$$\require{mhchem}$$
A 4X4 matrix of test tubes is set up. Zinc, lead, copper, and silver are placed in test tubes containing solutions of the nitrate salts of each of the four metals. When a metal is placed in a solution containing the salt of a metal that is less active than itself, the less active metal plates out.

The effectiveness of this demonstration can be enhanced by 1.) pairing a computer animation representing the dynamic events that occur at the surface of the metal, 2.) having students work in small groups on the activity series of metals activity, and 3.) extending the activity series of metals having students work the activity series of metals computer simulation and design their own activity series of metals.

Curriculum Notes 

This demonstration is best used when electrochemistry is being introduced. Students interacting with this demonstration observe an order of activity of metals and should identify the oxidation half-reaction and the reduction half-reaction occurring in each reaction.  Allow about twenty minutes for this demonstration. 

One day of lead time is required for this project.

Discussion 

Standard Reduction Potentials:

$$\ce{Zn^+2(aq) + 2e^{-} -> Zn(s);~-0.76 V}$$

$$\ce{Pb^+2(aq) + 2e^{-} -> Pb(s);~-0.13 V}$$

$$\ce{Cu^+2(aq) + 2e^{-} -> Cu(s);~+0.34 V}$$

$$\ce{Ag^{+}(aq) + e^{-} -> Ag(s);~~~+0.80 V}$$

Metals with a greater reduction potential will plate on on metals with a lesser reduction potential. That is to say, the metals with a greater reduction potential are more stable in their reduced (metallic) form than those with a lesser reduction potential. When you think about the properties of these metals, this makes intuitive sense. Silver has been used for money because its metallic form is very stable, whereas zinc is used for sacrificial anodes on ships because it is more "active" than the iron in the ships' steel hulls.

Materials 

• Sixteen 25X175 mm test tubes arranged in a 4X4 matrix in a test tube rack. Each test tube contains an aqueous solution to a depth of about 3 cm as follows:
  • The tubes on the first row contain 0.1 M silver nitrate.
  • The tubes on the second row contain 0.1 M copper(II) nitrate.
  • The tubes on the third row contain 0.1 M lead(II) nitrate.
  • The tubes on the fourth row contain 0.1 M zinc nitrate.
• Four 125 mL beakers each containing four about 60 X 8 mm strips of a certain metal as follows:
  • One beaker contains four silver strips.
  • One beaker contains four copper strips.
  • One beaker contains four lead strips.
  • One beaker contains four zinc strips.
• The test tubes in each row are marked with a piece of colored tape corresponding to the color of tape on the beaker containing the strips of corresponding metal, e.g., the row of test tubes containing zinc nitrate solution and the beaker containing zinc strips might both be marked with blue tape.
• Stoppers to fit the test tubes. (optional)

Procedure 

There are several ways to present this demo. Each experiment is initiated by dropping a metal strip into one of the solutions. The students can use the attached worksheet and standard reduction potentials to predict which test tubes a reaction will occur in. They can then suggest which strips should be dropped into which solutions. When all of the strips have been dropped in, it should form a 4X4 matrix with the rows consisting of particular metal nitrate salt solutions and the columns consisting of particular metals. A smaller class can come up around the demo table to observe the reaction. In a larger class, video projection can be used or the test tubes can be stoppered and passed around the class. 

Safety Precautions 

All nitrate solutions are oxidizers. Lead nitrate is toxic. If any of these liquids comes into contact with your skin, flush with copious amounts of water. If you get some in your eyes, flush with water for at least 15 minutes and seek medical attention. Wear goggles. 

Footnotes 

References

Bassam Z. Shakhashiri, “An Activity Series: Zinc, Copper, and Silver Half Cells,” Chemical Demonstrations: A Handbook for Teachers of Chemistry, Volume 4 (Madison: The University of Wisconsin Press, 1992), pp. 101−106.