Collision theory and activation energy

For a chemical reaction to occur, reacting particles must meet. Collision theory explains reaction rate by considering how often particles collide and how many of those collisions lead to a reaction.

Not every collision is successful. The particles must collide with enough energy to overcome the activation energy: the minimum energy needed for the reaction to occur. If they collide with less than this energy, they do not react in that collision.

Successful collisions and reaction rate

The reaction rate depends on both the frequency of collisions and the number of collisions with sufficient energy. More collisions can provide more opportunities for reaction, but a collision without enough energy is not successful. A faster reaction has a greater number of successful collisions each second.

Proportional reasoning in collision theory

A reaction happens when reacting particles collide successfully. The more successful collisions that occur each second, the faster the reaction. In a simple proportional model, if the number of collision opportunities each second increases and the same fraction of collisions remains successful, the reaction rate increases by the same factor.

For example, suppose there are twice as many particles of one reactant in the same volume, while the amount of the other reactant and the other conditions stay the same. In the simple model, particles of the two reactants have about twice as many opportunities to collide each second. If the same fraction of these collisions is successful, there are about twice as many successful collisions each second, so the reaction rate is about twice as high.

The relationship can also be expressed as a ratio. If collision opportunities increase in the ratio $$3:2$$, the model predicts that the reaction rate increases in the same ratio: it becomes $$1.5$$ times its original value.

This is a simplified model, not a rule that every change in conditions produces an equal change in rate. The proportional prediction depends on the fraction of collisions that are successful remaining unchanged.

Concentration and pressure effects on reaction rate

Reactant particles must collide for a reaction to occur. When there are more reactant particles in a given volume, collisions happen more frequently. This increases the number of successful collisions per second, making the reaction faster.

Concentration of solutions

Increasing the concentration of a reactant in solution means there are more of its particles in the same volume of solution. Other reactant particles are therefore more likely to collide with them each second. More frequent collisions lead to more successful collisions per second, so the reaction rate increases.

Pressure of reacting gases

Increasing the pressure of reacting gases by compressing them packs their particles more closely together. There are then more particles in each unit of volume, so collisions occur more frequently. The number of successful collisions per second increases, and the reaction becomes faster.

Not every collision produces a reaction. The rate increases because more frequent collisions give more opportunities for successful collisions.

Temperature effects on reaction rate

Raising the temperature gives reacting particles more kinetic energy. They move faster and collide more often, so the collision frequency increases.

A collision leads to a reaction only if the particles collide with enough energy to overcome the activation energy: the minimum energy needed for the reaction. At a higher temperature, collisions are more energetic, so a greater proportion have enough energy to be successful.

The reaction rate increases because particles collide more frequently and a greater proportion of those collisions are successful. This produces more successful collisions per unit time.

Surface area and reaction rate

A solid reactant can react only at its surface, where particles of another reactant can reach it. Particles inside the solid are not exposed to the other reactant.

Breaking the solid into smaller pieces exposes surfaces that were previously inside it. For the same amount of solid, smaller pieces have a greater total surface area than a large lump. A powder exposes even more surface area because it consists of many small particles.

With more surface exposed, particles of the other reactant can collide with the solid at more places. More collisions can therefore occur each second, increasing the reaction rate. A large lump has less exposed surface area, so fewer collisions occur each second and the reaction is slower.