Sciencemadness Discussion Board - Complex ions of copper

Complex ions of copper

The_Davster - 22-3-2004 at 18:06

Recently, I have been doing some work with some of the complex ions of copper.
I have prepared Cu(Cl)4 2+ tetrachlorocopper(II), Cu(NH3

42+ tetraaminocopper(II) and Cu(H20)6 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]

Blind Angel - 22-3-2004 at 18:23

for the subscript use /

Geomancer - 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.

darkflame89 - 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.

Nick F - 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!)

chemoleo - 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?

Nick F - 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!

chloric1 - 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!!

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

[Edited on 3/23/2004 by chloric1]

The_Davster - 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

Anyway, is it right to assume that tetraaminocopper(II) carbide would be more unstable and powerfull than copper(II) carbide?

darkflame89 - 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.

Nick F - 24-3-2004 at 04:28

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

unionised - 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)

Good point

chloric1 - 24-3-2004 at 20:01

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]

Pyrovus - 24-3-2004 at 22:13

How about Cu(NI3)4 ++, Cu(NCl3)4 ++ and the like? They should be very interesting . . .

unionised - 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).

[Cu(NH3)4](NO3)2

hodges - 25-3-2004 at 15:50

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.

Geomancer - 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.

The_Davster - 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.

Nick F - 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...).

Results

hodges - 26-3-2004 at 16:37

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.

The_Davster - 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]

hodges - 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.

The_Davster - 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.

Theoretic - 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...

Esplosivo - 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]

unionised - 29-3-2004 at 09:26

Practically speaking, ammonium hydroxide doesn't exist.

If I were going to try making this stuff I would try the same method that is generally used for tetramine copper sulphate ie use a concentrated solution of copper sulphate, add conc ammonia, then add alcohol to ppt the product.

The_Davster - 29-3-2004 at 11:21

Unionised, the problem for me with your method is that the only ammonia that I have acess to(other than generating the gas) is the dilute solution avaliable at the grocery store. Is there a way to concentrate it?

hodges - 29-3-2004 at 15:31

I tried this again, and had an even less favorable result. Because of the humid weather I decided to dry it in an oven at around 80C. The resulting product looked darker blue than copper nitrate, but it would not do a darn thing. I even heated a small amount of it with a blowtorch and all it did was sizzle a bit (may have been just water escaping). Looks like it had decomposed in the 80C heat of the oven.

The_Davster - 29-3-2004 at 16:23

I'm letting a sample of copper nitrate air dry. I have found that it seems to be slightly soluble in xylene, however the copper nitrate was not dry at the time of adding to the xylene so iy could have just been the water in the copper nitrate allowing it to dissolve. Will bubble the ammonia through tomorow.

[Cu(NH3)4]SO4 Results

hodges - 2-4-2004 at 16:09

I made some tetraamminecopper sulfate. I started with 12.5 grams (0.05 moles) of copper sulfate pentahydrate. I dissolved this in water. I then added an excess of ammonium hydroxide. This produced the characteristic deep blue color.

It took several weeks for all the water to evaporate from the solution at room temperature. I determined when it was dry by monitoring the weight until it was no longer falling. The result was a blue salt. It is not as dark as the solution, but considerably darker than copper sulfate. It is much more easily powdered than hydrated copper sulfate, but even when finely powdered it is hard to re-dissolve it in water. Heating caused a definite evolution of NH3 gas.

The resulting product weighted 10.9 grams. Doing the math, there are 3.22 molecules of NH3 for every Cu molecule. Theoretical is 4. Of course, it is also possible that there is less NH3 and more H2O, since the two have very similar masses.

t_Pyro - 2-4-2004 at 20:51

Tetrammine copper compounds have got me confused now. Till now, I always used to regard it as [Cu(NH3

4]++. Then, I came across a paragraph in my inorganic chem book which referred to it as [Cu(NH3

4(H2O)2]++, which gives it a coordination number of 6. So what is the coordination number of Cu? 4 or 6?

[Edited on 3-4-2004 by t_Pyro]

hodges - 5-4-2004 at 14:51

I'm not having too good of luck drying my tetraamminecopper nitrate. The dark blue color fades from the solution after a couple days, leaving a lighter blue color. I tried adding NH4OH again and the same thing happened after a couple of days. This problem did not occur when I was making tetraammine copper sulfate. The crystals were lighter than the original solution, but the solution itself remained dark until it all evaporated. Why should it matter whether the other ion is nitrate or sulfate as far as the rate of decomponsition?

I'm coming close to concluding that making tetraammine copper nitrate is impossible, at least in aqueus solution. I thought more about the previous observation that tissue paper soaked in this solution burns leaving behind copper. Probably what is happening is that copper nitrate is forming copper oxide as the paper burns. Then the remaining carbon from the paper reduces the copper oxide to copper. There probably isn't even any TACN at all (I assumed there was due to Cu being formed when it burned).

TACN success

The_Davster - 5-4-2004 at 15:09

Well as I did before, I made a concentrated solution of copper nitrate in a testube(15ml). Ammonia gas was bubbled through, first I got a tiny bit of dark blue precipitate, then it all slowly turned light blue due to a precipitate. I continued bubling the ammonia through for about 10min. At this time there was a mix of about 50/50 light blue and dark blue precipitates. some of this was filtered and the dark blue precipitate dissapeared for some reason. the rest of the solution I dident want to bother with filtering so I dumped it onto an icecrean bucket lid, placed a filter paper on top and absorbed all the water with paper towels. there was a mix of light and dark blue precipitate on the filter paper. this was allowed to dry and today I lit the filter paper....success :cool:

I got little green puffs of fire and some crackling as the crystalls burned.

chemoleo - 5-4-2004 at 15:28

Hmm, I don't understand the problem of it so much.

The light blue precipitate is Cu(OH)2 or copper hydroxide, which forms by adding the base ammonia. Continuous addition of NH3 eventually dissolves this again, to form the complex. Ammonia solution is better in this case as it has all the ammonia dissolved already, unlike the NH3 gas where it has to be absorbed first.

Why dont you try the following (similar to http://www.sciencemadness.org/talk/viewthread.php?tid=1778 ):
Make a saturated solution of Cu(NO3)2, and then add ethanol until copper nitrate starts to precipitate. I did this recently (by accident, with CuSO4), a 50% addition of EtOH or so should be fine. Just check it how much EtOH this solution can hold.
Then take the alcoholic solution of copper nitrate, and add a mixture of ammonia in water/ethanol (the proportions of the mixture determined by the result of the EtOH/Cu(NO3)2 mix, i.e. 50%NH3 solution/50% EtOH), and add this until you reach an excess of full saturation (calculate it).
At this point, you either get the deep blue precipitate of the nitrate complex, or everything remains in solution.
In case of precipitation, everything is good, you just purified the complex. In case of no precipitation, just add MORE ethanol until the complex finally precipitates.
Filter it, dry it, and test!!

Honestly, I don't see at all why this shouldnt work! Good luck!

[Edited on 7-4-2004 by chemoleo]

Theoretic - 14-6-2004 at 07:01

I have myself prepared Cu(NH3)4SO4.
I have added copper sulfate to a 25%solution of ammonia, a deep blue colour appeared above the crystals, and the reaction proceeded quite quickly. For some reason, it is now precipitating Cu(OH)2.
Hold on... shouldn't free copper ions destroy complexed copper ions:
Cu++ + Cu(NH3)4++ + 6H2O => 2Cu(OH)2 + 4NH4+
Otherwise there wouldn't be any precipitate when ammonia is added to copper salts. In my case there wasn't any, why are the textbooks stating it? Or does the precipitate redissolve if it forms?
I have dissolved copper in a (NH4)2SO4 solution, with aid of atmospheric oxygen. Right now the copper is still dissoving, and the solution is distinctly blue.
Dissolving copper in a NH4Cl solution should be fun, also an easy way to CuCl2 without HCl.., like so:
2Cu + O2 + 4NH4Cl => Cu(NH3)4CuCl4 + 2H2O.
Then heat to get rid of ammonia and you're done!

Edit: spelling, what else...

[Edited on 14-6-2004 by Theoretic]

The_Davster - 19-6-2004 at 12:16

I have been allowing some copper wire/strip/filings in an ammonium nitrate solution in the hopes of obtaining TACN. There were reports of success with this method on the E&W I think. I am in the process of testing this out and so far the solution has turned a deep blue. I think the reaction hapenning is 2Cu+4NH4NO3-->Cu(NH4)4(NO3)2 + Cu(NO3)2. Obviously this reaction produces two water soluble which would make extraction of pure TACN quite difficult. So I got to thinking on how to obtain a purer product of TACN and came up with this 2Cu+2NH4NO3+2NH4OH-->Cu(NH4)4(NO3)2+Cu(OH)2. Here there is the advantage of an insoluble product obtained which leaves just TACN in solution after filtering. However I am unsure if the second reaction will work. Also it would be advantageous I think to use a slight excess of copper to make sure that all the ammonium nitrate would be reacted to keep it from contaminating the final product.
Thoughts?

hodges - 19-6-2004 at 14:30

If you add excess NH4OH it will convert any Cu(NO3)2 to [Cu(NH4)3](NO3)2. So yes that should work. I never had much luck with TACN though, even starting with pure Cu(NO3)2. At best it would deflagrate, did have a pretty green flame though. Unfortunately it is also a powerful explosive if detonated, which to me makes it the worst of both worlds. It is hard to get to work and gives you a false sense of security becuase usually it does not do much. This might make it tempting to use larger amounts, which might actually detonate, and remember it is very powerful if detonated.

The_Davster - 19-6-2004 at 15:38

This summer I plan do some in depth work with the various methods of production of TACN and test the products relative to each other. Does anyone have the VoD of TACN?

Magius - 23-8-2004 at 14:20

Hey guys, I to have prepared some tetraaminocopper(II), I think. I set up an electrolysis experinemt with ammonia as the electrolyte, and copper annodes and cathodes. After running it for a few hours, I had a dark blue solution that appears to be tetraaminocopper(II). But I'm not sure what the anion of the solution is, anyone have any suggestions? No percipitate has formed, so OH is out of the question.

Geomancer - 23-8-2004 at 16:58

Isn't tetraamine copper (II) strong enough to be soluble even with hydroxide counter-ion?

Nitro-esteban - 3-5-2014 at 21:16

I made tetraamminecopper (II) perchlorate by adding a saturated solution of copper perchlorate to an excess of ammonia (30% concentrated). Adding the ammonia to the copper perchlorate only presipitates copper hydroxide.

DraconicAcid - 3-5-2014 at 21:35

Quote: Originally posted by Nitro-esteban

I made tetraamminecopper (II) perchlorate by adding a saturated solution of copper perchlorate to an excess of ammonia (30% concentrated). Adding the ammonia to the copper perchlorate only presipitates copper hydroxide.

You needed to add more ammonia. Copper(II) hydroxide should be plenty soluble in enough xs ammonia. Either that, or add some ammonium perchlorate as a buffer.

[Edited on 4-5-2014 by DraconicAcid]

nezza - 3-5-2014 at 23:45

The ethylenediamine perchlorate complex is an interesting dark purple compound which explodes erratically when heated giving a beautiful blue flame.

Nitro-esteban - 4-5-2014 at 13:22

Quote: Originally posted by DraconicAcid

Quote: Originally posted by Nitro-esteban

You needed to add more ammonia. Copper(II) hydroxide should be plenty soluble in enough xs ammonia. Either that, or add some ammonium perchlorate as a buffer.

[Edited on 4-5-2014 by DraconicAcid]

Copper (II) hydroxide forms when the ammonia is added to the copper perchlorate but if the reactants are mixed the other way around (copper perchlorate is added to the ammonia) an exothermic reaction occurs and tetraamminecopper (II) perchlorate precipitates as the temperature drops. In both attempts I used the same amount of ammonia and copper perchlorate, I just mixed them in different ways.
Here is a link to a thread about TACN: http://www.sciencemadness.org/talk/viewthread.php?tid=16220#...

[Edited on 4-5-2014 by Nitro-esteban]

The Volatile Chemist - 26-5-2014 at 14:05

Quote: Originally posted by chloric1

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]

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

[Edited on 3/23/2004 by chloric1]

When you make the bromide complex, this person said "acidic CuSO₄. Did they mean in-acid, or were they considering the copper sulfate 'acidic'?
Also, someone mentioned an acetate complex. What? I hadn't heard of this. How would one go about making it (w/o just making CuAc?
Thanks,
Nathan