Warning: Tin is slightly toxic. Copper compounds are toxic. Do not eat or drink them.
You will need:
Tin or high-tin alloy
Tin sits right between germanium, which forms a stable +4 oxidation state, and lead, which forms a stable +2 oxidation state. Because of this, its +2 oxidation state is stable, but a reducing agent. Tin can be oxidized twice.
Dissolve some copper sulfate in water. It is a blue solution. Add tin. The solution becomes more colorless. If it has a little excess sulfuric acid, the tin(II) sulfate will stay in solution. If the solution is slightly basic, tin(II) oxysulfate will precipitate as a white powder. Here, the tin(II) does not behave as a reductant. This is a picture of the reaction Sn + CuSO4 => SnSO4 + Cu:
Here is a picture of the final precipitate, filtered and dried:
Tin reacts differently with copper(II) chloride. The difference is that tin(II) sulfate cannot be oxidized, while tin(II) chloride can be.
Dissolve some copper(II) chloride in water. The solution should be green. Add tin. At first, the reaction Sn + CuCl2 => SnCl2 + Cu happens, and red copper precipitates. But then, the solution starts becoming cloudy as SnCl2 + 2 CuCl2 => SnCl4 + 2 CuCl happens. White CuCl precipitates. The SnCl4 hydrolyzes, creating white SnO2 precipitate. Then, the copper(I) chloride oxidizes: 4 CuCl + O2 => 2 CuCl2 + 2 CuO, creating black copper(II) oxide and soluble green copper(II) chloride, which starts the process over again. In my case, the CuCl2 ran out, leaving behind a solution of tin(II) chloride and a precipitate of tin(IV) oxide and copper(II) oxide. The latter's color predominates.
In this first picture, the tin has just begun reacting, and copper is visible on it.
In the second picture, the tin(II) chloride has been reducing the copper(II) chloride.
In the third picture, the reaction is complete.
The yellow color is the result of impurities in the tin. The white square is a stain cut out of the background.