ChemEstudent
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Precipitant that will drop Al that is dissolved in HCl??
I placed 8 grams of pure Al bars containing tantalum pins in HCl. I used the HCl to dissolve the aluminum so I could obtain the tantalum. I also wish
to recover the aluminum. I was thinking of neutralizing the acid with sodium bicarbonate and then adding magnesium to displace the aluminum and
recover it. Would this work?
I also thought about attacking the tantalum in a heated(~80c) nitrogen atmosphere to break the weld it has on the aluminum, without harming the Al.
I have also tried 50% NaOh(aq) but it dissolves both the Al and Ta, and I haven't the slightest clue on how to precipitate both of them from that
solution.
Al(s)+ OH-(aq)=> AlO2-(aq)+ H2O + H2(g)
[or Al(OH)4-(aq)\Al(OH)6 3-(aq)]
Does anyone have a clue as to how I can remove these pins without dissolving the aluminum or how to precipitate the Al out of the HCl??
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froot
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With HCl Al forms hydrated AlCl3 and with NaOH it forms Sodium Aluminate.
To retrieve the Al you'll need to reduce it back to elemental Al by electrolysis of a molten Al salt.
Heat might be the best way to go, if they don't release then melt the Al off them.
[Edited on 1-8-2013 by froot]
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blogfast25
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Quote: Originally posted by ChemEstudent  |
Does anyone have a clue as to how I can remove these pins without dissolving the aluminum or how to precipitate the Al out of the HCl??
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Add strong ammonia to precipitate the Al3+ as Al(OH)3. Ammonia is not alkaline enough to form soluble aluminates, so the
hydroxide crashes out instead.
But that's not exactly 'Al recovery' if by that you understand recovery as metal. Extracting Al metal from its compounds isn't workable at the DIY
level. Al is also too cheap to recover.
If you're doing what you're doing on a commercial scale you might want to look into Al recovery as alum or ammonium alum, both of which have
considerable commercial value.
[Edited on 1-8-2013 by blogfast25]
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AJKOER
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Ta melts at over 3,000 C while Al at 660 C.
So, in agreement with Froot, use a hot flame torch to melt away the Al from the Ta.
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MrHomeScientist
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There isn't any way to precipitate elemental aluminum from an aqueous solution of its salts, I'm nearly positive of that. Otherwise that would
probably be used by industry. It's really just too reactive - if any did form it would likely just immediately react with the water. Like froot said,
molten salt electrolysis is the way to go (the Hall-Heroult process). Not easy to do, though, as it requires very high temperatures.
EDIT: Apparently everybody replied as I was typing. Simply melting the aluminum off is a great idea.
[Edited on 8-1-2013 by MrHomeScientist]
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blogfast25
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| Quote: Originally posted by MrHomeScientist | There isn't any way to precipitate elemental aluminum from an aqueous solution of its salts, I'm nearly positive of that. Otherwise that would
probably be used by industry. It's really just too reactive - if any did form it would likely just immediately react with the water. Like froot said,
molten salt electrolysis is the way to go (the Hall-Heroult process). Not easy to do, though, as it requires very high temperatures.
EDIT: Apparently everybody replied as I was typing. Simply melting the aluminum off is a great idea.
[Edited on 8-1-2013 by MrHomeScientist] |
On a lab scale reduction of anhydrous AlCl3 with K or Na works. But it's not commercially viable: when that method was still used (pre
Hall-Heroult) the metal was reportedly as expensive as some precious metals. Apparently people with more money than sense used it for cutlery. Now
only backpackers would do that!
[Edited on 1-8-2013 by blogfast25]
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watson.fawkes
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What's the solubility of Ta in molten Al?
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AJKOER
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Actually, Ta apparently rapidly, per Wikipedia (http://en.wikipedia.org/wiki/Tantalum ), to quote:
"The tantalum is capable of capturing oxygen and nitrogen by forming nitrides and oxides and therefore helps to sustain the high vacuum needed for the
[electron] tubes"
does develop a protective oxide/nitride coating. So, the interaction is first between molten Al and these coatings before we have to address the
formation of the tantalum aluminides (like TaAl3 and Ta3Al).
However, I am assuming an air/fuel based torch. Why not H2/Cl2 (which produces a very hot flame) given the inert nature of Ta?
This is probably a case of trying various torches and seeing what works best.
[Edited on 1-8-2013 by AJKOER]
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phlogiston
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If you are doing this purely for the challenge you could try recovering the aluminum from the aqueous solution by electrolysis using mercury as the
cathode, and then distilling off the mercury.
-If- it works, it's not going to be worth it in terms of money, and it is definately a bad idea in terms of cost to your health.
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"If a rocket goes up, who cares where it comes down, that's not my concern said Wernher von Braun" - Tom Lehrer |
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