Chemical Reactions - Collision Theory

Collision Theory states that for a chemical reaction to occur, the reacting particles must collide with each other with sufficient energy and the correct orientation.

An Effective Collision is a collision that results in a reaction. This requires the particles to possess a minimum amount of energy known as the Activation Energy ($E_a$).

Activation Energy ($E_a$): The minimum energy required by colliding particles to break existing bonds and initiate a chemical reaction.

Factors affecting Reaction Rate: 1. Concentration: Higher concentration increases the number of particles per unit volume, leading to more frequent collisions. 2. Pressure (for gases): Increasing pressure crowds gas particles together, increasing collision frequency.

Temperature: Increasing temperature increases the average kinetic energy of particles. This leads to more frequent collisions and, more importantly, a higher proportion of particles having energy $\ge E_a$.

Surface Area: For solid reactants, increasing surface area (by grinding or crushing) exposes more particles to the other reactant, increasing the frequency of collisions.

Catalysts: A catalyst increases the rate of reaction by providing an alternative reaction pathway with a lower activation energy ($E_a$), allowing more collisions to be successful without the catalyst being consumed.