Concept of activation energy

Concept of Activation Energy

Activation energy is the minimum amount of energy required for a chemical reaction to occur. It refers to the minimum energy threshold that reacting molecules must surpass to convert into products. Even if collisions occur, a reaction won't proceed without this required energy.

Understanding Activation Energy

When molecules collide, they need a certain amount of energy to break existing bonds and form new ones. This minimum required energy is known as the activation energy (E_a). A chemical reaction occurs only when the colliding particles have energy equal to or exceeding E_a.

Energy Profile Diagram

An energy diagram provides a visual representation of the activation energy involved in a reaction. The peak of the curve represents the transition state, and the energy gap between the reactants and the peak is the activation energy.

Arrhenius Equation

The rate of a chemical reaction depends on the activation energy and temperature. This relationship is expressed by the Arrhenius equation:

k = A · e^–E_a/RT

Where:

• k = rate constant
• A = frequency factor (indicates collision frequency and orientation)
• E_a = activation energy
• R = universal gas constant
• T = temperature in Kelvin

Effect of Catalyst

By lowering the activation energy through an alternative route, catalysts enhance the reaction rate without being used up or altering the thermodynamic properties.

Key Characteristics

• Activation energy determines how fast or slow a reaction proceeds.
• Higher activation energy means slower reaction rate.
• Increasing temperature or using a catalyst can overcome high activation energy barriers.

Example

In the combustion of methane:

CH₄ + 2O₂ → CO₂ + 2H₂O

The reaction requires a spark to supply the activation energy. Once initiated, it proceeds rapidly and exothermically.