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elementcollector1
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Well, the sodium nitrate is going to evaporate over the course of a week or two... (I should really make a dessicator bag, but I used all my remaining
NaOH for this reaction...)
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blogfast25
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Quote: Originally posted by elementcollector1  | | Well, the sodium nitrate is going to evaporate over the course of a week or two... (I should really make a dessicator bag, but I used all my remaining
NaOH for this reaction...) |
You mean it's going to dry? I have force dried it [NaNO3] on a low light on a gas stove in a defunct stainless steel pan, no problems. Just keep
stirring to avoid sticking and it's a 10 - 15 min job.
Another one that works really well is CaSO4 + 8/3 Al === > CaS + 4/3 Al2O3. CaSO4 is of course... wallfiller! This mixture can also be ignited with
an Mg ribbon and burns like hell, possibly even hotter than chlorate/nitrate. It's best to dehydrate (by oven heating, 200 C does the trick) the
wallfiller (it's a hemihydrate) to anhydrous CaSO4 but for ignition mixes that might not even be strictly necessary.
[Edited on 23-6-2012 by blogfast25]
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elementcollector1
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Should I dehydrate the MnO2, or leave it as the brown hydrated form? (supposedly this is more active).
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blogfast25
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Dry it at 150 - 200 C. No need to calcinate. But it must be quite dry for aluminothermy...
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elementcollector1
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Well, it's drying fairly well in the open sun, so I could leave it to dry for two weeks while I'm in New York...
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rannyfash
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i have tried CuO/Mg thermite as a 'kewl' experiment, a possible source for atomised copper particles, can anyone explain why it deflagrates? Mg + CuO
==> MgO + Cu there is no difference in moles which usually causes the pressure difference in explosive reactions, does the extreme enthalpy change
cause the air around it to expand? or am i overlooking something simple
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elementcollector1
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Copper thermites are known for being explosive. >
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elementcollector1
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How long should the MnO2 be in the oven? Golden brown? XD
But seriously, how many hours at about 150 C?
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blogfast25
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Depends a bit on how wet it was to start with, of course. A nicely dripped dry filter cake: about 1 - 2 hours, I'd say.
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elementcollector1
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I heard mention of a molten MnCl2 - KCl eutectic a while back. What temp does this melt at? (Below 654 C is easily doable with a blowtorch, and
electrolysis is only mildly hard with that setup).
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blogfast25
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| Quote: Originally posted by elementcollector1 | | I heard mention of a molten MnCl2 - KCl eutectic a while back. What temp does this melt at? (Below 654 C is easily doable with a blowtorch, and
electrolysis is only mildly hard with that setup). |
The eutectic MP will be in that neighnourhood.
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elementcollector1
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Well, I gave the Mn-thermite another go.
Attempt one: Ignition with blowtorched plaster of paris and aluminum. Failed.
Attempt two: Ignition with magnesium. Failed, just burned on the top without igniting anything.
Attempt three: Ignition with blowtorch. Partial success?
I got bits and pieces of the powder to ignite and melt together, and I did get one particularly good-sized, gray lump. The other lumps are all
brittle-looking, gray to brown things, while this one looked more spherical while molten.
What should I do to reclaim a good-looking manganese sample (if the manganese is in there) from this mixture of Al2O3, Mn2O3, MnO2, Al, and hopefully
Mn?
EDIT: Ah, nope. Just crumbly gray powder. I washed it in bleach to remove some of the aluminum (I think), but I think I'll just discard this
particular experiment.
[Edited on 28-8-2012 by elementcollector1]
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elementcollector1
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Well, I had a potentially stupid but potentially game-changing idea. I read somewhere on this forum a while back (couldn't find OP even with search)
that if one melted aluminum, topped it up with manganese dioxide, added more aluminum, added more MnO2, added more aluminum and so on, they would
eventually have a container of manganese metal. (and aluminum, I suppose.) Is this a better idea? I have the materials to melt aluminum easily - I
just don't think my powdered aluminum works, and it takes forever to make new stuff. (Old stuff was made from Al foil, notoriously bad source of
aluminum.)
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blogfast25
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| Quote: Originally posted by elementcollector1 | | Well, I had a potentially stupid but potentially game-changing idea. I read somewhere on this forum a while back (couldn't find OP even with search)
that if one melted aluminum, topped it up with manganese dioxide, added more aluminum, added more MnO2, added more aluminum and so on, they would
eventually have a container of manganese metal. (and aluminum, I suppose.) Is this a better idea? I have the materials to melt aluminum easily - I
just don't think my powdered aluminum works, and it takes forever to make new stuff. (Old stuff was made from Al foil, notoriously bad source of
aluminum.) |
Well, it's worth a shot. Remember that the reduction reaction MnO2 + 4/3 Al === > Mn + 2/3 Al2O3 is very exothermic and now you'll carry out this
reaction at > 660 C! Chances are that your Mn will form in the vapour phase!
Add your MnO2 really slowly, pinch by pinch, to see how you get on and be well prepared in terms of fire and heat protection.
You'll also need to look up 'Al Mn alloys', to figure out what will happen to any formed Mn: will it dissolve in the Al bath? At first glance it seems
Mn solubility in Al is somewhat limited.
[Edited on 28-8-2012 by blogfast25]
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elementcollector1
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Although I believe MnAl (an intermetallic) was mentioned a while back as well. I think that if I add enough MnO2 slowly, then the Al will react with
that instead, leaving the Mn alone.
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blogfast25
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EC1:
Whatever you do with this, be really careful. Assuming the MP of Al is high enough to get the reaction going every time some MnO2 hits the Al melt,
then each time you will get a burst of enthalpy being released. Vaporous Mn is a real possibility!
Another problem is the slag (alumina): will it form in the molten state? Of course it will then solidify as its MP (2,054 C, IIRW) is much higher than
the MP of Al.
All in all this could be a really interesting experiment but with serious potential danger to the experimenter. Add the aliquots of MnO2 using a long
armed ladle or something like that.
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elementcollector1
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Ja, I know. That's why I mentioned the idea being possibly stupid. However, it would ensure a reaction, which is better than what I've been getting.
Also, tried electroplating of manganese again today. Surprise! A silvery coating formed. It's thin, but I think it could get thicker if I'm careful
about it.
Setup:
9V battery
Beaker 1: Dilute MnCl2 (anolyte)
Beaker 2: Saturated MnCl2 (catholyte)
Anode: 60/40 solder
Cathode: Copper PCB, unetched
I have to keep brushing it every few minutes, otherwise it builds up an oxide coating, but yay!
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elementcollector1
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So, how do I build up enough Mn to rip it off? (More importantly, how do I remove the copper foil from the PCB blank without etching?)
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blogfast25
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Hmmm... geting the experimental conditions for viable plating correct isn't so easy. Are you sure the build up is actually oxide? Search for patents
in that field, I'd say.
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elementcollector1
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I searched, and found that a concentrated catholyte is good, a cold solution is better, etc.
The plate is for the most part, bright, shiny silvery metal. Some black is visible, but that's to be expected.
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blogfast25
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| Quote: Originally posted by elementcollector1 | I searched, and found that a concentrated catholyte is good, a cold solution is better, etc.
The plate is for the most part, bright, shiny silvery metal. Some black is visible, but that's to be expected. |
Very nice. Buff it or polish it and put it in your collection. Recovering it would be hard, I guess industrially they obtain thick coatings/crystals
and somehow scrape off the metal, then ingot it.
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elementcollector1
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Well, I really just need to remove it from the PCB board.
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