Activation Energy
Definition of Activation Energy
Activation energy
Activation energy ($E_a$) is the minimum energy required for colliding particles to form an activated complex (also known as the transition state) and proceed to products.
- When particles collide, they must possess sufficient energy to overcome the energy barrier that separates reactants from products.
- This energy barrier corresponds to the activation energy.
- If the colliding particles do not have energy equal to or greater than $E_a$, the reaction will not proceed.
Think of activation energy as the “entry fee” that particles must pay to initiate a chemical reaction.
The Activated Complex (Transition State)
Activated complex
The activated complex is a high-energy, unstable arrangement of atoms that exists momentarily as reactants are transformed into products. It represents the "peak"; of the energy barrier in a reaction.
Graphical Representation: Maxwell–Boltzmann Energy Distribution
Maxwell–Boltzmann energy distribution curve
The Maxwell–Boltzmann energy distribution curve provides a visual representation of the distribution of kinetic energies among particles in a system.
It is particularly useful for understanding how temperature affects the proportion of particles with energy greater than or equal to the activation energy ($E_a$).
Key Features of the Maxwell–Boltzmann Curve
- X-axis: Represents the kinetic energy of particles.
- Y-axis: Represents the number of particles with a given energy.
- Shape: The curve is asymmetrical, with a peak corresponding to the most probable energy (the energy possessed by the largest number of particles).
- Area under the curve: Represents the total number of particles in the system.
Effect of Temperature on the Curve
- At lower temperatures: The curve is taller and narrower, with fewer particles having energy ≥ $E_a$.
- At higher temperatures: The curve flattens and spreads out, increasing the proportion of particles with energy ≥ $E_a$.
- Consider a reaction with an activation energy of $50 \text{ kJ mol}^{-1}$.
- At 300 K, only a small fraction of particles have enough energy to overcome this barrier.
- However, if the temperature is increased to 400 K, a significantly larger fraction of particles will have energy $≥ 50 \text{ kJ mol}^{-1}$, leading to a faster reaction rate.
Increasing temperature does not change the activation energy itself; it only increases the number of particles with sufficient energy to overcome it.
- Do not confuse the peak of the Maxwell–Boltzmann curve with the activation energy.
- The peak represents the most probable energy, not $E_a$.
Energy Profiles for Chemical Reactions
- Energy profiles illustrate the energy changes that occur during a chemical reaction.
- They show the energy difference between reactants, products, and the activation energy barrier.
Components of an Energy Profile
- Reactants: The starting substances of the reaction.
- Products: The substances formed after the reaction.
- Activation Energy ($E_a$): The energy required to reach the activated complex.
- Enthalpy Change ($\Delta H$): The difference in energy between reactants and products.
Endothermic vs. Exothermic Reactions
Exothermic Reactions
- Definition: Reactions that release energy to the surroundings.
- Energy Profile:
- The energy of the products is lower than that of the reactants.
- $\Delta H$ is negative.
Endothermic Reactions
- Definition: Reactions that absorb energy from the surroundings.
- Energy Profile:
- The energy of the products is higher than that of the reactants.
- $\Delta H$ is positive.
- Students often confuse the activation energy ($E_a$) with the enthalpy change ($\Delta H$).
- Remember, $E_a$ is the energy barrier, while $\Delta H$ is the net energy change of the reaction.
- What is the role of activation energy in a chemical reaction?
- How does increasing temperature affect the Maxwell–Boltzmann energy distribution curve?
- Compare the energy profiles of exothermic and endothermic reactions.
