Redox Titration (Edexcel A Level Chemistry)

Redox Titrations

• In a titration, the concentration of a solution is determined by titrating with a solution of known concentration.
• In redox titrations, an oxidizing 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 oxidation state
• There are two common redox titrations you should know about manganate(VII) titrations and iodine-thiosulfate titrations

• Iron has variable oxidation states and this characteristic can be exploited in redox titrations of Fe²⁺ with MnO₄^- or Cr₂O₇^2-
• Both the manganate(VII) ion and dichromate(VI) ions can be used as strong oxidising agents to find the amount of iron(II) ions in solution
• You need to be able to perform these separate redox titrations with Fe²⁺ and understand the colour changes in each example

Potassium manganate(VII) titrations

• In these redox titrations the manganate(VII) is the oxidising agent and is reduced to Mn²⁺(aq)
• The iron is the reducing agent and is oxidised to Fe²⁺(aq) and the reaction mixture must be acidified, to excess acid is added to the iron(II) ions before the reaction begins
• The choice of acid is important, as it must not react with the manganate(VII) ions, so the acid normally used is dilute sulfuric acid
  • As it does not oxidise under these conditions and does not react with the manganate(VII) ions

• You could be asked why other acids are not suitable for this redox titration in the exam so make sure you understand the suitability of dilute sulfuric acid

Table Explaining Why Other Acids are not Suitable for the Redox Titration

Indicator and end point

• Potassium permanganate acts as its own indicator, as the purple potassium permanganate solution is added to the titration flask from the burette and reacts rapidly with the Fe²⁺(aq)
  • The burette used in this practical should be one with white numbering not black, as you would struggle to read the values for your titres against the purple colour of the potassium permanganate if black numbering was used

• The manganese(II) ions, Mn²⁺(aq), have a very pale pink colour but they are present in such a low concentration that the solution looks colourless
• As soon as all of the iron(II), Fe²⁺(aq), ions have reacted with the added manganate(VII) ions, Mn⁷⁺(aq), a pale pink tinge appears in the flask due to an excess of manganate(VII) ions, Mn⁷⁺(aq)

Redox titration colour change for potassium permanganate and iron(II) ions

Answers:

Balance the electrons:

MnO₄^- (aq) + 5e^- + 8H⁺ (aq) → Mn²⁺ (aq) + 4H₂O (l)

5Fe²⁺ (aq) → 5Fe³⁺ (aq) + 5e^-

Add the two half equations:

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

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 end point
• 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 oxidizing agent, which oxidizes iodide ions to iodine molecules
• The amount of iodine is determined from titration against a known quantity of sodium thiosulfate solution

Redox titration calculations

• Common redox titration calculations can include:
  • Manganate(VII) titrations can be used to determine:
    • The percentage purity of iron supplements

Percentage purity = 

• • • The formula of a sample of hydrated ethanedioic acid
  • Iodine / thiosulfate reactions can be used to determine:
    • The amount of chlorate(I), ClO^-, in bleach
    • The amount of copper(II) ions, Cu²⁺, in copper(II) compounds
    • The copper content of alloys

Answer:

• • MnO₄^- (aq) + 8H⁺ (aq) + 5Fe²⁺ → Mn²⁺ (aq) + 5Fe³⁺ (aq) + 4H₂O (l)
  • 1 : 5 ratio of MnO₄^- : Fe²⁺
  • Number of moles of MnO₄^- (aq)5.13 x 10^-4 moles 
  • Moles of iron(II) = 5 x 5.13 x 10^-4 = 2.565 x 10^-3 moles
  • Mass of iron(II) = 56.0 x 2.565 x 10^-3 = 0.14364 g
  • Percentage by mass = 15.0%

Answer:

• • One mole of ClO^- (aq) produces one mole of I₂ (aq) which reacts with two moles of 2S₂O₃^2- (aq)
    • Therefore, 1 : 2 ratio of ClO^- (aq) : S₂O₃^2- (aq)
  • Number of moles of S₂O₃^2- (aq) 1.26 x 10^-3 moles 
  • Number of moles of I₂ (aq) and ClO^- (aq) in 25.0 cm³ 6.30 x 10^-4 moles 
  • Number of moles of ClO^- (aq) in 250.0 cm³ = 6.30 x 10^-4 x 10 = 6.30 x 10^-3 moles 
  • The 250.0 cm³ was prepared from 10.0 cm³ bleach
    • 10 cm³ bleach = 6.30 x 10^-3 moles of ClO^- ions
    • 1.0 dm³ bleach = 0.630 moles of ClO^- ions
    • Therefore, the concentration of ClO^- ions in the bleach is 0.630 mol dm^-3