16B2.Redox equlibria

• Potassium manganate (VII) is a powerful oxidising agent. It is used for the quantitative estimation of many reducing agents, especially compounds of iron(II) and for ethanedioic acid and its salts.
• Reduction of manganate(VII) ion in acidic solution


  MnO4(aq) + 8H+(aq) + 5e → Mn2+(aq) + 4H2O(l)
• In alkaline solution, potassium manganate(VII) yields manganese(IV) oxide as a brown precipitate, and this interferes with the end point colour. For this reason, potassium manganate(VII) is almost always used in solutions that are sufficiently acidic to prevent the formation of manganese(IV) oxide.
• As the titration proceeds, manganese(II) ions, Mn2+, accumulate but, at the dilution used, they give a colourless solution. As soon as the potassium manganate(VII) is in excess, the solution becomes pink. It therefore acts as its own indicator - the end point is the first permanent pink colour.
• Titration of potassium manganate(VII) with iron(II) ions

  MnO4(aq) + 8H+(aq) + 5Fe2+(aq) → Mn2+(aq) + 4H2O(l) + 5Fe3+(aq)
• Colours of ions in solutions

  • Fe2+(aq) pale green
  • Fe3+(aq) yellow
  • Mn2+(aq) pale pink
• Acidified potassium dichromate cannot be used as an oxidising agent for quantitative redox titrations because the colour change from orange to green has no clear end point.
• Titration of potassium manganate(VII) with ethanedioic acid

  2MnO4(aq) + 6H+(aq) + 5H2C2O4(aq) → 2Mn2+(aq) + 8H2O(l) + 10CO2(g)
• The reaction between potassium manganate(VII) and ethanedioic acid is very slow, so the ethanedioic acid is heated to around 60°C before starting the titration. The Mn2+ ions produced act as a catalyst and the reaction speeds up.
• Potassium manganate(VII) can also be used to estimate the amount of ethanedioate ion present in a solution.
• Autocatalysis


  A reaction in which a product acts as a catalyst
• The dilute sulfuric acid is boiled to remove dissolved oxygen that might otherwise oxidise some of the iron(II) ions to iron(III).
• In an exam, you may be asked to calculate the measurement uncertainty for a titration.
• The formula of hydrated iron(II) sulfate is FeSO4.7H2O
• Titration of iodine with sodium thiosulfate


  2S2O3(aq) + I2(aq) → S4O6(aq) + 2I-(aq)
• Starch solution indicator


  Produces a deep blue-black colour with free iodine, which disappears as sodium thiosulfate is added to react with the iodine
• Wear eye protection and avoid skin contact with the solutions
• Make sure the top of the burette is below eye level when you fill it
• FeSO¿.7H₂O


  Formula of hydrated iron(II) sulfate
• 176 16B.2 REDOX TITRATIONS
• Thiosulfate ions reduce iodine to iodide ions

  2S2O3(aq) + I2(aq) → S4O6(aq) + 2I-(aq)
• The reaction can be used for the direct estimation of iodine, or for the estimation of a substance that can take part in a reaction that produces iodine
• Starch solution


  Indicator used in titrating iodine with sodium thiosulfate
• Procedure for titrating iodine with sodium thiosulfate

  1. Add sodium thiosulfate solution until original brown colour changes to pale yellow

  2. Add starch solution to produce blue colouration

  3. Continue adding sodium thiosulfate until blue-black solution turns colourless
• If starch is added too early, it adsorbs some of the iodine and reduces the accuracy of the titration
• If starch is not added, it would be very difficult to detect the end point
• Procedure for redox titration of copper(II) sulfate

  1. Weigh sample of hydrated copper(II) sulfate crystals

  2. Transfer to volumetric flask and make up to 250 cm³

  3. Pipette 25.0 cm³ of solution, add potassium iodide

  4. Titrate iodine produced with sodium thiosulfate, adding starch near end point
• Wear eye protection and avoid skin contact with the solids and liquids, especially the hot acid
• Boil the acid gently for a short time in a larger beaker to avoid it boiling over and to cool it faster
• TOPIC 16 CHECKPOINT 16B.2 REDOX TITRATIONS
• SKILLS PROBLEM-SOLVING CREATIVITY