Concentration and Pressure
Simple Explanation
Increasing the concentration of reactants in solution (or the pressure of reacting gases) packs more particles into the same volume, which increases how often they collide — and therefore increases the reaction rate.
Why Do We Need It?
This is one of the most direct, controllable ways to speed up or slow down a reaction in both industrial and laboratory settings, and directly follows from collision theory rather than being a separate, unrelated rule to memorize.
Why Does This Work?
With more particles packed into the same volume, the average distance between particles shrinks, which makes collisions between them more frequent — since reaction rate depends on the number of successful collisions per second, more total collisions (at the same fraction of successful ones) directly means a faster rate.
Real-Life Example
Pure oxygen speeds up combustion
A glowing splint bursts back into flame when placed in a test tube of pure oxygen, but not in ordinary air.
Pure oxygen has a much higher concentration of O₂ than air (which is only about 21% oxygen), so combustion reactions collide far more frequently and proceed dramatically faster — vividly demonstrating the effect of concentration on rate.
Practice
Why does increasing the pressure of a gaseous reaction mixture increase its rate?
MediumCommon mistake
Thinking concentration changes the activation energy of a reaction — it does not; concentration only affects how FREQUENTLY particles collide, not what fraction of collisions have enough energy to succeed.
Quick Review
- Higher concentration/pressure → more frequent collisions → faster rate.
- This works by increasing collision frequency, not by lowering Ea.
- Pressure affects gas-phase reactions the same way concentration affects solutions.