Aspirin (acetylsalicylic acid) undergoes hydrolysis in aqueous solution to produce salicylic acid and ethanoic acid. The rate of this hydrolysis is strongly dependent on pH. In this investigation, you will use a colorimeter to track the appearance of salicylate ions (which form a purple complex with iron(III) ions) and measure how pH affects the rate of hydrolysis. This produces a fully continuous dataset in both IV and DV, making it an outstanding IB Chemistry IA.

This practical is suitable for IB Diploma Chemistry HL.

Background Theory

Aspirin hydrolyses according to the equation:

CH₃COOC₆H₄COOH + H₂O → C₆H₄(OH)COOH + CH₃COOH

This reaction is catalysed by both H⁺ (acid catalysis) and OH⁻ (base catalysis), meaning it proceeds faster at both low and high pH values, with a minimum rate near neutral pH. This U-shaped relationship between pH and rate is a distinctive and analytically rich feature for an IA investigation.

The salicylate product reacts with iron(III) chloride (FeCl₃) to form a purple-violet complex with strong absorbance around 530 nm. By measuring absorbance over time, the rate of hydrolysis at each pH can be determined as the gradient of an absorbance vs time graph (initial rate method).

Variables

• Independent variable (IV): pH of the buffer solution (continuous) — e.g. pH 2, 4, 6, 7, 8, 10, 12
• Dependent variable (DV): Initial rate of hydrolysis (continuous) — measured as Δabsorbance / Δtime (min⁻¹) using a colorimeter
• Controlled variables (CV): Concentration and volume of aspirin solution, concentration of FeCl₃ reagent, temperature (use a water bath), same colorimeter and wavelength throughout

Equipment

• Aspirin tablets or acetylsalicylic acid powder dissolved in ethanol to make a stock solution
• Buffer solutions at pH 2, 4, 6, 7, 8, 10, 12 (prepared or commercial)
• Iron(III) chloride solution, FeCl₃ (0.10 mol dm⁻³)
• Colorimeter with green filter (~530 nm) and cuvettes
• Water bath at 25 °C
• Stopwatch
• Pipettes and volumetric flasks
• pH meter or pH indicator paper to verify buffer pH

Method

1. Prepare a stock aspirin solution by dissolving a known mass (e.g. 0.180 g, approximately 1.0 mmol) in a small volume of ethanol, then making up to 100 cm³ with distilled water. This gives approximately 0.010 mol dm⁻³.
2. For each pH, pipette 5 cm³ of aspirin stock solution into a boiling tube. Add 5 cm³ of the appropriate buffer solution. Allow to equilibrate at 25 °C in a water bath for 5 minutes.
3. At time zero, add 1 cm³ of FeCl₃ solution and mix quickly. Transfer immediately to a cuvette in the colorimeter.
4. Record absorbance every 30 seconds for 10 minutes.
5. Plot absorbance vs time for each pH. Calculate the initial rate as the gradient of the linear portion of the curve (first 2–3 minutes).
6. Repeat each pH at least three times and calculate mean initial rates.

Results Table

pH	Initial Rate 1 (ΔAbs/min)	Initial Rate 2 (ΔAbs/min)	Initial Rate 3 (ΔAbs/min)	Mean Initial Rate (ΔAbs/min)
2
4
6
7
8
10
12

Analysis

1. Plot mean initial rate (y-axis) against pH (x-axis). You should observe a U-shaped curve with a minimum near neutral pH, reflecting acid and base catalysis at the extremes.

2. Identify the pH of minimum hydrolysis rate. This is where neither H⁺ nor OH⁻ catalysis dominates.

3. Compare the rates at pH 2 and pH 12. Is acid or base catalysis more effective? What does this suggest about the mechanism?

4. Consider plotting rate vs [H⁺] and rate vs [OH⁻] separately on log scales to linearise the relationship and test for first-order dependence on each catalyst.

Discussion Points

• Why does the rate of hydrolysis increase at both low and high pH? Explain in terms of acid and base catalysis mechanisms.
• Why is the FeCl₃ reagent added after incubation rather than at the start?
• Why is temperature controlled so carefully in this experiment?
• What does the U-shaped rate-pH curve tell you about the stability of aspirin in the human body (stomach pH ~2, blood pH ~7.4)?
• How does this relate to the shelf life and storage conditions of aspirin tablets?

• Practical Chemistry