7.5 Activation Energy

This topic explains activation energy — what it is, why it is needed, and how it affects the rate of a chemical reaction. It also describes the activated complex and potential energy diagrams.

Definition

Activation Energy (Eₐ): The minimum amount of energy, in addition to the average kinetic energy, which the particles must possess for effective collisions.

Activated Complex (Transition State): An unstable, high-energy species formed during an effective collision.

Main Idea

Chemical reactions involve the breaking and making of chemical bonds. These changes are accompanied by changes in energies. Collision theory has been proposed to explain the observed kinetics of reactions. For a chemical reaction to occur, the combining atoms or molecules must collide with one another. These collisions may be effective or ineffective depending upon the energy and orientation of the colliding particles. The effective collision can take place only if the energy of the colliding particles is high enough to overcome the repulsion between electrons around the reacting particles.

Core Concept

• Activation energy (Eₐ): Minimum energy for effective collision.
• Effective collision: Sufficient energy + correct orientation.
• Activated complex: High-energy, short-lived species.
• Higher Eₐ → slower reaction.
• Lower Eₐ → faster reaction.

Detailed Explanation

What is Activation Energy?

• Chemical reactions involve the breaking and making of chemical bonds.
• These changes are accompanied by changes in energies.
• Collision theory has been proposed to explain the observed kinetics of reactions.
• For a chemical reaction to occur, the combining atoms or molecules must collide with one another.
• These collisions may be effective or ineffective depending upon the energy and orientation of the colliding particles.
• The effective collision can take place only if the energy of the colliding particles is high enough to overcome the repulsion between electrons around the reacting particles.
• Proper orientation means that at the time of collision, the atoms which are required to make new bonds should collide with each other.

Do You Know?

Gas explosions in homes can be caused by switching on a light. If gas has been leaking, then a tiny spark from turning on a light can provide the activation energy to start the explosive reaction between the methane and the oxygen.

Definition of Activation Energy

- The minimum amount of energy, in addition to the average kinetic energy, which the particles must possess for effective collisions, is called activation energy.

Effect of Activation Energy on Reaction Rate

• No reaction occurs if the energy of the reacting particles is lower than the activation energy.
• Thus, the speed of a reaction depends on its activation energy.
• The higher the activation energy, the lower the reaction rate.
• This is because only a small fraction of the molecules have enough energy to react.
• On the other hand, if the activation energy is small, a large number of molecules can cause effective collisions.
• Therefore, the reaction rate is higher.

Activated Complex

• Consider the reaction between A₂ and B₂ molecules and a new molecule AB.
• If the energy of these molecules is equal to or greater than the activation energy, the collision breaks their bonds and new bonds are formed.
• In an effective collision, the molecules form an unstable species called an activated complex.
• Being a high-energy species, it is short-lived and quickly breaks down into products.
• The activated complex is also called the transition state.

Potential Energy Diagram

• In an efficient collision, the colliding molecules move closer to each other and slow down just before the collision.
• Their kinetic energy decreases, and this leads to a corresponding increase in their potential energy.
• Activation energy appears in the form of hills between reactants and products.
• Molecules must first climb over an energy barrier before they can roll down the hill to form products.
• Only molecular collisions with the right activation energy and orientation can do this.
• On the other hand, if they lack the right activation energy, they will not reach the top of the mountain and will fall back chemically unchanged.

Enthalpy Changes in Reactions

• A potential energy diagram can also be used to understand enthalpy changes in chemical reactions.
• The reaction enthalpy is equal to the difference between the energies of the reactants and the products.
• In an exothermic reaction, the products are at a lower energy level than the reactants.
• In an endothermic process, the products have a higher energy level than the reactants.
• In both reactions, the activation energy (Eₐ) is an energy barrier that must be overcome before products can be formed.
• If the activation energy is not available to the reacting particles, the reaction will not start.
• Endothermic reactions require a constant source of energy to complete the reaction.

Intermediate Level

Activation Energy

Potential Energy Diagram

• Hill = activation energy.
• Reactants at lower energy.
• Products at lower (exothermic) or higher (endothermic) energy.
• Eₐ is the barrier.

Exothermic vs. Endothermic

Key Points

• No reaction if energy < Eₐ.
• Higher Eₐ → lower rate.
• Lower Eₐ → higher rate.
• Endothermic reactions need constant energy.

Important Points

• 1. Activation energy: Minimum energy for effective collision.
• 2. Effective collision: Sufficient energy + correct orientation.
• 3. Activated complex: High-energy, short-lived species.
• 4. Higher Eₐ → slower reaction.
• 5. Lower Eₐ → faster reaction.
• 6. Potential energy diagram: Hill = Eₐ.
• 7. Exothermic: Products lower energy.
• 8. Endothermic: Products higher energy.
• 9. Endothermic reactions need constant energy.

Remember

• - Activation energy: Minimum energy for reaction
• - Effective collision: Energy + orientation
• - Activated complex: High-energy species
• - Higher Eₐ: Slower reaction
• - Lower Eₐ: Faster reaction
• - Exothermic: Products lower energy
• - Endothermic: Products higher energy

Common Mistakes

• 1. Thinking all collisions are effective: Only those with enough energy and correct orientation.
• 2. Forgetting the role of orientation: Molecules must collide correctly.
• 3. Confusing activation energy with enthalpy change: Eₐ is the barrier; ΔH is the net energy change.
• 4. Thinking endothermic reactions don't need energy: They need constant energy.

Quick Revision

• Activation energy: Minimum energy for effective collision.
• Effective collision: Sufficient energy + correct orientation.
• Activated complex: High-energy, short-lived species.
• Higher Eₐ → slower reaction.
• Lower Eₐ → faster reaction.
• Potential energy diagram: Hill = Eₐ.
• Exothermic: Products lower energy.
• Endothermic: Products higher energy.
• Endothermic reactions need constant energy.

60-Second Revision

• Activation energy: Minimum energy for reaction
• Effective collision: Energy + orientation
• Activated complex: High-energy species
• Higher Eₐ: Slower reaction
• Lower Eₐ: Faster reaction
• Exothermic: Products lower energy
• Endothermic: Products higher energy

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• Chapter: 7 (Chemical Kinetics)

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