Statement-1: If the activation energy of reaction is zero temperature will have no effect on the rate constant.Statement-2: Lower the activation energy fasten is the reaction.a)Statement-1 is True, Statement-2 is True; Statement-2 is a correct explanation for Statement-1.b)Statement-1 is True, Statement-2 is True; Statement-2 is NOT a correct explanation for Statement-1.c)Statement-1 is True, Statement-2 is False.d)Statement-1 is False, Statement-2 is TrueCorrect answer is option 'B'. Can you explain this answer?

Question

Answer

Statement-1: If the activation energy of reaction is zero temperature will have no effect on the rate constant.
Statement-2: Lower the activation energy fasten is the reaction.

Explanation:
The rate constant of a chemical reaction is determined by the Arrhenius equation, which relates the rate constant (k) to the activation energy (Ea) and the temperature (T):

k = A * exp(-Ea/RT)

Where:
- k is the rate constant
- A is the pre-exponential factor or frequency factor
- Ea is the activation energy
- R is the gas constant
- T is the temperature in Kelvin

Let's analyze each statement separately:

Statement-1: If the activation energy of reaction is zero temperature will have no effect on the rate constant.

This statement is not true. According to the Arrhenius equation, if the activation energy (Ea) is zero, the rate constant (k) will become temperature-dependent. The exponential term in the equation will reduce to 1, and the rate constant will be equal to the pre-exponential factor (A). Therefore, the rate constant will increase with increasing temperature.

Statement-2: Lower the activation energy fasten is the reaction.

This statement is true. The activation energy (Ea) is a measure of the energy barrier that must be overcome for a reaction to occur. A lower activation energy means that more reactant molecules have enough energy to overcome the barrier and react. As a result, the reaction will proceed at a faster rate.

Explanation of the correct option:

Option B is the correct answer: Statement-1 is True, Statement-2 is True; Statement-2 is NOT a correct explanation for Statement-1.

Explanation:
While both statements are individually true, Statement-2 does not provide a correct explanation for Statement-1. The activation energy affects the temperature dependence of the rate constant, but a zero activation energy does not imply that the rate constant is independent of temperature. Therefore, the two statements are not directly related, and Statement-2 does not explain why Statement-1 is true.

In conclusion, the activation energy of a reaction plays a crucial role in determining the rate constant and the temperature dependence of the reaction. A lower activation energy allows the reaction to proceed at a faster rate, but a zero activation energy does not eliminate the temperature dependence of the rate constant.

Similar Class 12 Doubts

Statement-1 :For a reaction A(g) → B(g) — re= 2.5 PA at 400K — re= 2.5 PA at 600Kactivation energy is 4135 J/mol.Statement-2 :Since for any reaction, values of rate constant at two different temp is same therefore activation energy of the reaction is zero.

At a certain temperature, the first order rate constant k1 is found to be smaller than the second order rate constant k2. If the energy of activation E1 of the first order reaction is greater than energy of activation E2 of the second order reaction, then with increase in temperature.

Two reactions with different activation energies have the same rate at room temperature. Which statement correctly describes the rates of these two reactions at the same higher temperature?

Which among the following statement is not true for rate constant of a reaction?

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