Which statement is usually true about the relationship between activation energy and reaction rates?

• A. A higher activation energy always results in a faster reaction rate.
• B. The reaction rate is independent of activation energy if temperature is constant.
• C. Activation energy does not affect the reaction rate.
• D. A lower activation energy generally leads to a faster reaction rate.

Answer: (D)

Verified step by step guidance

1. Understand that the activation energy (E_a) is the minimum energy required for reactants to transform into products during a chemical reaction.
2. Recall the Arrhenius equation, which relates the reaction rate constant (k) to the activation energy and temperature: \(k = A \times e^{-\frac{E_a}{RT}}\), where \(A\) is the frequency factor, \(R\) is the gas constant, and \(T\) is the temperature in Kelvin.
3. Note that in the Arrhenius equation, the exponential term \(e^{-\frac{E_a}{RT}}\) shows that as activation energy \(E_a\) decreases, the value of \(k\) increases, meaning the reaction rate increases.
4. Recognize that temperature is held constant in this problem, so changes in reaction rate are primarily influenced by changes in activation energy.
5. Conclude that a lower activation energy generally leads to a faster reaction rate because more molecules have sufficient energy to overcome the energy barrier, making the reaction proceed more quickly.