The activation energy (Ea) is the minimum amount of energy required for a chemical reaction to occur. The relationship between activation energy and the rate of reaction can be best summarized as follows:

1. **High Activation Energy:** Reactions with high activation energy tend to proceed slowly. This is because a larger energy barrier must be overcome for the reactants to form products, making it less likely for collisions between reactant molecules to result in a reaction.

2. **Low Activation Energy:** Conversely, reactions with low activation energy usually occur more rapidly. Since less energy is required to initiate the reaction, more reactant molecules can successfully collide with enough energy to break bonds and form new products.

Thus, generally, the lower the activation energy, the higher the rate of the reaction. This relationship is quantitatively expressed by the Arrhenius equation:
$$
k = A e^{-E_a/(RT)}
$$
where $ k $ is the rate constant, $ A $ is the pre-exponential factor, $ E_a $ is the activation energy, $ R $ is the universal gas constant, and $ T $ is the temperature in Kelvin. Hence, a lower $ E_a $ results in a larger value of $ k $, indicating a faster reaction rate.

Question

The activation energy (Ea) is the minimum amount of energy required for a chemical reaction to occur. The relationship between activation energy and the rate of reaction can be best summarized as follows:

1. **High Activation Energy:** Reactions with high activation energy tend to proceed slowly. This is because a larger energy barrier must be overcome for the reactants to form products, making it less likely for collisions between reactant molecules to result in a reaction.

2. **Low Activation Energy:** Conversely, reactions with low activation energy usually occur more rapidly. Since less energy is required to initiate the reaction, more reactant molecules can successfully collide with enough energy to break bonds and form new products.

Thus, generally, the lower the activation energy, the higher the rate of the reaction. This relationship is quantitatively expressed by the Arrhenius equation:
$$
k = A e^{-E_a/(RT)}
$$
where $ k $ is the rate constant, $ A $ is the pre-exponential factor, $ E_a $ is the activation energy, $ R $ is the universal gas constant, and $ T $ is the temperature in Kelvin. Hence, a lower $ E_a $ results in a larger value of $ k $, indicating a faster reaction rate.

Knowee AI · Accepted Answer

The activation energy (Ea) is the minimum amount of energy required for a chemical reaction to occur. The relationship between activation energy and the rate of reaction can be best summarized as follows:

1. **High Activation Energy:** Reactions with high activation energy tend to proceed slowly. This is because a larger energy barrier must be overcome for the reactants to form products, making it less likely for collisions between reactant molecules to result in a reaction.

2. **Low Activation Energy:** Conversely, reactions with low activation energy usually occur more rapidly. Since less energy is required to initiate the reaction, more reactant molecules can successfully collide with enough energy to break bonds and form new products.

Thus, generally, the lower the activation energy, the higher the rate of the reaction. This relationship is quantitatively expressed by the Arrhenius equation:
$$
k = A e^{-E_a/(RT)}
$$
where $ k $ is the rate constant, $ A $ is the pre-exponential factor, $ E_a $ is the activation energy, $ R $ is the universal gas constant, and $ T $ is the temperature in Kelvin. Hence, a lower $ E_a $ results in a larger value of $ k $, indicating a faster reaction rate.

Which statement best describes the relationship between activation energy and rate of reaction?

Question

Which statement best describes the relationship between activation energy and rate of reaction?

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