morsagh
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copper complex
I am thinking about reaction of calcium chloride and copper chloride to produce Ca[CuCl4 ] and then add it to
[Cu(NH3)4] to produce something like this: [Cu(NH3)4][CuCl4 ]
Can be this done or is it just a waste of time?
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morsagh
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*[Cu(NH3)4]SO4
(like to produce complex and precipitate CaSO4)
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Boffis
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Waste of time!
Copper has only slightly greater affinity of NH3 than it does for H2O therefore copper ammino complexes exist in equilibrium with free ammonia in
aqueous solution and free ammonia with ammonium and hydroxide ions. Since copper hydroxychlorides and hydroxides are extremly insoluble aqueous copper
chloride solution will instantly mop up any hydroxide ions around and for a precipitate of basic copper chloride (copper hydroxide will not for under
these conditions because the concentration of hydroxide ions will never reach a sufficient level as they exist inequilibrium with ammonia-ammonium
ion). This is also why there are so many basic copper chlorides in nature. The chlorocuprate ion is also very labile in aqueous solution and needs as
significant excss o f chloride to maintain the (CuCl4)2- ion, as you dilute the solution it unravels into ions with a lower chloride ion content.
What would be interesting would be to saturate say 25% aqueous ammonia solution with ammonium chloride to suppress hydroxide ion formation and then
add copper chloride solution to see what you get. The chloro complexes are yellow to orange brown while the ammino complexes deep blue.
Ammonia in water is most just ammonia in solution with only a little ionisation according to:
NH3 + H2O <<<-> NH4+ +OH-
If you add a large concentration of ammonium ions (as ammonium chloride) you will drive the equilibrium to the left and reduce the hydroxide ion
concentration, perhaps to the point where basic salts will not precipitate however, predicting the make up of the various complex ions will require a
lot of stability data and maths that I can't be bothered to do for you. I am sure this data is available as these are all well studied systems.
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aga
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You just want Calcium Sulphate ?
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blogfast25
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Quote: Originally posted by Boffis  | The chlorocuprate ion is also very labile in aqueous solution and needs as significant excss o f chloride to maintain the (CuCl4)2- ion, as you dilute
the solution it unravels into ions with a lower chloride ion content.
What would be interesting would be to saturate say 25% aqueous ammonia solution with ammonium chloride to suppress hydroxide ion formation and then
add copper chloride solution to see what you get. The chloro complexes are yellow to orange brown while the ammino complexes deep blue.
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The tetrachlorocuprate ion - [CuCl<sub>4</sub>]<sup>2-</sup> - in aqueous solution is green.
As regards basic copper chlorides, there are three trimorphs of the same composition Cu<sub>2</sub>Cl(OH)<sub>3</sub>:
atacamite, paratacamite and botallackite.
[Edited on 24-9-2015 by blogfast25]
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Amos
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Combining solutions of calcium chloride and copper(II) sulfate in a stiochiometric ratio will cause most of the calcium to precipitate out as calcium
sulfate, leaving you with a solution of mostly copper(II) and chloride ions; it will likely be green in color.
If you filter off the solids, and then treat your pure solution with aqueous ammonia, the first thing that will happen is a precipitate that is pale
blue to pale green in color, consisting of copper oxychloride and likely some copper(II) hydroxide as well. However, if you keep adding more ammonia,
eventually these compounds will dissolve completely to yield a royal blue solution containing the [Cu(NH3)4]2+ complex ion, as well as the original
chloride ions.
If you'd like to try isolating this in solid form, I recommend you use as concentrated a solution of copper(II) chloride as you can in the second
step, so this may entail boiling down your greeen solution until it is syrupy and possibly begins to crystallize. The idea here is to reduce the water
content as much as possible; I'll say why later. By using as concentrated ammonia as you can manage(I suspect even just household ammonia will be
acceptable but your yield will be lower), once again produce the royal blue solution containing the complex, stopping the addition of ammonia when no
solids remain. Finally, decant this solution into a larger container of acetone or isopropanol; the more water you have, the more acetone it will take
to precipitate the complex. I haven't made this myself, but this is the same procedure I have used to isolate tetraammine copper(II) sulfate several
times with ease; that complex can easily be dried in a stream of warm and then stored permanently in a sealed glass container.
*EDIT: So I just read your post saying you actually want the sulfate complex. Then just add ammonia to copper(II) sulfate solution until a precipitate
forms, and then continue adding until all of the precipitate dissolves. Dump the remaining solution into acetone, filter out the solid, you're done.
[Edited on 9-24-2015 by Amos]
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Boffis
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@blogfast25, I can assure you that the tetrachlorocuprate ion is yellow. If you add a suitable large cation such as Cs+ bright yellow prismatic
crystals form. If your aqueous solution is green then you haven't got a sufficient concentration of Cl- ions to convert all of the copper to the
tetrachloro complex. Interestingly if you isolate the yellow crystals from thier parent solution and then dissolve them in water you get a green
solution that deposits orange brown crystals. I haven't tried this with other alkali chloride salts.
As for the basic copper chlorides you have chosen only one stoichometric formula of which there are not three but possibly FIVE polymers: atacamite,
paratacamite, clinoatacamite, botallackite and anatacamite. There is some evidence, though, to suggest that all paratacamites contain a significant
and essential amount of another M2+ ion and so may not constitute a "pure" basic copper chloride. Anatacamite is supposed to be the triclinic
polymorph but is in doubt now. There are however, many other basic copper chlorides: eg bobkingite, calumetite, anthonyite, belloite etc.
I would concur with Amos' comment entirely.
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