Reduction-oxidation reaction demonstration

This exercise is a part of Educator Guide: Charging the Future / View Guide

Class time: About 15 to 20 minutes.

Purpose: This is a simple demonstration for introducing reduction-oxidation (redox) reactions.

Notes to the teacher: Please handle and dispose of copper (II) sulfate solution in accordance with federal, state and local environmental control regulations.


  • Copper(II) sulfate (available from chemical supply companies — sold as a hydrate, or sold as root killer at hardware stores)
  • A large steel or iron nail/screw/bolt (or more than one if you want)
  • A beaker or clear plastic cup
  • A stirring rod


1. Fill the beaker with hot water, pour in a few grams of copper(II) sulfate and stir until it dissolves. Show the students that the blue crystals of copper sulfate turn the water blue (due to Cu2+). Show the students the steel (iron alloy) or iron nail, then put it in the beaker. After a few minutes, pull out the nail and show it to the students. The nail should be coated with a thin layer of reddish-bronze copper atoms, which can be easily scraped off with a paper towel. If you use more than one nail, or have a dilute enough copper(II) sulfate solution, your students may be able to notice that the blue solution gets lighter in color due to the decreasing concentration of Cu2+ ions.

2. Explain that the copper starts off in the solution as ions with two positive charges (having donated two electrons to sulfate), and the iron starts off as a solid composed of neutral iron atoms with the same number of protons as electrons in each atom. Since the copper has a greater affinity for electrons than the iron does, the iron loses two electrons and gets dissolved into solution. Each Cu2+ ion gains two electrons and becomes a solid where the iron had been. In chemistry notation, the reactions are:

                                Oxidation:                       Fe(s) Fe2+(aq) + 2 e

                                Reduction:                       Cu2+(aq) + 2 eCu(s)

                                Net reaction:                  Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s)

3. The sulfate and the water were innocent bystanders in this reaction and did not directly participate. Explain that the ions, such as sulfate (SO42-), are called spectator ions.

4. Explain that this is just one example of an oxidation and reduction reaction, but there are many others. Reduction and oxidation reactions are a mouthful to say and always occur together (electrons have to come from someplace and go to someplace), so they are called “redox reactions” for short. Also, it is important to note that the number of electrons lost in a reaction must be equal to the number of electrons gained, so the chemical equation may need to be balanced. To remember the difference between oxidation and reduction reactions, use the mnemonic OIL RIG: Oxidation Is Loss of electrons, Reduction Is Gain of electrons.

5. Copper wanted electrons more than iron did in this demonstration. Scientists can make a list of atoms or molecular groups in order from those that are least determined to hang on to their electrons (or in other words, those most likely to be oxidized) to those most determined to keep their electrons (or in other words, those most likely to be reduced). This difference in affinity for electrons may be expressed in terms of volts of electric potential energy, or the standard reduction potential of each material. Show an example of your favorite reduction potential table such as this one given by California State University at Dominguez Hills.

6. Using the diagram above, explain that a battery is a device that converts the chemical energy of a redox reaction to usable electrical energy. A battery has an anode made of one material, a cathode made of another material, an electrolyte allowing ion flow between the anode and cathode and an external electrical circuit allowing electron flow between the anode and cathode. Ignoring real-world inefficiencies and imperfections, the voltage of the battery is the difference in the standard reduction potentials of the anode and cathode materials. You may want to demonstrate calculating the difference in reduction potentials from the demonstration materials. A Khan Academy video gives a brief summary of how to calculate the overall redox reaction potentials under standard conditions from a standard reduction potential table.

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