Sulfurous
Harmless
Posts: 3
Registered: 23-3-2006
Member Is Offline
Mood: Yellow
|
|
Sulfur from SO42-...maybe?
Hello, I'm new here so...yeah  . Ok down to the point. I was looking through
my massive electrochemical ]series chart and came across several reactions that might (if I am correct) lead to sulfur by electrolysis.
the first reaction at the cathode is as follows:
SO4 2- + 4H+ + 2e- ---> H2SO3 +
H2O E = 0.172V
and if I remeber correctly sulfurous acid should be a weak acid, thus not all of it breaks away to form HSO3-...so if that is true then this reaction should lead to sulfur:
H2SO3 + 4H+ + 4e- ---> S + 3H2O
E = 0.449V
and at the anode H2O ---> O2 + 4H+ +4e- E = -1.23V which I assume
could replace the lost acid.
So gentelmen what do you think? Is this possible or am I just a stupid fool? (don't answer that )
[Edited on 24-3-2006 by Sulfurous]
Who doesnt like Sulfur?
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
Hmm in general sulfate is very stable, there's no reason it can't be reduced as posted but it sure doesn't seem to want to?
Usual way to reduce sulfate is just blast the damn thing with charcoal and fire, dropping it down to sulfide (+6 to -2 oxidation state change!) by the
reaction R2SO4 + 4C = R2S + 4CO. (CO bubbled through a molten sulfate (Na2SO4, etc.) or passed over hot refractory sulfate (CaSO4, etc.) might work,
producing CO2. I don't know.) From there, an acid anhydride (e.g., SiO2, B2O3) can be added to release the more volatile sulfide anhydride,
otherwise known as sulfur. (Note that CaSO4 + 4C > CaS + 4CO, wait hmm, CaS + SiO2 = CaSiO3 + S(g) doesn't quite balance? Well anyway, note that
CaSiO3 (or Ca2SiO4) melts to glass at a good yellow heat plus some.)
Tim
|
|
|
Sulfurous
Harmless
Posts: 3
Registered: 23-3-2006
Member Is Offline
Mood: Yellow
|
|
Ah I see, that is quite a heat intensive process. But what would cause this reaction, as posted, to not happen? Perhaps it is possiable but too slow
or energy consuming to be used commerically.
Who doesnt like Sulfur?
|
|
|
BromicAcid
International Hazard
Posts: 3247
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
Honestly I don't know enough about the ins and outs of electrochemistry to explain why there are all of these electrochemical reactions listed that
simply don't occur, many of them are likely just so you can balance half reactions or something, but they're all supposed to be for aqueous solution
so why do they have reactions that don't occur in aqueous mediums? I am reminded of when I noticed there were reductions in my electrochemical chart
for phosphate going to elemental phosphorus.... but it doesn't happen under normal aqueous conditions, other reactions happen instead. Same for
sulfur, though it is comparatively eaiser to free from the sulfate then phosphorus is from the phosphate.
|
|
|
Sulfurous
Harmless
Posts: 3
Registered: 23-3-2006
Member Is Offline
Mood: Yellow
|
|
I see what your saying, but how would they get the E values for the half reactions? Also, just out of curiosity, why doesnt phosphate go to elemental
phosphorus (is the E value more negative than -0.83?)?
EDIT: oh btw if some of the values/half reactions dont really happen in an aqueous medium, it would be rather....useless dont you think?
[Edited on 24-3-2006 by Sulfurous]
Who doesnt like Sulfur?
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
Not necessarily. Na has a reduction potential. You need more stable ionic solvents is all.
Reactions in general are reversible, so you can find the potential from oxidation *or* reduction. Oxidation of course is much easier for sulfite, so
it would seem.
Tim
|
|
|
woelen
Super Administrator
Posts: 8013
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
You can make sulphur from sulfate in aqueous medium, but the yield is terrible. It is interesting from an academic point of view, not from a practical
point of view .
Add some zinc powder to dilute sulphuric acid, one can also use a solution of sodium sulfate in dilute hydrochloric acid. The zinc dissolves. Most
zinc is used for making hydrogen, but there are side reactions:
* sulfate --> sulfite (or better in the acidic liquid: sulphur dioxide)
* sulfate --> sulfide (or better in the acidic liquid: hydrogen sulfide)
When you smell the liquid with the zinc, then unmistakenly, the smell of rotten eggs is present. When all zinc has dissolves, then add a small amount
of dilute H2O2. A slight turbidity appears. The dissolved H2S is oxidized to S (H2O2 + H2S --> 2H2O + S).
|
|
|
praseodym
Hazard to Others
Posts: 137
Registered: 25-7-2005
Location: Schwarzschild Radius
Member Is Offline
Mood: crazy
|
|
Wouldnt the addition of zinc to sulphuric acid gives you zinc sulphate and hydrogen instead? Y is it tt sulphite and sulphide will b produced?
|
|
|
guy
National Hazard
Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline
Mood: Catalytic!
|
|
Yes it would be mostly making H2 and ZnSO4, but sulfate is a weak oxidizing agent in acidic conditions, and zinc is a pretty good reducing agent.
Evolution of H2 is most prominent becuase of its entropy.
|
|
|
neutrino
International Hazard
Posts: 1583
Registered: 20-8-2004
Location: USA
Member Is Offline
Mood: oscillating
|
|
We are overlooking nascent hydrogen. This is made when Zn reacts with an acid. Nascent hydrogen is a very strong reducing agent and I'd bet that this
is reducing the sulfate.
|
|
|
hodges
National Hazard
Posts: 525
Registered: 17-12-2003
Location: Midwest
Member Is Offline
|
|
If you add an iodine salt such as KI to concentrated H2SO4 the H2SO4 will be reduced to various products including SO2, H2S, and free S. In fact this
is used as a test for iodide.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
