Sciencemadness Discussion Board - Manganese Chloride Crystals

Manganese Chloride Crystals

Pages: 1 2

elementcollector1 - 24-5-2012 at 15:48

Reporting in; after using copious amounts of MnCO3 on a small amount of mother liquor, I have obtained a pink solution and excess carbonate (though the solution still smells like it's reacting). This is a more colored pink than the stuff I had last time, so I'm assuming either this is a very saturated solution or that there are still impurities in solution. Although that wouldn't explain such an excess of carbonate...

EDIT: Argh, how do I get this excess carbonate out? I threw four filters at it, and they're barely holding back half of the tan goop.

[Edited on 25-5-2012 by elementcollector1]

Arthur Dent - 25-5-2012 at 04:08

Quote: Originally posted by elementcollector1

how do I get this excess carbonate out? I threw four filters at it, and they're barely holding back half of the tan goop.

Yeah I know, coffee paper filters aren't the best in that situation. Ideally, I would suggest a Buchner funnel, but if you don't have one, perhaps some real lab paper filters would be more efficient, like some fine porosity Whatman paper filters, but they take an eternity if you use gravity filtration. Your best bet would be a Buchner funnel with vacuum.

However, using a Buchner funnel for heavy metals filtering means it should no longer be used for anything else but that. So if you use a Buchner funnel for food grade stuff like vegetal extracts or brewing/distilling, use it exclusively for that.

Robert

blogfast25 - 25-5-2012 at 05:10

Quote: Originally posted by elementcollector1

EDIT: Argh, how do I get this excess carbonate out? I threw four filters at it, and they're barely holding back half of the tan goop.

[Edited on 25-5-2012 by elementcollector1]

Don't 'throw filters at it'. Even with coffee filters your filtrate should eventually run clear if you just keep pouring it over the SAME filter over and over, until must pores are more or less clogged.

Alternatively, simmer the slurry gently for a bit: it will increase the grain size of the carbonate particles.

If all else fails, allow to stand for a couple of days until liquor is perfectly clear, then carefully decant it off.

elementcollector1 - 25-5-2012 at 17:14

Now it's much more pink, but still pretty colored (will filter a few more times).
This is working, though: I have a small test run of pure pink MnCl2 sitting next to me. It's nearly colorless, so no iron should be present (at least not in dangerous amounts).
Excellent!

elementcollector1 - 29-5-2012 at 09:51

Boiled it down and got a green solution. I'm assuming this is Fe 2+, so how did that escape the original neutralization with MnCO3?

blogfast25 - 29-5-2012 at 10:26

Quote: Originally posted by elementcollector1

Boiled it down and got a green solution. I'm assuming this is Fe 2+, so how did that escape the original neutralization with MnCO3?

Add some peroxide to a small sample. Does it fizz and go brown/red?

elementcollector1 - 29-5-2012 at 11:27

Not yet, but upon filtering through that carbonate-clogged filter it goes right back to pink (though I'm skeptical of that color now). I added peroxide to the unconcentrated, pink stuff with no effect; we'll see about the green variety.

blogfast25 - 29-5-2012 at 12:05

Quote: Originally posted by elementcollector1

Not yet, but upon filtering through that carbonate-clogged filter it goes right back to pink (though I'm skeptical of that color now). I added peroxide to the unconcentrated, pink stuff with no effect; we'll see about the green variety.

Photos would be great.

elementcollector1 - 20-6-2012 at 21:20

Quote: Originally posted by blogfast25

Quote: Originally posted by elementcollector1

Photos would be great.

Yes! Finally uploaded. Anyway, the first picture is that of the apparent green 'impurity', seconds before boiling down to the nice clean pic of the powdered salt of the second.

blogfast25 - 21-6-2012 at 05:05

Very nice end-product, EC1. A bit of a mystery that green...

elementcollector1 - 30-11-2012 at 23:01

I'm having trouble with my manganese stock again!
Oddly enough, my original solution looks pink in a translucent container, but yellow in a clear one.
I tried the NaOH method, and I'm waiting for results, but it seems that there's no iron contamination (precipitate had no red in it).
I started from old MnCO3 that I had leftover from some of my original purifications, which dissolved extraordinarily slowly (unusual, considering it was an acid-base reaction).
Is there hope for this solution?

IrC - 1-12-2012 at 00:10

Quote: Originally posted by 12AX7

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

I know this post was long ago but new to me. I have read stories of arsine causing harm from impure Zinc in acid. How much worry should we place on this possibility considering unknown purity Zn dust, i.e., typical fleabay purchase with buyer beware on quality?

12AX7 - 1-12-2012 at 09:47

Hmm, I wouldn't think it much of a problem, zinc is most often distilled, achieving purity comparable to electrolytic metals (for instance, copper, lead, silver, etc.). That said, arsenic is a volatile element itself. I haven't heard anything about it though. You'd certainly know by the smell! Last time I was dissolving zinc, it didn't really smell like anything, which is an excellent sign there is very little of any calchophile in it: sulfur, phosphorus, selenium, antimony, arsenic.

Tim

IrC - 1-12-2012 at 13:50

Good to know. I like to collect those turn of the century recipe books (on processes, etc.) and read a couple stories of deaths caused by dissolving ZN in acid. The coroner did tests looking for a deposition of a mirror surface and was positive in a few cases. I have been dissolving Zn for various reasons for decades and never had problems nor noticed any garlic like odor. Yet reading some of these old stories got me to thinking I should be less cavalier experimenting once in a while. I had assumed modern processes produced less of these problems. I do know it was a real danger back then depending upon where the starting materials were mined. At least that was the conclusion in one of the stories, that one about a doctor in England doing himself in playing around in his home lab. Was thinking about making some Zinc chloride from Zn dust and thought I should check into the subject first. Mainly because the Zn came from fleabay and one never knows where things originated when you buy from random private individuals sight unseen.

elementcollector1 - 1-12-2012 at 17:55

Guessing there's no hope for this solution, as the solution hasn't turned back to pink with the hydroxide method, the hydroxide precipitate is tan, perfectly dry (precipitated MnO2 is never dry, it's always a 'mud' of sorts) and doesn't change to black (the first precipitate did, oddly enough). Scrapping!

elementcollector1 - 3-1-2013 at 08:19

Gave this another shot, and I now find that my solution is an extremely dark purple. The last time I dissolved crude MnO2 into HCl, it was an extremely dark red. What on earth is going on here?
Pics will follow, if I can get them.

blogfast25 - 3-1-2013 at 13:26

Quote: Originally posted by elementcollector1

Gave this another shot, and I now find that my solution is an extremely dark purple. The last time I dissolved crude MnO2 into HCl, it was an extremely dark red. What on earth is going on here?
Pics will follow, if I can get them.

Pictures and more details about what you used and how you did it would be useful.

elementcollector1 - 3-1-2013 at 16:07

Well, here it is. The solution is so dark that the purple is only barely visible at the edge.

blogfast25 - 5-1-2013 at 06:43

EC1:

That colour is eerily reminicent of either KMnO4 or Mn(III) (not II) salts. Please describe in some detail what it is that lead to this colour. Even concentrated Mn(II) are much lighter.

Try also to add some concentrated HCl to a sample of the solution: both permanganate and Mn(III) oxidise chlorides to elemental chlorine. The solution would then clear and stink of chlorine.

elementcollector1 - 7-1-2013 at 10:02

As far as I can remember, I reacted raw MnO2 with contaminated HCl (Fe(II) contaminant), hoping to get a starting chemical for some MnCl2 or pure MnO2. (However, I recently got manganese metal, so I no longer have the need). I really don't think that this is permanganate, but I didn't know Mn(III) had this color. Please explain.

I guess I could try adding more HCl, just in case I really messed up on memory and was using sulfuric acid or some such instead... I'll just lean over and take a deep whiff of the solution, see if I can detect any chlorine.

blogfast25 - 7-1-2013 at 10:37

Quote: Originally posted by elementcollector1

As far as I can remember, I reacted raw MnO2 with contaminated HCl (Fe(II) contaminant), hoping to get a starting chemical for some MnCl2 or pure MnO2. (However, I recently got manganese metal, so I no longer have the need). I really don't think that this is permanganate, but I didn't know Mn(III) had this color. Please explain.

I guess I could try adding more HCl, just in case I really messed up on memory and was using sulfuric acid or some such instead... I'll just lean over and take a deep whiff of the solution, see if I can detect any chlorine.

Hmmm... with MnO2 and strong H2SO4, Mn2(SO4)3(aq) (Mn(III)) is what you get. Depending on concentration, the hydrated Mn3+ cation is a deep, wine red, somewhat similar to MnO4-.

Take a sample of that solution and add concentrated NaCl solution to it. Mn3+ immediately oxidises the chloride ions to Cl2:

Mn3+(aq) + Cl-(aq) === > Mn2+(aq) + 1/2 Cl2(g)

Thus you should observe, if it is Mn3+ that's causing the colour:

1. bubbles and smell of chlorine
2. solution clears to almost water colour
3. heat dissipation

elementcollector1 - 8-1-2013 at 07:42

Can that be replicated with HCl somehow? Maybe I have a solution of MnCl3 somehow...
You're right, I did some quick research and this solution, as well as the solution I made last summer, both bear a striking resemblance to Mn(III) in solution.

EDIT: Nevermind, calcium chloride seems to be a good distinguishing test between chlorides and sulfates. I'll use that.

[Edited on 8-1-2013 by elementcollector1]

blogfast25 - 8-1-2013 at 09:53

Quote: Originally posted by elementcollector1

Can that be replicated with HCl somehow? Maybe I have a solution of MnCl3 somehow...
You're right, I did some quick research and this solution, as well as the solution I made last summer, both bear a striking resemblance to Mn(III) in solution.

EDIT: Nevermind, calcium chloride seems to be a good distinguishing test between chlorides and sulfates. I'll use that.

[Edited on 8-1-2013 by elementcollector1]

No, no. No CaCl2: if there's sulphate in there then you'll get insoluble CaSO4! It's ruin everything. Unless you just want to test for sulpahtes, of course (BaCl2 is better for that)...

NaCl or HCl will do for the test. If you don't get immediate response, warm gently...

[Edited on 8-1-2013 by blogfast25]

elementcollector1 - 8-1-2013 at 22:22

I'm testing to see if I used sulfuric acid or hydrochloric. So, an insoluble precipitate of CaSO4 is exactly what I'm after.

I'll also test for chlorine, but I don't have too high hopes. If this was HCl, it's likely going to smell of chlorine anyway (due to the chlorine produced in the original reaction).

blogfast25 - 9-1-2013 at 09:03

Well, does it smell fo chlorine?

elementcollector1 - 10-1-2013 at 22:49

Yes. Probably was HCl.

blogfast25 - 11-1-2013 at 13:24

Quote: Originally posted by elementcollector1

Yes. Probably was HCl.

Test for sulphates to be sure?

cleaning with crude chromatography

mayko - 18-1-2013 at 15:46

I just ran this myself - lots of fun

I took care of chlorine by letting gas bubble through a test tube containing a stiff hydroxide solution and a bit of zinc, so that the chlorine got disproportionated into hypochlorite, and then decayed by the zinc. I also dropped in a bit of cabbage juice indicator to monitor for bleaching or pH change should either reactant run low. This was still not ideal (I still had a flaskful of chlorine at the end, and there would have been Problems if the setup had fallen over or something)

I had a bit of iron contamination; I was able to remove a lot of it by dripping acetone on the filtered crystals; the soluble iron salt was wicked away by the filter. I probably could have cleaned the whole pile if I'd stirred it around some, but as it was my first wash only cleaned off the surface of the pile.

Ultimately I scraped it into a flask and made a proper acetone wash - worked great!

elementcollector1 - 18-1-2013 at 15:49

It looks a little whitish to be pure. Pure MnCl2 is bubblegum-pink.

mayko - 18-1-2013 at 16:04

Quote: Originally posted by elementcollector1

It looks a little whitish to be pure. Pure MnCl2 is bubblegum-pink.

I agree; I thought it was a little pale, though there is a light pink color. What sort of impurity would remove the color?

elementcollector1 - 18-1-2013 at 20:18

I would hazard a guess at sodium, although I'd need to know more about how you made this MnCl2 to be sure.
Acetone was a good idea for iron, I never thought of that. Are you sure MnCl2 is insoluble in it?

blogfast25 - 19-1-2013 at 09:13

If you're going to do a solid/liquid extraction of FeCl3 with acetone you need to grind your starting material very finely.

mayko - 22-1-2013 at 23:01

Quote: Originally posted by elementcollector1

I would hazard a guess at sodium, although I'd need to know more about how you made this MnCl2 to be sure.
Acetone was a good idea for iron, I never thought of that. Are you sure MnCl2 is insoluble in it?

Flame test is negative for sodium, but seems distinctive. In an alcohol flame, there were small golden sparks, like you see when you burn iron filings, except that they were more glowwy and less sparkly. Once that died down, there was a faint bluish-green glow.

I made it MnO2 + HCl; the MnO2 is battery grade. I washed it, but might not have done a great job; I'm going to try again with a really clean batch.

My Merck described MnCl2 as 'soluble' in acetone, whereas FeCl3 was 'very soluble', so I chilled it and went cold; visually it seemed to leave what pink was there while quickly drawing out the yellow.

blogfast25 - 23-1-2013 at 08:54

To obtain iron free manganese salts there is a procedure described by ‘peach’ and also by nurdrage. In my own words (I tested it and it works very well):

Take about one quarter of the solution [Fe(III) contaminated Mn (II) solution, preferably only slightly acidic, pH >= 3] and set the rest (the stock) aside. Add enough of an alkali (NaOH, KOH or ammonia solution are all good but not Na2CO3) to precipitate all the manganese and iron as hydroxides. Filter this and wash the filter cake of manganese and iron hydroxides with plentiful small aliquots of clean water until the filtrate is almost neutral. This washes out the soluble cations.

Now scoop out most of the precipitated hydroxides on the filter, add them to the stock solution and leave this to stand overnight (do not discard the filter, instead cover it with cling film to keep it moist). During standing overnight any contaminating iron will be converted to highly insoluble iron (III) oxide. Now filter the stock plus precipitates, using the same filter used before. The obtained solution is now essentially free of any iron and can be used to re-precipitate the manganese as purified MnCO3 or to crystallise it as quite pure Mn(II) salt.

The method relies on the extreme insolubility of Fe(OH)3 in neutral conditions and uses the precipitate as a buffering agent to reach optimal pH at which the Mn(II) remains in solution and the Fe(III) precipitates 100.0 %.

[Edited on 23-1-2013 by blogfast25]

elementcollector1 - 23-1-2013 at 09:44

I've tried that method before, with about 80% success rate (there was that one time...)
It bears mentioning that you have to work fairly fast with the precipitated sludge, as the Mn(OH)2 formed will quickly oxidise in air to the brown-black MnO2 hydrate (which is significantly less reactive).

blogfast25 - 24-1-2013 at 05:58

The trick is to neutralise quite gently: the air oxidation of Mn(OH)2 is greatly accelerated in alkaline conditions:

Mn(OH)2 === > MnO2 +2 H+ + 2e
1/2 [O2 + 2 H+ + 4e === > 2 OH- ]

---------------------------------------------

Mn(OH)2 + 1/2 O2 === > MnO2 + H2O

But some loss of Mn(II) to MnO2 is probably unavoidable...

elementcollector1 - 24-1-2013 at 09:42

Neutralized the apparent Mn(III) solution yesterday to get the usual brown sluge. No apparent Fe contamination is visible, but I will have to check later. Probably going to make this into manganese sulfate for production of manganese dioxide electrodes.

Seriously, what

elementcollector1 - 14-6-2013 at 17:11

Today I took some pure potassium permanganate and drain-opener sulfuric acid, in an attempt to make some manganese sulfate solution for plating. I didn't weigh stuff out because I assumed that this would be a standard aqueous reduction, and would just turn pink when finished. However, it's been acting even more weird than my usual manganese solutions. At first, nothing really happened between the diluted sulfuric acid and the KMnO4. Then I added some alcohol, and this happened:

Opaque pink. Okay... maybe wait this one out?

Well, brown's kind of expected - that's obviously MnO2. Filtered this out. (Pic of alcohol and sulfuric acid too.)

...Orange?!

I think the intermediates may have been Mn(3+) and Mn(4+), but orange is ridiculous - what manganese compound is orange?

I was just trying to make manganese sulfate...

platedish29 - 14-6-2013 at 17:19

That may be over fu***d carbon dissolved in, try applying activated charcoal to purify the thing, if that doesn't work up I dunno I'm just to lazy to accomplish an further predictive analysis lol

elementcollector1 - 14-6-2013 at 19:26

Turns out it was something like that: A fine, orange-brown precipitate was visibly settled out an hour later, and the flask appeared to be water-clear. I have two solutions of what is presumably manganese sulfate (although it's hard to tell), and upon mixing and filtering, more particulate was observed. Not sure what this was, but at least I have a clear solution - one step forward towards having pink manganese sulfate solution! I thought I saw a pink color when I tilted the beaker a certain way, but it could easily have been a trick of the light. Will continue this stuff tomorrow.

Texium - 4-5-2014 at 12:03

Well, I hope nobody minds me bringing back this old thread.
I've been attempting to make a soluble manganese salt from manganese dioxide (from a battery) to use for plating out manganese metal, like elementcollector1. I was wondering whether manganese sulfate or chloride would be more practical to use for that. I attempted to make manganese chloride earlier using concentrated HCl on low heat on my hot plate. I don't yet have any sulfuric acid, but should be able to obtain some soon if it is needed.

elementcollector1 - 4-5-2014 at 14:33

Manganese chloride, cathode of your choice, anode of lead/tin solder was what worked for me. You might have to use a salt bridge or other means of separating the anolyte/catholyte MnO₂.

Texium - 4-5-2014 at 15:13

Alright, thanks. I'll try that if my MnO₂ + HCl ever finishes reacting…

How long should it take to react anyway? It's been going for a little over three hours, still bubbling away. I don't have it on heat, but the ambient temperature outside is pretty hot so I thought that it would work fairly well.

[Edited on 5-5-2014 by zts16]

elementcollector1 - 4-5-2014 at 18:01

Mine took quite a while - perhaps a day to finish. Give it some time, and above all, keep it out of the way of anything else - chlorine is very insidious.

Texium - 4-5-2014 at 18:18

Ok then. Right now I'm just trying a test tube sized quantity and I'm timing it to see about how long it will take. I've had it going for about six hours now, outside. No chance of rain or other weather, so it should be fine. I'll check again tomorrow morning before school if I have time!

Texium - 5-5-2014 at 15:29

Sorry if this is a silly question, but I was wondering (because I've never successfully done any electrochemistry) if a normal 9 volt battery would be sufficient for electrowinning the manganese. I realize that it would be much better to use a non-battery power source, but I don't have the confidence or electronics skills to set up something like that.

filter

bcp1211 - 2-10-2015 at 04:46

Dear all,
it's years after but maybe it's still interesting. I was also surprised to see the solution of MnCl2 to be orange-ish. Immediately after preparing it I filtered (0.22µm), and now it's completely colourless. So I guess the rest was just impurities, as it was a 97%purity MnCl2 flakes bottle.
Cheers

elementcollector1 - 2-10-2015 at 16:32

A quick note on electrowinning Mn: The conditions are extraordinarily finnicky. A 9V battery is not recommended, due to overvoltage of the cell (usually only about ~5v is required to start the half-reactions). It is also recommended to add a small amount of ammonium chloride to the catholyte (this I read years ago from various electroplating sources, though I've never tried it). A salt bridge is required to separate the cells.

Something I'd really like to try is plating a thick slab of Mn onto a piece of copper, then using myst32YT's trick to remove copper without harming more reactive metals to get a slab of pure Mn.

A project for another day...

JJay - 4-1-2016 at 12:28

I'd like to make some manganese chloride... do I understand correctly that all I have to do aside from working up the reaction is reduce manganese dioxide with hydrochloric acid and that this reaction occurs spontaneously?

Praxichys - 4-1-2016 at 13:02

Yes, but be aware that this produces copious amounts of chlorine gas, something like 70g for every 198g of the MnCl2 tetrahydrate. That's a little over 22 liters of gas - enough to need to evacuate a good-sized house. Try to use a scrubber, or at the very least you must do this outside.

MnO2 + 4 HCl → MnCl2 + Cl2 + 2 H2O

JJay - 4-1-2016 at 13:04

I'm sure I can put all that chlorine gas to good use for something....

morsagh - 7-1-2016 at 08:03

Your problem with color can be caused too by complexes of Mn(III) in cold conditions it is favoured product, they are mostly red to brown...

Pages: 1 2