Redox Reactions Explained

Redox is a cross-cutting theme: it appears in metal displacement, electrolysis, rusting, extraction of metals, and organic chemistry. Papers 2 and 4 test it through definitions, identifying oxidation/reduction in equations, writing half-equations, and assigning oxidation numbers (Supplement).

Three definitions of oxidation and reduction

The 0620 syllabus uses three equivalent definitions. All say the same thing differently:

	Oxidation	Reduction
In terms of oxygen	Gain of oxygen	Loss of oxygen
In terms of hydrogen	Loss of hydrogen	Gain of hydrogen
In terms of electrons	Loss of electrons	Gain of electrons

The electron definition (OIL RIG) is the most general and the one used for half-equations and oxidation numbers.

Examples using oxygen definition

CuO + H2 -> Cu + H2O

• CuO loses oxygen: CuO is reduced to Cu
• H2 gains oxygen: H2 is oxidised to H2O
• H2 is the reducing agent (it causes reduction of the CuO)
• CuO is the oxidising agent (it causes oxidation of the H2)

Examples using electron definition

Mg + CuSO4 -> MgSO4 + Cu

• Mg loses 2 electrons: Mg -> Mg2+ + 2e- (Mg is oxidised)
• Cu2+ gains 2 electrons: Cu2+ + 2e- -> Cu (Cu2+ is reduced)
• Mg is the reducing agent
• Cu2+ is the oxidising agent

Oxidising agents and reducing agents

• An oxidising agent causes another substance to be oxidised. It is itself reduced (gains electrons).
• A reducing agent causes another substance to be reduced. It is itself oxidised (loses electrons).

The key insight: the agent is named for what it does to the other substance, not what happens to itself.

Common oxidising agents: oxygen, chlorine, acidified potassium manganate(VII), hydrogen peroxide. Common reducing agents: metals (Mg, Zn, Fe), carbon, hydrogen, carbon monoxide.

Half-equations

Half-equations show what happens at each electrode during electrolysis, or they separate the oxidation and reduction halves of a redox reaction.

Rules for writing half-equations

1. Show only the species being oxidised or reduced
2. Balance atoms
3. Balance charge by adding electrons (e-) to the appropriate side
4. Electrons appear on the left for reduction, on the right for oxidation

Examples

Reduction (gain of electrons):

• Cu2+ + 2e- -> Cu
• Pb2+ + 2e- -> Pb
• 2H+ + 2e- -> H2
• Cl2 + 2e- -> 2Cl-

Oxidation (loss of electrons):

• Mg -> Mg2+ + 2e-
• 2Cl- -> Cl2 + 2e-
• 2Br- -> Br2 + 2e-
• Fe2+ -> Fe3+ + e-

Combining half-equations

To get the overall equation, make sure electrons cancel:

Mg -> Mg2+ + 2e- (oxidation) Cu2+ + 2e- -> Cu (reduction)

Overall: Mg + Cu2+ -> Mg2+ + Cu

The electrons cancel because the same number is lost and gained. If they do not cancel, multiply one equation.

Oxidation numbers (Supplement)

Oxidation number (oxidation state) is a way to track electron transfer using numbers rather than dot-and-cross diagrams.

Rules for assigning oxidation numbers

Rule	Oxidation number
Uncombined elements	0
Simple monatomic ions	Equal to the charge (Na+ = +1, Cl- = -1)
Oxygen in compounds	-2 (except in peroxides: -1)
Hydrogen in compounds	+1 (except in metal hydrides: -1)
Fluorine	Always -1
Sum in a neutral compound	0
Sum in a polyatomic ion	Equal to the charge

Finding unknown oxidation numbers

In H2SO4: H is +1, O is -2, so S is: 2(+1) + S + 4(-2) = 0, giving S = +6.

In MnO4-: O is -2, so Mn is: Mn + 4(-2) = -1, giving Mn = +7.

Using oxidation numbers to identify redox

• Oxidation number increases = the element is oxidised
• Oxidation number decreases = the element is reduced

Example: Fe2O3 + 3CO -> 2Fe + 3CO2

• Fe goes from +3 (in Fe2O3) to 0 (in Fe): decrease, so Fe is reduced
• C goes from +2 (in CO) to +4 (in CO2): increase, so C is oxidised
• CO is the reducing agent; Fe2O3 is the oxidising agent

Redox in context across the syllabus

Topic	Redox example
Reactivity series	Displacement: Zn + CuSO4 -> ZnSO4 + Cu (Zn oxidised, Cu2+ reduced)
Extraction of metals	Blast furnace: Fe2O3 + 3CO -> 2Fe + 3CO2
Electrolysis	Cathode: reduction. Anode: oxidation
Rusting	Fe -> Fe2+ + 2e- (iron oxidised)
Halogens	Displacement: Cl2 + 2KBr -> 2KCl + Br2
Organic chemistry	Combustion: C and H oxidised. Fermentation: glucose reduced

Common exam mistakes

1. Confusing oxidation and reduction: Use OIL RIG every time. Loss of electrons = oxidation.
2. Naming the agent wrongly: The reducing agent is oxidised. The oxidising agent is reduced. This is counterintuitive but crucial.
3. Forgetting that redox involves both processes: If one species is oxidised, another must be reduced. They always occur together.
4. Not balancing electrons in half-equations: The number of electrons must balance atoms and charge.

Worked exam questions

In the reaction Zn + 2AgNO3 -> Zn(NO3)2 + 2Ag, identify the substance oxidised and the substance reduced. Explain in terms of electron transfer. [4 marks]

• Zinc is oxidised [1]: Zn -> Zn2+ + 2e- (loses electrons) [1]
• Silver ions are reduced [1]: Ag+ + e- -> Ag (gains electrons) [1]

In the blast furnace, iron(III) oxide is reduced by carbon monoxide: Fe2O3 + 3CO -> 2Fe + 3CO2. What is the oxidation number change for iron and for carbon? [2 marks]

• Iron changes from +3 to 0 (reduced) [1]
• Carbon changes from +2 to +4 (oxidised) [1]

Write half-equations for the electrolysis of molten sodium chloride. Identify which electrode shows oxidation and which shows reduction. [4 marks]

• Cathode: Na+ + e- -> Na (reduction) [1][1]
• Anode: 2Cl- -> Cl2 + 2e- (oxidation) [1][1]