Collision Theory

Have you ever wondered how fast gas molecules move? It depends on lots of factors, such as their mass and the temperature. But as an example, scientists estimate that oxygen molecules in the air travel at room temperature at over .

Now, think about how many molecules of gas there are in the air around you. At room temperature and pressure, one mole of any gas takes up roughly . That’s molecules in an area just 60 x 20 x 20 cm!

If we put the two together, we get a very large number of molecules moving extremely quickly in a relatively small area. From this we can predict that there are a lot of collisions between the molecules each second.

But wait, the air around us is full of nitrogen and oxygen molecules. They can react to form harmful nitrous oxides! What if the molecules collide? Will they react? If they do, why isn’t the air around us filled with damaging nitrous oxides?

Before we take that thought any further, we need to look at collision theory.

• This article is about collision theory in physical chemistry.
• We'll begin by exploring the basic principles of collision theory.
• We'll then look at enthalpy diagrams, also known as energy profiles.

The principles of collision theory

As we mentioned above, collision theory is an explanation for the rate of reaction. It helps us predict whether two molecules will react or not.

Collision theory has two underlying components:

• Orientation

• Energy

First of all, let’s look at orientation.

Orientation

Molecules must firstly meet with just the right orientation in order for a collision to occur. Take the reaction between the gases hydrogen bromide and ethene, for example. This forms bromoethane. This reaction involves the hydrogen atom joining on to 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 collides with the double bond, or the hydrogen hits one of the carbon atoms instead of the double bond, nothing will happen - a reaction won’t occur.

Molecules must collide with the correct orientation to react. Anna Brewer, StudySmarter Originals

Energy

However, the correct orientation isn't the end of the story. In order to react, colliding molecules also need enough 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 on the species involved and the reaction itself. This energy is known as the activation energy.

This means 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 , 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:

The reactants start off with a medium amount of energy. In order to react, they need more energy, as shown by the peak in the graph - enough energy to break the bonds in the reactants and reach the transition state. This energy needed 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.

Because this is an exothermic reaction, the products have a lower energy than the reactants. However, we still come across the energy barrier in endothermic reactions, as shown below:

In this case, the only difference is that the products have a higher energy level than the reactants. 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.

A flow chart for collision theory. Anna Brewer, StudySmarter Originals

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 and 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 is no 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. 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.