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Author: Subject: Quest for the elements
Upsilon
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[*] posted on 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.
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[*] posted on 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?
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[*] posted on 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).



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[*] posted on 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.




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[*] posted on 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.
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[*] posted on 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.
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[*] posted on 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]




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[*] posted on 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.
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[*] posted on 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.





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[*] posted on 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.
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[*] posted on 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.




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[*] posted on 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.
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[*] posted on 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.




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[*] posted on 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]
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[*] posted on 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.




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[*] posted on 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]
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[*] posted on 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...




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[*] posted on 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]
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[*] posted on 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?




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[*] posted on 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.
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[*] posted on 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]




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[*] posted on 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]
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[*] posted on 24-10-2015 at 10:45


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



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[*] posted on 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?
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[*] posted on 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.




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