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woelen
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In personal communications, some years ago, it was suggested to me that the soluble complex could be something like Cl-Cu-(μ-Cl)-Cu-Cl. In the
other thread, guy also makes mention of such a structure. The Cl-Cu-(μ-Cl)-Cu-Cl then is symmetric, with copper in the +1½ oxidation state
(oxidation states need not be integer numbers, another example of non-integer oxidation state is the superoxide ion).
Unfortunately, I have no reference at all, so it remains speculation.
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CyrusGrey
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Wow. I know very little about these things like how complexes form, so you all have somewhat lost me on the exact chemistry of these things.
I did perform another experiment.
I put some of the brown liquid in a watch glass and some in a test tube. In the watch glass I placed a small bit of copper wire. This morning the test
tube had turned from the brown solution to the green [CuCl4]2- complex. However the liquid in the watch glass had not. The copper had a little blue
corrosion on it where it was sticking out of the liquid.
I then placed some copper wire in the green solution in the test tube and put mineral oil over it (to keep air out). A few hours later when I checked
back it was brown again.
I had tried heating this same solution to dryness, or adding H2O2, or adding permanganate, with no smell of Cl2 gas or white HCl fumes. Only with
hypochlorites did the solution evolve chlorine. This makes me think that there is no HCl left in solution to just form new brown complex from the
copper, and overpowering the green color.
So I think that adding excess copper to the [CuCl4]2- complex will form the brown complex. This would be a good method for storing it too.
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woelen
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Initially it indeed does form the brown complex, but reaction goes on. Finally, the solution becomes colorless and CuCl2(-) is formed.
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blogfast25
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Well, I had about half a liter of green Cu<sup>2+,1+</sup> solution left from various reductions and decided to convert it to oxide with
NaOH and boiling.
The precipitate formed was dark-khaki green, not blue as you would expect from Cu<sup>2+</sup> plus hydroxide.
The solution had been standing in my shed in a cut-in-half pop bottle for probably more than a month. Clearly much of the Cu<sup>+</sup>
had not been oxidised to Cu<sup>2+</sup>.
The dark-khaki green precipitate is also much less floccular than the blue Cu(OH)<sub>2</sub>.n H<sub>2</sub>O and on boiling
hardly changes colour. Altough this could simply be a physical mix of CuO and Cu<sub>2</sub>O, potentially it might be a mixed oxide:
Cu<sub>3</sub>O<sub>2</sub>.
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woelen
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The precipitate you get has an interesting color profile.
If your solution is purely based on copper(I), then you get a bright yellow/orange precipitate.
If it is mostly copper(I) with only a tiny amount of copper(II), then you get a dirty brown/yellow precipitate.
If it is an approximately equal mix of copper(I) and copper(II), then it becomes dark green/brown.
If it is mostly copper(II), then it becomes dirty green or dirty/blue.
If it is copper(II) only then it becomes bright blue in the absence of chloride, or green in the presence of chloride.
The absence or presence of chloride also has a marked effect. Try making a solution of copper(I) in ammonia (can be done by dissolving copper(II)
sulfate in ammonia and then adding a solution of sodium dithionite in ammonia). When a concentrated solution of NaOH is added, then bright
yellow/orange copper(I)oxide/hydroxide is precipitated. When the same experiment is repeated with chloride present, then the color hardly is
different. But as soon as copper(II) is present as well, then the presence of chloride does have a visible effect, which is stronger for higher
concentrations of copper(II).
I think that the precipitates you get are mixed (hydr)oxides and chlorides, with also mixes of copper(I) and copper(II) of indeterminate
stoichiometry. In simple words: complex stuff!
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blogfast25
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Yes, I've seen all these colours during various experiments with the [+I, +II] complex.
What I found interesting about this latest batch of precipitate is that, despite the solution having been exposed to air for so long, still
considerable Cu [+I] remains. The solution was an emerald kind of green and around 0.5 M in Cu (guesstimate). I thought the colour may by now be
entirely due to CuCl<sub>4</sub><sup>2-</sup>, i.e. Cu [+II], because there was plenty Cl<sup>1-</sup> present too
but clearly that was not the case: green solutions of CuCl<sub>4</sub><sup>2-</sup> precipitate as blue
Cu(OH)<sub>2</sub>.n H<sub>2</sub>O with NaOH.
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12AX7
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http://webpages.charter.net/dawill/tmoranwms/Chem_Cu.html#Ox...
The picture is reasonably representative of the color of copper oxychloride.
Mixed with some Cu2O, you'd get a drab color.
Tim
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CyrusGrey
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| Quote: |
Initially it indeed does form the brown complex, but reaction goes on. Finally, the solution becomes colorless and CuCl2(-) is formed.
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Well, I dunno Woelen. It has been 2 or 3 days and the brown color hasn't gotten any lighter, though I do have a layer of CuCl dust building up at the
bottom of the test tube.
Is the CuCl2(-) complex unstable?
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woelen
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I dumped the colorless solution of CuCl2(-) in 50 ml of cold water. No precipitate was formed at all. The solution remained colorless and clear, also
after stirring.
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blogfast25
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No oxidation either? Have you tried adding ammonia?
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CyrusGrey
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Sorry Woelen, I take my comment back. It is turning clearer. It took about a week for it to be noticeable, but now I can see my lab lights
through the solution in the test tube. I'm guessing that the copper was passivated somewhat by the layer of CuCl formed on it. Either that or the
reaction is just very slow.
When it becomes completely clear I'll take the solution and expose it to the atmosphere without copper and see if it changes back to the brown
complex.
My hypothesis is that the brown solution is stable on exposure to both metallic copper and atmospheric oxygen at once, but not one or the other.
[Edited on 18-3-2008 by CyrusGrey]
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