Assertion (A) : Rate of reaction will be doubled, when temperature increased from 298 k to 308 k.
Reason (R ): The activation energy of reaction decreases with increase in temperature.

Options

A. Both (A ) and (R) are true (R ) is the correct explanation to (A)
B. Both (A) and (R ) are true but (R) is not the correct explanation to (A)
C. (A)is true but (R) is false
D. Both (A) and (R) are false

Explanation

To solve the given question, we need to analyze both the assertion (A) and the reason (R) provided.

### Step 1: Analyze the Assertion (A)
The assertion states that the rate of reaction will be doubled when the temperature is increased from 298 K to 308 K.

- **Explanation**: It is a well-known principle in chemical kinetics that increasing the temperature generally increases the rate of reaction. This is primarily due to the increase in kinetic energy of the molecules, which leads to more frequent and effective collisions. Empirical observations suggest that for many reactions, a temperature increase of about 10 degrees Celsius (or 10 K) can approximately double the reaction rate.

### Step 2: Analyze the Reason (R)
The reason states that the activation energy of the reaction decreases with an increase in temperature.

- **Explanation**: This statement is incorrect. The activation energy (Ea) is a characteristic of the reaction and does not change with temperature. While the number of molecules that have enough energy to overcome the activation barrier increases with temperature, the activation energy itself remains constant. The Arrhenius equation supports this, as it shows that the rate constant (k) depends on the activation energy and temperature, but not that the activation energy decreases with temperature.

### Step 3: Conclusion
Based on the analysis:
- Assertion (A) is true.
- Reason (R) is false.

Therefore, the correct answer to the question is that A is true but R is false.

### Final Answer
The correct answer is: **C (A is true, R is false)**.

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