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

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

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

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

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]

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?

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

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

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]

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.

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]

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.

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.

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]

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.

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

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]

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!

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

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]

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!

[Edited on 8-1-2011 by peach]

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]

Quote: Originally posted by peach

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.

[Edited on 27-1-2011 by peach]

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

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

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.

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 MnO₂ 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 MnO₂ from a better source. I learn new thangs everyday!

Robert

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

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 MnO₂ 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!

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 H₂O₂, 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 H₂O₂ 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

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?

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...

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?

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?

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]

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.)

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 Mn²⁺ (and Fe³⁺ 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]

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 (Fe³⁺ 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(H₂O)₅OH]²⁺. 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 Fe³⁺ with KSCN or NH4SCN, because FeSCN²⁺ 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.

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]

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]

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.

Robert

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. Robert

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