Re: Question: Reaction between copper (II) and copper (0) in concentrated HCl

From: Wilco Oelen (photo_at_woelen.nl)
Date: 08/28/04 

Date: 28 Aug 2004 12:33:29 -0700

farooq_w@hotmail.com (Mohammed Farooq) wrote in message news:<66756669.0408270634.402ef95a@posting.google.com>...
> photo@woelen.nl (Wilco Oelen) wrote in message
> > I'll try this. See if I can get some crystals of K2[CuCl4].
>
> Please do. Do this at a mAcro-scale so that we can compare the
> properties of this one and your complex. In the meantime I will search
> for a more refined method of preparing K2[CuCl4].
>
I tried obtaining K2[CuCl4], but up to now I did not succeed. If the
concentration becomes too high, then I get very fine white crystals,
which dissolve very well in water. This can hardly be anything else
than KCl. During preparation of KCl from KOH and HCl I already
noticed, that this salt does not dissolve very well in conc. HCl.
Quite some water needs to be added to the mixture to keep the KCl in
solution and when much water is boiled away, then KCl crystallizes on
cooling down. The copper complex remains in solution.
Probably K2CuCl4 needs to be prepared in another way. If you can find
a way of preparing it in literature then I'd like to read about that.
In the meantime I'll look on Internet, whether I can something about
tetrachlorocuprate.

> > I prepared a CuCl solution with excess copper (as also described in
> > the last URL I have given) and added some CuCl2-solution. Mixing the
> > almost colorless CuCl-solution and a green CuCl2-solution results in a
> > very dark brown (almost black) liquid at once. This is precisely what
> > I expected.

>
> This is interesting, indicating that it is indeed a mixed valency
> complex. Our next target is somehow isolate the brown colored species.
> Can you do one more thing; To the brown solution add few mL of 33%
> H2O2 (to avoid dilution- hence any color change due to dilution) and
> note the color change. If there is no color change then heat slowly to
> avoid vigorous decomposition of peroxide. If there is no color change
> then you will rule out the existence of Cu(I) in it. I believe if any
> mixed valency Cu(I)-Cu(II) complex is formed it would not be so stable
> to resist oxidation by hydrogen peroxide.
The following experiment I have done:
Prepare a very concentrated dark brown solution from hot solution of
CuCl2 in conc. HCl and a fairly large amount of copper wire. Then pour
the liquid in a tenfold volume of water. A compact white precipitate,
consisting of little white crystals is formed. The liquid above the
precipitate is light blue. This precipitate can be rinsed with a small
amount of water to remove most of the acid and excess Cu(2+). CuCl
hardly dissolves in water. This again proves that the dark brown
liquid contains a lot of Cu(I) and also Cu(II).
>From this CuCl, I again made a dark brown solution, by simply adding
conc. HCl. The CuCl dissolves and the liquid becomes yellow/brown
(some oxygen already is in the HCl). On shaking with air contact, the
liquid quickly becomes dark brown again. To this, I added a small
amount of H2O2 (30%). The dark brown liquid _at once_ becomes green on
addition of the H2O2 and the liquid starts bubbling strongly and the
liquid becomes warm. A faint smell of Cl2 can also be observed,
apparently some of the HCl is oxidized to Cl2 by the H2O2. In order to
check that all copper (I) is really oxidized to copper (II) I diluted
again with water and now, no white precipitate was formed, there was
only a clear light blue solution. Pictures of this complete experiment
are shown at the following URL:

http://www.woelen.nl/chem/exp0006/exp0006.htm

Conclusion: The concentrated brown complex is quickly and completely
oxidized by H2O2 and is converted to concentrated dark green copper
(II) complex.

>
> Note that once you added a base to coper salt, first a blue ppt is
> formed, when more base is added it becomes dark brown, somthing
> similiar to the color your complex. Does that give us a clue?
If I add a base to a pure copper (II) salt, then I get a sky blue
precipitate of Cu(OH)2. Adding a STRONG base to a solution of a copper
(I) salt yields a yellow precipitate. Copper (I) salts, however, can
only be in solution as complex, so the comparison is not really good.

>
> In your last experiment you describe a colored precipitate which is
> soemthing like brown yellow solid. I will tell you of an interesting
> experiment that I used to do when I was in 8th grade. If you
> electrolyse a concentrated NaCl solution using thin copper wires as
> electrodes and use high voltage ie 9 Volt DC current, you would get
> exactly the same yellow colored voluminous ppt. sitting in the bottom
> of the electrolytic cell. You would see the anode getting green and
> after 15 minutes or so something yellow colored settles in the bottom.
> I used to prepare metal hydroxides this way, using the required metal
> as anode and carbon cathode (taken out from a dry cell, washed and
> boiled in water), conc. brine solution and high voltage. Zinc and iron
> behave similarly. In due course you might do some experiments with
> them. A modification that one can try is to electolyse a relatively
> concentrated HCl (in an open room) using copper electrodes, or try
> different metals and get nice and new results.
Yes, I also have noticed the yellow stuff with electrolysis. As a
young boy I also did that kind of experiments and at that time I also
was puzzled by the yellow stuff, formed at the copper anode. I indeed
can imagine that this yellow compound is the same or very similar to
the yellow compound, which I now created in the last experiment.

>
>
> > I'll also try mixing with some ethanol. I do not have pure ethanol
> > (this is incredibly expensive in the Netherlands, due to taxes), but
> > denatured ethanol. This denatured alcohol is colorless, mixes
> > perfectly well with water and on evaporation does not leave any
> > residue, just as pure ethanol. I don't know what is used for
> > denaturing it, I think some ester, because it has a somewhat fruity
> > odour. I hope it is suitable. I do have pure methanol, do you think
> > that is suitable?
>
> Methanol is quite toxic, its vapor makes my head ache and so does
> alcohol (perhaps allergic to organic solvents) so for the time being
> use denatured ethanol. Our purpose is reduce the polariy of solvent
> and since methanol is more polar than ethanol so it would not be as
> effective as the latter. You can very well use denatured alcohol for
> hastening the crystallization process.
The brown complex apparently is soluble in ethanol. I prepared a very
concentrated solution of the brown stuff and added a small amount of
this to the tenfold volume of ethanol. No crystals were formed, not
even after several minutes. I did the same with acetone. Again, no
crystals were formed.

As a counter test, I added some solid CuCl2.2H2O to some pure acetone,
and to my surprise the solid dissolves in acetone. The solution does
not become green or blue, but it becomes yellow/brown. This I can
explain, because I think that while the CuCl2.2H2O dissolves in
acetone, it is split in free water molecules and molecules of
anhydrous CuCl2. The copper does not dissolve in ionic form.

I did the same with CuSO4.5H2O, to see whether this dissolves in
acetone. This, however, does not dissolve at all, not the faintest
color can be observed, and even a very small pinch of solid does not
dissolve in acetone. The crystals of CuSO4.5H2O remain sky blue in
acetone.

Based on these experiments, I'm afraid that it will be very difficult
to isolate the dark brown compound and maybe it does not even exist
outside a concentrated solution of a chloride. As I already mentioned
above, on addition of water, it decomposes to solid crystalline CuCl
and the copper (II) goes in solution as blue Cu(2+)(aq).

Right now, I propose a complex of the following formulation:

   Cu Cl (x-)
     N M

I think that we have a multi-metal center complex with the coppers in
the complex having different oxidation states (that may explain the
very strong color) and some chloride-ligands around it. What the value
of N, M and x are, I don't know. Possibly x equals 0, explaining the
solubility in organic solvents.
This is just theory, I don't see a method to prove (or disprove) this
hypothesis.

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