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

6.2.6 Redox Titrations

Redox Potential and Influencing Factors

The redox potential of a transition metal ion, which measures its tendency to gain electrons (reduce), is affected by:

• pH: Redox potential varies with pH; a more acidic environment often favours reduction.
• Ligand Effects: The type of ligand bonded to the metal ion can alter the redox potential by stabilising certain oxidation states.

Practical Applications

Tollens' Reagent Test for Aldehydes

In this example, we'll use Tollens' reagent, $\text{Ag}(\text{NH}_3)_2 ]^{+}$, to test for the presence of aldehydes. This reagent, also known as the silver mirror test, takes advantage of silver's variable oxidation states to differentiate between aldehydes and ketones.

Aim

To determine whether a given organic compound is an aldehyde or a ketone by observing the reaction with Tollens' reagent.

Method

1. Prepare Tollens' Reagent:

• In a clean test tube, add a few drops of silver nitrate solution ($\text{AgNO}_3$)
• Add a few drops of dilute sodium hydroxide ($\text{NaOH}$) solution to form a brown precipitate of silver oxide ($\text{Ag}_2\text{O}$).
• Add dilute ammonia solution dropwise to the silver oxide precipitate until it completely dissolves, forming a clear solution of $[ \text{Ag}(\text{NH}_3)_2 ]^{+}$. This is Tollens' reagent.

2. Test the Sample:

• Add 1-2 mL of the Tollens' reagent to a clean test tube containing a few drops of the organic sample (suspected aldehyde or ketone).
• Gently warm the test tube in a hot water bath (about 50-60°C) for a few minutes.
• Observe any changes in the appearance of the solution and test tube walls.

Results

• Positive Test for Aldehyde: A silver mirror forms on the walls of the test tube as $[ \text{Ag}(\text{NH}_3)_2 ]^{+}$ is reduced to metallic silver. The solution may also become colourless.
• Negative Test for Ketone: No reaction occurs, and the solution remains clear without any mirror formation.

Explanation

Reaction Mechanism:

Aldehydes are reducing agents and can donate electrons to $[ \text{Ag}(\text{NH}_3)_2 ]^{+}$, which is reduced to metallic silver ($\text{Ag}$).

The half-equation for the reduction is:

$$[ \text{Ag}(\text{NH}_3)_2 ]^{+} + e^- \rightarrow \text{Ag} (s) + 2 \text{NH}_3$$

The aldehyde itself is oxidised to a carboxylate ion in the process, usually resulting in a colourless solution.

Why Ketones Don't React:

• Ketones lack the hydrogen attached to the carbonyl group, which is essential for the oxidation reaction.
• As a result, they cannot reduce $[ \text{Ag}(\text{NH}_3)_2 ]^{+}$, so no reaction is observed. This test provides a quick and visually striking way to identify aldehydes, making it particularly useful in organic chemistry labs for differentiating between these functional groups.

Redox Titrations with $\text{MnO}_4^-$

Potassium permanganate ($\text{KMnO}_4$) is commonly used in redox titrations due to its strong oxidising ability and distinct colour change:

Determining the Mass of Iron in Iron Tablets

Aim: To determine the mass of iron in an iron tablet by titrating $\text{Fe}^{2+}$ ions with potassium permanganate ($\text{MnO}_4^-$).

Method:

3. Dissolve a crushed iron tablet in dilute sulfuric acid in a conical flask, ensuring all $\text{Fe}^{2+}$ ions are fully dissolved.
4. Using a burette, add a standardised $\text{KMnO}_4$ solution to the flask.
5. Titrate until a faint pink colour persists, indicating the endpoint.
6. Record the volume of $\text{KMnO}_4$ used. Results:

Use the titration data to calculate the moles of ($\text{Fe}^{2+}$) reacted:

$$5\text{Fe}^{2+} + \text{MnO}_4^- + 8\text{H}^+ \rightarrow 5\text{Fe}^{3+} + \text{Mn}^{2+} + 4\text{H}_2\text{O}$$

Calculate the mass of iron by multiplying moles by the atomic mass of iron (55.85 g/mol).

Explanation:

Potassium permanganate ($\text{MnO}_4^-$) is a strong oxidising agent, which oxidises $\text{Fe}^{2+}$ to $\text{Fe}^{3+}$

The pink colour marks the endpoint, allowing you to calculate the mass of iron in the tablet.

Determining the Percentage of Iron in Steel

Aim: To determine the percentage of iron in a steel sample by titrating dissolved iron with potassium permanganate.

Method:

7. Dissolve a measured steel sample in sulfuric acid to produce $\text{Fe}^{2+}$ ions.
8. Transfer the solution to a conical flask.
9. Titrate with standardised $\text{KMnO}_4$ solution until a persistent pink colour appears.
10. Record the volume of $\text{KMnO}_4$ used. Results:

• Calculate the moles of $\text{Fe}^{2+}$ based on the reaction stoicheiometry.
• Use the moles of $\text{Fe}^{2+}$ to find the mass of iron and calculate the percentage of iron in the steel sample:

$$\text{Percentage of iron} = \frac{\text{mass of Fe}}{\text{mass of steel sample}} \times 100$$

Explanation:

This titration allows you to determine the iron content by oxidising $\text{Fe}^{2+}$ to $\text{Fe}^{3+}$. The percentage of iron in steel can then be calculated, which is useful in the quality control of steel products.

Calculating the $M_r$ of Hydrated Ammonium Iron(II) Sulphate and Ethanedioic Acid ($\text{C}_2\text{O}_4^{2-}$)

Aim: To determine the molar mass ($M_r$) of hydrated ammonium iron(II) sulphate and ethanedioic acid by titration.

Method:

11. Dissolve a known mass of hydrated ammonium iron(II) sulphate or ethanedioic acid in sulfuric acid to ensure full dissolution.
12. Titrate with standardised $\text{KMnO}_4$ solution, adding until a faint pink endpoint is reached.
13. Record the volume of $\text{KMnO}_4$ used. Results:

Calculate the moles of the substance based on the titration reaction:

For ethanedioic acid:

$$\text{C}_2\text{O}_4^{2-} + \text{MnO}_4^- + 8\text{H}^+ \rightarrow 2\text{CO}_2 + \text{Mn}^{2+} + 4\text{H}_2\text{O}$$

Calculate the $M_r$ by dividing the mass of the compound by the moles calculated from the titration data.

Explanation:

This method allows you to calculate the molar mass of the compounds by measuring how much permanganate is required to react with a known mass of the compound. This is useful for verifying the purity of a sample.

Finding the Concentration of Hydrogen Peroxide ($\text{H}_2\text{O}_2$) in Hair Bleach

Aim: To determine the concentration of hydrogen peroxide in a hair bleach solution using titration.

Method:

14. Add a measured volume of the $\text{H}_2\text{O}_2$ solution to a conical flask and add sulfuric acid.
15. Titrate with standardised $\text{KMnO}_4$ solution until a persistent pink colour indicates the endpoint.
16. Record the volume of $\text{KMnO}_4$ used. Results:

Calculate the concentration of $\text{H}_2\text{O}_2$ based on the reaction:

$$5\text{H}_2\text{O}_2 + 2\text{MnO}_4^- + 6\text{H}^+ \rightarrow 5\text{O}_2 + 2\text{Mn}^{2+} + 8\text{H}_2\text{O}$$

Use the volume and concentration of $\text{KMnO}_4$ to find the moles of $\text{H}_2\text{O}_2$ and calculate its concentration in the bleach.

Explanation:

In this titration, $\text{KMnO}_4$ oxidises $\text{H}_2\text{O}_2$ to oxygen gas. The titration results allow you to calculate the concentration of hydrogen peroxide, which is important in formulations like hair bleach.