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Arthur Dent
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@ blogfast25 : Yes sir this is the accurate color, they really look like hard candy! This morning however, I noticed that the edges of the beautiful
crystals above have lost their shine and have started to turn a very light opaque pink, which is the color of anhydrous MnCl<sub>2</sub>,
so it reacts a bit like Copper Sulphate crystals, who slowly turns white and anhydrous if left in a dry atmosphere.
It's very dry in my house, so I can imagine the manganese Chloride crystals have lost some of their hydration water. They do smell a bit less stingy.
The crystals have grown from a 150 ml solution down to 10 ml of liquid left when I pulled them out of the dessicator. Took a few months (3 or 4) in
the dessicator at low vacuum (0.9 atm).
Robert
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Random
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blogfast25
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@ Roger:
Next time perhaps boil in your soluion until boiling behaviour starts to change. On cooling and icing you should get a first crop of decent crystals
of MnCl2.4H2O (depending on solubility - temp. function), the rest can be obtained in dessicator.
Watch for a U2U.
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Chemistry Alchemist
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when MnCl2 Crystals form, what shape do grow?
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Arthur Dent
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Manganese Chloride forms crystals with rhombohedral symmetry, much like Cadmium Chloride. A typical crystal looks like a cube skewed at a 45 degree
angle.
--- Art is making something out of nothing and selling it. - Frank Zappa ---
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blogfast25
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The shape on the photos, LOL. Seriously, look up in Wiki about various crystalline structures. There should be pics and diagrams to your heart's
content if you drill down a bit...
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sxl168
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I was removing some small amounts of dissolved Iron today from MnCl2 prepared from alkaline batteries and I noticed this batch precipitated something
else along with the Iron. The resulting solids are a moderate grey color and looked like a dirty off white when it first precipitated. My solution is
a little acidic yet, so I wouldn't think any Mn had precipitated. Some of it is probably colloidal carbon that made it through the filter paper, but
not the quantities observed.
I was wondering what additives are added to alkaline batteries if anyone knows, which this stuff might be composed of. The quantity that I have
obtained so far is quite small <100 mg. Do they use Cd or Pb? Those additives would probably precipitate out in the conditions I have used. Could
it also be Hg as I have used a few rather old cells (circa late 1980's) in this batch? I was wondering if anyone else has seen this in their
precipitation reactions.
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woelen
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Thread Pruned
17-11-2011 at 05:27 |
blogfast25
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Quote: Originally posted by sxl168  | | I was removing some small amounts of dissolved Iron today from MnCl2 prepared from alkaline batteries and I noticed this batch precipitated something
else along with the Iron. |
Clues would lie in how exactly you removed the iron...
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Chemistry Alchemist
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So really the best way to remove the iron would be from the method NurdRage made a video about?
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blogfast25
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IMHO, yes.
See for instance, here:
http://www.sciencemadness.org/talk/viewthread.php?tid=17931#...
[Edited on 17-11-2011 by blogfast25]
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sxl168
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I use HCl as my acid as it's the only acid I can get ahold of here dirt cheap, the following may behave differently with other acids. To remove the
Iron, I oxidize it from Ferrous to Ferric with bleach or H2O2, doing this while the leaching solution is still highly acidic. Then I adjust the pH to
about 5 with baking soda or Sodium Carbonate. I do not have a pH meter, I use my nose and 5 is just a best guess. The solution will only have a faint
acidic smell to it and no alkaline smell. Zn and Mn will not drop out yet if done properly. While doing the pH adjustment, I have the solution heated
so that it is at least 50C. Ferric drops out readily when heated (but really friggin' slowly at room temp, if at all). It might take 2 tries at the
process if the solution is heavily contaminated with Iron. I get nice faint pink solutions once its done with no hints of Ferric colorization and no
additional colorization upon standing if you use enough oxidizer. I've leached whole AAAA 9V cells, casings and all and was able to extract clean
Mn/Zn with this method.
The hardest part IMO, is oxidizing the Iron but not the Mn. This is where having the acidic leaching solution works well. If MnO2 is formed while
adding oxidizer, which it usually does, it will redissolve with stirring and oxidize the Iron. I usually know the Iron is finished being oxidized if I
get a persistent bleach smell from the solution and MnO2 that won't redissolve. Just takes a few practice runs to get it right, IMO.
I saw NurdRage's video, but I don't have chemicals for his method. I thought everyone knew about method listed above as it's listed in many patents
and battery recycling papers I have come across. It's a somewhat slow process, but works.
This other stuff that drops out seems to come when adjusting the pH higher to kick out the last remaining bits of Iron, and just before Zn/Mn start
precipitating.
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blogfast25
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Yes, your method is a variant of 'Nurdrage's': it relies on selective precipitation of Fe3+ at pH =<5, but not Mn, Zn and some others.
Perhaps there's another variant in there. Add copious bleach to the battery gunge, this will oxidise all to Fe (III) and Mn (IV). Wash superficially
to get rid of the bleach. Dissolve washed gunge in HCl: MnCl2, FeCl3, ZnCl2 etc are leached out. Then treat the leachate like you do. That would
eliminate the tricky bit of oxidising Fe (II) but not Mn (II).
Nurdrage's method has one small advantage: once the pH is right and you filter, the filtrate is free of cations like Na+, provided the precipitate
that was used to 'buffer' the leachate had been washed properly.
You say you don't have the chemicals but for alkali you could use household ammonia, instead of NaOH. 'Washing soda' (Na2CO3) would also do.
[Edited on 17-11-2011 by blogfast25]
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sxl168
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My bad, I was thinking of his video where he extracts the Mn from battery paste via burning sulfur. I forgot about his other video for purification.
Yes indeed that method is the one that I use as a reference, but I just found it a bit lacking when I was dissolving whole cells with the casing (I
only dissolve whole cells as they exist in 9 Volt packs, other size cells I do material separation). That results in a lot of iron contamination and
it takes forever for it all to oxidize by exposure to air, which is why I was doing the oxidizer route. As for removing the alkali's, I precipitate
everything out as carbonates once the iron has been removed and wash thoroughly. I keep the iron precipitate for use later as Ferric Chloride.
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jsc
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That pic is pure chemical porn. 
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So its been epically hot recently and the solution which was some what hard to crystallize has now crystallized and i have nice long crystals of my
guess Manganese(II) Chloride, so would i just poor off the rest of the solution and then wash with really cold water (~0 degrees C) or would the
Manganese dissolve in the water even at that temperature? is there a solvent that can wash the remaining iron out and leave the manganese untouched?
did you want me to post pics of the crystals?
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blogfast25
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Wash with small amounts of iced water or with clean alcohol or acetone. Acetone would dissolve FeCl3 but only if the material is finely crushed.
Pic(s) please!
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i tried washing with a few crystals but it kinda looks like the Iron(III) Chloride has fused in with the crystals... and the crystals were dissolving
too fast to clean... the crystals are needle like as people said they should be if they are Manganese(II) Chloride
They are a bit damp still, but you can still see some of the needle crystal structure off the the left side of it, the colour looks like it has a lot
of iron in it
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blogfast25
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Actually they 'look' fairly pure to me. Did you crystallise them from a highly acidic solution or from a fairly neutral one? FeCl3 doesn't survive a
pH close to 7, or even 5 to 6: it hydrolyses and precipitates. If FeCl3 is present it is locked into the crustalline structure of the MnCl2.
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i started to slowly evaporate it straight away from the reaction of MnO2 and HCl after filtering... so fairly acetic but wouldn't the HCl evaporate
leaving neutral solution...? fairly puree MnCl2?
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blogfast25
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No. HCl forms an azeotrope with water at about 20 % HCl. No matter what strength HCl you started off from, by evaporation you always end up with
roughly 20 %. I tell a small lie: if you started from very weak HCl you'd probably not reach azeotropic composition.
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I geared from somewhere because it's just HCl gas dissolved in water, the whole thing will just evaporate to nothing... So do I have reasonably pure
MnCl2?
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blogfast25
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| Quote: Originally posted by Chemistry Alchemist | | I geared from somewhere because it's just HCl gas dissolved in water, the whole thing will just evaporate to nothing... So do I have reasonably pure
MnCl2? |
Yes because the H2O/HCl azeotrope completely eveporates, of course.
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