Making copper hydroxide?

[Theophrastus]

Finding the lovely, yet rather old thread, you guys have on making copper sulfate, (which actually yields copper hydroxide), I was curious as copper hydroxide is a useful chemical to have around, and as such performed the experiment as prescribed, using copper electrodes, and magnesium sulfate (epsom salt) as an electrolyte. I used about 18 volts initially, however, as batteries burned out, some were replaced, some weren't, I was left with about 12 volts, by the experiment's end, however the results, were not at all as I expected, as the solution turned a soft , pale shade of yellow, while the precipitate was orange. I thought this to be due to an excess of epsom salt, so I tested its solubility, and appropriately decreased its concentration upon my second attempt, however the yield was the same. Since there was naught but epsom salt, I think I can make the assumption, that the orange precipitate is copper. (this I shall test by heating it, when I get some more denatured alcohol, for my alcohol burner next week) But what happened. Many people seem to have used this method, with excellent results, however, why didn't my method work? For those unfamiliar with the forum in question, the theory is that the copper "excited" by the electricity will displace the magnesium momentarilly. In an aqueous solution, the displaced magnesium would then react with the water to produce magnesium hydroxide. The magnesium hydroxide would then react with the copper sulfate, the magnesium displacing the copper to once again attain magnesium sulfate and copper (II) hydroxide:

 

Cu* + MgSO4 > CuSO4 + Mg

 

Mg + 2H2O> Mg(OH)2 + H2

 

Mg(OH)2 + CuSO4 > MgSO4 + Cu(OH)2

 

Thoughts?

[Theophrastus]

(adjoin with previous post)

Having filtered out most of the solution, leaving the rest to boil off in the beaker, I found a rather interestinhg sight. I'm rather annoyed with my yield, and as such, have given up on making hydroxide, in such a fashion. Still, I remain interested in what I have prepared finally.

I think that the orange is likely to be copper, and the white crystals near the top of the beaker, leftover magnesium sulfate. The whitish bluish substance on the bottom, covered in copper, I think, is a mix of copper hydroxide, and carbonate, formed through further oxidation, though the colour worries me a bit, as it also looks a bit like copper oxychloride, however, as no sodium chloride salt was added, I doubt it.

I'll soon add some hydrochloric acid to the beaker, and if my guesses are correct, the solution should turn the blue colour of copper chloride, if not... well I haven't really thought of that. Any last- minute thoughts are quite welcome. If all goes well, I'll post my results this evening.

Edited May 9, 2009 by Theophrastus
Consecutive posts merged.

[UC]

By a very long shot, the easiest way to make copper salts is to buy a lot of copper sulfate from somewhere and work from that. If you're really desperate, bubbling air through a solution of ammonium sulfate in aqueous ammonia with copper metal sitting in it is pretty good at oxidizing the copper and dissolving it as tetraamminecopper hydroxide. From there, drying out the solution and heating to decompose the complex and drive off ammonia, then dissolving in your acid of choice is a decent, but very annoying way to go.

[Theophrastus]

While you may look at the thread, and notice its age, just today, I was digging through my chemical "storage room," to find the solution, I had from before, left for months untouched. With much precipitate encrusted on the bottom, I decided to finally clean this ancient mess up, and poured in a generous amount of hydrochloric acid. The solution slowly grew less and less turbid, until it turned a lovely golden colour, almost like honey, and suddenly I realised- iron chloride!

 

The precipitate that had looked black was actually brownish iron hydroxide, while the orangish precipitate was iron oxide! Looking back, I realise that I was using a set of electrodes I had bought which were copper- plated iron, which needless to say would have yielded little copper salt, but substantially more iron.

 

Needless, to say, this unholy mess settled out, my beaker's clean, and I'm hoping to boil down the solution, to reattain some iron chloride crystals, which I'll store for later use.

[Melvin]

You're on the right track with the electrolysis; I make all my Cu(OH)2 like that. You really don't need voltage that high; I use a 9v "wall-wart" power supply for this. Anyway, I'm pretty sure that the magnesium sulfate only is used to increase the conductivity of the solution, and doesn't play a role in the reaction. The most important part is to use a copper anode; the cathode could be almost any metal.

 

I believe the combined reaction is as follows:

Cu + 2H2O --> H2 + Cu(OH)2

[UC]

While you may look at the thread, and notice its age, just today, I was digging through my chemical "storage room," to find the solution, I had from before, left for months untouched. With much precipitate encrusted on the bottom, I decided to finally clean this ancient mess up, and poured in a generous amount of hydrochloric acid. The solution slowly grew less and less turbid, until it turned a lovely golden colour, almost like honey, and suddenly I realised- iron chloride!

 

The precipitate that had looked black was actually brownish iron hydroxide, while the orangish precipitate was iron oxide! Looking back, I realise that I was using a set of electrodes I had bought which were copper- plated iron, which needless to say would have yielded little copper salt, but substantially more iron.

 

Needless, to say, this unholy mess settled out, my beaker's clean, and I'm hoping to boil down the solution, to reattain some iron chloride crystals, which I'll store for later use.

 

You'll need to shield it from oxidation if you want ferrous chloride crystals. FeCl2 oxidizes to Fe(III) ion very easily and if you don't have an excess of HCl around to trap it as FeCl3 and keep it from hydrolysing, you will get nothing but rust.

 

If you want FeCl2, degrease some steel wool and throw into concentrated HCl in a covered beaker (It spatters a *lot*). A white powder settles out afterward, which is some form of FeCl2. I'm not sure if this would be anhydrous or a hydrate.

[Theophrastus]

Thanks for the response UC. Ironically, I had added excess (far too much given the solution's pH value of around 0.6) HCl, and tonight was thinking of a means to rid myself of it. I was thinking of adding some Fe2O3, to use up the acid, and then adding ammonia, to precipitate a iron (II) ammonia complex, at which point I can filter off the FeCl3, separating it from the complex, and then with mild heating decompose the complex, to attain ammonia gas and iron (II) chloride. This I would then, as you told me, add slight excess HCl, to decrease, oxidation to iron (III) chloride, leaving it in stored solution form.

By the way, is there any means to speed up the oxidation process of FeCl2, sufficiently so as to ignore it, and bypass the filteration?

As for the last note- yeah; there are a lot of easier ways to make FeCl2, and I'm not sure, but I would guess the resultant iron (II) chloride to be hydrous, as I believe that the dehydration process involves use of HCl gas- enoough said.

Edited July 17, 2009 by Theophrastus

[UC]

Fe (II) doesn't form complexes, at least not that I've heard of. You'll make lots of hydrous Fe(OH)2 which will oxidize before your eyes to rust. Fe (III) in solution will certainly form Fe(OH)3 of varying degrees of hydration the second ammonia hits it.

 

Since NH4Cl decomposes into fumes, you would be quite likely to get iron oxide if an iron ammine complex were heated.

 

Fe (II) is a crappy lewis acid. Strong lewis acids tend to be extremely hygroscopic, but crystallizing an anhydrous salt from aqueous solution is not unheard of. I would suspect that it is probably the dihydrate given that water is in relatively short supply in concentrated HCl.

 

Bubble chlorine gas through the solution. Instant Fe (III).

I'd tell you to add H2O2 to a solution of FeCl2 in HCl, but the iron ions catalyze the decomposition of H2O2 and you'll get a lot of hot oxygen and less chlorine than you'd like.

Adding bleach instead of H2O2 works if you don't have a problem with sodium ion contamination.

[Theophrastus]

My new plan (following suggestion):

 

First, as before, I'll add Fe2O3, in order to use up the excess hydrochloric acid, and forming an arbitrary mixture of iron (II) and (III) chloride.

 

I then have two paths, one: I could bubble in chlorine gas, in order to achieve a purely iron (III) chloride solution, and then boil it off to a attain a hydrous crystalline product.

 

My second option would be to then use potassium hexacyanoferrate to precipitate the iron (III) chloride, forming a dark blue complex. And the filter off the iron (II) chloride. (I realised that I was confusing iron 2+ and copper 2+. For the fomer, I suppose one would use something like potassium triflouride hexacyanide (forgive me if I'm mistaken))

 

Anyhow, My plan is to go with option 1 as of now, given a low stock of potassium hexacyanoferrate, though I'll have to purchase some PVC pipe first.

 

Thanks for all the help UC!

[UC]

My new plan (following suggestion):

 

First, as before, I'll add Fe2O3, in order to use up the excess hydrochloric acid, and forming an arbitrary mixture of iron (II) and (III) chloride.

 

 

Hopefully you mean freshly prepared Fe2O3. The pottery grade stuff is calcined and extremely hard to dissolve. If you have pure Fe2O3, you'll only have Fe(III) in the resulting solution.

 

You don't want to use up all the excess acid. FeCl3 is stabilized in solution by excess HCl.

 

In water, FeCl3 disassociates to HCl and ferric hydroxychlorides. In dilute enough and neutral/basic solution, it proceeds all the way to ferric hydroxide and you get precipitate. Adding extra HCl drives the equilibrium to the left toward FeCl3 (hydrated).

[Theophrastus]

Pretty ironic, because, even before I had read your post, I had already added some pure Fe2O3, and yet it seemed, it did absolutely nothing for the solution, as it remained just as it had before in colouration and acidity, and I simply filtered it off, and left it acidic, as you stated. (personally taking note that if I use it in an applicable chemical reaction, I'll be adding excess of the other reagent, in al likelihood).

 

As for making FeCl3, in contrast to my current FeCl2, I'l use up about half of the mixture, for that, as I made quite literally excessive amounts of FeCl2, and it wouldn't be bad to have both. I'll be getting some PVC pipe to channel the chlorine gas with. I'll post the results, when all has been completed. Thanks for all the help UC!