Rate of a general reaction A + B `to` products can be expressed as follows on the basis of collision theory.
Rate `= Z_(AB) e^(-E_(a)//RT)`
Which of the following statements is not correct for the above expression?

Options

A. Z is collision frequency andis equal to number of collisions per second per unit volume of the reaction mixture
B. `e^(E_(a)//RT)` is the fraction of molecules with kinetic energy equal to or greater than `E_(a)`
C. `E_(a)` is activation energy of the reaction.
D. All the molecules which collide with one other are effective collisions.

Explanation

To determine which statement is not correct regarding the rate expression for the reaction A + B → products, we will analyze each option based on the principles of collision theory.

### Step-by-Step Solution:

1. **Understanding the Rate Expression**:
The rate of the reaction is given by:
\[ \text{Rate} = Z_{AB} e^{-\frac{E_a}{RT}} \]
where:
- \( Z_{AB} \) is the collision frequency.
- \( E_a \) is the activation energy.
- \( R \) is the universal gas constant.
- \( T \) is the temperature in Kelvin.

2. **Analyzing the Statements**:
We need to evaluate the correctness of the following statements:

- **Option A**: "Z is collision frequency and is equal to the number of collisions per second per unit volume of the reaction mixture."
- This statement is correct. Collision frequency \( Z \) indeed represents the number of collisions per unit time per unit volume.

- **Option B**: "e^{-\frac{E_a}{RT}} is the fraction of molecules with kinetic energy equal to or greater than activation energy."
- This statement is also correct. The exponential term represents the fraction of molecules that have enough energy to overcome the activation energy barrier.

- **Option C**: "E_a is the activation energy of the reaction."
- This statement is correct as well. \( E_a \) is defined as the minimum energy required for a reaction to occur.

- **Option D**: "All the molecules which collide with one another are effective collisions."
- This statement is incorrect. Not all collisions are effective. Only those collisions that occur with sufficient energy (threshold energy) and proper orientation lead to a reaction.

3. **Conclusion**:
Based on the analysis, the statement that is not correct is **Option D**. Effective collisions are defined by both adequate energy and proper orientation, meaning not all collisions result in a reaction.

### Final Answer:
The statement that is not correct is **Option D**: "All the molecules which collide with one another are effective collisions."

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