Sciencemadness Discussion Board

Quest for the elements

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elementcollector1 - 17-10-2015 at 17:42

Quote: Originally posted by blogfast25  


Yeah, like when the beaker is empty: for each reduced Li<sup>+</sup> a nitrate ion must be oxidised! :D



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.

blogfast25 - 17-10-2015 at 18:33

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.



They only make beakers out of glass now? The things you 'learn' on this site! :D

Upsilon - 17-10-2015 at 18:56

Quote: Originally posted by blogfast25  


Yeah, like when the beaker is empty: for each reduced Li<sup>+</sup> a nitrate ion must be oxidised! :D

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?

blogfast25 - 17-10-2015 at 19:13

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]

Upsilon - 17-10-2015 at 19:31

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 :P)

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]

Upsilon - 18-10-2015 at 09:17

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. :P 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]

blogfast25 - 18-10-2015 at 17:02

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?



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.

Upsilon - 18-10-2015 at 17:24

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)

Upsilon - 18-10-2015 at 18:55

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.

j_sum1 - 18-10-2015 at 19:13

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).

Upsilon - 18-10-2015 at 19:19

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! :P

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]

j_sum1 - 18-10-2015 at 19:47

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.

Upsilon - 19-10-2015 at 14:37

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?

blogfast25 - 19-10-2015 at 15:01

Quote: Originally posted by Upsilon  
My best guess is:
2MnO2 + Na2S2O5 -> 2MnSO4 + Na2O
Can anyone confirm?


Na2S2O5 + 2 MnO2 + H2SO4 === > Na2SO4 + 2 MnSO4 + H2O

Note that this does need H2SO4 and probably some gentle heat.

Upsilon - 19-10-2015 at 15:07

Quote: Originally posted by blogfast25  


Na2S2O5 + 2 MnO2 + H2SO4 === > Na2SO4 + 2 MnSO4 + H2O

Note that this does need H2SO4 and probably some gentle heat.


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?

j_sum1 - 19-10-2015 at 15:52

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...

Upsilon - 19-10-2015 at 16:03

Well hell, at that rate I may as well just electrolyze molten manganese(II) acetate (only a 210C melting point!)

j_sum1 - 19-10-2015 at 16:16

I have no idea how acetate behaves in molten salt electrolysis. One way to find out.

Upsilon - 19-10-2015 at 16:28

Quote: Originally posted by j_sum1  
I have no idea how acetate behaves in molten salt electrolysis. One way to find out.


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.

blogfast25 - 19-10-2015 at 17:11

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.

elementcollector1 - 19-10-2015 at 17:11

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]

blogfast25 - 19-10-2015 at 17:48

Quote: Originally posted by elementcollector1  
It may also fail due to decomposition into carbonate, or some such.


Ermm... explain?

j_sum1 - 19-10-2015 at 17:53

Quote: Originally posted by blogfast25  


Correct on all counts.
<snip>

I have come a long way in a year of hanging around this place. Much of it due to reading your posts blogfast. Thanks.

blogfast25 - 19-10-2015 at 18:10

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]

Upsilon - 19-10-2015 at 18:29

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 :mad: 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 :P
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.

j_sum1 - 19-10-2015 at 18:41

You are on your own with the mercury. I know almost nothing.
Handle safely and all that.

BobD1001 - 19-10-2015 at 19:46

Quote: Originally posted by Upsilon  
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 :mad: 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 :P
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.


Upsilon,

There was a great write-up somewhere on this forum of a wet extraction of mercury from cinnabar. By far the best method I've yet seen. Although the mercury for my element collection just came from an old thermostat. Safe and sealed.

elementcollector1 - 19-10-2015 at 20:08

In fact, that write-up just so happens to be stickied in the Chemistry in General subforum.

Upsilon - 20-10-2015 at 07:17

Quote: Originally posted by elementcollector1  
In fact, that write-up just so happens to be stickied in the Chemistry in General subforum.


Yes, I saw that, looks like the general idea is basically what I considered - reacting it with an acid. Second thoughts about mercury nitrate though, the nitrate will oxidize the mercury as it is evolved. I'll have to use HgCl2 and heat it up so that a decent amount is dissolved

blogfast25 - 20-10-2015 at 09:10

Quote: Originally posted by Upsilon  
Second thoughts about mercury nitrate though, the nitrate will oxidize the mercury as it is evolved.


I'm not quite sure what method you're referring to but what you write may well be wrong.

The reduction of nitrate usually proceeds as follows:

NO<sub>3</sub><sup>-</sup> + 4 H<sup>+</sup> + 3 e<sup>-</sup> === > NO + 2 H<sub>2</sub>O.

Without acid (strictly speaking H<sub>3</sub>O<sup>+</sup>;), nitrate has almost no oxidising properties in aqueous solutions.

If it did, aqueous solutions of nitrates (other than nitric acid itself) could be used to dissolve metals but that isn't true.



[Edited on 20-10-2015 by blogfast25]

Upsilon - 20-10-2015 at 09:49

Quote: Originally posted by blogfast25  
Quote: Originally posted by Upsilon  


I'm not quite sure what method you're referring to but what you write may well be wrong.

The reduction of nitrate usually proceeds as follows:

NO<sub>3</sub><sup>-</sup> + 4 H<sup>+</sup> + 3 e<sup>-</sup> === > NO + 2 H<sub>2</sub>O.

Without acid (strictly speaking H<sub>3</sub>O<sup>+</sup>;), nitrate has almost no oxidising properties in aqueous solutions.

If it did, aqueous solutions of nitrates (other than nitric acid itself) could be used to dissolve metals but that isn't true.



[Edited on 20-10-2015 by blogfast25]


Ah, ok. I guess I was incorrectly relating the oxidizing properties of molten nitrates to aqueos ones. Looks like I'll use nitric acid to dissolve the HgS powder; it should be more soluble than HgCl2,

blogfast25 - 20-10-2015 at 10:27

Quote:
Ah, ok. I guess I was incorrectly relating the oxidizing properties of molten nitrates to aqueos ones. Looks like I'll use nitric acid to dissolve the HgS powder; it should be more soluble than HgCl2,


I'm really not sure whether that would work. HgS is one of these incredibly insoluble sulphides that really requires very strong acids to get it to dissolve. Unless the nitric acid manages to oxidise the sulphide ions, this method may prove a very slow boat to China... but I'm not putting my hand in the fire on this one.

HgS: K<sub>s</sub> = 2 x 10<sup>-53</sup>. Daaangng! Minus fiftythree...



[Edited on 20-10-2015 by blogfast25]

Upsilon - 20-10-2015 at 12:06

Quote: Originally posted by blogfast25  

I'm really not sure whether that would work. HgS is one of these incredibly insoluble sulphides that really requires very strong acids to get it to dissolve. Unless the nitric acid manages to oxidise the sulphide ions, this method may prove a very slow boat to China... but I'm not putting my hand in the fire on this one.

HgS: K<sub>s</sub> = 2 x 10<sup>-53</sup>. Daaangng! Minus fiftythree...



[Edited on 20-10-2015 by blogfast25]


Possibly. Wikipedia states that beta-HgS is "unreactive to all but concentrated acids". What I'm getting is alpha-HgS which it states nothing on, but the way it's worded hints that alpha-HgS is more reactive (not really solid proof but...). I'll just have to try several different acids; 50g is a hefty amount and I don't need nearly that much mass for a mercury sample so I can spare some to experiment.

EDIT: The oxidation of the sulfide ion is actually quite favorable in this case, but the problem is that there will be very few S2- ions to be oxidized - must of them will be locked up in the HgS. As long as some of it is able to dissolve then it should proceed. How fast is another story.

[Edited on 20-10-2015 by Upsilon]

blogfast25 - 20-10-2015 at 14:44

Have a look at this, if you haven't already:

http://www.sciencemadness.org/talk/viewthread.php?tid=18162&...

Upsilon - 20-10-2015 at 16:23

Manganese experiment failed before I even got to any electrolysis. The reaction between manganese dioxide, sodium metabisulfite, and sulfuric acid did not go as planned. I added 7.56g of manganese dioxide and 8.37g Na2S2O5 to a beaker, and added 16 ml of water to dissolve the metabisulfite. The contents instantly turned the color of manganese dioxide when I added the water; it is probably just the metabisulfite dissolving but it seemed to happen abnormally fast. After that I added 2.4 ml of 98% sulfuric acid drop by drop. At first I was seeing water instantaneously boil upon adding the acid, but towards the end of adding the acid it seemed to be fizzing like some kind of reaction was producing gas. Definitely was not boiling water. I suspect somehow that the metabisulfite was decomposing into sulfur dioxide and sodium sulfite, because the reaction was producing a pungent stinging smell (which may have been vaporized sulfuric acid as well). Note at this point there is still a large amount of unreacted manganese dioxide in the beaker. After that I put the beaker on a hot plate and let it get to around 70-80C; it was making periodic popping noises like it was boiling but no bubbles were apparent at the surface. I did not keep it on the plate very long and did not wait for it to stop making noise. There was still a lot of manganese dioxide in the solution, so I attempted to filter it off, but that pesky stuff slips right through my cheap grocery store coffee filters. At this point failure was inevitable so I started adding sodium carbonate to the solution; not much dissolved and no precipitation of MnCO3 was apparent (though it was difficult to see through the cloudy MnO2 suspended in the solution).

So looks like that route is a bust. Does anyone have a better suggestion? Apparently this can be done with sulfur dioxide but I'd rather try something else before doing that.

EDIT: There's also the oxalic acid + sulfuric acid method. I'll probably do that.

EDIT 2: Now that I'm thinking about it, can't I just use acetic acid instead of sulfuric acid in the aforementioned method since sulfate is just a spectator ion in that reaction? Since the actual redox reaction (the important part) is derived from:
MnO2 + 4H+ + 2e- -> Mn2+ + 2H2O
H2C2O4 -> CO2 + 2H+ + 2e-
Therefore the sulfate ion in the sulfuric acid only provides an anion for the Mn2+ to bind with. Acetate could play the same role, no? That way I would directly get to the desired product.

[Edited on 21-10-2015 by Upsilon]

[Edited on 21-10-2015 by Upsilon]

blogfast25 - 20-10-2015 at 16:41

Quote: Originally posted by Upsilon  

EDIT: There's also the oxalic acid + sulfuric acid method. I'll probably do that.



That works. Add the oxalic acid (or an oxalate) bit by bit because that thing foams a lot and is exothermic.

And yes, it should work with acetic acid instead of sulphuric but it'll be slower.


H2C2O4 -> CO2 + 2H+ + 2e- is incorrect though...


[Edited on 21-10-2015 by blogfast25]

Upsilon - 20-10-2015 at 16:55

Quote: Originally posted by blogfast25  


H2C2O4 -> CO2 + 2H+ + 2e- is incorrect though...


[Edited on 21-10-2015 by blogfast25]


Hah, good catch. Looks like the table I'm using has a mistake. Here's the right one:
H2C2O4 -> 2CO2 + 2H+ + 2e-

Anyway, I'm guessing household vinegar won't be feasible for this, since you said this occurs slowly even with concentrated acetic acid (I assume you were talking about concentrated acetic acid and not vinegar).

blogfast25 - 20-10-2015 at 17:01

Quote: Originally posted by Upsilon  
Quote: Originally posted by blogfast25  


H2C2O4 -> CO2 + 2H+ + 2e- is incorrect though...


[Edited on 21-10-2015 by blogfast25]


Hah, good catch. Looks like the table I'm using has a mistake. Here's the right one:
H2C2O4 -> 2CO2 + 2H+ + 2e-

Anyway, I'm guessing household vinegar won't be feasible for this, since you said this occurs slowly even with concentrated acetic acid (I assume you were talking about concentrated acetic acid and not vinegar).


Household vinegar is only about 0.8 M in HOAc. Why not try it on a small scale, using an excess vinegar, e.g. twice the stoichiometric amount?

Upsilon - 20-10-2015 at 17:06

Quote: Originally posted by blogfast25  

Household vinegar is only about 0.8 M in HOAc. Why not try it on a small scale, using an excess vinegar, e.g. twice the stoichiometric amount?


I do have some glacial acetic acid on the way, but I suppose it would be worth trying this out with vinegar for the sake of the more budget-minded home chemist. I'll give it a shot tomorrow I suppose.

blogfast25 - 20-10-2015 at 17:20

BTW, years ago I tried reducing MnO2 with bisulphite + acid too and it didn't work for me either.

Upsilon - 20-10-2015 at 18:46

Quote: Originally posted by blogfast25  
Have a look at this, if you haven't already:

http://www.sciencemadness.org/talk/viewthread.php?tid=18162&...


Back to this, I was actually wondering about nitrogen dioxide for this purpose. Perhaps the method described in that post can be made more efficient using liquid nitrogen dioxide? There's so much to try with with HgS. Some other things worth trying might be household bleach and hydrogen peroxide, since they too are good oxidizers. In an extreme case perchloric acid would probably work.

blogfast25 - 20-10-2015 at 18:56

Quote: Originally posted by Upsilon  
Quote: Originally posted by blogfast25  
Have a look at this, if you haven't already:

http://www.sciencemadness.org/talk/viewthread.php?tid=18162&...


Back to this, I was actually wondering about nitrogen dioxide for this purpose. Perhaps the method described in that post can be made more efficient using liquid nitrogen dioxide? There's so much to try with with HgS. Some other things worth trying might be household bleach and hydrogen peroxide, since they too are good oxidizers. In an extreme case perchloric acid would probably work.


Bleach will oxidise a sulphide to sulphur in a jiffy. Thin bleach is very dilute of course (4 to 5 % of hypochlorite, off the top of my head), but definitely worth trying, as is hydrogen peroxide.

Perchloric acid is a poor oxidiser in aqueous solution. Bizarre but true.

UC235 - 20-10-2015 at 19:01

If you're trying to use metabisulfite to reduce MnO2 to a soluble form, don't use too much. Manganous Sulfite trihydrate has poor water solubility and the granular crystals that form admixed with unreacted MnO2 often look like nothing has happened. Adding acid will produce copious amounts of SO2 and the solid will mostly go into solution.

Upsilon - 20-10-2015 at 19:03

Quote: Originally posted by blogfast25  

Perchloric acid is a poor oxidiser in aqueous solution. Bizarre but true.


Yeah, it seems sort of similar to sulfuric acid in this regard, not really exhibiting oxidizing properties unless heated. Hot perchloric acid is supposed to be a very good oxidizer, so if I were ever to try it then it would need heating.

Upsilon - 20-10-2015 at 19:07

Quote: Originally posted by UC235  
If you're trying to use metabisulfite to reduce MnO2 to a soluble form, don't use too much. Manganous Sulfite trihydrate has poor water solubility and the granular crystals that form admixed with unreacted MnO2 often look like nothing has happened. Adding acid will produce copious amounts of SO2 and the solid will mostly go into solution.


That's about what I did; the sulfuric acid I added caused some fizzing, which at this point I am pretty sure was sulfur dioxide. I did not know this was desired, though. Regardless I still ended up with a lot of suspended insolubles. I need better filter paper :mad: those cheapie brown paper towels worked extremely well for filtering MnO2 in Nurdrage's videos, I might pick some of those up.

Upsilon - 20-10-2015 at 19:39

Alright, using bleach or hydrogen peroxide may not get me anywhere. These would form HgO, which to my understanding is no easier to dissolve than HgS. However, I found this paper:
http://www.sciencedirect.com/science/article/pii/S0021979705...
That seems to suggest that HgO is unusually soluble at low pH. But it also says that HgO is very soluble at standard pH, so I don't know if it's got any credibility to it.

gdflp - 20-10-2015 at 19:58

I don't believe that is what they're saying. I think that they are discussing the rate of dissolution, not the solubility, at low pH. Also the "high solubility" at neutral pH is likely relative to other metal oxides.

Upsilon - 20-10-2015 at 20:56

Quote: Originally posted by gdflp  
I don't believe that is what they're saying. I think that they are discussing the rate of dissolution, not the solubility, at low pH. Also the "high solubility" at neutral pH is likely relative to other metal oxides.


Ah, you're right. I should have read that more closely. Well, regardless HgO may actually prove to be more compliant than HgS just because of its greater solubility. It would be more common than it is in the environment if it was as stable as HgS I think. Though that paper seems to explain why HgO is so rare in mineral form, but I don't completely understand it.

Upsilon - 21-10-2015 at 16:50

I'm in the middle of trying the manganese acetate reaction. I added ~300mL of household vinegar (stoichiometry only calls for about 170mL) to a beaker, and then added 6.26g of MnO2 and 9.08g of oxalic acid. There is definitely bubbling of CO2, but it is awfully slow (you can't tell that anything is happening until you get REALLY close). I'll let it sit for a while and I may try heating it gently tomorrow. I should theoretically get 25g of manganese acetate out of this (which I highly doubt I will). I'll also try it with glacial acetic acid once that arrives.

It would also be interesting to see if the CO2 being formed leads to any noticeable amounts of insoluble MnCO3 forming.

blogfast25 - 21-10-2015 at 17:10

Quote: Originally posted by Upsilon  

It would also be interesting to see if the CO2 being formed leads to any noticeable amounts of insoluble MnCO3 forming.


As long as pH < 7 that will not happen, trust me. MnCO3 even has a fairly high K<sub>S</sub>. At low pH the concentration of carbonate ions is basically zero.

[Edited on 22-10-2015 by blogfast25]

Upsilon - 21-10-2015 at 17:51

[Even after this relatively short amount of time, the manganese dioxide has thinned out significantly from an almost jet-black suspension to a more mild brown. This is going better than expected.

Upsilon - 22-10-2015 at 14:34

The solution is getting much clearer now. It is forming a light yellow-brown solution which I find strange - maybe some straggler manganese dioxide molecules in suspension? Regardless I'll be filtering it off soon to see what I get.

Also, while I'm waiting on my HgS, would As2S3 be feasible to practice on? If I can oxidize the sulfur in the As2S3, then it would probably work for HgS, right? The issue might be removing the sulfur from the mixture afterward. I have noticed in the past that sulfur tends to float to the top of a column of water, but I don't think I can depend on this to remove all of it. Perhaps I could burn it out? The problem with that though is that I fear that some of the arsenic trioxide could potentially vaporize since it has a fairly low boiling point - would burning sulfur be hot enough to warrant concern about this?

elementcollector1 - 22-10-2015 at 14:57

Burning sulfur most certainly would be hot enough to warrant your concern when it comes to arsenic fumes. Personally, I'd go for mercury first, as I consider it safer and easier to work with (in elemental form, at least).

blogfast25 - 22-10-2015 at 15:02

Quote: Originally posted by Upsilon  
The solution is getting much clearer now. It is forming a light yellow-brown solution which I find strange - maybe some straggler manganese dioxide molecules in suspension? Regardless I'll be filtering it off soon to see what I get.



MnO2 is often contaminated with Fe2O3. Depends on source and grade, of course.

Upsilon - 22-10-2015 at 15:10

Quote: Originally posted by elementcollector1  
Personally, I'd go for mercury first, as I consider it safer and easier to work with (in elemental form, at least).


Yeah, you're probably right. Im just antsy to see how oxidation of these low-solubility sulfides goes. I've been reading about arsenic compounds and it looks like it's going to be a pain in the arse (pun intended) to pull off safely. The only method in the realm of possibility for me would be to reduce arsenic acid, but reducing it too far produces arsine which I definitely do not want to mess with. I'm not even going to try it for a while yet but I'll be able to figure something out eventually.

Upsilon - 22-10-2015 at 15:27

Quote: Originally posted by blogfast25  

MnO2 is often contaminated with Fe2O3. Depends on source and grade, of course.


That would be it, then. I've had this manganese dioxide for a long time and I don't really remember what grade it was, but I got it from eBay. I wonder if this will affect the purity of the manganese metal; if the basic iron acetate compound melts along with the manganese acetate, then some iron metal would be present in the manganese.

blogfast25 - 22-10-2015 at 16:37

Quote: Originally posted by Upsilon  
I wonder if this will affect the purity of the manganese metal; if the basic iron acetate compound melts along with the manganese acetate, then some iron metal would be present in the manganese.


Most of the iron will not go into solution in your conditions (you will need to filter).

You can judge the final Mn(OAc)2 by its colour. Just try and use it as such, if your idea works you can always prepare pure Mn(+2) acetate later on.

You do realise that from watery solution you will get the tetrahydrate, right? Dehydrating that without hydrolysis and/or oxidation will be nigh impossible. And water is the mortal enemy of your melt electrolysis...


[Edited on 23-10-2015 by blogfast25]

Upsilon - 23-10-2015 at 06:07

Quote: Originally posted by blogfast25  

You do realise that from watery solution you will get the tetrahydrate, right? Dehydrating that without hydrolysis and/or oxidation will be nigh impossible. And water is the mortal enemy of your melt electrolysis...


Manganese(II) acetate can't be dehydrated by heat alone? I was aware that I would be making the hydrate but I thought I would be able to drive out the the water in melting it.

blogfast25 - 23-10-2015 at 06:52

Quote: Originally posted by Upsilon  

Manganese(II) acetate can't be dehydrated by heat alone? I was aware that I would be making the hydrate but I thought I would be able to drive out the the water in melting it.


Firstly, (ionic) acetates are prone to hydrolysis because OAc<sup>-</sup> is the conjugated base of a weak acid: HOAc. So:

OAc<sup>-</sup> + H2O < === > HOAc + OH<sup>-</sup>

And since as HOAc is volatile, you can guess the rest.

Furthermore, Mn(+2) compounds are very prone to air oxidation, especially with heat:

Mn<sup>2+</sup> + 2 H2O === > MnO2 + 4 H+ + 2 e<sup>-</sup>

To successfully dehydrate Mn(OAc)2 hydrate you probably need to boil it to dry from glacial acetic acid, in the absence of air.

Upsilon - 23-10-2015 at 08:47

Well my glacial acetic acid did arrive today so I can actually try that. Would the oxidation of oxalic acid be slower without water, though, since much less H+ is dissolved at one time? In that case, it might be easier to precipitate and dry manganese carbonate out of the solution I made, and then react that with the glacial acetic acid.

blogfast25 - 23-10-2015 at 09:18

Quote: Originally posted by Upsilon  
Well my glacial acetic acid did arrive today so I can actually try that. Would the oxidation of oxalic acid be slower without water, though, since much less H+ is dissolved at one time? In that case, it might be easier to precipitate and dry manganese carbonate out of the solution I made, and then react that with the glacial acetic acid.


Remember:

MnCO3 + 2 HOAc === > Mn(OAc)2 + H2O + CO2.

No escaping that water. But there will be less, so it's not a bad starting point.

Upsilon - 23-10-2015 at 10:03

Quote: Originally posted by blogfast25  

Remember:

MnCO3 + 2 HOAc === > Mn(OAc)2 + H2O + CO2.

No escaping that water. But there will be less, so it's not a bad starting point.


Since manganese(II) acetate forms a tetrahydrate, then based on that equation only a quarter of the product will be hydrated? That's probably good enough to try to electrolyze, since it has been mentioned earlier that manganese can be deposited even in aqueous solution with some difficulty.

blogfast25 - 23-10-2015 at 11:22

Quote: Originally posted by Upsilon  


Since manganese(II) acetate forms a tetrahydrate, then based on that equation only a quarter of the product will be hydrated? That's probably good enough to try to electrolyze, since it has been mentioned earlier that manganese can be deposited even in aqueous solution with some difficulty.


It's not that simple but roughly probably yes.

If you were to analyse such a melt (IN THE STRICT ABSENCE OF OXYGEN), expect the water to electrolyse off first.

Upsilon - 23-10-2015 at 13:44

Well, after filtering off the insolubles I added enough sodium carbonate to react with both the manganese(II) acetate and the excess acetic acid. When I added the sodium carbonate, it did nothing when it hit the solution, so I added a bunch more without waiting. It apparently only started reacting when it hit the bottom of the beaker so I got a large overflow. I imagine this caused a lot of loss. After letting all the sodium carbonate dissolve I filtered out the very small amount of insoluble substance, which is more of an orange-brown instead of the light pink like manganese(II) carbonate should be. It's probably due to impurities in the manganese dioxide but I can't be certain. I'll be trying this again sometime with more concentrated acetic acid.

Meanwhile my HgS has arrived; I'll be testing small amounts this weekend some time. I'm not quite sure where I need to take the waste when I'm done so I'll just bag it up tight for perpetual storage until I figure out what to do with it.

[Edited on 23-10-2015 by Upsilon]

blogfast25 - 23-10-2015 at 14:44

Quote: Originally posted by Upsilon  
Well, after filtering off the insolubles I added enough sodium carbonate to react with both the manganese(II) acetate and the excess acetic acid. When I added the sodium carbonate, it did nothing when it hit the solution, so I added a bunch more without waiting. It apparently only started reacting when it hit the bottom of the beaker so I got a large overflow. I imagine this caused a lot of loss. After letting all the sodium carbonate dissolve I filtered out the very small amount of insoluble substance, which is more of an orange-brown instead of the light pink like manganese(II) carbonate should be. It's probably due to impurities in the manganese dioxide but I can't be certain.


Yup, even the simplest of operations require planning! ;)

Your MnCO3 is Fe(OH)3 contaminated.

Upsilon - 23-10-2015 at 16:02

I just tested some HgS powder with bleach - nothing seemed to happen. There were (very few) small globules of the same color as the HgS floating to the surface; it's possible that these are sulfur coated in HgS, but I'm not for sure - could just be impurities. I imagine it would be reasonably difficult to point out HgO or sulfur among HgS, since they are all some combinatkon of red and yellow. As far as the legitimacy of the HgS goes, I am not completely sure, though just by holding up the bag I can tell that it is very dense.

Once I concentrate some hydrogen peroxide, I'll give that a shot on it.

UPDATE: I am starting to see yellow particles in the test tube, so I think it is working - though probably quite slowly since the bleach is so dilute. What exactly is the reaction here? Possibly this?
HgS + NaClO -> HgO + NaCl + S

[Edited on 24-10-2015 by Upsilon]

blogfast25 - 23-10-2015 at 17:34

ClO<sup>-</sup> + H<sub>2</sub>O + 2 e<sup>-</sup> === > Cl<sup>-</sup> + 2 OH<sup>-</sup>

How to balance redox reactions:

http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochem...

Upsilon - 23-10-2015 at 18:17

Quote: Originally posted by blogfast25  
ClO<sup>-</sup> + H<sub>2</sub>O + 2 e<sup>-</sup> === > Cl<sup>-</sup> + 2 OH<sup>-</sup>

How to balance redox reactions:

http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochem...


Thanks for that, I had completely forgotten about these methods for determining the outcome of a redox reaction. I guess the question remains, which anion is paired with Hg2+? Hg(OH)2 rapidly disproportionates into HgO and H2O much like AgOH.

To try and figure it out I looked towards the reaction of H2S and NaClO. Certainly, HCl is not produced in this reaction, so it should be as follows:
H2S + NaClO -> NaCl + S + H2O

So then mercury(II) oxide must be formed if it conforms to the H2S reaction.

Looking at HgO on Wikipedia, it claims that it reacts violently with reducing agents. It does not give any more detail than this, but I assume it means that Hg2+ will be reduced either to elemental mercury or Hg22+(I suspect it is this one since this reaction has the higher E° value):
2HgO + H2C2O4 -> Hg2O + 2CO2 + H2O

Hg2O then rapidly disproportionates into HgO and elemental Hg:
Hg2O -> Hg + HgO

[Edited on 24-10-2015 by Upsilon]

blogfast25 - 24-10-2015 at 05:52

Quote: Originally posted by Upsilon  

Looking at HgO on Wikipedia, it claims that it reacts violently with reducing agents. It does not give any more detail than this, but I assume it means that Hg2+ will be reduced either to elemental mercury or Hg22+(I suspect it is this one since this reaction has the higher E° value):
2HgO + H2C2O4 -> Hg2O + 2CO2 + H2O

Hg2O then rapidly disproportionates into HgO and elemental Hg:
Hg2O -> Hg + HgO

[Edited on 24-10-2015 by Upsilon]


I'm not sure where the oxalic acid comes into it?

Upsilon - 24-10-2015 at 07:29

Quote: Originally posted by blogfast25  
[
I'm not sure where the oxalic acid comes into it?


As a reductant. Quote from Wikipedia:
"Mercury(II) oxide reacts violently with reducing agents, chlorine, hydrogen peroxide, magnesium (when heated), disulfur dichloride and hydrogen trisulfide"

The reaction with oxalic acid is just a guess and I have no idea if it will actually work, but I don't really see a reason why it shouldn't if it really does react so vigorously with reducing agents.

blogfast25 - 24-10-2015 at 08:30

Quote: Originally posted by Upsilon  

The reaction with oxalic acid is just a guess and I have no idea if it will actually work, but I don't really see a reason why it shouldn't if it really does react so vigorously with reducing agents.


If there's a possibility of a violent reaction be very careful, given the nature of Hg and its compounds. Nothing is worth risking life and limb for.

[Edited on 24-10-2015 by blogfast25]

Upsilon - 24-10-2015 at 10:23

Quote: Originally posted by blogfast25  
If there's a possibility of a violent reaction be very careful, given the nature of Hg and its compounds. Nothing is worth risking life and limb for.


Before I can even try it I'm probably going to need a better oxidizer. I suppose bleach will eventually work but I'd need a lot of it and a lot of time since it's so dilute. I probably won't try hydrogen peroxide after all; on top of the "violent reaction" I have also read that the reaction of HgO and hydrogen peroxide forms an explosive compound. I'll probably try nitric acid and/or nitrogen dioxide but I won't be able to make either for a little while.

EDIT: Something I may do is experiment with lead(II) sulfide and getting solid results there before working too much with HgS - it will probably behave similarly and will at least be moderately less toxic, as well as being cheaper.

[Edited on 24-10-2015 by Upsilon]

blogfast25 - 24-10-2015 at 10:45

You could try 'Poolshock': solid calcium hypochlorite. Very OTC.

Upsilon - 25-10-2015 at 16:24

Quote: Originally posted by blogfast25  
You could try 'Poolshock': solid calcium hypochlorite. Very OTC.


Sounds promising. Looks like the OTC chemical contains a significant amount of calcium chloride but since CaCl2 is a product of the planned reaction then it shouldn't matter. I still would like some lead(II) sulfide to practice on, though. I have a broken lead-acid battery that I will be retrieving lead(II) nitrate from using nitric acid on its contents. To make PbS from this, I could use hydrogen sulfide, but I'd prefer not to. I'm thinking that I could create Na2S/NaHS by heating sulfur in sodium hydroxide solution, then adding the lead nitrate afterward to precipitate PbS?

blogfast25 - 25-10-2015 at 17:38

PbS wouldn't be the worst 'model', as it too is very insoluble.

Hypochlorite will kick Pb(+2) to (+4) I think (check the SRPs), that could be a complication.

woelen - 26-10-2015 at 02:38

If you want to create elements from Hg-compounds or As-compounds, stick to aqueous chemistry. Do not use hot chemistry. Breathing volatile Hg or As-oxides is something you must avoid at any cost.

Arsenic can be made in elemental form, using aqueous chemistry, but this is not straightforward. I myself have experimented with As2O5, which fairly easily can be reduced to As2O3 or arsenious acid, but the step to elemental arsenic is not easy. One safe method (no risk of formation of AsH3) is reduction with hypophosphite (which is a moderately strong reductor), but the reaction is slow and yield is low, resulting in a black powder of pure As. Using stronger reductors, such as NaBH4 or powdered zinc is very dangerous. You get quite a lot of hydrogen gas, contaminated with AsH3 as side product.

As2S3 can be converted to arsenic acid and sulfate by heating in conc. HNO3. The arsenic acid in turn can be reduced to arsenious acid and arsenic. The oxidation of As2S3 by nitric acid must be done outside on a breezy day! You get fumes of NOx and tiny droplets of liquid, which may contain dissolved arsenic acid. Loosely cover the beaker with a paper tissue. Gas can escape easily, but most tiny droplets are trapped by the paper.

Upsilon - 26-10-2015 at 14:12

Quote: Originally posted by woelen  

Arsenic can be made in elemental form, using aqueous chemistry, but this is not straightforward. I myself have experimented with As2O5, which fairly easily can be reduced to As2O3 or arsenious acid, but the step to elemental arsenic is not easy. One safe method (no risk of formation of AsH3) is reduction with hypophosphite (which is a moderately strong reductor), but the reaction is slow and yield is low, resulting in a black powder of pure As. Using stronger reductors, such as NaBH4 or powdered zinc is very dangerous. You get quite a lot of hydrogen gas, contaminated with AsH3 as side product.


Would it be possible to use tin(II) chloride as a reductant? SRPs make it look like it's feasible:
H3AsO3 + 3H+ + 3e- -> As + 3H2O (+0.24V)
Sn4+ + 2e- -> Sn2+ (+0.15V)
As + 3H+ + 3e- -> AsH3 (-0.23V)

A redox reaction theoretically occurs containing the first two half reactions, oxidizing tin(II) to tin(IV):
2H3AsO3 + 6H+ + 3Sn2+ -> 2As + 3H2O + 3Sn4+ (0.09V)

While the reduction of As to AsH3 shouldn't occur:
2As + 6H+ + 3Sn2+ -> 2AsH3 + 3Sn4+ (-0.38V)

So, overall:
2H3AsO3 + 6HCl + 3SnCl2 -> 3SnCl4+ 2As + 3H2O

Thoughts on this? It will probably be quite slow, but I would at least not have to buy a hypophosphite.

EDIT: Another possible reductant could be formaldehyde:
H2CO2 + 2H+ + 2e- -> H2CO + H2O (-0.03V)

2H3AsO3 + 6H+ + 2H2CO + 3H2O -> 2As + 6H2O + 3H2CO2 + 6H+ (+0.27V)

Simplified to:
2H3AsO3 + 3H2CO -> 2As + 3H2O + H2CO2


[Edited on 26-10-2015 by Upsilon]

[Edited on 26-10-2015 by Upsilon]

blogfast25 - 26-10-2015 at 14:45

Quote: Originally posted by Upsilon  

So, overall:
2H3AsO3 + 6HCl + 3SnCl2 -> 3SnCl4+ 2As + 3H2O

Thoughts on this? It will probably be quite slow, but I would at least not have to buy a hypophosphite.



Yes, I believe that should work.

There's something much simpler though: deposition of As (+3) or (+5) from acid solution as As(0) on copper wire or strip. This is still used as a forensic screening test for arsenic.

EC1 did that IIRW and there's a report on it.

Upsilon - 26-10-2015 at 15:06

Quote: Originally posted by blogfast25  


There's something much simpler though: deposition of As (+3) or (+5) from acid solution as As(0) on copper wire or strip. This is still used as a forensic screening test for arsenic.

EC1 did that IIRW and there's a report on it.


What exactly will allow these arsenic compounds to dissolve to form As3+ or As5+? As2O5 is apparently pretty soluble and can be made by the action of HNO3 on As2O3.

The use of formaldehyde seems pretty appealing to me, however (see edit above). I will theoretically get arsenous acid directly from the reaction of As2S3 and hypochlorite. The formaldehyde would then be added, leaving nothing with the arsenic but formic acid and water, which can easily be evaporated off.

blogfast25 - 26-10-2015 at 15:30

Quote: Originally posted by Upsilon  

What exactly will allow these arsenic compounds to dissolve to form As3+ or As5+? As2O5 is apparently pretty soluble and can be made by the action of HNO3 on As2O3.



Arsenic is one of these elements that's neither a metal nor a non-metal. It's also amphoteric.

woelen - 27-10-2015 at 04:27

I actually tried using tin(II) for making elemental arsenic. It does work, but the arsenic you get is very impure and I could not purify it without putting myself in high risk.

I dissolved SnCl2 in conc. HCl and added this to a solution of As2O5. When this is done you immediately get a dirty brown precipitate, which is a very intimate mix (solid solution?) of As and basic tin(IV) chloride. I read about this process and the only practical way of purification of the As is by sublimation. This is something which I do not want to do, I value my health too much.
The route with hypophosphite is much slower and the arsenic is formed as a compact fine crystalline powder. The arsenic looks black. With hypophosphite, however, the reaction only works reasonably fast in the presence of a lot of chloride ion (as HCl) and this adds a large risk: I saw tiny droplets of oily liquid on the surface, which must be AsCl3. When I saw the AsCl3 I immediately quit the experimenting (AsCl3 is volatile, is quite stable in the presence of humidity and acid, much more so than PCl3, and EXTREMELY poisonous). So, know what you are doing, work in well closed reaction vessels, a beaker is not OK.

You'll find that making elemental arsenic always is done through high temperature chemistry and sublimation of the element. Apparently there is no satisfying aqueous method with acceptable yield. The best I found is the hypophosphite route with yields of atmost a few percent. It is suitable for making a small sample of powdered arsenic when done carefully.

Upsilon - 27-10-2015 at 07:55

What about the possibility of formaldehyde as a reductant? Reposting from above:
H3AsO3 + 3H+ + 3e- -> As + 3H2O (+0.24V)
H2CO2 + 2H+ + 2e- -> H2CO + H2O (-0.03V)

Balancing for electrons, the protons conveniently cancel out:
2H3AsO3 + 6H+ + 2H2CO + 3H2O -> 2As + 6H2O + 3H2CO2 + 6H+ (+0.27V)

Simplified to:
2H3AsO3 + 3H2CO -> 2As + 3H2O + H2CO2

[Edited on 27-10-2015 by Upsilon]

blogfast25 - 27-10-2015 at 09:00

Quote: Originally posted by Upsilon  
What about the possibility of formaldehyde as a reductant? Reposting from above:
H3AsO3 + 3H+ + 3e- -> As + 3H2O (+0.24V)
H2CO2 + 2H+ + 2e- -> H2CO + H2O (-0.03V)

Balancing for electrons, the protons conveniently cancel out:
2H3AsO3 + 6H+ + 2H2CO + 3H2O -> 2As + 6H2O + 3H2CO2 + 6H+ (+0.27V)

Simplified to:
2H3AsO3 + 3H2CO -> 2As + 3H2O + H2CO2

[Edited on 27-10-2015 by Upsilon]


There's no shortage of reducers that will yield As(0) but all will yield powders, except Cu sheet/strips.

Upsilon - 27-10-2015 at 13:08

Quote: Originally posted by blogfast25  

There's no shortage of reducers that will yield As(0) but all will yield powders, except Cu sheet/strips.


I suppose that is true; I wouldn't really mind having a powder too much instead of a lump, but might as well go for the lump if it's not any more difficult. I'll get around to it when I'm ready.

I would like to discuss chromium next. I have chromium(III) oxide, but it is quite unreactive to both acids and bases. I did a small test thermite with it, and it was quite slow, did not burn to completion, and did not form nice pieces of metal (probably not hot enough to melt the chromium nicely). Will using a booster reaction provide better results? I figure that if I can at least get it to burn mostly to completion, then I could just dissolve everything in HCl an electrolyze the solution - chromium will be deposited but aluminum will not.

blogfast25 - 27-10-2015 at 13:27

Quote: Originally posted by Upsilon  

I would like to discuss chromium next. I have chromium(III) oxide, but it is quite unreactive to both acids and bases. I did a small test thermite with it, and it was quite slow, did not burn to completion, and did not form nice pieces of metal (probably not hot enough to melt the chromium nicely). Will using a booster reaction provide better results? I figure that if I can at least get it to burn mostly to completion, then I could just dissolve everything in HCl an electrolyze the solution - chromium will be deposited but aluminum will not.


The chromium thermite works very well with small amounts of booster like NaNO3 or K2Cr2O7 and provided you use at least 50 g (or so) of mixture. I have some nice 'large' round reguli from that. Always include a small amount of slag fluidizer (CaF2 or CaO) to improve yield.

Of course you can also electroplate Cr(0) from aqueous solutions, see the electrolytic chroming of objects.

Cr2O3 is usually totally unresponsive to HCl though...


[Edited on 27-10-2015 by blogfast25]

Upsilon - 27-10-2015 at 13:42

Quote: Originally posted by blogfast25  

The chromium thermite works very well with small amounts of booster like NaNO3 or K2Cr2O7 and provided you use at least 50 g (or so) of mixture. I have some nice 'large' round reguli from that.

Of course you can also electroplate Cr(0) from aqueous solutions, see the electrolytic chroming of objects.

Cr2O3 is usually totally unresponsive to HCl though...

[Edited on 27-10-2015 by blogfast25]


I wasn't intending to try and dissolve the Cr2O3 in HCl - I was referring to using it on the aftermath of the thermite if I could not get nice samples of Cr. The small grains of Cr would be dissolved in HCl along with unreacted aluminum and aluminum oxide, then Cr could be electroplated while aluminum would not be. However, since you say that the thermite alone can provide good results I may not do this; though if I am interested in further purifying the metal from the slag then I may go through with the electroplating anyways.

blogfast25 - 27-10-2015 at 15:55

Quote: Originally posted by Upsilon  
However, since you say that the thermite alone can provide good results I may not do this; though if I am interested in further purifying the metal from the slag then I may go through with the electroplating anyways.


In theory thermites can provide very pure metals, if you start from very pure oxide and Al powder and get the formulation right.

For very purely plated metal, start from a very pure water soluble Cr compound,

j_sum1 - 27-10-2015 at 16:39

I did a thermite with Cr2O3 and Al not long ago. I got a mixture of metallic Cr and Al2O3 for my efforts. Inseparable by mechanical means, and not going to do anything favourable with any acids. The lumps look kind of cool though.

I may take up bloggers' suggestion and add a nitrate booster.

Upsilon - 27-10-2015 at 17:03

Quote: Originally posted by j_sum1  
not going to do anything favourable with any acids


What do you mean by that? Treatment with HCl should dissolve everything but unreacted Cr2O3, then Cr3+ can be reduced to Cr by electrical means while Al3+ won't be reduced.

[Edited on 28-10-2015 by Upsilon]

blogfast25 - 27-10-2015 at 17:20

Quote: Originally posted by j_sum1  
I did a thermite with Cr2O3 and Al not long ago. I got a mixture of metallic Cr and Al2O3 for my efforts.

I may take up bloggers' suggestion and add a nitrate booster.


What can I say? Many feel called but few are chosen! ;)

Seriously though, getting aluminothermic reductions right is not as easy as it sounds. If I had a penny for each one that fails due to poor execution...

1) use right stoichiometry, including boosters if needed.

2) always use a slag fluidizer.

3) larger thermites tend to give much better yields.

Just a few rules of thumb that can go far...

elementcollector1 - 27-10-2015 at 18:52

Say, I've got a suddenly relevant question: I did a Vanadium thermite a few months back, and used the metal as my V element sample. However, I remember some thread I started a few years back stating that V/Al thermites contaminated the produced V with up to 10% Al. I know I starved the mix of Al when I lit it - did this help in any way?

blogfast25 - 27-10-2015 at 19:32

Quote: Originally posted by elementcollector1  
I know I starved the mix of Al when I lit it - did this help in any way?


Yes.

ALL chemical reactions have to be seen as equilibrium reactions, as is this one:

V2O5 + 10/3 Al < ==== > 2 V + 5/3 Al2O3

By increasing V2O5 in the mix you push the equilibrium to the right, thereby lowering the Al content in the metal obtained.

The process of adding extra target oxide to the mix is known in aluminothermic jargon as slagging. As I wrote above, aluminothermy has the potential to produce highly pure target metals. Slagging is one measure that contributes to it.

Slagging does have one disadvantage: it lowers the end-temperature somewhat, as the extra oxide also needs to be heated up.

[Edited on 28-10-2015 by blogfast25]

j_sum1 - 27-10-2015 at 19:34

Quote: Originally posted by Upsilon  
Quote: Originally posted by j_sum1  
not going to do anything favourable with any acids


What do you mean by that? Treatment with HCl should dissolve everything but unreacted Cr2O3, then Cr3+ can be reduced to Cr by electrical means while Al3+ won't be reduced.

[Edited on 28-10-2015 by Upsilon]

If I want Cr(0) and I don't want Al2O3, then I am not going to get there with my mix and an acid.

Upsilon - 27-10-2015 at 20:23

Quote: Originally posted by j_sum1  

If I want Cr(0) and I don't want Al2O3, then I am not going to get there with my mix and an acid.


You said that you were getting Cr metal that was highly contaminated with Al2O3. Putting the whole lot of this in acid will leave you with an AlCl3 and CrCl3 mixed solution:
2Cr + 6HCl -> 3H2 + 2CrCl3
Al2O3 + 6HCl -> 2AlCl3 + 3H2O
The mixed solution could then be electrolyzed, since Cr3+ can be reduced in aqueous solution while Al3+ cannot.

j_sum1 - 27-10-2015 at 21:27

I don't doubt I could do some electrolysis. However, at the moment I have some metallic Cr which I got from Cr(III). Acids are not going to purify my mixture without going back to Cr(III).
I also don't think that HCl will have a lot of effect on Al2O3 in any reasonable time frame. This link suggests 110 days at elevated temperatures.

The goal is samples for the element collection.
I have a nice lump I purchased.
I have some electrolysis samples and will do more.
I don't have a good piece obtained via thermite. Current sample has the alumina contamination.

Acids will not get me from where I am to what I want.

Upsilon - 28-10-2015 at 15:42

Quick question - what are these "complexed selenium" pills that I am seeing everywhere? I cannot find any sort of formulation for them, but if they contain selenium, then it would be interesting to have a go at extracting it.

blogfast25 - 28-10-2015 at 16:28

Quote: Originally posted by Upsilon  
Quick question - what are these "complexed selenium" pills that I am seeing everywhere? I cannot find any sort of formulation for them, but if they contain selenium, then it would be interesting to have a go at extracting it.


As with all these food supplements they contain very little of the advertised active ingredient.

For Se, that's a good thing, as otherwise pill poppers would suffer constant garlic breath! :D

https://en.wikipedia.org/wiki/Selenium#Toxicity

And 'complexed' is likely to be a marketing buzzword, rather than an objective, scientific description.

[Edited on 29-10-2015 by blogfast25]

Upsilon - 28-10-2015 at 18:03

Quote: Originally posted by blogfast25  


As with all these food supplements they contain very little of the advertised active ingredient.

For Se, that's a good thing, as otherwise pill poppers would suffer constant garlic breath! :D

https://en.wikipedia.org/wiki/Selenium#Toxicity

And 'complexed' is likely to be a marketing buzzword, rather than an objective, scientific description.

[Edited on 29-10-2015 by blogfast25]


Yeah I realized that not long after posting - most of these capsules contain only 200 ug of selenium. If I wanted a 1g selenium sample, I would need 5000 capsules. At around $4 USD for a bottle of 100 capsules, I would need 50 bottles which is $200 USD! I'm starting to grudgingly accept that selenium may just be one of those elements I have to buy - selenium compounds seem to be ridiculously expensive while samples cost only a few bucks - there's just no competition between the two.

elementcollector1 - 28-10-2015 at 18:10

Quote: Originally posted by Upsilon  

Yeah I realized that not long after posting - most of these capsules contain only 200 ug of selenium. If I wanted a 1g selenium sample, I would need 5000 capsules. At around $4 USD for a bottle of 100 capsules, I would need 50 bottles which is $200 USD! I'm starting to grudgingly accept that selenium may just be one of those elements I have to buy - selenium compounds seem to be ridiculously expensive while samples cost only a few bucks - there's just no competition between the two.


Fear not, young element collector, for I hath found thy solution!
https://en.wikipedia.org/wiki/Selenium_rectifier
This has been on my to-do list for quite a while - dissolve away everything else, and you're left with a few thin plates of pure Se.

Upsilon - 28-10-2015 at 18:26

Quote: Originally posted by elementcollector1  


Fear not, young element collector, for I hath found thy solution!
https://en.wikipedia.org/wiki/Selenium_rectifier
This has been on my to-do list for quite a while - dissolve away everything else, and you're left with a few thin plates of pure Se.


Interesting, I wonder just how much you could get out of one of those things? Too bad they're obsolete now and probably not as available as they used to be.

EDIT: Ah! I hadn't even thought about selenates yet! 20g of sodium selenate will run about $25 USD which is more than reasonable. I'll keep this option on the table as well.

[Edited on 29-10-2015 by Upsilon]
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