chemical reactions

Introduction

Activation energy (Ea) is the minimum amount of energy required to initiate a chemical reaction. Think of it as the energy “hump” that reactants must overcome to transform into products. Even exothermic reactions, which release energy overall, require an initial input of energy to get started. Exothermic and Endothermic Reactions

The Collision Theory

Activation energy is closely tied to the collision theory, which states that for a reaction to occur, reactant particles must:

1. Collide with each other.
2. Collide with the correct orientation.
3. Collide with sufficient energy to break existing bonds and form new ones.

Activation energy represents the minimum energy required to satisfy the third condition.

The Maxwell-Boltzmann Distribution

The Maxwell-Boltzmann distribution describes the distribution of kinetic energies among particles in a sample at a given temperature. It shows that only a fraction of particles possess enough kinetic energy to overcome the activation energy barrier. Maxwell-Boltzmann Distribution

Effect of Temperature on Reaction Rate

Increasing the temperature increases the average kinetic energy of the particles. This leads to a larger fraction of particles having energy greater than or equal to Ea, resulting in a higher reaction rate.

The Arrhenius Equation

The relationship between the rate constant (k) of a reaction, activation energy, and temperature is described by the Arrhenius equation:

$ k = Ae^{\frac{-E_a}{RT}} $

where:

• k is the rate constant
• A is the frequency factor (pre-exponential factor) Frequency Factor
• Ea is the activation energy
• R is the ideal gas constant (8.314 J/mol·K)
• T is the temperature in Kelvin

Cataly