Strange result from cuprous iodide experiment with EDTA

Quote: Originally posted by DrMario

I synthesized cuprous iodide (CuI) by adding a solution of potassium iodide to Cu(II) sulfate solution. I obtained the tell-tale dark brown solution of elemental iodine in water, and a precipitate that ended up being (after many decantations) off-white. When I reacted this precipitate with diluted ammonium hydroxide, the precipitate dissolved and I obtained a deep royal blue solution which I guess is some sort of complex with the cuprous ion (Cu(I)) I have never heard about before. But this is not what surprised me the most: when I added a solution of EDTA disodium salt to another sample of CuI precipitate, I obtained no significant dissolution, apart from a light green tint. The precipitate stayed mostly in the water, unaffected, except as I said, the water became very lightly green tinted. HOWEVER... what shocked me is that the solution started smelling rather pleasantly of lemon or lime. And it was quite distinct. What up with that?? What sort of organic compound did I obtain that it smells like lime? And perhaps contains iodine? Maybe?

Quote: Originally posted by DraconicAcid

The deep royal blue is actually the cupric ammine complex- the cuprous one is colourless, but oxidizes on exposure to air. The EDTA result is surprising. Then again, Ag(I) and Cu(I) tend to form two-coordinate complex ions, which EDTA doesn't really lend itself to. [Edited on 23-11-2014 by DraconicAcid]

Quote: Originally posted by DrMario

I would very carefully (and very, very respectfully) tend to disagree on the cupric ion formation... at least the mechanism you describe doesn't sound credible: the precipitate completely dissolved within a fraction of a second, in the NH4OH solution. There was no time for meaningful interaction with air.

Quote: Originally posted by DrMario

According to this ChemWiki: http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_... the [Cu(NH3)2]+ complex is stable!

Quote: Originally posted by DraconicAcid

Quote: Originally posted by DrMario   I would very carefully (and very, very respectfully) tend to disagree on the cupric ion formation... at least the mechanism you describe doesn't sound credible: the precipitate completely dissolved within a fraction of a second, in the NH4OH solution. There was no time for meaningful interaction with air. It dissolves quickly because it forms [Cu(NH3)2]+, but this (being a d10 system) is colourless. It turns blue because of oxidation- either through exposure to air, or oxygen dissolved in solution. [Edited on 23-11-2014 by DraconicAcid] [Edited on 23-11-2014 by DraconicAcid]

Quote: Originally posted by DrMario

Quote: Originally posted by DraconicAcid   The deep royal blue is actually the cupric ammine complex- the cuprous one is colourless, but oxidizes on exposure to air. The EDTA result is surprising. Then again, Ag(I) and Cu(I) tend to form two-coordinate complex ions, which EDTA doesn't really lend itself to. [Edited on 23-11-2014 by DraconicAcid] I would very carefully (and very, very respectfully) tend to disagree on the cupric ion formation... at least the mechanism you describe doesn't sound credible: the precipitate completely dissolved within a fraction of a second, in the NH4OH solution. There was no time for meaningful interaction with air. [Edited on 23-11-2014 by DrMario]

Quote: Originally posted by unionised

Did you consider the fact that the ammonia solution would contain air in solution before you added it to the CuI?

Quote: Originally posted by DrMario

Quote: Originally posted by unionised   Did you consider the fact that the ammonia solution would contain air in solution before you added it to the CuI? Yes - so let's assume that this oxygen will oxidize part of Cu(I) to Cu(II) and turn the solution somewhat blue. If I shake the container after this, more oxygen will dissolve and more Cu(I) should oxidize to Cu(II), creating a darker blue solution. But this latter phenomenon is not observed.

Trust me, DA et al are right about this: the copper(I) ammine complex (like basically all Cu(I) compounds) does oxidide quickly to the Cu(II) ammine complex.

Perhaps it helps to understand this as follows:

The Cu(I) ammine complex is in equilibrium:

Cu(NH₃₂⁺ < === > Cu⁺ + 2 NH₃, the formation equilibrium.

The free Cu⁺ can then be oxidised by air oxygen and the resulting cupric ions complex with the excess ammonia, to form the blue Cu(II) tetrammine complex ion. According to the Le Chatelier Principle the removal of cuprous ions further pulls the formation equilibrium to the right.

So the cuprous diammine complex is quite 'stable', yet still prone to air oxidation at the same time.

Also, the cupric tetrammine complex is so intensely coloured that even small concentrations with create the illusion that all the cuprous has disappeared.


[Edited on 23-11-2014 by blogfast25]

Quote: Originally posted by blogfast25

Quote: Originally posted by DrMario   Quote: Originally posted by unionised   Did you consider the fact that the ammonia solution would contain air in solution before you added it to the CuI? Yes - so let's assume that this oxygen will oxidize part of Cu(I) to Cu(II) and turn the solution somewhat blue. If I shake the container after this, more oxygen will dissolve and more Cu(I) should oxidize to Cu(II), creating a darker blue solution. But this latter phenomenon is not observed. Trust me, DA et al are right about this: the copper(I) ammine complex (like basically all Cu(I) compounds) does oxidide quickly to the Cu(II) ammine complex. Perhaps it helps to understand this as follows: The Cu(I) ammine complex is in equilibrium: Cu(NH<sub>3</sub><sub>2</sub><sup>+</sup> < === > Cu<sup>+</sup> + 2 NH<sub>3</sub>, the formation equilibrium. The free Cu<sup>+</sup> can then be oxidised by air oxygen and the resulting cupric ions complex with the excess ammonia, to form the blue Cu(II) tetrammine complex ion. According to the Le Chatelier Principle the removal of cuprous ions further pulls the formation equilibrium to the right. So the cuprous diammine complex is quite 'stable', yet still prone to air oxidation at the same time. Also, the cupric tetrammine complex is so intensely coloured that even small concentrations with create the illusion that all the cuprous has disappeared. [Edited on 23-11-2014 by blogfast25]

Quote: Originally posted by DrMario

I know the tetraamminecopper(II) ion is intensively blue. I don't dispute this. I don't _directly_ dispute any of the things you wrote, but am posing the question, and I repeat: why do I not see an increase in intensity of colour, while shaking the bottle with the solution? I have another question: If Cu+ ---> Cu2+, what is the cation we're talking about?

Quote: Originally posted by No Tears Only Dreams Now

Where did your EDTA and iodine come from?

Quote: Originally posted by DrMario

Quote: Originally posted by No Tears Only Dreams Now   Where did your EDTA and iodine come from? The iodine comes from the disproportionation of CuI2 into CuI + iodine: 2CuI2 ----> 2CuI + I2 And the EDTA:Na2 comes from Sigma Aldrich. Or maybe I don't understand your question.

Quote: Originally posted by No Tears Only Dreams Now

Quote: Originally posted by DrMario   Quote: Originally posted by No Tears Only Dreams Now   Where did your EDTA and iodine come from? The iodine comes from the disproportionation of CuI2 into CuI + iodine: 2CuI2 ----> 2CuI + I2 And the EDTA:Na2 comes from Sigma Aldrich. Or maybe I don't understand your question. I'm asking what source you obtained the starting reagents from, you answered on the EDTA already; I too am extremely puzzled by the smell you mentioned.