Sciencemadness Discussion Board

Manganese Chloride Crystals

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12AX7 - 2-9-2006 at 13:45

-- 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. :P 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

Chem_MnCl2_Sol.jpg - 15kB

chloric1 - 2-9-2006 at 14:27

I have had astoundingly pure rose pink manganous chloride by dissolving electrolytic Manganese in about 15% HCl. The reaction puts Alka-Seltzer to shame! :P

12AX7 - 2-9-2006 at 17:19

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. :o

Tim

Arthur Dent - 12-12-2010 at 12:56

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

Xenoid - 12-12-2010 at 17:55

Quote: Originally posted by Arthur Dent  
And how long is this reaction going to take before it stabilizes and does not evolve any more chlorine?
Robert


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.

gsd - 12-12-2010 at 20:15

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

Arthur Dent - 13-12-2010 at 05:03

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

Xenoid - 13-12-2010 at 10:54

Quote: Originally posted by Arthur Dent  
... the amount of Cl generated even at 4oC was quite impressive (and noxious).


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]

blogfast25 - 15-12-2010 at 09:53

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...

Arthur Dent - 2-1-2011 at 07:51

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

S.C. Wack - 2-1-2011 at 11:04

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]

UnintentionalChaos - 2-1-2011 at 11:16

With a large amount of sulfite around, manganese will precipitate as MnSO3*3H2O. Any idea if iron (II) sulfite is similarly poor in solubility?

blogfast25 - 2-1-2011 at 13:36

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


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]

Arthur Dent - 4-1-2011 at 13:04

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

blogfast25 - 4-1-2011 at 13:32

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]

peach - 5-1-2011 at 04:59

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. :D

This is where I wish I bought that' flame analyser and plasma spectroscope I saw.... :D

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]

Random - 5-1-2011 at 06:06

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?

peach - 5-1-2011 at 09:16

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]

peach - 5-1-2011 at 09:38

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.









Arthur Dent - 5-1-2011 at 16:02

Thanks for the info Peach! Lots of info to digest! ;)

So your pink solution boiled off to yellow crystals! That's... unexpected! :o 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]

DJF90 - 5-1-2011 at 16:24

Here's that paper you wanted John

Attachment: Separation of Iron from Manganese Ore Roast-leach Liquor.pdf (228kB)
This file has been downloaded 873 times


Random - 5-1-2011 at 16:43

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 :D 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.

peach - 6-1-2011 at 00:11

@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]

blogfast25 - 6-1-2011 at 04:36

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.



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]

peach - 6-1-2011 at 08:54

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.


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 :D

[Edited on 6-1-2011 by peach]

blogfast25 - 6-1-2011 at 09:48

Quote: Originally posted by peach  
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.



Try washing your crystals with acetone. FeCl3 in small amounts is yellow. Worked wonders for my ZrOCl2.8H2O which went from light yellow to snow white...

peach - 6-1-2011 at 10:14

Indeed.

I will give the acetone a go at some point - the stuff is bubbling away right now with the hydroxides method, so I have to see how that goes (need to give it time).

The problem with the ferric chloride idea is that material came out of an alkaline battery, and was treated with sulphuric acid.

Iron sulphate is green.

blogfast25 - 6-1-2011 at 14:15

That separation of MnO2 and Fe(OH)3.nH2O using vinegar (or other weak acid solution) worked, you know...

peach - 8-1-2011 at 04:56

Keep your jars and containers appropriately rinsed and cover them (and the filters) whenever possible to keep the purity up.

Small amount of contaminated solution decanted and mixed with KOH to produce the insoluble hydroxides of the metals

Filtered through coffee filters and thoroughly washed to remove excess base and contaminating ions

Cake retained

Cake mush returned to the original contaminated mess and thoroughly stirred

Solution aerated for approximately 4 hours with a nebulizer (you can leave it standing for ages if you don't have a pump. Don't use a fridge pump for this - grease - needs to be an aquarium style one; diaphragm)

Filtered through yet more coffee papers

Ta-da... bubble gum pink

Filter the insoluble hydroxides. I'm using plastic funnels, coffee filters and Kilner jars to demonstrate it can be done without anything special - indeed, that being the reason for me choosing the hydroxide method, as it can also be kept reasonably pure and requires only some NaOH or KOH to repeat verbatim

The brown gunk is from treating some of the contaminated solution with base

Note the nomograph over the filter to keep things clean - everything has been thoroughly washed



There's the cake, the insoluble hydroxides. I'm filtering them away from the base solution. I have poured some extra water through and emptied the cake out to give it a swirl in some fresh water and to swap the paper. I used approximately 2-3l of water to rinse - ALWAYS RINSE BEFORE AND AFTER CONDITIONING FOR THE HIGHEST GLOSSINESS AND MOST LUXURIOUS HAIR IN TOWN.


Bubble boggling the solution - I have emptied the cake out into the remaining original solution (contaminated brown / yellow) and stuck a thistle funnel in there to blow air through, wrapping the top of the jar with cling film to keep dust, hair and everything else from getting in there

I'm using the funnel as it's clean and will keep the mess off the nebulizer

The Kilner jar I got at ASDA for £2, and it's absolutely great for this kind of work when you a.) don't have or b.) don't want to be bothered cleaning lab glass. The hinged top and rubber seal means you can easily cover the work if you want to stop for a while or go to bed, or leave it on the shelf for months :)

I bought this in the middle of doing the work when I ran out to get some more papers, thinking "ooooo.... that looks handy to prove the point"

The glass isn't suitable for extreme heating like borosillicate, but is close in terms of reactivity and the rubber the seal is made from is the same as the natural rubber used for Suba-seals (septum seals), stoppers and caps on Sure-seal bottles from sigma (it's all natural rubber).

Natural rubber WILL react with really reactive acids for example, but the art is in not storing your jars upside down :D

You could also up the seal rating by replacing it with a bead of silicone sealant - silicone is the classy version of Suba-seals normal natural rubber seals





Filter it again and the manganese salt (the chloride in this instance) should now have fallen away from the insoluble iron forms. The salt you're after will drip out, collect, look after, love and then set free


Or boil it down, your choice

Using the chlorides, or testing the end result with hydrochloric, may be good, as the iron chloride is extremely bright and has a high pigment value, meaning it'll show up well as you do this

This idea came from some inspiration on the part of Arthur, who mentioned dissolving battery casings in hydrochloric and not seeing any yellow - which is great! As I was concerned they may contain some iron. I'd used sulphuric for mine, but his mention of trying hydrochloric and not seeing anything has subconsciously inspired this idea of using the same reasoning with the manganese.

My zinc sulphate was also pristine white, and the sulphate of iron is green.

Those two things make me hopeful it's a source of decent purity zinc, which I will need for the bulk of the special paint.



What are you doing now!? I'll tell you what you're not doing.... opening my dinner!


YUM!

Really does look a lot like Bubbalicious bubble gum - second hand, but still good



Up-skirt cam


Carefully and gently drying the last of the excess moisture off before recrystalization - cam

Not bad at all considering the starting point and that all you'll really need are some papers and a bit of base - no fine pH control, no special gas handling equipment or practice and no contamination by the use of kitchen condiments



Thanks go to nurd rage for doing the original work.

[Edited on 8-1-2011 by peach]

Arthur Dent - 8-1-2011 at 08:41

Splendid synthesis! Gives me some hope for my evil-looking tea-colored MnCl2/FeCl3 solution!

The results are indeed delicious-looking – like cotton candy – but I wouldn't eat a spoonful of that stuff! :o

I'll try to reproduce your excellent results soon, now if I can just find a place close by that has some KOH (sigh).

Robert

peach - 8-1-2011 at 11:20

Don't need KOH. :)

That's the beauty, I did this without using any of the special glass, equipment or chemicals.

NaOH will work - drain cleaner / lye / caustic soda

[Edited on 8-1-2011 by peach]

blogfast25 - 8-1-2011 at 11:30

All looks very interesting but thick as two short planks here doesn’t get it: you’re separating Fe(OH)3.nH2O from MnO2 by bubbling air through the slurry? MnO2 ends up as MnCl2? What reactions are involved?

It clearly works, so I’m interested…

peach - 8-1-2011 at 11:59

You've got most of it.

I'm starting with a big mess of the iron and manganese chlorides.

SOME of that gets split off and treated with base to produce the (insoluble) hydroxides of both - which are very well washed to remove the excess, soluble base.

The iron and manganese hydroxides go back into the original mess of iron and manganese chlorides, where they selectively convert the iron to the 3 hydroxide. The air is there to assist the production of the insoluble iron hydroxide.

For the sulphate, it looks like this

4Mn(OH)2 + 4FeSO4 + O2 + 2H2O -> 4MnSO4 + 4Fe(OH)3

The 4MnSO4 remains soluble, the Iron hydroxide goes insoluble, so filtering the result is all that's needed.

I doubt anyone is starting from something much worse than I did with this. Look back to the original photos, the stuff is on it's way to brown / black.

The thing I REALLY like about this is that it doesn't require pH control. I have a pH meter that will read to 0.001 (with a separate temperature input and a datalogger that I can make graphs with on the computer as it happens or store up hundreds of thousands of points for later), but very few others do. And I'm doing this to see if others can produce this special paint at home, so the method I use needs to be applicable to what others will have available and their level of ability - most of them are into electronics, not chemistry.

Have faith, you can yield the pink stuff from it - and it's not too difficult

[edit] I'd still be interested to know what that yellow was in the sulphuric treated beaker. Manganese sulphate... pink, iron sulphate, green..... yet yellow crystals. And no, it didn't go back to pink on re-hydrating it, but I bet the manganese is still in there.

[Edited on 8-1-2011 by peach]

blogfast25 - 8-1-2011 at 12:56

The bit I don’t get is why your in your mixed hydroxides (to be clear: Mn(OH)2 + Fe(OH)2/Fe(OH)2) Mn2+ doesn’t oxidise to MnO2:

Mn(OH)2 === > MnO2 + 2 H+ + 2 e-

Which definitely proceeds in alkaline/neutral conditions, as anyone who has ever precipitated an Mn2+ salt other than MnCO3 will testify to: the whitish Mn(OH)2 turns brown before your eyes, never mind blowing oxygen through it! Or even extensively washing it with water… I made some MnO once and the only way I know is precipitating as MnCO3, careful drying and ‘burn’ in inert atmosphere to MnO.

Of course you’ve got Mn2+ in the part you didn’t add base to. Now you mix it with well-washed Fe2+/Mn2+ hydroxides and start the blower. Now you selectively oxidise the Fe2+ to Fe(OH)3 while solubilising the Mn2+.

You must lose quite a bit of the Mn to MnO2, stuck in the Fe(OH)3.nH2O, right? Still, it’s quite a neat little trick…

peach - 8-1-2011 at 13:20

Well, if there's any acid left in the solution you add it to (a good possibility), the MnO2 is likely going back (partially) to the salt of that acid as it's re-added if it's formed in the cake. As will the iron, which means it should really be free of acid before re-adding the hydroxides.

I only used the supernatant from the battery extractions to produce the hydroxides, the bulk of the chlorides (the solid at the base of the jars) was left for treating with them.

This is likely a slightly lossy process. As is recrystalisation.

Quality is the issue here. It needs to be as pure as possible for someone who can't get anything more than papers and some general chemicals.

The manganese will be used as a dopant at below 1% to the bulk material, so a gram will produce over 100g of the finished paint.

The thinner the paint is applied, the better. Typical thickness are around a micron for commercial displays I think - not a lot at all.

I'm sure the materials could be made even purer, but if they don't work at this level, it's inaccessible to the people who want it. I am testing this accessibility.

I still have to find some easy, yet effective and pure method for reducing the sulphate for them. And then dope it, which means potentially a day or two of solid roasting the components to drive the manganese ions into the sulphide lattice.

Then spin coat it, and then test it.

The typical drive in temperature is around 600-900C, over the working temperature of borosillicate, so that needs sorting. And it's typically done under a reactive atmosphere using gases even I can't get, which also needs negotiating on their behalf.

LOTS to get on with! :D

[Edited on 8-1-2011 by peach]

blogfast25 - 8-1-2011 at 14:40

It probably works best if it's a real mix of Fe/Mn, not just 5 - 10 % Fe contamination. There, selective dissolution of the Fe(OH)3 (after complete oxydation of both to resp. III/IV) with weak sulphuric or acetic acid is simpler. The leachate will contain traces of Mn but the Mn part should be completely Fe free.

chemoleo - 8-1-2011 at 19:41

I haven't read all of the thread so forgive me if this is a repeat- but-
Manganese sulfate, MnSO4 (AR grade), forms very beautiful large rose crystals. However, it seems to dissolve itself in its own crystal water when the temperatures are too high, and its tendency to form very highly saturated solutions cause it to crystallise rapidly (freezing the solution contained in a whole beaker) rather than forming a nice looking single crystal (i.e. hanging off a nylon thread).
Nonetheless, one of the prettiest of crystals I have seen!


[Edited on 9-1-2011 by chemoleo]

peach - 25-1-2011 at 19:56

Quote: Originally posted by chemoleo  
I haven't read all of the thread so forgive me if this is a repeat- but-
Manganese sulfate, MnSO4 (AR grade), forms very beautiful large rose crystals. However, it seems to dissolve itself in its own crystal water when the temperatures are too high, and its tendency to form very highly saturated solutions cause it to crystallise rapidly (freezing the solution contained in a whole beaker) rather than forming a nice looking single crystal (i.e. hanging off a nylon thread).
Nonetheless, one of the prettiest of crystals I have seen!


You're correct.

I've had a go at recrystallising some of the results from the manganese work and it's not a particularly fun activity.

You're dead on with the mention of them either being in solution or hanging onto their water of crystallisation.

In practical terms, for those wishing to have a go, that means your solution will gradually drop with evaporation then, once almost all of it is gone, the crystals appear and you don't really have a lot to pour off. And what you are pouring off likely has a lot of the salt still in it.

It also makes filtering small volumes difficult, if the solution is anywhere near saturation and cools during filtering. Even though hot recrystallisation doesn't work at all well, you can find solutions going 'slushy' and not draining through filters if they've been moved from somewhere warm to cool.

But anyway, it can be done with some care and if you can put up with losses. You can always use the still manganese rich waste as a lower grade material.

I dissolved the slop from the hydroxide precipitation clean up in the smallest possible volume of hot distilled water, then set it on the radiator (45C / 30-40% RH) with a filter paper over the top. The next day, enough of the water should be gone that the crystals have reappeared. The remaining water may prove tricky to pour off. In fact, I could turn the dish upside down and it wouldn't drain. I had to give it a gentle tap over about ten minutes to encourage it away.

I didn't empty it out into a filter paper because that would entail mushing the crystals back in with the water, somewhat defeating the point.


Done! And perfectly doable by you with a spoonful of caustic soda.


These may appear browner than they actually are, as the surfaces and floors of the rooms are all wood, so the light is always yellow heavy. Not to mention, I'm one of 'those people' who still use incandescent bulbs, since the dimmers don't understand the word efficiency. The salt is pink in person.


The crystals will take forever to dry out alone, so I'm putting them under vacuum before warming them up. The vacuum helps the water leave but, of equal importance is that, the manganese will go back to the oxide if it gets hot in the open atmosphere.


After an initial drying, they get crushed up and go back in


I started with 21g of damp, potentially mucky stuff


And now have around 6 of recrystallised, filtered, drier salt. Which should theoretically be cleaner but, lacking the plasma chromatograph I wanted to buy, I can't actually test that with any high degree of accuracy


What John is using his for

I'm using around 0.03 to 0.01g of the manganese salt as the dopant for a zinc sulphide based pigment - which glows when exposed to an electric field (like the displays in a Betamax).

0.026g of the salt is less than you'd pick up by pressing your finger tip into a pile of it.

Those fractions of a gram need very, very thoroughly mixing with the zinc sulphide. Preferably so every microscopic grain of the sulphide is in contact with the manganese.

Back on the ball mill... for another 24h...


The resulting slurry, the next day


Moisture is driven off under vacuum and a low heat, then it's put under a stream of argon. Argon is usually heavier than air, this cylinder however, is possessed by Regan and hovering, thusly.


The pigment is roasted at 450C...


for about...


a long time. This is typically done in quartz boats, in a tube furnace, and an atmosphere of carbon disulphide or something similar - wherein it takes about two hours. Using lower temperatures and an inert atmosphere will mean it takes longer for the manganese to migrate into the lattice of the zinc sulphide. It's a diffusion process. The heat causes the lattice to vibrate, allowing the manganese to wiggle into place. The hotter it is, the more wiggling going on, the faster it happens - ain't it always the way!


Looks like some oxygen was left in there. START AGAIN!

The pigment isn't displaying any electroluminescence at present. Likely because this batch has oxidised.

Another possible candidate is my less than ideal electrical test setup. I'm simply wiping some of the result over a capacitor I made by laminating pieces of kitchen foil, which then get connected to the mains - sans fuse. I'm looking for any faint signs of glow around the overlap between the foil.

I have also tried powering the foils up using a TENS unit, which allows me to switch from an AC to a DC signal, as well to vary the frequency up to 100Hz (should be fine) and the pulse width (but only to 250uS, which may not be so great).

I would expect to see some signs of life, even if the conditions aren't perfect, as the pigment will respond over quite a large range I'm lead to believe.

Commercial, EL pigment is silk screened at a few tens of microns in thickness. It is also more commonly sandwiched between two layers of glass, with one slide being sputtered with aluminium and the other being a transparent & conductive coating, placing the pigment directly in the field.

I have made this pigment before and it was a light sunset orangey / salmon pink colour after driving it in, perhaps due to me using slightly more of the dopant. I also purged that a lot more thoroughly.

Once I get a spare hose tail in the post and manage to find the needle valve, I'll have another go.

I may also dispatch by courier pigeon samples to those nerds better equipped with the power supplies, frequency generators and conductive slides to have a go with it


The plate capacitor he speaks of. I have also seen these phosphors being illuminated with twisted magnet wire, which may be a better idea, as it's likely the pigment will come into more intense spots of the field than it will around the periphery of these plates


[Edited on 26-1-2011 by peach]

blogfast25 - 26-1-2011 at 09:52

Oxygen must be the great killer here, at higher temps it will oxidise your MnCl2 in a jiffy. Hard to get rid of the last traces of O2, considering also just how little Mn2+ you’ve got in there. Ever considered using MnS (the only ‘flesh coloured’ sulphide, apparently…) instead of MnCl2? Or start from anhydrous MnCl2 and just dry ball mill that with the ZnS till kingdom come? Then calcine in anaerobic conditions?

All in all this strikes me as a Real Boffin Project requiring Real Boffin Toys…


[Edited on 26-1-2011 by blogfast25]

mr.crow - 26-1-2011 at 10:51

Quote: Originally posted by peach  




Peach: I'm here to kick ass and chew bubble gum, and I'm all out of gum

peach - 27-1-2011 at 00:45

Finding information on how to do this is not particularly easy, as the majority of the people who are interested are making things that only really get sold to electronics companies, large amounts of it is in journal papers and they're far too expensive to be buying one by one. Also, the more recent the method is, the more likely it is to involve something too complex to do at home.

I was doing this after someone mentioned how it'd be good if electronics nerds could make it themselves, because then they could make all kinds of displays and visually exciting things. A lot of the phosphors require things you won't find at home. But zinc sulphide doped with manganese, both of those can be found in batteries, so it seemed like a possibility.

The first issue was obtaining both of those as simply and in the purest possible state using the most basic items and chemicals available. I've done that to the highest possible standard I think someone else with little or no chemistry knowledge could repeat.

But driving it in, as you point out, means that oxygen needs to go. Not only will oxidation ruin the manganese salt, it will start the zinc sulphide off going back to sulphate, which will then need washing out as well, as the sulphate isn't active. I haven't tried it with manganese sulphide, as all the information I originally found was related to use chlorides and sulphates, and getting those out of the battery paste wasn't difficult - particularly the sulphate since it doesn't involve the plumes of chlorine.

With those done, driving the manganese into the lattice is the only remaining problem. I am hoping, optimistically but it's better than being negative, that the diffusion rate will not be orders of magnitude lower at lower temperatures - meaning, doing it at 450C instead of 900C would mean it'd take days or weeks.

Obviously, borosillicate glass is not at all suitable for 900C work, and more so when it involves hours worth of exposure. And the average person wanting to make some of this pigment won't have quartz. I am also trying to avoid the use of special atmospheres beyond inert gas, as too many of them are extremely flammable or toxic.

Once I have the needle valve found, I will try again, but this time purging the sample much more thoroughly. Another problem with purging it is, there will likely be pockets of atmosphere distributed throughout the finely divided sample.

The sample is ball milled in a little methanol, which helps distribute the powders and encourages them to flow better. Otherwise, there's too high a chance you'll end up with the dopant staying in one place as it's milled. Removing the methanol is not tricky at all and other, easy, solvents can be used instead.

It may also be beneficial to remove the methanol once it's in the tube it'll be heated in, as the slurry will be much better packed once dry and less likely to be full of air pockets than a fluffed up and transferred sample. Another method to ensure it's well purged would be to suck the atmosphere out with a vacuum pump and then back flow the new atmosphere. Again, that's making it more complex for others.

If all that fails, I'd have to abandon wide spread repeatability and start with bits of quartz and hydrogen sulphide.

Finding useful things in failure, if my sample above has oxidised (which it likely has), it at least shows the manganese was well distributed, judging by the homogeneous colour tone.

[Edited on 27-1-2011 by peach]

blogfast25 - 27-1-2011 at 08:10

How about co-precipitating ZnS with MnS? To be honest, considering the ratio of Zn/Mn and the solubility constants of Mn(OH)2 (2 E-13) and MnS (3 E-11) I’m not even sure whether there is a pH range in which this is physically possible but that can be verified real easily. With such small quantities your Mn may truly be occluded into ZnS crystals, a finer dispersion you couldn’t dream of. Then carefully water wash, vacuum dry and the rest of it...

If you want to stick with MnCl2 then I think pulling high vacuum and backfilling with pure argon (perhaps even several times) is probably the only real possibility.

Something niggles at me that obtaining this electrical field sensitive material must be simpler than all this…


[Edited on 27-1-2011 by blogfast25]

peach - 27-1-2011 at 08:38

I doubt it.

The other materials used to do this include cadmium, blue diamond, indium phosphide, gallium arsenide & gallium phosphide. They're usually made on a nanoparticle scale using sol-gels, ultra high purity materials and by semiconductor companies with cylinders of ungodly toxic stuff. Quartz and tube furnaces are the basics for them.

I haven't seen a single reference to anyone doing it at home.

If anything, I'm trying to get away with using barely anything compared to how it's usually produced.

The precise mechanisms by which semiconductor fabrication functions are not often discussed in detail outside of journal entries. I've spent quite a lot of time wandering around university physics labs and studying it, so (outside of this forum) I'm interested in physics, electronics and machining as well. As an example, I consulted physicsforum.com (their version of SM) about bandgaps and doping and received a whopping one reply, from someone claiming to know it inside out, yet unable to explain what I was pointing out. Shuji Nakamura, who developed the blue LEDs, originally won $180 million from Nichia for the work - after he took them to court for giving him a $180 bonus for it. :o :D

[Edited on 27-1-2011 by peach]

blogfast25 - 27-1-2011 at 08:54

Well, in that case it sounds like you're trying to do the impossible. Good luck with that! ;)

peach - 27-1-2011 at 10:15

Quote:
The limits of the possible can only be defined by going beyond them into the impossible.

--Arthur C. Clarke

blogfast25 - 27-1-2011 at 12:36

'Impossible' was hardly le mot juste here: 'beyond your means' probably covers it...

peach - 27-1-2011 at 12:51

Beyond my audiences' means covers it better still, as I'm purposefully using the least toxic and most basic of things knowing they are more interested in the product than the chemistry - hence the Lego ball mill.

[Edited on 27-1-2011 by peach]

blogfast25 - 27-1-2011 at 14:13

You've kind of lost me there: your ball mill sure looks small but if it works, so what. Well, does it?

peach - 27-1-2011 at 15:14

It does.

I started trying to make this pigment not only for myself, but because other people are interested in being able to make it themselves (at home). These other people have spent most of their lives studying electronics and programming, so they are not at all well equipped in terms of chemistry - we're talking, they'll have to go out and buy test tubes, and they're going to have nothing in the way of chemicals.

Part of the goal was to see if I could make the pigment in a manner that they'd be able to repeat themselves. This is why I'm trying to use the absolute basics in terms of equipment, practical complexity and toxicity. The hydroxide precipitation method is an example of that, in that it is so simple compared to the others yet also functions so well.

I can't make the reagents much purer, given that this was supposed to start with a battery to make it accessible. That's done. I now need the driving in of the dopant to behave.

[Edited on 27-1-2011 by peach]

Arthur Dent - 17-3-2011 at 11:23

Update on my synthesis of Manganese Chloride...

A while ago, I decided to use the hydroxide method to precipitate the iron impurities in my Manganese chloride solution I made out of battery crud and HCl, so I put aside 1/5 of the MnCl<sub>2</sub> solution and mixed it with NaOH.

I got a lovely chocolate milk-colored solution which rapidly deposited into a brown mush. Washed and filtered thoroughly and I dried the mush which gave me a few grams of very fine brownish powder.

Then the stuff got put away on the top shelf for a few months until now. So yesterday, I dropped the tea-colored Manganese chloride solution in a flask and with a small funnel, poured the hydroxide dust in the flask. Instant blackness!

The solution right now is pitch-black! I know I have to "aerate" the solution a bit for the reaction to take place, but I don't have an air pump for now. More to come during the weekend.

Robert

blogfast25 - 18-3-2011 at 06:41

Quote: Originally posted by Arthur Dent  


[...]

Then the stuff got put away on the top shelf for a few months until now. So yesterday, I dropped the tea-colored Manganese chloride solution in a flask and with a small funnel, poured the hydroxide dust in the flask. Instant blackness!

The solution right now is pitch-black! I know I have to "aerate" the solution a bit for the reaction to take place, but I don't have an air pump for now. More to come during the weekend.

Robert


Well, that IS decidely strange and you may be ‘onto something’.

When neutralising MnCl2 with NaOH, firstly Mn(OH)2.nH2O (off white) precipitates, which upon contact with air quickly goes brown by oxidation:

Mn(OH)2 === > MnO2 + 2 H+ + 2 e
½ [ O2 + 4 H+ +4 e === > 2 H2O ]

Mn(OH)2 + ½ O2 === > MnO2 + H2O

Given a little time and air and stirring, this reaction proceeds to completion.

What doesn’t happen is IMMEDIATE precipitation of something black: freshly precipitated MnO2 is always very dark brown, rather than black anyway and made the way you did it starts off from off-white (Mn(OH)2). Even when adding a strong and alkaline oxidiser like hypochlorite (commercial bleach) you get that typical colour transition, only much faster.

I’m also surprised you describe the remainder of the MnCl2 solution as ‘tea coloured’: it should really be pink or if weak, almost colourless.

One possibility is that the solution contains iron, in mixed state of oxidation: Fe2+ AND Fe3+. Co-precipitated that gives black magnetite, Fe3O4 (see for instance ‘ferrofluids’)… The black would possibly mask anything else...


[Edited on 18-3-2011 by blogfast25]

Arthur Dent - 18-3-2011 at 09:32

Yeah I'm surprised also. Here is the list of the only reagents that were ever involved in this reaction:

1) Battery crud from Carbon Zinc batteries, thoroughly washed with distilled water.
2) Technical grade 32% Hydrochloric acid
3) Distilled water
4) Technical grade Sodium Hydroxide (small granules)

I can vouch for the relative purity of all the reagents. When I first prepared the manganese chloride some months ago, I expected to obtain a tea-colored solution just like the 1st post of this thread, due to the Fe impurities, so to get rid of the iron, I precipitated some of the solution into insoluble hydroxides with my NaOH, and I washed and dried the resulting precipitate. So far so good.

What I was expecting was that the dried Mn(OH)<sub>2</sub> would help precipitate the remaining Fe impurities and that with adequate aeration of the solution, I would obtain a pale pink chloride solution with a black/brown Fe/Mn precipitate at the bottom.

Now it looks like there's an inordinate amount of MnO<sub>2</sub> in suspension in the solution, but in that case it would be opaque and would have started settling, but it doesn't seem to settle at all. 2 days after the experiment, the solution is a very deep transparent brown, like cola, or even darker.

Robert




[Edited on 18-3-2011 by Arthur Dent]

blogfast25 - 18-3-2011 at 10:39

Quote: Originally posted by Arthur Dent  


[…]
When I first prepared the manganese chloride some months ago, I expected to obtain a tea-colored solution just like the 1st post of this thread, due to the Fe impurities, so to get rid of the iron, I precipitated some of the solution into insoluble hydroxides with my NaOH, and I washed and dried the resulting precipitate. So far so good.

What I was expecting was that the dried Mn(OH)<sub>2</sub> would help precipitate the remaining Fe impurities and that with adequate aeration of the solution, I would obtain a pale pink chloride solution with a black/brown Fe/Mn precipitate at the bottom.
Now it looks like there's an inordinate amount of MnO<sub>2</sub> in suspension in the solution, but in that case it would be opaque and would have started settling, but it doesn't seem to settle at all. 2 days after the experiment, the solution is a very deep transparent brown, like cola, or even darker.


Not sure what you mean by “dried Mn(OH)<sub>2</sub>”: on drying (in air) that turns immediately to MnO<sub>2</sub>. Did you mean you just allowed it to drip ‘dry’?

Have you tried filtering a small amount of the suspension? What does the filtrate look like?

I think peache’s method, assuming that’s what you’re trying to emulate to get rid of Fe, might be quite sensitive to experimental conditions…

Alternatively, precipitate everything with a small amount of hypochlorite, to a pH of 10 or so. This will ensure all Mn is as MnO<sub>2</sub> and all Fe as Fe<sub>2</sub>O<sub>3</sub> hydrate. Then acidify with dilute H<sub>2</sub>SO<sub>4</sub> (dilute HCl will also work, as will vinegar) to a pH of about 5. Allow to stand for a bit. The iron re-enters solution, the Mn stays as MnO<sub>2</sub>. Filter and wash and redissolve MnO<sub>2</sub> in strong HCl. Evaporate slowly to dryness.



[Edited on 18-3-2011 by blogfast25]

woelen - 18-3-2011 at 13:35

I never obtained a totally black precipitate from manganous solutions, not even at high pH in the presence of strong oxidizers. I have observed the brown color of hydrous MnO2 many times, but never pure black. There must be a strong contamination of the manganese in your case, but I am wondering what it could be. Alkaline iron(II)/iron(III) mixes also are not totally black, they tend to have a dark grey/blue color and on longer standing they turn brown.

Arthur Dent - 18-3-2011 at 14:24

Quote: Originally posted by blogfast25  

Not sure what you mean by “dried Mn(OH)<sub>2</sub>”: on drying (in air) that turns immediately to MnO<sub>2</sub>. Did you mean you just allowed it to drip ‘dry’?


Aw crud! There's my mistake. Yeah, I let it dry completely. Ugh.

My bad. So basically, I just chucked some MnO<sub>2</sub> and some iron oxide back into my solution. Back to square one. :(

I think my best bet will be to start from scratch with new MnO<sub>2</sub> from a better source. I learn new thangs everyday! ;)

Robert

blogfast25 - 19-3-2011 at 07:22

Quote: Originally posted by Arthur Dent  

Aw crud! There's my mistake. Yeah, I let it dry completely. Ugh.

My bad. So basically, I just chucked some MnO<sub>2</sub> and some iron oxide back into my solution. Back to square one. :(

I think my best bet will be to start from scratch with new MnO<sub>2</sub> from a better source. I learn new thangs everyday! ;)

Robert


No, no, no, no! Recycle the stuff you’ve got! There’s nothing wrong with it. And because it hasn’t been calcined yet, it will re-dissolve easily in HCl 32 %.

Alternatively try grinding it up and selectively leaching out the iron with a weak acid (or a weak solution of a strong acid) but the success of that method will depend on the state of dehydration of the Fe2O3 and its age. Fairly dry Fe2O3 will only dissolve in fairly strong acid (but that also dissolves MnO2!) It’s worth a shot: you can also try this selective leaching at BP to speed it up. Check the leachate for colour and with NaOH for any Fe(OH)3 dropping out. Easy.

-----

Or how about another route? Dissolve your stuff in 50 % H2SO4, applying heat. Both oxides will go into solution, Fe as Fe2(SO4)3 and Mn as ruby red Mn2(SO4)3 (yes, Mn [+III] sulphate, not [+II]!) Now add plenty of salt solution (NaCl). Mn3+ is as capable of oxidising Cl- as is Mn [+IV]:

Mn3+ + e === > Mn2+
Cl- === > ½ Cl2 + e

So your Mn [+III] is reduced back to MnCl2 with liberation of Cl2.

Then neutralise the solution carefully stirring with NaOH to about pH ≈ 8 - 9. Then further neutralise with Na2CO3. The stable MnCO3 (off white) precipitates. This DOES NOT OXIDISE in air (not even when wet) like Mn(OH)2 does, you can keep it for years. Filter and wash carefully. Dissolve in HCl to convert back to MnCl2. Of course that doesn’t get rid of the Fe but FeCl3 is highly soluble in acetone. Wash your MnCl2 crystals with plenty acetone to extract the soluble FeCl3…


[Edited on 19-3-2011 by blogfast25]

Arthur Dent - 27-3-2011 at 09:26

Since my solution of Manganese Chloride with Ferric impurities was in one of my nicest 500 ml boiling flasks, I decided to put it in a plain mason jar 'till I have time to experiment with it.

Since last week, the solution had turned back to its original dark tea-color, with a small black precipitate of MnO<sub>2</sub> at the bottom of the flask. So to rinse and clean off the flask, I decided to drop in a good 100 ml of 10% hydrogen peroxide and chuck the whole thing in the jar... Wow! Magic! :o

After a lot of fizzing and evolution of oxygen, the solution became water clear and all traces of the precipitate dissolved! It looked like plain water in the sun! That's pretty spiffy!

But very slowly, the liquid is getting a bit turbid and the Manganese Dioxide is slowly reforming a grayish suspension in the jar.

Now two intriguing questions...

1) If there are ferric impurities in my solution, the H<sub>2</sub>O<sub>2</sub>, shouldn't have affected the ferric chloride at all and the solution should have remained somewhat yellow after the peroxide addition, shouldn't it ?

2) My beautiful boiling flask shows grayish stains inside... so I filled the flask with H<sub>2</sub>O<sub>2</sub> and thought they would instantly dissapear, but they seem to be quite stubborn. I know that Manganese Dioxide is a bitch to clean off, so how should I deal with these stains?

Robert

Mixell - 27-3-2011 at 09:41

You can try pouring some on the solution to another vessel (if you don't have any manganese), and reduce the iron and the manganese in the solution to the metallic form using aluminium.
Take the manganese/iron powder and pour that into the first vessel and heat maybe to speed up the reaction (can be done with the aluminium too). And stop when all of the iron in the first vessel will be reduced and you should have a pure manganese chloride solution (presumingly you had enough iron/manganese powder).
I haven't done that myself, but by comparing the reduction potentials, it should work.

blogfast25 - 27-3-2011 at 10:11

Robert:

1) Yes. Fe3+ cannot affect H2O2 (neither oxidise nor reduce it). But don’t stare yourself blind to the colour: in dilute solutions the colour is quite deceptive. Just the water from your H2O2 solution could have diluted it enough to affect colour.

2) MnO2 is a bitch to clean off but strong HCl will get rid of it, each time and every time, fresh, wet, old, dry, whatever.

blogfast25 - 27-3-2011 at 10:13

Mixell:

Even if that does work, you're then contaminating your solution with aluminium and will have to separate these two (Mn and Al).

[Edited on 27-3-2011 by blogfast25]

Mixell - 27-3-2011 at 11:05

Well, I'll suggest using aluminium sheets (to easily separate from the power, this one worked fine for me).
Or to wash with concentrated alkali, either way, you will have an aluminium free powder.

blogfast25 - 27-3-2011 at 11:34

No, no, no. For Fe3+ to plate out, the Al has to go into solution:

Al(s) + Fe3+(aq) === > Al3+(aq) + Fe(s), a redox reaction: the Al supplies the electron needed to reduce the Fe3+, the Al therefore inevitably oxidises to aqueous Al3+. That's how it is...

Mixell - 27-3-2011 at 11:48

Well, thats what I meant...
You separate a part of the solution, insert aluminium- manganese and iron come out.
Insert the manganese/iron powder to the other part of the solution (not the aluminium one) - manganese goes in, iron comes out. And you receive a pure solution of manganese chloride.

Arthur Dent - 28-3-2011 at 04:14

@ blogfast25: You're right about the HCl to clean my flask, I filled it with 32% HCl and will leave it in for a couple of days, because I really want it ultra clean since i'll eventually use this flask in my home brewing and strong spirits distillation apparatus.

Just to make sure, i'll probably do a conc. sulphuric acid bath after that to be on the safe side. And much of my "distillation" glassware will also be treated in such manner. I do have some brand new glassware that has never been exposed to anything more than distilled water so that will also be used.

As for the Manganese solution... i'll put it on a shelf till I find some time to tinker with it.

Thanks guys!
Robert


blogfast25 - 28-3-2011 at 08:58

Robert: good!

Mixell: Al also plates out Mn2+:

Mn2+(aq) + 2e === > Mn(s) … -1.185 V (red)
Al(0) === > Al3+(aq) ... +1.662 V (ox)

Cell pot. = Ered + Eox = - 1.185 + 1.662 > 0, ergo following Nernst: ΔG < 0, reaction proceeds.

Mixell - 28-3-2011 at 09:30

Again, that is what I mean...
In the second vessel AL plates out Fe and Mn, then the Fe,Mn mix is added to the first vessel, where the Mn plates out Fe.
I hope thats clear enough.

blogfast25 - 28-3-2011 at 11:27

I see. I wouldn't put my noney on that working though...

Mixell - 31-3-2011 at 09:54

Well, I could try it, if I'll find the time and nerve to handle the awful smell of chlorine, I see no reason for why it would not work, but who knows...

elementcollector1 - 23-3-2012 at 20:41

Wait, slow down. Are you saying aluminum metal literally plates out manganese metal and iron metal? Or a compound of each? If it does the metal, can I melt it to a good-sized lump without oxygen?

blogfast25 - 24-3-2012 at 06:42

Quote: Originally posted by elementcollector1  
Wait, slow down. Are you saying aluminum metal literally plates out manganese metal and iron metal? Or a compound of each? If it does the metal, can I melt it to a good-sized lump without oxygen?


The problem is that manganese is actually quite reactive and freshly prepared metal reacts even with cold water (slowly). Actual plating out of Mn from aqueous media is therefore impractical.

elementcollector1 - 24-3-2012 at 10:21

Aw, darn. Is there another way to get manganese metal (short of thermite)?

blogfast25 - 24-3-2012 at 11:24

Quote: Originally posted by elementcollector1  
Aw, darn. Is there another way to get manganese metal (short of thermite)?


There's always the method used by E. Glatzel:

http://www.sciencemadness.org/talk/viewthread.php?tid=10249&...

But I think you won't fancy that much either... ;)

elementcollector1 - 24-3-2012 at 11:28

You're right, that sounds pretty beyond my reach. So, when the manganese reacts with cold water (assume ice water), what are the products?

blogfast25 - 24-3-2012 at 14:00

First it would react to Mn(OH)2, which would then quickly air-oxidise to MnO2.

Like with many other elements, there is no quick'n easy way to home make manganese. But there's always the shops...

If you do want to have a shot at it, a fairly large, slow burning (use coarse materials) MnO2 thermite with lots of CaF2 (to slow 'burn' rate even more) is probably your best bet. But even that make take many attempts to get it really right.


elementcollector1 - 24-3-2012 at 18:24

Durn. My aluminum be powdered too fine for a 'coarse' reaction.
I did make beautiful bubblegum-colored MnCl2 chunks just now, will aluminum reduce those?

For that manganese dioxide thingy, would that be a good way to make a MMO electrode?

blogfast25 - 25-3-2012 at 09:37

Quote: Originally posted by elementcollector1  
Durn. My aluminum be powdered too fine for a 'coarse' reaction.
I did make beautiful bubblegum-colored MnCl2 chunks just now, will aluminum reduce those?

For that manganese dioxide thingy, would that be a good way to make a MMO electrode?


Aluminium is not a very good reducing agent for chlorides because it's own chloride (AlCl3) is only part ionic and thus quite volatile (it sublimes quite easily). This option would only work in a bomb type reactor. Also, your MnCl2 is a hydrate. You need anhydrous MnCl2 for this and it is tricky to dehydrate the hydrate.

Trust me, I think I've more or less thought of/tried most conceivable ways of obtaining good quality Mn metal but w/o real success. I even tried anhydrous MnCl2 + Mg powder, only to find the MgCl2 (the 'slag') obtained is also too volatile and that this reaction would have to be carried out in a bomb, to avoid the slag from evaporating.

Regards MnO2 electrodes, search this forum...

elementcollector1 - 25-3-2012 at 12:57

So, I'll just make the hydroxide and then the dioxide. I heard from this (http://developing-your-web-presence.blogspot.com/2008/07/man...) that a mix of MnO and MnO2 may be suitable for good Mn production, but how do I make MnO? I have MnCO3, so I figure if I put that under mineral oil and heat it to 200C on the stovetop, that should make the stuff I'm looking for, correct?

blogfast25 - 25-3-2012 at 13:18

Quote: Originally posted by elementcollector1  
So, I'll just make the hydroxide and then the dioxide. I heard from this (http://developing-your-web-presence.blogspot.com/2008/07/man...) that a mix of MnO and MnO2 may be suitable for good Mn production, but how do I make MnO? I have MnCO3, so I figure if I put that under mineral oil and heat it to 200C on the stovetop, that should make the stuff I'm looking for, correct?


That blog post is actually mine. But it's also somewhat obsolete (one fine day I'll update it). NOTE the very low yields that I mentioned. For similar successful thermites with other metals oxides, 70 % yield and up, even in my fairly primitive conditions, are normal. With MnOx I've never gotten more than about 30 %.

The trouble with manganese thermites is that the boiling point of manganese is almost the same as the melting point of alumina. THAT's why it's difficult to obtain good Mn metal from thermite: much of the metal boils off! Using different types of Mn oxides does not really change that and further experimentation showed that trying to use MnO or MnO/MnO2 blends doesn't improve things much.

Heating MnCO3 under oil is a recipe for a mess: it will be near impossible to separate the MnO from the oil.

I did make MnO from MnCO3 by heating it in a stream of dry CO2 (nitrogen and argon will of course also work). Like I said it wasn't really worth the effort. Using an MnO/MnO2 blend makes the reaction run a bit cooler but doesn't (obviously!) alleviate the BP/MP problem mentioned.

A fairly large (> 200 g) MnO2 thermite with 20 - 30 % fluorite (CaF2) and using fairly coarse ingredients should at least leave you with some metal but probably not very high quality.

[Edited on 25-3-2012 by blogfast25]

elementcollector1 - 25-3-2012 at 14:32

Quote: Originally posted by blogfast25  
Quote: Originally posted by elementcollector1  
So, I'll just make the hydroxide and then the dioxide. I heard from this (http://developing-your-web-presence.blogspot.com/2008/07/man...) that a mix of MnO and MnO2 may be suitable for good Mn production, but how do I make MnO? I have MnCO3, so I figure if I put that under mineral oil and heat it to 200C on the stovetop, that should make the stuff I'm looking for, correct?


That blog post is actually mine. But it's also somewhat obsolete (one fine day I'll update it). NOTE the very low yields that I mentioned. For similar successful thermites with other metals oxides, 70 % yield and up, even in my fairly primitive conditions, are normal. With MnOx I've never gotten more than about 30 %.

The trouble with manganese thermites is that the boiling point of manganese is almost the same as the melting point of alumina. THAT's why it's difficult to obtain good Mn metal from thermite: much of the metal boils off! Using different types of Mn oxides does not really change that and further experimentation showed that trying to use MnO or MnO/MnO2 blends doesn't improve things much.

Heating MnCO3 under oil is a recipe for a mess: it will be near impossible to separate the MnO from the oil.

I did make MnO from MnCO3 by heating it in a stream of dry CO2 (nitrogen and argon will of course also work). Like I said it wasn't really worth the effort. Using an MnO/MnO2 blend makes the reaction run a bit cooler but doesn't (obviously!) alleviate the BP/MP problem mentioned.

A fairly large (> 200 g) MnO2 thermite with 20 - 30 % fluorite (CaF2) and using fairly coarse ingredients should at least leave you with some metal but probably not very high quality.

[Edited on 25-3-2012 by blogfast25]


Probably one of the most awkward moments I've had in the past few weeks...

So, just pump the thermite full of CaF2? Sounds like a plan.

blogfast25 - 26-3-2012 at 08:54

EC1:

The recommended amount of CaF2 is actually in the post linked to. I think it can probably be increased somewhat but above a certain level of CaF2, your mix won't light or will fizzle.

elementcollector1 - 8-4-2012 at 20:59

Sorry, the 0.27 mol or 0.24 mol?

blogfast25 - 9-4-2012 at 05:13

The amount of CaF2 is determined from the molar ratio of Al to CaF2 which I keep constant at 4.44, regardless of formulation. This way each formulation contains the same molar percentage of CaF2 in its slag.

So to calculate the number of moles of CaF2, divide the amount of moles of Al by 4.44.

For the Mn thermite you can probably lower that ratio a bit, perhaps go down to about 4.00.

LanthanumK - 9-5-2012 at 07:20

I dissolved some battery crud in the past and got a perfectly colorless solution from which beautiful pink manganese(II) chloride crystals precipitated. Just recently I used some more of the same crud and got a clean solution.

The batteries were carbon zinc batteries, probably Sunbeam brand. They can be bought at many discount stores. Here is it at Amazon: http://www.amazon.com/Sunbeam-Super-Heavy-Perfomance-Batteri...

Unfortunately, there is zinc chloride in the manganese dioxide. This makes the manganese chloride crystals super deliquescent because of the ZnCl2.

Whenever I use alkaline batteries I get a dark orange solution upon filtering. Maybe it is better to use these batteries for manganese salt production.

blogfast25 - 9-5-2012 at 08:12

Remove any zinc by washing the crud first with warm dilute acetic acid (vinegar) or dilute sulphuric acid. Simmer gently for an hour or so, then filter and wash. This should eliminate zinc, as well as other possible water soluble contaminants.

elementcollector1 - 9-5-2012 at 20:28

Hey boys, look what I found! Manganese metal electrodeposition from chloride at below-0 Celsius temperatures!
http://www.springerlink.com/content/tu633h877633826l/
Credible or not?

Back on main topic, I don't recommend using a diluted strong acid, that doesn't tend to work well for me. Just go with vinegar.

UPDATE: This was a triumph! Pure manganese plated out thickly onto my iron cathode! (Take that, blogfast25!)

[Edited on 11-5-2012 by elementcollector1]

nora_summers - 12-5-2012 at 10:42

wonderful! do you have any pictures of your manganese metal?

elementcollector1 - 12-5-2012 at 18:17

Sadly, I need to restock on HCl so I can get some more MnCl2, so I only have a little. I'll see if I can take a good pic...

elementcollector1 - 20-5-2012 at 13:04

Sigh... And we're back to square one, unknown foreign contamination. Despite repeated attempts to remove this from solution (MnCO3 bubbled but dissolved, NaOH made white precipitate and only partially worked even at high concentrations), this solution is stubbornly staying that deep, deep orange-red. Any ideas? (PS: It's probably contaminated with sodium ions too. Won't effect electrolysis, but might affect purification.)

blogfast25 - 21-5-2012 at 08:33

Quote: Originally posted by elementcollector1  
Sigh... And we're back to square one, unknown foreign contamination. Despite repeated attempts to remove this from solution (MnCO3 bubbled but dissolved, NaOH made white precipitate and only partially worked even at high concentrations), this solution is stubbornly staying that deep, deep orange-red. Any ideas? (PS: It's probably contaminated with sodium ions too. Won't effect electrolysis, but might affect purification.)


Summarise what you're doing to obtain the MnCl2. I may have an idea what the contamination is.

elementcollector1 - 21-5-2012 at 18:49

1) Dissolved in HCl (impure, hardware store grade). I KNOW this is where the impurity came from, as the MnO2 was painstakingly purified and dehydrated for months beforehand.
2) Obtained very, VERY dark red solution. Fe 3+? Attempted precipitation of hydroxides with first somewhat dilute NaOH, then strongly concentrated. Off-white precipitation formed, then quickly redissolved. (I'm probably going to boil this stuff down outside, to remove the pH issues.)
3) Attempted a few other methods, all to no avail. Solution remains as dark and unyielding of manganese as ever. What do?

blogfast25 - 22-5-2012 at 05:25

Quote: Originally posted by elementcollector1  
1) Dissolved in HCl (impure, hardware store grade). I KNOW this is where the impurity came from, as the MnO2 was painstakingly purified and dehydrated for months beforehand.
2) Obtained very, VERY dark red solution. Fe 3+? Attempted precipitation of hydroxides with first somewhat dilute NaOH, then strongly concentrated. Off-white precipitation formed, then quickly redissolved. (I'm probably going to boil this stuff down outside, to remove the pH issues.)
3) Attempted a few other methods, all to no avail. Solution remains as dark and unyielding of manganese as ever. What do?



Something’s wrong with point 2: if there really is Mn<sup>2+</sup> (and Fe<sup>3+</sup>;) in your solution then a PERMANENT precipitate with any strong alkali MUST form. Either you’re not adding enough alkali (check the pH of the solution after addition) or there something amphoteric in your solution.

The result also indicates that there is no iron (III) in your solution, as that starts precipitating as a fluffy, reddish brown precipitate from pH 4 – 5.

Dissolving purified MnO2 in dodgy HCl is asking for trouble. Either switch to a better grade of HCl (completely clear, non-yellow) or distil what you’ve got to better purity.


[Edited on 22-5-2012 by blogfast25]

elementcollector1 - 22-5-2012 at 15:39

I think I might just not be adding enough alkali. What I'll try to do is distill the liquid to get HCl fumes (which can be led into a beaker of distilled water, solving that problem) and a more saturated solution, and the salt color might help us out some more. FeCl3 is yellow, correct?
Alternatively, it was mentioned somewhere that the color impurity is sometimes due to organics.

blogfast25 - 23-5-2012 at 06:06

Quote: Originally posted by elementcollector1  
I think I might just not be adding enough alkali. What I'll try to do is distill the liquid to get HCl fumes (which can be led into a beaker of distilled water, solving that problem) and a more saturated solution, and the salt color might help us out some more. FeCl3 is yellow, correct?
Alternatively, it was mentioned somewhere that the color impurity is sometimes due to organics.


Yes, not enough alkali is the most likely cause for non/partial precipitaton.

The colour of the ferric ion (Fe<sup>3+</sup>;) depends strongly on concentration. Very dilute it's more or less colourless. More concentrated solutions start picking up colour due to hydrolysis forming species like [Fe(H<sub>2</sub>O)<sub>5</sub>OH]<sup>2+</sup>. Colour then ranges from yellow to amber-red to reddish-brown. Solutions are also slightly thermochromic: at higher temperature the equilibrium point of the hydrolysis shifts to the right and the solution darkens a little.

It's best to test for Fe<sup>3+</sup> with KSCN or NH4SCN, because FeSCN<sup>2+</sup> is a dark red complex ion, or with K4Fe(CN)6 with which you get Prussian Blue (Fe7(CN)18), a deep blue.

Test tentatively for organics by adding peroxide and heat, that tends to destroy the organics.

elementcollector1 - 23-5-2012 at 21:39

Don't have those compounds, and from what I've heard, the test is very sensitive.
Anyway, just threw some more base (MnCO3) in, and the color definitely lightened from a liquid-bromine darkness to a tan-orange, translucent color. Then, I threw some more carbonate in. Will check up tomorrow to see if it's that beautiful clear / rose-pink. (If so, I'm just going to mix my entire pound of manganese carbonate with my 1000 mLs of solution. Should work, right?

Arthur Dent - 24-5-2012 at 04:32

I recently crystallized another batch of Manganese Chloride, that one made with technical grade HCl and pottery store Manganese Carbonate... I noticed that the crystals are a lighter shade of pink than the original batch I made out of purified and thoroughly cleaned carbon/zinc battery crud.

The new batch looks exactly like the crystals you see if you google "Manganese Chloride" and click on images. My original batch was a definitely more intense shade of pink, with hints of magenta.

Here's a picture of batch one:



Maybe its the hydration level of the crystals, or maybe it's some inpurities that I couldn't separate from the original batch, but I think batch 2 is better. So that pottery store Manganese Carbonate wasn't that bad after all.

PS: the solution for batch 2 initially was brownish yellow, but after boiling it down to less than 1/4, went from yellow, to straw to nearly water clear. I left it to crystallize in the dessicator for 2 months.

Robert

[Edited on 24-5-2012 by Arthur Dent]

blogfast25 - 24-5-2012 at 04:58

Quote: Originally posted by Arthur Dent  

Maybe its the hydration level of the crystals, or maybe it's some inpurities that I couldn't separate from the original batch, but I think batch 2 is better. So that pottery store Manganese Carbonate wasn't that bad after all.



Robert

[Edited on 24-5-2012 by Arthur Dent]


Quality [of the MnCO3] will depend from one product to another. Mine definitely contains some soluble iron (III) but it's bearable.

Remember that FeCl3 is very highly soluble, so chances are very good that it will stay in the mother liquor and not crystallise, as long as you crystallise from acid solution.

Nice crystals, Robert.

[Edited on 24-5-2012 by blogfast25]

Arthur Dent - 24-5-2012 at 05:21

Quote: Originally posted by blogfast25  

Nice crystals, Robert.


Thanks! Yeah that first batch was very nice, but the latest batch looks much better and yielded a solid block of very light pink crystals. The remaining liquor was a bit yellowish and quite acidic.

I'll try to crush and dessicate the crystals further to obtain the anhydrous pink dust. I'll try to add a pic when time allows. :D

Robert

blogfast25 - 24-5-2012 at 07:33

Quote: Originally posted by Arthur Dent  


I'll try to crush and dessicate the crystals further to obtain the anhydrous pink dust. I'll try to add a pic when time allows. :D

Robert


Oh but that won't work, at least not if you mean 'anhydrous MnCl2' and not 'dry MnCl2 hydrate crystals'.

Anhydrous MnCl2 requires heating in a stream of dry HCl or calcining a mixture of MnCl2 hydrate and NH4Cl in a stream of a dry inert gas. Anything else basically yields MnO2: dry but not what you want.

If dry crystals of the hydrate is what you want then just air drying should be enough.

[Edited on 24-5-2012 by blogfast25]

Arthur Dent - 24-5-2012 at 09:40

Ugh, you're right... it's the tetrahydrate that are the common crystals we harvest, right? So if I heat it above 60/70 C, I should be able to obtain the dihydrate, correct?

Robert

blogfast25 - 24-5-2012 at 11:10

Quote: Originally posted by Arthur Dent  
Ugh, you're right... it's the tetrahydrate that are the common crystals we harvest, right? So if I heat it above 60/70 C, I should be able to obtain the dihydrate, correct?

Robert


I'm not sure. Partial hydrolysis and/or oxidation are risks if you don't at least exclude oxygen. Why would you want the dihydrate?

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