Sciencemadness Discussion Board

Copper solution mystery

electrokinetic - 6-6-2012 at 11:08

I'm adding sodium carbonate, in excess, to a solution of copper chloride and no precipitate is forming. I know there's copper still dissolved because the solution is still blue. Any ideas what's going on here?

MrTechGuy1995 - 6-6-2012 at 12:00

Must be the presence of an acid, most likely HCl. Under acidic conditions, CuCl2 is a blue solution, under basic, it's green.

The acid must be reacting with the Sodium Carbonate, Is it bubbling at all?

plante1999 - 6-6-2012 at 12:54

basic solution will make cuprate salt with copper ion.

woelen - 6-6-2012 at 22:30

With carbonate ion you won't get a cuprate salt in solution, but a precipitate of mixed copper hydroxide/carbonate. If your salt is CuCl2, then the precipitate also will contain chloride ions and you get a green basic copper chloride/carbonate.

Cuprate ion only is formed at very high pH over 14. This only is achieved in fairly concentrated solutions of NaOH or KOH. These cuprate solutions have a nice deep blue color.

Was the CuCl2 home-made? A very common impurity is HCl in such homemade material and then you can add quite some carbonate to such a solution, before precipitation occurs.

[Edited on 7-6-12 by woelen]

nezza - 7-6-2012 at 10:24

Does the solution contain any ions which complex with copper such as citrate. In the presence of sufficient citrate no precipitate will form with carbonate ion.

blogfast25 - 7-6-2012 at 12:17

Quote: Originally posted by electrokinetic  
I'm adding sodium carbonate, in excess, to a solution of copper chloride and no precipitate is forming. I know there's copper still dissolved because the solution is still blue. Any ideas what's going on here?


Keep goin' till it stops fizzin' And then some more...

electrokinetic - 7-6-2012 at 12:55

Quote: Originally posted by woelen  
With carbonate ion you won't get a cuprate salt in solution, but a precipitate of mixed copper hydroxide/carbonate. If your salt is CuCl2, then the precipitate also will contain chloride ions and you get a green basic copper chloride/carbonate.

Cuprate ion only is formed at very high pH over 14. This only is achieved in fairly concentrated solutions of NaOH or KOH. These cuprate solutions have a nice deep blue color.

Was the CuCl2 home-made? A very common impurity is HCl in such homemade material and then you can add quite some carbonate to such a solution, before precipitation occurs.

[Edited on 7-6-12 by woelen]


Okay, I see I've come to the right place. Let me be more specific. The copper chloride originally came from coins dipped in HCl (this is an experiment going on with some interested middle school students in my classroom). The students' goal was to copperplate some old "silver" dollars by electrolysis. In order to prepare the coins they were put into a dilute HCl solution. Some were left in too long and we got a nice green solution.

I didn't want to electrolyze that solution since I didn't want to produce any chlorine gas at the graphite anode. So the plan was to precipitate out the copper as basic copper carbonate, and then dissolve it using sulfuric acid to get a solution of copper sulfate. So I added sodium carbonate and originally got a lot of precipitate that was a dull green color. (I wondered why it was green and not blue and you answered that in your post--I'll come back to that later). But when the precipitate settled, the solution was still blue. So I added more sodium carbonate. And then more. And then more.

I finally filtered it out into a separate beaker and tried adding more sodium carbonate. The pH is currently at 9.9 and still no precipitate is forming. Could something be complexing the copper ions like nezza said? But what could it be? I know there's no citrate in there.

And if I could go back to the green precipitate. So if I understood correctly, at neutral pH rather than basic copper carbonate you get copper chloride carbonate if chloride ions are present? So what's the best way to separate the chloride from the copper?

Thank you!

electrokinetic - 7-6-2012 at 13:10

I forgot to add that when I was adding the sodium carbonate there wasn't any fizzing at all.

And as an update, I just added a lot of sulfuric acid, got a lot fizzing and the solution turned light green. I'm guessing I did have some kind of copper complex, but what was the ligand? OH? Maybe a complex like [Cu(OH)4]2-?

And I still need to know how to get rid of the chlorine. That was the problem in the first place!

Nathaniel - 8-6-2012 at 14:48

Uf, I just wrote a long post and then realized I misunderstood you somewhere in the middle :(

So let me just say that I've made the basic copper carbonate quite some times before and I could never bring the solution to colourless no matter how much carbonate I added...

Now I know that that's because the percipitate is formed in the first place, because the solubility product is reached. If you don't know what that is:
In solution where there's a salt with very low solubility (in this case basic copper carbonate-I'll use CuCO3 to make it more simple) the product of concentrations in solution: {Cu2+} x {CO32-} is constant (some very low number...usually 10-10 or lower)
That means that the copper carbonate percipitated because you added carbonate solution--> {CO32-} rises so {Cu2+} must fall to keep it constant (it leaves the solution as CuCO3)....But no matter how much carbonate you add, the {Cu2+} will never be zero. Never. It appears that in this case the constant is not that low and a bit more of the copper ions keeps remaining in solution but it still isn't much, believe me...It's just because the ions are strongly coloured that it looks like it :)



electrokinetic - 11-6-2012 at 07:58

Well, the solubility product of CuCO3 is 1.4x10-10, but that's for pure CuCO3. We have Cu2(OH)2CO3 and also it seems some other salt contaminated by chloride. Who knows what their SP's are; I couldn't find them.

After adding the sulfuric acid, the solution is clear, but you're right there are still copper ions floating around in there.

But what's more concerning to me now is how to get rid of the chloride, especially in the precipitate that originally formed. I suppose I could add silver nitrate and have AgCl precipitate out, but that is too expensive. I'm hoping someone could present a cheaper alternative.

phlogiston - 11-6-2012 at 12:35

1. Redisolve in minimal amount of acid (sulphuric/hydrochloric/nitric).
2. Add sodium hydroxide solution to precipitate copper hydroxide as a blue gel
3. Bring to a boil. The copper hydroxide will quickly convert to black CuO which settles readily and is easily filtered.
4. Wash the CuO precipitate repeatedly with water.
5. allow to settle, decant and redisolve in sulphuric acid

electrokinetic - 12-6-2012 at 15:31

Okay, did that, but the solution turned black and I got a reddish precipitate. No blue gel.

TheAMchemistry87 - 12-6-2012 at 18:24

Check the pH if its low its acidic and you still have HCl what you should do is add more copper (II) oxide until no more is dissolve then filter it and add the sodium carbonate and you should get a precipitate of copper (II) carbonate if you dont your solution is VERY impure

woelen - 12-6-2012 at 23:09

If I look at all the observations, mentioned by electrokinetic, then I get the impression that there is some strong contamination of the solutions. Such a contamination can severely affect the outcome of all kinds of experiments. The copper, used to make the copper chloride, or the acid, used for making it probably contained all kinds of impurities (such as iron or salts of iron, or organics).

phlogiston - 13-6-2012 at 15:17

I'm beginning to doubt whether it was a copper salt. I've made CuO this way numerous times, usually starting with copper nitrate made by dissolving metal (from electrical wiring) in acid. What coins did you start with?

Here's (someone elses) movie of what the gel-like copper hydroxide precipitate looks like:

http://www.flickr.com/photos/cmlburnett/4471334146/in/photos...

To confirm it's copper at all, you could do a quick flame test perhaps. Dip the tip of an wire (ideally platinum, but iron will do) in HCl(aq), and pick up a bit of your copper salt with the moist wire. Hold in a colorless/blue flame and you should get shades of blue and green.

AndersHoveland - 15-6-2012 at 02:45

Quote: Originally posted by woelen  
With carbonate ion you won't get a cuprate salt in solution,

I agree. Although Cu(OH)2 is very slightly amphoteric, sodium carbonate is not quite basic enough to form Na2CuO2, sodium hydroxide would be needed to form this.

[Edited on 15-6-2012 by AndersHoveland]

electrokinetic - 18-6-2012 at 13:30

Okay, as for the reddish stuff formed on the first go:
1. I noticed it turning black as it dried on the filter paper.
2. Flame test burned greenish.

Could this red stuff then have been copper (I) oxide? The copper chloride I started with was made when Eisenhower dollars, of the same copper clad make as quarters and dimes, was left overnight in HCl.

I mixed a separate batch of copper chloride with concentrated NaOH, and voila got exactly the results phlogiston described. Thanks phlogiston! I can't remember how that particular batch was made. I've washed it a few times, what's the best way to know when it's been washed enough?

I'm expecting great amounts of contamination, btw. I'm a middle school teacher and we do a lot of stuff with copper chloride and copper sulfate, and I got tired of just dunking Al foil into the used copper to "neutralize" it before pouring down the drain. So isolating copper sulfate and electroplating it out is more or less a game I made for myself. Thanks to everyone here for their help and suggestions.

Eddygp - 21-6-2012 at 13:19

Umm, maybe some sort of copper(II) chloride. Possibly... just thinking.