Collision theory
A reaction happens only when reactant particles collide with enough energy and in the correct direction. The least energy needed for a successful collision is the activation energy. Anything that increases the frequency or the energy of these effective collisions will raise the rate of reaction. Four factors are commonly studied.
The four factors
- Concentration — a higher concentration packs more particles into the same volume, so collisions happen more often and the rate increases.
- Temperature — heating makes particles move faster and gives more of them energy equal to or above the activation energy, so effective collisions become much more frequent.
- Surface area — breaking a solid into a powder exposes far more particles to attack, so a large surface area (small pieces) speeds up the reaction.
- Catalyst — a catalyst provides an alternative path with a lower activation energy, so a greater fraction of collisions succeed. It speeds up the reaction but is not used up itself, and only a small amount is needed.
Example
Adding manganese(IV) oxide to hydrogen peroxide makes it decompose quickly into water and oxygen. The manganese(IV) oxide is a catalyst — it is recovered unchanged at the end and can be used again.
Key idea
Higher concentration, higher temperature and larger surface area all increase the frequency of effective collisions; a catalyst increases the proportion of collisions that succeed by lowering the activation energy.
What a catalyst does not change
A catalyst only changes the speed of a reaction, not the amount of product formed. The same mass of product is made in the end — it is simply made faster. This is why catalysts are so valuable in industry, where they save time and energy and allow reactions to run at lower, cheaper temperatures. Catalysts are also often specific, meaning one catalyst speeds up one particular reaction and has no effect on others.
Remember
Powder reacts faster than a lump of the same mass because the powder has a much larger total surface area.