As some of you already may know, I have something with the element copper
Again, I made a special complex of this element. This time the red compound CsCuCl3. It really is remarkable that a combination of an alkali metal
ion, copper(II) and chloride can lead to a brightly colored orange/red compound. The following experiment shows how easy this can be made:
Will read it later but I see you're making good use of your new stockpile of Cs salts!
Very interesting. Do you have a reference for the existence of CsCuCl3 or is this your own deduction? If the latter couldn't it be Cs2CuCl4?
[Edited on 13-2-2011 by blogfast25]kmno4 - 13-2-2011 at 15:25
Literature says that CsCuCl3 is "garnet-red hexagonal prisms"
(doi: 10.1039/JR9470001662) and Cs2CuCl4 is "yellow-orange
needles" (doi: 10.1021/j100819a015, 10.1021/ja01125a012)
Unfortunately I did not see any elemental analysis of these compounds.....
See also intersting paper "A Redetermination of the Crystal Structure of CsCuCl3" (doi: 10.1021/ic50036a025), and from there:
Dark red-brown crystals of CsCuCl3, were obtained by evaporating to dryness a solution containing copper and cesium chlorides.
A molar excess of copper chloride was used to prevent formation of Cs2CuCl4. The crystals appeared as hexagonal prisms often capped by
hexagonal pyramids.
About KCuCl3 and NH4CuCl3 (both dark-red) you can read in another paper (doi: 10.1063/1.1733520)woelen - 13-2-2011 at 23:48
@kmno4: Many thanks for your comments. This is good to know. The books I have are very brief on this kind of complexes and you give me some new
insight. Cs2CuCl4 is yellow and this is exactly what I also observed. I accounted this for the size of the particles, but now I see that this is due
to excess CsCl. In my experiment I first added a big drop of the solution of CuCl2 in HCl to the solution of CsCl. This gives mainly yellow
precipitate. On adding the rest of the CuCl2 and shaking I get the red precipitate. With your info this nicely can be explained.
I'll change the page and mention the yellow Cs2CuCl4.
I also read about the K-salt, but apparently this cannot be made by simply mixing the solutions. KCl precipitates from the solution and all copper
remains in solution. As a home chemist, I unfortunately have no access to these doi: papers.
I understand that with my experiment I just raised the tip of the curtain and a whole new area of chemistry is to be explored over here. Some things I
have in mind:
- playing with concentrations
- slow evaporation of somewhat more dilute solutions of KCl and CuCl2 in dilute HCl.
- bromine analogues of these chemicals (I have anhydrous CuBr2, HBr (40%) and CsBr).
But for my website I first chose to do the chlorine-based experiments, because of their simplicity and the fact that very common chemicals can be
used.
[Edited on 14-2-11 by woelen]DJF90 - 14-2-2011 at 03:59
Papers can be requested in the references section.
I'll save you the hassle this one time...
Attachment: ic50036a025.pdf (488kB) This file has been downloaded 934 times
Attachment: j100819a015.pdf (278kB) This file has been downloaded 953 times
Attachment: ja01125a012.pdf (584kB) This file has been downloaded 921 times
Attachment: JR9470001662.pdf (852kB) This file has been downloaded 1022 times
blogfast25 - 14-2-2011 at 08:33
It shouldn't be too difficult to distinguish between both complexes in a semi-analytical way. Adding sulphide should precipitate all copper as CuS,
the released CsCl could be precipitated almost quantitatively as perchlorate or similar...woelen - 14-2-2011 at 11:35
@DJF90: Thanks for the articles. I used that information for updating the webpage. It now mentions the structure described in the 1966 article and it
mentions the yellow color of Cs2[CuCl4].
@blogfast25: I indeed think that this can be done, but for the sake of the webpage it suffices to refer to the literature. My observations now
perfectly fall in place with the information from kmno4 and DJF90.woelen - 26-2-2011 at 10:27
I have been experimenting further with similar complexes and I also found the following:
- A solution of CoCl2 in conc. HCl (dark blue) gives a microcrystalline dull blue precipitate of a cesium chlorocobaltate complex when solution of
CsCl is added.
- A solution of FeCl3 in conc. HCl (deep yellow/brown) gives a beautiful microcrystalline orange/yellow precipitate when a solution of CsCl is added.
I also did the experiment of dissolving CuBr2 in 40% HBr (this gives a very dark purple/red solution) and adding a solution of CsBr in water. This
leads to a crystalline precipitate with a very dark purple/black color. The reaction with bromide is not as successful as the reaction with chloride.
CuBr2 does not dissolve as well in HBr as CuCl2 does in HCl, so lower concentrations can be reached and hence only a small part of the copper(II)
precipitates as a complex.
There is another interesting observation. When the concentrated solution of FeCl3 in conc. HCl is added to a concentrated solution of CsCl in water,
then immediately a compact crystalline precipitate is formed, which quickly settles at the bottom, but the liquid also heats up noticeably! It changed
from cold (15 C or so) to luke warm (30 C or so). It find this remarkable, because the noticeable production of heat can only be due to the separation
of the cesium salt of a chloroferrate(III) complex.
Pictures of the cobalt and iron complex will follow soon.
[Edited on 26-2-11 by woelen]Mixell - 7-7-2011 at 11:49
I did a a small scale experiment:
Upon addition of concentrated NH4Cl solution to a copper chloride/hydrochloric acid solution the ammonium chloride precipitates. upon more addition of
ammonium chloride solution (about 5 time more in volume then the CuCl2*HCl) the NH4Cl redissolve's and a light blue precipitate settles.
Upon addition of concentrated NH4Cl solution to concentrated CuCl2, no precipitate appears.
Could the light blue solid be (NH4)2CuCl4?
Although it looks exactly like CuCl2*2H2O.Mixell - 7-7-2011 at 11:58
Nope, it dehydrates to form a brown solid, wrong call =\
The same experiment with BaCl2 yields nothing.
Are there any other cations that can form complexes with CuCl3 (-1) and CuCl4 (-2) (except cesium and rubidium)?
[Edited on 7-7-2011 by Mixell]Jor - 7-7-2011 at 12:14
Very nice, these compounds were also on my list, I found them in some book before.
I was going to use the dry method (IIRC it would be prepared by fusing anhydrous CuCl2 and CsCl.
Why don't you filter on a glass frit with vacuum? This is much faster than putting tissue under the frit. Optionally you can wash with ethanol and
ether, wich makes drying much easier.
[Edited on 7-7-2011 by Jor]blogfast25 - 7-7-2011 at 13:07
@mixell:
An interesting ammonium chloro complex is (NH4)2SnCl6 (there's a thread on this part of the forum on it). Also: (NH4)2PbCl6... Both quite easy to
make.DraconicAcid - 19-9-2020 at 11:08
Has anyone managed to make KCuCl3 or NH4CuCl3? According to the paper above, the latter can be made with ammonium chloride and copper(II) chloride in
anhydrous ethanol, but the ammonium chloride refused to dissolve or do anything at all with the solvent. Lion850 - 6-10-2020 at 14:39
Gents does anyone know if KFeCl4 is orange? See the photo of an orange salt that formed when I mixed a solution of FeCl3.6H20 in concentrated HCl with
a suspension of KCl in concentrated HCl. Most of the KCl did not react and is below the orange salt in the photo.
the ammonium chloride refused to dissolve or do anything at all with the solvent.
Ammonium chloride is basically insoluble in ethanol: 6g/L@20 C. There must be an error in communication somewhere; nobody would publish that
intentionally.
Maybe you gas CuCl2@EtOH with NH3? Followed by HCl, if necessary?woelen - 8-10-2020 at 01:20
I made the cesium salt CsFeCl4. That one is rusty orange, so that color you show in your picture matches quite well. It could very well be KFeCl4.
This kind of chemistry works better with cesium ions. Apparently, the bigger size of the Cs-ion makes crystal formation and precipitation easier. Some
complex anions also separate very well from the liquid with the [N(CH3)4](+) ion as counter-ion. In my experience, potassium ions and ammonium ions
work less well. Lion850 - 8-10-2020 at 16:20
Thanks for the feedback. I need to get some cesium salts....MidLifeChemist - 8-10-2020 at 18:27
