Copper Chemistry: a Mixed [+I], [+II] Oxidation State?

I am currently looking to produce relatively small amounts of copper (I) oxide (Cu₂O) in a quick and convenient way, for use in "backyard" copper thermite experiments. There are several routes open to choice. Electrolytic oxidation of copper scrap metal and reduction of Cu [+II] compounds with mild reduction agents like sodium metabisulfite or fructose are two possibilities.

Another possibility is the reduction of Cu [+II] compounds with scrap copper metal itself. For example, boiling up a copper (II) chloride (CuCl₂) solution with copper scrap, causes insoluble white copper (I) chloride (CuCl) to precipitate from the solution via Cu + CuCl₂ --> 2 CuCl. The CuCl can then be quickly and easily converted on the spot to the desired copper (I) oxide by treating it with a strong alkali, like NaOH.

I have no CuCl₂ in my possession at the moment but plenty of high quality copper scrap, as well as some decent 35 w% nitric acid (HNO₃). Copper readily dissolves in it to form copper (II) nitrate (Cu(NO₃)₂). A small excess of nitric acid is necessary to ensure complete dissolution of the copper metal. In this way a 0.58 M (mol/l) Cu(NO₃)₂ aqueous solution was made up.

Initial experimentation using various mixes of this solution and some kitchen salt (NaCl, a source of Cl^- anions, necessary to form the CuCl) led to some puzzling results and the suspicion that a mixed oxidation state of Cu [+I] and Cu [+II] might exist, as suggested also by a fellow backyard chemist. I decided to investigate.

I started from 2 solutions:

#1 = copper (II) nitrate ≈ 0.4 M; NaCl ≈ 0.8 M

and

#2 = copper (II) nitrate ≈ 0.4 M; NaCl ≈ 2 M.

Both were "neutralised" with NaOH ≈ 5 M, until a few blue blobs of Cu(OH)₂ no longer dissolved into the solution, the pH was then still slightly below 7.

To 75 ml of both solutions was added 1.9 g of wire copper. Both conical flasks where then gently boiled up for about 1 hour.

There is not much difference in how both solutions behave: both turn a very dark green colour on heating. From #1, some CuCl starts dropping out almost from the start, from #2 no precipitate at all forms during boiling.

On cooling more CuCl drops out of #1 and roughly (visual estimate) the same amount drops out of cooled #2. In both cases there was some copper wire left.

The supernatant liquids (SL) of #1 and #2, after boiling and cooling are indistinguishable to the naked eye, both are clear dark green (wine-bottle green is how I'd describe it).

Diluting or adding more NaCl to SL #1 has no effect. Adding 5 M NaOH caused a deep green, floccular precipitate to from (leaving clear liquid above).

Diluting SL #2 caused more CuCl to form. Adding 5 M NaOH caused a deep green, floccular precipitate to from (leaving clear liquid above).

And adding 5 M NaOH to the diluted SL #2 caused the same green precipitate to form, as well as small amounts or reddish-orange Cu2O.

Adding strong ammonia to the precipitated SL #2 caused the precipitate to dissolve and the typical very deep, intense blue of Cu(NH₃)₄²⁺ to appear.

Although I decanted SL #2 to have a CuCl-free sample, soon after that small, clear crystals began to appear in it (possibly NaCl, the #2 had indeed been concentrated quite a bit during the boiling). And the next day more CuCl had appeared in this sample (quite a bit, actually).

Whether the bottle green solution can actually be described as a "mixed Cu oxidation state" may be a little contentious at this point. But it's clear that Cu [+I] and Cu [+2] co-exist in it and in quite a stable manner too: almost 24 hours later and in the presence of air oxygen the bottle green solution hasn't changed at all. The dark green colour, quite unusual for Cu compounds may well be a [Cu(1+),Cu (2+)]Cl complex anion.

It would also appear that CuCl's complexing ability via CuCl + Cl^- ↔ CuCl₂^- overrides it tendency to form Cu2O (in near neutral conditions) and that this complexation would explain why #2 doesn't generate any precipitated CuCl right from the off, due to high NaCl content. Presumably the complexation constant (K_f=[CuCl₂^-]/[Cl^-]) is temperature dependent.

That still leaves the question of what constitutes the deep green precipitate. Copper (II) hydroxychloride (Cu(OH)Cl) springs immediately to mind but although I've never seen Cu(OH)Cl I somehow imagine it to be green but lighter in colour. And if it was indeed Cu(OH)Cl, this would still beg the question of what has happened to the Cu [+I], present also in the solution.

Presumably a little UV-VIS absorption spectroscopy could shed some light on this presumed mixed oxidation state and its associated Cu [+I], Cu [+II] chloride complex...

To be continued?

Update:

A solution #1, this time boiled and reduced with sodium metabisulfite (Na₂S₂O₅), also becomes dark green and CuCl drops out of it upon cooling and/or dilution. The same green precipitate forms with addition of NaOH.

Incidentally, the wine-bottle green solution, when strongly diluted to a light blueish-green coloured liquid, still produces the green Cu-OH-Cl compound when 5 M NaOH is added to it: the complex survives dilution.

Also, I've now got some 150 ml of 0.4 M CuCl₂ (probably about 0.1 to 0.01 M in HCl as well, from excess HCl). It's lighter than 0.4 M Cu (II) nitrate, greenish blue, I'd say.

Adding a few drops of 5 M NaOH to it causes blue Cu(OH)₂ to precipitate.

Saturating the CuCl₂ solution with NaCl causes it to shift to a bright green (from CuCl₄^2-) but 5 M NaOH still causes blue Cu(OH)₂ to drop out.

Heating the CuCl₂/saturated NaCl to BP doesn't change anything and on cooling 5 M NaOH still causes blue Cu(OH)₂ to form.

Boiling 75 ml of the CuCl₂ acidic solution with copper wire also caused the dark wine-bottle green to appear. No CuCl dropped out immediately but on cooling a great deal did. After cooling (and precipitation of the CuCl) the supernatant liquid was much lighter in colour. Reheating the cooled solution causes the CuCl to completely redissolve.

Adding 5 M NaOH to the cooled solution causes a khaki green flocculate to precipitate. Adding ammonia to this slurry, causes only a light blue supernatant liquid (Cu²⁺-ammonia complex?) to form, leaving behind reddish orange Cu₂O (I believe). The latter dissolves easily in an excess of ammonia.

Clearly the reductions causes a stable complex to form, which when neutralised causes some "Cu hydroxychloride" compound to precipitate. This complex has to be a species different from CuCl₄^2- and may combine the Cu²⁺ and Cu⁺ cations into a single hydroxy chloride anion.

Update 2:

It would appear that I'm far from the only one that has encountered and documented the unusual and so far unidentified complex that causes Cu [+II] bearing solutions to colour-shift to the unexpected dark green when exposed to a reducing agent and some heat. Over at Woelen.Scheikunde.net some experiments are described that are very similar to my own exploits and that are attributed to a "mixed oxidation state copper complex".

Hat tip to Tim over at ScienceMadness.org for pointing me to these pages.

Some thoughts on the possible composition of the unknown complex:

Assume we start from a solution of CuCl₂, acidified by a strong, monoprotonic, non-chlorine bearing acid, such as nitric acid (so that OH^- ions play no part). The solution now contains x mol of Cu (in whatever form) and 2x mol of chlorine (in whatever form). The molar chlorine to copper ration is thus 2.

Assume we add an excess of copper metal and allow the reduction of Cu²⁺ to proceed, via Cu + Cu²⁺ --> 2 Cu⁺ and that most of the Cu⁺ precipitates out of the solution as CuCl _(s). Assume y mol of copper metal reacted away, this would yield 2y mol of CuCl, reducing the amount of chlorine in the solution by 2y mol, leaving 2x - 2y.

Bear in mind that only half of the Cu needed to form the 2y mol of CuCl actually comes from the solution (the rest is supplied by the copper metal), then the amount of Cu precipitated from the solution is actually y mol and the remaining Cu in solution is x - y.

The new molar chlorine to copper ratio is thus (2x-2y)/(x-y) = 2.

The molar chlorine/copper ratio remains therefore unchanged, regardless whether some reduction has taken place or not.

The ionic species present in solution are probably H₃O⁺, Cu⁺, Cu²⁺, CuCl₂^-, CuCl₄^2-, the unknown Cu^+,2+ - Cl^- complex, Cl^- and a spectator anion (like nitrate, to comply with the neutrality requirement).

The first two species are likely to be small in concentration, due to complexation. [CuCl₂^-] is also likely to be quite small, since CuCl has a tendency to drop out of the solution, despite the formation of the assumed unknown complex. CuCl₂^- requires free Cl^- to be present in the solution and most of it is likely to be bound up in CuCl₄^2-and the unknown Cu^+,2+ - Cl^- complex.

This leaves CuCl₄^2- and the unknown Cu^+,2+ - Cl^- complex as the most likely dominant, copper and chlorine bearing species and would appear to suggest its Cl/Cu ratio to be close to 2.

Perhaps something like {Cl^--Cu⁺-Cu²⁺ - 3 Cl^-}^- could be envisaged. Simply put: Cu₂Cl₄^-?