Cutting a can - microscale version

by Alfredo Mateus

Electron transfer reactions, the so-called redox reactions, are an important part of understanding electrochemical processes, like those present in corrosion or in electrochemical cells and batteries.

The reactions between copper sulfate and a metal sheet, like zinc or iron, or even steel wool are commonly used in this context. In these experiments, although it is easy to see the copper deposit, it is not as easy to realize that the metal sheet is being oxidized and the metal is going into the solution. 

In this experiment, we can easily see the reduction of the copper ions and the oxidation of the aluminum in the can. And the best part is that, being a microscale experiment, we will spend a minimal ammount of reagents, will end up with a very small ammount of residues, and with a very easy clean up. Let's cut an aluminum can with Chemistry.

Cut a piece of the can and draw on it with a stencil

Use scissors to cut a square piece of the walls of a regular aluminum can. You may wish to find a white can to have a better contrast of the colors. Round up the corners of the square to avoid sharp edges. 
Find something with a sharp point, like a compass or a pin, and use it to scratch away the ink of the can. In our example, we used a stencil to draw a circle. It is very important to only scratch around the circle, and not inside it. That is a common mistake that some students do. We only want the reaction to take place at the line around the circle we are drawing..
Of course, you can draw any other shape you want. You can even make your own stencil using 3D printing. 

We need to make the part of the square with the drawing flat. To do that, just bend the four edges of the square up.

Now, place the square over a transparent plastic folder with a white paper sheet inside of it. The folder is just for having an easy to clean surface to work and to have a good contrast for the solution color after the reaction. The idea of using the plastic folder comes from Bob Worley's brilliant article on the Journal of Chemical Education. ( 

Add a few drops (around 5) of a copper chloride solution 1 mol/L. If you don't have copper chloride, or if you want to show the role of the chloride ions in the reaction, you can use copper sulfate and then add sodium chloride. 

Observe, film and take pictures

To observe the reaction closely, our suggestion is to let the students use their smartphone's camera. Provide a support for the phones that will let them frame the reaction well. 

In a few seconds we can see the formation of a brown solid over the exposed aluminum. We can also notice the formation of gas bubbles.

Keep following the reaction until the gas bubbles slow down. You can feel with the point of the compass if all the aluminum has reacted away and if you can lift one edge and remove the interior of the shape you've drawn. This takes around 2-3 minutes.

You can create you own stencil to cut the cans using a 3D printer. That would be a great way to get students involved to learn TinkerCAD. And if you don't have a compass, you can use a pin to scratch away the ink on the can.

What happened?

We have many clues that a chemical reaction happened. The blue solution with copper 2+ ions became colorless. A reddish brown solid formed over the part of the aluminum that was exposed. Gas bubbles appeared during the reaction. And, maybe the most surprising one, we managed to cut the aluminum can without touching it.

Copper 2+ ions give the solution its blue color. The blue color that disappears and the brown solid that formed are related. We can write an equation with the copper ions as a reagent and copper metal as the product.

The copper ions are receiving electrons and undergoing reduction to copper metal. The equation for this part of the process becomes like this:

If you add just copper sulfate to an aluminun sheet, you will notice that nothing happens. Use copper chloride or add chloride ions to the copper sulfate and the reaction happens immeadiately. Aluminum forms an oxide layer on its surface that protects the metal against oxidation. The chloride ions remove this protective layer and exposes the aluminum to the copper ions in solution.

But, where did the bubbles come from? What is this gas? The exposed aluminun is so reactive that it reacts not only with the copper ions but also with water! Hydrogen ions are reduced by the aluminum and change into hydrogen gas.

Last thoughts

This is a good example of a microscale reaction, where we can at the same time use less materials and make a better understanding of the reaction possible.

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