Redox & Electron Transfer (Cambridge International IGCSE Chemistry)

Spec point — 6.3.2 · Define oxidation in terms of: loss of electrons an increase in oxidation number Define reduction in terms of: gain of electrons a decrease in oxidation number Identify redox reactions as reactions involving gain and loss of electrons

Redox & electron transfer

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⁺

The Fe atom is oxidised (loses electrons) and the Cu²⁺ion is reduced (gains electrons)

Worked Example Which change in the following equation is oxidation?

V³⁺+ Fe³⁺→ V⁴⁺+ Fe²⁺

Answer:

• Step 1 - Identify the changes for each species:

  • V³⁺ to V⁴⁺

    • V³⁺ has lost 1 electron

  • Fe³⁺ to Fe²⁺

    • Fe³⁺ has gained 1 electron

• Step 2- Identify each change as either oxidation and reduction

  • V³⁺ to V⁴⁺ is oxidation
  • Fe³⁺ to Fe²⁺ is reduction

• Therefore, V³⁺ has been oxidised

Exam Hint Use the mnemonic OIL-RIG to remember oxidation and reduction in terms of the movement of electrons:

• Oxidation I s Loss
• Reduction Is Gain.

Identifying redox reactions

Spec point — 6.3.2 · Identify redox reactions by changes in oxidation number using: the oxidation number of elements in their uncombined state is zero the oxidation number of a monatomic ion is the same as the charge on the ion the sum of the oxidation numbers in a compound is zero the sum of the oxidation numbers in an ion is equal to the charge on the ion Identify redox reactions by the colour changes involved when using acidified aqueous potassium manganate(VII) or aqueous potassium iodide

Identifying redox reactions

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

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

Worked Example The equation for the reaction between chlorine and potassium iodide is shown below.

Cl₂ + 2KI → 2KCl + I₂

Identify which species has been:

1. Oxidised
2. Reduced

Answer:

1. The species that has been oxidised is iodine

• 2I^- → I₂ +2e^-
• The oxidation number of I^- is -1
• The oxidation number of iodine in I₂ is 0
• The oxidation number has increased so the iodide ions have been oxidised / lost electrons

1. The species that has been reduced is chloride ions

• Cl₂ + 2e^- → 2Cl^-
• The oxidation number of chlorine as Cl₂ is 0
• The oxidation number of Cl^- is -1
• The oxidation number has decreased so the Cl₂ has been reduced / gained electrons

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

Oxidising & reducing agents

Spec point — 6.3.2 · Define an oxidising agent as a substance that oxidises another substance and is itself reduced Define a reducing agent as a substance that reduces another substance and is itself oxidised Identify oxidising agents and reducing agents in redox reactions

Oxidising & reducing agents

Extended tier only

What is 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

Worked Example When iron reacts with bromine to form iron(II) bromide, a redox reaction reaction occurs:

Fe + Br₂ → FeBr₂

Which species is acting as the reducing agent in this reaction?

Answer

1. Step 1- Write half equations to work out what has gained/lost electrons

• Fe → Fe²⁺ + 2e^-
• Br₂ + 2e^- → 2Br^-
• Fe loses electrons; Br₂ gains electrons

1. Step 2 - Deduce what has been oxidised/reduced (remember OIL RIG)

• Fe has been oxidised as it has lost electrons
• Br₂ has been reduced as it has gained electrons

1. Step 3 - Identify the reducing agent

• Fe is the reducing agent as it has been oxidised by losing electrons and caused Br₂ to be reduced as it gained electrons