Redox Titrations (AQA A-Level Chemistry): Revision Notes

8.2.4 Redox Titrations

Aim:

To carry out redox titrations to determine the concentration of an unknown solution, such as:

• The mass of iron in an iron tablet
• The percentage of iron in steel
• The molar mass ($Mᵣ$) of hydrated ammonium iron(II) sulphate
• The molar mass ($Mᵣ$) of ethanedioic acid
• The concentration of hydrogen peroxide ($H₂O₂$) in hair bleach

What are redox titrations?

• Redox titrations are a type of titration where the chemical reaction involves the transfer of electrons between two species.
• One solution (the titrant) contains a known concentration of an oxidising or reducing agent, and the other solution contains the substance being analysed.
• Redox titrations often involve transition metal ions that change oxidation states during the reaction, leading to observable colour changes.

Key Concepts:

Transition Metals in Redox Reactions:

Transition metals are elements found in the d-block of the periodic table. They have partially filled d-orbitals and can participate in redox reactions by either gaining or losing electrons. In redox titrations, transition metals like iron ($Fe²⁺$/$Fe³⁺$) or copper ($Cu²⁺$/$Cu⁺$) commonly act as reducing or oxidising agents.

Redox Reactions:

A redox reaction involves two half-reactions: one for oxidation (loss of electrons) and one for reduction (gain of electrons). For example:

• Oxidation: $Fe²⁺ → Fe³⁺ + e⁻$
• Reduction: $MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O$

End Point and Colour Changes:

The end point in a redox titration is usually detected by a distinct colour change as the transition metal ion changes oxidation state. For example, in titrations involving permanganate ($MnO₄⁻$), the solution changes from purple to colourless when$MnO₄⁻$is reduced to $Mn²⁺$

Materials and Equipment:

• 1.0 mol dm$⁻³$ potassium permanganate ($KMnO₄$) solution (or another titrant)
• Iron (II) sulphate or other analyte
• 1.0 mol dm$⁻³$ sulfuric acid ($H₂SO₄$)
• Volumetric pipette and filler
• Burette and clamp stand
• Conical flask
• White tile (to improve visibility of the endpoint)
• Distilled water
• Measuring cylinder
• Indicator (optional, depends on the redox reaction)

Method:

Step 1: Prepare the Analyte Solution

• Using a volumetric pipette, measure 25 cm³ of the solution with the unknown concentration (e.g., iron tablet solution or hair bleach).
• Transfer it into a conical flask.

Step 2: Add Excess Sulfuric Acid

• Add 20 cm³ of dilute sulfuric acid ($H₂SO₄$) to the conical flask using a measuring cylinder.
• This ensures that the reaction proceeds fully in acidic conditions.
• Sulfuric acid is added in excess to maintain the correct conditions for the redox reaction.

Step 3: Set up the Burette

• Fill the burette with the titrant (e.g., potassium permanganate solution of known concentration) and record the initial volume in the burette.

Step 4: Perform the Titration

• Slowly add the titrant to the solution in the conical flask while continuously swirling.
• As the titrant is added, it reacts with the substance in the flask, and you should observe a gradual colour change.
• As you approach the endpoint, slow down the addition of titrant to dropwise.
• The endpoint is reached when a permanent colour change is observed in the solution.

Step 5: Record the Titre

• Record the final volume of the titrant in the burette when the colour change is stable.
• This is your titre volume.

Step 6: Repeat the Titration

• Repeat the titration until you have at least two concordant results (readings that are within 0.10 cm³ of each other).

Calculations:

1. Derive the Half-Equations:

• Write the half-equations for the oxidation and reduction processes occurring in the titration. For example, in the reaction between iron(II) and permanganate ions:
• Oxidation: $Fe²⁺ → Fe³⁺ + e⁻$
• Reduction: $MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O$

2. Use Stoichiometry:

• Use the balanced overall equation and the stoichiometric ratios from the half-equations to determine the moles of the unknown substance.

3. Calculate Concentration:

• Use the formula:

$$C_1 V_1 = C_2 V_2$$

Where:

• $C₁$ = concentration of the titrant
• $V₁$ = volume of the titrant used (titre)
• $C₂$ = concentration of the unknown solution (to be calculated)
• $V₂$ = volume of the analyte (the solution being titrated)

Common Errors and Troubleshooting:

• Indicator or Colour Change: Some redox titrations don't require an indicator because the substances themselves exhibit a colour change at the endpoint (e.g., purple $MnO₄⁻$becoming colourless $Mn²⁺$).
• Concentration of Acid: Ensure the correct acid concentration is used (e.g., sulfuric acid for iron(II) and permanganate titrations) to avoid side reactions.
• Contamination: Clean all apparatus thoroughly between titrations to prevent contamination and false results.

Conclusion:

Redox titrations are an essential method in analytical chemistry to determine unknown concentrations. By following these steps and calculating results from the half-equations, you can apply redox titrations to various contexts such as determining the composition of iron tablets, steel, or even hair bleach solutions.