Rate of Reaction and Equilibrium
Rate of reaction
The rate of reaction is how fast reactants turn into products. Collision theory explains it: reactions happen when particles collide with enough energy (the activation energy) and the correct orientation. Anything that increases the frequency or energy of collisions speeds up the reaction.
Factors affecting rate
| Factor | Why it increases the rate |
|---|---|
| Temperature | Particles move faster → collide more often and with more energy |
| Concentration (of a solution) | More particles in the same volume → more frequent collisions |
| Pressure (of gases) | Squeezes gas particles closer → more frequent collisions |
| Surface area (of a solid) | More exposed particles → more frequent collisions (powder reacts faster than a lump) |
| Catalyst | Provides a route with lower activation energy, so more collisions succeed |
Catalysts
A catalyst speeds up a reaction without being used up and without changing the products. It works by lowering the activation energy. Enzymes are biological catalysts. Because they aren't consumed, only a small amount is needed.
Measuring rate
You can measure rate by following how fast a product forms or a reactant is used up, e.g.:
- Volume of gas produced per unit time (gas syringe).
- Mass lost per unit time (balance, if a gas escapes).
- Time for a cross to disappear (a cloudy precipitate forming).
On a graph of product vs time: a steeper line = faster rate. The line levels off when the reaction finishes (a reactant is used up).
Reversible reactions and equilibrium
Some reactions are reversible — products can react to reform the reactants (shown by ⇌). In a closed system, a dynamic equilibrium is reached where the forward and backward reactions happen at the same rate, so the amounts of reactants and products stay constant (though both reactions are still occurring).
Le Chatelier's principle: if you change the conditions, the equilibrium shifts to oppose that change:
- Increase temperature → shifts in the endothermic direction.
- Increase pressure → shifts to the side with fewer gas molecules.
- Increase concentration of a reactant → shifts towards the products.
Worked example
Why does powdered calcium carbonate react faster with acid than a single lump of the same mass?
- The powder has a much larger surface area, so more particles are exposed to the acid → more frequent collisions → faster rate. ✓
Common mistakes
- Saying temperature only makes particles collide "more often" — it also gives them more energy (more successful collisions).
- Thinking a catalyst is used up or changes the products — it isn't and doesn't.
- Saying equilibrium means the reaction has stopped — both reactions continue at equal rates.
Exam tips
- Explain each factor using collision frequency/energy, not just "it's faster".
- On rate graphs, link steepness to rate and the plateau to the reaction finishing.
- For equilibrium, apply Le Chatelier: the system shifts to oppose the change.
Key facts to remember
- Rate depends on temperature, concentration, pressure, surface area and catalysts (via collision frequency/energy).
- Catalysts lower the activation energy, aren't used up, don't change products.
- Dynamic equilibrium = forward and backward rates equal in a closed system; conditions shift it per Le Chatelier's principle.