Chapter 6: Chemical Reactions — Part 3

Topic 6.2b · Collision theory and catalysts

Collision Theory

For a chemical reaction to occur, reactant particles must collide successfully with each other with energy equal to or greater than the activation energy (E_a).
The rate of reaction increases when there are more frequent successful collisions between reactant particles.
Explanations for Rate Factors:
- Concentration and Pressure: Increasing these results in more reactant particles in the same volume, leading to more frequent collisions.
- Surface Area: Increasing the surface area (e.g., fine powder) exposes more reactant particles at the surface, leading to more frequent collisions.
- Temperature: Increasing temperature provides particles with more kinetic energy, meaning they move faster and collide more frequently and more successfully.

Exam Traps

Do not say “particles collide more often” alone for temperature — also state more collisions have energy ≥ E_a.
Do not claim all collisions are successful — only those with sufficient energy (and correct orientation at Core level).

A catalyst increases the rate of reaction but remains unchanged (not used up) at the end.
Biological catalysts are known as enzymes.
Using collision theory, a catalyst works by decreasing the activation energy (E_a) needed for the reaction, providing an alternative pathway.
Because E_a is lowered, a higher proportion of reactant particles have sufficient energy to collide successfully.

Exam Traps

Do not say a catalyst is used up — it speeds up attainment of equilibrium but is not consumed.
Do not say a catalyst shifts equilibrium yield — it increases rate only; position of equilibrium is unchanged.

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