Redox Titrations Study Notes

Introduction to Redox Titrations

• Lesson suitable for A-level students.

• Objectives:

  • Write half equations for reduction of common oxidizing agents (e.g., MnO4⁻, Cr2O7²⁻).

  • Combine half equations for chemical calculations.

  • Explain the procedure of a redox titration and how to determine its endpoint.

Titration Principles

• Titrations can determine concentrations of acids, bases, oxidizing and reducing agents.

• Essential calculations involve finding concentration:

  • Concentration (mol/dm³) = $\frac{mass (g)}{molar \, mass (g/mol)} \times \frac{1000}{volume (cm³)}$

  • Moles = $\frac{concentration (mol/dm³) \times volume (cm³)}{1000}$

Redox Titrations

• Involve reduction and oxidation reactions.

• Common oxidizing agents: Manganate VII ion (MnO4⁻), Dichromate VI ion (Cr2O7²⁻).

• Oxidizing agents gain electrons (reduction).

• Reducing agents cause other substances to gain electrons.

Manganate VII Half Equations

• MnO4⁻ reduced to Mn²⁺:

  • $MnO<em>4^{-} + 8H^{+} + 5e^{-} \rightarrow Mn^{2+} + 4H</em>2O$

• Color change: purple (MnO4⁻) to colorless (Mn²⁺).

• Practice problem: Write half equation for BrO3⁻ reduction (Br2).

Combining Half Equations

• Ensure equal number of electrons in both half equations for cancellation.

• Example combining Fe²⁺ oxidation with MnO4⁻ reduction results in:

  • $Cr<em>2O</em>7^{2-} + 14H^{+} + 6Fe^{2+} \rightarrow 2Cr^{3+} + 7H_2O + 6Fe^{3+}$

Titration Procedure

• Use acidified potassium manganate VII solution to determine Fe²⁺ concentration.

• Process:

  • Measure Fe²⁺ sample in conical flask.

  • Add dilute sulfuric acid.

  • Titrate with KMnO4; observe color change.

• Endpoint recognition: color turns permanent pink when no Fe²⁺ remains.

Practice Questions and Review

• Electron configuration for Mn²⁺ and identification of endpoint.

• Oxidation state change from +7 (MnO4⁻) to +2 (Mn²⁺).

• Concentration setup for titration calculations, applying stoichiometric ratios effectively to deduce required values.