Collision Theory and Reaction Rate
Collision Theory
Collision theory explains the conditions for a chemical reaction to occur between reactants. It states that successful reactions occur when reactants collide with enough kinetic energy and in the correct orientation to break the bonds and allow the formation of products to occur. Within the scope of collision theory, the rate of reaction is a result of three main factors:
- Rate of collision
- Activation energy
- Molecular orientation
Rate of collision: For a chemical reaction to occur, reacting atoms or molecules have to come in contact with each other. The frequency at which this occurs is the rate of collision. The higher the rate of collision, the higher the reaction rate.
Activation energy: The reactants require a minimum amount of energy for a chemical reaction to occur. This is known as the activation energy.
Molecular orientation: The orientation of reactants can impact the way bonds interact and break. When they collide in the right orientation, there is a higher chance that bonds will break and allow products to form.
Factors Which Impact Collision Rate
Temperature: Higher temperatures result in an increase in the kinetic energy of the particles in a system. When the particles have higher kinetic energy, they will move around quicker and this will result in a higher rate of collision.
Concentration: Higher concentrations results in more particles in a given volume. Higher concentrations result in a higher rate of collision
Pressure and Volume: Pressure and volume are inversely proportional. Higher pressures (or smaller volumes) result in at higher rate of collision.
Factors Which Impact Activation Energy
Activation energy is the minimal amount of energy required for a reaction to take place. It is also illustrated as the enthalpy difference between the reactants and the activation complex:
If particles do not have enough energy to overcome the activation energy, a reaction will not take place. If the activation energy is lower, then more particles may be able to collide with sufficient energy to form products.
In a chemical system or reaction, both the forward and reverse reactions are considered.
- If the activation energy in either direction is too large than that reaction is not likely to proceed.
- If the activation energy in both directions is small than the reaction may occur in both directions simultaneously as some particles will have enough kinetic energy for a successful collision.
This is illustrated in the diagram below:
A Simple Mistake
In any chemical system, it is easy to incorrectly think that all the particles have the same energy at any given time and to then consider if the reaction occurs or not.
In actuality, the particles in the system have a range of energies known as the molecular energy. We can illustrate this along a molecular energy distribution curve:
We can mark the activation energy on this curve and this helps us visualise how many particles will or will not have the required energy to overcome the activation energy.
These graphs are also particularly useful for helping us visualise the impact of temperature changes or the addition of catalysts.
Increasing the Temperature – This shifts the molecular energy distribution curve to the right and results in more particles having the minimum energy to overcome the activation energy:
Impact of Catalysts – Catalysts reduce the activation energy of the reaction. Whilst the molecular energy distribution curve does not shift, the number of particles with the minimum energy to overcome the activation energy has increased: