peach
Bon Vivant
Posts: 1428
Registered: 14-11-2008
Member Is Offline
Mood: No Mood
|
|
Mercury Sulfate and decomposition
I've heard that the sulfate will decompose to Hg2SO4 on contact with water. I've searched around but not found much more information on that.
But I have something of a conundrum.
I tried running the Preparation Of Mercury Salts writeup by ZyGoat a few years ago. I don't plan to rerun it, but I have had a niggling itch about it
ever since; something I can't quite fully explain occurred, which annoys me.
Filtering the sulfate was, understandably, difficult. I also doubt it was even necessary; there'd only be some unreacted mercury left in there and
it'd be easier to remove at the chloride stage.
To pass the material through the paper, I had to pour around a liter of boiling water over 10g to dissolve it all. The sulfate seemed to pass through,
eventually. As soon as it cooled from boiling on the other side of the paper, it precipitated almost immediately.
After neutralizing, I obtained the dirty orange colored oxide.
I then treated it with HCl and obtained the white precipitate, the chloride salt.
From my memory, I seem to remember getting a pathetic yield of somewhere between 15 and 25%.
I'm now left with some questions.
I seem to have obtained at least some percentage of either Hg2Cl2 or HgCl2, the material was certainly corrosive, freely precipitated as the acid went
in and settled rapidly; meaning it's definitely not NaHSO4 or NaCl.
My main wondering is about the mercury sulfate decomposing in water. Taking this into account, there are two possibilities I can think of for how I
managed to obtain a chloride salt.
Firstly, the decomposition isn't complete, even in an excess of boiling water. So I could have still produced the oxide and then the chloride from the
small percentage of HgSO4 that didn't decompose when exposed.
Secondly, it did fully decompose and the material I collected was actually Hg2SO4. Which then makes me wonder, does this also go to the oxide on
neutralization? If so, perhaps my small yield was due to it not converting so efficiently to the oxide as the HgSO4 would.
I'm wondering if the decomposition only occurs (or gets closer to completion) on boiling the aqueous HgSO4 solution. I'm sure I've read somewhere that
the decomposition occurs when you try to dissolve (heat) it, not necessarily spontaneously on contact with cold water. I noticed not_important and
turd mentioned that it requires heat in a similar thread.
I did perform my neutralization using an accurate pH probe, and decanted the aqueous phase from the oxide; removing any NaHSO4 or surplus base (and
there won't have been much of that anyway given the accuracy with which I stopped after neutralization). I stopped my HCl addition once it was stable
just on the acidic side of neutral; meaning I didn't have an excess of acid present but had fully converted whatever oxide was present to the
chloride.
[Edited on 19-5-2010 by peach]
|
|
|
turd
National Hazard
Posts: 800
Registered: 5-3-2006
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by peach  | | I've heard that the sulfate will decompose to Hg2SO4 on contact with water. I've searched around but not found much more information on that.
|
First of all: bad style. "I've heard..." usually translates to "I made something up". Personally, I find it implausible that people constantly
remember facts they heard, but cannot remember where they did so.
Apart from that: This is of course nonsense. Treatment of HgSO4 with water will give some dissolution and some basic Hg(II) salts like
(HgSO4)2.Hg(OH)2.2H2O and ultimately HgSO4.HgO. Evidently it's an pH/equilibrium thing. If you want to make a HgSO4 solution slowly add sulfuric acid
until most is dissolved and filter off the remaining insolubles.
There is of course also a Hg(2+) + Hg <--> Hg2(2+) redox equilibrium, but this is not something you will encounter by when treating HgSO4 with
water.
Edit: How do you know you got Hg2SO4? What kind of colour did it have? I've made HgCl2 via a similar method and never had problems with Hg2(2+)-salts.
Edit2: If of course you still had metallic Hg in your reaction, than you could end up with a slow Hg + HgSO4 --> Hg2SO4 reaction. Obviously the
metallic Hg has to be removed.
[Edited on 19-5-2010 by turd]
[Edited on 19-5-2010 by turd]
|
|
|
not_important
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
The decomposition of HgSO4 is faster when hot and dilute, and reaches a stable stage with the formation of basic sulfate of mercury, HgSO4.2HgO. once
known as turpeth mineral or occasionally as yellow sulphate. In more concentrated solution the sulfuric acid formed reverses the reaction, leaving
some of the HgSO4 undecomposed and in solution. Prolonged boiling with large amounts of water will take it to the oxide, but rather slowly. No
reduction to Hg(I) / Hg2(2+) happens.
The basic sulfate when reacted with concentrated hydrochloric acid gives some HgCl2 : HgSO4.2HgO + HCl => HgCl2 + H2SO4 + 2HgO The oxide is not
quick to react with HCl, so you can get a ppt of it that contains about 2/3 of the mercury that was in the sulfate if that had completely hydrolyised
to the basic sulfate. No ppt of HgO with an excess of HCl means all the mercury is in solution as the chloride. You can always add a little chlorine
water to the solution to insure that there is no mercurous salts.
Most preparations of HgSO4 have you test for mercurous salts and continue boiling until that test is negative, or test and add HNO3 and boil. See
The Dispensatory of the United States of America by George Bacon Wood and Franklin Bache, and Thorp's Inorganic Chemical
Preparations (get the Microsoft scan at the Internet Archives, the Google one was quite poor) Both are on-line copyright free.
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
I have some HgSO4. This is a white solid, which as soon as it comes in contact with water forms a yellow compound. This yellow compound probably is
basic mercuric sulfate or a mix of mercuric sulfate and HgO. This is not a slow change, but it is immediate. The solid does not dissolve appreciably
in plain cold water.
When the solid is added to 20% H2SO4, then it dissolves, giving a colorless solution.
I also have Hg2(NO3)2 and when this is added to distilled water, then it dissolves, giving a colorless and perfectly clear solution. No formation of a
grey or black precipitate, just a colorless solution. So, mercury(I) ions can exist in aqueous solution and do not disproportionate, nor do they
hydrolyse.
|
|
|
peach
Bon Vivant
Posts: 1428
Registered: 14-11-2008
Member Is Offline
Mood: No Mood
|
|
A big thank you for the quick and detailed replies, they're really helpful! Some of you might want to edit the wiki articles on the salts to add a
little more detail about the decomposition; I edit wiki quite often but you obviously know more about what you're talking about on this topic.
| Quote: | | First of all: bad style. "I've heard..." usually translates to "I made something up". Personally, I find it implausible that people constantly
remember facts they heard, but cannot remember where they did so. |
Calm down turd!  That was a massive leap of unguided prediction, incorrect and
didn't need saying. I didn't reference you to the documents because you've already mentioned decomposition in your own posts on the sulfate and I
mentioned those posts in my original question, so what would be the point of referencing you to things you've already seen? The things I'd have been
referencing you to don't contain the information you've given. I knew that, so I asked you. I'm trying to say, I already know you know. I also didn't
say they were facts, the whole point of the post was that I was questioning them. Still, thank you for your help! I do appreciate it takes time and a
lot of motivation to be bothered replying to things like this given the attitudes of so many interested in mercury salts.
I'm genuinely not planning to rerun this or asking about it because I have a big batch of ketone waiting, I'm just annoyed I couldn't work out exactly
what had happened when I ran it a few years ago. At school, we were studying pKa's I think and I said to my friend "I just don't how ... works!", his
reply was "Don't bother trying to understand, just revise it and pass the exams". He's now a doctor and will be a surgeon very shortly. At the same
age, I was routinely taking appliances apart and breaking them in the process just to find out what was inside; the boxes had secrets inside, and I
wanted to hear all of them.
Edit: I forgot to say, my sulfate was white in the concentrated H2SO4. It turned into a solid yellow, very fine precipitate as the water was added.
Yes, I should have mentioned that; I had someone harassing me to put some panties on and go out for dinner. So I expect I lost most of the yield to
decomposition. I'm guessing that my concentrated sulfuric acid from the sulfate forming stage may have just been enough to squeeze some of the HgSO4
through to the neutralization and HCl stage even after it was diluted down during filtration; resulting in the small but measurable yield. Rather than
rely on watching for the ppt during the HCl addition, I stuck a temperature compensated, accurate pH probe in the flask, put the flask on a stir plate
and then dropped the acid until it I could leave it spinning for five or ten minutes and the pH would be sitting on the acidic side of neutral;
indicating the chloride had been formed. I did the same when I was neutralizing.
[Edited on 19-5-2010 by peach]
|
|
|
not_important
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
The yellow is the basic sulfate or basic sulfates plus oxide, that's the decomposition you'd get. It shouldn't have caused any problems or losses if
you just took it into the basification step.
If you filtered and washed once, there'd be almost no H2SO4 left. As many low solubility mercury compounds are pretty dense, it's relatively easy to
wash by decantation, just using the filter to catch any stray bits that drift along until the final collection of the product.
You also should have at least included a link to the old thread, better perhaps to have posted there.
http://sciencemadness.org/talk/viewthread.php?tid=597
reading that shows several attempts where really bad practices were done, massive excesses of reagents in places they shouldn't be used to that amount
and so on. It's like the n00b who thinks that crystallisation by evaporation means slapping the beaker on a burner on high and heating until dry and
frying. Temps me to dig out the wrench and open the mercury flask, do a step-by-step-don't-do-shortcuts write-up.
|
|
|
peach
Bon Vivant
Posts: 1428
Registered: 14-11-2008
Member Is Offline
Mood: No Mood
|
|
I certainly got rid of the excess H2SO4 and bisulfate with those gigantic aqueous washes.
Decanting is indeed very easy with these salts, given their remarkable densities; filtration is barely necessary.
I took the filtrate through to neutralization and then onto the chloride; carefully and with a pH probe continuously tracking the changes to a high
resolution.
This was a long, long time ago and I was a complete idiot back then. The only other possibility I can think of is that I went to the chloride and then
decanted the aqueous phase to speed up drying. But even then, I knew the chloride was far more soluble in water; I'm not entirely ruling it out as an
idiotic possibility though, given the poor yield.
I would STRONGLY recommend you make that write up. Literally five minutes ago I made another post on a different forum saying someone needs to rewrite
the so commonly referenced ZyGoat workup. I would also recommend you keep it in a scientific tone, such that people who shouldn't be working with
mercury in the first place don't start messing around with it or handing out their reactions for other people to eat. With any luck, it'll only be
themselves who eat the results and discover it hasn't worked out as planned. I didn't know exactly what I was doing when I read the ZyGoat workup, but
I also didn't ever eat the results or pass them on to anyone else. I didn't even bother with a reduction. I'd suggest a big section in bold about how
to wash out mercury salts, underlined and in red, at the very start; combined with some information about how much more toxic the chloride salt is
compared to elemental mercury (which so many people are scared of to begin with). I suspect some people are handing out the results of mercury based
reductions without even bothering to wash them; a truly horrendous possibility. I would consider the existence of those people as you write such an
article.
And yes! Frying = sublimation = mercury induced neurotoxicity or back to elemental mercury. Again, hopefully only for themselves.
I don't wish suffering on anyone. When I say only themselves, I only mean that it'd be worse if they were giving it out to other people.
I know what you're saying about the old thread, I should probably have posted in that, yep; in my wiki editing, I hate redundancy, which this thread
is (if any moderators see it, please splice it into the old thread). I'm chatting with a guy right now who has added a huge amount of base, failed to
decant or filter and then gone straight to HCl; without using indicator paper, let alone a probe.
[Edited on 19-5-2010 by peach]
|
|
|
turd
National Hazard
Posts: 800
Registered: 5-3-2006
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by woelen | | No formation of a grey or black precipitate, just a colorless solution. So, mercury(I) ions can exist in aqueous solution and do not disproportionate,
nor do they hydrolyse. |
Well of course they are stable. Actually you make Hg2(2+) compounds by reacting Hg2+ solutions with metallic Hg. You can for example dissolve Hg in a
Hg(NO3)2 solution. The equilibrium of the redox reaction I gave is typically far to the left. But of course you can also induce disproportionation,
e.g. the famous kalomel reaction: Hg2Cl2 + 2NH3 --> Hg(NH2)Cl + Hg + NH4Cl
@peach: Yah, maybe I'm a little bit rough when I'm in a hurry, but it seems I wasn't totally wrong. Looks like you *did* just make the Hg2SO4 part up.
The yellow thing is not a Hg2(2+) compound and I don't think it has a lot to do with time or heat, just with pH. But maybe you really got some Hg2SO4,
if you let it stand to long over metallic Hg.
|
|
|
peach
Bon Vivant
Posts: 1428
Registered: 14-11-2008
Member Is Offline
Mood: No Mood
|
|
| Quote: | | @peach: Yah, maybe I'm a little bit rough when I'm in a hurry, but it seems I wasn't totally wrong. Looks like you *did* just make the Hg2SO4 part up.
The yellow thing is not a Hg2(2+) compound and I don't think it has a lot to do with time or heat, just with pH. But maybe you really got some Hg2SO4,
if you let it stand to long over metallic Hg. |
Wires have become crossed turd! You're entirely right that I made my own gigantic leap to assuming Hg2SO4. I'd just assumed that from wiki. You could
have mentioned that specific jump in your first reply, but I understand why you might not have been bothered. I'm rapidly loosing all interest in
explaining those kinds of jumps myself. I love your name, I was laughing at it on the way back from the shop just now.
I'm fairly sure the majority of the elemental mercury was converted to the sulfate. I used a 500ml flask with a HUGE stir bar in it (I think it was
50-70mm long) and spinning at the plates maximum for around two or three hours. I'd adjusted the temperature so's that the white vapors of the acid
boiling just disappeared; indicating I was running as hot I could without a reflux. The flask had a wash head on it and was being vented.
I was using 10g of elemental mercury. I would guess, from ZyGoats writeup, that I was around 90ml of battery acid, which I then boiled down to 30ml of
concentrated sulfuric acid. I may have added a small excess to counteract any degradation in the battery acid.
I can't remember if I dissolved and then went straight to neutralization. I may have left it to stand overnight so's that I could decant the excess
H2SO4 and bisulfate the next day.
Again, thanks for replying with some more science! 
[Edited on 19-5-2010 by peach]
|
|
|
Sedit
International Hazard
Posts: 1939
Registered: 23-11-2008
Member Is Offline
Mood: Manic Expressive
|
|
| Quote: | | And yes! Frying = sublimation = mercury induced neurotoxicity or back to elemental mercury. Again, hopefully only for themselves
|
Careful, this seems to be a misconception and a simple quick test will prove that no matter what you do when evaporating HgCl2 solutions you are going
to get HgCl2 sublimination around the area.
Take an evaporation dish and place a piece of clean glass over the top of it. Place it on the lowest heat you can get it on and still get it to
evaporate. I used a low heat double boiler setup to warm the solution. Once condensation forms on the glass allow it to dry. You will see HgCl2 on the
glass.
This is abit disturbing.
Knowledge is useless to useless people...
"I see a lot of patterns in our behavior as a nation that parallel a lot of other historical processes. The fall of Rome, the fall of Germany — the
fall of the ruling country, the people who think they can do whatever they want without anybody else's consent. I've seen this story
before."~Maynard James Keenan
|
|
|
peach
Bon Vivant
Posts: 1428
Registered: 14-11-2008
Member Is Offline
Mood: No Mood
|
|
| Quote: | | This is abit disturbing. |
You're right. I dried mine in an evaporation dish by a window with a fan over it and stopped as soon as anything liquid had disappeared. Hopefully not
a lot of the chloride had left by then. If you're worried about me hurting other people, you need to multiply up toxic airborne contaminants by huge
orders of magnitude before they'll hurt anyone else. Nerve agent has to be thousands, tens or hundreds of thousands times more concentrated than the
LD50 once it's airborne before it'll hurt anyone; and it's lethal stuff! Thank the military for that piece of horrible information;
it's in their textbooks on how to kill people. It depends a lot on the weather outside. I also evaporated late at night, very few people would be
around; giving it a chance to dilute and blow off. But yes, I should have been using a flask, wash head and scrubber / neutralizing agent. I wouldn't
let the vapors leave the flask now, and that's something not_important needs to think about when writing his article.
I've been watching the MIT lab manual videos and lectures on chemistry for the last few years, I've rewatched them all recently; multiple times for
some of them. HOLY MOLY are they exceptionally good references on this kind of thing! They're videos that require ZERO effort to take in, they're
free, condensed and from people who know what they're talking about; you can't ask for anything more. You should have links to them here on the site.
I'm now routinely referencing people to them; they're ignoring me and then asking why they didn't yield as highly as they should have.
Did you get much sublimate on your glass before the aqueous phase had left?
[Edited on 19-5-2010 by peach]
|
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
