The iodine clock reaction is a classic chemistry experiment that produces a dramatic, sudden colour change from colourless to deep blue-black. In this investigation, you will explore how changing the temperature affects the rate of this reaction and use your results to determine the activation energy of the reaction.
This practical is suitable for IB Diploma Chemistry HL students and makes an excellent Internal Assessment (IA) topic due to the precise, quantitative data it generates.
Background Theory
The iodine clock reaction involves the oxidation of iodide ions (I⁻) by hydrogen peroxide (H₂O₂) in acidic solution. Sodium thiosulfate (Na₂S₂O₃) is added as a “clock” — it reacts instantly with any iodine produced, keeping the solution colourless. Once all the thiosulfate is consumed, free iodine accumulates and reacts with starch to give the characteristic blue-black colour.
The overall reaction being timed is:
H₂O₂ + 2I⁻ + 2H⁺ → I₂ + 2H₂O
According to the Arrhenius equation, increasing temperature increases the rate constant k, as more molecules have sufficient energy to overcome the activation energy barrier. By measuring reaction time at different temperatures, you can calculate the rate (1/t) and use an Arrhenius plot (ln k vs 1/T) to determine the activation energy Eₐ.
Variables
- Independent variable (IV): Temperature of the reaction mixture (°C / K)
- Dependent variable (DV): Time for the colour change to occur (seconds)
- Controlled variables (CV): Concentration of all reagents, volume of solutions, same starch solution, same water bath for temperature control
Equipment
- Hydrogen peroxide solution, H₂O₂ (0.050 mol dm⁻³)
- Potassium iodide solution, KI (0.010 mol dm⁻³)
- Sodium thiosulfate solution, Na₂S₂O₃ (0.001 mol dm⁻³)
- Sulfuric acid, H₂SO₄ (0.50 mol dm⁻³)
- Starch solution (1%)
- Water baths set at 10, 20, 30, 40, and 50 °C
- Thermometer (±0.5 °C)
- Stopwatch (±0.01 s)
- 5 × 25 cm³ measuring cylinders or burettes
- 250 cm³ beakers
- White tile
Safety
⚠️ Hydrogen peroxide is an oxidising agent and irritant — wear eye protection and gloves at all times. Sulfuric acid is corrosive. Dispose of all solutions into the appropriate waste disposal bottles provided.
Method
- Prepare five water baths at approximately 10, 20, 30, 40, and 50 °C. Record the exact temperature of each.
- For each temperature, measure out the following into separate beakers and allow to equilibrate in the water bath for at least 5 minutes: Beaker A: 20 cm³ KI + 5 cm³ Na₂S₂O₃ + 5 cm³ starch solution + 10 cm³ H₂SO₄ Beaker B: 10 cm³ H₂O₂
- When both beakers have reached the target temperature, pour Beaker B into Beaker A and start the stopwatch immediately. Swirl gently.
- Place the beaker on a white tile and stop the stopwatch the moment the solution turns blue-black.
- Record the time in seconds.
- Repeat each temperature at least three times and calculate a mean time.
- Rinse all equipment thoroughly between runs.
Results Table
| Temperature (°C) | Temperature (K) | 1/T (K⁻¹) | Time 1 (s) | Time 2 (s) | Time 3 (s) | Mean time (s) | Rate (1/t, s⁻¹) | ln(rate) |
|---|---|---|---|---|---|---|---|---|
| 10 | 283 | 0.003534 | ||||||
| 20 | 293 | 0.003413 | ||||||
| 30 | 303 | 0.003300 | ||||||
| 40 | 313 | 0.003195 | ||||||
| 50 | 323 | 0.003096 |
Analysis
1. Calculate the rate of reaction as 1/t (s⁻¹) for each temperature.
2. Plot a graph of ln(rate) on the y-axis against 1/T (K⁻¹) on the x-axis — this is an Arrhenius plot.
3. The gradient of the best-fit line = −Eₐ/R, where R = 8.314 J mol⁻¹ K⁻¹.
4. Therefore: Eₐ = −gradient × R
The activation energy for this reaction is typically around 40–60 kJ mol⁻¹. How does your calculated value compare?
Discussion Points
- Why does reaction rate increase with temperature? Refer to collision theory and the Maxwell-Boltzmann distribution.
- Why is 1/t used as a measure of rate rather than the actual rate constant k?
- What are the main sources of uncertainty? Consider temperature control, timing the colour change, and mixing.
- How could the design be improved? Consider using a colorimeter to give a more objective end-point.
IA Guidance
This experiment works very well as an IB Chemistry IA. To score highly:
- Research Design: Justify your choice of temperature range and number of trials. Explain how you will control variables precisely.
- Data Analysis: Include a full Arrhenius plot with error bars. Propagate uncertainties through your calculation of Eₐ.
- Conclusion: Compare your Eₐ value to a literature value and discuss any discrepancy.
- Evaluation: Identify systematic and random errors and suggest specific, realistic improvements.
Practical Science 