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j_sum1
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You are on your own with the mercury. I know almost nothing.
Handle safely and all that.
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BobD1001
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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  It's not really a big deal, it's barely got a gray
tinge to it, but I have the urge to clean it before using it.
Also I should note that my 50 grams of cinnabar (HgS) powder should be coming in soon. Wikipedia outlines roasting it in air and condensing the
mercury vapor...but uh, mercury vapor, not gonna happen 
Instead I will try reacting it with a concentrated acid and electrolysing the aqueous corresponding salt. HgCl2 isn't terribly soluble in water, so I
may use nitric acid instead of HCl. I don't know how soluble mercury nitrate is but nitrates are typically more soluble than chlorides so it's worth a
shot. |
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.
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elementcollector1
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In fact, that write-up just so happens to be stickied in the Chemistry in General subforum.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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Upsilon
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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
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blogfast25
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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]
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Upsilon
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| 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, |
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blogfast25
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| 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]
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Upsilon
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| 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]
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blogfast25
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Have a look at this, if you haven't already:
http://www.sciencemadness.org/talk/viewthread.php?tid=18162&...
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Upsilon
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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]
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blogfast25
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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]
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Upsilon
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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).
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blogfast25
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| Quote: Originally posted by Upsilon |
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?
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Upsilon
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| 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?
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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.
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blogfast25
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BTW, years ago I tried reducing MnO2 with bisulphite + acid too and it didn't work for me either.
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Upsilon
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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.
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blogfast25
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| Quote: Originally posted by Upsilon |
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.
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UC235
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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.
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Upsilon
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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.
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Upsilon
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| 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 those cheapie brown paper towels worked extremely well for filtering MnO2 in Nurdrage's videos, I might pick some
of those up.
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Upsilon
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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.
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gdflp
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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.
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Upsilon
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| 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.
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Upsilon
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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.
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blogfast25
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| 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]
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