Topic 8 · page 5 of 29
Oxidising and reducing agents
An oxidising agent takes electrons and is reduced; a reducing agent gives electrons and is oxidised. Group 1 and 2 metals are reducing agents and oxygen and the halogens are oxidising agents, in every s-block and p-block reaction you need.
Read first: Oxidation and reduction
The two agents 8.6
The agent does the opposite to the species it acts on. An oxidising agent oxidises something else, so it must be reduced itself. If you swap the two in your answer, the mark is lost.
Worked exampleZinc and copper(II) ions
. Identify what is oxidised and reduced, and the two agents.
- Zn: 0 → +2. The number increases, so zinc is oxidised: it loses 2 electrons.
- Cu: +2 → 0. The number decreases, so copper(II) is reduced: it gains 2 electrons.
- The reducing agent is Zn: it gives the electrons, and is the species oxidised.
- The oxidising agent is Cu²⁺: it takes the electrons, and is the species reduced.
| Typical oxidising agents (gain electrons) | Typical reducing agents (lose electrons) |
|---|---|
| O₂, Cl₂, Br₂ | Group 1 and 2 metals, Zn, Fe |
| MnO₄⁻ in acid, Cr₂O₇²⁻ in acid | I⁻, Br⁻ |
| Cu²⁺, Ag⁺, Fe³⁺ | Fe²⁺, SO₂ |
| concentrated H₂SO₄ | H₂, CO |
Hydrogen peroxide, H₂O₂, can be either. It is an oxidising agent towards iodide ions and a reducing agent towards acidified manganate(VII). Which role it takes depends on what it is reacting with.
Reactions of s-block and p-block elements 8.5
The metals of Groups 1 and 2 are the s-block. They have one or two outer electrons and lose them easily, so they are reducing agents in every reaction you need. The elements of Groups 6 and 7 (oxygen, the halogens) are in the p-block. They are short of electrons, so they are oxidising agents.
| Reaction | Oxidation numbers | Oxidised | Reduced |
|---|---|---|---|
| Na 0 → +1; Cl 0 → −1 | Na | Cl₂ | |
| Mg 0 → +2; O 0 → −2 | Mg | O₂ | |
| Na 0 → +1; H +1 → 0; O stays −2 | Na | H in H₂O | |
| Mg 0 → +2; H +1 → 0; Cl stays −1 | Mg | H in HCl | |
| Cl 0 → −1; Br −1 → 0 | Br⁻ | Cl₂ |
Two trends, to be explained in later pages, follow from this pattern. Down Groups 1 and 2 the metals become stronger reducing agents, because the outer electrons are lost more easily. Down Group 7 the halogens become weaker oxidising agents, because they gain electrons less easily.
Check yourself
- In , identify the species oxidised, the species reduced and the reducing agent.
- In , state the change in oxidation number of each element that changes and name the oxidising agent.
- Explain why Group 1 metals are always reducing agents.
- State the oxidation number of oxygen in the reaction before and after, and explain what has happened to the oxygen.
- Calcium reacts with water: . State, with numbers, what is oxidised and what is reduced.
- A student says "hydrogen peroxide is an oxidising agent". Give an example where this is not true, and explain what decides the role.
Answers cover these until you have tried the questions
- Zn: 0 → +2, so zinc is oxidised. H: +1 → 0, so hydrogen (in HCl) is reduced. Zinc gives the electrons, so zinc is the reducing agent. Cl is −1 on both sides.
- Cl: 0 → −1 (reduced). I: −1 → 0 (oxidised). The oxidising agent is Cl₂: it takes electrons from the iodide ions.
- A Group 1 atom has one outer electron, which it loses easily, so its oxidation number rises from 0 to +1. A species that loses electrons is a reducing agent.
- O: 0 in O₂ and −2 in CaO. Oxygen has gained electrons: it is reduced, and it is the oxidising agent.
- Calcium is oxidised: 0 → +2. Hydrogen in water is reduced: +1 → 0. Oxygen stays −2.
- With acidified potassium manganate(VII), H₂O₂ is oxidised to O₂ (oxygen −1 → 0), so it is the reducing agent. Whether it is oxidising or reducing depends on what it reacts with: it gains electrons from a stronger reducing agent such as I⁻, and loses electrons to a stronger oxidising agent such as MnO₄⁻.