In simple terms
A friendly intro before the formal notes — no formulas yet.
The Energy Hurdle for Reactions
For a reaction to happen, colliding particles must have enough energy to overcome a minimum barrier, known as the activation energy. Increasing the temperature gives more particles the energy they need to clear this hurdle, making the reaction go faster.
Imagine trying to kick a football over a high wall. The wall's height is the activation energy (). Most of your kicks might not be powerful enough to get the ball over. If you get a surge of energy (like increasing the temperature), you can kick harder and more of your attempts will succeed, getting the ball over the wall. A catalyst is like finding a lower section of the wall to kick the ball over - it's much easier, so more of your kicks are successful even without extra energy.
- 1
Particles must collide with energy ≥ activation energy Ea to react.
- 2
Higher temperature → greater fraction of collisions exceed Ea (Boltzmann distribution).
- 3
Rate approximately doubles per 10°C rise for many reactions (rule of thumb).
- 4
Catalysts provide alternative route with lower Ea - not consumed overall.
Explore the concept
Use the live diagram, PhET or GeoGebra sim, and synced steps — play it, drag controls, or tap a step.
Molecules have a spread of energies — few very slow or very fast, most in the middle.
At a glance — side by side
Compare key properties side by side — ideal for exam contrasts.
Comparison of Increasing Temperature vs. Adding a Catalyst
| Feature | Increasing Temperature | Adding a Catalyst |
|---|---|---|
| Effect on Activation Energy ($E_a$) | No effect. The activation energy remains unchanged. | Lowers the activation energy by providing an alternative reaction pathway. |
| Effect on Particle Kinetic Energy | Increases the average kinetic energy of all particles. | No effect. The average kinetic energy of particles remains the same. |
| Effect on Boltzmann Distribution Curve | The curve flattens and shifts to the right. The area representing particles with E ≥ increases. | The curve itself does not change. A new, lower activation energy () is marked on the axis, increasing the proportion of particles that can react. |
| Mechanism of Rate Increase | Increases the frequency and energy of collisions, so a greater proportion of collisions are successful. | Increases the proportion of successful collisions by lowering the energy requirement, without changing collision frequency or particle energy. |
Effect on Activation Energy ($E_a$)
Increasing Temperature
Adding a Catalyst
Effect on Particle Kinetic Energy
Increasing Temperature
Adding a Catalyst
Effect on Boltzmann Distribution Curve
Increasing Temperature
Adding a Catalyst
Mechanism of Rate Increase
Increasing Temperature
Adding a Catalyst
Full topic notes
Formal explanation with the rigour you need for the exam.
Revisiting Collision Theory and Energy
For a chemical reaction to occur, reactant particles must collide. However, not all collisions result in a reaction. Two crucial conditions must be met: the particles must collide with the correct orientation, and they must possess sufficient energy. This minimum required energy is known as the activation energy. Increasing the temperature of a system increases the average kinetic energy of all particles. This leads to two effects: particles move faster, increasing the frequency of all collisions, and more importantly, a significantly greater proportion of these collisions possess energy equal to or greater than the activation energy. The second effect is the primary reason for the substantial increase in reaction rate observed with a rise in temperature.
Two conditions for a successful collision: correct orientation and sufficient energy.
Increasing temperature increases the average kinetic energy of particles.
This results in more frequent collisions and, more significantly, a higher proportion of effective (successful) collisions.
Activation Energy (Ea) - The Energy Barrier
Activation energy, symbolised as , is the minimum amount of kinetic energy that colliding particles must have for a reaction to occur. It can be visualised as an energy barrier that must be overcome, similar to pushing a boulder over a hill before it can roll down the other side. This energy is required to break existing chemical bonds in the reactants, allowing new bonds to form in the products. Reactions with a high activation energy will have a slow rate at a given temperature because only a small fraction of particles possess enough energy to overcome the barrier. Conversely, reactions with a low proceed more rapidly as a larger fraction of collisions are successful.
Activation energy () is the minimum kinetic energy required for a successful collision.
It represents an energy barrier that must be surmounted.
The magnitude of determines the rate of reaction; high means a slow rate, low means a fast rate.
The value of is specific to a particular reaction and is unaffected by temperature.
The Boltzmann Distribution of Molecular Energies
At any given temperature, the particles in a gas or solution do not all have the same kinetic energy. Their energies are described by the Boltzmann distribution. A plot of the number of particles against their kinetic energy shows a characteristic curve that starts at the origin (no particles have zero energy), rises to a peak representing the most probable energy, and then tails off asymptotically towards the x-axis (a few particles have very high energies). The total area under the curve represents the total number of particles in the sample. This distribution is fundamental to understanding why temperature has such a profound effect on reaction rate, as it allows us to visualise the proportion of particles that possess the required activation energy.
The Boltzmann distribution shows the range of kinetic energies of particles at a constant temperature.
The y-axis is 'Number of particles' and the x-axis is 'Kinetic energy'.
Key features: starts at (0,0), has a peak (most probable energy), and is asymptotic to the energy axis.
The area under the curve is constant for a fixed number of particles.
How Temperature Changes the Boltzmann Distribution
When temperature is increased from T1 to a higher temperature T2, the Boltzmann distribution curve changes shape. The peak of the curve becomes lower and shifts to the right, indicating that the most probable energy has increased. The curve also becomes broader and more spread out, showing a wider range of energies. Although the shape changes, the total area under the curve remains constant because the total number of particles has not changed. By marking the activation energy () on the energy axis, we can see that the area under the curve to the right of is significantly larger at T2 than at T1. This shaded area represents the number of particles with sufficient energy to react, explaining the dramatic increase in reaction rate.
Increasing temperature lowers the peak of the curve and shifts it to the right.
The curve becomes more spread out at higher temperatures.
The area under the curve to the right of represents the number of particles capable of reacting.
This area increases significantly with a small increase in temperature, leading to a much higher frequency of successful collisions.
When drawing or interpreting Boltzmann distribution curves, always label the axes correctly ('Number of particles' vs 'Kinetic energy'). When comparing two temperatures, ensure the curve for the higher temperature (T2) has a lower peak, is shifted to the right, and crosses the lower temperature (T1) curve only once. Crucially, show that both curves start at the origin and neither touches the x-axis at high energy.
Activation Energy ($E_a$): The Energy Barrier
Not every collision results in a reaction. For a reaction to take place, colliding particles must possess a certain minimum amount of combined kinetic energy. This minimum energy is called the activation energy, symbolised as . Think of it as an energy 'hill' or barrier that reactants must get over to become products. This energy is required to break existing chemical bonds in the reactants, allowing new bonds to form in the products. Reactions with a high tend to be slow at room temperature, as few particles have enough energy to overcome the barrier, whereas reactions with a low are often very fast.
Activation energy () is the minimum kinetic energy required for a collision to be successful (i.e., to result in a reaction).
It represents the energy needed to break bonds and initiate the reaction.
The value of is specific to a particular reaction pathway.
The Effect of Temperature on Reaction Rate
Increasing the temperature of a reaction mixture increases the rate of reaction. This happens for two main reasons. Firstly, as particles gain thermal energy, their kinetic energy increases. This causes them to move faster, leading to more frequent collisions per unit time. Secondly, and more importantly, increasing the temperature leads to a significant increase in the proportion of particles that have kinetic energy equal to or greater than the activation energy (). This means a much higher fraction of the collisions that occur are 'successful' or 'effective' collisions.
When asked to explain the effect of temperature on rate, you must mention both increased collision frequency and the increased proportion of particles with . However, always state that the second factor is far more significant. A common mistake is to only mention one of the two points.
Visualising Energy: The Boltzmann Distribution
The range of kinetic energies of particles in a substance at a given temperature can be represented by a graph called the Boltzmann distribution curve. The x-axis represents kinetic energy and the y-axis represents the number of molecules with that energy. The curve starts at the origin (no molecules have zero energy), rises to a peak representing the most probable energy, and then tails off asymptotically towards the x-axis (a few molecules have very high energies).
The total area under the curve represents the total number of molecules in the sample.
At a higher temperature (), the curve is 'flatter' and broader, and the peak is shifted to the right compared to a lower temperature ().
The activation energy () is marked as a vertical line on the energy axis. The area under the curve to the right of this line represents the number of molecules with sufficient energy to react.
Increasing temperature does not change , but it significantly increases the area to the right of the line, explaining the sharp increase in reaction rate.
Worked examples
See the formulas applied — reveal one step at a time, like the exam.
Sketch a Boltzmann distribution curve for a sample of gas at temperature . On the same axes, sketch a second curve for the same sample at a higher temperature . Mark the activation energy, , on your graph. Use your sketch to explain why the rate of reaction is greater at .
- 1
The axes are labelled 'Number of molecules' (y-axis) and 'Kinetic energy' (x-axis).
The decomposition of hydrogen peroxide is catalysed by manganese(IV) oxide: . Using a single Boltzmann distribution curve, explain how a catalyst increases the rate of this reaction.
- 1
A catalyst, such as manganese(IV) oxide, increases the rate of reaction by providing an alternative reaction pathway with a lower activation energy.
How it all connects
The big idea sits in the middle — tap a linked idea to explore the link.
Tap a linked idea to see how it connects back to the main topic — that connection is what examiners reward.
Glossary
Try to recall each definition before you reveal it.
Quick check
Answer in your head first — then tap to check. No pressure.
Revision flashcards
Flip the card. Test yourself before the exam.
What is Activation Energy ()?
The minimum amount of kinetic energy that colliding particles must possess for a chemical reaction to occur.
Key takeaways
Review these before you close the topic — retrieval beats re-reading.
- ✓
Two conditions for a successful collision: correct orientation and sufficient energy.
- ✓
Increasing temperature increases the average kinetic energy of particles.
- ✓
This results in more frequent collisions and, more significantly, a higher proportion of effective (successful) collisions.
Practice — then mark it
The whole point: a real Cambridge question, marked mark-by-mark.
Test your knowledge on temperature, activation energy and Boltzmann distributions
Test your knowledge on temperature, activation energy and Boltzmann distributions
Extra simulations & links
PhET, GeoGebra and other curated tools — open in a new tab.
Frequently asked
Checkpoint
One marked question is worth ten re-reads — close the loop before you move on.
Reading it isn’t knowing it — prove it.
Before you move on: do Test your knowledge on temperature, activation energy and Boltzmann distributions on paper, snap a photo, and get examiner-style feedback on exactly where you win and lose marks.