tommy1641 - 13-7-2004 at 20:12
Ok, I tried something new and I think it worked. I put about a half cup of MgSO4 (Epsom Salt) and a 2 tablespoons of water in a cup then mixed them.
Next, I placed 2 wires in the cup. These had 18 volts going through them. Several different things seemed to be produced, but the one that caught my
attention may have been sulfur. The water was the exact same color as sulfur. So, I tried to filter the solution several times, and every time nothing
stayed in the filter. I expected some sulfur to be left behind because it is nonsoulable in water. Does anybody know: If sulfur really was produced,
and if it was, why didn't any stay behind? Thanks.
[Edited on 5-8-2004 by chemoleo]
BromicAcid - 13-7-2004 at 20:16
One key question, these two wires placed in solution, what material were they? Copper, iron, nickel...?
tommy1641 - 13-7-2004 at 20:18
The wires were copper. Sorry for not saying earlier.
darkflame89 - 14-7-2004 at 01:04
Copper wires? Then the color of the solution might have been due to the dissolved Copper..
Saerynide - 14-7-2004 at 02:20
The solution would turn bluish if it was copper, not yellow 
And how does 1/2 a cup of epsom salt dissolve in 2 table spoons of water? Or did you electrolyse a bowl of soggy crystals
frogfot - 14-7-2004 at 02:21
CuOH is yellow, and insoluble in water (it should dissolve in HCl, making complexes.. might be a good test).
[Edited on 14-7-2004 by frogfot]
Saerynide - 14-7-2004 at 02:26
Oops, my bad. Ive never managed to make or see CuOH, only Cu(OH)2 
hodges - 16-7-2004 at 14:29
I would expect magnesium hydroxide to be produced at the cathode (negative terminal); this would have the color of precipitated sulfur (white). Not
sure where the yellow would come from, though.
Sarevok - 17-7-2004 at 18:27
Aqueous electrolysis of this salt will merely produce oxygen and hydrogen. Sulfur will not be formed.
EDIT: The yellow color is probably due to the presence of cuprous hydroxide (Cu+ from the oxidation of your electrodes and OH- from the reduction of
water).
[Edited on 18/7/2004 by Sarevok]
chemoleo - 5-8-2004 at 00:02
It's very unlikely it is sulphur, indeed. I am not aware of an easy electrolytical process that reduces sulphate to sulphur.
I imagine all you'd get is H2 and O2, as the gasses coming off the electrodes.
As to the colour.... impurity!
Cuprous hydroxide CuOH was mentioned, but like the cupric hydroxide Cu(OH)2 both are insoluble. Additionally, CuOH undoubedly reacts either to Cu2O
(insoluble once again), or it is oxidised with AIR, to form green/turquoise Cu2+ salts (or rather, as Saerynide says, blue, as CuSO4 would be formed)
So I doubt its down to the copper. More likely it's some other impurity.
In addition it does help if you get your weight units sorted. Cups and spoons dont really help
WAs the final solution cloudy yellow, or clear yellow? Precipitates at all?
unionised - 5-8-2004 at 12:31
Copper (I) oxide is yellowish or reddish, and, like sulphur, would form a precipitate.
It's low solubillity means that suspensions of it in water are fairly stable.
My money is definitely on Cu2O, possibly hydrated.
Most impurities in copper wire are not popular because they reduce the conductivity.
If the MgSO4 was pharmaceutical grade it would be pretty pure too.
Why postulate some "impurity" when Cu2O is reasonably probable?
chemoleo - 5-8-2004 at 12:46
Yes i wonder that myself now.
I am just not aware of ever having prepared Cu2O with a colour 'exactly like sulphur'. The commercial batch I have is a deep red, and some
preps I did brought it up to orange red, but definitely not sulphur yellow. Who knows... I guess colour depends on the manner it was prepared 
Anyway, tommy, the issue can be easily solved if you subject your yellow precipitate (which you should filter off) to HCl(aq). If it is Cu2O, it will
react with it, forming initially a white precipitate, which, on air, will slowly solubilise and turn dark green (CuCl2).
Electrolysis
MadHatter - 5-8-2004 at 19:43
Tommy1641, if you decide to do it again I suggest using "gouging" rods,
then check your results. I use these for chlorate/perchlorate production.
Metal anodes and cathodes are likely to combine with the products
they produce.