Collision Theory and Activation Energy Revision Notes | A Level AQA Chemistry

1. Why Do Reactions Need a Trigger?

Picture a flask of hydrogen gas and oxygen gas sitting on a bench. Both gases are present, both are colliding constantly — billions of times per second — yet nothing happens. No water forms. No explosion. Why not?

The answer lies in a concept that sits at the heart of kinetics: not every collision counts. Particles need more than just contact — they need to arrive with enough energy to break existing bonds and start forming new ones. That energy threshold has a name, and understanding it unlocks almost everything about reaction rates.

2. Collision Theory: The Foundation

Collision theory states that a chemical reaction can only occur when reactant particles collide. At first glance this sounds obvious, but the theory goes further: collisions must satisfy two conditions simultaneously — sufficient energy and correct orientation — before products can form.

Think about orientation first. If two molecules approach each other from the wrong angle, the reactive parts of each molecule never meet. The collision is geometrically unfavourable and the particles bounce apart. This alone eliminates a large proportion of all collisions.

The energy condition is the one AQA focuses on most heavily at this stage. Even when orientation is perfect, a collision will fail if the kinetic energy of the colliding particles is below a critical minimum value. That minimum value is the activation energy.

A useful mental model: imagine rolling a ball over a hill to reach a valley on the other side. The valley (products) may be lower than the starting point (reactants), meaning energy is released overall. But the ball still has to get over the hill first. The height of that hill is the activation energy.

Notice in the profile above that the activation energy is measured from the reactants' energy level up to the peak of the curve — the transition state. The overall enthalpy change ($\Delta H$) is the difference between reactants and products, but $E_a$ is always measured from the reactants upward to the peak.

3. Why Most Collisions Fail

At any given temperature, the particles in a sample do not all move at the same speed. Some are moving slowly, some at moderate speeds, and a small number are moving very fast. Because kinetic energy depends on speed, this means particles have a wide spread of energies at any one moment.

This spread is described by the Maxwell-Boltzmann distribution. The curve shows the number of particles (y-axis) plotted against their kinetic energy (x-axis). The activation energy $E_a$ sits as a vertical marker on the x-axis. Only particles whose energy lies to the right of $E_a$ — the shaded tail of the distribution — have enough energy to react when they collide.

The key insight from the distribution is that the vast majority of particles cluster around the most probable energy — the peak of the curve — which is typically well below $E_a$ for most reactions at room temperature. This is precisely why most collisions do not lead to a reaction: most particles do not carry enough energy when they meet.

There is a second reason collisions fail even when energy is sufficient: incorrect orientation. If the reactive sites on two molecules are not aligned when they collide, the collision is unproductive regardless of how much kinetic energy the particles possess. Both conditions — energy and geometry — must be met simultaneously.

4. The Maxwell-Boltzmann Distribution in Detail

Reading the Maxwell-Boltzmann distribution accurately is an essential exam skill. Several features of the curve carry marks in their own right, so it is worth knowing each one precisely.

• The curve starts at the origin (0, 0): no particles have zero kinetic energy.
• The curve rises to a peak at the most probable energy, then falls away asymptotically — it never touches the x-axis again at high energies.
• The total area under the curve represents the total number of particles in the sample.
• The activation energy $E_a$ is marked on the x-axis; the shaded area to the right of $E_a$ represents the fraction of particles with sufficient energy to react.
• At higher temperatures, the peak shifts right and downward; the curve is broader and flatter, and the shaded area beyond $E_a$ increases significantly.

A common exam question asks you to sketch two Maxwell-Boltzmann curves on the same axes — one at temperature $T_1$ and one at a higher temperature $T_2$. The mark scheme will check that your higher-temperature curve has a lower, broader peak that is shifted to the right, and that both curves start at the origin and never cross the x-axis.

5. Connecting Activation Energy to Reaction Rate

Now that the two ideas — activation energy and the Maxwell-Boltzmann distribution — are established, they connect directly. The rate of a reaction depends on the frequency of successful collisions. A successful collision requires the colliding particles to have energy equal to or greater than $E_a$.

Reactions with a high activation energy are generally slow at room temperature because only a tiny fraction of particles in the tail of the distribution possess enough energy. Reactions with a low activation energy are faster because a much larger fraction of particles can react on collision.

This also explains why temperature has such a dramatic effect on rate. A modest rise in temperature — say 10 K — does not dramatically increase the total number of collisions, but it does substantially increase the proportion of particles with energy above $E_a$. The area in the tail of the distribution grows disproportionately, which is why a 10 K rise can roughly double the rate of many reactions.

6. Exam Technique and Common Mistakes

AQA questions on collision theory and activation energy tend to use the command words 'explain' and 'describe'. These demand different types of response, and mixing them up is one of the most reliable ways to drop marks.

When asked to 'explain why increasing temperature increases the rate of reaction', a full-mark answer must include a chain of reasoning: higher temperature gives particles greater kinetic energy; a greater proportion of particles now have energy equal to or greater than the activation energy; the frequency of successful collisions increases; therefore the rate increases. Each link in that chain is a separate marking point.

When asked to 'describe' the Maxwell-Boltzmann distribution, focus on observable features of the curve — shape, peak position, asymptotic tail — without explaining causes. Save the 'because' language for 'explain' questions.