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

From: Wilco Oelen (photo_at_woelen.nl)
Date: 08/24/04 

Date: 24 Aug 2004 13:51:10 -0700

farooq_w@hotmail.com (Mohammed Farooq) wrote in message news:<66756669.0408240645.6b592d6e@posting.google.com>...
> Wilco wrote
> >I'll see if I can do this. I do not have KCl. My chem supplier (just
> a
> >normal drugstore) does not carry that, because NaCl is present at low
> >cost and large volumes. Why KCl? Do you think this works better than
> >NaCl? I know that certain Na-salts are hygroscopic and K-salts are
> >not, but of course things can also be the other way around (e.g. NaF
> >vs. KF).
>
> It is said (I do not know the reason clearly) that larger cations
> accelerate the crystallization process rather the separation of
> complex ion from a solution is readily achieved when a large cation is
> added to the solution. The preference being NH4 > K .
I also read once about this. I also read that if cation and anion have
sizes very close to each other, that solubility is less, but I do not
exactly understand these 'rules' and I'm not quite sure that these
rules always hold.
I added some NH4Cl to a dark brown solution, prepared from CuCl2.2H2O,
Cu2O and HCl (30%). Surprisingly, the color of the liquid changes
considerably. It becomes MUCH lighter and sepia/green. NH4Cl hardly
dissolves in the liquid, I can understand that, because of the high
concentration of Cl(-). However, sufficient NH4Cl dissolves, such that
the color of the liquid changes a lot. Interesting... can this have
something to do with the fact that CuCl2.2H2O does not dissolve easily
in concentrated NH4Cl (see further below in this reply)?

> You must have
> used KMnO4 but not NaMnO4, due to high solubility of sodium
> permanganate.
This remark I do not understand. Did you do some experiment with
permanganate, related to the copper chemistry, we are investigating
now. I did not use MnO4(-) in any of these experiments.

> Add another salt of potassium if Kcl is not available
> but never KF (might attack the glass under acidic conditions).

I prepared solutions of CuCl2 of approximately the same concentration
(in copper) in different other solutions.
With saturated NH4Cl-solution I get a lime-green solution, the copper
chloride only dissolves with difficulty and some heating is needed to
get a reasonable amount in solution. When copper is added to this
liquid, then it quickly becomes dark brown.
With saturated NaCl-solution I get a bright green solution. The copper
chloride easily dissolves. When copper is added, I get a brown
solution, but this reaction does not go as fast as with the cooled
down NH4Cl/CuCl2 solution. I think this is due to the fact that
concentrated NH4Cl-solution has a higher concentration of Cl(-).
With conc. HCl (30%), the solution becomes yellow/green/brown,
somewhat hard to describe. Copper chloride dissolves very easily in
conc. HCl. With this solution, the formation of the brown compound on
addition of copper is fastest. This solution, however, also has by far
the highest Cl(-) content.

With K(+) cations, I'll need to do some preparative work.
I'll try KOH or K2CO3. I have both of these. From these I'll prepare
an acidic solution of KCl and use this for making the brown liquid. I,
however, do not expect much surprising results from this, given the
results, described above.

I think, that the color of the complex, both for only copper (II) and
for the mixed copper (I) / copper (II) complex does not depend
strongly on the cation in solution. I think that the concentration of
chloride is much more important. As a test, I also did an experiment
with 10% HCl and then the results are closer to the results with
saturated NaCl-solution.

>
> >The latter I do not understand. [CuCl4](-)? Isn't it [CuCl4](3-)?
>
> Correct, this was a typing mistake.
OK, can happen :-)

>
> >I indeed have observed the white crystals of CuCl on dilution of a
> >very concentrated dark brown solution.
>
> I think if you boil a copper chloride solution with excess of copper
> it should become colorless. Did you ever try boiling the brown
> solution, what do you get?
Yes, you get an almost colorless liquid. It is not even necessary to
boil the liquid. With some patience, you get a colorless liquid
without boiling. I did an experiment, with a large excess of copper in
a stoppered test tube. After a few days of standing, the lower part of
the liquid is colorless, only the upper part still is somewhat brown.
On shaking the test tube, while still being stoppered, the liquid
becomes light green/sepia and when it is allowed to stand for a while,
the liquid around the copper becomes colorless again within a few
hours.

>
> >This may be interesting. I'll try to make some plain CuCl4(2-) in
> >conc. HCl, also in NaCl-solution and add some K-salts or NH4Cl, but I
> >hardly can believe such a thing.
>
> K2[CuCl4} is well known. Addition of alcohol reduces the solubility of
> complex in water and hastens precipitation of complex ion with a
> suitable charge neutralizing ion. Also check the effect of simply
> cooling the brown solution down to -17 oC by a ice-salt mixture, does
> the color change?
I cooled down the brown solution in a refrigerator at -18 C, and the
color of the liquid did not change. After this, I heated the liquid
quickly to 60 C and still the color did not change. During this, the
test tube remained stoppered all the time in order to exclude the
possible influence of oxidation by oxygen from air. This stoppering
also is the reason, why I did not go further than 60 C on heating.

>
> >I can image such color dependencies
> >in the solid state but not in liquids. But nevertheless, I'll try and
> >let you know
> May be, but the author does not mention solid or solution state.
I tried some different cations (NH4(+), Na(+), H3O(+)) in solution,
see above, but comparison is not really fair, because concentrations
of chloride also differs considerably in the three experiments.

I think, that all these experiments, described above allow the
following conclusions:
The copper (I,II) / chloride complex does not depend strongly on the
cations present in the solution.
The complex does depend on the concentration of chloride in the
solution.
The complex does not strongly depend on temperature (at least not in
the range I checked).

Probably again, some literature search is needed to find more about
this complex. I'll look on internet on copper (I) and copper (II), but
I'm afraid, I will not find much relevant information.

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