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The_Davster
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[*] posted on 22-3-2004 at 18:06
Complex ions of copper


Recently, I have been doing some work with some of the complex ions of copper.
I have prepared Cu(Cl)<sub>4</sub> 2+ tetrachlorocopper(II), Cu(NH<sub>3</sub>;)<sub>4</sub>2+ tetraaminocopper(II) and Cu(H20)<sub>6</sub> 2+.
There seems to be very little info on the net about these types of ions so I was hoping that my questions could be answered here.
1) Is there any other complex ions of copper that I havent mentioned?- preferably cations
2) Is there a way to convert the tetrachlorocopper ions back to copper(II) ions?

Thanks


[Edited on 23-3-2004 by rogue chemist]
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[*] posted on 22-3-2004 at 18:23


for the subscript use < sub > / < /sub >



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[*] posted on 22-3-2004 at 19:10


Copper sould form all types of complexes. Commonly available ligands include acetate, oxalate and glycinate, as well as the other halogens. Not OTC, but accessable to determined people, are ethylenediamine, cyanide, salen, and perhaps EDTA. Many of these would form anions or uncharged complexes, though. Copper will form complexes with DMSO.
I'm away from the library right now, so I don't have much more info. I think cyanide is used to complex copper in electroplating. The glycine complexes are interesting in that they have geometric isomers, the cis isomer being crystalized from 30% ethanol, but converting to the trans isomer when heated at 220C for 15 minutes.
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[*] posted on 23-3-2004 at 02:35


Copper tetracloride(II) ions are interesting, i got some when i put a copper wire into a solution of vinegar and some salt.



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[*] posted on 23-3-2004 at 04:38


"Is there any other complex ions of copper that I havent mentioned?"

Just about anything negatively charged, or with a lone pair, can be a ligand if there's not too much steric hinderance. So yeah, there are quite a few more ;).

Azide springs to mind... (hexaazidocopper (II) perchlorate? Now that should be fun!)




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[*] posted on 23-3-2004 at 05:10


Yepp, and there are the Hydrazin nitrate/chlorate/perchlorate (which are primaries), ethylenediamine complexes too :)
Philous sent me the methods some time back, anyone interested?
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[*] posted on 23-3-2004 at 06:17


Why not, there are bound to be people who would like them.

Didn't anyone notice how silly I was in my last post? Hexaazidocopper will be negative... I must learn to think!




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[*] posted on 23-3-2004 at 11:16


Quote:

2) Is there a way to convert the tetrachlorocopper ions back to copper(II) ions?

Thanks

[Edited on 23-3-2004 by rogue chemist]


Yes there is, add to excess H2O! I have played around alot with CuSO4 and other common copper salts. When in excess chloride, especially acidic, you get bright green tetrachlorocuprate. Now you add water until you get a aquamarine color. If you want to really play, you should take acidic CuSO4 and add large portions of concentrated NaBr solution. Bromo complexes are Purple! My favorite color. If you boil this it is unstable and it will deposit cuprous bromide on dilution!! FUN FUN!!:D

OMG! Post 100!! I am hazardous to others now!! MAybe I have always been! LOL!:D

[Edited on 3/23/2004 by chloric1]




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The_Davster
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[*] posted on 23-3-2004 at 17:11


Wow thanks for all the replies.
Chemoleo: I would be interested in Philous methods.
Hexaazidocopper, sounds cool, anyone up for tetraaminecopper(II) Hexaazidocopper :D
Anyway, is it right to assume that tetraaminocopper(II) carbide would be more unstable and powerfull than copper(II) carbide?
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darkflame89
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[*] posted on 24-3-2004 at 01:12


Quote:

When in excess chloride, especially acidic, you get bright green tetrachlorocuprate


Bright green? I don't remember it was bright green. What i got was a dark green solution.




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Nick F
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[*] posted on 24-3-2004 at 04:28


Bright green, dark green - it will depend on concentration.



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[*] posted on 24-3-2004 at 14:22


I'm a bit suprised that nobody has mentioned the tartrate and citrate complexes (Fehlings and Benedict's reagents)
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[*] posted on 24-3-2004 at 20:01
Good point


Quote:
Originally posted by unionised
I'm a bit suprised that nobody has mentioned the tartrate and citrate complexes (Fehlings and Benedict's reagents)


I am glad you said that! This could be an avenue to explore alkaline non-cyanide copper plating with various additives and possibly pulse plating. Also, I have a Huge book on making patinas for copper, brass,bronze,and silver alloys and tartarate copper solutions are quite common. With various additional reagents for differing effects of coarse. [Edited on 3/25/2004 by chloric1]

[Edited on 3/25/2004 by chloric1]




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[*] posted on 24-3-2004 at 22:13


How about Cu(NI3)4 ++, Cu(NCl3)4 ++ and the like? They should be very interesting . . . :)
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[*] posted on 25-3-2004 at 14:05


Aren't nitrogen trihalides lewis acids? NI3 certainly complexes with ammonia, so don't think it would complex with copper (II).
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[*] posted on 25-3-2004 at 15:50
[Cu(NH3)4](NO3)2


This is also an interesting compound. I have heard it described as a high explosive, decomposing into water and nitrogen. Others say it is not too unstable and burns like black powder. Probably depends on how much water is left. For explosiveness I've heard it is best to prepare it without water (using alcohol and dry NH3 gas). I'm taking a much simpler approach - I'm making it in water solution and with NH4OH. For safety reasons I'm only making about 0.2 grams - will let this dry on a paper towel and then light the paper towel to see if I there is any deflagration.
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[*] posted on 25-3-2004 at 17:05


No personal experience here, but I believe that TACN is a highly brisant explosive, sensitive to shock but not to flame. When ignited, it simply burns. Be carefull.
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[*] posted on 25-3-2004 at 17:10


I am also going to try this reaction this weekend. I had been planning on trying it on the assumption it would be an explosive, your post confirmed my thougts of its explosive properties. I will use ammonia gas as oposed to solution because I have no concentrated NH4OH. Will post results soon.
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[*] posted on 26-3-2004 at 03:16


A lot of people at the E&W Forums worked on this and similar compounds; the general conclusion was "don't bother unless you work anhydrous." Unless you're using energetic ligands (nitroguanidine, 5-aminotetrazole...).



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[*] posted on 26-3-2004 at 16:37
Results


My solution was pretty dilute. The paper towel ended up adsorbing it. It dried completely to a light blue color. When burned, it did not seem to do anything special - in fact, it burned somewhat slower than an ordinary piece of paper towel. However, one interesting effect was that when the carbon finished burning, it left behind a bright red web of very fine copper in the same shape as the paper towel! I'm assuming the reaction is:
[Cu(NH3)4](NO3)2 ---> 6H2O + 3N2 + Cu
I left the rest of the solution to dry on a plastic lid. I got some small dark blue crystals around the edges, but the rest formed a light blue non-crystaline substance. The dark blue crystals appeared to melt when heated with a flame, then they would suddenly vaporize with a faint "poof". I was unable to test the light blue coating because it was very thin and hard and I didn't want to go scraping a potential high explosive. I tried adding water again and this light blue substance wouldn't even dissolve. I think it might be just ammonium carbonate (I used excessive NH4OH, and the solution was exposed to air). Or else perhaps an ammine molecule with fewer than four NH3 groups.

If anyone else tries this, I would be curious as to your results. I may try it again, making all crystals instead of putting on a paper towel, if I get time.
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[*] posted on 26-3-2004 at 19:13


I have now done some work with Cu(NH3)4(NO3)2. First some copper nitrate was made by the addition of excess copper to the 5-10mL of 70%nitric acid. Once NO2 stopped being emitted some of the copper nitrate had crystallized so 5mL of water were added to get it to dissolve. This was placed in a testube and ammonia gas was bubbled through(produced by reaction of ammonium nitrate with sodium hydroxide) using tubing with a one way valve on the end. There was an instant pale blue precipitate. The valve I was using did not work perfect so a little solution went into the valve. After the ammonia had been bubbled through the tube was filled with a pale blue precipitate. This was then filtered. It has not dried completly yet but a little bit of this precipitate when placed in the flame of an alcohol burner only turns the flame green, no puff of smoke no detonation, it was essentialy turned out to be a flame test :( . However when I was washing the equipment a very dark blue solution/precipitate was inside the valve. I carefully cracked open the valve and filtered what was inside. I got about 5-10 very small very dark blue crystalls. These crystals when burned made a puff of green flame and a sound-kinda like when a single crystall of AP is burned but with more of a "poof" cant describe the sound any better unfortunatly.
Conclusions:
As little water as possible should be used, or preferably an nearly anhydrous solvent-any ideas?
Also does anyone have any ideas on what the light blue precipate was?

[Edited on 27-3-2004 by rogue chemist]
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[*] posted on 27-3-2004 at 06:51


Quote:
Originally posted by rogue chemist
Also does anyone have any ideas on what the light blue precipate was?


Probably the result of not all the attached H2O being replaced by NH3. The original hydrated copper ion is Cu(H2O)4. The final desired ion is Cu(NH3)4. There are intermediates such as Cu(H2O)3(NH3), Cu(H2O)2(NH3)2, etc. I think you did not bubble enough NH3 through your solution for the amount of Cu(NO3)2 you had. You may have also had some remaining HNO3 which would also react with the NH3 and water to produce NH4NO3, resulting in a need to use more NH3 than expected.

I have some tetraammine copper sulfate solution which I have been drying for the past 3 weeks. Damp crystals have been forming along the sides of the container as the solution evaporates. Some of these are dark blue as expected, but some of them are lighter blue and in fact have some degree of white color as well. I though maybe the white was ammonium bicarbonate (from CO2 in the air). But apparently not. I let some NH4OH solution evaporate in air and I never saw a white substance form. Also, the NH4OH solution lost its smell in less than 24 hours, whereas my copper solution still has an NH3 smell after 3 weeks. I suspect that the crystals that form are slowly releasing their NH3. Once the rest of the solution evaporates (probably another week), I should be able to tell if this is the case by noting whether there is still an NH3 smell and whether the color lightens with time.

I made my copper nitrate by reacting copper sulfate with calcium nitrate (insoluable calcium sulphate precipitates eventually). I have time to make some more but the humidity is high here today so I doubt I would have much luck drying it. I may try drying a small amount (few tenths of a gram) on a paper plate in an oven on low heat.
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[*] posted on 27-3-2004 at 10:36


Hodges, I think you are right about the intermediates with both the water and ammonia complexes. There was no excess nitric acid, I used a large excess of copper and removed the unreacted copper before bubbling the ammonia through. I will repeat this but with a different solvent and another experiment with using water as the solvent and bubble the ammonia through for longer.
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thumbup.gif posted on 29-3-2004 at 06:15


I think the blue precipitate was Cu(OH)2, as ammonia dissolved, formed ammonium hydroxide and then:
2NH4OH + Cu(NO3)2 => Cu(OH)2 + 2NH4NO3.
Addition of more ammonia would give you the desired result, in your solution you had AN...
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[*] posted on 29-3-2004 at 06:23


What's the use of bubbling ammonia gas through water? The gas is readily soluble forming the hydroxide. Another solvent should be used, which at least contains a smaller percentage of water in it, say commonly available EtOH (which I buy from farmacies at approx 90% conc.). This should help in reducing the double displacement rxn stated below right?

2NH4OH + Cu(NO3)2 => Cu(OH)2 + 2NH4NO3

[Edited on 29-3-2004 by Esplosivo]
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