Why does the rate of reaction change with the energy of collisions?
There are more successful collisions that overcome the activation energy barrier when the collisions have more energy.
In every reaction there is an activation energy barrier. This activation energy barrier is the energy required to break the first bonds in a reaction.
Particles reacting must have enough energy to overcome this activation energy before they are able to fully react.
For example, the activation energy for the reaction of carbon monoxide and nitrogen dioxide to form carbon dioxide and nitrogen monoxide is 134 kJ/mol
CO(g) + NO₂(g) → CO₂(g) + NO(g)
For a successful reaction, the colliding molecules must have a total kinetic energy of 134 kJ/mol.
If the particles in a reaction have more energy (kinetic energy - they move faster) they will have more energetic collisions.
Having more energetic collisions means that when the particles collide they will have enough energy to overcome the activation energy barrier and bonds will successfully be broken.
This links to a faster reaction rate, as there are more successful collisions per second, and the rate that products are produced will be faster.
To increase the energy of collisions the temperature could be increased.