Collision theory and rates of reaction

In a nutshell

Particles require enough kinetic energy to react, this is called the activation energy. If they don't reach this amount of energy, they won't react.

Collision theory

definition

Particles constantly move in liquids and gases. For a reaction to occur, they must be in the right direction when they collide with at least a specific minimum amount of kinetic energy.

Particles require a specific minimum amount of kinetic energy to react, this is called the activation energy. This means that they need that much energy to break bonds in order to start the reaction.

The value of the activation energy depends on the reaction. Some reactions have a low activation energy, meaning the reaction can happen easily. However, some reactions have a high activation energy, this means it needs to be heated so the particles gain enough energy to react.

This information can all be shown in a reaction profile diagram. The peak will be the activation energy and all particles that have reached or surpassed this point will be able to react.

Before the peak, it will show the reactants' bonds being stretched to try to react. After the peak, it will show new bonds have formed, releasing energy, to form the products.

Chemistry; Kinetics I; KS5 Year 12; Collision theory and rates of reaction

The x-axis represents the progress of the reaction. At the right end, the reaction is finished and so the products are formed. The y-axis represents enthalpy, this is the energy of the particles. The enthalpy of the reactants and products are different because they are different molecules.

Energy of molecules

Not all molecules have the same amount of energy in liquids and gases. There's a distribution of energy. Some molecules have a lot of kinetic energy and some don't have a lot. This can be shown on a Maxwell-Boltzmann distribution curve. This is a graph that shows the distribution of kinetic energy between the molecules.

On the left side is when kinetic energy is the lowest. Before the peak, these molecules are moving very slowly. The peak shows that most of the molecules have that much kinetic energy, their speed is moderate. After the activation energy are the molecules able to react, they are moving the fastest.

Chemistry; Kinetics I; KS5 Year 12; Collision theory and rates of reaction

Every single molecule has energy, so the curve starts at ( $0,0$ ).

Temperature effects

Increasing the temperature will increase the kinetic energy of the molecules. They can move and collide faster. More collisions increases the chance of successful collisions. This also means more molecules will have enough kinetic energy to reach the activation energy and will be able to react. This change can be shown in the Maxwell-Boltzmann distribution curve.

Chemistry; Kinetics I; KS5 Year 12; Collision theory and rates of reaction

The curve shifts to the right as more molecules have higher kinetic energies. The area under the curve remains the same because the total number of molecules hasn't changed.

Reaction rate

The rate of reaction can be affected by temperature, concentration, pressure and a catalyst. Temperature increases the rate of reaction as more molecules have sufficient energy to reach the activation energy to react.

If the concentration increases, the rate of reaction also increases. This is because there are more particles in a given volume so there is less space around them meaning they are more likely to collide frequently. More collisions increases the chances to react.

If the pressure increases, the rate of reaction also increases. This is because there are more particles in a given volume so there is less space around them meaning they are more likely to collide frequently. More collisions increases the chances to react.

The presence of a catalyst also increases the rate of reaction. They provide an alternative reaction pathway with a lower activation energy. This means more molecules will have sufficient energy to reach the activation energy to react. Therefore, the reaction will occur faster.