lucavd - 8-5-2013 at 01:03
I'm trying to synthesize copper hydroxide in the following way:
- Two pure copper electrodes
- K2SO4 solution
- Electricity (from 5 to 12 V)
In theory in this way copper would precipitate from the basic solution while water is electrolysed. A white/pale blue precipitate of copper hydroxide
would form.
This is what I get but I get also an unknown dark precipitate forming at the NEGATIVE side of the cell.
The water is distilled water and the copper is pure and polished.
I thought about copper oxide but I have some doubts that it could be formed at the negative (so where hydrogen is forming) side.
Any idea?
woelen - 8-5-2013 at 01:15
Does the black material appear immediately when you start electrolysis, or does it form once the cathode area contains a lot of Cu(OH)2?
Another interesting thing to try may be performing the electrolysis while the liquid is stirred. This spreads heat and it spreads hydroxide ions,
formed at the cathode.
I can imagine that the black material is CuO, formed due to local heating of the area around the cathode. Cu(OH)2 is rather unstable, even weak heat
is capable of turning it black, decomposing it to CuO and H2O.
lucavd - 8-5-2013 at 01:22
Humm.. interesting..
It turns black locally and after a while. I'm doing it now on ice, I'll let u know
lucavd - 8-5-2013 at 02:06
So I did it @ 10-12°C, 0.22A (the former experiment was done at 1.02A) but I still have dark ppt (at this point I think it is CuO too) on the sharp
points of the negative electrode.
I got the same results with a titanium anode.
I found this video: http://www.youtube.com/watch?v=32q4NVEeJuQ
He doesn't get the dark ppt. The only difference in the experiment id the electrolyte. I'm using K2SO4 (dilute), he's using MgSO4 (saturated
solution). Could this be a factor that influence deposition of Cu(OH)2?
He uses also 9-10V.
[Edited on 8-5-2013 by lucavd]
[Edited on 8-5-2013 by lucavd]
woelen - 8-5-2013 at 04:09
Your setup with K2SO4 is much better. If you use MgSO4, then you get a precipitate of Mg(OH)2, which will cover the cathode. I think that the current
is limited with MgSO4 and that the there is less local heating. But he will have very impure material, containing a lot of Mg(OH)2.
You should use a lower voltage than 9 volts. Try with 5 volts. Higher voltage only leads to production of a lot of heat. If current becomes too low,
then add more K2SO4 to increase the conductivity of the liquid.
If you still obtain black precipitate, then it is time to think in another direction.
LanthanumK - 8-5-2013 at 04:59
For some reason, copper(II) hydroxide is unstable in water, and it has a variety of conditions under which it darkens by formation of copper(II)
oxide, as earlier mentioned. Copper could also be reduced. I used to electrolyze copper in magnesium sulfate solution, and quite a bit of clear blue
copper(II) aqua ions were formed. When I formed copper(II) hydroxide, I used sodium carbonate as an electrolyte, and about 5 volts for an extended
period of time. The only problem is clumping of copper(II) hydroxide on the anode after a while, which restricts current flow.
lucavd - 8-5-2013 at 06:51
So, 5.9V @ 0.4A in K2SO4 seem to work better (no dark ppt so far but the reaction is still on). So the temperature factor is proven to be fundamental
as suggested by woelen.
Tomorrow I try with Calcium Carbonate as electrolyte. Would Ca(OH)2 also come down? I think sodium carbonate would work better (NaOH is far more
soluble then Ca(OH)2 so it will not precipitate)
woelen - 8-5-2013 at 07:29
Calcium carbonate as electrolyte? This does not work, because calcium carbonate is insoluble (it is chalk).
Sodium carbonate also is not a really good alternative. With copper ions it will give basic copper carbonate, Cu(OH)2.CuCO3. Carbonate ions and copper
ions cannot coexist in aqueous solution. Basic copper carbonate has a nice light cyan color, but it probably is not what you want.
I would say, stick with K2SO4 or Na2SO4. Even better would be NaClO4. With sulfate ions there is a slight chance that there is reduction of sulfate at
the cathode, leading to impurities. Perchlorate ion is more inert in aqueous solution and is not reduced at the cathode.
Fantasma4500 - 8-5-2013 at 10:11
well Cu(OH)2 decomposes into CuO at 60*C or around that.. in solution..
i cant imagine anything else but it being CuO
electrolysis of HCl with Cu-Cu yields CuO and CuCl2, i know..
but lets see, the positive thats where the copper ions would be traveling, and before they can get on the cathode theyre immediately made into Cu(OH)2
and before even being visible it goes into CuO by the heat from this reaction.. (:
or well thats my hypothesis on this 'mystery'
but anyways, copper hydroxide im not familiar with methods producing this purely
if i remember correctly its decently soluble so it would be hard to isolate
CuSO4 + NaOH > Cu(OH)2 + Na2SO4
its solubility is apparently ~8 at about 40*C
perhaps NaOH + Cu-Cu would work? not sure if the NaOH would actually react but very likely it would speed up the production, where the water would
then if the NaOH does interfere react with the Na (if this is even possible to form) to again form NaOH rapidly
anyways for what are you gonna use the copper hydroxide anyways? CuCO3*2Cu(OH)2 is called basic copper carbonate, and its formed by CuSO4 +
NaHCO3/Na2CO3 and its insoluble so very easy to clean, also decomposes at pretty low temp. to CuO
12AX7 - 8-5-2013 at 15:11
Is the dark stuff clumpy?
It may be metallic copper, reduced from residual ions in solution or by hydrogen (or radical) reaction with the hydroxide floating about.
I don't think I've ever observed sulfate being reduced. It ought to be possible somehow or another, but may require certain solutions, catalyst,
special cathode, that sort of thing.
Unless there's a motivating need for Cu(OH)2, I'd just as well run it fast and hot, to get CuO directly. It's more stable and settles out of
suspension quickly.
Tim
lucavd - 9-5-2013 at 00:23
@woelen carbonate are slightly soluble but they can be used as electrolytes. Not a first line choice but still feasible. BTY, they
don't work in my case. NaClO4 looks great, I'll try to get some.
@Antiswat. Yes Cu(OH)2 decomposes @ 60°C even locally. NaOH doesn't work, only water hydrolysis occurs (with a bit of copper oxides
at the positive side).
In these experiments I noticed that Cu(OH)2 looks great at the beginning of the reaction in any condition but when it accumulates the current
efficiency drops and so the temperature rises, converting the hydroxide to the oxide.
A this point I think it is only a question of temperature (and local temperature). To support that I noticed that oxide forms only at the edges of the
electrodes and at the very top of the solution.
@Tim. It is not metallic copper because it is unlikely that it forms under these conditions and moreover, the dark ppt is soluble in
dilute sulfuric acid.
I'm trying to get Cu(OH)2 to see if it is possible to reduce it to metallic copper with ascorbic acid (vitamin C). CuO can be reduced by vitamin C
too.