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 effects on reaction rate

When a solid reacts with another substance, the other reacting particles can collide with the solid at its exposed surface. A reaction occurs when collisions are successful.

Breaking the same amount of solid into smaller pieces exposes more of the solid. Its total volume remains the same, but its total exposed surface area increases. This means its surface area to volume ratio also increases.

The larger exposed surface allows more reacting particles to collide with the solid each second. This increases the number of successful collisions per second, so the reaction rate increases.

Larger pieces of the same solid have a lower surface area to volume ratio and less exposed surface for collisions. Under otherwise unchanged conditions, they have fewer successful collisions per second and react more slowly.