Sciencemadness Discussion Board
Not logged in [Login ]
Go To Bottom

Printable Version  
 Pages:  1  ..  4    6    8  ..  14
Author: Subject: Quest for the elements
j_sum1
Administrator
********




Posts: 5641
Registered: 4-10-2014
Location: Oz
Member Is Offline

Mood: d jolly / dt > 0

[*] posted on 19-10-2015 at 18:41


You are on your own with the mercury. I know almost nothing.
Handle safely and all that.
View user's profile View All Posts By User
BobD1001
Hazard to Others
***




Posts: 182
Registered: 29-3-2013
Member Is Offline

Mood: No Mood

[*] posted on 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.
View user's profile Visit user's homepage View All Posts By User
elementcollector1
International Hazard
*****




Posts: 2684
Registered: 28-12-2011
Location: The Known Universe
Member Is Offline

Mood: Molten

[*] posted on 19-10-2015 at 20:08


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
View user's profile View All Posts By User
Upsilon
National Hazard
****




Posts: 392
Registered: 6-10-2013
Member Is Offline

Mood: No Mood

[*] posted on 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
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10340
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 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]




View user's profile Visit user's homepage View All Posts By User
Upsilon
National Hazard
****




Posts: 392
Registered: 6-10-2013
Member Is Offline

Mood: No Mood

[*] posted on 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,
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10340
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 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]




View user's profile Visit user's homepage View All Posts By User
Upsilon
National Hazard
****




Posts: 392
Registered: 6-10-2013
Member Is Offline

Mood: No Mood

[*] posted on 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]
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10340
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 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&...




View user's profile Visit user's homepage View All Posts By User
Upsilon
National Hazard
****




Posts: 392
Registered: 6-10-2013
Member Is Offline

Mood: No Mood

[*] posted on 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]
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10340
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 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]




View user's profile Visit user's homepage View All Posts By User
Upsilon
National Hazard
****




Posts: 392
Registered: 6-10-2013
Member Is Offline

Mood: No Mood

[*] posted on 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).
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10340
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 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?




View user's profile Visit user's homepage View All Posts By User
Upsilon
National Hazard
****




Posts: 392
Registered: 6-10-2013
Member Is Offline

Mood: No Mood

[*] posted on 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.
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10340
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 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.



View user's profile Visit user's homepage View All Posts By User
Upsilon
National Hazard
****




Posts: 392
Registered: 6-10-2013
Member Is Offline

Mood: No Mood

[*] posted on 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.
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10340
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 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.




View user's profile Visit user's homepage View All Posts By User
UC235
International Hazard
*****




Posts: 545
Registered: 28-12-2014
Member Is Offline

Mood: No Mood

[*] posted on 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.
View user's profile View All Posts By User
Upsilon
National Hazard
****




Posts: 392
Registered: 6-10-2013
Member Is Offline

Mood: No Mood

[*] posted on 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.
View user's profile View All Posts By User
Upsilon
National Hazard
****




Posts: 392
Registered: 6-10-2013
Member Is Offline

Mood: No Mood

[*] posted on 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.
View user's profile View All Posts By User
Upsilon
National Hazard
****




Posts: 392
Registered: 6-10-2013
Member Is Offline

Mood: No Mood

[*] posted on 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.
View user's profile View All Posts By User
gdflp
Super Moderator
*******




Posts: 1320
Registered: 14-2-2014
Location: NY, USA
Member Is Offline

Mood: Staring at code

[*] posted on 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.



View user's profile View All Posts By User
Upsilon
National Hazard
****




Posts: 392
Registered: 6-10-2013
Member Is Offline

Mood: No Mood

[*] posted on 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.
View user's profile View All Posts By User
Upsilon
National Hazard
****




Posts: 392
Registered: 6-10-2013
Member Is Offline

Mood: No Mood

[*] posted on 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.
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10340
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 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]




View user's profile Visit user's homepage View All Posts By User
 Pages:  1  ..  4    6    8  ..  14

  Go To Top