Collision Theory

Collision theory definition

Collision theory is the reason why we can live in an atmosphere full of nitrogen and oxygen molecules, without worrying about the danger of nitrous oxides. It helps us analyse the rate of reactions and work out how best to optimise a chemical process.

Collision theory principles

Collision theory has two underlying principles:

• Orientation

• Energy

First, let’s look at orientation.

Orientation

Molecules must firstly meet with the correct orientation in order for a collision to occur. Take the reaction between hydrogen bromide and ethene, for example. This forms bromoethane. The reaction involves the hydrogen atom joining onto the C=C double bond. To do this, the hydrogen end of the hydrogen bromide molecule must approach and collide with the double bond in ethene. If the bromine atom collides with the double bond, or the hydrogen atom hits one of the carbon atoms or C-H single bonds instead of the C=C double bond, nothing will happen - a reaction won’t occur.

Fig. 1 - Molecules must collide with the correct orientation to react

Energy

However, the correct orientation isn't the end of the story. In order to react, colliding molecules also need sufficient energy. This is because reactions all firstly involve breaking bonds, which is an endothermic process - it requires energy. The amount of energy needed varies depending both on the species involved and the reaction itself, and is known as the activation energy.

Collision theory tells us that even if molecules collide with the perfect orientation, they’ll only react if they meet or exceed the activation energy. If they don't have enough energy, they'll simply bounce off each other.

We can see the activation energy of a reaction using enthalpy diagrams. These are also known as energy profiles. Here's an example of an energy profile for an exothermic reaction:

Fig. 2 - An example of an energy profile for an exothermic reaction

Note the following:

• The x-axis shows the extent of the reaction, whilst the y-axis shows the energy of the species involved.
• In order to react, the reactants need to gain energy, as shown by the peak in the graph - they need enough energy to break the bonds in the reactants and reach the transition state. This energy is the activation energy, and we often call the peak in the graph the activation energy barrier.
• The molecules then lose energy as they form new bonds, turning into the products.

Because this is an exothermic reaction, the products have less energy than the reactants. Overall, the reaction releases energy. In contrast, in endothermic reactions, the products have more energy than the reactants, and overall, the reaction absorbs energy. However, we still come across the energy barrier in endothermic reactions, as shown below:

Fig. 3 - An example of an energy profile for an endothermic reaction

Note its similarities to an exothermic reaction:

• The x-axis still shows the extent of the reaction, and the y-axis shows the energy of the species involved.
• In order to react, the reactants need to gain energy in order to overcome the activation energy barrier.
• The molecules then lose energy as they form new bonds.

In the case of an endothermic reaction, the only difference is that the products have a higher energy level than the reactants. Overall, the reaction absorbs energy. However, we still need activation energy to get the reaction started.

Will they react?

We can think of the whole collision and reaction process like one big flow chart. Take two molecules. Firstly, do they collide? Secondly, are they orientated correctly? Thirdly, do they have enough energy? If the answer is 'no' at any stage, a reaction won’t occur.

Fig. 4 - A flow chart for collision theory

Collision theory example

Let’s go back to the problem at the start of the article. Although there may be many collisions between oxygen and nitrogen molecules in the air each second, there are hardly any reactions between them. Collision theory gives us a reason why. In this case, almost none of the molecules have sufficient energy to react. A reaction between nitrogen and oxygen would firstly require breaking the strong N≡N and O=O bonds within the molecules. This requires a lot of energy. In most cases, the nitrogen and oxygen molecules don't have enough energy to get over the activation energy barrier, so there isn't a reaction.

Collision theory and rate of reaction

We now know that in order to react, molecules must collide with the correct orientation and sufficient energy. We call any collisions that result in a reaction successful collisions or effective collisions. The more successful collisions we have per second, the faster the rate of reaction.

It is important to remember that only a small proportion of collisions result in a reaction. Most collisions are unsuccessful - they are either orientated incorrectly or don't have enough energy.

How can we use collision theory to increase the rate of a reaction? Well, we can't change the orientation of the molecules when they collide. However, we can influence how often they collide, and their overall energy requirements. We can do this in the following ways.

• Increasing the temperature of a system increases the kinetic energy of all the molecules within it. The molecules move faster, resulting in more collisions, and on average they have higher energy. This means that the molecules have an increased chance of meeting the activation energy requirements when they collide.

• Increasing the concentration of the reactants in a system, and increasing the pressure of a gaseous system, both increase the number of collisions per second.

• Increasing the surface area of the solid reactants increases the number of exposed particles that are able to react with a surrounding liquid or gas. This also increases the number of collisions.

• Adding a catalyst reduces the activation energy of the reaction. This means that an increased number of molecules meet or exceed the activation energy requirements when they collide.