Oxidation and reduction

Burning, rusting and browning

An iron nail rusts. A cut apple turns brown. A candle burns. A battery runs. Each of these is a redox reaction. The name joins two words, reduction and oxidation, because the two always happen together.

The old idea: oxygen and hydrogen

The word oxidation comes from oxygen. Early chemists used it for any change in which a substance burned in air or combined with oxygen. The opposite change, the removal of oxygen, they called reduction. Later the idea grew to include hydrogen and other elements.

• Oxidation is the addition of oxygen, or of another electronegative element, to a substance. It is also the removal of hydrogen, or of another electropositive element, from a substance.
• Reduction is the removal of oxygen, or of another electronegative element, from a substance. It is also the addition of hydrogen, or of an electropositive element, to a substance.

Here are some examples.

Magnesium gains oxygen, so magnesium is oxidised.

Copper(II) oxide loses oxygen, so it is reduced. Hydrogen gains oxygen, so it is oxidised.

Hydrogen sulfide loses hydrogen, so it is oxidised. Chlorine gains hydrogen, so it is reduced.

Iron(III) chloride loses chlorine, which is electronegative, so it is reduced. Hydrogen gains chlorine, so it is oxidised.

Look at the copper oxide reaction again. Something is oxidised and something is reduced in the same change. A substance cannot be oxidised unless another substance is reduced. That is why we say redox.

The better idea: electron transfer

The classical idea cannot explain a reaction such as sodium burning in chlorine, where no oxygen or hydrogen is involved. A simpler idea works for every case. Look at magnesium burning in oxygen. We can split the change into two parts.

Each part is a half-reaction. In the first, magnesium atoms lose electrons. In the second, oxygen molecules gain them. We now define the two words by electrons.

• Oxidation is the loss of electrons.
• Reduction is the gain of electrons.

A memory aid is OIL RIG: oxidation is loss, reduction is gain. Add the two half-reactions and the electrons cancel. The electrons lost by one species are exactly the electrons gained by the other.

The species that takes the electrons is the oxidising agent, or oxidant. It is itself reduced. The species that gives the electrons is the reducing agent, or reductant. It is itself oxidised. In the magnesium reaction, oxygen is the oxidising agent and magnesium is the reducing agent.

Zinc and copper

Dip a zinc rod into blue copper(II) sulfate solution. After a few minutes the blue colour fades. The rod becomes coated with a red-brown layer of copper, and the solution warms a little. The two half-reactions are:

Adding them gives the overall reaction:

Zinc is the reducing agent. Copper(II) is the oxidising agent. Now turn the experiment round. Dip a copper rod into zinc sulfate solution and nothing happens. Zinc has won the competition for electrons.

Try one more. Dip a copper rod into silver nitrate solution. Shiny crystals of silver grow on the rod, and the solution slowly turns blue:

So copper beats silver, and zinc beats copper. As reducing agents, the order is zinc, then copper, then silver. As oxidising agents, the ions go the other way: $\text{Ag}^+$ is strongest and $\text{Zn}^{2+}$ is weakest. A later course turns this competition into a number, the electrode potential.

Move the slider. How many moles of electrons pass for each mole of zinc? Is the copper gained heavier or lighter than the zinc lost?

Practise with new numbers

Try it in Python

The program counts the electrons and the masses when zinc dissolves in copper(II) sulfate.

Common mistakes

• Mixing up the agent and the change. The reducing agent is the one that is oxidised. The oxidising agent is the one that is reduced.
• Thinking one half-reaction can happen alone. Electrons must come from somewhere and go somewhere.
• Forgetting the number of electrons. Each zinc atom gives 2 electrons, so 1 mol of zinc gives 2 mol of electrons.
• Using the classical rule where there is no oxygen or hydrogen. Use electron transfer instead.

Remember these

• Oxidation is loss of electrons. Reduction is gain of electrons.
• They always happen together, and the electrons lost equal the electrons gained.
• A reducing agent gives electrons and is oxidised. An oxidising agent takes electrons and is reduced.
• A metal that wins the competition for electrons is the stronger reducing agent.

Check yourself

Try each one on your own first. If you are stuck, open a hint. Then check with Show the answer.

1. In $\text{CuO} + \text{H}_2 \rightarrow \text{Cu} + \text{H}_2\text{O}$, which substance is oxidised in the classical sense?

2. In $2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl}$, which species gains electrons?

3. A silver strip is dropped into copper(II) sulfate solution. What happens, and why?

4. $0.20$ mol of zinc dissolves in copper(II) sulfate. How many moles of electrons pass to copper(II) ions?