The energy of activation of positive catalyzed reaction as compared to that of an uncatalyzed reaction is

Options

A. more
B. less
C. same
D. may be more or less

Explanation

To solve the question regarding the energy of activation of a positively catalyzed reaction compared to that of an uncatalyzed reaction, we can follow these steps:

### Step-by-Step Solution:

1. **Understanding Activation Energy**:
- Activation energy (Ea) is the minimum energy required for a chemical reaction to occur. It is the energy barrier that reactants must overcome to form products.

2. **Comparing Catalyzed and Uncatalyzed Reactions**:
- In an uncatalyzed reaction, the reactants must reach a certain energy level (activation energy) to form the transition state and then proceed to products.
- In a catalyzed reaction, a catalyst provides an alternative pathway with a lower activation energy.

3. **Energy Profile Diagram**:
- Draw an energy profile diagram. On the y-axis, represent energy, and on the x-axis, represent the progress of the reaction.
- Mark the energy level of the reactants and products for both the uncatalyzed and catalyzed reactions.

4. **Identifying Activation Energies**:
- Identify the activation energy for the uncatalyzed reaction (Ea uncatalyzed) as the energy difference between the reactants and the transition state.
- Identify the activation energy for the catalyzed reaction (Ea catalyzed) as the energy difference between the reactants and the transition state in the presence of the catalyst.

5. **Conclusion**:
- Since the catalyst lowers the activation energy, we can conclude that:
\[ Ea \text{ (catalyzed)} < Ea \text{ (uncatalyzed)} \]
- Therefore, the energy of activation for a positively catalyzed reaction is less than that of an uncatalyzed reaction.

### Final Answer:
The energy of activation of a positively catalyzed reaction is less than that of an uncatalyzed reaction.

---