[Count of Monte Cristo]

Lately I have been looking into electrolysis experiments, and as a simple reaction I decided to electrolyze MgSO₄ with copper electrodes in order to produce a solution of Mg(OH)₂ and CuSO₄. I created a saturated solution of Epsom salt, placed my positive electrode in the bottom of the solution and my negative electrode above the positive. I let the reaction go for 45 minutes before taking out my electrodes, I observed that Cu(OH)₂ seemed to be produced and had collected at the bottom of the beaker with a blue solution above, and above that Mg(OH)₂ had collected around my negative electrode. Was Cu(OH)₂ produced because my Epsom salt solution was not saturated enough? and is there any way to tell when the maximum amount CuSO₄ has been produced?

Edit: I have several pictures of my set up and end result, as my descriptions are probably not very helpful

[rwiew]

Cool experiment! The problem is the solubility of the hydroxide and sulfate - the sulfate is pretty well soluble in water, while the hydroxide isn't. Your process is producing Cu²⁺ ions on the anode and reducing water on the cathode, which produces OH^- ions. Essentially you have a stoichiometric production of Cu(OH)₂, which of course will precipitate straight away. I see no way to produce the sulfate here, you can't produce an excess of Cu²⁺ without producing OH^- and getting Cu(OH)₂ from all of it (well there's clearly some hydrated copper ions in solution hence the blue color, but that's not much at all). What you can do is collect the hydroxide, treat it with not too dilute H₂SO₄ to take into solution and then crystallize CuSO₄.

[Count of Monte Cristo]

Then it would be simpler for me to electrolyze just water with copper electrodes, as then I wouldn't have to worry about separating Copper and Magnesium hydroxides produced with Epsom salt electrolysis. And if CuSO₄ is not produced, what happens to the SO₄^2- ions that are in solution as a result of the Epsom salt being dissolved?

[rwiew]

Actually I didn't think about magnesium hydroxide, good point. Well, if magnesium hydroxide is precipitating as well, then you will get a good amount of Cu²⁺ in the solution. So you are getting CuSO₄ in the solution, crystallization from that should give you an ok amount I guess. It will be much better to just electrolyze a solution of H₂SO₄ - you are avoiding the magnesium problem + you will not be getting Cu(OH)₂ as the solution will stay acidic enough. That should give you a clean solution of CuSO₄.

[Count of Monte Cristo]

Using my current method I would probably have to scale up in order to get a useful amount of CuSO₄, the reason I don't just use sulfuric acid is, quite embarrassingly, because I don't own any. Unlike hydrochloric acid, which is sold by the gallon, options for sulfuric acid seem to be rather limited, but with how useful the stuff is I'll probably just buy some conc. sulfuric acid off of amazon.

[rwiew]

Yeah, that is your best bet definitely.

[Mercuric cyanide]

Actually, magnesium hydroxide will not precipitate, as although its hardly soluble, it is much more so than copper hydroxide. Because copper hydroxide is produced initially, and it's less soluble, it will precipitate and Mg²⁺ and SO₄^2- will remain in solution.
This is a decent potential copper hydroxide production method, however due to the slight acidity of MgSO₄, yields will not be as good as they could be.
If your goal is copper sulfate, you will need to find a different way to get it. Without sulfuric acid, your options are limited. Find a hydroxide that is less soluble than that of copper, and displace it with the corresponding sulfate salt. Keep looking, it won't be too hard to find.

[billnotgatez]

Sulfuric acid is sold in some automotive parts store for use with some car batteries.
Be careful it is caustic.

[Count of Monte Cristo]

My local PepBoys is the first place I checked for the acid, however I wasn't able to find any. After asking the assistant at the store I found out that Sulfuric acid isn't sold anymore as to it's dangerous nature, and now some kind of colored distilled water is being introduced as its substitute.

Edit: I mentioned above that I have a quantity of HCl, would I be correct in assuming that the electrolysis of a conc. solution (10M) with copper electrodes would yield CuCl2 more efficiently than electrolysis of a chlorine salt would?

I apologize if this is slightly off topic, but I'm basing this off of what was mentioned above with Sulfuric acid.

[rwiew]

Which potential is more favorable though: oxidation of chloride or oxidation of copper? (not a tricky question, just writing a quick reply and don't have time to check).

[Mercuric cyanide]

would I be correct in assuming that the electrolysis of a concentrated. solution (10M) with copper electrodes would yield CuCl2 more efficiently than electrolysis of a chlorine salt would? 

To answer that one must define "efficiently". Your energy will certainly not be used efficiently, but your copper will. Some chloride will inevitably be oxidized to chlorine. If hydrochloric acid is cheap to you, then this is probably the best way for you.

Which potential is more favorable though: oxidation of chloride or oxidation of copper?

Oxidation of copper is of course much easier than that of chloride.
However, that proves nothing. Electrolysis of an HCl solution with an inert anode(I use MMO) yields chlorine-oxygen mixture at the anode, and hydrogen at the cathode. Unless you use very little current, or use a very porous copper electrode (large surface area) the copper wont react with the chlorine fast enough.
My guess is that either way, chlorine is oxidized, and as produced, it oxidizes copper and is reduced to it natural ^-1 state. I see no reason that the fact that copper is used as an electrode will change the fact that chlorine is initially oxidized, so its more a question of how fast will the warm, moist, possibly even partially radical chlorine react with copper. Rather than a question of whether copper or chlorine is more likely to get oxidized.

[billnotgatez]

MMO electrode
https://en.wikipedia.org/wiki/Mixed_metal_oxide_electrode
@Mercuric cyanide
Do you have a link to the kind of electrodes you use

[Borek]

Just comparing potentials is dangerous, as it can produce incorrect conclusions. Some reactions are slow and require high overpotentials to proceed fast enough, that in turn means before they start we can get into the potential range where other reactions start to dominate.

Not saying that's the case here, it is just a first thing that comes to my mind when I see "chlorine and oxygen produced on the electrode". Oxygen is the slow one, which is why we observe chlorine even if potentials suggest we shouldn't.

[Mercuric cyanide]

MMO electrode
https://en.wikipedia.org/wiki/Mixed_metal_oxide_electrode
@Mercuric cyanide
Do you have a link to the kind of electrodes you use

The exact one I bought was on eBay, and the deal has long since expired. From what I remember, it was a slightly cheaper  than platinum MMO anodes. Iridium and ruthenium oxide's, the latter of which (and perhaps the former as well) is[/are] very catalytically active in chlorate, hypochlorite and other chlorine oxianions production.
Borek, I have tested the concentration of the oxygen-chlorine mixtures produced at the anode under different conditions, (eg. Temperature, concentration of the sodium chloride solution etc.). Generally the ratio us 85-15 to 95-05 chlorine to oxygen.

BTW, does anyone know of a good way to scrub out the oxygen? I might use this method to isolate chlorine in a large scale. I usually use calcium or sodium hypochlorite and sulfuric acid to do it, but this is much cheaper.