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Collision Theory

Definition

Collision theory explains how chemical reactions occur and why reaction rates differ for different reactions. The theory states that for a reaction to occur, particles must collide with sufficient energy (activation energy) and correct geometric orientation.

Analogy

Imagine playing pool - not every hit results in sinking a ball into a pocket. For that successful shot, you need both enough force (akin to activation energy) and correct alignment (similar to geometric orientation).

Related terms

Transition State Theory: A theory which describes the activated complex during the transition phase of two molecules becoming one product. Like capturing that split second when your pool ball hits another just before it sinks into the pocket.

Catalyst: A substance that increases the rate of a chemical reaction by lowering its activation energy without being consumed itself - much like adding bumpers on pool table pockets making it easier for balls to go in.

Rate Law Expression: An equation showing relationship between the rate of a reaction and concentration of reactants. It's like calculating how your pool game success rate changes based on the number of balls and their positions.