bilcksneatff - 30-12-2007 at 17:54
I had been searching for information about sulfurous acid. I found this at http://www.ucc.ie/academic/chem/dolchem/html/comp/h2so3.html
"The solution [of sulfurous acid] when heated in a sealed tube deposits sulphur.
3H2SO3 ===> 2H2SO4 + H2O + S"
I did some more research and could not find anything to confirm this. Any other site said sulfurous acid decomposes to sulfur dioxide and water.
Which one is right? My guess is H2SO3 --> SO2 + H2O.
12AX7 - 30-12-2007 at 18:02
Over a long time, like... in a sealed tube in an autoclave, I could see that happenning.
In air, the equilibrium is forced towards SO2 gas (because, in its equilibrium gas form (SO2 + H2O <---> H2SO3), it is less soluble in hot water
(SO2(aq) <---> SO2(g)), which has a low partial pressure in the atmosphere so tends to leave the reaction.
Tim
bilcksneatff - 30-12-2007 at 18:08
Thank you very much
Pyrovus - 31-12-2007 at 02:13
I do recall it mentioned somewhere that UV light will cause the disproportionation of sulfurous acid; though not having tried it I cannot say if it is
a viable means of producing H2SO4.
Mumbles - 2-1-2008 at 19:09
I've heard the same thing about the heating in a glass tube producing solid sulfur. I could have sworn I've seen images of this. It was likely on a
teaching site, there were pictures of it in progress and how to do the demo. They probably explained how/why it happens, but it totally escapes me at
the moment. It is probably something that is pressure dependant
From thinking about how this might happen, I was reminded of a way to form colloidial sulfur from photographic hypo (sodium thiosulfate). Upon
acidifying and heating it, it forms sulfur and sulfur dioxide. Normal methods of preparing the thiosulfate ion is reaction of sulfites with elemental
sulfur at boiling temperatures, so this is a bit of a cyclical arguement.
Playing around a bit with a pen and paper I got another idea, Pyrosulfite. It exists in the following equilibrium:
2 HSO3- ----> S2O5(2-) + H2O
It's not exactly an analogue to thiosulfate, but it could perhaps be involved. I was thinking that the heat and pressure could perhaps provide the
energy, and force the gases into solution to provide whatever intermediate would not normally form at room temp/atomospheric pressure.
I found a paper that mentions it's decomposition via an 8 or 6 membered ring. This was done at low temps, and the abstract does not mention just
which products it decomposes too. I don't have my journal access available at the moment either.
Chemistry. 2002 Dec 16;8(24):5644-51