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Author: Subject: Manganese Chloride Crystals
elementcollector1
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[*] posted on 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]




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[*] posted on 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.

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[*] posted on 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.




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[*] posted on 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!




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[*] posted on 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?



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[*] posted on 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?




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[*] posted on 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.



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[*] posted on 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.




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[*] posted on 20-6-2012 at 21:20


Quote: Originally posted by blogfast25  
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.


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.

P1040948.JPG - 38kBP1040955.JPG - 31kB




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[*] posted on 21-6-2012 at 05:05


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



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[*] posted on 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?




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[*] posted on 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. :o

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?




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[*] posted on 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.

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[*] posted on 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.





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[*] posted on 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!



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[*] posted on 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.




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[*] posted on 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.




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[*] posted on 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.




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[*] posted on 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.




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[*] posted on 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. :D




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[*] posted on 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. :D


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






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[*] posted on 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]




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[*] posted on 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]




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[*] posted on 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. :P
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).




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[*] posted on 9-1-2013 at 09:03


Well, does it smell fo chlorine?



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