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Upsilon
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[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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Upsilon
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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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elementcollector1
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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).
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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blogfast25
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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.
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MnO2 is often contaminated with Fe2O3. Depends on source and grade, of course.
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Upsilon
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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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Upsilon
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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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blogfast25
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| 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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Upsilon
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| 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...
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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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blogfast25
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| 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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Upsilon
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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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blogfast25
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| 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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Upsilon
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| 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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blogfast25
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| 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.
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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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Upsilon
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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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blogfast25
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| 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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Upsilon
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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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blogfast25
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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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Upsilon
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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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blogfast25
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| 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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Upsilon
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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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blogfast25
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| 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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Upsilon
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| 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.
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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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blogfast25
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You could try 'Poolshock': solid calcium hypochlorite. Very OTC.
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Upsilon
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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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blogfast25
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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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