Exercise 9 page 71:

9 The oxidation of iodide ion I- with peroxydisulfate S2O82- is given according to the equation:

2 I- + S2O82- → I2 + 2 SO42-

Three experiments 1, 2, 3 are performed starting with the same initial concentration of reactants, but varying the temperature in each experiment, and adding a few drops of iron (III) sulfate Fe2(SO4)3 in one of them.

The results are represented in the following table:

Experiment | Temperature (°C) | r0 (mol.L-1.min-1)
1 | T1 | 2.05 x 10^-4
2 | T2 T1 | 3.65 x 10^-4
3 | T2 (presence of Fe3+ ions) | 9.75 x 10^-4

Explain the role of the kinetic factors in the values of r0 for each experiment.

Q: Exercise 9 page 71: 9 The oxidation of iodide ion I- with peroxydisulfate S2O82- is given according to the equation: 2 I- + S2O82- → I2 + 2 SO42- Three experiments 1, 2, 3 are performed starting with the same initial concentration of reactants, but varying the temperature in each experiment, and adding a few drops of iron (III) sulfate Fe2(SO4)3 in one of them. The results are represented in the following table: Experiment | Temperature (°C) | r0 (mol.L-1.min-1) 1 | T1 | 2.05 x 10^-4 2 | T2 > T1 | 3.65 x 10^-4 3 | T2 (presence of Fe3+ ions) | 9.75 x 10^-4 Explain the role of the kinetic factors in the values of r0 for each experiment.

The exercise involves the oxidation of iodide ions $ \text{I}^- $ with peroxydisulfate ions $ \text{S}_2\text{O}_8^{2-} $. The reaction is as follows: \[ 2 \text{I}^- + \text{S}_2\text{O}_8^{2-} \rightarrow \text{I}_2 + 2 \text{SO}_4^{2-} \] Three experiments were conducted with varying temperatures and the addition of iron (III) sulfate in one experiment. The results show different initial rates $ r_0 $. ### Explanation of Kinetic Factors: 1. **Experiment 1 (Temperature $ T_1 $):** - Initial rate $ r_0 = 2.05 \times 10^{-4} \, \text{mol.L}^{-1}\text{min}^{-1} $ - This is the baseline rate at temperature $ T_1 $. 2. **Experiment 2 (Higher Temperature $ T_2 > T_1 $):** - Initial rate $ r_0 = 3.65 \times 10^{-4} \, \text{mol.L}^{-1}\text{min}^{-1} $ - The increase in temperature $ T_2 $ results in a higher rate. According to the Arrhenius equation, increasing temperature increases the kinetic energy of the molecules, leading to more frequent and energetic collisions, thus increasing the reaction rate. 3. **Experiment 3 (Presence of $ \text{Fe}^{3+} $ ions at $ T_2 $):** - Initial rate $ r_0 = 9.75 \times 10^{-4} \, \text{mol.L}^{-1}\text{min}^{-1} $ - The presence of $ \text{Fe}^{3+} $ ions acts as a catalyst, further increasing the rate significantly. Catalysts provide an alternative pathway with a lower activation energy for the reaction, increasing the reaction rate without being consumed. ### Summary: - **Temperature Increase:** Raises the reaction rate by increasing molecular collisions. - **Catalyst Addition ($ \text{Fe}^{3+} $):** Significantly enhances the reaction rate by lowering activation energy.

Question

Exercise 9 page 71:

9 The oxidation of iodide ion I- with peroxydisulfate S2O82- is given according to the equation:

2 I- + S2O82- → I2 + 2 SO42-

Three experiments 1, 2, 3 are performed starting with the same initial concentration of reactants, but varying the temperature in each experiment, and adding a few drops of iron (III) sulfate Fe2(SO4)3 in one of them.

The results are represented in the following table:

Experiment  | Temperature (°C) | r0 (mol.L-1.min-1) 
1           | T1               | 2.05 x 10^-4 
2           | T2 > T1        | 3.65 x 10^-4 
3           | T2 (presence of Fe3+ ions) | 9.75 x 10^-4

Explain the role of the kinetic factors in the values of r0 for each experiment.

Accepted Answer

The exercise involves the oxidation of iodide ions $ \text{I}^- $ with peroxydisulfate ions $ \text{S}_2\text{O}_8^{2-} $. The reaction is as follows:

$$ 2 \text{I}^- + \text{S}_2\text{O}_8^{2-} \rightarrow \text{I}_2 + 2 \text{SO}_4^{2-} $$

Three experiments were conducted with varying temperatures and the addition of iron (III) sulfate in one experiment. The results show different initial rates $ r_0 $.

### Explanation of Kinetic Factors:

1. **Experiment 1 (Temperature $ T_1 $):**
   - Initial rate $ r_0 = 2.05 \times 10^{-4} \, \text{mol.L}^{-1}\text{min}^{-1} $
   - This is the baseline rate at temperature $ T_1 $.

2. **Experiment 2 (Higher Temperature $ T_2 > T_1 $):**
   - Initial rate $ r_0 = 3.65 \times 10^{-4} \, \text{mol.L}^{-1}\text{min}^{-1} $
   - The increase in temperature $ T_2 $ results in a higher rate. According to the Arrhenius equation, increasing temperature increases the kinetic energy of the molecules, leading to more frequent and energetic collisions, thus increasing the reaction rate.

3. **Experiment 3 (Presence of $ \text{Fe}^{3+} $ ions at $ T_2 $):**
   - Initial rate $ r_0 = 9.75 \times 10^{-4} \, \text{mol.L}^{-1}\text{min}^{-1} $
   - The presence of $ \text{Fe}^{3+} $ ions acts as a catalyst, further increasing the rate significantly. Catalysts provide an alternative pathway with a lower activation energy for the reaction, increasing the reaction rate without being consumed.

### Summary:
- **Temperature Increase:** Raises the reaction rate by increasing molecular collisions.
- **Catalyst Addition ($ \text{Fe}^{3+} $):** Significantly enhances the reaction rate by lowering activation energy.

Exercise 9 page 71: 9 The oxidation of iodide ion I- with p...

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