Half Equations (HL IB Chemistry)

Redox titrations

• In titrations, the concentration of a solution is determined by titrating with a solution of known concentration.
• In redox titrations, an oxidising agent is titrated against a reducing agent
  • Electrons are transferred from one species to the other
• Indicators are sometimes used to show the endpoint of the titration
• However, most transition metal ions naturally change colour when changing the oxidation state
• There are two common redox titrations you should know about:
  • Manganate(VII) titrations
  • Iodine-thiosulfate titrations

Manganate(VII) titrations

• A redox reaction occurs between acidified manganate(VII) ions and iron(II) ions:

MnO₄^– (aq) +  8H⁺ (aq) +  5Fe²⁺ (aq)  → Mn²⁺ (aq) +  5Fe³⁺ (aq) + 4H₂O (l)

• This reaction needs no indicator as the manganate(VII) is a strong purple colour which disappears at the endpoint, so the titration is self-indicating
• This reaction is often used for the analysis of iron for example in iron tablets (health supplements)

Iodine-thiosulfate titrations

• A redox reaction occurs between iodine and thiosulfate ions:

2S₂O₃^2– (aq) + I₂ (aq) → 2I^–(aq) + S₄O₆^2– (aq)

• The light brown/yellow colour of the iodine turns paler as it is converted to colourless iodide ions
• When the solution is a straw colour, starch is added to clarify the endpoint
• The solution turns blue/black until all the iodine reacts, at which point the colour disappears.
• This titration can be used to determine the concentration of an oxidising agent, which oxidises iodide ions to iodine molecules
• The amount of iodine is determined from titration against a known quantity of sodium thiosulfate solution
• This reaction can be used for the analysis of chlorine in bleach