Sciencemadness Discussion Board - Process to dissolve Cu but not Zn

Process to dissolve Cu but not Zn

pantone159 - 6-7-2007 at 18:48

What chemistry could be used to dissolve copper but not zinc?

I wonder about this, because I saw a column by Theodore Gray about the 'dissolve the inner Zn part of a US cent leaving Cu foil' trick...
http://www.popsci.com/popsci/how20/0091804df3c83110vgnvcm100...

The normal part isn't that interesting, just HCl, but he vaguely described a process to dissolve the outer copper layer while leaving the zinc intact. All he said was it involved cyanide and persulfate, but no further details were given. Anyone have any ideas on what specifics would work?

not_important - 6-7-2007 at 19:17

Cyanide complexes copper, among other metals. Treating copper, silver, and so on, with a solution containing cyanide ions and hydrogen peroxide as an oxidiser will dissolve the metal. The main complex is M(1+)2[Cu(CN)4], M being Na, K, ...

However zinc also forms such complexes. I'm assuming that what was done was a carefully regulated dissolution of the copper and a bit of the zinc, stopping when all of the copper had been dissolved. By using the cyanide/persulfate agent a fairly uniform attack on the the copper occurs, and the attack on zinc is not too vigorous (I'm guessing) Also, by complexing the copper the rate of its plating out onto the exposed zinc is reduced a good deal. If HNO3, or HCl/H2O2, had been used the copper originally dissolved would have attacked any exposed zinc, Cu(2+) + Zn => Cu + Zn(2+), resulting in all the zinc dissolving while the copper keep being reduced to metal.

[Edited on 7-7-2007 by not_important]

JohnWW - 8-7-2007 at 21:34

The problem is that Zn is significantly above Cu in the electrochemical series, with Zn being above H and Cu below H in the series, meaning that non-oxidizing acids dissolve Zn but not Cu. This makes it extremely difficult to find a way to dissolve Zn but not Cu. It would have to be some sort of complexing ligand that preferentially bonds very strongly to Cu but not Zn.

[Edited on 10-7-07 by JohnWW]

not_important - 8-7-2007 at 22:56

Doesn't matter if the zinc is also complexed, just so long as the copper stays complexed. And copper does form complexes such as K2[Cu(CN)4].

evil_lurker - 9-7-2007 at 00:36

Make a nick in the coin, and nitric acid will do the rest... the nitric will attack the zinc much faster than the copper, leaving a shell.

Nerro - 9-7-2007 at 00:54

I vaguely remember about Zn being susceptible to attack from strong alkali's. It forms zincate I believe. Does copper also do this? If it doesn't, give it a go, boil a nicked coin in concentrated NaOH soln.

YT2095 - 9-7-2007 at 01:04

Dilute sulphuric will do the same and the product is much more soluble.

Nerro - 9-7-2007 at 05:35

But then you have the problem of oxygen dissolving in the solution which will dissolve the Cu surprisingly rapidly.

YT2095 - 9-7-2007 at 05:53

where does the Oxygen come from?
the Zn is in the +2 oxidation state and simply replaces the H2, in H2SO4.

Zn(s) + H2SO4(aq) → Zn2+(aq) + SO42-(aq) + H2(g)

Nerro - 9-7-2007 at 06:11

Air.

YT2095 - 9-7-2007 at 06:21

Interesting, although I`ve never had any luck at all trying to dissolve copper in sulphuric acid, the addition of H2O2 may do the job though, and I guess Air given enough time would a little impact, but nothing of any significance I wouldn`t imagine.

even HCl won`t attach copper properly in any appreciable way without the addition of a Chlorate of H2O2, after that the reaction IS self sustaining from ariel sources.

HNO3 isn`t a great idea either, although the Zinc will afford a little cathodic protection, as soon as it`s gone the copper will be too in minutes, I dump all my scrap copper in a HNO3 bucket, it`s great for dissolving the stuff.

Nerro - 9-7-2007 at 06:32

I've dissolved the copper layer off of a coin in half an hour once by dumping the coin in dilute sulphuric acid. The effect of oxygen in the air really is significant.

YT2095 - 9-7-2007 at 06:43

that`s quite amazing, I`ve never been able to dissolve copper in H2SO4 directly, even the Oxide (CuO) needs heating in order to dissolve at an appreciable rate.

did you use a Bubbler or something?

pantone159 - 9-7-2007 at 07:01

Quote:

Originally posted by evil_lurker
Make a nick in the coin, and nitric acid will do the rest... the nitric will attack the zinc much faster than the copper, leaving a shell.

The idea is to dissolve the copper but NOT the zinc.

The other way around, dissolving the inner zinc to make a shell, is no problem. Still kind of fun, but no real difficulty.

Nerro - 9-7-2007 at 07:16

Quote:

Originally posted by YT2095
that`s quite amazing, I`ve never been able to dissolve copper in H2SO4 directly, even the Oxide (CuO) needs heating in order to dissolve at an appreciable rate.

did you use a Bubbler or something?

No just the occasional stir.

As to the real topic, I still don't see how persulphate would be of any use in this case. The persulphate would oxidize the copper to 2+ and the Zn to +2 (or even +3 I believe). The cyanide makes more sense, there is a stable Cu(CN)42- ion but it seems quite likely that something similar like Zn(CN)42- could also exist...

[Edited on 9-7-2007 by Nerro]

YT2095 - 9-7-2007 at 07:23

I can`t think of a way either Zn is far more reactive, the only other exploit I can think of at a push, would be to cut the coin in half and heat it, copper melts at a higher temp than Zinc does and might be possible to melt it out.

Nerro I can`t think why it works for You but not for me? without passing a current through the copper, it remains fairly inactive in H2SO4 here????

Nerro - 9-7-2007 at 13:08

I still have the pictures on my photobucket.

These are the coins in sulphuric acid after about four hours of standing outside, it got occasionally swirled.

These are the same kinds of coins (a €0.05 coin and a €0.02 coin which are both composed of Cu plated Fe) in 10% HCl after 4 hours with the same treatment.

As you can see in the second picture the Cu plating had by then been penetrated and the Fe was starting to dissolve which noticably retarded the dissolution of the Cu. In the first picture the Cu layer that was left was only microns in thickness. I could scratch through it with a nail quite easily after I recovered the coins. The condensation on the erlenmeyer was because it was a fairly hot day (~25°C probably, it was a hot summer).

[Edited on 9-7-2007 by Nerro]

Polverone - 9-7-2007 at 19:50

To remove the copper and leave zinc intact, add the pennies to a jar containing sodium polysulfide and give them a periodic good shaking. You may need to scrub with a toothbrush or finger to remove the last bits. The copper is converted to poorly-adherent sulfide, while the zinc (or zinc rich alloy? it's shiny silver, either way) seems mostly unaffected.

Nerro - 9-7-2007 at 22:41

Care to explain why the Zn wouldn't react with the polysulfide?

12AX7 - 10-7-2007 at 14:35

Does Zn form a tightly-adhering sulfide?

not_important - 12-7-2007 at 23:59

While Polverone's suggestion sounds safer, the cyanide method seems to be well grounded. From the 2nd edition of Cotton & Wilkinson Advanced Inorganic Chemistry

Quote:

Cuprous Cyanide ... The cyanide is soluble in solutions of complexing ions; cyanide gives [Cu(CN)4](3-) as the main species. The latter has such a large formation constant the copper metal will dissolve in potassium cyanide solution with evolution of hydrogen

The complex cyanides of zinc do not seem to be strongly favoured. It would seem that the strength of the cuprous complex is so great that copper does not plat out on the zinc, and that it may be working as sacrificial protection for the zinc.

Mr. Wizard - 13-7-2007 at 07:24

There was a discussion on this already at:
http://www.sciencemadness.org/talk/viewthread.php?tid=2172&a...

There might be some more ideas on that link.

Polverone - 13-7-2007 at 08:14

Thanks, Mr. Wizard. I couldn't remember where I'd read of the polysulfide method -- apparently it was from you.

Waffles - 14-7-2007 at 15:38

When Theo first did this, he proudly emailed me with the process and pictures:

Quote:

Dissolved the inside with HCl, and the outside with a combination of potassium cyanide and potassium persulfate. The cyanide strongly complexes copper, reducing its redox potential to just below that of zinc, and the persulfate oxidizes it. My friend Tryggvi worked this out, and it's quite delicate. You have to keep the solution cold otherwise the persulfate starts attacking the zinc too.

not_important - 14-7-2007 at 20:17

And possibly the reason the polysulfide method works:

Ksp CuS 8 x 10^-37

Ksp ZnS (alpha) 2 x 10^-25
Ksp ZnS (beta) 3 x 10^-23

and suggesting an experiment with high copper Cu-Ni alloys
Ksp NiS (alpha) 4 x 10^-20
Ksp NiS (beta) 1,3 x 10^-25

12AX7 - 14-7-2007 at 20:29

It's damn insoluble, but that doesn't suggest why the zinc isn't likewise reacted. It's most likely a rusting vs. passivation effect.

And yeah, it looks like sulfides are a stinky but possible way to seperate nickel. 10^-20 is pretty damn insoluble despite the difference, though. Probably a lot of coprecipitation in the nonequilibrium solutions during mixing when all the crap suddenly precipitates...

Alternately, you could process it at the rate of a few coins worth per swimming pool. Then the concentration would be low enough to see some good seperation.

Tim