# Collision theory

For a reaction to occur, reactant particles must collide. For a successful collision to occur, there has to be

• sufficient energy to start breaking reactant bonds;
• the proper geometry (arrangement of the atoms).

Recall that temperature is a measure of the average kinetic energy of the particles in a substance.

In this graph, two temperatures $\displaystyle T_{1}$ and $\displaystyle T_{2}$ are shown. For each one, there are particles with all different kinetic energies, but there are more particles at some kinetic energies then there are at others. The area below each curve is the same (they have the same total number of particles), but with $\displaystyle T_{2}$, the shape is shifted to the right, so there are more particles with high kinetic energies, meaning the temperature of the object is greater.

For a chemical reaction to occur, a minimum energy barrier must be surpassed. This is called the activation energy ($\displaystyle E_{\text{a}}$). Because increasing temperature increases the number of collisions (more kinetic energy leads to more frequent collisions) and the fraction of collisions that are successful (because the kinetic energies are more likely to reach $\displaystyle E_{\text{a}}$), a 10 ºC rise in temperature can actually double or triple the rate.

If a collision is successful, then a higher energy, unstable, transitory particle if formed: the activated complex. This is a middle step (neither reactant nor product) that forms, has partial bonds, is highly reactive, exists for a very short duration, and then either breaks down to form products or re-forms the reactants. It exists between the time when reactant bonds break (exothermic) and product bonds form (endothermic), which is why it has so much energy.