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elementcollector1
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I hope nobody's planning on actually using a beaker - molten lithium doesn't play nice with glass, as detailed by Zan Divine in his thread on
ampouling lithium.
I like this idea so far - might be tempted to get some lithium salts and try it out for myself when I get the chance.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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blogfast25
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Quote: Originally posted by elementcollector1  |
I hope nobody's planning on actually using a beaker - molten lithium doesn't play nice with glass, as detailed by Zan Divine in his thread on
ampouling lithium.
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They only make beakers out of glass now? The things you 'learn' on this site! 
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Upsilon
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| Quote: Originally posted by blogfast25 |
Yeah, like when the beaker is empty: for each reduced Li<sup>+</sup> a nitrate ion must be oxidised!
I did tell you someone had tried this before, I just couldn't find the reference.
Search for related posts with 'WGTR' as author and you might find what he did precisely.
Remember: molten nitrates will oxidise anything that's oxidisable. |
I'm still not following; say you have 3 LiNO3 molecules. When the first one is electrolyzed, Li is formed at the cathode and something (NO2 and O2?)
is formed at the anode. This leaves 2 molecules of LiNO3. One of them reacts with the lithium metal to form an LiNO3 molecule. But then this leaves 2
NO3- ions and 3 Li+ ions? How can this be?
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blogfast25
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| Quote: Originally posted by Upsilon |
I'm still not following; say you have 3 LiNO3 molecules. When the first one is electrolyzed, Li is formed at the cathode and something (NO2 and O2?)
is formed at the anode. This leaves 2 molecules of LiNO3. One of them reacts with the lithium metal to form an LiNO3 molecule. But then this leaves 2
NO3- ions and 3 Li+ ions? How can this be? |
You need to work out the EXACT redox reaction for the oxidation of Li(0) to Li(+1) by nitrate ions. Assume the nitrate is reduced to NO.
Hint: nitrogen goes from N(+5) to N(+2), Li from Li(0) to Li(+1).
[Edited on 18-10-2015 by blogfast25]
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Upsilon
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AH! Ok, I see now, lithium nitrate isn't formed, lithium oxide is:
4Li + NO3- -> NO + 2Li2O
(EDIT: And I even got it before I saw the hint  )
And thus this lithium oxide layer isn't soluble in molten lithium nitrate, protecting the bulk of the lithium metal from further oxidation.
And thanks for letting me know about the incompatibility of lithium with glass, though it shouldn't matter for me; the LiNO3 and KNO3 eutectic melts
at around 120C which is much lower than lithium's melting point. I do need to make some KNO3 though.
[Edited on 18-10-2015 by Upsilon]
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Upsilon
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I want to get to the TiO2 thermite today but it's taking a while to get rid of the water in the NaNO3 solution. Calcium nitrate should technically
work as well, right?
UPDATE: So calcium nitrate did not work. I didn't have high hopes because it was in hydrated form (I don't see a quick effective way to dehydrate it).
It sparked a little bit when the Mg ribbon hit the pile but it quickly went out. You couldn't even tell if anything happened by looking at it.
So, I did a regular Fe2O3 thermite to get an iron sample (5 pages of the thread already and still no elements officially collected!!!). I used 2.5g of
Al and roughly 7g of Fe2O3 and got a couple blobs of iron. I let them soak in vinegar for a while and then scrubbed them with CLR and they cleaned up
a bit, but they're still a little dull. I have them stored in a little vial under mineral oil; I would post a picture but I'm too lazy. If anyone has any suggestions to make them shine then please, suggest away. But hey,
my first officially collected element, woohoo!
Also, one of the pieces is much more magnetically responsive than the other. I'm thinking the other one has a high alumina slag content; I may just
throw it out - it's the uglier of the two anyway. Unless someone has a good suggestion for dissolving the alumina without harming the iron too much.
[Edited on 18-10-2015 by Upsilon]
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blogfast25
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| Quote: Originally posted by Upsilon | I want to get to the TiO2 thermite today but it's taking a while to get rid of the water in the NaNO3 solution. Calcium nitrate should technically
work as well, right?
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Calcium nitrate would work if it wasn't so G-d damn hygroscopic.
NaNO3 solutions you can simply boil down, then oven dry the product.
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Upsilon
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| Quote: Originally posted by blogfast25 |
Calcium nitrate would work if it wasn't so G-d damn hygroscopic.
NaNO3 solutions you can simply boil down, then oven dry the product. |
That's what I'm having to do. However once I boiled the majority of the water off, it seemed like mini explosions started happening (extremely loud
POPs that had enough force to move a 1000mL beaker with ~200mL of liquid; had I not seen it happening it would have eventually migrated off of the hot
plate and smashed on the floor). So now I have to leave the rest to oven drying, which is taking quite a while (my drying oven is actually just a fish
tank wrapped in aluminum foil, wrapped in a blanket, with a couple 65W incandescent bulbs in the lid. I'm quite pleased with how well it works - it's
currently bordering 90C inside! Problem is that there's very little ventilation and it stays fairly humid inside, slowing evaporation rate)
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Upsilon
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For gallium and indium, I plan on purchasing a Galinstan alloy sample (68% Ga, 22% In, and 10% Sn). To separate them, I am thinking of burning the
alloy in chlorine gas to produce GaCl3, InCl3, and SnCl2. The GaCl3 will likely be vaporized during this process and will need to be recondensed.
Afterward the InCl3 and SnCl2 mixture will be heated to drive off any remaining GaCl3. The remaining InCl3 and SnCl2 will be dissolved in water,
because SnCl2 apparently hydrolyzes when diluted in hot water to form insoluble Sn(OH)Cl (I will likely throw this out, this small amount of tin isn't
really worth saving). This will precipitate out, leaving an InCl3 solution. Ga and In could probably be obtained by electrolysis of solutions of their
respective chlorides.
That's the theory anyway. This is the best approach I could come up with using readily available information - there is not a whole lot out there on
gallium and indium compounds. The Galinstan alloy is also the best "bang for your buck" approach I see - a 10g alloy sample is $30 USD and will
theoretically give me 6.8g of gallium and 2.2g of indium.
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j_sum1
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| Quote: | | a 10g alloy sample is $30 USD and will theoretically give me 6.8g of gallium and 2.2g of indium. |
You can buy gallium a lot cheaper than this. Indium is still expensive (and will continue to be so.) I would be surprised if this was really the
best bang for your buck. Especially when you consider the costs of the process you describe (and the losses).
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Upsilon
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| Quote: Originally posted by j_sum1 |
You can buy gallium a lot cheaper than this. Indium is still expensive (and will continue to be so.) I would be surprised if this was really the
best bang for your buck. Especially when you consider the costs of the process you describe (and the losses). |
I know it's probably cheaper to buy them on their own. However a lot of places won't sell the metals in these small amounts, so I'll end up paying
more for more than I need. I'm sure there's somewhere I can buy small samples, but you're forgetting the OCD factor of needing to produce them myself!
I would need to do more research on this before actually carrying this out. I really don't know how the alloy will react to chlorine - whether it will
spontaneously ignite upon exposure or if it needs heating. If it doesn't spontaneously ignite, then loss shouldn't be too bad - I would fill the flask
with chlorine, then move it to a condenser setup and apply heat. If it does ignite on contact, then I would probably lose some GaCl3 through the
chlorine input tube.
EDIT: It might actually be easier to use bromine instead of chlorine; more predictable and easier to carry out (essentially reacting 2 liquids; no
burning required). The only issue with that is that I don't know if the tin(ii) bromide hydrolyzes like the chloride; in that case I could run
chlorine through the InBr3 + SnBr2 solution to convert them to chlorides and continue from here as outlined above. Still uses chlorine gas but at
least will avoid losses associated with burning.
[Edited on 19-10-2015 by Upsilon]
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j_sum1
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Good luck. I have not forgotten the synthesising challenge. That is where the fun is. i just know what I am capable of and this looks like a bit of
a headache. Even playing with something as gunky as SnCl2 is offputting.
I like your thinking though. This might be a good way of purifying an already relatively clean sample.
It seems to me like gallinstan might be an ideal candidate for zone-refining. Totally different process, but it might really work well.
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Upsilon
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I wonder if the alloy would react with conc. HCl at any appreciable rate? That would certainly make things much easier. I'm not sure I want to order
the alloy just yet, but when I finally get around to to, I'll have to test a tiny amount with some HCl to see if it works. Unless someone knows for
sure whether or not it will work.
I think I'm going to go for manganese next (made from manganese dioxide). I'd rather avoid manganese dioxide thermite (it is certainly cool but not
really practical for collecting samples), so I have to try something else. I'm thinking of reacting manganese dioxide with aqueous sodium
metabisulfite - manganese should be reduced to Mn2+ and metabisulfite reduced to sulfate affording manganese sulfate, but I'm not totally sure what
the reaction here is. My best guess is:
2MnO2 + Na2S2O5 -> 2MnSO4 + Na2O
Can anyone confirm?
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blogfast25
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Na2S2O5 + 2 MnO2 + H2SO4 === > Na2SO4 + 2 MnSO4 + H2O
Note that this does need H2SO4 and probably some gentle heat.
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Upsilon
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Ah, ok. That's not a problem. To extract metallic manganese, would it be feasible to reduce the MnSO4 solution with aluminum? Would the manganese
react too badly with water?
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j_sum1
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A quick scan of reduction potentials Link gives a potential of -1.185V for reduction of Mn. This is the wrong side of -0.8227 for electrolysis of water. In other words, Mn reduction
is not going to work easily in an aqueous environment. (The reduction of water at -0.83V forms a convenient dividing line in the table. Metals on
one side can be electroplated in solution. Metals on the other side can't.)
Aluminium is certainly a powerful enough reducing agent. I think you could come up with a nice thermite for this. You would want to crystallise the
manganese sulfate -- you will probably get the heptahydrate. You then want to drive off the water of hydration. Then set up for a thermite.
An alternative, and probably more feasible would be to convert back to MnO2. This, freshly precipitated and nice and pure will be easy to work with
and react well.
Blogfast will undoubtedly be along soon to tell us that Mn thermites tend to have low yields and a lot of slag mixed with the metal. The Mn tends to
boil off. This is correct. I don't know if working with the sulfate improves this situation at all. I suspect not. You should probably consider
adding a flux and something to absorb a bit of the heat if you want to keep your product.
An alternative is carbon reduction. If only you had a source of nicely mixed MnO2 and carbon powder...
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Upsilon
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Well hell, at that rate I may as well just electrolyze molten manganese(II) acetate (only a 210C melting point!)
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j_sum1
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I have no idea how acetate behaves in molten salt electrolysis. One way to find out.
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Upsilon
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Well technically any molten ionic compound can be electrolyzed. Manganese is immediately formed at the cathode, and god knows what at the anode - but
what happens at the anode shouldn't really matter as long as it is kept separate from the cathode. The only issue I could imagine is some
incompatibility of manganese with molten manganese acetate - I don't think it would attack the manganese since acetate isn't a very good oxidizer.
There's the possibility that this could fail for the same reason KOH electrolysis fails, but I find that unlikely since manganese isn't even close to
liquefying at these temperatures - but hey, I really don't know either. I'll definitely be trying this.
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blogfast25
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| Quote: Originally posted by j_sum1 | A quick scan of reduction potentials Link gives a potential of -1.185V for reduction of Mn. This is the wrong side of -0.8227 for electrolysis of water. In other words, Mn reduction
is not going to work easily in an aqueous environment. (The reduction of water at -0.83V forms a convenient dividing line in the table. Metals on
one side can be electroplated in solution. Metals on the other side can't.)
Aluminium is certainly a powerful enough reducing agent. I think you could come up with a nice thermite for this. You would want to crystallise the
manganese sulfate -- you will probably get the heptahydrate. You then want to drive off the water of hydration. Then set up for a thermite.
An alternative, and probably more feasible would be to convert back to MnO2. This, freshly precipitated and nice and pure will be easy to work with
and react well.
Blogfast will undoubtedly be along soon to tell us that Mn thermites tend to have low yields and a lot of slag mixed with the metal. The Mn tends to
boil off. This is correct. I don't know if working with the sulfate improves this situation at all. I suspect not. You should probably consider
adding a flux and something to absorb a bit of the heat if you want to keep your product.
An alternative is carbon reduction. If only you had a source of nicely mixed MnO2 and carbon powder... |
Correct on all counts.
The MnO2 thermite is possible, gives fairly low yields and has been reported on by moi on this forum. Let me know if you can't find it.
Electrodeposition of Mn from aqueous solution appears possible (despite the SRPs being unfavourable) but it ain't easy, I think.
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elementcollector1
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It may also fail due to decomposition into carbonate, or some such. Be sure to check stability at your desired temperature.
I managed electroplating of Mn using MnCl2, a 60/40 solder anode, and a Cu cathode. The Mn easily oxidized, and required scrubbing every
now and then to keep shiny and silver. Yield was pitiful, though that may have been more the fault of my electrical setup than anything.
[Edited on 10-20-2015 by elementcollector1]
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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blogfast25
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Ermm... explain?
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j_sum1
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I have come a long way in a year of hanging around this place. Much of it due to reading your posts blogfast. Thanks.
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blogfast25
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| Quote: Originally posted by j_sum1 |
I have come a long way in a year of hanging around this place. Much of it due to reading your posts blogfast. Thanks. |
I prefer to think whatever progress anyone makes is largely due to their own efforts.
Having said that, flattery works for me too! 
A word on the MnO2 thermite. It's a rare case where instead of boosting, cooling (with CaF2 or CaO) is needed. That's because the BP of Mn
and the MP of alumina are very close together, leading to some of the formed Mn boiling off. That effect can be reduced a bit by slowing the
burn with inert heat sinks like CaF2 or the cheaper CaO. Personal experience, that. I spent an awful lot of time experimenting that thermite.
[Edited on 20-10-2015 by blogfast25]
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Upsilon
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I'll be trying the manganese acetate molten electrolysis when I have the time, then. Problem is that I need to clean up my sulfuric acid first from
being stored for too long  It's not really a big deal, it's barely got a gray
tinge to it, but I have the urge to clean it before using it.
Also I should note that my 50 grams of cinnabar (HgS) powder should be coming in soon. Wikipedia outlines roasting it in air and condensing the
mercury vapor...but uh, mercury vapor, not gonna happen
Instead I will try reacting it with a concentrated acid and electrolysing the aqueous corresponding salt. HgCl2 isn't terribly soluble in water, so I
may use nitric acid instead of HCl. I don't know how soluble mercury nitrate is but nitrates are typically more soluble than chlorides so it's worth a
shot.
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