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Author: Subject: Manganese Chloride Crystals
elementcollector1
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[*] posted on 10-1-2013 at 22:49


Yes. Probably was HCl.



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[*] posted on 11-1-2013 at 13:24


Quote: Originally posted by elementcollector1  
Yes. Probably was HCl.


Test for sulphates to be sure?




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[*] posted on 18-1-2013 at 15:46
cleaning with crude chromatography


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!





MnCl2-1.jpg - 139kB MnCl2-2.jpg - 123kB MnCl2-3.jpg - 113kB
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elementcollector1
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[*] posted on 18-1-2013 at 15:49


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



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[*] posted on 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?
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elementcollector1
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[*] posted on 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?




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[*] posted on 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.



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[*] posted on 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.
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[*] posted on 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]




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[*] posted on 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).




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[*] posted on 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...




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[*] posted on 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.



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[*] posted on 14-6-2013 at 17:11
Seriously, what


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...




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[*] posted on 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
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[*] posted on 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.



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[*] posted on 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.
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[*] posted on 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.



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[*] posted on 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]
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[*] posted on 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.



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[*] posted on 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!
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[*] posted on 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.
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[*] posted on 2-10-2015 at 04:46
filter


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
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[*] posted on 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...




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[*] posted on 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?
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[*] posted on 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




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