When I reacted copper chloride and aluminum foil instead of getting the brown color I got green color. Can anyone share their procedure where the experiment actually succeeded?

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    $\begingroup$ This is the first one that popped up (using Google): melscience.com/US-en/articles/battle-metals-experiment . Make sure your copper chloride solution is not real dilute. Alternatively, use copper sulfate and sodium chloride. You will know it works when the reaction really gets going: it gets hot and very obvious! $\endgroup$
    – Ed V
    Commented Feb 19, 2020 at 2:15
  • $\begingroup$ You asked "Can anyone share their procedure where the experiment actually succeeded?" I did exactly that in my answer. Please consider accepting the answer. $\endgroup$
    – Ed V
    Commented Feb 19, 2020 at 21:19

2 Answers 2


This is an easy experiment to do at home, so I did it about an hour ago. The first figure shows the starting chemicals and "equipment":

Starting items

I did not have any copper chloride, so I used copper sulfate pentahydrate, $\ce{ CuSO4.5H2O}$, and $\ce{NaCl}$. (The sodium and sulfate ions are just spectator ions here.)

I added the two chemicals to the $\pu{25 mL}$ of water in the little beaker, stirred until they were fully dissolved, then poured the solution onto the aluminum foil. The second figure shows what happens about every 2 minutes:

Reaction progress shots

The third figure shows what happens around the 10 minute mark: the solution has eaten through the aluminum and the brown copper is evident:

Fig. 3 Near the end

Finally, Fig. 4 shows the result after the solution has leaked away:

Fig. 4 Final result

So, the redox reaction works as expected and as shown in various videos on the web. Note that copper sulfate pentahydrate crystals are blue, as are aqueous solution of it. The copper chloro complex ion, $\ce{CuCl4^2-}$, is green.

EDIT: Per the comment from @AJKOER, I ran the experiment again, with powdered iron (Fisher, 100 mesh, electrolytic) and copper sulfate solution. The next figure shows the starting materials. I put a strong magnet under the little beaker to prevent iron, which was deliberately in excess, from going over in the decantation that would follow.

Iron and copper sulfate

Pouring the copper sulfate solution into the beaker resulted in a vigorous reaction and quite a bit of heat. I stirred the reaction mixture and let it go to completion. The magnet still stuck strongly to the bottom of the beaker, indicating that there was still substantial powdered iron remaining.

The mixture was decanted to another beaker, with some mostly useless filtering (no good filter paper at hand), and allowed to settle for about an hour. The copper particles produced in the reaction are very small and settle out very slowly. Some of the supernatant solution was transferred to a sample cell, illuminated from the left via an LED flashlight and photographed. This is shown in the next figure:

Ferrous io green color

Despite the light scattering, the green color of the ferrous sulfate solution is evident. The photo will be updated after more particulate has settled out.

  • $\begingroup$ Upgrade, nice experiment! However, employing CuSO4 implies if there is any Fe presence in your heavy-duty foil, per my analysis below, some green FeSO4 formation. This may also be present per your picture. $\endgroup$
    – AJKOER
    Commented Feb 19, 2020 at 18:59
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    $\begingroup$ Good point, but the blue and green colors from the copper species make it hard to know, without a spectrophotometer at home, what the green ferrous ion contribution happens to be. I have a bottle of powdered iron, so I will do the before and after experiment, using an excess of iron powder to reduce the copper ion to negligible levels. I will add the pics here ASAP. Too bad the question is closed: I see value in actually doing simple experiments. $\endgroup$
    – Ed V
    Commented Feb 19, 2020 at 19:11
  • $\begingroup$ Good answer and fun experiment. I'm sorry I contributed to the closing but at that time, it was granted because no details on the question. $\endgroup$ Commented Mar 20, 2021 at 22:17
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    $\begingroup$ @MathewMahindaratne No problem at all! You are one of the best folks here, and that is saying something, given how really good the high rep members are! By thge way, I am still trying to get that cell cleaned! ;-) $\endgroup$
    – Ed V
    Commented Mar 21, 2021 at 0:40

My guess is that the heavy duty aluminum foil you employed is likely alloyed with iron. To quote a source:

Reynolds Wrap® Aluminum Foil is 98.5% aluminum. The balance is made up of other elements: primarily iron and silicon.

As such, your reaction likely proceeds as follows:

$\ce{2 Al + 3 CuCl2 (aq) -> 2 AlCl3 (aq) + 3 Cu (s)}$

$\ce{Fe + CuCl2 (aq) -> FeCl2 (aq) + 2 Cu (s)}$

Per Wikipedia on Aluminium chloride, to quote:

Aluminium chloride (AlCl3), also known as aluminium trichloride, is the main compound of aluminium and chlorine. It is white, but samples are often contaminated with iron(III) chloride, giving it a yellow color.

However, in the current case, the contamination is ferrous chloride which per Wikipedia:

FeCl2 crystallizes from water as the greenish tetrahydrate, which is the form that is most commonly encountered in commerce and the laboratory. There is also a dihydrate. The compound is highly soluble in water, giving pale green solutions.

Unclear on your targeted product, but you can repeat with a pure Al foil to remove the coloration.


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