Ok, I am purely curious because I think this would be an excellent learning exercise.
I was just thinking and wondered if there was a way to measure dissolved CU(1+) ions in solution, in other words, how much CU(1+) remains Un-oxidized
to the CU(2+) ion?
Is this possible without intricate and expensive measuring equipment?
thanks.m1tanker78 - 20-7-2012 at 14:33
What anion? Copper(I) chloride is practically insoluble in water and dissolves slow enough in dilute HCl so that you can qualitatively see how much is
there. The unintended samples of this compound I've produced (by electrolysis) are off white/khaki colored at first. It eventually oxidizes to
copper(II) chloride (in dilute HCl) which is very soluble.
TankCHRIS25 - 20-7-2012 at 15:15
Hi, yes I know that copper1 chloride is insoluble in water. But not when there is an excess of HCl in the solution. this is a copper chloride
etchant and does not work the same way as if there is just water. I don't know how to explain things, so please see other posts regarding the etchant
solution, then you will understand. There are plenty of free CU(1+) ions in my solution, and I know they can be measured trouble is I do not know how
to do it. I wanted to learn something new relevant to what I am busy with.
I have to say this: that the HCl Keeps the CU(1+) dissloved and then oxidizes it to CU(2+), at which point it is this that dissolves the actual metal
copper and in the process the resultant reaction reverts back to CU(1+) and the whole cycle continues again. There never is any precipitate of
CU(1+)chloride unless the HCl becomes absolutely exhausted. Also the CU(1+) ions are interlocked with the chloride ions in such a complex way the
precipitation is impossible, So they Can be Measured.
[Edited on 20-7-2012 by CHRIS25]m1tanker78 - 20-7-2012 at 15:46
and I know they can be measured trouble is I do not know how to do it.
???
Hmm, sorry can't help you with that. Try looking at some careful experiments using acetylene gas maybe? The interaction with chloride is heavily pH
dependent. If you're eager to learn something, the 'separation of US nickel' thread contains a lot of good info and experiments posted by SM user
"blogfast" concerning copper solutions/complexes and precipitation of the hydroxychloride. It's a bit scattered but worth it IMO.
thanks tanker but I have to stay on target, too many digressions recently, useful digressions but I can not possibly take on any more projects. This
one has to be simpler than ploughing through the nickel and spending hours dissecting a tiny piece of info from science talk, but thanks for the input
anyway.bbartlog - 21-7-2012 at 10:03
Quote:
the HCl Keeps the CU(1+) dissloved and then oxidizes it to CU(2+)
I believe it's normally atmospheric oxygen that oxidizes the Cu(I). HCl isn't an oxidizer. It just reacts with the cupric oxide(s) to form water and
CuCl2.
This does suggest one way to estimate the Cu(I) present in the solution: draw off a sample, make sure that there is excess HCl present (enough so that
all CuCl can be converted to CuCl2), and then see how much oxygen the resulting solution will take up. If you have a good scale, measuring the weight
increase of the solution over time might be easier than measuring the volume of gas taken up.kristofvagyok - 21-7-2012 at 12:09
Add some potassium iodide. The Cu(2+) will form CuI2 what will immediately decompose to Cu(1+) and to elemental iodine. And the elemental iodine could
be titrated easily on various ways.
[Edited on 21-7-2012 by kristofvagyok]bbartlog - 21-7-2012 at 13:02
That works as a way to estimate Cu(II), but would leave you in the dark as to how much Cu(I) was in solution, unless you either already knew the total
amount of dissolved copper *or* also dried and measured the precipitated CuI (again to determine the total copper and by subtraction the amount that
was in the +1 oxidation state).CHRIS25 - 22-7-2012 at 01:02
@Kristof...and @bbart.... Thanks. Both ideas give me something to work on and research into. At least now you have given me a foundation from
which I can look into.
