Redox & Electron Transfer (CIE IGCSE Chemistry)

Extended tier only

• 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-

• For example, when iron reacts with a compound of copper such as copper sulfate, a displacement reaction occurs

iron + copper sulfate → iron(II) sulfate + copper 

Fe + CuSO₄ → FeSO₄ + Cu

• We can write this as an ionic equation

Fe + Cu²⁺ + SO₄^2– → Fe²⁺ + SO₄^2– + Cu

• We can then remove the spectator ions to see the overall change

Fe + Cu²⁺→ Fe²⁺ + Cu

• The iron atom has lost electrons to become a positive ion, so has been oxidised
• The positive copper ion has gained electrons to become an atom, so have been reduced

The redox reaction between Fe and Cu²⁺

Extended tier only

Identifying redox reactions using oxidation numbers

• The oxidation number is a number assigned to an atom or ion in a compound 
• It shows the number of electrons that an atom has lost, gained or shared in forming a compound
  • So, 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
  • Positive oxidation number = loss of electrons
  • Negative oxidation number = gain of electrons
  • For example, aluminium in a compound usually has the oxidation number of +3 indicating it has lost 3 electrons
  • Careful: It is easy to confuse oxidation number with charge which is written by a number followed by a +/- sign)
• A few simple rules help guide you through the process of determining the oxidation number of any element

Rules for assigning oxidation numbers

	Rule	Example
1	The oxidation number of any uncombined element is zero	H2  Zn O2
2	Many atoms or ions have fixed oxidation numbers in compounds	Group 1 elements are always +1 Group 2 elements are always +2 Fluorine is always –1 Hydrogen is +1, except in hydrides like NaH where it is –1 Oxygen is –2, except in peroxides where it is in –1 and in F2O where it is +2
3	The oxidation number of an element in a monoatomic ion is always the same as the charge	Zn2+ = +2 Fe3+ = +3 Cl– = –1
4	The sum of the oxidation numbers in a compound is zero	NaCl Na = +1 Cl = –1 Sum of oxidation numbers = 1 – 1 = 0
5	The sum of the oxidation numbers in an ion is equal to the charge on the ion	SO42– S = +6 Four O atoms = 4 x (–2) = –8 Sum of oxidation numbers = 6 – 8 = –2
6	In either a compound or an ion, the more electronegative element is given the negative oxidation number	F2O Two F atoms = 2 x (–1) = –2 O = +2

• 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

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

• The potassium iodide is oxidised as it loses electrons
• The hydrogen peroxide is reduced
• Therefore, potassium iodide is acting as a reducing agent 

Extended tier only

What is an an oxidising agent?

• An oxidising agent is 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

What is a reducing agent?

• A reducing agent is 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 

Identifying oxidising and reducing agents

CuO + H₂ → Cu + H₂O

• Hydrogen is reducing the CuO
• Hydrogen is itself oxidised as it has gained oxygen / lost electrons
• So, the reducing agent is hydrogen:

H₂ → 2H⁺ + 2e^-

• CuO is reduced by hydrogen
  • This means that the hydrogen is oxidised by CuO
• CuO is reduced as it has lost oxygen / gained electrons
• So, the oxidising agent is copper oxide

Cu²⁺ +2e^- →  Cu