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Author: Subject: Why can't I dissolve my CuI in acetonitrile?
michalJenco
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[*] posted on 19-11-2019 at 13:53
Why can't I dissolve my CuI in acetonitrile?


Hello chemist friends.

I made some copper iodide by adding solution of potassium iodide to a solution of copper sulfate in 1:1 stoichiometric ratio. I filtered and washed the precipitate with water and then with ethanol to dissolve all elemental iodine that formed. I then let it thoroughly dry out outside. The product is completely insoluble in water and alcohol.

I wanted to make crystals and sources (wiki, researchgate) suggest it should be 7-8% soluble in acetonitrile.

However, I couldn't even get 0.5g of it dissolved in 50g of boiling acetonitrile. What could be the problem here? After filtering the acetonitrile "solution" no crystals or precipitate formed on cooling down.

My acetonitrile is pure and from a chem supplier.

Attached is a photo of my supposed CuI.

IMG_20191119_224841.jpg - 934kB
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DraconicAcid
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[*] posted on 19-11-2019 at 14:27


I wonder if it started photodecomposing like silver iodide would. It should be white.



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[*] posted on 19-11-2019 at 15:24


Quote: Originally posted by DraconicAcid  
I wonder if it started photodecomposing like silver iodide would. It should be white.


it should be white if pure, but most samples have an off white color as copper iodide easily adsorbs free iodine from the solution it was made from, a simple washing may be not enough.
as to why it didn't dissolve in acetonitrile could it be because the solvent was wet?
if it is so insoluble in water i assume that adding water to a solution of AcCN and CuI would crash it out of solution pretty quickly.





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fusso
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[*] posted on 19-11-2019 at 20:09


Did you try adding CuI to MeCN in very small increments, ie only a few grains of the powder? That way you can see if it's really that insoluble.

[Edited on 191120 by fusso]




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woelen
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[*] posted on 20-11-2019 at 00:28


Was the solution from which you made CuI neutral or acidic. I also made CuI once, but this worked best when the solution was a little acidic. Otherwise I obtained a brown/orange impurity, which I could not separate from the CuI and which did not disappear on adding a reductor like Na2SO3. I think that the brown impurity is Cu2O or a hydrated variation of that.

I made my CuI by adding a copper sulfate solution to a solution of KI (I used approximately stoichiometric amounts, one KI for one CuSO4.5H2O, with a little excess KI). I added an acidified solution of Na2SO3 to this. This works as a reductor. I dissolved the KI and acidified Na2SO3 in water (giving a pale yellow solution, due to coordination of I(-) and SO2) and dissolved the CuSO4 in another amount of water, and then mixed the solutions. An off-white precipitate is formed. Do not add too much Na2SO3! My material is not white, but it is greyish-tan, not reddish-brown like yours. I used it in an experiment with pyridine, making a fluorescent complex. I did not dry it, I just filtered it and rinsed with water and then stored under water. CuI, especially when humid, is air-sensitive, and is easily oxidized by oxygen from air.




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[*] posted on 20-11-2019 at 03:51


Quote: Originally posted by michalJenco  
After filtering the acetonitrile "solution" no crystals or precipitate formed on cooling down.

How does the solubility curve look like? There are systems where the solubility of a solid in a solvent does not change too much according temperature e.g. NaCl - H2O. There are even systems where solubility decreases as temperature increases (inverted solubility curve) e.g. CaSO4 - H2O, MnSO4 - H2O, Ce2(SO4)3 - H2O. Perhaps the some could exist for organic solvents, not only for H2O.
Cu(I) binds with acetonitrile into a complex (also with pyridine as woelen pointed out). I believe even if you are able to obtaining crystals from acetonitrile solution they won't be CuI anymore but a complex (and this complex could be very soluble in acetonitrile). Did you try to distill out the solvent after you did not obtain crystals on cooling down?
For a crystallization it could be helpful to introduce crystal seeds or at least scratch walls of supersaturated solution. Decades ago when I crystallized abietic acid it last few weeks, the system had consistency of honey, few days latency when no crystals, then the amount of crystals increased very slowly every day.
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[*] posted on 20-11-2019 at 04:05


and here some experiments how to grow CuI crystals and discussion about their discoloration - they used NH4I aqueous solution
https://sci-hub.tw/https://pubs.acs.org/doi/pdf/10.1021/acs....
https://sci-hub.tw/https://aip.scitation.org/doi/10.1063/1.5...
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rockyit98
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[*] posted on 20-11-2019 at 05:19


Quote: Originally posted by michalJenco  
Hello chemist friends.

I made some copper iodide by adding solution of potassium iodide to a solution of copper sulfate in 1:1 stoichiometric ratio.



dude! watch your chemistry.the ratio is 2:1 ( 4KI +2CuSO4 -------->Cu2I2 +I2 +2K2SO4)
Cu2I2 (Copper(1+)diiodide) is CuI but written differently.
if it helps, i made pure CuCl (Cu2Cl2) by passing a current through a Con. NaCl solution by using Cu wire which are 99.9% pure. 2NaCl + 2Cu +H2O ------> Cu2Cl2 +2NaOH +H2) i think KI might do better given Cu2I2 has much lower solubility.




far as i know everything is trying to be an equilibrium.
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[*] posted on 20-11-2019 at 08:04


Quote: Originally posted by Fery  
and here some experiments how to grow CuI crystals and discussion about their discoloration - they used NH4I aqueous solution
https://sci-hub.tw/https://pubs.acs.org/doi/pdf/10.1021/acs....
https://sci-hub.tw/https://aip.scitation.org/doi/10.1063/1.5...


Another possibility might be to dissolve the CuI in an aqueous iodide solution and then diffuse through an iodide-free gel, with crystallization taking place when the complexing agent drops in concentration. There's a mention of this technique as applied to cuprous chloride here:
https://www.sciencemadness.org/whisper/viewthread.php?tid=75...




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[*] posted on 20-11-2019 at 08:06


Quote: Originally posted by rockyit98  
[...]dude! watch your chemistry.the ratio is 2:1 ( 4KI +2CuSO4 -------->Cu2I2 +I2 +2K2SO4)[...]

Depends . . .
If you add a suitable reductor, you can use all iodine for binding to Cu and then you can indeed use a 1 : 1 ratio. I did that in my experiments with formation of CuI-pyridine complexes. Iodine/iodide is relatively expensive, while Na2SO3 (the reductor) is dirt cheap.




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michalJenco
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[*] posted on 20-11-2019 at 10:02


Wow, thanks for all the responses! I will answer some of your questions and give you updates.

My solvent is not wet, label says <2% H2O.

Adding water to the filtered acetonitrile "solution" really did give a white precipitate, it was even crystalline (tiny flakes which reflected light). I needed to add about the same volume of water to make it crash out, however, not just a tiny bit. I heated this new solution with precipitate to almost boiling and even added about 50% further acetonitrile, but it didn't all dissolve back. Right now I am waiting for the filtrate to cool down and maybe crystallise out whatever re-dissolved.

I figure some of the CuI I made really decomposed into an insoluble oxide and I didn't judge correctly whether anything dissolved at all.

I didn't add any acid when originally making the CuI, just CuSO4.5H2O and KI. Apparently it worked anyway. But thanks for your experimental experience, woelen.

I did not try to distill the solvent after the precipitate did not form. I did scratch the glass beaker with a glass rod and introduced seeds of the crude CuI, however. Didn't do anything.





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[*] posted on 20-11-2019 at 10:14


Quote: Originally posted by woelen  
Quote: Originally posted by rockyit98  
[...]dude! watch your chemistry.the ratio is 2:1 ( 4KI +2CuSO4 -------->Cu2I2 +I2 +2K2SO4)[...]

Depends . . .
If you add a suitable reductor, you can use all iodine for binding to Cu and then you can indeed use a 1 : 1 ratio. I did that in my experiments with formation of CuI-pyridine complexes. Iodine/iodide is relatively expensive, while Na2SO3 (the reductor) is dirt cheap.


So that is why so much elemental iodine was formed in my synthesis .. next time I will use Na2SO3 as a reductor and see how I2 formation reduces!
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[*] posted on 20-11-2019 at 23:40


Exactly, and that also is the reason why I added some acid, because Na2SO3 is somewhat alkaline and without acid you get all kinds of side reactions, with formation of oxides and hydroxides. As an alternative, you could add sodium metabisulfite instead of sodium sulfite. Just be sure that the solution remains on the acidic side of the pH-scale.



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