I've been making some unusual observations while trying to reduce Cu [+II] bearing solutions to Cu [+I] (in attempts to make copper (I) oxide (Cu2O)).
In the presence of chloride (Cl-), these reductions invariably lead to a strong shift in colour towards a deep "wine-bottle green".
I used solutions that were ≈ 0.4 M copper (II) nitrate and ≈ 0.8 M NaCl and pH ≈ 1 - 2. This is a typical deep blue. I've also used
CuCl2 ≈ 0.4 M as a starting solution.
As reduction agent I've used mainly copper metal itself (via Cu [+II] + Cu [0] --> 2 Cu [+I]) and also sodium metabisulfite (Na2S2O5).
Regardless of starting solution and reducing agent, this deep green colour always forms on boiling the solution with the reducing agent (RA).
Afterwards, adding NaOH to it causes a similarly coloured flocculate to precipitate.
A couple more observations can be added. While using Cu as a RA the reduction always appears very incomplete (assuming the green species contains Cu
[+II], of course), with metabisulphite as RA, it's possible to push the reduction further by going strongly non-stoichiometric (excess sulfite). The
resulting solution is then much lighter in colour and when NaOH is added the precipitate is actually much darker: from khaki green to
actually brown depending on how far the reduction has been "pushed".
I think it's safe to say that the green solution contains CuCl2 (1-) and CuCl4 (2-) but that doesn't explain the wine bottle green or the various
precipitates.
Could there be a {Cu(1+) - Cu(2+) - Cl(1-)} complex present, responsible for that particular green colour?not_important - 3-2-2008 at 09:00
[Edited on 3-2-2008 by not_important]blogfast25 - 3-2-2008 at 09:19
Well, that's the green alright. But how could that form in an acid solution? Not impossible of course, depending on complexation constants. Something
that forms in natural conditions can surely form in my conical flask... 12AX7 - 3-2-2008 at 09:30
[Edited on 9.1.14 by bfesser]pyrochem - 3-2-2008 at 09:46
Unless I'm misreading the formula, doesn't atacamite only contain Cu<sup>2+</sup>?
Does the color appear on boiling without a reducing agent?
It is possible that it could be a Cu<sup>2+</sup> hydroxide and chloride complex. You say on your blog that you added NaOH until the pH
was slightly below 7. The Cu<sup>2+</sup> should have some hydroxide ligands at this pH, as Cu<sup>2+</sup> salts will
complex with hydroxide ions from neutral water to leave an acidic solution.guy - 3-2-2008 at 12:15
there are Cu(II)-Cl-Cu(I) complexes, the dark color is due to a charge transfer band which is usually pretty intense.blogfast25 - 3-2-2008 at 13:10
Quote:
Originally posted by pyrochem
Unless I'm misreading the formula, doesn't atacamite only contain Cu<sup>2+</sup>?
Does the color appear on boiling without a reducing agent?
Yes, Atacamite is Cu [+II] only.
No, the colour only appears when the RA is present.
I'll reproduce the reduction with Zn too...
Tim, thanks for the Woelen pages, they're beautiful and really clear.
This one in particular is very close to what I observed in slightly different solutions. It's safe to say that the colour I observed is the Cu(1+)
- Cu(2+) complex Woelen also surmises.
[Edited on 9.1.14 by bfesser]Random - 8-1-2014 at 19:08
I wonder if Cu2+ can oxidize H2SO3 to H2SO4 and itself get reduced to Cu+. It would be interesting way of making H2SO4 with CuSO4 and H2SO3AJKOER - 10-1-2014 at 06:50
"Cupric oxide is somewhat soluble in fused caustic alkalies, and in their concentrated solutions. Several coloured hydrates have been described,
including the blue or green CuO,H2O, the olive-green CuO,0.8H2O, and the brown CuO,0.35H2O. A magnetic, crystalline cupric ferrite, CuO,Fe2O3, is
formed by precipitating a mixture of cupric and ferric salts with alkali. "
blogfast25 - 10-1-2014 at 09:31
AJ:
These have nothing to do with the interesting Cu(II)/Cu(I) mixed oxidation state complex woelen and me describe.
