Physical Chemistry Practicals (A Level only) (AQA A-Level Chemistry): Revision Notes

8.1.10 Making Buffers

Objective:

To prepare and test buffer solutions using both direct and indirect methods. The aim is to understand how buffers resist pH changes upon the addition of small amounts of acid or base and to explore how the Henderson-Hasselbalch equation applies to buffer calculations.

Theory Overview:

Buffers are solutions that resist changes in pH when small amounts of acid or base are added. They are made from a weak acid and its conjugate base or a weak base and its conjugate acid. In this practical, ethanoic acid (a weak acid) and sodium ethanoate (its conjugate base) will be used to create an acidic buffer solution.

The Henderson-Hasselbalch equation is central to buffer calculations:

$$pH = pK_a + \log \left( \frac{[\text{salt}]}{[\text{acid}]} \right)$$

This equation allows for the calculation of pH based on the ratio of salt to acid in the buffer.

Apparatus & Chemicals:

• 1.0 mol dm$⁻³$ ethanoic acid ($CH₃COOH$)
• Sodium ethanoate ($CH₃COONa$)
• 1.0 mol dm$⁻³$ sodium hydroxide ($NaOH$)
• 1.0 mol dm$⁻³$ hydrochloric acid ($HCl$)
• pH metre or calibrated pH probe
• Volumetric pipette (50 cm³)
• Volumetric flask (250 cm³)
• Weighing balance
• Beakers
• Funnel
• Stirring rod
• Distilled water

Key Procedure Steps:

Direct Method:

1. Preparation of Ethanoic Acid/Ethanoate Buffer:

• Measure 50 cm³ of 0.1 mol dm⁻³ ethanoic acid solution and pour it into a clean beaker.
• Weigh 2.05 g of sodium ethanoate ($CH₃COONa$), which corresponds to 0.025 mol, and add it to the beaker.
• Stir the solution until the sodium ethanoate dissolves completely.

2. Transfer to Volumetric Flask:

• Using a funnel, transfer the solution into a 250 cm³ volumetric flask.
• Rinse the beaker and funnel with distilled water, and pour the rinsings into the volumetric flask to ensure no residue is left behind.
• Fill the volumetric flask to the 250 cm³ mark with more 0.1 mol dm⁻³ ethanoic acid solution to ensure accurate concentrations.

3. Measure the pH:

• Calibrate the pH meter and then use it to measure the pH of the buffer solution.
• The expected pH for a 1:1 ratio of ethanoic acid to sodium ethanoate is approximately 4.75.

4. Testing the Buffer:

• To test the buffer's capacity, add 1 cm³ of 1.0 mol dm⁻³ HCl to 100 cm³ of the buffer solution and measure the pH.
• The pH should change very little.
• Repeat by adding 1 cm³ of 1.0 mol dm⁻³ NaOH to the buffer solution and measuring the pH again.
• A minimal change in pH is expected.

Indirect Method:

1. Partial Neutralisation of Ethanoic Acid:

• Add 25 cm³ of 0.1 mol dm⁻³ sodium hydroxide ($NaOH$) to 50 cm³ of 0.1 mol dm⁻³ ethanoic acid in a beaker.
• This neutralises half of the ethanoic acid, forming an ethanoic acid/sodium ethanoate buffer in the process.

2. Make Up to 250 cm³:

• Transfer the solution to a 250 cm³ volumetric flask and make up the volume to 250 cm³ using distilled water.
• Measure and record the pH of this buffer using the pH meter.

Analysis:

Using the Henderson-Hasselbalch Equation:

• The buffer contains 0.025 mol of sodium ethanoate and an equal concentration of ethanoic acid (since 2.05 g of sodium ethanoate were used).
• The pKa of ethanoic acid is 4.75. When the concentrations of the acid and salt are equal, the equation simplifies to:

$$pH = pK_a + \log \left( \frac{[\text{salt}]}{[\text{acid}]} \right)$$

Since the ratio of salt to acid is 1:1, $\log 1 = 0$, so:

$$pH = pK_a = 4.75$$

Adjusting pH to 4.25:

If a pH of 4.25 is required, use the Henderson-Hasselbalch equation to adjust the ratio of salt to acid:

$$4.25 = 4.75 + \log \left( \frac{[\text{salt}]}{[\text{acid}]} \right)$$

Solving this gives:

$$\log \left( \frac{[\text{salt}]}{[\text{acid}]} \right) = -0.50$$ $$\frac{[\text{salt}]}{[\text{acid}]} = 0.50$$

This means the ratio of salt to acid should be 0.5:1, i.e., half the amount of sodium ethanoate should be used. Therefore, to prepare a buffer with pH 4.25, 1.025 g of sodium ethanoate should be dissolved in the same volume of ethanoic acid solution.

Conclusion:

By preparing a buffer solution using both the direct and indirect methods, students gain insight into how buffers are made and tested. They learn how to apply the Henderson-Hasselbalch equation to calculate pH based on the ratio of weak acid to its conjugate base. Buffers play a crucial role in maintaining pH stability, which is essential in biological systems and industrial applications.

Practical Tips for Success:

• Accurate Calibration: Ensure the pH metre is well-calibrated before use, as precise pH readings are crucial for buffer experiments.
• Thorough Mixing: When preparing buffer solutions, ensure that all components are completely dissolved and the solution is well-mixed before taking pH measurements.
• Minimising Errors: When transferring solutions to volumetric flasks, make sure to include rinsings to avoid loss of material and maintain concentration accuracy.