CIE IGCSE Chemistry

Revision Notes

7.3 Redox Reactions

Oxidation & Reduction

  • Oxidation and reduction take place together at the same time in the same reaction
  • These are called redox reactions
  • There are three definitions of oxidation. It is a reaction in which:
    • Oxygen is added to an element or a compound
    • An element, ion or compound loses electrons
    • The oxidation state of an element is increased
  • There are three definitions of reduction. It is a reaction in which:
    • Oxygen is removed from an element or a compound
    • An element, ion or compound gains electrons
    • The oxidation state of an element is decreased

Oxidation state

  • The oxidation state (also called oxidation number) is a number assigned to an atom or ion in a compound which indicates the degree of oxidation (or reduction)
  • The oxidation state helps you to keep track of the movement of electrons in a redox process
  • It is written as a +/- sign followed by a number.
  • Eg O2- means that it is an atom of oxygen that has an oxidation state of -2. It is not written as O2- as this refers to the ion and its charge

Assigning the oxidation number

  • Oxidation number refers to a single atom or ion only
  • The oxidation number of a compound is 0 and of an element (for example Br in Br2) is also 0
  • The oxidation number of oxygen in a compound is always -2 (except in peroxide R-O-O-R, where it is -1)
  • For example in FeO, oxygen is -2 then Fe must have an oxidation number of +2 as the overall oxidation number for the compound must be 0

Ionic Equations

  • Ionic equations are used to show only the particles that actually take part in a reaction
  • These equations show only the ions that change their status during a chemical process, i.e: their bonding or physical state changes
  • The other ions present are not involved and are called spectator ions

Writing ionic equations

  • For the neutralisation reaction between hydrochloric acid and sodium hydroxide:
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
  • If we write out all of the ions present in the equation and include the state symbols, we get:
H+(aq) + Cl- (aq)+ Na+(aq) + OH-(aq) → Na+ (aq)+ Cl-(aq) + H2O(l)
  • The spectator ions are thus Na+ and Cl. Removing these from the previous equation leaves the overall net ionic equation:
H+(aq) + OH-(aq) →H2O(l)
  • This ionic equation is the same for all acid-base neutralisation reactions

Example redox equation: oxygen loss/gain

Zinc oxide + carbon → zinc + carbon monoxide
ZnO + C → Zn + CO
  • In this reaction the zinc oxide has been reduced since it has lost The carbon atom has been oxidised since it has gained oxygen
Extended Only

Redox & Electron Transfer

Example redox equation: electron loss/gain and oxidation state

Zinc + copper sulphate → zinc sulphate + copper
Zn + CuSO4 → ZnSO4 + Cu
  • Writing this as an ionic equation:
Zn(s) + Cu2+(aq) + SO42-(aq) →Zn2+(aq) + SO42-(aq) + Cu(s)
  • By analysing the ionic equation, it becomes clear that zinc has become oxidised as its oxidation state has increased and it has lost electrons:
Zn(s) →Zn2+(aq)
  • Copper has been reduced as its oxidation state has decreased and it has gained electrons:
Cu2+(aq) → Cu(s)

Exam Tip

Use the mnemonic OIL-RIG to remember oxidation and reduction in terms of the movement of electrons: Oxidation Is Loss –  Reduction Is Gain.

Extended Only

Oxidising & Reducing Agents

Oxidising agent

  • A substance that oxidises another substance, in so doing becoming itself reduced
  • Common examples include hydrogen peroxide, fluorine and chlorine

Reducing agent

  • A substance that reduces another substance, in so doing becoming itself oxidised
  • Common examples include carbon and hydrogen
  • The process of reduction is very important in the chemical industry as a means of extracting metals from their ores 


CuO + H2 →Cu + H2O
  • In the above reaction, hydrogen is reducing the CuO and is itself oxidised, so the reducing agent is therefore hydrogen
  • The CuO is reduced to Cu and has oxidised the hydrogen, so the oxidising agent is therefore copper oxide
Extended Only

Redox Reactions

Identifying redox reactions

  • Redox reactions can be identified by the changes in the oxidation states when a reactant goes to a product


Chlorine + potassium iodide → potassium chloride + iodine
Cl2 + 2KI → 2KCl + I2
  • Chlorine has become reduced as its oxidation state has decreased from 0 to -1 on changing from the chlorine molecule to chloride ions:
Cl2(g) → 2Cl-(aq)
  • Iodine has been oxidised as its oxidation state has increased from -1 to 0 on changing from iodide ions to the iodine molecule:
2I-(aq) → I2(s)


Identifying redox reactions by colour changes

  • The tests for redox reactions involve the observation of a colour change in the solution being analyse
  • Two common examples are acidified potassium manganate(VII), and potassium iodide
  • Potassium manganate (VII), KMnO4, is an oxidising agent which is often used to test for the presence of reducing agents
  • When acidified potassium manganate (VII) is added to a reducing agent its colour changes from pink-purple to colourless
  • Potassium iodide, KI, is a reducing agent which is often used to test for the presence of oxidising agents
  • When added to an acidified solution of an oxidising agent such as aqueous chlorine or hydrogen peroxide, the solution turns a brown colour due to the formation of iodine

Author: Morgan

Morgan’s passion for the Periodic Table begun on his 10th birthday when he received his first Chemistry set. After studying the subject at university he went on to become a fully fledged Chemistry teacher, and now works in an international school in Madrid! In his spare time he helps create our fantastic resources to help you ace your exams.

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