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jamit
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silver iodide or copper iodide?
I was given a sample of either copper or silver iodide, but I don't know which. The chemical is in a water solution, and since it is insoluble in
water, it just sits there on the bottom. The color of the chemical is a a light brownish yellow (tan) color, which can fit both copper iodide and
silver iodide.
I want to get the silver. How would I go about identifying it and turning it into something like silver chloride?
I though I add hydrochloric acid and if it's silver iodide than it should produce a white ppt along with hydroiodic acid? right? And if it's copper
iodide then it should produce a blue/green solution of copper chloride? right? Am i going in the right direction? thanks.
[Edited on 8-11-2014 by jamit]
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DraconicAcid
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Try dissolving it in conc. ammonia. Silver iodide will not dissolve; copper(I) iodide will, and the solution will turn blue.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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DrMario
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Quote: Originally posted by DraconicAcid  | | Try dissolving it in conc. ammonia. Silver iodide will not dissolve; copper(I) iodide will, and the solution will turn blue. |
Wow, cool! I'll make some copper iodide just to try this.
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jamit
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| Quote: Originally posted by DraconicAcid | | Try dissolving it in conc. ammonia. Silver iodide will not dissolve; copper(I) iodide will, and the solution will turn blue. |
I'll try it and report back. Thanks.
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gdflp
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I believe that silver iodide will dissolve. Silver(I) forms an amine complex as well as copper, and the CRC lists silver iodide as being soluble in
NH4OH.
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DraconicAcid
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| Quote: Originally posted by gdflp | | I believe that silver iodide will dissolve. Silver(I) forms an amine complex as well as copper, and the CRC lists silver iodide as being soluble in
NH4OH. |
Silver does form an amine complex, but the iodide is really insoluble. I was under the impression that silver chloride dissolves easily in dilute
ammonia, silver bromide requires concentrated ammonia, and silver iodide stays put. If it is soluble in conc. ammonia, it's not very.
Anyway, the silver complex is colourless and stays that way. The copper(I) complex is colourless, but rapidly turns dark blue on exposure to air.
[Edited on 9-11-2014 by DraconicAcid]
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Metacelsus
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http://alevelchem.com/aqa_a_level_chemistry/unit3.2/sub3205/...
| Quote: | | Silver iodide does not dissolve in even concentrated ammonia. |
Edit: Didn't notice above post (facepalm)
[Edited on 10-11-2014 by Cheddite Cheese]
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gdflp
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I apologize then. Maybe I need to find a newer CRC.
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The Volatile Chemist
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| Quote: | | Wow, cool! I'll make some copper iodide just to try this. |
It works with any Copper(II) ion salt.
Edit: most salts, eg. CuCl2
[Edited on 11-9-2014 by The Volatile Chemist]
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DraconicAcid
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The blue is from the ammonia complex with copper(II)- as the volatile one says, it will work with any copper(II) salt. Copper(I) iodide isn't worth
making specifically for it.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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DrMario
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| Quote: Originally posted by The Volatile Chemist | | Quote: | | Wow, cool! I'll make some copper iodide just to try this. |
It works with any Copper(II) ion salt.
Edit: most salts, eg. CuCl2
[Edited on 11-9-2014 by The Volatile Chemist] |
I've tried this with most Cu(II) salts as a youngster, but never with copper iodide.
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The Volatile Chemist
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| Quote: Originally posted by DrMario | | Quote: Originally posted by The Volatile Chemist | | Quote: | | Wow, cool! I'll make some copper iodide just to try this. |
It works with any Copper(II) ion salt.
Edit: most salts, eg. CuCl2
[Edited on 11-9-2014 by The Volatile Chemist] |
I've tried this with most Cu(II) salts as a youngster, but never with copper iodide. |
I see. I suppose it would be a bit more dramatic with the CuI2.
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DrMario
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Totally worth it!
A few comments: I synthesized the CuI2 via a simple substitution reaction between CuSO4 and KI - but to my surprise, elemental iodine was also
forming. The iodine formation was slow but after a quarter of an hour it made the liquid very clearly brown-yellow and the iodine smell (one of my
favorites) was unmistakable. What up with that?
The other thing is the colour of the precipitate: it was light grey, not "light brownish yellow" as jamit describes. I assume his sample is still
contaminated with some iodine. I have, on my part, carefully decanted my product multiple times.
The change from light gray precipitate to splendid royal blue with the addition of ammonia, was stark and, as I said, totally worth the effort.
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gdflp
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The elemental iodine is due to the fact that you didn't synthesize CuI2, you synthesized CuI. CuI2 is unstable and rapidly decomposes 2CuI2 -->
2CuI + I2
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DrMario
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| Quote: Originally posted by gdflp | | The elemental iodine is due to the fact that you didn't synthesize CuI2, you synthesized CuI. CuI2 is unstable and rapidly decomposes 2CuI2 -->
2CuI + I2 |
jamit, you hear that? So now we at least know which compound you do NOT have.
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DraconicAcid
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I think it was a given that it wasn't CuI2; it was CuI that was considered a possibility.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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unionised
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Is it just my imagination, or is this a rather long thread to see if someone has silver iodide (which is yellow) or copper iodide (which is white)?
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DrMario
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It wasn't necessarily a given, since CuI2 is indeed yellow. 
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DrMario
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| Quote: Originally posted by unionised | | Is it just my imagination, or is this a rather long thread to see if someone has silver iodide (which is yellow) or copper iodide (which is white)?
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It's not your imagination, but you have to consider that some of us (me, for instance) didn't know that CuI2 is so unstable. But I suspect
there may be other thread participants that were blissfully unaware of this.
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DJF90
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| Quote: Originally posted by gdflp | | The elemental iodine is due to the fact that you didn't synthesize CuI2, you synthesized CuI. CuI2 is unstable and rapidly decomposes 2CuI2 -->
2CuI + I2 |
This technically isn't true. CuI2 cannot exist because Cu2+ oxidises I- to iodine (look at the SRPs), and so the two ionic species cannot co-exist (as
both are kinetically fast reactants). It is a redox reaction rather than a decomposition.
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unionised
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How can it be yellow when it doesn't exist?
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DJF90
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Another indication is that silver iodide photolyses in light...
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DrMario
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| Quote: Originally posted by DJF90 | | Quote: Originally posted by gdflp | | The elemental iodine is due to the fact that you didn't synthesize CuI2, you synthesized CuI. CuI2 is unstable and rapidly decomposes 2CuI2 -->
2CuI + I2 |
This technically isn't true. CuI2 cannot exist because Cu2+ oxidises I- to iodine (look at the SRPs), and so the two ionic species cannot co-exist (as
both are kinetically fast reactants). It is a redox reaction rather than a decomposition. |
It clearly exists for a while - the reaction where 2CuI2 --> 2CuI + 2I is relatively slow. You should try it, it's educative.
Now the fact that the OP does not have CuI2 because of its instability, is another question. This instability is new information that we learned in
this thread (or from some Wikipedia page).
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DrMario
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It clearly exists for a while - the reaction where 2CuI2 --> 2CuI + 2I is relatively slow. You should try it, it's educative.
Now the fact that the OP does not have CuI2 because of its instability, is another question. Etc.
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unionised
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I see no evidence of any CuI2
The colour of I2 forms immediately
http://www.youtube.com/watch?v=6SQus2Ikqqs
about 03:35
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