Collision Theory and Factors Affecting Reaction Rates

Why some collisions produce a reaction

The rate of reaction describes how quickly reactants are used up or products are formed. To explain why changing the conditions changes this rate, we need to consider what happens to the reacting particles.

Particles in gases and solutions move around and collide. According to collision theory, a chemical reaction can occur only when reactant particles collide with sufficient energy. Simply meeting is not enough: if the particles have too little energy, they separate without reacting.

The activation energy is the minimum energy that reacting particles must have for a collision to result in a reaction. A collision with sufficient energy can be successful: atoms rearrange to form products. Different reactions have different activation energies.

The important link is between reaction rate and successful collisions per second. More successful collisions each second means that products form more quickly. Conditions can increase the rate by making collisions more frequent, making a greater proportion of collisions energetic enough to react, or both.

Concentration: more particles in the same volume

The concentration of a solution tells us how much dissolved substance there is in a given volume. A more concentrated reactant solution contains more reactant particles per unit volume.

Compare two equal volumes of acid, one dilute and one more concentrated, reacting with identical pieces of the same solid. In the more concentrated acid, more acid particles are available to collide with the solid's surface. Collisions happen more frequently, so there are more successful collisions per second and the reaction is faster.

At the same temperature, increasing concentration does not make the particles more energetic. It increases the number of opportunities for a successful collision, rather than increasing the proportion of collisions with sufficient energy.

Diluting a reactant solution has the opposite effect. There are fewer reactant particles per unit volume, so successful collisions occur less frequently and the reaction is slower.

Pressure: bringing reacting gas particles closer together

Pressure affects the rate of reactions involving gases. Consider compressing reacting gases into a smaller volume while keeping the temperature the same. The number of particles has not changed, but there are now more reactant particles per unit volume.

The particles are closer together, so they collide more frequently. This produces more successful collisions per second and increases the reaction rate. Reducing the pressure by allowing the gases to occupy a larger volume has the opposite effect.

This is similar to increasing the concentration of a solution: both changes put more reacting particles into each unit of volume. Neither explanation depends on an increase in particle energy when temperature is unchanged.

Temperature: more frequent and more energetic collisions

Increasing temperature increases the average kinetic energy of the particles. In gases and solutions, the particles move faster and collide more frequently.

There is also a second effect. The particles have a range of energies, and at a higher temperature a greater proportion of collisions have energy at least equal to the activation energy. More collisions can therefore lead to a reaction.

Together, the increased collision frequency and the greater proportion of sufficiently energetic collisions produce more successful collisions per second. The reaction rate increases. Heating does not lower the activation energy; it enables more particles to meet that energy requirement.

Cooling reverses both effects. Particles move more slowly, collisions become less frequent, and a smaller proportion have sufficient energy to react.

Surface area: exposing more of a solid

When a solid reacts with a solution or gas, collisions occur at its exposed surface. Breaking the same mass of solid into smaller pieces increases the total surface area available to the other reactant.

For example, calcium carbonate powder reacts faster with hydrochloric acid than an equal mass of calcium carbonate chips under otherwise identical conditions. More solid particles are exposed, allowing more collisions and therefore more successful collisions each second. The solid particles do not gain extra energy simply because the pieces are smaller.

Using collision theory to predict a change

A useful explanation links the changed condition to what happens to particles, then to successful collisions per second, and finally to the reaction rate.

Suppose a reaction between calcium carbonate chips and hydrochloric acid is repeated using more dilute acid, with temperature and the chips unchanged. The prediction is a slower reaction: there are fewer acid particles per unit volume, so collisions with the chips occur less frequently and there are fewer successful collisions per second.

If instead the reaction is cooled while concentration and surface area stay the same, the prediction is also a slower reaction—but the explanation includes energy as well as collision frequency. The particles move more slowly, and fewer collisions have enough energy to overcome the activation energy.

Collision theory

Reactant particles must collide with energy at least equal to the activation energy to react. Reaction rate increases when there are more successful collisions per second.

Effects of increasing each factor