Oxidation & Reduction (Cambridge (CIE) O Level Chemistry): Revision Note

Oxidation & Reduction

Redox reactions

• Oxidation and reduction take place together at the same time in the same reaction

• These are called redox reactions

• Oxidation can be defined as:

  • The gain of oxygen

  • The loss of electrons

  • An increase in oxidation number

• Reduction can be defined as:

  • The loss of oxygen

  • The gain of electrons

  • A decrease in oxidation number

Names using oxidation numbers

• Transition elements can bond in different ways by forming ions with different charges

• When naming, the charge on the ion is shown by using a Roman numeral after the element's name

  • e.g. iron can form ions with a 2+ charge, called iron(II) ions or a 3+ charge, called iron(III) ions

• The Roman numeral is the oxidation number of the element

• When iron reacts with oxygen to form iron oxide, the formula depends on the oxidation state of the iron ions

  • The compound where iron has a 2+ charge has the formula FeO and is called iron(II) oxide

  • The compound where iron has a 3+ charge has the formula Fe₂O₃ and is called iron(III) oxide

Oxidation and reduction in terms of oxygen

• Oxidation is a reaction in which oxygen is added to an element or a compound

• Reduction is a reaction in which oxygen is removed from an element or compound

Example: Identifying the loss and gain of oxygen in an equation

zinc oxide + carbon → zinc + carbon monoxide

ZnO + C → Zn + CO

• In this reaction, the zinc oxide has been reduced since it has lost oxygen

• The carbon atom has been oxidised since it has gained oxygen

Oxidation and reduction in terms of electrons and oxidation number

• Redox reactions can also be defined in terms of electron transfer

• Oxidation is a reaction in which an element, ion or compound loses electrons

  • The oxidation number of the element is increased

  • This can be shown in a half equation, e.g. when silver reacts with chlorine, silver is oxidised to silver ions:

Ag → Ag⁺ + e^-

• Reduction is a reaction in which an element, ion or compound gains electrons

  • The oxidation number of the element is decreased

  • This can be shown in a half equation, e.g. when oxygen reacts with magnesium, oxygen is reduced to oxide ions:

O₂ + 4e^- → 2O^2-

Example: Identifying Redox Reactions

zinc + copper sulphate → zinc sulphate + copper

Zn + CuSO₄ → ZnSO₄ + Cu

• The ions present (with state symbols) in the equation are:

Zn(s) + Cu²⁺(aq) + SO₄^2-(aq) →Zn²⁺(aq) + SO₄^2-(aq) + Cu(s)

• The spectator ions (those that do not change) are SO₄^2-(aq)

• These  can be removed and the ionic equation written as:

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

• By analysing the ionic equation, we can split the reaction into two half equations by adding in the electrons to show how the changes in charge have occurred:

Zn(s) → Zn²⁺(aq) + 2e^-

Cu²⁺(aq) +2e^- → Cu(s)

• It then becomes clear that zinc has been oxidised as it has lost electrons 

• Copper ions have been reduced as they have gained electrons

Identifying Redox Reactions

Oxidation Number

• The oxidation number (also called oxidation state) is a number assigned to an atom or ion in a compound which indicates the degree of oxidation (or reduction)

• It shows the number of electrons that an atom has lost, gained or shared in forming a compound

• The oxidation number helps you to keep track of the movement of electrons in a redox process

• It is written as a +/- sign followed by a number (not to be confused with charge which is written by a number followed by a +/- sign)

• E.g. aluminium in a compound usually has the oxidation state +3

• A few simple rules help guide you through the process of determining the oxidation number of any element

Table of Rules for Assigning Oxidation Numbers 

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

Identifying Redox Reactions by Colour Changes

• The tests for redox reactions involve the observation of a colour change in the solution being analysed

• 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 purple to colourless

Diagram to show the colour change when potassium manganate(VII) is added to a reducing agent

• 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 (H₂O₂), the solution turns a red-brown colour due to the formation of iodine, I₂:

2KI (aq) + H₂SO₄ (aq) + H₂O₂ (aq) →  I₂ (aq) + K₂SO₄ (aq) + 2H₂0 (l)

• The potassium iodide is oxidised as it loses electrons and hydrogen peroxide is reduced, therefore potassium iodide is acting as a reducing agent as it will itself be oxidised:

2I^- →  I₂ + 2e^-

Diagram to show the colour change when potassium iodide is added to an oxidising agent

Oxidising & Reducing Agents

Oxidising agent

• A substance that oxidises another substance, and becomes reduced in the process

• An oxidising agent gains electrons as another substance loses electrons

• Common examples include hydrogen peroxide, fluorine and chlorine

Reducing agent

• A substance that reduces another substance, and becomes oxidised in the process

• A reducing agent loses electrons as another substance gains electrons

• 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 

Example

CuO + H₂ → Cu + H₂O

• In the above reaction, hydrogen is reducing the CuO and is itself oxidised as it has lost electrons, so the reducing agent is therefore hydrogen:

H₂ → 2H⁺ + 2e^-

• The CuO is reduced to Cu by gaining electrons and has oxidised the hydrogen, so the oxidising agent is therefore copper oxide

Cu²⁺ +2e^- →  Cu