Rate of Reaction & Collision Theory (Cambridge (CIE) AS Chemistry): Revision Note

Rate of Reaction & Collision Theory

Reaction rate

• The rate of a reaction is the speed at which a chemical reaction takes place and has units mol dm^-3 s^-1

• The rate of a reaction can be calculated by:

Rate of reaction = $\frac{\mathbf{change}\mathbf{in}\mathbf{the}\mathbf{amount}\mathbf{of}\mathbf{reactants}\mathbf{or}\mathbf{products}\left(\mathrm{mol}{\mathrm{dm}}^{-3}\right)}{\mathbf{time}\left(s\right)}$

Collision theory

• The collision theory states that for a chemical reaction to take place, the particles need to collide with each other in the correct orientation and with enough energy

• The minimum energy that colliding particles must have for a collision to be successful and a reaction to take place is called the activation energy (E_a)

Collision theory table

• An ineffective collision is when particles collide in the wrong orientation or when they don’t have enough energy and bounce off each other without causing a chemical reaction

 Effective and ineffective collisions

Increase in reaction rate

• The collision frequency is the number of collisions per unit time

• When there are more collisions per unit time, the number of particles with energy greater than the E_a increases

• This causes an increase in the rate of reaction

• A catalyst is a substance that increases the rate of reaction without taking part in the chemical reaction by providing the particles with an alternative mechanism with a lower activation energy

How catalysts affect reaction pathways

Concentration

• The more concentrated a solution is, the greater the number of particles in a given volume of solvent

• An increase in concentration causes an increased collision frequency and therefore an increased rate of reaction

How increasing concentration affects collisions

Pressure

• An increase in pressure in reactions that involve gases has the same effect as an increased concentration of solutions

• When the pressure is increased, the molecules have less space in which they can move

• This means that the number of effective collisions increases due to an increased collision frequency

• An increase in pressure therefore increases the rate of reaction

How increasing pressure affects collisions