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blogfast25
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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? |
Aluminium is not a very good reducing agent for chlorides because it's own chloride (AlCl3) is only part ionic and thus quite volatile (it sublimes
quite easily). This option would only work in a bomb type reactor. Also, your MnCl2 is a hydrate. You need anhydrous MnCl2 for this and it is tricky
to dehydrate the hydrate.
Trust me, I think I've more or less thought of/tried most conceivable ways of obtaining good quality Mn metal but w/o real success. I even tried
anhydrous MnCl2 + Mg powder, only to find the MgCl2 (the 'slag') obtained is also too volatile and that this reaction would have to be carried out in
a bomb, to avoid the slag from evaporating.
Regards MnO2 electrodes, search this forum...
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elementcollector1
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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?
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blogfast25
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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]
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elementcollector1
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| Quote: Originally posted by blogfast25 |
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] |
Probably one of the most awkward moments I've had in the past few weeks...
So, just pump the thermite full of CaF2? Sounds like a plan.
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blogfast25
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EC1:
The recommended amount of CaF2 is actually in the post linked to. I think it can probably be increased somewhat but above a certain level of CaF2,
your mix won't light or will fizzle.
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elementcollector1
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Sorry, the 0.27 mol or 0.24 mol?
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blogfast25
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The amount of CaF2 is determined from the molar ratio of Al to CaF2 which I keep constant at 4.44, regardless of formulation. This way each
formulation contains the same molar percentage of CaF2 in its slag.
So to calculate the number of moles of CaF2, divide the amount of moles of Al by 4.44.
For the Mn thermite you can probably lower that ratio a bit, perhaps go down to about 4.00.
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LanthanumK
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I dissolved some battery crud in the past and got a perfectly colorless solution from which beautiful pink manganese(II) chloride crystals
precipitated. Just recently I used some more of the same crud and got a clean solution.
The batteries were carbon zinc batteries, probably Sunbeam brand. They can be bought at many discount stores. Here is it at Amazon: http://www.amazon.com/Sunbeam-Super-Heavy-Perfomance-Batteri...
Unfortunately, there is zinc chloride in the manganese dioxide. This makes the manganese chloride crystals super deliquescent because of the ZnCl2.
Whenever I use alkaline batteries I get a dark orange solution upon filtering. Maybe it is better to use these batteries for manganese salt
production.
hibernating...
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blogfast25
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Remove any zinc by washing the crud first with warm dilute acetic acid (vinegar) or dilute sulphuric acid. Simmer gently for an hour or so, then
filter and wash. This should eliminate zinc, as well as other possible water soluble contaminants.
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elementcollector1
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Hey boys, look what I found! Manganese metal electrodeposition from chloride at below-0 Celsius temperatures!
http://www.springerlink.com/content/tu633h877633826l/
Credible or not?
Back on main topic, I don't recommend using a diluted strong acid, that doesn't tend to work well for me. Just go with vinegar.
UPDATE: This was a triumph! Pure manganese plated out thickly onto my iron cathode! (Take that, blogfast25!)
[Edited on 11-5-2012 by elementcollector1]
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nora_summers
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wonderful! do you have any pictures of your manganese metal?
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elementcollector1
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Sadly, I need to restock on HCl so I can get some more MnCl2, so I only have a little. I'll see if I can take a good pic...
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elementcollector1
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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.)
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blogfast25
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| 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.)
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Summarise what you're doing to obtain the MnCl2. I may have an idea what the contamination is.
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elementcollector1
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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?
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blogfast25
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| 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 Mn<sup>2+</sup> (and Fe<sup>3+</sup> 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]
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elementcollector1
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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.
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blogfast25
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| 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 (Fe<sup>3+</sup> 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(H<sub>2</sub>O)<sub>5</sub>OH]<sup>2+</sup>. 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 Fe<sup>3+</sup> with KSCN or NH4SCN, because FeSCN<sup>2+</sup> 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.
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elementcollector1
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Don't have those compounds, and from what I've heard, the test is very sensitive.
Anyway, just threw some more base (MnCO3) in, and the color definitely lightened from a liquid-bromine darkness to a tan-orange, translucent color.
Then, I threw some more carbonate in. Will check up tomorrow to see if it's that beautiful clear / rose-pink. (If so, I'm just going to mix my entire
pound of manganese carbonate with my 1000 mLs of solution. Should work, right?
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Arthur Dent
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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]
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blogfast25
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| 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] |
Quality [of the MnCO3] will depend from one product to another. Mine definitely contains some soluble iron (III) but it's bearable.
Remember that FeCl3 is very highly soluble, so chances are very good that it will stay in the mother liquor and not crystallise, as long as you
crystallise from acid solution.
Nice crystals, Robert.
[Edited on 24-5-2012 by blogfast25]
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Arthur Dent
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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
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blogfast25
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| 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
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Oh but that won't work, at least not if you mean 'anhydrous MnCl2' and not 'dry MnCl2 hydrate crystals'.
Anhydrous MnCl2 requires heating in a stream of dry HCl or calcining a mixture of MnCl2 hydrate and NH4Cl in a stream of a dry inert gas. Anything
else basically yields MnO2: dry but not what you want.
If dry crystals of the hydrate is what you want then just air drying should be enough.
[Edited on 24-5-2012 by blogfast25]
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Arthur Dent
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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
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blogfast25
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| 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
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I'm not sure. Partial hydrolysis and/or oxidation are risks if you don't at least exclude oxygen. Why would you want the dihydrate?
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