Rate of Reaction and Equilibrium
Collision theory
A reaction happens when particles collide with enough energy (the activation energy). More frequent and more energetic collisions → faster rate.
Factors that increase rate
| Factor | Why |
|---|---|
| Higher temperature | Particles move faster, collide more often and harder |
| Higher concentration / pressure | More particles in a space → more collisions |
| Larger surface area (smaller pieces) | More area exposed to collide |
| Catalyst | Lowers activation energy; not used up |
Measuring rate
Measure gas volume produced or mass lost over time.
mean rate = amount of product formed ÷ time
= amount of reactant used ÷ timeA steeper graph = faster rate; the line levels off when a reactant runs out.
Reversible reactions & equilibrium
Some reactions go both ways (⇌). In a closed system they reach dynamic equilibrium — forward and backward rates are equal, concentrations stay constant.
- Le Chatelier's principle: change a condition and the equilibrium shifts to oppose the change (higher-tier).
Exam tip
A catalyst speeds a reaction by providing a route with lower activation energy and is not used up. The gradient of a rate graph shows how fast the reaction is going.