Redox Titrations

• Half-Equations: The overall reaction is derived by combining the reduction half-equation and the oxidation half-equation, ensuring that the number of electrons lost equals the number of electrons gained.

• Manganate(VII) Reduction: In acidic conditions, the purple $MnO_4^-$ ion is reduced to the nearly colorless $Mn^{2+}$ ion according to: $MnO_4^-(aq) + 8H^+(aq) + 5e^- \rightarrow Mn^{2+}(aq) + 4H_2O(l)$

• Dichromate(VI) Reduction: The orange $Cr_2O_7^{2-}$ ion is reduced to green $Cr^{3+}$ ions: $Cr_2O_7^{2-}(aq) + 14H^+(aq) + 6e^- \rightarrow 2Cr^{3+}(aq) + 7H_2O(l)$

• Stoichiometric Ratios: The molar ratio between the oxidant and reductant (e.g., 1:5 for $MnO_4^-$ and $Fe^{2+}$) is critical for calculating the unknown concentration from the titer volume.

Manganate(VII) Titration Procedure

• Acidification: The reaction must occur in the presence of excess dilute sulfuric acid ($H_2SO_4$) to provide the $H^+$ ions necessary for the reduction of $MnO_4^-$.

• Burette Preparation: Potassium manganate(VII) is typically placed in the burette; because of its deep purple color, the upper meniscus is often read instead of the lower one.

• Endpoint Identification: The endpoint is reached at the first appearance of a permanent pale pink color, indicating a slight excess of $MnO_4^-$ ions.

Dichromate(VI) Titration Procedure

• Indicator Requirement: Unlike manganate, the color change from orange to green is difficult to detect visually, so an indicator like diphenylaminesulfonate is used, which turns from colorless to purple at the endpoint.

• Comparison of Common Oxidants:

• Acid Selection: Dilute sulfuric acid is the only suitable acid; hydrochloric acid ($HCl$) is avoided because $MnO_4^-$ can oxidize $Cl^-$ to $Cl_2$, while nitric acid ($HNO_3$) is an oxidizing agent itself that would interfere with the reaction.