Study of reaction rates of Reaction of iodide ions with hydrogen peroxide (varying iodide concentration)

Understanding reaction rates is fundamental to chemical kinetics, and the reaction between iodide ions and hydrogen peroxide serves as an excellent model for studying how concentration affects reaction velocity. This detailed guide covers the pitch lab experiment for investigating the effect of iodide concentration on reaction rate.

Aim

To investigate the effect of varying iodide ion concentration on the rate of reaction between iodide ions and hydrogen peroxide, and to determine the order of reaction with respect to iodide ions.

Apparatus Required

• Burette (25 mL) - 2 nos.
• Pipette (10 mL) - 1 no.
• Conical flasks (250 mL) - 5 nos.
• Measuring cylinder (50 mL) - 1 no.
• Stopwatch or digital timer
• White tile with cross mark
• Beakers (100 mL) - several
• Wash bottle
• Test tube stand

Chemicals Required

• Potassium iodide (KI) solution (0.1 M)
• Hydrogen peroxide (H₂O₂) solution (0.1 M)
• Sodium thiosulfate (Na₂S₂O₃) solution (0.05 M)
• Starch solution (freshly prepared 1%)
• Dilute sulfuric acid (H₂SO₄) (1 M)

Theory

The reaction between iodide ions and hydrogen peroxide in acidic medium can be represented as:

H₂O₂ + 2I⁻ + 2H⁺ → I₂ + 2H₂O

This reaction is monitored using the iodine clock reaction method. A known quantity of sodium thiosulfate is added to the reaction mixture, which immediately reacts with the liberated iodine:

I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻

Once all thiosulfate is consumed, the excess iodine reacts with starch indicator, producing a characteristic blue-black color. The time taken for this color change indicates the reaction rate.

Rate Law Expression

Rate = k[H₂O₂]^m[I⁻]^n

Where:

• k = rate constant
• m, n = orders of reaction with respect to respective reactants

By keeping [H₂O₂] constant and varying [I⁻], we can determine 'n'.

Concentration-Time Relationship

The reciprocal of time (1/t) is directly proportional to the rate of reaction.

Rate ∝ 1/t

Procedure

1. Preparation of Solutions:

   • Prepare five different concentrations of KI solution by dilution
   • Keep H₂O₂ concentration constant

2. Standard Mixture Preparation:

   • Take 10 mL of 0.05 M Na₂S₂O₃ in each conical flask
   • Add 5 mL of starch solution to each flask
   • Add required volume of dilute H₂SO₄ (5 mL)

3. Experiment Execution:

   • Add different volumes of KI solution to each flask
   • Make up total volume to 50 mL with distilled water
   • Add 10 mL of H₂O₂ solution rapidly and start the timer
   • Shake the flask and place it over white tile with cross mark
   • Stop the timer when blue-black color appears
   • Record the time for each trial

4. Repeat the Experiment:

   • Conduct each concentration trial in triplicate for accuracy
   • Calculate average time for each concentration

Observation Table

Result

1. Graphical Analysis:

   • Plot a graph between [I⁻] vs Rate (1/t)
   • If straight line passing through origin → First order reaction
   • If curved graph → Other order reaction

2. Order Determination:

   • From the experimental data, the rate increases proportionally with iodide ion concentration
   • Order with respect to I⁻ = 1 (First order)

3. Rate Law: Rate = k[H₂O₂][I⁻]

4. Rate Constant Calculation: Using the rate equation and experimental data, calculate the average rate constant.

Precautions

1. Temperature Control:

   • Maintain constant temperature throughout the experiment
   • Perform all trials at room temperature

2. Timing Accuracy:

   • Start timer immediately after adding H₂O₂
   • Stop timer precisely at the appearance of blue color

3. Mixing Efficiency:

   • Mix solutions thoroughly but gently
   • Avoid splashing during rapid addition

4. Chemical Handling:

   • Handle H₂O₂ carefully as it's a strong oxidizing agent
   • Use freshly prepared starch solution for sharp endpoint

5. Glassware:

   • Rinse all apparatus with respective solutions before use
   • Ensure accurate measurement of volumes

Viva Voce Questions and Answers

Q1: What is the role of sodium thiosulfate in this experiment?

A: Sodium thiosulfate acts as a titrant that reacts with liberated iodine, preventing immediate color formation. Once all thiosulfate is consumed, excess iodine reacts with starch, indicating reaction completion.

Q2: Why is starch solution used as an indicator?

A: Starch forms a deep blue-black complex with iodine, providing a sharp and easily detectable color change endpoint for accurate timing.

Q3: How does varying iodide concentration affect the reaction rate?

A: Increasing iodide concentration increases the reaction rate proportionally, as more iodide ions are available to react with hydrogen peroxide, demonstrating first-order kinetics.

Q4: What is the significance of maintaining constant H₂O₂ concentration?

A: Keeping H₂O₂ concentration constant allows isolation of variables, enabling accurate determination of the effect of iodide ion concentration alone on reaction rate.

Q5: What are the units of rate constant for this reaction?

A: For a second-order overall reaction (first order in both H₂O₂ and I⁻), the rate constant units are L mol⁻¹ s⁻¹.

Q6: Why is the reaction performed in acidic medium?

A: Acidic conditions (H⁺ ions) are essential for the reaction to proceed, as hydrogen ions participate in the overall reaction mechanism and enhance the reaction rate.

Q7: How can you confirm the order of reaction graphically?

A: Plot [I⁻] vs Rate (1/t) – a straight line through origin confirms first-order dependence.

Conclusion

This experiment successfully demonstrates the direct proportionality between iodide ion concentration and reaction rate, confirming first-order kinetics with respect to iodide ions. The systematic approach using the iodine clock reaction provides accurate and reproducible results for kinetic studies. Understanding these fundamental principles is crucial for advanced studies in chemical kinetics and reaction mechanism analysis.

The experiment highlights the importance of proper experimental technique, data analysis, and mathematical interpretation in determining reaction orders and rate laws – essential skills for chemistry students and researchers.