Unit 1 — Chemical Reactions and Their Environmental Impacts

Science — Third Preparatory

Unit 1 — Chemical Reactions and Their Environmental Impacts

1.6 Oxidation–reduction (redox) reactions

In earlier years you learned a definition based on oxygen and hydrogen:

Key points

Oxidationincrease in the percentage of oxygen, or decrease in the percentage of hydrogen, in a substance.

Reductionincrease in the percentage of hydrogen, or decrease in the percentage of oxygen.

But in many reactions (for example zinc in copper sulphate solution) there is no oxygen or hydrogen at all, yet clearly something is oxidised and something reduced.
We need a definition in terms of electrons.
Look again at Zn + CuSO₄ → ZnSO₄ + Cu. The ionic form is

Zn + Cu²⁺ → Zn²⁺ + Cu

(The sulphate ion SO₄²⁻ is unchanged and is called a spectator ion.)
Before the reaction, zinc consists of neutral atoms and copper is present as positive ions (Cu²⁺) in the solution.
During the reaction:

Comparison aspect	Main idea	Keywords
Oxidation half-reaction	a zinc atom loses two electrons and becomes a zinc ion: Zn → Zn²⁺ + 2e⁻	zincatomloseselectrons
Reduction half-reaction	a copper ion gains those two electrons and becomes a neutral copper atom: Cu²⁺ + 2e⁻ → Cu	coppergainsthoseelectrons
•	Therefore	Therefore

Oxidation is a process in which an atom or ion loses one or more electrons.
Reduction is a process in which an atom or ion gains one or more electrons.
A reducing agent is the substance that loses electrons (it is itself oxidised and it causes the other substance to be reduced).
An oxidising agent is the substance that gains electrons (it is itself reduced and it causes the other to be oxidised).

Oxidation and reduction always occur together and in equal amounts of electrons; one cannot happen without the other.
Hence the name "redox".
In the reaction above, zinc is oxidised and is the reducing agent; copper ions (the copper sulphate) are reduced and copper sulphate is the oxidising agent.
A useful memory aid OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons).

Worked example 1: magnesium and bromine

Mg + Br₂ → MgBr₂→Mg → Mg²⁺ + 2e⁻. Each bromine atom gains one electron: Br₂ + 2e⁻ → 2Br⁻.

Key points

Oxidisedmagnesium.

Reducedbromine.

Oxidising agentbromine (it gains electrons).

Reducing agentmagnesium (it loses electrons).

Worked example 2: magnesium and hydrochloric acid

Mg + 2HCl → MgCl₂ + H₂→oxidation, so Mg is the reducing agent. The hydrogen ions of the acid gain electrons and become hydrogen molecules: 2H⁺ + 2e⁻ → H₂ : reduction, so the acid (the hydrogen ions) is the oxidising agent.

Worked example 3: sodium and water

Na + 2H₂O → 2NaOH + H₂.
Sodium atoms each lose one electron (Na → Na⁺ + e⁻): sodium is oxidised, the reducing agent.
Hydrogen of water gains electrons to make H₂: it is reduced; water is the oxidising agent.

Note that every single replacement reaction of a metal is a redox reaction, because a metal atom loses electrons while another ion gains them.
Combination reactions between a metal and a non-metal, and the combustion of a metal in oxygen (2Mg + O₂ → 2MgO: Mg loses electrons, oxygen gains) are also redox.

Why the reactivity series matters. The more reactive a metal, the more readily its atoms lose electrons.
That is why potassium and sodium react so violently, and why gold hardly reacts at all.

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