a percent of
manganese). Were you doing this Tim? 
The source is zinc and alkaline batteries.
I have had precisely the same brown results from both Zinc batteries (which use slightly acidic zinc or ammonia chloride in their electrolyte and a
zinc casing with a carbon electrode) and alkaline batteries (which use KOH as the electrolyte and a zinc paste for the centre terminal, with a steel
can).
Universally, when exposing the washed pastes to hydrochloric, the result is so dark brown it is unacceptable as a clean source.
The impurity is ferric chloride, and a lot more than I was hoping would be in there.
Selective precipitation of heavy metals, like manganese, I believe is done when recycling ferric chloride commercially by adding finely divided iron
to the solution and churning it - causing the heavy metal contaminants to fall out of solution. The churning is because nickel sticks it's self onto
the iron and passivates the particle; the churning knocks bits of it back off to keep the process going, there is a japanese patent on a method /
apparatus for churning the solutions in a manner to effect good removal.
I have been repeating these experiments, as I too remembered gaining pink salt from a black paste.
According to my note book, this was the chloride. But, as usual, I haven't included enough details in that.
I figured 'John has probably missed some detail out, he may have made the sulphate and then treated THAT with hydrochloric', and repeated with another
Energizer battery and battery strength sulphuric.
Bingo... pink salt.
I am now considering the purity issue.
When I was dissolving the zinc casings in sulphuric, zinc is slightly more reactive than iron, so using weak sulphuric and a slight deficit, I could
bias the zinc away from the iron - hopefully.
When doing the manganese, I wondered if I was approaching it from the wrong direction, going straight to the chlorides and then trying to separate
them.
I wondered if by using sulphuric I could, again, bias it from the beginning.
I have yet to test the results, but hydrated iron sulphate is green, and the solution is a clean pink. Given how much iron chloride appeared in the
hydrochloric tests, I would expect to see a dirty / off colour pink if the same amount of iron sulphate had come through.
Search youtube for a video by nurdrage of him opening lantern batteries. It's probably on his page under manganese something. He uses the sulphate
method and goes as far as gassing the paste with sulphur dioxide. He claims this increases the purity. I am 100% willing to believe him, but I'll have
to ask if he's checked this.
In a twist of fate, after posting photos to my facebook and people commenting, I thought I'd make a video as I filtered the batteries and sat around
speculating what might be the easiest method for people who've not got anything to hand. I was recording it last night at five past eleven. Then
clicked 'todays posts' and found this. I'll post it up in the thread once I render it - on my 256mb of ram. 
This is where I wish I bought that' flame analyser and plasma spectroscope I saw....
I will also email Energizer and see if I can speak to a chemist or someone who's willing to tell me what the impurities in the source are, as they
publish detailed PDFs and specifically say they are trying to lower the heavy metal contaminants in these chemistries.
Their lithium cells also don't use the standard, toxic, electrolytes like thionyl. This is worth remembering when you see the videos about opening
energizer lithiums, if you then find another brand and plan to open that - as you may be in for quite a different experience.
Chloride salts
The solution on the right came from a Zinc lantern battery, the one on the left came from a single AA energizer alkaline.
In both instances, the paste was washed with water, filtered, the cake retained and treated with hydrochloric (fumes chlorine - lots) and then
filtered, retaining the filtrate.
The sulphates
Same again with an energizer battery, but this one was done with sulphuric. I've added some foil and a strip of paper as colour references.
This solution is weak. There is only one AA energizer alkaline in there. I'm sure it'll look much nicer once I boil it down. I'll post another photo
when done (probably tonight).
Paper that may be of use (I can't get them anymore)
Separation of iron from manganese ore roast-leach liquor
"The leach liquor obtained after the extraction of manganese from a low grade manganese ore contained both manganese and iron. In this paper an
attempt has been made to remove iron by aerial oxidation. Parameters such as pH, temperature and time were varied in order to arrive at an optimum
condition. It was observed that at a temperature of 333 K, pH 5.5 and aeration time of 2.25 h the iron could be completely removed at an air-flow rate
of 1000 cm3 min−1 without any loss of manganese."
^^^ this sound easy and, according to them, like it could achieve a high separation.
Can anyone get this?
From a site on drinking water
"High levels of dissolved or oxidized iron and manganese greater than 10 mg/l can be treated by chemical oxidation, using an oxidizing chemical such
as chlorine, followed by a sand trap filter to remove the precipitated material. Iron or manganese also can be oxidized from the dissolved to solid
form by adding potassium permanganate or hydrogen peroxide to untreated water.
The ideal pH range for chlorine bleach to oxidize iron is 6.5 to 7.5. Chlorination is not the method of choice for high manganese levels since a pH
greater than 9.5 is required for complete oxidation. "
Allergy buyers club
[Edited on 5-1-2011 by peach]
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Random
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How do you get rid of all that chlorine? It's heavier than air so it should stay on the place where it is released longer time, that could make a
cloud of deadly gas hanging around my house. I am scared to mix my battery mno2 with HCl, maybe if I would take some old clothes, soak them in sodium
bicarbonate solution and then put them on top of the jar for reaction could neutralize the chlorine?
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peach
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I would not recommend you do this inside - particularly if you're not used to dealing with that kind of thing in dangerous quantities.
Please heed that warning - as I rarely bother giving them.
One AA cell's worth + the hydrochloric will be enough to fill a room in the house with enough that it'll start to hurt.
Chlorine is more of a problem than gases like hydrogen chloride or sulphur dioxide and the others, as they sting immediately when you inhale them, and
they will dissolve very quickly in moisture in the air or surfaces around them.
Chlorine, on the other hand, it is quite a lot easier to stand in the room thinking "This isn't so bad", but hours later it gets a lot worse - even if
you left the room a long time ago.
Chloride ions control the thickness of the mucus lining in your lungs. By breathing so much in, you trick your lungs into thinking the mucus needs
thinning down. They will do that to an extreme, effectively drowning you - this is called Pulmonary Edema (pulmonary meaning your lungs, edema meaning
'abnormal water build up') when it becomes a problem.
It's unlikely you'll die from one AA's worth, but the feeling of being suffocated for hours is not nice at all!
To get rid of it I simply do it outside in the garage, open the windows and doors and leave the room. It will dilute down a huge amount once it gets
blown around in the air over the house to swimming pool levels and then nothingness.
But inside the house, it's going to concentrate.
This is why I think it would be better to go to the sulphate first, as you don't have the mountains of chlorine to deal with.
The stuff will actually fizz, for quite a while, as the hydrochloric goes onto it otherwise. With a lantern battery's worth, it will give off enough
to turn the glassware green as it hangs around. - if you were doing that much D cells and up, you would need the correct glass and / or a fume hood if
it wasn't a good distance from other people.
In the video I made, you can see me adding it to the hydrochloric, so that may be of interest to you - I'll have to sort the rendering out (take
AAAAAAAges and the video is just showing me experimenting with different batteries and ideas for the people also experimenting, it's not a guide or
how to).
[Edited on 5-1-2011 by peach]
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peach
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Updates on the pink
I tried boiling the sulphate down just now and got the following photos.
WHERE THE HELL IS IT?
I've had pink solid from one of these before.
And yellow crystals? Iron sulphate would be green from memory. And the manganese sulphate is pink. So what's that?
After roasting it on the plate (?350C+?), I failed to see anything other than fainter yellow. I had to move it outside to get the last of the
sulphuric off.
I have now emailed Energizer about the specific content of the paste and I have dug out a big collection of alkalines to give them a go.
I think I'll try Panasonic, that name seems to ring bells - but those bells were ringing months ago.
I'll try a few of these and probably try gassing one with SO2. I'll also have to have a look at the electrochemical methods.
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Arthur Dent
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Thanks for the info Peach! Lots of info to digest!
So your pink solution boiled off to yellow crystals! That's... unexpected!
Have you tried rehydrating it to see if it would give a pink solution again?
I agreee with you that the reaction of MnO2 battery paste with HCl produces insane amounts of chlorine, like an eerie green cloud hovering over the
beaker. I did this outside and even at that, got a whiff or two of chlorine gas even though I tried to avoid getting close to the reaction vessel
while stirring and it is indeed suffocating.
I'll experiment further with this stuff over the weekend and post the results.
Robert
[Edited on 6-1-2011 by Arthur Dent]
--- Art is making something out of nothing and selling it. - Frank Zappa ---
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DJF90
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Here's that paper you wanted John
Attachment: Separation of Iron from Manganese Ore Roast-leach Liquor.pdf (228kB)
This file has been downloaded 1375 times
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Random
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I will try the test tube amounts of mno2 and hcl to become friendly with chlorine as much as I can be before mixing bigger amounts Thanks for your advice, maybe I could also mix mno2 with potassium metabisulfite to
dissolve it and then precipitate mn as carbonate salt. This could be used to avoid chlorine.
That yellow crystals could be decomposed mnso4 or something, I saw that nurdrage concentrated the solution of mnso4 by boiling and then dried it in
desiccator to pink crystals to maybe avoid that.
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peach
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@DJF
Thanks! I've sent you a new email for my msn, since the old one was lost when reformating
@All
I have a vacuum desiccator and some of the salts do go back to the oxide if overheated - which usually shows up as brown.
I was trying to get mine clean in a manner that others without chemistry equipment or many chemicals would be able to emulate.
As such, I have started trying to clean them up using nurd's hydroxide method.
You decant some of your contaminate muck into another container, then add some strong base (I used KOH).
The other metals go to their hydroxides and the solution will go all lumpy and gel like - I added a bit more water to break it up into a fine
suspension.
You dump all that out into a filter and drain off the excess base in solution - saving the cake of iron and manganese hydroxides. Then thoroughly wash
the cake to remove any surplus Na or K ions from the cake.
Once it's well washed, you add some of the cake back to the main stock of your contaminated muck and, hopefully, the manganese remains in solution and
the iron drops out. It's like the pH leaching but without the need for a pH meter.
It's sitting downstairs in the filter, so I'll give it a try today and see if it will clean mine up.
[Edited on 6-1-2011 by peach]
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blogfast25
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This all seems a bit like reinventing the wheel, if you ask me. Threads on Mn2+ compounds from battry gunge galore on this forum. The self-proclaimed
‘manganese nut’ here used to be 'DerAlte', search for ‘manganese’ threads by him.
MnO2 dissolution in HCl, H2SO4, gassing with SO2 or reduction of MnO2 with H2SO4 – oxalic acid (Nurdrage), it’s all good and not very hard to do.
| Quote: Originally posted by peach | I have now emailed Energizer about the specific content of the paste and I have dug out a big collection of alkalines to give them a go.
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John, I hope you don’t seriously expect Energizer to reveal their production secrets to you, do you? All these electrolyte past compositions
probably resemble each other but all are probably also rather fine-tuned… Not something they’ll tell any Tom, Dick or Harry, IMHO.
Have you at least checked your pretty yellow crystals for manganese?
[Edited on 6-1-2011 by blogfast25]
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peach
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| Quote: | | MnO2 dissolution in HCl, H2SO4, gassing with SO2 or reduction of MnO2 with H2SO4 – oxalic acid (Nurdrage), it’s all good and not very hard to do.
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Pouring acid over the washed paste is as easy as making a cup of coffee to drink in the meantime.
The problem is that it often yields massively contaminated results, and separating the contaminants from the manganese, with a decent to high level of
purity at the end, is less simple.
| Quote: | | John, I hope you don’t seriously expect Energizer to reveal their production secrets to you, do you? |
There is a phrase I have come to well associate myself with, and put into practice, that being... "If you don't ask, you don't get" and it's sister
"xxx, you'll never know".
And since you ask.... :p
The electrolytes are explained in their PDF's, as well as many other details about the battery construction.
This is likely because they have patents and / or design rights on the construction and chemistry.
The thing I'm quizzing them over is the contamination in the manganese source. I am playing their game as well, as they specifically state in many of
their PDF's that they're attempting to lower health hazards and things like heavy metal contamination, so I am hopeful they'll reveal more information
based on that.
I'm not after the workings or magic of their chemistry, just the tolerances. I expect they have to declare the impurities for legal reasons, as it's
hazardous waste. Whether or not they're willing to do that without a prod from THE LOW-AH is another thing.
Anyway, they've put me on to someone to speak to, so I'll get on that.
I know the manganese is there. It's a manganese chemistry and the solution is pink, and I've had pink solid back from it before. It's the purity
issue.
I'm using this for doping a semiconductor style material, so pure is a good thing if possible - as I'll be using less than 1% to the bulk material and
it's effects are dramatic.
I was also pottering around outside filtering the gunk when I began wondering if all the sublimation and boiling points discussed in regards to AlCl3
would be of any use.
But for now, right now, I'm filtering the hydroxides and will be treating the contaminated extract tonight.
Fingers crossed.
Some 'magic paintings' will be on the way later.
All the best,
Not Tom, Dick or Harry
[Edited on 6-1-2011 by peach]
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