Redox Titration Calculations

Potassium Manganate(VII) Titrations

• Self-Indicating: The $MnO_4^-$ ion is deep purple, while the reduced $Mn^{2+}$ ion is virtually colorless; the first permanent pale pink tinge indicates the endpoint.
• Acidification: Dilute sulfuric acid ($H_2SO_4$) must be added to provide the $H^+$ ions required for the reduction: $MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O$

Iodine-Thiosulfate Titrations

• Indirect Method: An oxidising agent reacts with excess iodide ($I^-$) to produce iodine ($I_2$), which is then titrated against sodium thiosulfate ($S_2O_3^{2-}$).
• Starch Indicator: Added near the endpoint (when the solution is straw-colored) to turn the solution blue-black; the endpoint is reached when the blue-black color disappears.

• Check the Ratio: Always derive the full equation from half-equations; for example, $MnO_4^-$ reacting with $Fe^{2+}$ has a $1:5$ ratio, but reacting with $C_2O_4^{2-}$ has a $2:5$ ratio.

• Scaling Factors: If a sample was diluted (e.g., $25.0$ cm$^3$ taken from a $250$ cm$^3$ volumetric flask), remember to multiply the moles by $10$ to find the amount in the original sample.

• Unit Consistency: Ensure all volumes are converted to dm$^3$ (divide cm$^3$ by $1000$) before using the $n = C \times V$ formula.

• Sanity Check: Percentage purity should never exceed $100\%$; if it does, re-check your molar ratios or scaling factors.

• Incorrect Acid Choice: Using $HCl$ in manganate titrations is a common error; the $Cl^-$ ions would be oxidised to $Cl_2$ by the $MnO_4^-$, leading to an artificially high titre.

• Starch Timing: Adding starch too early in an iodine titration results in the iodine becoming trapped in the starch complex, making the endpoint difficult to see and inaccurate.

• Ignoring Hydration: When calculating the mass of a solid (like hydrated iron(II) sulfate), ensure the molar mass includes the water of crystallisation ($xH_2O$).