Temperature, concentration, surface area, pressure and catalysts — how each factor affects rate and why, using collision theory.
Rates of reaction questions appear on every GCSE Chemistry paper. They test both your understanding of collision theory and your ability to interpret graphs of concentration or volume against time. The good news is that the same explanations apply every time — once you understand collision theory properly, you can explain any factor's effect on rate in a consistent, mark-scoring way.
For a chemical reaction to occur, particles must collide with each other with sufficient energy to break existing bonds and form new ones. The minimum energy required for a successful collision is called the activation energy.
Two conditions must both be met for a collision to be successful:
Reaction rate depends on the frequency of successful collisions per unit time. Anything that increases the frequency of collisions, or the proportion of collisions that have sufficient energy, increases the reaction rate.
Every explanation of a factor affecting rate must reference collision theory. "Temperature increases rate" scores zero. "Increasing temperature gives particles more kinetic energy, so they collide more frequently and a greater proportion of collisions have energy greater than or equal to the activation energy, so the rate increases" scores full marks.
Increasing temperature increases reaction rate significantly. Two things happen simultaneously:
The second point is actually more significant than the first. A 10°C rise in temperature roughly doubles the reaction rate for many reactions — not because particles collide twice as often, but because the proportion of successful collisions increases dramatically.
On a graph of rate against time, a higher temperature produces a steeper initial gradient (faster initial rate) and the reaction is completed sooner. The final amount of product is the same — temperature doesn't change how much product forms, only how quickly.
Increasing the concentration of a dissolved reactant increases the reaction rate. With more solute particles in the same volume of solution, there are more particles per unit volume. Particles are closer together and therefore collide more frequently. The activation energy is unchanged, so the same proportion of collisions are successful — but there are simply more collisions happening per second.
❌ Common error: saying higher concentration means particles "have more energy". Concentration does not affect the energy of particles — only temperature does. Concentration affects frequency of collisions only. Stating that concentration increases energy will lose marks.
For reactions involving a solid reactant, increasing surface area (by using smaller pieces or a powder instead of a lump) increases reaction rate. Breaking a solid into smaller pieces exposes more of its surface to the surrounding reactant particles. More surface area means more particles at the surface available for collision — collision frequency increases and rate increases.
This is why flour mills and coal mines can be dangerous — fine powder suspended in air has an enormous surface area and can combust explosively with a single spark, even though the same material in a solid block burns much more slowly.
Increasing pressure on a gaseous reaction increases the reaction rate. Higher pressure squeezes the same number of gas particles into a smaller volume. The particles are more crowded, so they collide more frequently. Rate increases.
Pressure only affects reactions involving gases — it has no effect on reactions in solution.
A catalyst increases reaction rate without being used up in the reaction. This is a crucial distinction — catalysts can be recovered unchanged at the end of a reaction.
Catalysts work by providing an alternative reaction pathway with a lower activation energy. With a lower activation energy, a greater proportion of collisions have sufficient energy to be successful — so the rate increases even though the temperature and concentration are unchanged.
On an energy profile diagram, a catalyst lowers the peak of the curve (the activation energy) but does not change the energy levels of the reactants or products — it does not affect the overall energy change of the reaction.
Homogeneous catalysts are in the same physical state as the reactants (e.g. both in solution). Heterogeneous catalysts are in a different physical state — typically a solid catalyst used with liquid or gaseous reactants. Heterogeneous catalysts work by providing a surface on which reactants adsorb, bringing them into close contact. Examples: iron catalyst in the Haber process, platinum/rhodium in catalytic converters, vanadium(V) oxide in the Contact process.
Rate can be measured in several ways depending on the reaction:
The AQA rates of reaction specification is at the AQA GCSE Chemistry specification page.
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