Suggested languages for you:

Americas

Europe

|
|

# 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 400 ms-1.

Now, think about how many molecules of gas there are in the surrounding air. At room temperature and pressure, one mole of any gas takes up roughly 24 dm3. That’s 6.022 x 1023 molecules in an area of 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 surrounding air 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 surrounding air filled with damaging nitrous oxides?

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

• We'll begin by exploring the basic principles of collision theory.
• We'll then look at enthalpy diagrams, also known as energy profiles.

## Collision theory definition

Collision theory is an explanation for the rates of many reactions. It proposes two key ideas: molecules must collide with the correct orientation, and sufficient energy, in order for a reaction to occur.

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.

Activation energy is the minimum amount of energy needed to start a chemical reaction. It takes the symbol Ea and is typically measured in kJ mol-1.

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.

You can explore energy profiles in more depth, including transition states, in Chemical Kinetics.

## 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.

If you want to find out more about how these factors change reaction rates, check out Factors Affecting Reaction Rates.

## Collision theory and enzymes

Collision theory can help explain how enzymes work. Enzymes are biological catalysts, meaning that they increase the rate of reaction. There are a few different hypotheses explaining how. One idea is that they could simply lower the reaction's activation energy.

Another idea looks at the shapes of enzymes. Enzymes have specific shapes. Scientists hypothesise that they hold the reactants in just the right position, so when two reactants collide, there is an increased chance of the pair being orientated correctly. This increases the chance of a reaction.

## Collision Theory - Key takeaways

• Collision theory is an explanation for the rates of many reactions.
• Collision theory has two key ideas: molecules must collide with the correct orientation and sufficient energy in order for a reaction to occur.
• The minimum amount of energy needed for a reaction to occur is known as the activation energy.
• We can show activation energy using energy profiles.
• Most collisions are unsuccessful because they have an incorrect orientation or lack enough energy to react.
• Collision theory tells us how we can increase the rate of a reaction. Factors affecting the rate of a reaction include temperature, pressure, concentration, the surface area of the particles, and the presence of a catalyst.

Collision theory is an explanation for the rates of many reactions. It proposes two key ideas: molecules must collide with the correct orientation and sufficient energy in order for a reaction to occur.

There are three important parts to collision theory. First, the reacting substances must collide. Secondly, they must collide with the correct orientation. Thirdly, they must collide with enough energy. If all of this occurs, then the molecules will react.

Collision theory states that molecules must collide with the correct orientation and sufficient energy in order for a reaction to occur.

Collision theory is important because it helps us influence the rate of reaction. By changing how often molecules collide and their average energy, we can increase the rate of a reaction.

## Final Collision Theory Quiz

Question

What is collision theory?

For a reaction to occur, the reactant particles, atoms or molecules, must collide and interact with each other in some way: this is collision theory. Collision theory states that the rate of a chemical reaction is proportional to the number of collisions between reactant molecules. The more often reactant molecules collide, the more often they react with one another, and the faster the reaction is.

Show question

Question

What is an effective collision?

Only a small fraction of collisions result in a chemical reaction: these are called effective collisions. Inversely, ineffective collisions do not result in a chemical reaction.

Show question

Question

How does temperature affect a chemical reaction?

At lower temperatures molecules have less kinetic energy. As the temperature of the molecules increases, so does their kinetic energy, meaning the molecules are moving faster. A slight increase in temperature causes a large increase in the number of molecules with kinetic energy greater than the activation energy, meaning that the rate of reaction increases.

Show question

Question

What is an exothermic reaction?

When energy is transferred out to the surroundings, this is called an exothermic reaction, and the temperature of the surroundings increases. The energy level decreases in an exothermic reaction. This is because heat energy is given out to the surroundings. Because of this, a reaction is also said to be exothermic when the energy of the product(s) is less than that of the reactants.

Show question

Question

What is an endothermic reaction?

When energy is taken in from the surroundings, this is called an endothermic reaction and the temperature of the surroundings decreases. The energy level increases in an endothermic reaction. This is because heat energy is taken in from the surroundings. Because of this, a reaction is also said to be endothermic when the energy of the product(s) is greater than that of the reactants.

Show question

Question

What is enthalpy?

Enthalpy is a measure of the amount of heat absorbed or released during a chemical reaction. It is calculated by subtracting the energy of the reactants from the energy of the products.

Show question

Question

What is a catalyst ?

Catalysis is the process of increasing the rate of a chemical reaction by adding a substance known as a catalyst. Catalysts are not consumed in the reaction and remain unchanged after it.

Show question

Question

How does a catalyst lower the activation energy of a reaction?

A catalyst can lower the activation energy for a reaction by:

• Providing a surface for reactions to take place on. Reactant molecules are then oriented at a favourable angle for collisions to occur, increasing the likelihood of successful collisions, thus speeding up the reaction.
• Reacting with one or more reactants to form intermediates that require lower energy to form the final reaction product, in the process regenerating the catalyst, and increasing the rate of reaction.

Show question

Question

Molecules always react when they collide. True or false?

False

Show question

Question

What are the two main principles of collision theory?

Molecules must collide with the correct orientation and enough energy for a reaction to occur.

Show question

Question

Exothermic reactions have no activation energy. True or false?

False

Show question

Question

What factors affect the rate of reaction?

• Pressure
• Concentration
• Temperature
• Surface area
• Catalysts

Show question

Question

How can we increase the rate of a reaction?

• Increase the pressure of a gaseous system.
• Increase the concentration of a reactant.
• Increase the surface area of the reactants.

Show question

Question

In endothermic reactions, the products have ____ energy than the reactants.

More

Show question

Question

In exothermic reactions, the products have ____ energy than the reactants.

Less

Show question

60%

of the users don't pass the Collision Theory quiz! Will you pass the quiz?

Start Quiz

## Study Plan

Be perfectly prepared on time with an individual plan.

## Quizzes

Test your knowledge with gamified quizzes.

## Flashcards

Create and find flashcards in record time.

## Notes

Create beautiful notes faster than ever before.

## Study Sets

Have all your study materials in one place.

## Documents

Upload unlimited documents and save them online.

## Study Analytics

Identify your study strength and weaknesses.

## Weekly Goals

Set individual study goals and earn points reaching them.

## Smart Reminders

Stop procrastinating with our study reminders.

## Rewards

Earn points, unlock badges and level up while studying.

## Magic Marker

Create flashcards in notes completely automatically.

## Smart Formatting

Create the most beautiful study materials using our templates.